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Newton can also dial a phone number through the built-in speaker of the Newton device by simply holding a telephone handset up to the speaker and transmitting the appropriate tones. Fax and printing support is also built in at the operating system level, although it requires peripherals such as parallel adapters, PCMCIA cards, or serial modems, the most notable of which is the lightweight Newton Fax Modem released by Apple in 1993. It is powered by 2 AA batteries, and can also be used with a power adapter. It provides data transfer at 2,400 bit/s, and can also send and receive fax messages at 9,600 and 4,800 bit/s respectively. Power options The original Apple MessagePad and MessagePad 100 used four AAA batteries. They were eventually replaced by AA batteries with the release of the Apple MessagePad 110. The use of 4 AA NiCd (MessagePad 110, 120 and 130) and 4x AA NiMH cells (MP2x00 series, eMate 300) give a runtime of up to 30 hours (MP2100 with two 20 MB Linear Flash memory PC Cards, no backlight usage) and up to 24 hours with backlight on. While adding more weight to the handheld Newton devices than AAA batteries or custom battery packs, the choice of an easily replaceable/rechargeable cell format gives the user a still unsurpassed runtime and flexibility of power supply. This, together with the flash memory used as internal storage starting with the Apple MessagePad 120 (if all cells lost their power, no data was lost due to the non-volatility of this storage), gave birth to the slogan "Newton never dies, it only gets new batteries". Later efforts and improvements The Apple MessagePad 2000/2100, with a vastly improved handwriting recognition system, 162 MHz StrongARM SA-110 RISC processor, Newton OS 2.1, and a better, clearer, backlit screen, attracted critical plaudits. eMate 300 The eMate 300 was a Newton device in a laptop form factor offered to schools in 1997 as an inexpensive ($799 US, originally sold to education markets only) and durable computer for classroom use. However, in order to achieve its low price, the eMate 300 did not have all the speed and features of the contemporary MessagePad equivalent, the MessagePad 2000. The eMate was cancelled along with the rest of the Newton products in 1998. It is the only Newton device to use the ARM710 microprocessor (running at 25 MHz), have an integrated keyboard, use Newton OS 2.2 (officially numbered 2.1), and its batteries are officially irreplaceable, although several users replaced them with longer-lasting ones without any damage to the eMate hardware whatsoever. Prototypes Many prototypes of additional Newton devices were spotted. Most notable was a Newton tablet or "slate", a large, flat screen that could be written on.

Newton can also dial a phone number through the built-in speaker of the Newton device by simply holding a telephone handset up to the speaker and transmitting the appropriate tones. Fax and printing support is also built in at the operating system level, although it requires peripherals such as parallel adapters, PCMCIA cards, or serial modems, the most notable of which is the lightweight Newton Fax Modem released by Apple in 1993. It is powered by 2 AA batteries, and can also be used with a power adapter. It provides data transfer at 2,400 bit/s, and can also send and receive fax messages at 9,600 and 4,800 bit/s respectively. Power options The original Apple MessagePad and MessagePad 100 used four AAA batteries. They were eventually replaced by AA batteries with the release of the Apple MessagePad 110. The use of 4 AA NiCd (MessagePad 110, 120 and 130) and 4x AA NiMH cells (MP2x00 series, eMate 300) give a runtime of up to 30 hours (MP2100 with two 20 MB Linear Flash memory PC Cards, no backlight usage) and up to 24 hours with backlight on. While adding more weight to the handheld Newton devices than AAA batteries or custom battery packs, the choice of an easily replaceable/rechargeable cell format gives the user a still unsurpassed runtime and flexibility of power supply. This, together with the flash memory used as internal storage starting with the Apple MessagePad 120 (if all cells lost their power, no data was lost due to the non-volatility of this storage), gave birth to the slogan "Newton never dies, it only gets new batteries". Later efforts and improvements The Apple MessagePad 2000/2100, with a vastly improved handwriting recognition system, 162 MHz StrongARM SA-110 RISC processor, Newton OS 2.1, and a better, clearer, backlit screen, attracted critical plaudits. eMate 300 The eMate 300 was a Newton device in a laptop form factor offered to schools in 1997 as an inexpensive ($799 US, originally sold to education markets only) and durable computer for classroom use. However, in order to achieve its low price, the eMate 300 did not have all the speed and features of the contemporary MessagePad equivalent, the MessagePad 2000. The eMate was cancelled along with the rest of the Newton products in 1998. It is the only Newton device to use the ARM710 microprocessor (running at 25 MHz), have an integrated keyboard, use Newton OS 2.2 (officially numbered 2.1), and its batteries are officially irreplaceable, although several users replaced them with longer-lasting ones without any damage to the eMate hardware whatsoever. Prototypes Many prototypes of additional Newton devices were spotted. Most notable was a Newton tablet or "slate", a large, flat screen that could be written on.

Others included a "Kids Newton" with side handgrips and buttons, "VideoPads" which would have incorporated a video camera and screen on their flip-top covers for two-way communications, the "Mini 2000" which would have been very similar to a Palm Pilot, and the NewtonPhone developed by Siemens, which incorporated a handset and a keyboard. Market reception Fourteen months after Sculley demoed it at the May 1992, Chicago CES, the MessagePad was first offered for sale on August 2, 1993, at the Boston Macworld Expo. The hottest item at the show, it cost $900. 50,000 MessagePads were sold in the device's first three months on the market. The original Apple MessagePad and MessagePad 100 were limited by the very short lifetime of their inadequate AAA batteries. Critics also panned the handwriting recognition that was available in the debut models, which had been trumpeted in the Newton's marketing campaign. It was this problem that was skewered in the Doonesbury comic strips in which a written text entry is (erroneously) translated as "Egg Freckles? ", as well as in the animated series The Simpsons. However, the word 'freckles' was not included in the Newton dictionary, although a user could add it themselves. Difficulties were in part caused by the long time requirements for the Calligrapher handwriting recognition software to "learn" the user's handwriting; this process could take from two weeks to two months. Another factor which limited the early Newton devices' appeal was that desktop connectivity was not included in the basic retail package, a problem that was later solved with 2.x Newton devices - these were bundled with a serial cable and the appropriate Newton Connection Utilities software. Later versions of Newton OS offered improved handwriting recognition, quite possibly a leading reason for the continued popularity of the devices among Newton users. Even given the age of the hardware and software, Newtons still demand a sale price on the used market far greater than that of comparatively aged PDAs produced by other companies. In 2006, CNET compared an Apple MessagePad 2000 to a Samsung Q1, and the Newton was declared better. In 2009, CNET compared an Apple MessagePad 2000 to an iPhone 3GS, and the Newton was declared more innovative at its time of release. A chain of dedicated Newton only stores called Newton Source existed from 1994 until 1998. Locations included New York, Los Angeles, San Francisco, Chicago and Boston. The Westwood Village, California, near U.C.L.A. featured the trademark red and yellow light bulb Newton logo in neon. The stores provided an informative educational venue to learn about the Newton platform in a hands on relaxed fashion. The stores had no traditional computer retail counters and featured oval desktops where interested users could become intimately involved with the Newton product range. The stores were a model for the later Apple Stores.

Others included a "Kids Newton" with side handgrips and buttons, "VideoPads" which would have incorporated a video camera and screen on their flip-top covers for two-way communications, the "Mini 2000" which would have been very similar to a Palm Pilot, and the NewtonPhone developed by Siemens, which incorporated a handset and a keyboard. Market reception Fourteen months after Sculley demoed it at the May 1992, Chicago CES, the MessagePad was first offered for sale on August 2, 1993, at the Boston Macworld Expo. The hottest item at the show, it cost $900. 50,000 MessagePads were sold in the device's first three months on the market. The original Apple MessagePad and MessagePad 100 were limited by the very short lifetime of their inadequate AAA batteries. Critics also panned the handwriting recognition that was available in the debut models, which had been trumpeted in the Newton's marketing campaign. It was this problem that was skewered in the Doonesbury comic strips in which a written text entry is (erroneously) translated as "Egg Freckles? ", as well as in the animated series The Simpsons. However, the word 'freckles' was not included in the Newton dictionary, although a user could add it themselves. Difficulties were in part caused by the long time requirements for the Calligrapher handwriting recognition software to "learn" the user's handwriting; this process could take from two weeks to two months. Another factor which limited the early Newton devices' appeal was that desktop connectivity was not included in the basic retail package, a problem that was later solved with 2.x Newton devices - these were bundled with a serial cable and the appropriate Newton Connection Utilities software. Later versions of Newton OS offered improved handwriting recognition, quite possibly a leading reason for the continued popularity of the devices among Newton users. Even given the age of the hardware and software, Newtons still demand a sale price on the used market far greater than that of comparatively aged PDAs produced by other companies. In 2006, CNET compared an Apple MessagePad 2000 to a Samsung Q1, and the Newton was declared better. In 2009, CNET compared an Apple MessagePad 2000 to an iPhone 3GS, and the Newton was declared more innovative at its time of release. A chain of dedicated Newton only stores called Newton Source existed from 1994 until 1998. Locations included New York, Los Angeles, San Francisco, Chicago and Boston. The Westwood Village, California, near U.C.L.A. featured the trademark red and yellow light bulb Newton logo in neon. The stores provided an informative educational venue to learn about the Newton platform in a hands on relaxed fashion. The stores had no traditional computer retail counters and featured oval desktops where interested users could become intimately involved with the Newton product range. The stores were a model for the later Apple Stores.

Others included a "Kids Newton" with side handgrips and buttons, "VideoPads" which would have incorporated a video camera and screen on their flip-top covers for two-way communications, the "Mini 2000" which would have been very similar to a Palm Pilot, and the NewtonPhone developed by Siemens, which incorporated a handset and a keyboard. Market reception Fourteen months after Sculley demoed it at the May 1992, Chicago CES, the MessagePad was first offered for sale on August 2, 1993, at the Boston Macworld Expo. The hottest item at the show, it cost $900. 50,000 MessagePads were sold in the device's first three months on the market. The original Apple MessagePad and MessagePad 100 were limited by the very short lifetime of their inadequate AAA batteries. Critics also panned the handwriting recognition that was available in the debut models, which had been trumpeted in the Newton's marketing campaign. It was this problem that was skewered in the Doonesbury comic strips in which a written text entry is (erroneously) translated as "Egg Freckles? ", as well as in the animated series The Simpsons. However, the word 'freckles' was not included in the Newton dictionary, although a user could add it themselves. Difficulties were in part caused by the long time requirements for the Calligrapher handwriting recognition software to "learn" the user's handwriting; this process could take from two weeks to two months. Another factor which limited the early Newton devices' appeal was that desktop connectivity was not included in the basic retail package, a problem that was later solved with 2.x Newton devices - these were bundled with a serial cable and the appropriate Newton Connection Utilities software. Later versions of Newton OS offered improved handwriting recognition, quite possibly a leading reason for the continued popularity of the devices among Newton users. Even given the age of the hardware and software, Newtons still demand a sale price on the used market far greater than that of comparatively aged PDAs produced by other companies. In 2006, CNET compared an Apple MessagePad 2000 to a Samsung Q1, and the Newton was declared better. In 2009, CNET compared an Apple MessagePad 2000 to an iPhone 3GS, and the Newton was declared more innovative at its time of release. A chain of dedicated Newton only stores called Newton Source existed from 1994 until 1998. Locations included New York, Los Angeles, San Francisco, Chicago and Boston. The Westwood Village, California, near U.C.L.A. featured the trademark red and yellow light bulb Newton logo in neon. The stores provided an informative educational venue to learn about the Newton platform in a hands on relaxed fashion. The stores had no traditional computer retail counters and featured oval desktops where interested users could become intimately involved with the Newton product range. The stores were a model for the later Apple Stores.

Newton device models {| class="wikitable" |+ !Brand | colspan="2" |Apple |Sharp |Siemens | colspan="2" |Apple |Sharp |Apple |Digital Ocean |Motorola |Harris |Digital Ocean | colspan="4" |Apple | colspan="3" |Harris |Siemens |Schlumberger |- !Device |OMP (Original Newton MessagePad) |Newton "Dummy" |ExpertPad PI-7000 |Notephone. [better source needed] |MessagePad 100 |MessagePad 110 |Sharp ExpertPad PI-7100 |MessagePad 120 |Tarpon |Marco |SuperTech 2000 |Seahorse |MessagePad 130 |eMate 300 |MessagePad 2000 |MessagePad 2100 |Access Device 2000 |Access Device, GPS |Access Device, Wireline |Online Terminal, also known as Online Access Device(OAD) |Watson |- !Introduced |August 3, 1993 (US) December 1993 (Germany) |? |August 3, 1993(US), ? (Japan) |1993? | colspan="2" |March 1994 |April 1994 |October 1994 (Germany), January 1995 (US) | colspan="2" |January 1995 (US) |August 1995 in the US |January 1996 in the US |March 1996 | colspan="2" |March 1997 |November 1997 | colspan="3" |1998 |Announced 1997 |? |- !Discontinued | colspan="3" |March 1994 |? | colspan="2" |April 1995 |late 1994 |June 1996 |? |? |? |? |April 1997 | colspan="3" |February 1998 | | | | | |- !Code name |Junior | |? |? |Junior |Lindy |? |Gelato |? |? |? |? |Dante |? |Q |? | | | | | |- !Model No. |H1000 | |? |? |H1000 |H0059 |? |H0131 |? |? |? |? |H0196 |H0208 |H0136 |H0149 | | | | | |- !Processor | colspan="13" |ARM 610 (20 MHz) |ARM 710a (25 MHz) | colspan="7" |StrongARM SA-110 (162 MHz) |- !ROM | colspan="7" |4 MB | colspan="2" |4 MB (OS 1.3) or 8 MB (OS 2.0) |5 MB |4 MB | colspan="5" |8 MB | | | | | |- !System Memory (RAM) | colspan="5" |490 KB* SRAM |544 KB SRAM |490 KB* SRAM | colspan="2" |639/687 KB DRAM |544 KB SRAM |639 KB DRAM | colspan="2" |1199 KB DRAM |1 MB DRAM (Upgradable) |1 MB DRAM |4 MB DRAM | colspan="3" |1 MB DRAM |? |1 MB DRAM |- !User Storage | colspan="5" |150 KB* SRAM |480 KB SRAM |150 KB* SRAM | colspan="2" |385/1361 KB Flash RAM |480 KB SRAM |385 KB Flash RAM | colspan="2" |1361 KB Flash RAM |2 MB Flash RAM(Upgradable) | colspan="5" |4 MB Flash RAM |? |4 MB Flash RAM |- !Total RAM | colspan="5" |640 KB |1 MB |640 KB | colspan="2" |1.0/2.0 MB | colspan="2" |1 MB | colspan="2" |2.5 MB |3 MB (Upgradable via Internal Expansion) |5 MB |8 MB | colspan="3" |5 MB |?

Newton device models {| class="wikitable" |+ !Brand | colspan="2" |Apple |Sharp |Siemens | colspan="2" |Apple |Sharp |Apple |Digital Ocean |Motorola |Harris |Digital Ocean | colspan="4" |Apple | colspan="3" |Harris |Siemens |Schlumberger |- !Device |OMP (Original Newton MessagePad) |Newton "Dummy" |ExpertPad PI-7000 |Notephone. [better source needed] |MessagePad 100 |MessagePad 110 |Sharp ExpertPad PI-7100 |MessagePad 120 |Tarpon |Marco |SuperTech 2000 |Seahorse |MessagePad 130 |eMate 300 |MessagePad 2000 |MessagePad 2100 |Access Device 2000 |Access Device, GPS |Access Device, Wireline |Online Terminal, also known as Online Access Device(OAD) |Watson |- !Introduced |August 3, 1993 (US) December 1993 (Germany) |? |August 3, 1993(US), ? (Japan) |1993? | colspan="2" |March 1994 |April 1994 |October 1994 (Germany), January 1995 (US) | colspan="2" |January 1995 (US) |August 1995 in the US |January 1996 in the US |March 1996 | colspan="2" |March 1997 |November 1997 | colspan="3" |1998 |Announced 1997 |? |- !Discontinued | colspan="3" |March 1994 |? | colspan="2" |April 1995 |late 1994 |June 1996 |? |? |? |? |April 1997 | colspan="3" |February 1998 | | | | | |- !Code name |Junior | |? |? |Junior |Lindy |? |Gelato |? |? |? |? |Dante |? |Q |? | | | | | |- !Model No. |H1000 | |? |? |H1000 |H0059 |? |H0131 |? |? |? |? |H0196 |H0208 |H0136 |H0149 | | | | | |- !Processor | colspan="13" |ARM 610 (20 MHz) |ARM 710a (25 MHz) | colspan="7" |StrongARM SA-110 (162 MHz) |- !ROM | colspan="7" |4 MB | colspan="2" |4 MB (OS 1.3) or 8 MB (OS 2.0) |5 MB |4 MB | colspan="5" |8 MB | | | | | |- !System Memory (RAM) | colspan="5" |490 KB* SRAM |544 KB SRAM |490 KB* SRAM | colspan="2" |639/687 KB DRAM |544 KB SRAM |639 KB DRAM | colspan="2" |1199 KB DRAM |1 MB DRAM (Upgradable) |1 MB DRAM |4 MB DRAM | colspan="3" |1 MB DRAM |? |1 MB DRAM |- !User Storage | colspan="5" |150 KB* SRAM |480 KB SRAM |150 KB* SRAM | colspan="2" |385/1361 KB Flash RAM |480 KB SRAM |385 KB Flash RAM | colspan="2" |1361 KB Flash RAM |2 MB Flash RAM(Upgradable) | colspan="5" |4 MB Flash RAM |? |4 MB Flash RAM |- !Total RAM | colspan="5" |640 KB |1 MB |640 KB | colspan="2" |1.0/2.0 MB | colspan="2" |1 MB | colspan="2" |2.5 MB |3 MB (Upgradable via Internal Expansion) |5 MB |8 MB | colspan="3" |5 MB |?

Newton device models {| class="wikitable" |+ !Brand | colspan="2" |Apple |Sharp |Siemens | colspan="2" |Apple |Sharp |Apple |Digital Ocean |Motorola |Harris |Digital Ocean | colspan="4" |Apple | colspan="3" |Harris |Siemens |Schlumberger |- !Device |OMP (Original Newton MessagePad) |Newton "Dummy" |ExpertPad PI-7000 |Notephone. [better source needed] |MessagePad 100 |MessagePad 110 |Sharp ExpertPad PI-7100 |MessagePad 120 |Tarpon |Marco |SuperTech 2000 |Seahorse |MessagePad 130 |eMate 300 |MessagePad 2000 |MessagePad 2100 |Access Device 2000 |Access Device, GPS |Access Device, Wireline |Online Terminal, also known as Online Access Device(OAD) |Watson |- !Introduced |August 3, 1993 (US) December 1993 (Germany) |? |August 3, 1993(US), ? (Japan) |1993? | colspan="2" |March 1994 |April 1994 |October 1994 (Germany), January 1995 (US) | colspan="2" |January 1995 (US) |August 1995 in the US |January 1996 in the US |March 1996 | colspan="2" |March 1997 |November 1997 | colspan="3" |1998 |Announced 1997 |? |- !Discontinued | colspan="3" |March 1994 |? | colspan="2" |April 1995 |late 1994 |June 1996 |? |? |? |? |April 1997 | colspan="3" |February 1998 | | | | | |- !Code name |Junior | |? |? |Junior |Lindy |? |Gelato |? |? |? |? |Dante |? |Q |? | | | | | |- !Model No. |H1000 | |? |? |H1000 |H0059 |? |H0131 |? |? |? |? |H0196 |H0208 |H0136 |H0149 | | | | | |- !Processor | colspan="13" |ARM 610 (20 MHz) |ARM 710a (25 MHz) | colspan="7" |StrongARM SA-110 (162 MHz) |- !ROM | colspan="7" |4 MB | colspan="2" |4 MB (OS 1.3) or 8 MB (OS 2.0) |5 MB |4 MB | colspan="5" |8 MB | | | | | |- !System Memory (RAM) | colspan="5" |490 KB* SRAM |544 KB SRAM |490 KB* SRAM | colspan="2" |639/687 KB DRAM |544 KB SRAM |639 KB DRAM | colspan="2" |1199 KB DRAM |1 MB DRAM (Upgradable) |1 MB DRAM |4 MB DRAM | colspan="3" |1 MB DRAM |? |1 MB DRAM |- !User Storage | colspan="5" |150 KB* SRAM |480 KB SRAM |150 KB* SRAM | colspan="2" |385/1361 KB Flash RAM |480 KB SRAM |385 KB Flash RAM | colspan="2" |1361 KB Flash RAM |2 MB Flash RAM(Upgradable) | colspan="5" |4 MB Flash RAM |? |4 MB Flash RAM |- !Total RAM | colspan="5" |640 KB |1 MB |640 KB | colspan="2" |1.0/2.0 MB | colspan="2" |1 MB | colspan="2" |2.5 MB |3 MB (Upgradable via Internal Expansion) |5 MB |8 MB | colspan="3" |5 MB |?

|5 MB |- !Display | colspan="5" |336 × 240 (B&W) |320 × 240 (B&W) |336 × 240 (B&W) |320 × 240 (B&W) |320 × 240 (B&W) w/ backlight |320 × 240 (B&W) | colspan="3" |320 × 240 (B&W) w/ backlight | colspan="6" |480 × 320 grayscale (16 shades) w/ backlight | |480 × 320 greyscale (16 shades) w/ backlight |- !Newton OS version | colspan="3" |1.0 to 1.05, or 1.10 to 1.11 |1.11 | colspan="2" |1.2 or 1.3 |1.3 | colspan="2" |1.3 or 2.0 | colspan="2" |1.3 | colspan="2" |2.0 |2.1 (2.2) | colspan="2" |2.1 | colspan="5" |2.1 |- !Newton OS languages |English or German | |English or Japanese |German |English, German or French |English or French |English or Japanese |English, German or French | colspan="4" |English |English or German | colspan="2" |English |English or German | colspan="3" |English |German |French |- !Connectivity | colspan="3" |RS422, LocalTalk & SHARP ASK Infrared |Modem and Telephone dock Attachment | colspan="4" |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk, Infrared, ARDIS Network |RS232, LocalTalk WLAN, V.22bis modem, Analog/Digital Cellular, CDPD, RAM, ARDIS , Trunk Radio |RS232, LocalTalk, CDPD, WLAN, Optional dGPS, GSM, or IR via modular attachments |RS422, LocalTalk & SHARP ASK Infrared |IrDA, headphone port, Interconnect port, LocalTalk, Audio I/O, Autodock |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |Dual-mode IR; IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock | colspan="3" |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |? |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |- !PCMCIA | colspan="13" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot I/II/III, 5v | colspan="2" |2 PCMCIA-slot II, 5v or 12v | colspan="2" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot II, 5v or 12v, 2nd slot Propriety Rado Card | colspan="2" |1 PCMCIA-slot II, 5v or 12v, 1 Smart Card Reader |- !Power | colspan="5" |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply | colspan="2" |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiMH battery pack (built-in) or external power supply | colspan="2" |4 AA or NiMH rechargeable or external power supply | colspan="3" |Custom NiMH rechargeable or external power supply |? Unknown, but likely external power supply |4 AA or NiMH rechargeable or external power supply |- !Dimensions (HxWxD) | | | (lid open) | colspan="2" | | | (lid open) | | | |? | | | | colspan="2" | |? |? |? |9 x 14.5 x 5.1 inches (23 x 37 x 13 cm) |? |- !Weight | | | with batteries installed | | | with batteries installed | with batteries installed |with batteries installed | | |?

|5 MB |- !Display | colspan="5" |336 × 240 (B&W) |320 × 240 (B&W) |336 × 240 (B&W) |320 × 240 (B&W) |320 × 240 (B&W) w/ backlight |320 × 240 (B&W) | colspan="3" |320 × 240 (B&W) w/ backlight | colspan="6" |480 × 320 grayscale (16 shades) w/ backlight | |480 × 320 greyscale (16 shades) w/ backlight |- !Newton OS version | colspan="3" |1.0 to 1.05, or 1.10 to 1.11 |1.11 | colspan="2" |1.2 or 1.3 |1.3 | colspan="2" |1.3 or 2.0 | colspan="2" |1.3 | colspan="2" |2.0 |2.1 (2.2) | colspan="2" |2.1 | colspan="5" |2.1 |- !Newton OS languages |English or German | |English or Japanese |German |English, German or French |English or French |English or Japanese |English, German or French | colspan="4" |English |English or German | colspan="2" |English |English or German | colspan="3" |English |German |French |- !Connectivity | colspan="3" |RS422, LocalTalk & SHARP ASK Infrared |Modem and Telephone dock Attachment | colspan="4" |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk, Infrared, ARDIS Network |RS232, LocalTalk WLAN, V.22bis modem, Analog/Digital Cellular, CDPD, RAM, ARDIS , Trunk Radio |RS232, LocalTalk, CDPD, WLAN, Optional dGPS, GSM, or IR via modular attachments |RS422, LocalTalk & SHARP ASK Infrared |IrDA, headphone port, Interconnect port, LocalTalk, Audio I/O, Autodock |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |Dual-mode IR; IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock | colspan="3" |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |? |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |- !PCMCIA | colspan="13" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot I/II/III, 5v | colspan="2" |2 PCMCIA-slot II, 5v or 12v | colspan="2" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot II, 5v or 12v, 2nd slot Propriety Rado Card | colspan="2" |1 PCMCIA-slot II, 5v or 12v, 1 Smart Card Reader |- !Power | colspan="5" |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply | colspan="2" |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiMH battery pack (built-in) or external power supply | colspan="2" |4 AA or NiMH rechargeable or external power supply | colspan="3" |Custom NiMH rechargeable or external power supply |? Unknown, but likely external power supply |4 AA or NiMH rechargeable or external power supply |- !Dimensions (HxWxD) | | | (lid open) | colspan="2" | | | (lid open) | | | |? | | | | colspan="2" | |? |? |? |9 x 14.5 x 5.1 inches (23 x 37 x 13 cm) |? |- !Weight | | | with batteries installed | | | with batteries installed | with batteries installed |with batteries installed | | |?

|5 MB |- !Display | colspan="5" |336 × 240 (B&W) |320 × 240 (B&W) |336 × 240 (B&W) |320 × 240 (B&W) |320 × 240 (B&W) w/ backlight |320 × 240 (B&W) | colspan="3" |320 × 240 (B&W) w/ backlight | colspan="6" |480 × 320 grayscale (16 shades) w/ backlight | |480 × 320 greyscale (16 shades) w/ backlight |- !Newton OS version | colspan="3" |1.0 to 1.05, or 1.10 to 1.11 |1.11 | colspan="2" |1.2 or 1.3 |1.3 | colspan="2" |1.3 or 2.0 | colspan="2" |1.3 | colspan="2" |2.0 |2.1 (2.2) | colspan="2" |2.1 | colspan="5" |2.1 |- !Newton OS languages |English or German | |English or Japanese |German |English, German or French |English or French |English or Japanese |English, German or French | colspan="4" |English |English or German | colspan="2" |English |English or German | colspan="3" |English |German |French |- !Connectivity | colspan="3" |RS422, LocalTalk & SHARP ASK Infrared |Modem and Telephone dock Attachment | colspan="4" |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk & SHARP ASK Infrared |RS422, LocalTalk, Infrared, ARDIS Network |RS232, LocalTalk WLAN, V.22bis modem, Analog/Digital Cellular, CDPD, RAM, ARDIS , Trunk Radio |RS232, LocalTalk, CDPD, WLAN, Optional dGPS, GSM, or IR via modular attachments |RS422, LocalTalk & SHARP ASK Infrared |IrDA, headphone port, Interconnect port, LocalTalk, Audio I/O, Autodock |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |Dual-mode IR; IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock | colspan="3" |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |? |Dual-mode IR;IrDA & SHARP ASK, LocalTalk, Audio I/O, Autodock, Phone I/O |- !PCMCIA | colspan="13" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot I/II/III, 5v | colspan="2" |2 PCMCIA-slot II, 5v or 12v | colspan="2" |1 PCMCIA-slot II, 5v or 12v |1 PCMCIA-slot II, 5v or 12v, 2nd slot Propriety Rado Card | colspan="2" |1 PCMCIA-slot II, 5v or 12v, 1 Smart Card Reader |- !Power | colspan="5" |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply |4 AAA or NiCd rechargeable or external power supply |4 AA or NiCd rechargeable or external power supply | colspan="2" |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiCd battery pack or external power supply |4 AA or NiCd rechargeable or external power supply |NiMH battery pack (built-in) or external power supply | colspan="2" |4 AA or NiMH rechargeable or external power supply | colspan="3" |Custom NiMH rechargeable or external power supply |? Unknown, but likely external power supply |4 AA or NiMH rechargeable or external power supply |- !Dimensions (HxWxD) | | | (lid open) | colspan="2" | | | (lid open) | | | |? | | | | colspan="2" | |? |? |? |9 x 14.5 x 5.1 inches (23 x 37 x 13 cm) |? |- !Weight | | | with batteries installed | | | with batteries installed | with batteries installed |with batteries installed | | |?

| | with batteries installed | | colspan="2" | |? |? |? |? |? |} * Varies with Installed OS Notes: The eMate 300 actually has ROM chips silk screened with 2.2 on them. Stephanie Mak on her website discusses this: If one removes all patches to the eMate 300 (by replacing the ROM chip, and then putting in the original one again, as the eMate and the MessagePad 2000/2100 devices erase their memory completely after replacing the chip), the result will be the Newton OS saying that this is version 2.2.00. Also, the Original MessagePad and the MessagePad 100 share the same model number, as they only differ in the ROM chip version. (The OMP has OS versions 1.0 to 1.05, or 1.10 to 1.11, while the MP100 has 1.3 that can be upgraded with various patches.) Other uses There were a number of projects that used the Newton as a portable information device in cultural settings such as museums. For example, Visible Interactive created a walking tour in San Francisco's Chinatown but the most significant effort took place in Malaysia at the Petronas Discovery Center, known as Petrosains. In 1995, an exhibit design firm, DMCD Inc., was awarded the contract to design a new science museum in the Petronas Towers in Kuala Lumpur. A major factor in the award was the concept that visitors would use a Newton device to access additional information, find out where they were in the museum, listen to audio, see animations, control robots and other media, and to bookmark information for printout at the end of the exhibit. The device became known as the ARIF, a Malay word for "wise man" or "seer" and it was also an acronym for A Resourceful Informative Friend. Some 400 ARIFS were installed and over 300 are still in use today. The development of the ARIF system was extremely complex and required a team of hardware and software engineers, designers, and writers. ARIF is an ancestor of the PDA systems used in museums today and it boasted features that have not been attempted since. Anyway & Company firm was involved with the Petronas Discovery Center project back in 1998 and NDAs were signed which prevents getting to know more information about this project. It was confirmed that they purchased of MP2000u or MP2100's by this firm on the behalf of the project under the name of "Petrosains Project Account". By 1998 they had invested heavily into the R&D of this project with the Newton at the center. After Apple officially cancelled the Newton in 1998 they had to acquire as many Newtons as possible for this project. It was estimated initially 1000 Newtons, but later readjusted the figure to possibly 750 Newtons. They placed an “Internet Call” for Newtons. They purchased them in large and small quantities. The Newton was also used in healthcare applications, for example in collecting data directly from patients.

| | with batteries installed | | colspan="2" | |? |? |? |? |? |} * Varies with Installed OS Notes: The eMate 300 actually has ROM chips silk screened with 2.2 on them. Stephanie Mak on her website discusses this: If one removes all patches to the eMate 300 (by replacing the ROM chip, and then putting in the original one again, as the eMate and the MessagePad 2000/2100 devices erase their memory completely after replacing the chip), the result will be the Newton OS saying that this is version 2.2.00. Also, the Original MessagePad and the MessagePad 100 share the same model number, as they only differ in the ROM chip version. (The OMP has OS versions 1.0 to 1.05, or 1.10 to 1.11, while the MP100 has 1.3 that can be upgraded with various patches.) Other uses There were a number of projects that used the Newton as a portable information device in cultural settings such as museums. For example, Visible Interactive created a walking tour in San Francisco's Chinatown but the most significant effort took place in Malaysia at the Petronas Discovery Center, known as Petrosains. In 1995, an exhibit design firm, DMCD Inc., was awarded the contract to design a new science museum in the Petronas Towers in Kuala Lumpur. A major factor in the award was the concept that visitors would use a Newton device to access additional information, find out where they were in the museum, listen to audio, see animations, control robots and other media, and to bookmark information for printout at the end of the exhibit. The device became known as the ARIF, a Malay word for "wise man" or "seer" and it was also an acronym for A Resourceful Informative Friend. Some 400 ARIFS were installed and over 300 are still in use today. The development of the ARIF system was extremely complex and required a team of hardware and software engineers, designers, and writers. ARIF is an ancestor of the PDA systems used in museums today and it boasted features that have not been attempted since. Anyway & Company firm was involved with the Petronas Discovery Center project back in 1998 and NDAs were signed which prevents getting to know more information about this project. It was confirmed that they purchased of MP2000u or MP2100's by this firm on the behalf of the project under the name of "Petrosains Project Account". By 1998 they had invested heavily into the R&D of this project with the Newton at the center. After Apple officially cancelled the Newton in 1998 they had to acquire as many Newtons as possible for this project. It was estimated initially 1000 Newtons, but later readjusted the figure to possibly 750 Newtons. They placed an “Internet Call” for Newtons. They purchased them in large and small quantities. The Newton was also used in healthcare applications, for example in collecting data directly from patients.

| | with batteries installed | | colspan="2" | |? |? |? |? |? |} * Varies with Installed OS Notes: The eMate 300 actually has ROM chips silk screened with 2.2 on them. Stephanie Mak on her website discusses this: If one removes all patches to the eMate 300 (by replacing the ROM chip, and then putting in the original one again, as the eMate and the MessagePad 2000/2100 devices erase their memory completely after replacing the chip), the result will be the Newton OS saying that this is version 2.2.00. Also, the Original MessagePad and the MessagePad 100 share the same model number, as they only differ in the ROM chip version. (The OMP has OS versions 1.0 to 1.05, or 1.10 to 1.11, while the MP100 has 1.3 that can be upgraded with various patches.) Other uses There were a number of projects that used the Newton as a portable information device in cultural settings such as museums. For example, Visible Interactive created a walking tour in San Francisco's Chinatown but the most significant effort took place in Malaysia at the Petronas Discovery Center, known as Petrosains. In 1995, an exhibit design firm, DMCD Inc., was awarded the contract to design a new science museum in the Petronas Towers in Kuala Lumpur. A major factor in the award was the concept that visitors would use a Newton device to access additional information, find out where they were in the museum, listen to audio, see animations, control robots and other media, and to bookmark information for printout at the end of the exhibit. The device became known as the ARIF, a Malay word for "wise man" or "seer" and it was also an acronym for A Resourceful Informative Friend. Some 400 ARIFS were installed and over 300 are still in use today. The development of the ARIF system was extremely complex and required a team of hardware and software engineers, designers, and writers. ARIF is an ancestor of the PDA systems used in museums today and it boasted features that have not been attempted since. Anyway & Company firm was involved with the Petronas Discovery Center project back in 1998 and NDAs were signed which prevents getting to know more information about this project. It was confirmed that they purchased of MP2000u or MP2100's by this firm on the behalf of the project under the name of "Petrosains Project Account". By 1998 they had invested heavily into the R&D of this project with the Newton at the center. After Apple officially cancelled the Newton in 1998 they had to acquire as many Newtons as possible for this project. It was estimated initially 1000 Newtons, but later readjusted the figure to possibly 750 Newtons. They placed an “Internet Call” for Newtons. They purchased them in large and small quantities. The Newton was also used in healthcare applications, for example in collecting data directly from patients.

Newtons were used as electronic diaries, with patients entering their symptoms and other information concerning their health status on a daily basis. The compact size of the device and its ease of use made it possible for the electronic diaries to be carried around and used in the patients' everyday life setting. This was an early example of electronic patient-reported outcomes (ePRO) See also Newton (platform) Newton OS eMate 300 NewtonScript Orphaned technology Pen computing References Bibliography Apple's press release on the debut of the MessagePad 2100 Apple's overview of features & limitations of Newton Connection Utilities Newton overview at Newton Source archived from Apple Newton FAQ Pen Computing's First Look at Newton OS 2.0 Newton Gallery Birth of the Newton The Newton Hall of Fame: People behind the Newton Pen Computing's Why did Apple kill the Newton? Pen Computing's Newton Notes column archive A.I. Magazine article by Yaeger on Newton HWR design, algorithms, & quality and associated slides Info on Newton HWR from Apple's HWR Technical Lead External links Additional resources and information Defying Gravity: The Making of Newton, by Kounalakis & Menuez (Hardcover) Hardcover: 192 pages Publisher: Beyond Words Publishing (October 1993) Complete Developer's manual for the StrongARM SA-110 Beginner's overview of the StrongARM SA-110 Microprocessor Reviews MessagePad 2000 review at "The History and Macintosh Society" Prof. Wittmann's collection of Newton & MessagePad reviews Apple Newton Products introduced in 1993 Apple Inc. personal digital assistants

Newtons were used as electronic diaries, with patients entering their symptoms and other information concerning their health status on a daily basis. The compact size of the device and its ease of use made it possible for the electronic diaries to be carried around and used in the patients' everyday life setting. This was an early example of electronic patient-reported outcomes (ePRO) See also Newton (platform) Newton OS eMate 300 NewtonScript Orphaned technology Pen computing References Bibliography Apple's press release on the debut of the MessagePad 2100 Apple's overview of features & limitations of Newton Connection Utilities Newton overview at Newton Source archived from Apple Newton FAQ Pen Computing's First Look at Newton OS 2.0 Newton Gallery Birth of the Newton The Newton Hall of Fame: People behind the Newton Pen Computing's Why did Apple kill the Newton? Pen Computing's Newton Notes column archive A.I. Magazine article by Yaeger on Newton HWR design, algorithms, & quality and associated slides Info on Newton HWR from Apple's HWR Technical Lead External links Additional resources and information Defying Gravity: The Making of Newton, by Kounalakis & Menuez (Hardcover) Hardcover: 192 pages Publisher: Beyond Words Publishing (October 1993) Complete Developer's manual for the StrongARM SA-110 Beginner's overview of the StrongARM SA-110 Microprocessor Reviews MessagePad 2000 review at "The History and Macintosh Society" Prof. Wittmann's collection of Newton & MessagePad reviews Apple Newton Products introduced in 1993 Apple Inc. personal digital assistants

Newtons were used as electronic diaries, with patients entering their symptoms and other information concerning their health status on a daily basis. The compact size of the device and its ease of use made it possible for the electronic diaries to be carried around and used in the patients' everyday life setting. This was an early example of electronic patient-reported outcomes (ePRO) See also Newton (platform) Newton OS eMate 300 NewtonScript Orphaned technology Pen computing References Bibliography Apple's press release on the debut of the MessagePad 2100 Apple's overview of features & limitations of Newton Connection Utilities Newton overview at Newton Source archived from Apple Newton FAQ Pen Computing's First Look at Newton OS 2.0 Newton Gallery Birth of the Newton The Newton Hall of Fame: People behind the Newton Pen Computing's Why did Apple kill the Newton? Pen Computing's Newton Notes column archive A.I. Magazine article by Yaeger on Newton HWR design, algorithms, & quality and associated slides Info on Newton HWR from Apple's HWR Technical Lead External links Additional resources and information Defying Gravity: The Making of Newton, by Kounalakis & Menuez (Hardcover) Hardcover: 192 pages Publisher: Beyond Words Publishing (October 1993) Complete Developer's manual for the StrongARM SA-110 Beginner's overview of the StrongARM SA-110 Microprocessor Reviews MessagePad 2000 review at "The History and Macintosh Society" Prof. Wittmann's collection of Newton & MessagePad reviews Apple Newton Products introduced in 1993 Apple Inc. personal digital assistants

A. E. van Vogt Alfred Elton van Vogt (; April 26, 1912 – January 26, 2000) was a Canadian-born science fiction author. His fragmented, bizarre narrative style influenced later science fiction writers, notably Philip K. Dick. He was one of the most popular and influential practitioners of science fiction in the mid-twentieth century, the genre's so-called Golden Age, and one of the most complex. The Science Fiction Writers of America named him their 14th Grand Master in 1995 (presented 1996). Early life Alfred Vogt (both "Elton" and "van" were added much later) was born on April 26, 1912, on his grandparents' farm in Edenburg, Manitoba, a tiny (and now defunct) Russian Mennonite community east of Gretna, Manitoba, Canada, in the Mennonite West Reserve. He was the third of six children born to Heinrich "Henry" Vogt and Aganetha "Agnes" Vogt (née Buhr), both of whom were born in Manitoba and grew up in heavily immigrant communities. Until age four, van Vogt and his family spoke only Plautdietsch at home. For the first dozen or so years of his life, van Vogt's father, Henry Vogt, a lawyer, moved his family several times within western Canada, moving to Neville, Saskatchewan; Morden, Manitoba; and finally Winnipeg, Manitoba. Alfred Vogt found these moves difficult, later remarking: By the 1920s, living in Winnipeg, father Henry worked as an agent for a steamship company, but the stock market crash of 1929 proved financially disastrous, and the family could not afford to send Alfred to college. During his teen years, Alfred worked as a farmhand and a truck driver, and by the age of 19, he was working in Ottawa for the Canadian Census Bureau. He began his writing career with stories in the true confession style of pulp magazines such as True Story. Most of these stories were published anonymously, with the first-person narratives allegedly being written by people (often women) in extraordinary, emotional, and life-changing circumstances. After a year in Ottawa, he moved back to Winnipeg, where he sold newspaper advertising space and continued to write. While continuing to pen melodramatic "true confessions" stories through 1937, he also began writing short radio dramas for local radio station CKY, as well as conducting interviews published in trade magazines. He added the middle name "Elton" at some point in the mid-1930s, and at least one confessional story (1937's "To Be His Keeper") was sold to the Toronto Star, who misspelled his name "Alfred Alton Bogt" in the byline. Shortly thereafter, he added the "van" to his surname, and from that point forward he used the name "A. E. van Vogt" both personally and professionally. Career By 1938, van Vogt decided to switch to writing science fiction, a genre he enjoyed reading. He was inspired by the August 1938 issue of Astounding Science Fiction, which he picked up at a newsstand. John W. Campbell's novelette "Who Goes There?"

A. E. van Vogt Alfred Elton van Vogt (; April 26, 1912 – January 26, 2000) was a Canadian-born science fiction author. His fragmented, bizarre narrative style influenced later science fiction writers, notably Philip K. Dick. He was one of the most popular and influential practitioners of science fiction in the mid-twentieth century, the genre's so-called Golden Age, and one of the most complex. The Science Fiction Writers of America named him their 14th Grand Master in 1995 (presented 1996). Early life Alfred Vogt (both "Elton" and "van" were added much later) was born on April 26, 1912, on his grandparents' farm in Edenburg, Manitoba, a tiny (and now defunct) Russian Mennonite community east of Gretna, Manitoba, Canada, in the Mennonite West Reserve. He was the third of six children born to Heinrich "Henry" Vogt and Aganetha "Agnes" Vogt (née Buhr), both of whom were born in Manitoba and grew up in heavily immigrant communities. Until age four, van Vogt and his family spoke only Plautdietsch at home. For the first dozen or so years of his life, van Vogt's father, Henry Vogt, a lawyer, moved his family several times within western Canada, moving to Neville, Saskatchewan; Morden, Manitoba; and finally Winnipeg, Manitoba. Alfred Vogt found these moves difficult, later remarking: By the 1920s, living in Winnipeg, father Henry worked as an agent for a steamship company, but the stock market crash of 1929 proved financially disastrous, and the family could not afford to send Alfred to college. During his teen years, Alfred worked as a farmhand and a truck driver, and by the age of 19, he was working in Ottawa for the Canadian Census Bureau. He began his writing career with stories in the true confession style of pulp magazines such as True Story. Most of these stories were published anonymously, with the first-person narratives allegedly being written by people (often women) in extraordinary, emotional, and life-changing circumstances. After a year in Ottawa, he moved back to Winnipeg, where he sold newspaper advertising space and continued to write. While continuing to pen melodramatic "true confessions" stories through 1937, he also began writing short radio dramas for local radio station CKY, as well as conducting interviews published in trade magazines. He added the middle name "Elton" at some point in the mid-1930s, and at least one confessional story (1937's "To Be His Keeper") was sold to the Toronto Star, who misspelled his name "Alfred Alton Bogt" in the byline. Shortly thereafter, he added the "van" to his surname, and from that point forward he used the name "A. E. van Vogt" both personally and professionally. Career By 1938, van Vogt decided to switch to writing science fiction, a genre he enjoyed reading. He was inspired by the August 1938 issue of Astounding Science Fiction, which he picked up at a newsstand. John W. Campbell's novelette "Who Goes There?"

(later adapted into The Thing from Another World and The Thing) inspired van Vogt to write "Vault of the Beast", which he submitted to that same magazine. Campbell, who edited Astounding (and had written the story under a pseudonym), sent van Vogt a rejection letter, but one which encouraged van Vogt to try again. Van Vogt sent another story, entitled "Black Destroyer", which was accepted. It featured a fierce, carnivorous alien stalking the crew of a spaceship, and served as the inspiration for multiple science fiction movies, including Alien (1979). A revised version of "Vault of the Beast" was published in 1940. While still living in Winnipeg, in 1939 van Vogt married Edna Mayne Hull, a fellow Manitoban. Hull, who had previously worked as a private secretary, went on to act as van Vogt's typist, and was credited with writing several SF stories of her own throughout the early 1940s. The outbreak of World War II in September 1939 caused a change in van Vogt's circumstances. Ineligible for military service due to his poor eyesight, he accepted a clerking job with the Canadian Department of National Defence. This necessitated a move back to Ottawa, where he and his wife stayed for the next year and a half. Meanwhile, his writing career continued. "Discord in Scarlet" was van Vogt's second story to be published, also appearing as the cover story. It was accompanied by interior illustrations created by Frank Kramer and Paul Orban. (Van Vogt and Kramer thus debuted in the issue of Astounding that is sometimes identified as the start of the Golden Age of Science Fiction.) Among his most famous works of this era, "Far Centaurus" appeared in the January 1944 edition of Astounding. Van Vogt's first completed novel, and one of his most famous, is Slan (Arkham House, 1946), which Campbell serialized in Astounding (September to December 1940). Using what became one of van Vogt's recurring themes, it told the story of a nine-year-old superman living in a world in which his kind are slain by Homo sapiens. Others saw van Vogt's talent from his first story, and in May 1941 van Vogt decided to become a full-time writer, quitting his job at the Canadian Department of National Defence. Freed from the necessity of living in Ottawa, he and his wife lived for a time in the Gatineau region of Quebec before moving to Toronto in the fall of 1941. Prolific throughout this period, van Vogt wrote many of his more famous short stories and novels in the years from 1941 through 1944. The novels The Book of Ptath and The Weapon Makers both appeared in magazines in serial form during this period; they were later published in book form after World War II. As well, several (though not all) of the stories that were compiled to make up the novels The Weapon Shops of Isher, The Mixed Men and The War Against the Rull were published during this time.

(later adapted into The Thing from Another World and The Thing) inspired van Vogt to write "Vault of the Beast", which he submitted to that same magazine. Campbell, who edited Astounding (and had written the story under a pseudonym), sent van Vogt a rejection letter, but one which encouraged van Vogt to try again. Van Vogt sent another story, entitled "Black Destroyer", which was accepted. It featured a fierce, carnivorous alien stalking the crew of a spaceship, and served as the inspiration for multiple science fiction movies, including Alien (1979). A revised version of "Vault of the Beast" was published in 1940. While still living in Winnipeg, in 1939 van Vogt married Edna Mayne Hull, a fellow Manitoban. Hull, who had previously worked as a private secretary, went on to act as van Vogt's typist, and was credited with writing several SF stories of her own throughout the early 1940s. The outbreak of World War II in September 1939 caused a change in van Vogt's circumstances. Ineligible for military service due to his poor eyesight, he accepted a clerking job with the Canadian Department of National Defence. This necessitated a move back to Ottawa, where he and his wife stayed for the next year and a half. Meanwhile, his writing career continued. "Discord in Scarlet" was van Vogt's second story to be published, also appearing as the cover story. It was accompanied by interior illustrations created by Frank Kramer and Paul Orban. (Van Vogt and Kramer thus debuted in the issue of Astounding that is sometimes identified as the start of the Golden Age of Science Fiction.) Among his most famous works of this era, "Far Centaurus" appeared in the January 1944 edition of Astounding. Van Vogt's first completed novel, and one of his most famous, is Slan (Arkham House, 1946), which Campbell serialized in Astounding (September to December 1940). Using what became one of van Vogt's recurring themes, it told the story of a nine-year-old superman living in a world in which his kind are slain by Homo sapiens. Others saw van Vogt's talent from his first story, and in May 1941 van Vogt decided to become a full-time writer, quitting his job at the Canadian Department of National Defence. Freed from the necessity of living in Ottawa, he and his wife lived for a time in the Gatineau region of Quebec before moving to Toronto in the fall of 1941. Prolific throughout this period, van Vogt wrote many of his more famous short stories and novels in the years from 1941 through 1944. The novels The Book of Ptath and The Weapon Makers both appeared in magazines in serial form during this period; they were later published in book form after World War II. As well, several (though not all) of the stories that were compiled to make up the novels The Weapon Shops of Isher, The Mixed Men and The War Against the Rull were published during this time.

(later adapted into The Thing from Another World and The Thing) inspired van Vogt to write "Vault of the Beast", which he submitted to that same magazine. Campbell, who edited Astounding (and had written the story under a pseudonym), sent van Vogt a rejection letter, but one which encouraged van Vogt to try again. Van Vogt sent another story, entitled "Black Destroyer", which was accepted. It featured a fierce, carnivorous alien stalking the crew of a spaceship, and served as the inspiration for multiple science fiction movies, including Alien (1979). A revised version of "Vault of the Beast" was published in 1940. While still living in Winnipeg, in 1939 van Vogt married Edna Mayne Hull, a fellow Manitoban. Hull, who had previously worked as a private secretary, went on to act as van Vogt's typist, and was credited with writing several SF stories of her own throughout the early 1940s. The outbreak of World War II in September 1939 caused a change in van Vogt's circumstances. Ineligible for military service due to his poor eyesight, he accepted a clerking job with the Canadian Department of National Defence. This necessitated a move back to Ottawa, where he and his wife stayed for the next year and a half. Meanwhile, his writing career continued. "Discord in Scarlet" was van Vogt's second story to be published, also appearing as the cover story. It was accompanied by interior illustrations created by Frank Kramer and Paul Orban. (Van Vogt and Kramer thus debuted in the issue of Astounding that is sometimes identified as the start of the Golden Age of Science Fiction.) Among his most famous works of this era, "Far Centaurus" appeared in the January 1944 edition of Astounding. Van Vogt's first completed novel, and one of his most famous, is Slan (Arkham House, 1946), which Campbell serialized in Astounding (September to December 1940). Using what became one of van Vogt's recurring themes, it told the story of a nine-year-old superman living in a world in which his kind are slain by Homo sapiens. Others saw van Vogt's talent from his first story, and in May 1941 van Vogt decided to become a full-time writer, quitting his job at the Canadian Department of National Defence. Freed from the necessity of living in Ottawa, he and his wife lived for a time in the Gatineau region of Quebec before moving to Toronto in the fall of 1941. Prolific throughout this period, van Vogt wrote many of his more famous short stories and novels in the years from 1941 through 1944. The novels The Book of Ptath and The Weapon Makers both appeared in magazines in serial form during this period; they were later published in book form after World War II. As well, several (though not all) of the stories that were compiled to make up the novels The Weapon Shops of Isher, The Mixed Men and The War Against the Rull were published during this time.

California and post-war writing (1944–1950) In November 1944, van Vogt and Hull moved to Hollywood; van Vogt would spend the rest of his life in California. He had been using the name "A. E. van Vogt" in his public life for several years, and as part of the process of obtaining American citizenship in 1945 he finally and formally changed his legal name from Alfred Vogt to Alfred Elton van Vogt. To his friends in the California science fiction community, he was known as "Van". Method and themes Van Vogt systematized his writing method, using scenes of 800 words or so where a new complication was added or something resolved. Several of his stories hinge on temporal conundra, a favorite theme. He stated that he acquired many of his writing techniques from three books: Narrative Technique by Thomas Uzzell, The Only Two Ways to Write a Story by John Gallishaw, and Twenty Problems of the Fiction Writer by Gallishaw. He also claimed many of his ideas came from dreams; throughout his writing life he arranged to be awakened every 90 minutes during his sleep period so he could write down his dreams. Van Vogt was also always interested in the idea of all-encompassing systems of knowledge (akin to modern meta-systems). The characters in his very first story used a system called "Nexialism" to analyze the alien's behavior. Around this time, he became particularly interested in the general semantics of Alfred Korzybski. He subsequently wrote a novel merging these overarching themes, The World of Ā, originally serialized in Astounding in 1945. Ā (often rendered as Null-A), or non-Aristotelian logic, refers to the capacity for, and practice of, using intuitive, inductive reasoning (compare fuzzy logic), rather than reflexive, or conditioned, deductive reasoning. The novel recounts the adventures of an individual living in an apparent Utopia, where those with superior brainpower make up the ruling class... though all is not as it seems. A sequel, The Players of Ā (later re-titled The Pawns of Null-A) was serialized in 1948–49. At the same time, in his fiction, van Vogt was consistently sympathetic to absolute monarchy as a form of government. This was the case, for instance, in the Weapon Shop series, the Mixed Men series, and in single stories such as "Heir Apparent" (1945), whose protagonist was described as a "benevolent dictator". These sympathies were the subject of much critical discussion during van Vogt's career, and afterwards. Van Vogt published "Enchanted Village" in the July 1950 issue of Other Worlds Science Stories. It was reprinted in over 20 collections or anthologies, and appeared many times in translation. Dianetics and fix-ups (1950–1961) In 1950, van Vogt was briefly appointed as head of L. Ron Hubbard's Dianetics operation in California. Van Vogt had first met Hubbard in 1945, and became interested in his Dianetics theories, which were published shortly thereafter. Dianetics was the secular precursor to Hubbard's Church of Scientology; van Vogt would have no association with Scientology, as he did not approve of its mysticism.

California and post-war writing (1944–1950) In November 1944, van Vogt and Hull moved to Hollywood; van Vogt would spend the rest of his life in California. He had been using the name "A. E. van Vogt" in his public life for several years, and as part of the process of obtaining American citizenship in 1945 he finally and formally changed his legal name from Alfred Vogt to Alfred Elton van Vogt. To his friends in the California science fiction community, he was known as "Van". Method and themes Van Vogt systematized his writing method, using scenes of 800 words or so where a new complication was added or something resolved. Several of his stories hinge on temporal conundra, a favorite theme. He stated that he acquired many of his writing techniques from three books: Narrative Technique by Thomas Uzzell, The Only Two Ways to Write a Story by John Gallishaw, and Twenty Problems of the Fiction Writer by Gallishaw. He also claimed many of his ideas came from dreams; throughout his writing life he arranged to be awakened every 90 minutes during his sleep period so he could write down his dreams. Van Vogt was also always interested in the idea of all-encompassing systems of knowledge (akin to modern meta-systems). The characters in his very first story used a system called "Nexialism" to analyze the alien's behavior. Around this time, he became particularly interested in the general semantics of Alfred Korzybski. He subsequently wrote a novel merging these overarching themes, The World of Ā, originally serialized in Astounding in 1945. Ā (often rendered as Null-A), or non-Aristotelian logic, refers to the capacity for, and practice of, using intuitive, inductive reasoning (compare fuzzy logic), rather than reflexive, or conditioned, deductive reasoning. The novel recounts the adventures of an individual living in an apparent Utopia, where those with superior brainpower make up the ruling class... though all is not as it seems. A sequel, The Players of Ā (later re-titled The Pawns of Null-A) was serialized in 1948–49. At the same time, in his fiction, van Vogt was consistently sympathetic to absolute monarchy as a form of government. This was the case, for instance, in the Weapon Shop series, the Mixed Men series, and in single stories such as "Heir Apparent" (1945), whose protagonist was described as a "benevolent dictator". These sympathies were the subject of much critical discussion during van Vogt's career, and afterwards. Van Vogt published "Enchanted Village" in the July 1950 issue of Other Worlds Science Stories. It was reprinted in over 20 collections or anthologies, and appeared many times in translation. Dianetics and fix-ups (1950–1961) In 1950, van Vogt was briefly appointed as head of L. Ron Hubbard's Dianetics operation in California. Van Vogt had first met Hubbard in 1945, and became interested in his Dianetics theories, which were published shortly thereafter. Dianetics was the secular precursor to Hubbard's Church of Scientology; van Vogt would have no association with Scientology, as he did not approve of its mysticism.

California and post-war writing (1944–1950) In November 1944, van Vogt and Hull moved to Hollywood; van Vogt would spend the rest of his life in California. He had been using the name "A. E. van Vogt" in his public life for several years, and as part of the process of obtaining American citizenship in 1945 he finally and formally changed his legal name from Alfred Vogt to Alfred Elton van Vogt. To his friends in the California science fiction community, he was known as "Van". Method and themes Van Vogt systematized his writing method, using scenes of 800 words or so where a new complication was added or something resolved. Several of his stories hinge on temporal conundra, a favorite theme. He stated that he acquired many of his writing techniques from three books: Narrative Technique by Thomas Uzzell, The Only Two Ways to Write a Story by John Gallishaw, and Twenty Problems of the Fiction Writer by Gallishaw. He also claimed many of his ideas came from dreams; throughout his writing life he arranged to be awakened every 90 minutes during his sleep period so he could write down his dreams. Van Vogt was also always interested in the idea of all-encompassing systems of knowledge (akin to modern meta-systems). The characters in his very first story used a system called "Nexialism" to analyze the alien's behavior. Around this time, he became particularly interested in the general semantics of Alfred Korzybski. He subsequently wrote a novel merging these overarching themes, The World of Ā, originally serialized in Astounding in 1945. Ā (often rendered as Null-A), or non-Aristotelian logic, refers to the capacity for, and practice of, using intuitive, inductive reasoning (compare fuzzy logic), rather than reflexive, or conditioned, deductive reasoning. The novel recounts the adventures of an individual living in an apparent Utopia, where those with superior brainpower make up the ruling class... though all is not as it seems. A sequel, The Players of Ā (later re-titled The Pawns of Null-A) was serialized in 1948–49. At the same time, in his fiction, van Vogt was consistently sympathetic to absolute monarchy as a form of government. This was the case, for instance, in the Weapon Shop series, the Mixed Men series, and in single stories such as "Heir Apparent" (1945), whose protagonist was described as a "benevolent dictator". These sympathies were the subject of much critical discussion during van Vogt's career, and afterwards. Van Vogt published "Enchanted Village" in the July 1950 issue of Other Worlds Science Stories. It was reprinted in over 20 collections or anthologies, and appeared many times in translation. Dianetics and fix-ups (1950–1961) In 1950, van Vogt was briefly appointed as head of L. Ron Hubbard's Dianetics operation in California. Van Vogt had first met Hubbard in 1945, and became interested in his Dianetics theories, which were published shortly thereafter. Dianetics was the secular precursor to Hubbard's Church of Scientology; van Vogt would have no association with Scientology, as he did not approve of its mysticism.

The California Dianetics operation went broke nine months later, but never went bankrupt, due to van Vogt's arrangements with creditors. Very shortly after that, van Vogt and his wife opened their own Dianetics center, partly financed by his writings, until he "signed off" around 1961. From 1951 until 1961, van Vogt's focus was on Dianetics, and no new story ideas flowed from his typewriter. Fix-ups However, during the 1950s, van Vogt retrospectively patched together many of his previously published stories into novels, sometimes creating new interstitial material to help bridge gaps in the narrative. Van Vogt referred to the resulting books as "fix-ups", a term that entered the vocabulary of science-fiction criticism. When the original stories were closely related this was often successful, although some van Vogt fix-ups featured disparate stories thrown together that bore little relation to each other, generally making for a less coherent plot. One of his best-known (and well-regarded) novels, The Voyage of the Space Beagle (1950) was a fix-up of four short stories including "Discord in Scarlet"; it was published in at least five European languages by 1955. Although Van Vogt averaged a new book title every ten months from 1951 to 1961, none of them were new stories; they were all fix-ups, collections of previously published stories, expansions of previously published short stories to novel length, or republications of previous books under new titles and all based on story material written and originally published between 1939 and 1950. Examples include The Weapon Shops of Isher (1951), The Mixed Men (1952), The War Against the Rull (1959), and the two "Clane" novels, Empire of the Atom (1957) and The Wizard of Linn (1962), which were inspired (like Asimov's Foundation series) by Roman imperial history; specifically, as Damon Knight wrote, the plot of Empire of the Atom was "lifted almost bodily" from that of Robert Graves' I, Claudius. (Also, one non-fiction work, The Hypnotism Handbook, appeared in 1956, though it had apparently been written much earlier.) After more than a decade of running their Dianetics center, Hull and van Vogt closed it in 1961. Nevertheless, van Vogt maintained his association with the organization and was still president of the Californian Association of Dianetic Auditors into the 1980s. Return to writing and later career (1962–1986) Though the constant re-packaging of his older work meant that he had never really been away from the book publishing world, van Vogt had not published any wholly new fiction for almost 12 years when he decided to return to writing in 1962. He did not return immediately to science fiction, but instead wrote the only mainstream, non-sf novel of his career. Van Vogt was profoundly affected by revelations of totalitarian police states that emerged after World War II. Accordingly, he wrote a mainstream novel that he set in Communist China, The Violent Man (1962). Van Vogt explained that to research this book he had read 100 books about China.

The California Dianetics operation went broke nine months later, but never went bankrupt, due to van Vogt's arrangements with creditors. Very shortly after that, van Vogt and his wife opened their own Dianetics center, partly financed by his writings, until he "signed off" around 1961. From 1951 until 1961, van Vogt's focus was on Dianetics, and no new story ideas flowed from his typewriter. Fix-ups However, during the 1950s, van Vogt retrospectively patched together many of his previously published stories into novels, sometimes creating new interstitial material to help bridge gaps in the narrative. Van Vogt referred to the resulting books as "fix-ups", a term that entered the vocabulary of science-fiction criticism. When the original stories were closely related this was often successful, although some van Vogt fix-ups featured disparate stories thrown together that bore little relation to each other, generally making for a less coherent plot. One of his best-known (and well-regarded) novels, The Voyage of the Space Beagle (1950) was a fix-up of four short stories including "Discord in Scarlet"; it was published in at least five European languages by 1955. Although Van Vogt averaged a new book title every ten months from 1951 to 1961, none of them were new stories; they were all fix-ups, collections of previously published stories, expansions of previously published short stories to novel length, or republications of previous books under new titles and all based on story material written and originally published between 1939 and 1950. Examples include The Weapon Shops of Isher (1951), The Mixed Men (1952), The War Against the Rull (1959), and the two "Clane" novels, Empire of the Atom (1957) and The Wizard of Linn (1962), which were inspired (like Asimov's Foundation series) by Roman imperial history; specifically, as Damon Knight wrote, the plot of Empire of the Atom was "lifted almost bodily" from that of Robert Graves' I, Claudius. (Also, one non-fiction work, The Hypnotism Handbook, appeared in 1956, though it had apparently been written much earlier.) After more than a decade of running their Dianetics center, Hull and van Vogt closed it in 1961. Nevertheless, van Vogt maintained his association with the organization and was still president of the Californian Association of Dianetic Auditors into the 1980s. Return to writing and later career (1962–1986) Though the constant re-packaging of his older work meant that he had never really been away from the book publishing world, van Vogt had not published any wholly new fiction for almost 12 years when he decided to return to writing in 1962. He did not return immediately to science fiction, but instead wrote the only mainstream, non-sf novel of his career. Van Vogt was profoundly affected by revelations of totalitarian police states that emerged after World War II. Accordingly, he wrote a mainstream novel that he set in Communist China, The Violent Man (1962). Van Vogt explained that to research this book he had read 100 books about China.

The California Dianetics operation went broke nine months later, but never went bankrupt, due to van Vogt's arrangements with creditors. Very shortly after that, van Vogt and his wife opened their own Dianetics center, partly financed by his writings, until he "signed off" around 1961. From 1951 until 1961, van Vogt's focus was on Dianetics, and no new story ideas flowed from his typewriter. Fix-ups However, during the 1950s, van Vogt retrospectively patched together many of his previously published stories into novels, sometimes creating new interstitial material to help bridge gaps in the narrative. Van Vogt referred to the resulting books as "fix-ups", a term that entered the vocabulary of science-fiction criticism. When the original stories were closely related this was often successful, although some van Vogt fix-ups featured disparate stories thrown together that bore little relation to each other, generally making for a less coherent plot. One of his best-known (and well-regarded) novels, The Voyage of the Space Beagle (1950) was a fix-up of four short stories including "Discord in Scarlet"; it was published in at least five European languages by 1955. Although Van Vogt averaged a new book title every ten months from 1951 to 1961, none of them were new stories; they were all fix-ups, collections of previously published stories, expansions of previously published short stories to novel length, or republications of previous books under new titles and all based on story material written and originally published between 1939 and 1950. Examples include The Weapon Shops of Isher (1951), The Mixed Men (1952), The War Against the Rull (1959), and the two "Clane" novels, Empire of the Atom (1957) and The Wizard of Linn (1962), which were inspired (like Asimov's Foundation series) by Roman imperial history; specifically, as Damon Knight wrote, the plot of Empire of the Atom was "lifted almost bodily" from that of Robert Graves' I, Claudius. (Also, one non-fiction work, The Hypnotism Handbook, appeared in 1956, though it had apparently been written much earlier.) After more than a decade of running their Dianetics center, Hull and van Vogt closed it in 1961. Nevertheless, van Vogt maintained his association with the organization and was still president of the Californian Association of Dianetic Auditors into the 1980s. Return to writing and later career (1962–1986) Though the constant re-packaging of his older work meant that he had never really been away from the book publishing world, van Vogt had not published any wholly new fiction for almost 12 years when he decided to return to writing in 1962. He did not return immediately to science fiction, but instead wrote the only mainstream, non-sf novel of his career. Van Vogt was profoundly affected by revelations of totalitarian police states that emerged after World War II. Accordingly, he wrote a mainstream novel that he set in Communist China, The Violent Man (1962). Van Vogt explained that to research this book he had read 100 books about China.

Into this book he incorporated his view of "the violent male type", which he described as a "man who had to be right", a man who "instantly attracts women" and who he said were the men who "run the world". Contemporary reviews were lukewarm at best, and van Vogt thereafter returned to science fiction. From 1963 through the mid-1980s, van Vogt once again published new material on a regular basis, though fix-ups and reworked material also appeared relatively often. His later novels included fix-ups such as The Beast (also known as Moonbeast) (1963), Rogue Ship (1965), Quest for the Future (1970) and Supermind (1977). He also wrote novels by expanding previously published short stories; works of this type include The Darkness on Diamondia (1972) and Future Glitter (also known as Tyranopolis; 1973). Novels that were written simply as novels, and not serialized magazine pieces or fix-ups, had been very rare in van Vogt's oeuvre, but began to appear regularly beginning in the 1970s. Van Vogt's original novels included Children of Tomorrow (1970), The Battle of Forever (1971) and The Anarchistic Colossus (1977). Over the years, many sequels to his classic works were promised, but only one appeared: Null-A Three (1984; originally published in French). Several later books were initially published in Europe, and at least one novel only ever appeared in foreign language editions and was never published in its original English. Final years When the 1979 film Alien appeared, it was noted that the plot closely matched the plots of both Black Destroyer and Discord in Scarlet, both published in Astounding magazine in 1939, and then later published in the 1950 book Voyage of the Space Beagle. Van Vogt sued the production company for plagiarism, and eventually collected an out-of-court settlement of $50,000 from 20th Century Fox. In increasingly frail health, van Vogt published his final short story in 1986. Personal life Van Vogt's first wife, Edna Mayne Hull, died in 1975. Van Vogt married Lydia Bereginsky in 1979; they remained together until his death. Death On January 26, 2000, A. E. van Vogt died in Los Angeles from Alzheimer's disease. He was survived by his second wife. Critical reception Critical opinion about the quality of van Vogt's work is sharply divided. An early and articulate critic was Damon Knight. In a 1945 chapter-long essay reprinted in In Search of Wonder, entitled "Cosmic Jerrybuilder: A. E. van Vogt", Knight described van Vogt as "no giant; he is a pygmy who has learned to operate an overgrown typewriter". Knight described The World of Null-A as "one of the worst allegedly adult science fiction stories ever published". Concerning van Vogt's writing, Knight said: About Empire of the Atom Knight wrote: Knight also expressed misgivings about van Vogt's politics. He noted that van Vogt's stories almost invariably present absolute monarchy in a favorable light.

Into this book he incorporated his view of "the violent male type", which he described as a "man who had to be right", a man who "instantly attracts women" and who he said were the men who "run the world". Contemporary reviews were lukewarm at best, and van Vogt thereafter returned to science fiction. From 1963 through the mid-1980s, van Vogt once again published new material on a regular basis, though fix-ups and reworked material also appeared relatively often. His later novels included fix-ups such as The Beast (also known as Moonbeast) (1963), Rogue Ship (1965), Quest for the Future (1970) and Supermind (1977). He also wrote novels by expanding previously published short stories; works of this type include The Darkness on Diamondia (1972) and Future Glitter (also known as Tyranopolis; 1973). Novels that were written simply as novels, and not serialized magazine pieces or fix-ups, had been very rare in van Vogt's oeuvre, but began to appear regularly beginning in the 1970s. Van Vogt's original novels included Children of Tomorrow (1970), The Battle of Forever (1971) and The Anarchistic Colossus (1977). Over the years, many sequels to his classic works were promised, but only one appeared: Null-A Three (1984; originally published in French). Several later books were initially published in Europe, and at least one novel only ever appeared in foreign language editions and was never published in its original English. Final years When the 1979 film Alien appeared, it was noted that the plot closely matched the plots of both Black Destroyer and Discord in Scarlet, both published in Astounding magazine in 1939, and then later published in the 1950 book Voyage of the Space Beagle. Van Vogt sued the production company for plagiarism, and eventually collected an out-of-court settlement of $50,000 from 20th Century Fox. In increasingly frail health, van Vogt published his final short story in 1986. Personal life Van Vogt's first wife, Edna Mayne Hull, died in 1975. Van Vogt married Lydia Bereginsky in 1979; they remained together until his death. Death On January 26, 2000, A. E. van Vogt died in Los Angeles from Alzheimer's disease. He was survived by his second wife. Critical reception Critical opinion about the quality of van Vogt's work is sharply divided. An early and articulate critic was Damon Knight. In a 1945 chapter-long essay reprinted in In Search of Wonder, entitled "Cosmic Jerrybuilder: A. E. van Vogt", Knight described van Vogt as "no giant; he is a pygmy who has learned to operate an overgrown typewriter". Knight described The World of Null-A as "one of the worst allegedly adult science fiction stories ever published". Concerning van Vogt's writing, Knight said: About Empire of the Atom Knight wrote: Knight also expressed misgivings about van Vogt's politics. He noted that van Vogt's stories almost invariably present absolute monarchy in a favorable light.

Into this book he incorporated his view of "the violent male type", which he described as a "man who had to be right", a man who "instantly attracts women" and who he said were the men who "run the world". Contemporary reviews were lukewarm at best, and van Vogt thereafter returned to science fiction. From 1963 through the mid-1980s, van Vogt once again published new material on a regular basis, though fix-ups and reworked material also appeared relatively often. His later novels included fix-ups such as The Beast (also known as Moonbeast) (1963), Rogue Ship (1965), Quest for the Future (1970) and Supermind (1977). He also wrote novels by expanding previously published short stories; works of this type include The Darkness on Diamondia (1972) and Future Glitter (also known as Tyranopolis; 1973). Novels that were written simply as novels, and not serialized magazine pieces or fix-ups, had been very rare in van Vogt's oeuvre, but began to appear regularly beginning in the 1970s. Van Vogt's original novels included Children of Tomorrow (1970), The Battle of Forever (1971) and The Anarchistic Colossus (1977). Over the years, many sequels to his classic works were promised, but only one appeared: Null-A Three (1984; originally published in French). Several later books were initially published in Europe, and at least one novel only ever appeared in foreign language editions and was never published in its original English. Final years When the 1979 film Alien appeared, it was noted that the plot closely matched the plots of both Black Destroyer and Discord in Scarlet, both published in Astounding magazine in 1939, and then later published in the 1950 book Voyage of the Space Beagle. Van Vogt sued the production company for plagiarism, and eventually collected an out-of-court settlement of $50,000 from 20th Century Fox. In increasingly frail health, van Vogt published his final short story in 1986. Personal life Van Vogt's first wife, Edna Mayne Hull, died in 1975. Van Vogt married Lydia Bereginsky in 1979; they remained together until his death. Death On January 26, 2000, A. E. van Vogt died in Los Angeles from Alzheimer's disease. He was survived by his second wife. Critical reception Critical opinion about the quality of van Vogt's work is sharply divided. An early and articulate critic was Damon Knight. In a 1945 chapter-long essay reprinted in In Search of Wonder, entitled "Cosmic Jerrybuilder: A. E. van Vogt", Knight described van Vogt as "no giant; he is a pygmy who has learned to operate an overgrown typewriter". Knight described The World of Null-A as "one of the worst allegedly adult science fiction stories ever published". Concerning van Vogt's writing, Knight said: About Empire of the Atom Knight wrote: Knight also expressed misgivings about van Vogt's politics. He noted that van Vogt's stories almost invariably present absolute monarchy in a favorable light.

In 1974, Knight retracted some of his criticism after finding out about Vogt's writing down his dreams as a part of his working methods: Knight's criticism greatly damaged van Vogt's reputation. On the other hand, when science fiction author Philip K. Dick was asked which science fiction writers had influenced his work the most, he replied: Dick also defended van Vogt against Damon Knight's criticisms: In a review of Transfinite: The Essential A. E. van Vogt, science fiction writer Paul Di Filippo said: In The John W. Campbell Letters, Campbell says, "The son-of-a-gun gets hold of you in the first paragraph, ties a knot around you, and keeps it tied in every paragraph thereafter—including the ultimate last one". Harlan Ellison (who had begun reading van Vogt as a teenager) wrote, "Van was the first writer to shine light on the restricted ways in which I had been taught to view the universe and the human condition". Writing in 1984, David Hartwell said: The literary critic Leslie A. Fiedler said something similar: American literary critic Fredric Jameson says of van Vogt: Van Vogt still has his critics. For example, Darrell Schweitzer, writing to The New York Review of Science Fiction in 1999, quoted a passage from the original van Vogt novelette "The Mixed Men", which he was then reading, and remarked: Recognition In 1946, van Vogt and his first wife, Edna Mayne Hull, were Guests of Honor at the fourth World Science Fiction Convention. In 1980, van Vogt received a "Casper Award" (precursor to the Canadian Prix Aurora Awards) for Lifetime Achievement. In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, he was inducted as an inaugural member of the Science Fiction and Fantasy Hall of Fame. The Science Fiction Writers of America (SFWA) named him its 14th Grand Master in 1995 (presented 1996). Great controversy within SFWA accompanied its long wait in bestowing its highest honor (limited to living writers, no more than one annually). Writing an obituary of van Vogt, Robert J. Sawyer, a fellow Canadian writer of science fiction, remarked: It is generally held that a key factor in the delay was "damnable SFWA politics" reflecting the concerns of Damon Knight, the founder of the SFWA, who abhorred van Vogt's style and politics and thoroughly demolished his literary reputation in the 1950s. Harlan Ellison was more explicit in 1999 introduction to Futures Past: The Best Short Fiction of A. E. van Vogt: In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, the Science Fiction and Fantasy Hall of Fame inducted him in its inaugural class of two deceased and two living persons, along with writer Jack Williamson (also living) and editors Hugo Gernsback and John W. Campbell.

In 1974, Knight retracted some of his criticism after finding out about Vogt's writing down his dreams as a part of his working methods: Knight's criticism greatly damaged van Vogt's reputation. On the other hand, when science fiction author Philip K. Dick was asked which science fiction writers had influenced his work the most, he replied: Dick also defended van Vogt against Damon Knight's criticisms: In a review of Transfinite: The Essential A. E. van Vogt, science fiction writer Paul Di Filippo said: In The John W. Campbell Letters, Campbell says, "The son-of-a-gun gets hold of you in the first paragraph, ties a knot around you, and keeps it tied in every paragraph thereafter—including the ultimate last one". Harlan Ellison (who had begun reading van Vogt as a teenager) wrote, "Van was the first writer to shine light on the restricted ways in which I had been taught to view the universe and the human condition". Writing in 1984, David Hartwell said: The literary critic Leslie A. Fiedler said something similar: American literary critic Fredric Jameson says of van Vogt: Van Vogt still has his critics. For example, Darrell Schweitzer, writing to The New York Review of Science Fiction in 1999, quoted a passage from the original van Vogt novelette "The Mixed Men", which he was then reading, and remarked: Recognition In 1946, van Vogt and his first wife, Edna Mayne Hull, were Guests of Honor at the fourth World Science Fiction Convention. In 1980, van Vogt received a "Casper Award" (precursor to the Canadian Prix Aurora Awards) for Lifetime Achievement. In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, he was inducted as an inaugural member of the Science Fiction and Fantasy Hall of Fame. The Science Fiction Writers of America (SFWA) named him its 14th Grand Master in 1995 (presented 1996). Great controversy within SFWA accompanied its long wait in bestowing its highest honor (limited to living writers, no more than one annually). Writing an obituary of van Vogt, Robert J. Sawyer, a fellow Canadian writer of science fiction, remarked: It is generally held that a key factor in the delay was "damnable SFWA politics" reflecting the concerns of Damon Knight, the founder of the SFWA, who abhorred van Vogt's style and politics and thoroughly demolished his literary reputation in the 1950s. Harlan Ellison was more explicit in 1999 introduction to Futures Past: The Best Short Fiction of A. E. van Vogt: In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, the Science Fiction and Fantasy Hall of Fame inducted him in its inaugural class of two deceased and two living persons, along with writer Jack Williamson (also living) and editors Hugo Gernsback and John W. Campbell.

In 1974, Knight retracted some of his criticism after finding out about Vogt's writing down his dreams as a part of his working methods: Knight's criticism greatly damaged van Vogt's reputation. On the other hand, when science fiction author Philip K. Dick was asked which science fiction writers had influenced his work the most, he replied: Dick also defended van Vogt against Damon Knight's criticisms: In a review of Transfinite: The Essential A. E. van Vogt, science fiction writer Paul Di Filippo said: In The John W. Campbell Letters, Campbell says, "The son-of-a-gun gets hold of you in the first paragraph, ties a knot around you, and keeps it tied in every paragraph thereafter—including the ultimate last one". Harlan Ellison (who had begun reading van Vogt as a teenager) wrote, "Van was the first writer to shine light on the restricted ways in which I had been taught to view the universe and the human condition". Writing in 1984, David Hartwell said: The literary critic Leslie A. Fiedler said something similar: American literary critic Fredric Jameson says of van Vogt: Van Vogt still has his critics. For example, Darrell Schweitzer, writing to The New York Review of Science Fiction in 1999, quoted a passage from the original van Vogt novelette "The Mixed Men", which he was then reading, and remarked: Recognition In 1946, van Vogt and his first wife, Edna Mayne Hull, were Guests of Honor at the fourth World Science Fiction Convention. In 1980, van Vogt received a "Casper Award" (precursor to the Canadian Prix Aurora Awards) for Lifetime Achievement. In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, he was inducted as an inaugural member of the Science Fiction and Fantasy Hall of Fame. The Science Fiction Writers of America (SFWA) named him its 14th Grand Master in 1995 (presented 1996). Great controversy within SFWA accompanied its long wait in bestowing its highest honor (limited to living writers, no more than one annually). Writing an obituary of van Vogt, Robert J. Sawyer, a fellow Canadian writer of science fiction, remarked: It is generally held that a key factor in the delay was "damnable SFWA politics" reflecting the concerns of Damon Knight, the founder of the SFWA, who abhorred van Vogt's style and politics and thoroughly demolished his literary reputation in the 1950s. Harlan Ellison was more explicit in 1999 introduction to Futures Past: The Best Short Fiction of A. E. van Vogt: In 1996, van Vogt received a Special Award from the World Science Fiction Convention "for six decades of golden age science fiction". That same year, the Science Fiction and Fantasy Hall of Fame inducted him in its inaugural class of two deceased and two living persons, along with writer Jack Williamson (also living) and editors Hugo Gernsback and John W. Campbell.

The works of van Vogt were translated into French by the surrealist Boris Vian (The World of Null-A as Le Monde des Å in 1958), and van Vogt's works were "viewed as great literature of the surrealist school". In addition, Slan was published in French, translated by Jean Rosenthal, under the title À la poursuite des Slans, as part of the paperback series 'Editions J'ai Lu: Romans-Texte Integral' in 1973. This edition also listing the following works by van Vogt as having been published in French as part of this series: Le Monde des Å, La faune de l'espace, Les joueurs du Å, L'empire de l'atome, Le sorcier de Linn, Les armureries d'Isher, Les fabricants d'armes, and Le livre de Ptath. Works Novels and novellas Special works published as books Planets For Sale by E. Mayne Hull (1954). A fix-up of five stories by Hull, originally published 1942 to 1946. Certain later editions (from 1965) credit both authors. The Enchanted Village (1979). A 25-page chapbook of a short story originally published in 1950. Slan Hunter by Kevin J. Anderson (2007). A sequel to Slan, based an unfinished draft by van Vogt. Null-A Continuum by John C. Wright (2008). An authorized continuation of the Null-A series which ignored the events of Null-A Three. Collections Out of the Unknown (1948), with Edna Mayne Hull Masters of Time (1950) (a.k.a. Recruiting Station) [also includes The Changeling, both works were later published separately] Triad (1951) omnibus of The World of Null A, The Voyage of the Space Beagle, Slan. Away and Beyond (1952) (abridged in paperback in 1959; abridged (differently) in paperback in 1963) Destination: Universe! (1952) The Twisted Men (1964) Monsters (1965) (later as SF Monsters (1967)) abridged as The Blal (1976) A Van Vogt Omnibus (1967), omnibus of Planets for Sale (with Edna Mayne Hull), The Beast, The Book of Ptath The Far Out Worlds of Van Vogt (1968) The Sea Thing and Other Stories (1970) (expanded from Out of the Unknown by adding an original story by Hull; later abridged in paperback as Out of the Unknown by removing 2 of the stories) M33 in Andromeda (1971) More Than Superhuman (1971) The Proxy Intelligence and Other Mind Benders, ), with Edna Mayne Hull (1971), revised as The Gryb (1976) Van Vogt Omnibus 2 (1971), omnibus of The Mind Cage, The Winged Man (with Edna Mayne Hull), Slan. The Book of Van Vogt (1972), also published as Lost: Fifty Suns (1979) The Three Eyes of Evil Including Earth's Last Fortress (1973) The Best of A. E. van Vogt (1974) later split into 2 volumes The Worlds of A. E. van Vogt (1974) (expanded from The Far Out Worlds of Van Vogt by adding 3 stories) The Best of A. E. van Vogt (1976) [differs to 1974 edition] Away and Beyond (1977) Pendulum (1978) (almost all original stories and articles) Futures Past: The Best Short Fiction of A.E. Van Vogt (1999) Transfinite: The Essential A.E.

The works of van Vogt were translated into French by the surrealist Boris Vian (The World of Null-A as Le Monde des Å in 1958), and van Vogt's works were "viewed as great literature of the surrealist school". In addition, Slan was published in French, translated by Jean Rosenthal, under the title À la poursuite des Slans, as part of the paperback series 'Editions J'ai Lu: Romans-Texte Integral' in 1973. This edition also listing the following works by van Vogt as having been published in French as part of this series: Le Monde des Å, La faune de l'espace, Les joueurs du Å, L'empire de l'atome, Le sorcier de Linn, Les armureries d'Isher, Les fabricants d'armes, and Le livre de Ptath. Works Novels and novellas Special works published as books Planets For Sale by E. Mayne Hull (1954). A fix-up of five stories by Hull, originally published 1942 to 1946. Certain later editions (from 1965) credit both authors. The Enchanted Village (1979). A 25-page chapbook of a short story originally published in 1950. Slan Hunter by Kevin J. Anderson (2007). A sequel to Slan, based an unfinished draft by van Vogt. Null-A Continuum by John C. Wright (2008). An authorized continuation of the Null-A series which ignored the events of Null-A Three. Collections Out of the Unknown (1948), with Edna Mayne Hull Masters of Time (1950) (a.k.a. Recruiting Station) [also includes The Changeling, both works were later published separately] Triad (1951) omnibus of The World of Null A, The Voyage of the Space Beagle, Slan. Away and Beyond (1952) (abridged in paperback in 1959; abridged (differently) in paperback in 1963) Destination: Universe! (1952) The Twisted Men (1964) Monsters (1965) (later as SF Monsters (1967)) abridged as The Blal (1976) A Van Vogt Omnibus (1967), omnibus of Planets for Sale (with Edna Mayne Hull), The Beast, The Book of Ptath The Far Out Worlds of Van Vogt (1968) The Sea Thing and Other Stories (1970) (expanded from Out of the Unknown by adding an original story by Hull; later abridged in paperback as Out of the Unknown by removing 2 of the stories) M33 in Andromeda (1971) More Than Superhuman (1971) The Proxy Intelligence and Other Mind Benders, ), with Edna Mayne Hull (1971), revised as The Gryb (1976) Van Vogt Omnibus 2 (1971), omnibus of The Mind Cage, The Winged Man (with Edna Mayne Hull), Slan. The Book of Van Vogt (1972), also published as Lost: Fifty Suns (1979) The Three Eyes of Evil Including Earth's Last Fortress (1973) The Best of A. E. van Vogt (1974) later split into 2 volumes The Worlds of A. E. van Vogt (1974) (expanded from The Far Out Worlds of Van Vogt by adding 3 stories) The Best of A. E. van Vogt (1976) [differs to 1974 edition] Away and Beyond (1977) Pendulum (1978) (almost all original stories and articles) Futures Past: The Best Short Fiction of A.E. Van Vogt (1999) Transfinite: The Essential A.E.

The works of van Vogt were translated into French by the surrealist Boris Vian (The World of Null-A as Le Monde des Å in 1958), and van Vogt's works were "viewed as great literature of the surrealist school". In addition, Slan was published in French, translated by Jean Rosenthal, under the title À la poursuite des Slans, as part of the paperback series 'Editions J'ai Lu: Romans-Texte Integral' in 1973. This edition also listing the following works by van Vogt as having been published in French as part of this series: Le Monde des Å, La faune de l'espace, Les joueurs du Å, L'empire de l'atome, Le sorcier de Linn, Les armureries d'Isher, Les fabricants d'armes, and Le livre de Ptath. Works Novels and novellas Special works published as books Planets For Sale by E. Mayne Hull (1954). A fix-up of five stories by Hull, originally published 1942 to 1946. Certain later editions (from 1965) credit both authors. The Enchanted Village (1979). A 25-page chapbook of a short story originally published in 1950. Slan Hunter by Kevin J. Anderson (2007). A sequel to Slan, based an unfinished draft by van Vogt. Null-A Continuum by John C. Wright (2008). An authorized continuation of the Null-A series which ignored the events of Null-A Three. Collections Out of the Unknown (1948), with Edna Mayne Hull Masters of Time (1950) (a.k.a. Recruiting Station) [also includes The Changeling, both works were later published separately] Triad (1951) omnibus of The World of Null A, The Voyage of the Space Beagle, Slan. Away and Beyond (1952) (abridged in paperback in 1959; abridged (differently) in paperback in 1963) Destination: Universe! (1952) The Twisted Men (1964) Monsters (1965) (later as SF Monsters (1967)) abridged as The Blal (1976) A Van Vogt Omnibus (1967), omnibus of Planets for Sale (with Edna Mayne Hull), The Beast, The Book of Ptath The Far Out Worlds of Van Vogt (1968) The Sea Thing and Other Stories (1970) (expanded from Out of the Unknown by adding an original story by Hull; later abridged in paperback as Out of the Unknown by removing 2 of the stories) M33 in Andromeda (1971) More Than Superhuman (1971) The Proxy Intelligence and Other Mind Benders, ), with Edna Mayne Hull (1971), revised as The Gryb (1976) Van Vogt Omnibus 2 (1971), omnibus of The Mind Cage, The Winged Man (with Edna Mayne Hull), Slan. The Book of Van Vogt (1972), also published as Lost: Fifty Suns (1979) The Three Eyes of Evil Including Earth's Last Fortress (1973) The Best of A. E. van Vogt (1974) later split into 2 volumes The Worlds of A. E. van Vogt (1974) (expanded from The Far Out Worlds of Van Vogt by adding 3 stories) The Best of A. E. van Vogt (1976) [differs to 1974 edition] Away and Beyond (1977) Pendulum (1978) (almost all original stories and articles) Futures Past: The Best Short Fiction of A.E. Van Vogt (1999) Transfinite: The Essential A.E.

van Vogt (2002) Transgalactic (2006) Nonfiction The Hypnotism Handbook (1956, Griffin Publishing Company, with Charles Edward Cooke) The Money Personality (1972, Parker Publishing Company Inc., West Nyack, NY, ) Reflections of A. E. Van Vogt: The Autobiography of a Science Fiction Giant (1979, Fictioneer Books Ltd., Lakemont, GA) A Report on the Violent Male (1992, Paupers' Press, UK, ) See also Notes References Bibliography External links Sevagram, the A.E. van Vogt information site Obituary at LocusOnline'' (Locus Publications) "Writers: A. E. van Vogt (1912–2000, Canada)" – bibliography at SciFan A. E. van Vogt Papers (MS 322) at the Kenneth Spencer Research Library, University of Kansas A. E. van Vogt's fiction at Free Speculative Fiction Online 1912 births 2000 deaths 20th-century American novelists American male novelists American science fiction writers Canadian male novelists Canadian science fiction writers Canadian male short story writers Canadian emigrants to the United States Neurological disease deaths in California Deaths from Alzheimer's disease SFWA Grand Masters Science Fiction Hall of Fame inductees Writers from Manitoba Mennonite writers Canadian Mennonites American male short story writers 20th-century Canadian short story writers 20th-century American short story writers 20th-century Canadian male writers Weird fiction writers Pulp fiction writers Writers from Winnipeg 20th-century American male writers

van Vogt (2002) Transgalactic (2006) Nonfiction The Hypnotism Handbook (1956, Griffin Publishing Company, with Charles Edward Cooke) The Money Personality (1972, Parker Publishing Company Inc., West Nyack, NY, ) Reflections of A. E. Van Vogt: The Autobiography of a Science Fiction Giant (1979, Fictioneer Books Ltd., Lakemont, GA) A Report on the Violent Male (1992, Paupers' Press, UK, ) See also Notes References Bibliography External links Sevagram, the A.E. van Vogt information site Obituary at LocusOnline'' (Locus Publications) "Writers: A. E. van Vogt (1912–2000, Canada)" – bibliography at SciFan A. E. van Vogt Papers (MS 322) at the Kenneth Spencer Research Library, University of Kansas A. E. van Vogt's fiction at Free Speculative Fiction Online 1912 births 2000 deaths 20th-century American novelists American male novelists American science fiction writers Canadian male novelists Canadian science fiction writers Canadian male short story writers Canadian emigrants to the United States Neurological disease deaths in California Deaths from Alzheimer's disease SFWA Grand Masters Science Fiction Hall of Fame inductees Writers from Manitoba Mennonite writers Canadian Mennonites American male short story writers 20th-century Canadian short story writers 20th-century American short story writers 20th-century Canadian male writers Weird fiction writers Pulp fiction writers Writers from Winnipeg 20th-century American male writers

van Vogt (2002) Transgalactic (2006) Nonfiction The Hypnotism Handbook (1956, Griffin Publishing Company, with Charles Edward Cooke) The Money Personality (1972, Parker Publishing Company Inc., West Nyack, NY, ) Reflections of A. E. Van Vogt: The Autobiography of a Science Fiction Giant (1979, Fictioneer Books Ltd., Lakemont, GA) A Report on the Violent Male (1992, Paupers' Press, UK, ) See also Notes References Bibliography External links Sevagram, the A.E. van Vogt information site Obituary at LocusOnline'' (Locus Publications) "Writers: A. E. van Vogt (1912–2000, Canada)" – bibliography at SciFan A. E. van Vogt Papers (MS 322) at the Kenneth Spencer Research Library, University of Kansas A. E. van Vogt's fiction at Free Speculative Fiction Online 1912 births 2000 deaths 20th-century American novelists American male novelists American science fiction writers Canadian male novelists Canadian science fiction writers Canadian male short story writers Canadian emigrants to the United States Neurological disease deaths in California Deaths from Alzheimer's disease SFWA Grand Masters Science Fiction Hall of Fame inductees Writers from Manitoba Mennonite writers Canadian Mennonites American male short story writers 20th-century Canadian short story writers 20th-century American short story writers 20th-century Canadian male writers Weird fiction writers Pulp fiction writers Writers from Winnipeg 20th-century American male writers

Anna Kournikova Anna Sergeyevna Kournikova (; born 7 June 1981) is a Russian former professional tennis player and American television personality. Her appearance and celebrity status made her one of the best known tennis stars worldwide. At the peak of her fame, fans looking for images of Kournikova made her name one of the most common search strings on Google Search. Despite never winning a singles title, she reached No. 8 in the world in 2000. She achieved greater success playing doubles, where she was at times the world No. 1 player. With Martina Hingis as her partner, she won Grand Slam titles in Australia in 1999 and 2002, and the WTA Championships in 1999 and 2000. They referred to themselves as the "Spice Girls of Tennis". Kournikova retired from professional tennis in 2003 due to serious back and spinal problems, including a herniated disk. She lives in Miami Beach, Florida, and played in occasional exhibitions and in doubles for the St. Louis Aces of World Team Tennis before the team folded in 2011. She was a new trainer for season 12 of the television show The Biggest Loser, replacing Jillian Michaels, but did not return for season 13. In addition to her tennis and television work, Kournikova serves as a Global Ambassador for Population Services International's "Five & Alive" program, which addresses health crises facing children under the age of five and their families. Early life Kournikova was born in Moscow, Russia on 7 June 1981. Her father, Sergei Kournikov (born 1961), a former Greco-Roman wrestling champion, eventually earned a PhD and was a professor at the University of Physical Culture and Sport in Moscow. As of 2001, he was still a part-time martial arts instructor there. Her mother Alla (born 1963) had been a 400-metre runner. Her younger half-brother, Allan, is a youth golf world champion who was featured in the 2013 documentary film The Short Game. Sergei Kournikov has said, "We were young and we liked the clean, physical life, so Anna was in a good environment for sport from the beginning". Kournikova received her first tennis racquet as a New Year gift in 1986 at the age of five. Describing her early regimen, she said, "I played two times a week from age six. It was a children's program. And it was just for fun; my parents didn't know I was going to play professionally, they just wanted me to do something because I had lots of energy. It was only when I started playing well at seven that I went to a professional academy. I would go to school, and then my parents would take me to the club, and I'd spend the rest of the day there just having fun with the kids." In 1986, Kournikova became a member of the Spartak Tennis Club, coached by Larissa Preobrazhenskaya.

Anna Kournikova Anna Sergeyevna Kournikova (; born 7 June 1981) is a Russian former professional tennis player and American television personality. Her appearance and celebrity status made her one of the best known tennis stars worldwide. At the peak of her fame, fans looking for images of Kournikova made her name one of the most common search strings on Google Search. Despite never winning a singles title, she reached No. 8 in the world in 2000. She achieved greater success playing doubles, where she was at times the world No. 1 player. With Martina Hingis as her partner, she won Grand Slam titles in Australia in 1999 and 2002, and the WTA Championships in 1999 and 2000. They referred to themselves as the "Spice Girls of Tennis". Kournikova retired from professional tennis in 2003 due to serious back and spinal problems, including a herniated disk. She lives in Miami Beach, Florida, and played in occasional exhibitions and in doubles for the St. Louis Aces of World Team Tennis before the team folded in 2011. She was a new trainer for season 12 of the television show The Biggest Loser, replacing Jillian Michaels, but did not return for season 13. In addition to her tennis and television work, Kournikova serves as a Global Ambassador for Population Services International's "Five & Alive" program, which addresses health crises facing children under the age of five and their families. Early life Kournikova was born in Moscow, Russia on 7 June 1981. Her father, Sergei Kournikov (born 1961), a former Greco-Roman wrestling champion, eventually earned a PhD and was a professor at the University of Physical Culture and Sport in Moscow. As of 2001, he was still a part-time martial arts instructor there. Her mother Alla (born 1963) had been a 400-metre runner. Her younger half-brother, Allan, is a youth golf world champion who was featured in the 2013 documentary film The Short Game. Sergei Kournikov has said, "We were young and we liked the clean, physical life, so Anna was in a good environment for sport from the beginning". Kournikova received her first tennis racquet as a New Year gift in 1986 at the age of five. Describing her early regimen, she said, "I played two times a week from age six. It was a children's program. And it was just for fun; my parents didn't know I was going to play professionally, they just wanted me to do something because I had lots of energy. It was only when I started playing well at seven that I went to a professional academy. I would go to school, and then my parents would take me to the club, and I'd spend the rest of the day there just having fun with the kids." In 1986, Kournikova became a member of the Spartak Tennis Club, coached by Larissa Preobrazhenskaya.

In 1989, at the age of eight, Kournikova began appearing in junior tournaments, and by the following year, was attracting attention from tennis scouts across the world. She signed a management deal at age ten and went to Bradenton, Florida, to train at Nick Bollettieri's celebrated tennis academy. Tennis career 1989–1997: Early years and breakthrough Following her arrival in the United States, she became prominent on the tennis scene. At the age of 14, she won the European Championships and the Italian Open Junior tournament. In December 1995, she became the youngest player to win the 18-and-under division of the Junior Orange Bowl tennis tournament. By the end of the year, Kournikova was crowned the ITF Junior World Champion U-18 and Junior European Champion U-18. Earlier, in September 1995, Kournikova, still only 14 years of age, debuted in the WTA Tour, when she received a wildcard into the qualifications at the WTA tournament in Moscow, the Moscow Ladies Open, and qualified before losing in the second round of the main draw to third-seeded Sabine Appelmans. She also reached her first WTA Tour doubles final in that debut appearance — partnering with 1995 Wimbledon girls' champion in both singles and doubles Aleksandra Olsza, she lost the title match to Meredith McGrath and Larisa Savchenko-Neiland. In February–March 1996, Kournikova won two ITF titles, in Midland, Michigan and Rockford, Illinois. Still only 14 years of age, in April 1996 she debuted at the Fed Cup for Russia, the youngest player ever to participate and win a match. In 1996, she started playing under a new coach, Ed Nagel. Her six-year association with Nagel was successful. At 15, she made her Grand Slam debut, reaching the fourth round of the 1996 US Open, losing to Steffi Graf, the eventual champion. After this tournament, Kournikova's ranking jumped from No. 144 to debut in the Top 100 at No. 69. Kournikova was a member of the Russian delegation to the 1996 Olympic Games in Atlanta, Georgia. In 1996, she was named WTA Newcomer of the Year, and she was ranked No. 57 in the end of the season. Kournikova entered the 1997 Australian Open as world No. 67, where she lost in the first round to world No. 12, Amanda Coetzer. At the Italian Open, Kournikova lost to Amanda Coetzer in the second round. She reached the semi-finals in the doubles partnering with Elena Likhovtseva, before losing to the sixth seeds Mary Joe Fernández and Patricia Tarabini. At the French Open, Kournikova made it to the third round before losing to world No. 1, Martina Hingis. She also reached the third round in doubles with Likhovtseva. At the Wimbledon Championships, Kournikova became only the second woman in the open era to reach the semi-finals in her Wimbledon debut, the first being Chris Evert in 1972. There she lost to eventual champion Martina Hingis. At the US Open, she lost in the second round to the eleventh seed Irina Spîrlea.

In 1989, at the age of eight, Kournikova began appearing in junior tournaments, and by the following year, was attracting attention from tennis scouts across the world. She signed a management deal at age ten and went to Bradenton, Florida, to train at Nick Bollettieri's celebrated tennis academy. Tennis career 1989–1997: Early years and breakthrough Following her arrival in the United States, she became prominent on the tennis scene. At the age of 14, she won the European Championships and the Italian Open Junior tournament. In December 1995, she became the youngest player to win the 18-and-under division of the Junior Orange Bowl tennis tournament. By the end of the year, Kournikova was crowned the ITF Junior World Champion U-18 and Junior European Champion U-18. Earlier, in September 1995, Kournikova, still only 14 years of age, debuted in the WTA Tour, when she received a wildcard into the qualifications at the WTA tournament in Moscow, the Moscow Ladies Open, and qualified before losing in the second round of the main draw to third-seeded Sabine Appelmans. She also reached her first WTA Tour doubles final in that debut appearance — partnering with 1995 Wimbledon girls' champion in both singles and doubles Aleksandra Olsza, she lost the title match to Meredith McGrath and Larisa Savchenko-Neiland. In February–March 1996, Kournikova won two ITF titles, in Midland, Michigan and Rockford, Illinois. Still only 14 years of age, in April 1996 she debuted at the Fed Cup for Russia, the youngest player ever to participate and win a match. In 1996, she started playing under a new coach, Ed Nagel. Her six-year association with Nagel was successful. At 15, she made her Grand Slam debut, reaching the fourth round of the 1996 US Open, losing to Steffi Graf, the eventual champion. After this tournament, Kournikova's ranking jumped from No. 144 to debut in the Top 100 at No. 69. Kournikova was a member of the Russian delegation to the 1996 Olympic Games in Atlanta, Georgia. In 1996, she was named WTA Newcomer of the Year, and she was ranked No. 57 in the end of the season. Kournikova entered the 1997 Australian Open as world No. 67, where she lost in the first round to world No. 12, Amanda Coetzer. At the Italian Open, Kournikova lost to Amanda Coetzer in the second round. She reached the semi-finals in the doubles partnering with Elena Likhovtseva, before losing to the sixth seeds Mary Joe Fernández and Patricia Tarabini. At the French Open, Kournikova made it to the third round before losing to world No. 1, Martina Hingis. She also reached the third round in doubles with Likhovtseva. At the Wimbledon Championships, Kournikova became only the second woman in the open era to reach the semi-finals in her Wimbledon debut, the first being Chris Evert in 1972. There she lost to eventual champion Martina Hingis. At the US Open, she lost in the second round to the eleventh seed Irina Spîrlea.

In 1989, at the age of eight, Kournikova began appearing in junior tournaments, and by the following year, was attracting attention from tennis scouts across the world. She signed a management deal at age ten and went to Bradenton, Florida, to train at Nick Bollettieri's celebrated tennis academy. Tennis career 1989–1997: Early years and breakthrough Following her arrival in the United States, she became prominent on the tennis scene. At the age of 14, she won the European Championships and the Italian Open Junior tournament. In December 1995, she became the youngest player to win the 18-and-under division of the Junior Orange Bowl tennis tournament. By the end of the year, Kournikova was crowned the ITF Junior World Champion U-18 and Junior European Champion U-18. Earlier, in September 1995, Kournikova, still only 14 years of age, debuted in the WTA Tour, when she received a wildcard into the qualifications at the WTA tournament in Moscow, the Moscow Ladies Open, and qualified before losing in the second round of the main draw to third-seeded Sabine Appelmans. She also reached her first WTA Tour doubles final in that debut appearance — partnering with 1995 Wimbledon girls' champion in both singles and doubles Aleksandra Olsza, she lost the title match to Meredith McGrath and Larisa Savchenko-Neiland. In February–March 1996, Kournikova won two ITF titles, in Midland, Michigan and Rockford, Illinois. Still only 14 years of age, in April 1996 she debuted at the Fed Cup for Russia, the youngest player ever to participate and win a match. In 1996, she started playing under a new coach, Ed Nagel. Her six-year association with Nagel was successful. At 15, she made her Grand Slam debut, reaching the fourth round of the 1996 US Open, losing to Steffi Graf, the eventual champion. After this tournament, Kournikova's ranking jumped from No. 144 to debut in the Top 100 at No. 69. Kournikova was a member of the Russian delegation to the 1996 Olympic Games in Atlanta, Georgia. In 1996, she was named WTA Newcomer of the Year, and she was ranked No. 57 in the end of the season. Kournikova entered the 1997 Australian Open as world No. 67, where she lost in the first round to world No. 12, Amanda Coetzer. At the Italian Open, Kournikova lost to Amanda Coetzer in the second round. She reached the semi-finals in the doubles partnering with Elena Likhovtseva, before losing to the sixth seeds Mary Joe Fernández and Patricia Tarabini. At the French Open, Kournikova made it to the third round before losing to world No. 1, Martina Hingis. She also reached the third round in doubles with Likhovtseva. At the Wimbledon Championships, Kournikova became only the second woman in the open era to reach the semi-finals in her Wimbledon debut, the first being Chris Evert in 1972. There she lost to eventual champion Martina Hingis. At the US Open, she lost in the second round to the eleventh seed Irina Spîrlea.

Partnering with Likhovtseva, she reached the third round of the women's doubles event. Kournikova played her last WTA Tour event of 1997 at Porsche Tennis Grand Prix in Filderstadt, losing to Amanda Coetzer in the second round of singles, and in the first round of doubles to Lindsay Davenport and Jana Novotná partnering with Likhovtseva. She broke into the top 50 on 19 May, and was ranked No. 32 in singles and No. 41 in doubles at the end of the season. 1998–2000: Success and stardom In 1998, Kournikova broke into the WTA's top 20 rankings for the first time, when she was ranked No. 16. At the Australian Open, Kournikova lost in the third round to world No. 1 player, Martina Hingis. She also partnered with Larisa Savchenko-Neiland in women's doubles, and they lost to eventual champions Hingis and Mirjana Lučić in the second round. Although she lost in the second round of the Paris Open to Anke Huber in singles, Kournikova reached her second doubles WTA Tour final, partnering with Larisa Savchenko-Neiland. They lost to Sabine Appelmans and Miriam Oremans. Kournikova and Savchenko-Neiland reached their second consecutive final at the Linz Open, losing to Alexandra Fusai and Nathalie Tauziat. At the Miami Open, Kournikova reached her first WTA Tour singles final, before losing to Venus Williams in the final. Kournikova then reached two consecutive quarterfinals, at Amelia Island and the Italian Open, losing respectively to Lindsay Davenport and Martina Hingis. At the German Open, she reached the semi-finals in both singles and doubles, partnering with Larisa Savchenko-Neiland. At the French Open Kournikova had her best result at this tournament, making it to the fourth round before losing to Jana Novotná. She also reached her first Grand Slam doubles semi-finals, losing with Savchenko-Neiland to Lindsay Davenport and Natasha Zvereva. During her quarterfinals match at the grass-court Eastbourne Open versus Steffi Graf, Kournikova injured her thumb, which would eventually force her to withdraw from the 1998 Wimbledon Championships. However, she won that match, but then withdrew from her semi-finals match against Arantxa Sánchez Vicario. Kournikova returned for the Du Maurier Open and made it to the third round, before losing to Conchita Martínez. At the US Open Kournikova reached the fourth round before losing to Arantxa Sánchez Vicario. Her strong year qualified her for the year-end 1998 WTA Tour Championships, but she lost to Monica Seles in the first round. However, with Seles, she won her first WTA doubles title, in Tokyo, beating Mary Joe Fernández and Arantxa Sánchez Vicario in the final. At the end of the season, she was ranked No. 10 in doubles. At the start of the 1999 season, Kournikova advanced to the fourth round in singles before losing to Mary Pierce. However, Kournikova won her first doubles Grand Slam title, partnering with Martina Hingis. The two defeated Lindsay Davenport and Natasha Zvereva in the final. At the Tier I Family Circle Cup, Kournikova reached her second WTA Tour final, but lost to Martina Hingis.

Partnering with Likhovtseva, she reached the third round of the women's doubles event. Kournikova played her last WTA Tour event of 1997 at Porsche Tennis Grand Prix in Filderstadt, losing to Amanda Coetzer in the second round of singles, and in the first round of doubles to Lindsay Davenport and Jana Novotná partnering with Likhovtseva. She broke into the top 50 on 19 May, and was ranked No. 32 in singles and No. 41 in doubles at the end of the season. 1998–2000: Success and stardom In 1998, Kournikova broke into the WTA's top 20 rankings for the first time, when she was ranked No. 16. At the Australian Open, Kournikova lost in the third round to world No. 1 player, Martina Hingis. She also partnered with Larisa Savchenko-Neiland in women's doubles, and they lost to eventual champions Hingis and Mirjana Lučić in the second round. Although she lost in the second round of the Paris Open to Anke Huber in singles, Kournikova reached her second doubles WTA Tour final, partnering with Larisa Savchenko-Neiland. They lost to Sabine Appelmans and Miriam Oremans. Kournikova and Savchenko-Neiland reached their second consecutive final at the Linz Open, losing to Alexandra Fusai and Nathalie Tauziat. At the Miami Open, Kournikova reached her first WTA Tour singles final, before losing to Venus Williams in the final. Kournikova then reached two consecutive quarterfinals, at Amelia Island and the Italian Open, losing respectively to Lindsay Davenport and Martina Hingis. At the German Open, she reached the semi-finals in both singles and doubles, partnering with Larisa Savchenko-Neiland. At the French Open Kournikova had her best result at this tournament, making it to the fourth round before losing to Jana Novotná. She also reached her first Grand Slam doubles semi-finals, losing with Savchenko-Neiland to Lindsay Davenport and Natasha Zvereva. During her quarterfinals match at the grass-court Eastbourne Open versus Steffi Graf, Kournikova injured her thumb, which would eventually force her to withdraw from the 1998 Wimbledon Championships. However, she won that match, but then withdrew from her semi-finals match against Arantxa Sánchez Vicario. Kournikova returned for the Du Maurier Open and made it to the third round, before losing to Conchita Martínez. At the US Open Kournikova reached the fourth round before losing to Arantxa Sánchez Vicario. Her strong year qualified her for the year-end 1998 WTA Tour Championships, but she lost to Monica Seles in the first round. However, with Seles, she won her first WTA doubles title, in Tokyo, beating Mary Joe Fernández and Arantxa Sánchez Vicario in the final. At the end of the season, she was ranked No. 10 in doubles. At the start of the 1999 season, Kournikova advanced to the fourth round in singles before losing to Mary Pierce. However, Kournikova won her first doubles Grand Slam title, partnering with Martina Hingis. The two defeated Lindsay Davenport and Natasha Zvereva in the final. At the Tier I Family Circle Cup, Kournikova reached her second WTA Tour final, but lost to Martina Hingis.

Partnering with Likhovtseva, she reached the third round of the women's doubles event. Kournikova played her last WTA Tour event of 1997 at Porsche Tennis Grand Prix in Filderstadt, losing to Amanda Coetzer in the second round of singles, and in the first round of doubles to Lindsay Davenport and Jana Novotná partnering with Likhovtseva. She broke into the top 50 on 19 May, and was ranked No. 32 in singles and No. 41 in doubles at the end of the season. 1998–2000: Success and stardom In 1998, Kournikova broke into the WTA's top 20 rankings for the first time, when she was ranked No. 16. At the Australian Open, Kournikova lost in the third round to world No. 1 player, Martina Hingis. She also partnered with Larisa Savchenko-Neiland in women's doubles, and they lost to eventual champions Hingis and Mirjana Lučić in the second round. Although she lost in the second round of the Paris Open to Anke Huber in singles, Kournikova reached her second doubles WTA Tour final, partnering with Larisa Savchenko-Neiland. They lost to Sabine Appelmans and Miriam Oremans. Kournikova and Savchenko-Neiland reached their second consecutive final at the Linz Open, losing to Alexandra Fusai and Nathalie Tauziat. At the Miami Open, Kournikova reached her first WTA Tour singles final, before losing to Venus Williams in the final. Kournikova then reached two consecutive quarterfinals, at Amelia Island and the Italian Open, losing respectively to Lindsay Davenport and Martina Hingis. At the German Open, she reached the semi-finals in both singles and doubles, partnering with Larisa Savchenko-Neiland. At the French Open Kournikova had her best result at this tournament, making it to the fourth round before losing to Jana Novotná. She also reached her first Grand Slam doubles semi-finals, losing with Savchenko-Neiland to Lindsay Davenport and Natasha Zvereva. During her quarterfinals match at the grass-court Eastbourne Open versus Steffi Graf, Kournikova injured her thumb, which would eventually force her to withdraw from the 1998 Wimbledon Championships. However, she won that match, but then withdrew from her semi-finals match against Arantxa Sánchez Vicario. Kournikova returned for the Du Maurier Open and made it to the third round, before losing to Conchita Martínez. At the US Open Kournikova reached the fourth round before losing to Arantxa Sánchez Vicario. Her strong year qualified her for the year-end 1998 WTA Tour Championships, but she lost to Monica Seles in the first round. However, with Seles, she won her first WTA doubles title, in Tokyo, beating Mary Joe Fernández and Arantxa Sánchez Vicario in the final. At the end of the season, she was ranked No. 10 in doubles. At the start of the 1999 season, Kournikova advanced to the fourth round in singles before losing to Mary Pierce. However, Kournikova won her first doubles Grand Slam title, partnering with Martina Hingis. The two defeated Lindsay Davenport and Natasha Zvereva in the final. At the Tier I Family Circle Cup, Kournikova reached her second WTA Tour final, but lost to Martina Hingis.

She then defeated Jennifer Capriati, Lindsay Davenport and Patty Schnyder on her route to the Bausch & Lomb Championships semi-finals, losing to Ruxandra Dragomir. At The French Open, Kournikova reached the fourth round before losing to eventual champion Steffi Graf. Once the grass-court season commenced in England, Kournikova lost to Nathalie Tauziat in the semi-finals in Eastbourne. At Wimbledon, Kournikova lost to Venus Williams in the fourth round. She also reached the final in mixed doubles, partnering with Jonas Björkman, but they lost to Leander Paes and Lisa Raymond. Kournikova again qualified for year-end WTA Tour Championships, but lost to Mary Pierce in the first round, and ended the season as World No. 12. While Kournikova had a successful singles season, she was even more successful in doubles. After their victory at the Australian Open, she and Martina Hingis won tournaments in Indian Wells, Rome, Eastbourne and the WTA Tour Championships, and reached the final of The French Open where they lost to Serena and Venus Williams. Partnering with Elena Likhovtseva, Kournikova also reached the final in Stanford. On 22 November 1999 she reached the world No. 1 ranking in doubles, and ended the season at this ranking. Anna Kournikova and Martina Hingis were presented with the WTA Award for Doubles Team of the Year. Kournikova opened her 2000 season winning the Gold Coast Open doubles tournament partnering with Julie Halard. She then reached the singles semi-finals at the Medibank International Sydney, losing to Lindsay Davenport. At the Australian Open, she reached the fourth round in singles and the semi-finals in doubles. That season, Kournikova reached eight semi-finals (Sydney, Scottsdale, Stanford, San Diego, Luxembourg, Leipzig and Tour Championships), seven quarterfinals (Gold Coast, Tokyo, Amelia Island, Hamburg, Eastbourne, Zürich and Philadelphia) and one final. On 20 November 2000 she broke into top 10 for the first time, reaching No. 8. She was also ranked No. 4 in doubles at the end of the season. Kournikova was once again, more successful in doubles. She reached the final of the US Open in mixed doubles, partnering with Max Mirnyi, but they lost to Jared Palmer and Arantxa Sánchez Vicario. She also won six doubles titles – Gold Coast (with Julie Halard), Hamburg (with Natasha Zvereva), Filderstadt, Zürich, Philadelphia and the Tour Championships (with Martina Hingis). 2001–2003: Injuries and final years Her 2001 season was plagued by injuries, including a left foot stress fracture which made her withdraw from 12 tournaments, including the French Open and Wimbledon. She underwent surgery in April. She reached her second career grand slam quarterfinals, at the Australian Open. Kournikova then withdrew from several events due to continuing problems with her left foot and did not return until Leipzig. With Barbara Schett, she won the doubles title in Sydney. She then lost in the finals in Tokyo, partnering with Iroda Tulyaganova, and at San Diego, partnering with Martina Hingis. Hingis and Kournikova also won the Kremlin Cup. At the end of the 2001 season, she was ranked No.

She then defeated Jennifer Capriati, Lindsay Davenport and Patty Schnyder on her route to the Bausch & Lomb Championships semi-finals, losing to Ruxandra Dragomir. At The French Open, Kournikova reached the fourth round before losing to eventual champion Steffi Graf. Once the grass-court season commenced in England, Kournikova lost to Nathalie Tauziat in the semi-finals in Eastbourne. At Wimbledon, Kournikova lost to Venus Williams in the fourth round. She also reached the final in mixed doubles, partnering with Jonas Björkman, but they lost to Leander Paes and Lisa Raymond. Kournikova again qualified for year-end WTA Tour Championships, but lost to Mary Pierce in the first round, and ended the season as World No. 12. While Kournikova had a successful singles season, she was even more successful in doubles. After their victory at the Australian Open, she and Martina Hingis won tournaments in Indian Wells, Rome, Eastbourne and the WTA Tour Championships, and reached the final of The French Open where they lost to Serena and Venus Williams. Partnering with Elena Likhovtseva, Kournikova also reached the final in Stanford. On 22 November 1999 she reached the world No. 1 ranking in doubles, and ended the season at this ranking. Anna Kournikova and Martina Hingis were presented with the WTA Award for Doubles Team of the Year. Kournikova opened her 2000 season winning the Gold Coast Open doubles tournament partnering with Julie Halard. She then reached the singles semi-finals at the Medibank International Sydney, losing to Lindsay Davenport. At the Australian Open, she reached the fourth round in singles and the semi-finals in doubles. That season, Kournikova reached eight semi-finals (Sydney, Scottsdale, Stanford, San Diego, Luxembourg, Leipzig and Tour Championships), seven quarterfinals (Gold Coast, Tokyo, Amelia Island, Hamburg, Eastbourne, Zürich and Philadelphia) and one final. On 20 November 2000 she broke into top 10 for the first time, reaching No. 8. She was also ranked No. 4 in doubles at the end of the season. Kournikova was once again, more successful in doubles. She reached the final of the US Open in mixed doubles, partnering with Max Mirnyi, but they lost to Jared Palmer and Arantxa Sánchez Vicario. She also won six doubles titles – Gold Coast (with Julie Halard), Hamburg (with Natasha Zvereva), Filderstadt, Zürich, Philadelphia and the Tour Championships (with Martina Hingis). 2001–2003: Injuries and final years Her 2001 season was plagued by injuries, including a left foot stress fracture which made her withdraw from 12 tournaments, including the French Open and Wimbledon. She underwent surgery in April. She reached her second career grand slam quarterfinals, at the Australian Open. Kournikova then withdrew from several events due to continuing problems with her left foot and did not return until Leipzig. With Barbara Schett, she won the doubles title in Sydney. She then lost in the finals in Tokyo, partnering with Iroda Tulyaganova, and at San Diego, partnering with Martina Hingis. Hingis and Kournikova also won the Kremlin Cup. At the end of the 2001 season, she was ranked No.

She then defeated Jennifer Capriati, Lindsay Davenport and Patty Schnyder on her route to the Bausch & Lomb Championships semi-finals, losing to Ruxandra Dragomir. At The French Open, Kournikova reached the fourth round before losing to eventual champion Steffi Graf. Once the grass-court season commenced in England, Kournikova lost to Nathalie Tauziat in the semi-finals in Eastbourne. At Wimbledon, Kournikova lost to Venus Williams in the fourth round. She also reached the final in mixed doubles, partnering with Jonas Björkman, but they lost to Leander Paes and Lisa Raymond. Kournikova again qualified for year-end WTA Tour Championships, but lost to Mary Pierce in the first round, and ended the season as World No. 12. While Kournikova had a successful singles season, she was even more successful in doubles. After their victory at the Australian Open, she and Martina Hingis won tournaments in Indian Wells, Rome, Eastbourne and the WTA Tour Championships, and reached the final of The French Open where they lost to Serena and Venus Williams. Partnering with Elena Likhovtseva, Kournikova also reached the final in Stanford. On 22 November 1999 she reached the world No. 1 ranking in doubles, and ended the season at this ranking. Anna Kournikova and Martina Hingis were presented with the WTA Award for Doubles Team of the Year. Kournikova opened her 2000 season winning the Gold Coast Open doubles tournament partnering with Julie Halard. She then reached the singles semi-finals at the Medibank International Sydney, losing to Lindsay Davenport. At the Australian Open, she reached the fourth round in singles and the semi-finals in doubles. That season, Kournikova reached eight semi-finals (Sydney, Scottsdale, Stanford, San Diego, Luxembourg, Leipzig and Tour Championships), seven quarterfinals (Gold Coast, Tokyo, Amelia Island, Hamburg, Eastbourne, Zürich and Philadelphia) and one final. On 20 November 2000 she broke into top 10 for the first time, reaching No. 8. She was also ranked No. 4 in doubles at the end of the season. Kournikova was once again, more successful in doubles. She reached the final of the US Open in mixed doubles, partnering with Max Mirnyi, but they lost to Jared Palmer and Arantxa Sánchez Vicario. She also won six doubles titles – Gold Coast (with Julie Halard), Hamburg (with Natasha Zvereva), Filderstadt, Zürich, Philadelphia and the Tour Championships (with Martina Hingis). 2001–2003: Injuries and final years Her 2001 season was plagued by injuries, including a left foot stress fracture which made her withdraw from 12 tournaments, including the French Open and Wimbledon. She underwent surgery in April. She reached her second career grand slam quarterfinals, at the Australian Open. Kournikova then withdrew from several events due to continuing problems with her left foot and did not return until Leipzig. With Barbara Schett, she won the doubles title in Sydney. She then lost in the finals in Tokyo, partnering with Iroda Tulyaganova, and at San Diego, partnering with Martina Hingis. Hingis and Kournikova also won the Kremlin Cup. At the end of the 2001 season, she was ranked No.

74 in singles and No. 26 in doubles. Kournikova regained some success in 2002. She reached the semi-finals of Auckland, Tokyo, Acapulco and San Diego, and the final of the China Open, losing to Anna Smashnova. This was Kournikova's last singles final. With Martina Hingis, she lost in the final at Sydney, but they won their second Grand Slam title together, the Australian Open. They also lost in the quarterfinals of the US Open. With Chanda Rubin, Kournikova played the semi-finals of Wimbledon, but they lost to Serena and Venus Williams. Partnering with Janet Lee, she won the Shanghai title. At the end of 2002 season, she was ranked No. 35 in singles and No. 11 in doubles. In 2003, Anna Kournikova achieved her first Grand Slam match victory in two years at the Australian Open. She defeated Henrieta Nagyová in the first round, and then lost to Justine Henin-Hardenne in the 2nd round. She withdrew from Tokyo due to a sprained back suffered at the Australian Open and did not return to Tour until Miami. On 9 April, in what would be the final WTA match of her career, Kournikova dropped out in the first round of the Family Circle Cup in Charleston, due to a left adductor strain. Her singles world ranking was 67. She reached the semi-finals at the ITF tournament in Sea Island, before withdrawing from a match versus Maria Sharapova due to the adductor injury. She lost in the first round of the ITF tournament in Charlottesville. She did not compete for the rest of the season due to a continuing back injury. At the end of the 2003 season and her professional career, she was ranked No. 305 in singles and No. 176 in doubles. Kournikova's two Grand Slam doubles titles came in 1999 and 2002, both at the Australian Open in the Women's Doubles event with partner Martina Hingis. Kournikova proved a successful doubles player on the professional circuit, winning 16 tournament doubles titles, including two Australian Opens and being a finalist in mixed doubles at the US Open and at Wimbledon, and reaching the No. 1 ranking in doubles in the WTA Tour rankings. Her pro career doubles record was 200–71. However, her singles career plateaued after 1999. For the most part, she managed to retain her ranking between 10 and 15 (her career high singles ranking was No.8), but her expected finals breakthrough failed to occur; she only reached four finals out of 130 singles tournaments, never in a Grand Slam event, and never won one. Her singles record is 209–129. Her final playing years were marred by a string of injuries, especially back injuries, which caused her ranking to erode gradually. As a personality Kournikova was among the most common search strings for both articles and images in her prime. 2004–present: Exhibitions and World Team Tennis Kournikova has not played on the WTA Tour since 2003, but still plays exhibition matches for charitable causes.

74 in singles and No. 26 in doubles. Kournikova regained some success in 2002. She reached the semi-finals of Auckland, Tokyo, Acapulco and San Diego, and the final of the China Open, losing to Anna Smashnova. This was Kournikova's last singles final. With Martina Hingis, she lost in the final at Sydney, but they won their second Grand Slam title together, the Australian Open. They also lost in the quarterfinals of the US Open. With Chanda Rubin, Kournikova played the semi-finals of Wimbledon, but they lost to Serena and Venus Williams. Partnering with Janet Lee, she won the Shanghai title. At the end of 2002 season, she was ranked No. 35 in singles and No. 11 in doubles. In 2003, Anna Kournikova achieved her first Grand Slam match victory in two years at the Australian Open. She defeated Henrieta Nagyová in the first round, and then lost to Justine Henin-Hardenne in the 2nd round. She withdrew from Tokyo due to a sprained back suffered at the Australian Open and did not return to Tour until Miami. On 9 April, in what would be the final WTA match of her career, Kournikova dropped out in the first round of the Family Circle Cup in Charleston, due to a left adductor strain. Her singles world ranking was 67. She reached the semi-finals at the ITF tournament in Sea Island, before withdrawing from a match versus Maria Sharapova due to the adductor injury. She lost in the first round of the ITF tournament in Charlottesville. She did not compete for the rest of the season due to a continuing back injury. At the end of the 2003 season and her professional career, she was ranked No. 305 in singles and No. 176 in doubles. Kournikova's two Grand Slam doubles titles came in 1999 and 2002, both at the Australian Open in the Women's Doubles event with partner Martina Hingis. Kournikova proved a successful doubles player on the professional circuit, winning 16 tournament doubles titles, including two Australian Opens and being a finalist in mixed doubles at the US Open and at Wimbledon, and reaching the No. 1 ranking in doubles in the WTA Tour rankings. Her pro career doubles record was 200–71. However, her singles career plateaued after 1999. For the most part, she managed to retain her ranking between 10 and 15 (her career high singles ranking was No.8), but her expected finals breakthrough failed to occur; she only reached four finals out of 130 singles tournaments, never in a Grand Slam event, and never won one. Her singles record is 209–129. Her final playing years were marred by a string of injuries, especially back injuries, which caused her ranking to erode gradually. As a personality Kournikova was among the most common search strings for both articles and images in her prime. 2004–present: Exhibitions and World Team Tennis Kournikova has not played on the WTA Tour since 2003, but still plays exhibition matches for charitable causes.

74 in singles and No. 26 in doubles. Kournikova regained some success in 2002. She reached the semi-finals of Auckland, Tokyo, Acapulco and San Diego, and the final of the China Open, losing to Anna Smashnova. This was Kournikova's last singles final. With Martina Hingis, she lost in the final at Sydney, but they won their second Grand Slam title together, the Australian Open. They also lost in the quarterfinals of the US Open. With Chanda Rubin, Kournikova played the semi-finals of Wimbledon, but they lost to Serena and Venus Williams. Partnering with Janet Lee, she won the Shanghai title. At the end of 2002 season, she was ranked No. 35 in singles and No. 11 in doubles. In 2003, Anna Kournikova achieved her first Grand Slam match victory in two years at the Australian Open. She defeated Henrieta Nagyová in the first round, and then lost to Justine Henin-Hardenne in the 2nd round. She withdrew from Tokyo due to a sprained back suffered at the Australian Open and did not return to Tour until Miami. On 9 April, in what would be the final WTA match of her career, Kournikova dropped out in the first round of the Family Circle Cup in Charleston, due to a left adductor strain. Her singles world ranking was 67. She reached the semi-finals at the ITF tournament in Sea Island, before withdrawing from a match versus Maria Sharapova due to the adductor injury. She lost in the first round of the ITF tournament in Charlottesville. She did not compete for the rest of the season due to a continuing back injury. At the end of the 2003 season and her professional career, she was ranked No. 305 in singles and No. 176 in doubles. Kournikova's two Grand Slam doubles titles came in 1999 and 2002, both at the Australian Open in the Women's Doubles event with partner Martina Hingis. Kournikova proved a successful doubles player on the professional circuit, winning 16 tournament doubles titles, including two Australian Opens and being a finalist in mixed doubles at the US Open and at Wimbledon, and reaching the No. 1 ranking in doubles in the WTA Tour rankings. Her pro career doubles record was 200–71. However, her singles career plateaued after 1999. For the most part, she managed to retain her ranking between 10 and 15 (her career high singles ranking was No.8), but her expected finals breakthrough failed to occur; she only reached four finals out of 130 singles tournaments, never in a Grand Slam event, and never won one. Her singles record is 209–129. Her final playing years were marred by a string of injuries, especially back injuries, which caused her ranking to erode gradually. As a personality Kournikova was among the most common search strings for both articles and images in her prime. 2004–present: Exhibitions and World Team Tennis Kournikova has not played on the WTA Tour since 2003, but still plays exhibition matches for charitable causes.

In late 2004, she participated in three events organized by Elton John and by fellow tennis players Serena Williams and Andy Roddick. In January 2005, she played in a doubles charity event for the Indian Ocean tsunami with John McEnroe, Andy Roddick, and Chris Evert. In November 2005, she teamed up with Martina Hingis, playing against Lisa Raymond and Samantha Stosur in the WTT finals for charity. Kournikova is also a member of the St. Louis Aces in the World Team Tennis (WTT), playing doubles only. In September 2008, Kournikova showed up for the 2008 Nautica Malibu Triathlon held at Zuma Beach in Malibu, California. The Race raised funds for children's Hospital Los Angeles. She won that race for women's K-Swiss team. On 27 September 2008, Kournikova played exhibition mixed doubles matches in Charlotte, North Carolina, partnering with Tim Wilkison and Karel Nováček. Kournikova and Wilkison defeated Jimmy Arias and Chanda Rubin, and then Kournikova and Novacek defeated Rubin and Wilkison. On 12 October 2008, Anna Kournikova played one exhibition match for the annual charity event, hosted by Billie Jean King and Elton John, and raised more than $400,000 for the Elton John AIDS Foundation and Atlanta AIDS Partnership Fund. She played doubles with Andy Roddick (they were coached by David Chang) versus Martina Navratilova and Jesse Levine (coached by Billie Jean King); Kournikova and Roddick won. Kournikova competed alongside John McEnroe, Tracy Austin and Jim Courier at the "Legendary Night", which was held on 2 May 2009, at the Turning Stone Event Center in Verona, New York. The exhibition included a mixed doubles match of McEnroe and Austin against Courier and Kournikova. In 2008, she was named a spokesperson for K-Swiss. In 2005, Kournikova stated that if she were 100% fit, she would like to come back and compete again. In June 2010, Kournikova reunited with her doubles partner Martina Hingis to participate in competitive tennis for the first time in seven years in the Invitational Ladies Doubles event at Wimbledon. On 29 June 2010 they defeated the British pair Samantha Smith and Anne Hobbs. Playing style Kournikova plays right-handed with a two-handed backhand. She is a great player at the net. She can hit forceful groundstrokes and also drop shots. Her playing style fits the profile for a doubles player, and is complemented by her height. She has been compared to such doubles specialists as Pam Shriver and Peter Fleming. Personal life Kournikova was in a relationship with fellow Russian, Pavel Bure, an NHL ice hockey player. The two met in 1999, when Kournikova was still linked to Bure's former Russian teammate Sergei Fedorov. Bure and Kournikova were reported to have been engaged in 2000 after a reporter took a photo of them together in a Florida restaurant where Bure supposedly asked Kournikova to marry him. As the story made headlines in Russia, where they were both heavily followed in the media as celebrities, Bure and Kournikova both denied any engagement.

In late 2004, she participated in three events organized by Elton John and by fellow tennis players Serena Williams and Andy Roddick. In January 2005, she played in a doubles charity event for the Indian Ocean tsunami with John McEnroe, Andy Roddick, and Chris Evert. In November 2005, she teamed up with Martina Hingis, playing against Lisa Raymond and Samantha Stosur in the WTT finals for charity. Kournikova is also a member of the St. Louis Aces in the World Team Tennis (WTT), playing doubles only. In September 2008, Kournikova showed up for the 2008 Nautica Malibu Triathlon held at Zuma Beach in Malibu, California. The Race raised funds for children's Hospital Los Angeles. She won that race for women's K-Swiss team. On 27 September 2008, Kournikova played exhibition mixed doubles matches in Charlotte, North Carolina, partnering with Tim Wilkison and Karel Nováček. Kournikova and Wilkison defeated Jimmy Arias and Chanda Rubin, and then Kournikova and Novacek defeated Rubin and Wilkison. On 12 October 2008, Anna Kournikova played one exhibition match for the annual charity event, hosted by Billie Jean King and Elton John, and raised more than $400,000 for the Elton John AIDS Foundation and Atlanta AIDS Partnership Fund. She played doubles with Andy Roddick (they were coached by David Chang) versus Martina Navratilova and Jesse Levine (coached by Billie Jean King); Kournikova and Roddick won. Kournikova competed alongside John McEnroe, Tracy Austin and Jim Courier at the "Legendary Night", which was held on 2 May 2009, at the Turning Stone Event Center in Verona, New York. The exhibition included a mixed doubles match of McEnroe and Austin against Courier and Kournikova. In 2008, she was named a spokesperson for K-Swiss. In 2005, Kournikova stated that if she were 100% fit, she would like to come back and compete again. In June 2010, Kournikova reunited with her doubles partner Martina Hingis to participate in competitive tennis for the first time in seven years in the Invitational Ladies Doubles event at Wimbledon. On 29 June 2010 they defeated the British pair Samantha Smith and Anne Hobbs. Playing style Kournikova plays right-handed with a two-handed backhand. She is a great player at the net. She can hit forceful groundstrokes and also drop shots. Her playing style fits the profile for a doubles player, and is complemented by her height. She has been compared to such doubles specialists as Pam Shriver and Peter Fleming. Personal life Kournikova was in a relationship with fellow Russian, Pavel Bure, an NHL ice hockey player. The two met in 1999, when Kournikova was still linked to Bure's former Russian teammate Sergei Fedorov. Bure and Kournikova were reported to have been engaged in 2000 after a reporter took a photo of them together in a Florida restaurant where Bure supposedly asked Kournikova to marry him. As the story made headlines in Russia, where they were both heavily followed in the media as celebrities, Bure and Kournikova both denied any engagement.

In late 2004, she participated in three events organized by Elton John and by fellow tennis players Serena Williams and Andy Roddick. In January 2005, she played in a doubles charity event for the Indian Ocean tsunami with John McEnroe, Andy Roddick, and Chris Evert. In November 2005, she teamed up with Martina Hingis, playing against Lisa Raymond and Samantha Stosur in the WTT finals for charity. Kournikova is also a member of the St. Louis Aces in the World Team Tennis (WTT), playing doubles only. In September 2008, Kournikova showed up for the 2008 Nautica Malibu Triathlon held at Zuma Beach in Malibu, California. The Race raised funds for children's Hospital Los Angeles. She won that race for women's K-Swiss team. On 27 September 2008, Kournikova played exhibition mixed doubles matches in Charlotte, North Carolina, partnering with Tim Wilkison and Karel Nováček. Kournikova and Wilkison defeated Jimmy Arias and Chanda Rubin, and then Kournikova and Novacek defeated Rubin and Wilkison. On 12 October 2008, Anna Kournikova played one exhibition match for the annual charity event, hosted by Billie Jean King and Elton John, and raised more than $400,000 for the Elton John AIDS Foundation and Atlanta AIDS Partnership Fund. She played doubles with Andy Roddick (they were coached by David Chang) versus Martina Navratilova and Jesse Levine (coached by Billie Jean King); Kournikova and Roddick won. Kournikova competed alongside John McEnroe, Tracy Austin and Jim Courier at the "Legendary Night", which was held on 2 May 2009, at the Turning Stone Event Center in Verona, New York. The exhibition included a mixed doubles match of McEnroe and Austin against Courier and Kournikova. In 2008, she was named a spokesperson for K-Swiss. In 2005, Kournikova stated that if she were 100% fit, she would like to come back and compete again. In June 2010, Kournikova reunited with her doubles partner Martina Hingis to participate in competitive tennis for the first time in seven years in the Invitational Ladies Doubles event at Wimbledon. On 29 June 2010 they defeated the British pair Samantha Smith and Anne Hobbs. Playing style Kournikova plays right-handed with a two-handed backhand. She is a great player at the net. She can hit forceful groundstrokes and also drop shots. Her playing style fits the profile for a doubles player, and is complemented by her height. She has been compared to such doubles specialists as Pam Shriver and Peter Fleming. Personal life Kournikova was in a relationship with fellow Russian, Pavel Bure, an NHL ice hockey player. The two met in 1999, when Kournikova was still linked to Bure's former Russian teammate Sergei Fedorov. Bure and Kournikova were reported to have been engaged in 2000 after a reporter took a photo of them together in a Florida restaurant where Bure supposedly asked Kournikova to marry him. As the story made headlines in Russia, where they were both heavily followed in the media as celebrities, Bure and Kournikova both denied any engagement.

Kournikova, 10 years younger than Bure, was 18 years old at the time. Fedorov claimed that he and Kournikova were married in 2001, and divorced in 2003. Kournikova's representatives deny any marriage to Fedorov; however, Fedorov's agent Pat Brisson claims that although he does not know when they got married, he knew "Fedorov was married". Kournikova started dating singer Enrique Iglesias in late 2001 after she had appeared in his music video for "Escape". She has consistently refused to directly confirm or deny the status of her personal relationships. In June 2008, Iglesias was quoted by the Daily Star as having married Kournikova the previous year. They reportedly split in October 2013 but reconciled. The couple have a son and daughter, Nicholas and Lucy, who are fraternal twins born on 16 December 2017. On 30 January 2020, their third child, a daughter, Mary, was born. It was reported in 2010 that Kournikova had become an American citizen. Media publicity In 2000, Kournikova became the new face for Berlei's shock absorber sports bras, and appeared in the "only the ball should bounce" billboard campaign. Following that, she was cast by the Farrelly brothers for a minor role in the 2000 film Me, Myself & Irene starring Jim Carrey and Renée Zellweger. Photographs of her have appeared on covers of various publications, including men's magazines, such as one in the much-publicized 2004 Sports Illustrated Swimsuit Issue, where she posed in bikinis and swimsuits, as well as in FHM and Maxim. Kournikova was named one of Peoples 50 Most Beautiful People in 1998 and was voted "hottest female athlete" on ESPN.com. In 2002, she also placed first in FHM's 100 Sexiest Women in the World in US and UK editions. By contrast, ESPN – citing the degree of hype as compared to actual accomplishments as a singles player – ranked Kournikova 18th in its "25 Biggest Sports Flops of the Past 25 Years". Kournikova was also ranked No. 1 in the ESPN Classic series "Who's number 1?" when the series featured sport's most overrated athletes. She continued to be the most searched athlete on the Internet through 2008 even though she had retired from the professional tennis circuit years earlier. After slipping from first to sixth among athletes in 2009, she moved back up to third place among athletes in terms of search popularity in 2010. In October 2010, Kournikova headed to NBC's The Biggest Loser where she led the contestants in a tennis-workout challenge. In May 2011, it was announced that Kournikova would join The Biggest Loser as a regular celebrity trainer in season 12. She did not return for season 13. Legacy and influence on popular culture A variation of a White Russian made with skim milk is known as an Anna Kournikova. A video game featuring Kournikova's licensed appearance, titled Anna Kournikova's Smash Court Tennis, was developed by Namco and released for the PlayStation in Japan and Europe in November 1998.

Kournikova, 10 years younger than Bure, was 18 years old at the time. Fedorov claimed that he and Kournikova were married in 2001, and divorced in 2003. Kournikova's representatives deny any marriage to Fedorov; however, Fedorov's agent Pat Brisson claims that although he does not know when they got married, he knew "Fedorov was married". Kournikova started dating singer Enrique Iglesias in late 2001 after she had appeared in his music video for "Escape". She has consistently refused to directly confirm or deny the status of her personal relationships. In June 2008, Iglesias was quoted by the Daily Star as having married Kournikova the previous year. They reportedly split in October 2013 but reconciled. The couple have a son and daughter, Nicholas and Lucy, who are fraternal twins born on 16 December 2017. On 30 January 2020, their third child, a daughter, Mary, was born. It was reported in 2010 that Kournikova had become an American citizen. Media publicity In 2000, Kournikova became the new face for Berlei's shock absorber sports bras, and appeared in the "only the ball should bounce" billboard campaign. Following that, she was cast by the Farrelly brothers for a minor role in the 2000 film Me, Myself & Irene starring Jim Carrey and Renée Zellweger. Photographs of her have appeared on covers of various publications, including men's magazines, such as one in the much-publicized 2004 Sports Illustrated Swimsuit Issue, where she posed in bikinis and swimsuits, as well as in FHM and Maxim. Kournikova was named one of Peoples 50 Most Beautiful People in 1998 and was voted "hottest female athlete" on ESPN.com. In 2002, she also placed first in FHM's 100 Sexiest Women in the World in US and UK editions. By contrast, ESPN – citing the degree of hype as compared to actual accomplishments as a singles player – ranked Kournikova 18th in its "25 Biggest Sports Flops of the Past 25 Years". Kournikova was also ranked No. 1 in the ESPN Classic series "Who's number 1?" when the series featured sport's most overrated athletes. She continued to be the most searched athlete on the Internet through 2008 even though she had retired from the professional tennis circuit years earlier. After slipping from first to sixth among athletes in 2009, she moved back up to third place among athletes in terms of search popularity in 2010. In October 2010, Kournikova headed to NBC's The Biggest Loser where she led the contestants in a tennis-workout challenge. In May 2011, it was announced that Kournikova would join The Biggest Loser as a regular celebrity trainer in season 12. She did not return for season 13. Legacy and influence on popular culture A variation of a White Russian made with skim milk is known as an Anna Kournikova. A video game featuring Kournikova's licensed appearance, titled Anna Kournikova's Smash Court Tennis, was developed by Namco and released for the PlayStation in Japan and Europe in November 1998.

Kournikova, 10 years younger than Bure, was 18 years old at the time. Fedorov claimed that he and Kournikova were married in 2001, and divorced in 2003. Kournikova's representatives deny any marriage to Fedorov; however, Fedorov's agent Pat Brisson claims that although he does not know when they got married, he knew "Fedorov was married". Kournikova started dating singer Enrique Iglesias in late 2001 after she had appeared in his music video for "Escape". She has consistently refused to directly confirm or deny the status of her personal relationships. In June 2008, Iglesias was quoted by the Daily Star as having married Kournikova the previous year. They reportedly split in October 2013 but reconciled. The couple have a son and daughter, Nicholas and Lucy, who are fraternal twins born on 16 December 2017. On 30 January 2020, their third child, a daughter, Mary, was born. It was reported in 2010 that Kournikova had become an American citizen. Media publicity In 2000, Kournikova became the new face for Berlei's shock absorber sports bras, and appeared in the "only the ball should bounce" billboard campaign. Following that, she was cast by the Farrelly brothers for a minor role in the 2000 film Me, Myself & Irene starring Jim Carrey and Renée Zellweger. Photographs of her have appeared on covers of various publications, including men's magazines, such as one in the much-publicized 2004 Sports Illustrated Swimsuit Issue, where she posed in bikinis and swimsuits, as well as in FHM and Maxim. Kournikova was named one of Peoples 50 Most Beautiful People in 1998 and was voted "hottest female athlete" on ESPN.com. In 2002, she also placed first in FHM's 100 Sexiest Women in the World in US and UK editions. By contrast, ESPN – citing the degree of hype as compared to actual accomplishments as a singles player – ranked Kournikova 18th in its "25 Biggest Sports Flops of the Past 25 Years". Kournikova was also ranked No. 1 in the ESPN Classic series "Who's number 1?" when the series featured sport's most overrated athletes. She continued to be the most searched athlete on the Internet through 2008 even though she had retired from the professional tennis circuit years earlier. After slipping from first to sixth among athletes in 2009, she moved back up to third place among athletes in terms of search popularity in 2010. In October 2010, Kournikova headed to NBC's The Biggest Loser where she led the contestants in a tennis-workout challenge. In May 2011, it was announced that Kournikova would join The Biggest Loser as a regular celebrity trainer in season 12. She did not return for season 13. Legacy and influence on popular culture A variation of a White Russian made with skim milk is known as an Anna Kournikova. A video game featuring Kournikova's licensed appearance, titled Anna Kournikova's Smash Court Tennis, was developed by Namco and released for the PlayStation in Japan and Europe in November 1998.

A computer virus named after her spread worldwide beginning on 12 February 2001 infecting computers through email in a matter of hours. The Texas hold 'em opening hand of Ace-King is sometimes referred to as an Anna Kournikova, both for the initials on the cards and because the hand looks better than it performs. Career statistics and awards Doubles performance timeline Grand Slam tournament finals Doubles: 3 (2–1) Mixed doubles: 2 (0–2) Awards 1996: WTA Newcomer of the Year 1999: WTA Doubles Team of the Year (with Martina Hingis) Books Anna Kournikova by Susan Holden (2001) ( / ) Anna Kournikova by Connie Berman (2001) (Women Who Win) ( / ) References External links 1981 births Living people Australian Open (tennis) champions Grand Slam (tennis) champions in women's doubles Olympic tennis players of Russia Participants in American reality television series Russian emigrants to the United States Russian female tennis players Russian female models Russian socialites Sportspeople from Miami-Dade County, Florida Tennis players from Moscow Tennis players at the 1996 Summer Olympics People with acquired American citizenship Iglesias family

A computer virus named after her spread worldwide beginning on 12 February 2001 infecting computers through email in a matter of hours. The Texas hold 'em opening hand of Ace-King is sometimes referred to as an Anna Kournikova, both for the initials on the cards and because the hand looks better than it performs. Career statistics and awards Doubles performance timeline Grand Slam tournament finals Doubles: 3 (2–1) Mixed doubles: 2 (0–2) Awards 1996: WTA Newcomer of the Year 1999: WTA Doubles Team of the Year (with Martina Hingis) Books Anna Kournikova by Susan Holden (2001) ( / ) Anna Kournikova by Connie Berman (2001) (Women Who Win) ( / ) References External links 1981 births Living people Australian Open (tennis) champions Grand Slam (tennis) champions in women's doubles Olympic tennis players of Russia Participants in American reality television series Russian emigrants to the United States Russian female tennis players Russian female models Russian socialites Sportspeople from Miami-Dade County, Florida Tennis players from Moscow Tennis players at the 1996 Summer Olympics People with acquired American citizenship Iglesias family

A computer virus named after her spread worldwide beginning on 12 February 2001 infecting computers through email in a matter of hours. The Texas hold 'em opening hand of Ace-King is sometimes referred to as an Anna Kournikova, both for the initials on the cards and because the hand looks better than it performs. Career statistics and awards Doubles performance timeline Grand Slam tournament finals Doubles: 3 (2–1) Mixed doubles: 2 (0–2) Awards 1996: WTA Newcomer of the Year 1999: WTA Doubles Team of the Year (with Martina Hingis) Books Anna Kournikova by Susan Holden (2001) ( / ) Anna Kournikova by Connie Berman (2001) (Women Who Win) ( / ) References External links 1981 births Living people Australian Open (tennis) champions Grand Slam (tennis) champions in women's doubles Olympic tennis players of Russia Participants in American reality television series Russian emigrants to the United States Russian female tennis players Russian female models Russian socialites Sportspeople from Miami-Dade County, Florida Tennis players from Moscow Tennis players at the 1996 Summer Olympics People with acquired American citizenship Iglesias family

Alfons Maria Jakob Alfons Maria Jakob (2 July 1884 – 17 October 1931) was a German neurologist who worked in the field of neuropathology. He was born in Aschaffenburg, Bavaria and educated in medicine at the universities of Munich, Berlin, and Strasbourg, where he received his doctorate in 1908. During the following year, he began clinical work under the psychiatrist Emil Kraepelin and did laboratory work with Franz Nissl and Alois Alzheimer in Munich. In 1911, by way of an invitation from Wilhelm Weygandt, he relocated to Hamburg, where he worked with Theodor Kaes and eventually became head of the laboratory of anatomical pathology at the psychiatric State Hospital Hamburg-Friedrichsberg. Following the death of Kaes in 1913, Jakob succeeded him as prosector. During World War I he served as an army physician in Belgium, and afterwards returned to Hamburg. In 1919, he obtained his habilitation for neurology and in 1924 became a professor of neurology. Under Jakob's guidance the department grew rapidly. He made significant contributions to knowledge on concussion and secondary nerve degeneration and became a doyen of neuropathology. Jakob was the author of five monographs and nearly 80 scientific papers. His neuropathological research contributed greatly to the delineation of several diseases, including multiple sclerosis and Friedreich's ataxia. He first recognised and described Alper's disease and Creutzfeldt–Jakob disease (named along with Munich neuropathologist Hans Gerhard Creutzfeldt). He gained experience in neurosyphilis, having a 200-bed ward devoted entirely to that disorder. Jakob made a lecture tour of the United States (1924) and South America (1928), of which, he wrote a paper on the neuropathology of yellow fever. He suffered from chronic osteomyelitis for the last seven years of his life. This eventually caused a retroperitoneal abscess and paralytic ileus from which he died following operation. Associated eponym Creutzfeldt–Jakob disease: A very rare and incurable degenerative neurological disease. It is the most common form of transmissible spongiform encephalopathies caused by prions. Eponym introduced by Walther Spielmeyer in 1922. Bibliography Die extrapyramidalen Erkrankungen. In: Monographien aus dem Gesamtgebiete der Neurologie und Psychiatry, Berlin, 1923 Normale und pathologische Anatomie und Histologie des Grosshirns. Separate printing of Handbuch der Psychiatry. Leipzig, 1927–1928 Das Kleinhirn. In: Handbuch der mikroskopischen Anatomie, Berlin, 1928 Die Syphilis des Gehirns und seiner Häute. In: Oswald Bumke (edit. ): Handbuch der Geisteskrankheiten, Berlin, 1930. References People from Aschaffenburg University of Hamburg faculty German neurologists German neuroscientists 1884 births 1931 deaths

Alfons Maria Jakob Alfons Maria Jakob (2 July 1884 – 17 October 1931) was a German neurologist who worked in the field of neuropathology. He was born in Aschaffenburg, Bavaria and educated in medicine at the universities of Munich, Berlin, and Strasbourg, where he received his doctorate in 1908. During the following year, he began clinical work under the psychiatrist Emil Kraepelin and did laboratory work with Franz Nissl and Alois Alzheimer in Munich. In 1911, by way of an invitation from Wilhelm Weygandt, he relocated to Hamburg, where he worked with Theodor Kaes and eventually became head of the laboratory of anatomical pathology at the psychiatric State Hospital Hamburg-Friedrichsberg. Following the death of Kaes in 1913, Jakob succeeded him as prosector. During World War I he served as an army physician in Belgium, and afterwards returned to Hamburg. In 1919, he obtained his habilitation for neurology and in 1924 became a professor of neurology. Under Jakob's guidance the department grew rapidly. He made significant contributions to knowledge on concussion and secondary nerve degeneration and became a doyen of neuropathology. Jakob was the author of five monographs and nearly 80 scientific papers. His neuropathological research contributed greatly to the delineation of several diseases, including multiple sclerosis and Friedreich's ataxia. He first recognised and described Alper's disease and Creutzfeldt–Jakob disease (named along with Munich neuropathologist Hans Gerhard Creutzfeldt). He gained experience in neurosyphilis, having a 200-bed ward devoted entirely to that disorder. Jakob made a lecture tour of the United States (1924) and South America (1928), of which, he wrote a paper on the neuropathology of yellow fever. He suffered from chronic osteomyelitis for the last seven years of his life. This eventually caused a retroperitoneal abscess and paralytic ileus from which he died following operation. Associated eponym Creutzfeldt–Jakob disease: A very rare and incurable degenerative neurological disease. It is the most common form of transmissible spongiform encephalopathies caused by prions. Eponym introduced by Walther Spielmeyer in 1922. Bibliography Die extrapyramidalen Erkrankungen. In: Monographien aus dem Gesamtgebiete der Neurologie und Psychiatry, Berlin, 1923 Normale und pathologische Anatomie und Histologie des Grosshirns. Separate printing of Handbuch der Psychiatry. Leipzig, 1927–1928 Das Kleinhirn. In: Handbuch der mikroskopischen Anatomie, Berlin, 1928 Die Syphilis des Gehirns und seiner Häute. In: Oswald Bumke (edit. ): Handbuch der Geisteskrankheiten, Berlin, 1930. References People from Aschaffenburg University of Hamburg faculty German neurologists German neuroscientists 1884 births 1931 deaths

Agnosticism Agnosticism is the view or belief that the existence of God, of the divine or the supernatural is unknown or unknowable. Another definition provided is the view that "human reason is incapable of providing sufficient rational grounds to justify either the belief that God exists or the belief that God does not exist." The English biologist Thomas Henry Huxley coined the word agnostic in 1869, and said "It simply means that a man shall not say he knows or believes that which he has no scientific grounds for professing to know or believe." Earlier thinkers, however, had written works that promoted agnostic points of view, such as Sanjaya Belatthaputta, a 5th-century BCE Indian philosopher who expressed agnosticism about any afterlife; and Protagoras, a 5th-century BCE Greek philosopher who expressed agnosticism about the existence of "the gods". Defining agnosticism Being a scientist, above all else, Huxley presented agnosticism as a form of demarcation. A hypothesis with no supporting, objective, testable evidence is not an objective, scientific claim. As such, there would be no way to test said hypotheses, leaving the results inconclusive. His agnosticism was not compatible with forming a belief as to the truth, or falsehood, of the claim at hand. Karl Popper would also describe himself as an agnostic. According to philosopher William L. Rowe, in this strict sense, agnosticism is the view that human reason is incapable of providing sufficient rational grounds to justify either the belief that God exists or the belief that God does not exist. George H. Smith, while admitting that the narrow definition of atheist was the common usage definition of that word, and admitting that the broad definition of agnostic was the common usage definition of that word, promoted broadening the definition of atheist and narrowing the definition of agnostic. Smith rejects agnosticism as a third alternative to theism and atheism and promotes terms such as agnostic atheism (the view of those who do not hold a belief in the existence of any deity, but claim that the existence of a deity is unknown or inherently unknowable) and agnostic theism (the view of those who believe in the existence of a deity(s), but claim that the existence of a deity is unknown or inherently unknowable). Etymology Agnostic () was used by Thomas Henry Huxley in a speech at a meeting of the Metaphysical Society in 1869 to describe his philosophy, which rejects all claims of spiritual or mystical knowledge. Early Christian church leaders used the Greek word gnosis (knowledge) to describe "spiritual knowledge". Agnosticism is not to be confused with religious views opposing the ancient religious movement of Gnosticism in particular; Huxley used the term in a broader, more abstract sense. Huxley identified agnosticism not as a creed but rather as a method of skeptical, evidence-based inquiry.

Agnosticism Agnosticism is the view or belief that the existence of God, of the divine or the supernatural is unknown or unknowable. Another definition provided is the view that "human reason is incapable of providing sufficient rational grounds to justify either the belief that God exists or the belief that God does not exist." The English biologist Thomas Henry Huxley coined the word agnostic in 1869, and said "It simply means that a man shall not say he knows or believes that which he has no scientific grounds for professing to know or believe." Earlier thinkers, however, had written works that promoted agnostic points of view, such as Sanjaya Belatthaputta, a 5th-century BCE Indian philosopher who expressed agnosticism about any afterlife; and Protagoras, a 5th-century BCE Greek philosopher who expressed agnosticism about the existence of "the gods". Defining agnosticism Being a scientist, above all else, Huxley presented agnosticism as a form of demarcation. A hypothesis with no supporting, objective, testable evidence is not an objective, scientific claim. As such, there would be no way to test said hypotheses, leaving the results inconclusive. His agnosticism was not compatible with forming a belief as to the truth, or falsehood, of the claim at hand. Karl Popper would also describe himself as an agnostic. According to philosopher William L. Rowe, in this strict sense, agnosticism is the view that human reason is incapable of providing sufficient rational grounds to justify either the belief that God exists or the belief that God does not exist. George H. Smith, while admitting that the narrow definition of atheist was the common usage definition of that word, and admitting that the broad definition of agnostic was the common usage definition of that word, promoted broadening the definition of atheist and narrowing the definition of agnostic. Smith rejects agnosticism as a third alternative to theism and atheism and promotes terms such as agnostic atheism (the view of those who do not hold a belief in the existence of any deity, but claim that the existence of a deity is unknown or inherently unknowable) and agnostic theism (the view of those who believe in the existence of a deity(s), but claim that the existence of a deity is unknown or inherently unknowable). Etymology Agnostic () was used by Thomas Henry Huxley in a speech at a meeting of the Metaphysical Society in 1869 to describe his philosophy, which rejects all claims of spiritual or mystical knowledge. Early Christian church leaders used the Greek word gnosis (knowledge) to describe "spiritual knowledge". Agnosticism is not to be confused with religious views opposing the ancient religious movement of Gnosticism in particular; Huxley used the term in a broader, more abstract sense. Huxley identified agnosticism not as a creed but rather as a method of skeptical, evidence-based inquiry.

The term Agnostic is also cognate with the Sanskrit word Ajñasi which translates literally to "not knowable", and relates to the ancient Indian philosophical school of Ajñana, which proposes that it is impossible to obtain knowledge of metaphysical nature or ascertain the truth value of philosophical propositions; and even if knowledge was possible, it is useless and disadvantageous for final salvation. In recent years, scientific literature dealing with neuroscience and psychology has used the word to mean "not knowable". In technical and marketing literature, "agnostic" can also mean independence from some parameters—for example, "platform agnostic" (referring to cross-platform software) or "hardware-agnostic". Qualifying agnosticism Scottish Enlightenment philosopher David Hume contended that meaningful statements about the universe are always qualified by some degree of doubt. He asserted that the fallibility of human beings means that they cannot obtain absolute certainty except in trivial cases where a statement is true by definition (e.g. tautologies such as "all bachelors are unmarried" or "all triangles have three corners"). Types Strong agnosticism (also called "hard", "closed", "strict", or "permanent agnosticism") The view that the question of the existence or nonexistence of a deity or deities, and the nature of ultimate reality is unknowable by reason of our natural inability to verify any experience with anything but another subjective experience. A strong agnostic would say, "I cannot know whether a deity exists or not, and neither can you." Weak agnosticism (also called "soft", "open", "empirical", or "temporal agnosticism") The view that the existence or nonexistence of any deities is currently unknown but is not necessarily unknowable; therefore, one will withhold judgment until evidence, if any, becomes available. A weak agnostic would say, "I don't know whether any deities exist or not, but maybe one day, if there is evidence, we can find something out." Apathetic agnosticism The view that no amount of debate can prove or disprove the existence of one or more deities, and if one or more deities exist, they do not appear to be concerned about the fate of humans. Therefore, their existence has little to no impact on personal human affairs and should be of little interest. An apathetic agnostic would say, "I don't know whether any deity exists or not, and I don't care if any deity exists or not." History Hindu philosophy Throughout the history of Hinduism there has been a strong tradition of philosophic speculation and skepticism. The Rig Veda takes an agnostic view on the fundamental question of how the universe and the gods were created. Nasadiya Sukta (Creation Hymn) in the tenth chapter of the Rig Veda says: Hume, Kant, and Kierkegaard Aristotle, Anselm, Aquinas, Descartes, and Gödel presented arguments attempting to rationally prove the existence of God. The skeptical empiricism of David Hume, the antinomies of Immanuel Kant, and the existential philosophy of Søren Kierkegaard convinced many later philosophers to abandon these attempts, regarding it impossible to construct any unassailable proof for the existence or non-existence of God.

The term Agnostic is also cognate with the Sanskrit word Ajñasi which translates literally to "not knowable", and relates to the ancient Indian philosophical school of Ajñana, which proposes that it is impossible to obtain knowledge of metaphysical nature or ascertain the truth value of philosophical propositions; and even if knowledge was possible, it is useless and disadvantageous for final salvation. In recent years, scientific literature dealing with neuroscience and psychology has used the word to mean "not knowable". In technical and marketing literature, "agnostic" can also mean independence from some parameters—for example, "platform agnostic" (referring to cross-platform software) or "hardware-agnostic". Qualifying agnosticism Scottish Enlightenment philosopher David Hume contended that meaningful statements about the universe are always qualified by some degree of doubt. He asserted that the fallibility of human beings means that they cannot obtain absolute certainty except in trivial cases where a statement is true by definition (e.g. tautologies such as "all bachelors are unmarried" or "all triangles have three corners"). Types Strong agnosticism (also called "hard", "closed", "strict", or "permanent agnosticism") The view that the question of the existence or nonexistence of a deity or deities, and the nature of ultimate reality is unknowable by reason of our natural inability to verify any experience with anything but another subjective experience. A strong agnostic would say, "I cannot know whether a deity exists or not, and neither can you." Weak agnosticism (also called "soft", "open", "empirical", or "temporal agnosticism") The view that the existence or nonexistence of any deities is currently unknown but is not necessarily unknowable; therefore, one will withhold judgment until evidence, if any, becomes available. A weak agnostic would say, "I don't know whether any deities exist or not, but maybe one day, if there is evidence, we can find something out." Apathetic agnosticism The view that no amount of debate can prove or disprove the existence of one or more deities, and if one or more deities exist, they do not appear to be concerned about the fate of humans. Therefore, their existence has little to no impact on personal human affairs and should be of little interest. An apathetic agnostic would say, "I don't know whether any deity exists or not, and I don't care if any deity exists or not." History Hindu philosophy Throughout the history of Hinduism there has been a strong tradition of philosophic speculation and skepticism. The Rig Veda takes an agnostic view on the fundamental question of how the universe and the gods were created. Nasadiya Sukta (Creation Hymn) in the tenth chapter of the Rig Veda says: Hume, Kant, and Kierkegaard Aristotle, Anselm, Aquinas, Descartes, and Gödel presented arguments attempting to rationally prove the existence of God. The skeptical empiricism of David Hume, the antinomies of Immanuel Kant, and the existential philosophy of Søren Kierkegaard convinced many later philosophers to abandon these attempts, regarding it impossible to construct any unassailable proof for the existence or non-existence of God.

The term Agnostic is also cognate with the Sanskrit word Ajñasi which translates literally to "not knowable", and relates to the ancient Indian philosophical school of Ajñana, which proposes that it is impossible to obtain knowledge of metaphysical nature or ascertain the truth value of philosophical propositions; and even if knowledge was possible, it is useless and disadvantageous for final salvation. In recent years, scientific literature dealing with neuroscience and psychology has used the word to mean "not knowable". In technical and marketing literature, "agnostic" can also mean independence from some parameters—for example, "platform agnostic" (referring to cross-platform software) or "hardware-agnostic". Qualifying agnosticism Scottish Enlightenment philosopher David Hume contended that meaningful statements about the universe are always qualified by some degree of doubt. He asserted that the fallibility of human beings means that they cannot obtain absolute certainty except in trivial cases where a statement is true by definition (e.g. tautologies such as "all bachelors are unmarried" or "all triangles have three corners"). Types Strong agnosticism (also called "hard", "closed", "strict", or "permanent agnosticism") The view that the question of the existence or nonexistence of a deity or deities, and the nature of ultimate reality is unknowable by reason of our natural inability to verify any experience with anything but another subjective experience. A strong agnostic would say, "I cannot know whether a deity exists or not, and neither can you." Weak agnosticism (also called "soft", "open", "empirical", or "temporal agnosticism") The view that the existence or nonexistence of any deities is currently unknown but is not necessarily unknowable; therefore, one will withhold judgment until evidence, if any, becomes available. A weak agnostic would say, "I don't know whether any deities exist or not, but maybe one day, if there is evidence, we can find something out." Apathetic agnosticism The view that no amount of debate can prove or disprove the existence of one or more deities, and if one or more deities exist, they do not appear to be concerned about the fate of humans. Therefore, their existence has little to no impact on personal human affairs and should be of little interest. An apathetic agnostic would say, "I don't know whether any deity exists or not, and I don't care if any deity exists or not." History Hindu philosophy Throughout the history of Hinduism there has been a strong tradition of philosophic speculation and skepticism. The Rig Veda takes an agnostic view on the fundamental question of how the universe and the gods were created. Nasadiya Sukta (Creation Hymn) in the tenth chapter of the Rig Veda says: Hume, Kant, and Kierkegaard Aristotle, Anselm, Aquinas, Descartes, and Gödel presented arguments attempting to rationally prove the existence of God. The skeptical empiricism of David Hume, the antinomies of Immanuel Kant, and the existential philosophy of Søren Kierkegaard convinced many later philosophers to abandon these attempts, regarding it impossible to construct any unassailable proof for the existence or non-existence of God.

In his 1844 book, Philosophical Fragments, Kierkegaard writes: Hume was Huxley's favourite philosopher, calling him "the Prince of Agnostics". Diderot wrote to his mistress, telling of a visit by Hume to the Baron D'Holbach, and describing how a word for the position that Huxley would later describe as agnosticism didn't seem to exist, or at least wasn't common knowledge, at the time. United Kingdom Charles Darwin Raised in a religious environment, Charles Darwin (1809–1882) studied to be an Anglican clergyman. While eventually doubting parts of his faith, Darwin continued to help in church affairs, even while avoiding church attendance. Darwin stated that it would be "absurd to doubt that a man might be an ardent theist and an evolutionist". Although reticent about his religious views, in 1879 he wrote that "I have never been an atheist in the sense of denying the existence of a God. – I think that generally ... an agnostic would be the most correct description of my state of mind." Thomas Henry Huxley Agnostic views are as old as philosophical skepticism, but the terms agnostic and agnosticism were created by Huxley (1825–1895) to sum up his thoughts on contemporary developments of metaphysics about the "unconditioned" (William Hamilton) and the "unknowable" (Herbert Spencer). Though Huxley began to use the term "agnostic" in 1869, his opinions had taken shape some time before that date. In a letter of September 23, 1860, to Charles Kingsley, Huxley discussed his views extensively: And again, to the same correspondent, May 6, 1863: Of the origin of the name agnostic to describe this attitude, Huxley gave the following account: In 1889, Huxley wrote:Therefore, although it be, as I believe, demonstrable that we have no real knowledge of the authorship, or of the date of composition of the Gospels, as they have come down to us, and that nothing better than more or less probable guesses can be arrived at on that subject. William Stewart Ross William Stewart Ross (1844–1906) wrote under the name of Saladin. He was associated with Victorian Freethinkers and the organization the British Secular Union. He edited the Secular Review from 1882; it was renamed Agnostic Journal and Eclectic Review and closed in 1907. Ross championed agnosticism in opposition to the atheism of Charles Bradlaugh as an open-ended spiritual exploration. In Why I am an Agnostic (c. 1889) he claims that agnosticism is "the very reverse of atheism". Bertrand Russell Bertrand Russell (1872–1970) declared Why I Am Not a Christian in 1927, a classic statement of agnosticism. He calls upon his readers to "stand on their own two feet and look fair and square at the world with a fearless attitude and a free intelligence". In 1939, Russell gave a lecture on The existence and nature of God, in which he characterized himself as an atheist. He said: However, later in the same lecture, discussing modern non-anthropomorphic concepts of God, Russell states: In Russell's 1947 pamphlet, Am I An Atheist or an Agnostic?

In his 1844 book, Philosophical Fragments, Kierkegaard writes: Hume was Huxley's favourite philosopher, calling him "the Prince of Agnostics". Diderot wrote to his mistress, telling of a visit by Hume to the Baron D'Holbach, and describing how a word for the position that Huxley would later describe as agnosticism didn't seem to exist, or at least wasn't common knowledge, at the time. United Kingdom Charles Darwin Raised in a religious environment, Charles Darwin (1809–1882) studied to be an Anglican clergyman. While eventually doubting parts of his faith, Darwin continued to help in church affairs, even while avoiding church attendance. Darwin stated that it would be "absurd to doubt that a man might be an ardent theist and an evolutionist". Although reticent about his religious views, in 1879 he wrote that "I have never been an atheist in the sense of denying the existence of a God. – I think that generally ... an agnostic would be the most correct description of my state of mind." Thomas Henry Huxley Agnostic views are as old as philosophical skepticism, but the terms agnostic and agnosticism were created by Huxley (1825–1895) to sum up his thoughts on contemporary developments of metaphysics about the "unconditioned" (William Hamilton) and the "unknowable" (Herbert Spencer). Though Huxley began to use the term "agnostic" in 1869, his opinions had taken shape some time before that date. In a letter of September 23, 1860, to Charles Kingsley, Huxley discussed his views extensively: And again, to the same correspondent, May 6, 1863: Of the origin of the name agnostic to describe this attitude, Huxley gave the following account: In 1889, Huxley wrote:Therefore, although it be, as I believe, demonstrable that we have no real knowledge of the authorship, or of the date of composition of the Gospels, as they have come down to us, and that nothing better than more or less probable guesses can be arrived at on that subject. William Stewart Ross William Stewart Ross (1844–1906) wrote under the name of Saladin. He was associated with Victorian Freethinkers and the organization the British Secular Union. He edited the Secular Review from 1882; it was renamed Agnostic Journal and Eclectic Review and closed in 1907. Ross championed agnosticism in opposition to the atheism of Charles Bradlaugh as an open-ended spiritual exploration. In Why I am an Agnostic (c. 1889) he claims that agnosticism is "the very reverse of atheism". Bertrand Russell Bertrand Russell (1872–1970) declared Why I Am Not a Christian in 1927, a classic statement of agnosticism. He calls upon his readers to "stand on their own two feet and look fair and square at the world with a fearless attitude and a free intelligence". In 1939, Russell gave a lecture on The existence and nature of God, in which he characterized himself as an atheist. He said: However, later in the same lecture, discussing modern non-anthropomorphic concepts of God, Russell states: In Russell's 1947 pamphlet, Am I An Atheist or an Agnostic?

In his 1844 book, Philosophical Fragments, Kierkegaard writes: Hume was Huxley's favourite philosopher, calling him "the Prince of Agnostics". Diderot wrote to his mistress, telling of a visit by Hume to the Baron D'Holbach, and describing how a word for the position that Huxley would later describe as agnosticism didn't seem to exist, or at least wasn't common knowledge, at the time. United Kingdom Charles Darwin Raised in a religious environment, Charles Darwin (1809–1882) studied to be an Anglican clergyman. While eventually doubting parts of his faith, Darwin continued to help in church affairs, even while avoiding church attendance. Darwin stated that it would be "absurd to doubt that a man might be an ardent theist and an evolutionist". Although reticent about his religious views, in 1879 he wrote that "I have never been an atheist in the sense of denying the existence of a God. – I think that generally ... an agnostic would be the most correct description of my state of mind." Thomas Henry Huxley Agnostic views are as old as philosophical skepticism, but the terms agnostic and agnosticism were created by Huxley (1825–1895) to sum up his thoughts on contemporary developments of metaphysics about the "unconditioned" (William Hamilton) and the "unknowable" (Herbert Spencer). Though Huxley began to use the term "agnostic" in 1869, his opinions had taken shape some time before that date. In a letter of September 23, 1860, to Charles Kingsley, Huxley discussed his views extensively: And again, to the same correspondent, May 6, 1863: Of the origin of the name agnostic to describe this attitude, Huxley gave the following account: In 1889, Huxley wrote:Therefore, although it be, as I believe, demonstrable that we have no real knowledge of the authorship, or of the date of composition of the Gospels, as they have come down to us, and that nothing better than more or less probable guesses can be arrived at on that subject. William Stewart Ross William Stewart Ross (1844–1906) wrote under the name of Saladin. He was associated with Victorian Freethinkers and the organization the British Secular Union. He edited the Secular Review from 1882; it was renamed Agnostic Journal and Eclectic Review and closed in 1907. Ross championed agnosticism in opposition to the atheism of Charles Bradlaugh as an open-ended spiritual exploration. In Why I am an Agnostic (c. 1889) he claims that agnosticism is "the very reverse of atheism". Bertrand Russell Bertrand Russell (1872–1970) declared Why I Am Not a Christian in 1927, a classic statement of agnosticism. He calls upon his readers to "stand on their own two feet and look fair and square at the world with a fearless attitude and a free intelligence". In 1939, Russell gave a lecture on The existence and nature of God, in which he characterized himself as an atheist. He said: However, later in the same lecture, discussing modern non-anthropomorphic concepts of God, Russell states: In Russell's 1947 pamphlet, Am I An Atheist or an Agnostic?

(subtitled A Plea For Tolerance in the Face of New Dogmas), he ruminates on the problem of what to call himself: In his 1953 essay, What Is An Agnostic? Russell states: Later in the essay, Russell adds: Leslie Weatherhead In 1965, Christian theologian Leslie Weatherhead (1893–1976) published The Christian Agnostic, in which he argues: Although radical and unpalatable to conventional theologians, Weatherhead's agnosticism falls far short of Huxley's, and short even of weak agnosticism: United States Robert G. Ingersoll Robert G. Ingersoll (1833–1899), an Illinois lawyer and politician who evolved into a well-known and sought-after orator in 19th-century America, has been referred to as the "Great Agnostic". In an 1896 lecture titled Why I Am An Agnostic, Ingersoll related why he was an agnostic: In the conclusion of the speech he simply sums up the agnostic position as: In 1885, Ingersoll explained his comparative view of agnosticism and atheism as follows: Bernard Iddings Bell Canon Bernard Iddings Bell (1886–1958), a popular cultural commentator, Episcopal priest, and author, lauded the necessity of agnosticism in Beyond Agnosticism: A Book for Tired Mechanists, calling it the foundation of "all intelligent Christianity." Agnosticism was a temporary mindset in which one rigorously questioned the truths of the age, including the way in which one believed God. His view of Robert Ingersoll and Thomas Paine was that they were not denouncing true Christianity but rather "a gross perversion of it." Part of the misunderstanding stemmed from ignorance of the concepts of God and religion. Historically, a god was any real, perceivable force that ruled the lives of humans and inspired admiration, love, fear, and homage; religion was the practice of it. Ancient peoples worshiped gods with real counterparts, such as Mammon (money and material things), Nabu (rationality), or Ba'al (violent weather); Bell argued that modern peoples were still paying homage—with their lives and their children's lives—to these old gods of wealth, physical appetites, and self-deification. Thus, if one attempted to be agnostic passively, he or she would incidentally join the worship of the world's gods. In Unfashionable Convictions (1931), he criticized the Enlightenment's complete faith in human sensory perception, augmented by scientific instruments, as a means of accurately grasping Reality. Firstly, it was fairly new, an innovation of the Western World, which Aristotle invented and Thomas Aquinas revived among the scientific community. Secondly, the divorce of "pure" science from human experience, as manifested in American Industrialization, had completely altered the environment, often disfiguring it, so as to suggest its insufficiency to human needs. Thirdly, because scientists were constantly producing more data—to the point where no single human could grasp it all at once—it followed that human intelligence was incapable of attaining a complete understanding of universe; therefore, to admit the mysteries of the unobserved universe was to be actually scientific. Bell believed that there were two other ways that humans could perceive and interact with the world.

(subtitled A Plea For Tolerance in the Face of New Dogmas), he ruminates on the problem of what to call himself: In his 1953 essay, What Is An Agnostic? Russell states: Later in the essay, Russell adds: Leslie Weatherhead In 1965, Christian theologian Leslie Weatherhead (1893–1976) published The Christian Agnostic, in which he argues: Although radical and unpalatable to conventional theologians, Weatherhead's agnosticism falls far short of Huxley's, and short even of weak agnosticism: United States Robert G. Ingersoll Robert G. Ingersoll (1833–1899), an Illinois lawyer and politician who evolved into a well-known and sought-after orator in 19th-century America, has been referred to as the "Great Agnostic". In an 1896 lecture titled Why I Am An Agnostic, Ingersoll related why he was an agnostic: In the conclusion of the speech he simply sums up the agnostic position as: In 1885, Ingersoll explained his comparative view of agnosticism and atheism as follows: Bernard Iddings Bell Canon Bernard Iddings Bell (1886–1958), a popular cultural commentator, Episcopal priest, and author, lauded the necessity of agnosticism in Beyond Agnosticism: A Book for Tired Mechanists, calling it the foundation of "all intelligent Christianity." Agnosticism was a temporary mindset in which one rigorously questioned the truths of the age, including the way in which one believed God. His view of Robert Ingersoll and Thomas Paine was that they were not denouncing true Christianity but rather "a gross perversion of it." Part of the misunderstanding stemmed from ignorance of the concepts of God and religion. Historically, a god was any real, perceivable force that ruled the lives of humans and inspired admiration, love, fear, and homage; religion was the practice of it. Ancient peoples worshiped gods with real counterparts, such as Mammon (money and material things), Nabu (rationality), or Ba'al (violent weather); Bell argued that modern peoples were still paying homage—with their lives and their children's lives—to these old gods of wealth, physical appetites, and self-deification. Thus, if one attempted to be agnostic passively, he or she would incidentally join the worship of the world's gods. In Unfashionable Convictions (1931), he criticized the Enlightenment's complete faith in human sensory perception, augmented by scientific instruments, as a means of accurately grasping Reality. Firstly, it was fairly new, an innovation of the Western World, which Aristotle invented and Thomas Aquinas revived among the scientific community. Secondly, the divorce of "pure" science from human experience, as manifested in American Industrialization, had completely altered the environment, often disfiguring it, so as to suggest its insufficiency to human needs. Thirdly, because scientists were constantly producing more data—to the point where no single human could grasp it all at once—it followed that human intelligence was incapable of attaining a complete understanding of universe; therefore, to admit the mysteries of the unobserved universe was to be actually scientific. Bell believed that there were two other ways that humans could perceive and interact with the world.

(subtitled A Plea For Tolerance in the Face of New Dogmas), he ruminates on the problem of what to call himself: In his 1953 essay, What Is An Agnostic? Russell states: Later in the essay, Russell adds: Leslie Weatherhead In 1965, Christian theologian Leslie Weatherhead (1893–1976) published The Christian Agnostic, in which he argues: Although radical and unpalatable to conventional theologians, Weatherhead's agnosticism falls far short of Huxley's, and short even of weak agnosticism: United States Robert G. Ingersoll Robert G. Ingersoll (1833–1899), an Illinois lawyer and politician who evolved into a well-known and sought-after orator in 19th-century America, has been referred to as the "Great Agnostic". In an 1896 lecture titled Why I Am An Agnostic, Ingersoll related why he was an agnostic: In the conclusion of the speech he simply sums up the agnostic position as: In 1885, Ingersoll explained his comparative view of agnosticism and atheism as follows: Bernard Iddings Bell Canon Bernard Iddings Bell (1886–1958), a popular cultural commentator, Episcopal priest, and author, lauded the necessity of agnosticism in Beyond Agnosticism: A Book for Tired Mechanists, calling it the foundation of "all intelligent Christianity." Agnosticism was a temporary mindset in which one rigorously questioned the truths of the age, including the way in which one believed God. His view of Robert Ingersoll and Thomas Paine was that they were not denouncing true Christianity but rather "a gross perversion of it." Part of the misunderstanding stemmed from ignorance of the concepts of God and religion. Historically, a god was any real, perceivable force that ruled the lives of humans and inspired admiration, love, fear, and homage; religion was the practice of it. Ancient peoples worshiped gods with real counterparts, such as Mammon (money and material things), Nabu (rationality), or Ba'al (violent weather); Bell argued that modern peoples were still paying homage—with their lives and their children's lives—to these old gods of wealth, physical appetites, and self-deification. Thus, if one attempted to be agnostic passively, he or she would incidentally join the worship of the world's gods. In Unfashionable Convictions (1931), he criticized the Enlightenment's complete faith in human sensory perception, augmented by scientific instruments, as a means of accurately grasping Reality. Firstly, it was fairly new, an innovation of the Western World, which Aristotle invented and Thomas Aquinas revived among the scientific community. Secondly, the divorce of "pure" science from human experience, as manifested in American Industrialization, had completely altered the environment, often disfiguring it, so as to suggest its insufficiency to human needs. Thirdly, because scientists were constantly producing more data—to the point where no single human could grasp it all at once—it followed that human intelligence was incapable of attaining a complete understanding of universe; therefore, to admit the mysteries of the unobserved universe was to be actually scientific. Bell believed that there were two other ways that humans could perceive and interact with the world.

Artistic experience was how one expressed meaning through speaking, writing, painting, gesturing—any sort of communication which shared insight into a human's inner reality. Mystical experience was how one could "read" people and harmonize with them, being what we commonly call love. In summary, man was a scientist, artist, and lover. Without exercising all three, a person became "lopsided." Bell considered a humanist to be a person who cannot rightly ignore the other ways of knowing. However, humanism, like agnosticism, was also temporal, and would eventually lead to either scientific materialism or theism. He lays out the following thesis: Truth cannot be discovered by reasoning on the evidence of scientific data alone. Modern peoples' dissatisfaction with life is the result of depending on such incomplete data. Our ability to reason is not a way to discover Truth but rather a way to organize our knowledge and experiences somewhat sensibly. Without a full, human perception of the world, one's reason tends to lead them in the wrong direction. Beyond what can be measured with scientific tools, there are other types of perception, such as one's ability know another human through loving. One's loves cannot be dissected and logged in a scientific journal, but we know them far better than we know the surface of the sun. They show us an undefinable reality that is nevertheless intimate and personal, and they reveal qualities lovelier and truer than detached facts can provide. To be religious, in the Christian sense, is to live for the Whole of Reality (God) rather than for a small part (gods). Only by treating this Whole of Reality as a person—good and true and perfect—rather than an impersonal force, can we come closer to the Truth. An ultimate Person can be loved, but a cosmic force cannot. A scientist can only discover peripheral truths, but a lover is able to get at the Truth. There are many reasons to believe in God but they are not sufficient for an agnostic to become a theist. It is not enough to believe in an ancient holy book, even though when it is accurately analyzed without bias, it proves to be more trustworthy and admirable than what we are taught in school. Neither is it enough to realize how probable it is that a personal God would have to show human beings how to live, considering they have so much trouble on their own. Nor is it enough to believe for the reason that, throughout history, millions of people have arrived at this Wholeness of Reality only through religious experience. The aforementioned reasons may warm one toward religion, but they fall short of convincing. However, if one presupposes that God is in fact a knowable, loving person, as an experiment, and then lives according that religion, he or she will suddenly come face to face with experiences previously unknown. One's life becomes full, meaningful, and fearless in the face of death. It does not defy reason but exceeds it.

Artistic experience was how one expressed meaning through speaking, writing, painting, gesturing—any sort of communication which shared insight into a human's inner reality. Mystical experience was how one could "read" people and harmonize with them, being what we commonly call love. In summary, man was a scientist, artist, and lover. Without exercising all three, a person became "lopsided." Bell considered a humanist to be a person who cannot rightly ignore the other ways of knowing. However, humanism, like agnosticism, was also temporal, and would eventually lead to either scientific materialism or theism. He lays out the following thesis: Truth cannot be discovered by reasoning on the evidence of scientific data alone. Modern peoples' dissatisfaction with life is the result of depending on such incomplete data. Our ability to reason is not a way to discover Truth but rather a way to organize our knowledge and experiences somewhat sensibly. Without a full, human perception of the world, one's reason tends to lead them in the wrong direction. Beyond what can be measured with scientific tools, there are other types of perception, such as one's ability know another human through loving. One's loves cannot be dissected and logged in a scientific journal, but we know them far better than we know the surface of the sun. They show us an undefinable reality that is nevertheless intimate and personal, and they reveal qualities lovelier and truer than detached facts can provide. To be religious, in the Christian sense, is to live for the Whole of Reality (God) rather than for a small part (gods). Only by treating this Whole of Reality as a person—good and true and perfect—rather than an impersonal force, can we come closer to the Truth. An ultimate Person can be loved, but a cosmic force cannot. A scientist can only discover peripheral truths, but a lover is able to get at the Truth. There are many reasons to believe in God but they are not sufficient for an agnostic to become a theist. It is not enough to believe in an ancient holy book, even though when it is accurately analyzed without bias, it proves to be more trustworthy and admirable than what we are taught in school. Neither is it enough to realize how probable it is that a personal God would have to show human beings how to live, considering they have so much trouble on their own. Nor is it enough to believe for the reason that, throughout history, millions of people have arrived at this Wholeness of Reality only through religious experience. The aforementioned reasons may warm one toward religion, but they fall short of convincing. However, if one presupposes that God is in fact a knowable, loving person, as an experiment, and then lives according that religion, he or she will suddenly come face to face with experiences previously unknown. One's life becomes full, meaningful, and fearless in the face of death. It does not defy reason but exceeds it.

Artistic experience was how one expressed meaning through speaking, writing, painting, gesturing—any sort of communication which shared insight into a human's inner reality. Mystical experience was how one could "read" people and harmonize with them, being what we commonly call love. In summary, man was a scientist, artist, and lover. Without exercising all three, a person became "lopsided." Bell considered a humanist to be a person who cannot rightly ignore the other ways of knowing. However, humanism, like agnosticism, was also temporal, and would eventually lead to either scientific materialism or theism. He lays out the following thesis: Truth cannot be discovered by reasoning on the evidence of scientific data alone. Modern peoples' dissatisfaction with life is the result of depending on such incomplete data. Our ability to reason is not a way to discover Truth but rather a way to organize our knowledge and experiences somewhat sensibly. Without a full, human perception of the world, one's reason tends to lead them in the wrong direction. Beyond what can be measured with scientific tools, there are other types of perception, such as one's ability know another human through loving. One's loves cannot be dissected and logged in a scientific journal, but we know them far better than we know the surface of the sun. They show us an undefinable reality that is nevertheless intimate and personal, and they reveal qualities lovelier and truer than detached facts can provide. To be religious, in the Christian sense, is to live for the Whole of Reality (God) rather than for a small part (gods). Only by treating this Whole of Reality as a person—good and true and perfect—rather than an impersonal force, can we come closer to the Truth. An ultimate Person can be loved, but a cosmic force cannot. A scientist can only discover peripheral truths, but a lover is able to get at the Truth. There are many reasons to believe in God but they are not sufficient for an agnostic to become a theist. It is not enough to believe in an ancient holy book, even though when it is accurately analyzed without bias, it proves to be more trustworthy and admirable than what we are taught in school. Neither is it enough to realize how probable it is that a personal God would have to show human beings how to live, considering they have so much trouble on their own. Nor is it enough to believe for the reason that, throughout history, millions of people have arrived at this Wholeness of Reality only through religious experience. The aforementioned reasons may warm one toward religion, but they fall short of convincing. However, if one presupposes that God is in fact a knowable, loving person, as an experiment, and then lives according that religion, he or she will suddenly come face to face with experiences previously unknown. One's life becomes full, meaningful, and fearless in the face of death. It does not defy reason but exceeds it.

Because God has been experienced through love, the orders of prayer, fellowship, and devotion now matter. They create order within one's life, continually renewing the "missing piece" that had previously felt lost. They empower one to be compassionate and humble, not small-minded or arrogant. No truth should be denied outright, but all should be questioned. Science reveals an ever-growing vision of our universe that should not be discounted due to bias toward older understandings. Reason is to be trusted and cultivated. To believe in God is not to forego reason or to deny scientific facts, but to step into the unknown and discover the fullness of life. Demographics Demographic research services normally do not differentiate between various types of non-religious respondents, so agnostics are often classified in the same category as atheists or other non-religious people. A 2010 survey published in Encyclopædia Britannica found that the non-religious people or the agnostics made up about 9.6% of the world's population. A November–December 2006 poll published in the Financial Times gives rates for the United States and five European countries. The rates of agnosticism in the United States were at 14%, while the rates of agnosticism in the European countries surveyed were considerably higher: Italy (20%), Spain (30%), Great Britain (35%), Germany (25%), and France (32%). A study conducted by the Pew Research Center found that about 16% of the world's people, the third largest group after Christianity and Islam, have no religious affiliation. According to a 2012 report by the Pew Research Center, agnostics made up 3.3% of the US adult population. In the U.S. Religious Landscape Survey, conducted by the Pew Research Center, 55% of agnostic respondents expressed "a belief in God or a universal spirit", whereas 41% stated that they thought that they felt a tension "being non-religious in a society where most people are religious". According to the 2011 Australian Bureau of Statistics, 22% of Australians have "no religion", a category that includes agnostics. Between 64% and 65% of Japanese and up to 81% of Vietnamese are atheists, agnostics, or do not believe in a god. An official European Union survey reported that 3% of the EU population is unsure about their belief in a god or spirit. Criticism Agnosticism is criticized from a variety of standpoints. Some atheists criticize the use of the term agnosticism as functionally indistinguishable from atheism; this results in frequent criticisms of those who adopt the term as avoiding the atheist label. Theistic Theistic critics claim that agnosticism is impossible in practice, since a person can live only either as if God did not exist (etsi deus non-daretur), or as if God did exist (etsi deus daretur). Christian According to Pope Benedict XVI, strong agnosticism in particular contradicts itself in affirming the power of reason to know scientific truth. He blames the exclusion of reasoning from religion and ethics for dangerous pathologies such as crimes against humanity and ecological disasters.

Because God has been experienced through love, the orders of prayer, fellowship, and devotion now matter. They create order within one's life, continually renewing the "missing piece" that had previously felt lost. They empower one to be compassionate and humble, not small-minded or arrogant. No truth should be denied outright, but all should be questioned. Science reveals an ever-growing vision of our universe that should not be discounted due to bias toward older understandings. Reason is to be trusted and cultivated. To believe in God is not to forego reason or to deny scientific facts, but to step into the unknown and discover the fullness of life. Demographics Demographic research services normally do not differentiate between various types of non-religious respondents, so agnostics are often classified in the same category as atheists or other non-religious people. A 2010 survey published in Encyclopædia Britannica found that the non-religious people or the agnostics made up about 9.6% of the world's population. A November–December 2006 poll published in the Financial Times gives rates for the United States and five European countries. The rates of agnosticism in the United States were at 14%, while the rates of agnosticism in the European countries surveyed were considerably higher: Italy (20%), Spain (30%), Great Britain (35%), Germany (25%), and France (32%). A study conducted by the Pew Research Center found that about 16% of the world's people, the third largest group after Christianity and Islam, have no religious affiliation. According to a 2012 report by the Pew Research Center, agnostics made up 3.3% of the US adult population. In the U.S. Religious Landscape Survey, conducted by the Pew Research Center, 55% of agnostic respondents expressed "a belief in God or a universal spirit", whereas 41% stated that they thought that they felt a tension "being non-religious in a society where most people are religious". According to the 2011 Australian Bureau of Statistics, 22% of Australians have "no religion", a category that includes agnostics. Between 64% and 65% of Japanese and up to 81% of Vietnamese are atheists, agnostics, or do not believe in a god. An official European Union survey reported that 3% of the EU population is unsure about their belief in a god or spirit. Criticism Agnosticism is criticized from a variety of standpoints. Some atheists criticize the use of the term agnosticism as functionally indistinguishable from atheism; this results in frequent criticisms of those who adopt the term as avoiding the atheist label. Theistic Theistic critics claim that agnosticism is impossible in practice, since a person can live only either as if God did not exist (etsi deus non-daretur), or as if God did exist (etsi deus daretur). Christian According to Pope Benedict XVI, strong agnosticism in particular contradicts itself in affirming the power of reason to know scientific truth. He blames the exclusion of reasoning from religion and ethics for dangerous pathologies such as crimes against humanity and ecological disasters.

Because God has been experienced through love, the orders of prayer, fellowship, and devotion now matter. They create order within one's life, continually renewing the "missing piece" that had previously felt lost. They empower one to be compassionate and humble, not small-minded or arrogant. No truth should be denied outright, but all should be questioned. Science reveals an ever-growing vision of our universe that should not be discounted due to bias toward older understandings. Reason is to be trusted and cultivated. To believe in God is not to forego reason or to deny scientific facts, but to step into the unknown and discover the fullness of life. Demographics Demographic research services normally do not differentiate between various types of non-religious respondents, so agnostics are often classified in the same category as atheists or other non-religious people. A 2010 survey published in Encyclopædia Britannica found that the non-religious people or the agnostics made up about 9.6% of the world's population. A November–December 2006 poll published in the Financial Times gives rates for the United States and five European countries. The rates of agnosticism in the United States were at 14%, while the rates of agnosticism in the European countries surveyed were considerably higher: Italy (20%), Spain (30%), Great Britain (35%), Germany (25%), and France (32%). A study conducted by the Pew Research Center found that about 16% of the world's people, the third largest group after Christianity and Islam, have no religious affiliation. According to a 2012 report by the Pew Research Center, agnostics made up 3.3% of the US adult population. In the U.S. Religious Landscape Survey, conducted by the Pew Research Center, 55% of agnostic respondents expressed "a belief in God or a universal spirit", whereas 41% stated that they thought that they felt a tension "being non-religious in a society where most people are religious". According to the 2011 Australian Bureau of Statistics, 22% of Australians have "no religion", a category that includes agnostics. Between 64% and 65% of Japanese and up to 81% of Vietnamese are atheists, agnostics, or do not believe in a god. An official European Union survey reported that 3% of the EU population is unsure about their belief in a god or spirit. Criticism Agnosticism is criticized from a variety of standpoints. Some atheists criticize the use of the term agnosticism as functionally indistinguishable from atheism; this results in frequent criticisms of those who adopt the term as avoiding the atheist label. Theistic Theistic critics claim that agnosticism is impossible in practice, since a person can live only either as if God did not exist (etsi deus non-daretur), or as if God did exist (etsi deus daretur). Christian According to Pope Benedict XVI, strong agnosticism in particular contradicts itself in affirming the power of reason to know scientific truth. He blames the exclusion of reasoning from religion and ethics for dangerous pathologies such as crimes against humanity and ecological disasters.

"Agnosticism", said Benedict, "is always the fruit of a refusal of that knowledge which is in fact offered to man ... The knowledge of God has always existed". He asserted that agnosticism is a choice of comfort, pride, dominion, and utility over truth, and is opposed by the following attitudes: the keenest self-criticism, humble listening to the whole of existence, the persistent patience and self-correction of the scientific method, a readiness to be purified by the truth. The Catholic Church sees merit in examining what it calls "partial agnosticism", specifically those systems that "do not aim at constructing a complete philosophy of the unknowable, but at excluding special kinds of truth, notably religious, from the domain of knowledge". However, the Church is historically opposed to a full denial of the capacity of human reason to know God. The Council of the Vatican declares, "God, the beginning and end of all, can, by the natural light of human reason, be known with certainty from the works of creation". Blaise Pascal argued that even if there were truly no evidence for God, agnostics should consider what is now known as Pascal's Wager: the infinite expected value of acknowledging God is always greater than the finite expected value of not acknowledging his existence, and thus it is a safer "bet" to choose God. Atheistic According to Richard Dawkins, a distinction between agnosticism and atheism is unwieldy and depends on how close to zero a person is willing to rate the probability of existence for any given god-like entity. About himself, Dawkins continues, "I am agnostic only to the extent that I am agnostic about fairies at the bottom of the garden." Dawkins also identifies two categories of agnostics; "Temporary Agnostics in Practice" (TAPs), and "Permanent Agnostics in Principle" (PAPs). He states that "agnosticism about the existence of God belongs firmly in the temporary or TAP category. Either he exists or he doesn't. It is a scientific question; one day we may know the answer, and meanwhile we can say something pretty strong about the probability" and considers PAP a "deeply inescapable kind of fence-sitting". Ignosticism A related concept is ignosticism, the view that a coherent definition of a deity must be put forward before the question of the existence of a deity can be meaningfully discussed. If the chosen definition is not coherent, the ignostic holds the noncognitivist view that the existence of a deity is meaningless or empirically untestable. A. J. Ayer, Theodore Drange, and other philosophers see both atheism and agnosticism as incompatible with ignosticism on the grounds that atheism and agnosticism accept the statement "a deity exists" as a meaningful proposition that can be argued for or against. See also References Further reading Alexander, Nathan G. "An Atheist with a Tall Hat On: The Forgotten History of Agnosticism." The Humanist, February 19, 2019. Annan, Noel. Leslie Stephen: The Godless Victorian (U of Chicago Press, 1984) Cockshut, A.O.J. The Unbelievers, English Thought, 1840–1890 (1966). Dawkins, Richard.

"Agnosticism", said Benedict, "is always the fruit of a refusal of that knowledge which is in fact offered to man ... The knowledge of God has always existed". He asserted that agnosticism is a choice of comfort, pride, dominion, and utility over truth, and is opposed by the following attitudes: the keenest self-criticism, humble listening to the whole of existence, the persistent patience and self-correction of the scientific method, a readiness to be purified by the truth. The Catholic Church sees merit in examining what it calls "partial agnosticism", specifically those systems that "do not aim at constructing a complete philosophy of the unknowable, but at excluding special kinds of truth, notably religious, from the domain of knowledge". However, the Church is historically opposed to a full denial of the capacity of human reason to know God. The Council of the Vatican declares, "God, the beginning and end of all, can, by the natural light of human reason, be known with certainty from the works of creation". Blaise Pascal argued that even if there were truly no evidence for God, agnostics should consider what is now known as Pascal's Wager: the infinite expected value of acknowledging God is always greater than the finite expected value of not acknowledging his existence, and thus it is a safer "bet" to choose God. Atheistic According to Richard Dawkins, a distinction between agnosticism and atheism is unwieldy and depends on how close to zero a person is willing to rate the probability of existence for any given god-like entity. About himself, Dawkins continues, "I am agnostic only to the extent that I am agnostic about fairies at the bottom of the garden." Dawkins also identifies two categories of agnostics; "Temporary Agnostics in Practice" (TAPs), and "Permanent Agnostics in Principle" (PAPs). He states that "agnosticism about the existence of God belongs firmly in the temporary or TAP category. Either he exists or he doesn't. It is a scientific question; one day we may know the answer, and meanwhile we can say something pretty strong about the probability" and considers PAP a "deeply inescapable kind of fence-sitting". Ignosticism A related concept is ignosticism, the view that a coherent definition of a deity must be put forward before the question of the existence of a deity can be meaningfully discussed. If the chosen definition is not coherent, the ignostic holds the noncognitivist view that the existence of a deity is meaningless or empirically untestable. A. J. Ayer, Theodore Drange, and other philosophers see both atheism and agnosticism as incompatible with ignosticism on the grounds that atheism and agnosticism accept the statement "a deity exists" as a meaningful proposition that can be argued for or against. See also References Further reading Alexander, Nathan G. "An Atheist with a Tall Hat On: The Forgotten History of Agnosticism." The Humanist, February 19, 2019. Annan, Noel. Leslie Stephen: The Godless Victorian (U of Chicago Press, 1984) Cockshut, A.O.J. The Unbelievers, English Thought, 1840–1890 (1966). Dawkins, Richard.

"Agnosticism", said Benedict, "is always the fruit of a refusal of that knowledge which is in fact offered to man ... The knowledge of God has always existed". He asserted that agnosticism is a choice of comfort, pride, dominion, and utility over truth, and is opposed by the following attitudes: the keenest self-criticism, humble listening to the whole of existence, the persistent patience and self-correction of the scientific method, a readiness to be purified by the truth. The Catholic Church sees merit in examining what it calls "partial agnosticism", specifically those systems that "do not aim at constructing a complete philosophy of the unknowable, but at excluding special kinds of truth, notably religious, from the domain of knowledge". However, the Church is historically opposed to a full denial of the capacity of human reason to know God. The Council of the Vatican declares, "God, the beginning and end of all, can, by the natural light of human reason, be known with certainty from the works of creation". Blaise Pascal argued that even if there were truly no evidence for God, agnostics should consider what is now known as Pascal's Wager: the infinite expected value of acknowledging God is always greater than the finite expected value of not acknowledging his existence, and thus it is a safer "bet" to choose God. Atheistic According to Richard Dawkins, a distinction between agnosticism and atheism is unwieldy and depends on how close to zero a person is willing to rate the probability of existence for any given god-like entity. About himself, Dawkins continues, "I am agnostic only to the extent that I am agnostic about fairies at the bottom of the garden." Dawkins also identifies two categories of agnostics; "Temporary Agnostics in Practice" (TAPs), and "Permanent Agnostics in Principle" (PAPs). He states that "agnosticism about the existence of God belongs firmly in the temporary or TAP category. Either he exists or he doesn't. It is a scientific question; one day we may know the answer, and meanwhile we can say something pretty strong about the probability" and considers PAP a "deeply inescapable kind of fence-sitting". Ignosticism A related concept is ignosticism, the view that a coherent definition of a deity must be put forward before the question of the existence of a deity can be meaningfully discussed. If the chosen definition is not coherent, the ignostic holds the noncognitivist view that the existence of a deity is meaningless or empirically untestable. A. J. Ayer, Theodore Drange, and other philosophers see both atheism and agnosticism as incompatible with ignosticism on the grounds that atheism and agnosticism accept the statement "a deity exists" as a meaningful proposition that can be argued for or against. See also References Further reading Alexander, Nathan G. "An Atheist with a Tall Hat On: The Forgotten History of Agnosticism." The Humanist, February 19, 2019. Annan, Noel. Leslie Stephen: The Godless Victorian (U of Chicago Press, 1984) Cockshut, A.O.J. The Unbelievers, English Thought, 1840–1890 (1966). Dawkins, Richard.

"The poverty of agnosticism", in The God Delusion, Black Swan, 2007 (). Lightman, Bernard. The Origins of Agnosticism (1987). Royle, Edward. Radicals, Secularists, and Republicans: Popular Freethought in Britain, 1866–1915 (Manchester UP, 1980). External links Albert Einstein on Religion Shapell Manuscript Foundation Why I Am An Agnostic by Robert G. Ingersoll, [1896]. Dictionary of the History of Ideas: Agnosticism Agnosticism from INTERS – Interdisciplinary Encyclopedia of Religion and Science Agnosticism – from ReligiousTolerance.org What do Agnostics Believe? – A Jewish perspective Fides et Ratio  – the relationship between faith and reason Karol Wojtyla [1998] The Natural Religion by Dr Brendan Connolly, 2008 Epistemological theories Philosophy of religion Skepticism Irreligion Doubt

"The poverty of agnosticism", in The God Delusion, Black Swan, 2007 (). Lightman, Bernard. The Origins of Agnosticism (1987). Royle, Edward. Radicals, Secularists, and Republicans: Popular Freethought in Britain, 1866–1915 (Manchester UP, 1980). External links Albert Einstein on Religion Shapell Manuscript Foundation Why I Am An Agnostic by Robert G. Ingersoll, [1896]. Dictionary of the History of Ideas: Agnosticism Agnosticism from INTERS – Interdisciplinary Encyclopedia of Religion and Science Agnosticism – from ReligiousTolerance.org What do Agnostics Believe? – A Jewish perspective Fides et Ratio  – the relationship between faith and reason Karol Wojtyla [1998] The Natural Religion by Dr Brendan Connolly, 2008 Epistemological theories Philosophy of religion Skepticism Irreligion Doubt

"The poverty of agnosticism", in The God Delusion, Black Swan, 2007 (). Lightman, Bernard. The Origins of Agnosticism (1987). Royle, Edward. Radicals, Secularists, and Republicans: Popular Freethought in Britain, 1866–1915 (Manchester UP, 1980). External links Albert Einstein on Religion Shapell Manuscript Foundation Why I Am An Agnostic by Robert G. Ingersoll, [1896]. Dictionary of the History of Ideas: Agnosticism Agnosticism from INTERS – Interdisciplinary Encyclopedia of Religion and Science Agnosticism – from ReligiousTolerance.org What do Agnostics Believe? – A Jewish perspective Fides et Ratio  – the relationship between faith and reason Karol Wojtyla [1998] The Natural Religion by Dr Brendan Connolly, 2008 Epistemological theories Philosophy of religion Skepticism Irreligion Doubt

Argon Argon is a chemical element with the symbol Ar and atomic number 18. It is in group 18 of the periodic table and is a noble gas. Argon is the third-most abundant gas in the Earth's atmosphere, at 0.934% (9340 ppmv). It is more than twice as abundant as water vapor (which averages about 4000 ppmv, but varies greatly), 23 times as abundant as carbon dioxide (400 ppmv), and more than 500 times as abundant as neon (18 ppmv). Argon is the most abundant noble gas in Earth's crust, comprising 0.00015% of the crust. Nearly all of the argon in the Earth's atmosphere is radiogenic argon-40, derived from the decay of potassium-40 in the Earth's crust. In the universe, argon-36 is by far the most common argon isotope, as it is the most easily produced by stellar nucleosynthesis in supernovas. The name "argon" is derived from the Greek word , neuter singular form of meaning 'lazy' or 'inactive', as a reference to the fact that the element undergoes almost no chemical reactions. The complete octet (eight electrons) in the outer atomic shell makes argon stable and resistant to bonding with other elements. Its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990. Argon is extracted industrially by the fractional distillation of liquid air. Argon is mostly used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily unreactive substances become reactive; for example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. Argon is also used in incandescent, fluorescent lighting, and other gas-discharge tubes. Argon makes a distinctive blue-green gas laser. Argon is also used in fluorescent glow starters. Characteristics Argon has approximately the same solubility in water as oxygen and is 2.5 times more soluble in water than nitrogen. Argon is colorless, odorless, nonflammable and nontoxic as a solid, liquid or gas. Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature. Although argon is a noble gas, it can form some compounds under various extreme conditions. Argon fluorohydride (HArF), a compound of argon with fluorine and hydrogen that is stable below , has been demonstrated. Although the neutral ground-state chemical compounds of argon are presently limited to HArF, argon can form clathrates with water when atoms of argon are trapped in a lattice of water molecules. Ions, such as , and excited-state complexes, such as ArF, have been demonstrated. Theoretical calculation predicts several more argon compounds that should be stable but have not yet been synthesized. History Argon (Greek , neuter singular form of meaning "lazy" or "inactive") is named in reference to its chemical inactivity. This chemical property of this first noble gas to be discovered impressed the namers. An unreactive gas was suspected to be a component of air by Henry Cavendish in 1785.

Argon Argon is a chemical element with the symbol Ar and atomic number 18. It is in group 18 of the periodic table and is a noble gas. Argon is the third-most abundant gas in the Earth's atmosphere, at 0.934% (9340 ppmv). It is more than twice as abundant as water vapor (which averages about 4000 ppmv, but varies greatly), 23 times as abundant as carbon dioxide (400 ppmv), and more than 500 times as abundant as neon (18 ppmv). Argon is the most abundant noble gas in Earth's crust, comprising 0.00015% of the crust. Nearly all of the argon in the Earth's atmosphere is radiogenic argon-40, derived from the decay of potassium-40 in the Earth's crust. In the universe, argon-36 is by far the most common argon isotope, as it is the most easily produced by stellar nucleosynthesis in supernovas. The name "argon" is derived from the Greek word , neuter singular form of meaning 'lazy' or 'inactive', as a reference to the fact that the element undergoes almost no chemical reactions. The complete octet (eight electrons) in the outer atomic shell makes argon stable and resistant to bonding with other elements. Its triple point temperature of 83.8058 K is a defining fixed point in the International Temperature Scale of 1990. Argon is extracted industrially by the fractional distillation of liquid air. Argon is mostly used as an inert shielding gas in welding and other high-temperature industrial processes where ordinarily unreactive substances become reactive; for example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. Argon is also used in incandescent, fluorescent lighting, and other gas-discharge tubes. Argon makes a distinctive blue-green gas laser. Argon is also used in fluorescent glow starters. Characteristics Argon has approximately the same solubility in water as oxygen and is 2.5 times more soluble in water than nitrogen. Argon is colorless, odorless, nonflammable and nontoxic as a solid, liquid or gas. Argon is chemically inert under most conditions and forms no confirmed stable compounds at room temperature. Although argon is a noble gas, it can form some compounds under various extreme conditions. Argon fluorohydride (HArF), a compound of argon with fluorine and hydrogen that is stable below , has been demonstrated. Although the neutral ground-state chemical compounds of argon are presently limited to HArF, argon can form clathrates with water when atoms of argon are trapped in a lattice of water molecules. Ions, such as , and excited-state complexes, such as ArF, have been demonstrated. Theoretical calculation predicts several more argon compounds that should be stable but have not yet been synthesized. History Argon (Greek , neuter singular form of meaning "lazy" or "inactive") is named in reference to its chemical inactivity. This chemical property of this first noble gas to be discovered impressed the namers. An unreactive gas was suspected to be a component of air by Henry Cavendish in 1785.

Argon was first isolated from air in 1894 by Lord Rayleigh and Sir William Ramsay at University College London by removing oxygen, carbon dioxide, water, and nitrogen from a sample of clean air. They first accomplished this by replicating an experiment of Henry Cavendish's. They trapped a mixture of atmospheric air with additional oxygen in a test-tube (A) upside-down over a large quantity of dilute alkali solution (B), which in Cavendish's original experiment was potassium hydroxide, and conveyed a current through wires insulated by U-shaped glass tubes (CC) which sealed around the platinum wire electrodes, leaving the ends of the wires (DD) exposed to the gas and insulated from the alkali solution. The arc was powered by a battery of five Grove cells and a Ruhmkorff coil of medium size. The alkali absorbed the oxides of nitrogen produced by the arc and also carbon dioxide. They operated the arc until no more reduction of volume of the gas could be seen for at least an hour or two and the spectral lines of nitrogen disappeared when the gas was examined. The remaining oxygen was reacted with alkaline pyrogallate to leave behind an apparently non-reactive gas which they called argon. Before isolating the gas, they had determined that nitrogen produced from chemical compounds was 0.5% lighter than nitrogen from the atmosphere. The difference was slight, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall and W. N. Hartley. Each observed new lines in the emission spectrum of air that did not match known elements. Until 1957, the symbol for argon was "A", but now it is "Ar". Occurrence Argon constitutes 0.934% by volume and 1.288% by mass of the Earth's atmosphere. Air is the primary industrial source of purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon. The Earth's crust and seawater contain 1.2 ppm and 0.45 ppm of argon, respectively. Isotopes The main isotopes of argon found on Earth are (99.6%), (0.34%), and (0.06%). Naturally occurring , with a half-life of 1.25 years, decays to stable (11.2%) by electron capture or positron emission, and also to stable (88.8%) by beta decay. These properties and ratios are used to determine the age of rocks by K–Ar dating. In the Earth's atmosphere, is made by cosmic ray activity, primarily by neutron capture of followed by two-neutron emission. In the subsurface environment, it is also produced through neutron capture by , followed by proton emission. is created from the neutron capture by followed by an alpha particle emission as a result of subsurface nuclear explosions. It has a half-life of 35 days. Between locations in the Solar System, the isotopic composition of argon varies greatly.

Argon was first isolated from air in 1894 by Lord Rayleigh and Sir William Ramsay at University College London by removing oxygen, carbon dioxide, water, and nitrogen from a sample of clean air. They first accomplished this by replicating an experiment of Henry Cavendish's. They trapped a mixture of atmospheric air with additional oxygen in a test-tube (A) upside-down over a large quantity of dilute alkali solution (B), which in Cavendish's original experiment was potassium hydroxide, and conveyed a current through wires insulated by U-shaped glass tubes (CC) which sealed around the platinum wire electrodes, leaving the ends of the wires (DD) exposed to the gas and insulated from the alkali solution. The arc was powered by a battery of five Grove cells and a Ruhmkorff coil of medium size. The alkali absorbed the oxides of nitrogen produced by the arc and also carbon dioxide. They operated the arc until no more reduction of volume of the gas could be seen for at least an hour or two and the spectral lines of nitrogen disappeared when the gas was examined. The remaining oxygen was reacted with alkaline pyrogallate to leave behind an apparently non-reactive gas which they called argon. Before isolating the gas, they had determined that nitrogen produced from chemical compounds was 0.5% lighter than nitrogen from the atmosphere. The difference was slight, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall and W. N. Hartley. Each observed new lines in the emission spectrum of air that did not match known elements. Until 1957, the symbol for argon was "A", but now it is "Ar". Occurrence Argon constitutes 0.934% by volume and 1.288% by mass of the Earth's atmosphere. Air is the primary industrial source of purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon. The Earth's crust and seawater contain 1.2 ppm and 0.45 ppm of argon, respectively. Isotopes The main isotopes of argon found on Earth are (99.6%), (0.34%), and (0.06%). Naturally occurring , with a half-life of 1.25 years, decays to stable (11.2%) by electron capture or positron emission, and also to stable (88.8%) by beta decay. These properties and ratios are used to determine the age of rocks by K–Ar dating. In the Earth's atmosphere, is made by cosmic ray activity, primarily by neutron capture of followed by two-neutron emission. In the subsurface environment, it is also produced through neutron capture by , followed by proton emission. is created from the neutron capture by followed by an alpha particle emission as a result of subsurface nuclear explosions. It has a half-life of 35 days. Between locations in the Solar System, the isotopic composition of argon varies greatly.

Argon was first isolated from air in 1894 by Lord Rayleigh and Sir William Ramsay at University College London by removing oxygen, carbon dioxide, water, and nitrogen from a sample of clean air. They first accomplished this by replicating an experiment of Henry Cavendish's. They trapped a mixture of atmospheric air with additional oxygen in a test-tube (A) upside-down over a large quantity of dilute alkali solution (B), which in Cavendish's original experiment was potassium hydroxide, and conveyed a current through wires insulated by U-shaped glass tubes (CC) which sealed around the platinum wire electrodes, leaving the ends of the wires (DD) exposed to the gas and insulated from the alkali solution. The arc was powered by a battery of five Grove cells and a Ruhmkorff coil of medium size. The alkali absorbed the oxides of nitrogen produced by the arc and also carbon dioxide. They operated the arc until no more reduction of volume of the gas could be seen for at least an hour or two and the spectral lines of nitrogen disappeared when the gas was examined. The remaining oxygen was reacted with alkaline pyrogallate to leave behind an apparently non-reactive gas which they called argon. Before isolating the gas, they had determined that nitrogen produced from chemical compounds was 0.5% lighter than nitrogen from the atmosphere. The difference was slight, but it was important enough to attract their attention for many months. They concluded that there was another gas in the air mixed in with the nitrogen. Argon was also encountered in 1882 through independent research of H. F. Newall and W. N. Hartley. Each observed new lines in the emission spectrum of air that did not match known elements. Until 1957, the symbol for argon was "A", but now it is "Ar". Occurrence Argon constitutes 0.934% by volume and 1.288% by mass of the Earth's atmosphere. Air is the primary industrial source of purified argon products. Argon is isolated from air by fractionation, most commonly by cryogenic fractional distillation, a process that also produces purified nitrogen, oxygen, neon, krypton and xenon. The Earth's crust and seawater contain 1.2 ppm and 0.45 ppm of argon, respectively. Isotopes The main isotopes of argon found on Earth are (99.6%), (0.34%), and (0.06%). Naturally occurring , with a half-life of 1.25 years, decays to stable (11.2%) by electron capture or positron emission, and also to stable (88.8%) by beta decay. These properties and ratios are used to determine the age of rocks by K–Ar dating. In the Earth's atmosphere, is made by cosmic ray activity, primarily by neutron capture of followed by two-neutron emission. In the subsurface environment, it is also produced through neutron capture by , followed by proton emission. is created from the neutron capture by followed by an alpha particle emission as a result of subsurface nuclear explosions. It has a half-life of 35 days. Between locations in the Solar System, the isotopic composition of argon varies greatly.

Where the major source of argon is the decay of in rocks, will be the dominant isotope, as it is on Earth. Argon produced directly by stellar nucleosynthesis is dominated by the alpha-process nuclide . Correspondingly, solar argon contains 84.6% (according to solar wind measurements), and the ratio of the three isotopes 36Ar : 38Ar : 40Ar in the atmospheres of the outer planets is 8400 : 1600 : 1. This contrasts with the low abundance of primordial in Earth's atmosphere, which is only 31.5 ppmv (= 9340 ppmv × 0.337%), comparable with that of neon (18.18 ppmv) on Earth and with interplanetary gasses, measured by probes. The atmospheres of Mars, Mercury and Titan (the largest moon of Saturn) contain argon, predominantly as , and its content may be as high as 1.93% (Mars). The predominance of radiogenic is the reason the standard atomic weight of terrestrial argon is greater than that of the next element, potassium, a fact that was puzzling when argon was discovered. Mendeleev positioned the elements on his periodic table in order of atomic weight, but the inertness of argon suggested a placement before the reactive alkali metal. Henry Moseley later solved this problem by showing that the periodic table is actually arranged in order of atomic number (see History of the periodic table). Compounds Argon's complete octet of electrons indicates full s and p subshells. This full valence shell makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. The first argon compound with tungsten pentacarbonyl, W(CO)5Ar, was isolated in 1975. However it was not widely recognised at that time. In August 2000, another argon compound, argon fluorohydride (HArF), was formed by researchers at the University of Helsinki, by shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride with caesium iodide. This discovery caused the recognition that argon could form weakly bound compounds, even though it was not the first. It is stable up to 17 kelvins (−256 °C). The metastable dication, which is valence-isoelectronic with carbonyl fluoride and phosgene, was observed in 2010. Argon-36, in the form of argon hydride (argonium) ions, has been detected in interstellar medium associated with the Crab Nebula supernova; this was the first noble-gas molecule detected in outer space. Solid argon hydride (Ar(H2)2) has the same crystal structure as the MgZn2 Laves phase. It forms at pressures between 4.3 and 220 GPa, though Raman measurements suggest that the H2 molecules in Ar(H2)2 dissociate above 175 GPa. Production Industrial Argon is extracted industrially by the fractional distillation of liquid air in a cryogenic air separation unit; a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K, and liquid oxygen, which boils at 90.2 K. About 700,000 tonnes of argon are produced worldwide every year.

Where the major source of argon is the decay of in rocks, will be the dominant isotope, as it is on Earth. Argon produced directly by stellar nucleosynthesis is dominated by the alpha-process nuclide . Correspondingly, solar argon contains 84.6% (according to solar wind measurements), and the ratio of the three isotopes 36Ar : 38Ar : 40Ar in the atmospheres of the outer planets is 8400 : 1600 : 1. This contrasts with the low abundance of primordial in Earth's atmosphere, which is only 31.5 ppmv (= 9340 ppmv × 0.337%), comparable with that of neon (18.18 ppmv) on Earth and with interplanetary gasses, measured by probes. The atmospheres of Mars, Mercury and Titan (the largest moon of Saturn) contain argon, predominantly as , and its content may be as high as 1.93% (Mars). The predominance of radiogenic is the reason the standard atomic weight of terrestrial argon is greater than that of the next element, potassium, a fact that was puzzling when argon was discovered. Mendeleev positioned the elements on his periodic table in order of atomic weight, but the inertness of argon suggested a placement before the reactive alkali metal. Henry Moseley later solved this problem by showing that the periodic table is actually arranged in order of atomic number (see History of the periodic table). Compounds Argon's complete octet of electrons indicates full s and p subshells. This full valence shell makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. The first argon compound with tungsten pentacarbonyl, W(CO)5Ar, was isolated in 1975. However it was not widely recognised at that time. In August 2000, another argon compound, argon fluorohydride (HArF), was formed by researchers at the University of Helsinki, by shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride with caesium iodide. This discovery caused the recognition that argon could form weakly bound compounds, even though it was not the first. It is stable up to 17 kelvins (−256 °C). The metastable dication, which is valence-isoelectronic with carbonyl fluoride and phosgene, was observed in 2010. Argon-36, in the form of argon hydride (argonium) ions, has been detected in interstellar medium associated with the Crab Nebula supernova; this was the first noble-gas molecule detected in outer space. Solid argon hydride (Ar(H2)2) has the same crystal structure as the MgZn2 Laves phase. It forms at pressures between 4.3 and 220 GPa, though Raman measurements suggest that the H2 molecules in Ar(H2)2 dissociate above 175 GPa. Production Industrial Argon is extracted industrially by the fractional distillation of liquid air in a cryogenic air separation unit; a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K, and liquid oxygen, which boils at 90.2 K. About 700,000 tonnes of argon are produced worldwide every year.

Where the major source of argon is the decay of in rocks, will be the dominant isotope, as it is on Earth. Argon produced directly by stellar nucleosynthesis is dominated by the alpha-process nuclide . Correspondingly, solar argon contains 84.6% (according to solar wind measurements), and the ratio of the three isotopes 36Ar : 38Ar : 40Ar in the atmospheres of the outer planets is 8400 : 1600 : 1. This contrasts with the low abundance of primordial in Earth's atmosphere, which is only 31.5 ppmv (= 9340 ppmv × 0.337%), comparable with that of neon (18.18 ppmv) on Earth and with interplanetary gasses, measured by probes. The atmospheres of Mars, Mercury and Titan (the largest moon of Saturn) contain argon, predominantly as , and its content may be as high as 1.93% (Mars). The predominance of radiogenic is the reason the standard atomic weight of terrestrial argon is greater than that of the next element, potassium, a fact that was puzzling when argon was discovered. Mendeleev positioned the elements on his periodic table in order of atomic weight, but the inertness of argon suggested a placement before the reactive alkali metal. Henry Moseley later solved this problem by showing that the periodic table is actually arranged in order of atomic number (see History of the periodic table). Compounds Argon's complete octet of electrons indicates full s and p subshells. This full valence shell makes argon very stable and extremely resistant to bonding with other elements. Before 1962, argon and the other noble gases were considered to be chemically inert and unable to form compounds; however, compounds of the heavier noble gases have since been synthesized. The first argon compound with tungsten pentacarbonyl, W(CO)5Ar, was isolated in 1975. However it was not widely recognised at that time. In August 2000, another argon compound, argon fluorohydride (HArF), was formed by researchers at the University of Helsinki, by shining ultraviolet light onto frozen argon containing a small amount of hydrogen fluoride with caesium iodide. This discovery caused the recognition that argon could form weakly bound compounds, even though it was not the first. It is stable up to 17 kelvins (−256 °C). The metastable dication, which is valence-isoelectronic with carbonyl fluoride and phosgene, was observed in 2010. Argon-36, in the form of argon hydride (argonium) ions, has been detected in interstellar medium associated with the Crab Nebula supernova; this was the first noble-gas molecule detected in outer space. Solid argon hydride (Ar(H2)2) has the same crystal structure as the MgZn2 Laves phase. It forms at pressures between 4.3 and 220 GPa, though Raman measurements suggest that the H2 molecules in Ar(H2)2 dissociate above 175 GPa. Production Industrial Argon is extracted industrially by the fractional distillation of liquid air in a cryogenic air separation unit; a process that separates liquid nitrogen, which boils at 77.3 K, from argon, which boils at 87.3 K, and liquid oxygen, which boils at 90.2 K. About 700,000 tonnes of argon are produced worldwide every year.

In radioactive decays 40Ar, the most abundant isotope of argon, is produced by the decay of 40K with a half-life of 1.25 years by electron capture or positron emission. Because of this, it is used in potassium–argon dating to determine the age of rocks. Applications Argon has several desirable properties: Argon is a chemically inert gas. Argon is the cheapest alternative when nitrogen is not sufficiently inert. Argon has low thermal conductivity. Argon has electronic properties (ionization and/or the emission spectrum) desirable for some applications. Other noble gases would be equally suitable for most of these applications, but argon is by far the cheapest. Argon is inexpensive, since it occurs naturally in air and is readily obtained as a byproduct of cryogenic air separation in the production of liquid oxygen and liquid nitrogen: the primary constituents of air are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful by far. The bulk of argon applications arise simply because it is inert and relatively cheap. Industrial processes Argon is used in some high-temperature industrial processes where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in some types of arc welding such as gas metal arc welding and gas tungsten arc welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium. Argon is used in the poultry industry to asphyxiate birds, either for mass culling following disease outbreaks, or as a means of slaughter more humane than electric stunning. Argon is denser than air and displaces oxygen close to the ground during inert gas asphyxiation. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life. Argon is sometimes used for extinguishing fires where valuable equipment may be damaged by water or foam. Scientific research Liquid argon is used as the target for neutrino experiments and direct dark matter searches. The interaction between the hypothetical WIMPs and an argon nucleus produces scintillation light that is detected by photomultiplier tubes. Two-phase detectors containing argon gas are used to detect the ionized electrons produced during the WIMP–nucleus scattering. As with most other liquefied noble gases, argon has a high scintillation light yield (about 51 photons/keV), is transparent to its own scintillation light, and is relatively easy to purify. Compared to xenon, argon is cheaper and has a distinct scintillation time profile, which allows the separation of electronic recoils from nuclear recoils. On the other hand, its intrinsic beta-ray background is larger due to contamination, unless one uses argon from underground sources, which has much less contamination.

In radioactive decays 40Ar, the most abundant isotope of argon, is produced by the decay of 40K with a half-life of 1.25 years by electron capture or positron emission. Because of this, it is used in potassium–argon dating to determine the age of rocks. Applications Argon has several desirable properties: Argon is a chemically inert gas. Argon is the cheapest alternative when nitrogen is not sufficiently inert. Argon has low thermal conductivity. Argon has electronic properties (ionization and/or the emission spectrum) desirable for some applications. Other noble gases would be equally suitable for most of these applications, but argon is by far the cheapest. Argon is inexpensive, since it occurs naturally in air and is readily obtained as a byproduct of cryogenic air separation in the production of liquid oxygen and liquid nitrogen: the primary constituents of air are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful by far. The bulk of argon applications arise simply because it is inert and relatively cheap. Industrial processes Argon is used in some high-temperature industrial processes where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in some types of arc welding such as gas metal arc welding and gas tungsten arc welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium. Argon is used in the poultry industry to asphyxiate birds, either for mass culling following disease outbreaks, or as a means of slaughter more humane than electric stunning. Argon is denser than air and displaces oxygen close to the ground during inert gas asphyxiation. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life. Argon is sometimes used for extinguishing fires where valuable equipment may be damaged by water or foam. Scientific research Liquid argon is used as the target for neutrino experiments and direct dark matter searches. The interaction between the hypothetical WIMPs and an argon nucleus produces scintillation light that is detected by photomultiplier tubes. Two-phase detectors containing argon gas are used to detect the ionized electrons produced during the WIMP–nucleus scattering. As with most other liquefied noble gases, argon has a high scintillation light yield (about 51 photons/keV), is transparent to its own scintillation light, and is relatively easy to purify. Compared to xenon, argon is cheaper and has a distinct scintillation time profile, which allows the separation of electronic recoils from nuclear recoils. On the other hand, its intrinsic beta-ray background is larger due to contamination, unless one uses argon from underground sources, which has much less contamination.

In radioactive decays 40Ar, the most abundant isotope of argon, is produced by the decay of 40K with a half-life of 1.25 years by electron capture or positron emission. Because of this, it is used in potassium–argon dating to determine the age of rocks. Applications Argon has several desirable properties: Argon is a chemically inert gas. Argon is the cheapest alternative when nitrogen is not sufficiently inert. Argon has low thermal conductivity. Argon has electronic properties (ionization and/or the emission spectrum) desirable for some applications. Other noble gases would be equally suitable for most of these applications, but argon is by far the cheapest. Argon is inexpensive, since it occurs naturally in air and is readily obtained as a byproduct of cryogenic air separation in the production of liquid oxygen and liquid nitrogen: the primary constituents of air are used on a large industrial scale. The other noble gases (except helium) are produced this way as well, but argon is the most plentiful by far. The bulk of argon applications arise simply because it is inert and relatively cheap. Industrial processes Argon is used in some high-temperature industrial processes where ordinarily non-reactive substances become reactive. For example, an argon atmosphere is used in graphite electric furnaces to prevent the graphite from burning. For some of these processes, the presence of nitrogen or oxygen gases might cause defects within the material. Argon is used in some types of arc welding such as gas metal arc welding and gas tungsten arc welding, as well as in the processing of titanium and other reactive elements. An argon atmosphere is also used for growing crystals of silicon and germanium. Argon is used in the poultry industry to asphyxiate birds, either for mass culling following disease outbreaks, or as a means of slaughter more humane than electric stunning. Argon is denser than air and displaces oxygen close to the ground during inert gas asphyxiation. Its non-reactive nature makes it suitable in a food product, and since it replaces oxygen within the dead bird, argon also enhances shelf life. Argon is sometimes used for extinguishing fires where valuable equipment may be damaged by water or foam. Scientific research Liquid argon is used as the target for neutrino experiments and direct dark matter searches. The interaction between the hypothetical WIMPs and an argon nucleus produces scintillation light that is detected by photomultiplier tubes. Two-phase detectors containing argon gas are used to detect the ionized electrons produced during the WIMP–nucleus scattering. As with most other liquefied noble gases, argon has a high scintillation light yield (about 51 photons/keV), is transparent to its own scintillation light, and is relatively easy to purify. Compared to xenon, argon is cheaper and has a distinct scintillation time profile, which allows the separation of electronic recoils from nuclear recoils. On the other hand, its intrinsic beta-ray background is larger due to contamination, unless one uses argon from underground sources, which has much less contamination.

Most of the argon in the Earth's atmosphere was produced by electron capture of long-lived ( + e− → + ν) present in natural potassium within the Earth. The activity in the atmosphere is maintained by cosmogenic production through the knockout reaction (n,2n) and similar reactions. The half-life of is only 269 years. As a result, the underground Ar, shielded by rock and water, has much less contamination. Dark-matter detectors currently operating with liquid argon include DarkSide, WArP, ArDM, microCLEAN and DEAP. Neutrino experiments include ICARUS and MicroBooNE, both of which use high-purity liquid argon in a time projection chamber for fine grained three-dimensional imaging of neutrino interactions. At Linköping University, Sweden, the inert gas is being utilized in a vacuum chamber in which plasma is introduced to ionize metallic films. This process results in a film usable for manufacturing computer processors. The new process would eliminate the need for chemical baths and use of expensive, dangerous and rare materials. Preservative Argon is used to displace oxygen- and moisture-containing air in packaging material to extend the shelf-lives of the contents (argon has the European food additive code E938). Aerial oxidation, hydrolysis, and other chemical reactions that degrade the products are retarded or prevented entirely. High-purity chemicals and pharmaceuticals are sometimes packed and sealed in argon. In winemaking, argon is used in a variety of activities to provide a barrier against oxygen at the liquid surface, which can spoil wine by fueling both microbial metabolism (as with acetic acid bacteria) and standard redox chemistry. Argon is sometimes used as the propellant in aerosol cans. Argon is also used as a preservative for such products as varnish, polyurethane, and paint, by displacing air to prepare a container for storage. Since 2002, the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to inhibit their degradation. Argon is preferable to the helium that had been used in the preceding five decades, because helium gas escapes through the intermolecular pores in most containers and must be regularly replaced. Laboratory equipment Argon may be used as the inert gas within Schlenk lines and gloveboxes. Argon is preferred to less expensive nitrogen in cases where nitrogen may react with the reagents or apparatus. Argon may be used as the carrier gas in gas chromatography and in electrospray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon gas is also commonly used for sputter deposition of thin films as in microelectronics and for wafer cleaning in microfabrication. Medical use Cryosurgery procedures such as cryoablation use liquid argon to destroy tissue such as cancer cells. It is used in a procedure called "argon-enhanced coagulation", a form of argon plasma beam electrosurgery. The procedure carries a risk of producing gas embolism and has resulted in the death of at least one patient.

Most of the argon in the Earth's atmosphere was produced by electron capture of long-lived ( + e− → + ν) present in natural potassium within the Earth. The activity in the atmosphere is maintained by cosmogenic production through the knockout reaction (n,2n) and similar reactions. The half-life of is only 269 years. As a result, the underground Ar, shielded by rock and water, has much less contamination. Dark-matter detectors currently operating with liquid argon include DarkSide, WArP, ArDM, microCLEAN and DEAP. Neutrino experiments include ICARUS and MicroBooNE, both of which use high-purity liquid argon in a time projection chamber for fine grained three-dimensional imaging of neutrino interactions. At Linköping University, Sweden, the inert gas is being utilized in a vacuum chamber in which plasma is introduced to ionize metallic films. This process results in a film usable for manufacturing computer processors. The new process would eliminate the need for chemical baths and use of expensive, dangerous and rare materials. Preservative Argon is used to displace oxygen- and moisture-containing air in packaging material to extend the shelf-lives of the contents (argon has the European food additive code E938). Aerial oxidation, hydrolysis, and other chemical reactions that degrade the products are retarded or prevented entirely. High-purity chemicals and pharmaceuticals are sometimes packed and sealed in argon. In winemaking, argon is used in a variety of activities to provide a barrier against oxygen at the liquid surface, which can spoil wine by fueling both microbial metabolism (as with acetic acid bacteria) and standard redox chemistry. Argon is sometimes used as the propellant in aerosol cans. Argon is also used as a preservative for such products as varnish, polyurethane, and paint, by displacing air to prepare a container for storage. Since 2002, the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to inhibit their degradation. Argon is preferable to the helium that had been used in the preceding five decades, because helium gas escapes through the intermolecular pores in most containers and must be regularly replaced. Laboratory equipment Argon may be used as the inert gas within Schlenk lines and gloveboxes. Argon is preferred to less expensive nitrogen in cases where nitrogen may react with the reagents or apparatus. Argon may be used as the carrier gas in gas chromatography and in electrospray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon gas is also commonly used for sputter deposition of thin films as in microelectronics and for wafer cleaning in microfabrication. Medical use Cryosurgery procedures such as cryoablation use liquid argon to destroy tissue such as cancer cells. It is used in a procedure called "argon-enhanced coagulation", a form of argon plasma beam electrosurgery. The procedure carries a risk of producing gas embolism and has resulted in the death of at least one patient.

Most of the argon in the Earth's atmosphere was produced by electron capture of long-lived ( + e− → + ν) present in natural potassium within the Earth. The activity in the atmosphere is maintained by cosmogenic production through the knockout reaction (n,2n) and similar reactions. The half-life of is only 269 years. As a result, the underground Ar, shielded by rock and water, has much less contamination. Dark-matter detectors currently operating with liquid argon include DarkSide, WArP, ArDM, microCLEAN and DEAP. Neutrino experiments include ICARUS and MicroBooNE, both of which use high-purity liquid argon in a time projection chamber for fine grained three-dimensional imaging of neutrino interactions. At Linköping University, Sweden, the inert gas is being utilized in a vacuum chamber in which plasma is introduced to ionize metallic films. This process results in a film usable for manufacturing computer processors. The new process would eliminate the need for chemical baths and use of expensive, dangerous and rare materials. Preservative Argon is used to displace oxygen- and moisture-containing air in packaging material to extend the shelf-lives of the contents (argon has the European food additive code E938). Aerial oxidation, hydrolysis, and other chemical reactions that degrade the products are retarded or prevented entirely. High-purity chemicals and pharmaceuticals are sometimes packed and sealed in argon. In winemaking, argon is used in a variety of activities to provide a barrier against oxygen at the liquid surface, which can spoil wine by fueling both microbial metabolism (as with acetic acid bacteria) and standard redox chemistry. Argon is sometimes used as the propellant in aerosol cans. Argon is also used as a preservative for such products as varnish, polyurethane, and paint, by displacing air to prepare a container for storage. Since 2002, the American National Archives stores important national documents such as the Declaration of Independence and the Constitution within argon-filled cases to inhibit their degradation. Argon is preferable to the helium that had been used in the preceding five decades, because helium gas escapes through the intermolecular pores in most containers and must be regularly replaced. Laboratory equipment Argon may be used as the inert gas within Schlenk lines and gloveboxes. Argon is preferred to less expensive nitrogen in cases where nitrogen may react with the reagents or apparatus. Argon may be used as the carrier gas in gas chromatography and in electrospray ionization mass spectrometry; it is the gas of choice for the plasma used in ICP spectroscopy. Argon is preferred for the sputter coating of specimens for scanning electron microscopy. Argon gas is also commonly used for sputter deposition of thin films as in microelectronics and for wafer cleaning in microfabrication. Medical use Cryosurgery procedures such as cryoablation use liquid argon to destroy tissue such as cancer cells. It is used in a procedure called "argon-enhanced coagulation", a form of argon plasma beam electrosurgery. The procedure carries a risk of producing gas embolism and has resulted in the death of at least one patient.

Blue argon lasers are used in surgery to weld arteries, destroy tumors, and correct eye defects. Argon has also been used experimentally to replace nitrogen in the breathing or decompression mix known as Argox, to speed the elimination of dissolved nitrogen from the blood. Lighting Incandescent lights are filled with argon, to preserve the filaments at high temperature from oxidation. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with pure argon provide lilac/violet light; with argon and some mercury, blue light. Argon is also used for blue and green argon-ion lasers. Miscellaneous uses Argon is used for thermal insulation in energy-efficient windows. Argon is also used in technical scuba diving to inflate a dry suit because it is inert and has low thermal conductivity. Argon is used as a propellant in the development of the Variable Specific Impulse Magnetoplasma Rocket (VASIMR). Compressed argon gas is allowed to expand, to cool the seeker heads of some versions of the AIM-9 Sidewinder missile and other missiles that use cooled thermal seeker heads. The gas is stored at high pressure. Argon-39, with a half-life of 269 years, has been used for a number of applications, primarily ice core and ground water dating. Also, potassium–argon dating and related argon-argon dating is used to date sedimentary, metamorphic, and igneous rocks. Argon has been used by athletes as a doping agent to simulate hypoxic conditions. In 2014, the World Anti-Doping Agency (WADA) added argon and xenon to the list of prohibited substances and methods, although at this time there is no reliable test for abuse. Safety Although argon is non-toxic, it is 38% more dense than air and therefore considered a dangerous asphyxiant in closed areas. It is difficult to detect because it is colorless, odorless, and tasteless. A 1994 incident, in which a man was asphyxiated after entering an argon-filled section of oil pipe under construction in Alaska, highlights the dangers of argon tank leakage in confined spaces and emphasizes the need for proper use, storage and handling. See also Industrial gas Oxygen–argon ratio, a ratio of two physically similar gases, which has importance in various sectors. References Further reading On triple point pressure at 69 kPa. On triple point pressure at 83.8058 K. External links Argon at The Periodic Table of Videos (University of Nottingham) USGS Periodic Table – Argon Diving applications: Why Argon? Chemical elements E-number additives Noble gases Industrial gases

Blue argon lasers are used in surgery to weld arteries, destroy tumors, and correct eye defects. Argon has also been used experimentally to replace nitrogen in the breathing or decompression mix known as Argox, to speed the elimination of dissolved nitrogen from the blood. Lighting Incandescent lights are filled with argon, to preserve the filaments at high temperature from oxidation. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with pure argon provide lilac/violet light; with argon and some mercury, blue light. Argon is also used for blue and green argon-ion lasers. Miscellaneous uses Argon is used for thermal insulation in energy-efficient windows. Argon is also used in technical scuba diving to inflate a dry suit because it is inert and has low thermal conductivity. Argon is used as a propellant in the development of the Variable Specific Impulse Magnetoplasma Rocket (VASIMR). Compressed argon gas is allowed to expand, to cool the seeker heads of some versions of the AIM-9 Sidewinder missile and other missiles that use cooled thermal seeker heads. The gas is stored at high pressure. Argon-39, with a half-life of 269 years, has been used for a number of applications, primarily ice core and ground water dating. Also, potassium–argon dating and related argon-argon dating is used to date sedimentary, metamorphic, and igneous rocks. Argon has been used by athletes as a doping agent to simulate hypoxic conditions. In 2014, the World Anti-Doping Agency (WADA) added argon and xenon to the list of prohibited substances and methods, although at this time there is no reliable test for abuse. Safety Although argon is non-toxic, it is 38% more dense than air and therefore considered a dangerous asphyxiant in closed areas. It is difficult to detect because it is colorless, odorless, and tasteless. A 1994 incident, in which a man was asphyxiated after entering an argon-filled section of oil pipe under construction in Alaska, highlights the dangers of argon tank leakage in confined spaces and emphasizes the need for proper use, storage and handling. See also Industrial gas Oxygen–argon ratio, a ratio of two physically similar gases, which has importance in various sectors. References Further reading On triple point pressure at 69 kPa. On triple point pressure at 83.8058 K. External links Argon at The Periodic Table of Videos (University of Nottingham) USGS Periodic Table – Argon Diving applications: Why Argon? Chemical elements E-number additives Noble gases Industrial gases

Blue argon lasers are used in surgery to weld arteries, destroy tumors, and correct eye defects. Argon has also been used experimentally to replace nitrogen in the breathing or decompression mix known as Argox, to speed the elimination of dissolved nitrogen from the blood. Lighting Incandescent lights are filled with argon, to preserve the filaments at high temperature from oxidation. It is used for the specific way it ionizes and emits light, such as in plasma globes and calorimetry in experimental particle physics. Gas-discharge lamps filled with pure argon provide lilac/violet light; with argon and some mercury, blue light. Argon is also used for blue and green argon-ion lasers. Miscellaneous uses Argon is used for thermal insulation in energy-efficient windows. Argon is also used in technical scuba diving to inflate a dry suit because it is inert and has low thermal conductivity. Argon is used as a propellant in the development of the Variable Specific Impulse Magnetoplasma Rocket (VASIMR). Compressed argon gas is allowed to expand, to cool the seeker heads of some versions of the AIM-9 Sidewinder missile and other missiles that use cooled thermal seeker heads. The gas is stored at high pressure. Argon-39, with a half-life of 269 years, has been used for a number of applications, primarily ice core and ground water dating. Also, potassium–argon dating and related argon-argon dating is used to date sedimentary, metamorphic, and igneous rocks. Argon has been used by athletes as a doping agent to simulate hypoxic conditions. In 2014, the World Anti-Doping Agency (WADA) added argon and xenon to the list of prohibited substances and methods, although at this time there is no reliable test for abuse. Safety Although argon is non-toxic, it is 38% more dense than air and therefore considered a dangerous asphyxiant in closed areas. It is difficult to detect because it is colorless, odorless, and tasteless. A 1994 incident, in which a man was asphyxiated after entering an argon-filled section of oil pipe under construction in Alaska, highlights the dangers of argon tank leakage in confined spaces and emphasizes the need for proper use, storage and handling. See also Industrial gas Oxygen–argon ratio, a ratio of two physically similar gases, which has importance in various sectors. References Further reading On triple point pressure at 69 kPa. On triple point pressure at 83.8058 K. External links Argon at The Periodic Table of Videos (University of Nottingham) USGS Periodic Table – Argon Diving applications: Why Argon? Chemical elements E-number additives Noble gases Industrial gases

Arsenic Arsenic is a chemical element with the symbol As and atomic number 33. Arsenic occurs in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. Arsenic is a metalloid. It has various allotropes, but only the gray form, which has a metallic appearance, is important to industry. The primary use of arsenic is in alloys of lead (for example, in car batteries and ammunition). Arsenic is a common n-type dopant in semiconductor electronic devices. It is also a component of the III-V compound semiconductor gallium arsenide. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining with the increasing recognition of the toxicity of arsenic and its compounds. A few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic are an essential dietary element in rats, hamsters, goats, chickens, and presumably other species. A role in human metabolism is not known. However, arsenic poisoning occurs in multicellular life if quantities are larger than needed. Arsenic contamination of groundwater is a problem that affects millions of people across the world. The United States' Environmental Protection Agency states that all forms of arsenic are a serious risk to human health. The United States' Agency for Toxic Substances and Disease Registry ranked arsenic as number 1 in its 2001 Priority List of Hazardous Substances at Superfund sites. Arsenic is classified as a Group-A carcinogen. Characteristics Physical characteristics The three most common arsenic allotropes are gray, yellow, and black arsenic, with gray being the most common. Gray arsenic (α-As, space group Rm No. 166) adopts a double-layered structure consisting of many interlocked, ruffled, six-membered rings. Because of weak bonding between the layers, gray arsenic is brittle and has a relatively low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 5.73 g/cm3. Gray arsenic is a semimetal, but becomes a semiconductor with a bandgap of 1.2–1.4 eV if amorphized. Gray arsenic is also the most stable form. Yellow arsenic is soft and waxy, and somewhat similar to tetraphosphorus (). Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense, and most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, . It is rapidly transformed into gray arsenic by light. The yellow form has a density of 1.97 g/cm3. Black arsenic is similar in structure to black phosphorus. Black arsenic can also be formed by cooling vapor at around 100–220 °C and by crystallization of amorphous arsenic in the presence of mercury vapors. It is glassy and brittle.

Arsenic Arsenic is a chemical element with the symbol As and atomic number 33. Arsenic occurs in many minerals, usually in combination with sulfur and metals, but also as a pure elemental crystal. Arsenic is a metalloid. It has various allotropes, but only the gray form, which has a metallic appearance, is important to industry. The primary use of arsenic is in alloys of lead (for example, in car batteries and ammunition). Arsenic is a common n-type dopant in semiconductor electronic devices. It is also a component of the III-V compound semiconductor gallium arsenide. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining with the increasing recognition of the toxicity of arsenic and its compounds. A few species of bacteria are able to use arsenic compounds as respiratory metabolites. Trace quantities of arsenic are an essential dietary element in rats, hamsters, goats, chickens, and presumably other species. A role in human metabolism is not known. However, arsenic poisoning occurs in multicellular life if quantities are larger than needed. Arsenic contamination of groundwater is a problem that affects millions of people across the world. The United States' Environmental Protection Agency states that all forms of arsenic are a serious risk to human health. The United States' Agency for Toxic Substances and Disease Registry ranked arsenic as number 1 in its 2001 Priority List of Hazardous Substances at Superfund sites. Arsenic is classified as a Group-A carcinogen. Characteristics Physical characteristics The three most common arsenic allotropes are gray, yellow, and black arsenic, with gray being the most common. Gray arsenic (α-As, space group Rm No. 166) adopts a double-layered structure consisting of many interlocked, ruffled, six-membered rings. Because of weak bonding between the layers, gray arsenic is brittle and has a relatively low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 5.73 g/cm3. Gray arsenic is a semimetal, but becomes a semiconductor with a bandgap of 1.2–1.4 eV if amorphized. Gray arsenic is also the most stable form. Yellow arsenic is soft and waxy, and somewhat similar to tetraphosphorus (). Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense, and most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, . It is rapidly transformed into gray arsenic by light. The yellow form has a density of 1.97 g/cm3. Black arsenic is similar in structure to black phosphorus. Black arsenic can also be formed by cooling vapor at around 100–220 °C and by crystallization of amorphous arsenic in the presence of mercury vapors. It is glassy and brittle.

It is also a poor electrical conductor. As arsenic's triple point is at 3.628 MPa (35.81 atm), it does not have a melting point at standard pressure but instead sublimes from solid to vapor at 887 K (615 °C or 1137 °F). Isotopes Arsenic occurs in nature as a monoisotopic element, composed of one stable isotope, 75As. As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.30 days. All other isotopes have half-lives of under one day, with the exception of 71As (t1/2=65.30 hours), 72As (t1/2=26.0 hours), 74As (t1/2=17.77 days), 76As (t1/2=1.0942 days), and 77As (t1/2=38.83 hours). Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Chemistry Arsenic has a similar electronegativity and ionization energies to its lighter congener phosphorus and accordingly readily forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which eventually becomes a black surface layer. When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer. It burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, and in fluorine to give arsenic pentafluoride. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at . The triple point is 3.63 MPa and . Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid; however, it does not react with water, alkalis, or non-oxidising acids. Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be highly endothermic and even the group 1 arsenides have properties of intermetallic compounds. Like germanium, selenium, and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the group oxidation state of +5 than its vertical neighbors phosphorus and antimony, and hence arsenic pentoxide and arsenic acid are potent oxidizers. Compounds Compounds of arsenic resemble in some respects those of phosphorus which occupies the same group (column) of the periodic table. The most common oxidation states for arsenic are: −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, and +5 in the arsenates and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square As ions in the mineral skutterudite.

It is also a poor electrical conductor. As arsenic's triple point is at 3.628 MPa (35.81 atm), it does not have a melting point at standard pressure but instead sublimes from solid to vapor at 887 K (615 °C or 1137 °F). Isotopes Arsenic occurs in nature as a monoisotopic element, composed of one stable isotope, 75As. As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.30 days. All other isotopes have half-lives of under one day, with the exception of 71As (t1/2=65.30 hours), 72As (t1/2=26.0 hours), 74As (t1/2=17.77 days), 76As (t1/2=1.0942 days), and 77As (t1/2=38.83 hours). Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Chemistry Arsenic has a similar electronegativity and ionization energies to its lighter congener phosphorus and accordingly readily forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which eventually becomes a black surface layer. When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer. It burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, and in fluorine to give arsenic pentafluoride. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at . The triple point is 3.63 MPa and . Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid; however, it does not react with water, alkalis, or non-oxidising acids. Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be highly endothermic and even the group 1 arsenides have properties of intermetallic compounds. Like germanium, selenium, and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the group oxidation state of +5 than its vertical neighbors phosphorus and antimony, and hence arsenic pentoxide and arsenic acid are potent oxidizers. Compounds Compounds of arsenic resemble in some respects those of phosphorus which occupies the same group (column) of the periodic table. The most common oxidation states for arsenic are: −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, and +5 in the arsenates and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square As ions in the mineral skutterudite.

It is also a poor electrical conductor. As arsenic's triple point is at 3.628 MPa (35.81 atm), it does not have a melting point at standard pressure but instead sublimes from solid to vapor at 887 K (615 °C or 1137 °F). Isotopes Arsenic occurs in nature as a monoisotopic element, composed of one stable isotope, 75As. As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.30 days. All other isotopes have half-lives of under one day, with the exception of 71As (t1/2=65.30 hours), 72As (t1/2=26.0 hours), 74As (t1/2=17.77 days), 76As (t1/2=1.0942 days), and 77As (t1/2=38.83 hours). Isotopes that are lighter than the stable 75As tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds. Chemistry Arsenic has a similar electronegativity and ionization energies to its lighter congener phosphorus and accordingly readily forms covalent molecules with most of the nonmetals. Though stable in dry air, arsenic forms a golden-bronze tarnish upon exposure to humidity which eventually becomes a black surface layer. When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer. It burns in oxygen to form arsenic trioxide and arsenic pentoxide, which have the same structure as the more well-known phosphorus compounds, and in fluorine to give arsenic pentafluoride. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at . The triple point is 3.63 MPa and . Arsenic makes arsenic acid with concentrated nitric acid, arsenous acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid; however, it does not react with water, alkalis, or non-oxidising acids. Arsenic reacts with metals to form arsenides, though these are not ionic compounds containing the As3− ion as the formation of such an anion would be highly endothermic and even the group 1 arsenides have properties of intermetallic compounds. Like germanium, selenium, and bromine, which like arsenic succeed the 3d transition series, arsenic is much less stable in the group oxidation state of +5 than its vertical neighbors phosphorus and antimony, and hence arsenic pentoxide and arsenic acid are potent oxidizers. Compounds Compounds of arsenic resemble in some respects those of phosphorus which occupies the same group (column) of the periodic table. The most common oxidation states for arsenic are: −3 in the arsenides, which are alloy-like intermetallic compounds, +3 in the arsenites, and +5 in the arsenates and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square As ions in the mineral skutterudite.

In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons. Inorganic compounds One of the simplest arsenic compound is the trihydride, the highly toxic, flammable, pyrophoric arsine (AsH3). This compound is generally regarded as stable, since at room temperature it decomposes only slowly. At temperatures of 250–300 °C decomposition to arsenic and hydrogen is rapid. Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition. It oxidises readily in air to form arsenic trioxide and water, and analogous reactions take place with sulfur and selenium instead of oxygen. Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic") and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid and the salts are called arsenates, the most common arsenic contamination of groundwater, and a problem that affects many people. Synthetic arsenates include Scheele's Green (cupric hydrogen arsenate, acidic copper arsenate), calcium arsenate, and lead hydrogen arsenate. These three have been used as agricultural insecticides and poisons. The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorous acid, arsenous acid is genuinely tribasic, with the formula As(OH)3. A broad variety of sulfur compounds of arsenic are known. Orpiment (As2S3) and realgar (As4S4) are somewhat abundant and were formerly used as painting pigments. In As4S10, arsenic has a formal oxidation state of +2 in As4S4 which features As-As bonds so that the total covalency of As is still 3. Both orpiment and realgar, as well as As4S3, have selenium analogs; the analogous As2Te3 is known as the mineral kalgoorlieite, and the anion As2Te− is known as a ligand in cobalt complexes. All trihalides of arsenic(III) are well known except the astatide, which is unknown. Arsenic pentafluoride (AsF5) is the only important pentahalide, reflecting the lower stability of the +5 oxidation state; even so, it is a very strong fluorinating and oxidizing agent. (The pentachloride is stable only below −50 °C, at which temperature it decomposes to the trichloride, releasing chlorine gas.) Alloys Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different which results in distinct bulk properties. Other arsenic alloys include the II-V semiconductor cadmium arsenide. Organoarsenic compounds A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry. Cacodyl was the first organometallic compound known (even though arsenic is not a true metal) and was named from the Greek κακωδία "stink" for its offensive odor; it is very poisonous.

In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons. Inorganic compounds One of the simplest arsenic compound is the trihydride, the highly toxic, flammable, pyrophoric arsine (AsH3). This compound is generally regarded as stable, since at room temperature it decomposes only slowly. At temperatures of 250–300 °C decomposition to arsenic and hydrogen is rapid. Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition. It oxidises readily in air to form arsenic trioxide and water, and analogous reactions take place with sulfur and selenium instead of oxygen. Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic") and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid and the salts are called arsenates, the most common arsenic contamination of groundwater, and a problem that affects many people. Synthetic arsenates include Scheele's Green (cupric hydrogen arsenate, acidic copper arsenate), calcium arsenate, and lead hydrogen arsenate. These three have been used as agricultural insecticides and poisons. The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorous acid, arsenous acid is genuinely tribasic, with the formula As(OH)3. A broad variety of sulfur compounds of arsenic are known. Orpiment (As2S3) and realgar (As4S4) are somewhat abundant and were formerly used as painting pigments. In As4S10, arsenic has a formal oxidation state of +2 in As4S4 which features As-As bonds so that the total covalency of As is still 3. Both orpiment and realgar, as well as As4S3, have selenium analogs; the analogous As2Te3 is known as the mineral kalgoorlieite, and the anion As2Te− is known as a ligand in cobalt complexes. All trihalides of arsenic(III) are well known except the astatide, which is unknown. Arsenic pentafluoride (AsF5) is the only important pentahalide, reflecting the lower stability of the +5 oxidation state; even so, it is a very strong fluorinating and oxidizing agent. (The pentachloride is stable only below −50 °C, at which temperature it decomposes to the trichloride, releasing chlorine gas.) Alloys Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different which results in distinct bulk properties. Other arsenic alloys include the II-V semiconductor cadmium arsenide. Organoarsenic compounds A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry. Cacodyl was the first organometallic compound known (even though arsenic is not a true metal) and was named from the Greek κακωδία "stink" for its offensive odor; it is very poisonous.

In the +3 oxidation state, arsenic is typically pyramidal owing to the influence of the lone pair of electrons. Inorganic compounds One of the simplest arsenic compound is the trihydride, the highly toxic, flammable, pyrophoric arsine (AsH3). This compound is generally regarded as stable, since at room temperature it decomposes only slowly. At temperatures of 250–300 °C decomposition to arsenic and hydrogen is rapid. Several factors, such as humidity, presence of light and certain catalysts (namely aluminium) facilitate the rate of decomposition. It oxidises readily in air to form arsenic trioxide and water, and analogous reactions take place with sulfur and selenium instead of oxygen. Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic") and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid and the salts are called arsenates, the most common arsenic contamination of groundwater, and a problem that affects many people. Synthetic arsenates include Scheele's Green (cupric hydrogen arsenate, acidic copper arsenate), calcium arsenate, and lead hydrogen arsenate. These three have been used as agricultural insecticides and poisons. The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorous acid, arsenous acid is genuinely tribasic, with the formula As(OH)3. A broad variety of sulfur compounds of arsenic are known. Orpiment (As2S3) and realgar (As4S4) are somewhat abundant and were formerly used as painting pigments. In As4S10, arsenic has a formal oxidation state of +2 in As4S4 which features As-As bonds so that the total covalency of As is still 3. Both orpiment and realgar, as well as As4S3, have selenium analogs; the analogous As2Te3 is known as the mineral kalgoorlieite, and the anion As2Te− is known as a ligand in cobalt complexes. All trihalides of arsenic(III) are well known except the astatide, which is unknown. Arsenic pentafluoride (AsF5) is the only important pentahalide, reflecting the lower stability of the +5 oxidation state; even so, it is a very strong fluorinating and oxidizing agent. (The pentachloride is stable only below −50 °C, at which temperature it decomposes to the trichloride, releasing chlorine gas.) Alloys Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different which results in distinct bulk properties. Other arsenic alloys include the II-V semiconductor cadmium arsenide. Organoarsenic compounds A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry. Cacodyl was the first organometallic compound known (even though arsenic is not a true metal) and was named from the Greek κακωδία "stink" for its offensive odor; it is very poisonous.

Occurrence and production Arsenic comprises about 1.5 ppm (0.00015%) of the Earth's crust, and is the 53rd most abundant element. Typical background concentrations of arsenic do not exceed 3 ng/m3 in the atmosphere; 100 mg/kg in soil; 400 μg/kg in vegetation; 10 μg/L in freshwater and 1.5 μg/L in seawater. Minerals with the formula MAsS and MAs2 (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native (elemental) arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment. In 2014, China was the top producer of white arsenic with almost 70% world share, followed by Morocco, Russia, and Belgium, according to the British Geological Survey and the United States Geological Survey. Most arsenic refinement operations in the US and Europe have closed over environmental concerns. Arsenic is found in the smelter dust from copper, gold, and lead smelters, and is recovered primarily from copper refinement dust. On roasting arsenopyrite in air, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of gray arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum, in a hydrogen atmosphere, or by distillation from molten lead-arsenic mixture. History The word arsenic has its origin in the Syriac word (al) zarniqa, from Arabic al-zarnīḵ 'the orpiment’, based on Persian zar 'gold' from the word zarnikh, meaning "yellow" (literally "gold-colored") and hence "(yellow) orpiment". It was adopted into Greek as arsenikon (), a form that is folk etymology, being the neuter form of the Greek word arsenikos (), meaning "male", "virile". The Greek word was adopted in Latin as arsenicum, which in French became arsenic, from which the English word arsenic is taken. Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenic trioxide), which he then reduces to gray arsenic. As the symptoms of arsenic poisoning are not very specific, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the "poison of kings" and the "king of poisons". During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). The isolation of arsenic was described by Jabir ibn Hayyan before 815 AD. Albertus Magnus (Albert the Great, 1193–1280) later isolated the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found in nature, although rare.

Occurrence and production Arsenic comprises about 1.5 ppm (0.00015%) of the Earth's crust, and is the 53rd most abundant element. Typical background concentrations of arsenic do not exceed 3 ng/m3 in the atmosphere; 100 mg/kg in soil; 400 μg/kg in vegetation; 10 μg/L in freshwater and 1.5 μg/L in seawater. Minerals with the formula MAsS and MAs2 (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native (elemental) arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment. In 2014, China was the top producer of white arsenic with almost 70% world share, followed by Morocco, Russia, and Belgium, according to the British Geological Survey and the United States Geological Survey. Most arsenic refinement operations in the US and Europe have closed over environmental concerns. Arsenic is found in the smelter dust from copper, gold, and lead smelters, and is recovered primarily from copper refinement dust. On roasting arsenopyrite in air, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of gray arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum, in a hydrogen atmosphere, or by distillation from molten lead-arsenic mixture. History The word arsenic has its origin in the Syriac word (al) zarniqa, from Arabic al-zarnīḵ 'the orpiment’, based on Persian zar 'gold' from the word zarnikh, meaning "yellow" (literally "gold-colored") and hence "(yellow) orpiment". It was adopted into Greek as arsenikon (), a form that is folk etymology, being the neuter form of the Greek word arsenikos (), meaning "male", "virile". The Greek word was adopted in Latin as arsenicum, which in French became arsenic, from which the English word arsenic is taken. Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenic trioxide), which he then reduces to gray arsenic. As the symptoms of arsenic poisoning are not very specific, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the "poison of kings" and the "king of poisons". During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). The isolation of arsenic was described by Jabir ibn Hayyan before 815 AD. Albertus Magnus (Albert the Great, 1193–1280) later isolated the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found in nature, although rare.

Occurrence and production Arsenic comprises about 1.5 ppm (0.00015%) of the Earth's crust, and is the 53rd most abundant element. Typical background concentrations of arsenic do not exceed 3 ng/m3 in the atmosphere; 100 mg/kg in soil; 400 μg/kg in vegetation; 10 μg/L in freshwater and 1.5 μg/L in seawater. Minerals with the formula MAsS and MAs2 (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native (elemental) arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment. In 2014, China was the top producer of white arsenic with almost 70% world share, followed by Morocco, Russia, and Belgium, according to the British Geological Survey and the United States Geological Survey. Most arsenic refinement operations in the US and Europe have closed over environmental concerns. Arsenic is found in the smelter dust from copper, gold, and lead smelters, and is recovered primarily from copper refinement dust. On roasting arsenopyrite in air, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of gray arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum, in a hydrogen atmosphere, or by distillation from molten lead-arsenic mixture. History The word arsenic has its origin in the Syriac word (al) zarniqa, from Arabic al-zarnīḵ 'the orpiment’, based on Persian zar 'gold' from the word zarnikh, meaning "yellow" (literally "gold-colored") and hence "(yellow) orpiment". It was adopted into Greek as arsenikon (), a form that is folk etymology, being the neuter form of the Greek word arsenikos (), meaning "male", "virile". The Greek word was adopted in Latin as arsenicum, which in French became arsenic, from which the English word arsenic is taken. Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenic trioxide), which he then reduces to gray arsenic. As the symptoms of arsenic poisoning are not very specific, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the "poison of kings" and the "king of poisons". During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). The isolation of arsenic was described by Jabir ibn Hayyan before 815 AD. Albertus Magnus (Albert the Great, 1193–1280) later isolated the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found in nature, although rare.

Cadet's fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide. In the Victorian era, "arsenic" ("white arsenic" or arsenic trioxide) was mixed with vinegar and chalk and eaten by women to improve the complexion of their faces, making their skin paler to show they did not work in the fields. The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in 21 deaths. Wallpaper production also began to use dyes made from arsenic, which was thought to increase the pigment's brightness. Two arsenic pigments have been widely used since their discovery – Paris Green and Scheele's Green. After the toxicity of arsenic became widely known, these chemicals were used less often as pigments and more often as insecticides. In the 1860s, an arsenic byproduct of dye production, London Purple, was widely used. This was a solid mixture of arsenic trioxide, aniline, lime, and ferrous oxide, insoluble in water and very toxic by inhalation or ingestion But it was later replaced with Paris Green, another arsenic-based dye. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942. Applications Agricultural The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1930s, a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades, this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic led to a ban of CCA in consumer products in 2004, initiated by the European Union and United States. However, CCA remains in heavy use in other countries (such as on Malaysian rubber plantations). Arsenic was also used in various agricultural insecticides and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – replaced lead arsenate in agriculture. These organic arsenicals were in turn phased out by 2013 in all agricultural activities except cotton farming. The biogeochemistry of arsenic is complex and includes various adsorption and desorption processes. The toxicity of arsenic is connected to its solubility and is affected by pH. Arsenite () is more soluble than arsenate () and is more toxic; however, at a lower pH, arsenate becomes more mobile and toxic. It was found that addition of sulfur, phosphorus, and iron oxides to high-arsenite soils greatly reduces arsenic phytotoxicity.

Cadet's fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide. In the Victorian era, "arsenic" ("white arsenic" or arsenic trioxide) was mixed with vinegar and chalk and eaten by women to improve the complexion of their faces, making their skin paler to show they did not work in the fields. The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in 21 deaths. Wallpaper production also began to use dyes made from arsenic, which was thought to increase the pigment's brightness. Two arsenic pigments have been widely used since their discovery – Paris Green and Scheele's Green. After the toxicity of arsenic became widely known, these chemicals were used less often as pigments and more often as insecticides. In the 1860s, an arsenic byproduct of dye production, London Purple, was widely used. This was a solid mixture of arsenic trioxide, aniline, lime, and ferrous oxide, insoluble in water and very toxic by inhalation or ingestion But it was later replaced with Paris Green, another arsenic-based dye. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942. Applications Agricultural The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1930s, a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades, this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic led to a ban of CCA in consumer products in 2004, initiated by the European Union and United States. However, CCA remains in heavy use in other countries (such as on Malaysian rubber plantations). Arsenic was also used in various agricultural insecticides and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – replaced lead arsenate in agriculture. These organic arsenicals were in turn phased out by 2013 in all agricultural activities except cotton farming. The biogeochemistry of arsenic is complex and includes various adsorption and desorption processes. The toxicity of arsenic is connected to its solubility and is affected by pH. Arsenite () is more soluble than arsenate () and is more toxic; however, at a lower pH, arsenate becomes more mobile and toxic. It was found that addition of sulfur, phosphorus, and iron oxides to high-arsenite soils greatly reduces arsenic phytotoxicity.

Cadet's fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide. In the Victorian era, "arsenic" ("white arsenic" or arsenic trioxide) was mixed with vinegar and chalk and eaten by women to improve the complexion of their faces, making their skin paler to show they did not work in the fields. The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in 21 deaths. Wallpaper production also began to use dyes made from arsenic, which was thought to increase the pigment's brightness. Two arsenic pigments have been widely used since their discovery – Paris Green and Scheele's Green. After the toxicity of arsenic became widely known, these chemicals were used less often as pigments and more often as insecticides. In the 1860s, an arsenic byproduct of dye production, London Purple, was widely used. This was a solid mixture of arsenic trioxide, aniline, lime, and ferrous oxide, insoluble in water and very toxic by inhalation or ingestion But it was later replaced with Paris Green, another arsenic-based dye. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942. Applications Agricultural The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1930s, a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades, this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic led to a ban of CCA in consumer products in 2004, initiated by the European Union and United States. However, CCA remains in heavy use in other countries (such as on Malaysian rubber plantations). Arsenic was also used in various agricultural insecticides and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – replaced lead arsenate in agriculture. These organic arsenicals were in turn phased out by 2013 in all agricultural activities except cotton farming. The biogeochemistry of arsenic is complex and includes various adsorption and desorption processes. The toxicity of arsenic is connected to its solubility and is affected by pH. Arsenite () is more soluble than arsenate () and is more toxic; however, at a lower pH, arsenate becomes more mobile and toxic. It was found that addition of sulfur, phosphorus, and iron oxides to high-arsenite soils greatly reduces arsenic phytotoxicity.

Arsenic is used as a feed additive in poultry and swine production, in particular in the U.S. to increase weight gain, improve feed efficiency, and prevent disease. An example is roxarsone, which had been used as a broiler starter by about 70% of U.S. broiler growers. Alpharma, a subsidiary of Pfizer Inc., which produces roxarsone, voluntarily suspended sales of the drug in response to studies showing elevated levels of inorganic arsenic, a carcinogen, in treated chickens. A successor to Alpharma, Zoetis, continues to sell nitarsone, primarily for use in turkeys. Arsenic is intentionally added to the feed of chickens raised for human consumption. Organic arsenic compounds are less toxic than pure arsenic, and promote the growth of chickens. Under some conditions, the arsenic in chicken feed is converted to the toxic inorganic form. A 2006 study of the remains of the Australian racehorse, Phar Lap, determined that the 1932 death of the famous champion was caused by a massive overdose of arsenic. Sydney veterinarian Percy Sykes stated, "In those days, arsenic was quite a common tonic, usually given in the form of a solution (Fowler's Solution) ... It was so common that I'd reckon 90 per cent of the horses had arsenic in their system." Medical use During the 17th, 18th, and 19th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine, as well as neosalvarsan, was indicated for syphilis, but has been superseded by modern antibiotics. However, arsenicals such as melarsoprol are still used for the treatment of trypanosomiasis, since although these drugs have the disadvantage of severe toxicity, the disease is almost uniformly fatal if untreated. Arsenic trioxide has been used in a variety of ways over the past 500 years, most commonly in the treatment of cancer, but also in medications as diverse as Fowler's solution in psoriasis. The US Food and Drug Administration in the year 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to all-trans retinoic acid. A 2008 paper reports success in locating tumors using arsenic-74 (a positron emitter). This isotope produces clearer PET scan images than the previous radioactive agent, iodine-124, because the body tends to transport iodine to the thyroid gland producing signal noise. Nanoparticles of arsenic have shown ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations. In subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid-18th to 19th centuries; its use as a stimulant was especially prevalent as sport animals such as race horses or with work dogs. Alloys The main use of arsenic is in alloying with lead. Lead components in car batteries are strengthened by the presence of a very small percentage of arsenic. Dezincification of brass (a copper-zinc alloy) is greatly reduced by the addition of arsenic.

Arsenic is used as a feed additive in poultry and swine production, in particular in the U.S. to increase weight gain, improve feed efficiency, and prevent disease. An example is roxarsone, which had been used as a broiler starter by about 70% of U.S. broiler growers. Alpharma, a subsidiary of Pfizer Inc., which produces roxarsone, voluntarily suspended sales of the drug in response to studies showing elevated levels of inorganic arsenic, a carcinogen, in treated chickens. A successor to Alpharma, Zoetis, continues to sell nitarsone, primarily for use in turkeys. Arsenic is intentionally added to the feed of chickens raised for human consumption. Organic arsenic compounds are less toxic than pure arsenic, and promote the growth of chickens. Under some conditions, the arsenic in chicken feed is converted to the toxic inorganic form. A 2006 study of the remains of the Australian racehorse, Phar Lap, determined that the 1932 death of the famous champion was caused by a massive overdose of arsenic. Sydney veterinarian Percy Sykes stated, "In those days, arsenic was quite a common tonic, usually given in the form of a solution (Fowler's Solution) ... It was so common that I'd reckon 90 per cent of the horses had arsenic in their system." Medical use During the 17th, 18th, and 19th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine, as well as neosalvarsan, was indicated for syphilis, but has been superseded by modern antibiotics. However, arsenicals such as melarsoprol are still used for the treatment of trypanosomiasis, since although these drugs have the disadvantage of severe toxicity, the disease is almost uniformly fatal if untreated. Arsenic trioxide has been used in a variety of ways over the past 500 years, most commonly in the treatment of cancer, but also in medications as diverse as Fowler's solution in psoriasis. The US Food and Drug Administration in the year 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to all-trans retinoic acid. A 2008 paper reports success in locating tumors using arsenic-74 (a positron emitter). This isotope produces clearer PET scan images than the previous radioactive agent, iodine-124, because the body tends to transport iodine to the thyroid gland producing signal noise. Nanoparticles of arsenic have shown ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations. In subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid-18th to 19th centuries; its use as a stimulant was especially prevalent as sport animals such as race horses or with work dogs. Alloys The main use of arsenic is in alloying with lead. Lead components in car batteries are strengthened by the presence of a very small percentage of arsenic. Dezincification of brass (a copper-zinc alloy) is greatly reduced by the addition of arsenic.

Arsenic is used as a feed additive in poultry and swine production, in particular in the U.S. to increase weight gain, improve feed efficiency, and prevent disease. An example is roxarsone, which had been used as a broiler starter by about 70% of U.S. broiler growers. Alpharma, a subsidiary of Pfizer Inc., which produces roxarsone, voluntarily suspended sales of the drug in response to studies showing elevated levels of inorganic arsenic, a carcinogen, in treated chickens. A successor to Alpharma, Zoetis, continues to sell nitarsone, primarily for use in turkeys. Arsenic is intentionally added to the feed of chickens raised for human consumption. Organic arsenic compounds are less toxic than pure arsenic, and promote the growth of chickens. Under some conditions, the arsenic in chicken feed is converted to the toxic inorganic form. A 2006 study of the remains of the Australian racehorse, Phar Lap, determined that the 1932 death of the famous champion was caused by a massive overdose of arsenic. Sydney veterinarian Percy Sykes stated, "In those days, arsenic was quite a common tonic, usually given in the form of a solution (Fowler's Solution) ... It was so common that I'd reckon 90 per cent of the horses had arsenic in their system." Medical use During the 17th, 18th, and 19th centuries, a number of arsenic compounds were used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine, as well as neosalvarsan, was indicated for syphilis, but has been superseded by modern antibiotics. However, arsenicals such as melarsoprol are still used for the treatment of trypanosomiasis, since although these drugs have the disadvantage of severe toxicity, the disease is almost uniformly fatal if untreated. Arsenic trioxide has been used in a variety of ways over the past 500 years, most commonly in the treatment of cancer, but also in medications as diverse as Fowler's solution in psoriasis. The US Food and Drug Administration in the year 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to all-trans retinoic acid. A 2008 paper reports success in locating tumors using arsenic-74 (a positron emitter). This isotope produces clearer PET scan images than the previous radioactive agent, iodine-124, because the body tends to transport iodine to the thyroid gland producing signal noise. Nanoparticles of arsenic have shown ability to kill cancer cells with lesser cytotoxicity than other arsenic formulations. In subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid-18th to 19th centuries; its use as a stimulant was especially prevalent as sport animals such as race horses or with work dogs. Alloys The main use of arsenic is in alloying with lead. Lead components in car batteries are strengthened by the presence of a very small percentage of arsenic. Dezincification of brass (a copper-zinc alloy) is greatly reduced by the addition of arsenic.

"Phosphorus Deoxidized Arsenical Copper" with an arsenic content of 0.3% has an increased corrosion stability in certain environments. Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made from GaAs are much faster (but also much more expensive) than those made from silicon. Unlike silicon, GaAs has a direct bandgap, and can be used in laser diodes and LEDs to convert electrical energy directly into light. Military After World War I, the United States built a stockpile of 20,000 tons of weaponized lewisite (ClCH=CHAsCl2), an organoarsenic vesicant (blister agent) and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico in the 1950s. During the Vietnam War, the United States used Agent Blue, a mixture of sodium cacodylate and its acid form, as one of the rainbow herbicides to deprive North Vietnamese soldiers of foliage cover and rice. Other uses Copper acetoarsenite was used as a green pigment known under many names, including Paris Green and Emerald Green. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets. Arsenic is used in bronzing and pyrotechnics. As much as 2% of produced arsenic is used in lead alloys for lead shot and bullets. Arsenic is added in small quantities to alpha-brass to make it dezincification-resistant. This grade of brass is used in plumbing fittings and other wet environments. Arsenic is also used for taxonomic sample preservation. Arsenic was used as an opacifier in ceramics, creating white glazes. Until recently, arsenic was used in optical glass. Modern glass manufacturers, under pressure from environmentalists, have ceased using both arsenic and lead. Biological role Bacteria Some species of bacteria obtain their energy in the absence of oxygen by oxidizing various fuels while reducing arsenate to arsenite. Under oxidative environmental conditions some bacteria use arsenite as fuel, which they oxidize to arsenate. The enzymes involved are known as arsenate reductases (Arr). In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just as ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that, over the course of history, these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain PHS-1 has been isolated and is related to the gammaproteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues. In 2011, it was postulated that a strain of Halomonadaceae could be grown in the absence of phosphorus if that element were substituted with arsenic, exploiting the fact that the arsenate and phosphate anions are similar structurally. The study was widely criticised and subsequently refuted by independent researcher groups. Essential trace element in higher animals Some evidence indicates that arsenic is an essential trace mineral in birds (chickens), and in mammals (rats, hamsters, and goats). However, the biological function is not known.

"Phosphorus Deoxidized Arsenical Copper" with an arsenic content of 0.3% has an increased corrosion stability in certain environments. Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made from GaAs are much faster (but also much more expensive) than those made from silicon. Unlike silicon, GaAs has a direct bandgap, and can be used in laser diodes and LEDs to convert electrical energy directly into light. Military After World War I, the United States built a stockpile of 20,000 tons of weaponized lewisite (ClCH=CHAsCl2), an organoarsenic vesicant (blister agent) and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico in the 1950s. During the Vietnam War, the United States used Agent Blue, a mixture of sodium cacodylate and its acid form, as one of the rainbow herbicides to deprive North Vietnamese soldiers of foliage cover and rice. Other uses Copper acetoarsenite was used as a green pigment known under many names, including Paris Green and Emerald Green. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets. Arsenic is used in bronzing and pyrotechnics. As much as 2% of produced arsenic is used in lead alloys for lead shot and bullets. Arsenic is added in small quantities to alpha-brass to make it dezincification-resistant. This grade of brass is used in plumbing fittings and other wet environments. Arsenic is also used for taxonomic sample preservation. Arsenic was used as an opacifier in ceramics, creating white glazes. Until recently, arsenic was used in optical glass. Modern glass manufacturers, under pressure from environmentalists, have ceased using both arsenic and lead. Biological role Bacteria Some species of bacteria obtain their energy in the absence of oxygen by oxidizing various fuels while reducing arsenate to arsenite. Under oxidative environmental conditions some bacteria use arsenite as fuel, which they oxidize to arsenate. The enzymes involved are known as arsenate reductases (Arr). In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just as ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that, over the course of history, these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain PHS-1 has been isolated and is related to the gammaproteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues. In 2011, it was postulated that a strain of Halomonadaceae could be grown in the absence of phosphorus if that element were substituted with arsenic, exploiting the fact that the arsenate and phosphate anions are similar structurally. The study was widely criticised and subsequently refuted by independent researcher groups. Essential trace element in higher animals Some evidence indicates that arsenic is an essential trace mineral in birds (chickens), and in mammals (rats, hamsters, and goats). However, the biological function is not known.

"Phosphorus Deoxidized Arsenical Copper" with an arsenic content of 0.3% has an increased corrosion stability in certain environments. Gallium arsenide is an important semiconductor material, used in integrated circuits. Circuits made from GaAs are much faster (but also much more expensive) than those made from silicon. Unlike silicon, GaAs has a direct bandgap, and can be used in laser diodes and LEDs to convert electrical energy directly into light. Military After World War I, the United States built a stockpile of 20,000 tons of weaponized lewisite (ClCH=CHAsCl2), an organoarsenic vesicant (blister agent) and lung irritant. The stockpile was neutralized with bleach and dumped into the Gulf of Mexico in the 1950s. During the Vietnam War, the United States used Agent Blue, a mixture of sodium cacodylate and its acid form, as one of the rainbow herbicides to deprive North Vietnamese soldiers of foliage cover and rice. Other uses Copper acetoarsenite was used as a green pigment known under many names, including Paris Green and Emerald Green. It caused numerous arsenic poisonings. Scheele's Green, a copper arsenate, was used in the 19th century as a coloring agent in sweets. Arsenic is used in bronzing and pyrotechnics. As much as 2% of produced arsenic is used in lead alloys for lead shot and bullets. Arsenic is added in small quantities to alpha-brass to make it dezincification-resistant. This grade of brass is used in plumbing fittings and other wet environments. Arsenic is also used for taxonomic sample preservation. Arsenic was used as an opacifier in ceramics, creating white glazes. Until recently, arsenic was used in optical glass. Modern glass manufacturers, under pressure from environmentalists, have ceased using both arsenic and lead. Biological role Bacteria Some species of bacteria obtain their energy in the absence of oxygen by oxidizing various fuels while reducing arsenate to arsenite. Under oxidative environmental conditions some bacteria use arsenite as fuel, which they oxidize to arsenate. The enzymes involved are known as arsenate reductases (Arr). In 2008, bacteria were discovered that employ a version of photosynthesis in the absence of oxygen with arsenites as electron donors, producing arsenates (just as ordinary photosynthesis uses water as electron donor, producing molecular oxygen). Researchers conjecture that, over the course of history, these photosynthesizing organisms produced the arsenates that allowed the arsenate-reducing bacteria to thrive. One strain PHS-1 has been isolated and is related to the gammaproteobacterium Ectothiorhodospira shaposhnikovii. The mechanism is unknown, but an encoded Arr enzyme may function in reverse to its known homologues. In 2011, it was postulated that a strain of Halomonadaceae could be grown in the absence of phosphorus if that element were substituted with arsenic, exploiting the fact that the arsenate and phosphate anions are similar structurally. The study was widely criticised and subsequently refuted by independent researcher groups. Essential trace element in higher animals Some evidence indicates that arsenic is an essential trace mineral in birds (chickens), and in mammals (rats, hamsters, and goats). However, the biological function is not known.

Heredity Arsenic has been linked to epigenetic changes, heritable changes in gene expression that occur without changes in DNA sequence. These include DNA methylation, histone modification, and RNA interference. Toxic levels of arsenic cause significant DNA hypermethylation of tumor suppressor genes p16 and p53, thus increasing risk of carcinogenesis. These epigenetic events have been studied in vitro using human kidney cells and in vivo using rat liver cells and peripheral blood leukocytes in humans. Inductively coupled plasma mass spectrometry (ICP-MS) is used to detect precise levels of intracellular arsenic and other arsenic bases involved in epigenetic modification of DNA. Studies investigating arsenic as an epigenetic factor can be used to develop precise biomarkers of exposure and susceptibility. The Chinese brake fern (Pteris vittata) hyperaccumulates arsenic from the soil into its leaves and has a proposed use in phytoremediation. Biomethylation Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized through a process of methylation. For example, the mold Scopulariopsis brevicaulis produces trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms, but there is little danger in eating fish because this arsenic compound is nearly non-toxic. Environmental issues Exposure Naturally occurring sources of human exposure include volcanic ash, weathering of minerals and ores, and mineralized groundwater. Arsenic is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is inhalation of atmospheric gases and dusts. During the Victorian era, arsenic was widely used in home decor, especially wallpapers. Occurrence in drinking water Extensive arsenic contamination of groundwater has led to widespread arsenic poisoning in Bangladesh and neighboring countries. It is estimated that approximately 57 million people in the Bengal basin are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion (ppb). However, a study of cancer rates in Taiwan suggested that significant increases in cancer mortality appear only at levels above 150 ppb. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater, caused by the anoxic conditions of the subsurface. This groundwater was used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacteria-contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in Southeast Asia, such as Vietnam and Cambodia, have geological environments that produce groundwater with a high arsenic content.

Heredity Arsenic has been linked to epigenetic changes, heritable changes in gene expression that occur without changes in DNA sequence. These include DNA methylation, histone modification, and RNA interference. Toxic levels of arsenic cause significant DNA hypermethylation of tumor suppressor genes p16 and p53, thus increasing risk of carcinogenesis. These epigenetic events have been studied in vitro using human kidney cells and in vivo using rat liver cells and peripheral blood leukocytes in humans. Inductively coupled plasma mass spectrometry (ICP-MS) is used to detect precise levels of intracellular arsenic and other arsenic bases involved in epigenetic modification of DNA. Studies investigating arsenic as an epigenetic factor can be used to develop precise biomarkers of exposure and susceptibility. The Chinese brake fern (Pteris vittata) hyperaccumulates arsenic from the soil into its leaves and has a proposed use in phytoremediation. Biomethylation Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized through a process of methylation. For example, the mold Scopulariopsis brevicaulis produces trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms, but there is little danger in eating fish because this arsenic compound is nearly non-toxic. Environmental issues Exposure Naturally occurring sources of human exposure include volcanic ash, weathering of minerals and ores, and mineralized groundwater. Arsenic is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is inhalation of atmospheric gases and dusts. During the Victorian era, arsenic was widely used in home decor, especially wallpapers. Occurrence in drinking water Extensive arsenic contamination of groundwater has led to widespread arsenic poisoning in Bangladesh and neighboring countries. It is estimated that approximately 57 million people in the Bengal basin are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion (ppb). However, a study of cancer rates in Taiwan suggested that significant increases in cancer mortality appear only at levels above 150 ppb. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater, caused by the anoxic conditions of the subsurface. This groundwater was used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacteria-contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in Southeast Asia, such as Vietnam and Cambodia, have geological environments that produce groundwater with a high arsenic content.

Heredity Arsenic has been linked to epigenetic changes, heritable changes in gene expression that occur without changes in DNA sequence. These include DNA methylation, histone modification, and RNA interference. Toxic levels of arsenic cause significant DNA hypermethylation of tumor suppressor genes p16 and p53, thus increasing risk of carcinogenesis. These epigenetic events have been studied in vitro using human kidney cells and in vivo using rat liver cells and peripheral blood leukocytes in humans. Inductively coupled plasma mass spectrometry (ICP-MS) is used to detect precise levels of intracellular arsenic and other arsenic bases involved in epigenetic modification of DNA. Studies investigating arsenic as an epigenetic factor can be used to develop precise biomarkers of exposure and susceptibility. The Chinese brake fern (Pteris vittata) hyperaccumulates arsenic from the soil into its leaves and has a proposed use in phytoremediation. Biomethylation Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized through a process of methylation. For example, the mold Scopulariopsis brevicaulis produces trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms, but there is little danger in eating fish because this arsenic compound is nearly non-toxic. Environmental issues Exposure Naturally occurring sources of human exposure include volcanic ash, weathering of minerals and ores, and mineralized groundwater. Arsenic is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is inhalation of atmospheric gases and dusts. During the Victorian era, arsenic was widely used in home decor, especially wallpapers. Occurrence in drinking water Extensive arsenic contamination of groundwater has led to widespread arsenic poisoning in Bangladesh and neighboring countries. It is estimated that approximately 57 million people in the Bengal basin are drinking groundwater with arsenic concentrations elevated above the World Health Organization's standard of 10 parts per billion (ppb). However, a study of cancer rates in Taiwan suggested that significant increases in cancer mortality appear only at levels above 150 ppb. The arsenic in the groundwater is of natural origin, and is released from the sediment into the groundwater, caused by the anoxic conditions of the subsurface. This groundwater was used after local and western NGOs and the Bangladeshi government undertook a massive shallow tube well drinking-water program in the late twentieth century. This program was designed to prevent drinking of bacteria-contaminated surface waters, but failed to test for arsenic in the groundwater. Many other countries and districts in Southeast Asia, such as Vietnam and Cambodia, have geological environments that produce groundwater with a high arsenic content.

Arsenicosis was reported in Nakhon Si Thammarat, Thailand in 1987, and the Chao Phraya River probably contains high levels of naturally occurring dissolved arsenic without being a public health problem because much of the public uses bottled water. In Pakistan, more than 60 million people are exposed to arsenic polluted drinking water indicated by a recent report of Science. Podgorski's team investigated more than 1200 samples and more than 66% exceeded the WHO minimum contamination level. Since the 1980s, residents of the Ba Men region of Inner Mongolia, China have been chronically exposed to arsenic through drinking water from contaminated wells. A 2009 research study observed an elevated presence of skin lesions among residents with well water arsenic concentrations between 5 and 10 µg/L, suggesting that arsenic induced toxicity may occur at relatively low concentrations with chronic exposure. Overall, 20 of China's 34 provinces have high arsenic concentrations in the groundwater supply, potentially exposing 19 million people to hazardous drinking water. In the United States, arsenic is most commonly found in the ground waters of the southwest. Parts of New England, Michigan, Wisconsin, Minnesota and the Dakotas are also known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 part per billion drinking water standard. According to a recent film funded by the US Superfund, millions of private wells have unknown arsenic levels, and in some areas of the US, more than 20% of the wells may contain levels that exceed established limits. Low-level exposure to arsenic at concentrations of 100 parts per billion (i.e., above the 10 parts per billion drinking water standard) compromises the initial immune response to H1N1 or swine flu infection according to NIEHS-supported scientists. The study, conducted in laboratory mice, suggests that people exposed to arsenic in their drinking water may be at increased risk for more serious illness or death from the virus. Some Canadians are drinking water that contains inorganic arsenic. Private-dug–well waters are most at risk for containing inorganic arsenic. Preliminary well water analysis typically does not test for arsenic. Researchers at the Geological Survey of Canada have modeled relative variation in natural arsenic hazard potential for the province of New Brunswick. This study has important implications for potable water and health concerns relating to inorganic arsenic. Epidemiological evidence from Chile shows a dose-dependent connection between chronic arsenic exposure and various forms of cancer, in particular when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated at contaminations less than 50 ppb. Arsenic is itself a constituent of tobacco smoke. Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable increase in risk for bladder cancer at 10 ppb. According to Peter Ravenscroft of the Department of Geography at the University of Cambridge, roughly 80 million people worldwide consume between 10 and 50 ppb arsenic in their drinking water.

Arsenicosis was reported in Nakhon Si Thammarat, Thailand in 1987, and the Chao Phraya River probably contains high levels of naturally occurring dissolved arsenic without being a public health problem because much of the public uses bottled water. In Pakistan, more than 60 million people are exposed to arsenic polluted drinking water indicated by a recent report of Science. Podgorski's team investigated more than 1200 samples and more than 66% exceeded the WHO minimum contamination level. Since the 1980s, residents of the Ba Men region of Inner Mongolia, China have been chronically exposed to arsenic through drinking water from contaminated wells. A 2009 research study observed an elevated presence of skin lesions among residents with well water arsenic concentrations between 5 and 10 µg/L, suggesting that arsenic induced toxicity may occur at relatively low concentrations with chronic exposure. Overall, 20 of China's 34 provinces have high arsenic concentrations in the groundwater supply, potentially exposing 19 million people to hazardous drinking water. In the United States, arsenic is most commonly found in the ground waters of the southwest. Parts of New England, Michigan, Wisconsin, Minnesota and the Dakotas are also known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 part per billion drinking water standard. According to a recent film funded by the US Superfund, millions of private wells have unknown arsenic levels, and in some areas of the US, more than 20% of the wells may contain levels that exceed established limits. Low-level exposure to arsenic at concentrations of 100 parts per billion (i.e., above the 10 parts per billion drinking water standard) compromises the initial immune response to H1N1 or swine flu infection according to NIEHS-supported scientists. The study, conducted in laboratory mice, suggests that people exposed to arsenic in their drinking water may be at increased risk for more serious illness or death from the virus. Some Canadians are drinking water that contains inorganic arsenic. Private-dug–well waters are most at risk for containing inorganic arsenic. Preliminary well water analysis typically does not test for arsenic. Researchers at the Geological Survey of Canada have modeled relative variation in natural arsenic hazard potential for the province of New Brunswick. This study has important implications for potable water and health concerns relating to inorganic arsenic. Epidemiological evidence from Chile shows a dose-dependent connection between chronic arsenic exposure and various forms of cancer, in particular when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated at contaminations less than 50 ppb. Arsenic is itself a constituent of tobacco smoke. Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable increase in risk for bladder cancer at 10 ppb. According to Peter Ravenscroft of the Department of Geography at the University of Cambridge, roughly 80 million people worldwide consume between 10 and 50 ppb arsenic in their drinking water.

Arsenicosis was reported in Nakhon Si Thammarat, Thailand in 1987, and the Chao Phraya River probably contains high levels of naturally occurring dissolved arsenic without being a public health problem because much of the public uses bottled water. In Pakistan, more than 60 million people are exposed to arsenic polluted drinking water indicated by a recent report of Science. Podgorski's team investigated more than 1200 samples and more than 66% exceeded the WHO minimum contamination level. Since the 1980s, residents of the Ba Men region of Inner Mongolia, China have been chronically exposed to arsenic through drinking water from contaminated wells. A 2009 research study observed an elevated presence of skin lesions among residents with well water arsenic concentrations between 5 and 10 µg/L, suggesting that arsenic induced toxicity may occur at relatively low concentrations with chronic exposure. Overall, 20 of China's 34 provinces have high arsenic concentrations in the groundwater supply, potentially exposing 19 million people to hazardous drinking water. In the United States, arsenic is most commonly found in the ground waters of the southwest. Parts of New England, Michigan, Wisconsin, Minnesota and the Dakotas are also known to have significant concentrations of arsenic in ground water. Increased levels of skin cancer have been associated with arsenic exposure in Wisconsin, even at levels below the 10 part per billion drinking water standard. According to a recent film funded by the US Superfund, millions of private wells have unknown arsenic levels, and in some areas of the US, more than 20% of the wells may contain levels that exceed established limits. Low-level exposure to arsenic at concentrations of 100 parts per billion (i.e., above the 10 parts per billion drinking water standard) compromises the initial immune response to H1N1 or swine flu infection according to NIEHS-supported scientists. The study, conducted in laboratory mice, suggests that people exposed to arsenic in their drinking water may be at increased risk for more serious illness or death from the virus. Some Canadians are drinking water that contains inorganic arsenic. Private-dug–well waters are most at risk for containing inorganic arsenic. Preliminary well water analysis typically does not test for arsenic. Researchers at the Geological Survey of Canada have modeled relative variation in natural arsenic hazard potential for the province of New Brunswick. This study has important implications for potable water and health concerns relating to inorganic arsenic. Epidemiological evidence from Chile shows a dose-dependent connection between chronic arsenic exposure and various forms of cancer, in particular when other risk factors, such as cigarette smoking, are present. These effects have been demonstrated at contaminations less than 50 ppb. Arsenic is itself a constituent of tobacco smoke. Analyzing multiple epidemiological studies on inorganic arsenic exposure suggests a small but measurable increase in risk for bladder cancer at 10 ppb. According to Peter Ravenscroft of the Department of Geography at the University of Cambridge, roughly 80 million people worldwide consume between 10 and 50 ppb arsenic in their drinking water.

If they all consumed exactly 10 ppb arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation. Early (1973) evaluations of the processes for removing dissolved arsenic from drinking water demonstrated the efficacy of co-precipitation with either iron or aluminum oxides. In particular, iron as a coagulant was found to remove arsenic with an efficacy exceeding 90%. Several adsorptive media systems have been approved for use at point-of-service in a study funded by the United States Environmental Protection Agency (US EPA) and the National Science Foundation (NSF). A team of European and Indian scientists and engineers have set up six arsenic treatment plants in West Bengal based on in-situ remediation method (SAR Technology). This technology does not use any chemicals and arsenic is left in an insoluble form (+5 state) in the subterranean zone by recharging aerated water into the aquifer and developing an oxidation zone that supports arsenic oxidizing micro-organisms. This process does not produce any waste stream or sludge and is relatively cheap. Another effective and inexpensive method to avoid arsenic contamination is to sink wells 500 feet or deeper to reach purer waters. A recent 2011 study funded by the US National Institute of Environmental Health Sciences' Superfund Research Program shows that deep sediments can remove arsenic and take it out of circulation. In this process, called adsorption, arsenic sticks to the surfaces of deep sediment particles and is naturally removed from the ground water. Magnetic separations of arsenic at very low magnetic field gradients with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals have been demonstrated in point-of-use water purification. Using the high specific surface area of Fe3O4 nanocrystals, the mass of waste associated with arsenic removal from water has been dramatically reduced. Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of all leading causes of mortality. The literature indicates that arsenic exposure is causative in the pathogenesis of diabetes. Chaff-based filters have recently been shown to reduce the arsenic content of water to 3 µg/L. This may find applications in areas where the potable water is extracted from underground aquifers. San Pedro de Atacama For several centuries, the people of San Pedro de Atacama in Chile have been drinking water that is contaminated with arsenic, and some evidence suggests they have developed some immunity. Hazard maps for contaminated groundwater Around one-third of the world's population drinks water from groundwater resources. Of this, about 10 percent, approximately 300 million people, obtains water from groundwater resources that are contaminated with unhealthy levels of arsenic or fluoride. These trace elements derive mainly from minerals and ions in the ground.

If they all consumed exactly 10 ppb arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation. Early (1973) evaluations of the processes for removing dissolved arsenic from drinking water demonstrated the efficacy of co-precipitation with either iron or aluminum oxides. In particular, iron as a coagulant was found to remove arsenic with an efficacy exceeding 90%. Several adsorptive media systems have been approved for use at point-of-service in a study funded by the United States Environmental Protection Agency (US EPA) and the National Science Foundation (NSF). A team of European and Indian scientists and engineers have set up six arsenic treatment plants in West Bengal based on in-situ remediation method (SAR Technology). This technology does not use any chemicals and arsenic is left in an insoluble form (+5 state) in the subterranean zone by recharging aerated water into the aquifer and developing an oxidation zone that supports arsenic oxidizing micro-organisms. This process does not produce any waste stream or sludge and is relatively cheap. Another effective and inexpensive method to avoid arsenic contamination is to sink wells 500 feet or deeper to reach purer waters. A recent 2011 study funded by the US National Institute of Environmental Health Sciences' Superfund Research Program shows that deep sediments can remove arsenic and take it out of circulation. In this process, called adsorption, arsenic sticks to the surfaces of deep sediment particles and is naturally removed from the ground water. Magnetic separations of arsenic at very low magnetic field gradients with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals have been demonstrated in point-of-use water purification. Using the high specific surface area of Fe3O4 nanocrystals, the mass of waste associated with arsenic removal from water has been dramatically reduced. Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of all leading causes of mortality. The literature indicates that arsenic exposure is causative in the pathogenesis of diabetes. Chaff-based filters have recently been shown to reduce the arsenic content of water to 3 µg/L. This may find applications in areas where the potable water is extracted from underground aquifers. San Pedro de Atacama For several centuries, the people of San Pedro de Atacama in Chile have been drinking water that is contaminated with arsenic, and some evidence suggests they have developed some immunity. Hazard maps for contaminated groundwater Around one-third of the world's population drinks water from groundwater resources. Of this, about 10 percent, approximately 300 million people, obtains water from groundwater resources that are contaminated with unhealthy levels of arsenic or fluoride. These trace elements derive mainly from minerals and ions in the ground.

If they all consumed exactly 10 ppb arsenic in their drinking water, the previously cited multiple epidemiological study analysis would predict an additional 2,000 cases of bladder cancer alone. This represents a clear underestimate of the overall impact, since it does not include lung or skin cancer, and explicitly underestimates the exposure. Those exposed to levels of arsenic above the current WHO standard should weigh the costs and benefits of arsenic remediation. Early (1973) evaluations of the processes for removing dissolved arsenic from drinking water demonstrated the efficacy of co-precipitation with either iron or aluminum oxides. In particular, iron as a coagulant was found to remove arsenic with an efficacy exceeding 90%. Several adsorptive media systems have been approved for use at point-of-service in a study funded by the United States Environmental Protection Agency (US EPA) and the National Science Foundation (NSF). A team of European and Indian scientists and engineers have set up six arsenic treatment plants in West Bengal based on in-situ remediation method (SAR Technology). This technology does not use any chemicals and arsenic is left in an insoluble form (+5 state) in the subterranean zone by recharging aerated water into the aquifer and developing an oxidation zone that supports arsenic oxidizing micro-organisms. This process does not produce any waste stream or sludge and is relatively cheap. Another effective and inexpensive method to avoid arsenic contamination is to sink wells 500 feet or deeper to reach purer waters. A recent 2011 study funded by the US National Institute of Environmental Health Sciences' Superfund Research Program shows that deep sediments can remove arsenic and take it out of circulation. In this process, called adsorption, arsenic sticks to the surfaces of deep sediment particles and is naturally removed from the ground water. Magnetic separations of arsenic at very low magnetic field gradients with high-surface-area and monodisperse magnetite (Fe3O4) nanocrystals have been demonstrated in point-of-use water purification. Using the high specific surface area of Fe3O4 nanocrystals, the mass of waste associated with arsenic removal from water has been dramatically reduced. Epidemiological studies have suggested a correlation between chronic consumption of drinking water contaminated with arsenic and the incidence of all leading causes of mortality. The literature indicates that arsenic exposure is causative in the pathogenesis of diabetes. Chaff-based filters have recently been shown to reduce the arsenic content of water to 3 µg/L. This may find applications in areas where the potable water is extracted from underground aquifers. San Pedro de Atacama For several centuries, the people of San Pedro de Atacama in Chile have been drinking water that is contaminated with arsenic, and some evidence suggests they have developed some immunity. Hazard maps for contaminated groundwater Around one-third of the world's population drinks water from groundwater resources. Of this, about 10 percent, approximately 300 million people, obtains water from groundwater resources that are contaminated with unhealthy levels of arsenic or fluoride. These trace elements derive mainly from minerals and ions in the ground.

Redox transformation of arsenic in natural waters Arsenic is unique among the trace metalloids and oxyanion-forming trace metals (e.g. As, Se, Sb, Mo, V, Cr, U, Re). It is sensitive to mobilization at pH values typical of natural waters (pH 6.5–8.5) under both oxidizing and reducing conditions. Arsenic can occur in the environment in several oxidation states (−3, 0, +3 and +5), but in natural waters it is mostly found in inorganic forms as oxyanions of trivalent arsenite [As(III)] or pentavalent arsenate [As(V)]. Organic forms of arsenic are produced by biological activity, mostly in surface waters, but are rarely quantitatively important. Organic arsenic compounds may, however, occur where waters are significantly impacted by industrial pollution. Arsenic may be solubilized by various processes. When pH is high, arsenic may be released from surface binding sites that lose their positive charge. When water level drops and sulfide minerals are exposed to air, arsenic trapped in sulfide minerals can be released into water. When organic carbon is present in water, bacteria are fed by directly reducing As(V) to As(III) or by reducing the element at the binding site, releasing inorganic arsenic. The aquatic transformations of arsenic are affected by pH, reduction-oxidation potential, organic matter concentration and the concentrations and forms of other elements, especially iron and manganese. The main factors are pH and the redox potential. Generally, the main forms of arsenic under oxic conditions are H3AsO4, H2AsO4−, HAsO42−, and AsO43− at pH 2, 2–7, 7–11 and 11, respectively. Under reducing conditions, H3AsO4 is predominant at pH 2–9. Oxidation and reduction affects the migration of arsenic in subsurface environments. Arsenite is the most stable soluble form of arsenic in reducing environments and arsenate, which is less mobile than arsenite, is dominant in oxidizing environments at neutral pH. Therefore, arsenic may be more mobile under reducing conditions. The reducing environment is also rich in organic matter which may enhance the solubility of arsenic compounds. As a result, the adsorption of arsenic is reduced and dissolved arsenic accumulates in groundwater. That is why the arsenic content is higher in reducing environments than in oxidizing environments. The presence of sulfur is another factor that affects the transformation of arsenic in natural water. Arsenic can precipitate when metal sulfides form. In this way, arsenic is removed from the water and its mobility decreases. When oxygen is present, bacteria oxidize reduced sulfur to generate energy, potentially releasing bound arsenic. Redox reactions involving Fe also appear to be essential factors in the fate of arsenic in aquatic systems. The reduction of iron oxyhydroxides plays a key role in the release of arsenic to water. So arsenic can be enriched in water with elevated Fe concentrations. Under oxidizing conditions, arsenic can be mobilized from pyrite or iron oxides especially at elevated pH. Under reducing conditions, arsenic can be mobilized by reductive desorption or dissolution when associated with iron oxides. The reductive desorption occurs under two circumstances.

Redox transformation of arsenic in natural waters Arsenic is unique among the trace metalloids and oxyanion-forming trace metals (e.g. As, Se, Sb, Mo, V, Cr, U, Re). It is sensitive to mobilization at pH values typical of natural waters (pH 6.5–8.5) under both oxidizing and reducing conditions. Arsenic can occur in the environment in several oxidation states (−3, 0, +3 and +5), but in natural waters it is mostly found in inorganic forms as oxyanions of trivalent arsenite [As(III)] or pentavalent arsenate [As(V)]. Organic forms of arsenic are produced by biological activity, mostly in surface waters, but are rarely quantitatively important. Organic arsenic compounds may, however, occur where waters are significantly impacted by industrial pollution. Arsenic may be solubilized by various processes. When pH is high, arsenic may be released from surface binding sites that lose their positive charge. When water level drops and sulfide minerals are exposed to air, arsenic trapped in sulfide minerals can be released into water. When organic carbon is present in water, bacteria are fed by directly reducing As(V) to As(III) or by reducing the element at the binding site, releasing inorganic arsenic. The aquatic transformations of arsenic are affected by pH, reduction-oxidation potential, organic matter concentration and the concentrations and forms of other elements, especially iron and manganese. The main factors are pH and the redox potential. Generally, the main forms of arsenic under oxic conditions are H3AsO4, H2AsO4−, HAsO42−, and AsO43− at pH 2, 2–7, 7–11 and 11, respectively. Under reducing conditions, H3AsO4 is predominant at pH 2–9. Oxidation and reduction affects the migration of arsenic in subsurface environments. Arsenite is the most stable soluble form of arsenic in reducing environments and arsenate, which is less mobile than arsenite, is dominant in oxidizing environments at neutral pH. Therefore, arsenic may be more mobile under reducing conditions. The reducing environment is also rich in organic matter which may enhance the solubility of arsenic compounds. As a result, the adsorption of arsenic is reduced and dissolved arsenic accumulates in groundwater. That is why the arsenic content is higher in reducing environments than in oxidizing environments. The presence of sulfur is another factor that affects the transformation of arsenic in natural water. Arsenic can precipitate when metal sulfides form. In this way, arsenic is removed from the water and its mobility decreases. When oxygen is present, bacteria oxidize reduced sulfur to generate energy, potentially releasing bound arsenic. Redox reactions involving Fe also appear to be essential factors in the fate of arsenic in aquatic systems. The reduction of iron oxyhydroxides plays a key role in the release of arsenic to water. So arsenic can be enriched in water with elevated Fe concentrations. Under oxidizing conditions, arsenic can be mobilized from pyrite or iron oxides especially at elevated pH. Under reducing conditions, arsenic can be mobilized by reductive desorption or dissolution when associated with iron oxides. The reductive desorption occurs under two circumstances.

Redox transformation of arsenic in natural waters Arsenic is unique among the trace metalloids and oxyanion-forming trace metals (e.g. As, Se, Sb, Mo, V, Cr, U, Re). It is sensitive to mobilization at pH values typical of natural waters (pH 6.5–8.5) under both oxidizing and reducing conditions. Arsenic can occur in the environment in several oxidation states (−3, 0, +3 and +5), but in natural waters it is mostly found in inorganic forms as oxyanions of trivalent arsenite [As(III)] or pentavalent arsenate [As(V)]. Organic forms of arsenic are produced by biological activity, mostly in surface waters, but are rarely quantitatively important. Organic arsenic compounds may, however, occur where waters are significantly impacted by industrial pollution. Arsenic may be solubilized by various processes. When pH is high, arsenic may be released from surface binding sites that lose their positive charge. When water level drops and sulfide minerals are exposed to air, arsenic trapped in sulfide minerals can be released into water. When organic carbon is present in water, bacteria are fed by directly reducing As(V) to As(III) or by reducing the element at the binding site, releasing inorganic arsenic. The aquatic transformations of arsenic are affected by pH, reduction-oxidation potential, organic matter concentration and the concentrations and forms of other elements, especially iron and manganese. The main factors are pH and the redox potential. Generally, the main forms of arsenic under oxic conditions are H3AsO4, H2AsO4−, HAsO42−, and AsO43− at pH 2, 2–7, 7–11 and 11, respectively. Under reducing conditions, H3AsO4 is predominant at pH 2–9. Oxidation and reduction affects the migration of arsenic in subsurface environments. Arsenite is the most stable soluble form of arsenic in reducing environments and arsenate, which is less mobile than arsenite, is dominant in oxidizing environments at neutral pH. Therefore, arsenic may be more mobile under reducing conditions. The reducing environment is also rich in organic matter which may enhance the solubility of arsenic compounds. As a result, the adsorption of arsenic is reduced and dissolved arsenic accumulates in groundwater. That is why the arsenic content is higher in reducing environments than in oxidizing environments. The presence of sulfur is another factor that affects the transformation of arsenic in natural water. Arsenic can precipitate when metal sulfides form. In this way, arsenic is removed from the water and its mobility decreases. When oxygen is present, bacteria oxidize reduced sulfur to generate energy, potentially releasing bound arsenic. Redox reactions involving Fe also appear to be essential factors in the fate of arsenic in aquatic systems. The reduction of iron oxyhydroxides plays a key role in the release of arsenic to water. So arsenic can be enriched in water with elevated Fe concentrations. Under oxidizing conditions, arsenic can be mobilized from pyrite or iron oxides especially at elevated pH. Under reducing conditions, arsenic can be mobilized by reductive desorption or dissolution when associated with iron oxides. The reductive desorption occurs under two circumstances.

One is when arsenate is reduced to arsenite which adsorbs to iron oxides less strongly. The other results from a change in the charge on the mineral surface which leads to the desorption of bound arsenic. Some species of bacteria catalyze redox transformations of arsenic. Dissimilatory arsenate-respiring prokaryotes (DARP) speed up the reduction of As(V) to As(III). DARP use As(V) as the electron acceptor of anaerobic respiration and obtain energy to survive. Other organic and inorganic substances can be oxidized in this process. Chemoautotrophic arsenite oxidizers (CAO) and heterotrophic arsenite oxidizers (HAO) convert As(III) into As(V). CAO combine the oxidation of As(III) with the reduction of oxygen or nitrate. They use obtained energy to fix produce organic carbon from CO2. HAO cannot obtain energy from As(III) oxidation. This process may be an arsenic detoxification mechanism for the bacteria. Equilibrium thermodynamic calculations predict that As(V) concentrations should be greater than As(III) concentrations in all but strongly reducing conditions, i.e. where SO42− reduction is occurring. However, abiotic redox reactions of arsenic are slow. Oxidation of As(III) by dissolved O2 is a particularly slow reaction. For example, Johnson and Pilson (1975) gave half-lives for the oxygenation of As(III) in seawater ranging from several months to a year. In other studies, As(V)/As(III) ratios were stable over periods of days or weeks during water sampling when no particular care was taken to prevent oxidation, again suggesting relatively slow oxidation rates. Cherry found from experimental studies that the As(V)/As(III) ratios were stable in anoxic solutions for up to 3 weeks but that gradual changes occurred over longer timescales. Sterile water samples have been observed to be less susceptible to speciation changes than non-sterile samples. Oremland found that the reduction of As(V) to As(III) in Mono Lake was rapidly catalyzed by bacteria with rate constants ranging from 0.02 to 0.3-day−1. Wood preservation in the US As of 2002, US-based industries consumed 19,600 metric tons of arsenic. Ninety percent of this was used for treatment of wood with chromated copper arsenate (CCA). In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the voluntary phasing-out of arsenic in production of consumer products and residential and general consumer construction products began on 31 December 2003, and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole. Although discontinued, this application is also one of the most concerning to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber.

One is when arsenate is reduced to arsenite which adsorbs to iron oxides less strongly. The other results from a change in the charge on the mineral surface which leads to the desorption of bound arsenic. Some species of bacteria catalyze redox transformations of arsenic. Dissimilatory arsenate-respiring prokaryotes (DARP) speed up the reduction of As(V) to As(III). DARP use As(V) as the electron acceptor of anaerobic respiration and obtain energy to survive. Other organic and inorganic substances can be oxidized in this process. Chemoautotrophic arsenite oxidizers (CAO) and heterotrophic arsenite oxidizers (HAO) convert As(III) into As(V). CAO combine the oxidation of As(III) with the reduction of oxygen or nitrate. They use obtained energy to fix produce organic carbon from CO2. HAO cannot obtain energy from As(III) oxidation. This process may be an arsenic detoxification mechanism for the bacteria. Equilibrium thermodynamic calculations predict that As(V) concentrations should be greater than As(III) concentrations in all but strongly reducing conditions, i.e. where SO42− reduction is occurring. However, abiotic redox reactions of arsenic are slow. Oxidation of As(III) by dissolved O2 is a particularly slow reaction. For example, Johnson and Pilson (1975) gave half-lives for the oxygenation of As(III) in seawater ranging from several months to a year. In other studies, As(V)/As(III) ratios were stable over periods of days or weeks during water sampling when no particular care was taken to prevent oxidation, again suggesting relatively slow oxidation rates. Cherry found from experimental studies that the As(V)/As(III) ratios were stable in anoxic solutions for up to 3 weeks but that gradual changes occurred over longer timescales. Sterile water samples have been observed to be less susceptible to speciation changes than non-sterile samples. Oremland found that the reduction of As(V) to As(III) in Mono Lake was rapidly catalyzed by bacteria with rate constants ranging from 0.02 to 0.3-day−1. Wood preservation in the US As of 2002, US-based industries consumed 19,600 metric tons of arsenic. Ninety percent of this was used for treatment of wood with chromated copper arsenate (CCA). In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the voluntary phasing-out of arsenic in production of consumer products and residential and general consumer construction products began on 31 December 2003, and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole. Although discontinued, this application is also one of the most concerning to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber.

One is when arsenate is reduced to arsenite which adsorbs to iron oxides less strongly. The other results from a change in the charge on the mineral surface which leads to the desorption of bound arsenic. Some species of bacteria catalyze redox transformations of arsenic. Dissimilatory arsenate-respiring prokaryotes (DARP) speed up the reduction of As(V) to As(III). DARP use As(V) as the electron acceptor of anaerobic respiration and obtain energy to survive. Other organic and inorganic substances can be oxidized in this process. Chemoautotrophic arsenite oxidizers (CAO) and heterotrophic arsenite oxidizers (HAO) convert As(III) into As(V). CAO combine the oxidation of As(III) with the reduction of oxygen or nitrate. They use obtained energy to fix produce organic carbon from CO2. HAO cannot obtain energy from As(III) oxidation. This process may be an arsenic detoxification mechanism for the bacteria. Equilibrium thermodynamic calculations predict that As(V) concentrations should be greater than As(III) concentrations in all but strongly reducing conditions, i.e. where SO42− reduction is occurring. However, abiotic redox reactions of arsenic are slow. Oxidation of As(III) by dissolved O2 is a particularly slow reaction. For example, Johnson and Pilson (1975) gave half-lives for the oxygenation of As(III) in seawater ranging from several months to a year. In other studies, As(V)/As(III) ratios were stable over periods of days or weeks during water sampling when no particular care was taken to prevent oxidation, again suggesting relatively slow oxidation rates. Cherry found from experimental studies that the As(V)/As(III) ratios were stable in anoxic solutions for up to 3 weeks but that gradual changes occurred over longer timescales. Sterile water samples have been observed to be less susceptible to speciation changes than non-sterile samples. Oremland found that the reduction of As(V) to As(III) in Mono Lake was rapidly catalyzed by bacteria with rate constants ranging from 0.02 to 0.3-day−1. Wood preservation in the US As of 2002, US-based industries consumed 19,600 metric tons of arsenic. Ninety percent of this was used for treatment of wood with chromated copper arsenate (CCA). In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the voluntary phasing-out of arsenic in production of consumer products and residential and general consumer construction products began on 31 December 2003, and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole. Although discontinued, this application is also one of the most concerning to the general public. The vast majority of older pressure-treated wood was treated with CCA. CCA lumber is still in widespread use in many countries, and was heavily used during the latter half of the 20th century as a structural and outdoor building material. Although the use of CCA lumber was banned in many areas after studies showed that arsenic could leach out of the wood into the surrounding soil (from playground equipment, for instance), a risk is also presented by the burning of older CCA timber.

The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber are not consistent throughout the world. Widespread landfill disposal of such timber raises some concern, but other studies have shown no arsenic contamination in the groundwater. Mapping of industrial releases in the US One tool that maps the location (and other information) of arsenic releases in the United States is TOXMAP. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the United States National Library of Medicine (NLM) funded by the US Federal Government. With marked-up maps of the United States, TOXMAP enables users to visually explore data from the United States Environmental Protection Agency's (EPA) Toxics Release Inventory and Superfund Basic Research Programs. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET), PubMed, and from other authoritative sources. Bioremediation Physical, chemical, and biological methods have been used to remediate arsenic contaminated water. Bioremediation is said to be cost-effective and environmentally friendly. Bioremediation of ground water contaminated with arsenic aims to convert arsenite, the toxic form of arsenic to humans, to arsenate. Arsenate (+5 oxidation state) is the dominant form of arsenic in surface water, while arsenite (+3 oxidation state) is the dominant form in hypoxic to anoxic environments. Arsenite is more soluble and mobile than arsenate. Many species of bacteria can transform arsenite to arsenate in anoxic conditions by using arsenite as an electron donor. This is a useful method in ground water remediation. Another bioremediation strategy is to use plants that accumulate arsenic in their tissues via phytoremediation but the disposal of contaminated plant material needs to be considered. Bioremediation requires careful evaluation and design in accordance with existing conditions. Some sites may require the addition of an electron acceptor while others require microbe supplementation (bioaugmentation). Regardless of the method used, only constant monitoring can prevent future contamination. Toxicity and precautions Arsenic and many of its compounds are especially potent poisons. Classification Elemental arsenic and arsenic sulfate and trioxide compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC. The International Agency for Research on Cancer (IARC) recognizes arsenic and inorganic arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide, and arsenate salts as category 1 carcinogens. Arsenic is known to cause arsenicosis when present in drinking water, "the most common species being arsenate [; As(V)] and arsenite [H3AsO3; As(III)]". Legal limits, food, and drink In the United States since 2006, the maximum concentration in drinking water allowed by the Environmental Protection Agency (EPA) is 10 ppb and the FDA set the same standard in 2005 for bottled water.

The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber are not consistent throughout the world. Widespread landfill disposal of such timber raises some concern, but other studies have shown no arsenic contamination in the groundwater. Mapping of industrial releases in the US One tool that maps the location (and other information) of arsenic releases in the United States is TOXMAP. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the United States National Library of Medicine (NLM) funded by the US Federal Government. With marked-up maps of the United States, TOXMAP enables users to visually explore data from the United States Environmental Protection Agency's (EPA) Toxics Release Inventory and Superfund Basic Research Programs. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET), PubMed, and from other authoritative sources. Bioremediation Physical, chemical, and biological methods have been used to remediate arsenic contaminated water. Bioremediation is said to be cost-effective and environmentally friendly. Bioremediation of ground water contaminated with arsenic aims to convert arsenite, the toxic form of arsenic to humans, to arsenate. Arsenate (+5 oxidation state) is the dominant form of arsenic in surface water, while arsenite (+3 oxidation state) is the dominant form in hypoxic to anoxic environments. Arsenite is more soluble and mobile than arsenate. Many species of bacteria can transform arsenite to arsenate in anoxic conditions by using arsenite as an electron donor. This is a useful method in ground water remediation. Another bioremediation strategy is to use plants that accumulate arsenic in their tissues via phytoremediation but the disposal of contaminated plant material needs to be considered. Bioremediation requires careful evaluation and design in accordance with existing conditions. Some sites may require the addition of an electron acceptor while others require microbe supplementation (bioaugmentation). Regardless of the method used, only constant monitoring can prevent future contamination. Toxicity and precautions Arsenic and many of its compounds are especially potent poisons. Classification Elemental arsenic and arsenic sulfate and trioxide compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC. The International Agency for Research on Cancer (IARC) recognizes arsenic and inorganic arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide, and arsenate salts as category 1 carcinogens. Arsenic is known to cause arsenicosis when present in drinking water, "the most common species being arsenate [; As(V)] and arsenite [H3AsO3; As(III)]". Legal limits, food, and drink In the United States since 2006, the maximum concentration in drinking water allowed by the Environmental Protection Agency (EPA) is 10 ppb and the FDA set the same standard in 2005 for bottled water.

The direct or indirect ingestion of wood ash from burnt CCA lumber has caused fatalities in animals and serious poisonings in humans; the lethal human dose is approximately 20 grams of ash. Scrap CCA lumber from construction and demolition sites may be inadvertently used in commercial and domestic fires. Protocols for safe disposal of CCA lumber are not consistent throughout the world. Widespread landfill disposal of such timber raises some concern, but other studies have shown no arsenic contamination in the groundwater. Mapping of industrial releases in the US One tool that maps the location (and other information) of arsenic releases in the United States is TOXMAP. TOXMAP is a Geographic Information System (GIS) from the Division of Specialized Information Services of the United States National Library of Medicine (NLM) funded by the US Federal Government. With marked-up maps of the United States, TOXMAP enables users to visually explore data from the United States Environmental Protection Agency's (EPA) Toxics Release Inventory and Superfund Basic Research Programs. TOXMAP's chemical and environmental health information is taken from NLM's Toxicology Data Network (TOXNET), PubMed, and from other authoritative sources. Bioremediation Physical, chemical, and biological methods have been used to remediate arsenic contaminated water. Bioremediation is said to be cost-effective and environmentally friendly. Bioremediation of ground water contaminated with arsenic aims to convert arsenite, the toxic form of arsenic to humans, to arsenate. Arsenate (+5 oxidation state) is the dominant form of arsenic in surface water, while arsenite (+3 oxidation state) is the dominant form in hypoxic to anoxic environments. Arsenite is more soluble and mobile than arsenate. Many species of bacteria can transform arsenite to arsenate in anoxic conditions by using arsenite as an electron donor. This is a useful method in ground water remediation. Another bioremediation strategy is to use plants that accumulate arsenic in their tissues via phytoremediation but the disposal of contaminated plant material needs to be considered. Bioremediation requires careful evaluation and design in accordance with existing conditions. Some sites may require the addition of an electron acceptor while others require microbe supplementation (bioaugmentation). Regardless of the method used, only constant monitoring can prevent future contamination. Toxicity and precautions Arsenic and many of its compounds are especially potent poisons. Classification Elemental arsenic and arsenic sulfate and trioxide compounds are classified as "toxic" and "dangerous for the environment" in the European Union under directive 67/548/EEC. The International Agency for Research on Cancer (IARC) recognizes arsenic and inorganic arsenic compounds as group 1 carcinogens, and the EU lists arsenic trioxide, arsenic pentoxide, and arsenate salts as category 1 carcinogens. Arsenic is known to cause arsenicosis when present in drinking water, "the most common species being arsenate [; As(V)] and arsenite [H3AsO3; As(III)]". Legal limits, food, and drink In the United States since 2006, the maximum concentration in drinking water allowed by the Environmental Protection Agency (EPA) is 10 ppb and the FDA set the same standard in 2005 for bottled water.

The Department of Environmental Protection for New Jersey set a drinking water limit of 5 ppb in 2006. The IDLH (immediately dangerous to life and health) value for arsenic metal and inorganic arsenic compounds is 5 mg/m3 (5 ppb). The Occupational Safety and Health Administration has set the permissible exposure limit (PEL) to a time-weighted average (TWA) of 0.01 mg/m3 (0.01 ppb), and the National Institute for Occupational Safety and Health (NIOSH) has set the recommended exposure limit (REL) to a 15-minute constant exposure of 0.002 mg/m3 (0.002 ppb). The PEL for organic arsenic compounds is a TWA of 0.5 mg/m3. (0.5 ppb). In 2008, based on its ongoing testing of a wide variety of American foods for toxic chemicals, the U.S. Food and Drug Administration set the "level of concern" for inorganic arsenic in apple and pear juices at 23 ppb, based on non-carcinogenic effects, and began blocking importation of products in excess of this level; it also required recalls for non-conforming domestic products. In 2011, the national Dr. Oz television show broadcast a program highlighting tests performed by an independent lab hired by the producers. Though the methodology was disputed (it did not distinguish between organic and inorganic arsenic) the tests showed levels of arsenic up to 36 ppb. In response, FDA tested the worst brand from the Dr. Oz show and found much lower levels. Ongoing testing found 95% of the apple juice samples were below the level of concern. Later testing by Consumer Reports showed inorganic arsenic at levels slightly above 10 ppb, and the organization urged parents to reduce consumption. In July 2013, on consideration of consumption by children, chronic exposure, and carcinogenic effect, the FDA established an "action level" of 10 ppb for apple juice, the same as the drinking water standard. Concern about arsenic in rice in Bangladesh was raised in 2002, but at the time only Australia had a legal limit for food (one milligram per kilogram). Concern was raised about people who were eating U.S. rice exceeding WHO standards for personal arsenic intake in 2005. In 2011, the People's Republic of China set a food standard of 150 ppb for arsenic. In the United States in 2012, testing by separate groups of researchers at the Children's Environmental Health and Disease Prevention Research Center at Dartmouth College (early in the year, focusing on urinary levels in children) and Consumer Reports (in November) found levels of arsenic in rice that resulted in calls for the FDA to set limits. The FDA released some testing results in September 2012, and as of July 2013, is still collecting data in support of a new potential regulation. It has not recommended any changes in consumer behavior.

The Department of Environmental Protection for New Jersey set a drinking water limit of 5 ppb in 2006. The IDLH (immediately dangerous to life and health) value for arsenic metal and inorganic arsenic compounds is 5 mg/m3 (5 ppb). The Occupational Safety and Health Administration has set the permissible exposure limit (PEL) to a time-weighted average (TWA) of 0.01 mg/m3 (0.01 ppb), and the National Institute for Occupational Safety and Health (NIOSH) has set the recommended exposure limit (REL) to a 15-minute constant exposure of 0.002 mg/m3 (0.002 ppb). The PEL for organic arsenic compounds is a TWA of 0.5 mg/m3. (0.5 ppb). In 2008, based on its ongoing testing of a wide variety of American foods for toxic chemicals, the U.S. Food and Drug Administration set the "level of concern" for inorganic arsenic in apple and pear juices at 23 ppb, based on non-carcinogenic effects, and began blocking importation of products in excess of this level; it also required recalls for non-conforming domestic products. In 2011, the national Dr. Oz television show broadcast a program highlighting tests performed by an independent lab hired by the producers. Though the methodology was disputed (it did not distinguish between organic and inorganic arsenic) the tests showed levels of arsenic up to 36 ppb. In response, FDA tested the worst brand from the Dr. Oz show and found much lower levels. Ongoing testing found 95% of the apple juice samples were below the level of concern. Later testing by Consumer Reports showed inorganic arsenic at levels slightly above 10 ppb, and the organization urged parents to reduce consumption. In July 2013, on consideration of consumption by children, chronic exposure, and carcinogenic effect, the FDA established an "action level" of 10 ppb for apple juice, the same as the drinking water standard. Concern about arsenic in rice in Bangladesh was raised in 2002, but at the time only Australia had a legal limit for food (one milligram per kilogram). Concern was raised about people who were eating U.S. rice exceeding WHO standards for personal arsenic intake in 2005. In 2011, the People's Republic of China set a food standard of 150 ppb for arsenic. In the United States in 2012, testing by separate groups of researchers at the Children's Environmental Health and Disease Prevention Research Center at Dartmouth College (early in the year, focusing on urinary levels in children) and Consumer Reports (in November) found levels of arsenic in rice that resulted in calls for the FDA to set limits. The FDA released some testing results in September 2012, and as of July 2013, is still collecting data in support of a new potential regulation. It has not recommended any changes in consumer behavior.

The Department of Environmental Protection for New Jersey set a drinking water limit of 5 ppb in 2006. The IDLH (immediately dangerous to life and health) value for arsenic metal and inorganic arsenic compounds is 5 mg/m3 (5 ppb). The Occupational Safety and Health Administration has set the permissible exposure limit (PEL) to a time-weighted average (TWA) of 0.01 mg/m3 (0.01 ppb), and the National Institute for Occupational Safety and Health (NIOSH) has set the recommended exposure limit (REL) to a 15-minute constant exposure of 0.002 mg/m3 (0.002 ppb). The PEL for organic arsenic compounds is a TWA of 0.5 mg/m3. (0.5 ppb). In 2008, based on its ongoing testing of a wide variety of American foods for toxic chemicals, the U.S. Food and Drug Administration set the "level of concern" for inorganic arsenic in apple and pear juices at 23 ppb, based on non-carcinogenic effects, and began blocking importation of products in excess of this level; it also required recalls for non-conforming domestic products. In 2011, the national Dr. Oz television show broadcast a program highlighting tests performed by an independent lab hired by the producers. Though the methodology was disputed (it did not distinguish between organic and inorganic arsenic) the tests showed levels of arsenic up to 36 ppb. In response, FDA tested the worst brand from the Dr. Oz show and found much lower levels. Ongoing testing found 95% of the apple juice samples were below the level of concern. Later testing by Consumer Reports showed inorganic arsenic at levels slightly above 10 ppb, and the organization urged parents to reduce consumption. In July 2013, on consideration of consumption by children, chronic exposure, and carcinogenic effect, the FDA established an "action level" of 10 ppb for apple juice, the same as the drinking water standard. Concern about arsenic in rice in Bangladesh was raised in 2002, but at the time only Australia had a legal limit for food (one milligram per kilogram). Concern was raised about people who were eating U.S. rice exceeding WHO standards for personal arsenic intake in 2005. In 2011, the People's Republic of China set a food standard of 150 ppb for arsenic. In the United States in 2012, testing by separate groups of researchers at the Children's Environmental Health and Disease Prevention Research Center at Dartmouth College (early in the year, focusing on urinary levels in children) and Consumer Reports (in November) found levels of arsenic in rice that resulted in calls for the FDA to set limits. The FDA released some testing results in September 2012, and as of July 2013, is still collecting data in support of a new potential regulation. It has not recommended any changes in consumer behavior.

Consumer Reports recommended: That the EPA and FDA eliminate arsenic-containing fertilizer, drugs, and pesticides in food production; That the FDA establish a legal limit for food; That industry change production practices to lower arsenic levels, especially in food for children; and That consumers test home water supplies, eat a varied diet, and cook rice with excess water, then draining it off (reducing inorganic arsenic by about one third along with a slight reduction in vitamin content). Evidence-based public health advocates also recommend that, given the lack of regulation or labeling for arsenic in the U.S., children should eat no more than 1.5 servings per week of rice and should not drink rice milk as part of their daily diet before age 5. They also offer recommendations for adults and infants on how to limit arsenic exposure from rice, drinking water, and fruit juice. A 2014 World Health Organization advisory conference was scheduled to consider limits of 200–300 ppb for rice. Reducing arsenic content in rice In 2020, scientists assessed multiple preparation procedures of rice for their capacity to reduce arsenic content and preserve nutrients, recommending a procedure involving parboiling and water-absorption. Occupational exposure limits Ecotoxicity Arsenic is bioaccumulative in many organisms, marine species in particular, but it does not appear to biomagnify significantly in food webs. In polluted areas, plant growth may be affected by root uptake of arsenate, which is a phosphate analog and therefore readily transported in plant tissues and cells. In polluted areas, uptake of the more toxic arsenite ion (found more particularly in reducing conditions) is likely in poorly-drained soils. Toxicity in animals Biological mechanism Arsenic's toxicity comes from the affinity of arsenic(III) oxides for thiols. Thiols, in the form of cysteine residues and cofactors such as lipoic acid and coenzyme A, are situated at the active sites of many important enzymes. Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits lipoic acid, which is a cofactor for pyruvate dehydrogenase. By competing with phosphate, arsenate uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration and ATP synthesis. Hydrogen peroxide production is also increased, which, it is speculated, has potential to form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure. The organ failure is presumed to be from necrotic cell death, not apoptosis, since energy reserves have been too depleted for apoptosis to occur. Exposure risks and remediation Occupational exposure and arsenic poisoning may occur in persons working in industries involving the use of inorganic arsenic and its compounds, such as wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry. The conversion between As(III) and As(V) is a large factor in arsenic environmental contamination. According to Croal, Gralnick, Malasarn and Newman, "[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal).

Consumer Reports recommended: That the EPA and FDA eliminate arsenic-containing fertilizer, drugs, and pesticides in food production; That the FDA establish a legal limit for food; That industry change production practices to lower arsenic levels, especially in food for children; and That consumers test home water supplies, eat a varied diet, and cook rice with excess water, then draining it off (reducing inorganic arsenic by about one third along with a slight reduction in vitamin content). Evidence-based public health advocates also recommend that, given the lack of regulation or labeling for arsenic in the U.S., children should eat no more than 1.5 servings per week of rice and should not drink rice milk as part of their daily diet before age 5. They also offer recommendations for adults and infants on how to limit arsenic exposure from rice, drinking water, and fruit juice. A 2014 World Health Organization advisory conference was scheduled to consider limits of 200–300 ppb for rice. Reducing arsenic content in rice In 2020, scientists assessed multiple preparation procedures of rice for their capacity to reduce arsenic content and preserve nutrients, recommending a procedure involving parboiling and water-absorption. Occupational exposure limits Ecotoxicity Arsenic is bioaccumulative in many organisms, marine species in particular, but it does not appear to biomagnify significantly in food webs. In polluted areas, plant growth may be affected by root uptake of arsenate, which is a phosphate analog and therefore readily transported in plant tissues and cells. In polluted areas, uptake of the more toxic arsenite ion (found more particularly in reducing conditions) is likely in poorly-drained soils. Toxicity in animals Biological mechanism Arsenic's toxicity comes from the affinity of arsenic(III) oxides for thiols. Thiols, in the form of cysteine residues and cofactors such as lipoic acid and coenzyme A, are situated at the active sites of many important enzymes. Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits lipoic acid, which is a cofactor for pyruvate dehydrogenase. By competing with phosphate, arsenate uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration and ATP synthesis. Hydrogen peroxide production is also increased, which, it is speculated, has potential to form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure. The organ failure is presumed to be from necrotic cell death, not apoptosis, since energy reserves have been too depleted for apoptosis to occur. Exposure risks and remediation Occupational exposure and arsenic poisoning may occur in persons working in industries involving the use of inorganic arsenic and its compounds, such as wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry. The conversion between As(III) and As(V) is a large factor in arsenic environmental contamination. According to Croal, Gralnick, Malasarn and Newman, "[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal).

Consumer Reports recommended: That the EPA and FDA eliminate arsenic-containing fertilizer, drugs, and pesticides in food production; That the FDA establish a legal limit for food; That industry change production practices to lower arsenic levels, especially in food for children; and That consumers test home water supplies, eat a varied diet, and cook rice with excess water, then draining it off (reducing inorganic arsenic by about one third along with a slight reduction in vitamin content). Evidence-based public health advocates also recommend that, given the lack of regulation or labeling for arsenic in the U.S., children should eat no more than 1.5 servings per week of rice and should not drink rice milk as part of their daily diet before age 5. They also offer recommendations for adults and infants on how to limit arsenic exposure from rice, drinking water, and fruit juice. A 2014 World Health Organization advisory conference was scheduled to consider limits of 200–300 ppb for rice. Reducing arsenic content in rice In 2020, scientists assessed multiple preparation procedures of rice for their capacity to reduce arsenic content and preserve nutrients, recommending a procedure involving parboiling and water-absorption. Occupational exposure limits Ecotoxicity Arsenic is bioaccumulative in many organisms, marine species in particular, but it does not appear to biomagnify significantly in food webs. In polluted areas, plant growth may be affected by root uptake of arsenate, which is a phosphate analog and therefore readily transported in plant tissues and cells. In polluted areas, uptake of the more toxic arsenite ion (found more particularly in reducing conditions) is likely in poorly-drained soils. Toxicity in animals Biological mechanism Arsenic's toxicity comes from the affinity of arsenic(III) oxides for thiols. Thiols, in the form of cysteine residues and cofactors such as lipoic acid and coenzyme A, are situated at the active sites of many important enzymes. Arsenic disrupts ATP production through several mechanisms. At the level of the citric acid cycle, arsenic inhibits lipoic acid, which is a cofactor for pyruvate dehydrogenase. By competing with phosphate, arsenate uncouples oxidative phosphorylation, thus inhibiting energy-linked reduction of NAD+, mitochondrial respiration and ATP synthesis. Hydrogen peroxide production is also increased, which, it is speculated, has potential to form reactive oxygen species and oxidative stress. These metabolic interferences lead to death from multi-system organ failure. The organ failure is presumed to be from necrotic cell death, not apoptosis, since energy reserves have been too depleted for apoptosis to occur. Exposure risks and remediation Occupational exposure and arsenic poisoning may occur in persons working in industries involving the use of inorganic arsenic and its compounds, such as wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry. The conversion between As(III) and As(V) is a large factor in arsenic environmental contamination. According to Croal, Gralnick, Malasarn and Newman, "[the] understanding [of] what stimulates As(III) oxidation and/or limits As(V) reduction is relevant for bioremediation of contaminated sites (Croal).

The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic. Treatment Treatment of chronic arsenic poisoning is possible. British anti-lewisite (dimercaprol) is prescribed in doses of 5 mg/kg up to 300 mg every 4 hours for the first day, then every 6 hours for the second day, and finally every 8 hours for 8 additional days. However the USA's Agency for Toxic Substances and Disease Registry (ATSDR) states that the long-term effects of arsenic exposure cannot be predicted. Blood, urine, hair, and nails may be tested for arsenic; however, these tests cannot foresee possible health outcomes from the exposure. Long-term exposure and consequent excretion through urine has been linked to bladder and kidney cancer in addition to cancer of the liver, prostate, skin, lungs, and nasal cavity. See also Aqua Tofana Arsenic and Old Lace Arsenic biochemistry Arsenic compounds Arsenic poisoning Arsenic toxicity Arsenic trioxide Fowler's solution GFAJ-1 Grainger challenge Hypothetical types of biochemistry Organoarsenic chemistry Toxic heavy metal White arsenic References Bibliography Further reading External links Arsenic Cancer Causing Substances, U.S. National Cancer Institute. CTD's Arsenic page and CTD's Arsenicals page from the Comparative Toxicogenomics Database Arsenic intoxication: general aspects and chelating agents, by Geir Bjørklund, Massimiliano Peana et al. Archives of Toxicology (2020) 94:1879–1897. A Small Dose of Toxicology Arsenic in groundwater Book on arsenic in groundwater by IAH's Netherlands Chapter and the Netherlands Hydrological Society Contaminant Focus: Arsenic by the EPA. Environmental Health Criteria for Arsenic and Arsenic Compounds, 2001 by the WHO. National Institute for Occupational Safety and Health – Arsenic Page Arsenic at The Periodic Table of Videos (University of Nottingham) Chemical elements Metalloids Hepatotoxins Pnictogens Biology and pharmacology of chemical elements Endocrine disruptors IARC Group 1 carcinogens Trigonal minerals Minerals in space group 166 Teratogens Fetotoxicants Suspected testicular toxicants Native element minerals Chemical elements with rhombohedral structure

The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic. Treatment Treatment of chronic arsenic poisoning is possible. British anti-lewisite (dimercaprol) is prescribed in doses of 5 mg/kg up to 300 mg every 4 hours for the first day, then every 6 hours for the second day, and finally every 8 hours for 8 additional days. However the USA's Agency for Toxic Substances and Disease Registry (ATSDR) states that the long-term effects of arsenic exposure cannot be predicted. Blood, urine, hair, and nails may be tested for arsenic; however, these tests cannot foresee possible health outcomes from the exposure. Long-term exposure and consequent excretion through urine has been linked to bladder and kidney cancer in addition to cancer of the liver, prostate, skin, lungs, and nasal cavity. See also Aqua Tofana Arsenic and Old Lace Arsenic biochemistry Arsenic compounds Arsenic poisoning Arsenic toxicity Arsenic trioxide Fowler's solution GFAJ-1 Grainger challenge Hypothetical types of biochemistry Organoarsenic chemistry Toxic heavy metal White arsenic References Bibliography Further reading External links Arsenic Cancer Causing Substances, U.S. National Cancer Institute. CTD's Arsenic page and CTD's Arsenicals page from the Comparative Toxicogenomics Database Arsenic intoxication: general aspects and chelating agents, by Geir Bjørklund, Massimiliano Peana et al. Archives of Toxicology (2020) 94:1879–1897. A Small Dose of Toxicology Arsenic in groundwater Book on arsenic in groundwater by IAH's Netherlands Chapter and the Netherlands Hydrological Society Contaminant Focus: Arsenic by the EPA. Environmental Health Criteria for Arsenic and Arsenic Compounds, 2001 by the WHO. National Institute for Occupational Safety and Health – Arsenic Page Arsenic at The Periodic Table of Videos (University of Nottingham) Chemical elements Metalloids Hepatotoxins Pnictogens Biology and pharmacology of chemical elements Endocrine disruptors IARC Group 1 carcinogens Trigonal minerals Minerals in space group 166 Teratogens Fetotoxicants Suspected testicular toxicants Native element minerals Chemical elements with rhombohedral structure

The study of chemolithoautotrophic As(III) oxidizers and the heterotrophic As(V) reducers can help the understanding of the oxidation and/or reduction of arsenic. Treatment Treatment of chronic arsenic poisoning is possible. British anti-lewisite (dimercaprol) is prescribed in doses of 5 mg/kg up to 300 mg every 4 hours for the first day, then every 6 hours for the second day, and finally every 8 hours for 8 additional days. However the USA's Agency for Toxic Substances and Disease Registry (ATSDR) states that the long-term effects of arsenic exposure cannot be predicted. Blood, urine, hair, and nails may be tested for arsenic; however, these tests cannot foresee possible health outcomes from the exposure. Long-term exposure and consequent excretion through urine has been linked to bladder and kidney cancer in addition to cancer of the liver, prostate, skin, lungs, and nasal cavity. See also Aqua Tofana Arsenic and Old Lace Arsenic biochemistry Arsenic compounds Arsenic poisoning Arsenic toxicity Arsenic trioxide Fowler's solution GFAJ-1 Grainger challenge Hypothetical types of biochemistry Organoarsenic chemistry Toxic heavy metal White arsenic References Bibliography Further reading External links Arsenic Cancer Causing Substances, U.S. National Cancer Institute. CTD's Arsenic page and CTD's Arsenicals page from the Comparative Toxicogenomics Database Arsenic intoxication: general aspects and chelating agents, by Geir Bjørklund, Massimiliano Peana et al. Archives of Toxicology (2020) 94:1879–1897. A Small Dose of Toxicology Arsenic in groundwater Book on arsenic in groundwater by IAH's Netherlands Chapter and the Netherlands Hydrological Society Contaminant Focus: Arsenic by the EPA. Environmental Health Criteria for Arsenic and Arsenic Compounds, 2001 by the WHO. National Institute for Occupational Safety and Health – Arsenic Page Arsenic at The Periodic Table of Videos (University of Nottingham) Chemical elements Metalloids Hepatotoxins Pnictogens Biology and pharmacology of chemical elements Endocrine disruptors IARC Group 1 carcinogens Trigonal minerals Minerals in space group 166 Teratogens Fetotoxicants Suspected testicular toxicants Native element minerals Chemical elements with rhombohedral structure

Antimony Antimony is a chemical element with the symbol Sb (from ) and atomic number 51. A lustrous gray metalloid, it is found in nature mainly as the sulfide mineral stibnite (Sb2S3). Antimony compounds have been known since ancient times and were powdered for use as medicine and cosmetics, often known by the Arabic name kohl. The earliest known description of the metal in the West was written in 1540 by Vannoccio Biringuccio. China is the largest producer of antimony and its compounds, with most production coming from the Xikuangshan Mine in Hunan. The industrial methods for refining antimony from stibnite are roasting followed by reduction with carbon, or direct reduction of stibnite with iron. The largest applications for metallic antimony are in alloys with lead and tin, which have improved properties for solders, bullets, and plain bearings. It improves the rigidity of lead-alloy plates in lead–acid batteries. Antimony trioxide is a prominent additive for halogen-containing flame retardants. Antimony is used as a dopant in semiconductor devices. Characteristics Properties Antimony is a member of group 15 of the periodic table, one of the elements called pnictogens, and has an electronegativity of 2.05. In accordance with periodic trends, it is more electronegative than tin or bismuth, and less electronegative than tellurium or arsenic. Antimony is stable in air at room temperature, but reacts with oxygen if heated to produce antimony trioxide, Sb2O3. Antimony is a silvery, lustrous gray metalloid with a Mohs scale hardness of 3, which is too soft to make hard objects. Coins of antimony were issued in China's Guizhou province in 1931; durability was poor, and minting was soon discontinued. Antimony is resistant to attack by acids. Four allotropes of antimony are known: a stable metallic form, and three metastable forms (explosive, black, and yellow). Elemental antimony is a brittle, silver-white, shiny metalloid. When slowly cooled, molten antimony crystallizes into a trigonal cell, isomorphic with the gray allotrope of arsenic. A rare explosive form of antimony can be formed from the electrolysis of antimony trichloride. When scratched with a sharp implement, an exothermic reaction occurs and white fumes are given off as metallic antimony forms; when rubbed with a pestle in a mortar, a strong detonation occurs. Black antimony is formed upon rapid cooling of antimony vapor. It has the same crystal structure as red phosphorus and black arsenic; it oxidizes in air and may ignite spontaneously. At 100 °C, it gradually transforms into the stable form. The yellow allotrope of antimony is the most unstable; it has been generated only by oxidation of stibine (SbH3) at −90 °C. Above this temperature and in ambient light, this metastable allotrope transforms into the more stable black allotrope. Elemental antimony adopts a layered structure (space group Rm No. 166) whose layers consist of fused, ruffled, six-membered rings. The nearest and next-nearest neighbors form an irregular octahedral complex, with the three atoms in each double layer slightly closer than the three atoms in the next.

Antimony Antimony is a chemical element with the symbol Sb (from ) and atomic number 51. A lustrous gray metalloid, it is found in nature mainly as the sulfide mineral stibnite (Sb2S3). Antimony compounds have been known since ancient times and were powdered for use as medicine and cosmetics, often known by the Arabic name kohl. The earliest known description of the metal in the West was written in 1540 by Vannoccio Biringuccio. China is the largest producer of antimony and its compounds, with most production coming from the Xikuangshan Mine in Hunan. The industrial methods for refining antimony from stibnite are roasting followed by reduction with carbon, or direct reduction of stibnite with iron. The largest applications for metallic antimony are in alloys with lead and tin, which have improved properties for solders, bullets, and plain bearings. It improves the rigidity of lead-alloy plates in lead–acid batteries. Antimony trioxide is a prominent additive for halogen-containing flame retardants. Antimony is used as a dopant in semiconductor devices. Characteristics Properties Antimony is a member of group 15 of the periodic table, one of the elements called pnictogens, and has an electronegativity of 2.05. In accordance with periodic trends, it is more electronegative than tin or bismuth, and less electronegative than tellurium or arsenic. Antimony is stable in air at room temperature, but reacts with oxygen if heated to produce antimony trioxide, Sb2O3. Antimony is a silvery, lustrous gray metalloid with a Mohs scale hardness of 3, which is too soft to make hard objects. Coins of antimony were issued in China's Guizhou province in 1931; durability was poor, and minting was soon discontinued. Antimony is resistant to attack by acids. Four allotropes of antimony are known: a stable metallic form, and three metastable forms (explosive, black, and yellow). Elemental antimony is a brittle, silver-white, shiny metalloid. When slowly cooled, molten antimony crystallizes into a trigonal cell, isomorphic with the gray allotrope of arsenic. A rare explosive form of antimony can be formed from the electrolysis of antimony trichloride. When scratched with a sharp implement, an exothermic reaction occurs and white fumes are given off as metallic antimony forms; when rubbed with a pestle in a mortar, a strong detonation occurs. Black antimony is formed upon rapid cooling of antimony vapor. It has the same crystal structure as red phosphorus and black arsenic; it oxidizes in air and may ignite spontaneously. At 100 °C, it gradually transforms into the stable form. The yellow allotrope of antimony is the most unstable; it has been generated only by oxidation of stibine (SbH3) at −90 °C. Above this temperature and in ambient light, this metastable allotrope transforms into the more stable black allotrope. Elemental antimony adopts a layered structure (space group Rm No. 166) whose layers consist of fused, ruffled, six-membered rings. The nearest and next-nearest neighbors form an irregular octahedral complex, with the three atoms in each double layer slightly closer than the three atoms in the next.

This relatively close packing leads to a high density of 6.697 g/cm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony. Isotopes Antimony has two stable isotopes: 121Sb with a natural abundance of 57.36% and 123Sb with a natural abundance of 42.64%. It also has 35 radioisotopes, of which the longest-lived is 125Sb with a half-life of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is 120m1Sb with a half-life of 5.76 days. Isotopes that are lighter than the stable 123Sb tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. Occurrence The abundance of antimony in the Earth's crust is estimated to be 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in more than 100 mineral species. Antimony is sometimes found natively (e.g. on Antimony Peak), but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Compounds Antimony compounds are often classified according to their oxidation state: Sb(III) and Sb(V). The +5 oxidation state is more stable. Oxides and hydroxides Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide () can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (), which features both Sb(III) and Sb(V). Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form well-defined oxoacids, and react with acids to form antimony salts. Antimonous acid is unknown, but the conjugate base sodium antimonite () forms upon fusing sodium oxide and . Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing this anion is dehydrated, the precipitate contains mixed oxides. Many antimony ores are sulfides, including stibnite (), pyrargyrite (), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is non-stoichiometric and features antimony in the +3 oxidation state and S–S bonds. Several thioantimonides are known, such as and . Halides Antimony forms two series of halides: and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry. The trifluoride is prepared by the reaction of with HF: + 6 HF → 2 + 3 It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten is a weak electrical conductor. The trichloride is prepared by dissolving in hydrochloric acid: + 6 HCl → 2 + 3 The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric. is a powerful Lewis acid used to make the superacid fluoroantimonic acid ("H2SbF7").

This relatively close packing leads to a high density of 6.697 g/cm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony. Isotopes Antimony has two stable isotopes: 121Sb with a natural abundance of 57.36% and 123Sb with a natural abundance of 42.64%. It also has 35 radioisotopes, of which the longest-lived is 125Sb with a half-life of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is 120m1Sb with a half-life of 5.76 days. Isotopes that are lighter than the stable 123Sb tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. Occurrence The abundance of antimony in the Earth's crust is estimated to be 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in more than 100 mineral species. Antimony is sometimes found natively (e.g. on Antimony Peak), but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Compounds Antimony compounds are often classified according to their oxidation state: Sb(III) and Sb(V). The +5 oxidation state is more stable. Oxides and hydroxides Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide () can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (), which features both Sb(III) and Sb(V). Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form well-defined oxoacids, and react with acids to form antimony salts. Antimonous acid is unknown, but the conjugate base sodium antimonite () forms upon fusing sodium oxide and . Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing this anion is dehydrated, the precipitate contains mixed oxides. Many antimony ores are sulfides, including stibnite (), pyrargyrite (), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is non-stoichiometric and features antimony in the +3 oxidation state and S–S bonds. Several thioantimonides are known, such as and . Halides Antimony forms two series of halides: and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry. The trifluoride is prepared by the reaction of with HF: + 6 HF → 2 + 3 It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten is a weak electrical conductor. The trichloride is prepared by dissolving in hydrochloric acid: + 6 HCl → 2 + 3 The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric. is a powerful Lewis acid used to make the superacid fluoroantimonic acid ("H2SbF7").

This relatively close packing leads to a high density of 6.697 g/cm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony. Isotopes Antimony has two stable isotopes: 121Sb with a natural abundance of 57.36% and 123Sb with a natural abundance of 42.64%. It also has 35 radioisotopes, of which the longest-lived is 125Sb with a half-life of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is 120m1Sb with a half-life of 5.76 days. Isotopes that are lighter than the stable 123Sb tend to decay by β+ decay, and those that are heavier tend to decay by β− decay, with some exceptions. Occurrence The abundance of antimony in the Earth's crust is estimated to be 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in more than 100 mineral species. Antimony is sometimes found natively (e.g. on Antimony Peak), but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Compounds Antimony compounds are often classified according to their oxidation state: Sb(III) and Sb(V). The +5 oxidation state is more stable. Oxides and hydroxides Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide () can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (), which features both Sb(III) and Sb(V). Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form well-defined oxoacids, and react with acids to form antimony salts. Antimonous acid is unknown, but the conjugate base sodium antimonite () forms upon fusing sodium oxide and . Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing this anion is dehydrated, the precipitate contains mixed oxides. Many antimony ores are sulfides, including stibnite (), pyrargyrite (), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is non-stoichiometric and features antimony in the +3 oxidation state and S–S bonds. Several thioantimonides are known, such as and . Halides Antimony forms two series of halides: and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry. The trifluoride is prepared by the reaction of with HF: + 6 HF → 2 + 3 It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten is a weak electrical conductor. The trichloride is prepared by dissolving in hydrochloric acid: + 6 HCl → 2 + 3 The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric. is a powerful Lewis acid used to make the superacid fluoroantimonic acid ("H2SbF7").

Oxyhalides are more common for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and . Antimonides, hydrides, and organoantimony compounds Compounds in this class generally are described as derivatives of Sb3−. Antimony forms antimonides with metals, such as indium antimonide (InSb) and silver antimonide (). The alkali metal and zinc antimonides, such as Na3Sb and Zn3Sb2, are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine, : + 3 → Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly. Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents. A large variety of compounds are known with both Sb(III) and Sb(V) centers, including mixed chloro-organic derivatives, anions, and cations. Examples include Sb(C6H5)3 (triphenylstibine), Sb2(C6H5)4 (with an Sb-Sb bond), and cyclic [Sb(C6H5)]n. Pentacoordinated organoantimony compounds are common, examples being Sb(C6H5)5 and several related halides. History Antimony(III) sulfide, Sb2S3, was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented. An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' (artifact mentioned above) must represent the lost art of rendering antimony malleable." The British archaeologist Roger Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), "attempted to relate the metal to Transcaucasian natural antimony" (i.e. native metal) and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable." The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony. The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony. The intentional isolation of antimony is described by Jabir ibn Hayyan before 815 AD.

Oxyhalides are more common for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and . Antimonides, hydrides, and organoantimony compounds Compounds in this class generally are described as derivatives of Sb3−. Antimony forms antimonides with metals, such as indium antimonide (InSb) and silver antimonide (). The alkali metal and zinc antimonides, such as Na3Sb and Zn3Sb2, are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine, : + 3 → Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly. Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents. A large variety of compounds are known with both Sb(III) and Sb(V) centers, including mixed chloro-organic derivatives, anions, and cations. Examples include Sb(C6H5)3 (triphenylstibine), Sb2(C6H5)4 (with an Sb-Sb bond), and cyclic [Sb(C6H5)]n. Pentacoordinated organoantimony compounds are common, examples being Sb(C6H5)5 and several related halides. History Antimony(III) sulfide, Sb2S3, was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented. An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' (artifact mentioned above) must represent the lost art of rendering antimony malleable." The British archaeologist Roger Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), "attempted to relate the metal to Transcaucasian natural antimony" (i.e. native metal) and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable." The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony. The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony. The intentional isolation of antimony is described by Jabir ibn Hayyan before 815 AD.

Oxyhalides are more common for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and . Antimonides, hydrides, and organoantimony compounds Compounds in this class generally are described as derivatives of Sb3−. Antimony forms antimonides with metals, such as indium antimonide (InSb) and silver antimonide (). The alkali metal and zinc antimonides, such as Na3Sb and Zn3Sb2, are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine, : + 3 → Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly. Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents. A large variety of compounds are known with both Sb(III) and Sb(V) centers, including mixed chloro-organic derivatives, anions, and cations. Examples include Sb(C6H5)3 (triphenylstibine), Sb2(C6H5)4 (with an Sb-Sb bond), and cyclic [Sb(C6H5)]n. Pentacoordinated organoantimony compounds are common, examples being Sb(C6H5)5 and several related halides. History Antimony(III) sulfide, Sb2S3, was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented. An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' (artifact mentioned above) must represent the lost art of rendering antimony malleable." The British archaeologist Roger Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), "attempted to relate the metal to Transcaucasian natural antimony" (i.e. native metal) and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable." The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony. The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony. The intentional isolation of antimony is described by Jabir ibn Hayyan before 815 AD.

A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio, predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio. The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface. With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides, and other compounds, as do other metals. The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden. Etymology The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French antimoine, still has adherents; this would mean "monk-killer", and is explained by many early alchemists being monks, and antimony being poisonous. However, the low toxicity of antimony (see below) makes this unlikely. Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed". Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean "floret", and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence. The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek. The standard chemical symbol for antimony (Sb) is credited to Jöns Jakob Berzelius, who derived the abbreviation from stibium. The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony. The Egyptians called antimony mśdmt; in hieroglyphs, the vowels are uncertain, but the Coptic form of the word is ⲥⲧⲏⲙ (stēm). Egyptian stm: O34:D46-G17-F21:D4 The Greek word, στίμμι (stimmi) is used by Attic tragic poets of the 5th century BC, and is possibly a loan word from Arabic or from Egyptian stm.

A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio, predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio. The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface. With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides, and other compounds, as do other metals. The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden. Etymology The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French antimoine, still has adherents; this would mean "monk-killer", and is explained by many early alchemists being monks, and antimony being poisonous. However, the low toxicity of antimony (see below) makes this unlikely. Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed". Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean "floret", and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence. The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek. The standard chemical symbol for antimony (Sb) is credited to Jöns Jakob Berzelius, who derived the abbreviation from stibium. The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony. The Egyptians called antimony mśdmt; in hieroglyphs, the vowels are uncertain, but the Coptic form of the word is ⲥⲧⲏⲙ (stēm). Egyptian stm: O34:D46-G17-F21:D4 The Greek word, στίμμι (stimmi) is used by Attic tragic poets of the 5th century BC, and is possibly a loan word from Arabic or from Egyptian stm.

A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio, predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio. The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface. With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides, and other compounds, as do other metals. The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden. Etymology The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French antimoine, still has adherents; this would mean "monk-killer", and is explained by many early alchemists being monks, and antimony being poisonous. However, the low toxicity of antimony (see below) makes this unlikely. Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed". Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean "floret", and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence. The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek. The standard chemical symbol for antimony (Sb) is credited to Jöns Jakob Berzelius, who derived the abbreviation from stibium. The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony. The Egyptians called antimony mśdmt; in hieroglyphs, the vowels are uncertain, but the Coptic form of the word is ⲥⲧⲏⲙ (stēm). Egyptian stm: O34:D46-G17-F21:D4 The Greek word, στίμμι (stimmi) is used by Attic tragic poets of the 5th century BC, and is possibly a loan word from Arabic or from Egyptian stm.

Later Greeks also used στἰβι stibi, as did Celsus and Pliny, writing in Latin, in the first century AD. Pliny also gives the names stimi, larbaris, alabaster, and the "very common" platyophthalmos, "wide-eye" (from the effect of the cosmetic). Later Latin authors adapted the word to Latin as stibium. The Arabic word for the substance, as opposed to the cosmetic, can appear as إثمد ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi. Production Process The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lower-grade ores are concentrated by froth flotation, while higher-grade ores are heated to 500–600 °C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron: + 3 Fe → 2 Sb + 3 FeS The sulfide is converted to an oxide; the product is then roasted, sometimes for the purpose of vaporizing the volatile antimony(III) oxide, which is recovered. This material is often used directly for the main applications, impurities being arsenic and sulfide. Antimony is isolated from the oxide by a carbothermal reduction: 2 + 3 C → 4 Sb + 3 The lower-grade ores are reduced in blast furnaces while the higher-grade ores are reduced in reverberatory furnaces. Top producers and production volumes The British Geological Survey (BGS) reported that in 2005 China was the top producer of antimony with approximately 84% of the world share, followed at a distance by South Africa, Bolivia and Tajikistan. Xikuangshan Mine in Hunan province has the largest deposits in China with an estimated deposit of 2.1 million metric tons. In 2016, according to the US Geological Survey, China accounted for 76.9% of total antimony production, followed in second place by Russia with 6.9% and Tajikistan with 6.2%. Chinese production of antimony is expected to decline in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher. According to the National Bureau of Statistics in China, by September 2015 50% of antimony production capacity in the Hunan province (the province with biggest antimony reserves in China) had not been used. Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted.

Later Greeks also used στἰβι stibi, as did Celsus and Pliny, writing in Latin, in the first century AD. Pliny also gives the names stimi, larbaris, alabaster, and the "very common" platyophthalmos, "wide-eye" (from the effect of the cosmetic). Later Latin authors adapted the word to Latin as stibium. The Arabic word for the substance, as opposed to the cosmetic, can appear as إثمد ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi. Production Process The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lower-grade ores are concentrated by froth flotation, while higher-grade ores are heated to 500–600 °C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron: + 3 Fe → 2 Sb + 3 FeS The sulfide is converted to an oxide; the product is then roasted, sometimes for the purpose of vaporizing the volatile antimony(III) oxide, which is recovered. This material is often used directly for the main applications, impurities being arsenic and sulfide. Antimony is isolated from the oxide by a carbothermal reduction: 2 + 3 C → 4 Sb + 3 The lower-grade ores are reduced in blast furnaces while the higher-grade ores are reduced in reverberatory furnaces. Top producers and production volumes The British Geological Survey (BGS) reported that in 2005 China was the top producer of antimony with approximately 84% of the world share, followed at a distance by South Africa, Bolivia and Tajikistan. Xikuangshan Mine in Hunan province has the largest deposits in China with an estimated deposit of 2.1 million metric tons. In 2016, according to the US Geological Survey, China accounted for 76.9% of total antimony production, followed in second place by Russia with 6.9% and Tajikistan with 6.2%. Chinese production of antimony is expected to decline in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher. According to the National Bureau of Statistics in China, by September 2015 50% of antimony production capacity in the Hunan province (the province with biggest antimony reserves in China) had not been used. Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted.

Later Greeks also used στἰβι stibi, as did Celsus and Pliny, writing in Latin, in the first century AD. Pliny also gives the names stimi, larbaris, alabaster, and the "very common" platyophthalmos, "wide-eye" (from the effect of the cosmetic). Later Latin authors adapted the word to Latin as stibium. The Arabic word for the substance, as opposed to the cosmetic, can appear as إثمد ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi. Production Process The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lower-grade ores are concentrated by froth flotation, while higher-grade ores are heated to 500–600 °C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron: + 3 Fe → 2 Sb + 3 FeS The sulfide is converted to an oxide; the product is then roasted, sometimes for the purpose of vaporizing the volatile antimony(III) oxide, which is recovered. This material is often used directly for the main applications, impurities being arsenic and sulfide. Antimony is isolated from the oxide by a carbothermal reduction: 2 + 3 C → 4 Sb + 3 The lower-grade ores are reduced in blast furnaces while the higher-grade ores are reduced in reverberatory furnaces. Top producers and production volumes The British Geological Survey (BGS) reported that in 2005 China was the top producer of antimony with approximately 84% of the world share, followed at a distance by South Africa, Bolivia and Tajikistan. Xikuangshan Mine in Hunan province has the largest deposits in China with an estimated deposit of 2.1 million metric tons. In 2016, according to the US Geological Survey, China accounted for 76.9% of total antimony production, followed in second place by Russia with 6.9% and Tajikistan with 6.2%. Chinese production of antimony is expected to decline in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher. According to the National Bureau of Statistics in China, by September 2015 50% of antimony production capacity in the Hunan province (the province with biggest antimony reserves in China) had not been used. Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted.

The world's largest antimony producers, according to Roskill, are listed below: Reserves Supply risk For antimony-importing regions such as Europe and the U.S., antimony is considered to be a critical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China (74%), Tajikistan(8%), and Russia(4%), these sources are critical to supply. European Union: Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100% imported, coming mainly from Turkey (62%), Bolivia (20%) and Guatemala (7%). United Kingdom: The British Geological Survey's 2015 risk list ranks antimony second highest (after rare earth elements) on the relative supply risk index. United States: Antimony is a mineral commodity considered critical to the economic and national security. In 2021, no antimony was mined in the U.S. Applications About 60% of antimony is consumed in flame retardants, and 20% is used in alloys for batteries, plain bearings, and solders. Flame retardants Antimony is mainly used as the trioxide for flame-proofing compounds, always in combination with halogenated flame retardants except in halogen-containing polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flame-retardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed. Alloys Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. For most applications involving lead, varying amounts of antimony are used as alloying metal. In lead–acid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardness and tensile strength of the lead keel, antimony is mixed with lead between 2% and 5% by volume. Antimony is used in antifriction alloys (such as Babbitt metal), in bullets and lead shot, electrical cable sheathing, type metal (for example, for linotype printing machines), solder (some "lead-free" solders contain 5% Sb), in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes. Other applications Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens - antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments. In 1990s antimony was increasingly being used in semiconductors as a dopant in n-type silicon wafers for diodes, infrared detectors, and Hall-effect devices.

The world's largest antimony producers, according to Roskill, are listed below: Reserves Supply risk For antimony-importing regions such as Europe and the U.S., antimony is considered to be a critical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China (74%), Tajikistan(8%), and Russia(4%), these sources are critical to supply. European Union: Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100% imported, coming mainly from Turkey (62%), Bolivia (20%) and Guatemala (7%). United Kingdom: The British Geological Survey's 2015 risk list ranks antimony second highest (after rare earth elements) on the relative supply risk index. United States: Antimony is a mineral commodity considered critical to the economic and national security. In 2021, no antimony was mined in the U.S. Applications About 60% of antimony is consumed in flame retardants, and 20% is used in alloys for batteries, plain bearings, and solders. Flame retardants Antimony is mainly used as the trioxide for flame-proofing compounds, always in combination with halogenated flame retardants except in halogen-containing polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flame-retardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed. Alloys Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. For most applications involving lead, varying amounts of antimony are used as alloying metal. In lead–acid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardness and tensile strength of the lead keel, antimony is mixed with lead between 2% and 5% by volume. Antimony is used in antifriction alloys (such as Babbitt metal), in bullets and lead shot, electrical cable sheathing, type metal (for example, for linotype printing machines), solder (some "lead-free" solders contain 5% Sb), in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes. Other applications Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens - antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments. In 1990s antimony was increasingly being used in semiconductors as a dopant in n-type silicon wafers for diodes, infrared detectors, and Hall-effect devices.

The world's largest antimony producers, according to Roskill, are listed below: Reserves Supply risk For antimony-importing regions such as Europe and the U.S., antimony is considered to be a critical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China (74%), Tajikistan(8%), and Russia(4%), these sources are critical to supply. European Union: Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100% imported, coming mainly from Turkey (62%), Bolivia (20%) and Guatemala (7%). United Kingdom: The British Geological Survey's 2015 risk list ranks antimony second highest (after rare earth elements) on the relative supply risk index. United States: Antimony is a mineral commodity considered critical to the economic and national security. In 2021, no antimony was mined in the U.S. Applications About 60% of antimony is consumed in flame retardants, and 20% is used in alloys for batteries, plain bearings, and solders. Flame retardants Antimony is mainly used as the trioxide for flame-proofing compounds, always in combination with halogenated flame retardants except in halogen-containing polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flame-retardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed. Alloys Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. For most applications involving lead, varying amounts of antimony are used as alloying metal. In lead–acid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardness and tensile strength of the lead keel, antimony is mixed with lead between 2% and 5% by volume. Antimony is used in antifriction alloys (such as Babbitt metal), in bullets and lead shot, electrical cable sheathing, type metal (for example, for linotype printing machines), solder (some "lead-free" solders contain 5% Sb), in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes. Other applications Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens - antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments. In 1990s antimony was increasingly being used in semiconductors as a dopant in n-type silicon wafers for diodes, infrared detectors, and Hall-effect devices.

In the 1950s, the emitters and collectors of n-p-n alloy junction transistors were doped with tiny beads of a lead-antimony alloy. Indium antimonide is used as a material for mid-infrared detectors. Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an anti-schistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals. Antimony-based drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Besides having low therapeutic indices, the drugs have minimal penetration of the bone marrow, where some of the Leishmania amastigotes reside, and curing the disease – especially the visceral form – is very difficult. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination. Antimony(III) sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony-124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony-124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources. Historically, the powder derived from crushed antimony (kohl) has been applied to the eyes with a metal rod and with one's spittle, thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries. Precautions The effects of antimony and its compounds on human and environmental health differ widely. Elemental antimony metal does not affect human and environmental health. Inhalation of antimony trioxide (and similar poorly soluble Sb(III) dust particles such as antimony dust) is considered harmful and suspected of causing cancer. However, these effects are only observed with female rats and after long-term exposure to high dust concentrations. The effects are hypothesized to be attributed to inhalation of poorly soluble Sb particles leading to impaired lung clearance, lung overload, inflammation and ultimately tumour formation, not to exposure to antimony ions (OECD, 2008). Antimony chlorides are corrosive to skin. The effects of antimony are not comparable to those of arsenic; this might be caused by the significant differences of uptake, metabolism, and excretion between arsenic and antimony. For oral absorption, ICRP (1994) has recommended values of 10% for tartar emetic and 1% for all other antimony compounds. Dermal absorption for metals is estimated to be at most 1% (HERAG, 2007).

In the 1950s, the emitters and collectors of n-p-n alloy junction transistors were doped with tiny beads of a lead-antimony alloy. Indium antimonide is used as a material for mid-infrared detectors. Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an anti-schistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals. Antimony-based drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Besides having low therapeutic indices, the drugs have minimal penetration of the bone marrow, where some of the Leishmania amastigotes reside, and curing the disease – especially the visceral form – is very difficult. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination. Antimony(III) sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony-124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony-124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources. Historically, the powder derived from crushed antimony (kohl) has been applied to the eyes with a metal rod and with one's spittle, thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries. Precautions The effects of antimony and its compounds on human and environmental health differ widely. Elemental antimony metal does not affect human and environmental health. Inhalation of antimony trioxide (and similar poorly soluble Sb(III) dust particles such as antimony dust) is considered harmful and suspected of causing cancer. However, these effects are only observed with female rats and after long-term exposure to high dust concentrations. The effects are hypothesized to be attributed to inhalation of poorly soluble Sb particles leading to impaired lung clearance, lung overload, inflammation and ultimately tumour formation, not to exposure to antimony ions (OECD, 2008). Antimony chlorides are corrosive to skin. The effects of antimony are not comparable to those of arsenic; this might be caused by the significant differences of uptake, metabolism, and excretion between arsenic and antimony. For oral absorption, ICRP (1994) has recommended values of 10% for tartar emetic and 1% for all other antimony compounds. Dermal absorption for metals is estimated to be at most 1% (HERAG, 2007).

In the 1950s, the emitters and collectors of n-p-n alloy junction transistors were doped with tiny beads of a lead-antimony alloy. Indium antimonide is used as a material for mid-infrared detectors. Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an anti-schistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals. Antimony-based drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Besides having low therapeutic indices, the drugs have minimal penetration of the bone marrow, where some of the Leishmania amastigotes reside, and curing the disease – especially the visceral form – is very difficult. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination. Antimony(III) sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony-124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony-124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources. Historically, the powder derived from crushed antimony (kohl) has been applied to the eyes with a metal rod and with one's spittle, thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries. Precautions The effects of antimony and its compounds on human and environmental health differ widely. Elemental antimony metal does not affect human and environmental health. Inhalation of antimony trioxide (and similar poorly soluble Sb(III) dust particles such as antimony dust) is considered harmful and suspected of causing cancer. However, these effects are only observed with female rats and after long-term exposure to high dust concentrations. The effects are hypothesized to be attributed to inhalation of poorly soluble Sb particles leading to impaired lung clearance, lung overload, inflammation and ultimately tumour formation, not to exposure to antimony ions (OECD, 2008). Antimony chlorides are corrosive to skin. The effects of antimony are not comparable to those of arsenic; this might be caused by the significant differences of uptake, metabolism, and excretion between arsenic and antimony. For oral absorption, ICRP (1994) has recommended values of 10% for tartar emetic and 1% for all other antimony compounds. Dermal absorption for metals is estimated to be at most 1% (HERAG, 2007).

Inhalation absorption of antimony trioxide and other poorly soluble Sb(III) substances (such as antimony dust) is estimated at 6.8% (OECD, 2008), whereas a value <1% is derived for Sb(V) substances. Antimony(V) is not quantitatively reduced to antimony(III) in the cell, and both species exist simultaneously. Antimony is mainly excreted from the human body via urine. Antimony and its compounds do not cause acute human health effects, with the exception of antimony potassium tartrate ("tartar emetic"), a prodrug that is intentionally used to treat leishmaniasis patients. Prolonged skin contact with antimony dust may cause dermatitis. However, it was agreed at the European Union level that the skin rashes observed are not substance-specific, but most probably due to a physical blocking of sweat ducts (ECHA/PR/09/09, Helsinki, 6 July 2009). Antimony dust may also be explosive when dispersed in the air; when in a bulk solid it is not combustible. Antimony is incompatible with strong acids, halogenated acids, and oxidizers; when exposed to newly formed hydrogen it may form stibine (SbH3). The 8-hour time-weighted average (TWA) is set at 0.5 mg/m3 by the American Conference of Governmental Industrial Hygienists and by the Occupational Safety and Health Administration (OSHA) as a legal permissible exposure limit (PEL) in the workplace. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of 0.5 mg/m3 as an 8-hour TWA. Antimony compounds are used as catalysts for polyethylene terephthalate (PET) production. Some studies report minor antimony leaching from PET bottles into liquids, but levels are below drinking water guidelines. Antimony concentrations in fruit juice concentrates were somewhat higher (up to 44.7 µg/L of antimony), but juices do not fall under the drinking water regulations. The drinking water guidelines are: World Health Organization: 20 µg/L Japan: 15 µg/L United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 µg/L EU and German Federal Ministry of Environment: 5 µg/L The tolerable daily intake (TDI) proposed by WHO is 6 µg antimony per kilogram of body weight. The immediately dangerous to life or health (IDLH) value for antimony is 50 mg/m3. Toxicity Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans. Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay. See also Phase change memory Notes References Bibliography Edmund Oscar von Lippmann (1919) Entstehung und Ausbreitung der Alchemie, teil 1. Berlin: Julius Springer (in German).

Inhalation absorption of antimony trioxide and other poorly soluble Sb(III) substances (such as antimony dust) is estimated at 6.8% (OECD, 2008), whereas a value <1% is derived for Sb(V) substances. Antimony(V) is not quantitatively reduced to antimony(III) in the cell, and both species exist simultaneously. Antimony is mainly excreted from the human body via urine. Antimony and its compounds do not cause acute human health effects, with the exception of antimony potassium tartrate ("tartar emetic"), a prodrug that is intentionally used to treat leishmaniasis patients. Prolonged skin contact with antimony dust may cause dermatitis. However, it was agreed at the European Union level that the skin rashes observed are not substance-specific, but most probably due to a physical blocking of sweat ducts (ECHA/PR/09/09, Helsinki, 6 July 2009). Antimony dust may also be explosive when dispersed in the air; when in a bulk solid it is not combustible. Antimony is incompatible with strong acids, halogenated acids, and oxidizers; when exposed to newly formed hydrogen it may form stibine (SbH3). The 8-hour time-weighted average (TWA) is set at 0.5 mg/m3 by the American Conference of Governmental Industrial Hygienists and by the Occupational Safety and Health Administration (OSHA) as a legal permissible exposure limit (PEL) in the workplace. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of 0.5 mg/m3 as an 8-hour TWA. Antimony compounds are used as catalysts for polyethylene terephthalate (PET) production. Some studies report minor antimony leaching from PET bottles into liquids, but levels are below drinking water guidelines. Antimony concentrations in fruit juice concentrates were somewhat higher (up to 44.7 µg/L of antimony), but juices do not fall under the drinking water regulations. The drinking water guidelines are: World Health Organization: 20 µg/L Japan: 15 µg/L United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 µg/L EU and German Federal Ministry of Environment: 5 µg/L The tolerable daily intake (TDI) proposed by WHO is 6 µg antimony per kilogram of body weight. The immediately dangerous to life or health (IDLH) value for antimony is 50 mg/m3. Toxicity Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans. Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay. See also Phase change memory Notes References Bibliography Edmund Oscar von Lippmann (1919) Entstehung und Ausbreitung der Alchemie, teil 1. Berlin: Julius Springer (in German).

Inhalation absorption of antimony trioxide and other poorly soluble Sb(III) substances (such as antimony dust) is estimated at 6.8% (OECD, 2008), whereas a value <1% is derived for Sb(V) substances. Antimony(V) is not quantitatively reduced to antimony(III) in the cell, and both species exist simultaneously. Antimony is mainly excreted from the human body via urine. Antimony and its compounds do not cause acute human health effects, with the exception of antimony potassium tartrate ("tartar emetic"), a prodrug that is intentionally used to treat leishmaniasis patients. Prolonged skin contact with antimony dust may cause dermatitis. However, it was agreed at the European Union level that the skin rashes observed are not substance-specific, but most probably due to a physical blocking of sweat ducts (ECHA/PR/09/09, Helsinki, 6 July 2009). Antimony dust may also be explosive when dispersed in the air; when in a bulk solid it is not combustible. Antimony is incompatible with strong acids, halogenated acids, and oxidizers; when exposed to newly formed hydrogen it may form stibine (SbH3). The 8-hour time-weighted average (TWA) is set at 0.5 mg/m3 by the American Conference of Governmental Industrial Hygienists and by the Occupational Safety and Health Administration (OSHA) as a legal permissible exposure limit (PEL) in the workplace. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of 0.5 mg/m3 as an 8-hour TWA. Antimony compounds are used as catalysts for polyethylene terephthalate (PET) production. Some studies report minor antimony leaching from PET bottles into liquids, but levels are below drinking water guidelines. Antimony concentrations in fruit juice concentrates were somewhat higher (up to 44.7 µg/L of antimony), but juices do not fall under the drinking water regulations. The drinking water guidelines are: World Health Organization: 20 µg/L Japan: 15 µg/L United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment: 6 µg/L EU and German Federal Ministry of Environment: 5 µg/L The tolerable daily intake (TDI) proposed by WHO is 6 µg antimony per kilogram of body weight. The immediately dangerous to life or health (IDLH) value for antimony is 50 mg/m3. Toxicity Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans. Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay. See also Phase change memory Notes References Bibliography Edmund Oscar von Lippmann (1919) Entstehung und Ausbreitung der Alchemie, teil 1. Berlin: Julius Springer (in German).

Public Health Statement for Antimony External links International Antimony Association vzw (i2a) Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Antimony Antimony at The Periodic Table of Videos (University of Nottingham) CDC – NIOSH Pocket Guide to Chemical Hazards – Antimony Antimony Mineral data and specimen images Chemical elements Metalloids Native element minerals Nuclear materials Pnictogens Trigonal minerals Minerals in space group 166 Materials that expand upon freezing Chemical elements with rhombohedral structure

Public Health Statement for Antimony External links International Antimony Association vzw (i2a) Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Antimony Antimony at The Periodic Table of Videos (University of Nottingham) CDC – NIOSH Pocket Guide to Chemical Hazards – Antimony Antimony Mineral data and specimen images Chemical elements Metalloids Native element minerals Nuclear materials Pnictogens Trigonal minerals Minerals in space group 166 Materials that expand upon freezing Chemical elements with rhombohedral structure

Public Health Statement for Antimony External links International Antimony Association vzw (i2a) Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Antimony Antimony at The Periodic Table of Videos (University of Nottingham) CDC – NIOSH Pocket Guide to Chemical Hazards – Antimony Antimony Mineral data and specimen images Chemical elements Metalloids Native element minerals Nuclear materials Pnictogens Trigonal minerals Minerals in space group 166 Materials that expand upon freezing Chemical elements with rhombohedral structure

Actinium Actinium is a chemical element with the symbol Ac and atomic number 89. It was first isolated by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance André-Louis Debierne found in 1899 and called actinium. Actinium gave the name to the actinide series, a group of 15 similar elements between actinium and lawrencium in the periodic table. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be isolated. A soft, silvery-white radioactive metal, actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that prevents further oxidation. As with most lanthanides and many actinides, actinium assumes oxidation state +3 in nearly all its chemical compounds. Actinium is found only in traces in uranium and thorium ores as the isotope 227Ac, which decays with a half-life of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 228Ac, which is beta active with a half-life of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium-227, and one tonne of thorium contains about 5 nanograms of actinium-228. The close similarity of physical and chemical properties of actinium and lanthanum makes separation of actinium from the ore impractical. Instead, the element is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. Owing to its scarcity, high price and radioactivity, actinium has no significant industrial use. Its current applications include a neutron source and an agent for radiation therapy. History André-Louis Debierne, a French chemist, announced the discovery of a new element in 1899. He separated it from pitchblende residues left by Marie and Pierre Curie after they had extracted radium. In 1899, Debierne described the substance as similar to titanium and (in 1900) as similar to thorium. Friedrich Oskar Giesel found in 1902 a substance similar to lanthanum and called it "emanium" in 1904. After a comparison of the substances' half-lives determined by Debierne, Harriet Brooks in 1904, and Otto Hahn and Otto Sackur in 1905, Debierne's chosen name for the new element was retained because it had seniority, despite the contradicting chemical properties he claimed for the element at different times. Articles published in the 1970s and later suggest that Debierne's results published in 1904 conflict with those reported in 1899 and 1900. Furthermore, the now-known chemistry of actinium precludes its presence as anything other than a minor constituent of Debierne's 1899 and 1900 results; in fact, the chemical properties he reported make it likely that he had, instead, accidentally identified protactinium, which would not be discovered for another fourteen years, only to have it disappear due to its hydrolysis and adsorption onto his laboratory equipment. This has led some authors to advocate that Giesel alone should be credited with the discovery. A less confrontational vision of scientific discovery is proposed by Adloff.

Actinium Actinium is a chemical element with the symbol Ac and atomic number 89. It was first isolated by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance André-Louis Debierne found in 1899 and called actinium. Actinium gave the name to the actinide series, a group of 15 similar elements between actinium and lawrencium in the periodic table. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be isolated. A soft, silvery-white radioactive metal, actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that prevents further oxidation. As with most lanthanides and many actinides, actinium assumes oxidation state +3 in nearly all its chemical compounds. Actinium is found only in traces in uranium and thorium ores as the isotope 227Ac, which decays with a half-life of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 228Ac, which is beta active with a half-life of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium-227, and one tonne of thorium contains about 5 nanograms of actinium-228. The close similarity of physical and chemical properties of actinium and lanthanum makes separation of actinium from the ore impractical. Instead, the element is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. Owing to its scarcity, high price and radioactivity, actinium has no significant industrial use. Its current applications include a neutron source and an agent for radiation therapy. History André-Louis Debierne, a French chemist, announced the discovery of a new element in 1899. He separated it from pitchblende residues left by Marie and Pierre Curie after they had extracted radium. In 1899, Debierne described the substance as similar to titanium and (in 1900) as similar to thorium. Friedrich Oskar Giesel found in 1902 a substance similar to lanthanum and called it "emanium" in 1904. After a comparison of the substances' half-lives determined by Debierne, Harriet Brooks in 1904, and Otto Hahn and Otto Sackur in 1905, Debierne's chosen name for the new element was retained because it had seniority, despite the contradicting chemical properties he claimed for the element at different times. Articles published in the 1970s and later suggest that Debierne's results published in 1904 conflict with those reported in 1899 and 1900. Furthermore, the now-known chemistry of actinium precludes its presence as anything other than a minor constituent of Debierne's 1899 and 1900 results; in fact, the chemical properties he reported make it likely that he had, instead, accidentally identified protactinium, which would not be discovered for another fourteen years, only to have it disappear due to its hydrolysis and adsorption onto his laboratory equipment. This has led some authors to advocate that Giesel alone should be credited with the discovery. A less confrontational vision of scientific discovery is proposed by Adloff.

He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry: highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89. The name actinium originates from the Ancient Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde. Properties Actinium is a soft, silvery-white, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation. The first element of the actinides, actinium gave the group its name, much as lanthanum had done for the lanthanides. The group of elements is more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements (although it had been proposed as early as 1892 by British chemist Henry Bassett). Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state +3, and the Ac3+ ions are colorless in solutions. The oxidation state +3 originates from the [Rn]6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. The rare oxidation state +2 is only known for actinium dihydride (AcH2); even this may in reality be an electride compound like its lighter congener LaH2 and thus have actinium(III). Ac3+ is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 ± 0.5 water molecules. Chemical compounds Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include: AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and Ac(NO3)3. Except for AcPO4, they are all similar to the corresponding lanthanum compounds. They all contain actinium in the oxidation state +3. In particular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.

He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry: highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89. The name actinium originates from the Ancient Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde. Properties Actinium is a soft, silvery-white, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation. The first element of the actinides, actinium gave the group its name, much as lanthanum had done for the lanthanides. The group of elements is more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements (although it had been proposed as early as 1892 by British chemist Henry Bassett). Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state +3, and the Ac3+ ions are colorless in solutions. The oxidation state +3 originates from the [Rn]6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. The rare oxidation state +2 is only known for actinium dihydride (AcH2); even this may in reality be an electride compound like its lighter congener LaH2 and thus have actinium(III). Ac3+ is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 ± 0.5 water molecules. Chemical compounds Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include: AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and Ac(NO3)3. Except for AcPO4, they are all similar to the corresponding lanthanum compounds. They all contain actinium in the oxidation state +3. In particular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.

He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry: highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89. The name actinium originates from the Ancient Greek aktis, aktinos (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde. Properties Actinium is a soft, silvery-white, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation. The first element of the actinides, actinium gave the group its name, much as lanthanum had done for the lanthanides. The group of elements is more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements (although it had been proposed as early as 1892 by British chemist Henry Bassett). Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state +3, and the Ac3+ ions are colorless in solutions. The oxidation state +3 originates from the [Rn]6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. The rare oxidation state +2 is only known for actinium dihydride (AcH2); even this may in reality be an electride compound like its lighter congener LaH2 and thus have actinium(III). Ac3+ is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 ± 0.5 water molecules. Chemical compounds Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include: AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and Ac(NO3)3. Except for AcPO4, they are all similar to the corresponding lanthanum compounds. They all contain actinium in the oxidation state +3. In particular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.

Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters. Oxides Actinium oxide (Ac2O3) can be obtained by heating the hydroxide at 500 °C or the oxalate at 1100 °C, in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth metals. Halides Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at 700 °C in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at 900–1000 °C yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at 800 °C for an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product. AcF3 + 2 NH3 + H2O → AcOF + 2 NH4F Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above 960 °C. Similar to oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at 1000 °C. However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia. Reaction of aluminium bromide and actinium oxide yields actinium tribromide: Ac2O3 + 2 AlBr3 → 2 AcBr3 + Al2O3 and treating it with ammonium hydroxide at 500 °C results in the oxybromide AcOBr. Other compounds Actinium hydride was obtained by reduction of actinium trichloride with potassium at 300 °C, and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain. Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at 1400 °C for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at 1000 °C. Isotopes Naturally occurring actinium is composed of two radioactive isotopes; (from the radioactive family of ) and (a granddaughter of ). decays mainly as a beta emitter with a very small energy, but in 1.38% of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirty-six radioisotopes have been identified, the most stable being with a half-life of 21.772 years, with a half-life of 10.0 days and with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is (half-life of 69 nanoseconds) which decays through alpha decay. Actinium also has two known meta states.

Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters. Oxides Actinium oxide (Ac2O3) can be obtained by heating the hydroxide at 500 °C or the oxalate at 1100 °C, in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth metals. Halides Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at 700 °C in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at 900–1000 °C yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at 800 °C for an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product. AcF3 + 2 NH3 + H2O → AcOF + 2 NH4F Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above 960 °C. Similar to oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at 1000 °C. However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia. Reaction of aluminium bromide and actinium oxide yields actinium tribromide: Ac2O3 + 2 AlBr3 → 2 AcBr3 + Al2O3 and treating it with ammonium hydroxide at 500 °C results in the oxybromide AcOBr. Other compounds Actinium hydride was obtained by reduction of actinium trichloride with potassium at 300 °C, and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain. Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at 1400 °C for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at 1000 °C. Isotopes Naturally occurring actinium is composed of two radioactive isotopes; (from the radioactive family of ) and (a granddaughter of ). decays mainly as a beta emitter with a very small energy, but in 1.38% of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirty-six radioisotopes have been identified, the most stable being with a half-life of 21.772 years, with a half-life of 10.0 days and with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is (half-life of 69 nanoseconds) which decays through alpha decay. Actinium also has two known meta states.

Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters. Oxides Actinium oxide (Ac2O3) can be obtained by heating the hydroxide at 500 °C or the oxalate at 1100 °C, in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth metals. Halides Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at 700 °C in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at 900–1000 °C yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at 800 °C for an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product. AcF3 + 2 NH3 + H2O → AcOF + 2 NH4F Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above 960 °C. Similar to oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at 1000 °C. However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia. Reaction of aluminium bromide and actinium oxide yields actinium tribromide: Ac2O3 + 2 AlBr3 → 2 AcBr3 + Al2O3 and treating it with ammonium hydroxide at 500 °C results in the oxybromide AcOBr. Other compounds Actinium hydride was obtained by reduction of actinium trichloride with potassium at 300 °C, and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain. Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at 1400 °C for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at 1000 °C. Isotopes Naturally occurring actinium is composed of two radioactive isotopes; (from the radioactive family of ) and (a granddaughter of ). decays mainly as a beta emitter with a very small energy, but in 1.38% of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirty-six radioisotopes have been identified, the most stable being with a half-life of 21.772 years, with a half-life of 10.0 days and with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is (half-life of 69 nanoseconds) which decays through alpha decay. Actinium also has two known meta states.

The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac. Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life emitting mostly beta (98.62%) and some alpha particles (1.38%); the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles (maximum 44.8 keV) and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products. The isotopes of actinium range in atomic weight from 205 u () to 236 u (). Occurrence and synthesis Actinium is found only in traces in uranium ores – one tonne of uranium in ore contains about 0.2 milligrams of 227Ac – and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac is a transient member of the uranium-actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with thallium (205Tl) and near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out reactions on natural 238U. The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actinium-bearing ores, render separation of actinium from the ore impractical, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. ^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ \ce{min}] ^{227}_{89}Ac The reaction yield is about 2% of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetone-benzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH (to about 6.0 for actinium). An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking cation exchange resin and nitric acid as eluant. 225Ac was first produced artificially at the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and at St George Hospital in Sydney using a linac in 2000.

The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac. Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life emitting mostly beta (98.62%) and some alpha particles (1.38%); the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles (maximum 44.8 keV) and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products. The isotopes of actinium range in atomic weight from 205 u () to 236 u (). Occurrence and synthesis Actinium is found only in traces in uranium ores – one tonne of uranium in ore contains about 0.2 milligrams of 227Ac – and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac is a transient member of the uranium-actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with thallium (205Tl) and near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out reactions on natural 238U. The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actinium-bearing ores, render separation of actinium from the ore impractical, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. ^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ \ce{min}] ^{227}_{89}Ac The reaction yield is about 2% of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetone-benzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH (to about 6.0 for actinium). An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking cation exchange resin and nitric acid as eluant. 225Ac was first produced artificially at the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and at St George Hospital in Sydney using a linac in 2000.

The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac. Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life emitting mostly beta (98.62%) and some alpha particles (1.38%); the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles (maximum 44.8 keV) and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products. The isotopes of actinium range in atomic weight from 205 u () to 236 u (). Occurrence and synthesis Actinium is found only in traces in uranium ores – one tonne of uranium in ore contains about 0.2 milligrams of 227Ac – and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac is a transient member of the uranium-actinium series decay chain, which begins with the parent isotope 235U (or 239Pu) and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope (225Ac) is transiently present in the neptunium series decay chain, beginning with 237Np (or 233U) and ending with thallium (205Tl) and near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out reactions on natural 238U. The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actinium-bearing ores, render separation of actinium from the ore impractical, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. ^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ \ce{min}] ^{227}_{89}Ac The reaction yield is about 2% of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetone-benzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH (to about 6.0 for actinium). An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking cation exchange resin and nitric acid as eluant. 225Ac was first produced artificially at the Institute for Transuranium Elements (ITU) in Germany using a cyclotron and at St George Hospital in Sydney using a linac in 2000.

This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between 1100 and 1300 °C. Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile. Applications Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies. 227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs. In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The 227AcBe neutron sources can be applied in a neutron probe – a standard device for measuring the quantity of water present in soil, as well as moisture/density for quality control in highway construction. Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations. 225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy (TAT). This isotope has a half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than stable but toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid (EDTA) or diethylene triamine pentaacetic acid (DTPA). This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating agents as HEHA () or DOTA () coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu receptor.

This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between 1100 and 1300 °C. Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile. Applications Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies. 227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs. In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The 227AcBe neutron sources can be applied in a neutron probe – a standard device for measuring the quantity of water present in soil, as well as moisture/density for quality control in highway construction. Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations. 225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy (TAT). This isotope has a half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than stable but toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid (EDTA) or diethylene triamine pentaacetic acid (DTPA). This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating agents as HEHA () or DOTA () coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu receptor.

This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between 1100 and 1300 °C. Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile. Applications Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies. 227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs. In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The 227AcBe neutron sources can be applied in a neutron probe – a standard device for measuring the quantity of water present in soil, as well as moisture/density for quality control in highway construction. Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations. 225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy (TAT). This isotope has a half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than stable but toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid (EDTA) or diethylene triamine pentaacetic acid (DTPA). This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating agents as HEHA () or DOTA () coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu receptor.

The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers. The medium half-life of 227Ac (21.77 years) makes it very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities (of the order 50 meters per year). However, evaluation of the concentration depth-profiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows: oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa depth profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior. There are theoretical predictions that AcHx hydrides (in this case with very high pressure) are a candidate for a near room-temperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K. Precautions 227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33% of actinium is deposited into the bones and 50% into the liver. Its toxicity is comparable to, but slightly lower than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary. See also Actinium series Notes References Bibliography Meyer, Gerd and Morss, Lester R. (1991) Synthesis of lanthanide and actinide compounds, Springer. External links Actinium at The Periodic Table of Videos (University of Nottingham) NLM Hazardous Substances Databank – Actinium, Radioactive Actinium in Chemical elements Actinides

The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers. The medium half-life of 227Ac (21.77 years) makes it very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities (of the order 50 meters per year). However, evaluation of the concentration depth-profiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows: oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa depth profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior. There are theoretical predictions that AcHx hydrides (in this case with very high pressure) are a candidate for a near room-temperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K. Precautions 227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33% of actinium is deposited into the bones and 50% into the liver. Its toxicity is comparable to, but slightly lower than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary. See also Actinium series Notes References Bibliography Meyer, Gerd and Morss, Lester R. (1991) Synthesis of lanthanide and actinide compounds, Springer. External links Actinium at The Periodic Table of Videos (University of Nottingham) NLM Hazardous Substances Databank – Actinium, Radioactive Actinium in Chemical elements Actinides

The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers. The medium half-life of 227Ac (21.77 years) makes it very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities (of the order 50 meters per year). However, evaluation of the concentration depth-profiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows: oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa depth profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior. There are theoretical predictions that AcHx hydrides (in this case with very high pressure) are a candidate for a near room-temperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K. Precautions 227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33% of actinium is deposited into the bones and 50% into the liver. Its toxicity is comparable to, but slightly lower than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary. See also Actinium series Notes References Bibliography Meyer, Gerd and Morss, Lester R. (1991) Synthesis of lanthanide and actinide compounds, Springer. External links Actinium at The Periodic Table of Videos (University of Nottingham) NLM Hazardous Substances Databank – Actinium, Radioactive Actinium in Chemical elements Actinides

Americium Americium is a synthetic radioactive chemical element with the symbol Am and atomic number 95. It is a transuranic member of the actinide series, in the periodic table located under the lanthanide element europium, and thus by analogy was named after the Americas. Americium was first produced in 1944 by the group of Glenn T. Seaborg from Berkeley, California, at the Metallurgical Laboratory of the University of Chicago, as part of the Manhattan Project. Although it is the third element in the transuranic series, it was discovered fourth, after the heavier curium. The discovery was kept secret and only released to the public in November 1945. Most americium is produced by uranium or plutonium being bombarded with neutrons in nuclear reactors – one tonne of spent nuclear fuel contains about 100 grams of americium. It is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges. Several unusual applications, such as nuclear batteries or fuel for space ships with nuclear propulsion, have been proposed for the isotope 242mAm, but they are as yet hindered by the scarcity and high price of this nuclear isomer. Americium is a relatively soft radioactive metal with silvery appearance. Its most common isotopes are 241Am and 243Am. In chemical compounds, americium usually assumes the oxidation state +3, especially in solutions. Several other oxidation states are known, ranging from +2 to +7, and can be identified by their characteristic optical absorption spectra. The crystal lattice of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by self-irradiation with alpha particles, which accumulates with time; this can cause a drift of some material properties over time, more noticeable in older samples. History Although americium was likely produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60-inch cyclotron at the University of California, Berkeley. The element was chemically identified at the Metallurgical Laboratory (now Argonne National Laboratory) of the University of Chicago. Following the lighter neptunium, plutonium, and heavier curium, americium was the fourth transuranium element to be discovered. At the time, the periodic table had been restructured by Seaborg to its present layout, containing the actinide row below the lanthanide one. This led to americium being located right below its twin lanthanide element europium; it was thus by analogy named after the Americas: "The name americium (after the Americas) and the symbol Am are suggested for the element on the basis of its position as the sixth member of the actinide rare-earth series, analogous to europium, Eu, of the lanthanide series." The new element was isolated from its oxides in a complex, multi-step process.

Americium Americium is a synthetic radioactive chemical element with the symbol Am and atomic number 95. It is a transuranic member of the actinide series, in the periodic table located under the lanthanide element europium, and thus by analogy was named after the Americas. Americium was first produced in 1944 by the group of Glenn T. Seaborg from Berkeley, California, at the Metallurgical Laboratory of the University of Chicago, as part of the Manhattan Project. Although it is the third element in the transuranic series, it was discovered fourth, after the heavier curium. The discovery was kept secret and only released to the public in November 1945. Most americium is produced by uranium or plutonium being bombarded with neutrons in nuclear reactors – one tonne of spent nuclear fuel contains about 100 grams of americium. It is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges. Several unusual applications, such as nuclear batteries or fuel for space ships with nuclear propulsion, have been proposed for the isotope 242mAm, but they are as yet hindered by the scarcity and high price of this nuclear isomer. Americium is a relatively soft radioactive metal with silvery appearance. Its most common isotopes are 241Am and 243Am. In chemical compounds, americium usually assumes the oxidation state +3, especially in solutions. Several other oxidation states are known, ranging from +2 to +7, and can be identified by their characteristic optical absorption spectra. The crystal lattice of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by self-irradiation with alpha particles, which accumulates with time; this can cause a drift of some material properties over time, more noticeable in older samples. History Although americium was likely produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60-inch cyclotron at the University of California, Berkeley. The element was chemically identified at the Metallurgical Laboratory (now Argonne National Laboratory) of the University of Chicago. Following the lighter neptunium, plutonium, and heavier curium, americium was the fourth transuranium element to be discovered. At the time, the periodic table had been restructured by Seaborg to its present layout, containing the actinide row below the lanthanide one. This led to americium being located right below its twin lanthanide element europium; it was thus by analogy named after the Americas: "The name americium (after the Americas) and the symbol Am are suggested for the element on the basis of its position as the sixth member of the actinide rare-earth series, analogous to europium, Eu, of the lanthanide series." The new element was isolated from its oxides in a complex, multi-step process.

First plutonium-239 nitrate (239PuNO3) solution was coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide (PuO2) by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium (from Greek for all demons or hell) and delirium (from Latin for madness). Initial experiments yielded four americium isotopes: 241Am, 242Am, 239Am and 238Am. Americium-241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as years but then corrected to 432.2 years. The times are half-lives The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h. The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor. The initial americium samples weighed a few micrograms; they were barely visible and were identified by their radioactivity. The first substantial amounts of metallic americium weighing 40–200 micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C. Occurrence The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium (americium that was present on Earth during its formation) should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of nuclear reactions, though this has not been confirmed. Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides including americium; but due to military secrecy, this result was not published until later, in 1956.

First plutonium-239 nitrate (239PuNO3) solution was coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide (PuO2) by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium (from Greek for all demons or hell) and delirium (from Latin for madness). Initial experiments yielded four americium isotopes: 241Am, 242Am, 239Am and 238Am. Americium-241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as years but then corrected to 432.2 years. The times are half-lives The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h. The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor. The initial americium samples weighed a few micrograms; they were barely visible and were identified by their radioactivity. The first substantial amounts of metallic americium weighing 40–200 micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C. Occurrence The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium (americium that was present on Earth during its formation) should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of nuclear reactions, though this has not been confirmed. Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides including americium; but due to military secrecy, this result was not published until later, in 1956.

First plutonium-239 nitrate (239PuNO3) solution was coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide (PuO2) by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium (from Greek for all demons or hell) and delirium (from Latin for madness). Initial experiments yielded four americium isotopes: 241Am, 242Am, 239Am and 238Am. Americium-241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as years but then corrected to 432.2 years. The times are half-lives The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h. The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor. The initial americium samples weighed a few micrograms; they were barely visible and were identified by their radioactivity. The first substantial amounts of metallic americium weighing 40–200 micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C. Occurrence The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium (americium that was present on Earth during its formation) should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of nuclear reactions, though this has not been confirmed. Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, (1 November 1952, Enewetak Atoll), revealed high concentrations of various actinides including americium; but due to military secrecy, this result was not published until later, in 1956.

Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland. In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuries/g (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils. Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to short-lived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium. The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so. Americium is also one of the elements that have been detected in Przybylski's Star. Synthesis and extraction Isotope nucleosynthesis Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now. Nevertheless, since it was first offered for sale in 1962, its price, about US$1,500 per gram of 241Am, remains almost unchanged owing to the very complex separation procedure. The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order 100,000–160,000 USD/g. Americium is not synthesized directly from uranium – the most common reactor material – but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process: ^{238}_{92}U ->[\ce{(n,\gamma)}] ^{239}_{92}U ->[\beta^-][23.5 \ \ce{min}] ^{239}_{93}Np ->[\beta^-][2.3565 \ \ce{d}] ^{239}_{94}Pu The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am: ^{239}_{94}Pu ->[\ce{2(n,\gamma)}] ^{241}_{94}Pu ->[\beta^-][14.35 \ \ce{yr}] ^{241}_{95}Am The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it spontaneously converts to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow: half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years. The obtained 241Am can be used for generating heavier americium isotopes by further neutron capture inside a nuclear reactor.

Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland. In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuries/g (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils. Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to short-lived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium. The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so. Americium is also one of the elements that have been detected in Przybylski's Star. Synthesis and extraction Isotope nucleosynthesis Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now. Nevertheless, since it was first offered for sale in 1962, its price, about US$1,500 per gram of 241Am, remains almost unchanged owing to the very complex separation procedure. The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order 100,000–160,000 USD/g. Americium is not synthesized directly from uranium – the most common reactor material – but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process: ^{238}_{92}U ->[\ce{(n,\gamma)}] ^{239}_{92}U ->[\beta^-][23.5 \ \ce{min}] ^{239}_{93}Np ->[\beta^-][2.3565 \ \ce{d}] ^{239}_{94}Pu The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am: ^{239}_{94}Pu ->[\ce{2(n,\gamma)}] ^{241}_{94}Pu ->[\beta^-][14.35 \ \ce{yr}] ^{241}_{95}Am The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it spontaneously converts to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow: half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years. The obtained 241Am can be used for generating heavier americium isotopes by further neutron capture inside a nuclear reactor.

Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland. In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuries/g (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils. Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to short-lived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium. The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so. Americium is also one of the elements that have been detected in Przybylski's Star. Synthesis and extraction Isotope nucleosynthesis Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now. Nevertheless, since it was first offered for sale in 1962, its price, about US$1,500 per gram of 241Am, remains almost unchanged owing to the very complex separation procedure. The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order 100,000–160,000 USD/g. Americium is not synthesized directly from uranium – the most common reactor material – but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process: ^{238}_{92}U ->[\ce{(n,\gamma)}] ^{239}_{92}U ->[\beta^-][23.5 \ \ce{min}] ^{239}_{93}Np ->[\beta^-][2.3565 \ \ce{d}] ^{239}_{94}Pu The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am: ^{239}_{94}Pu ->[\ce{2(n,\gamma)}] ^{241}_{94}Pu ->[\beta^-][14.35 \ \ce{yr}] ^{241}_{95}Am The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it spontaneously converts to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow: half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years. The obtained 241Am can be used for generating heavier americium isotopes by further neutron capture inside a nuclear reactor.

In a light water reactor (LWR), 79% of 241Am converts to 242Am and 10% to its nuclear isomer 242mAm: Americium-242 has a half-life of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux: ^{239}_{94}Pu ->[\ce{4(n,\gamma)}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ \ce{h}] ^{243}_{95}Am Metal generation Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of americium can be separated. In a typical procedure, the spent reactor fuel (e.g. MOX fuel) is dissolved in nitric acid, and the bulk of uranium and plutonium is removed using a PUREX-type extraction (Plutonium–URanium EXtraction) with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue (raffinate) by a diamide-based extraction, to give, after stripping, a mixture of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multi-step chromatographic and centrifugation techniques with an appropriate reagent. A large amount of work has been done on the solvent extraction of americium. For example, a 2003 EU-funded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bis-triazinyl bipyridine complex was proposed in 2009 as such a reagent is highly selective to americium (and curium). Separation of americium from the highly similar curium can be achieved by treating a slurry of their hydroxides in aqueous sodium bicarbonate with ozone, at elevated temperatures. Both Am and Cm are mostly present in solutions in the +3 valence state; whereas curium remains unchanged, americium oxidizes to soluble Am(IV) complexes which can be washed away. Metallic americium is obtained by reduction from its compounds. Americium(III) fluoride was first used for this purpose. The reaction was conducted using elemental barium as reducing agent in a water- and oxygen-free environment inside an apparatus made of tantalum and tungsten. An alternative is the reduction of americium dioxide by metallic lanthanum or thorium: Physical properties In the periodic table, americium is located to the right of plutonium, to the left of curium, and below the lanthanide europium, with which it shares many physical and chemical properties. Americium is a highly radioactive element. When freshly prepared, it has a silvery-white metallic lustre, but then slowly tarnishes in air. With a density of 12 g/cm3, americium is less dense than both curium (13.52 g/cm3) and plutonium (19.8 g/cm3); but has a higher density than europium (5.264 g/cm3)—mostly because of its higher atomic mass. Americium is relatively soft and easily deformable and has a significantly lower bulk modulus than the actinides before it: Th, Pa, U, Np and Pu. Its melting point of 1173 °C is significantly higher than that of plutonium (639 °C) and europium (826 °C), but lower than for curium (1340 °C).

In a light water reactor (LWR), 79% of 241Am converts to 242Am and 10% to its nuclear isomer 242mAm: Americium-242 has a half-life of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux: ^{239}_{94}Pu ->[\ce{4(n,\gamma)}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ \ce{h}] ^{243}_{95}Am Metal generation Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of americium can be separated. In a typical procedure, the spent reactor fuel (e.g. MOX fuel) is dissolved in nitric acid, and the bulk of uranium and plutonium is removed using a PUREX-type extraction (Plutonium–URanium EXtraction) with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue (raffinate) by a diamide-based extraction, to give, after stripping, a mixture of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multi-step chromatographic and centrifugation techniques with an appropriate reagent. A large amount of work has been done on the solvent extraction of americium. For example, a 2003 EU-funded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bis-triazinyl bipyridine complex was proposed in 2009 as such a reagent is highly selective to americium (and curium). Separation of americium from the highly similar curium can be achieved by treating a slurry of their hydroxides in aqueous sodium bicarbonate with ozone, at elevated temperatures. Both Am and Cm are mostly present in solutions in the +3 valence state; whereas curium remains unchanged, americium oxidizes to soluble Am(IV) complexes which can be washed away. Metallic americium is obtained by reduction from its compounds. Americium(III) fluoride was first used for this purpose. The reaction was conducted using elemental barium as reducing agent in a water- and oxygen-free environment inside an apparatus made of tantalum and tungsten. An alternative is the reduction of americium dioxide by metallic lanthanum or thorium: Physical properties In the periodic table, americium is located to the right of plutonium, to the left of curium, and below the lanthanide europium, with which it shares many physical and chemical properties. Americium is a highly radioactive element. When freshly prepared, it has a silvery-white metallic lustre, but then slowly tarnishes in air. With a density of 12 g/cm3, americium is less dense than both curium (13.52 g/cm3) and plutonium (19.8 g/cm3); but has a higher density than europium (5.264 g/cm3)—mostly because of its higher atomic mass. Americium is relatively soft and easily deformable and has a significantly lower bulk modulus than the actinides before it: Th, Pa, U, Np and Pu. Its melting point of 1173 °C is significantly higher than that of plutonium (639 °C) and europium (826 °C), but lower than for curium (1340 °C).

In a light water reactor (LWR), 79% of 241Am converts to 242Am and 10% to its nuclear isomer 242mAm: Americium-242 has a half-life of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux: ^{239}_{94}Pu ->[\ce{4(n,\gamma)}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ \ce{h}] ^{243}_{95}Am Metal generation Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of americium can be separated. In a typical procedure, the spent reactor fuel (e.g. MOX fuel) is dissolved in nitric acid, and the bulk of uranium and plutonium is removed using a PUREX-type extraction (Plutonium–URanium EXtraction) with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue (raffinate) by a diamide-based extraction, to give, after stripping, a mixture of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multi-step chromatographic and centrifugation techniques with an appropriate reagent. A large amount of work has been done on the solvent extraction of americium. For example, a 2003 EU-funded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bis-triazinyl bipyridine complex was proposed in 2009 as such a reagent is highly selective to americium (and curium). Separation of americium from the highly similar curium can be achieved by treating a slurry of their hydroxides in aqueous sodium bicarbonate with ozone, at elevated temperatures. Both Am and Cm are mostly present in solutions in the +3 valence state; whereas curium remains unchanged, americium oxidizes to soluble Am(IV) complexes which can be washed away. Metallic americium is obtained by reduction from its compounds. Americium(III) fluoride was first used for this purpose. The reaction was conducted using elemental barium as reducing agent in a water- and oxygen-free environment inside an apparatus made of tantalum and tungsten. An alternative is the reduction of americium dioxide by metallic lanthanum or thorium: Physical properties In the periodic table, americium is located to the right of plutonium, to the left of curium, and below the lanthanide europium, with which it shares many physical and chemical properties. Americium is a highly radioactive element. When freshly prepared, it has a silvery-white metallic lustre, but then slowly tarnishes in air. With a density of 12 g/cm3, americium is less dense than both curium (13.52 g/cm3) and plutonium (19.8 g/cm3); but has a higher density than europium (5.264 g/cm3)—mostly because of its higher atomic mass. Americium is relatively soft and easily deformable and has a significantly lower bulk modulus than the actinides before it: Th, Pa, U, Np and Pu. Its melting point of 1173 °C is significantly higher than that of plutonium (639 °C) and europium (826 °C), but lower than for curium (1340 °C).

At ambient conditions, americium is present in its most stable α form which has a hexagonal crystal symmetry, and a space group P63/mmc with cell parameters a = 346.8 pm and c = 1124 pm, and four atoms per unit cell. The crystal consists of a double-hexagonal close packing with the layer sequence ABAC and so is isotypic with α-lanthanum and several actinides such as α-curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, α-Am transforms to the β modification, which has a face-centered cubic (fcc) symmetry, space group Fmm and lattice constant a = 489 pm. This fcc structure is equivalent to the closest packing with the sequence ABC. Upon further compression to 23 GPa, americium transforms to an orthorhombic γ-Am structure similar to that of α-uranium. There are no further transitions observed up to 52 GPa, except for an appearance of a monoclinic phase at pressures between 10 and 15 GPa. There is no consistency on the status of this phase in the literature, which also sometimes lists the α, β and γ phases as I, II and III. The β-γ transition is accompanied by a 6% decrease in the crystal volume; although theory also predicts a significant volume change for the α-β transition, it is not observed experimentally. The pressure of the α-β transition decreases with increasing temperature, and when α-americium is heated at ambient pressure, at 770 °C it changes into an fcc phase which is different from β-Am, and at 1075 °C it converts to a body-centered cubic structure. The pressure-temperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium. As with many other actinides, self-damage of the crystal structure due to alpha-particle irradiation is intrinsic to americium. It is especially noticeable at low temperatures, where the mobility of the produced structure defects is relatively low, by broadening of X-ray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as electrical resistivity. So for americium-241, the resistivity at 4.2 K increases with time from about 2 µOhm·cm to 10 µOhm·cm after 40 hours, and saturates at about 16 µOhm·cm after 140 hours. This effect is less pronounced at room temperature, due to annihilation of radiation defects; also heating to room temperature the sample which was kept for hours at low temperatures restores its resistivity. In fresh samples, the resistivity gradually increases with temperature from about 2 µOhm·cm at liquid helium to 69 µOhm·cm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which show a rapid rise up to 60 K followed by saturation. The room temperature value for americium is lower than that of neptunium, plutonium and curium, but higher than for uranium, thorium and protactinium. Americium is paramagnetic in a wide temperature range, from that of liquid helium, to room temperature and above.

At ambient conditions, americium is present in its most stable α form which has a hexagonal crystal symmetry, and a space group P63/mmc with cell parameters a = 346.8 pm and c = 1124 pm, and four atoms per unit cell. The crystal consists of a double-hexagonal close packing with the layer sequence ABAC and so is isotypic with α-lanthanum and several actinides such as α-curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, α-Am transforms to the β modification, which has a face-centered cubic (fcc) symmetry, space group Fmm and lattice constant a = 489 pm. This fcc structure is equivalent to the closest packing with the sequence ABC. Upon further compression to 23 GPa, americium transforms to an orthorhombic γ-Am structure similar to that of α-uranium. There are no further transitions observed up to 52 GPa, except for an appearance of a monoclinic phase at pressures between 10 and 15 GPa. There is no consistency on the status of this phase in the literature, which also sometimes lists the α, β and γ phases as I, II and III. The β-γ transition is accompanied by a 6% decrease in the crystal volume; although theory also predicts a significant volume change for the α-β transition, it is not observed experimentally. The pressure of the α-β transition decreases with increasing temperature, and when α-americium is heated at ambient pressure, at 770 °C it changes into an fcc phase which is different from β-Am, and at 1075 °C it converts to a body-centered cubic structure. The pressure-temperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium. As with many other actinides, self-damage of the crystal structure due to alpha-particle irradiation is intrinsic to americium. It is especially noticeable at low temperatures, where the mobility of the produced structure defects is relatively low, by broadening of X-ray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as electrical resistivity. So for americium-241, the resistivity at 4.2 K increases with time from about 2 µOhm·cm to 10 µOhm·cm after 40 hours, and saturates at about 16 µOhm·cm after 140 hours. This effect is less pronounced at room temperature, due to annihilation of radiation defects; also heating to room temperature the sample which was kept for hours at low temperatures restores its resistivity. In fresh samples, the resistivity gradually increases with temperature from about 2 µOhm·cm at liquid helium to 69 µOhm·cm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which show a rapid rise up to 60 K followed by saturation. The room temperature value for americium is lower than that of neptunium, plutonium and curium, but higher than for uranium, thorium and protactinium. Americium is paramagnetic in a wide temperature range, from that of liquid helium, to room temperature and above.

At ambient conditions, americium is present in its most stable α form which has a hexagonal crystal symmetry, and a space group P63/mmc with cell parameters a = 346.8 pm and c = 1124 pm, and four atoms per unit cell. The crystal consists of a double-hexagonal close packing with the layer sequence ABAC and so is isotypic with α-lanthanum and several actinides such as α-curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, α-Am transforms to the β modification, which has a face-centered cubic (fcc) symmetry, space group Fmm and lattice constant a = 489 pm. This fcc structure is equivalent to the closest packing with the sequence ABC. Upon further compression to 23 GPa, americium transforms to an orthorhombic γ-Am structure similar to that of α-uranium. There are no further transitions observed up to 52 GPa, except for an appearance of a monoclinic phase at pressures between 10 and 15 GPa. There is no consistency on the status of this phase in the literature, which also sometimes lists the α, β and γ phases as I, II and III. The β-γ transition is accompanied by a 6% decrease in the crystal volume; although theory also predicts a significant volume change for the α-β transition, it is not observed experimentally. The pressure of the α-β transition decreases with increasing temperature, and when α-americium is heated at ambient pressure, at 770 °C it changes into an fcc phase which is different from β-Am, and at 1075 °C it converts to a body-centered cubic structure. The pressure-temperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium. As with many other actinides, self-damage of the crystal structure due to alpha-particle irradiation is intrinsic to americium. It is especially noticeable at low temperatures, where the mobility of the produced structure defects is relatively low, by broadening of X-ray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as electrical resistivity. So for americium-241, the resistivity at 4.2 K increases with time from about 2 µOhm·cm to 10 µOhm·cm after 40 hours, and saturates at about 16 µOhm·cm after 140 hours. This effect is less pronounced at room temperature, due to annihilation of radiation defects; also heating to room temperature the sample which was kept for hours at low temperatures restores its resistivity. In fresh samples, the resistivity gradually increases with temperature from about 2 µOhm·cm at liquid helium to 69 µOhm·cm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which show a rapid rise up to 60 K followed by saturation. The room temperature value for americium is lower than that of neptunium, plutonium and curium, but higher than for uranium, thorium and protactinium. Americium is paramagnetic in a wide temperature range, from that of liquid helium, to room temperature and above.

This behavior is markedly different from that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to along the shorter a axis and for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is , from which the standard enthalpy change of formation (ΔfH°) of aqueous Am3+ ion is . The standard potential Am3+/Am0 is . Chemical properties Americium metal readily reacts with oxygen and dissolves in aqueous acids. The most stable oxidation state for americium is +3,. The chemistry of americium(III) has many similarities to the chemistry of lanthanide(III) compounds. For example, trivalent americium forms insoluble fluoride, oxalate, iodate, hydroxide, phosphate and other salts. Compounds of americium in oxidation states 2, 4, 5, 6 and 7 have also been studied. This is the widest range that has been observed with actinide elements. The color of americium compounds in aqueous solution is as follows: Am3+ (yellow-reddish), Am4+ (yellow-reddish), AmV; (yellow), AmVI (brown) and AmVII (dark green). The absorption spectra have sharp peaks, due to f-f transitions' in the visible and near-infrared regions. Typically, Am(III) has absorption maxima at ca. 504 and 811 nm, Am(V) at ca. 514 and 715 nm, and Am(VI) at ca. 666 and 992 nm. Americium compounds with oxidation state +4 and higher are strong oxidizing agents, comparable in strength to the permanganate ion () in acidic solutions. Whereas the Am4+ ions are unstable in solutions and readily convert to Am3+, compounds such as americium dioxide (AmO2) and americium(IV) fluoride (AmF4) are stable in the solid state. The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the ion is unstable with respect to disproportionation. The reaction 3[AmO2]+ + 4H+ -> 2[AmO2]2+ + Am3+ + 2H2O is typical. The chemistry of Am(V) and Am(VI) is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like Li3AmO4 and Li6AmO6 are comparable to uranates and the ion AmO22+ is comparable to the uranyl ion, UO22+. Such compounds can be prepared by oxidation of Am(III) in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silver(I) oxide, ozone and sodium persulfate. Chemical compounds Oxygen compounds Three americium oxides are known, with the oxidation states +2 (AmO), +3 (Am2O3) and +4 (AmO2). Americium(II) oxide was prepared in minute amounts and has not been characterized in detail. Americium(III) oxide is a red-brown solid with a melting point of 2205 °C. Americium(IV) oxide is the main form of solid americium which is used in nearly all its applications. As most other actinide dioxides, it is a black solid with a cubic (fluorite) crystal structure. The oxalate of americium(III), vacuum dried at room temperature, has the chemical formula Am2(C2O4)3·7H2O. Upon heating in vacuum, it loses water at 240 °C and starts decomposing into AmO2 at 300 °C, the decomposition completes at about 470 °C.

This behavior is markedly different from that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to along the shorter a axis and for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is , from which the standard enthalpy change of formation (ΔfH°) of aqueous Am3+ ion is . The standard potential Am3+/Am0 is . Chemical properties Americium metal readily reacts with oxygen and dissolves in aqueous acids. The most stable oxidation state for americium is +3,. The chemistry of americium(III) has many similarities to the chemistry of lanthanide(III) compounds. For example, trivalent americium forms insoluble fluoride, oxalate, iodate, hydroxide, phosphate and other salts. Compounds of americium in oxidation states 2, 4, 5, 6 and 7 have also been studied. This is the widest range that has been observed with actinide elements. The color of americium compounds in aqueous solution is as follows: Am3+ (yellow-reddish), Am4+ (yellow-reddish), AmV; (yellow), AmVI (brown) and AmVII (dark green). The absorption spectra have sharp peaks, due to f-f transitions' in the visible and near-infrared regions. Typically, Am(III) has absorption maxima at ca. 504 and 811 nm, Am(V) at ca. 514 and 715 nm, and Am(VI) at ca. 666 and 992 nm. Americium compounds with oxidation state +4 and higher are strong oxidizing agents, comparable in strength to the permanganate ion () in acidic solutions. Whereas the Am4+ ions are unstable in solutions and readily convert to Am3+, compounds such as americium dioxide (AmO2) and americium(IV) fluoride (AmF4) are stable in the solid state. The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the ion is unstable with respect to disproportionation. The reaction 3[AmO2]+ + 4H+ -> 2[AmO2]2+ + Am3+ + 2H2O is typical. The chemistry of Am(V) and Am(VI) is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like Li3AmO4 and Li6AmO6 are comparable to uranates and the ion AmO22+ is comparable to the uranyl ion, UO22+. Such compounds can be prepared by oxidation of Am(III) in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silver(I) oxide, ozone and sodium persulfate. Chemical compounds Oxygen compounds Three americium oxides are known, with the oxidation states +2 (AmO), +3 (Am2O3) and +4 (AmO2). Americium(II) oxide was prepared in minute amounts and has not been characterized in detail. Americium(III) oxide is a red-brown solid with a melting point of 2205 °C. Americium(IV) oxide is the main form of solid americium which is used in nearly all its applications. As most other actinide dioxides, it is a black solid with a cubic (fluorite) crystal structure. The oxalate of americium(III), vacuum dried at room temperature, has the chemical formula Am2(C2O4)3·7H2O. Upon heating in vacuum, it loses water at 240 °C and starts decomposing into AmO2 at 300 °C, the decomposition completes at about 470 °C.

This behavior is markedly different from that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to along the shorter a axis and for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is , from which the standard enthalpy change of formation (ΔfH°) of aqueous Am3+ ion is . The standard potential Am3+/Am0 is . Chemical properties Americium metal readily reacts with oxygen and dissolves in aqueous acids. The most stable oxidation state for americium is +3,. The chemistry of americium(III) has many similarities to the chemistry of lanthanide(III) compounds. For example, trivalent americium forms insoluble fluoride, oxalate, iodate, hydroxide, phosphate and other salts. Compounds of americium in oxidation states 2, 4, 5, 6 and 7 have also been studied. This is the widest range that has been observed with actinide elements. The color of americium compounds in aqueous solution is as follows: Am3+ (yellow-reddish), Am4+ (yellow-reddish), AmV; (yellow), AmVI (brown) and AmVII (dark green). The absorption spectra have sharp peaks, due to f-f transitions' in the visible and near-infrared regions. Typically, Am(III) has absorption maxima at ca. 504 and 811 nm, Am(V) at ca. 514 and 715 nm, and Am(VI) at ca. 666 and 992 nm. Americium compounds with oxidation state +4 and higher are strong oxidizing agents, comparable in strength to the permanganate ion () in acidic solutions. Whereas the Am4+ ions are unstable in solutions and readily convert to Am3+, compounds such as americium dioxide (AmO2) and americium(IV) fluoride (AmF4) are stable in the solid state. The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the ion is unstable with respect to disproportionation. The reaction 3[AmO2]+ + 4H+ -> 2[AmO2]2+ + Am3+ + 2H2O is typical. The chemistry of Am(V) and Am(VI) is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like Li3AmO4 and Li6AmO6 are comparable to uranates and the ion AmO22+ is comparable to the uranyl ion, UO22+. Such compounds can be prepared by oxidation of Am(III) in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silver(I) oxide, ozone and sodium persulfate. Chemical compounds Oxygen compounds Three americium oxides are known, with the oxidation states +2 (AmO), +3 (Am2O3) and +4 (AmO2). Americium(II) oxide was prepared in minute amounts and has not been characterized in detail. Americium(III) oxide is a red-brown solid with a melting point of 2205 °C. Americium(IV) oxide is the main form of solid americium which is used in nearly all its applications. As most other actinide dioxides, it is a black solid with a cubic (fluorite) crystal structure. The oxalate of americium(III), vacuum dried at room temperature, has the chemical formula Am2(C2O4)3·7H2O. Upon heating in vacuum, it loses water at 240 °C and starts decomposing into AmO2 at 300 °C, the decomposition completes at about 470 °C.

The initial oxalate dissolves in nitric acid with the maximum solubility of 0.25 g/L. Halides Halides of americium are known for the oxidation states +2, +3 and +4, where the +3 is most stable, especially in solutions. Reduction of Am(III) compounds with sodium amalgam yields Am(II) salts – the black halides AmCl2, AmBr2 and AmI2. They are very sensitive to oxygen and oxidize in water, releasing hydrogen and converting back to the Am(III) state. Specific lattice constants are: Orthorhombic AmCl2: a = , b = and c = Tetragonal AmBr2: a = and c = . They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X = Cl, Br or I: {Am} + \underset{mercury\ halide}{HgX2} ->[{} \atop 400 - 500 ^\circ \ce C] {AmX2} + {Hg} Americium(III) fluoride (AmF3) is poorly soluble and precipitates upon reaction of Am3+ and fluoride ions in weak acidic solutions: Am^3+ + 3F^- -> AmF3(v) The tetravalent americium(IV) fluoride (AmF4) is obtained by reacting solid americium(III) fluoride with molecular fluorine: 2AmF3 + F2 -> 2AmF4 Another known form of solid tetravalent americium fluoride is KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black Am(OH)4 was dissolved in 15-M NH4F with the americium concentration of 0.01 M. The resulting reddish solution had a characteristic optical absorption spectrum which is similar to that of AmF4 but differed from other oxidation states of americium. Heating the Am(IV) solution to 90 °C did not result in its disproportionation or reduction, however a slow reduction was observed to Am(III) and assigned to self-irradiation of americium by alpha particles. Most americium(III) halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride (AmCl3) is reddish and has a structure isotypic to uranium(III) chloride (space group P63/m) and the melting point of 715 °C. The fluoride is isotypic to LaF3 (space group P63/mmc) and the iodide to BiI3 (space group R). The bromide is an exception with the orthorhombic PuBr3-type structure and space group Cmcm. Crystals of americium hexahydrate (AmCl3·6H2O) can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellow-reddish color and a monoclinic crystal structure. Oxyhalides of americium in the form AmVIO2X2, AmVO2X, AmIVOX2 and AmIIIOX can be obtained by reacting the corresponding americium halide with oxygen or Sb2O3, and AmOCl can also be produced by vapor phase hydrolysis: AmCl3 + H2O -> AmOCl + 2HCl Chalcogenides and pnictides The known chalcogenides of americium include the sulfide AmS2, selenides AmSe2 and Am3Se4, and tellurides Am2Te3 and AmTe2. The pnictides of americium (243Am) of the AmX type are known for the elements phosphorus, arsenic, antimony and bismuth. They crystallize in the rock-salt lattice. Silicides and borides Americium monosilicide (AmSi) and "disilicide" (nominally AmSix with: 1.87 < x < 2.0) were obtained by reduction of americium(III) fluoride with elementary silicon in vacuum at 1050 °C (AmSi) and 1150−1200 °C (AmSix).

The initial oxalate dissolves in nitric acid with the maximum solubility of 0.25 g/L. Halides Halides of americium are known for the oxidation states +2, +3 and +4, where the +3 is most stable, especially in solutions. Reduction of Am(III) compounds with sodium amalgam yields Am(II) salts – the black halides AmCl2, AmBr2 and AmI2. They are very sensitive to oxygen and oxidize in water, releasing hydrogen and converting back to the Am(III) state. Specific lattice constants are: Orthorhombic AmCl2: a = , b = and c = Tetragonal AmBr2: a = and c = . They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X = Cl, Br or I: {Am} + \underset{mercury\ halide}{HgX2} ->[{} \atop 400 - 500 ^\circ \ce C] {AmX2} + {Hg} Americium(III) fluoride (AmF3) is poorly soluble and precipitates upon reaction of Am3+ and fluoride ions in weak acidic solutions: Am^3+ + 3F^- -> AmF3(v) The tetravalent americium(IV) fluoride (AmF4) is obtained by reacting solid americium(III) fluoride with molecular fluorine: 2AmF3 + F2 -> 2AmF4 Another known form of solid tetravalent americium fluoride is KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black Am(OH)4 was dissolved in 15-M NH4F with the americium concentration of 0.01 M. The resulting reddish solution had a characteristic optical absorption spectrum which is similar to that of AmF4 but differed from other oxidation states of americium. Heating the Am(IV) solution to 90 °C did not result in its disproportionation or reduction, however a slow reduction was observed to Am(III) and assigned to self-irradiation of americium by alpha particles. Most americium(III) halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride (AmCl3) is reddish and has a structure isotypic to uranium(III) chloride (space group P63/m) and the melting point of 715 °C. The fluoride is isotypic to LaF3 (space group P63/mmc) and the iodide to BiI3 (space group R). The bromide is an exception with the orthorhombic PuBr3-type structure and space group Cmcm. Crystals of americium hexahydrate (AmCl3·6H2O) can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellow-reddish color and a monoclinic crystal structure. Oxyhalides of americium in the form AmVIO2X2, AmVO2X, AmIVOX2 and AmIIIOX can be obtained by reacting the corresponding americium halide with oxygen or Sb2O3, and AmOCl can also be produced by vapor phase hydrolysis: AmCl3 + H2O -> AmOCl + 2HCl Chalcogenides and pnictides The known chalcogenides of americium include the sulfide AmS2, selenides AmSe2 and Am3Se4, and tellurides Am2Te3 and AmTe2. The pnictides of americium (243Am) of the AmX type are known for the elements phosphorus, arsenic, antimony and bismuth. They crystallize in the rock-salt lattice. Silicides and borides Americium monosilicide (AmSi) and "disilicide" (nominally AmSix with: 1.87 < x < 2.0) were obtained by reduction of americium(III) fluoride with elementary silicon in vacuum at 1050 °C (AmSi) and 1150−1200 °C (AmSix).

The initial oxalate dissolves in nitric acid with the maximum solubility of 0.25 g/L. Halides Halides of americium are known for the oxidation states +2, +3 and +4, where the +3 is most stable, especially in solutions. Reduction of Am(III) compounds with sodium amalgam yields Am(II) salts – the black halides AmCl2, AmBr2 and AmI2. They are very sensitive to oxygen and oxidize in water, releasing hydrogen and converting back to the Am(III) state. Specific lattice constants are: Orthorhombic AmCl2: a = , b = and c = Tetragonal AmBr2: a = and c = . They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X = Cl, Br or I: {Am} + \underset{mercury\ halide}{HgX2} ->[{} \atop 400 - 500 ^\circ \ce C] {AmX2} + {Hg} Americium(III) fluoride (AmF3) is poorly soluble and precipitates upon reaction of Am3+ and fluoride ions in weak acidic solutions: Am^3+ + 3F^- -> AmF3(v) The tetravalent americium(IV) fluoride (AmF4) is obtained by reacting solid americium(III) fluoride with molecular fluorine: 2AmF3 + F2 -> 2AmF4 Another known form of solid tetravalent americium fluoride is KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black Am(OH)4 was dissolved in 15-M NH4F with the americium concentration of 0.01 M. The resulting reddish solution had a characteristic optical absorption spectrum which is similar to that of AmF4 but differed from other oxidation states of americium. Heating the Am(IV) solution to 90 °C did not result in its disproportionation or reduction, however a slow reduction was observed to Am(III) and assigned to self-irradiation of americium by alpha particles. Most americium(III) halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride (AmCl3) is reddish and has a structure isotypic to uranium(III) chloride (space group P63/m) and the melting point of 715 °C. The fluoride is isotypic to LaF3 (space group P63/mmc) and the iodide to BiI3 (space group R). The bromide is an exception with the orthorhombic PuBr3-type structure and space group Cmcm. Crystals of americium hexahydrate (AmCl3·6H2O) can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellow-reddish color and a monoclinic crystal structure. Oxyhalides of americium in the form AmVIO2X2, AmVO2X, AmIVOX2 and AmIIIOX can be obtained by reacting the corresponding americium halide with oxygen or Sb2O3, and AmOCl can also be produced by vapor phase hydrolysis: AmCl3 + H2O -> AmOCl + 2HCl Chalcogenides and pnictides The known chalcogenides of americium include the sulfide AmS2, selenides AmSe2 and Am3Se4, and tellurides Am2Te3 and AmTe2. The pnictides of americium (243Am) of the AmX type are known for the elements phosphorus, arsenic, antimony and bismuth. They crystallize in the rock-salt lattice. Silicides and borides Americium monosilicide (AmSi) and "disilicide" (nominally AmSix with: 1.87 < x < 2.0) were obtained by reduction of americium(III) fluoride with elementary silicon in vacuum at 1050 °C (AmSi) and 1150−1200 °C (AmSix).

AmSi is a black solid isomorphic with LaSi, it has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice (space group I41/amd), it is isomorphic with PuSi2 and ThSi2. Borides of americium include AmB4 and AmB6. The tetraboride can be obtained by heating an oxide or halide of americium with magnesium diboride in vacuum or inert atmosphere. Organoamericium compounds Analogous to uranocene, americium forms the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula (η8-C8H8)2Am. A cyclopentadienyl complex is also known that is likely to be stoichiometrically AmCp3. Formation of the complexes of the type Am(n-C3H7-BTP)3, where BTP stands for 2,6-di(1,2,4-triazin-3-yl)pyridine, in solutions containing n-C3H7-BTP and Am3+ ions has been confirmed by EXAFS. Some of these BTP-type complexes selectively interact with americium and therefore are useful in its selective separation from lanthanides and another actinides. Biological aspects Americium is an artificial element of recent origin, and thus does not have a biological requirement. It is harmful to life. It has been proposed to use bacteria for removal of americium and other heavy metals from rivers and streams. Thus, Enterobacteriaceae of the genus Citrobacter precipitate americium ions from aqueous solutions, binding them into a metal-phosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi. Fission The isotope 242mAm (half-life 141 years) has the largest cross sections for absorption of thermal neutrons (5,700 barns), that results in a small critical mass for a sustained nuclear chain reaction. The critical mass for a bare 242mAm sphere is about 9–14 kg (the uncertainty results from insufficient knowledge of its material properties). It can be lowered to 3–5 kg with a metal reflector and should become even smaller with a water reflector. Such small critical mass is favorable for portable nuclear weapons, but those based on 242mAm are not known yet, probably because of its scarcity and high price. The critical masses of two other readily available isotopes, 241Am and 243Am, are relatively high – 57.6 to 75.6 kg for 241Am and 209 kg for 243Am. Scarcity and high price yet hinder application of americium as a nuclear fuel in nuclear reactors. There are proposals of very compact 10-kW high-flux reactors using as little as 20 grams of 242mAm. Such low-power reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals. Isotopes About 19 isotopes and 8 nuclear isomers are known for americium. There are two long-lived alpha-emitters; 243Am has a half-life of 7,370 years and is the most stable isotope, and 241Am has a half-life of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long half-life of 141 years. The half-lives of other isotopes and isomers range from 0.64 microseconds for 245m1Am to 50.8 hours for 240Am. As with most other actinides, the isotopes of americium with odd number of neutrons have relatively high rate of nuclear fission and low critical mass.

AmSi is a black solid isomorphic with LaSi, it has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice (space group I41/amd), it is isomorphic with PuSi2 and ThSi2. Borides of americium include AmB4 and AmB6. The tetraboride can be obtained by heating an oxide or halide of americium with magnesium diboride in vacuum or inert atmosphere. Organoamericium compounds Analogous to uranocene, americium forms the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula (η8-C8H8)2Am. A cyclopentadienyl complex is also known that is likely to be stoichiometrically AmCp3. Formation of the complexes of the type Am(n-C3H7-BTP)3, where BTP stands for 2,6-di(1,2,4-triazin-3-yl)pyridine, in solutions containing n-C3H7-BTP and Am3+ ions has been confirmed by EXAFS. Some of these BTP-type complexes selectively interact with americium and therefore are useful in its selective separation from lanthanides and another actinides. Biological aspects Americium is an artificial element of recent origin, and thus does not have a biological requirement. It is harmful to life. It has been proposed to use bacteria for removal of americium and other heavy metals from rivers and streams. Thus, Enterobacteriaceae of the genus Citrobacter precipitate americium ions from aqueous solutions, binding them into a metal-phosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi. Fission The isotope 242mAm (half-life 141 years) has the largest cross sections for absorption of thermal neutrons (5,700 barns), that results in a small critical mass for a sustained nuclear chain reaction. The critical mass for a bare 242mAm sphere is about 9–14 kg (the uncertainty results from insufficient knowledge of its material properties). It can be lowered to 3–5 kg with a metal reflector and should become even smaller with a water reflector. Such small critical mass is favorable for portable nuclear weapons, but those based on 242mAm are not known yet, probably because of its scarcity and high price. The critical masses of two other readily available isotopes, 241Am and 243Am, are relatively high – 57.6 to 75.6 kg for 241Am and 209 kg for 243Am. Scarcity and high price yet hinder application of americium as a nuclear fuel in nuclear reactors. There are proposals of very compact 10-kW high-flux reactors using as little as 20 grams of 242mAm. Such low-power reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals. Isotopes About 19 isotopes and 8 nuclear isomers are known for americium. There are two long-lived alpha-emitters; 243Am has a half-life of 7,370 years and is the most stable isotope, and 241Am has a half-life of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long half-life of 141 years. The half-lives of other isotopes and isomers range from 0.64 microseconds for 245m1Am to 50.8 hours for 240Am. As with most other actinides, the isotopes of americium with odd number of neutrons have relatively high rate of nuclear fission and low critical mass.

AmSi is a black solid isomorphic with LaSi, it has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice (space group I41/amd), it is isomorphic with PuSi2 and ThSi2. Borides of americium include AmB4 and AmB6. The tetraboride can be obtained by heating an oxide or halide of americium with magnesium diboride in vacuum or inert atmosphere. Organoamericium compounds Analogous to uranocene, americium forms the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula (η8-C8H8)2Am. A cyclopentadienyl complex is also known that is likely to be stoichiometrically AmCp3. Formation of the complexes of the type Am(n-C3H7-BTP)3, where BTP stands for 2,6-di(1,2,4-triazin-3-yl)pyridine, in solutions containing n-C3H7-BTP and Am3+ ions has been confirmed by EXAFS. Some of these BTP-type complexes selectively interact with americium and therefore are useful in its selective separation from lanthanides and another actinides. Biological aspects Americium is an artificial element of recent origin, and thus does not have a biological requirement. It is harmful to life. It has been proposed to use bacteria for removal of americium and other heavy metals from rivers and streams. Thus, Enterobacteriaceae of the genus Citrobacter precipitate americium ions from aqueous solutions, binding them into a metal-phosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi. Fission The isotope 242mAm (half-life 141 years) has the largest cross sections for absorption of thermal neutrons (5,700 barns), that results in a small critical mass for a sustained nuclear chain reaction. The critical mass for a bare 242mAm sphere is about 9–14 kg (the uncertainty results from insufficient knowledge of its material properties). It can be lowered to 3–5 kg with a metal reflector and should become even smaller with a water reflector. Such small critical mass is favorable for portable nuclear weapons, but those based on 242mAm are not known yet, probably because of its scarcity and high price. The critical masses of two other readily available isotopes, 241Am and 243Am, are relatively high – 57.6 to 75.6 kg for 241Am and 209 kg for 243Am. Scarcity and high price yet hinder application of americium as a nuclear fuel in nuclear reactors. There are proposals of very compact 10-kW high-flux reactors using as little as 20 grams of 242mAm. Such low-power reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals. Isotopes About 19 isotopes and 8 nuclear isomers are known for americium. There are two long-lived alpha-emitters; 243Am has a half-life of 7,370 years and is the most stable isotope, and 241Am has a half-life of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long half-life of 141 years. The half-lives of other isotopes and isomers range from 0.64 microseconds for 245m1Am to 50.8 hours for 240Am. As with most other actinides, the isotopes of americium with odd number of neutrons have relatively high rate of nuclear fission and low critical mass.

Americium-241 decays to 237Np emitting alpha particles of 5 different energies, mostly at 5.486 MeV (85.2%) and 5.443 MeV (12.8%). Because many of the resulting states are metastable, they also emit gamma rays with the discrete energies between 26.3 and 158.5 keV. Americium-242 is a short-lived isotope with a half-life of 16.02 h. It mostly (82.7%) converts by β-decay to 242Cm, but also by electron capture to 242Pu (17.3%). Both 242Cm and 242Pu transform via nearly the same decay chain through 238Pu down to 234U. Nearly all (99.541%) of 242m1Am decays by internal conversion to 242Am and the remaining 0.459% by α-decay to 238Np. The latter subsequently decays to 238Pu and then to 234U. Americium-243 transforms by α-emission into 239Np, which converts by β-decay to 239Pu, and the 239Pu changes into 235U by emitting an α-particle. Applications Ionization-type smoke detector Americium is used in the most common type of household smoke detector, which uses 241Am in the form of americium dioxide as its source of ionizing radiation. This isotope is preferred over 226Ra because it emits 5 times more alpha particles and relatively little harmful gamma radiation. The amount of americium in a typical new smoke detector is 1 microcurie (37 kBq) or 0.29 microgram. This amount declines slowly as the americium decays into neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life of 432.2 years, the americium in a smoke detector includes about 3% neptunium after 19 years, and about 5% after 32 years. The radiation passes through an ionization chamber, an air-filled space between two electrodes, and permits a small, constant current between the electrodes. Any smoke that enters the chamber absorbs the alpha particles, which reduces the ionization and affects this current, triggering the alarm. Compared to the alternative optical smoke detector, the ionization smoke detector is cheaper and can detect particles which are too small to produce significant light scattering; however, it is more prone to false alarms. Radionuclide As 241Am has a roughly similar half-life to 238Pu (432.2 years vs. 87 years), it has been proposed as an active element of radioisotope thermoelectric generators, for example in spacecraft. Although americium produces less heat and electricity – the power yield is 114.7 mW/g for 241Am and 6.31 mW/g for 243Am (cf. 390 mW/g for 238Pu) – and its radiation poses more threat to humans owing to neutron emission, the European Space Agency is considering using americium for its space probes. Another proposed space-related application of americium is a fuel for space ships with nuclear propulsion. It relies on the very high rate of nuclear fission of 242mAm, which can be maintained even in a micrometer-thick foil. Small thickness avoids the problem of self-absorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alpha-particles. The fission products of 242mAm can either directly propel the spaceship or they can heat a thrusting gas.

Americium-241 decays to 237Np emitting alpha particles of 5 different energies, mostly at 5.486 MeV (85.2%) and 5.443 MeV (12.8%). Because many of the resulting states are metastable, they also emit gamma rays with the discrete energies between 26.3 and 158.5 keV. Americium-242 is a short-lived isotope with a half-life of 16.02 h. It mostly (82.7%) converts by β-decay to 242Cm, but also by electron capture to 242Pu (17.3%). Both 242Cm and 242Pu transform via nearly the same decay chain through 238Pu down to 234U. Nearly all (99.541%) of 242m1Am decays by internal conversion to 242Am and the remaining 0.459% by α-decay to 238Np. The latter subsequently decays to 238Pu and then to 234U. Americium-243 transforms by α-emission into 239Np, which converts by β-decay to 239Pu, and the 239Pu changes into 235U by emitting an α-particle. Applications Ionization-type smoke detector Americium is used in the most common type of household smoke detector, which uses 241Am in the form of americium dioxide as its source of ionizing radiation. This isotope is preferred over 226Ra because it emits 5 times more alpha particles and relatively little harmful gamma radiation. The amount of americium in a typical new smoke detector is 1 microcurie (37 kBq) or 0.29 microgram. This amount declines slowly as the americium decays into neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life of 432.2 years, the americium in a smoke detector includes about 3% neptunium after 19 years, and about 5% after 32 years. The radiation passes through an ionization chamber, an air-filled space between two electrodes, and permits a small, constant current between the electrodes. Any smoke that enters the chamber absorbs the alpha particles, which reduces the ionization and affects this current, triggering the alarm. Compared to the alternative optical smoke detector, the ionization smoke detector is cheaper and can detect particles which are too small to produce significant light scattering; however, it is more prone to false alarms. Radionuclide As 241Am has a roughly similar half-life to 238Pu (432.2 years vs. 87 years), it has been proposed as an active element of radioisotope thermoelectric generators, for example in spacecraft. Although americium produces less heat and electricity – the power yield is 114.7 mW/g for 241Am and 6.31 mW/g for 243Am (cf. 390 mW/g for 238Pu) – and its radiation poses more threat to humans owing to neutron emission, the European Space Agency is considering using americium for its space probes. Another proposed space-related application of americium is a fuel for space ships with nuclear propulsion. It relies on the very high rate of nuclear fission of 242mAm, which can be maintained even in a micrometer-thick foil. Small thickness avoids the problem of self-absorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alpha-particles. The fission products of 242mAm can either directly propel the spaceship or they can heat a thrusting gas.

Americium-241 decays to 237Np emitting alpha particles of 5 different energies, mostly at 5.486 MeV (85.2%) and 5.443 MeV (12.8%). Because many of the resulting states are metastable, they also emit gamma rays with the discrete energies between 26.3 and 158.5 keV. Americium-242 is a short-lived isotope with a half-life of 16.02 h. It mostly (82.7%) converts by β-decay to 242Cm, but also by electron capture to 242Pu (17.3%). Both 242Cm and 242Pu transform via nearly the same decay chain through 238Pu down to 234U. Nearly all (99.541%) of 242m1Am decays by internal conversion to 242Am and the remaining 0.459% by α-decay to 238Np. The latter subsequently decays to 238Pu and then to 234U. Americium-243 transforms by α-emission into 239Np, which converts by β-decay to 239Pu, and the 239Pu changes into 235U by emitting an α-particle. Applications Ionization-type smoke detector Americium is used in the most common type of household smoke detector, which uses 241Am in the form of americium dioxide as its source of ionizing radiation. This isotope is preferred over 226Ra because it emits 5 times more alpha particles and relatively little harmful gamma radiation. The amount of americium in a typical new smoke detector is 1 microcurie (37 kBq) or 0.29 microgram. This amount declines slowly as the americium decays into neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life of 432.2 years, the americium in a smoke detector includes about 3% neptunium after 19 years, and about 5% after 32 years. The radiation passes through an ionization chamber, an air-filled space between two electrodes, and permits a small, constant current between the electrodes. Any smoke that enters the chamber absorbs the alpha particles, which reduces the ionization and affects this current, triggering the alarm. Compared to the alternative optical smoke detector, the ionization smoke detector is cheaper and can detect particles which are too small to produce significant light scattering; however, it is more prone to false alarms. Radionuclide As 241Am has a roughly similar half-life to 238Pu (432.2 years vs. 87 years), it has been proposed as an active element of radioisotope thermoelectric generators, for example in spacecraft. Although americium produces less heat and electricity – the power yield is 114.7 mW/g for 241Am and 6.31 mW/g for 243Am (cf. 390 mW/g for 238Pu) – and its radiation poses more threat to humans owing to neutron emission, the European Space Agency is considering using americium for its space probes. Another proposed space-related application of americium is a fuel for space ships with nuclear propulsion. It relies on the very high rate of nuclear fission of 242mAm, which can be maintained even in a micrometer-thick foil. Small thickness avoids the problem of self-absorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alpha-particles. The fission products of 242mAm can either directly propel the spaceship or they can heat a thrusting gas.

They can also transfer their energy to a fluid and generate electricity through a magnetohydrodynamic generator. One more proposal which utilizes the high nuclear fission rate of 242mAm is a nuclear battery. Its design relies not on the energy of the emitted by americium alpha particles, but on their charge, that is the americium acts as the self-sustaining "cathode". A single 3.2 kg 242mAm charge of such battery could provide about 140 kW of power over a period of 80 days. Even with all the potential benefits, the current applications of 242mAm are as yet hindered by the scarcity and high price of this particular nuclear isomer. In 2019, researchers at the UK National Nuclear Laboratory and the University of Leicester demonstrated the use of heat generated by americium to illuminate a small light bulb. This technology could lead to systems to power missions with durations up to 400 years into interstellar space, where solar panels do not function. Neutron source The oxide of 241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{241}_{95}Am -> ^{237}_{93}Np + ^{4}_{2}He + \gamma ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The most widespread use of 241AmBe neutron sources is a neutron probe – a device used to measure the quantity of water present in soil, as well as moisture/density for quality control in highway construction. 241Am neutron sources are also used in well logging applications, as well as in neutron radiography, tomography and other radiochemical investigations. Production of other elements Americium is a starting material for the production of other transuranic elements and transactinides – for example, 82.7% of 242Am decays to 242Cm and 17.3% to 242Pu. In the nuclear reactor, 242Am is also up-converted by neutron capture to 243Am and 244Am, which transforms by β-decay to 244Cm: ^{243}_{95}Am ->[\ce{(n,\gamma)}] ^{244}_{95}Am ->[\beta^-][10.1 \ \ce{h}] ^{244}_{96}Cm Irradiation of 241Am by 12C or 22Ne ions yields the isotopes 247Es (einsteinium) or 260Db (dubnium), respectively. Furthermore, the element berkelium (243Bk isotope) had been first intentionally produced and identified by bombarding 241Am with alpha particles, in 1949, by the same Berkeley group, using the same 60-inch cyclotron. Similarly, nobelium was produced at the Joint Institute for Nuclear Research, Dubna, Russia, in 1965 in several reactions, one of which included irradiation of 243Am with 15N ions. Besides, one of the synthesis reactions for lawrencium, discovered by scientists at Berkeley and Dubna, included bombardment of 243Am with 18O. Spectrometer Americium-241 has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. The 59.5409 keV gamma ray emissions from 241Am in such sources can be used for indirect analysis of materials in radiography and X-ray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers.

They can also transfer their energy to a fluid and generate electricity through a magnetohydrodynamic generator. One more proposal which utilizes the high nuclear fission rate of 242mAm is a nuclear battery. Its design relies not on the energy of the emitted by americium alpha particles, but on their charge, that is the americium acts as the self-sustaining "cathode". A single 3.2 kg 242mAm charge of such battery could provide about 140 kW of power over a period of 80 days. Even with all the potential benefits, the current applications of 242mAm are as yet hindered by the scarcity and high price of this particular nuclear isomer. In 2019, researchers at the UK National Nuclear Laboratory and the University of Leicester demonstrated the use of heat generated by americium to illuminate a small light bulb. This technology could lead to systems to power missions with durations up to 400 years into interstellar space, where solar panels do not function. Neutron source The oxide of 241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{241}_{95}Am -> ^{237}_{93}Np + ^{4}_{2}He + \gamma ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The most widespread use of 241AmBe neutron sources is a neutron probe – a device used to measure the quantity of water present in soil, as well as moisture/density for quality control in highway construction. 241Am neutron sources are also used in well logging applications, as well as in neutron radiography, tomography and other radiochemical investigations. Production of other elements Americium is a starting material for the production of other transuranic elements and transactinides – for example, 82.7% of 242Am decays to 242Cm and 17.3% to 242Pu. In the nuclear reactor, 242Am is also up-converted by neutron capture to 243Am and 244Am, which transforms by β-decay to 244Cm: ^{243}_{95}Am ->[\ce{(n,\gamma)}] ^{244}_{95}Am ->[\beta^-][10.1 \ \ce{h}] ^{244}_{96}Cm Irradiation of 241Am by 12C or 22Ne ions yields the isotopes 247Es (einsteinium) or 260Db (dubnium), respectively. Furthermore, the element berkelium (243Bk isotope) had been first intentionally produced and identified by bombarding 241Am with alpha particles, in 1949, by the same Berkeley group, using the same 60-inch cyclotron. Similarly, nobelium was produced at the Joint Institute for Nuclear Research, Dubna, Russia, in 1965 in several reactions, one of which included irradiation of 243Am with 15N ions. Besides, one of the synthesis reactions for lawrencium, discovered by scientists at Berkeley and Dubna, included bombardment of 243Am with 18O. Spectrometer Americium-241 has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. The 59.5409 keV gamma ray emissions from 241Am in such sources can be used for indirect analysis of materials in radiography and X-ray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers.

They can also transfer their energy to a fluid and generate electricity through a magnetohydrodynamic generator. One more proposal which utilizes the high nuclear fission rate of 242mAm is a nuclear battery. Its design relies not on the energy of the emitted by americium alpha particles, but on their charge, that is the americium acts as the self-sustaining "cathode". A single 3.2 kg 242mAm charge of such battery could provide about 140 kW of power over a period of 80 days. Even with all the potential benefits, the current applications of 242mAm are as yet hindered by the scarcity and high price of this particular nuclear isomer. In 2019, researchers at the UK National Nuclear Laboratory and the University of Leicester demonstrated the use of heat generated by americium to illuminate a small light bulb. This technology could lead to systems to power missions with durations up to 400 years into interstellar space, where solar panels do not function. Neutron source The oxide of 241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large cross-section for the (α,n) nuclear reaction: ^{241}_{95}Am -> ^{237}_{93}Np + ^{4}_{2}He + \gamma ^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma The most widespread use of 241AmBe neutron sources is a neutron probe – a device used to measure the quantity of water present in soil, as well as moisture/density for quality control in highway construction. 241Am neutron sources are also used in well logging applications, as well as in neutron radiography, tomography and other radiochemical investigations. Production of other elements Americium is a starting material for the production of other transuranic elements and transactinides – for example, 82.7% of 242Am decays to 242Cm and 17.3% to 242Pu. In the nuclear reactor, 242Am is also up-converted by neutron capture to 243Am and 244Am, which transforms by β-decay to 244Cm: ^{243}_{95}Am ->[\ce{(n,\gamma)}] ^{244}_{95}Am ->[\beta^-][10.1 \ \ce{h}] ^{244}_{96}Cm Irradiation of 241Am by 12C or 22Ne ions yields the isotopes 247Es (einsteinium) or 260Db (dubnium), respectively. Furthermore, the element berkelium (243Bk isotope) had been first intentionally produced and identified by bombarding 241Am with alpha particles, in 1949, by the same Berkeley group, using the same 60-inch cyclotron. Similarly, nobelium was produced at the Joint Institute for Nuclear Research, Dubna, Russia, in 1965 in several reactions, one of which included irradiation of 243Am with 15N ions. Besides, one of the synthesis reactions for lawrencium, discovered by scientists at Berkeley and Dubna, included bombardment of 243Am with 18O. Spectrometer Americium-241 has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. The 59.5409 keV gamma ray emissions from 241Am in such sources can be used for indirect analysis of materials in radiography and X-ray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers.

For example, the element has been employed to gauge glass thickness to help create flat glass. Americium-241 is also suitable for calibration of gamma-ray spectrometers in the low-energy range, since its spectrum consists of nearly a single peak and negligible Compton continuum (at least three orders of magnitude lower intensity). Americium-241 gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is however obsolete. Health concerns As a highly radioactive element, americium and its compounds must be handled only in an appropriate laboratory under special arrangements. Although most americium isotopes predominantly emit alpha particles which can be blocked by thin layers of common materials, many of the daughter products emit gamma-rays and neutrons which have a long penetration depth. If consumed, most of the americium is excreted within a few days, with only 0.05% absorbed in the blood, of which roughly 45% goes to the liver and 45% to the bones, and the remaining 10% is excreted. The uptake to the liver depends on the individual and increases with age. In the bones, americium is first deposited over cortical and trabecular surfaces and slowly redistributes over the bone with time. The biological half-life of 241Am is 50 years in the bones and 20 years in the liver, whereas in the gonads (testicles and ovaries) it remains permanently; in all these organs, americium promotes formation of cancer cells as a result of its radioactivity. Americium often enters landfills from discarded smoke detectors. The rules associated with the disposal of smoke detectors are relaxed in most jurisdictions. In 1994, 17-year-old David Hahn extracted the americium from about 100 smoke detectors in an attempt to build a breeder nuclear reactor. There have been a few cases of exposure to americium, the worst case being that of chemical operations technician Harold McCluskey, who at the age of 64 was exposed to 500 times the occupational standard for americium-241 as a result of an explosion in his lab. McCluskey died at the age of 75 of unrelated pre-existing disease. See also Actinides in the environment :Category:Americium compounds Notes References Bibliography Penneman, R. A. and Keenan T. K. The radiochemistry of americium and curium, University of California, Los Alamos, California, 1960 Further reading Nuclides and Isotopes – 14th Edition, GE Nuclear Energy, 1989. External links Americium at The Periodic Table of Videos (University of Nottingham) ATSDR – Public Health Statement: Americium World Nuclear Association – Smoke Detectors and Americium Chemical elements Actinides Carcinogens Synthetic elements

For example, the element has been employed to gauge glass thickness to help create flat glass. Americium-241 is also suitable for calibration of gamma-ray spectrometers in the low-energy range, since its spectrum consists of nearly a single peak and negligible Compton continuum (at least three orders of magnitude lower intensity). Americium-241 gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is however obsolete. Health concerns As a highly radioactive element, americium and its compounds must be handled only in an appropriate laboratory under special arrangements. Although most americium isotopes predominantly emit alpha particles which can be blocked by thin layers of common materials, many of the daughter products emit gamma-rays and neutrons which have a long penetration depth. If consumed, most of the americium is excreted within a few days, with only 0.05% absorbed in the blood, of which roughly 45% goes to the liver and 45% to the bones, and the remaining 10% is excreted. The uptake to the liver depends on the individual and increases with age. In the bones, americium is first deposited over cortical and trabecular surfaces and slowly redistributes over the bone with time. The biological half-life of 241Am is 50 years in the bones and 20 years in the liver, whereas in the gonads (testicles and ovaries) it remains permanently; in all these organs, americium promotes formation of cancer cells as a result of its radioactivity. Americium often enters landfills from discarded smoke detectors. The rules associated with the disposal of smoke detectors are relaxed in most jurisdictions. In 1994, 17-year-old David Hahn extracted the americium from about 100 smoke detectors in an attempt to build a breeder nuclear reactor. There have been a few cases of exposure to americium, the worst case being that of chemical operations technician Harold McCluskey, who at the age of 64 was exposed to 500 times the occupational standard for americium-241 as a result of an explosion in his lab. McCluskey died at the age of 75 of unrelated pre-existing disease. See also Actinides in the environment :Category:Americium compounds Notes References Bibliography Penneman, R. A. and Keenan T. K. The radiochemistry of americium and curium, University of California, Los Alamos, California, 1960 Further reading Nuclides and Isotopes – 14th Edition, GE Nuclear Energy, 1989. External links Americium at The Periodic Table of Videos (University of Nottingham) ATSDR – Public Health Statement: Americium World Nuclear Association – Smoke Detectors and Americium Chemical elements Actinides Carcinogens Synthetic elements

For example, the element has been employed to gauge glass thickness to help create flat glass. Americium-241 is also suitable for calibration of gamma-ray spectrometers in the low-energy range, since its spectrum consists of nearly a single peak and negligible Compton continuum (at least three orders of magnitude lower intensity). Americium-241 gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is however obsolete. Health concerns As a highly radioactive element, americium and its compounds must be handled only in an appropriate laboratory under special arrangements. Although most americium isotopes predominantly emit alpha particles which can be blocked by thin layers of common materials, many of the daughter products emit gamma-rays and neutrons which have a long penetration depth. If consumed, most of the americium is excreted within a few days, with only 0.05% absorbed in the blood, of which roughly 45% goes to the liver and 45% to the bones, and the remaining 10% is excreted. The uptake to the liver depends on the individual and increases with age. In the bones, americium is first deposited over cortical and trabecular surfaces and slowly redistributes over the bone with time. The biological half-life of 241Am is 50 years in the bones and 20 years in the liver, whereas in the gonads (testicles and ovaries) it remains permanently; in all these organs, americium promotes formation of cancer cells as a result of its radioactivity. Americium often enters landfills from discarded smoke detectors. The rules associated with the disposal of smoke detectors are relaxed in most jurisdictions. In 1994, 17-year-old David Hahn extracted the americium from about 100 smoke detectors in an attempt to build a breeder nuclear reactor. There have been a few cases of exposure to americium, the worst case being that of chemical operations technician Harold McCluskey, who at the age of 64 was exposed to 500 times the occupational standard for americium-241 as a result of an explosion in his lab. McCluskey died at the age of 75 of unrelated pre-existing disease. See also Actinides in the environment :Category:Americium compounds Notes References Bibliography Penneman, R. A. and Keenan T. K. The radiochemistry of americium and curium, University of California, Los Alamos, California, 1960 Further reading Nuclides and Isotopes – 14th Edition, GE Nuclear Energy, 1989. External links Americium at The Periodic Table of Videos (University of Nottingham) ATSDR – Public Health Statement: Americium World Nuclear Association – Smoke Detectors and Americium Chemical elements Actinides Carcinogens Synthetic elements

Astatine Astatine is a chemical element with the symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity. The bulk properties of astatine are not known with certainty. Many of them have been estimated based on the element's position on the periodic table as a heavier analog of iodine, and a member of the halogens (the group of elements including fluorine, chlorine, bromine, and iodine). However, astatine also falls roughly along the dividing line between metals and nonmetals, and some metallic behavior has also been observed and predicted for it. Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal. Chemically, several anionic species of astatine are known and most of its compounds resemble those of iodine, but it also sometimes displays metallic characteristics and shows some similarities to silver. The first synthesis of the element was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley, who named it from the Ancient Greek () 'unstable'. Four isotopes of astatine were subsequently found to be naturally occurring, although much less than one gram is present at any given time in the Earth's crust. Neither the most stable isotope astatine-210, nor the medically useful astatine-211, occur naturally; they can only be produced synthetically, usually by bombarding bismuth-209 with alpha particles. Characteristics Astatine is an extremely radioactive element; all its isotopes have half-lives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium, or radon. Most of its isotopes are very unstable, with half-lives of one second or less. Of the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than francium are in any case synthetic and do not occur in nature. The bulk properties of astatine are not known with any certainty. Research is limited by its short half-life, which prevents the creation of weighable quantities. A visible piece of astatine would immediately vaporize itself because of the heat generated by its intense radioactivity. It remains to be seen if, with sufficient cooling, a macroscopic quantity of astatine could be deposited as a thin film. Astatine is usually classified as either a nonmetal or a metalloid; metal formation has also been predicted. Physical Most of the physical properties of astatine have been estimated (by interpolation or extrapolation), using theoretically or empirically derived methods. For example, halogens get darker with increasing atomic weight – fluorine is nearly colorless, chlorine is yellow green, bromine is red brown, and iodine is dark gray/violet.

Astatine Astatine is a chemical element with the symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity. The bulk properties of astatine are not known with certainty. Many of them have been estimated based on the element's position on the periodic table as a heavier analog of iodine, and a member of the halogens (the group of elements including fluorine, chlorine, bromine, and iodine). However, astatine also falls roughly along the dividing line between metals and nonmetals, and some metallic behavior has also been observed and predicted for it. Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal. Chemically, several anionic species of astatine are known and most of its compounds resemble those of iodine, but it also sometimes displays metallic characteristics and shows some similarities to silver. The first synthesis of the element was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley, who named it from the Ancient Greek () 'unstable'. Four isotopes of astatine were subsequently found to be naturally occurring, although much less than one gram is present at any given time in the Earth's crust. Neither the most stable isotope astatine-210, nor the medically useful astatine-211, occur naturally; they can only be produced synthetically, usually by bombarding bismuth-209 with alpha particles. Characteristics Astatine is an extremely radioactive element; all its isotopes have half-lives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium, or radon. Most of its isotopes are very unstable, with half-lives of one second or less. Of the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than francium are in any case synthetic and do not occur in nature. The bulk properties of astatine are not known with any certainty. Research is limited by its short half-life, which prevents the creation of weighable quantities. A visible piece of astatine would immediately vaporize itself because of the heat generated by its intense radioactivity. It remains to be seen if, with sufficient cooling, a macroscopic quantity of astatine could be deposited as a thin film. Astatine is usually classified as either a nonmetal or a metalloid; metal formation has also been predicted. Physical Most of the physical properties of astatine have been estimated (by interpolation or extrapolation), using theoretically or empirically derived methods. For example, halogens get darker with increasing atomic weight – fluorine is nearly colorless, chlorine is yellow green, bromine is red brown, and iodine is dark gray/violet.

Astatine is sometimes described as probably being a black solid (assuming it follows this trend), or as having a metallic appearance (if it is a metalloid or a metal). Astatine sublimes less readily than does iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions. The structure of solid astatine is unknown. As an analogue of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV). Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic face-centered cubic structure; in this structure it may well be a superconductor, like the similar high-pressure phase of iodine. Evidence for (or against) the existence of diatomic astatine (At2) is sparse and inconclusive. Some sources state that it does not exist, or at least has never been observed, while other sources assert or imply its existence. Despite this controversy, many properties of diatomic astatine have been predicted; for example, its bond length would be , dissociation energy , and heat of vaporization (∆Hvap) 54.39 kJ/mol. Many values have been predicted for the melting and boiling points of astatine, but only for At2. Chemical The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity by-products, and other unwanted nano-scale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, such as plating onto a cathode, and coprecipitating with metal sulfides in hydrochloric acid. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects astatine in the +1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine. It has been suggested that astatine can form a stable monatomic cation in aqueous solution, but electromigration evidence suggests that the cationic At(I) species is protonated hypoastatous acid (H2OAt+), showing analogy to iodine. Astatine has an electronegativity of 2.2 on the revised Pauling scale – lower than that of iodine (2.66) and the same as hydrogen. In hydrogen astatide (HAt), the negative charge is predicted to be on the hydrogen atom, implying that this compound could be referred to as astatine hydride according to certain nomenclatures. That would be consistent with the electronegativity of astatine on the Allred–Rochow scale (1.9) being less than that of hydrogen (2.2).

Astatine is sometimes described as probably being a black solid (assuming it follows this trend), or as having a metallic appearance (if it is a metalloid or a metal). Astatine sublimes less readily than does iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions. The structure of solid astatine is unknown. As an analogue of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV). Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic face-centered cubic structure; in this structure it may well be a superconductor, like the similar high-pressure phase of iodine. Evidence for (or against) the existence of diatomic astatine (At2) is sparse and inconclusive. Some sources state that it does not exist, or at least has never been observed, while other sources assert or imply its existence. Despite this controversy, many properties of diatomic astatine have been predicted; for example, its bond length would be , dissociation energy , and heat of vaporization (∆Hvap) 54.39 kJ/mol. Many values have been predicted for the melting and boiling points of astatine, but only for At2. Chemical The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity by-products, and other unwanted nano-scale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, such as plating onto a cathode, and coprecipitating with metal sulfides in hydrochloric acid. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects astatine in the +1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine. It has been suggested that astatine can form a stable monatomic cation in aqueous solution, but electromigration evidence suggests that the cationic At(I) species is protonated hypoastatous acid (H2OAt+), showing analogy to iodine. Astatine has an electronegativity of 2.2 on the revised Pauling scale – lower than that of iodine (2.66) and the same as hydrogen. In hydrogen astatide (HAt), the negative charge is predicted to be on the hydrogen atom, implying that this compound could be referred to as astatine hydride according to certain nomenclatures. That would be consistent with the electronegativity of astatine on the Allred–Rochow scale (1.9) being less than that of hydrogen (2.2).

Astatine is sometimes described as probably being a black solid (assuming it follows this trend), or as having a metallic appearance (if it is a metalloid or a metal). Astatine sublimes less readily than does iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions. The structure of solid astatine is unknown. As an analogue of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV). Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic face-centered cubic structure; in this structure it may well be a superconductor, like the similar high-pressure phase of iodine. Evidence for (or against) the existence of diatomic astatine (At2) is sparse and inconclusive. Some sources state that it does not exist, or at least has never been observed, while other sources assert or imply its existence. Despite this controversy, many properties of diatomic astatine have been predicted; for example, its bond length would be , dissociation energy , and heat of vaporization (∆Hvap) 54.39 kJ/mol. Many values have been predicted for the melting and boiling points of astatine, but only for At2. Chemical The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity by-products, and other unwanted nano-scale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, such as plating onto a cathode, and coprecipitating with metal sulfides in hydrochloric acid. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects astatine in the +1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine. It has been suggested that astatine can form a stable monatomic cation in aqueous solution, but electromigration evidence suggests that the cationic At(I) species is protonated hypoastatous acid (H2OAt+), showing analogy to iodine. Astatine has an electronegativity of 2.2 on the revised Pauling scale – lower than that of iodine (2.66) and the same as hydrogen. In hydrogen astatide (HAt), the negative charge is predicted to be on the hydrogen atom, implying that this compound could be referred to as astatine hydride according to certain nomenclatures. That would be consistent with the electronegativity of astatine on the Allred–Rochow scale (1.9) being less than that of hydrogen (2.2).

However, official IUPAC stoichiometric nomenclature is based on an idealized convention of determining the relative electronegativities of the elements by the mere virtue of their position within the periodic table. According to this convention, astatine is handled as though it is more electronegative than hydrogen, irrespective of its true electronegativity. The electron affinity of astatine, at 233 kJ mol−1, is 21% less than that of iodine. In comparison, the value of Cl (349) is 6.4% higher than F (328); Br (325) is 6.9% less than Cl; and I (295) is 9.2% less than Br. The marked reduction for At was predicted as being due to spin–orbit interactions. The first ionisation energy of astatine is about 899 kJ mol−1, which continues the trend of decreasing first ionisation energies down the halogen group (fluorine, 1681; chlorine, 1251; bromine, 1140; iodine, 1008). Compounds Less reactive than iodine, astatine is the least reactive of the halogens. Its compounds have been synthesized in microscopic amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques (such as filtration and precipitation) to work. Like iodine, astatine has been shown to adopt odd-numbered oxidation states ranging from −1 to +7. Only a few compounds with metals have been reported, in the form of astatides of sodium, palladium, silver, thallium, and lead. Some characteristic properties of silver and sodium astatide, and the other hypothetical alkali and alkaline earth astatides, have been estimated by extrapolation from other metal halides. The formation of an astatine compound with hydrogen – usually referred to as hydrogen astatide – was noted by the pioneers of astatine chemistry. As mentioned, there are grounds for instead referring to this compound as astatine hydride. It is easily oxidized; acidification by dilute nitric acid gives the At0 or At+ forms, and the subsequent addition of silver(I) may only partially, at best, precipitate astatine as silver(I) astatide (AgAt). Iodine, in contrast, is not oxidized, and precipitates readily as silver(I) iodide. Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C bonds. Astatine can replace a hydrogen atom in benzene to form astatobenzene C6H5At; this may be oxidized to C6H5AtCl2 by chlorine. By treating this compound with an alkaline solution of hypochlorite, C6H5AtO2 can be produced. The dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate (a non-coordinating anion) and with nitrate, [At(C5H5N)2]NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatine(I) centre with each of the pyridine rings via their nitrogen atoms.

However, official IUPAC stoichiometric nomenclature is based on an idealized convention of determining the relative electronegativities of the elements by the mere virtue of their position within the periodic table. According to this convention, astatine is handled as though it is more electronegative than hydrogen, irrespective of its true electronegativity. The electron affinity of astatine, at 233 kJ mol−1, is 21% less than that of iodine. In comparison, the value of Cl (349) is 6.4% higher than F (328); Br (325) is 6.9% less than Cl; and I (295) is 9.2% less than Br. The marked reduction for At was predicted as being due to spin–orbit interactions. The first ionisation energy of astatine is about 899 kJ mol−1, which continues the trend of decreasing first ionisation energies down the halogen group (fluorine, 1681; chlorine, 1251; bromine, 1140; iodine, 1008). Compounds Less reactive than iodine, astatine is the least reactive of the halogens. Its compounds have been synthesized in microscopic amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques (such as filtration and precipitation) to work. Like iodine, astatine has been shown to adopt odd-numbered oxidation states ranging from −1 to +7. Only a few compounds with metals have been reported, in the form of astatides of sodium, palladium, silver, thallium, and lead. Some characteristic properties of silver and sodium astatide, and the other hypothetical alkali and alkaline earth astatides, have been estimated by extrapolation from other metal halides. The formation of an astatine compound with hydrogen – usually referred to as hydrogen astatide – was noted by the pioneers of astatine chemistry. As mentioned, there are grounds for instead referring to this compound as astatine hydride. It is easily oxidized; acidification by dilute nitric acid gives the At0 or At+ forms, and the subsequent addition of silver(I) may only partially, at best, precipitate astatine as silver(I) astatide (AgAt). Iodine, in contrast, is not oxidized, and precipitates readily as silver(I) iodide. Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C bonds. Astatine can replace a hydrogen atom in benzene to form astatobenzene C6H5At; this may be oxidized to C6H5AtCl2 by chlorine. By treating this compound with an alkaline solution of hypochlorite, C6H5AtO2 can be produced. The dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate (a non-coordinating anion) and with nitrate, [At(C5H5N)2]NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatine(I) centre with each of the pyridine rings via their nitrogen atoms.

However, official IUPAC stoichiometric nomenclature is based on an idealized convention of determining the relative electronegativities of the elements by the mere virtue of their position within the periodic table. According to this convention, astatine is handled as though it is more electronegative than hydrogen, irrespective of its true electronegativity. The electron affinity of astatine, at 233 kJ mol−1, is 21% less than that of iodine. In comparison, the value of Cl (349) is 6.4% higher than F (328); Br (325) is 6.9% less than Cl; and I (295) is 9.2% less than Br. The marked reduction for At was predicted as being due to spin–orbit interactions. The first ionisation energy of astatine is about 899 kJ mol−1, which continues the trend of decreasing first ionisation energies down the halogen group (fluorine, 1681; chlorine, 1251; bromine, 1140; iodine, 1008). Compounds Less reactive than iodine, astatine is the least reactive of the halogens. Its compounds have been synthesized in microscopic amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques (such as filtration and precipitation) to work. Like iodine, astatine has been shown to adopt odd-numbered oxidation states ranging from −1 to +7. Only a few compounds with metals have been reported, in the form of astatides of sodium, palladium, silver, thallium, and lead. Some characteristic properties of silver and sodium astatide, and the other hypothetical alkali and alkaline earth astatides, have been estimated by extrapolation from other metal halides. The formation of an astatine compound with hydrogen – usually referred to as hydrogen astatide – was noted by the pioneers of astatine chemistry. As mentioned, there are grounds for instead referring to this compound as astatine hydride. It is easily oxidized; acidification by dilute nitric acid gives the At0 or At+ forms, and the subsequent addition of silver(I) may only partially, at best, precipitate astatine as silver(I) astatide (AgAt). Iodine, in contrast, is not oxidized, and precipitates readily as silver(I) iodide. Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C bonds. Astatine can replace a hydrogen atom in benzene to form astatobenzene C6H5At; this may be oxidized to C6H5AtCl2 by chlorine. By treating this compound with an alkaline solution of hypochlorite, C6H5AtO2 can be produced. The dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate (a non-coordinating anion) and with nitrate, [At(C5H5N)2]NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatine(I) centre with each of the pyridine rings via their nitrogen atoms.

With oxygen, there is evidence of the species AtO− and AtO+ in aqueous solution, formed by the reaction of astatine with an oxidant such as elemental bromine or (in the last case) by sodium persulfate in a solution of perchloric acid: the latter species might also be protonated astatous acid, . The species previously thought to be has since been determined to be , a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH). The well characterized anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of , such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion ; this is only stable in neutral or alkaline solutions. Astatine is also thought to be capable of forming cations in salts with oxyanions such as iodate or dichromate; this is based on the observation that, in acidic solutions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silver(I) iodate or thallium(I) dichromate. Astatine may form bonds to the other chalcogens; these include S7At+ and with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium colloid with tellurium. Astatine is known to react with its lighter homologs iodine, bromine, and chlorine in the vapor state; these reactions produce diatomic interhalogen compounds with formulas AtI, AtBr, and AtCl. The first two compounds may also be produced in water – astatine reacts with iodine/iodide solution to form AtI, whereas AtBr requires (aside from astatine) an iodine/iodine monobromide/bromide solution. The excess of iodides or bromides may lead to and ions, or in a chloride solution, they may produce species like or via equilibrium reactions with the chlorides. Oxidation of the element with dichromate (in nitric acid solution) showed that adding chloride turned the astatine into a molecule likely to be either AtCl or AtOCl. Similarly, or may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed to have been precipitated. In a plasma ion source mass spectrometer, the ions [AtI]+, [AtBr]+, and [AtCl]+ have been formed by introducing lighter halogen vapors into a helium-filled cell containing astatine, supporting the existence of stable neutral molecules in the plasma ion state. No astatine fluorides have been discovered yet. Their absence has been speculatively attributed to the extreme reactivity of such compounds, including the reaction of an initially formed fluoride with the walls of the glass container to form a non-volatile product. Thus, although the synthesis of an astatine fluoride is thought to be possible, it may require a liquid halogen fluoride solvent, as has already been used for the characterization of radon fluoride.

With oxygen, there is evidence of the species AtO− and AtO+ in aqueous solution, formed by the reaction of astatine with an oxidant such as elemental bromine or (in the last case) by sodium persulfate in a solution of perchloric acid: the latter species might also be protonated astatous acid, . The species previously thought to be has since been determined to be , a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH). The well characterized anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of , such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion ; this is only stable in neutral or alkaline solutions. Astatine is also thought to be capable of forming cations in salts with oxyanions such as iodate or dichromate; this is based on the observation that, in acidic solutions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silver(I) iodate or thallium(I) dichromate. Astatine may form bonds to the other chalcogens; these include S7At+ and with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium colloid with tellurium. Astatine is known to react with its lighter homologs iodine, bromine, and chlorine in the vapor state; these reactions produce diatomic interhalogen compounds with formulas AtI, AtBr, and AtCl. The first two compounds may also be produced in water – astatine reacts with iodine/iodide solution to form AtI, whereas AtBr requires (aside from astatine) an iodine/iodine monobromide/bromide solution. The excess of iodides or bromides may lead to and ions, or in a chloride solution, they may produce species like or via equilibrium reactions with the chlorides. Oxidation of the element with dichromate (in nitric acid solution) showed that adding chloride turned the astatine into a molecule likely to be either AtCl or AtOCl. Similarly, or may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed to have been precipitated. In a plasma ion source mass spectrometer, the ions [AtI]+, [AtBr]+, and [AtCl]+ have been formed by introducing lighter halogen vapors into a helium-filled cell containing astatine, supporting the existence of stable neutral molecules in the plasma ion state. No astatine fluorides have been discovered yet. Their absence has been speculatively attributed to the extreme reactivity of such compounds, including the reaction of an initially formed fluoride with the walls of the glass container to form a non-volatile product. Thus, although the synthesis of an astatine fluoride is thought to be possible, it may require a liquid halogen fluoride solvent, as has already been used for the characterization of radon fluoride.

With oxygen, there is evidence of the species AtO− and AtO+ in aqueous solution, formed by the reaction of astatine with an oxidant such as elemental bromine or (in the last case) by sodium persulfate in a solution of perchloric acid: the latter species might also be protonated astatous acid, . The species previously thought to be has since been determined to be , a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH). The well characterized anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of , such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion ; this is only stable in neutral or alkaline solutions. Astatine is also thought to be capable of forming cations in salts with oxyanions such as iodate or dichromate; this is based on the observation that, in acidic solutions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silver(I) iodate or thallium(I) dichromate. Astatine may form bonds to the other chalcogens; these include S7At+ and with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium colloid with tellurium. Astatine is known to react with its lighter homologs iodine, bromine, and chlorine in the vapor state; these reactions produce diatomic interhalogen compounds with formulas AtI, AtBr, and AtCl. The first two compounds may also be produced in water – astatine reacts with iodine/iodide solution to form AtI, whereas AtBr requires (aside from astatine) an iodine/iodine monobromide/bromide solution. The excess of iodides or bromides may lead to and ions, or in a chloride solution, they may produce species like or via equilibrium reactions with the chlorides. Oxidation of the element with dichromate (in nitric acid solution) showed that adding chloride turned the astatine into a molecule likely to be either AtCl or AtOCl. Similarly, or may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed to have been precipitated. In a plasma ion source mass spectrometer, the ions [AtI]+, [AtBr]+, and [AtCl]+ have been formed by introducing lighter halogen vapors into a helium-filled cell containing astatine, supporting the existence of stable neutral molecules in the plasma ion state. No astatine fluorides have been discovered yet. Their absence has been speculatively attributed to the extreme reactivity of such compounds, including the reaction of an initially formed fluoride with the walls of the glass container to form a non-volatile product. Thus, although the synthesis of an astatine fluoride is thought to be possible, it may require a liquid halogen fluoride solvent, as has already been used for the characterization of radon fluoride.

History In 1869, when Dmitri Mendeleev published his periodic table, the space under iodine was empty; after Niels Bohr established the physical basis of the classification of chemical elements, it was suggested that the fifth halogen belonged there. Before its officially recognized discovery, it was called "eka-iodine" (from Sanskrit eka – "one") to imply it was one space under iodine (in the same manner as eka-silicon, eka-boron, and others). Scientists tried to find it in nature; given its extreme rarity, these attempts resulted in several false discoveries. The first claimed discovery of eka-iodine was made by Fred Allison and his associates at the Alabama Polytechnic Institute (now Auburn University) in 1931. The discoverers named element 85 "alabamine", and assigned it the symbol Ab, designations that were used for a few years. In 1934, H. G. MacPherson of University of California, Berkeley disproved Allison's method and the validity of his discovery. There was another claim in 1937, by the chemist Rajendralal De. Working in Dacca in British India (now Dhaka in Bangladesh), he chose the name "dakin" for element 85, which he claimed to have isolated as the thorium series equivalent of radium F (polonium-210) in the radium series. The properties he reported for dakin do not correspond to those of astatine; moreover, astatine is not found in the thorium series, and the true identity of dakin is not known. In 1936, the team of Romanian physicist Horia Hulubei and French physicist Yvette Cauchois claimed to have discovered element 85 via X-ray analysis. In 1939, they published another paper which supported and extended previous data. In 1944, Hulubei published a summary of data he had obtained up to that time, claiming it was supported by the work of other researchers. He chose the name "dor", presumably from the Romanian for "longing" [for peace], as World War II had started five years earlier. As Hulubei was writing in French, a language which does not accommodate the "ine" suffix, dor would likely have been rendered in English as "dorine", had it been adopted. In 1947, Hulubei's claim was effectively rejected by the Austrian chemist Friedrich Paneth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine, his means to detect it were too weak, by current standards, to enable correct identification. He had also been involved in an earlier false claim as to the discovery of element 87 (francium) and this is thought to have caused other researchers to downplay his work. In 1940, the Swiss chemist Walter Minder announced the discovery of element 85 as the beta decay product of radium A (polonium-218), choosing the name "helvetium" (from , the Latin name of Switzerland). Berta Karlik and Traude Bernert were unsuccessful in reproducing his experiments, and subsequently attributed Minder's results to contamination of his radon stream (radon-222 is the parent isotope of polonium-218).

History In 1869, when Dmitri Mendeleev published his periodic table, the space under iodine was empty; after Niels Bohr established the physical basis of the classification of chemical elements, it was suggested that the fifth halogen belonged there. Before its officially recognized discovery, it was called "eka-iodine" (from Sanskrit eka – "one") to imply it was one space under iodine (in the same manner as eka-silicon, eka-boron, and others). Scientists tried to find it in nature; given its extreme rarity, these attempts resulted in several false discoveries. The first claimed discovery of eka-iodine was made by Fred Allison and his associates at the Alabama Polytechnic Institute (now Auburn University) in 1931. The discoverers named element 85 "alabamine", and assigned it the symbol Ab, designations that were used for a few years. In 1934, H. G. MacPherson of University of California, Berkeley disproved Allison's method and the validity of his discovery. There was another claim in 1937, by the chemist Rajendralal De. Working in Dacca in British India (now Dhaka in Bangladesh), he chose the name "dakin" for element 85, which he claimed to have isolated as the thorium series equivalent of radium F (polonium-210) in the radium series. The properties he reported for dakin do not correspond to those of astatine; moreover, astatine is not found in the thorium series, and the true identity of dakin is not known. In 1936, the team of Romanian physicist Horia Hulubei and French physicist Yvette Cauchois claimed to have discovered element 85 via X-ray analysis. In 1939, they published another paper which supported and extended previous data. In 1944, Hulubei published a summary of data he had obtained up to that time, claiming it was supported by the work of other researchers. He chose the name "dor", presumably from the Romanian for "longing" [for peace], as World War II had started five years earlier. As Hulubei was writing in French, a language which does not accommodate the "ine" suffix, dor would likely have been rendered in English as "dorine", had it been adopted. In 1947, Hulubei's claim was effectively rejected by the Austrian chemist Friedrich Paneth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine, his means to detect it were too weak, by current standards, to enable correct identification. He had also been involved in an earlier false claim as to the discovery of element 87 (francium) and this is thought to have caused other researchers to downplay his work. In 1940, the Swiss chemist Walter Minder announced the discovery of element 85 as the beta decay product of radium A (polonium-218), choosing the name "helvetium" (from , the Latin name of Switzerland). Berta Karlik and Traude Bernert were unsuccessful in reproducing his experiments, and subsequently attributed Minder's results to contamination of his radon stream (radon-222 is the parent isotope of polonium-218).

History In 1869, when Dmitri Mendeleev published his periodic table, the space under iodine was empty; after Niels Bohr established the physical basis of the classification of chemical elements, it was suggested that the fifth halogen belonged there. Before its officially recognized discovery, it was called "eka-iodine" (from Sanskrit eka – "one") to imply it was one space under iodine (in the same manner as eka-silicon, eka-boron, and others). Scientists tried to find it in nature; given its extreme rarity, these attempts resulted in several false discoveries. The first claimed discovery of eka-iodine was made by Fred Allison and his associates at the Alabama Polytechnic Institute (now Auburn University) in 1931. The discoverers named element 85 "alabamine", and assigned it the symbol Ab, designations that were used for a few years. In 1934, H. G. MacPherson of University of California, Berkeley disproved Allison's method and the validity of his discovery. There was another claim in 1937, by the chemist Rajendralal De. Working in Dacca in British India (now Dhaka in Bangladesh), he chose the name "dakin" for element 85, which he claimed to have isolated as the thorium series equivalent of radium F (polonium-210) in the radium series. The properties he reported for dakin do not correspond to those of astatine; moreover, astatine is not found in the thorium series, and the true identity of dakin is not known. In 1936, the team of Romanian physicist Horia Hulubei and French physicist Yvette Cauchois claimed to have discovered element 85 via X-ray analysis. In 1939, they published another paper which supported and extended previous data. In 1944, Hulubei published a summary of data he had obtained up to that time, claiming it was supported by the work of other researchers. He chose the name "dor", presumably from the Romanian for "longing" [for peace], as World War II had started five years earlier. As Hulubei was writing in French, a language which does not accommodate the "ine" suffix, dor would likely have been rendered in English as "dorine", had it been adopted. In 1947, Hulubei's claim was effectively rejected by the Austrian chemist Friedrich Paneth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine, his means to detect it were too weak, by current standards, to enable correct identification. He had also been involved in an earlier false claim as to the discovery of element 87 (francium) and this is thought to have caused other researchers to downplay his work. In 1940, the Swiss chemist Walter Minder announced the discovery of element 85 as the beta decay product of radium A (polonium-218), choosing the name "helvetium" (from , the Latin name of Switzerland). Berta Karlik and Traude Bernert were unsuccessful in reproducing his experiments, and subsequently attributed Minder's results to contamination of his radon stream (radon-222 is the parent isotope of polonium-218).

In 1942, Minder, in collaboration with the English scientist Alice Leigh-Smith, announced the discovery of another isotope of element 85, presumed to be the product of thorium A (polonium-216) beta decay. They named this substance "anglo-helvetium", but Karlik and Bernert were again unable to reproduce these results. Later in 1940, Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè isolated the element at the University of California, Berkeley. Instead of searching for the element in nature, the scientists created it by bombarding bismuth-209 with alpha particles in a cyclotron (particle accelerator) to produce, after emission of two neutrons, astatine-211. The discoverers, however, did not immediately suggest a name for the element. The reason for this was that at the time, an element created synthetically in "invisible quantities" that had not yet been discovered in nature was not seen as a completely valid one; in addition, chemists were reluctant to recognize radioactive isotopes as legitimately as stable ones. In 1943, astatine was found as a product of two naturally occurring decay chains by Berta Karlik and Traude Bernert, first in the so-called uranium series, and then in the actinium series. (Since then, astatine was also found in a third decay chain, the neptunium series.) Friedrich Paneth in 1946 called to finally recognize synthetic elements, quoting, among other reasons, recent confirmation of their natural occurrence, and proposed that the discoverers of the newly discovered unnamed elements name these elements. In early 1947, Nature published the discoverers' suggestions; a letter from Corson, MacKenzie, and Segrè suggested the name "astatine" coming from the Greek astatos (αστατος) meaning "unstable", because of its propensity for radioactive decay, with the ending "-ine", found in the names of the four previously discovered halogens. The name was also chosen to continue the tradition of the four stable halogens, where the name referred to a property of the element. Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry. Subsequent investigators reported iodine-like, cationic, or amphoteric behavior. In a 2003 retrospective, Corson wrote that "some of the properties [of astatine] are similar to iodine … it also exhibits metallic properties, more like its metallic neighbors Po and Bi." Isotopes There are 39 known isotopes of astatine, with atomic masses (mass numbers) of 191–229. Theoretical modeling suggests that 37 more isotopes could exist. No stable or long-lived astatine isotope has been observed, nor is one expected to exist. Astatine's alpha decay energies follow the same trend as for other heavy elements. Lighter astatine isotopes have quite high energies of alpha decay, which become lower as the nuclei become heavier. Astatine-211 has a significantly higher energy than the previous isotope, because it has a nucleus with 126 neutrons, and 126 is a magic number corresponding to a filled neutron shell. Despite having a similar half-life to the previous isotope (8.1 hours for astatine-210 and 7.2 hours for astatine-211), the alpha decay probability is much higher for the latter: 41.81% against only 0.18%.

In 1942, Minder, in collaboration with the English scientist Alice Leigh-Smith, announced the discovery of another isotope of element 85, presumed to be the product of thorium A (polonium-216) beta decay. They named this substance "anglo-helvetium", but Karlik and Bernert were again unable to reproduce these results. Later in 1940, Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè isolated the element at the University of California, Berkeley. Instead of searching for the element in nature, the scientists created it by bombarding bismuth-209 with alpha particles in a cyclotron (particle accelerator) to produce, after emission of two neutrons, astatine-211. The discoverers, however, did not immediately suggest a name for the element. The reason for this was that at the time, an element created synthetically in "invisible quantities" that had not yet been discovered in nature was not seen as a completely valid one; in addition, chemists were reluctant to recognize radioactive isotopes as legitimately as stable ones. In 1943, astatine was found as a product of two naturally occurring decay chains by Berta Karlik and Traude Bernert, first in the so-called uranium series, and then in the actinium series. (Since then, astatine was also found in a third decay chain, the neptunium series.) Friedrich Paneth in 1946 called to finally recognize synthetic elements, quoting, among other reasons, recent confirmation of their natural occurrence, and proposed that the discoverers of the newly discovered unnamed elements name these elements. In early 1947, Nature published the discoverers' suggestions; a letter from Corson, MacKenzie, and Segrè suggested the name "astatine" coming from the Greek astatos (αστατος) meaning "unstable", because of its propensity for radioactive decay, with the ending "-ine", found in the names of the four previously discovered halogens. The name was also chosen to continue the tradition of the four stable halogens, where the name referred to a property of the element. Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry. Subsequent investigators reported iodine-like, cationic, or amphoteric behavior. In a 2003 retrospective, Corson wrote that "some of the properties [of astatine] are similar to iodine … it also exhibits metallic properties, more like its metallic neighbors Po and Bi." Isotopes There are 39 known isotopes of astatine, with atomic masses (mass numbers) of 191–229. Theoretical modeling suggests that 37 more isotopes could exist. No stable or long-lived astatine isotope has been observed, nor is one expected to exist. Astatine's alpha decay energies follow the same trend as for other heavy elements. Lighter astatine isotopes have quite high energies of alpha decay, which become lower as the nuclei become heavier. Astatine-211 has a significantly higher energy than the previous isotope, because it has a nucleus with 126 neutrons, and 126 is a magic number corresponding to a filled neutron shell. Despite having a similar half-life to the previous isotope (8.1 hours for astatine-210 and 7.2 hours for astatine-211), the alpha decay probability is much higher for the latter: 41.81% against only 0.18%.

In 1942, Minder, in collaboration with the English scientist Alice Leigh-Smith, announced the discovery of another isotope of element 85, presumed to be the product of thorium A (polonium-216) beta decay. They named this substance "anglo-helvetium", but Karlik and Bernert were again unable to reproduce these results. Later in 1940, Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè isolated the element at the University of California, Berkeley. Instead of searching for the element in nature, the scientists created it by bombarding bismuth-209 with alpha particles in a cyclotron (particle accelerator) to produce, after emission of two neutrons, astatine-211. The discoverers, however, did not immediately suggest a name for the element. The reason for this was that at the time, an element created synthetically in "invisible quantities" that had not yet been discovered in nature was not seen as a completely valid one; in addition, chemists were reluctant to recognize radioactive isotopes as legitimately as stable ones. In 1943, astatine was found as a product of two naturally occurring decay chains by Berta Karlik and Traude Bernert, first in the so-called uranium series, and then in the actinium series. (Since then, astatine was also found in a third decay chain, the neptunium series.) Friedrich Paneth in 1946 called to finally recognize synthetic elements, quoting, among other reasons, recent confirmation of their natural occurrence, and proposed that the discoverers of the newly discovered unnamed elements name these elements. In early 1947, Nature published the discoverers' suggestions; a letter from Corson, MacKenzie, and Segrè suggested the name "astatine" coming from the Greek astatos (αστατος) meaning "unstable", because of its propensity for radioactive decay, with the ending "-ine", found in the names of the four previously discovered halogens. The name was also chosen to continue the tradition of the four stable halogens, where the name referred to a property of the element. Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry. Subsequent investigators reported iodine-like, cationic, or amphoteric behavior. In a 2003 retrospective, Corson wrote that "some of the properties [of astatine] are similar to iodine … it also exhibits metallic properties, more like its metallic neighbors Po and Bi." Isotopes There are 39 known isotopes of astatine, with atomic masses (mass numbers) of 191–229. Theoretical modeling suggests that 37 more isotopes could exist. No stable or long-lived astatine isotope has been observed, nor is one expected to exist. Astatine's alpha decay energies follow the same trend as for other heavy elements. Lighter astatine isotopes have quite high energies of alpha decay, which become lower as the nuclei become heavier. Astatine-211 has a significantly higher energy than the previous isotope, because it has a nucleus with 126 neutrons, and 126 is a magic number corresponding to a filled neutron shell. Despite having a similar half-life to the previous isotope (8.1 hours for astatine-210 and 7.2 hours for astatine-211), the alpha decay probability is much higher for the latter: 41.81% against only 0.18%.

The two following isotopes release even more energy, with astatine-213 releasing the most energy. For this reason, it is the shortest-lived astatine isotope. Even though heavier astatine isotopes release less energy, no long-lived astatine isotope exists, because of the increasing role of beta decay (electron emission). This decay mode is especially important for astatine; as early as 1950 it was postulated that all isotopes of the element undergo beta decay, though nuclear mass measurements indicate that 215At is in fact beta-stable, as it has the lowest mass of all isobars with A = 215. A beta decay mode has been found for all other astatine isotopes except for astatine-213, astatine-214, and astatine-216m. Astatine-210 and lighter isotopes exhibit beta plus decay (positron emission), astatine-216 and heavier isotopes exhibit beta minus decay, and astatine-212 decays via both modes, while astatine-211 undergoes electron capture. The most stable isotope is astatine-210, which has a half-life of 8.1 hours. The primary decay mode is beta plus, to the relatively long-lived (in comparison to astatine isotopes) alpha emitter polonium-210. In total, only five isotopes have half-lives exceeding one hour (astatine-207 to -211). The least stable ground state isotope is astatine-213, with a half-life of 125 nanoseconds. It undergoes alpha decay to the extremely long-lived bismuth-209. Astatine has 24 known nuclear isomers, which are nuclei with one or more nucleons (protons or neutrons) in an excited state. A nuclear isomer may also be called a "meta-state", meaning the system has more internal energy than the "ground state" (the state with the lowest possible internal energy), making the former likely to decay into the latter. There may be more than one isomer for each isotope. The most stable of these nuclear isomers is astatine-202m1, which has a half-life of about 3 minutes, longer than those of all the ground states bar those of isotopes 203–211 and 220. The least stable is astatine-214m1; its half-life of 265 nanoseconds is shorter than those of all ground states except that of astatine-213. Natural occurrence Astatine is the rarest naturally occurring element. The total amount of astatine in the Earth's crust (quoted mass 2.36 × 1025 grams) is estimated by some to be less than one gram at any given time. Other sources estimate the amount of ephemeral astatine, present on earth at any given moment, to be up to one ounce (about 28 grams). Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes (astatine-215, -217, -218 and -219) are instead continuously produced as a result of the decay of radioactive thorium and uranium ores, and trace quantities of neptunium-237. The landmass of North and South America combined, to a depth of 16 kilometers (10 miles), contains only about one trillion astatine-215 atoms at any given time (around 3.5 × 10−10 grams). Astatine-217 is produced via the radioactive decay of neptunium-237. Primordial remnants of the latter isotope—due to its relatively short half-life of 2.14 million years—are no longer present on Earth.

The two following isotopes release even more energy, with astatine-213 releasing the most energy. For this reason, it is the shortest-lived astatine isotope. Even though heavier astatine isotopes release less energy, no long-lived astatine isotope exists, because of the increasing role of beta decay (electron emission). This decay mode is especially important for astatine; as early as 1950 it was postulated that all isotopes of the element undergo beta decay, though nuclear mass measurements indicate that 215At is in fact beta-stable, as it has the lowest mass of all isobars with A = 215. A beta decay mode has been found for all other astatine isotopes except for astatine-213, astatine-214, and astatine-216m. Astatine-210 and lighter isotopes exhibit beta plus decay (positron emission), astatine-216 and heavier isotopes exhibit beta minus decay, and astatine-212 decays via both modes, while astatine-211 undergoes electron capture. The most stable isotope is astatine-210, which has a half-life of 8.1 hours. The primary decay mode is beta plus, to the relatively long-lived (in comparison to astatine isotopes) alpha emitter polonium-210. In total, only five isotopes have half-lives exceeding one hour (astatine-207 to -211). The least stable ground state isotope is astatine-213, with a half-life of 125 nanoseconds. It undergoes alpha decay to the extremely long-lived bismuth-209. Astatine has 24 known nuclear isomers, which are nuclei with one or more nucleons (protons or neutrons) in an excited state. A nuclear isomer may also be called a "meta-state", meaning the system has more internal energy than the "ground state" (the state with the lowest possible internal energy), making the former likely to decay into the latter. There may be more than one isomer for each isotope. The most stable of these nuclear isomers is astatine-202m1, which has a half-life of about 3 minutes, longer than those of all the ground states bar those of isotopes 203–211 and 220. The least stable is astatine-214m1; its half-life of 265 nanoseconds is shorter than those of all ground states except that of astatine-213. Natural occurrence Astatine is the rarest naturally occurring element. The total amount of astatine in the Earth's crust (quoted mass 2.36 × 1025 grams) is estimated by some to be less than one gram at any given time. Other sources estimate the amount of ephemeral astatine, present on earth at any given moment, to be up to one ounce (about 28 grams). Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes (astatine-215, -217, -218 and -219) are instead continuously produced as a result of the decay of radioactive thorium and uranium ores, and trace quantities of neptunium-237. The landmass of North and South America combined, to a depth of 16 kilometers (10 miles), contains only about one trillion astatine-215 atoms at any given time (around 3.5 × 10−10 grams). Astatine-217 is produced via the radioactive decay of neptunium-237. Primordial remnants of the latter isotope—due to its relatively short half-life of 2.14 million years—are no longer present on Earth.

The two following isotopes release even more energy, with astatine-213 releasing the most energy. For this reason, it is the shortest-lived astatine isotope. Even though heavier astatine isotopes release less energy, no long-lived astatine isotope exists, because of the increasing role of beta decay (electron emission). This decay mode is especially important for astatine; as early as 1950 it was postulated that all isotopes of the element undergo beta decay, though nuclear mass measurements indicate that 215At is in fact beta-stable, as it has the lowest mass of all isobars with A = 215. A beta decay mode has been found for all other astatine isotopes except for astatine-213, astatine-214, and astatine-216m. Astatine-210 and lighter isotopes exhibit beta plus decay (positron emission), astatine-216 and heavier isotopes exhibit beta minus decay, and astatine-212 decays via both modes, while astatine-211 undergoes electron capture. The most stable isotope is astatine-210, which has a half-life of 8.1 hours. The primary decay mode is beta plus, to the relatively long-lived (in comparison to astatine isotopes) alpha emitter polonium-210. In total, only five isotopes have half-lives exceeding one hour (astatine-207 to -211). The least stable ground state isotope is astatine-213, with a half-life of 125 nanoseconds. It undergoes alpha decay to the extremely long-lived bismuth-209. Astatine has 24 known nuclear isomers, which are nuclei with one or more nucleons (protons or neutrons) in an excited state. A nuclear isomer may also be called a "meta-state", meaning the system has more internal energy than the "ground state" (the state with the lowest possible internal energy), making the former likely to decay into the latter. There may be more than one isomer for each isotope. The most stable of these nuclear isomers is astatine-202m1, which has a half-life of about 3 minutes, longer than those of all the ground states bar those of isotopes 203–211 and 220. The least stable is astatine-214m1; its half-life of 265 nanoseconds is shorter than those of all ground states except that of astatine-213. Natural occurrence Astatine is the rarest naturally occurring element. The total amount of astatine in the Earth's crust (quoted mass 2.36 × 1025 grams) is estimated by some to be less than one gram at any given time. Other sources estimate the amount of ephemeral astatine, present on earth at any given moment, to be up to one ounce (about 28 grams). Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes (astatine-215, -217, -218 and -219) are instead continuously produced as a result of the decay of radioactive thorium and uranium ores, and trace quantities of neptunium-237. The landmass of North and South America combined, to a depth of 16 kilometers (10 miles), contains only about one trillion astatine-215 atoms at any given time (around 3.5 × 10−10 grams). Astatine-217 is produced via the radioactive decay of neptunium-237. Primordial remnants of the latter isotope—due to its relatively short half-life of 2.14 million years—are no longer present on Earth.

However, trace amounts occur naturally as a product of transmutation reactions in uranium ores. Astatine-218 was the first astatine isotope discovered in nature. Astatine-219, with a half-life of 56 seconds, is the longest lived of the naturally occurring isotopes. Isotopes of astatine are sometimes not listed as naturally occurring because of misconceptions that there are no such isotopes, or discrepancies in the literature. Astatine-216 has been counted as a naturally occurring isotope but reports of its observation (which were described as doubtful) have not been confirmed. Synthesis Formation Astatine was first produced by bombarding bismuth-209 with energetic alpha particles, and this is still the major route used to create the relatively long-lived isotopes astatine-209 through astatine-211. Astatine is only produced in minuscule quantities, with modern techniques allowing production runs of up to 6.6 giga becquerels (about 86 nanograms or 2.47 × 1014 atoms). Synthesis of greater quantities of astatine using this method is constrained by the limited availability of suitable cyclotrons and the prospect of melting the target. Solvent radiolysis due to the cumulative effect of astatine decay is a related problem. With cryogenic technology, microgram quantities of astatine might be able to be generated via proton irradiation of thorium or uranium to yield radon-211, in turn decaying to astatine-211. Contamination with astatine-210 is expected to be a drawback of this method. The most important isotope is astatine-211, the only one in commercial use. To produce the bismuth target, the metal is sputtered onto a gold, copper, or aluminium surface at 50 to 100 milligrams per square centimeter. Bismuth oxide can be used instead; this is forcibly fused with a copper plate. The target is kept under a chemically neutral nitrogen atmosphere, and is cooled with water to prevent premature astatine vaporization. In a particle accelerator, such as a cyclotron, alpha particles are collided with the bismuth. Even though only one bismuth isotope is used (bismuth-209), the reaction may occur in three possible ways, producing astatine-209, astatine-210, or astatine-211. In order to eliminate undesired nuclides, the maximum energy of the particle accelerator is set to a value (optimally 29.17 MeV) above that for the reaction producing astatine-211 (to produce the desired isotope) and below the one producing astatine-210 (to avoid producing other astatine isotopes). Separation methods Since astatine is the main product of the synthesis, after its formation it must only be separated from the target and any significant contaminants. Several methods are available, "but they generally follow one of two approaches—dry distillation or [wet] acid treatment of the target followed by solvent extraction." The methods summarized below are modern adaptations of older procedures, as reviewed by Kugler and Keller. Pre-1985 techniques more often addressed the elimination of co-produced toxic polonium; this requirement is now mitigated by capping the energy of the cyclotron irradiation beam. Dry The astatine-containing cyclotron target is heated to a temperature of around 650 °C. The astatine volatilizes and is condensed in (typically) a cold trap.

However, trace amounts occur naturally as a product of transmutation reactions in uranium ores. Astatine-218 was the first astatine isotope discovered in nature. Astatine-219, with a half-life of 56 seconds, is the longest lived of the naturally occurring isotopes. Isotopes of astatine are sometimes not listed as naturally occurring because of misconceptions that there are no such isotopes, or discrepancies in the literature. Astatine-216 has been counted as a naturally occurring isotope but reports of its observation (which were described as doubtful) have not been confirmed. Synthesis Formation Astatine was first produced by bombarding bismuth-209 with energetic alpha particles, and this is still the major route used to create the relatively long-lived isotopes astatine-209 through astatine-211. Astatine is only produced in minuscule quantities, with modern techniques allowing production runs of up to 6.6 giga becquerels (about 86 nanograms or 2.47 × 1014 atoms). Synthesis of greater quantities of astatine using this method is constrained by the limited availability of suitable cyclotrons and the prospect of melting the target. Solvent radiolysis due to the cumulative effect of astatine decay is a related problem. With cryogenic technology, microgram quantities of astatine might be able to be generated via proton irradiation of thorium or uranium to yield radon-211, in turn decaying to astatine-211. Contamination with astatine-210 is expected to be a drawback of this method. The most important isotope is astatine-211, the only one in commercial use. To produce the bismuth target, the metal is sputtered onto a gold, copper, or aluminium surface at 50 to 100 milligrams per square centimeter. Bismuth oxide can be used instead; this is forcibly fused with a copper plate. The target is kept under a chemically neutral nitrogen atmosphere, and is cooled with water to prevent premature astatine vaporization. In a particle accelerator, such as a cyclotron, alpha particles are collided with the bismuth. Even though only one bismuth isotope is used (bismuth-209), the reaction may occur in three possible ways, producing astatine-209, astatine-210, or astatine-211. In order to eliminate undesired nuclides, the maximum energy of the particle accelerator is set to a value (optimally 29.17 MeV) above that for the reaction producing astatine-211 (to produce the desired isotope) and below the one producing astatine-210 (to avoid producing other astatine isotopes). Separation methods Since astatine is the main product of the synthesis, after its formation it must only be separated from the target and any significant contaminants. Several methods are available, "but they generally follow one of two approaches—dry distillation or [wet] acid treatment of the target followed by solvent extraction." The methods summarized below are modern adaptations of older procedures, as reviewed by Kugler and Keller. Pre-1985 techniques more often addressed the elimination of co-produced toxic polonium; this requirement is now mitigated by capping the energy of the cyclotron irradiation beam. Dry The astatine-containing cyclotron target is heated to a temperature of around 650 °C. The astatine volatilizes and is condensed in (typically) a cold trap.

However, trace amounts occur naturally as a product of transmutation reactions in uranium ores. Astatine-218 was the first astatine isotope discovered in nature. Astatine-219, with a half-life of 56 seconds, is the longest lived of the naturally occurring isotopes. Isotopes of astatine are sometimes not listed as naturally occurring because of misconceptions that there are no such isotopes, or discrepancies in the literature. Astatine-216 has been counted as a naturally occurring isotope but reports of its observation (which were described as doubtful) have not been confirmed. Synthesis Formation Astatine was first produced by bombarding bismuth-209 with energetic alpha particles, and this is still the major route used to create the relatively long-lived isotopes astatine-209 through astatine-211. Astatine is only produced in minuscule quantities, with modern techniques allowing production runs of up to 6.6 giga becquerels (about 86 nanograms or 2.47 × 1014 atoms). Synthesis of greater quantities of astatine using this method is constrained by the limited availability of suitable cyclotrons and the prospect of melting the target. Solvent radiolysis due to the cumulative effect of astatine decay is a related problem. With cryogenic technology, microgram quantities of astatine might be able to be generated via proton irradiation of thorium or uranium to yield radon-211, in turn decaying to astatine-211. Contamination with astatine-210 is expected to be a drawback of this method. The most important isotope is astatine-211, the only one in commercial use. To produce the bismuth target, the metal is sputtered onto a gold, copper, or aluminium surface at 50 to 100 milligrams per square centimeter. Bismuth oxide can be used instead; this is forcibly fused with a copper plate. The target is kept under a chemically neutral nitrogen atmosphere, and is cooled with water to prevent premature astatine vaporization. In a particle accelerator, such as a cyclotron, alpha particles are collided with the bismuth. Even though only one bismuth isotope is used (bismuth-209), the reaction may occur in three possible ways, producing astatine-209, astatine-210, or astatine-211. In order to eliminate undesired nuclides, the maximum energy of the particle accelerator is set to a value (optimally 29.17 MeV) above that for the reaction producing astatine-211 (to produce the desired isotope) and below the one producing astatine-210 (to avoid producing other astatine isotopes). Separation methods Since astatine is the main product of the synthesis, after its formation it must only be separated from the target and any significant contaminants. Several methods are available, "but they generally follow one of two approaches—dry distillation or [wet] acid treatment of the target followed by solvent extraction." The methods summarized below are modern adaptations of older procedures, as reviewed by Kugler and Keller. Pre-1985 techniques more often addressed the elimination of co-produced toxic polonium; this requirement is now mitigated by capping the energy of the cyclotron irradiation beam. Dry The astatine-containing cyclotron target is heated to a temperature of around 650 °C. The astatine volatilizes and is condensed in (typically) a cold trap.

Higher temperatures of up to around 850 °C may increase the yield, at the risk of bismuth contamination from concurrent volatilization. Redistilling the condensate may be required to minimize the presence of bismuth (as bismuth can interfere with astatine labeling reactions). The astatine is recovered from the trap using one or more low concentration solvents such as sodium hydroxide, methanol or chloroform. Astatine yields of up to around 80% may be achieved. Dry separation is the method most commonly used to produce a chemically useful form of astatine. Wet The irradiated bismuth (or sometimes bismuth trioxide) target is first dissolved in, for example, concentrated nitric or perchloric acid. Following this first step, the acid can be distilled away to leave behind a white residue that contains both bismuth and the desired astatine product. This residue is then dissolved in a concentrated acid, such as hydrochloric acid. Astatine is extracted from this acid using an organic solvent such as butyl or isopropyl ether, diisopropylether (DIPE), or thiosemicarbazide. Using liquid-liquid extraction, the astatine product can be repeatedly washed with an acid, such as HCl, and extracted into the organic solvent layer. A separation yield of 93% using nitric acid has been reported, falling to 72% by the time purification procedures were completed (distillation of nitric acid, purging residual nitrogen oxides, and redissolving bismuth nitrate to enable liquid–liquid extraction). Wet methods involve "multiple radioactivity handling steps" and have not been considered well suited for isolating larger quantities of astatine. However, wet extraction methods are being examined for use in production of larger quantities of astatine-211, as it is thought that wet extraction methods can provide more consistency. They can enable the production of astatine in a specific oxidation state and may have greater applicability in experimental radiochemistry. Uses and precautions {| class="wikitable" |+ Several 211At-containing molecules and their experimental uses ! Agent ! Applications |- | [211At]astatine-tellurium colloids | Compartmental tumors |- | 6-[211At]astato-2-methyl-1,4-naphtaquinol diphosphate | Adenocarcinomas |- | 211At-labeled methylene blue | Melanomas |- | Meta-[211At]astatobenzyl guanidine | Neuroendocrine tumors |- | 5-[211At]astato-2'-deoxyuridine | Various |- | 211At-labeled biotin conjugates | Various pretargeting |- | 211At-labeled octreotide | Somatostatin receptor |- | 211At-labeled monoclonal antibodies and fragments | Various |- | 211At-labeled bisphosphonates | Bone metastases |} Newly formed astatine-211 is the subject of ongoing research in nuclear medicine. It must be used quickly as it decays with a half-life of 7.2 hours; this is long enough to permit multistep labeling strategies. Astatine-211 has potential for targeted alpha-particle therapy, since it decays either via emission of an alpha particle (to bismuth-207), or via electron capture (to an extremely short-lived nuclide, polonium-211, which undergoes further alpha decay), very quickly reaching its stable granddaughter lead-207. Polonium X-rays emitted as a result of the electron capture branch, in the range of 77–92 keV, enable the tracking of astatine in animals and patients. Although astatine-210 has a slightly longer half-life, it is wholly unsuitable because it usually undergoes beta plus decay to the extremely toxic polonium-210.

Higher temperatures of up to around 850 °C may increase the yield, at the risk of bismuth contamination from concurrent volatilization. Redistilling the condensate may be required to minimize the presence of bismuth (as bismuth can interfere with astatine labeling reactions). The astatine is recovered from the trap using one or more low concentration solvents such as sodium hydroxide, methanol or chloroform. Astatine yields of up to around 80% may be achieved. Dry separation is the method most commonly used to produce a chemically useful form of astatine. Wet The irradiated bismuth (or sometimes bismuth trioxide) target is first dissolved in, for example, concentrated nitric or perchloric acid. Following this first step, the acid can be distilled away to leave behind a white residue that contains both bismuth and the desired astatine product. This residue is then dissolved in a concentrated acid, such as hydrochloric acid. Astatine is extracted from this acid using an organic solvent such as butyl or isopropyl ether, diisopropylether (DIPE), or thiosemicarbazide. Using liquid-liquid extraction, the astatine product can be repeatedly washed with an acid, such as HCl, and extracted into the organic solvent layer. A separation yield of 93% using nitric acid has been reported, falling to 72% by the time purification procedures were completed (distillation of nitric acid, purging residual nitrogen oxides, and redissolving bismuth nitrate to enable liquid–liquid extraction). Wet methods involve "multiple radioactivity handling steps" and have not been considered well suited for isolating larger quantities of astatine. However, wet extraction methods are being examined for use in production of larger quantities of astatine-211, as it is thought that wet extraction methods can provide more consistency. They can enable the production of astatine in a specific oxidation state and may have greater applicability in experimental radiochemistry. Uses and precautions {| class="wikitable" |+ Several 211At-containing molecules and their experimental uses ! Agent ! Applications |- | [211At]astatine-tellurium colloids | Compartmental tumors |- | 6-[211At]astato-2-methyl-1,4-naphtaquinol diphosphate | Adenocarcinomas |- | 211At-labeled methylene blue | Melanomas |- | Meta-[211At]astatobenzyl guanidine | Neuroendocrine tumors |- | 5-[211At]astato-2'-deoxyuridine | Various |- | 211At-labeled biotin conjugates | Various pretargeting |- | 211At-labeled octreotide | Somatostatin receptor |- | 211At-labeled monoclonal antibodies and fragments | Various |- | 211At-labeled bisphosphonates | Bone metastases |} Newly formed astatine-211 is the subject of ongoing research in nuclear medicine. It must be used quickly as it decays with a half-life of 7.2 hours; this is long enough to permit multistep labeling strategies. Astatine-211 has potential for targeted alpha-particle therapy, since it decays either via emission of an alpha particle (to bismuth-207), or via electron capture (to an extremely short-lived nuclide, polonium-211, which undergoes further alpha decay), very quickly reaching its stable granddaughter lead-207. Polonium X-rays emitted as a result of the electron capture branch, in the range of 77–92 keV, enable the tracking of astatine in animals and patients. Although astatine-210 has a slightly longer half-life, it is wholly unsuitable because it usually undergoes beta plus decay to the extremely toxic polonium-210.

Higher temperatures of up to around 850 °C may increase the yield, at the risk of bismuth contamination from concurrent volatilization. Redistilling the condensate may be required to minimize the presence of bismuth (as bismuth can interfere with astatine labeling reactions). The astatine is recovered from the trap using one or more low concentration solvents such as sodium hydroxide, methanol or chloroform. Astatine yields of up to around 80% may be achieved. Dry separation is the method most commonly used to produce a chemically useful form of astatine. Wet The irradiated bismuth (or sometimes bismuth trioxide) target is first dissolved in, for example, concentrated nitric or perchloric acid. Following this first step, the acid can be distilled away to leave behind a white residue that contains both bismuth and the desired astatine product. This residue is then dissolved in a concentrated acid, such as hydrochloric acid. Astatine is extracted from this acid using an organic solvent such as butyl or isopropyl ether, diisopropylether (DIPE), or thiosemicarbazide. Using liquid-liquid extraction, the astatine product can be repeatedly washed with an acid, such as HCl, and extracted into the organic solvent layer. A separation yield of 93% using nitric acid has been reported, falling to 72% by the time purification procedures were completed (distillation of nitric acid, purging residual nitrogen oxides, and redissolving bismuth nitrate to enable liquid–liquid extraction). Wet methods involve "multiple radioactivity handling steps" and have not been considered well suited for isolating larger quantities of astatine. However, wet extraction methods are being examined for use in production of larger quantities of astatine-211, as it is thought that wet extraction methods can provide more consistency. They can enable the production of astatine in a specific oxidation state and may have greater applicability in experimental radiochemistry. Uses and precautions {| class="wikitable" |+ Several 211At-containing molecules and their experimental uses ! Agent ! Applications |- | [211At]astatine-tellurium colloids | Compartmental tumors |- | 6-[211At]astato-2-methyl-1,4-naphtaquinol diphosphate | Adenocarcinomas |- | 211At-labeled methylene blue | Melanomas |- | Meta-[211At]astatobenzyl guanidine | Neuroendocrine tumors |- | 5-[211At]astato-2'-deoxyuridine | Various |- | 211At-labeled biotin conjugates | Various pretargeting |- | 211At-labeled octreotide | Somatostatin receptor |- | 211At-labeled monoclonal antibodies and fragments | Various |- | 211At-labeled bisphosphonates | Bone metastases |} Newly formed astatine-211 is the subject of ongoing research in nuclear medicine. It must be used quickly as it decays with a half-life of 7.2 hours; this is long enough to permit multistep labeling strategies. Astatine-211 has potential for targeted alpha-particle therapy, since it decays either via emission of an alpha particle (to bismuth-207), or via electron capture (to an extremely short-lived nuclide, polonium-211, which undergoes further alpha decay), very quickly reaching its stable granddaughter lead-207. Polonium X-rays emitted as a result of the electron capture branch, in the range of 77–92 keV, enable the tracking of astatine in animals and patients. Although astatine-210 has a slightly longer half-life, it is wholly unsuitable because it usually undergoes beta plus decay to the extremely toxic polonium-210.

The principal medicinal difference between astatine-211 and iodine-131 (a radioactive iodine isotope also used in medicine) is that iodine-131 emits high-energy beta particles, and astatine does not. Beta particles have much greater penetrating power through tissues than do the much heavier alpha particles. An average alpha particle released by astatine-211 can travel up to 70 µm through surrounding tissues; an average-energy beta particle emitted by iodine-131 can travel nearly 30 times as far, to about 2 mm. The short half-life and limited penetrating power of alpha radiation through tissues offers advantages in situations where the "tumor burden is low and/or malignant cell populations are located in close proximity to essential normal tissues." Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine-211 atoms bound per cell. Several obstacles have been encountered in the development of astatine-based radiopharmaceuticals for cancer treatment. World War II delayed research for close to a decade. Results of early experiments indicated that a cancer-selective carrier would need to be developed and it was not until the 1970s that monoclonal antibodies became available for this purpose. Unlike iodine, astatine shows a tendency to dehalogenate from molecular carriers such as these, particularly at sp3 carbon sites (less so from sp2 sites). Given the toxicity of astatine accumulated and retained in the body, this emphasized the need to ensure it remained attached to its host molecule. While astatine carriers that are slowly metabolized can be assessed for their efficacy, more rapidly metabolized carriers remain a significant obstacle to the evaluation of astatine in nuclear medicine. Mitigating the effects of astatine-induced radiolysis of labeling chemistry and carrier molecules is another area requiring further development. A practical application for astatine as a cancer treatment would potentially be suitable for a "staggering" number of patients; production of astatine in the quantities that would be required remains an issue. Animal studies show that astatine, similarly to iodine – although to a lesser extent, perhaps because of its slightly more metallic nature  – is preferentially (and dangerously) concentrated in the thyroid gland. Unlike iodine, astatine also shows a tendency to be taken up by the lungs and spleen, possibly because of in-body oxidation of At– to At+. If administered in the form of a radiocolloid it tends to concentrate in the liver. Experiments in rats and monkeys suggest that astatine-211 causes much greater damage to the thyroid gland than does iodine-131, with repetitive injection of the nuclide resulting in necrosis and cell dysplasia within the gland. Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue; this conclusion remained controversial for many years. General agreement was later reached that this was likely caused by the effect of breast tissue irradiation combined with hormonal changes due to irradiation of the ovaries. Trace amounts of astatine can be handled safely in fume hoods if they are well-aerated; biological uptake of the element must be avoided.

The principal medicinal difference between astatine-211 and iodine-131 (a radioactive iodine isotope also used in medicine) is that iodine-131 emits high-energy beta particles, and astatine does not. Beta particles have much greater penetrating power through tissues than do the much heavier alpha particles. An average alpha particle released by astatine-211 can travel up to 70 µm through surrounding tissues; an average-energy beta particle emitted by iodine-131 can travel nearly 30 times as far, to about 2 mm. The short half-life and limited penetrating power of alpha radiation through tissues offers advantages in situations where the "tumor burden is low and/or malignant cell populations are located in close proximity to essential normal tissues." Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine-211 atoms bound per cell. Several obstacles have been encountered in the development of astatine-based radiopharmaceuticals for cancer treatment. World War II delayed research for close to a decade. Results of early experiments indicated that a cancer-selective carrier would need to be developed and it was not until the 1970s that monoclonal antibodies became available for this purpose. Unlike iodine, astatine shows a tendency to dehalogenate from molecular carriers such as these, particularly at sp3 carbon sites (less so from sp2 sites). Given the toxicity of astatine accumulated and retained in the body, this emphasized the need to ensure it remained attached to its host molecule. While astatine carriers that are slowly metabolized can be assessed for their efficacy, more rapidly metabolized carriers remain a significant obstacle to the evaluation of astatine in nuclear medicine. Mitigating the effects of astatine-induced radiolysis of labeling chemistry and carrier molecules is another area requiring further development. A practical application for astatine as a cancer treatment would potentially be suitable for a "staggering" number of patients; production of astatine in the quantities that would be required remains an issue. Animal studies show that astatine, similarly to iodine – although to a lesser extent, perhaps because of its slightly more metallic nature  – is preferentially (and dangerously) concentrated in the thyroid gland. Unlike iodine, astatine also shows a tendency to be taken up by the lungs and spleen, possibly because of in-body oxidation of At– to At+. If administered in the form of a radiocolloid it tends to concentrate in the liver. Experiments in rats and monkeys suggest that astatine-211 causes much greater damage to the thyroid gland than does iodine-131, with repetitive injection of the nuclide resulting in necrosis and cell dysplasia within the gland. Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue; this conclusion remained controversial for many years. General agreement was later reached that this was likely caused by the effect of breast tissue irradiation combined with hormonal changes due to irradiation of the ovaries. Trace amounts of astatine can be handled safely in fume hoods if they are well-aerated; biological uptake of the element must be avoided.

The principal medicinal difference between astatine-211 and iodine-131 (a radioactive iodine isotope also used in medicine) is that iodine-131 emits high-energy beta particles, and astatine does not. Beta particles have much greater penetrating power through tissues than do the much heavier alpha particles. An average alpha particle released by astatine-211 can travel up to 70 µm through surrounding tissues; an average-energy beta particle emitted by iodine-131 can travel nearly 30 times as far, to about 2 mm. The short half-life and limited penetrating power of alpha radiation through tissues offers advantages in situations where the "tumor burden is low and/or malignant cell populations are located in close proximity to essential normal tissues." Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine-211 atoms bound per cell. Several obstacles have been encountered in the development of astatine-based radiopharmaceuticals for cancer treatment. World War II delayed research for close to a decade. Results of early experiments indicated that a cancer-selective carrier would need to be developed and it was not until the 1970s that monoclonal antibodies became available for this purpose. Unlike iodine, astatine shows a tendency to dehalogenate from molecular carriers such as these, particularly at sp3 carbon sites (less so from sp2 sites). Given the toxicity of astatine accumulated and retained in the body, this emphasized the need to ensure it remained attached to its host molecule. While astatine carriers that are slowly metabolized can be assessed for their efficacy, more rapidly metabolized carriers remain a significant obstacle to the evaluation of astatine in nuclear medicine. Mitigating the effects of astatine-induced radiolysis of labeling chemistry and carrier molecules is another area requiring further development. A practical application for astatine as a cancer treatment would potentially be suitable for a "staggering" number of patients; production of astatine in the quantities that would be required remains an issue. Animal studies show that astatine, similarly to iodine – although to a lesser extent, perhaps because of its slightly more metallic nature  – is preferentially (and dangerously) concentrated in the thyroid gland. Unlike iodine, astatine also shows a tendency to be taken up by the lungs and spleen, possibly because of in-body oxidation of At– to At+. If administered in the form of a radiocolloid it tends to concentrate in the liver. Experiments in rats and monkeys suggest that astatine-211 causes much greater damage to the thyroid gland than does iodine-131, with repetitive injection of the nuclide resulting in necrosis and cell dysplasia within the gland. Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue; this conclusion remained controversial for many years. General agreement was later reached that this was likely caused by the effect of breast tissue irradiation combined with hormonal changes due to irradiation of the ovaries. Trace amounts of astatine can be handled safely in fume hoods if they are well-aerated; biological uptake of the element must be avoided.

See also Radiation protection Notes References Bibliography External links Astatine at The Periodic Table of Videos (University of Nottingham) Astatine: Halogen or Metal? Halogens Metalloids Chemical elements

See also Radiation protection Notes References Bibliography External links Astatine at The Periodic Table of Videos (University of Nottingham) Astatine: Halogen or Metal? Halogens Metalloids Chemical elements

See also Radiation protection Notes References Bibliography External links Astatine at The Periodic Table of Videos (University of Nottingham) Astatine: Halogen or Metal? Halogens Metalloids Chemical elements

Atom An atom is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are extremely small, typically around 100 picometers across. They are so small that accurately predicting their behavior using classical physics—as if they were tennis balls, for example—is not possible due to quantum effects. Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number of neutrons. Only the most common variety of hydrogen has no neutrons. More than 99.94% of an atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of protons and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively – such atoms are called ions. The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus splits and leaves behind different elements. This is a form of nuclear decay. The number of protons in the nucleus is the atomic number and it defines to which chemical element the atom belongs. For example, any atom that contains 29 protons is copper. The number of neutrons defines the isotope of the element. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to associate and dissociate is responsible for most of the physical changes observed in nature. Chemistry is the discipline that studies these changes. History of atomic theory In philosophy The basic idea that matter is made up of tiny, indivisible particles appears in many ancient cultures such as those of Greece and India. The word atom is derived from the ancient Greek word atomos (a combination of the negative term "a-" and "τομή," the term for "cut") that means "uncuttable". This ancient idea was based in philosophical reasoning rather than scientific reasoning; modern atomic theory is not based on these old concepts. Nonetheless, the term "atom" was used throughout the ages by thinkers who suspected that matter was ultimately granular in nature. It has since been discovered that "atoms" can be split, but the misnomer is still used. Dalton's law of multiple proportions In the early 1800s, the English chemist John Dalton compiled experimental data gathered by himself and other scientists and discovered a pattern now known as the "law of multiple proportions".

Atom An atom is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are extremely small, typically around 100 picometers across. They are so small that accurately predicting their behavior using classical physics—as if they were tennis balls, for example—is not possible due to quantum effects. Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number of neutrons. Only the most common variety of hydrogen has no neutrons. More than 99.94% of an atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of protons and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively – such atoms are called ions. The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus splits and leaves behind different elements. This is a form of nuclear decay. The number of protons in the nucleus is the atomic number and it defines to which chemical element the atom belongs. For example, any atom that contains 29 protons is copper. The number of neutrons defines the isotope of the element. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to associate and dissociate is responsible for most of the physical changes observed in nature. Chemistry is the discipline that studies these changes. History of atomic theory In philosophy The basic idea that matter is made up of tiny, indivisible particles appears in many ancient cultures such as those of Greece and India. The word atom is derived from the ancient Greek word atomos (a combination of the negative term "a-" and "τομή," the term for "cut") that means "uncuttable". This ancient idea was based in philosophical reasoning rather than scientific reasoning; modern atomic theory is not based on these old concepts. Nonetheless, the term "atom" was used throughout the ages by thinkers who suspected that matter was ultimately granular in nature. It has since been discovered that "atoms" can be split, but the misnomer is still used. Dalton's law of multiple proportions In the early 1800s, the English chemist John Dalton compiled experimental data gathered by himself and other scientists and discovered a pattern now known as the "law of multiple proportions".

He noticed that in chemical compounds which contain a particular chemical element, the content of that element in these compounds will differ by ratios of small whole numbers. This pattern suggested to Dalton that each chemical element combines with other elements by some basic and consistent unit of mass. For example, there are two types of tin oxide: one is a black powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the black oxide there is about 13.5 g of oxygen for every 100 g of tin, and in the white oxide there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. In these oxides, for every tin atom there are one or two oxygen atoms respectively (SnO and SnO2). As a second example, Dalton considered two iron oxides: a black powder which is 78.1% iron and 21.9% oxygen, and a red powder which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black oxide there is about 28 g of oxygen for every 100 g of iron, and in the red oxide there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. In these respective oxides, for every two atoms of iron, there are two or three atoms of oxygen (Fe2O2 and Fe2O3). As a final example: nitrous oxide is 63.3% nitrogen and 36.7% oxygen, nitric oxide is 44.05% nitrogen and 55.95% oxygen, and nitrogen dioxide is 29.5% nitrogen and 70.5% oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 1:2:4. The respective formulas for these oxides are N2O, NO, and NO2. Kinetic theory of gases In the late 18th century, a number of scientists found that they could better explain the behavior of gases by describing them as collections of sub-microscopic particles and modelling their behavior using statistics and probability. Unlike Dalton's atomic theory, the kinetic theory of gases describes not how gases react chemically with each other to form compounds, but how they behave physically: diffusion, viscosity, conductivity, pressure, etc. Brownian motion In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905, Albert Einstein proved the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion.

He noticed that in chemical compounds which contain a particular chemical element, the content of that element in these compounds will differ by ratios of small whole numbers. This pattern suggested to Dalton that each chemical element combines with other elements by some basic and consistent unit of mass. For example, there are two types of tin oxide: one is a black powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the black oxide there is about 13.5 g of oxygen for every 100 g of tin, and in the white oxide there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. In these oxides, for every tin atom there are one or two oxygen atoms respectively (SnO and SnO2). As a second example, Dalton considered two iron oxides: a black powder which is 78.1% iron and 21.9% oxygen, and a red powder which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black oxide there is about 28 g of oxygen for every 100 g of iron, and in the red oxide there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. In these respective oxides, for every two atoms of iron, there are two or three atoms of oxygen (Fe2O2 and Fe2O3). As a final example: nitrous oxide is 63.3% nitrogen and 36.7% oxygen, nitric oxide is 44.05% nitrogen and 55.95% oxygen, and nitrogen dioxide is 29.5% nitrogen and 70.5% oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 1:2:4. The respective formulas for these oxides are N2O, NO, and NO2. Kinetic theory of gases In the late 18th century, a number of scientists found that they could better explain the behavior of gases by describing them as collections of sub-microscopic particles and modelling their behavior using statistics and probability. Unlike Dalton's atomic theory, the kinetic theory of gases describes not how gases react chemically with each other to form compounds, but how they behave physically: diffusion, viscosity, conductivity, pressure, etc. Brownian motion In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905, Albert Einstein proved the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion.

He noticed that in chemical compounds which contain a particular chemical element, the content of that element in these compounds will differ by ratios of small whole numbers. This pattern suggested to Dalton that each chemical element combines with other elements by some basic and consistent unit of mass. For example, there are two types of tin oxide: one is a black powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the black oxide there is about 13.5 g of oxygen for every 100 g of tin, and in the white oxide there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. In these oxides, for every tin atom there are one or two oxygen atoms respectively (SnO and SnO2). As a second example, Dalton considered two iron oxides: a black powder which is 78.1% iron and 21.9% oxygen, and a red powder which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black oxide there is about 28 g of oxygen for every 100 g of iron, and in the red oxide there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. In these respective oxides, for every two atoms of iron, there are two or three atoms of oxygen (Fe2O2 and Fe2O3). As a final example: nitrous oxide is 63.3% nitrogen and 36.7% oxygen, nitric oxide is 44.05% nitrogen and 55.95% oxygen, and nitrogen dioxide is 29.5% nitrogen and 70.5% oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 1:2:4. The respective formulas for these oxides are N2O, NO, and NO2. Kinetic theory of gases In the late 18th century, a number of scientists found that they could better explain the behavior of gases by describing them as collections of sub-microscopic particles and modelling their behavior using statistics and probability. Unlike Dalton's atomic theory, the kinetic theory of gases describes not how gases react chemically with each other to form compounds, but how they behave physically: diffusion, viscosity, conductivity, pressure, etc. Brownian motion In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905, Albert Einstein proved the reality of these molecules and their motions by producing the first statistical physics analysis of Brownian motion.

French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of molecules, thereby providing physical evidence for the particle nature of matter. Discovery of the electron In 1897, J. J. Thomson discovered that cathode rays are not electromagnetic waves but made of particles that are 1,800 times lighter than hydrogen (the lightest atom). Thomson concluded that these particles came from the atoms within the cathode — they were subatomic particles. He called these new particles corpuscles but they were later renamed electrons. Thomson also showed that electrons were identical to particles given off by photoelectric and radioactive materials. It was quickly recognized that electrons are the particles that carry electric currents in metal wires. Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as the name atomos suggests. Discovery of the nucleus J. J. Thomson thought that the negatively-charged electrons were distributed throughout the atom in a sea of positive charge that was distributed across the whole volume of the atom. This model is sometimes known as the plum pudding model. Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the charge-to-mass ratio of alpha particles (these are positively-charged particles emitted by certain radioactive substances such as radium). The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a near-perfect vacuum in his instruments. Rutherford didn't think he'd run into this same problem because alpha particles are much heavier than electrons. According to Thomson's model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully. Between 1908 and 1913, Rutheford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They spotted alpha particles being deflected by angles greater than 90°. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom's volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed. Discovery of isotopes While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table.

French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of molecules, thereby providing physical evidence for the particle nature of matter. Discovery of the electron In 1897, J. J. Thomson discovered that cathode rays are not electromagnetic waves but made of particles that are 1,800 times lighter than hydrogen (the lightest atom). Thomson concluded that these particles came from the atoms within the cathode — they were subatomic particles. He called these new particles corpuscles but they were later renamed electrons. Thomson also showed that electrons were identical to particles given off by photoelectric and radioactive materials. It was quickly recognized that electrons are the particles that carry electric currents in metal wires. Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as the name atomos suggests. Discovery of the nucleus J. J. Thomson thought that the negatively-charged electrons were distributed throughout the atom in a sea of positive charge that was distributed across the whole volume of the atom. This model is sometimes known as the plum pudding model. Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the charge-to-mass ratio of alpha particles (these are positively-charged particles emitted by certain radioactive substances such as radium). The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a near-perfect vacuum in his instruments. Rutherford didn't think he'd run into this same problem because alpha particles are much heavier than electrons. According to Thomson's model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully. Between 1908 and 1913, Rutheford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They spotted alpha particles being deflected by angles greater than 90°. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom's volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed. Discovery of isotopes While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table.

French physicist Jean Perrin used Einstein's work to experimentally determine the mass and dimensions of molecules, thereby providing physical evidence for the particle nature of matter. Discovery of the electron In 1897, J. J. Thomson discovered that cathode rays are not electromagnetic waves but made of particles that are 1,800 times lighter than hydrogen (the lightest atom). Thomson concluded that these particles came from the atoms within the cathode — they were subatomic particles. He called these new particles corpuscles but they were later renamed electrons. Thomson also showed that electrons were identical to particles given off by photoelectric and radioactive materials. It was quickly recognized that electrons are the particles that carry electric currents in metal wires. Thomson concluded that these electrons emerged from the very atoms of the cathode in his instruments, which meant that atoms are not indivisible as the name atomos suggests. Discovery of the nucleus J. J. Thomson thought that the negatively-charged electrons were distributed throughout the atom in a sea of positive charge that was distributed across the whole volume of the atom. This model is sometimes known as the plum pudding model. Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden came to have doubts about the Thomson model after they encountered difficulties when they tried to build an instrument to measure the charge-to-mass ratio of alpha particles (these are positively-charged particles emitted by certain radioactive substances such as radium). The alpha particles were being scattered by the air in the detection chamber, which made the measurements unreliable. Thomson had encountered a similar problem in his work on cathode rays, which he solved by creating a near-perfect vacuum in his instruments. Rutherford didn't think he'd run into this same problem because alpha particles are much heavier than electrons. According to Thomson's model of the atom, the positive charge in the atom is not concentrated enough to produce an electric field strong enough to deflect an alpha particle, and the electrons are so lightweight they should be pushed aside effortlessly by the much heavier alpha particles. Yet there was scattering, so Rutherford and his colleagues decided to investigate this scattering carefully. Between 1908 and 1913, Rutheford and his colleagues performed a series of experiments in which they bombarded thin foils of metal with alpha particles. They spotted alpha particles being deflected by angles greater than 90°. To explain this, Rutherford proposed that the positive charge of the atom is not distributed throughout the atom's volume as Thomson believed, but is concentrated in a tiny nucleus at the center. Only such an intense concentration of charge could produce an electric field strong enough to deflect the alpha particles as observed. Discovery of isotopes While experimenting with the products of radioactive decay, in 1913 radiochemist Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table.

The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J. J. Thomson created a technique for isotope separation through his work on ionized gases, which subsequently led to the discovery of stable isotopes. Bohr model In 1913, the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization was used to explain why the electrons' orbits are stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see synchrotron radiation) and why elements absorb and emit electromagnetic radiation in discrete spectra. Later in the same year Henry Moseley provided additional experimental evidence in favor of Niels Bohr's theory. These results refined Ernest Rutherford's and Antonius van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today. Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus. The Bohr model of the atom was the first complete physical model of the atom. It described the overall structure of the atom, how atoms bond to each other, and predicted the spectral lines of hydrogen. Bohr's model was not perfect and was soon superseded by the more accurate Schrödinger model, but it was sufficient to evaporate any remaining doubts that matter is composed of atoms. For chemists, the idea of the atom had been a useful heuristic tool, but physicists had doubts as to whether matter really is made up of atoms as nobody had yet developed a complete physical model of the atom. The Schrödinger model The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom's angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction.

The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J. J. Thomson created a technique for isotope separation through his work on ionized gases, which subsequently led to the discovery of stable isotopes. Bohr model In 1913, the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization was used to explain why the electrons' orbits are stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see synchrotron radiation) and why elements absorb and emit electromagnetic radiation in discrete spectra. Later in the same year Henry Moseley provided additional experimental evidence in favor of Niels Bohr's theory. These results refined Ernest Rutherford's and Antonius van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today. Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus. The Bohr model of the atom was the first complete physical model of the atom. It described the overall structure of the atom, how atoms bond to each other, and predicted the spectral lines of hydrogen. Bohr's model was not perfect and was soon superseded by the more accurate Schrödinger model, but it was sufficient to evaporate any remaining doubts that matter is composed of atoms. For chemists, the idea of the atom had been a useful heuristic tool, but physicists had doubts as to whether matter really is made up of atoms as nobody had yet developed a complete physical model of the atom. The Schrödinger model The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom's angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction.

The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J. J. Thomson created a technique for isotope separation through his work on ionized gases, which subsequently led to the discovery of stable isotopes. Bohr model In 1913, the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization was used to explain why the electrons' orbits are stable (given that normally, charges in acceleration, including circular motion, lose kinetic energy which is emitted as electromagnetic radiation, see synchrotron radiation) and why elements absorb and emit electromagnetic radiation in discrete spectra. Later in the same year Henry Moseley provided additional experimental evidence in favor of Niels Bohr's theory. These results refined Ernest Rutherford's and Antonius van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today. Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner. Groups of electrons were thought to occupy a set of electron shells about the nucleus. The Bohr model of the atom was the first complete physical model of the atom. It described the overall structure of the atom, how atoms bond to each other, and predicted the spectral lines of hydrogen. Bohr's model was not perfect and was soon superseded by the more accurate Schrödinger model, but it was sufficient to evaporate any remaining doubts that matter is composed of atoms. For chemists, the idea of the atom had been a useful heuristic tool, but physicists had doubts as to whether matter really is made up of atoms as nobody had yet developed a complete physical model of the atom. The Schrödinger model The Stern–Gerlach experiment of 1922 provided further evidence of the quantum nature of atomic properties. When a beam of silver atoms was passed through a specially shaped magnetic field, the beam was split in a way correlated with the direction of an atom's angular momentum, or spin. As this spin direction is initially random, the beam would be expected to deflect in a random direction.

Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field. In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics). One year earlier, Louis de Broglie had proposed the de Broglie hypothesis: that all particles behave like waves to some extent, and in 1926 Erwin Schrödinger used this idea to develop the Schrödinger equation, a mathematical model of the atom (wave mechanics) that described the electrons as three-dimensional waveforms rather than point particles. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed. Discovery of the neutron The development of the mass spectrometer allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. The explanation for these different isotopes awaited the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus. Fission, high-energy physics and condensed matter In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. In 1944, Hahn received the Nobel Prize in Chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized. In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks.

Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field. In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics). One year earlier, Louis de Broglie had proposed the de Broglie hypothesis: that all particles behave like waves to some extent, and in 1926 Erwin Schrödinger used this idea to develop the Schrödinger equation, a mathematical model of the atom (wave mechanics) that described the electrons as three-dimensional waveforms rather than point particles. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed. Discovery of the neutron The development of the mass spectrometer allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. The explanation for these different isotopes awaited the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus. Fission, high-energy physics and condensed matter In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. In 1944, Hahn received the Nobel Prize in Chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized. In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks.

Instead, the beam was split into two directional components, corresponding to the atomic spin being oriented up or down with respect to the magnetic field. In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics). One year earlier, Louis de Broglie had proposed the de Broglie hypothesis: that all particles behave like waves to some extent, and in 1926 Erwin Schrödinger used this idea to develop the Schrödinger equation, a mathematical model of the atom (wave mechanics) that described the electrons as three-dimensional waveforms rather than point particles. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed. Discovery of the neutron The development of the mass spectrometer allowed the mass of atoms to be measured with increased accuracy. The device uses a magnet to bend the trajectory of a beam of ions, and the amount of deflection is determined by the ratio of an atom's mass to its charge. The chemist Francis William Aston used this instrument to show that isotopes had different masses. The atomic mass of these isotopes varied by integer amounts, called the whole number rule. The explanation for these different isotopes awaited the discovery of the neutron, an uncharged particle with a mass similar to the proton, by the physicist James Chadwick in 1932. Isotopes were then explained as elements with the same number of protons, but different numbers of neutrons within the nucleus. Fission, high-energy physics and condensed matter In 1938, the German chemist Otto Hahn, a student of Rutherford, directed neutrons onto uranium atoms expecting to get transuranium elements. Instead, his chemical experiments showed barium as a product. A year later, Lise Meitner and her nephew Otto Frisch verified that Hahn's result were the first experimental nuclear fission. In 1944, Hahn received the Nobel Prize in Chemistry. Despite Hahn's efforts, the contributions of Meitner and Frisch were not recognized. In the 1950s, the development of improved particle accelerators and particle detectors allowed scientists to study the impacts of atoms moving at high energies. Neutrons and protons were found to be hadrons, or composites of smaller particles called quarks.

The standard model of particle physics was developed that so far has successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions. Structure Subatomic particles Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron. The electron is by far the least massive of these particles at , with a negative electrical charge and a size that is too small to be measured using available techniques. It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details. Protons have a positive charge and a mass 1,836 times that of the electron, at . The number of protons in an atom is called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton. Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or . Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of —although the 'surface' of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick. In the Standard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +) and one down quark (with a charge of −). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The quarks are held together by the strong interaction (or strong force), which is mediated by gluons. The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.

The standard model of particle physics was developed that so far has successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions. Structure Subatomic particles Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron. The electron is by far the least massive of these particles at , with a negative electrical charge and a size that is too small to be measured using available techniques. It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details. Protons have a positive charge and a mass 1,836 times that of the electron, at . The number of protons in an atom is called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton. Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or . Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of —although the 'surface' of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick. In the Standard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +) and one down quark (with a charge of −). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The quarks are held together by the strong interaction (or strong force), which is mediated by gluons. The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.

The standard model of particle physics was developed that so far has successfully explained the properties of the nucleus in terms of these sub-atomic particles and the forces that govern their interactions. Structure Subatomic particles Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton and the neutron. The electron is by far the least massive of these particles at , with a negative electrical charge and a size that is too small to be measured using available techniques. It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details. Protons have a positive charge and a mass 1,836 times that of the electron, at . The number of protons in an atom is called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton. Neutrons have no electrical charge and have a free mass of 1,839 times the mass of the electron, or . Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of —although the 'surface' of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick. In the Standard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +) and one down quark (with a charge of −). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The quarks are held together by the strong interaction (or strong force), which is mediated by gluons. The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.

Nucleus All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to  femtometres, where is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other. Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud. A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus. The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element. If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass-energy equivalence formula, , where is the mass loss and is the speed of light.

Nucleus All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to  femtometres, where is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other. Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud. A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus. The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element. If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass-energy equivalence formula, , where is the mass loss and is the speed of light.

Nucleus All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to  femtometres, where is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other. Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud. A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus. The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element. If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass-energy equivalence formula, , where is the mass loss and is the speed of light.

This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. Electron cloud The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations. Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured. Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines. The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons.

This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. Electron cloud The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations. Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured. Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines. The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons.

This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. Electron cloud The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations. Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured. Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines. The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons.

Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. Properties Nuclear properties By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible. About 339 nuclides occur naturally on Earth, of which 252 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 162 (bringing the total to 252) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 34 radioactive nuclides have half-lives longer than 100 million years, and are long-lived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14). For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.2 stable isotopes per element. Twenty-six elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes. Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 252 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10 and nitrogen-14.

Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. Properties Nuclear properties By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible. About 339 nuclides occur naturally on Earth, of which 252 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 162 (bringing the total to 252) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 34 radioactive nuclides have half-lives longer than 100 million years, and are long-lived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14). For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.2 stable isotopes per element. Twenty-six elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes. Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 252 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10 and nitrogen-14.

Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. Properties Nuclear properties By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible. About 339 nuclides occur naturally on Earth, of which 252 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 162 (bringing the total to 252) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 34 radioactive nuclides have half-lives longer than 100 million years, and are long-lived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14). For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.2 stable isotopes per element. Twenty-six elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes. Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 252 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10 and nitrogen-14.

Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects. Mass The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons). The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately . Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da), but this number will not be exactly an integer except (by definition) in the case of carbon-12. The heaviest stable atom is lead-208, with a mass of . As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about ). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. Shape and size Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right).

Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects. Mass The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons). The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately . Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da), but this number will not be exactly an integer except (by definition) in the case of carbon-12. The heaviest stable atom is lead-208, with a mass of . As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about ). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. Shape and size Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right).

Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects. Mass The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons). The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately . Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da), but this number will not be exactly an integer except (by definition) in the case of carbon-12. The heaviest stable atom is lead-208, with a mass of . As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about ). This number was chosen so that if an element has an atomic mass of 1 u, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. Shape and size Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right).

Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm. When subjected to external forces, like electrical fields, the shape of an atom may deviate from spherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites. Significant ellipsoidal deformations have been shown to occur for sulfur ions and chalcogen ions in pyrite-type compounds. Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700 nm) so they cannot be viewed using an optical microscope, although individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. A single drop of water contains about 2 sextillion () atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of contains about 10 sextillion (1022) atoms of carbon. If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple. Radioactive decay Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm. The most common forms of radioactive decay are: Alpha decay: this process is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number. Beta decay (and electron capture): these processes are regulated by the weak force, and result from a transformation of a neutron into a proton, or a proton into a neutron. The neutron to proton transition is accompanied by the emission of an electron and an antineutrino, while proton to neutron transition (except in electron capture) causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one. Electron capture is more common than positron emission, because it requires less energy. In this type of decay, an electron is absorbed by the nucleus, rather than a positron emitted from the nucleus. A neutrino is still emitted in this process, and a proton changes to a neutron. Gamma decay: this process results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation.