Violet<unk>  vapour in a flask.<unk>  is the fourth halogen, being a member of group 17 in the periodic table, below fluorine, chlorine, and bromine; it is the heaviest stable member of its group. (The fifth and sixth halogens, the radioactive astatine and tennessine, are not well-studied due to their expense and inaccessibility in large quantities, but appear to show various unusual properties for the group due to relativistic effects.)<unk>  has an electron configuration of Kr4d105s25p5, with the seven electrons in the fifth and outermost shell being its valence electrons. Like the other halogens, it is one electron short of a full octet and is hence an oxidising agent, reacting with many elements in order to complete its outer shell, although in keeping with periodic trends, it is the weakest oxidising agent among the stable halogens: it has the lowest electronegativity among them, just 2.66 on the Pauling scale (compare fluorine, chlorine, and bromine at 3.98, 3.16, and 2.96 respectively; astatine continues the trend with an electronegativity of 2.2). Elemental<unk>  hence forms diatomic molecules with chemical formula<unk> , where two<unk>  atoms share a pair of electrons in order to each achieve a stable octet for themselves; at high temperatures, these diatomic molecules reversibly dissociate a pair of<unk>  atoms. Similarly, the<unk>  anion, I−, is the strongest reducing agent among the stable halogens, being the most easily oxidised back to diatomic<unk> . (Astatine goes further, being indeed unstable as At− and readily oxidised to At0 or At+, although the existence of At2 is not settled.) The halogens darken in colour as the group is descended: fluorine is a very pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, and<unk>  is violet. Elemental<unk>  is slightly soluble in water, with one gram dissolving in 3450 ml at 20 °C and 1280 ml at 50 °C; potassium<unk>  may be added to increase solubility via formation of triiodide ions, among other polyiodides. Nonpolar solvents such as hexane and carbon tetrachloride provide a higher solubility. Polar solutions, such as aqueous solutions, are brown, reflecting the role of these solvents as Lewis bases; on the other hand, nonpolar solutions are violet, the color of<unk>  vapour. Charge-transfer complexes form when<unk>  is dissolved in polar solvents, hence changing the colour.<unk>  is violet when dissolved in carbon tetrachloride and saturated hydrocarbons but deep brown in alcohols and amines, solvents that form charge-transfer adducts. PPh3 charge-transfer complexes in CH2Cl2. From left to right: (1)<unk>  dissolved in dichloromethane – no CT complex. (2) A few seconds after excess PPh3 was added – CT complex is forming. (3) One minute later after excess PPh3 was added, the CT complex Ph3PI+I− has been formed. (4) Immediately after excess<unk>  was added, which contains Ph3PI+I3−. The melting and boiling points of<unk>  are the highest among the halogens, conforming to the increasing trend down the group, since<unk>  has the largest electron cloud among them that is the most easily polarised, resulting in its molecules having the strongest van der Waals interactions among the halogens. Similarly,<unk>  is the least volatile of the halogens, though the solid still can be observed to give off purple vapor. Due to this property<unk>  is commonly used to demonstrate sublimation directly from solid to gas, which gives rise to a misconception that it does not melt in atmospheric pressure. Because it has the largest atomic radius among the halogens,<unk>  has the lowest first ionisation energy, lowest electron affinity, lowest electronegativity and lowest reactivity of the halogens. Structure of solid<unk>  The interhalogen bond in diiodine is the weakest of all the halogens. As such, 1% of a sample of gaseous<unk>  at atmospheric pressure is dissociated into<unk>  atoms at 575 °C. Temperatures greater than 750 °C are required for fluorine, chlorine, and bromine to dissociate to a similar extent. Most bonds to<unk>  are weaker than the analogous bonds to the lighter halogens. Gaseous<unk>  is composed of<unk>  molecules with an I–I bond length of 266.6 pm. The I–I bond is one of the longest single bonds known. It is even longer (271.5 pm) in solid orthorhombic crystalline<unk> , which has the same crystal structure as chlorine and bromine. (The record is held by<unk> 's neighbour xenon: the Xe–Xe bond length is 308.71 pm.) As such, within the<unk>  molecule, significant electronic interactions occur with the two next-nearest neighbours of each atom, and these interactions give rise, in bulk<unk> , to a shiny appearance and semiconducting properties.<unk>  is a two-dimensional semiconductor with a band gap of 1.3 eV (125 kJ/mol): it is a semiconductor in the plane of its crystalline layers and an insulator in the perpendicular direction. Isotopes Of the thirty-seven known isotopes of<unk> , only one occurs in nature,<unk> -127. The others are radioactive and have half-lives too short to be primordial. As such,<unk>  is both monoisotopic and mononuclidic and its atomic weight is known to great precision, as it is a constant of nature. The longest-lived of the radioactive isotopes of<unk>  is<unk> -129, which has a half-life of 15.7 million years, decaying via beta decay to stable xenon-129. Some<unk> -129 was formed along with<unk> -127 before the formation of the Solar System, but it has by now completely decayed away, making it an extinct radionuclide that is nevertheless still useful in dating the history of the early Solar System or very old groundwaters, due to its mobility in the environment. Its former presence may be determined from an excess of its daughter xenon-129. Traces of<unk> -129 still exist today, as it is also a cosmogenic nuclide, formed from cosmic ray spallation of atmospheric xenon: these traces make up 10−14 to 10−10 of all terrestrial<unk> . It also occurs from open-air nuclear testing, and is not hazardous because of its very long half-life, the longest of all fission products. At the peak of thermonuclear testing in the 1960s and 1970s,<unk> -129 still made up only about 10−7 of all terrestrial<unk> . Excited states of<unk> -127 and<unk> -129 are often used in Mössbauer spectroscopy. The other<unk>  radioisotopes have much shorter half-lives, no longer than days. Some of them have medical applications involving the thyroid gland, where the<unk>  that enters the body is stored and concentrated.<unk> -123 has a half-life of thirteen hours and decays by electron capture to tellurium-123, emitting gamma radiation; it is used in nuclear medicine imaging, including single photon emission computed tomography (SPECT) and X-ray computed tomography (X-Ray CT) scans.<unk> -125 has a half-life of fifty-nine days, decaying by electron capture to tellurium-125 and emitting low-energy gamma radiation; the second-longest-lived<unk>  radioisotope, it has uses in biological assays, nuclear medicine imaging and in radiation therapy as brachytherapy to treat a number of conditions, including prostate cancer, uveal melanomas, and brain tumours. Finally,<unk> -131, with a half-life of eight days, beta decays to an excited state of stable xenon-131 that then converts to the ground state by emitting gamma radiation. It is a common fission product and thus is present in high levels in radioactive fallout. It may then be absorbed through contaminated food, and will also accumulate in the thyroid. As it decays, it may cause damage to the thyroid. The primary risk from exposure to high levels of<unk> -131 is the chance occurrence of radiogenic thyroid cancer in later life. Other risks include the possibility of non-cancerous growths and thyroiditis. The usual means of protection against the negative effects of<unk> -131 is by saturating the thyroid gland with stable<unk> -127 in the form of potassium<unk>  tablets, taken daily for optimal prophylaxis. However,<unk> -131 may also be used for medicinal purposes in radiation therapy for this very reason, when tissue destruction is desired after<unk>  uptake by the tissue.<unk> -131 is also used as a radioactive tracer.
Iodine
Lindane
Copper 
Chorus
Fluoride
Amine fluoride
Hydrogen fluoride
Sodium fluoride