alt=A flame with a small metal rod penetrating it; the flame near the rod is pale blue. Bulk<unk>  exposed to moist air forms a protective layer of varying composition.<unk> (II) carbonate is a common constituent; the sulfate or chloride may also be present in urban or maritime settings. This layer makes bulk<unk>  effectively chemically inert in the air. Finely powdered<unk> , as with many metals, is pyrophoric, and burns with a bluish-white flame. Fluorine reacts with<unk>  at room temperature, forming<unk> (II) fluoride. The reaction with chlorine is similar but requires heating, as the resulting chloride layer diminishes the reactivity of the elements. Molten<unk>  reacts with the chalcogens to give<unk> (II) chalcogenides.<unk>  metal resists sulfuric and phosphoric acid but not hydrochloric or nitric acid; the outcome depends on insolubility and subsequent passivation of the product salt. Organic acids, such as acetic acid, dissolve<unk>  in the presence of oxygen. Concentrated alkalis will dissolve<unk>  and form plumbites. Inorganic compounds<unk>  shows two main oxidation states: +4 and +2. The tetravalent state is common for the carbon group. The divalent state is rare for carbon and silicon, minor for germanium, important (but not prevailing) for tin, and is the more important of the two oxidation states for<unk> . This is attributable to relativistic effects, specifically the inert pair effect, which manifests itself when there is a large difference in electronegativity between<unk>  and oxide, halide, or nitride anions,<unk> ing to a significant partial positive charge on<unk> . The result is a stronger contraction of the<unk>  6s orbital than is the case for the 6p orbital, making it rather inert in ionic compounds. The inert pair effect is less applicable to compounds in which<unk>  forms covalent bonds with elements of similar electronegativity, such as carbon in organolead compounds. In these, the 6s and 6p orbitals remain similarly sized and sp3 hybridization is still energetically favorable.<unk> , like carbon, is predominantly tetravalent in such compounds. There is a relatively large difference in the electronegativity of<unk> (II) at 1.87 and<unk> (IV) at 2.33. This difference marks the reversal in the trend of increasing stability of the +4 oxidation state going down the carbon group; tin, by comparison, has values of 1.80 in the +2 oxidation state and 1.96 in the +4 state.<unk> (II) alt=Cream powderLead(II) compounds are characteristic of the inorganic chemistry of<unk> . Even strong oxidizing agents like fluorine and chlorine react with<unk>  to give only<unk> F2 and<unk> Cl2.<unk> (II) ions are usually colorless in solution, and partially hydrolyze to form<unk> (OH)+ and finally<unk> 4(OH)44+ (in which the hydroxyl ions act as bridging ligands), but are not reducing agents as tin(II) ions are. Techniques for identifying the presence of the<unk> 2+ ion in water generally rely on the precipitation of<unk> (II) chloride using dilute hydrochloric acid. As the chloride salt is sparingly soluble in water, in very dilute solutions the precipitation of<unk> (II) sulfide is achieved by bubbling hydrogen sulfide through the solution.<unk>  monoxide exists in two polymorphs, litharge α-PbO (red) and massicot β-PbO (yellow), the latter being stable only above around 488 °C. Litharge is the most commonly used inorganic compound of<unk> . There is no<unk> (II) hydroxide; increasing the pH of solutions of<unk> (II) salts<unk> s to hydrolysis and condensation.<unk>  commonly reacts with heavier chalcogens.<unk>  sulfide is a semiconductor, a photoconductor, and an extremely sensitive infrared radiation detector. The other two chalcogenides,<unk>  selenide and<unk>  telluride, are likewise photoconducting. They are unusual in that their color becomes lighter going down the group.<unk>  and oxygen in a tetragonal unit cell of<unk> (II,alt=Alternating dark gray and red balls connected by dark gray-red cylinders<unk>  dihalides are well-characterized; this includes the diastatide and mixed halides, such as<unk> FCl. The relative insolubility of the latter forms a useful basis for the gravimetric determination of fluorine. The difluoride was the first solid ionically conducting compound to be discovered (in 1834, by Michael Faraday). The other dihalides decompose on exposure to ultraviolet or visible light, especially the diiodide. Many<unk> (II) pseudohalides are known, such as the cyanide, cyanate, and thiocyanate.<unk> (II) forms an extensive variety of halide coordination complexes, such as<unk> Cl42−,<unk> Cl64−, and the<unk> 2Cl9n5n− chain anion.<unk> (II) sulfate is insoluble in water, like the sulfates of other heavy divalent cations.<unk> (II) nitrate and<unk> (II) acetate are very soluble, and this is exploited in the synthesis of other<unk>  compounds.<unk> (IV) Few inorganic<unk> (IV) compounds are known. They are only formed in highly oxidizing solutions and do not normally exist under standard conditions.<unk> (II) oxide gives a mixed oxide on further oxidation,<unk> 3O4. It is described as<unk> (II,IV) oxide, or structurally 2PbO·PbO2, and is the best-known mixed valence<unk>  compound.<unk>  dioxide is a strong oxidizing agent, capable of oxidizing hydrochloric acid to chlorine gas. This is because the expected<unk> Cl4 that would be produced is unstable and spontaneously decomposes to<unk> Cl2 and Cl2. Analogously to<unk>  monoxide,<unk>  dioxide is capable of forming plumbate anions.<unk>  disulfide and<unk>  diselenide are only stable at high pressures.<unk>  tetrafluoride, a yellow crystalline powder, is stable, but less so than the difluoride.<unk>  tetrachloride (a yellow oil) decomposes at room temperature,<unk>  tetrabromide is less stable still, and the existence of<unk>  tetraiodide is questionable. Other oxidation states capped square antiprismatic anion<unk> 94− from K(18-crown-6)2K2Pb9·(en)1.alt=Nine dark gray spheres connected by cylinders of the same color forming a convex shape Some<unk>  compounds exist in formal oxidation states other than +4 or +2.<unk> (III) may be obtained, as an intermediate between<unk> (II) and<unk> (IV), in larger organolead complexes; this oxidation state is not stable, as both the<unk> (III) ion and the larger complexes containing it are radicals. The same applies for<unk> (I), which can be found in such radical species. Numerous mixed<unk> (II,IV) oxides are known. When<unk> O2 is heated in air, it becomes<unk> 12O19 at 293 °C,<unk> 12O17 at 351 °C,<unk> 3O4 at 374 °C, and finally<unk> O at 605 °C. A further sesquioxide,<unk> 2O3, can be obtained at high pressure, along with several non-stoichiometric phases. Many of them show defective fluorite structures in which some oxygen atoms are replaced by vacancies:<unk> O can be considered as having such a structure, with every alternate layer of oxygen atoms absent. Negative oxidation states can occur as Zintl phases, as either free<unk>  anions, as in Ba2Pb, with<unk>  formally being<unk> (−IV), or in oxygen-sensitive ring-shaped or polyhedral cluster ions such as the trigonal bipyramidal<unk> 52− ion, where two<unk>  atoms are<unk> (−I) and three are<unk> (0). In such anions, each atom is at a polyhedral vertex and contributes two electrons to each covalent bond along an edge from their sp3 hybrid orbitals, the other two being an external lone pair. They may be made in liquid ammonia via the reduction of<unk>  by sodium. Organolead Structure of a tetraethyllead molecule: Carbon Hydrogen alt=A gray-green sphere linked to four black spheres, each, in turn, linked also to three white ones<unk>  can form multiply-bonded chains, a property it shares with its lighter homologs in the carbon group. Its capacity to do so is much less because the<unk> –Pb bond energy is over three and a half times lower than that of the C–C bond. With itself,<unk>  can build metal–metal bonds of an order up to three. With carbon,<unk>  forms organolead compounds similar to, but generally less stable than, typical organic compounds (due to the<unk> –C bond being rather weak). This makes the organometallic chemistry of<unk>  far less wide-ranging than that of tin.<unk>  predominantly forms organolead(IV) compounds, even when starting with inorganic<unk> (II) reactants; very few organolead(II) compounds are known. The most well-characterized exceptions are<unk> CH(SiMe3)22 and<unk> (η5-C5H5)2. The<unk>  analog of the simplest organic compound, methane, is plumbane. Plumbane may be obtained in a reaction between metallic<unk>  and atomic hydrogen. Two simple derivatives, tetramethyllead and tetraethyllead, are the best-known organolead compounds. These compounds are relatively stable: tetraethyllead only starts to decompose if heated or if exposed to sunlight or ultraviolet light. With sodium metal,<unk>  readily forms an equimolar alloy that reacts with alkyl halides to form organometallic compounds such as tetraethyllead. The oxidizing nature of many organolead compounds is usefully exploited:<unk>  tetraacetate is an important laboratory reagent for oxidation in organic synthesis. Tetraethyllead, once added to gasoline, was produced in larger quantities than any other organometallic compound. Other organolead compounds are less chemically stable. For many organic compounds, a<unk>  analog does not exist.
Lead
Basilicon
Mercury
Arsenic
Metaldehyde
Sybil 
Vinyl chloride
Lindane