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Chemistry and compounds

Crystalline bulk silicon is rather inert, but becomes more reactive at high temperatures. Like its neighbour aluminium, silicon forms a thin, continuous surface layer of silicon dioxide (SiO) that protects the metal from oxidation. Thus silicon does not measurably react with the air below 900 °C, but formation of the vitreous dioxide rapidly increases between 950 °C and 1160 °C and when 1400 °C is reached, atmospheric nitrogen also reacts to give the nitrides SiN and Si N. Silicon reacts with gaseous sulfur at 600 °C and gaseous phosphorus at 1000 °C. This oxide layer nevertheless does not prevent reaction with the halogens; fluorine attacks silicon vigorously at room temperature, chlorine does so at about 300 °C, and bromine and iodine at about 500 °C. Silicon does not react with most aqueous acids, but is oxidised and complexed by hydrofluoric acid mixtures containing either chlorine or nitric acid to form hexafluorosilicates. It readily dissolves in hot aqueous alkali to form silicates. At high temperatures, silicon also reacts with alkyl halides; this reaction may be catalysed by copper to directly synthesise organosilicon chlorides as precursors to silicone polymers. Upon melting, silicon becomes extremely reactive, alloying with most metals to form silicides, and reducing most metal oxides because the heat of formation of silicon dioxide is so large. In fact, molten silicon reacts virtually with every known kind of crucible material (except its own oxide, SiO). This happens due to silicon's high binding forces for the light elements and to its high dissolving power for most elements. As a result, containers for liquid silicon must be made of refractory, unreactive materials such as zirconium dioxide or group 4, 5, and 6 borides.

Tetrahedral coordination is a major structural motif in silicon chemistry just as it is for carbon chemistry. However, the 3p subshell is rather more diffuse than the 2p subshell and does not hybridise so well with the 3s subshell. As a result, the chemistry of silicon and its heavier congeners shows significant differences from that of carbon, and thus octahedral coordination is also significant. For example, the electronegativity of silicon (1.90) is much less than that of carbon (2.55), because the valence electrons of silicon are further from the nucleus than those of carbon and hence experience smaller electrostatic forces of attraction from the nucleus. The poor overlap of 3p orbitals also results in a much lower tendency toward catenation (formation of Si–Si bonds) for silicon than for carbon, due to the concomitant weakening of the Si–Si bond compared to the C–C bond: the average Si–Si bond energy is approximately 226 kJ/mol, compared to a value of 356 kJ/mol for the C–C bond. This results in multiply bonded silicon compounds generally being much less stable than their carbon counterparts, an example of the double bond rule. On the other hand, the presence of radial nodes in the 3p orbitals of silicon suggests the possibility of hypervalence, as seen in five and six-coordinate derivatives of silicon such as SiX and SiF. Lastly, because of the increasing energy gap between the valence s and p orbitals as the group is descended, the divalent state grows in importance from carbon to lead, so that a few unstable divalent compounds are known for silicon; this lowering of the main oxidation state, in tandem with increasing atomic radii, results in an increase of metallic character down the group. Silicon already shows some incipient metallic behavior, particularly in the behavior of its oxide compounds and its reaction with acids as well as bases (though this takes some effort), and is hence often referred to as a metalloid rather than a nonmetal. Germanium shows more, and tin is generally considered a metal.

Silicon shows clear differences from carbon. For example, organic chemistry has very few analogies with silicon chemistry, while silicate minerals have a structural complexity unseen in oxocarbons. Silicon tends to resemble germanium far more than it does carbon, and this resemblance is enhanced by the d-block contraction, resulting in the size of the germanium atom being much closer to that of the silicon atom than periodic trends would predict. Nevertheless, there are still some differences because of the growing importance of the divalent state in germanium compared to silicon. Additionally, the lower Ge–O bond strength compared to the Si–O bond strength results in the absence of "germanone" polymers that would be analogous to silicone polymers.