Silicon compounds

Chemical compounds with at least one silicon atom

Silicon compounds are compounds containing the element silicon (Si). As a carbon group element, silicon often forms compounds in the +4 oxidation state, though many unusual compounds have been discovered that differ from expectations based on its valence electrons, including the silicides and some silanes. Metal silicides, silicon halides, and similar inorganic compounds can be prepared by directly reacting elemental silicon or silicon dioxide with stable metals or with halogens. Silanes, compounds of silicon and hydrogen, are often used as strong reducing agents, and can be prepared from aluminum–silicon alloys and hydrochloric acid. Several inorganic compounds have been formed with silicon and other nonmetals such as sulfur and nitrogen; most of these compounds are highly incompatible with water. One of the most useful and successfully marketed inorganic silicon compounds is silicon carbide. Naturally occurring silicon is found in silicate and aluminosilicate minerals. One of the most common silicon compounds found in the Earth's crust is silicon dioxide or silica, which often occurs as quartz. Organosilicon compounds are fairly stable due to the similarity in strength of the Si–C bond to the C–C bond. Organosilicates include silicone polymers.

Silicon and silicon carbide readily react with all four stable halogens, forming the colourless, reactive, and volatile silicon tetrahalides. Silicon tetrafluoride also may be made by fluorinating the other silicon halides, and is produced by the attack of hydrofluoric acid on glass. The reaction of silicon tetrafluoride with excess hydrofluoric acid produces the octahedral hexafluorosilicate anion SiF62 Pure silicon crystals are very rarely found in nature, but notable exceptions are crystals as large as 0.3 mm across found during sampling gases from the Kudriavy volcano on Iturup, one of the Kuril Islands. Silicate and aluminosilicate minerals have many different structures and varying stoichiometry, but they may be classified following some general principles. Tetrahedral SiO|4 units are common to almost all these compounds, either as discrete structures, or combined into larger units by the sharing of corner oxygen atoms. These may be divided into neso silicates (discrete SiO|4 units) sharing no oxygen atoms, soro silicates (discrete Si|2|O|7 units) sharing one, cyclo silicates (closed ring structures) and ino silicates (continuous chain or ribbon structures) both sharing two, phyllo silicates (continuous sheets) sharing three, and tecto silicates (continuous three dimensional frameworks) sharing four. The lattice of oxygen atoms that results is usually close packed, or close to it, with the charge being balanced by other cations in various different polyhedral sites according to size. The orthosilicates MSiO (M = Be, Mg, Mn, Fe, Zn) and are neso silicates. (phenacite) is unusual as both BeII and SiIV occupy tetrahedral four coordinated sites; the other divalent cations instead occupy six coordinated octahedral sites and often isomorphously replace each other as in olivine, Zircon, , demands eight coordination of the ZrIV cations due to stoichiometry and because of their larger ionic radius (84 pm). Also significant are the garnets, [MM(SiO)], in which the divalent cations (e. g. Ca, Mg, Fe) are eight coordinated and the trivalent ones are six coordinated (e. g. Al, Cr, Fe). Regular coordination is not always present: for example, it is not found in , which mixes six and eight coordinate sites for CaII. Soro silicates, involving discrete double or triple tetrahedral units, are quite rare: metasilicates involving cyclic 2n− units of corner abutting tetrahedra forming a polygonal ring are also known. Chain metasilicates, {SiO}, form by corner sharing of an indefinite chain of linked SiO|4 tetrahedra. Many differences arise due to the differing repeat distances of conformation across the line of tetrahedra. A repeat distance of two is most common, as in most pyroxene minerals, but repeat distances of one, three, four, five, six, seven, nine, and twelve are also known. These chains may then link across each other to form double chains and ribbons, as in the asbestos minerals, involving repeated chains of cyclic tetrahedron rings. Layer silicates, such as the clay minerals and the micas, are very common, and often are formed by horizontal cross linking of metasilicate chains or planar condensation of smaller units. An example is kaolinite []; in many of these minerals cation and anion replacement is common, so that for example tetrahedral SiIV may be replaced by AlIII, octahedral AlIII by MgII, and by Three dimensional framework aluminosilicates are structurally very complex; they may be conceived of as starting from the structure, but having replaced up to one half of the SiIV atoms with AlIII, they require more cations to be included in the structure to balance charge. Examples include feldspars (the most abundant minerals on the Earth), zeolites, and ultramarines. Many feldspars can be thought of as forming part of the ternary system Their lattice is destroyed by high pressure prompting AlIII to undergo six coordination rather than four coordination, and this reaction destroying feldspars may be a reason for the Mohorovičić discontinuity, which would imply that the crust and mantle have the same chemical composition, but different lattices, although this is not a universally held view. Zeolites have many polyhedral cavities in their frameworks (truncated cuboctahedra being most common, but other polyhedra also are known as zeolite cavities), allowing them to include loosely bound molecules such as water in their structure. Ultramarines alternate silicon and aluminium atoms and include a variety of other anions such as , SO42 , and S22 , but are otherwise similar to the feldspars.

Because the Si–C bond is close in strength to the C–C bond, organosilicon compounds tend to be markedly thermally and chemically stable. For example, tetraphenylsilane may be distilled in air even at its boiling point of 428 °C, and so may its substituted derivatives Ph3SiCl and , which boil at 378 °C and 305 °C respectively. Furthermore, since carbon and silicon are chemical congeners, organosilicon chemistry shows some significant similarities with carbon chemistry, for example in the propensity of such compounds for catenation and forming multiple bonds. However, significant differences also arise: since silicon is more electropositive than carbon, bonds to more electronegative elements are generally stronger with silicon than with carbon, and vice versa. Thus the Si–F bond is significantly stronger than even the C–F bond and is one of the strongest single bonds, while the Si–H bond is much weaker than the C–H bond and is readily broken. Furthermore, the ability of silicon to expand its octet is not shared by carbon, and hence some organosilicon reactions have no organic analogues. For example, nucleophilic attack on silicon does not proceed by the SN2 or SN1 processes, but instead goes through a negatively charged true pentacoordinate intermediate and appears like a substitution at a hindered tertiary atom. This works for silicon, unlike for carbon, because the long Si–C bonds reduce the steric hindrance and there are no geometric constraints for nucleophilic attack, unlike for example a C–O σ* antibonding orbital. Nevertheless, despite these differences, the mechanism is still often called "SN2 at silicon" for simplicity. One of the most useful silicon containing groups is trimethylsilyl, Me3Si–. The Si–C bond connecting it to the rest of the molecule is reasonably strong, allowing it to remain while the rest of the molecule undergoes reactions, but is not so strong that it cannot be removed specifically when needed, for example by the fluoride ion, which is a very weak nucleophile for carbon compounds but a very strong one for organosilicon compounds. It may be compared to acidic protons; while trimethylsilyl is removed by hard nucleophiles instead of bases, both removals usually promote elimination. As a general rule, while saturated carbon is best attacked by nucleophiles that are neutral compounds, those based on nonmetals far down on the periodic table (e. g. sulfur, selenium, or iodine), or even both, silicon is best attacked by charged nucleophiles, particularly those involving such highly electronegative nonmetals as oxygen, fluorine, or chlorine. For example, enolates react at the carbon in haloalkanes, but at the oxygen in silyl chlorides; and when trimethylsilyl is removed from an organic molecule using hydroxide as a nucleophile, the product of the reaction is not the silanol as one would expect from using carbon chemistry as an analogy, because the siloxide is strongly nucleophilic and attacks the original molecule to yield the silyl ether hexamethyldisiloxane, Conversely, while the SN2 reaction is mostly unaffected by the presence of a partial positive charge (δ+) at the carbon, the analogous "SN2" reaction at silicon is so affected. Thus, for example, the silyl triflates are so electrophilic that they react 108 to 109 times faster than silyl chlorides with oxygen containing nucleophiles. Trimethylsilyl triflate is in particular a very good Lewis acid and is used to convert carbonyl compounds to acetals and silyl enol ethers, reacting them together analogously to the aldol reaction. Si–C bonds are commonly formed in three ways. In the laboratory, preparation is often carried out in small quantities by reacting tetrachlorosilane (silicon tetrachloride) with organolithium, Grignard, or organoaluminium reagents, or by catalytic addition of Si–H across C=C double bonds. The second route has the drawback of not being applicable to the most important silanes, the methyl and phenyl silanes. Organosilanes are made industrially by directly reacting alkyl or aryl halides with silicon with 10% by weight metallic copper as a catalyst. Standard organic reactions suffice to produce many derivatives; the resulting organosilanes are often significantly more reactive than their carbon congeners, readily undergoing hydrolysis, ammonolysis, alcoholysis, and condensation to form cyclic oligomers or linear polymers.

The word "silicone" was first used by Frederic Kipping in 1901. He invented the word to illustrate the similarity of chemical formulae between Ph2SiO and benzophenone, Ph2CO, although he also stressed the lack of chemical resemblance due to the polymeric structure of Ph2SiO, which is not shared by Ph2CO. Silicones may be considered analogous to mineral silicates, in which the methyl groups of the silicones correspond to the isoelectronic < of the silicates. They are quite stable to extreme temperatures, oxidation, and water, and have useful dielectric, antistick, and antifoam properties. Furthermore, they are resistant over long periods of time to ultraviolet radiation and weathering, and are inert physiologically. They are fairly unreactive, but do react with concentrated solutions bearing the hydroxide ion and fluorinating agents, and occasionally, may even be used as mild reagents for selective syntheses. For example, is valuable for the preparation of derivatives of molybdenum and tungsten oxyhalides, converting a tungsten hexachloride suspension in dichloroethane solution quantitatively to in under an hour at room temperature, and then to yellow at 100 °C in light petroleum at a yield of 95% overnight.