Silicon

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Silicon is a chemical element with symbol Si and atomic number 14. A hard and brittle crystalline solid with a blue-grey metallic lustre, it is a tetravalent metalloid and semiconductor. It is a member of group 14 in the periodic table, along with carbon above it and germanium, tin, and lead below. It is rather unreactive, though less so than germanium, and has a very large chemical affinity for oxygen; as such, it was first prepared and characterized in pure form only in 1823 by Jöns Jakob Berzelius.

Silicon is the eighth most common element in the universe by mass, but very rarely occurs as the pure element in the Earth’s crust. It is most widely distributed in dusts, sands, planetoids, and planets as various forms of silicon dioxide (silica) or silicates. Over 90% of the Earth’s crust is composed of silicate minerals, making silicon the second most abundant element in the Earth’s crust (about 28% by mass) after oxygen.^[9]

Most silicon is used commercially without being separated, and often with little processing of the natural minerals. Such use includes industrial construction with clays, silica sand, and stone. Silicates are used in Portland cement for mortar and stucco, and mixed with silica sand and gravel to make concrete for walkways, foundations, and roads. They are also used in whiteware ceramics such as porcelain, and in traditional quartz-based soda-lime glass and many other specialty glasses. Silicon compounds such as silicon carbideare used as abrasives and components of high-strength ceramics.

Elemental silicon also has a large impact on the modern world economy. Most free silicon is used in the steel refining, aluminium-casting, and fine chemical industries (often to make fumed silica). Even more visibly, the relatively small portion of very highly purified silicon used in semiconductor electronics (< 10%) is essential to integrated circuits — most computers, cell phones, and modern technology depend on it. Silicon is the basis of the widely used synthetic polymers called silicones.

Silicon is an essential element in biology, although only tiny traces are required by animals.^[10] However, various sea sponges and microorganisms, such as diatoms and radiolaria, secrete skeletal structures made of silica. Silica is deposited in many plant tissues, such as in the bark and wood of Chrysobalanaceae and the silica cells and silicified trichomes of Cannabis sativa, horsetails and many grasses.

History

Jöns Jacob Berzelius, discoverer of silicon

In 1787 Antoine Lavoisier suspected that silica might be an oxide of a fundamental chemical element,^[12] but the chemical affinity of silicon for oxygen is high enough that he had no means to reduce the oxide and isolate the element.^[13] After an attempt to isolate silicon in 1808, Sir Humphry Davy proposed the name “silicium” for silicon, from the Latin silex, silicis for flint, and adding the “-ium” ending because he believed it to be a metal.^[14] Most other languages use transliterated forms of Davy’s name, sometimes adapted to local phonology (e.g. GermanSilizium, French silicium). A few others use instead a calque of the Latin root (e.g. Russianкремний, from кремень “flint”; Greek πυριτιο from πυρ “fire”; Finnish pii from piikivi “flint”).^[15]

In 1811, Gay-Lussac and Thénard are thought to have prepared impure amorphous silicon, through the heating of recently isolated potassium metal with silicon tetrafluoride, but they did not purify and characterize the product, nor identify it as a new element.^[16] Silicon was given its present name in 1817 by Scottish chemist Thomas Thomson. He retained part of Davy’s name but added “-on” because he believed that silicon was a nonmetal similar to boron and carbon.^[17]In 1823, Jöns Jacob Berzelius prepared amorphous silicon using approximately the same method as Gay-Lussac (reducing potassium fluorosilicate with molten potassium metal), but purifying the product to a brown powder by repeatedly washing it.^[18] As a result, he is usually given credit for the element’s discovery.^[19][20] The same year, Berzelius became the first to prepare silicon tetrachloride; silicon tetrafluoride had already been prepared long before in 1771 by Carl Wilhelm Scheele by dissolving silica in hydrofluoric acid.^[13]

Silicon in its more common crystalline form was not prepared until 31 years later, by Deville.^[21][22] By electrolyzing a mixture of sodium chloride and aluminium chloride containing approximately 10% silicon, he was able to obtain a slightly impure allotrope of silicon in 1854.^[23] Later, more cost-effective methods have been developed to isolate several allotrope forms, the most recent being silicene in 2010.^[24][25] Meanwhile, research on the chemistry of silicon continued; Friedrich Wöhler discovered the first volatile hydrides of silicon, synthesising trichlorosilane in 1857 and silane itself in 1858, but a detailed investigation of the silanes was only carried out in the early 20th century by Alfred Stock, despite early speculation on the matter dating as far back as the beginnings of synthetic organic chemistry in the 1830s.^[26] Similarly, the first organosilicon compound, tetraethylsilane, was synthesised by Charles Friedel and James Crafts in 1863, but detailed characterisation of organosilicon chemisry was only done in the early 20th century by Frederick Kipping.^[13]

Starting in the 1920s, the work of William Lawrence Bragg on X-ray crystallography successfully elucidated the compositions of the silicates, which had previously been known from analytical chemistry but had not yet been understood, together with Linus Pauling‘s development of crystal chemistry and Victor Goldschmidt‘s development of geochemistry. The middle of the 20th century saw the development of the chemistry and industrial use of siloxanes and the growing use of silicone polymers, elastomers, and resins. In the late 20th century, the complexity of the crystal chemistry of silicides was mapped, along with the solid-state chemistry of dopedsemiconductors.^[13]

Because silicon is an important element in high-technology semiconductor devices, many places in the world bear its name. For example, Santa Clara Valley in California acquired the nickname Silicon Valley since the element is the base material used in the semiconductor industry located there. Since then, many other locations have been nicknamed for similar reasons.^[27]

Characteristics

Physical and atomic

Silicon crystallizes in a diamond cubic crystal structure

A silicon atom has fourteen electrons. In the ground state, they are arranged in the electron configuration [Ne]3s²3p². Of these, four are valence electrons, occupying the 3s orbital and two of the 3p orbitals. Like the other members of its group, the lighter carbon and the heavier germanium, tin, and lead, it has the same number of valence electrons as valence orbitals: hence, it can complete its octet and obtain the stable noble gas configuration of argon by forming sp³ hybrid orbitals, forming tetrahedral SiX₄ derivatives where the central silicon atom shares an electron pair with each of the four atoms it is bonded to.^[28] The first four ionisation energies of silicon are 786.3, 1576.5, 3228.3, and 4354.4 kJ/mol respectively; these figures are high enough to preclude the possibility of simple cationic chemistry for the element. Following periodic trends, its single-bond covalent radius of 117.6 pm is intermediate between those of carbon (77.2 pm) and germanium (122.3 pm). The hexacoordinate ionic radius of silicon may be considered to be 40 pm, although this must be taken as a purely notional figure given the lack of a simple Si⁴⁺ cation in reality.^[29]

At standard temperature and pressure, silicon is a shiny semiconductor with a bluish-grey metallic lustre; as typical for semiconductors, its resistivity drops as temperature rises. This arises because silicon has a small energy gap between its highest occupied energy levels (the valence band) and the lowest unoccupied ones (the conduction band). The Fermi level is about halfway between these and is the energy at which a state is as likely to be occupied by an electron as not. Hence pure silicon is an insulator at room temperature. However, doping silicon with a pnictogen such as phosphorus, arsenic, or antimony introduces one extra electron per dopant and these may then be excited into the conduction band either thermally or photolytically, creating an n-type semiconductor. Similarly, doping silicon with a group 13 element such as boron, aluminium, and gallium results in the introduction of acceptor levels that trap electrons that may be excited from the filled valence band, creating a p-type semiconductor. Joining n-type silicon to p-type silicon creates a p-n junction with a common Fermi level; electrons flow from n to p, while holes flow from p to n, creating a voltage drop. This p-n junction thus acts as a diode that can rectify alternating curent that allows current to pass more easily one way than the other. A transistor is an n-p-n junction, with a thin layer of weakly p-type silicon between two n-type regions. Biasing the emitter through a small forward voltage and the collector through a large reverse voltage allows the transistor to act as a triode amplifier.^[30]

Silicon crystallises in a giant covalent structure at standard conditions, specifically in a diamond cubic lattice. It thus has a high melting point of 1414 °C. It is not known to have any allotropes at standard pressure, but several other crystal structures are known at higher pressures. The general trend is one of increasing coordination number with pressure, culminating in a hexagonal close-packed allotrope at about 40 gigapascals known as Si–VII (the standard modification being Si–I). Silicon boils at 3265 °C: this, while high, is still lower than the temperature at which its lighter congener carbon sublimes (3642 °C) and silicon similarly has a lower heat of vaporisation than carbon, consistent with the fact that the Si–Si bond is weaker than the C–C bond.^[30]

Isotopes

Main article: isotopes of silicon

Naturally occurring silicon is composed of three stable isotopes, ²⁸Si (92.23%), ²⁹Si (4.67%), and ³⁰Si (3.10%), with ²⁸Si being the most abundant.^[31] Out of these, only ²⁹Si is of use in NMR and EPR spectroscopy,^[32] as it is the only one with a nuclear spin (I = 1/2).^[33] All three are produced in stars through the oxygen-burning process.^[34]

Twenty radioisotopes have been characterized, with the most stable being ³²Si with a half-life of about 150 years, and ³¹Si with a half-life of 2.62 hours.^[31] All of the remaining radioactive isotopes have half-lives that are less than seven seconds, and the majority of these have half-lives that are less than one tenth of a second.^[31] Silicon does not have any known nuclear isomers.^[31] 32Si undergoes low-energy beta decay to ³²P and then stable ³²S. ³¹Si may be produced by the neutron activation of natural silicon and is thus useful for quantitative analysis; it can be easily detected by its characteristic beta decay to stable ³¹P, in which the emitted electron carries up to 1.48 MeV of energy.^[33]

The known isotopes of silicon range in mass number from 22 to 44.^[31] The most common decay mode of the isotopes with mass numbers lower than the three stable isotopes is inverse beta decay, primarily forming aluminium isotopes (13 protons) as decay products.^[31] The most common decay mode for the heavier unstable isotopes is beta decay, primarily forming phosphorus isotopes (15 protons) as decay products.^[31]

Chemistry and compounds

C–X and Si–X bond energies (kJ/mol)^[26]

X =	C	Si	H	F	Cl	Br	I	O–	N<
C–X	368	360	435	453	351	293	216	~360	~305
Si–X	360	340	393	565	381	310	234	452	322

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 (SiO₂) 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 Si₃N₄. 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 fluorinated by a mixture of concentrated nitric acid and hydrofluoric acid; it readily dissolves in hot aqueous alkali to form silicates. At high temperatures, silicon also reacts with alkyl halides; this reaction can be catalysed by copper to directly synthesise organosilicon chlorides as precursors to silicone polymers. Upon melting, silicon becomes extremely reactive in the liquid state, alloying with most metals to form silicides, and reducing most metal oxides because the heat of formation of silicon dioxide is so large. 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.^[30]

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 as well with the 3s subshell. As a result, the chemistry of silicon and its heavier congeners shows significant differences from that of carbon,^[35] and thus octahedral coordination is also significant.^[30] 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 towards 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:^[36] the average Si–Si bond energy is approximately 226 kJ/mol, compared to a value of 356 kJ/mol for the C–C bond.^[37] 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 3d orbitals in the valence shell of silicon suggests the possibility of hypervalence, as seen in five- and six-coordinate derivatives of silicon such as SiX−
5 and SiF2−
6.^[36] 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 behaviour, particularly in the behaviour 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.^[36] However, metallicity does not become clear until germanium and dominant until tin, with the growing importance of the lower +2 oxidation state.^[13]

Silicon shows clear differences with carbon. For example, organic chemistry has very few analogies with silicon chemistry, while silicate minerals have a structural complexity unseen in oxocarbons.^[13] 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.^[29] Nevertheless, there are still some differences because of the growing importance of the divalent state in germanium compared to silicon, that result in germanium being significantly more metallic than 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.^[37]

Silicides

Main articles: Silicide and Binary compounds of silicon

Phase diagram of the Fe–Si system

Many metal silicides are known, most of which have formulae that cannot be explained through simple appeals to valence: their bonding ranges from metallic to ionic and covalent. Some known stoichiometries are M₆Si, M₅Si, M₄Si, M₁₅Si₄, M₃Si, M₅Si₂, M₂Si, M₅Si₃, M₃Si₂, MSi, M₂Si₃, MSi₂, MSi₃, and MSi₆. They are structurally more similar to the borides than the carbides, in keeping with the diagonal relationship between boron and silicon, although the larger size of silicon than boron means that exact structural analogies are few and far between. The heats of formation of the silicides are usually similar to those of the borides and carbides of the same elements, but they usually melt at lower temperatures.^[38] Silicides are known for all stable elements in groups 1–10, with the exception of beryllium: in particular, uranium and the transition metals of groups 4–10 show the widest range of stoichiometries. Except for copper, the metals in groups 11–15 do not form silicides. Most instead form eutectic mixtures, although the heaviest post-transition metals mercury, thallium, lead, and bismuth are completely immiscible with liquid silicon.^[38]

Silicides are usually prepared by direct reaction of the elements. For example, the alkali metals and alkaline earth metals react with silicon or silicon oxide to give silicides. Nevertheless, even with these highly electropositive elements true silicon anions are not obtainable, and most of these compounds are semiconductors. For example, the alkali metal silicides (M+
)
4(E4−
4) contain pyramidal tricoordinate silicon in the Si4−
4 anion, isoeelctronic with white phosphorus, P₄.^[38][39] Metal-rich silicides tend to have isolated silicon atoms (e.g. Cu₅Si); with increasing silicon content, catenation increases, resulting in isolated clusters of two (e.g. U₃Si₂) or four silicon atoms (e.g. [K⁺]₄[Si₄]⁴⁻) at first, followed by chains (e.g. CaSi), layers (e.g. CaSi₂), or three-dimensional networks of silicon atoms spanning space (e.g. α-ThSi₂) as the silicon content rises even higher.^[38]

The silicides of the group 1 and 2 metals are usually more reactive than the transition metal silicides. The latter usually do not react with aqueous reagents, except for hydrofluoric acid; however, they do react with much more aggressive reagents like liquid potassium hydroxide, or gaseous fluorine or chlorine when red-hot. The pre-transition metal silicides instead readily react with water and aqueous acids, usually producing hydrogen or silanes:^[38]

Na₂Si + 3 H₂O → Na₂SiO₃ + 3 H₂

Mg₂Si + 2 H₂SO₄ → 2 MgSO₄ + SiH₄

Products often vary with the stoichiometry of the silicide reactant. For example, Ca₂Si is polar and non-conducting and has the anti-PbCl₂ structure with single isolated silicon atoms, and reacts with water to produce calcium hydroxide, hydrated silicon dioxide, and hydrogen gas. CaSi with its zigzag chains of silicon atoms instead reacts to give silanes and polymeric SiH₂, while CaSi₂ with its puckered layers of silicon atoms does not react with water, but will react with dilute hydrochloric acid: the product is a yellow polymeric solid with stoichiometry Si₂H₂O.^[38]

Silanes

Main article: Silanes

Speculation on silicon hydride chemistry started from the 1830s, contemporary with the development of synthetic organic chemistry. Silane itself, as well as trichlorosilane, were first synthesised by Friedrich Wöhler and Heinrich Buff in 1857 by reacting aluminium–silicon alloys with hydrochloric acid, and characterised as SiH₄ and SiHCl₃ by Charles Friedel and Albert Ladenburg in 1867. Disilane (Si₂H₆) followed in 1902, when it was first made by Henri Moissan and Samuel Smiles by the protonolysis of magnesium silicides. Further investigation had to wait until 1916 because of the great reactivity and thermal instability of the silanes; it was then that Alfred Stock began the study of silicon hydrides in earnest with new greaseless vacuum techniques, as they were found as contaminants of his focus, the boron hydrides. The names silanes and boranes are due to him, based on analogy with the alkanes.^[26][40][41] Moissan and Smiles’ method of preparation of silanes and silane derivatives via protonolysis of metal silicides is still used, although the yield is lowered by the hydrolysis of the products that occurs simultaneously, and thus the preferred route today is to treat substituted silanes with hydride reducing agents such as lithium aluminium hydride in etheric solutions at low temperatures. Direct reaction of HX or RX with silicon, possibly with a catalyst such as copper, is also a viable method to produce substituted silanes.^[26]

The silanes comprise a homologous series of silicon hydrides with the general formula Si_nH_2n + 2. They are all strong reducing agents. Unbranched and branched trains are known up to n = 8, and the cycles Si₅H₁₀ and Si₆H₁₂ are also known. The first two, silane and disilane, are colourless gases; the heavier members of the series are volatile liquids. All silanes are very reactive and catch fire or explode spontaneously in air. They become less thermally stable with room temperature, so that only silane is indefinitely stable at room temperature, although disilane does not decompose very quickly (only 2.5% of a sample decomposes after eight months have passed).^[26] They decompose to form polymeric polysilicon hydrideand hydrogen gas.^[42][43] As expected from the difference in atomic weight, the silanes are less volatile than the corresponding alkanes and boranes, but more volatile than the corresponding germanes. They are much more reactive than the corresponding alkanes, because the larger radius of silicon compared to carbon facilitates nucleophilic attack at the silicon, the greater polarity of the Si–H bond compared to the C–H bond, and the ability of silicon to expand its octet and hence form adducts and lower the reaction’s activation energy.^[26]

Silane pyrolysis gives polymeric species and finally elemental silicon and hydrogen; indeed ultrapure silicon is commercially produced by the pyrolysis of silane. While the thermal decomposition of alkanes starts by the breaking of a C–H or C–C bond and the formation of radical intermediates, polysilanes decompose by eliminating silenes :SiH₂ or :SiHR, as the activation energy of this process (~210 kJ/mol) is much less than the Si–Si and Si–H bond energies. While pure silanes do not react with pure water or dilute acids, traces of alkali catalyse immediate hydrolysis to hydrated silicon dioxide. If the reaction is carried out in methanol, controlled solvolysis results in the products SiH₂(OMe)₂, SiH(OMe)₃, and Si(OMe)₄. The Si–H bond also adds to alkenes, a reaction which proceeds slowly and speeds up with increasing substitution of the silane involved. At 450 °C, silane participates in an addition reaction with acetone, as well as a ring-opening reaction with ethylene oxide. Direct reaction of the silanes with chlorine or bromine results in explosions at room temperature, but the reaction of silane with bromine at −80 °C is controlled and yields bromosilane and dibromosilane. The monohalosilanes may be formed by reacting silane with the appropriate hydrogen halide with an Al₂X₆ catalyst, or by reacting silane with a solid silver halide in a heated flow reactor:^[26]

SiH₄ + 2 AgCl 260 °C→  SiH₃Cl + HCl + 2 Ag

Among the derivatives of silane, iodosilane (SiH₃I) and potassium silanide (KSiH₃) are very useful synthetic intermediates in the production of more complicated silicon-containing compounds: the latter is a colourless crystalline ionic solid containing K⁺ cations and SiH−
3 anions in the NaCl structure, and is made by the reduction of silane by potassium metal.^[44]Additionally, the reactive hypervalent species SiH−
5 is also known.^[26] With suitable organic substituents it is possible to produce stable polysilanes: they have surprisingly high electric conductivities, arising from sigma delocalisation of the electrons in the chain.^[45]

Halides

Silicon and silicon carbide readily react with all four stable halogens, forming the colourless, reactive and volatile silicon tetrahalides.^[46] Silicon tetrafluoride may also be made by fluorinating the other silicon halides, and is produced by the attack of hydrofluoric acid on glass.^[47] Heating two different tetrahalides together also produce a random mixture of mixed halides, which may also be produced by halogen exchange reactions. The melting and boiling points of these species usually rise with increasing atomic weight, though there are many exceptions: for example, the melting and boiling points drop as one passes from SiFBr₃ through SiFClBr₂ to SiFCl₂Br. The shift from the hypoelectronic elements in group 13 and earlier to the group 14 elements is illustrated by the change from an infinite ionic structure in aluminium fluoride to a lattice of simple covalent silicon tetrafluoride molecules, as dictated by the lower electronegativity of aluminium than silicon, the stoichiometry (the +4 oxidation state being too high for true ionicity), and the smaller size of the silicon atom compared to the aluminium atom.^[46] Silicon tetrachloride is manufactured on a huge scale as a precursor to the production of pure silicon, silicon dioxide, and some silicon esters.^[46] The silicon tetrahalides hydrolyse readily in water, unlike the carbon tetrahalides, again because of the larger size of the silicon atom rendering it more open to nucleophilic attack and the ability of the silicon atom to expand its octet which carbon lacks.^[47] The reaction of silicon fluoride with excess hydrofluoric acid produces the octahedral hexafluorosilicate anion SiF2−
6.^[47]

Analogous to the silanes, halopolysilanes Si_nX_2n + 2 are also known. While catenation in carbon compounds is maximised in the hydrogen compounds rather than the halides, the opposite is true for silicon, so that the halopolysilanes are known up to at least Si₁₄F₃₀, Si₆Cl₁₄, and Si₄Br₁₀. A suggested explanation for this phenomenon is the compensation for the electron loss of silicon to the more electronegative halogen atoms by pi backbonding from the filled p_π orbitals on the halogen atoms to the empty d_π orbitals on silicon: this is similar to the situation of carbon monoxide in metal carbonyl complexes and explains their stability. These halopolysilanes may be produced by comproportionation of silicon tetrahalides with elemental silicon, or by condensation of lighter halopolysilanes (trimethylammonium being a useful catalyst for this reaction).^[46]

Silica

Silicon dioxide (SiO₂), also known as silica, is one of the most well-studied compounds, second only to water. Twelve different crystal modifications of silica are known, the most common being α-quartz, a major constituent of many rocks such as granite and sandstone. It is also known to occur pure as rock crystal; impure forms are known as rose quartz, smoky quartz, morion, amethyst, and citrine. Some poorly crystalline forms of quartz are also known, such as chalcedony, chrysoprase, carnelian, agate, onyx, jasper, heliotrope, and flint. Other modifications of silicon dioxide are known in some other minerals such as tridymite and cristobalite, as well as the much less common coesite and stishovite. Biologically generated forms are also known as kieselguhr and diatomaceous earth. Vitreous silicon dioxide is known as tectites, and obsidian, and rarely as lechatelierite. Some synthetic forms are known as keatite and W-silica. Opals are composed of complicated crystalline aggregates of partially hydrated silicon dioxide.^[48]

• 

  Quartz

• 

  Sandstone

• 

  Smoky quartz

• 

  Amethyst

• 

  Citrine

• 

  Agate

• 

  Chrysoprase

• 

  Onyx

• 

  Jasper

• 

  Heliotrope

• 

  Tridymite

• 

  Cristobalite

• 

  Coesite

• 

  Lechatelierite

Most crystalline forms of silica are made of infinite arrangements of {SiO₄} tetrahedra connected at their corners, with each oxygen atom linked to two silicon atoms. In the thermodynamically stable room-temperature form, α-quartz, these tetrahedra are linked in intertwined helical chains with two different Si–O distances (159.7 and 161.7 pm) with a Si–O–Si angle of 144°. These helices can be either left- or right-handed, so that individual α-quartz crystals are optically active. At 537 °C, this transforms quickly and reversibly into the similar β-quartz, with a change of the Si–O–Si angle to 155° but a retention of handedness. Further heating to 867 °C results in another reversible phase transition to β-tridymite, in which some Si–O bonds are broken to allow for the arrangement of the {SiO₄} tetrahedra into a more open and less dense hexagonal structure. This transition is slow and hence tridymite occurs as a metastable mineral even below this transition temperature; when cooled to about 120 °C it quickly and reversibly transforms by slight displacements of individual silicon and oxygen atoms to α-tridymite, similarly to the transition from α-quartz to β-quartz. β-tridymite slowly transforms to cubic β-cristobalite at about 1470 °C, which once again exists metastably below this transition temperature and transforms at 200–280 °C to α-cristobalite via small atomic displacements. β-cristobalite melts at 1713 °C; the freezing of silica from themelt is quite slow and vitrification, or the formation of a glass, is likely to occur instead. In vitreous silica, the {SiO₄} tetrahedra remain corner-connected, but the symmetry and periodicity of the crystalline forms are lost. Because of the slow conversions between these three forms, it is possible upon rapid heating to melt β-quartz (1550 °C) or β-tridymite (1703 °C). Silica boils at approximately 2800 °C. Other high-pressure forms of silica are known, such as coesite and stishovite: these are known in nature, formed under the shock pressure of a meteorite impact and then rapidly quenched to preserve the crystal structure. Similar melting and cooling of silica occurs following lightning strikes, forming glassy lechatelierite. W-silica is an unstable low-density form involving {SiO₄} tetrahedra sharing opposite edges instead of corners, forming parallel chains similarly to silicon disulfide (SiS₂) and silicon diselenide (SiSe₂): it quickly returns to forming amorphous silica with heat or traces of water.^[48]

Condensed polysilicic acid

Silica is rather inert chemically. It is not attacked by any acids other than hydrofluoric acid. However, it slowly dissolves in hot concentrated alkalis, and does so rather quickly in fused metal hydroxides or carbonates to give metal silicates. Among the elements, it is attacked only by fluorine at room temperature to form silicon tetrafluoride: hydrogen and carbon also react, but require temperatures over 1000 °C to do so. Silica nevertheless reacts with many metal and metalloid oxides to form a wide variety of compounds important in the glass and ceramic industries above all, but also have many other uses: for example, sodium silicate is often used in detergents due to its buffering, saponifying, and emulsifying properties.^[48]

Silicic acids

Adding water to silica drops its melting point by around 800 °C due to the breaking of the structure by replacing Si–O–Si linkages with terminating Si–OH groups. Increasing water concentration results in the formation of hydrated silica gels and colloidal silica dispersions. Many hydrates and silicic acids exist in the most dilute of aqueous solutions, but these are rather insoluble and quickly precipitate and condense and cross-link to form various polysilicic acids of variable combinations following the formula [SiO_x(OH)_4−2x]_n, similar to the behaviour of boron, aluminium, and iron, among other elements. Hence, although some simple silicic acids have been identified in dilute solutions, such as orthosilicic acid Si(OH)₄ and metasilicic acid SiO(OH)₂, none of these are likely to exist in the solid state.^[48]