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What is silicon?
Click:1552 Date:2012-11-7 15:16:48

Silicon is an word known to the public, at least in the phrase "Silicon Valley," and is associated with computers, although the public probably does not know what a chemical element is, or how silicon is used in computers, or where it comes from, or how it behaves. The earth is constructed of silicon, and we use it every day in the form of glass and pottery. What carbon is to the living world, silicon is to the non-living. Its study is, therefore, interesting and wide-ranging, leading into many areas of technology and science. In this article, I shall strive to present the fundamentals of the most important of these areas, so that a useful understanding of silicon can be developed in those who seek it. Not only is there a great amount of information about silicon, but study brings up new questions at every turn. I am certain that somebody knows the answers to most of these questions, but I do not, and must leave them for further research. Indeed, the authorities often do not agree, figures may vary, and ignorance is glossed over in many cases. I shall do the best I can, and admit my ignorance where it appears.


In this large article, I shall first discuss the physical and chemical properties of silicon and its oxide silica. If it were not for the semiconductor applications of pure silicon, silica would be by far the more important topic. The earth is more precisely constructed of silica; except possibly for the core, the whole earth is mainly silica, with a few impurities. Then, the use of elemental silica as a semiconductor will be discussed, and how several familiar semiconductor devices made from it work. After this, the large field of silica minerals and their structures will be explained. Then, the relation of silica to life will be explored, concentrating on the diatom and the diatomite it has left, which is an important industrial mineral. Silicate minerals weather to clay, which not only forms the soils in which plants grow, but also gives us pottery, bricks and cement. Silicon also gives us glass, a substance that is probably always in our view in daily life, and which is not only useful, but also beautiful. The lore of silica is a rich garden that yields many tasty fruits. The links in the Contents will allow you to sample any subject that interests you. This article is brand new, and will be revised to eliminate errors and misprints, and to include new things that I learn.


The name silicon comes from the Latin silex, silicis, the word for "paving stone." The genitive is given to show where the "c" comes from. Flints made a very good, durable paving stone, so the word came to be associated with flints. Flints are made of silicon dioxide, so the name was acquired by this important substance, "silica." Silica is very difficult to decompose into silicon and oxygen, so was long supposed to be a simple substance, an "element," along with similar compounds, such as alumina, Al2O3, or magnesia, MgO. As sand, it has been known from time immemorial. The Greek for sand, psammites, and the Latin, arena, have not found their ways into chemical terminology as the flint has.


When Humphrey Davy decomposed soda and potash into metallic sodium and potassium by electrolysis just after 1800, powerful reducing agents became available. A reducing agent is a reagent that supplies electrons, which can turn a metal ion into the metal itself, as Fe++ + 2e- → Fe. Potassium is just about the most powerful reducing agent that exists, and can be used to reduce nearly every other element. Davy, and in Paris Gay-Lussac and Thénard, competed to decompose previously undecomposable substances and discover new elements. Davy won with boron in 1809, closely followed by the French pair. Though Davy suspected silica was a compound, Gay-Lussac and Thénard won the race for silicon in 1811. What they got was not all that impressive, and Berzelius, in 1824, finally exhibited relatively pure silicon, and is credited with the discovery.


What these workers did was dissolve silica in hydrofluoric acid (the only acid that attacks it) to get silicon fluoride: SiO2 + 4HF → SiF4 + 2H2O. The gaseous silicon fluoride was then passed through molten potassium, where the silicon was reduced: SiF4 + 4K → Si + 4KF. The potassium fluoride could then be washed away, leaving a brown powder, amorphous silicon. It is not actually amorphous, just microcrystalline. The actual reactions are not this neat, and other compounds are formed, such as fluosilicates, which means that the silicon is not very pure. It was not until 1854 that pure, crystalline silicon was produced. This is a silvery-gray, metallic-looking substance that is, nevertheless, rather brittle and low in density.


Most silicon now is made by reduction of SiO2 with C in the electric furnace. With carefully selected pure sand, the result is commercial brown silicon of 97% purity or better. This is the silicon used for semiconductors, but it must be further purified to bring impurities below the parts-per-billion level. Both sand and the brown silicon are the starting points for the synthesis of silicon compounds.


Silicon belongs to the IVA family in the periodic table, which consists of carbon, silicon, germanium, tin and lead. Its atomic number is 14, and its atomic weight is 28.086. Its isotope Si28 has an abundance of 92%. It melts at 1440°C, and boils at 2355°C. Its density is 2.36 g/cc. The outer electron configuration of each of these atoms is s2p2. Carbon and silicon are nonmetals, white tin and lead are metals, and germanium and gray tin are typical semiconductors. Of course, the important use of elemental silicon is as a semiconductor, so this matter will be discussed in its own section below. These five atoms are very different in their physical and chemical behaviors, despite being in the same column in the periodic table.


Silicon is often compared to carbon, and differences are what is significant, not similarities. Each has a valence of 4, but there the similarity ends. Carbon will share one, two or three electrons with another carbon, forming single, double and triple covalent bonds. The lengths of these bonds are typically 0.154, 0.133 and 0.120 nm. The radius of a carbon atom can be taken as 0.077, based on the single bond. Multiple bonds are also formed with other atoms, such as N, O, S and P. This leads to the great richness of carbon chemistry. Carbon will not let go of its four valence electrons, but is willing to share them.


 Silicon, on the other hand, will share an electron with another silicon, showing a radius of 0.117 nm, but will not make multiple bonds. The fact that it is about 50% larger than a carbon atom makes the difference. Silicon, likewise, will not let go of its four valence electrons, and does not even like to share them very much if the alternative of sharing them with an oxygen atom is possible. Both carbon and silicon find that hybrid sp3 tetrahedral orbitals are the most stable, with the angle between bonds equal to 109° 28', the tetrahedral angle. Silicon prefers above all else to surround itself with four oxygen atoms as if in the orthosilicate ion (SiO4)----. This ion actually exists in water solution, forming the weakly ionized orthosilicic acid, H4SiO4, with ionization constant K1 = 2 x 10-10. The orthosilicate ion, which we shall call the silica tetrahedron, is shown in the diagram at the right. All you can see of it are the oxygen atoms; the silicon is safe in the middle.


Molecules of orthosilicic acid can condense at their vertices by the elimination of H2O to form Si-O-Si linkages. Each silcon atom is surrounded by four oxygens, in each of which it has a half-interest, so the composition is expressed by SiO2, the usual formula given for silica. Note carefully that there are no SiO2 molecules here, just a giant macromolecule that is the crystalline silica. If not given a lot of time, the tetrahedra will condense in a sort of mess that is not actually crystalline, but a glass. With time and luck, or if allowed to grow on a pattern, the tetrahedra will arrange themselves in the form of a crystal, perhaps the hexagonal crystal of the mineral quartz. The density of quartz is 2.65 g/cc. Silicon specializes in these large macromolecules, of which the silica tetrahedron is the building block.


The analogous compounds CO2 and SiO2 are very different. The first is a gas, or a soft solid at low temperatures, always consisting of CO2 molecules, in which the carbon and oxygen are connected by double bonds. It is soluble in water, forming carbonate ions. The second is a hard, crystalline solid not softened by heat below 1700°C, and very resistant to chemical attack. Silica is one large molecule, not an assembly of SiO2 molecules. I have not heard exactly what happens in the gaseous state.


If there is no oxygen around, the silicon atom must be satisfied by others of its own kind. The Si-Si covalent bond is 0.117 nm long, and is strong. The crystal formed is the same as in diamond, a face-centered cubic lattice. Its density is 2.36 g/cc. Large crystals are not as reactive as the brown powder of amorphous silicon (which is also fcc, but finely divided). The crystal is silvery gray, and shiny like a metal, but is lighter that you would expect from its appearance. One handbook mentions a "graphitic" silicon, but all other sources I have consulted are silent on this subject. If it were not for the electrons, pure crystalline silica would be transparent.


Silicon can be convinced to combine with halogens and hydrogen if there is no oxygen around. This combination must usually be brought about indirectly, and water must be excluded. Silicon forms silane, SiH4, which is analogous to methane, CH4. However, the two gases are quite different. Methane can be bubbled through water, but silane is immediately hydrolyzed: SiH4 + 3H2O → H2SiO3 + 3H2, since silicon has a hunger for oxygen. There are analogues to the alkanes, such as Si2H6, Si3H8, and so on, where there are Si-Si bonds. The last-mentioned compound, and all higher ones, are pyrophoric, bursting into flame even with only atmospheric oxygen available.


Treatment of brown silicon with chlorine makes SiCl4. This substance melts at -70°C, and boils at 60°C. It is tetrahedral, with bond length 0.201 nm, slightly less than expected. The contraction of about 0.016 nm in the bond length is evidence of additional stability conferred by resonance with slightly ionic structures in which electrons are moved from chlorine to chlorine. In water, SiCl4 + 3H2O → H2SiO3 + 4HCl. If SiCl4 is mixed with NH4OH, we get a dense white smoke of the metasilicic acid and ammonium chloride, that is used for military purposes and in skywriting. SiCl4 liquid is nonconducting, showing that it is not ionic.


SiF4 is curious in that its melting point is -90°C, but its "boiling point" is -95°C. What this means is that the solid sublimes to SiF4 molecules before it melts. The solid does not consist of SiF4 molecules, but of a giant molecule in silicon style. SiF4 is covalently bonded, and does not give fluoride ions in solution--instead, the molecule hydrolyzes, as we would expect, making metasilicic acid and fluosilicic acid (instead of hydrofluoric). Fluosilicic acid, H2SiF6, is a strong acid where the silicon is in the center of an octahedron of fluorines, instead of a tetrahedron of oxygens. The silicofluoride ions are stable in solution, and form salts. The silicofluoride octahedron is the only thing that can be on equal terms with the silica tetrahedron. SiF4 is evolved when silica dissolves in HF.


There are also numerous compounds with both hydrogen and halogens, such as trichlorosilane, SiHCl3. This compound is made when brown silicon is treated with anydrous HCl. It decomposes on a hot surface to elemental silicon, chlorine and hydrochloric acid. In the purification of silicon, it can be distilled to high purity, and then decomposed to silicon with impurities at the parts-per-billion level, satisfactory for many purposes.

Sand will also react directly with carbon in the electric furnace to produce silicon carbide, SiC, or carborundum: SiO2 + 3C → SiC + 2CO. This reaction proceeds because the gaseous carbon monoxide expelled. Commercial carborundum is black and impure, but still is very hard, 9.2 Mohs. When pure, it is a transparent, greenish crystal. It occurs in two modifications, α-SiC which is hexagonal in structure, and β-SiC, which is face-centered cubic, the diamond structure, and apparently the hardest form. Its electrical conductivity is 107-200 Ω-cm, so it can be used for electrodes. Its density is 3.217 g/cc, and it sublimes at 2700°C. Carborundum was discovered by E. G. Acheson in 1891, and was the first of the artificial abrasives.


Ferrosilicon, an alloy of iron and silicon, is produced if iron oxide is added to the sand and coke in the electric furnace. This alloy does not have to be especally pure, since it is thrown into molten iron to deoxidize it. The silicon combines with any oxygen present, as we might expect. Extra silicon remains and alloys with the iron. A few percent gives an excellent steel for magnetic cores, and in cast iron it aids the castability.


Sodium oxide, Na2O, reacts rapidly with water to make the strong base NaOH: Na2O + H2O → 2NaOH. NaOH dissociates completely in water to give OH- ion, and because of this is called a base. Sulphur trioxide, SO3, also reacts rapidly with water to make the strong acid H2SO4: SO3 + H2O → H2SO4. This molecule dissociates in water to give H+, and because of this is called an acid. Chemistry texts often point out that this ion is hydrated. This is no surprise: all ions in water are hydrated. If we mix the two, we get sodium sulphate and water, which is neither acidic nor basic. Therefore, we call sodium basic and sulphur acidic, representing two principles that will react with each other.


Calcium oxide, lime, CaO, gives Ca(OH)2 in water. Incidentally, this is called slaking the lime and releases much heat. Calcium hydroxide is a weak base, but we have no trouble classifying calcium as basic. Silicon oxide, SiO2, does not actually react with water, but we can imagine that it would form H2SiO3, a feeble acid, and so classify silica as acid. When the two meet in a fused state, they react to form Ca2SiO3, which has a lower melting point than either reagent. In general, acidic and basic oxides react to give a slag that is usually of relatively low melting point. This reaction is used to remove silicates in the smelting of iron. Limestone is added, which burns to lime, evolving carbon dioxide, and then the lime reacts with the silica to form a slag in which the iron is insoluble, and which can be drawn off separately. The use of the terms acid and base in this case, where there is no water to be seen, can be made reasonable by the preceding argument.


Orthosilicic acid can lose water to form metasilicic acid, H2SiO3, which we have encountered several times already. Silica is attacked by alkalis, SiO2 + 2NaOH → Na2SiO3 + H2O. If we now treat this with HCl, we get Na2SiO3 + 2Hcl → H2SiO3 + 2NaCl. The metasilicic acid is here called "water glass," since it forms a gelatinous precipitate that is insoluble in water. Water glass is actually formed from chains of silica tetrahedra in an irregular and indefinite pattern. The formula gives only the composition, not the molecule. This is, as we realize, typical of silicon. This gel can be used to seal the pores in eggs to preserve them longer, or as an adhesive. If heated, water is driven out of the abundant pore space, and the result is the familiar silica gel used as a desiccant. When it gets full of water, it can be renewed by simple heating. Sodium silicate is made commercially by fusing sodium carbonate, sand and carbon. Water glass is then produced by acidifying the product.


Silicon makes giant molecules, as we have seen. To get some perspective on atomic sizes, consider a picogram (a millionth of a microgram) of silica. This would fit in a cube 722 nm on a side, so it would be at the limit of the optical microscope if you wanted to see it, and almost as small as a colloidal particle. However, this speck would still contain 1010 SiO2 units, 10,000 million of them! Silica vapor could easily consist of clumps of SiO2 units, say hundreds of them in each fragment, but I have not heard of anyone who has investigated. The boiling point of 2230-2590°C makes measurements a little difficult. Crystalline silicon has 5 x 1022 atoms per cc.


There is another oxide of silicon, silicon monoxide, SiO. It forms cubic crystals of density 2.13 g/cc that are transparent, with index of refraction 2.0. Polycrystalline SiO is used as an optical thin film, usually to improve reflectivity. It melts somewhere above 1700°C, and boils at 1880°C. I have seen little information on its structure and technology. Apparently it can autooxidize according to 2SiO → Si + SiO2.

A silicone is a silicon-oxygen chain with hydrocarbon radicals attached to the silicons, as in methylsilicone, (CH3)3Si-O-(CH3)2Si-O-...-O-Si(CH3)3. The hydrocarbons make the molecule look like a hydrocarbon to its surroundings, while the strong silicon-oxygen chain makes it very stable at high temperatures. The viscosity of a silicone oil increases much less rapidly than that of a hydrocarbon oil when the temperature decreases, say a factor of only 70 from 100°C to -35°C, while a hydrocarbon oil's would increase by a factor of 1800. The chains can be cross-linked and polymerized by oxygen or other means to give rubbery solids that are equally inert. These remarkable compounds are strictly artificial, never found in nature.


Silicon and its compounds are not poisonous or otherwise hazardous, except for rare compounds like the silanes. Most, indeed, are almost completely insoluble. One exception is the industrial disease silicosis, which affects workers exposed to silica dust over a considerable period. The workers involved are mainly stoneworkers and miners, and the disease can be prevented by wearing respiratory filters. Some years ago, a Denver TV reporter found to her horror that children's playgrounds often contained deadly silica, and that nobody was doing anything about it.


Learn more about silicon from Wikipedia Please click here


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