Silicon – facts and properties

Position in the periodic table

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Occurrence and extraction

Silicon is the second most abundant element in the Earth's crust after oxygen, accounting for more than a quarter of its mass. It never occurs uncombined. Most of it is bound up in quartz, which is silicon dioxide and the main component of sand, and in the silicates, a huge family of minerals that includes feldspars, micas, clays and many gemstones. Living things use it too: diatoms build their delicate shells from silicon dioxide, and horsetails stiffen their stems with it.

Industrially, quartz is reduced with carbon in electric arc furnaces at very high temperatures. This gives metallurgical-grade silicon, pure enough for alloys but far too impure for electronics. To reach the purity that chips require, the silicon is converted into trichlorosilane, a liquid that can be purified by distillation, and then deposited again as ultra-pure polycrystalline silicon. Large single crystals are pulled from the melt and sawn into thin slices called wafers. Jöns Jacob Berzelius obtained impure, powdery silicon in the eighteen-twenties, and Henri Sainte-Claire Deville prepared crystalline silicon in 1854. The name comes from the Latin silex, meaning flint.

Properties and reactions

Silicon has the electron configuration [Ne] 3s² 3p² and four valence electrons. It lies in period 3 and group 14, directly below carbon, and like carbon it prefers to form four bonds arranged in a tetrahedron. In the crystal every silicon atom is linked to four neighbours in the same pattern as the carbon atoms in diamond. Its electronegativity is 1.9.

Pure silicon is dark grey with a metallic sheen, hard and brittle. Its density is 2.33 g/cm³; it melts at 1414 °C and boils at 3265 °C. Electrically it is a semiconductor: at low temperature it barely conducts, but heat or light frees more electrons and its conductivity rises. Adding tiny amounts of other elements, a process called doping, changes this dramatically. Boron produces p-type silicon and phosphorus produces n-type silicon, and the junctions between the two are the basis of diodes, transistors and solar cells.

At room temperature silicon is fairly unreactive, protected by a thin film of silicon dioxide. A mixture of hydrofluoric and nitric acids, as well as hot alkalis, attacks it, and at high temperatures it burns to silicon dioxide and reacts with chlorine to give silicon tetrachloride. The bond between silicon and oxygen is exceptionally strong, which explains why silicon on Earth is almost always found combined with oxygen.

Uses in everyday life and technology

By tonnage, most silicon goes into metallurgy. Ferrosilicon removes oxygen from molten steel, and aluminium-silicon alloys flow so well when molten that they are ideal for casting engine blocks and wheels. Another large share is used by the chemical industry, above all to make silicones.

The best-known application is in electronics. Processors and memory chips containing billions of tiny transistors are built on wafers of highly purified silicon. Silicon solar cells convert sunlight into electricity and supply a fast-growing share of the world's energy.

Silicon and silicone are easy to confuse but quite different. Silicones are synthetic polymers whose backbone consists of alternating silicon and oxygen atoms carrying organic side groups. They are found in sealants, bakeware, gaskets, lubricants and cosmetics, but they are compounds, not the element. Other silicon compounds shape everyday life as well: glass is made by melting quartz sand, cement, bricks and ceramics consist largely of silicates, and silicon carbide is a hard abrasive and a material for power electronics.

Molecules containing this element

• Silane SiH₄View in 3D
• Silicon carbide SiCView in 3D
• Disiloxane H₃SiOSiH₃View in 3D

Quick question

Which of these elements shows the strongest metallic character?

FAQ