Although silicon is ubiquitous in nature as silica (sand, quartz), it was isolated late. In 1787, Antoine Lavoisier (1743-1794) suspected that silica was the oxide of an unknown element. In 1808, Humphry Davy (1778-1829) unsuccessfully attempted to isolate this element by electrolysis. It was in 1823 that Jöns Jacob Berzelius (1779-1848), a Swedish chemist, succeeded in obtaining amorphous silicon by reducing silicon tetrafluoride (SiF₄) with potassium. He named this element silicon (from the Latin silex = pebble). In 1854, Henri Sainte-Claire Deville (1818-1881) produced crystalline silicon, paving the way for the study of its semiconductor properties that would revolutionize the 20th century.
Silicon (symbol Si, atomic number 14) is a metalloid belonging to group 14 of the periodic table, in the same column as carbon. Its atom has 14 protons, 14 electrons, and usually 14 neutrons in its most abundant isotope (\(\,^{28}\mathrm{Si}\)). Three stable isotopes exist: silicon-28 (\(\,^{28}\mathrm{Si}\)), silicon-29 (\(\,^{29}\mathrm{Si}\)), and silicon-30 (\(\,^{30}\mathrm{Si}\)).
At room temperature, pure crystalline silicon is a hard, brittle solid with a metallic gray-blue color (density ≈ 2.33 g/cm³). The melting point of silicon: 1,687 K (1,414 °C). The boiling point: 3,538 K (3,265 °C). Silicon has a diamond-like crystalline structure and exhibits essential semiconductor properties for modern electronics. Its electrical conductivity increases with temperature, unlike metals.
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Silicon-28 — \(\,^{28}\mathrm{Si}\,\) | 14 | 14 | 27.976927 u | ≈ 92.23 % | Stable | Most abundant isotope; basis of the semiconductor industry. |
| Silicon-29 — \(\,^{29}\mathrm{Si}\) | 14 | 15 | 28.976495 u | ≈ 4.67 % | Stable | Used in NMR and quantum computing research. |
| Silicon-30 — \(\,^{30}\mathrm{Si}\) | 14 | 16 | 29.973770 u | ≈ 3.10 % | Stable | Enriched isotope for the redefinition of the kilogram (Avogadro sphere). |
| Silicon-32 — \(\,^{32}\mathrm{Si}\) | 14 | 18 | 31.974148 u | Cosmogenic trace | 153 years | Radioactive β\(^-\) decaying into phosphorus-32. Used to date groundwater and polar ice. |
| Other isotopes — \(\,^{22}\mathrm{Si}\) to \(\,^{44}\mathrm{Si}\) | 14 | 8 — 30 | — (variable) | Not natural | Milliseconds to hours | Unstable isotopes produced artificially; nuclear physics research. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Silicon has 14 electrons distributed across three electron shells. Its full electronic configuration is: 1s² 2s² 2p⁶ 3s² 3p², or simplified as: [Ne] 3s² 3p². This configuration can also be written as: K(2) L(8) M(4).
K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 8 electrons distributed as 2s² 2p⁶. This shell is also complete, forming a noble gas configuration (neon).
M Shell (n=3): contains 4 electrons distributed as 3s² 3p². The 3s orbitals are complete, while the 3p orbitals contain only 2 out of 6 possible electrons. Thus, 4 electrons are missing to saturate this outer shell.
The 4 electrons in the outer shell (3s² 3p²) are the valence electrons of silicon. This configuration explains its chemical properties:
By losing 4 electrons, silicon forms the Si⁴⁺ ion (oxidation state +4), its most common oxidation state, particularly in silica SiO₂ and silicates.
By gaining 4 electrons, silicon would theoretically form the Si⁴⁻ ion (oxidation state -4), a very rare state observed only in certain metallic silicides.
Silicon can also exhibit intermediate oxidation states such as +2, but +4 remains by far the most stable and widespread.
The electronic configuration of silicon, with 4 electrons in its valence shell, places it in group 14 of the periodic table, just below carbon. This structure gives it characteristic properties: the ability to form four covalent bonds by sharing its valence electrons, essential semiconductor properties in electronics, and a tendency to form tetrahedral structures similar to carbon but with weaker Si-Si bonds than C-C bonds. Silicon mainly forms polar covalent bonds, especially with oxygen, creating a vast variety of silicates that make up the bulk of the Earth's crust. Unlike carbon, which easily forms long chains, silicon prefers to form three-dimensional structures with oxygen. Its technological importance is considerable: ultra-pure silicon is the basic material of the electronics and computer industry (chips, processors, solar panels), while its natural compounds (quartz, sand) are used in the manufacture of glass, cement, and ceramics. Silicon is the second most abundant element in the Earth's crust after oxygen.
Pure silicon is relatively unreactive at room temperature due to the layer of silicon oxide (SiO₂) that forms on its surface. At high temperatures, it reacts with oxygen, halogens, and some metals. Silicon does not react with most acids (except hydrofluoric acid, which dissolves silica) but dissolves in strong bases, forming silicates. It mainly forms compounds in the +IV oxidation state, including silica (SiO₂), silicates, silanes (silicon analogs of hydrocarbons), and silicones (organic polymers of silicon). Silicon can form Si-Si, Si-O, Si-C, and Si-H bonds, giving rise to a very rich organosilicon chemistry.
Silicon is the fundamental element of the electronic and digital revolution of the 20th and 21st centuries. Its ability to be precisely doped (controlled addition of impurities) allows its electrical conductivity to be modulated, thus creating transistors and integrated circuits. Silicon Valley takes its name from this material, which enabled the progressive miniaturization of electronic components, following Moore's Law. A modern microprocessor can contain several billion transistors etched into ultra-pure silicon (99.9999999% purity). Silicon has made possible computers, smartphones, the Internet, and all the information technologies that shape our contemporary world.
Silicon is the second most abundant element in the Earth's crust (about 27.7% by mass), just after oxygen. It is never found in its pure state in nature but always combined, mainly as silica (SiO₂) in sand, quartz, and silicate rocks. Silicates make up the majority of the minerals forming terrestrial rocks (feldspars, micas, clays). Metallurgical-grade silicon is produced by reducing silica with carbon in electric arc furnaces. For electronics, ultra-pure silicon is required, obtained through complex purification and crystal growth processes (Czochralski method).
Silicon is synthesized in massive stars during the fusion of oxygen and carbon in the deep layers. During type II supernova explosions, silicon is ejected into the interstellar medium, contributing to the chemical enrichment of subsequent generations of stars and planets. Spectroscopy reveals the presence of silicon in many stars and nebulae. In the solar system, silicon is a major constituent of terrestrial planets (Mercury, Venus, Earth, Mars) and rocky asteroids. On Earth, silicon is the dominant element of the terrestrial mantle in the form of silicates, playing a crucial role in plate tectonics and geodynamics.
N.B.:
The ultra-high purity silicon required for electronics is one of the purest materials ever produced by humanity. To manufacture electronic chips, silicon must reach a purity of 99.9999999% (nine nines after the decimal), meaning it contains only one foreign atom per billion silicon atoms. This extraordinary level of purity is achieved through successive chemical purification processes, notably the distillation of trichlorosilane and the growth of single crystals by the Czochralski method, where a perfect crystal is slowly pulled from a bath of molten silicon.
