Silicon
| Silicon | |
|---|---|
| |
| Atomic number | 14 |
| Symbol | Si |
| Group | 14 (carbon group) |
| Electron configuration | [Ne] 3s2 3p2 |
| Period | 3 |
| Crystal structure | Diamond cubic |
| Phase STP | Solid |
| Block | p |
| Oxidation states | −4, +2, +4 |
| Wikidata | Q670 |
Silicon: Overview and Key Facts
Silicon (symbol Si) is the chemical element with atomic number 14. It is classified as a metalloid – an element with properties intermediate between those of metals and nonmetals. Silicon is found in Group 14 (the carbon group) of the periodic table, in period 3 and the p-block. At ordinary temperature and pressure, it is a shiny gray crystalline solid that resembles a metal but is brittle. Silicon’s valence electron configuration is [Ne] 3s²3p², meaning it has four electrons available for bonding. This tetravalence underlies its common oxidation states of +4 (as in silica, SiO₂) and −4 (as in some metal silicides), with +2 and +3 states rarely encountered. Silicon constitutes about 28% of the Earth’s crust by mass (second only to oxygen) and is also abundant in cosmic dust and meteorites. Key facts include: atomic weight ~28.09 g/mol, density ~2.33 g/cm³ (solid), melting point ~1687 K, and boiling point ~3538 K. Commonly, silicon appears in a cubic crystal structure identical to the diamond lattice of carbon, which contributes to its semiconducting properties.
Atomic Structure and Periodic Trends
A silicon atom has 14 protons and (usually) 14 neutrons in its nucleus, with 14 electrons orbiting in shells. Its full electron configuration is 1s²2s²2p⁶3s²3p²; the first ten electrons fill the neon core Ne and the outer four occupy the 3s and 3p orbitals. The four valence electrons in the third shell allow silicon to form up to four covalent bonds with other atoms, similar to carbon. As a Group 14 element, its atomic radius and electronegativity are between those of aluminum (Group 13) and phosphorus (Group 15). The covalent atomic radius of Si is about 111 picometers. On the electronegativity (Pauling) scale, silicon is about 1.90 (versus 1.61 for aluminum and 2.19 for phosphorus), reflecting its moderate ability to attract bonding electrons. Silicon’s first ionization energy (the energy needed to remove one electron) is about 786 kJ/mol (8.15 eV), again intermediate in the period. These periodic trends (radius, electronegativity, ionization energy) influence silicon’s chemical behavior: it is far less reactive than aluminum, but more so than nonmetals like carbon or sulfur.
Isotopes and Nuclear Properties
Naturally occurring silicon consists of three stable isotopes: ^28Si (about 92.2% of natural silicon), ^29Si (~4.7%), and ^30Si (~3.1%). ^28Si and ^30Si have nuclear spin 0, while ^29Si has spin-½, making ^29Si NMR-active. This nuclear spin of ^29Si is used in nuclear magnetic resonance (NMR) and quantum computing research (for example, dopants in silicon qubits often involve ^29Si spins). The abundance of ^28Si is so high that isotopically purified silicon (enriched ^28Si) is used in precision metrology; a near-pure ^28Si sphere was famously used to determine the Avogadro constant and refine the definition of the kilogram.
In addition to these stable isotopes, silicon has several radioactive isotopes. The longest-lived is ^32Si (half-life ~153 years), produced in the atmosphere by cosmic rays (through argon spallation) and by cosmic ray spallation in the upper atmosphere. ^32Si decays by beta emission to ^32P. It has been used as a tracer to date marine sediments and atmospheric processes, since it is incorporated into biogenic silica (for example, in mollusk shells and diatoms). Other radioactive isotopes like ^31Si (half-life ~2.6 hours) decay by beta^- to ^31P and are produced in reactor environments. These short-lived isotopes have niche applications in research, for instance as tracers. The long-lived ^32Si and the stable isotopes of silicon have no significant roles in natural radioactivity concerns (silicon is not a major source of background radiation).
Allotropes and Typical Compounds
Unlike carbon, which has well-known forms like diamond and graphite, elemental silicon has only one stable crystalline allotrope at ambient pressure: the diamond cubic lattice. In this structure, each Si atom is tetrahedrally bonded to four neighbors, forming a rigid 3D network. This diamond-cubic silicon is a hard, brittle semiconductor crystal (with an indirect band gap of about 1.12 eV at room temperature). There is also amorphous silicon, a non-crystalline form obtained by vapor deposition or rapid cooling, used in thin-film solar cells and electronics. Under extreme conditions (high pressure or unusual preparation), other silicon phases appear (for example, BC8, R8, or clathrate forms), but these are of academic interest and not encountered in ordinary use. Two-dimensional analogs (“silicene”) have been synthesized on metallic substrates, though they are unstable in air and exist only in research settings.
Silicon’s most common compounds are its oxides and silicates. Silicon dioxide (SiO₂) occurs in nature as quartz, sand, and many minerals; it is a robust, tetrahedral network solid that is the basis of glass and ceramics. Silicates (compounds with Si–O frameworks) make up most rocks and soils (clay, feldspar, etc.). Silicon carbide (SiC) is another major binary compound: a very hard, refractory ceramic with several crystal forms (notably 3C-SiC, 4H-SiC, 6H-SiC). SiC is used as an abrasive and as a high-voltage, high-temperature semiconductor. Silicon forms halides like silicon tetrachloride (SiCl₄) and tetrafluoride (SiF₄). SiCl₄ is a volatile liquid used in chemical synthesis and purification of silicon. Silicon hydrides (silanes) are also important compounds: silane (SiH₄) and disilane (Si₂H₆) are analogous to methane and ethane, though silanes are less stable and pyrophoric. Many organosilicon compounds exist (silicones, silanes with carbon groups), but these belong to organic chemistry. Metal silicides (e.g., FeSi₂, HfSi₂) are intermetallic compounds where silicon acts like an anion (oxidation state −4) combined with metals.
Physical Properties
Crystalline silicon is a metallic-gray solid with a shiny luster. Its density is about 2.329 g/cm³ at room temperature. As noted, it crystallizes in the diamond cubic structure. This structure and strong covalent bonding give silicon a high hardness (Mohs hardness around 6.5) and a melting point of 1687 K (1414°C). It boils at about 3538 K (3265°C) under standard pressure. Silicon is a semiconductor: pure crystalline silicon has an indirect band gap of ~1.12 eV at 300 K. This gap means that at room temperature, pure silicon conducts only very poorly; however, when doped or at elevated temperatures its conductivity rises. Silicon’s thermal conductivity at room temperature is ~148 W·m⁻¹·K⁻¹, relatively high for a non-metal. Its electrical conductivity can vary over many orders of magnitude depending on doping (with impurities) and temperature. In the infrared and visible range, silicon is opaque except at longer wavelengths (it is transparent in the near IR, which is why infrared detectors often use silicon).
Spectroscopically, atomic silicon has several emission/absorption lines in the ultraviolet; however, in everyday contexts one usually deals with silicon as a solid, where it does not exhibit well-defined spectral lines like gases do. In solid-state devices, silicon luminesces weakly when excited (usually via impurity levels), but it is not used as a direct band-gap light emitter. Instead, its indirect band nature means it emits infrared radiation inefficiently, making silicon poor for LEDs. (By contrast, silicon carbide does have a direct band gap and is used in blue LEDs.)
Chemical Reactivity and Trends
Elemental silicon is relatively unreactive at room temperature. It does not dissolve in water or react with ordinary acids such as hydrochloric or sulfuric acid under standard conditions. However, it does react with hydrofluoric acid (HF); in HF it forms fluorosilicate complexes. The net reaction with concentrated aqueous HF is approximately:
<code>Si(s) + 6 HF(aq) → H2SiF6(aq) + 2 H2(g). </code>
This gives hexafluorosilicic acid (H₂SiF₆) and hydrogen gas. The reaction proceeds because the silicon(IV) fluoride complex [SiF₆]²⁻ is very stable. With strong oxidizers at high temperature, silicon will burn to form silicon dioxide:
<code>Si + O2 → SiO2. </code>
The native oxide that forms on silicon surfaces (a few nanometers thick) is passivating: it protects the material from further corrosion (similar to aluminum oxide on aluminum). Silicon is also attacked by molten alkali: hot concentrated sodium hydroxide reacts, for example, to form soluble silicates. One simplified equation is:
<code>Si + 2 NaOH + H2O → Na2SiO3 + 2 H2. </code>
This yields sodium silicate and hydrogen gas. Such chemistry (alkaline attack) is why glass (mostly SiO2) is etched by strong alkali.
In oxidation states, silicon typically appears as +4 in oxides (SiO₂ and silicates) and +2 or +3 in suboxides (rare polymeric SiO compounds), though the latter are not common substances. Silicon can also form a formal −4 state in some metal silicides (e.g. Ca₂Si), where it behaves like an anion (similar to carbon in carbides). Silicon shows little metallic character: it does not form simple silicon cations like Si²⁺ or Si⁴⁺ in the way metals do; instead, its chemistry is largely covalent. It does not participate in acid-base chemistry as an individual element (silicates and silica are amphoteric in their reactions). Silicon can form complexes with fluoride (e.g., hexafluorosilicate) and with oxygen bridges in siloxane polymers, but it rarely forms coordination complexes like transition metals do. In the context of the reactivity series of metals, silicon would be placed far to the right (toward nonmetals): it does not displace hydrogen from acids (except HF) and is not used as a reducing metal.
Occurrence and Production
Silicon is extremely abundant on Earth. It is the second most abundant element in the Earth’s crust by weight (after oxygen), making up about 27–28%. It never occurs as a free element in nature; instead, it is found almost entirely in oxidized form. The most common sources are silica (SiO₂) found in quartz, sand, and sandstone, and silicate minerals (such as feldspars, clay minerals, and quartzites). These materials include quartz, opal, feldspar, mica, amphiboles, and many soils and rocks.
In space, silicon is one of the more abundant elements, produced by oxygen-burning processes in massive stars and dispersed in supernovae; it is a moderate-abundance element in the solar system and is common in meteorites (often as silicon carbide or troilite with iron).
Commercially, silicon is obtained by reducing high-purity silica with carbon in electric furnaces at high temperature. The carbothermic reduction process produces metallurgical-grade silicon (about 98–99% pure) and carbon monoxide gas:
<code>SiO2 + 2 C → Si + 2 CO (g). </code>
This reaction requires temperatures above ~2000°C (kiln or arc furnace). The metal produced (often called “ferrosilicon” when alloyed with iron for steelmaking, though ferrosilicon usually has only 15–90% Si) is crude and contains carbon and impurities. For semiconductor or specialty use, this metallurgical silicon is further refined. One common refinement route is to convert it to a volatile compound (silicon tetrachloride, SiCl₄) or silane (SiH₄) and then distill or react to get pure silicon again. For example, the Siemens process converts metallurgical silicon to trichlorosilane (HSiCl₃) and then decomposes it on hot rods to deposit extremely pure polycrystalline silicon. Another method is zone refining, where a polycrystalline rod is passed through a high-temperature zone that melts a small region; impurities concentrate in the melt and can be moved to one end, leaving behind ultra-pure single crystals. These techniques are used to make “electronic grade” silicon (99.9999% pure or higher).
Major producers of silicon in the world are China (by far the largest, producing several million metric tons per year, mostly metallurgical silicon and ferrosilicon), followed by Russia, the United States, Norway, Brazil, and Kazakhstan. China’s output greatly exceeds that of any other country. Producers typically focus on either metallurgical-grade silicon (for alloys, foundries, and chemical feedstock) or polysilicon (for electronics and photovoltaic industries), which earns a high price.
Applications and Technology
Silicon’s combination of abundance, physical properties, and chemistry has made it the cornerstone of modern technology, particularly in electronics and energy. About half of the silicon produced (by mass) goes into metal alloys. Ferrosilicon (iron-silicon alloy) and silicon metal (often blended with aluminum) are essential in steelmaking to deoxidize steel and improve mechanical properties. Silicon is used in cast irons to strengthen the iron matrix. Silicon-aluminum alloys are used in engine parts and other structural components. Silicon also finds use in aluminum-silicon alloys for casting.
Another huge category is construction materials. Silica and silicates are the basis of glass (windows, containers), cement (calcium silicate), ceramics, and concrete.. Sand (silicon dioxide) is melted to form glass fibers and optical fibers for data transmission. Silicates are also used in refractories (materials resisting high temperatures).
In chemicals, silicon leads to a broad class of compounds: silicones (polysiloxanes) are polymers with Si–O backbones used in sealants, lubricants, adhesives, and medical implants. They are prized for thermal stability and water resistance. Organosilicon compounds serve as intermediates in producing silicon polyurethane and epoxy materials. As a chemical feedstock, silicon yields silicon tetrachloride (used to make optical fiber and silicon for semiconductors) and fumed silica (used as a thickener and anti-caking agent).
The most societally prominent application is in electronics. Silicon is the primary semiconductor material for integrated circuits (microchips) and transistors. High-purity single-crystal silicon (the dark gray material seen in silicon wafers) forms the substrate on which billions of circuits are fabricated. Silicon’s semiconducting behavior can be precisely controlled by doping it with tiny amounts of group III elements (e.g. boron or gallium) to create “p-type” silicon (with mobile positive holes) or group V elements (e.g. phosphorus or arsenic) to create “n-type” silicon (with extra electrons). p–n junctions of these materials form diodes, transistors, and other fundamental components of electronic devices. This silicon technology underlies computers, smartphones, and most modern electronics; it also spawned the name “Silicon Valley” for the tech-centric region of California.
Silicon is also key to renewable energy: it is the main material for photovoltaic (solar) cells. Silicon solar cells (made from either single-crystal or polycrystalline silicon) convert sunlight into electricity. Its band gap is well-matched to the solar spectrum. The photovoltaic industry uses both the metallurgical and the electronic-grade silicon, though efficiencies require high purity. Research into silicon-based photovoltaic technology continues actively, including work on thin-film amorphous silicon and micromorph tandem cells.
Other applications of silicon and its compounds include: silicon carbide (SiC) semiconductor devices for high-power/high-frequency electronics, abrasives (grinding wheels, sandpaper), and LEDs (blue/gallium nitride LEDs often have a SiC substrate). Silicon nitride (Si₃N₄) is a ceramic used in bearings and engine components for its strength and heat resistance. Silicon photonics (using silicon chips to manipulate light for data communication) is a growing field. Also, silane gas (SiH₄) and reaction with silicon surfaces are used in chemical vapor deposition processes to produce thin films of silicon.
Biology, Environment, and Safety
Silicon is not considered an essential element for most animals and humans, though it is abundant in nature. It enters organisms primarily as silicic acid (H₄SiO₄) dissolved from soil and plants. Certain organisms do use silicon biologically: many plants (particularly grasses and cereals) accumulate silica in their tissues as protective stiffness; freshwater diatoms and some sponges build microscopic structures out of biogenic silica (amorphous SiO₂). In humans and animals, silicon is present in trace amounts in connective tissues and bones, and some studies suggest it may play a role in bone health or collagen formation, but it is not classified as a nutrient. The frontiers of biology consider silicon as a potential alternative biochemistry (silicon-based life), but no such life is known on Earth, and carbon remains vastly more versatile for biochemistry.
In the environment, silicon cycles through rock weathering and sedimentation. Silicate minerals slowly dissolve in water, releasing silicic acid into rivers and oceans, where it is taken up by marine organisms. When these organisms die, the silica may settle and eventually form sedimentary rocks like radiolarian ooze. Industrially, mining silica sand for glass or polysilicon production alters landscapes and can produce particulate dust.
Safety-wise, elemental silicon is quite inert and has low toxicity; it is neither flammable nor particularly hazardous in bulk. However, inhalation of fine silica dust (from quartz, sandblasting, stone cutting, or production of microelectronic wafers) is a serious health risk. Respirable crystalline silica can cause silicosis, a progressive lung disease, and is a known carcinogen. Occupational exposure limits for silica dust are very low (parts per million by mass). Powdered pure silicon (fine powder) can also be hazard: silicon powder can ignite when dispersed in air (dust explosion hazard) or react with strong oxidizers. Silane (SiH₄) gas, used in chip manufacturing, is pyrophoric (ignites on contact with air) and toxic. Some silicon compounds (like silicon tetrachloride) are corrosive and release HCl upon hydrolysis. Therefore, handling of reactive silicon compounds requires precautions (gloves, ventilation, etc.). Overall, while bulk silicon materials (crystals, blocks) are safe to handle, one must control dust and reactive chemicals in industrial and lab settings.
History and Etymology
The conjecture that silica is an oxide of a fundamental element dates back to Antoine Lavoisier in 1787. French chemist Joseph-Louis Gay-Lussac and Louis-Jacques Thénard in 1811 heated potassium and silicon tetrachloride and produced amorphous silicon, but they only identified it as a new substance. In 1823–1824, Scottish chemist Sir Humphry Davy attempted to isolate silicon by electrolysis (unsuccessfully). That same year, Swedish chemist Jöns Jacob Berzelius and French chemist Gay-Lussac and Thénard independently succeeded: Berzelius heated potassium metal with silicon tetrafluoride (though he incorrectly named it “silicium”) and obtained amorphous material. The element was named “silicium” by Berzelius after the Latin word silex or silicis (flint). In 1814, English chemist Thomas Thomson suggested the modern name “silicon” to conform to the “-on” used for elements like carbon.
Pure crystalline silicon was first prepared by Deville and Troost in 1854 by heating magnesium with silicon tetrachloride. Work on silicon accelerated in the late 19th and early 20th centuries, especially as physicists and engineers began to explore semiconductors. The native oxide layer on silicon was noted early (traces seen when silicon metal was prepared). In the 20th century, Germanium was the first semiconductor used in transistors (band gap ~0.67 eV), but silicon took over after the 1950s because its oxide, SiO₂, could be grown into an excellent insulating layer – a key advantage for making reliable transistors and integrated circuits (the “oxide” in MOSFET and CMOS stands for silicon dioxide). In 1954, Bell Labs announced the first working silicon transistor. Soon, chemists perfected methods like the Czochralski process (growing single silicon crystals from the melt) and zone refining to produce extremely pure silicon crystals.
Throughout the late 20th century, silicon became synonymous with technology: microprocessors, memory chips, and solar cells. The term “Silicon Valley” (coined in the 1970s for Santa Clara Valley, California) reflects silicon’s role. Silicon’s name and symbol derive from “silex” and the suffix “-on”, emphasizing it as an element. It has appeared culturally in phrases like “silicon-based life” (as a hypothetical alternative to carbon-based life) and in company and product names (e.g., Silicon Graphics, Silicon wafer).
Data Table
| Property | Value and Description |
|---|---|
| Atomic number (Z) | 14 |
| Symbol | Si |
| Atomic weight | 28.0855 g/mol |
| Group / Period / Block | 14 (carbon group) / Period 3 / p-block |
| Element category | Metalloid |
| Standard state | Solid (gray crystalline) |
| Crystal structure | Diamond cubic (Si each tetrahedrally coordinated) |
| Density (at 293K) | 2.3296 g·cm⁻³ |
| Melting point | 1687 K (1414°C) |
| Boiling point | 3538 K (3265°C) |
| Specific heat (300K) | ≈ 0.71 J·g⁻¹·K⁻¹ |
| Thermal conductivity | ≈ 148 W·m⁻¹·K⁻¹ (crystalline, 300K) |
| Electron configuration | [Ne] 3s² 3p² |
| Valence electrons | 4 |
| Atomic radius (covalent) | ≈ 111 pm |
| Electronegativity (Pauling) | 1.90 |
| First ionization energy | 786 kJ·mol⁻¹ (8.15 eV) |
| Common oxidation states | –4 (in silicides), +2 (rare; suboxides), +4 (oxides) |
| Stable isotopes | ^28Si (~92.23%), ^29Si (~4.67%), ^30Si (~3.10%) |
| Notable radioisotope | ^32Si (t₁/₂ ≈ 153 years; β⁻ decay to ^32P) |
| Magnetic spins (nuclear) | ^28Si (spin 0), ^29Si (½), ^30Si (0) |
| Band gap (indirect) | ~1.12 eV (at 300 K) |
Each entry above gives a concise fact or range for quick reference. The stable isotope abundances and nuclear spins are included because they are relevant for certain technologies (e.g. NMR and quantum applications). The data table summarizes silicon’s core constants without reference citations, as requested.