Editing Silicon (section)

===Physical and atomic===
[[File:Silicon-unit-cell-3D-balls.png|thumb|upright=0.8|Silicon crystallizes in a [[diamond]] [[Cubic crystal system|cubic crystal structure]] by forming [[orbital hybridization|sp<sup>3</sup> hybrid orbitals]].<ref>{{Cite web|title=Silicon and Germanium|url=http://hyperphysics.phy-astr.gsu.edu/hbase/Solids/sili2.html|access-date=2021-06-07|website=hyperphysics.phy-astr.gsu.edu}}</ref>
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A silicon atom has fourteen [[electrons]]. In the ground state, they are arranged in the electron configuration [Ne]3s<sup>2</sup>3p<sup>2</sup>. Of these, four are [[valence electron]]s, 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 rule|octet]] and obtain the stable [[noble gas]] configuration of [[argon]] by forming [[orbital hybridization|sp<sup>3</sup> hybrid orbitals]], forming tetrahedral {{chem|SiX|4}} derivatives where the central silicon atom shares an electron pair with each of the four atoms it is bonded to.<ref name="King0">{{harvnb|King|1995|pp=xiii–xviii}}</ref> The first four [[ionisation energy|ionisation energies]] of silicon are 786.3, 1576.5, 3228.3, and 4354.4&nbsp;kJ/mol respectively; these figures are high enough to preclude the possibility of simple cationic chemistry for the element. Following [[periodic trend]]s, its single-bond covalent radius of 117.6&nbsp;pm is intermediate between those of carbon (77.2&nbsp;pm) and germanium (122.3&nbsp;pm). The hexacoordinate ionic radius of silicon may be considered to be 40&nbsp;pm, although this must be taken as a purely notional figure given the lack of a simple {{chem|Si|4+}} cation in reality.{{sfn|Greenwood|Earnshaw|1997|pp=372}}

====Electrical====
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 ([[band 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 the [[valence and conduction bands]] and is the energy at which a state is as likely to be occupied by an electron as not. Hence pure silicon is effectively an insulator at room temperature. However, [[Doping (semiconductor)|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 [[Extrinsic semiconductor#N-type semiconductors|n-type semiconductor]]. Similarly, doping silicon with a [[boron group|group 13 element]] such as [[boron]], [[aluminium]], or [[gallium]] results in the introduction of acceptor levels that trap electrons that may be excited from the filled valence band, creating a [[Extrinsic semiconductor#P-type semiconductors|p-type semiconductor]].{{sfn|Greenwood|Earnshaw|1997|p=331}} <!-- this could be moved to electronic applications --> 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 current 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.{{sfn|Greenwood|Earnshaw|1997|p=331}}

====Crystal structure====
{{Main|Allotropes of silicon}}
Silicon crystallises in a giant covalent structure at standard conditions, specifically in a [[diamond cubic]] crystal lattice ([[:Category:Crystals in space group 227|space group 227]]). It thus has a high melting point of 1414&nbsp;°C, as a lot of energy is required to break the strong covalent bonds and melt the solid. Upon melting silicon contracts as the long-range tetrahedral network of bonds breaks up and the voids in that network are filled in, similar to water ice when hydrogen bonds are broken upon melting. It does not have any thermodynamically stable 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&nbsp;[[gigapascal]]s known as Si–VII (the standard modification being Si–I). An allotrope called BC8 (or bc8), having a [[body-centred cubic]] lattice with eight atoms per primitive unit cell ([[:Category:Crystals in space group 206|space group 206]]), can be created at high pressure and remains metastable at low pressure. Its properties have been studied in detail.<ref>{{cite journal |last1=Vladimir E. Dmitrienko and Viacheslav A. Chizhikov |title=An infinite family of bc8-like metastable phases in silicon |journal=Phys. Rev. B |year=2020 |volume=101 |issue=24 |page=245203 |doi=10.1103/PhysRevB.101.245203 |arxiv=1912.10672 |bibcode=2020PhRvB.101x5203D |s2cid=209444444 }}</ref>

Silicon boils at 3265&nbsp;°C: this, while high, is still lower than the temperature at which its lighter congener [[carbon]] sublimes (3642&nbsp;°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.{{sfn|Greenwood|Earnshaw|1997|p=331}}

It is also possible to construct [[silicene]] layers analogous to [[graphene]].<ref name="Aufray2010" /><ref name="Lalmi2010" />