Editing Silicon (section)

==Chemistry and compounds==
{{Main|Silicon compounds}}

{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ style="margin-bottom: 5px;" | C–X and Si–X bond energies (kJ/mol){{sfn|Greenwood|Earnshaw|1997|p=330}}
|-
! {{nowrap|X {{=}}}}
! C
! Si
! H
! F
! Cl
! Br
! I
! O–
! N<
|-
! {{nowrap|C–X}}
| 368
| 360
| 435
| 453
| 351
| 293
| 216
| ~360
| ~305
|-
! {{nowrap|Si–X}}
| 360
| 340
| 393
| 565
| 381
| 310
| 234
| 452
| 322
|}
Crystalline bulk silicon is rather inert, but becomes more reactive at high temperatures. Like its neighbour aluminium, silicon forms a thin, continuous surface layer of [[silicon dioxide]] ({{chem|SiO|2}}) that protects the material beneath from oxidation. Because of this, silicon does not measurably react with the air below 900&nbsp;°C. Between 950&nbsp;°C and 1160&nbsp;°C, the formation rate of the [[vitreous lustre|vitreous]] dioxide rapidly increases, and when 1400&nbsp;°C is reached, atmospheric [[nitrogen]] also reacts to give the nitrides SiN and {{chem|Si|3|N|4}}. Silicon reacts with gaseous [[sulfur]] at 600&nbsp;°C and gaseous [[phosphorus]] at 1000&nbsp;°C. This oxide layer nevertheless does not prevent reaction with the [[halogen]]s; [[fluorine]] attacks silicon vigorously at room temperature, [[chlorine]] does so at about 300&nbsp;°C, and [[bromine]] and [[iodine]] at about 500&nbsp;°C. Silicon does not react with most aqueous acids, but is oxidised and complexed by [[hydrofluoric acid]] mixtures containing either [[chlorine]] or [[nitric acid]] to form [[Hexafluorosilicic acid|hexafluorosilicates]]. It readily dissolves in hot aqueous alkali to form [[silicate]]s.<ref>{{Citation|last1=Stapf|first1=André|title=Wafer Cleaning, Etching, and Texturization|date=2019|url=http://link.springer.com/10.1007/978-3-662-56472-1_17|work=Handbook of Photovoltaic Silicon|pages=311–358|editor-last=Yang|editor-first=Deren|place=Berlin, Heidelberg|publisher=Springer Berlin Heidelberg|language=en|doi=10.1007/978-3-662-56472-1_17|isbn=978-3-662-56471-4|access-date=2021-03-07|last2=Gondek|first2=Christoph|last3=Kroke|first3=Edwin|last4=Roewer|first4=Gerhard|s2cid=226945433}}</ref> At high temperatures, silicon also reacts with [[alkyl halide]]s; this reaction may be catalysed by [[copper]] to directly synthesise [[organosilicon]] chlorides as precursors to [[silicone]] polymers. Upon melting, silicon becomes extremely reactive, alloying with most metals to form [[silicide]]s, and reducing most metal oxides because the [[heat of formation]] of silicon dioxide is so large. In fact, molten silicon reacts virtually with every known kind of crucible material (except its own oxide, {{chem|SiO|2}}).<ref name="Grabmaier-1982">{{Cite book |last=Grabmaier |first=J. |url=https://www.worldcat.org/oclc/840294227 |title=Silicon Chemical Etching |date=1982 |publisher=Springer Berlin Heidelberg |isbn=978-3-642-68765-5 |location=Berlin, Heidelberg |oclc=840294227}}</ref>{{Rp|page=13}} This happens due to silicon's high binding forces for the light elements and to its high dissolving power for most elements.<ref name="Grabmaier-1982" />{{Rp|page=13}} As a result, containers for liquid silicon must be made of [[refractory]], unreactive materials such as [[zirconium dioxide]] or group 4, 5, and 6 borides.{{sfn|Greenwood|Earnshaw|1997|p=331}}<ref>{{harvnb|Greenwood|Earnshaw|1997|pp=331–5}}</ref>

Tetrahedral coordination is a major structural motif in silicon chemistry just as it is for carbon chemistry. However, the 3p subshell is rather more diffuse than the 2p subshell and does not hybridise so well with the 3s subshell. As a result, the chemistry of silicon and its heavier congeners shows significant differences from that of carbon,<ref name="Kaupp">{{cite journal |last=Kaupp |first=Martin |date=1 December 2006 |title=The role of radial nodes of atomic orbitals for chemical bonding and the periodic table |url=http://depa.fquim.unam.mx/amyd/archivero/LecturasobreNodosRadiales_12854.pdf |journal=Journal of Computational Chemistry |volume=28 |issue=1 |pages=320–325 |doi=10.1002/jcc.20522 |pmid=17143872 |s2cid=12677737 |access-date=14 October 2016 |doi-access=free }}</ref> and thus octahedral coordination is also significant.{{sfn|Greenwood|Earnshaw|1997|p=331}} For example, the [[electronegativity]] of silicon (1.90) is much less than that of carbon (2.55), because the valence electrons of silicon are further from the nucleus than those of carbon and hence experience smaller electrostatic forces of attraction from the nucleus. The poor overlap of 3p orbitals also results in a much lower tendency toward [[catenation]] (formation of Si–Si bonds) for silicon than for carbon, due to the concomitant weakening of the Si–Si bond compared to the C–C bond:<ref name="King43">{{harvnb|King|1995|pp=43–44}}</ref> the average Si–Si bond energy is approximately 226&nbsp;kJ/mol, compared to a value of 356&nbsp;kJ/mol for the C–C bond.{{sfn|Greenwood|Earnshaw|1997|pp=374}} This results in multiply bonded silicon compounds generally being much less stable than their carbon counterparts, an example of the [[double bond rule]]. On the other hand, the presence of radial nodes in the 3p orbitals of silicon suggests the possibility of [[hypervalence]], as seen in five and six-coordinate derivatives of silicon such as {{chem|SiX|5|-}} and {{chem|SiF|6|2-}}.<ref>{{cite journal
| last1   =Kaupp
| first1  =Martin
| title   =The role of radial nodes of atomic orbitals for chemical bonding and the periodic table
| journal =Journal of Computational Chemistry
| volume  =28
| issue   =1
| year    =2007
| pages   =320–325
| issn    =0192-8651
| doi     =10.1002/jcc.20522
| pmid  =17143872
| doi-access=
}}</ref><ref name="King43" /> Lastly, because of the increasing energy gap between the valence s and p orbitals as the group is descended, the divalent state grows in importance from carbon to lead, so that a few unstable divalent compounds are known for silicon; this lowering of the main oxidation state, in tandem with increasing atomic radii, results in an increase of metallic character down the group. Silicon already shows some incipient metallic behavior, particularly in the behavior of its oxide compounds and its reaction with acids as well as bases (though this takes some effort), and is hence often referred to as a [[metalloid]] rather than a nonmetal.<ref name="King43" /> Germanium shows more, and tin is generally considered a metal.{{sfn|Greenwood|Earnshaw|1997|p=328}}

Silicon shows clear differences from carbon. For example, [[organic chemistry]] has very few analogies with silicon chemistry, while [[silicate]] minerals have a structural complexity unseen in [[oxocarbon]]s.{{sfn|Greenwood|Earnshaw|1997|pp=327–328}} Silicon tends to resemble germanium far more than it does carbon, and this resemblance is enhanced by the [[d-block contraction]], resulting in the size of the germanium atom being much closer to that of the silicon atom than periodic trends would predict.{{sfn|Greenwood|Earnshaw|1997|pp=359-361}} Nevertheless, there are still some differences because of the growing importance of the divalent state in germanium compared to silicon. Additionally, the lower Ge–O bond strength compared to the [[Silicon–oxygen bond|Si–O bond]] strength results in the absence of "germanone" polymers that would be analogous to silicone polymers.{{sfn|Greenwood|Earnshaw|1997|pp=374}}