Silicon–oxygen bond

On the Pauling electronegativity scale, silicon has an electronegativity of 1.90 and oxygen 3.44. The electronegativity difference between the elements is therefore 1.54. Because of this moderately large difference in electronegativities, the {{chem2|Si\sO}} bond is polar but not fully ionic. Carbon has an electronegativity of 2.55 so carbon–oxygen bonds have an electronegativity difference of 0.89 and are less polar than silicon–oxygen bonds. Silicon–oxygen bonds are therefore covalent and polar, with a partial positive charge on silicon and a partial negative charge on oxygen: Si^δ+—O^δ−.{{ cite journal | title = Siloxane Coordination Revisited: Si􏰉–O Bond Character, Reactivity and Magnificent Molecular Shapes | first1 = Fabian | last1 = Dankert | first2 = Carsten | last2 = von Hänisch | journal = Eur. J. Inorg. Chem. | year = 2021 | volume = 2021 | issue = 29 | pages = 2907–2927 | doi = 10.1002/ejic.202100275 | s2cid = 239645449 }}

Silicon–oxygen single bonds are longer (1.6 vs 1.4 Å) but stronger (452 vs. about 360 kJ mol⁻¹) than carbon–oxygen single bonds. However, silicon–oxygen double bonds are weaker than carbon–oxygen double bonds (590 vs. 715 kJ mol⁻¹) due to a better overlap of p orbitals forming a stronger pi bond in the latter. This is an example of the double bond rule. For these reasons, carbon dioxide is a molecular gas containing two C=O double bonds per carbon atom whereas silicon dioxide is a polymeric solid containing four Si–O single bonds per silicon atom; molecular SiO₂ containing two Si=O double bonds would polymerise.{{ cite book | title = Periodicity and the s- and p-Block Elements | author = N. C. Norman | publisher = Oxford University Press | year = 1997 | isbn = 978-0-19-855961-0 | pages = 50–52, 65–67 }} Other compounds containing Si=O double bonds are normally very reactive and unstable with respect to polymerisation or oligomerization. Silanones oligomerise to siloxanes unless they are stabilised,{{ cite journal | title = Chemical Tricks To Stabilize Silanones and Their Heavier Homologues with EO Bonds (E=Si–Pb): From Elusive Species to Isolable Building Blocks | first1 = Y. | last1 = Xiong | first2 = S. | last2 = Yao | first3 = M. | last3 = Driess | journal = Angew. Chem. Int. Ed. | year = 2013 | volume = 52 | issue = 16 | pages = 4302–4311 | doi = 10.1002/anie.201209766 | pmid = 23450830 }} for example by coordination to a metal centre,{{ cite journal | title = A Stable Silanone with a Three-Coordinate Silicon Atom: A Century-Long Wait is Over | first1 = S. S. | last1 = Sen | journal = Angew. Chem. Int. Ed. | year = 2014 | volume = 53 | issue = 34 | pages = 8820–8822 | doi = 10.1002/anie.201404793 | pmid = 24990653 }} coordination to Lewis acids or bases,{{ cite journal | first1 = T. | last1 = Sun | first2 = J. | last2 = Li | first3 = H. | last3 = Wang | title = Recent Advances in the Chemistry of Heavier Group 14 Analogues of Carbonyls | journal = Chem. Asian J. | year = 2022 | volume = 17 | issue = 18 | pages = e202200611 | doi = 10.1002/asia.202200611 | pmid = 35883252 | s2cid = 251104394 }} or by steric shielding.{{ cite journal | title = An Isolable Silicon Analogue of a Ketone that Contains an Unperturbed Si=O Double Bond | first1 = Ryo | last1 = Kobayashi | first2 = Shintaro | last2 = Ishida | first3 = Takeaki | last3 = Iwamoto | journal = Angew. Chem. Int. Ed. | volume = 58 | issue = 28 | year = 2019 | pages = 9425–9428 | doi = 10.1002/anie.201905198 | pmid = 31095845 | s2cid = 157056381 | doi-access = free }}

Disiloxane groups, Si–O–Si, tend to have larger bond angles than their carbon counterparts, C–O–C. The Si–O–Si angle ranges from about 130–180°, whereas the C–O–C angle in ethers is typically 107–113°. Si–O–C groups are intermediate, tending to have bond angles smaller than Si–O–Si but larger than C–O–C. The main reasons are hyperconjugation (donation from an oxygen p orbital to an Si–R σ* sigma antibonding molecular orbital, for example) and ionic effects (such as electrostatic repulsion between the two neighbouring partially positive silicon atoms). Recent calculations suggest π backbonding from an oxygen 2p orbital to a silicon 3d orbital makes only a minor contribution to bonding as the Si 3d orbital is too high in energy.

The Si–O–Si angle is 144° in α-quartz, 155° in β-quartz, 147° in α-cristobalite and (153±20)° in vitreous silica. It is 180° in coesite (another polymorph of SiO₂), in Ph₃Si–O–SiPh₃,{{ cite journal | first1 = C. | last1 = Glidewell | first2 = D. C. | last2 = Liles | title = The crystal and molecular structure of oxobis[triphenylsilicon(IV)] | journal = Acta Crystallogr. B | year = 1978 | volume = 34 | pages = 124–128 | doi = 10.1107/S0567740878002435 | s2cid = 98347658 | doi-access = }} and in the [O₃Si–O–SiO₃]⁶⁻ ion in thortveitite, Sc₂Si₂O₇. It increases progressively from 133° to 180° in Ln₂Si₂O₇ as the size and coordination number of the lanthanide decreases from neodymium to lutetium. It is 150° in hemimorphite and 134° in lithium metasilicate and sodium metasilicate.

In silicate minerals, silicon often forms single bonds to four oxygen atoms in a tetrahedral molecular geometry, forming a silicon–oxygen tetrahedron. At high pressures, silicon can increase its coordination number to six, as in stishovite.

• Organosilicon compound
• Carbon–hydrogen bond
• Carbon–carbon bond
• Carbon–fluorine bond
• Bonding in solids

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Category:Chemical bonding