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Silicon–oxygen bond

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A silicon–oxygen bond (Si−O bond) is a chemical bond between silicon and oxygen atoms that can be found in many inorganic and organic compounds.[1] In a silicon–oxygen bond, electrons are shared unequally between the two atoms, with oxygen taking the larger share due to its greater electronegativity. This polarisation means Si–O bonds show characteristics of both covalent and ionic bonds.[2] Compounds containing silicon–oxygen bonds include materials of major geological and industrial significance such as silica, silicate minerals and silicone polymers like polydimethylsiloxane.[1][3]

Bond polarity, length and strength

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 Si−O 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: Siδ+—Oδ−.[2]

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.[1] 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.[4] 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,[5] for example by coordination to a metal centre,[6] coordination to Lewis acids or bases,[7] or by steric shielding.[8]

Comparison of C–O and Si–O bonds
Bond Carbon–oxygen Silicon–oxygen
E C Si
Pauling electronegativity of E 2.55 1.90
Pauling electronegativity difference between E and O 0.89 1.54
H3E–O–EH3 Bond angle / ° 111[9] 142[10]
Typical sp3 E–O single bond length / Å 1.43[11] 1.63[12]
Typical sp2 E–O single bond length / Å 1.34[11]
Typical sp2 E=O double bond length / Å 1.21[11] 1.52[8][13]
Typical sp E=O double bond length / Å 1.16[14] 1.48[15][16]
Typical E–O single bond strength / kJ mol−1 ~360[1] 452[1]
Typical E=O double bond strength / kJ mol−1 715[4] 590[4]

Bond angles

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.[2]

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,[17] 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.[1]

Coordination number

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.[1]

See also

References

  1. ^ a b c d e f g Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 342–366. ISBN 978-0-08-037941-8.
  2. ^ a b c Dankert, Fabian; von Hänisch, Carsten (2021). "Siloxane Coordination Revisited: Si􏰉–O Bond Character, Reactivity and Magnificent Molecular Shapes". Eur. J. Inorg. Chem. 2021 (29): 2907–2927. doi:10.1002/ejic.202100275. S2CID 239645449.
  3. ^ Housecroft, C. E.; Sharpe, A. G. (2008). Inorganic Chemistry (3rd ed.). Prentice Hall. pp. 413–424. ISBN 978-0-13-175553-6.
  4. ^ a b c N. C. Norman (1997). Periodicity and the s- and p-Block Elements. Oxford University Press. pp. 50–52, 65–67. ISBN 978-0-19-855961-0.
  5. ^ Xiong, Y.; Yao, S.; Driess, M. (2013). "Chemical Tricks To Stabilize Silanones and Their Heavier Homologues with EO Bonds (E=Si–Pb): From Elusive Species to Isolable Building Blocks". Angew. Chem. Int. Ed. 52 (16): 4302–4311. doi:10.1002/anie.201209766. PMID 23450830.
  6. ^ Sen, S. S. (2014). "A Stable Silanone with a Three-Coordinate Silicon Atom: A Century-Long Wait is Over". Angew. Chem. Int. Ed. 53 (34): 8820–8822. doi:10.1002/anie.201404793. PMID 24990653.
  7. ^ Sun, T.; Li, J.; Wang, H. (2022). "Recent Advances in the Chemistry of Heavier Group 14 Analogues of Carbonyls". Chem. Asian J.: e202200611. doi:10.1002/asia.202200611. PMID 35883252. S2CID 251104394.
  8. ^ a b Kobayashi, Ryo; Ishida, Shintaro; Iwamoto, Takeaki (2019). "An Isolable Silicon Analogue of a Ketone that Contains an Unperturbed Si=O Double Bond". Angew. Chem. Int. Ed. 58 (28): 9425–9428. doi:10.1002/anie.201905198. PMID 31095845. S2CID 157056381.
  9. ^ Vojinović, Krunoslav; Losehand, Udo; Mitzel, Nobert W. (2004). "Dichlorosilane–dimethyl ether aggregation: a new motif in halosilane adduct formation". Dalton Trans. (16): 2578–2581. doi:10.1039/B405684A. PMID 15303175.
  10. ^ Barrow, M. J.; Ebsworth, E. A. V.; Harding, M. M. (1979). "The crystal and molecular structures of disiloxane (at 108 K) and hexamethyldisiloxane (at 148 K)". Acta Crystallogr. B. 35 (9): 2093–2099. doi:10.1107/S0567740879008529.
  11. ^ a b c Smith, Michael B.; March, Jerry (2007). March's Advanced Organic Chemistry (6th ed.). John Wiley & Sons. pp. 24–25. ISBN 978-0-471-72091-1.
  12. ^ Kaftory, Menahem; Kapon, Moshe; Botoshansky, Mark (1998). "The Structural Chemistry of Organosilicon Compounds". In Rappoport, Zvi; Apeloig, Yitzhak (eds.). The Chemistry of Organic Silicon Compounds, Volume 2. PATAI'S Chemistry of Functional Groups. John Wiley & Sons, Ltd. doi:10.1002/0470857250. ISBN 9780471967576.
  13. ^ Bogey, Marcel; Delcroix, Bruno; Jean-Claude Guillemin, Adam Walters (1996). "Experimentally Determined Structure of H2SiO by Rotational Spectroscopy and Isotopic Substitution". J. Mol. Spectrosc. 175 (2): 421–428. Bibcode:1996JMoSp.175..421B. doi:10.1006/jmsp.1996.0048.
  14. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 292, 304–314. ISBN 978-0-08-037941-8.
  15. ^ Schnöckel, Hansgeorg (1978). "IR Spectroscopic Detection of Molecular SiO2". Angew. Chem. Int. Ed. 17 (8): 616–617. doi:10.1002/anie.197806161.
  16. ^ Jutzi, Peter; Schubert, Ulrich (2003). Silicon Chemistry: From the Atom to Extended Systems. Wiley-VCH. pp. 27–28. ISBN 9783527306473.
  17. ^ Glidewell, C.; Liles, D. C. (1978). "The crystal and molecular structure of oxobis[triphenylsilicon(IV)]". Acta Crystallogr. B. 34: 124–128. doi:10.1107/S0567740878002435. S2CID 98347658.
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