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]
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* sigmaantibonding 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]
^Housecroft, C. E.; Sharpe, A. G. (2008). Inorganic Chemistry (3rd ed.). Prentice Hall. pp. 413–424. ISBN978-0-13-175553-6.
^ abcN. C. Norman (1997). Periodicity and the s- and p-Block Elements. Oxford University Press. pp. 50–52, 65–67. ISBN978-0-19-855961-0.
^Xiong, Y.; Yao, S.; Driess, M. (2013). "Chemical Tricks To Stabilize Silanones and Their Heavier Homologues with EO Bonds (E=Si–Pb): From Elusive Species to Isolable Building Blocks". Angew. Chem. Int. Ed.52 (16): 4302–4311. doi:10.1002/anie.201209766. PMID23450830.
^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. PMID15303175.
^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.
^ abcSmith, Michael B.; March, Jerry (2007). March's Advanced Organic Chemistry (6th ed.). John Wiley & Sons. pp. 24–25. ISBN978-0-471-72091-1.