Silicon dioxide has a huge variety of structures. Most of them are built up from connected $\ce{SiO4}$ units — the $\ce{O–Si–O}$ angle is $109.5°$, accordingly. The VB/hybridization approach to this tetrahedron would consequently assign $4\text{ }\mathrm{sp^3}$ orbitals to the silicon atom.

Now, the interesting part about silicon dioxide is the flexibility of its $\ce{Si–O–Si}$ angles, i.e. the bond angle of the bridging oxygen atoms; they vary from $100$ to $180$ degrees (the latter occurs in hexagonal tridymite, see e.g. here), depending on the polymorph. In vitreous samples, these angles show a distribution centered around $145$ degrees or so (source). In common forms of crystalline $\ce{SiO2}$, the angle also is close to $145$ degrees (e.g. this source). How is this possible? Since the oxygen bonds to two silicon atoms, one would expect a $\mathrm{sp^3}$ hybridization, with two orbitals filled with lone pairs — hence, a $104°$ angle like in water, or a slightly bigger one (since silicon is less electronegative than hydrogen — thus, the $\ce{Si–O}$ bond is more polar than the H–O bond and the angle bigger). But never an angle as large as $170°$: that would almost point at an $\mathrm{sp}$ hybridization of the oxygen atoms. How is such an angle energetically more convenient than the typical bent one?

  • $\begingroup$ Most of the silicon dioxide polymorphs show deviation from the ideal (predicted) angles, but they're metastable and will eventually revert to $\alpha$-quartz. See The Quartz Page for a more thorough discussion. Also note that the angles are for Si-O-Si and not O-Si-O. $\endgroup$ Dec 27, 2015 at 19:07
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    $\begingroup$ That the Si–O–Si bond angles arise from the arrangement of the SiO4 tetrahedrons is obvious, as well as the fact that the O–Si–O angles are always 109.5° in stable polymorphs (since we're talking about tetrahedrons). The question is: how does VB theory represent the bonding between oxygen and silicon with these extreme variety of angles? How does it justify the existence of non-109° bonds around an oxygen molecule? Why do these bond angles differ so drastically from other X–O–X bonds such as those in water, according to VB theory? $\endgroup$
    – Corundum
    Dec 27, 2015 at 23:34
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    $\begingroup$ There is also the possibility that valence bond theory fails for a silicon dioxide. There are many compounds for which VB can't explain the structure. $\endgroup$
    – Tyberius
    Aug 3, 2017 at 2:51
  • $\begingroup$ After adding the bounty, I realized this question had a 150 point bounty two years ago with no takers. I included a link to a paper giving the bond angles in common polymorphs of crystalline $\ce{SiO2}$. I also added an citation for an article on bond angle distribution in glass to the question - maybe that helps. $\endgroup$
    – Karsten
    Feb 1, 2021 at 22:41
  • $\begingroup$ In this paper, they modeled liquid $\ce{SiO2}$ and cristobalite using an ab initio approach treating it as an ionic system: researchgate.net/publication/… $\endgroup$
    – Karsten
    Feb 1, 2021 at 23:17

1 Answer 1


Tl;dr version: the O-Si bonds utilize sp hybrid orbitals so that “lone pairs” are in nearly pure p orbitals and, rather than being true lone pairs, engage in some sort of $\pi$ bonding, the exact nature of which is unknown. Other hypotheses are that the bond is highly ionic or that it is favored due to oxygen-oxygen repulsion.

Long version: First, let me point out a slight inaccuracy in the posted question:

Since the oxygen bonds to two silicon atoms, one would expect a sp3 hybridization, with two orbitals filled with lone pairs —

This is an incorrect interpretation of VB theory. Although we might approximate a molecule such as $\ce{H2O}$ as having an sp3-hybridized O atom, quantitative approaches to VB theory account for the fact that the contribution of s and p atomic orbitals to localized orbitals differs if the group differs. That is, lone pairs and bond pairs can have (and usually do have) different contributions from the s and p orbitals, ie different hybridization. A commonly cited example of this is $\ce{SH2}$, where the S-H bond orbitals are nearly entirely comprised of p orbitals, so the lone pairs are necessarily essentially sp hybrids.

Furthermore, one quantitative approach calculates the p orbital contribution based on the bond angle (Coulson's Theorem). In silicate, the Si-O-Si bond angle varies between ~140$^\circ$ and 180$^\circ$, depending on polymorph. We can calculate that the hybridization of these bond orbitals thus varies from sp (180 degree angle) to sp1.3 (140 degree angle), which is actually a relatively small variation in the s orbital fraction (43-50%). The lone pairs, therefore, are nearly pure p (93-100%) orbitals.

Although it is known that the $\alpha$-quartz polymorph (bond angle 144 degrees) is the most stable at low temperature, the observation of many polymorphs suggests that there is not much of an energetic penalty for the variation in bond angle within this range.

For simplicity, let’s consider the extreme case of the 180 degree bond angle, so that the O-Si bonds are formed from sp orbitals on O, and the lone pairs are in pure p orbitals. Both VB theory and MO theory indicate that this arrangement is only energetically favorable relative to a bent conformation if either the p orbitals are empty (as in $\ce{BeH2}$ for example) or if the electrons in the p orbitals are involved in $\pi$ bonding, for example the carbon atom of $\ce{CO2}$ or $\ce{HCN}$.

COnsistent with this idea, researchers have argued since at least as far back as the 1960's that the “lone pairs” on oxygen in silicate are in fact not lone pairs, but are involved in $\pi$ bonding to the silicon. This is supported by the short Si-O bond lengths and high bond strength which suggest a bond order greater than 1.

Historically, this interaction has been proposed to involve empty d orbitals on silicon, but more recently there has been a shift away from d orbital involvement, for similar reasons that d orbital involvement in phosphorus bonding has generally been disregarded. Interaction with empty $\sigma^*$ antibonding orbitals on Si, however, also does not appear satisfactory, as that would weaken the Si-O bonds (assuming that all Si-O bonds are equivalent, so all would have increased bond order from $\pi$ bonding, but decreased from addition of density to $\sigma^*$). Essentially, there would be a significant contribution of a resonance structure in which each Si has two Si-O bonds and one Si=O bond, similar to a carbonate.

Nonetheless, the idea that a large bond angle is favored due to some sort of $\pi$ bonding seems to be the prevalent view. Two other hypotheses, however, are still supported by some.

Other hypotheses

Ionic interactions

One alternative explanation is that there is very little difference in bond orbitals and lone pairs, which could result from a strongly ionic structure of $\ce{O^2-}$ anions and $\ce{Si^4+}$ cations[1]. On the oxygen, all of the electrons then are lone pairs, and the “bond angle” (which isn’t between any true bonds) is simply optimized to minimize repulsion between negatively charged oxygen atoms and between the positively charged silicon atoms. This would result in a tetrahedral arrangement of the negative charges around Si (as is observed) and a linear arrangement of the two Si atoms around each oxygen (the occasionally observed 180 degree angle). Since there aren’t any bonds, VB theory doesn’t really apply, and we consider the oxyanion as a lone atom with completely filled valence shell.

The problem with this argument is that free $\ce{Si^4+}$ is not observed under any circumstances, and most computational approaches suggest a high degree of covalency in the Si-O bond.

Oxygen-oxygen repulsion

A third hypothesis is that VB theory does not apply here because the minimization of energy of the Si-O-Si bond angle is not the only factor. Specifically, oxygen-oxygen repulsion favors a larger angle so as to increase the distance between silicon atoms and, therefore, the oxygens attached to adjacent silicon atoms.

Even at the small scale of a single molecule of $\ce{(HO)3Si-O-Si(OH)3}$, it has been argued that oxygen-oxygen repulsion is sufficient to cause a large Si-O-Si bond angle without any need for $\pi$ bonding[2]. The observed angle is a compromise between repulsion, which favors a 180 degree angle, and bond orbital energy, which favors a bent structure with bond angle closer to 110 degrees.

A problem with this hypothesis is that the Si-O-Si bond angle in disiloxane ($\ce{H3Si-O-SiH3}$) is also observed to be 144 degrees, even though there should be no oxygen atom repulsion effects and general steric effects should be much less than in silicates.

[1]Gillespie, RJ and Johnson, SA (1997) Inorg Chem 36:3031-3039. doi:10.1021/ic961381d

[2]Noritake F (2019) J Comput Chem Jpn 5:1. doi:10.2477/jccjie.2018-0016

  • $\begingroup$ "the occasionally observed 180 degree angle". Skimming the papers, it was not clear to me whether this was an artifact (disorder of say 175 degrees and 185 degrees, so to speak, averaged to 180 degrees) or real. $\endgroup$
    – Karsten
    Feb 2, 2021 at 18:47
  • $\begingroup$ I appreciate that the answer discusses constraints on the angle other than the local bonding (it is a solid, after all). $\endgroup$
    – Karsten
    Feb 2, 2021 at 18:49
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    $\begingroup$ @KarstenTheis - with respect to the 180 degree, I'm pretty sure I saw a paper with a table of polymorphs and one had 180 degrees, but until I can find that again, take a look at scripts.iucr.org/cgi-bin/paper?a19115. And w.r.t. other constraints, I was actually surprised that there wasn't more discussion in the lit about the long range order of a solid, perhaps because these wide bond angles are observed even in isolated small molecules like $\ce{(HO)3Si-O-Si(OH)3}$, and the computational work often focuses on them. $\endgroup$
    – Andrew
    Feb 2, 2021 at 20:25

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