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I have three factors in mind:

1) vdW repulsions

2) Bond length

3) Electron-electron repulsions.

I know that vdW repulsions between ligand atoms push the atoms apart. And bond length obviously influences bond angle; if the bonds were very long, then the bond angle can shrink without bringing the ligand atoms too close together.

Are there any other factors or are these it?

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    $\begingroup$ This is not a full answer, so I'll leave it as a comment. You also need to consider various electronic effects distinct from sterics. Overlap of $\pi$-orbitals, for example, introduces planarity constraints. There are also phenomena such as the anomeric effect and the gauche effect, which are often explained in terms of hyperconjugation. These are factors in determining lowest-energy conformers for given molecules, and sometimes have subtle effects on bond lengths and bond angles (and I mean beyond the obvious impact on dihedral angles). $\endgroup$
    – Greg E.
    Jul 13, 2014 at 1:37
  • $\begingroup$ These are just a few of the many complex electronic phenomena that may be relevant. In organometallic chem, there's also $\pi$-backbonding, for example, which shortens the bond between the metal and ligand, and can therefore slightly alter bond angles in asymmetric complexes. $\endgroup$
    – Greg E.
    Jul 13, 2014 at 1:44
  • $\begingroup$ Thank you. I was curious as to the statement my prof made. It was something along the lines of that bond angles can be rationalized using vdW repulsions alone. $\endgroup$
    – Dissenter
    Jul 13, 2014 at 1:50
  • $\begingroup$ I don't know what to call it, but ring formation could influence bond angles, such as cyclobutane. $\endgroup$
    – LDC3
    Jul 13, 2014 at 1:52
  • $\begingroup$ @Dissenter, yeah, I would charitably describe that as an oversimplification on your prof's part. I suspect your prof may have wanted to avoid getting into certain gray areas (which are numerous in chemistry); I've certainly seen this plenty of times with my own profs. Steric effects (which is kind of a catch-all term mostly referring to Pauli electron-electron repulsion) are certainly quite dominant in many cases, especially in simple and highly symmetrical structures, but they're not the whole story. $\endgroup$
    – Greg E.
    Jul 13, 2014 at 2:21

1 Answer 1

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Bond angles mainly depend on the following three factors:

  1. Hybridization: Bond angle depends on the state of hybridization of the central atom
    Hybridization: $\ce{sp^3}$, Bond angle: $109^\circ$, Example: $\ce{CH4}$
    Hybridization: $\ce{sp^2}$, Bond angle: $120^\circ$, Example: $\ce{BCl3}$
    Hybridization: $\ce{sp}$, Bond angle: $180^\circ$, Example: $\ce{BeCl2}$
    Generally s- character increase in the hybrid bond, the bond angle increases.

  2. Lone pair repulsion: Bond angle is affected by the presence of lone pair of electrons at the central atom. A lone pair of electrons at the central atom always tries to repel the shared pair (bonded pair) of electrons. Due to this, the bonds are displaced slightly inside resulting in a decrease of bond angle.

  3. Electronegativity: If the electronegativity of the central atom decreases, bond angle decreases.

Good Side note:

Triple bonds repel other bonding-electrons more strongly than double bonds.
Double bonds repel other bonding-electrons more strongly than single bonds.

Heres a nice read into the different components (extra)

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    $\begingroup$ Hybridization is a man-made concept that we use to try and approximate reality. Being a concept, hybridization has nothing to do with the actual energetics of why a molecule adopts a certain geometry. In fact, there are many cases were hybridization fails to explain reality, the water molecule being a case in point. It is often taught that the lone pairs residing on the central oxygen in water are ~sp3 hybridized, while in fact these two orbitals are largely unhybridized. P.S. Could you elaborate on what you mean by "repel" when you write, "Triple bonds repel other bonding-electrons more" $\endgroup$
    – ron
    Jul 13, 2014 at 18:13
  • $\begingroup$ @ron you're exactly right about hybridization. But what it allows us is classification. It allows us to associate particular arrangements with particular geometries. And with Chemistry, there are a ton of exceptions, so it's definitely not perfect. But when thinking about bond angles its helpful to associate like things so we can see patterns. And by Repel, I mean simply to push away, as in more often a more significant change in bond angle. $\endgroup$
    – David BasedMathematician Coven
    Jul 13, 2014 at 18:56
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    $\begingroup$ I am strongly agreeing with @ron. Not only is hybridisation, in this way of using it, a man-made concept (of purely mathematical considerations) - it is also always only the result of a geometrical arrangement of atoms in a molecule. And as this is true, it can only be applied to molecules composed out of the first and second period. A popular example for the failure of this concept is $\ce{PH3}$, which has bonding angles close to $90^\circ$. (Apart from the obvious water example ron already mentioned.) $\endgroup$
    – Martin - マーチン ♦
    Jul 14, 2014 at 11:33
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    $\begingroup$ The extra reading link is not available.. $\endgroup$
    – Stanislav Bashkyrtsev
    Apr 18, 2017 at 21:26
  • $\begingroup$ Downvoting because angles dominated by geometric constraints and dispersion forces are absent, consider paracyclophanes en.wikipedia.org/wiki/Cyclophane and dx.doi.org/10.1002/anie.201304674 $\endgroup$
    – TAR86
    Apr 19, 2017 at 4:55

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