What are the possible resonance structures for $\ce{ClO_2^-}$? Assigning one double bond to the structure makes for formal charges of $\ce{O(-1)-Cl(0)=O(0) <-> O(-1)=Cl(0)-O(-1)}$. It appears that the same set of formal charges can be achieved with two double bonds (expanding the octet on $\ce{Cl}$ to 12 electron capacity, using the d orbital). That is, $\ce{O(0)=Cl(-1)=O(0)}$.

Is this feasible, and if not, why not?


1 Answer 1


The here used methods are also only approximations and I do not claim that the provided numbers are accurate. However, they are used here to make the point, that bonding situations involving resonance are not easily explained.

Oxygen is slightly more electronegative than chlorine, therefore the bonds will always be slightly polarised towards the oxygen.

The following resonance scheme is based on an NBO analysis on the DF-BP86/def2-SVP level of theory. The percentage value gives an estimate how close the given conformation is to an idealised Lewis structure. From this you could also derive how much such a structure would contribute to a resonance situation (in terms of Valence Bond theory).
What we do see here is, that the electronic structure is better described with more formal charges. The contribution of a π bond is best described by an ionic donor-acceptor structure, rather than breaking the octet rule at chlorine.
Noteworthy is also that the doubly ionic donor-acceptor configuration is as equally good a lewis structure than any that breaks the octet rule.
The worst representation is the one with a four electron three centre bond (which is also not part of the Lewis bonding picture anyway).
resonance structures of chlorite anion

Videos like the one mentioned in the Wikipedia article are easy to follow, but often they are too simplified to be correct.

  • $\begingroup$ This is quite a curious answer, I had no idea calculations such as these were possible. It took me a minute to understand the bottom three structures though, because the arrows look almost like dative covalent bonds rather than electrostatic attraction between the ions. $\endgroup$ Commented Apr 2, 2015 at 12:08
  • $\begingroup$ @Nicolau In principle they are dative covalent bonds. These are just representations of the same electron density. So in general all the presented structures are different viewpoints of the same situation. NBO localises and then partitions the electron density - in principle you could go about and construct valence bond configuration from that and actually calculate the mixing, but that is tedious and takes a lot of time and I cannot do it. $\endgroup$ Commented Apr 3, 2015 at 3:16
  • $\begingroup$ Then I'm not quite sure how chlorine has a positive formal charge even though it's being given electron pairs from the oxygen atoms. The bottom-right representation looks particularly strange to me. The charges seem to be calculated as if there were no covalent bond present at all, dative or otherwise. $\endgroup$ Commented Apr 3, 2015 at 3:34
  • $\begingroup$ @Nicolau In the original Lewis picture there is no differentiating between bonds and in this sense only the very first structure is a true Lewis structure. The next two structures break the octet rule. The last in the first row has delocalised bonds. The dative bond does not exist in the Lewis picture, but extending this formalism, it only counts as electrons from where it originates. In the Lewis picture it does not matter if there is a bond in that case or not. $\endgroup$ Commented Apr 3, 2015 at 8:25
  • $\begingroup$ Ahh, now I see. I thought that might be it, but counting electrons from where they originate is rather unusual, at least to me. Thanks for taking the time to explain. $\endgroup$ Commented Apr 3, 2015 at 11:51

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