Is science sometimes in danger of getting tunnel vision? Recently published ebook author, Ian Miller, looks at other possible theories arising from data that we think we understand. Can looking problems in a different light give scientists a different perspective?

Hypervalency: What Is Involved With The Bonding In Perchlorate And Sulphate?

In the latest "Chemistry World" there is an article arguing there is a controversy relating to the nature of the bonding in molecules such as the perchlorate anion, which appears now to be describable as having the chlorine atom with a positive charge of three, the four oxygen atoms with a charge of minus one each. The bonding is therefore one of four equal single bonds. Presumably, sulphate has the same issues, and according to Wikipedia, computational chemists put a charge of 2.45 on the sulfur atom. Crystal structures apparently indicate the four bonds are equal. Why got to these extremes? The problem is that chlorine has seven outer electrons, but six of them are usually regarded as residing in three pairs, and hence should be inert. Accordingly, chlorine has a valence of 1. Now many chlorine compounds do, but perchlorate, by definition, can be considered as the adduct of water on Cl2O7, i.e. all the outer electrons are involved. In principle, upon electron pairing, that gives 14 electrons in the outer valence shell. How can that be? The sulphur in sulphate has six outer electrons, four of which are paired. To get the required valence of six, again all electrons have to be unpaired if electron pairing is relevant.
The traditional method was to invoke 3d orbitals. These are empty, so they may be available for hybridization, BUT, according to the article, "quantum chemists have shown that it is energetically unfeasible to use d orbitals for extra bonds". It was asserted that this undermines a quantum mechanical account of Lewis bonding. My immediate problem with this assertion is, "how do we know?" The 3d orbital energies are obviously higher than 3p for chlorine, but how much higher, and does the energy difference remain if the orbitals are used for bonding? I am not arguing the statement is wrong, but merely that I would like to know why everyone thinks it is right. The output of computations is insufficient, because computations, according to Pople's Nobel lecture, are heavily dependent on validation, and we are a little short of the requirement to validate this statement. We can go further. The 2p orbitals are clearly at a higher energy than the 2s orbitals when we excite to them, yet boron almost never forms a B – X molecule, other than in highly energetic experiments, and not only does it use all three electrons, but it tries harder to achieve a tetrahedral configuration. So, if boron can do this, why cannot sulphur do it with 3d orbitals.
The article suggests that the answer might come from putting large negative charge on the oxygen atoms, and strong positive charge on the chlorine. The perchlorate anion is therefore an anion with four oxygen atoms with nearly a negative charge on each, and nearly three positive charges on the chlorine atom. The question then is, why does not this positive charge attract and polarize towards it the negative charge. If it does, we are back to the original problem.
What we need are data, and there are some. Consider only sulphate. We can form stable esters, such as dimethyl sulphate. If we do, the structure is consistent with two S=O and two S-O bonds. The  S – O bond length is 156.7 pm, the S = O bond length 141.7 pm. (J Mol. Str. 73, 99 – 104) while the infrared spectrum (Spectrochim Acta 28A, 1889 – 1898) gives the symmetric and asymmetric stretches of two pairs: the double bonds at 1389 and 1199 cm-1, with the single bonds at 829 and 757 cm-1. The infrared spectra of sulphates as a whole typically have medium to strong signals around 645 cm-1, and very strong signals at1110 cm-1, yet the S – O bonds in the anions all have the same length, so what does that mean? Obviously, even this common molecule still needs further work. I don't know the answer, but I would very much prefer it if the theoreticians would publish the reasons, and assumptions used, when they publish a statement saying the central atom has an extremely high positive charge. Their model might work for the sulphate anion, but it does not appear to for dimethyl sulphate, so the problem with how to explain hypervalency remains.

Posted by Ian Miller on Oct 26, 2015 1:53 AM Europe/London

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