I recently looked up hydrogen sulfate ($\ce{HSO4-}$) to see how it's put together, and found this image of methyl hydrogen sulfate:

lewis structure of methylsulfate

I'd assume that hydrogen sulfate looks the same, but without the methyl group attached. However I'm confused, how can it be constructed like that? This shows the sulfur having a total of 12 valence electrons, but shouldn't the maximum be 8?

I might be asking something that'll be obvious to me later on in chemistry, but this (and other groups that act as atoms) appear early on. I probably don't need to know it yet but I'd prefer to know how it works if I can as opposed to just accepting it.


2 Answers 2


Hypervalency is an obsolete concept - it was used to explain bonding situations of molecules, that seem to exceed the octet rule. It was accomplished by the help of $\ce{spd}$ hybrid orbitals. It is now known, that for almost all molecules that were described this way the contribution of $\ce{d}$ orbitals is negligible (<1%).

Focussing on the question at hand, the bonding situation of the core $\ce{SO4^{2-}}$ anion is isoelectronic to the phosphate anion $\ce{PO4^{3-}}$.

For the methylsulfate this situation may be extended. A qualitative analysis of the natural charges (BP86/cc-pVDZ) reveals no difference. All of the oxygen atoms carry a negative charge, while sulfur is highly positive charged. Further more, NBO analysis confirms that the terminal oxygens have three lone pairs, while the other oxygens have two lone pairs. A more accurate Lewis structure would represent that fact and only use single bonds.

chargeslewis structure

  • $\begingroup$ While I can understand the hypervalency concept, if that's obsolete then this might be a bit too advanced for me. My current study material hints towards using the d orbital but still doesn't explain it, simply states that "some" atoms can hold up to 18 electrons, including sulphur and phosphorus, and that others like boron and beryllium are content with less than 8. My main issue is because they are exceptions to the octet rule when drawing Lewis structure, but at this point I don't think there's much I can do besides memorising them. All the same, thanks a lot! $\endgroup$
    – puggsoy
    Commented Jun 26, 2014 at 0:19

I think it's time for you to know about Octet Expansion rule. The molecules that expand their octet are known as Hypervalent Molecules.

Normally, atoms in group 2 have 8 electrons and can thus form 4 bonds. This is mainly because they have s and p orbitals. However, in elements of group 3, they have extra orbitals called d orbitals that can be used to form more bonds. when forming extra bonds, the atom promotes electrons into the d orbitals and allows the atom to from more than 4 bonds.

Same thing happens in the case of sulfur. Sulfur belongs to 3rd period and therefore it has a vacant 3d orbital and thus can accommodate more than 8 electrons in its valence cell.

  • $\begingroup$ I see. I know about orbitals so I suppose that makes sense. I hadn't heard of hypervalent molecules before though, so thanks! $\endgroup$
    – puggsoy
    Commented Jun 19, 2014 at 7:50
  • 3
    $\begingroup$ The involvement of d orbitals in sulfur compounds is marginal (below 1%). Hypervalency and hybridisation are the most famous misunderstood concepts in chemistry. $\endgroup$ Commented Jun 19, 2014 at 8:07
  • $\begingroup$ @Martin : So what are you proposing here? o.O $\endgroup$
    – ashu
    Commented Jun 19, 2014 at 8:11
  • 2
    $\begingroup$ As $\ce{SO4^{2-}}$ is isoelectronic with $\ce{PO4^{3-}}$ I would simply suggest an analogous bonding picture. I.e. Only sigma single bonds to the oxygens, each carrying a negative charge and sulfur being positively charged. The bonding picture of $\ce{HSO4-}$ can then be derived accordingly. $\endgroup$ Commented Jun 19, 2014 at 8:21
  • 1
    $\begingroup$ No definitely not. The concept you provide is/was, however outdated it is now, in use. Unfortunately it will almost certainly still be taught at some schools and it is hard to get rid of these myths - but then again some teach creation theory. $\endgroup$ Commented Jun 19, 2014 at 8:45

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy.

Not the answer you're looking for? Browse other questions tagged or ask your own question.