My lecturer said that in a $\ce{C-S}$ bond, sulfur is slightly $\delta^{-}$ and the carbon slightly $\delta^{+}$, although they have (almost) the same electronegativity. What is the cause of this phenomena if it is not electronegativity?


I think your lecturer would have been more correct had he said that the carbon-sulfur bond reacts as if the sulfur is slightly $\delta^{-}$ and the carbon slightly $\delta^{+}$.

Sulphur is a larger atom so it has more, loosely held electrons than carbon. This means that a sulfur atom is more polarizable than a carbon atom. Although carbon and sulfur have similar electronegativities, the effect of this increased polarizability for sulfur is that sulfur behaves or reacts as if the bond were polar. Take the case of nucleophilic attack on a $\ce{C-S}$ bond. As the nucleophile approaches, the $\ce{C-S}$ bond will tend to polarize because of sulfur's polarizability. And in which direction will the bond polarize? Electron density will prefer to shift such that it resides on the larger atom (charge to size effect), sulfur.

This is the same argument used in the case of a carbon-iodine bond. Again the two atoms have similar electronegativities (carbon = 2.55, iodine = 2.66), so although the bond is not polar based on electronegativity considerations, the polarizability of the iodine allows the bond to react as if it were polar when a nucleophile approaches.

  • $\begingroup$ Although I think everything you said is correct, I still guess it depends on what kind of bond it is (single, double, triple) and what saturates the carbon. But in the most simplest case you are absolutely right. $\endgroup$ Sep 3 '14 at 17:51
  • $\begingroup$ This was exactly what I thought, but then why isn't there a $\delta^{-}$ on sulfur if charge can be distributed there better? It seems to me that if the electrons shift towards the sulfur atom, they will become more stable because there is less interaction with the other electrons of the atom on sulfur. $\endgroup$
    – Jori
    Sep 3 '14 at 18:56
  • $\begingroup$ Oh wait, perhaps I get it. The electrons are of course confined to the C-S bonding MO (which is symmetric I guess?). So not until some reaction happens that breaks this bond, can the size of the sulfur atom (and its polarizability) come into play and stabilize these electrons (thus behaving as if it was $\delta^{-}$). Is that right? $\endgroup$
    – Jori
    Sep 3 '14 at 19:13
  • $\begingroup$ Yes, you've got it. The electronegativities are similar, so the bond is not polar. But when a reactant approaches, the polarizable bond will polarize such that electron density shifts towards sulfur (or iodine). $\endgroup$
    – ron
    Sep 3 '14 at 19:21

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