It is fairly well-known that the acidity of a proton alpha to a sulfide/sulfoxide/sulfone increases in that order. Some pKa values (in DMSO), courtesy of the Evans pKa table:

acidities of sulfide vs sulfoxide vs sulfone

Traditionally this has been explained by resonance, or delocalisation of the carbon-based negative charge onto oxygen (left, below). But this necessarily invokes hypervalency and a S=O double bond. The major resonance form in a sulfoxide should be the charge-separated, octet-compliant form (right, below).

resonance forms of sulfoxide and sulfoxide anion

What is the proper explanation for the acidity trend above, then?

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    $\begingroup$ So if we were just considering inductive effects, how much would we expect the $\mathrm{p}K_{\mathrm{a}}$ value to drop for every extra oxygen in a $\beta$-position? $\endgroup$ – Zhe Sep 6 '17 at 17:29
  • $\begingroup$ @Zhe I actually suspect it is perhaps more to do with the positive charge on sulfur, rather than the oxygen (don't think the oxygen really does much, cf. acidity of typical alkane vs typical ether) But I don't have a good source, it's just a hunch. $\endgroup$ – orthocresol Sep 6 '17 at 17:34
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    $\begingroup$ Isn't the positive charge on sulfur just caused by the electronegative oxygens? $\endgroup$ – Zhe Sep 6 '17 at 19:20
  • $\begingroup$ @Zhe the positive formal charge, I mean. $\endgroup$ – orthocresol Sep 7 '17 at 6:38

The oxidation state of sulphur and its therefore increased inductive effect should be enough to explain this effect.

Besides the number of oxygens, the principle difference between the three structures is the oxidation state of sulphur that rises from $\mathrm{-II}$ to $\pm0$ to $\mathrm{+II}$. The higher an element’s oxidation state, the stronger it draws electrons towards itself since a higher oxidation state typically correlates with less electron density.

Unfortunately, we cannot draw good comparisons. Phosphorus would suggest itself, but neither the Evans table nor the Bordwell data have enough data points for scrutinisation. Furthermore, phosphorus is less electronegative than carbon, meaning the bonds would be polarised towards carbon.

Chlorine might be an option being on the other side of sulphur in the periodic table, but to the best of my knowledge no ‘chlorone’ or ‘chloroxide’ organic compounds are known let alone their acidities.

Nitrogen in its higher oxidation state (a nitro group) has a double bond which can resonate with the negative charge so it is not a good comparison. And the elements in the fourth period and below are again too electropositive in general. If anything, we may be able to draw conclusions based on the β-acidity of an ether, the α-acidity of an acetal/ketal or the α-acidity of an — if these are even stable enough to be deprotonated.

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    $\begingroup$ Well, there's for example Perchlorylbenzene, so some analogies might be drawn. $\endgroup$ – Mithoron Sep 7 '17 at 12:30
  • $\begingroup$ @Mithoron Didn’t know it existed! $\endgroup$ – Jan Sep 7 '17 at 12:35
  • $\begingroup$ Found this in J. Phys. Chem. A 2005, 109, 4966: "The acidity trend apparently stems from other factors [than resonance], such as the electrical stabilization of the anionic carbon center adjacent to polarized sulfenyl (RS), sulfinyl (RSO), and sulfonyl (RSO2) groups." I agree that sulfur electronegativity is probably the best explanation. For analogies, amine vs. N-oxide might work, but I wonder if the data exists. $\endgroup$ – orthocresol Sep 7 '17 at 16:31
  • $\begingroup$ @orthocresol Oooh, I wonder whether someone wants to add a literature-suppored answer to this Q. $\endgroup$ – Jan Sep 8 '17 at 9:41

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