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After recent question about non-existent pnictide(V) iodides, I stumbled upon a similar issue in oxygen group. For example, $\ce{SI2}$ was reported to be found… but at the temperature as low as $\pu{9 K}$. Well, OK, I heard that sulfur prefers disulfides, but $\ce{S2I2}$ was characterised at $\pu{-90 °C}$. Moreover, situation with selenium iodides is even worse!

One of the reasons for that is weakness of $\ce{S-I}$ and $\ce{Se-I}$ bonds, which causes the compounds to be endothermic.

Klapötke and Passmore [1, p. 234] provide an explanation for the weakness of these bonds:

The instability of the $\ce{S-I}$ and $\ce{Se-I}$ bonds can be attributed to their very low ionic resonance stabilization energies as the electronegativity of iodine is about the same as that of sulfur and selenium.

I think this is a partial explanation, at best - ionic component sure makes bonds stronger, but by this logic all non-polar covalent bonds should be weak. I’m not saying I’d expect bonds created by iodine to be strong, but it’s not like we need cryogenic temperatures to get methyl iodide!

Could anyone provide a good explanation as to why these bonds are so weak, maybe describing stereo-electronic effects, or using computational methods?

References

  1. T. Klapötke, J. Passmore, J. Acc Chem. Res. 1989, 22, 234–240. DOI:10.1021/AR00163A002.
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  • $\begingroup$ I know it's not what you're looking for, but what's wrong with the good ol' "atom size" explanation? They're all pretty large atoms with minute differences in electronegativity. Or are you looking for a fully non-empirical explanation? $\endgroup$
    – TheRelentlessNucleophile
    Commented Jun 24, 2023 at 16:59
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    $\begingroup$ Well, phosphorus has similar size and electronegativity, but PI3 is a common reagent. $\endgroup$
    – Mithoron
    Commented Jun 24, 2023 at 19:16
  • $\begingroup$ @Mithoron It is worth reworking. To me, sulfur and selenium bond energies in general seem extremely interesting. Iodine may not be exceptional, but still an example. $\endgroup$
    – Paul Kolk
    Commented Apr 4 at 20:09
  • $\begingroup$ @PaulKolk I edited it just side-stepping this bit. Why are you so interested in sulfur and selenium? $\endgroup$
    – Mithoron
    Commented Jun 18 at 17:58
  • $\begingroup$ @Mithoron They happen to be promising candidates for use in composite materials involving mixed carbides containing aluminum. Suitable oxides are too reactive to aluminum. Too long story to tell anyway. Before I start experimenting, I need answers to my questions. $\endgroup$
    – Paul Kolk
    Commented Jun 18 at 18:31

1 Answer 1

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It isn't just sulfur iodide. Sulfur in low oxidation states does not form any $\ce{SX2}$ molecules that are stable under ambient or near-ambient conditions. Such molecules tend to dimerize, with loss of some halogen atoms in the case of chlorine or a heavier halogens, forming disulfur compounds. (In the case of $\ce{S2I2}$, Wikipedia reports stability up to a somewhat higher temperature than suggested in the question, $\pu{-30 °C}).$

What happens is sulfur in the disulfur compounds can act as a pi donor, thereby strengthening the sulfur–sulfur bond while only slightly weakening the sulfur–halogen bond. The specific case of disulfur dichloride is discussed here.

In addition to the halides, the structure of disulfur monoxide, which appears to exist in the Jovian moon Io, allows a similar interaction between the sulfur–sulfur bond and oxygen.

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