7
$\begingroup$

On my most recent high-school chemistry exam I had a longform question where I was supposed to compare the solubilities in room-temperature water under normal pressure of a few molecules. According to the question $\ce{H2S}$ is very weakly soluble in water, and I needed to explain why.

I suggested that the difference in electronegativity between $\ce{H}$ and $\ce{S}$ is around $0.38$ (on the Pauling scale), which is very close to the threshold we use to categorize polar and non-polar bonds $(0.4).$ So, although the $\ce{H-S}$ bond is technically non-polar, making the whole molecule non-polar, it is very close to being polar, which might make the molecule a very slight dipole, favorizing solubility in water, a polar solvent

The teacher accepted this explanation, although they took away points because I did not mention the hydrogen bonds that supposedly form between the hydrogen sulfide and water molecules.

This confused me because I thought that hydrogen bonds can only form between hydrogen atoms in a $\ce{H-O},$ $\ce{H-N}$ or $\ce{H-F}$ bond and a very electronegative atom in a covalent bond such as $\ce{N},$ $\ce{O},$ or $\ce{F},$ which is what I find in most high-school level resources.

What am I missing out on? Can $\ce{H2S}$ form hydrogen bonds (even very weak) with water when dissolved?

$\endgroup$
7
  • $\begingroup$ Then your teacher would seem pretty good, but solubility of H2S is by no means very weak. Was this his idea or yours? $\endgroup$
    – Mithoron
    Mar 23 at 18:43
  • $\begingroup$ @Mithoron Well, the data I’m given is 2.5 mol/L for H2S, which I’m to compare to 0.0013 mol/L for diiodine and 50 mol/L for NH3. She did not say that my remark that H2S is weakly soluble is wrong, but just said that the part about the hydrogen bond is missing. $\endgroup$ Mar 23 at 18:50
  • $\begingroup$ This source says H2S does not form hydrogen bonds, and given that is a publication of the American Chemical Society, I am inclined to believe it acs.org/content/dam/acsorg/education/publications/ch8.pdf $\endgroup$
    – Waylander
    Mar 23 at 19:19
  • $\begingroup$ @Waylander yes but this is "pure" H2S. My question was not whether H2S could form hydrogen bonds with itself, but rather if it can with H2O... although the two might be related? I’m not sure. Thanks for this resource. $\endgroup$ Mar 23 at 19:21
  • 1
    $\begingroup$ The important thing is that H2S is an acid - that's what makes it way more soluble then CH4. Then again NH3 has much more pronounced acid-base reaction with water. Another issue is boiling point, or rather equilibrium vapor pressure. $\endgroup$
    – Mithoron
    Mar 23 at 19:59

1 Answer 1

6
$\begingroup$

Hydrogen sulfide is actually relatively strongly soluble in water compared with other low-polarity gases. In Earth Science SE, this answer explains that water-ice hydrates are most likely to be found with gases having low solubility in liquid water such as methane, and in particular methane hydrates are more common than hydrogen-sulfide hydrates because (on a molar basis) hydrogen sulfide is 80 times as soluble as methane.

Hydrogen sulfide is too weak as a hydrogen-bonding agent and an acid for either of these attributes to play much of a role in its relatively large solubility. Rather, the relatively large sulfur atom is more easily polarized than the hydrogen and carbon of methane, so hydrogen sulfide can interact more strongly with water through a dipole-induced dipole interaction. The induced polarity imparted to the hydrogen sulfide also contributes to its being more acidic than methane despite carbon and sulfur having nearly identical electronegativities.

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

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