When I first learned about acid strength trends in general chemistry, I remember it was explained in terms of the character of the H-A bond. The more polar the H-A bond, the easier it was to heterolytically split it into H$^{+}$ and A$^{-}$. This explained why acid strength increased from left to right across the periodic table: as you went from left to right, electronegativity increases, so the H-A bond becomes more polar, hence more acidic. The strength of the H-A bond also contributed to the acidity; the weaker the bond, the easier to break it into A$^{-}$ and H$^{+}$, so the higher the acidity. This explained why acidity increased as you went down a group; the H-A bond became weaker, leading to higher acidity.

However, I have also seen acidity strength explained solely through the stability of the conjugate base, without mention of the nature of the H-A bond. When viewing acidity through the base stability point of view, we can explain the increase in acidity from left to right as a result of charge stabilization: in the conjugate base, a more electronegative atom is better able to stabilize the excess negative charge, therefore the conjugate base becomes more stable as the electronegativity of the atom attached to the acidic proton increases. The trend down a group can also be explained in terms of charge stabilization: as the size of an atom increases, any excess charge on the atom is spread out due to the diffuse orbitals, leading to charge stabilization. I think because I so often saw acid strength just explained in terms of base stability, I started to think that the H-A bond explanation was just a simplified version of the base stability argument.

However, I recently saw this question, which made me reassess my thinking: enter image description here

Here, you can not use base stability to explain the acidity difference, because both acids have the same conjugate base. So in this case, any difference in acidity must come from the stability of the acid, not the base. In other words, the differences in the $\Delta$G's for the two dissociations must be due to a difference in the Gibbs Free Energy of the acids. This makes me think that the discussions in general chemistry about the H-A bond were really explaining the $\Delta$G of acid dissociation, which was independent of the stability of the base. Now, I am starting to think that weak and polar H-A bonds lead to more acidity because they tend to create a more exergonic reaction (if this is the case, could someone explain why?).

This leads to my overall question: is explaining acidity through the nature of the H-A bond just another way of explaining it through conjugate base stability, or does the H-A bond tell us something about the $\Delta$G of the acid dissociation that we would not know from looking at conjugate base stability? Is it just coincidental that the nature of the H-A bond and the stability of the conjugate base contribute to the same trends in acid strength?

  • $\begingroup$ That is an interesting and a somewhat silly question because in aqueous solution the two compounds are the same. If Either were initially pure and added to water the eventual equilibrium would be identical giving the same [H3O+]. The deciding factor would be which of the two tautomeric forms of the acid is present in greater concentration in the solution. A Raman or IR spectrum might give the answer. $\endgroup$
    – jimchmst
    Commented Aug 3, 2022 at 19:23
  • $\begingroup$ I looked up thioacetic acid it is a considerably stronger acid than acetic and exists almost entirely as the thiol tautomer; the difference from the stronger carbonyl than thionyl bond. This raises question about the acidity of the much less stable tautomer. $\endgroup$
    – jimchmst
    Commented Aug 3, 2022 at 19:47

1 Answer 1


This is a good question, and Ill try my best!

The acidity of an acid is solely dependent on the nature of the HA bond. When learning about acidity and basicity, the concept of conjugate base stability is invariably taught to help explain the strength of various acids. But, molecules don't have brains. They arent thinking, "When I lose my proton, I'll be super stable, so I wanna be really acidic!"

The electronics of the starting acid are what determines its acidity, and as a result, things that are more acidic have relatively stable conjugate bases. The stability of the CB only helps explain its acidity (and 99% of the time, it works).

OH vs CH? Oxygen is more electronegative, pulling electron density away from H, so the electron pair done (the base) can use its electrons to grab that hydrogen more easily.

Carboxylic acid vs alcohol? Oxygen is donating electron density to the carbonyl, resulting in the oxygen of the OH bond being more electron deficient. But electronegative oxygen wants electrons, so it pulls even MORE electron density from H, making the proton even more grabbable than in an alcohol.

Thiol vs alcohol? Yeah sure, oxygen is more electronegative, but the orbital overlap between sulfur and hydrogen is garbage, that is, the energy gap between H's s orbital and sulfur sp3 orbital is much larger than that of an alcohol. Thus the S-H bond is less stable kinetically, and so bases come and snatch that hydrogen.

To answer your question completely, it is NOT coincidental that the acidity of HA correlates with the stability of the conjugate base. However, rather than thinking that HA is more acidic because the CB is more stable, think about how the two concepts go hand in hand and influence each other.

Here's another example: The iodide in tert-butyl iodide doesn't leave because it forms a more stable tertiary carbocation; the molecule doesn't know that it can form a stable tertiary carbocation. Rather, it leaves because the adjacent sigma C-H bonds donate electrons density to the C-I antibonding orbital, significantly labilizing the C-I bond. Thus, with enough energy, iodine leaves, forming the tertiary carbocation.


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