Why is acetylsalicylic acid still more acid than benzoic acid?

Question

In this question, it is explained why salicylic acid is a stronger acid than benzoic acid. So, why acetylsalicylic acid (aspirin), whose conjugate base cannot hydrogen bond since there is no hydrogen, is still more acid than benzoic acid?

• aspirin: $\mathrm{p}K_\text{a}$ = 3.5
• benzoic acid: $\mathrm{p}K_\text{a}$ = 4.2

Acetoxyl is a moderate activating group for the aromatic electrophilic substitution, so it should mean that the resonance-driven electron-donating effect prevails over the inductive electron-withdrawing effect, therefore it should raise the $\mathrm{p}K_\text{a}$ of the acid.

Is it different between electrophilic substitution reactivity and acidity? Does the inductive effect prevail in this case?

Answer 1

The acetyl moiety in acetylsalicylic acid has a $(-I)$-effect, i.e. compared to a proton attached to the aromatic ring (as in the case of benzoic acid), it withdraws electron density out of the aromatic ring. This provides an additional stabilization of the conjugated base, that is equivalent with rendering acetylsalicylic acid a stronger Brønsted acid than benzoic acid. Hence, the $pK_a$ of acetylsalicylic acid is lower than the one of benzoic acid.

Regarding the second paragraph, please notice $pK_a$ is a thermodynamic property -- where time does not play a role. In contrast, the substitution reactivity you likely refer to (Hammett parameters) refers to reaction kinetics and rate constants. Here, time does play a role.