Why is acetic acid more acidic than phenol?

Question

Acetic acid (ethanoic acid, $\\mathrm pK_\\mathrm a \\approx 5$) is more acidic than phenol ($\\mathrm pK_\\mathrm a \\approx 10$), which is reflected in their reactivity with a weak base such as sodium carbonate (adding acetic acid leads to effervescence of carbon dioxide). However, if we look at the structures of acetate and phenoxide ions, the delocalization of negative charge should be greater in the phenoxide ion as the charge moves inside the benzene ring. On the other hand, in the acetate ion, the negative charge is delocalized between two oxygen atoms only. How is it possible that acetic acid is still a stronger acid, even when its conjugate base has fewer resonance structures?

Answer 1

If you look at acetate, you have 2 nice resonance structures where the negative charge is spread over the 2 oxygen atoms. Acetate is a resonance-stabilized anion, which makes acetic acid tend to be acidic. Phenol, similarly, forms the resonance-stabilized anion phenolate. Phenolate also has multiple resonance structures. On first glance, the negative charge looks more delocalized than in acetate: not only can it be on the oxygen, it can also be on the ortho or para carbons!

I think the detail here is that oxygen is electronegative, much more than carbon. In acetate, the negative charge is delocalized onto both oxygens, which happen to be the electronegative atoms. In phenolate, the charge is trying to delocalize onto carbons which just aren't that electronegative. So phenolate is less stabilized by resonance, and phenol is a weaker acid.