# Why do most carboxylic acids have high pKa (~5) in spite of having a conjugate base ion that is stabilized by resonance?

This is from my textbook:

Carboxylic acids owe their acidity ($\mathrm pK_\mathrm a$ of about $5$) to the resonance-stabilized carboxylate anions formed by deprotonation.

Why are they such weak acids? If the anions are resonance-stabilized, I would expect them to be much stronger acids.

Stated another way: this screenshot of the textbook page shows a representative protonation/deprotonation equilibrium for a carboxylic acid:

Why is the reaction favoured in the reverse direction if the carboxylate ion is more stable than the reactant?

• Resonance is not all that important in the grand scheme of things. If you really want to focus on just one factor, it should be the H-X bond strength. – orthocresol Jun 7 '18 at 18:42
• Note also that you're asking about the aqueous stability of the anion. Now, solvation spheres and dipole effects and dielectric effects are all in play and make this a very complicated problem. – Zhe Jun 7 '18 at 19:40
• About 4-5 is typical pKa for acid with charge of 1 delocalised over 2 oxygen atoms strong acids need need delocalisation over 3 O atoms. – Mithoron Jun 7 '18 at 20:07

Everything is relative in acid-base reactions; the equilibrium will favor the side with the weaker acid. A carboxylic acid with $\ce{pKa = 5}$ will not protonate $\ce{H2O}$ significantly because $\ce{pKa(H3O+) = -0.7}$, but it will protonate something more basic like $\ce{NH3}$ because $\ce{pKa(NH4+) = 9.24}$.
Note: the pKa of $\ce{H3O+}$ is debated, see What is the pKa of the hydronium, or oxonium, ion (H3O+)? for more discussion.