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?
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.