[minijumbuk]

Question:
If you have a pH 3 acetic acid solution and a pH 3 HCl solution, predict which one will require more alkaline solution to neutralise it in a titration.

My answer is that the acetic acid will consume more alkaline solution than HCl, because it's a weak acid.
CH₃COOH + H₂O <--> CH₃COO^- + H₃O⁺

So according to this equation, when more Hydronium ions are reacted to form water, the equilibrium equation shifts to the right by Le Chatelier's Principle, and hence replaces the lost hydronium ions. This will result in more base needed to neutralise it.
This does not occur in HCl because it's a strong acid, and hence the number of hydronium ions is fixed from the start.

Is that correct, or is there some technical flaw here?

[jansenwrasse]

Do some reading on ionization constants,  and dissociation in water of acids.  But what you have written is correct

[MikeS]

If there are the same mols of monoprotic acids, they will both consume the same amount of base to reach neutralization.

Now buffers like acetic acid will change pH more slowly at first when you titrate it with a strong base, but both acetic acid and HCl will be neutralized by an equal amount of NaOH if you start off with an equal amount of mols of both.

In your question, however, it doesn't say how many mols of HCl or acetic acid. It only says that both are at pH 3. My guess is that since HCl is a strong acid, it will take only a small amount to get the pH to drop to 3. Acetic acid is a weak acid so it may take a large amount of it to reach pH 3 (pKa of acetic acid 4.75)

So my answer is the same as yours. At pH 3, acetic acid will take more NaOH to neutralize