why is it that the strength of an acid/base does not affect the volume required for titration (rather, the concentration of the acids affects the volumes)? E.g. if both acid/base solutions are 0.1M it does not matter whether one or the other is weak or strong - the same volumes for neutralisation will be required.
This makes sense calculation wise for tritrations to me (n=cv), but, I thought that for neutralisation, it's based on H+ + OH- --> H2O, and the concentration of H+ in weak acids is less than the concentration of the acid, wouldn't this affect the volume neeeded for neutralisation?
but it doesn't. ._.
This is the exact same reaction that occurs in all titrations. Which is why we never use a weak base and weak acid because that will cause concentrations to matter too much.
At the instant one of the species is strong, things happen. Note that the weak one, say, weak acid, does not fully ionise by itself.
But what happens when you add a strong base? The OH-'s there are fully ionised. Clearly the OH- will react with the H+.
Thus the weak acid lost H+. Hence by Le chatelier's principle, the system will want to replace the most H+, thus more of it gets ionised.
But this H+ also gets used by the OH-. Hence the weak acid is ultimately forced to keep ionising until all the OH- is gone. So ultimately the H+ gets forced into ionising completely because the base is strong and is, in a way, "hydroxide abundant"