# Why is the vertical part of a titration curve longer for a strong acid by a strong base than for a weak?

As the title says, I would like to understand chemically why this occurs. My theory behind it is that in the former case both are completely dissociated, and so at the equivalence point (and approaching it), everything is controlled increasingly by the autoionization of water, whereas in weak acids there is also a much more sizable contribution of $H^+$ by the acid itself, so the change is not as drastic. If this is correct, I'd like further explanation as to why it is, since I don't quite feel comfortable enough with it intuitively.

This is not a duplicate. I am not asking why the strong acid curve is steep, but rather why the strong acid curve is steep while the weak acid one is not.!

Weak acid and strong acid titration curves with addition of a strong base

• Could you provide an example of titration of a strong acid by a strong base? Usually titrations are of weak acids and bases Mar 9, 2015 at 5:37
• This is irrelevant. The tendency is the same for all such titrations. However, consider the titration of 0.1 M HCl by 0.2 M NaOH, for instance. Mar 9, 2015 at 15:29
• In regards to your edit, is this an example of what you're talking about? Mar 13, 2015 at 7:10
• @JohnSnow, yes, it is exactly the example. Apr 12, 2015 at 19:13

Compare (1) the titration of $\ce{NaOH}$ with $\ce{HCl}$ and (2) that of $\ce{NaOH}$ with $\ce{AcOH}$.
At the equivalence point of each titration, solution (1) contains $\ce{Na+}$, $\ce{Cl-}$, and $\ce{H2O}$, with negligible concentration of $\ce{H+}$ and $\ce{OH-}$. In contrast, solution (2) contains $\ce{Na+}$, $\ce{AcO-}$, and $\ce{H2O}$, and in particular $\ce{AcO-}$ will react with water to form a buffer solution of $\ce{AcO-/AcOH}$.