What is the difference between the titration of a strong acid with a strong base and that of the titration of a weak acid with a strong base?
The answer is A.

Let's look at a weak acid: acetic acid. It has a $K_a$ of $1.8\cdot  10^{-5}$. From this you can calculate how much it will dissociate in water. Let's say we have a concentration of $1~\mathrm{M}$:

\begin{align}
K_a &amp;= \frac{[\ce{CH3COO^{-}}][\ce{H+}]}{[\ce{CH3COOH}]}\\
1.8 \cdot 10^{-5} &amp;= \frac{x \cdot x}{1~\mathrm{M} - x}\\
x &amp;= 4.233 \cdot 10^{-3}\\
\end{align}

This means, at equilibrium, $[\ce{CH3COOH}] = 1~\mathrm{M}- 4.233 \cdot 10^{-3} = 0.995~\mathrm{M}$. As you can see, by far the largest part is not dissociated!

B is not true, you can still observe the equivalence point. Have a look at the following picture, that shows the titration curve for weak / strong acid titrated with strong base. When choosing an indicator for colorimetric titration select one so that the pH jump at the equivalence point contains the interval $\mathrm{p}K_a\pm1$. Phenolphthalein has a $\mathrm{p}K_a \approx 9$, so to decide if it is a suitable indicator you would check if the pH jumps from 8 to 10 at the equivalence point. It does:



C is not true either. If it was a very slow reaction, you could not titrate accurately, as you would have to wait a long time between every drop of $\ce{NaOH}$.