What is causing the buffer region in a weak acid - strong base titration?
The linked question's answers were too long for me, so I'll attempt to summarize quickly.

The equation that tells you everything you need to know is the Henderson-Hasselbalch equation:

$$\mathrm{pH} = \mathrm{p}K_{\mathrm a} - \log\frac{\ce{[A-]}}{\ce{[HA]}}$$

Here $\ce{HA}$ is a weak acid, and it's conjugate base is also weak. But we can imagine also a weak base with a weak conjugate acid. The magic happens when you've half neutralized a weak acid or base, creating a mixture of a weak acid and a weak base. The fraction on the right is fairly close to 1 (within an order of magnitude, let's say).

As you add more acid, the weak base can neutralize it. As you add more base, the weak acid can neutralize it. Assuming these amounts are reasonably small, the logarithm of the fraction on the right doesn't change by much, so the overall pH of the solution does not change by much. For example, changing the fraction on the right by a factor of 10 only affects the pH value by $\pm 1$.

Of course, if you add enough acid or base to neutralize (or close to neutralize) one of the buffer components, then the buffer stops working. This is also when the fraction gets very large or goes to zero and even the logarithm can't change that.