An acid-base titration curve plots solution pH against the amount or volume of titrant added. Its shape records which chemical regime controls the pH before, near, and after the equivalence point.

Consider titrating a weak monoprotic acid $\mathrm{HA}$ with a strong base.

Before titrant is added

The pH is set by the weak-acid equilibrium and its initial concentration.

Before equivalence

Added hydroxide reacts essentially stoichiometrically with the weak acid:

$$\mathrm{HA+OH^-\rightarrow A^-+H_2O}.$$

As long as appreciable $\mathrm{HA}$ and $\mathrm{A^-}$ remain, the mixture is a buffer. Its pH is conveniently described by the conjugate-pair ratio.

At the half-equivalence point, exactly half of the original acid has been converted to conjugate base, so

$$n_{\mathrm{A^-}}=n_{\mathrm{HA}}$$

and therefore

$$\boxed{\mathrm{pH}=pK_a}.$$

This provides a direct experimental way to estimate the $pK_a$ of a weak monoprotic acid.

At equivalence

All initial $\mathrm{HA}$ has been converted stoichiometrically to $\mathrm{A^-}$. The conjugate base hydrolyzes water,

$$\mathrm{A^-+H_2O\rightleftharpoons HA+OH^-},$$

so the equivalence-point pH is greater than 7 at $25,^\circ\mathrm C$ for a weak-acid/strong-base titration.

After equivalence

Excess strong base determines the hydroxide concentration after accounting for dilution by the total solution volume.

A strong-acid/strong-base titration is simpler: before equivalence pH comes from excess strong acid, at equivalence the solution is approximately neutral at $25,^\circ\mathrm C$, and after equivalence pH comes from excess strong base.

Weak-base/strong-acid titrations are the conjugate mirror image: they contain a buffer region and have an acidic equivalence point.

The curve therefore cannot be calculated using one formula everywhere. The correct method changes with chemical regime: equilibrium before reaction, stoichiometric neutralization, buffer equilibrium, conjugate-ion hydrolysis at equivalence, and excess titrant after equivalence.