An acid-base neutralization transfers protons between an acid and a base. For a strong acid and strong base in water, the essential net ionic reaction is

$$\mathrm{H_3O^+(aq)+OH^-(aq)\rightarrow2H_2O(l)}.$$

In quantitative work, the balanced reaction determines how many moles of acid and base react.

For a monoprotic acid $\mathrm{HA}$ reacting 1:1 with hydroxide,

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

If $n_{\mathrm{HA}}$ moles of acid are present, exactly $n_{\mathrm{HA}}$ moles of hydroxide are required for stoichiometric neutralization.

The equivalence point is the composition at which the amount of titrant added is exactly stoichiometrically equivalent to the amount of analyte according to the balanced reaction. It is a mole relation, not a particular pH.

Example

A $25.00,\mathrm{mL}$ sample of $0.1000,\mathrm M$ monoprotic acid contains

$$n=(0.1000)(0.02500)=2.500\times10^{-3},\mathrm{mol}.$$

If titrated with $0.1250,\mathrm M$ $\mathrm{NaOH}$, equivalence requires the same amount of hydroxide, so

$$V=\frac{n}{c}=\frac{2.500\times10^{-3}}{0.1250} =0.02000,\mathrm L =20.00,\mathrm{mL}.$$

The equivalence-point pH depends on what species remain after neutralization. A strong acid with a strong base is approximately neutral at $25,^\circ\mathrm C$, but a weak acid titrated by a strong base leaves its conjugate base at equivalence and therefore has a basic equivalence-point pH. Likewise, a weak base titrated by a strong acid gives an acidic equivalence point.

This distinction is fundamental: equivalence is defined by stoichiometry; pH is determined by the acid-base chemistry of the resulting solution.