Dissolving a salt can produce an acidic, basic, or approximately neutral solution because the ions themselves may react with water as Brønsted-Lowry acids or bases. This reaction of an ion with water is often called salt hydrolysis.

The key question is whether each ion is the conjugate partner of a strong or weak acid or base.

Ions with negligible acid-base behavior

The conjugate base of a strong acid is extremely weak. Ions such as $\mathrm{Cl^-}$ and $\mathrm{NO_3^-}$ therefore have little tendency to accept protons from water.

Likewise, cations derived from strong bases such as $\mathrm{Na^+}$ and $\mathrm{K^+}$ have negligible Brønsted acidity in ordinary introductory aqueous chemistry.

A salt such as $\mathrm{NaCl}$ therefore produces an approximately neutral solution.

Basic anions

The conjugate base of a weak acid can accept a proton from water. Acetate behaves as

$$\mathrm{CH_3COO^-+H_2O\rightleftharpoons CH_3COOH+OH^-}.$$

Its base constant is related to the acid constant of acetic acid by

$$K_b(\mathrm{CH_3COO^-})=\frac{K_w}{K_a(\mathrm{CH_3COOH})}.$$

A soluble acetate salt such as sodium acetate therefore gives a basic solution.

Acidic cations

The conjugate acid of a weak base can donate a proton to water. Ammonium behaves as

$$\mathrm{NH_4^++H_2O\rightleftharpoons NH_3+H_3O^+},$$

with

$$K_a(\mathrm{NH_4^+})=\frac{K_w}{K_b(\mathrm{NH_3})}.$$

An ammonium salt such as $\mathrm{NH_4Cl}$ therefore gives an acidic solution.

When both ions have appreciable acid-base character, the outcome depends on their relative strengths rather than on a universal salt rule.

The pH can be calculated by treating the relevant ion exactly as a weak acid or weak base with the concentration supplied by salt dissociation. Thus salt hydrolysis is not a separate kind of equilibrium: it is ordinary conjugate acid-base chemistry applied to ions released by an electrolyte.