An equilibrium is heterogeneous when the reacting species occupy more than one physical phase. Pure solids and pure liquids require special treatment in equilibrium expressions because changing how much of a pure phase is present does not change its effective concentration while that phase remains present.

Consider dissolution of silver chloride:

$$\mathrm{AgCl(s)\rightleftharpoons Ag^+(aq)+Cl^-(aq)}.$$

The introductory equilibrium expression is

$$K_c=[\mathrm{Ag^+}][\mathrm{Cl^-}],$$

not

$$\frac{[\mathrm{Ag^+}][\mathrm{Cl^-}]}{[\mathrm{AgCl}]},$$

because pure solid $\mathrm{AgCl}$ has a fixed composition and its contribution is absorbed into the equilibrium constant.

The same principle applies to a pure liquid such as water when it acts as a separate bulk phase: its activity is effectively constant and is incorporated into $K$ rather than written as a variable concentration factor.

This does not mean solids and liquids are chemically irrelevant. They can be essential reactants or products. It means that, while a pure phase remains present, adding more of that same pure phase does not by itself alter the equilibrium relation among the variable species.

For example, placing extra solid $\mathrm{AgCl}$ into an already saturated solution does not increase the equilibrium concentrations of $\mathrm{Ag^+}$ and $\mathrm{Cl^-}$ merely because more solid is present. The amount of solid changes, but its pure-phase composition does not.

If the solid is completely consumed, however, the situation changes because that phase is no longer available to participate in the equilibrium.

Heterogeneous equilibrium expressions are the foundation for quantitative solubility equilibria such as $K_{sp}$ and for many phase and gas-solid chemical equilibria.