Equilibrium is based upon the chemical activities of the substances involved.

Chemical activity [local] is often expressed as the molar concentration of a solute or pressure of a gas.

In some cases the chemical activity is quite different than the actual concentration, especially when dealing with electrolytes in solution. In these cases, the concentration is multiplied by a factor, the activity coefficient, to convert it into a chemical activity.

At the high school level, the effects of activities usually are disregarded.

In a chemical reaction, some species may not have their activities change much as the result of reaction.

Consider the ionization of a solution of acetic acid in water:

CH₃COOH(aq) + H₂O(l) CH₃COO^-(aq) + H₃O⁺(aq)

In one liter of water, there are 55.5 moles of water; the concentration of water itself is 55.5M. This is a very large concentration compared to the other concentrations in the equation, which usually are 1 M or less. The change in molarity of water is too small in comparison to other quantities to change the value significantly. As a result, by convention, it is considered to be included in the constant (Keq).

The above reaction is an example of a homogeneous equilibrium in which all the reactants and products are in the same phase. Consider:

CaCO₃(s) CaO(s) + CO₂(g)

In writing an equilbrium expression for the decomposition of CaCO_3, the terms for the activities of the solid phases are included in the constant.

As a rule of thumb solids and liquids are not part of the equilibrium equations, whereas solutions and gases are included.

Terms conventionally dropped when evaluating K_eq are also dropped when evaluating the reaction quotient, Q.