Chemists use equilibrium constants, or Keq values, to express whether a reversible chemical reactions favours the formation of products or reactants. Acetic acid, or CH3COOH, for example, undergoes a partial reaction with water according to CH3COOH + H2O <---> CH3COO- + H3O+. An equilibrium constant greater than one would indicate that most of the acetic acid molecules convert to CH3COO-. An equilibrium constant less than one would indicate that most of the acetic acid molecules remain intact. When specifically applied to acids, chemists refer to equilibrium constants as acid-dissociation constants, or Ka values. These numbers are typically very small. The value for acetic acid, for example, is 0.000018, or 1.8 x 10^-5. Consequently, chemists sometimes express these quantities as pKa values, where pKa = -log Ka. In this case, the pKa of acetic acid becomes pKa = -log (0.000018) = 4.74.

Refer to a table of pKa values and locate the pKa of the compound under investigation. For example, the pKa of benzoic acid, a common food preservative, is about 4.20.

Multiply the pKa value by negative one to invert its sign. In the case of benzoic acid, 4.20 x (-1) = -4.20.

Calculate the Ka or Keq value by using a calculator to raise 10 to the power of the negative pKa. Continuing the previous example, if benzoic acid exhibits a pKa of 4.20, then its Ka = 10^(-4.20) = 6.31 x 10^-5, or 0.0000631. And because benzoic acid is acidic, its Keq and Ka values are numerically identical.