Self-Ionization of Water and the pH Scale

The observation that pure water has low levels of H⁺ and OH⁻ is easily understood by the Brønsted–Lowry theory. In this self-ionization behavior, one molecule of water acts as the proton donor (acid) and a second molecule acts as the proton acceptor (base) to generate hydroxide and a hydronium ion (H₃O⁺).

  As noted in the equation above, thermodynamic equilibrium constants (K_eq) are formally defined by the activities (a_i) of each species. However, for dilute solutions where the solvent is in vast excess, the activity of the solvent (water in this case) is unity (a_H2O = 1) and the activities for the dilute ions are essentially equal to their molar concentrations (a_H3O⁺ ~ [H₃O⁺], a_OH⁻ ~ [:OH⁻]). As a result, the equilibrium dissociation constant for the self-ionization reaction of water (its acid dissociation constant) is simplified to

K_w = [H₃O⁺][: OH⁻]

At 25°C, K_w = 10⁻¹⁴ M², and since the self-ionization reaction yields equal amounts of H₃O⁺ and OH⁻, both are present at 10⁻⁷ M in pure water. As other acids or bases are dissolved in and react with water, the concentrations of H₃O⁺ and OH⁻ re-equilibrate such that their product always equals 10⁻¹⁴. Thus when [H₃O⁺] = 1 M, [OH⁻] = 10⁻¹⁴ M and vice versa. To simplify discussions of the relative acidity of aqueous solutions, we use the pH scale, where

pH = – log₁₀[H₃O⁺]

For pure water with [H₃O⁺] = 10⁻⁷ M, the pH = 7, which is also referred to as neutral pH because the concentrations of H₃O⁺ and OH⁻ ions are equal. Acidic solutions have higher concentrations of H₃O⁺ ([H₃O⁺] > [OH⁻]) and pH values < 7, whereas basic solutions have lower concentrations of H₃O⁺ ([H₃O⁺] < [OH⁻]) and pH values > 7.