pH

pH, term indicating the hydrogen ion concentration of a solution, a measure of the solution's acidity. The term (from French pouvoir hydrogène,”hydrogen power”) is defined as the negative logarithm of the concentration of H⁺ ions (protons): pH = -log₁₀[H⁺], where [H⁺] is the concentration of H⁺ ions in moles per litre. Because H⁺ ions associate with water molecules to form hydronium (H₃O⁺) ions (see Acids and Bases), pH is also often expressed in terms of the concentration of hydronium ions.

In pure water at 22° C (72° F), H₃O⁺ and hydroxyl (OH^-) ions exist in equal quantities; the concentration of each is 0.10⁷ moles/litre. Consequently, the pH of pure water is -log (0.10⁷), which equals log 10⁷, or 7. If an acid is added to water, however, an excess of H₃O⁺ ions is formed; their concentration can range between 0.10⁶ and 0.10 moles/litre, depending on the strength and amount of the acid. Therefore, acid solutions have a pH ranging from 6 (for a weak acid) to 1 (for a strong acid). Inversely, a basic solution has a low concentration of H₃O⁺ ions and an excess of OH^- ions, and the pH ranges from 8 (for a weak base) to 14 (for a strong base).

The pH of a solution can be measured by titration, which consists of the neutralization of the acid (or base) by a measured quantity of base (or acid) of known concentration, in the presence of an indicator (a compound the colour of which depends on the pH). The pH of a solution can also be determined directly by measuring the electric potential arising at special electrodes immersed in the solution (see Chemical Analysis).