D24.2 Autoionization of Water

H₂O is the base of its conjugate acid H₃O⁺, and H₂O is also the acid of its conjugate base OH^–. (However, H₃O⁺ is not the conjugate acid of OH^–; these two species are not a conjugate acid-base pair because their structures do not differ by a single H⁺.)

Hence, H₂O can react as an acid or a base depending on the other species involved in the reaction. In pure water, H₂O acts as both acid and base—a very small fraction of water molecules donate protons to other water molecules:

This type of reaction, in which a substance ionizes when one molecule of the substance reacts with another molecule of the same substance, is referred to as autoionization.

Pure water undergoes autoionization to a very slight extent at 25 °C. The equilibrium concentrations of H₃O⁺ and OH^– give an autoionization constant for water, K_w = 1.0 × 10⁻¹⁴ at 25 °C.

Because it is the mathematical product of concentrations of two ions, it is also called the ion-product constant for water.

Water is an example of an amphiprotic chemical species, a molecule that can either gain a proton or lose a proton in a Brønsted-Lowry acid-base reaction. Amphiprotic species are also amphoteric, a more general term for a species that may act either as an acid or a base by any definition (not just the Brønsted-Lowry definition). For example, the bicarbonate ion is also amphoteric: