BrønstedLowryDefinition

Brønsted–Lowry theory, developed independently by Johannes Brønsted and Thomas Martin Lowry in 1923, is a fundamental model of acid–base chemistry. It defines acids as substances that donate protons (H+) and bases as substances that accept protons. An acid–base reaction proceeds through the transfer of a proton from the Brønsted acid to the Brønsted base, yielding a conjugate base and a conjugate acid. The pair formed by the acid and its conjugate base, and the pair formed by the base and its conjugate acid, are called conjugate acid–base pairs.

In aqueous solution, common examples include the reaction of hydrogen chloride with water, where HCl donates a proton to water to form Cl− and H3O+. Ammonia can act as a base by accepting a proton from water, yielding NH4+ and OH−. Water itself is amphoteric, able to donate a proton to become OH− or accept a proton to become H3O+. The strength of an acid is quantified by its dissociation constant (Ka); the strength of a base is given by Kb. For a given conjugate pair in the same solvent, Ka × Kb = Kw, the ion-product of water, linking acid and base strengths.

The Brønsted–Lowry framework is widely applicable, extending beyond water to non-aqueous solvents and to many biological and catalytic processes. It emphasizes the role of proton transfer rather than the solvent, metal, or electron-pair donation alone. Limitations include cases where proton transfer does not occur or is not the rate-determining step, and the influence of solvent on acid and base strength.

Overall, the theory provides a simple and general lens for understanding acid–base reactions across chemistry and biology.