Brønstedtype

Brønsted-type refers to the Brønsted–Lowry theory of acids and bases, formulated independently by Johan Brønsted and Thomas Lowry in 1923. In this framework, an acid is any substance that donates a proton (H+) and a base is any substance that accepts a proton. A reaction is viewed as proton transfer from an acid to a base, AH + B → A− + BH+.

Acids and bases form conjugate pairs: after donating a proton, an acid becomes its conjugate base, and after accepting a proton, a base becomes its conjugate acid. The strength of an acid in a given solvent is determined by its tendency to donate protons; the strength of a base is the tendency to accept protons. In aqueous solution, this is quantified by dissociation constants Ka and Kb, with Ka × Kb = Kw, and by pKa and pKb values. The Henderson–Hasselbalch equation relates pH to pKa for buffer systems.

Brønsted-type theory is not limited to water; it applies to proton transfers in many solvents and in the gas phase. Some substances, like water, ammonia, and amino acids, are amphoteric—they can act as acids or bases depending on the reaction partner and conditions.

Limitations include solvent effects and the fact that extremely strong acids or bases can drive reactions through mechanisms beyond simple proton transfer. Nevertheless, the Brønsted framework remains a foundational tool in organic, inorganic, and biochemical contexts for understanding acid–base behavior, catalysis, and buffering.

Common examples illustrate the concept: hydrochloric acid acts as a Brønsted acid, ammonia as a Brønsted base, and their conjugate bases chloride and ammonium embody the conjugate-pair relationship.