Brønsted–Lowry theory
/ BRUN-sted LOW-ree /
Picture a simple game of catch where one player throws a ball and another catches it. Brønsted–Lowry theory says that is exactly what acids and bases are doing, except the ball is a single proton (a hydrogen ion, H+). The thrower is the acid; the catcher is the base.
Formally, a Brønsted–Lowry acid is any species that donates a proton, and a Brønsted–Lowry base is any species that accepts one. This view is broader and cleaner than the older idea that bases must release hydroxide: now anything that can grab a proton counts as a base, even if it contains no OH at all. Every acid–base reaction becomes a proton hopping from a donor to an acceptor, and the same molecule can play either role depending on its partner.
This matters because it explains acid–base behaviour in any solvent, not just water, and it naturally pairs up reactants with products as conjugate partners. Its limit is that it is all about the proton: reactions where electron pairs, not protons, change hands need the broader Lewis picture instead.
When ammonia dissolves in water, NH3 grabs a proton from H2O to become NH4+, leaving OH-. Here ammonia is the base and water is the acid — even though water is usually thought of as neutral.
Acid–base reaction = a proton thrown from a donor (acid) to an acceptor (base).
Because acid or base depends on the partner, a substance like water is amphoteric — it can donate a proton to ammonia yet accept one from hydrogen chloride. Roles are relative, not fixed.