Chemistry 1923

The Proton Theory of Acids and Bases

Johannes Brønsted & Thomas Lowry

An acid is anything that gives up a proton; a base, anything that takes one.

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In depth · the introduction

An acid and a base aren't so much two kinds of thing as two halves of one move: one hands over a proton, the other takes it.

The big idea

For a long time, acids were defined by what they did in water: an acid made hydrogen ions (H⁺), a base made hydroxide ions (OH⁻). In 1923 two chemists, working separately, found a simpler and deeper definition. Forget the water and the hydroxide; keep just the proton. An acid is anything that can give away a proton — a single hydrogen ion. A base is anything that can take one.

That small shift has a big consequence. Since a proton can't simply float free, an acid can only let go of its proton if a base is there to catch it. So every acid needs a base, every reaction is a hand-off, and the same molecule can play either part depending on its partner. Water, famously, plays both: it catches a proton from an acid and gives one to a base.

How it came about

The reigning theory, from Svante Arrhenius in the 1880s, worked well for simple acids in water but creaked elsewhere — it could not neatly explain why ammonia, which contains no hydroxide, behaves as a base, or what happens in solvents other than water. In 1923 Johannes Brønsted in Copenhagen and Thomas Lowry in Cambridge each published the proton view within months of one another, unaware they were saying the same thing.

The same year, in America, Gilbert Lewis offered yet another definition, built on shared electron pairs. Chemistry ended up keeping both: the Danish–English proton theory as the everyday language of acids in water, and the Lewis theory for the broader cases. Lowry's contribution is often the one left out of the textbook shorthand 'Brønsted–Lowry', but the two reached the idea independently and at the same moment.

Why it mattered

By making acidity a relationship rather than a fixed property, the proton theory tied a sprawl of facts together under one rule — and, crucially, gave them numbers. The familiar pH scale, the pKa that ranks how readily an acid lets go of its proton, titration curves, and the buffers that keep your blood and every careful experiment at a steady acidity — all of it is the Brønsted–Lowry picture in daily use.

A way to picture it

Think of a proton as a hot potato. An acid is a hand willing to toss it; a base is a hand willing to catch it. No throw can happen with only one hand — you always need a tosser and a catcher. 'Strong' and 'weak' just describe how eager a hand is to let the potato go. And some hands, like water's, will throw or catch depending on who they are playing with.

Where it sits

It refined the Arrhenius theory that came before it and runs in parallel with the Lewis electron-pair theory introduced the same year. Downstream, it underlies the chemistry of life — the bicarbonate buffer that steadies your blood — and of the planet: the slow acidification of the oceans as they absorb carbon dioxide is a Brønsted–Lowry equilibrium playing out across the whole sea.

The original document

Original source text

J. N. Brønsted · Recueil des Travaux Chimiques des Pays-Bas 42 (1923): 718–728 (in German) · T. M. Lowry · J. Soc. Chem. Ind. 42 (1923): 43–47

Brønsted, Copenhagen — acids and bases as a conjugate pair

Brønsted's paper sets out to free the concepts of acid and base from the solvent. Where the older theory called an acid a source of hydrogen ions and a base a source of hydroxide ions in water, he keeps only the proton: an acid is any species — a neutral molecule, a cation, or an anion — that can give up a hydrogen ion, and a base is any species that can take one up. (Paraphrase of his definition.)

The two are bound together. Strip the proton from an acid and what remains is its base; add it back and the base becomes the acid again. Brønsted writes this relationship as the scheme below, the conjugate acid–base pair — the organising idea of the whole paper.

Acid ⇌ Base + H⁺

Because a free proton cannot exist on its own in solution, an acid can only act when a base is present to receive the proton. Every acid–base reaction is therefore a proton transfer between two such pairs; the solvent is usually one of the partners. Water is both at once — it takes a proton to become H₃O⁺ when it meets an acid, and gives one to become OH⁻ when it meets a base. (Paraphrase.)

[ … ]

Lowry, Cambridge — the uniqueness of hydrogen

Working independently, Lowry reached the same proton-centred view, framed around the peculiar nature of the hydrogen nucleus — a particle so small and so reluctant to exist alone that 'acidity' is really about whether a molecule will release it to something else.

It is a remarkable fact that strong acidity is apparently developed only in mixtures and never in pure compounds. Even hydrogen chloride only becomes an acid when mixed with water.

That single observation captures the new picture: pure hydrogen chloride is not yet an acid; it becomes one only when water is there to accept its proton. Acidity is a relationship, not a fixed property of a substance.