An acid gives a proton and a base takes it: Arrhenius, Brønsted–Lowry and Lewis

Arrhenius: hydrogen ions and hydroxide ions in water

By 1890 the Swedish chemist Svante Arrhenius had set out the first useful theory of acids. In his terms an Arrhenius acid is a substance whose molecules carry a hydrogen atom that comes away in water as a hydrogen ion. An Arrhenius base likewise releases hydroxide ions it already holds when it dissolves. In this section, “acid” and “base” carry those meanings.

Strictly, then, an Arrhenius acid has to contain hydrogen, though not every hydrogen counts. The three hydrogens on the CH₃ group of acetic acid are non-acidic. Sulfuric acid and phosphoric acid each have more than one hydrogen that counts, which makes them polyprotic acids.

Hydrogen chloride fits the acid definition, containing hydrogen and giving hydrogen ions in water. Sodium hydroxide fits the base definition: its solid is two interpenetrating lattices, one of sodium ions and one of hydroxide ions, so the hydroxide is there before any water arrives.

An acid and a base react to give water and a salt, an ionic compound; the reaction is neutralisation. Hydrogen chloride and sodium hydroxide give water and sodium chloride. PubChem’s Laboratory Chemical Safety Summaries list the hazard classifications of both. Arrhenius’s scheme cannot explain why ammonia, with no hydroxide ion in it, behaves as a base rather than an acid.

Brønsted–Lowry: a proton needs somewhere to go

Arrhenius’s hydrogen ion has a second problem: it is no more than a proton, the bare nucleus of a hydrogen atom that has lost its one electron.

The proton’s charge density, its charge for the space it takes up, is so high that it is drawn to any nearby excess of negative charge. In water that is a lone pair on the oxygen, outer electrons in no bond, as counted in valence-shell electron-pair repulsion (VSEPR) shapes. The proton buries itself in the pair, bonding to it and making the hydronium ion, H₃O⁺. A bare proton does not survive in water.

In 1923 the Danish chemist J. N. Brønsted proposed a new definition, with similar ideas from T. M. Lowry that year. A Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base a proton acceptor; from here on, “acid” and “base” alone mean these. Neither can act without the other: an acid needs a base present to accept its proton.

In water, one HCl molecule and one H₂O molecule become a chloride ion, Cl⁻, and a hydronium ion, so water is the base. What Arrhenius called yielding hydrogen ions is a proton passing to a water molecule.

Ammonia takes a proton from water instead, giving an ammonium ion, NH₄⁺, and a hydroxide ion, OH⁻. The solution is alkaline, holding an excess of hydroxide ions, but the hydroxide came from water. Ammonia holds no hydroxide of its own to release, so it is a Brønsted–Lowry base but not an Arrhenius base; its nitrogen’s lone pair accepts the proton.

Ammonia’s figure holds one nitrogen and three hydrogens: no oxygen, so no hydroxide.

Water, the base towards hydrogen chloride, is the acid towards ammonia.

Conjugate pairs, and water on both sides

Two species are conjugate when they differ by the gain or loss of a single proton. Read backwards, both reactions above are built from such pairs. Cl⁻ is the conjugate base of HCl, and H₃O⁺ is the conjugate acid of water. On the ammonia side, NH₄⁺ is the conjugate acid of ammonia and OH⁻ the conjugate base of water.

So water stands on both sides: a base towards hydrogen chloride and an acid towards ammonia. A substance able to act as either is amphiprotic. What water’s lone pairs do between neighbouring molecules is the subject of intermolecular forces.

Towards ammonia a proton leaves one of water’s O–H bonds, and from hydrogen chloride one lands on its oxygen. Water’s figure shows the three atoms and two O–H bonds behind both roles.

Water can also play both parts with itself. One molecule acting as an acid passes a proton to another acting as a base, giving H₃O⁺ and OH⁻; this is the autoprotolysis of water. The hydrogen carbonate ion of sodium bicarbonate is amphiprotic as well.

Strong and weak: how completely the proton leaves

In water, two bases compete for an acid’s proton: water, and the acid’s own conjugate base. If water wins overwhelmingly, the acid is a strong acid; the word names the outcome and nothing more. Hydrogen chloride and nitric acid are effectively 100 per cent dissociated: the proton has gone to water.

Chloride, the conjugate base of HCl, cannot hold onto the proton against water or hydronium. In water H₃O⁺ is the strongest acid that can exist, and the strong acids, though they differ as proton donors, are all completely dissociated. Hydroxide is the strongest base possible in water, and sodium hydroxide, a lattice of its ions, is a strong base.

Hydrogen chloride’s strength has no single cause; the factors include how hard the bond to hydrogen is to break and the energy released as the ions are hydrated. Electronegativity alone tends not to explain it; what it predicts for the H–Cl bond is in polarity.

Most organic acids, acetic acid among them, lose the contest: only a small fraction is dissociated. Each is a weak acid, its conjugate base keeping the proton better than water can. Hydrogen cyanide, HCN, is weak because its proton shares the cyanide ion’s lone pair more effectively than water’s.

Acetic acid’s figure carries four hydrogens, three on carbon and one on oxygen; only the oxygen’s goes to water.

Sodium chloride, the salt of hydrogen chloride and sodium hydroxide, gives a neutral solution. Sodium ions have no acidic character, and chloride, conjugate to a strong acid, barely tends to take a proton.

Lewis: the electron pair that makes the transfer possible

In 1916 G. N. Lewis proposed that the electron pair is the main actor in acid–base chemistry, though the idea was not well known until about 1923. A Lewis acid accepts a pair of electrons and so forms a covalent bond with whatever supplies them. A Lewis base donates a lone pair to a species that can share it. Here, “acid” and “base” alone still carry their Brønsted–Lowry meaning.

The hydronium ion was a Lewis case all along: the proton took one of the oxygen’s lone pairs, and a shared bond formed. On Chem1’s account, it is the chance to share an electron pair that allows a proton to be transferred at all. Ammonia is a Brønsted–Lowry base and a Lewis base at once, owing to the lone pair on its nitrogen.

Every Arrhenius acid is also a Brønsted–Lowry acid. Every Brønsted–Lowry acid counts as a Lewis acid only as the source of the proton that accepts the pair: hydrogen chloride itself accepts no electron pair. With that qualification, the three definitions nest.

What the Lewis definition adds is reach beyond proton transfer. Its classic case is boron trifluoride, BF₃, a compound of boron, combining with ammonia with no proton moving at all. The product is an adduct, the name for what a Lewis acid and base form together.