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In [[chemistry]], the '''Brønsted-Lowry theory''' is an [[Acid-base reaction|acid-base theory]], proposed independently by [[Johannes Nicolaus Brønsted]] and [[Thomas Martin Lowry]] in 1923.<ref>R.H. Petrucci, W.S. Harwood, and F.G. Herring, General Chemistry (8th edn, Prentice-Hall 2002), p.666</ref><ref>G.L. Miessler and D.A. Tarr, Inorganic Chemistry (2nd edn, Prentice-Hall 1998), p.154</ref> In this system, '''Brønsted acids''' and '''Brønsted bases''' are defined, by which an [[acid]] is a molecule or ion that is able to lose, or "donate" a [[hydrogen]] cation ([[proton]], H<sup>+</sup>), and a [[Base (chemistry)|base]] is a species with the ability to gain or "accept" a hydrogen cation (proton).
[[File:Autoprotolyse eau.svg|thumb|upright=1.5|Water as both base and acid. One H<sub>2</sub>O acts as a base and gains H<sup>+</sup> to become H<sub>3</sub>O<sup>+</sup>; the other H<sub>2</sub>O acts as an acid and loses H<sup>+</sup> to become OH<sup>-</sup>.]] 
== Properties of acids and bases ==

It follows that if a compound is to behave as an acid, donating a proton, there must be a base to accept the proton. So the Brønsted–Lowry concept can be defined by the reaction:
- acid + base {{eqm}} conjugate base + conjugate acid.
The conjugate base is the ion or molecule remaining after the acid has lost a proton, and the conjugate acid is the species created when the base accepts the proton. The reaction can proceed in either forward or backward direction; in each case the acid donates a proton to the base. 

Water is [[amphoteric]] and can act as an acid or as a base. In the reaction between [[acetic acid]], CH<sub>3</sub>CO<sub>2</sub>H, and water, H<sub>2</sub>O, water acts as a base.
:CH<sub>3</sub>CO<sub>2</sub>H + H<sub>2</sub>O {{eqm}} CH<sub>3</sub>CO<sub>2</sub><sup>-</sup> + H<sub>3</sub>O<sup>+</sup>
The [[acetate]] ion, CH<sub>3</sub>CO<sub>2</sub><sup>-</sup>, is the conjugate base of acetic acid and the [[hydronium]] ion, H<sub>3</sub>O<sup>+</sup>, is the conjugate acid of the base, water.

Water can also act as an acid, for instance when it reacts with ammonia. The equation given for this reaction is:
:H<sub>2</sub>O + NH<sub>3</sub> {{eqm}} OH<sup>-</sup> + NH<sub>4</sub><sup>+</sup>
in which H<sub>2</sub>O donates a proton to NH<sub>3</sub>. The [[hydroxide]] ion is the conjugate base of water acting as an acid and the [[ammonium]] ion is the conjugate acid of the base, ammonia.

A strong acid, such as hydrochloric acid, dissociates completely. A weak acid, such as acetic acid, may be partially dissociated; the [[acid dissociation constant]], p''K''<sub>a</sub>, is a quantitative measure of the strength of the acid. 

A wide range of compounds can be classified in the Brønsted–Lowry framework: [[mineral acid]]s and derivatives such as [[sulfonate]]s, [[phosphonate]]s, etc., [[carboxylic acid]]s, [[amine]]s, [[carbon acid]]s, 1,3-[[diketone]]s such as [[acetylacetone]], [[ethyl acetoacetate]] or [[Meldrum's acid]] and many more.

A [[Lewis base]], defined as an electron-pair donor, can act as a Brønsted–Lowry base as the pair of electrons can be donated to a proton. This means that the Brønsted–Lowry concept is not limited to aqueous solutions. Any donor solvent, S, can act as a proton acceptor.
:AH + S: {{eqm}} A<sup>-</sup> + SH<sup>+</sup>
Typical donor solvents used in acid-base chemistry, such as [[dimethyl sulfoxide]] or liquid [[ammonia]]  have an oxygen or nitrogen atom with a lone pair of electrons that can be used to form a bond with a proton.

==Brønsted acidity of some Lewis acids==
Some [[Lewis acid]]s, defined as electron-pair acceptors, also act as Brønsted–Lowry acids. For example, the aluminium ion, Al<sup>3+</sup> can accept electron pairs from water molecules, as in the reaction
:Al<sup>3+</sup> + 6H<sub>2</sub>O &rarr; Al(H<sub>2</sub>O)<sub>6</sub><sup>3+</sup>
The aqua ion formed is a weak Brønsted–Lowry acid.
:Al(H<sub>2</sub>O)<sub>6</sub><sup>3+</sup> + H<sub>2</sub>O {{eqm}} Al(H<sub>2</sub>O)<sub>5</sub>OH<sup>2+</sup> + H<sub>3</sub>O<sup>+</sup> .....[[Acid dissociation constant|K<sub>a</sub>]] = 1.7 x 10<sup>-5</sup>
The overall reaction is described as [[Hydrolysis#Hydrolysis of metal aqua ions|acid hydrolysis]] of the aluminium ion.

However not all [[Lewis acid]]s generate Brønsted–Lowry acidity. The magnesium ion similarly reacts as a Lewis acid with six water molecules
:Mg<sup>2+</sup> + 6H<sub>2</sub>O &rarr; Mg(H<sub>2</sub>O)<sub>6</sub><sup>2+</sup>
but here no protons are exchanged since the Brønsted–Lowry acidity of the aqua ion is negligible (K<sub>a</sub> ~ 10<sup>-12</sup>).

[[Boric acid]] also exemplifies the usefulness of the Brønsted–Lowry concept for an acid which does not dissociate, but does effectively donate a proton to the base, water. The reaction is
:B(OH)<sub>3</sub> + 2H<sub>2</sub>O {{eqm}} B(OH)<sub>4</sub><sup>-</sup> + H<sub>3</sub>O<sup>+</sup>
Here boric acid acts as a Lewis acid and accepts an electron pair from the oxygen of one water molecule, which in turn donates a proton to a second water molecule and therefore acts as a Brønsted acid.

==See also==
*[[Acid]]
*[[Acid-base reaction]]
*[[Base (chemistry)]]
*[[Lewis acids and bases]]

==References==
{{reflist}}

{{DEFAULTSORT:Bronsted-Lowry acid-base theory}}
[[Category:Acid-base chemistry]]
[[Category:Equilibrium chemistry]]

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