In chemistry, the Brønsted-Lowry theory is an acid-base theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. In this system, Brønsted-Lowry acids and Brønsted-Lowry bases are defined as follows: an acid is a molecule or ion that is able to lose, or "donate," a hydrogen cation (H+, a proton); a base is a species with the ability to gain, or "accept," a hydrogen cation (a proton). It follows that, if a compound is to behave as an acid by 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
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 the forward or backward direction; in each case, the acid donates a proton to the base.
Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid (CH3COOH), the organic acid that gives vinegar its characteristic taste:
$CH_3COOH + H_2O \rightleftharpoons CH_3COO^− + H_3O^+$ $CH_3COOH + NH_3 \rightleftharpoons CH_3COO^− + NH_4^+$
Both theories easily describe the first reaction: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, in this case donating a proton to ammonia (NH3), but this cannot be described using the Arrhenius definition of an acid because the reaction does not produce hydronium.
Water is amphoteric, which means it can act as an acid or as a base. In the reaction between acetic acid, CH3CO2H, and water, H2O, water acts as a base. The acetate ion, CH3CO2-, is the conjugate base of acetic acid, and the hydronium ion, H3O+, is the conjugate acid of the base, water.
Water can also act as an acid, as when it reacts with ammonia. The equation given for this reaction is:
Here, H2O donates a proton to NH3. The hydroxide ion is the conjugate base of water, which is acting as an acid, and the ammonium ion is the conjugate acid of the base, ammonia.
Brønsted-Lowry theory can also be used to describe molecular compounds, whereas Arrhenius acids must be ionic compounds. Hydrogen chloride (HCl) and ammonia combine under several different conditions to form ammonium chloride, NH4Cl. In aqueous solution, HCl behaves as hydrochloric acid and exists as hydronium and chloride ions. The following reactions illustrate the limitations of Arrhenius's definition:
$H_3O^+ (aq) + Cl^− (aq) + NH_3 \rightarrow Cl^− (aq) + NH_4^+ (aq)$ $HCl (benzene) + NH_3 (benzene) \rightarrow NH_4Cl (s)$ $HCl (g) + NH_3 (g) \rightarrow NH_4Cl (s)$
As for the acetic acid reactions, both definitions work for the first example, where water is the solvent and a hydronium ion is formed. The next two reactions do not involve the formation of ions but are still proton transfer reactions. In the second reaction, hydrogen chloride and ammonia (dissolved in benzene) react to form solid ammonium chloride in a benzene solvent, and in the third, gaseous HCl and NH3 combine to form the solid.
A wide range of compounds can be classified in the Brønsted-Lowry framework: mineral acids and derivatives such as sulfonates and phosphonates; carboxylic acids; amines; carbon acids; 1,3-diketones such as acetylacetone, ethyl acetoacetate, and Meldrum's acid and many more.
A Lewis base, defined as an electron-pair donor, can act as a Brønsted-Lowry base since 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.
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.
Some Lewis acids, defined as electron-pair acceptors, also act as Brønsted-Lowry acids. For example, the aluminium ion, Al3+, can accept electron pairs from water molecules, as in the following reaction:
The aqua ion formed is a weak Brønsted-Lowry acid.
The overall reaction is described as acid hydrolysis of the aluminium ion.
but here, very few protons are exchanged since the Brønsted-Lowry acidity of the aqua ion is negligible (Ka = 3.0 × 10-12).