Brønsted-Lowry theory

Brønsted-Lowry theory
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In chemistry, the Brønsted–Lowry acid-base theory is an acid-base theory. It was proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923. In this system, an acid is defined as any chemical species (molecule or ion) that is able to lose, or "donate" a hydrogen ion (proton), and a base is a species with the ability to gain or "accept" a hydrogen ion (proton). 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 $\rightleftarrows$ 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, CH₃CO₂H, and water, H₂O, water acts as a base.

CH₃CO₂H + H₂O $\rightleftharpoons$ CH₃CO₂^- + H₃O⁺

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, for instance when it reacts with ammonia. The equation given for this reaction is:

H₂O + NH₃ $\rightleftharpoons$ OH^- + NH₄⁺

in which H₂O donates a proton to NH₃. The hydroxide ion is the conjugate base of water acting as an acid.

A strong acid, such as hydrochloric acid, dissociates completely. A weak acid, such as acetic acid, may be partially dissociated; the acid dissociation constant, pK_a, is a quantitative measure of the strength of the acid.

A wide range of compounds can be classified in the Brønsted–Lowry framework: mineral acids and derivatives such as sulfonates, phosphonates, etc., carboxylic acids, amines, carbon acids, 1,3-diketones 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: $\rightleftarrows$ A^- + SH⁺

Typical donor solvents used in acid-base chemistry, such as dimethyl sulphoxide or liquid ammonia have an oxygen or nitrogen atom with a lone pair of electrons that can used to form a bond with a proton.

Brønsted acidity of some Lewis acids

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 reaction

Al³⁺ + 6H₂O → Al(H₂O)₆³⁺

The aqua ion formed is a weak Brønsted–Lowry acid.

Al(H₃O)₆²⁺ + H₂O $\rightleftharpoons$ Al(H₂O)₅OH²⁺ + H₃O⁺ .....K_a = 1.7 x 10^-5

The overall reaction is described as acid hydrolysis of the aluminium ion.

However not all Lewis acids generate Brønsted–Lowry acidity. The magnesium ion similarly reacts as a Lewis acid with six water molecules

Mg²⁺ + 6H₂O → Mg(H₂O)₆²⁺

but here no protons are exchanged since the Brønsted–Lowry acidity of the aqua ion is negligible (K_a ~ 10^-12).

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)₃ + 2H₂O $\rightleftharpoons$ B(OH)₄^- + H₃O⁺

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.

Brønsted acidity of some Lewis acids

See also