Brønsted–Lowry acid–base theory

In chemistry, the Brønsted–Lowry theory is an acid-base theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923.^[1][2] 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⁺), and a base is a species with the ability to gain, or "accept," a hydrogen cation (proton).

Water as both base and acid. One H₂O acts as a base and gains H⁺ to become H₃O⁺; the other H₂O acts as an acid and loses H⁺ to become OH^-.

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 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.

CH₃COOH + H₂O CH₃COO⁻ + 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 OH⁻ + NH₄⁺

in which H₂O donates a proton to NH₃. 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, 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, and 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: A⁻ + SH⁺

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 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(H₂O)₆³⁺ + H₂O Al(H₂O)₅OH²⁺ + H₃O⁺ .....K_a = 1.2 × 10^{−5 [3]}

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 very few protons are exchanged since the Brønsted–Lowry acidity of the aqua ion is negligible (K_a = 3.0 × 10^-12).^[3]

Boric acid also exemplifies the usefulness of the Brønsted–Lowry concept for an acid that does not dissociate but does effectively donate a proton to the base, water. The reaction is

B(OH)₃ + 2H₂O 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.

See also

• Acid
• Acid-base reaction
• Base (chemistry)
• Lewis acids and bases

References

1. ^ R.H. Petrucci, W.S. Harwood, and F.G. Herring, General Chemistry (8th edn, Prentice-Hall 2002), p.666
2. ^ G.L. Miessler and D.A. Tarr, Inorganic Chemistry (2nd edn, Prentice-Hall 1998), p.154
3. ^ ^{a b} K.W. Whitten, K.D. Gailey and R.E. Davis, "General Chemistry" (4th edn., Saunders College Publishing 1992) p.750