Brønsted–Lowry acid–base theory

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-.

Contents

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 is in equilibrium with 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 is in equilibrium with 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 is in equilibrium with 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, pKa, 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: is in equilibrium with 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 is in equilibrium with Al(H2O)5OH2+ + H3O+ .....Ka = 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 (Ka = 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 is in equilibrium with 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

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

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