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).
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 15px 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 15px 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 15px 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 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: 15px 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.
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 (Ka ~ 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 15px 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
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