Acid-base Equilibria
Brønsted and Lowry characterized an acid-base equilibrium as involving a proton exchange reaction:
- acid + base conjugate base + conjugate acid.
An acid is a proton donor; the proton is transferred to the base, a proton acceptor, creating a conjugate acid. For aqueous solutions of an acid HA, the base is water; the conjugate base is A− and the conjugate acid is the solvated hydrogen ion. In solution chemistry, it is usual to use H+ as an abbreviation for the solvated hydrogen ion, regardless of the solvent. In aqueous solution H+ denotes a solvated hydronium ion.
The Brønsted–Lowry definition applies to other solvents, such as dimethyl sulfoxide: the solvent S acts as a base, accepting a proton and forming the conjugate acid SH+. A broader definition of acid dissociation includes hydrolysis, in which protons are produced by the splitting of water molecules. For example, boric acid, B(OH)3, acts as a weak acid, even though it is not a proton donor, because of the hydrolysis equilibrium
- B(OH)3 + H2O B(OH)4− + H+.
Similarly, metal ion hydrolysis causes ions such as 3+ to behave as weak acids:
- 3+ 2+ + H+.
Acid-base equilibria are important in a very wide range of applications, such as acid-base homeostasis, ocean acidification, pharmacology and analytical chemistry.
Read more about this topic: Equilibrium Chemistry