There are two major classes of acids: Brønsted-Lowry acids and Lewis acids.
For Brønsted acids, acidity is the tendency of a compound to act as an H+ donor.
The acidity of a Brønsted acid can be quantitatively expressed by the acid dissociation constant of the compound in water or some other specified medium. The most familiar quantitative measure of acidity is the pH scale.
Typical examples of Brønsted acids are acetic acid and sulfuric acid.
For Lewis acids, acidity relates to the compound's ability to accept an electron pair. When a Lewis acid accepts an electron pair, it forms a covalent bond with the donor of the electron pair. The electron pair donor is a Lewis base. Lewis acidity can be quantitatively expressed by the association constants of the resulting Lewis adducts and π-adducts.
Typical examples of Lewis acids are CO2 and K+.
The use of this term is mainly restricted to a medium containing Brønsted acids, where it means the tendency of the medium to add H+ to a specific reference base. It is quantitatively expressed by the appropriate acidity function.
