Acid

Acids are generally defined as those chemical substances which release hydrogen ions on dissolving in water. Although acids are most often thought of as proton (H⁺) or hydronium ion ([H₃O]⁺) donors, as defined by the Brønstad-Lowry theory, the Lewis acid-base theory is more comprehensive and defines acids as electron pair acceptors. All acids have a pH less than 7.

Acids play an integral role in the metabolic functions of animals, in chemistry, material science and the food industry. Acids, such as the twenty common amino acids and carboxylic acids (including valeric acid, and propionic acid), are necessary for life.

A distinction is made between strong and weak acids. Strong acids, e.g. hydrochloric acid (HCl), sulfuric acid (H₂SO₄), are completely dissociated in a solution. This means that there will be a higher concentration of H⁺ ions for a strong acid. Weak acids, e.g. ethanoic acid or acetic acid (CH₃COOH), do not dissociate completely but release only some of their protons as H⁺ ions.

Brønstad-Lowry acids

According to the theory proposed by Brønstad and Lowry, an acid is defined as a species which donates a proton (H⁺ ion) or a hydronium ([H₃O]⁺) ion. For example, hydrochloric acid (HCl) donates a H⁺ ion. Therefore it is an acid. Sodium hydroxide (NaOH), by contrast, donates a OH^- ion, and so, it is not considered an acid.
There is a drawback in this theory: acids which do not release protons or hydronium ions are not considered acids. For example, in the following reaction, boron trifluoride does not donate a proton or hydronium ion, so it is not a Bronstad-Lowry acid.

$BF_{3}\;+\;:NH_{3}\;\longrightarrow \;F_{3}B:NH_{3}$

Yet by nature BF₃ is an acid.
Consequently, this definition of acids has been discarded in favor of Lewis' definition.

Lewis acids

For a discussion of Lewis acids and Lewis bases, see Lewis acid-base theory.