Acids and Bases

Although acids and bases have been known since prehistoric times (vinegar, for example, is an acid), the first attempt to define what makes a compound an acid or a base was made by the Swedish chemist Svante Arrhenius (1859-1927), who proposed the definition that an acid was any compound that produced hydrogen ions, H⁺, when dissolved in water, and a base was any compound that produced hydroxide ions, OH^-, when dissolved in water. Although this was and still is a very useful definition, it has two major limitations. First, it was limited to water, or aqueous, solutions. Second, it practically limited acids and bases to ionic compounds that contained the H⁺ ion or the OH^- ion (compounds like hydrochloric acid, HCl, or sodium hydroxide, NaOH). Limited though it might be, it was an important step in the understanding of chemistry in solutions, and for his work on solution chemistry Arrhenius was awarded the 1903 Nobel Prize in chemistry.

Many common acids and bases are consistent with the Arrhenius definition. The following table shows a few common acids and bases and their uses. In all cases it is assumed that the acid or base is dissolved in water.

Many acids release only a single hydrogen ion per molecule into solution. Such acids are called monoprotic. Examples include hydrochloric acid, HCl, and nitric acid, HNO₃. Diprotic acids can release two hydrogen ions per molecule. H₂SO₄ is an example. Triprotic acids, like H₃PO₄, can release three hydrogen ions into solution. Acetic acid has the formula HC₂H₃O₂ and is a monoprotic acid because it is composed of one H⁺ion and one acetate ion, C₂H₃O₂-. The three hydrogen atoms in the acetate ion do not act as acids.