Introduced by Svante Arrhenius in 1884, the Arrhenius definition of a base is a cornerstone of classical acid-base chemistry. Under this theory, a base is any chemical compound that dissociates or ionizes in water to release hydroxide ions into the solution. This influx of hydroxide ions raises the overall pH of the mixture above 7, giving the solution its characteristic basic, or alkaline, properties.
The most common examples of Arrhenius bases are ionic compounds containing alkali metals or alkaline earth metals bonded to hydroxide groups. Sodium hydroxide and potassium hydroxide are classic examples. When these solid salts are placed in water, the polar water molecules surround and pull the ionic bonds apart.
The metal cations and the hydroxide anions separate and disperse evenly throughout the liquid, immediately and dramatically increasing the overall hydroxide concentration in the solvent.
The hallmark of an Arrhenius acid-base reaction is neutralization. When an Arrhenius acid, which produces hydrogen ions, is mixed with an Arrhenius base, which produces hydroxide ions, the two ions rapidly combine. This combination forms pure liquid water, while the remaining metal and nonmetal ions form an ionic salt. This predictable formation of water and salt is a key feature of Arrhenius chemistry.
Like the Arrhenius acid definition, the Arrhenius base definition is fundamentally limited because it requires water as a solvent. It also struggles to explain why compounds that do not contain a hydroxide group, such as ammonia, act as strong bases in water. To explain how ammonia makes a solution basic without directly releasing its own hydroxide ions, chemists rely on the more comprehensive Bronsted-Lowry theory.