An acid (often represented by the generic formula HA 1) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a hydrogen ion activity greater than in pure water, i.e. a pH less than 7.0. That approximates the modern definition of Johannes Nicolaus Brønsted and Martin Lowry, who independently defined an acid as a compound which donates a hydrogen ion (H⁺) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries). Acid/base systems are different from redox reactions in that there is no change in oxidation state.

Definitions

main: acid-base reaction theories The word "acid" comes from the Latin acidus meaning "sour," but in chemistry the term acid has a more specific meaning. There are four common ways to define an acid:

• Arrhenius: According to this definition developed by the Swedish chemist Svante Arrhenius, an acid is a substance that increases the concentration of hydrogen ions (H⁺), which are carried as hydronium ions (H₃O⁺) when dissolved in water, while bases are substances that increase the concentration of hydroxide ions (OH^-). This definition limits acids and bases to substances that can dissolve in water. Around 1800, many French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. Indeed the modern German word for oxygen is Sauerstoff (lit. sour substance), as is the Afrikaans word for oxygen suurstof, with the same meaning. English chemists, including Sir Humphry Davy, at the same time believed all acids contained hydrogen. Arrhenius used this belief to develop this definition of acid.
• Brønsted-Lowry: According to this definition, an acid is a proton (hydrogen nucleus) donor and a base is a proton acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as conjugate acid-base pairs. Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition. Acids according to this definition are variously referred to as Brønsted acids, Brønsted-Lowry acids, proton acids, protic acids, or protonic acids.
• Solvent-system definition: According to this definition, an acid is a substance that, when dissolved in an autodissociating solvent, increases the concentration of the solvonium cations, such as H₃O⁺ in water, NH₄⁺ in liquid ammonia, NO⁺ in liquid N₂O₄, SbCl₂⁺ in SbCl₃, etc. Base is defined as the substance that increases the concentration of the solvate anions, respectively OH^-, NH₂^-, NO₃^-, or SbCl₄^-. This definition extends acid-base reactions to non-aqueous systems and even some aprotic systems, where no hydrogen nuclei are involved in the reactions. This definition is not absolute, a compound acting as acid in one solvent may act as a base in another.
• Lewis: According to this definition developed by Gilbert N. Lewis, an acid is an electron-pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in coordination chemistry.) Lewis acids include substances with no transferable protons (ie H⁺ hydrogen ions), such as iron(III) chloride, and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition. In fact, the term Lewis acid is often used to exclude protic (Brønsted-Lowry) acids. The Lewis definition can also be explained with molecular orbital theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base. That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital.