Acid

An acid is a molecule or ion capable of either donating a proton, known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.
The first category of acids are the proton donors, or Brønsted–Lowry acids. In the special case of aqueous solutions, proton donors form the hydronium ion H₃O⁺ and are known as Arrhenius acids. Brønsted and Lowry generalized the Arrhenius theory to include non-aqueous solvents. A Brønsted–Lowry or Arrhenius acid usually contains a hydrogen atom bonded to a chemical structure that is still energetically favorable after loss of H⁺.
Aqueous Arrhenius acids have characteristic properties that provide a practical description of an acid. Acids form aqueous solutions with a sour taste, can turn blue litmus red, and react with bases and certain metals to form salts. The word acid is derived from the Latin acidus, meaning 'sour'. An aqueous solution of an acid has a pH less than 7 and is colloquially also referred to as "acid", while the strict definition refers only to the solute. A lower pH means a higher acidity, and thus a higher concentration of hydrogen cations in the solution. Chemicals or substances having the property of an acid are said to be acidic.
Common aqueous acids include hydrochloric acid, acetic acid, sulfuric acid, and citric acid. As these examples show, acids can be solutions or pure substances, and can be derived from acids that are solids, liquids, or gases. Strong acids and some concentrated weak acids are corrosive, but there are exceptions such as carboranes and boric acid.
The second category of acids are Lewis acids, which form a covalent bond with an electron pair. An example is boron trifluoride, whose boron atom has a vacant orbital that can form a covalent bond by sharing a lone pair of electrons on an atom in a base, for example the nitrogen atom in ammonia. Lewis considered this as a generalization of the Brønsted definition, so that an acid is a chemical species that accepts electron pairs either directly or by releasing protons into the solution, which then accept electron pairs. Hydrogen chloride, acetic acid, and most other Brønsted–Lowry acids cannot form a covalent bond with an electron pair, however, and are therefore not Lewis acids. Conversely, many Lewis acids are not Arrhenius or Brønsted–Lowry acids. In modern terminology, an acid is implicitly a Brønsted acid and not a Lewis acid, since chemists almost always refer to a Lewis acid explicitly as such.

Definitions and concepts

Modern definitions are concerned with the fundamental chemical reactions common to all acids.
Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, so the Arrhenius and Brønsted–Lowry definitions are the most relevant.
The Brønsted–Lowry definition is the most widely used definition; unless otherwise specified, acid–base reactions are assumed to involve the transfer of a proton from an acid to a base.
Hydronium ions are acids according to all three definitions. Although alcohols and amines can be Brønsted–Lowry acids, they can also function as Lewis bases due to the lone pairs of electrons on their oxygen and nitrogen atoms.

Arrhenius acids

In 1884, Svante Arrhenius attributed the properties of acidity to hydrogen cations, later described as protons or hydrons. An Arrhenius acid is a substance that, when added to water, increases the concentration of H⁺ ions in the water. Chemists often write H⁺ and refer to the hydrogen cation when describing acid–base reactions but the free hydrogen nucleus, a proton, does not exist alone in water, it exists as the hydronium ion or other forms. Thus, an Arrhenius acid can also be described as a substance that increases the concentration of hydronium ions when added to water. Examples include molecular substances such as hydrogen chloride and acetic acid.
An Arrhenius base, on the other hand, is a substance that increases the concentration of hydroxide ions when dissolved in water. This decreases the concentration of hydronium because the ions react to form H₂O molecules:
Due to this equilibrium, any increase in the concentration of hydronium is accompanied by a decrease in the concentration of hydroxide. Thus, an Arrhenius acid could also be said to be one that decreases hydroxide concentration, while an Arrhenius base increases it.
In an acidic solution, the concentration of hydronium ions is greater than 10⁻⁷ moles per liter. Since pH is defined as the negative logarithm of the concentration of hydronium ions, acidic solutions thus have a pH of less than 7.

Brønsted–Lowry acids

While the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. In 1923, chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid–base reactions involve the transfer of a proton. A Brønsted–Lowry acid is a species that donates a proton to a Brønsted–Lowry base. Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid, the organic acid that gives vinegar its characteristic taste:
Both theories easily describe the first reaction: CH₃COOH acts as an Arrhenius acid because it acts as a source of H₃O⁺ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH₃COOH undergoes the same transformation, in this case donating a proton to ammonia, but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH₃COOH is both an Arrhenius and a Brønsted–Lowry acid.
Brønsted–Lowry theory can be used to describe reactions of molecular compounds in nonaqueous solution or the gas phase. Hydrogen chloride and ammonia combine under several different conditions to form ammonium chloride, NH₄Cl. In aqueous solution HCl behaves as hydrochloric acid and exists as hydronium and chloride ions. The following reactions illustrate the limitations of Arrhenius's definition:

H₃O + Cl + NH₃ → Cl + NH + H₂O
HCl + NH₃ → NH₄Cl
HCl + NH₃ → NH₄Cl

As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and hydronium ion is formed by the HCl solute. The next two reactions do not involve the formation of ions but are still proton-transfer reactions. In the second reaction hydrogen chloride and ammonia react to form solid ammonium chloride in a benzene solvent and in the third gaseous HCl and NH₃ combine to form the solid.

Lewis acids

A third, only marginally related concept was proposed in 1923 by Gilbert N. Lewis, which includes reactions with acid–base characteristics that do not involve a proton transfer. A Lewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor. Brønsted acid–base reactions are proton transfer reactions while Lewis acid–base reactions are electron pair transfers. Many Lewis acids are not Brønsted–Lowry acids. Contrast how the following reactions are described in terms of acid–base chemistry:
In the first reaction a fluoride ion, F⁻, gives up an electron pair to boron trifluoride to form the product tetrafluoroborate. Fluoride "loses" a pair of valence electrons because the electrons shared in the B—F bond are located in the region of space between the two atomic nuclei and are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF₃ is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Brønsted theory because there is no proton transfer.
The second reaction can be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted base; alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form a bond with a hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H₃O⁺ gains a pair of electrons when one of the H—O bonds is broken and the electrons shared in the bond become localized on oxygen.
Depending on the context, a Lewis acid may also be described as an oxidizer or an electrophile. Organic Brønsted acids, such as acetic, citric, or oxalic acid, are not Lewis acids. They dissociate in water to produce a Lewis acid, H⁺, but at the same time, they also yield an equal amount of a Lewis base. This article deals mostly with Brønsted acids rather than Lewis acids.

Dissociation and equilibrium

Reactions of acids are often generalized in the form, where HA represents the acid and A⁻ is the conjugate base. This reaction is referred to as protolysis. The protonated form of an acid is also sometimes referred to as the free acid.
Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton. The acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written as. In solution there exists an equilibrium between the acid and its conjugate base. The equilibrium constant K is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that means the concentration of H₂O. The acid dissociation constant K_a is generally used in the context of acid–base reactions. The numerical value of K_a is equal to the product of the concentrations of the products divided by the concentration of the reactants, where the reactant is the acid and the products are the conjugate base and H⁺.
The stronger of two acids will have a higher K_a than the weaker acid; the ratio of hydrogen cations to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for K_a spans many orders of magnitude, a more manageable constant, pK_a is more frequently used, where pK_a = −log₁₀ K_a. Stronger acids have a smaller pK_a than weaker acids. Experimentally determined pK_a at 25 °C in aqueous solution are often quoted in textbooks and reference material.