Editing Acid (section)

==Definitions and concepts==
{{main|Acid–base reaction}}
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 (H<sup>+</sup>) 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 base]]s due to the lone pairs of electrons on their oxygen and nitrogen atoms.

===Arrhenius acids===
[[File:Arrhenius2.jpg|thumb|150px|Svante Arrhenius]]

In 1884, [[Svante Arrhenius]] attributed the properties of acidity to hydrogen cations (H<sup>+</sup>), later described as [[Proton#Hydrogen ion|protons]] or [[Hydron (chemistry)|hydron]]s. An '''Arrhenius acid''' is a substance that, when added to water, increases the concentration of H<sup>+</sup> ions in the water.<ref name="Oxtoby8th"/><ref name="Ebbing"/> Chemists often write H<sup>+</sup>(''aq'') 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''' (H<sub>3</sub>O<sup>+</sup>) or other forms (H<sub>5</sub>O<sub>2</sub><sup>+</sup>, H<sub>9</sub>O<sub>4</sub><sup>+</sup>). 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 (chemistry)|base]], on the other hand, is a substance that increases the concentration of [[hydroxide]] (OH<sup>−</sup>) ions when dissolved in water. This decreases the concentration of hydronium because the ions react to form H<sub>2</sub>O molecules:

:H<sub>3</sub>O{{su|p=+|b=(aq)}} + OH{{su|p=−|b=(aq)}} ⇌ H<sub>2</sub>O<sub>(liq)</sub> + H<sub>2</sub>O<sub>(liq)</sub>

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<sup>−7</sup> [[Mole (unit)|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{{anchor|Brønsted acids}}===
{{Main|Brønsted–Lowry acid–base theory}}

[[File:Acetic-acid-dissociation-3D-balls.png|thumb|350px|alt=Acetic acid, CH<sub>3</sub>COOH, is composed of a methyl group, CH<sub>3</sub>, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H<sub>2</sub>0, leaving behind an acetate anion CH<sub>3</sub>COO- and creating a hydronium cation H<sub>3</sub>O. This is an equilibrium reaction, so the reverse process can also take place.|[[Acetic acid]], a [[weak acid]], donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the [[acetate]] ion and the [[hydronium]] ion. Red: oxygen, black: carbon, white: hydrogen.]]

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''' (or simply Brønsted acid) is a species that donates a proton to a Brønsted–Lowry base.<ref name="Ebbing" /> Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions of [[acetic acid]] (CH<sub>3</sub>COOH), the [[organic acid]] that gives vinegar its characteristic taste:

:{{chem2|CH3COOH + H2O <-> CH3COO- + H3O+}}
:{{chem2|CH3COOH + NH3 <-> CH3COO− + NH4+}}

Both theories easily describe the first reaction: CH<sub>3</sub>COOH acts as an Arrhenius acid because it acts as a source of H<sub>3</sub>O<sup>+</sup> when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH<sub>3</sub>COOH undergoes the same transformation, in this case donating a proton to ammonia (NH<sub>3</sub>), but does not relate to the Arrhenius definition of an acid because the reaction does not produce hydronium. Nevertheless, CH<sub>3</sub>COOH is both an Arrhenius and a Brønsted–Lowry acid.

Brønsted–Lowry theory can be used to describe reactions of [[molecule|molecular compounds]] in nonaqueous solution or the gas phase. [[Hydrogen chloride]] (HCl) and ammonia combine under several different conditions to form [[ammonium chloride]], NH<sub>4</sub>Cl. 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:
# H<sub>3</sub>O{{su|p=+|b=(aq)}} + Cl{{su|p=−|b=(aq)}} + NH<sub>3</sub> → Cl{{su|p=−|b=(aq)}} + NH{{su|b=4|p=+}}<sub>(aq)</sub> + H<sub>2</sub>O
# HCl<sub>(benzene)</sub> + NH<sub>3(benzene)</sub> → NH<sub>4</sub>Cl<sub>(s)</sub>
# HCl<sub>(g)</sub> + NH<sub>3(g)</sub> → NH<sub>4</sub>Cl<sub>(s)</sub>

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 (dissolved in [[benzene]]) react to form solid ammonium chloride in a benzene solvent and in the third gaseous HCl and NH<sub>3</sub> combine to form the solid.

===Lewis acids===
{{main|Lewis acids and bases}}
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.<ref name="Ebbing" /> 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:
:[[File:LewisAcid.png|374px]]
In the first reaction a [[fluoride|fluoride ion]], F<sup>−</sup>, gives up an [[lone pair|electron pair]] to [[boron trifluoride]] to form the product [[tetrafluoroborate]]. Fluoride "loses" a pair of [[valence electron]]s because the electrons shared in the B—F bond are located in the region of space between the two atomic [[atomic nucleus|nuclei]] and are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF<sub>3</sub> 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 H<sub>3</sub>O<sup>+</sup> 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 [[Oxidizing agent|oxidizer]] or an [[electrophile]]. Organic Brønsted acids, such as acetic, citric, or oxalic acid, are not Lewis acids.<ref name="Oxtoby8th" /> They dissociate in water to produce a Lewis acid, H<sup>+</sup>, but at the same time, they also yield an equal amount of a Lewis base (acetate, citrate, or oxalate, respectively, for the acids mentioned). This article deals mostly with Brønsted acids rather than Lewis acids.