An acid (often represented by the generic formula HA) is traditionally considered any chemical compound that when dissolved in water, gives a solution with a pH of less than 7. That approximates the modern definition of Bronsted and Lowry, who 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). Acids generally taste sour; however, tasting acids, particularly concentrated acids, can be dangerous and is not recommended.

Definitions of acids and bases

The word "acid" comes from the Latin acidus meaning "sour," but in chemistry the term acid has a more specific meaning. There are three common ways to define an acid, namely, the Arrhenius, the Bronsted-Lowry and the Lewis definitions, in order of increasing generality.

- Arrhenius: According to this definition, an acid is a substance that increases the concentration of hydronium ion (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. English chemists, including Sir Humphry Davy at the same time believed all acids contained hydrogen. The Swedish chemist Svante Arrhenius used this belief to develop this definition of acid.
- Bronsted-Lowry: According to this definition, an acid is a proton 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 conjugate acid-base pairs. Bronsted and Lowry formulated this definition, which includes water-insoluble substances not in the Arrhenius definition.
- Lewis: According to this definition, 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 protons (i.e. hydrogen), such as iron(III) chloride. 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. This definition was developed by Gilbert N. Lewis.

Although not the most general theory, the Bronsted-Lowry definition is the most widely used definition. The strength of an acid may be understood by this defintion by the stability of hydronium and the solvated conjugate base upon dissociation. Increasing stability of the conjugate base will increase the acidity of a compound. This concept of acidity is used frequently for organic acids such as carboxylic acid. The molecular orbital description, where the unfilled proton orbital overlaps with a lone pair, is connected to the Lewis definition.

Solutions of weak acids and salts of their conjugate bases form buffer solutions.

Acid/base systems are different from redox reactions in that there is no change in oxidation state.

Generally, acids have the following chemical and physical properties:

- Taste: Acids generally are sour when dissolved in water.
- Touch: Acids produce a stinging feeling, particularly strong acids.
- Reactivity: Acids react aggressively with or corrode most metals.
- Electrical conductivity: Acids are electrolytes.

Strong acids are dangerous, causing severe burns for even minor contact. Generally, acid burns are treated by rinsing the affected area abundantly with water and followed up with immediate medical attention.


Acids are named according to the ending of their anion. That ionic ending is dropped and replaced with a new suffix according to the table below. For example, HCl has chloride as its anion, so the -ide suffix makes it take the form hydrochloric acid.

Anion Ending: per-anion-ate
Acid Prefix: per
Acid Suffix:
ic acid

Anion Ending: ate
Acid Prefix: -
Acid Suffix:
ic acid

Anion Ending: ite
Acid Prefix: -
Acid Suffix:
ous acid

Anion Ending: hypo-anion-ite
Acid Prefix: hypo
Acid Suffix:
ous acid

Anion Ending: ide
Acid Prefix: ide
Acid Suffix:
ic acid

Chemical characteristics

In water the following equilibrium occurs between an acid (HA) and water, which acts as a base:

HA(aq) <--> H3O+(aq) + A-(aq)

The acidity constant (or acid dissociation constant) is the equilibrium constant for the reaction of HA with water:

Strong acids have large Ka values (i.e. the reaction equilibrium lies far to the right; the acid is almost completely dissociated to H3O+ and A-). Strong acids include the heavier hydrohalic acids: hydrochloric acid (HCl), hydrobromic acid (HBr), and hydroiodic acid (HI). (However, hydrofluoric acid, HF, is relatively weak.) For example, the Ka value for hydrochloric acid (HCl) is 107.

Weak acids have small Ka values (i.e. at equilibrium significant amounts of HA and A- exist together in solution; modest levels of H3O+ are present; the acid is only partially dissociated). For example, the Ka value for acetic acid is 1.8 x 10-5. Most organic acids are weak acids. Oxoacids, which tend to contain central atoms in high oxidation states surrounded by oxygen may be quite strong or weak. Nitric acid, sulfuric acid, and perchloric acid are all strong acids, whereas nitrous acid, sulfurous acid and hypochlorous acid are all weak.

Note the following:

- The terms "hydrogen ion" and "proton" are used interchangeably; both refer to H+.
- In aqueous solution, the water is protonated to form hydronium ion, H3O+(aq). This is often abbreviated as H+(aq) even though the symbol is not chemically correct.
- The strength of an acid is measured by its acid dissociation constant (Ka) or equivalently its pKa (pKa= - log(Ka).
- The pH of a solution is a measurement of the concentration of hydronium. This will depend of the concentration and nature of acids and bases in solution.

Polyprotic acids

Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic acid (two potential protons to donate) and triprotic acid (three potential protons to donate)

A monoprotic acid can undergo one dissociation (sometimes called ionization) as follows and simply has one acid dissociation constant (Ka):

HA(aq) + H2O(l) <--> H3O+(aq) + A-(aq)

A diprotic acid (here symbolized by H2A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, Ka1 and Ka2.

H2A(aq) + H2O(l) <--> H3O+(aq) + HA-(aq)

HA-(aq) + H2O(l) <--> H3O+(aq) + A2-(aq)

The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2 . For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion (HSO4-), for which Ka1 is very large; then it can donate a second proton to form the sulfate anion (SO42-), wherein the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable carbonic acid (H2CO3) can lose one proton to form bicarbonate anion (HCO3-) and lose a second to form carbonate anion (CO32-). Both Ka values are small, but Ka1 > Ka2 .

A triprotic acid (H3A) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3 .

H3A(aq) + H2O(l) <--> H3O+(aq) + H2A-(aq)        Ka1

H2A-(aq) + H2O(l) <--> H3O+(aq) + HA2-(aq)       Ka2

HA2-(aq) + H2O(l) <--> H3O+(aq) + A3-(aq)         Ka3

An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three protons can be successively lost to yield H2PO4-, then HPO42-, and finally PO43- , the orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. Even though the positions of the protons on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.


Neutralization is the reaction between equal amounts of an acid and a base, producing a salt and water; for example, hydrochloric acid and sodium hydroxide form sodium chloride and water:

HCl(aq) + NaOH(aq) -> H2O(l) + NaCl(aq)

Neutralization is the basis of titration, where a pH indicator shows equivalence point when the equivalent number of moles of a base have been added to an acid.

Common acids

Strong inorganic acids

- Hydrobromic acid
- Hydrochloric acid
- Hydroiodic acid
- Nitric acid
- Sulfuric acid
- Perchloric acid

Medium to weak inorganic acids

- Boric acid
- Carbonic acid
- Chloric acid
- Hydrofluoric acid
- Phosphoric acid
- Pyrophosphoric acid

Weak organic acids

- Acetic acid
- Benzoic acid
- Butyric acid
- Citric acid
- Formic acid
- Lactic acid
- Malic acid
- Mandelic acid
- Methanethiol
- Propionic acid
- Pyruvic acid
- Valeric acid


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