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For alternative meanings see acid (disambiguation).

An acid (often represented by the generic formula AH) is typically a water-soluble, sour-tasting chemical compound. In common usage an acid is a species that, when dissolved in water, gives a solution with a pH of less than 7. In general scientific usage an acid is a molecule or ion that is able to give up a proton (H+ ion) to a base, or accept an unshared pair of electrons from a base. An acid reacts with a base in a neutralization reaction to form a salt.


Chemical characteristics

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

\mbox{AH} +\mbox{H}_2\mbox{O} \leftrightarrow \mbox{A}^- + \mbox{H}_3\mbox{O}^+

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

K_a = {[A^-]\cdot[\mbox{H}_3\mbox{O}^+] \over [AH]}

Strong acids have large Ka values (i.e. the reaction equilibrium lies far to the right, lots of H3O+ present; the acid is almost completely dissociated). For example, the Ka value for hydrochloric acid (HCl) is 107.

Weak acids have small Ka values (i.e. at equilibrium significant amounts of AH 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.

Strong acids include the hydrohalic acids - HCl, HBr, and HI. (However, hydrofluoric acid, HF, is relatively weak.) Oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, are also quite strong and include HNO3, H2SO4, HClO4. Most organic acids are weak acids.

A few clarifications:

  • The terms "hydrogen ion" and "proton" are used interchangebly; both refer to H+.
  • In chemical equations H+ is often written, although in water it will actually be H3O+.
  • The strength of an acid is measured by its Ka value. pH measures how many hydrogen ions are present, which depends on both the type of acid (or base) and how much is there.
  • Acid strength is also defined by pKa= - log(Ka).

Number of acid dissociations

Some acid molecules are able give up more than one H+ ion (proton). Those acids which can give up only one H+ ion per molecule are called monoprotic acids, those acid molecules that can give up two H+ ions are diprotic acids, those that can give up three are triprotic acids, etc. A monoprotic acid can undergo one dissociation (sometimes called ionization) as follows and simply has one acid dissociation constant as shown above:

AH + H20 → A- + H3O+         Ka

A diprotic acid (here symbolized by AH2) can undergo one or two dissociations depending on the conditions (namely pH). Each dissociation has its own dissociation constant, Ka1 and Ka2.

AH2 + H20 → AH- + H3O+       Ka1
AH- + H20 → A-2 + H3O+        Ka2

The first dissociation constant is typically greater than the second; i. e. Ka1 > Ka2 . For example, sulfuric acid (H2SO4) can give up one H+ to form the singly charged bisulfate anion (HSO4-), for which Ka1 is very large; then it can give up a second H+ to form the doubly charged sulfate anion (SO4-2) where the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. Similarly, the weak unstable carbonic acid (H2CO3) can lose one H+ to form a singly charged bicarbonate anion (HCO3-) and lose a second to form a doubly charged carbonate anion (CO3-2). Both Ka values are small, but Ka1 > Ka2 .

Analogously, a triprotic acid (AH3) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3 .

AH3 + H20 → AH2- + H3O+        Ka1
AH2- + H20 → AH-2 + H3O+       Ka2
AH-2 + H20 → A-3 + H3O+         Ka3

An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three of its H atoms can be successively lost as H+ (or H3O+ in water) to yield H2PO4-, then HPO4-2, and finally PO4-3 , the triply charged orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three H+ ions to finally form the triply charged citrate ion. Even though the positions of the H atoms on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a positive H+ if the ion is more negatively charged.


Acids are generally:

  • Taste: sour when dissolved in water
  • Touch: strong acids have a stinging feeling
  • Reactivity: acids react aggresively with many metals
  • Electrical conductivity: acids are electrolytes

Different definitions of acid/base

The word acid comes from the Latin acidus meaning sour. In chemistry the term acid has a more specific meaning.

The Swedish chemist Svante Arrhenius defined an acid to be a substance that gives up hydrogen ions (H+) when dissolved in water, while bases are substances that give up hydroxide ions (OH-). This definition limits acids and bases to substances that can dissolve in water. Later on, Brønsted and Lowry defined an acid to be a proton donor and a base to be a proton acceptor. In this definition, even substances that are insoluble in water can be acids and bases. The most general definition of acids and bases is the Lewis definition, given by the American chemist Gilbert N. Lewis. Lewis theory defines a "Lewis acid" as an electron-pair acceptor and a "Lewis base" as an electron-pair donor. It can include acids that do not contain any hydrogen atoms, such as iron(III) chloride. Acid/base systems are different from redox reactions in that there is no change in oxidation state. 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.

The Brønsted-Lowry definition, where an acid is treated as a proton donor, is sufficient for many situations. In this case, the proton (H+) is the actual acid and the acidity of the proton-donating-compound, such as an organic acid, is determined by its stability when it donates protons to the solution it is embedded in. So if the organic acid likes letting protons go, it has high acidity because it donates protons with empty molecular orbitals to the solution. This is how organic acids such as carboxylic acids work, here the Brønsted definition is nice for calculations while the Lewis definition is good for understanding.

Acid number

This is used to quantify the amount of acid present, for example in a sample of biodiesel. It is the quantity of base, expressed in milligrams of potassium hydroxide, that is required to neutralize the acidic constituents in 1 g of sample.

AN = (Veq-beq)×N×56.1/Woil.

Veq is the amount of titrant (ml) consumed by the crude oil sample and 1ml spiking solution at the equivalent point, and beqbeq is the amount of titrant (ml) consumed by 1ml spiking solution at the equivalent point.

The molarity concentration of titrant (N) is calculated as such: N = 1000×WKHP/(204.23×Veq).

In which, WKHP is the amount (g) of KHP in 50ml of KHP standard solution, and Veq is the amount of titrant (ml) consumed by 50ml KHP standard solution at the equivalent point.

Acid number (mgKOH/g oil) for biodiesel is preferred to be lower than 3.


Neutralization is a type of reaction between an acid and a base. The products include a salt and water. So, it is also called a water forming reaction acid + base \rarr water + salt
Example: HCl + NaOH \rarr H_2O + NaCl

This type of reaction forms the basis of titration methods for analysing acids, where a pH indicator shows the point of neutralization.

Naming acids

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 takes the form hydrochloric acid.

Anion Ending Acid Prefex Acid Suffix
ate ic acid
ite ous acid
ide Hydro ic acid

Common acids

Strong inorganic acids

Medium to weak inorganic acids

Weak organic acids

Acids in food


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