In chemistry and biochemistry, acid dissociation constant, the acidity constant, or the acid-ionization constant (<math>K_a<math>) is a specific type of equilibrium constant that indicates the extent of dissociation of hydrogen ions from an acid. While strong acids dissociate more or less completely in solution and consequently have large acidity constants, weak acids do not fully dissociate and generally have acidity constants significantly less than 1. Because this constant differs for each acid and varies over many degrees of magnitude, the acidity constant is often represented by the inverse of its common logarithm, represented by the symbol <math>pK_a<math> (similar to the concept of pH, though not related directly).

Given a weak acid HA, its dissolution into water is subject to the following Chemical_equilibrium:

HA + H₂O ↔ H₃O⁺ + A^– (it is also acceptable to write this as: HA ↔ H⁺ + A^–, the difference being only what theory of acids/bases you are applying. See the Bronsted-Lowry Theory and the Arrhenius Theory for more information)

The acidity constant for the acid HA is the dissociation constant for this equilibrium. In other words,

<math>K_a = \frac{[\mbox{H}_3\mbox{O}^+][\mbox{A}^- ]} {[\mbox{HA}]}<math>, where [X] denotes the molar concentration of X in the solution

Using this definition, chemists can quickly and easily determine the concentrations of various chemicals in an equilibrium. For example, to determine the pH of a solution of sodium hydroxide and hydrofluoric acid, if you know the K_a of the acid at a given temperature (which is easily attainable information) you can determine the concentration of hydrogen ions, which will allow the determination of the pH after taking into account the neutralization due to the base.

Basicity constant of the conjugate base

By analogy, one can define the basicity constant (<math>K_b<math>) and the <math>pK_b<math> of the conjugate base A^–:

<math>K_b = \frac{[\mbox{HA}][\mbox{OH}^-]} {[\mbox{A}^-]}<math>

This is the dissociation constant for the equilibrium

A^– + H₂O ↔ HA + OH^–

Analogously to <math>K_a<math>, the magnitude of <math>K_b<math> indicates the relative strength of the base, with <math>K_b >> 1<math> indicating a strong base.

Relationship between acidity and basicity constants

There exists a relationship between the value of <math>K_a<math> for an acid HA and the value of <math>K_b<math> for its conjugate base A^–. Since adding the ionization reaction for HA and the ionization reaction of A^– always gives the reaction for the self-ionization of water, the product of the acidity and basicity constants gives the dissociation constant of water (<math>K_w<math>), which is 1.0 × 10^-14 mol².l^-2 at 20°C. In other words,

<math>K_aK_b = K_w<math>

<math>pK_a + pK_b = pK_w<math>

As the product of K_a and K_b remains constant, it follows that stronger acids have weaker conjugate bases, and vice versa.

pKa of some common substances

Measurements are at 25ºC:

• 3.75: Formic acid
• 4.19: Succinic acid
• 4.20: Benzoic acid
• 4.63: Aniline
• 4.75: Acetic
• 4.76: Citrate
• 5.21: Pyridine
• 6.37: Carbonic acid
• 6.40: Citrate
• 6.99: Ethylenediamine
• 7.00: Imidazole
• 9.14: Borate
• 9.25: Ammonia
• 9.33: Benzylamine
• 9.81: Trimethylamine
• 9.99: Phenol
• 10.08: Ethylenediamine
• 10.66: Methylamine
• 10.73: Dimethylamine
• 10.81: Ethylamine
• 11.01: Triethylamine
• 11.09: Diethylamine
• 12.67: Phosphate
• 12.74: Borate

Many more are available here: [1] (http://www.uaf.edu/chem/321Fa04/pkas.html) and [2] (http://daecr1.harvard.edu/pdf/evans_pKa_table.pdf)