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In chemistry and biochemistry, acid dissociation constant, the acidity constant, or the acid-ionization constant (Ka) 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 pKa (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,
- , 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 Ka 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 (Kb) and the pKb of the conjugate base A–:
This is the dissociation constant for the equilibrium
- A– + H2O ↔ HA + OH–
Analogously to Ka, the magnitude of Kb indicates the relative strength of the base, with Kb > > 1 indicating a strong base.
Relationship between acidity and basicity constants There exists a relationship between the value of Ka for an acid HA and the value of Kb 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 (Kw), which is 1.0 × 10-14 mol2.l-2 at 20°C. In other words, - KaKb = Kw
- pKa + pKb = pKw
As the product of Ka and Kb remains constant, it follows that stronger acids have weaker conjugate bases, and vice versa.
pKa of some common substances Measurements are at 25ºC: 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) | 
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