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In chemistry and biochemistry, the acidity constant, acid-ionization constant or acid-dissocation 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 (cf. pH).
Given a weak acid HA, its dissolution into water is subject to the following equilibrium:
The acidity constant for the acid HA is the dissociation constant for this equilibrium. In other words,
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
Analogously to Ka, the magnitude of Kb indicates the relative strength of the base, with Kb > > 1 indicating a strong base.
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,
As the product of Ka and Kb remains constant, it follows that stronger acids have weaker conjugate bases, and vice versa.
Measurements are at 25ºC:
Many more are available here:  http://www.uaf.edu/chem/321Fa04/pkas.html