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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 practically completely in solution and consequently have large acidity constants, weak acids do not fully dissociate and generally have acidity constants far 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 additive inverse of its common logarithm, represented by the symbol pKa (similar to the concept of pH, though not related directly).
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- pKa = −log10Ka
Given a weak acid HA, its dissociation in 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,
![K_a = \frac{[\mbox{H}_3\mbox{O}^+][\mbox{A}^- ]} {[\mbox{HA}]}](/math/4d89813dee5901e2f4a863fdbdb99053.png) , 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 known molar concentrations of sodium hydroxide and hydrofluoric acid, if the Ka of the acid at the given temperature (which is easily attainable information) is known, the concentration of hydrogen ions can be calculated, which will allow the determination of the pH after taking into account the neutralization due to the base. This procedure is summed up in the Henderson-Hasselbalch equation.
Basicity constant of the conjugate base
By analogy, one can define the basicity constant (Kb) and the pKb of the conjugate base A–:
![K_b = \frac{[\mbox{HA}][\mbox{OH}^-]} {[\mbox{A}^-]}](/math/b82014abb8acb2fdcb171f98420bb2c6.png)
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- pKb = −log10Kb
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 M2 at 25°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 in water:
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