Acidity constant

This article is in need of attention. Please improve it and then remove this notice and the listing on .


In chemistry and biochemistry, the acidity constant or acid dissocation constant (Ka) is a specific type of dissociation constant that indicates the extent of dissociation of hydrogen ions from an acid. An acid with an acidity constant near 1 is almost completely dissociated when dissolved in water; conversely, an acid with an acidity constant near 0 remains almost completely undissociated. In lay terms, the higher the Ka, the stronger the acid in question.

Because this number varies over many degrees of magnitude, the acidity constant is often represented by the inverse of its common logarithm, represented by pKa. (cf. pH).

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

HA + H2O ↔ H3O+ + A–

This is often written as:

HA ↔ H+ + A–

The species A– is referred to as the conjugate base of the acid. The acidity constant for the acid HA is the dissociation constant for this equilibrium. Thus:

K_a = \frac{[\mbox{H}_3\mbox{O}^+][\mbox{A}^- ]} {[\mbox{HA}]}

where [X] denotes the molar concentration of X in the solution.

By analogy, one can define the basicity constant (Kb, and similarly pKb) of the conjugate base A–:

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

For the equilibrium:

A– + H2O ↔ HA + OH–

Analogously to Ka, the magnitude of Kb indicates the relative strength of the base, with Kb closer to 1 indicating a much stronger base.

The relation between Ka and Kb is:

Kw = KaKb
pKw = pKa + pKb

where Kw is the dissociation constant of water, which is 1.0x10-14 mol2 dm-6at 20 °C.

As the product of Ka and Kb must remain a constant, it follows that stronger acids will have weaker conjugate bases, and vice versa.


This article is licensed under the GNU Free Documentation License. It uses material from Wikipedia article. Browse Wikipedia for more information.