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acid dissociation constant

alt. (context chemistry English) The equilibrium constant for the dissociation of an acid, and thus a measure of its strength n. (context chemistry English) The equilibrium constant for the dissociation of an acid, and thus a measure of its strength

Acid dissociation constant

An acid dissociation constant, K, (also known as acidity constant, or acid-ionization constant) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid–base reactions. In aqueous solution, the equilibrium of acid dissociation can be written symbolically as:

HA +H_2O \rightleftharpoons A^- + H_3O^+

where HA is a generic acid that dissociates into A, known as the conjugate base of the acid and a hydrogen ion which combines with a water molecule to make an hydronium ion. In the example shown in the figure, HA represents acetic acid, and A represents the acetate ion, the conjugate base.

The chemical species HA, A and HO are said to be in equilibrium when their concentrations do not change with the passing of time. The dissociation constant is usually written as a quotient of the equilibrium concentrations (in mol/L), denoted by [HA], [A] and [HO]

$$K_{\mathrm a} = \mathrm{\frac{[A^-] [H_3O^+]}{[HA][H_2O]}}$$

In all but the most concentrated aqueous solutions of an acid the concentration of water can be taken as constant and can be ignored. The definition can then be written more simply

$$\mathrm{HA \rightleftharpoons A^- + H^+}:K_{\mathrm a} = \mathrm{\frac{[A^-][H^+]}{[HA]}}$$
This is the definition in common usage. For many practical purposes it is more convenient to discuss the logarithmic constant, pK

$$\ \mathrm{p}K_{\mathrm a} = - \log_{10}K_{\mathrm a}$$

The larger the value of pK, the smaller the extent of dissociation at any given pH (see Henderson–Hasselbalch equation)—that is, the weaker the acid. A weak acid has a pK value in the approximate range −2 to 12 in water. Acids with a pK value of less than about −2 are said to be strong acids; the dissociation of a strong acid is effectively complete such that concentration of the undissociated acid is too small to be measured. pK values for strong acids can, however, be estimated by theoretical means.

The definition can be extended to non-aqueous solvents, such as acetonitrile and dimethylsulfoxide. Denoting a solvent molecule by S

$$\mathrm{HA +S \rightleftharpoons A^- + SH^+}; K_{\mathrm a} = \mathrm{\frac{[A^-] [SH^+]}{[HA][S]}}$$
When the concentration of solvent molecules can be taken to be constant, $K_{\mathrm a} = \mathrm{\frac{[A^-][H^+]}{[HA]}}$, as before.