An acid dissociation constant, Ka, (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. The equilibrium can be written symbolically as:
- HA A− + H+,
where HA is a generic acid that dissociates by splitting into A−, known as the conjugate base of the acid, and the hydrogen ion or proton, H+, which, in the case of aqueous solutions, exists as a solvated hydronium ion. In the example shown in the figure, HA represents acetic acid, and A− the acetate ion. The chemical species HA, A− and H+ 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, and :
Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice. The logarithmic constant, pKa, which is equal to −log10 Ka, is sometimes also (but incorrectly) referred to as an acid dissociation constant:
The larger the value of pKa, the smaller the extent of dissociation. A weak acid has a pKa value in the approximate range −2 to 12 in water. Acids with a pKa value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable. pKa values for strong acids can, however, be estimated by theoretical means or by extrapolating from measurements in non-aqueous solvents in which the dissociation constant is smaller, such as acetonitrile and dimethylsulfoxide.
Read more about Acid Dissociation Constant: Theoretical Background, Definitions, Equilibrium Constant, Acidity in Nonaqueous Solutions, Factors That Affect PKa Values, Experimental Determination, Applications and Significance, Values For Common Substances
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