User:EagleFalconn/ADC Lead

An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization constant) is the equilibrium constant for the acid-base equilibrium of an acid with its conjugate base.

Ka is defined, subject to constant temperature, as

where [HA], [A&minus;] and [H+] are equilibrium concentrations of the acid, its conjugate base, and the hydrogen ion respectively at equilibrium. The term acid dissociation constant is also used for pKa, which is equal to −log10 Ka. While the standard enthalpy change for a weak acid dissociation reaction may be positive (endothermic reaction) or negative (exothermic reaction), the standard entropy change is always negative. pKa values for endothermic reactions increase with increasing temperature; the opposite is true for exothermic reactions. This is in accord with Le Chatelier's principle.

In aqueous solutions, monoprotic acids, such as acetic acid, are partially dissociated to an appreciable extent in the pH range pKa ± 2. This constitutes the buffer region for the acid. In buffer solutions at a lower pH the acid is effectively undissociated and at higher pH it is effectively fully dissociated. Polyprotic acids, such as oxalic acid or citric acid, have a pKa value for each non-simultaneous deprotonation. The concentrations of all the species in a solution of known composition, containing one or more acids or bases, can be calculated if all the pKa values are known. Acidic behaviour can also be characterised in non-aqueous solutions. pKa can be experimentally determined by potentiometric (pH) titration, but for values of pKa less than about 2 or more than about 11 spectrophotometric or NMR measurements may be required.

Factors that determine the magnitude of pKa values include Pauling's rules for acidity constants and, for organic acids and bases, inductive effects and mesomeric effects; these effects are summarised in the Hammett equation. Structural effects, such as intra-molecular hydrogen bonding, can also be important.

The quantitative behaviour of acids and bases in solution can only be understood if their pKa values are known. For example, many compounds used for medication are weak acids or bases, so a knowledge of the pKa and log p values is essential for an understanding of the extent to which the compound enters the blood stream. There are many other applications, including aquatic chemistry, chemical oceanography, buffer solutions, acid-base homeostasis and enzyme kinetics. A knowledge of pKa values is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes in solution.