Acid dissociation constant
In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction
"Acid–base equilibrium" redirects here. For acid–base balance in physiology, see Acid–base homeostasis.
known as dissociation in the context of acid–base reactions. The chemical species HA is an acid that dissociates into A−, called the conjugate base of the acid, and a hydrogen ion, H+.[a] The system is said to be in equilibrium when the concentrations of its components do not change over time, because both forward and backward reactions are occurring at the same rate.[1]
The dissociation constant is defined by[b]
where quantities in square brackets represent the molar concentrations of the species at equilibrium.[c][2] For example, a hypothetical weak acid having Ka = 10−5, the value of log Ka is the exponent (−5), giving pKa = 5. For acetic acid, Ka = 1.8 x 10−5, so pKa is about 5. A higher Ka corresponds to a stronger acid (an acid that is more dissociated at equilibrium). The form pKa is often used because it provides a convenient logarithmic scale, where a lower pKa corresponds to a stronger acid.
Theoretical background[edit]
The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pKa value is directly proportional to the standard Gibbs free energy change for the reaction. The value of the pKa changes with temperature and can be understood qualitatively based on Le Châtelier's principle: when the reaction is endothermic, Ka increases and pKa decreases with increasing temperature; the opposite is true for exothermic reactions.
The value of pKa also depends on molecular structure of the acid in many ways. For example, Pauling proposed two rules: one for successive pKa of polyprotic acids (see Polyprotic acids below), and one to estimate the pKa of oxyacids based on the number of =O and −OH groups (see Factors that affect pKa values below). Other structural factors that influence the magnitude of the acid dissociation constant include inductive effects, mesomeric effects, and hydrogen bonding. Hammett type equations have frequently been applied to the estimation of pKa.[3][4]
The quantitative behaviour of acids and bases in solution can be understood only if their pKa values are known. In particular, the pH of a solution can be predicted when the analytical concentration and pKa values of all acids and bases are known; conversely, it is possible to calculate the equilibrium concentration of the acids and bases in solution when the pH is known. These calculations find application in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the pKa values, together with the octanol-water partition coefficient, can be used for estimating the extent to which the compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a fundamental role. In living organisms, acid–base homeostasis and enzyme kinetics are dependent on the pKa values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of pKa values is necessary for the preparation of buffer solutions and is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes. Experimentally, pKa values can be 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 due to practical difficulties with pH measurements.
Strong acids and bases[edit]
An acid is classified as "strong" when the concentration of its undissociated species is too low to be measured.[6] Any aqueous acid with a pKa value of less than 0 is almost completely deprotonated and is considered a strong acid.[20] All such acids transfer their protons to water and form the solvent cation species (H3O+ in aqueous solution) so that they all have essentially the same acidity, a phenomenon known as solvent leveling.[21][22] They are said to be fully dissociated in aqueous solution because the amount of undissociated acid, in equilibrium with the dissociation products, is below the detection limit. Likewise, any aqueous base with an association constant pKb less than about 0, corresponding to pKa greater than about 14, is leveled to OH− and is considered a strong base.[22]
Nitric acid, with a pK value of around −1.7, behaves as a strong acid in aqueous solutions with a pH greater than 1.[23] At lower pH values it behaves as a weak acid.
pKa values for strong acids have been estimated by theoretical means.[24] For example, the pKa value of aqueous HCl has been estimated as −9.3.
After rearranging the expression defining Ka, and putting pH = −log10[H+], one obtains[25]
This is the Henderson–Hasselbalch equation, from which the following conclusions can be drawn.
In water, measurable pKa values range from about −2 for a strong acid to about 12 for a very weak acid (or strong base).
A buffer solution of a desired pH can be prepared as a mixture of a weak acid and its conjugate base. In practice, the mixture can be created by dissolving the acid in water, and adding the requisite amount of strong acid or base. When the pKa and analytical concentration of the acid are known, the extent of dissociation and pH of a solution of a monoprotic acid can be easily calculated using an ICE table.