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Concept
Acid dissociation constant
Summary
In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted K_a) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction
HA A^- + H^+
known as dissociation in the context of acid–base reactions. The chemical species HA is an acid that dissociates into , the conjugate base of the acid and a hydrogen ion, . The system is said to be in equilibrium when the concentrations of its components will not change over time, because both forward and backward reactions are occurring at the same rate.
The dissociation constant is defined by
or
where quantities in square brackets represent the concentrations of the species at equilibrium. As a simple example for a weak acid with Ka = 10−5, log Ka is the exponent which is -5, so that pKa = 5. And for acetic acid with Ka = 1.8 x 10−5, pKa is close to 5. A higher Ka corresponds to a stronger acid which is more dissociated at equilibrium. For the more convenient logarithmic scale, a lower pKa means a stronger acid.
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.
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Acid strength is the tendency of an acid, symbolised by the chemical formula HA, to dissociate into a proton, H+, and an anion, A-. The dissociation of a strong acid in solution is effectively complete, except in its most concentrated solutions. HA -> H+ + A- Examples of strong acids are hydrochloric acid (HCl), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4). A weak acid is only partially dissociated, with both the undissociated acid and its dissociation products being present, in solution, in equilibrium with each other.
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A buffer solution (more precisely, pH buffer or hydrogen ion buffer) is an acid or a base aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many living systems that use buffering for pH regulation.
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