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An acid dissociation constant, ''K''a, (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: : 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 H3O+ 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 () : 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 : This is the definition in common usage. For many practical purposes it is more convenient to discuss the logarithmic constant, p''K''a :〔p''K''a, is sometimes referred to as an acid dissociation constant. Strictly speaking this is incorrect: it refers to the logarithm of a stability constant.〕 The larger the value of p''K''a, 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 p''K''a value in the approximate range −2 to 12 in water. Acids with a p''K''a 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. p''K''a 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 : When the concentration of solvent molecules can be taken to be constant, , as before. == Theoretical background == The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the p''K''a value is directly proportional to the standard Gibbs energy change for the reaction. The value of the p''K''a changes with temperature and can be understood qualitatively based on Le Chatelier's principle: when the reaction is endothermic, the p''K''a decreases with increasing temperature; the opposite is true for exothermic reactions. The value of p''K''a also depends on molecular structure in many ways. For example, Pauling proposed two rules: one for successive p''K''a of polyprotic acids (see Polyprotic acids below), and one to estimate the p''K''a of oxyacids based on the number of =O and −OH groups (see Factors that affect p''K''a values below). Other structural factors that influence the magnitude of the acid dissociation constant include inductive effects, mesomeric effects, and hydrogen bonding. The quantitative behaviour of acids and bases in solution can be understood only if their p''K''a values are known. In particular, the pH of a solution can be predicted when the analytical concentration and p''K''a 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 p''K''a values, together with the water–octanol 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 p''K''a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of p''K''a 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, p''K''a values can be determined by potentiometric (pH) titration, but for values of p''K''a less than about 2 or more than about 11, spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements. 抄文引用元・出典: フリー百科事典『 ウィキペディア（Wikipedia）』 ■ウィキペディアで「Acid dissociation constant」の詳細全文を読む スポンサード リンク
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