Image:Acetic-acid-dissociation-3D-balls.png|thumb|305px|alt=Acetic acid, CH₃COOH, is composed of a methyl group, CH₃, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H₂0, leaving behind an acetate anion CH₃COO- and creating a hydronium cation H₃O . This is an equilibrium reaction, so the reverse process can also take place.|Acetic acid, a weak acid, donates a proton (hydrogen ion, green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen. 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. The equilibrium can be written symbolically as: 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 HA, A⁻ and H⁺: Due to the many orders of magnitude spanned by K_a values, a logarithmic measure of the acid dissociation constant is more commonly used in practice. pK_a, which is equal to −log₁₀ K_a, may also be referred to as an acid dissociation constant: The larger the value of pK_a, the smaller the extent of dissociation. A weak acid has a pK_a value in the approximate range −2 to 12 in water. Acids with a pK_a 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. pK_a 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. The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pK_a value is directly proportional to the standard Gibbs energy change for the reaction. The value of the pK_a changes with temperature and can be understood qualitatively based on Le Chatelier's principle: when the reaction is endothermic, the pK_a decreases with increasing temperature; the opposite is true for exothermic reactions. The underlying structural factors that influence the magnitude of the acid dissociation constant include Pauling's rules for acidity constants, inductive effects, mesomeric effects, and hydrogen bonding. The quantitative behaviour of acids and bases in solution can be understood only if their pK_a values are known. In particular, the pH of a solution can be predicted when the analytical concentration and pK_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 pK_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 pK_a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of pK_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, pK_a values can be determined by potentiometric (pH) titration, but for values of pK_a less than about 2 or more than about 11 spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements.

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