To understand the acid dissociation constant, it is first important to understand the equilibrium equation for acid dissocation. Like all equilibria, an acid dissociation will have a particular equilibrium constant that will determine the extent of the reaction (whether it lies to the left or right of the equation). An example of an acid in equilibrium can be seen in .
Acetic Acid Dissociation
A ball-and-stick model of the dissociation of acetic acid to acetate. A water molecule is protonated to form a hydronium ion in the process. The acidic proton that is transferred from acetic acid to water is labelled in green.
As the equilibrium constant approaches zero, the reaction tends to form 100 percent reactants. As the equilibrium constant approaches infinity, the reaction tends to form 100 percent products. The equilibrium constant K = 1 states that there will be 50 percent products and 50 percent reactants.
Because the equilibrium constant is used for calculating the concentrations of weak acids, very little water actually reacts relative to the total concentration. The concentration of water during the reaction is, therefore, a constant, and can be excluded from the expression for K. This gives rise to a special equilibrium constant known as theacid dissociation constant, Ka. It is simply K multiplied by the concentration of water.
Ka 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. Due to the many orders of magnitude spanned by Ka values, a logarithmic measure of the acid dissociation constant is more commonly used in practice. The logarithmic constant, pKa, which is equal to −log10 (Ka), is sometimes incorrectly referred to as an acid dissociation constant as well.
Smaller Ka values yield larger pKa values. Therefore, the larger the value of pKa, the smaller the extent of dissociation. A weak acid has a pKa value in the approximate range of -2 to 12 in water. Acids with a pKa 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, which makes experimental calculations difficult. pKa 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.
Importance of Acid Dissociation Constants
A knowledge of pKa values is important for the quantitative evaluation of systems involving acid-base equilibria in solution. Many applications exist in biochemistry. For example, the pKa values of proteins and amino acid side chains are important for the activity of enzymes and the stability of proteins.
An understanding of Ka is also essential for working with buffers; the design of these solutions depends on a knowledge of the pKa values of their components. Buffers are used whenever there is a need to fix the pH of a solution at a particular value. Buffering is an essential part of in-vitro biochemical studies and acid-base physiology and plays a key role in analytical chemistry. Compared with an aqueous solution, the pH of a buffer solution is relatively insensitive to the addition of a small amount of strong acid or strong base.
Another important application of Ka is with pH indicators. A pH indicator is a weak acid or weak base that changes color in the transition pH range, which is approximately pKa ± 1. A universal indicator to test any pH would require a mixture of indicators whose adjacent pKa values differ by about two, so that their transition pH ranges just overlap.