Question

With relevance to Organic Chemistry and acids and bases, can somebody teach me everything there is to know about pka?

Answer 1

pK a

/pK a/ the negative logarithm of the ionization constant (K) of an acid, the pH of a solution in which half of the acid molecules are ionized.

PKA,

abbreviation for protein kinase.

p,

n a mathematical function analogous to the calculation of pH. It is the negative log (p) of the constant of acid dissociation (
K_a). When the p
K_a of a buffering agent equals the pH of the solution to be buffered, the buffering system is most effective. According to the Henderson Hasselbach equation, p
K_a = pH + log ([HOAc]/[OAc-]).

pK_a

The negative decadic logarithm of the ionization constant (K_a) of an acid; equal to the pH value at which equal concentrations of the acid and conjugate base forms of a substance (often a buffer) are present.

pK_a

The negative decadic logarithm of the ionization constant (K_a) of an acid; equal to the pH value at which equal concentrations of the acid and conjugate base forms of a substance (often a buffer) are present.

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. The logarithmic constant, pK_a, which is equal to ?log₁₀K_a, is sometimes also (but incorrectly) referred to as an acid dissociation constant:

The larger the value of pK_a, the smaller the extent of dissociation at any given pH that is, the weaker the acid. 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 value of pK_a also depends on molecular structure in many ways. For example, Pauling proposed two rules: one for successive pK_a of polyprotic acids

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 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