Titration of amino acids

The isoelectric point (pI) the isoelectric point, pI, is the pH of an aqueous solution of an amino acid at which the molecules have no net charge. In other words, the positively charged groups are exactly balanced by the negatively charged groups. For simple amino acids such as alanine, the pI is an average of the pKa's of the carboxyl (2.34) and ammonium (9.69) groups. Thus, the pI for alanine is calculated to be: (2.34 + 9.69)/2 = 6.02. If additional acidic or basic groups are present as side-chain functions, the pI is the average of the pKa's of the two most similar acids. The net charge on an amino acid—the algebraic sum of all the positively and negatively charged groups present—depends upon the pKa values of its functional groups and on the pH of the surrounding medium.

Titration Titration curves are produced by monitoring the pH of given volume of a sample solution after successive addition of acid or alkali The curves are usually plots of pH against the volume of titrant added or more correctly against the number of equivalents added per mole of the sample When an amino acid is dissolved in water it exists predominantly in the isoelectric form. Upon titration with acid, it acts as a base, and upon titration with base, it acts as an acid( a compound that can act as either an acid or a base is known as an amphoteric compound).

+H3N-CH2-COO- + HCl +H3N-CH2-COOH + Cl- (base) (acid) (1) +H3N-CH2-COO- + NaOH H2N-CH2-COO- + Na+ +H2O (acid) (base) (2) These ionizations follow the Henderson-Hasselbalch equation: pH= pKa+log [unprotonated form(base)] [protonated form (acid) ] When the conc of the unprotonated form equals that of the protonated form, the ratio of their concentrations equals 1, and log 1=0. Hence, pKa can be defined as the pH at which the concentrations of the protonated and unprotonated forms of a particular ionizable species are equal. .

Acid–base properties All of the amino acids have an acidic group (COOH) and a basic group (NH2) attached to the α carbon. Two of the amino acids have acidic side chains: aspartate and glutamate. Three of the amino acids have basic side chains: arginine, histidine , and lysine. All amino acids contain ionizable groups that act as weak acids or bases, giving off or taking on protons when the pH is altered.

Consider applying the Henderson-Hasselbalch equation to the titration of glycine with acid and base. Glycine has two ionizable groups: a corboxyl group and an amino group, with pKa values of 2.4 and 9.6 respectively. In water at pH 6, glycine exists as a dipolar ion, or zwitterion, in which the carboxyl group is unprotonated(-COO- ) and the amino group is protonated to give the substituted ammonium ion(-NH3+). Addition of acid to the solution lowers the pH rapidly at first and then more slowly. At pH 2.4 the pKa is reached, one-half the acid has been consumed, and the carboxyl group is half ionized and is most effective as a buffer. Titration of the amino group with base follows a similar curve into the alkaline region. The intersection between the titration of the carboxyl group and the titration of the amino group describes in this case the point at which glycine has no net charge, and is called the isoelectric point (pI).

Cont.. (pI) In the case of aspartic acid, the similar acids are the alpha-carboxyl function (pKa = 2.1) and the side-chain carboxyl function (pKa = 3.9), so pI = (2.1 + 3.9)/2 = 3.0. For arginine, the similar acids are the guanidinium species on the side-chain (pKa = 12.5) and the alpha-ammonium function (pKa = 9.0), so the calculated pI = (12.5 + 9.0)/2 = 10.75.