A buffer is a solution that resists change in pH following the addition of an acid or base. A buffer can be created by mixing a weak acid (HA) with its conjugate base (A−). If an acid such as HCl is added to a buffer, A− can neutralize it, being converted to HA in the process. If a base is added, HA can likewise neutralize it, being converted to A− in the process. Maximum buffering capacity occurs at a pH equal to the pKa, but a conjugate acid-base pair can still serve as an effective buffer when the pH of a solution is within approximately ±1 pH unit of the pKa. If the amounts of HA and A− are equal, the pH is equal to the pKa. As shown in Figure 1.9, a solution containing acetic acid (HA = CH3 – COOH) and acetate (A− = CH3 – COO−) with a pKa of 4.8 resists a change in pH from pH 3.8 to 5.8, with maximum buffering at pH 4.8. At pH values less than the pKa, the protonated acid form (CH3 – COOH) is the predominant species in solution.
At pH values greater than the pKa, the deprotonated base form (CH3 – COO−) is the predominant species.
C. Amino acid titration
The titration curve of an amino acid can be analyzed in the same way as described for acetic acid.
1.
Carboxyl group dissociation: Consider alanine, for example, which contains an ionizable α-carboxyl and α-amino group. [Note: Its –CH3 R group is nonionizable.] At a low (acidic) pH, both of these groups are protonated (Fig. 1.10). As the pH of the solution is raised, the −COOH group of form I can dissociate by donating a H+ to the medium. The release of a H+ results in the formation of the carboxylate group, −COO−. This structure is shown as form II, which is the dipolar form of the molecule (see Fig. 1.10). This form, also called a zwitterion (from the German word for “hybrid”), is the isoelectric form of alanine, that is, it has an overall (net) charge of zero.
2.
Application of the Henderson-Hasselbalch equation: The dissociation constant of the carboxyl group of an amino acid is called K1, rather than
Ka, because the molecule contains a second titratable group. The Henderson-Hasselbalch equation can be used to analyze the dissociation of the carboxyl group of alanine in the same way as described for acetic acid: where I is the fully protonated form of alanine and II is the isoelectric form of alanine (see Fig. 1.10). This equation can be rearranged and converted to its logarithmic form to yield: 3.
Amino group dissociation: The second titratable group of alanine is the amino (−NH3+) group shown in Figure 1.10. Because this is a much weaker acid than the –COOH group, it has a much smaller dissociation constant, K2. [Note: Its pKa is, therefore, larger.] Release of a H+ from the protonated amino group of form II results in the fully deprotonated form of alanine, form III (see Fig. 1.10).
4.
Alanine pKs: The sequential dissociation of H+ from the carboxyl and amino groups of alanine is summarized in Figure 1.10. Each titratable group has a pKa that is numerically equal to the pH at which exactly one half of the H+ have been removed from that group. The pKa for the most acidic group (−COOH) is pK1, whereas the pKa for the next most acidic group (−NH3+) is pK2. [Note: The pKa of the α-carboxyl group of amino acids is ~2, whereas that of the α-amino group is ~9.] 5.