MCAT Biochemistry Review

Chapter 1: Amino Acids, Peptides, and Proteins

1.2 Acid–Base Chemistry of Amino Acids

The AAMC loves to use amino acids to test your understanding of acid–base chemistry because they have both an acidic carboxylic acid group and a basic amino group. That makes them amphoteric species, as they can either accept a proton or donate a proton; how they react depends on the pH of their environment. The key to understanding the behavior of amino acids is to remember two facts:

·        Ionizable groups tend to gain protons under acidic conditions and lose them under basic conditions. So, in general, at low pH, ionizable groups tend to be protonated; at high pH, they tend to be deprotonated.

·        The pKa of a group is the pH at which, on average, half of the molecules of that species are deprotonated; that is, [protonated version of the ionizable group] = [deprotonated version of the ionizable group] or [HA] = [A]. If the pH is less than the pKa, a majority of the species will be protonated. If the pH is higher than the pKa, a majority of the species will be deprotonated.


You are not expected to know exact values for isoelectric points or side chain pKa values. The MCAT will either provide you with numerical data to use in answering a question, or test clear-cut distinctions that can be made without additional information (such as, Which is more hydrophilic, valine or lysine? but not Which is more hydrophobic, valine or alanine?).


Because all amino acids have at least two groups that can be deprotonated, they all have at least two pKa values. The first one, pKa1, is the pKa for the carboxyl group, and is usually around 2. For most amino acids, the second pKa value, pKa2, is the pKa for the amino group, which is usually between 9 and 10. For amino acids with an ionizable side chain, there will be three pKa values, but we'll come back to that later. As an example, let's take glycine, which doesn't have an ionizable side chain.


Like all polyprotic acids, each ionizable proton has its own pKa at which it is “half-deprotonated.”

Positively Charged Under Acidic Conditions

At pH 1 (below even the pH of the stomach), there are plenty of protons in solution. Because we're far below the pKa of the amino group, the amino group will be fully protonated (–NH3+), and thus positively charged. Because we're also below the pKa of the carboxylic acid group, it too will be fully protonated (–COOH), and thus neutral. Therefore, at very acidic pH values, amino acids tend to be positively charged, as shown in Figure 1.5.

Figure 1.5. Amino Acid Structure at Acidic pH

Zwitterions at Intermediate pH

If we increase the pH of the amino acid solution from pH 1 to pH 7.4, the normal pH of human blood, we've moved far above the pKa of the carboxylic acid group. At physiological pH, you will not find amino acids with the carboxylate group protonated (–COOH) and the amino group unprotonated (–NH2). Under these conditions, the carboxyl group will be in its conjugate base form and be deprotonated, becoming –COO. Conversely, we're still well below the pKa of the basic amino group, so it will remain fully protonated and in its conjugate acid form (–NH3+). Thus we have a molecule that has both a positive charge and a negative charge, but overall, the molecule is electrically neutral. We call such molecules dipolar ions, or zwitterions (from the German zwitter, or “hybrid”), as depicted in Figure 1.6. The two charges neutralize one another, and zwitterions exist in water as internal salts.

Figure 1.6. Carboxylic Acids Become Deprotonated at Neutral pH, Forming a Zwitterion

Negatively Charged Under Basic Conditions

Milk of magnesia, which is often used as an antacid, has a pH around 10.5. At that pH, the carboxylate group is already deprotonated and thus remains –COO. On the other hand, we are now well above the pKa for the amino group, so it deprotonates too, becoming –NH2. So, at highly basic pH, glycine is now negatively charged, as depicted in Figure 1.7.

Figure 1.7. Amino Groups Become Deprotonated at Basic pH, Forming an Anion


At highly acidic pH values, amino acids tend to be positively charged. At highly alkaline pH values, amino acids tend to be negatively charged.


Because of these acid–base properties, amino acids are great candidates for titrations. We assume that the titration of each proton occurs as a distinct step, resembling that of a simple monoprotic acid. Thus, the titration curve looks like a combination of two monoprotic acid titration curves (or three curves, if the side chain is charged). Figure 1.8 shows the titration curve for glycine. After we inspect this curve, we'll look at the differences for the amino acids with charged side chains.

Figure 1.8. Titration Curve for Glycine

Imagine an acidic 1 M glycine solution. At low pH values, glycine exists predominantly as +NH3CH2COOH; it is fully protonated, with a positive charge. As the solution is titrated with NaOH, the carboxyl group will deprotonate first because it is more acidic than the amino group. When 0.5 equivalents of base have been added to the solution, the concentrations of the fully protonated glycine and its zwitterion, +NH3CH2COO, are equal; that is, [+NH3CH2COOH] = [+NH3CH2COO]. At this point, the pH equals pKa1. Remember: when the pH is close to the pKa value of a solute, a solution is acting as a buffer, and the titration curve is relatively flat, as demonstrated in the blue box in the diagram.


When the pH of a solution is approximately equal to the pKa of the solute, the solution acts as a buffer.

As we add more base, the carboxylate group goes from half-deprotonated to fully deprotonated. The amino acid stops acting like a buffer, and pH starts to increase rapidly during this phase. When we've added 1.0 equivalent of base, glycine exists exclusively as the zwitterion form (remember, we started with 1.0 equivalent of glycine). This means that every molecule is now electrically neutral, and thus the pH equals the isoelectric point (pI) of glycine. This is true of all amino acids: the isoelectric point is the pH at which the molecule is electrically neutral. For neutral amino acids, it can be calculated by averaging the two pKa values for the carboxylic acid and amino groups:

Equation 1.1

For glycine, the pI value is (2.34 + 9.60) ÷ 2 = 5.97. Remember that when the molecule is neutral, it is especially sensitive to pH changes, and the titration curve is nearly vertical.


When the pH of an amino acid solution equals the isoelectric point (pI) of the amino acid, it exists as electrically neutral molecules. The pI is calculated as the average of the two nearest pKa values. For amino acids with non-ionizable side chains, the pI is usually around 6.

As we continue adding base, glycine passes through a second buffering phase as the amino group deprotonates; again, the pH remains relatively constant. When 1.5 equivalents of base have been added, the concentration of the zwitterion form equals the concentration of the fully deprotonated form; that is, [+NH3CH2COO] = [NH2CH2COO], and the pH equals pKa2. Once again, the titration curve is nearly horizontal. Finally, when we've added 2.0 equivalents of base, the amino acid has become fully deprotonated, and all that remains is NH2CH2COO; additional base will only increase the pH further.

Amino Acids with Charged Side Chains

For amino acids with charged side chains, such as glutamic acid and lysine, the titration curve has an extra “step,” but works along the same principles as described above.

Because glutamic acid has two carboxyl groups and one amino group, its charge in its fully protonated state is still +1. It undergoes the first deprotonation, losing the proton from its main carboxyl group, just as glycine does. At that point, it is electrically neutral. When it loses its second proton, just as with glycine, its overall charge will be –1. However, the second proton that is removed in this case comes from the side chain carboxyl group, not the amino group! This is a relatively acidic group, with a pKa of around 4.2. The result is that the pI of glutamic acid is much lower than that of glycine, around 3.2. The isoelectric point for an acidic amino acid can be calculated as follows:

Equation 1.2

Lysine, on the other hand, has two amino groups and one carboxyl group. Thus, its charge in its fully protonated state is +2, not +1. Losing the carboxyl proton, which still happens around pH 2, brings the charge down to +1. Lysine does not become electrically neutral until it loses the proton from its main amino group, which happens around pH 9. It gets a negative charge when it loses the proton on the amino group in its side chain, which happens around pH 10.5. Thus, the isoelectric point of lysine is the average of the pKa values for the amino group and side chain; the pI is around 9.75. The isoelectric point for a basic amino acid can be calculated as follows:

Equation 1.3


Amino acids with acidic side chains have pI values well below 6; amino acids with basic side chains have pI values well above 6.

The take-home message: amino acids with acidic side chains have relatively low isoelectric points, while those with basic side chains have relatively high ones.

MCAT Concept Check 1.2:

Before you move on, assess your understanding of the material with these questions.

1.    For a generic amino acid, NH2CRHCOOH, with an uncharged side chain, what would be the predominant form at each of the following pH values?

·        pH = 1:

·        pH = 7:

·        pH = 11:

2.    Given the following pKa values, what is the value of the pI for each of the amino acids listed below?

·        Aspartic acid (pKa1 = 1.88, pKa2 = 3.65, pKa3 = 9.60): pI = 

·        Arginine (pKa1 = 2.17, pKa2 = 9.04, pKa3 = 12.48): pI = 

·        Valine (pKa1 = 2.32, pKa2 = 9.62): pI =