Amino acids are building blocks of proteins. These compounds contain at least one amino group \(\left(-\mathrm{NH}_{2}\right)\) and one carboxyl group \((-\mathrm{COOH})\) Consider glycine \(\left(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{COOH}\right) .\) Depending on the pH of the solution, glycine can exist in one of three possible forms: $$ \begin{array}{l} \text { Fully protonated: } \mathrm{NH}_{3}-\mathrm{CH}_{2}-\mathrm{COOH} \\ \text { Dipolar ion: } \mathrm{NH}_{3}-\mathrm{CH}_{2}-\mathrm{COO}^{-} \\ \text {Fully ionized: } \mathrm{NH}_{2}-\mathrm{CH}_{2}-\mathrm{COO}^{-} \end{array} $$ Predict the predominant form of glycine at \(\mathrm{pH} 1.0\), 7.0, and 12.0. The \(\mathrm{p} K_{\mathrm{a}}\) of the carboxyl group is 2.3 and that of the ammonium group \(\left(-\mathrm{NH}_{3}^{+}\right)\) is 9.6

Understanding the pKa value of a molecule, like an amino acid, is crucial for predicting its behavior in different pH environments. The pKa value tells us the pH at which half of the species is in the protonated form, and half is in the deprotonated form. For amino acids, there are at least two important pKa values to consider: one for the carboxyl group and one for the amino group.

For instance, in our glycine example, the pKa for the carboxyl group is 2.3, and the pKa for the ammonium group is 9.6. When the pH of the surroundings is lower than the pKa value, the functional group tends to be protonated. Conversely, if the pH is higher than the pKa value, the group is usually deprotonated. It's this delicate balance between protonated and deprotonated forms regulated by pKa values that allows for the accurate prediction of an amino acid's state at any given pH.

The Henderson-Hasselbalch equation is like a GPS for chemists to navigate the acid-base landscape. It provides a relationship between the pH, pKa, and the ratio of the protonated to the deprotonated form of a group in a species like an amino acid. Mathematically represented as \(\text{pH} = \text{pKa} + \log \left(\frac{\text{[A-]}}{\text{[HA]}}\right)\), where \([A-]\) is the concentration of the deprotonated form and \([HA]\) is the concentration of the protonated form.

This equation effectively helps predict how much of the compound will be in a certain ionization state at a specified pH. For example, by using the Henderson-Hasselbalch equation and the pKa values provided for glycine, students can determine that at pH 7, the zwitterion form of glycine, which has the carboxyl group deprotonated and the amino group protonated, will be prevalent.

A zwitterion is a molecule that carries both positive and negative charges, but which is overall electrically neutral. This peculiar form is common in amino acids, like glycine, under certain pH conditions. When an amino acid is in its zwitterionic form, it typically has a protonated amine (\(\text{NH}_3^{+}\)) and a deprotonated carboxylate (\(\text{COO}^{-}\)) group.

In the case of glycine, at a pH close to the average of its two pKa values (around 6.0, which is between 2.3 and 9.6), the molecule is most likely to be found in its zwitterionic form. This dual nature of zwitterions plays an important role in the physical properties of amino acids, such as their solubility in water and their ability to act as buffers in biological systems.