An acid is a substance capable of donating a proton, or hydrogen ion (\(\text{H}^+\)), a definition central to the Brønsted-Lowry theory. Triprotic acids are a specific class of proton-donating substances characterized by their ability to donate three separate protons per molecule. This structure allows them to undergo three distinct stages of ionization, making them chemically versatile.
Stepwise Dissociation The Three Ionization Stages
The defining characteristic of a triprotic acid is that it loses its three protons sequentially, a process known as stepwise dissociation. Each removal of a proton represents a distinct chemical equilibrium. For a general triprotic acid, \(\text{H}_3\text{A}\), the process begins with the loss of the first proton, forming \(\text{H}_2\text{A}^-\).
The tendency for an acid to release a proton is quantified by the acid dissociation constant (\(\text{K}_a\)). Because there are three stages of dissociation, there are three distinct acid dissociation constants: \(\text{K}_{a1}\), \(\text{K}_{a2}\), and \(\text{K}_{a3}\). The first proton is always the easiest to remove, resulting in the largest value for \(\text{K}_{a1}\).
The removal of the second and third protons is significantly more difficult than the first because of the increasing negative charge on the remaining molecule. Removing a positively charged proton from the already negative ion (\(\text{H}_2\text{A}^-\)) requires more energy, effectively weakening the acid in the second step. This difficulty is reflected in the much smaller value of \(\text{K}_{a2}\).
The trend continues with the final proton, which is removed from a doubly negative ion (\(\text{HA}^{2-}\)) to form the fully deprotonated ion (\(\text{A}^{3-}\)). Therefore, the acid dissociation constants follow a predictable relationship where \(\text{K}_{a1} > \text{K}_{a2} > \text{K}_{a3}\). These successive constants often differ by a factor of \(10^5\) to \(10^6\), meaning the first ionization is vastly more extensive than the later ones.
Key Examples and Real-World Applications
Phosphoric acid (\(\text{H}_3\text{PO}_4\)) is the most commonly cited example of a triprotic acid. This inorganic acid plays a significant role in various industries and biological systems. It is used extensively in the production of fertilizers, where it supplies the nutrient phosphorus to crops.
A familiar application of phosphoric acid is its use as an acidulant in carbonated beverages, such as cola-flavored sodas. In these drinks, it provides a sharp flavor and acts as a preservative by inhibiting the growth of mold and bacteria. The three distinct \(\text{K}_a\) values for phosphoric acid are \(\text{K}_{a1} \approx 7.5 \times 10^{-3}\), \(\text{K}_{a2} \approx 6.2 \times 10^{-8}\), and \(\text{K}_{a3} \approx 4.2 \times 10^{-13}\).
Another biologically relevant example is citric acid, a weak organic acid found naturally in citrus fruits. It possesses three carboxylic acid groups, allowing it to donate three protons. Citric acid is a natural preservative and flavoring agent in foods and beverages, and it is a central molecule in the metabolic citric acid cycle, which drives energy production in living organisms. The fully deprotonated forms of phosphoric acid are also fundamental to life, as the phosphate group forms the backbone of deoxyribonucleic acid (\(\text{DNA}\)) and ribonucleic acid (\(\text{RNA}\)).
How Triprotic Acids Function as Buffers
A buffer system is a solution that resists changes in \(\text{pH}\) when small amounts of acid or base are added, achieved by having a mixture of a weak acid and its conjugate base. Triprotic acids are effective buffers because their three dissociation stages create three separate buffering zones.
Each dissociation step of a triprotic acid has a unique \(\text{pKa}\) value (the negative logarithm of the \(\text{K}_a\) value). A buffer works best when the \(\text{pH}\) of the solution is close to the \(\text{pKa}\) of the weak acid component. Since a triprotic acid has three different \(\text{pKa}\) values, it can maintain \(\text{pH}\) stability over a much wider range than a monoprotic acid.
The phosphate buffer system, derived from phosphoric acid, is one of the most important buffers in the human body, operating in the intracellular fluid. The second dissociation step, involving the conversion of \(\text{H}_2\text{PO}_4^-\) to \(\text{HPO}_4^{2-}\), has a \(\text{pKa}\) value near \(7.2\). This value is close to the physiological \(\text{pH}\) of \(7.4\), making the phosphate system an efficient regulator of \(\text{pH}\) inside cells.