Acids are substances that donate a hydrogen ion, often called a proton (\(\text{H}^+\)), when dissolved in a solution like water. The release of this proton gives acids their characteristic properties. Most acids, known as monoprotic acids (e.g., \(\text{HCl}\) or \(\text{HNO}_3\)), possess only one proton to donate. Polyprotic acids, however, are capable of releasing two or more protons per molecule. Their behavior is more complex, as the release of these multiple hydrogens is a controlled, sequential process dictated by fundamental chemical forces.
What Defines a Polyprotic Acid
A polyprotic acid contains multiple ionizable hydrogen atoms within its molecular structure. These hydrogens are covalently bonded and are released when the acid reacts with a base or water. The total number of removable protons determines the classification; for example, an acid donating two protons is diprotic, and one donating three is triprotic.
When a polyprotic acid releases a proton, the remaining structure is called its conjugate base. This conjugate base retains one or more acidic hydrogens, allowing it to act as an acid in a subsequent reaction step.
The Stepwise Removal of Hydrogens
The defining characteristic of polyprotic acid dissociation is that hydrogens are never removed simultaneously. The process occurs in distinct, successive stages, releasing one proton at a time. In the first step, the neutral acid molecule (\(\text{H}_2\text{A}\)) donates its initial proton to form a singly charged conjugate base (\(\text{HA}^-\)).
This conjugate base (\(\text{HA}^-\)) then acts as the acid for the second stage. It donates a second proton to form a new conjugate base, which carries a double negative charge (\(\text{A}^{2-}\)). For triprotic acids, this sequence continues for a third step, creating a more negatively charged conjugate base with each reaction. Each ionization step is an equilibrium process.
The Chemical Reason for Decreasing Acidity
The sequential removal of protons is governed by a measurable decrease in the acid’s strength after each step. This reduction is quantified by the acid dissociation constant (\(\text{K}_\text{a}\)), which measures how readily an acid gives up its proton. For a polyprotic acid, each step has its own constant (\(\text{K}_{a1}\), \(\text{K}_{a2}\), etc.). The initial constant, \(\text{K}_{a1}\), is always significantly larger than \(\text{K}_{a2}\), which is larger than \(\text{K}_{a3}\).
This consistent trend (\(\text{K}_{a1} > \text{K}_{a2} > \text{K}_{a3}\)) results from electrostatic attraction. The first proton is removed easily from a neutral molecule. Once the first proton is gone, the molecule becomes a negatively charged conjugate base. Removing a second, positively charged proton (\(\text{H}^+\)) from this negative species requires overcoming a strong electrostatic force.
The negative charge strongly holds onto the remaining protons, making their removal much more difficult. This difficulty is reflected in the \(\text{K}_\text{a}\) values, which typically decrease by a factor of \(10^5\) to \(10^6\) between successive steps. Consequently, the first proton is the most acidic and readily released, while the last proton is the least acidic and hardest to remove.
Common Examples in Chemistry and Biology
The principle of stepwise deprotonation is observed in many molecules central to chemistry and biological systems. Phosphoric acid (\(\text{H}_3\text{PO}_4\)), a triprotic acid, is a common example found in soft drinks and DNA. Its first dissociation step creates the dihydrogen phosphate ion (\(\text{H}_2\text{PO}_4^-\)), a relatively strong acid.
The second step removes a proton from \(\text{H}_2\text{PO}_4^-\) to form the monohydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), a much less favorable process. The third proton is removed from the doubly charged \(\text{HPO}_4^{2-}\), a reaction that barely proceeds in water. Carbonic acid (\(\text{H}_2\text{CO}_3\)), a diprotic acid, is another prominent example responsible for the body’s primary blood buffering system.
This stepwise release allows polyprotic acids to exist in multiple forms simultaneously. This enables them to effectively resist changes in \(\text{pH}\) across a wide range of acidity and maintain the stability of biological environments.