Animated Solution for Chemistry - Ionic Equilibrium: Which one of the following species acts as both Bronsted acid and base?

Bronsted Acids and Bases

When we dive into the world of ionic equilibrium, the definitions of acids and bases become incredibly dynamic. The Bronsted-Lowry theory revolutionized our understanding by defining an acid simply as a proton ($H^{+}$) donor and a base as a proton ($H^{+}$) acceptor.

But what happens when a molecule is caught in the middle? Some species have the remarkable ability to play both roles depending on the situation. These chemical chameleons are known as amphiprotic species. In this problem, we are on a hunt to find which of the given phosphorus-based anions possesses this dual nature.

Before we analyze the options, we must address a classic trap in chemistry: Not all hydrogens are created equal.

In the oxoacids of phosphorus, only the hydrogen atoms that are bonded to the highly electronegative oxygen atoms (forming $P-O-H$ bonds) are acidic and can be ionized. Hydrogens bonded directly to the central phosphorus atom ($P-H$ bonds) are essentially non-polar and cannot be donated as protons. This structural secret is the key to solving our mystery.

Let's interrogate each option based on its parent acid:

1. The $H_{2}P{O}_2^{-}$ Ion: This ion is derived from hypophosphorous acid ($H_{3}P{O}_2$). Despite having three hydrogens, $H_{3}P{O}_2$ is a monobasic acid. It has only one $P-O-H$ bond. Once it donates that single proton, it becomes $H_{2}P{O}_2^{-}$. The two remaining hydrogens are stuck directly to the phosphorus atom. Therefore, $H_{2}P{O}_2^{-}$ has no more protons to give. It can only accept a proton to revert to $H_{3}P{O}_2$. It is strictly a Bronsted base.

2. The $HP{O}_3^{2-}$ Ion: This ion comes from phosphorous acid ($H_{3}P{O}_3$), which is a dibasic acid with two $P-O-H$ bonds. The $HP{O}_3^{2-}$ ion has already lost both of its acidic protons. The single hydrogen it has left is bonded directly to phosphorus. Just like the previous case, it cannot donate this proton. It can only accept protons, making it strictly a Bronsted base.

3. The $HP{O}_4^{2-}$ Ion: Finally, we look at the derivative of phosphoric acid ($H_{3}P{O}_4$). Phosphoric acid is a tribasic acid; all three of its hydrogens are attached to oxygen atoms. The $HP{O}_4^{2-}$ ion has lost two protons, but crucially, it still retains one acidic $P-O-H$ bond.

This means it can act as an acid by donating its last proton to water:

Simultaneously, because it carries a negative charge, it can act as a base by accepting a proton from water:

Because the $HP{O}_4^{2-}$ ion can seamlessly switch between donating and accepting a proton, it perfectly fits the definition of an amphiprotic species. It acts as both a Bronsted acid and a Bronsted base, making it the correct answer to our problem.