Amphiproticity

Amphiprotic substances, or amphoteric species, are chemical entities that can act as either proton donors (acids) or proton acceptors (bases) in Bronsted-Lowry acid–base reactions. This dual capability arises when a species has a reasonably stable conjugate acid and conjugate base in solution, allowing it to participate in proton transfer in either direction.

Water is the classic example. It can donate a proton to form OH−, or accept a proton to form H3O+. Bicarbonate (HCO3−) is another well-known amphiprotic species: it can lose a proton to become CO3^2− or gain a proton to form H2CO3. Many amino acids are amphiprotic in aqueous solution, often existing as zwitterions near physiological pH and acting as acids or bases depending on the surrounding pH.

In buffer systems, amphiprotic species help stabilize pH over a range between their two relevant pKa values. Their acid–base behavior is dictated by these equilibria and their relative strengths, determining the pH at which they predominantly exist as one form or another. This property makes amphiprotic substances important in biological contexts, environmental chemistry, and analytical applications such as buffering and electrophoresis.

It is important to distinguish amphiproticity from amphiphilicity. Amphiproticity concerns proton exchange in acid–base chemistry, whereas amphiphilicity describes molecules with both hydrophilic and hydrophobic regions. Some solvents are amphiprotic; water’s autoprotolysis is a fundamental example that underpins many pH and pKa calculations in chemistry.