James Richard Fromm
The term amphiprotic in modern acid-base chemistry is the replacement for the older term amphoteric. An amphiprotic substance is a substance which can act both as an acid and as a base, participating in two acid-base equilibria, because it contains at least one proton which can be given up and at least one site at which a proton can be acquired. Most polyprotic acids have at least one amphiprotic ion. Using phosphoric acid as an example, the monohydrogen phosphate ion and the dihydrogen phosphate ion are both amphiprotic while phosphoric acid itself can only be an acid and the phosphate ion can only be a base.
Organic compounds which contain both a carboxylic acid group and an amine group on the same molecule are called amino acids. Many amino acids are essential to living organisms and will be discussed later in this course. When an amino acid such as glycine, H2NCH2COOH, is dissolved in water. the carboxylic acid group loses a proton which is gained by the more basic amine group. This produces an ionic structure with opposite charges on both ends, a zwitterion. The zwitterion structure of glycine is +H3NCH2COO-. The protonated form of this amphiprotic zwitterion, +H3NCH2COOH, is the glycinium ion.
When an amphiprotic substance alone is dissolved in water, it will act both as an acid and as a base.
Example. Dissolution of sodium hydrogen carbonate, NaHCO3, will involve the two equilibria
H2O + H2CO3 HCO3- + H3O+ and H2O + HCO3- CO32- + H3O+.
Two acid ionization equilibria are involved and both equilibrium constants must be used to make a realistic calculation of the pH of the solution of the amphiprotic substance:
K1 = [H3O+][HCO32-]/[H2CO3]
K2 = [H3O+][CO32-]/[HCO3-]
K1K2 = [H3O+]2[CO32-]/[H2CO3]
The dissociation of the hydrogen carbonate ion HCO3- to the carbonate ion CO32- will release a proton to protonate the hydrogen carbonate ion HCO3- to hydrogen carbonate ("carbonic acid", actually CO2(aq) + H2O), H2CO3. This reaction will occur stoichiometrically since H3O+ is a strong acid while HCO3- is not. As a consequence [H2CO3] = [CO32-] and
[H3O+]2 = K1K2, [H3O+] = 4.52 x 10-9, pH = 8.35.
The pH of the solution is independent of the concentration of the amphiprotic substance in the above example. Similar calculations can be done for any amphiprotic substance. Solutions of amphiprotic substances are encountered less frequently than are solutions of weak acids and weak bases. They do not hold pH well and are not effective as buffer solutions.
In all cases it will be found that the solution pH is dependent on the two ionization constants of the amphiprotic substance and independent of the concentration of the amphiprotic substance in the solution. Polyprotic acids which have not lost all of their protons are the most commonly encountered amphiprotic substances.
Example. Let us calculate the pH of a 0.01 molar aqueous solution of disodium hydrogen phosphate, Na2HPO4. The two equilibria involved and their acid ionization constants are:
K2 = [H3O+][HPO42-]/[H2PO4-]
K3 = [H3O+][PO43-]/[HPO42-]
K2K3 = [H3O+]2[PO43-]/[H2PO4-] = [H3O+]2
[H3O+]2 = (6.23 x 10-8)(4.55 x 10-13); [H3O+] = 1.68 x 10-10, pH = 9.77