Most acid-base titrations are performed in an aqueous medium. In aqueous titrations, water competes with weaker acids or bases for proton donation or acceptance, leading to ambiguous endpoints in the titration curve. Water also affects the partial ionization of weak acids or bases. For example, water accepts a proton from acetic acid to form hydronium and acetate ions. The hydronium ion formed is a stronger acid than acetic acid, and the acetate ion is a stronger base than water. As a result, they react to give back the reactants. The effect of this process on weak acids and bases means that calculations based purely on pKa values may yield inaccurate results.
Using a non-aqueous solvent like ammonia, which is a stronger base than water, enables the complete ionization of acetic acid into acetate ion, effectively turning acetic acid into a strong acid in ammonia. Put another way, ammonia has a higher dissociation constant (Ks) than water, thereby increasing the equivalence point and sharpening the endpoint in the titration curve of acetic acid. This is commonly observed in the titration of weak acids and bases in non-aqueous solvents. Note that Ks of water is denoted by Kw.
There are four types of non-aqueous solvents: aprotic–unable to donate protons, protophilic–able to accept protons, protogenic–able to donate protons, and amphoteric–able to donate and accept protons. In addition to their uses in the titration of weak acids or bases, some of these solvents can also be used effectively in titrating organic analytes, which have poor solubility in water. The reactions occurring in non-aqueous titrations are explained by the Bronsted-Lowry theory of acids and bases. Here, while an acid behaves as a proton donor, the base is a proton acceptor.
