Bisulfate

Bisulfate, commonly called hydrogen sulfate, denotes the anion HSO4− and the salts that contain it. It is the conjugate base of sulfuric acid (H2SO4) and the conjugate acid of sulfate (SO4^2−). In common usage, bisulfate refers to HSO4− in solution or to salts such as sodium bisulfate (NaHSO4), potassium bisulfate (KHSO4), and ammonium bisulfate (NH4HSO4). The term is distinct from bisulfite (HSO3−).

In aqueous solution, bisulfate is amphiprotic: it can donate a proton to form sulfate (SO4^2−) or accept a proton to form sulfuric acid (H2SO4) in highly acidic media. Its second dissociation constant is about Ka2 ≈ 1.2×10−2 (pKa ≈ 1.99), reflecting its moderate acidity when acting as an acid. The anion adopts a tetrahedral arrangement around sulfur, and many bisulfate salts are highly soluble in water. On heating, bisulfates decompose to sulfates, often via the formation of pyrosulfates; for example, 2 HSO4− → S2O7^2− + H2O in the molten state.

Preparation and occurrence: bisulfate salts are formed by partial neutralization of sulfuric acid with one equivalent of base. They are encountered in various chemical processes and are isolated as practical salts for use in industry and agriculture.

Uses: sodium bisulfate is used as a pH-lowering agent in cleaning products and in pools; potassium and ammonium bisulfates serve as sources of sulfate and, where appropriate, potassium or ammonium in fertilizers. These salts provide acidity and good water solubility for various applications.

Safety: bisulfates are corrosive and should be handled with appropriate protection. They can release sulfuric acid fumes when heated and may irritate skin, eyes, and the respiratory tract.