Sulfate

Sulfate, or the sulfate ion, is SO4^2−. The sulfur atom is in the +6 oxidation state and the ion adopts a tetrahedral geometry. Through resonance, the negative charge is delocalized over the four oxygen atoms, resulting in two equivalent S–O bonds and two longer bonds.

Naturally occurring sulfates appear in minerals such as gypsum (CaSO4·2H2O), anhydrite (CaSO4), and barite (BaSO4). Industrially, sulfates are produced by oxidizing sulfur dioxide to sulfur trioxide and then hydrating to sulfuric acid; neutralization yields sulfate salts like sodium sulfate (Na2SO4) and calcium sulfate (CaSO4). Gypsum and anhydrite are also mined as minerals.

Common sulfate compounds include sodium sulfate, calcium sulfate, and magnesium sulfate. Sulfate salts have wide uses: sulfuric acid is a foundational chemical in industry; ammonium sulfate and other sulfates provide fertilizers; sodium lauryl sulfate and related esters are widely used in detergents and personal care products; calcium sulfate is used in plaster and drywall; magnesium sulfate (Epsom salt) has medical and agricultural applications. Organic sulfates form sulfate esters (R–OSO3−) found in various drugs and biomolecules.

Environmental and health aspects: atmospheric sulfur dioxide oxidizes to sulfates, contributing to acid rain and sulfate aerosols that affect climate and air quality. In water, sulfates are common and can influence hardness, corrosion, and taste at higher concentrations. In biology, sulfate is a nutrient that organisms assimilate and reduce to incorporate sulfur into biomolecules such as cysteine and methionine; sulfate-reducing bacteria can generate sulfide under anaerobic conditions.

Overall, sulfate is a versatile and widespread chemical motif with both inorganic and organic significance.