SO4

SO4 commonly refers to the sulfate ion, SO4^2−, a tetrahedral oxoanion of sulfur in oxidation state +6. In the ion, sulfur is surrounded by four oxygen atoms in a tetrahedral arrangement. Resonance among several S–O structures makes all four bonds equivalent, giving an average S–O bond length around 1.48–1.49 Å and O–S–O angles close to 109.5 degrees. The sulfate ion carries a −2 charge and forms salts with many cations.

Sulfates occur broadly in nature and industry. In minerals, examples include gypsum (CaSO4·2H2O) and anhydrite (CaSO4), as well as barite (BaSO4) and celestine (SrSO4). They are abundant in seawater and soils as dissolved ions. Industrially, sulfates are encountered as salts such as sodium sulfate (Na2SO4), ammonium sulfate ((NH4)2SO4), and magnesium sulfate (MgSO4). They arise or persist at various stages of sulfuric acid production and metallurgical processes, and are common byproducts or feedstocks in chemical manufacturing.

Applications span agriculture, construction, medicine, and industry. Fertilizers such as ammonium sulfate are widely used; gypsum serves as a building material and soil amendment; magnesium sulfate is used medically as a laxative and in obstetrics, and sodium sulfate finds use as a drying agent and in detergents. Sulfates also appear in paper and textile processing and as components of various industrial salts. Environmentally, sulfate salts are generally of low acute toxicity, though high intake can cause gastrointestinal effects; atmospheric sulfate aerosols influence air quality and can contribute to acid rain upon deposition. In biogeochemical cycles, sulfate can be reduced to sulfide by sulfate-reducing bacteria under anaerobic conditions.