Thiosulfat

Thiosulfate, also called thiosulfat in some languages, denotes the thiosulfate anion S2O3^2− and the salts derived from it, such as sodium thiosulfate (Na2S2O3) and potassium thiosulfate (K2S2O3). In aqueous solution, thiosulfate acts as a mild reducing agent and as a sulfur-bearing oxoanion with two sulfur atoms linked by an S–S bond.

Formation and structure: The ion is the conjugate base of thiosulfuric acid H2S2O3. In solution it can be oxidized to sulfite or sulfate; in acid solution it decomposes with sulfur precipitation: S2O3^2− + 2 H+ → S(s) + SO2 + H2O. The thiosulfate ion participates in redox chemistry and can form complexes with metals, and under certain conditions can be oxidized to tetrathionate (S4O6^2−).

Occurrence and production: Thiosulfate occurs commercially as salts and is produced industrially by treating sulfite solutions with sulfur, or by other sulfur-oxide interconversions. It is also formed in some biological and environmental processes. In water treatment, thiosulfate is used as a dechlorinating agent to neutralize residual chlorine.

Applications: Thiosulfates are employed in photography as fixers to dissolve silver halides after exposure, in chemistry as reducing agents and quenchers, and in mining for thiosulfate leaching of gold and silver. In medicine, sodium thiosulfate is used as an antidote in cyanide poisoning and as a mucolytic agent; it also serves as a source of sulfur in various enzymatic reactions (for example via the rhodanese enzyme).

Safety: Thiosulfate salts are generally of low toxicity but can be irritating; handle with standard laboratory precautions and avoid contact with strong oxidizers.