Sodium thiosulfate (sodium thiosulphate) is an inorganic compound with the formula Na2S2O3·(H2O)(x). Typically it is available as the white or colorless pentahydrate (x = 5). It is a white solid that dissolves well in water. The compound is a reducing agent and a ligand, and these properties underpin its applications.[2]
Sodium thiosulfate is used predominantly in dyeing. It converts some dyes to their soluble colorless "leuco" forms. It is also used to bleach "wool, cotton, silk, ...soaps, glues, clay, sand, bauxite, and... edible oils, edible fats, and gelatin."[2]
In photography, sodium thiosulfate is known as a fixer, sometimes still called 'hypo' from the original chemical name, hyposulphite of soda.[12] It functions to dissolve silver halides, e.g., AgBr, components of photographic emulsions. It is used for both film and photographic paper processing. Ammonium thiosulfate is typically preferred to sodium thiosulfate for this application.[2] The ability of thiosulfate to dissolve silver ions is related to its ability to dissolve gold ions described below.
It is used to dechlorinate tap water including lowering chlorine levels for use in aquariums, swimming pools, and spas (e.g., following superchlorination) and within water treatment plants to treat settled backwash water prior to release into rivers.[2] The reduction reaction is analogous to the iodine reduction reaction.
InpH testing of bleach substances, sodium thiosulfate neutralizes the color-removing effects of bleach and allows one to test the pH of bleach solutions with liquid indicators. The relevant reaction is akin to the iodine reaction: thiosulfate reduces the hypochlorite (the active ingredient in bleach) and in so doing becomes oxidized to sulfate. The complete reaction is:
Similarly, sodium thiosulfate reacts with bromine, removing the free bromine from the solution. Solutions of sodium thiosulfate are commonly used as a precaution in chemistry laboratories when working with bromine and for the safe disposal of bromine, iodine, or other strong oxidizers.
Structure of sodium thiosulfate according to X-ray crystallography, showing the tetrahedral thiosulfate anion embedded in a network of sodium ions. Color code: red = O, yellow = S
Two polymorphs are known as pentahydrate. The anhydrous salt exists in several polymorphs.[2] In the solid state, the thiosulfateanion is tetrahedral in shape and is notionally derived by replacing one of the oxygen atoms by a sulfur atom in a sulfate anion. The S-S distance indicates a single bond, implying that the terminal sulfur holds a significant negative charge and the S-O interactions have more double-bond character.
This salt can also be prepared by boiling aqueous sodium hydroxide and sulfur according to the following equation.[14][15] However, this is not recommended outside of a laboratory, as exposure to hydrogen sulfide can result if improperly handled.
Thiosulfate salts characteristically decompose upon treatment with acids. Initial protonation occurs at sulfur. When the protonation is conducted in diethyl ether at −78 °C, H2S2O3 (thiosulfuric acid) can be obtained. It is a somewhat strong acid with pKas of 0.6 and 1.7 for the first and second dissociations, respectively. Under normal conditions, acidification of solutions of this salt excess with even dilute acids results in complete decomposition to sulfur, sulfur dioxide, and water:[13]
Due to the quantitative nature of this reaction, as well as because Na2S2O3·5H2O has an excellent shelf-life, it is used as a titrantiniodometry. Na2S2O3·5H2O is also a component of iodine clock experiments.
This particular use can be set up to measure the oxygen content of water through a long series of reactions in the Winkler test for dissolved oxygen. It is also used in estimating volumetrically the concentrations of certain compounds in solution (hydrogen peroxide, for instance) and in estimating the chlorine content in commercial bleaching powder and water.
Alkylation of sodium thiosulfate gives S-alkylthiosulfates, which are called Bunte salts.[16] The alkylthiosulfates are susceptible to hydrolysis, affording the thiol. This reaction is illustrated by one synthesis of thioglycolic acid:
^ abStuart MC, Kouimtzi M, Hill SR, eds. (2009). WHO Model Formulary 2008. World Health Organization. p. 66. hdl:10665/44053. ISBN978-92-4-154765-9.
^World Health Organization model list of essential medicines: 21st list 2019. Geneva: World Health Organization. 2019. hdl:10665/325771. WHO/MVP/EMP/IAU/2019.06. License: CC BY-NC-SA 3.0 IGO.
^World Health Organization model list of essential medicines: 22nd list (2021). Geneva: World Health Organization. 2021. hdl:10665/345533. WHO/MHP/HPS/EML/2021.02.
Y
