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Sodium thiosulfate

Sodium thiosulfate
Sodium thiosulfate
Names
IUPAC name
Sodium thiosulfate
Other names
Sodium hyposulphite
Hyposulphite of soda
Hypo
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.970 Edit this at Wikidata
EC Number
  • anhydrous: 231-867-5
E number E539 (acidity regulators, ...)
RTECS number
  • anhydrous: XN6476000
UNII
  • InChI=1S/2Na.H2O3S2/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 checkY
    Key: AKHNMLFCWUSKQB-UHFFFAOYSA-L checkY
  • anhydrous: InChI=1/2Na.H2O3S2/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2
    Key: AKHNMLFCWUSKQB-NUQVWONBAM
  • pentahydrate: InChI=1S/2Na.H2O3S2.5H2O/c;;1-5(2,3)4;;;;;/h;;(H2,1,2,3,4);5*1H2/q2*+1;;;;;;/p-2
    Key: PODWXQQNRWNDGD-UHFFFAOYSA-L
  • anhydrous: [Na+].[Na+].[O-]S(=O)(=O)[S-]
  • pentahydrate: O.O.O.O.O.[Na+].[Na+].[O-]S(=O)(=O)[S-]
Properties
Na2S2O3
Molar mass 158.11 g/mol (anhydrous)
248.18 g/mol (pentahydrate)
Appearance White crystals
Odor Odorless
Density 1.667 g/cm3
Melting point 48.3 °C (118.9 °F; 321.4 K) (pentahydrate)
Boiling point 100 °C (212 °F; 373 K) (pentahydrate, - 5H2O decomposition)
70.1 g/100 mL (20 °C)[1]
231 g/100 mL (100 °C)
Solubility negligible in alcohol
1.489
Structure
monoclinic
Hazards
GHS labelling:
GHS07: Exclamation mark
Warning
H315, H319, H335
P261, P264, P271, P280, P302+P352, P304+P340, P305+P351+P338, P312, P321, P332+P313, P337+P313, P362, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
1
0
0
Flash point Non-flammable
Safety data sheet (SDS) External MSDS
Related compounds
Other cations
 Thiosulfuric acid
Lithium thiosulfate
Potassium thiosulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

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), which 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]

Uses

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]

Medical uses

Main article: Sodium thiosulfate (medical use)

Sodium thiosulfate is used in the treatment of cyanide poisoning.[3] It is on the World Health Organization's List of Essential Medicines.[4][5] Other uses include topical treatment of ringworm and tinea versicolor,[3][6] and treating some side effects of hemodialysis[7] and chemotherapy.[8][9] In September 2022, the U.S. Food and Drug Administration (FDA) approved sodium thiosulfate under the trade name Pedmark to lessen the risk of ototoxicity and hearing loss in infant, child, and adolescent cancer patients receiving the chemotherapy medication cisplatin.[10][11]

Photographic processing

See also: Collodion § Wet-plate collodion photography

In photography, sodium thiosulfate is used in both film and photographic paper processing 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. 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.

Neutralizing chlorinated water

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.

In pH 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:

4 NaClO + Na2S2O3 + 2 NaOH → 4 NaCl + 2 Na2SO4 + H2O

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

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 thiosulfate anion 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.

Production

Sodium thiosulfate is prepared by oxidation of sodium sulfite with sulfur.[2] It is also produced from waste sodium sulfide from the manufacture of sulfur dyes.[13]

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.

6 NaOH + 4 S → 2 Na2S + Na2S2O3 + 3 H2O

Principal reactions

Upon heating to 300 °C, it decomposes to sodium sulfate and sodium polysulfide:

4 Na2S2O3 → 3 Na2SO4 + Na2S5

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]

8 Na2S2O3 + 16 HCl → 16 NaCl + S8 + 8 SO2 + 8 H2O

Coordination chemistry

Thiosulfate forms complexes with transition metal ions. One such complex is [Au(S2O3)2]3−.

Iodometry

Some analytical procedures exploit the oxidizability of thiosulfate anion by iodine. The reaction produces tetrathionate:

2 S2O2−3 + I2 → S4O2−6 + 2 I−

Due to the quantitative nature of this reaction, as well as because Na2S2O3·5H2O has an excellent shelf-life, it is used as a titrant in iodometry. 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.

Organic chemistry

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:

ClCH2CO2H + Na2S2O3 → Na[O3S2CH2CO2H] + NaCl
Na[O3S2CH2CO2H] + H2O → HSCH2CO2H + NaHSO4

Safety

Sodium thiosulfate has low toxicity. LDLo for rabbits is 4000 mg/kg.[2]

References

  1. ^ Record in the GESTIS Substance Database of the Institute for Occupational Safety and Health
  2. ^ a b c d e f g Barbera JJ, Metzger A, Wolf M (2012). "Sulfites, Thiosulfates, and Dithionites". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_477. ISBN 978-3-527-30673-2.
  3. ^ a b Stuart MC, Kouimtzi M, Hill SR, eds. (2009). WHO Model Formulary 2008. World Health Organization. p. 66. hdl:10665/44053. ISBN 978-92-4-154765-9.
  4. ^ 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.
  5. ^ 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.
  6. ^ Sunenshine PJ, Schwartz RA, Janniger CK (2002). "Tinea versicolor". Int. J. Dermatol. 37 (9): 648–55. doi:10.1046/j.1365-4362.1998.00441.x. PMID 9762812. S2CID 75657768.
  7. ^ Auriemma M, Carbone A, Di Liberato L, et al. (2011). "Treatment of Cutaneous Calciphylaxis with Sodium Thiosulfate: Two Case Reports and a Review of the Literature". Am. J. Clin. Dermatol. 12 (5): 339–46. doi:10.2165/11587060-000000000-00000. PMID 21834598. S2CID 28366905.
  8. ^ Orgel E, Villaluna D, Krailo MD, Esbenshade A, Sung L, Freyer DR (May 2022). "Sodium thiosulfate for prevention of cisplatin-induced hearing loss: updated survival from ACCL0431". The Lancet. Oncology. 23 (5): 570–572. doi:10.1016/S1470-2045(22)00155-3. PMC 9635495. PMID 35489339.
  9. ^ Dickey DT, Wu YJ, Muldoon LL, et al. (2005). "Protection against Cisplatin-Induced Toxicities by N-Acetylcysteine and Sodium Thiosulfate as Assessed at the Molecular, Cellular, and in Vivo Levels". J. Pharmacol. Exp. Ther. 314 (3): 1052–8. doi:10.1124/jpet.105.087601. PMID 15951398. S2CID 11381393.
  10. ^ Winstead, Edward (October 6, 2022). "Sodium Thiosulfate Reduces Hearing Loss in Kids with Cancer". National Cancer Institute. Retrieved March 9, 2023.
  11. ^ "FDA approves sodium thiosulfate to reduce the risk of ototoxicity associated with cisplatin in pediatric patients with localized, non-metastatic solid tumors". U.S. Food and Drug Administration. 20 September 2022. Retrieved 9 March 2023.
  12. ^ Gibson CR (1908). The Romance of Modern Photography, Its Discovery & Its Achievements. Seeley & Co. pp. 37. hyposulphite-of-soda herschel fixer hypo.
  13. ^ a b Holleman AF, Wiberg E, Wiberg N (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 978-0-12-352651-9.
  14. ^ Gordin HM (1913). Elementary Chemistry. Vol. 1. Inorganic Chemistry. Chicago: Medico-Dental Publishing Co. pp. 162 & 287–288.
  15. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  16. ^ Alonso ME, Aragona H (1978). "Sulfide Synthesis in Preparation of Unsymmetrical Dialkyl Disulfides: Sec-butyl Isopropyl Disulfide". Org. Synth. 58: 147. doi:10.15227/orgsyn.058.0147. Free access icon
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