Sodium thiosulfate

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Sodium thiosulfate
Sodium thiosulfate
Crystal structure of sodium thiosulfate pentahydrate
Sodium thiosulfate.jpg
Names
IUPAC name
Sodium thiosulfate
Other names
Sodium hyposulfite
Hyposulphite of soda
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.970
E number E539 (acidity regulators, ...)
RTECS number XN6476000
UNII
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
Safety data sheet External MSDS
R-phrases (outdated) R21 R36 R37 R38
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentineReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
1
0
Flash point Non-flammable
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is ☑Y☒N ?)
Infobox references

Sodium thiosulfate (sodium thiosulphate) is a chemical and medication. As a medication it is used to treat cyanide poisoning and pityriasis versicolor.[2]

It is an inorganic compound with the formula Na2S2O3.xH2O. Typically it is available as the white or colorless pentahydrate, Na2S2O3·5H2O. The solid is an efflorescent (loses water readily) crystalline substance that dissolves well in water. It is also called sodium hyposulfite or "hypo".[3]

It is on the World Health Organization's List of Essential Medicines, the most effective and safe medicines needed in a health system.[4]

Uses[edit]

Medical uses[edit]

Iodometry[edit]

In analytical chemistry, the most important use comes because the thiosulfate anion reacts stoichiometrically with iodine in aqueous solution, reducing it to iodide as it is oxidized to tetrathionate:

2 S2O32− + I2 → S4O62− + 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.

Photographic processing[edit]

Silver halides, e.g., AgBr, typical components of photographic emulsions, dissolve upon treatment with aqueous thiosulfate:

2 S2O32− + AgBr → [Ag(S2O3)2]3− + Br

This application as a photographic fixer was discovered by John Herschel. It is used for both film and photographic paper processing; the sodium thiosulfate is known as a photographic fixer, and is often referred to as 'hypo', from the original chemical name, hyposulphite of soda.[13] Ammonium thiosulfate is typically preferred to sodium thiosulfate for this application.[3]

Gold extraction[edit]

Sodium thiosulfate and ammonium thiosulfate area component of an alternative lixiviants to cyanide for extraction of gold.[14][3] Thiosulfate forms strong soluble complexes with gold(I) ions, [Au(S2O3)2]3−. The advantagea of this approach are that (i) thiosulfate is essentially nontoxic and (ii) that ore types that are refractory to gold cyanidation (e.g. carbonaceous or Carlin-type ores) can be leached by thiosulfate. Some problems with this alternative process include the high consumption of thiosulfate, and the lack of a suitable recovery technique, since [Au(S2O3)2]3− does not adsorb to activated carbon, which is the standard technique used in gold cyanidation to separate the gold complex from the ore slurry.

Neutralizing chlorinated water[edit]

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.[3] 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 (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 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[edit]

Two polymorphs are known of the pentahydrate. The anhydrous salt exists in several polymorphs.[3] 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 bears significant negative charge and the S-O interactions have more double-bond character.

Production[edit]

On an industrial scale, sodium thiosulfate is produced chiefly from liquid waste products of sodium sulfide or sulfur dye manufacture.[15]

In the laboratory, this salt can be prepared by heating an aqueous solution of sodium sulfite with sulfur or by boiling aqueous sodium hydroxide and sulfur according to this equation:[16]

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

Principal reactions[edit]

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:[15]

Na2S2O3 + 2 HCl → 2 NaCl + S + SO2 + H2O

This reaction is known as a "clock reaction", because when the sulfur reaches a certain concentration, the solution turns from colourless to a pale yellow. This reaction has been employed to generate colloidal sulfur. This process is used to demonstrate the concept of reaction rate in chemistry classes.

Aluminium cation reaction[edit]

Sodium thiosulfate is also used in analytical chemistry.[citation needed] It can, when heated with a sample containing aluminium cations, produce a white precipitate:

2 Al3+ + 3 S2O32− + 3 H2O → 3 SO2 + 3 S + 2 Al(OH)3

Organic chemistry[edit]

Alkylation of sodium thiosulfate gives S-alkylthiosulfates, which are called Bunte salts.[17] 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

References[edit]

  1. ^ Record in the GESTIS Substance Database of the Institute for Occupational Safety and Health
  2. ^ a b WHO Model Formulary 2008 (PDF). World Health Organization. 2009. p. 66. ISBN 9789241547659. Retrieved 8 January 2017.
  3. ^ a b c d e J. J. Barbera, A. Metzger, M. Wolf (2012). "Sulfites, Thiosulfates, and Dithionites". Ullmann's Encyclopedia of Industrial Chemistry. Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_477. ISBN 978-3527306732.
  4. ^ "WHO Model List of Essential Medicines (19th List)" (PDF). World Health Organization. April 2015. Retrieved 8 December 2016.
  5. ^ "Toxicity, Cyanide: Overview - eMedicine". Retrieved 2009-01-01.
  6. ^ Hall AH, Dart R, Bogdan G (June 2007). "Sodium thiosulfate or hydroxocobalamin for the empiric treatment of cyanide poisoning?". Ann Emerg Med. 49 (6): 806–13. doi:10.1016/j.annemergmed.2006.09.021. PMID 17098327.
  7. ^ Cicone JS, Petronis JB, Embert CD, Spector DA (June 2004). "Successful treatment of calciphylaxis with intravenous sodium thiosulfate". Am. J. Kidney Dis. 43 (6): 1104–8. doi:10.1053/j.ajkd.2004.03.018. PMID 15168392.
  8. ^ http://www.medscape.com/viewarticle/762244
  9. ^ Selk N, Rodby RA (Jan–Feb 2011). "Unexpectedly severe metabolic acidosis associated with sodium thiosulfate therapy in a patient with calcific uremic arteriolopathy". Semin. Dial. 24 (1): 85–8. doi:10.1111/j.1525-139X.2011.00848.x. PMID 21338397.
  10. ^ "Sodium thiosulfate" at Dorland's Medical Dictionary
  11. ^ Journal of Pharmacology and Experimental Therapeutics (September 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.
  12. ^ Brock, Penelope R.; Maibach, Rudolf; Childs, Margaret; Rajput, Kaukab; Roebuck, Derek; Sullivan, Michael J.; Laithier, Véronique; Ronghe, Milind; Dall'Igna, Patrizia; Hiyama, Eiso; Brichard, Bénédicte; Skeen, Jane; Mateos, M. Elena; Capra, Michael; Rangaswami, Arun A.; Ansari, Marc; Rechnitzer, Catherine; Veal, Gareth J.; Covezzoli, Anna; Brugières, Laurence; Perilongo, Giorgio; Czauderna, Piotr; Morland, Bruce; Neuwelt, Edward A. (2018). "Sodium Thiosulfate for Protection from Cisplatin-Induced Hearing Loss". New England Journal of Medicine. 378 (25): 2376–2385. doi:10.1056/NEJMoa1801109. PMC 6117111. PMID 29924955.
  13. ^ Charles Robert Gibson (1908). The Romance of Modern Photography, Its Discovery & Its Achievements. Seeley & Co. p. 37.
  14. ^ Aylmore MG, Muir DM (2001). "Thiosulfate Leaching of Gold - a Review". Minerals Engineering. 14 (2): 135–174. doi:10.1016/s0892-6875(00)00172-2.
  15. ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
  16. ^ Gordin, H. M. (1913). Elementary Chemistry, Volume I. Inorganic Chemistry. Chicago: Medico-Dental Publishing Co. pp. 162 & 287–288.
  17. ^ M. E. Alonso and H. Aragona (1978). "Sulfide Synthesis in Preparation of Unsymmetrical Dialkyl Disulfides: Sec-butyl Isopropyl Disulfide". Org. Synth. 58: 147. doi:10.15227/orgsyn.058.0147.