Sodium thiosulfate

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Sodium thiosulfate
Identifiers
CAS number7772-98-7 YesY, 10102-17-7 (pentahydrate) YesY
PubChem24477
ChemSpider22885 YesY
UNIIL0IYT1O31N YesY
ChEMBLCHEMBL2096650 N
RTECS numberXN6476000
Jmol-3D imagesImage 1
Properties
Molecular formulaNa2S2O3
Molar mass158.11 g/mol (anhydrous)
248.18 g/mol (pentahydrate)
AppearanceWhite crystals
OdorOdorless
Density1.667 g/cm3
Melting point48.3 °C (pentahydrate)
Boiling point100 °C (pentahydrate, - 5H2O decomposition)
Solubility in water70.1 g/100 mL (20 °C)[1]
231 g/100 mL (100 °C)
Solubilitynegligible in alcohol
Refractive index (nD)1.489
Structure
Crystal structuremonoclinic
Hazards
MSDSExternal MSDS
NFPA 704
NFPA 704.svg
0
1
0
Flash pointNon-flammable
 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references
 
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Sodium thiosulfate
Identifiers
CAS number7772-98-7 YesY, 10102-17-7 (pentahydrate) YesY
PubChem24477
ChemSpider22885 YesY
UNIIL0IYT1O31N YesY
ChEMBLCHEMBL2096650 N
RTECS numberXN6476000
Jmol-3D imagesImage 1
Properties
Molecular formulaNa2S2O3
Molar mass158.11 g/mol (anhydrous)
248.18 g/mol (pentahydrate)
AppearanceWhite crystals
OdorOdorless
Density1.667 g/cm3
Melting point48.3 °C (pentahydrate)
Boiling point100 °C (pentahydrate, - 5H2O decomposition)
Solubility in water70.1 g/100 mL (20 °C)[1]
231 g/100 mL (100 °C)
Solubilitynegligible in alcohol
Refractive index (nD)1.489
Structure
Crystal structuremonoclinic
Hazards
MSDSExternal MSDS
NFPA 704
NFPA 704.svg
0
1
0
Flash pointNon-flammable
 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references

Sodium thiosulfate (Na2S2O3), also spelled sodium thiosulphate, is a colorless crystalline compound that is more familiar as the pentahydrate, Na2S2O3·5H2O, an efflorescent, monoclinic crystalline substance also called sodium hyposulfite or “hypo.”

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 sulfur bears significant negative charge and the S-O interactions have more double bond character. The first protonation of thiosulfate occurs at sulfur.

Industrial production and laboratory synthesis[edit]

On an industrial scale, sodium thiosulfate is produced chiefly from liquid waste products of sodium sulfide or sulfur dye manufacture.[2]

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 with the following equation:[3]

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

Upon cooling sodium thiosulfate crystallizes out of solution.

Principal reactions and applications[edit]

Thiosulfate anions characteristically react with dilute acids to produce sulfur, sulfur dioxide and water:[2]

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. 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 dissociation respectively.

Iodometry[edit]

In analytical chemistry, the most important use comes from the fact that 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 the fact that 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]

The terminal sulfur atom in S2O32− binds to soft metals with high affinity. Thus, silver halides, e.g. AgBr, typical components of photographic emulsions, dissolve upon treatment with aqueous thiosulfate:

2 S2O32− + AgBr → [Ag(S2O3)2]3− + Br

In this application to photographic processing, discovered by John Herschel and 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.[4]

Gold extraction[edit]

Sodium thiosulfate is one component of an alternative lixiviant to cyanide for extraction of gold.[5] However, It forms a strong complex with gold(I) ions, [Au(S2O3)2]3−. The advantage of this approach is that thiosulfate is essentially non-toxic and 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.

Aluminium cation reaction[edit]

Sodium thiosulfate is also used in analytical chemistry. 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

Medical[edit]

Other uses[edit]

Sodium thiosulfate is also used:

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

References[edit]

  1. ^ Record in the GESTIS Substance Database from the IFA
  2. ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
  3. ^ Gordin, H. M. (1913). Elementary Chemistry, Volume I. Inorganic Chemistry. Chicago: Medico-Dental Publishing Co. pp. 162 & 287–288. 
  4. ^ Charles Robert Gibson (1908). The Romance of Modern Photography, Its Discovery & Its Achievements. Seeley & Co. p. 37. 
  5. ^ Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", Minerals Engineering, 2001, 14, 135-174
  6. ^ "Toxicity, Cyanide: Overview - eMedicine". Retrieved 2009-01-01. 
  7. ^ 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. 
  8. ^ 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. 
  9. ^ "Sodium thiosulfate" at Dorland's Medical Dictionary
  10. ^ http://www.ncbi.nlm.nih.gov/pmc/articles/PMC357669/?page=1
  11. ^ http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2217764/?page=1