Sodium thiosulfate

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Sodium thiosulfate
Sodium thiosulfate
Crystal structure of sodium thiosulfate pentahydrate
Sodium thiosulfate.jpg
Identifiers
CAS number 7772-98-7 YesY, 10102-17-7 (pentahydrate) YesY
PubChem 24477
ChemSpider 22885 YesY
UNII L0IYT1O31N YesY
ChEMBL CHEMBL2096650 N
RTECS number XN6476000
Jmol-3D images Image 1
Properties
Molecular formula 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)
Solubility in water 70.1 g/100 mL (20 °C)[1]
231 g/100 mL (100 °C)
Solubility negligible in alcohol
Refractive index (nD) 1.489
Structure
Crystal structure monoclinic
Hazards
MSDS External MSDS
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 N (verify) (what is: YesY/N?)
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”.[2]

Structure[edit]

Two polymorphs are known of the 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 sulfur bears significant negative charge and the S-O interactions have more double-bond character.

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.[3]

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:[4]

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

Principal reactions and applications[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:[3]

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.

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 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.[5]

Gold extraction[edit]

Sodium thiosulfate is a component of an alternative lixiviant to cyanide for extraction of gold.[6] However, it forms a strong soluble complex with gold(I) ions, [Au(S2O3)2]3−. The advantage of this approach is that thiosulfate is essentially not 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.[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 thiosulphate gives S-alkylthiosulonates, which are called Bunte salts. This reaction is employed in one synthesis of the industrial reagent thioglycolic acid:

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

Neutralizing bleach, chlorinated water, and related treatments[edit]

It is used to dechlorinate tap water for myriad purposes notable water treatment plants prior to release into rivers.[2] The reduction reaction is analogous to the iodine reduction reaction. Treatment of tap water lowers chlorine levels for use in aquaria and swimming pools and spas (e.g., following superchlorination)

In pH testing of bleach substances, sodium thiosulfate neutralize the color-removing effects of bleach and allow 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.

Other uses[edit]

Sodium thiosulfate is also used:

  • As a component in hand warmers and other chemical heating pads that produce heat by exothermic crystallization of a supercooled solution.
  • In bacteriological water assessment, as it promotes the survival of coliform organisms by neutralizing residual chlorine.[7][8]
  • In the tanning of leather
  • To demonstrate the concept of supercooling in physics classes: Sodium thiosulfate, when heated, dissolves in its own water of crystallisation. This solution can be cooled to room temperature without recrystallisation. When crystallisation is induced by the addition of a small seed crystal, the sudden temperature rise can be experienced by touch.
  • As part of patina recipes for copper alloys
  • Often used in pharmaceutical preparations as an anionic surfactant to aid in dispersion
  • As an ingredient to table salt, e.g. Sysco Corporation's small packets of iodized salt.

Medical[edit]

References[edit]

  1. ^ Record in the GESTIS Substance Database from the IFA
  2. ^ a b c J. J. Barbera, A. Metzger, M. Wolf "Sulfites, Thiosulfates, and Dithionites" in Ullmann's Encyclopedia of Industrial Chemistry 2012, Wiley-VCH, Weinheim. doi:10.1002/14356007.a25_477
  3. ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
  4. ^ Gordin, H. M. (1913). Elementary Chemistry, Volume I. Inorganic Chemistry. Chicago: Medico-Dental Publishing Co. pp. 162 & 287–288. 
  5. ^ Charles Robert Gibson (1908). The Romance of Modern Photography, Its Discovery & Its Achievements. Seeley & Co. p. 37. 
  6. ^ Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", Minerals Engineering, 2001, 14, 135-174
  7. ^ http://www.ncbi.nlm.nih.gov/pmc/articles/PMC357669/?page=1
  8. ^ http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2217764/?page=1
  9. ^ 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. 
  10. ^ http://www.medscape.com/viewarticle/762244
  11. ^ 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. 
  12. ^ "Toxicity, Cyanide: Overview - eMedicine". Retrieved 2009-01-01. 
  13. ^ 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. 
  14. ^ "Sodium thiosulfate" at Dorland's Medical Dictionary