View source for Sodium thiosulfate

{{chembox
| Watchedfields = changed
| verifiedrevid = 265548656
|   Name = Sodium thiosulfate
|   ImageFile = Sodium-thiosulfate-3D-vdW.png
<!-- | ImageSize = 200px -->
|   ImageName = Sodium thiosulfate
|   ImageFile1 = Sodium-thiosulfate-xtal-3D-balls.png
|   ImageName1 = Crystal structure of sodium thiosulfate pentahydrate
|   IUPACName = Sodium thiosulfate
|   OtherNames = Sodium hyposulfite<br/>Hyposulphite of soda
| Section1 = {{Chembox Identifiers
|   CASNo = 7772-98-7
|    CASNo_Ref = {{cascite}}
|   CASOther = <br/>10102-17-7 (pentahydrate)
|   PubChem = 24477
|   RTECS = XN6476000
  }}
| Section2 = {{Chembox Properties
|   Formula = Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub>
|   MolarMass = 158.108 g/mol
|   Appearance = white crystals
|   Odor = odorless
|   Density = 1.667 g/cm<sup>3</sup>
|   Solubility = 20.9 g/100 ml (20 °C)
|   MeltingPt = 48.3 °C (pentahydrate)
|   BoilingPt = 100 °C (pentahydrate, - 5H<sub>2</sub>O decomp)
|   pKb = 
  }}
| Section3 = {{Chembox Structure
|   Coordination = 
|   CrystalStruct = 
|   Dipole =
  }}
| Section7 = {{Chembox Hazards
|   ExternalMSDS = [http://www.jtbaker.com/msds/englishhtml/S5230.htm External MSDS]
|   EUIndex = Not listed
|   RPhrases = 
|   SPhrases = 
|   NFPA-H = 1
|   NFPA-F = 0
|   NFPA-R = 0
|   FlashPt = Non-flammable
  }}
}}

'''Sodium thiosulfate''' ([[Sodium|Na]]<sub>2</sub>[[Sulfur|S]]<sub>2</sub>[[Oxygen|O]]<sub>3</sub>), also spelled '''sodium thiosulphate''', is a colorless [[crystal]]line compound that is more familiar as the [[pentahydrate]], [[Sodium|Na]]<sub>2</sub>[[Sulfur|S]]<sub>2</sub>[[Oxygen|O]]<sub>3</sub>•5[[Hydrogen|H]]<sub>2</sub>[[Oxygen|O]], 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==

On an industrial scale, sodium thiosulfate is produced chiefly from liquid waste products of [[sodium sulfide]] or sulfur [[dye]] manufacture.<ref name="holleman">Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5</ref>  In the laboratory, this salt can be prepared by heating an aqueous solution of [[sodium sulfite]] with [[sulfur]].

==Principal reactions and applications==
[[File:Thiosíran sodný.JPG|thumb|left|180px|Sodium thiosulfate]]
Thiosulfate anion characteristically reacts with [[dilution|dilute]] [[acid]]s to produce sulfur, [[sulfur dioxide]] and [[water]]:<ref name="holleman" />
:Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub> + 2HCl → 2NaCl + S + SO<sub>2</sub> + H<sub>2</sub>O
This reaction is known as a "[[Chemical clock|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 at low temperatures, H<sub>2</sub>S<sub>2</sub>O<sub>3</sub> ([[thiosulfuric acid]]) can be obtained. It is a somewhat strong acid with pK<sub>a</sub>s of 0.6 and 1.7 for the first and second dissociation respectively. 

===Iodometry===
In analytical chemistry, the most important use comes from the fact that the thiosulfate anion reacts [[stoichiometrically]] with [[iodine]], reducing it to [[iodide]] as it is oxidized to [[tetrathionate]]:
:2 S<sub>2</sub>O<sub>3</sub><sup>2−</sup><small>(aq)</small> + [[iodine|I]]<sub>2</sub><small>(aq)</small> → S<sub>4</sub>O<sub>6</sub><sup>2−</sup><small>(aq)</small> + 2 I<sup>−</sup><small>(aq)</small>

Due to the quantitative nature of this reaction, as well as the fact that Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub>•5H<sub>2</sub>O has an excellent shelf-life, it is used as a [[titrant]] in [[iodometry]].  Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub>•5H<sub>2</sub>O is also a component of [[iodine clock reaction|''iodine clock'' experiments]].

This particular use can be set up to measure the oxygen content of water through a long series of reactions. 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===

The terminal sulfur atom in S<sub>2</sub>O<sub>3</sub><sup>2−</sup> binds to [[HSAB concept|soft metals]] with high [[Chemical affinity|affinity]]. Thus, [[silver]] halides, e.g. [[silver bromide|AgBr]], typical components of photographic emulsions, dissolve upon treatment with aqueous thiosulfate:
:2 S<sub>2</sub>O<sub>3</sub><sup>2−</sup>  +  AgBr  →  [Ag(S<sub>2</sub>O<sub>3</sub>)<sub>2</sub>]<sup>3−</sup>  +  Br<sup>-</sup>

In this application to [[photographic processing]], discovered by [[John Herschel]] and used for both [[photographic film|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.''<ref>{{cite book | title = The Romance of Modern Photography, Its Discovery & Its Achievements | author = Charles Robert Gibson | publisher = Seeley & Co | year = 1908 | page = 37 | url = http://books.google.com/?id=whYaAAAAMAAJ&pg=PA37&dq=hyposulphite-of-soda+herschel+fixer+hypo }}</ref>

===Gold extraction===
Sodium thiosulfate is one component of an alternative [[lixiviant]] to [[cyanide]] for extraction of [[gold]].<ref name="aylmore">Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", ''Minerals Engineering'', 2001, 14, 135-174</ref>  It forms a strong complex with gold(I) ions, [Au(S<sub>2</sub>O<sub>3</sub>)<sub>2</sub>]<sup>3-</sup>.  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 gold deposit|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(S<sub>2</sub>O<sub>3</sub>)<sub>2</sub>]<sup>3-</sup> does not [[adsorb]] to [[activated carbon]], which is the standard technique used in [[gold cyanidation]] to separate the gold complex from the ore slurry.

===Analytical chemistry ===
Sodium thiosulfate is also used in analytical chemistry. It can, when heated with a sample containing aluminum cation, produce a white precipitation:
:2 Al<sup>3+</sup> + 3 S<sub>2</sub>O<sub>3</sub><sup>2-</sup> + 3 H<sub>2</sub>O → 3 SO<sub>2</sub> + 3 S + 2 Al(OH)<sub>3</sub>

==Medical==

* It is used as an antidote to [[cyanide poisoning]].<ref name="urlToxicity, Cyanide: Overview - eMedicine">{{cite web |url=http://emedicine.medscape.com/article/814287-overview |title=Toxicity, Cyanide: Overview - eMedicine |format= |work= |accessdate=2009-01-01}}</ref><ref name="pmid17098327">{{cite journal |author=Hall AH, Dart R, Bogdan G |title=Sodium thiosulfate or hydroxocobalamin for the empiric treatment of cyanide poisoning? |journal=Ann Emerg Med |volume=49 |issue=6 |pages=806–13 |year=2007 |month=June |pmid=17098327 |doi=10.1016/j.annemergmed.2006.09.021 }}</ref>  Thiosulfate acts as a sulfur donor for the conversion for cyanide to thiocyanate (which can then be safely excreted in the urine), catalyzed by the enzyme [[rhodanase]].

* It has also been used as treatment of [[calciphylaxis]] in [[hemodialysis]] patients with [[end-stage renal disease]].<ref>{{cite journal |author=Cicone JS, Petronis JB, Embert CD, Spector DA |title=Successful treatment of calciphylaxis with intravenous sodium thiosulfate |journal=Am. J. Kidney Dis. |volume=43 |issue=6 |pages=1104–8 |year=2004 |month=June |pmid=15168392 |doi=10.1053/j.ajkd.2004.03.018}}</ref>

* It is used in the management of [[Extravasation of urine|extravasation]]s during [[chemotherapy]].  Sodium thiosulfate prevents alkylation and tissue destruction by providing a substrate for the alkylating agents that have invaded the subcutaneous tisuses.  The dose may be 2mL of 0.17M (a solution of 4mL 10% sodium thiosulfate and 6mL sterile water for injection).  It may be instilled subcutaneously into multiple sites using a small gauge needle.  There is limited data on this method with few recommendations.{{Citation needed|date=June 2010}}

* in foot baths for prophylaxis of [[ringworm]], and as a topical [[antifungal]] agent for [[tinea versicolor]].
* in measuring the volume of [[extracellular fluid|extracellular body fluid]] and the [[renal]] [[glomerular filtration rate]]. <ref>{{DorlandsDict|seven/000098060|Sodium thiosulfate}}</ref>

==Other uses==

Sodium thiosulfate is also used:

*As a component in [[hand warmer]]s and other chemical [[heating pad]]s that produce heat by exothermic crystallization of a [[supercooling|supercooled]] solution.
*In Bleach
*In pH testing of bleach substances. The [[universal indicator]] and any other liquid pH indicator are destroyed by bleach, rendering them useless for testing the [[pH]]. If one first adds sodium thiosulfate to such solutions, it will 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  +  Na<sub>2</sub>S<sub>2</sub>O<sub>3</sub>  +  2 NaOH  →  4 NaCl  +  2 Na<sub>2</sub>SO<sub>4</sub>  +  H<sub>2</sub>O

*To dechlorinate tap water for aquariums or treat effluent from waste water treatments prior to release into rivers.  The reduction reaction is analogous to the iodine reduction reaction.  Treatment of tap water requires between 0.1 grams and 0.3 grams of pentahydrated (crystalline) sodium thiosulfate per 10 liters of water.
*To lower chlorine levels in swimming pools and spas following super chlorination.
*To remove iodine stains, e.g. after the explosion of [[nitrogen triiodide]].
*Similarly, sodium thiosulfate reacts with [[bromine]] to render harmless products. Solutions of sodium thiosulfate are commonly used as a precaution in chemistry laboratories when working with bromine.
*In [[bacteria|bacteriological]] water assessment.
*In the [[tanning]] of [[leather]].
*To demonstrate the concept of [[reaction rate]] in chemistry classes. The [[thiosulfate]] ion can decompose into the [[sulfite]] ion and a colloidal suspension of sulfur, which is opaque.  The equation for this acid-catalysed reaction is as follows:<br />S<sub>2</sub>O<sub>3</sub><sup>2−</sup>(aq) → SO<sub>3</sub><sup>2−</sup>(aq) + S(s)
*To demonstrate the concept of [[supercooling]] in physics classes. Melted sodium thiosulfate is very easy to overcool to room temperature and when crystallization is forced, the sudden temperature jump to 48.3°C 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.
*It can also be used as a very interesting solute in supersaturation experiments.
*An experiment performed in schools that shows that when mixed with [[Hydrochloric acid]], the solution will go cloudy.

==References==
{{reflist|2}}

{{CommonsCat}}

{{Sodium compounds}}
{{Antifungals}}
{{Antidotes}}

[[Category:Thiosulfates]]
[[Category:Sodium compounds]]
[[Category:Photographic chemicals]]
[[Category:Antidotes]]
[[Category:World Health Organization essential medicines]]

[[ar:ثيوكبريتات الصوديوم]]
[[bs:Natrijum tiosulfat]]
[[ca:Tiosulfat de sodi]]
[[cs:Thiosíran sodný]]
[[da:Natriumthiosulfat]]
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[[fr:Thiosulfate de sodium]]
[[it:Tiosolfato di sodio]]
[[he:נתרן תיוסולפט]]
[[hu:Nátrium-tioszulfát]]
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[[ja:チオ硫酸ナトリウム]]
[[pl:Tiosiarczan sodu]]
[[pt:Tiossulfato de sódio]]
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[[ru:Тиосульфат натрия]]
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[[Category:2Fix]]