Sodium thiosulfate

Sodium thiosulfate
IUPAC name Sodium thiosulfate
Other names Sodium hyposulfite Hyposulphite of soda
Identifiers
CAS number	7772-98-7 Y,  10102-17-7 (pentahydrate)
PubChem	24477
ChemSpider	22885 Y
UNII	L0IYT1O31N Y
ChEMBL	CHEMBL1201157 N
RTECS number	XN6476000
Jmol-3D images	Image 1
SMILES [Na+].[Na+].[O-]S([O-])(=O)=S
InChI InChI=1S/2Na.H2O3S2/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 Y Key: AKHNMLFCWUSKQB-UHFFFAOYSA-L Y InChI=1/2Na.H2O3S2/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 Key: AKHNMLFCWUSKQB-NUQVWONBAM
Properties
Molecular formula	Na2S2O3
Molar mass	158.11 g/mol
Appearance	white crystals
Odor	odorless
Density	1.667 g/cm3
Melting point	48.3 °C (pentahydrate)
Boiling point	100 °C (pentahydrate, - 5H2O decomp)
Solubility in water	76.4 g/100 g H20(20 °C)
Hazards
MSDS	External MSDS
EU Index	Not listed
NFPA 704	0 1 0
Flash point	Non-flammable
N (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 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

On an industrial scale, sodium thiosulfate is produced chiefly from liquid waste products of sodium sulfide or sulfur dye manufacture.^[1] In the laboratory, this salt can be prepared by heating an aqueous solution of sodium sulfite with sulfur.

Principal reactions and applications

Thiosulfate anion characteristically reacts with dilute acids to produce sulfur, sulfur dioxide and water:^[1]

Na₂S₂O₃ + 2HCl → 2NaCl + 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 at low temperatures, H₂S₂O₃ (thiosulfuric acid) can be obtained. It is a somewhat strong acid with pK_as 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 S₂O₃²⁻(aq) + I₂(aq) → S₄O₆²⁻(aq) + 2 I⁻(aq)

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. 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 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.^[2]

Gold extraction

Sodium thiosulfate is one component of an alternative lixiviant to cyanide for extraction of gold.^[3] 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.

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

Medical

• It is used as an antidote to cyanide poisoning.^[4][5] Thiosulfate acts as a sulfur donor for the conversion of 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.^[6]

• It is used in the management of extravasations during chemotherapy. Sodium thiosulfate prevents alkylation and tissue destruction by providing a substrate for the alkylating agents that have invaded the subcutaneous tissues. 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]}

• in foot baths for prophylaxis of ringworm, and as a topical antifungal agent for tinea versicolor.
• in measuring the volume of extracellular body fluid and the renal glomerular filtration rate. ^[7]

Other uses

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 bleach
• In pH testing of bleach substances. The universal indicator and any other liquid pH indicators 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 + Na₂S₂O₃ + 2 NaOH → 4 NaCl + 2 Na₂SO₄ + H₂O

• 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 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:
  S₂O₃²⁻(aq) → SO₃²⁻(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.

References

1. ^ ^{a b} Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
2. ^ Charles Robert Gibson (1908). The Romance of Modern Photography, Its Discovery & Its Achievements. Seeley & Co. p. 37. http://books.google.com/?id=whYaAAAAMAAJ&pg=PA37&dq=hyposulphite-of-soda+herschel+fixer+hypo. 
3. ^ Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", Minerals Engineering, 2001, 14, 135-174
4. ^ "Toxicity, Cyanide: Overview - eMedicine". http://emedicine.medscape.com/article/814287-overview. Retrieved 2009-01-01. 
5. ^ 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. 
6. ^ 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. 
7. ^ "Sodium thiosulfate" at Dorland's Medical Dictionary