Sodium thiosulfate
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| Sodium thiosulfate | |
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| General | |
| Systematic name | Sodium thiosulphate (Sodium thiosulphate) |
| Other names | Sodium hyposulfite Hyposulphite of soda |
| Molecular formula | Na2S2O3 |
| Molar mass | 158.09774 g/mol |
| Appearance | White crystals |
| CAS number | [7772-98-7] |
| Properties | |
| Density and phase | 1.667 g/cm³, solid |
| Solubility in water | Very Soluble |
| Melting point | 48.3 °C |
| Boiling point | N/A |
| Basicity (pKb) | N/A |
| Structure | |
| Coordination geometry |
tetrahedral anion |
| Crystal structure | ? |
| Dipole moment | ? D |
| Hazards | |
| MSDS | External MSDS |
| EU classification | Non-toxic. |
| R-phrases | R35 |
| S-phrases | S1/2, S26, S37/39, S45 |
| NFPA 704 |
0
1
0
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| Flash point | Non flammable |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data |
Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Sodium thiosulfate (Na2S2O3) (sometimes spelled 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.
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[edit] 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]
Small scale synthesis is by boiling an aqueous solution of sodium sulfite with sulfur.
As such, the anion S2O32− represents a water-soluble form of elemental sulfur.
[edit] Principal reactions and applications
Thiosulfate anion characteristically reacts with dilute acids to produce sulfur, sulfur dioxide and water[1]:
- S2O32−(aq) + 2H+ (aq) → S(s) + SO2(g) + H2O(l)
This reaction has been employed to generate colloidal sulfur. When the protonation is conducted at low temperatures, H2S2O3 (Hydrogen thiosulfate) can be obtained. It is a strong acid pKa = 0.6, 1.7.
[edit] Iodometry
Perhaps most notably in the laboratory, the thiosulfate anion reacts stoichiometrically with iodine, reducing it to iodide as it is oxidized to tetrathionate:
- 2S2O32−(aq) + I2(aq) → S4O62−(aq) + 2I−(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.
[edit] Photographic processing
The terminal sulfur atom in S2O32− binds to soft metals with high affinity. Thus it dissolves silver halides, e.g. AgBr, which is a component of photographic emulsions:
- 2 S2O32− + AgBr → [Ag(S2O3)2]3−) + Br-
In this application to photographic processing, discovered by John Herschel and used for both film and paper processing, sodium thiosulfate is known as a photographic fixer.
[edit] Gold extraction
Sodium thiosulfate is one component of an alternative lixiviant to cyanide for extraction of gold.[2] 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.
[edit] 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 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:
- 4NaClO + Na2S2O3 + 2NaOH → 4NaCl + 2Na2SO4 + 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.
- As an antidote to cyanide poisoning. Thiosulfate acts as a sulfur donor for the conversion for cyanide to thiocyanate, catalyzed by the enzyme rhodanese.
- 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) - As part of patina recipes for copper alloys.
- Often used in pharmaceutical preparations as an anionic surfactant to aid in dispersion.
[edit] References
- ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
- ^ Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", Minerals Engineering, 2001, 14, 135-174
