Sodium thiosulfate

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Sodium thiosulfate
Image:Sodium-thiosulfate-3D-vdW.png Image:Sodium-thiosulfate.png
General
Systematic name	Sodium thiosulfate (Sodium thiosulphate)
Other names	Sodium hyposulfite Hyposulfate 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
Solubility in water	Very Soluble
Melting point	48.3 °C
Boiling point	N/A
Basicity (pK_b)	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	Image:NFPA 704.svg 0 1 0
Flash point	Non flammable
Supplementary data page
Structure and properties	n, ε_r, 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

Sodium thiosulfate (Na₂S₂O₃) 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.

1 Production and chemical synthesis<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5</ref>
2 Principal reactions and applications<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5</ref>
- 2.1 Iodometry
- 2.2 Photographic processing
3 Other uses
4 References
5 External links

[edit] Production and chemical synthesis<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5</ref>

Sodium thiosulfate is produced chiefly from liquid waste products of sodium sulfide or sulfur dye manufacture. It is also arises from the action of sodium sulfite on sulfur in aqueous solution. As such, the anion S₂O₃²⁻ represents a water-soluble form of elemental sulfur.

[edit] Principal reactions and applications<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5</ref>

Thiosulfate anion characteristically reacts with dilute acids to produce sulfur, sulfur dioxide and water:

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 can be obtained. It is a strong acid pK_a = 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.

[edit] Photographic processing

The "planetary" sulfur atom in S₂O₃^2- binds to "soft" metals with high affinity. Thus it dissolves silver halides, e.g. AgBr, which is a component of photographic emulsions:

2 S₂O₃^2- + AgBr → [Ag(S₂O₃)₂]^3-) + Br^-

In this application, discovered by John Herschel and used for both film and paper processing, sodium thiosulfate is known as a Photographic fixer.

[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 (crystaline) 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)