sodium thiosulfate

Sodium thiosulfate
IUPAC name	Sodium thiosulfate
Other names	Sodium hyposulfite Hyposulphite of soda
Identifiers
CAS number	7772-98-7
Properties
Molecular formula	Na2S2O3
Molar mass	158.09774 g/mol
Appearance	White crystals
Density	1.667 g/cm³, solid
Melting point	48.3 °C
Boiling point	N/A
Solubility in water	Very Soluble
Basicity (pKb)	N/A
Structure
Coordination geometry	Tetrahedral anion
Hazards
MSDS	External MSDS
EU classification	Non-toxic.
NFPA 704	0 1 0
R-phrases	R35
S-phrases	(S1/2) S26 S37/39 S45
Flash point	Non flammable
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) (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.

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.

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 (thiosulfuric acid) can be obtained. It is a strong acid pK_a = 0.6, 1.7.

Iodometry

Perhaps most notably in the laboratory^{[citation needed]}, 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 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.

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.

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 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 + 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.
• As an antidote to cyanide poisoning. Thiosulfate acts as a sulfur donor for the conversion for cyanide to thiocyanate, catalyzed by the enzyme rhodanase.
• 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.

References

1. ^ ^{a b} Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5
2. ^ Aylmore, M. G.; Muir, D. M. "Thiosulfate Leaching of Gold - a Review", Minerals Engineering, 2001, 14, 135-174

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