Sodium thiosulfate
[CAS 7772-98-7]

Identification

• Classification: Inorganic chemical industry >> Inorganic salt >> Metal sulfides and sulfates
• Name: Sodium thiosulfate
• Synonyms: Sodium hyposulfite; Sodium subsulfite
• Molecular Formula: Na₂S₂O₃
• Molecular Weight: 158.10
• CAS Registry Number: 7772-98-7
• EC Number: 231-867-5
• SMILES: [O-]S(=O)(=S)[O-].[Na+].[Na+]

Properties

• Density: Density Data source g/mL (Expl.)
• Melting point: 48 °C (Decomposes) (Expl.)
• Boiling point: 100 °C (Expl.)
• Solubility: water: 70.0% w/w in 20 °C (Expl.)

Safety Data

• Hazard Symbols: GHS07 Warning
• Risk Statements: H302-H315-H319-H335
• Safety Statements: P261-P301+P312-P302+P352-P304+P340-P305+P351+P338

Hazard Classification

Hazard	Class	Category Code	Hazard Statement
Specific target organ toxicity - single exposure	STOT SE	3	H335
Eye irritation	Eye Irrit.	2	H319
Skin irritation	Skin Irrit.	2	H315
Acute toxicity	Acute Tox.	4	H302

chemBlink Chemical Story

Sodium thiosulfate, CAS 7772-98-7, is an inorganic salt commonly represented by the formula Na₂S₂O₃. It is the anhydrous form of a compound also widely encountered as sodium thiosulfate pentahydrate. Although sodium thiosulfate has applications ranging from photography to water treatment, one of the clearest demonstrations of its chemistry occurs when it is added to iodine: the characteristic iodine color disappears.

This apparently simple color change became the basis of one of classical analytical chemistry's most useful quantitative reactions.

Molecular iodine, I₂, produces a yellow-brown to brown color in aqueous systems, depending on concentration and the presence of iodide. Thiosulfate reacts rapidly with iodine according to the overall reaction:

2 S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2 I⁻

In this reaction, iodine is reduced to iodide while thiosulfate is oxidized to tetrathionate. Iodide is colorless in ordinary dilute solution, so the visible iodine color vanishes as the reaction proceeds.

The stoichiometry is especially useful. One mole of iodine consumes two moles of thiosulfate. If the concentration of a sodium thiosulfate solution is accurately known, measuring how much is required to consume iodine provides a quantitative way to determine the amount of iodine present.

This is the foundation of iodometric titration.

In many iodometric analyses, the substance being measured is not iodine itself. Instead, an oxidizing analyte is allowed to react with excess iodide, liberating an amount of iodine related stoichiometrically to the original analyte. The liberated iodine is then titrated with standardized sodium thiosulfate solution.

The analytical sequence can therefore be thought of as a chemical translation. An oxidizing substance that may be difficult to measure directly is first translated into a measurable quantity of iodine. Sodium thiosulfate then measures that iodine.

The method becomes even more visually striking when starch is used as an indicator. Iodine forms an intensely colored blue complex with starch. Near the end of a thiosulfate titration, only a small amount of iodine remains, and addition of starch produces a deep blue color that is much easier to detect than the faint yellow color of dilute iodine.

As the final iodine is consumed by thiosulfate, the blue color disappears.

That sudden transition from blue to colorless is one of the classic endpoints of analytical chemistry. What the observer sees is not a dye being bleached arbitrarily. It is the visible consequence of the last detectable iodine being converted to iodide.

Starch is normally added relatively near the endpoint rather than at the beginning of a strongly iodine-containing titration. A very stable iodine-starch association at high iodine concentration can make the endpoint less satisfactory. The timing of indicator addition is therefore part of obtaining a sharp and reliable analytical result.

Thiosulfate solutions themselves also require some care in quantitative work. Sodium thiosulfate solutions are generally standardized rather than assumed to have an exact concentration simply from preparation. Their stability can be affected by acidity, microorganisms, light, and storage conditions. Classical analytical chemistry therefore developed detailed procedures for preparing, storing, and standardizing thiosulfate titrants.

The behavior in acidic solution is particularly interesting because thiosulfate is not indefinitely stable in strong acid. Acidification can lead to decomposition with formation of sulfur and sulfur dioxide:

S₂O₃²⁻ + 2 H⁺ → S + SO₂ + H₂O

A clear solution can consequently become cloudy as finely divided elemental sulfur appears. This is another reminder that the thiosulfate ion occupies an unusual position in sulfur chemistry: its sulfur atoms are not chemically equivalent, and the ion can participate readily in oxidation-reduction transformations.

The same reducing ability that removes iodine also allows thiosulfate to react with chlorine and related oxidizing disinfectants. This is why sodium thiosulfate can be used for dechlorination.

Municipal drinking water is commonly disinfected with chlorine or chlorine-containing species because they suppress harmful microorganisms. That residual disinfectant is desirable while water is being distributed for human use, but it can be undesirable in some laboratory procedures, aquaculture systems, or environmental sampling.

Thiosulfate can consume oxidizing chlorine species and convert them into chloride while sulfur species are oxidized. In appropriately controlled applications, this provides a convenient chemical method for neutralizing residual chlorine.

Aquarium keepers may therefore encounter thiosulfate chemistry without ever performing an iodometric titration. Some dechlorination products use thiosulfate chemistry to neutralize free chlorine before tap water contacts aquatic organisms. Modern water supplies may also use chloramine, however, and treatment of chloraminated water involves additional considerations; a simple statement that thiosulfate alone solves every aquarium-water problem would be an oversimplification.

The connection between titration and dechlorination is nevertheless fundamental. In both cases, thiosulfate encounters an oxidizing substance and is itself oxidized while converting the other species into a more reduced form.

Photography provides another famous application, although the chemistry is somewhat different. Sodium thiosulfate dissolves unexposed silver halide by forming soluble silver-thiosulfate complexes. This allows residual light-sensitive silver salts to be removed from photographic film or paper after development, preventing continued darkening. The photographic name "hypo," historically used for thiosulfate fixer, became familiar to generations of photographers.

Thus one ion can perform useful chemistry in several different ways. It can act as a reducing agent toward iodine and chlorine, while also serving as a ligand capable of forming soluble complexes with silver ions. Both properties arise from the unusual bonding and reactivity of the thiosulfate ion.

Sodium thiosulfate therefore provides an excellent example of chemistry becoming visible. Add it to iodine and brown color disappears. Use starch and the analytical endpoint becomes an abrupt change from blue to colorless. Add acid under suitable conditions and a clear solution can become cloudy with elemental sulfur.

Each observation corresponds to a different chemical transformation.

Among these, the iodine reaction remains especially elegant because it turns color into measurement. The disappearance of the last trace of blue tells the chemist that a stoichiometric reaction has reached its endpoint.

Sometimes quantitative chemistry ends with a number on an instrument. In iodometric analysis, it can end with something much simpler: the moment a blue solution suddenly becomes colorless.

References

1. Harris, D. C. Quantitative Chemical Analysis. Iodometric and iodimetric titrations, thiosulfate standardization, and starch indicators.

2. Mendham, J.; Denney, R. C.; Barnes, J. D.; Thomas, M. J. K. Vogel's Textbook of Quantitative Chemical Analysis. Classical iodine-thiosulfate analytical methods.

3. PubChem. Sodium Thiosulfate, CAS 7772-98-7. Identity, properties, uses, and safety information.

4. Published analytical literature on iodine-thiosulfate reactions and standardization of sodium thiosulfate solutions.

5. Water-treatment literature on the use of thiosulfate for neutralization of residual chlorine and other oxidizing disinfectants.