Difference between revisions of "Thiosulfate"
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==Properties== | ==Properties== | ||
| − | The thiosulfate ion has a charge of -2. [[Sodium thiosulfate]], the primary source of the ion, has an interesting | + | The thiosulfate ion has a charge of -2. [[Sodium thiosulfate]], the primary source of the ion, has an interesting hexagonal crystal structure. Because of its three oxygen molecules, the thiosulfate ion can complex, forming a thiosulfato complex. Thiosulfate isn't stable in acidic conditions, only neutral and basic. [[Ammonium thiosulfate]] and potassium thiosulfate are also sometimes encountered. |
Thiosulfuric acid does not exist at temperatures obtainable by the amateur chemist. When protonated, the ion forms sulfur and sulfur dioxide. The sulfur produced forms a colloidial suspension with the water, and can be used to demonstrate Rayleigh scattering. | Thiosulfuric acid does not exist at temperatures obtainable by the amateur chemist. When protonated, the ion forms sulfur and sulfur dioxide. The sulfur produced forms a colloidial suspension with the water, and can be used to demonstrate Rayleigh scattering. | ||
| − | Thiosulfates are capable of reducing halogens to halides. When this occurs, the thiosulfate is oxidized to | + | Thiosulfates are capable of reducing halogens to halides. When this occurs, the thiosulfate is oxidized to tetrathionate (S<sub>4</sub>O<sub>6</sub><sup>2-</sup>) in the presence of iodine, or [[sulfate]] in the presence of bromine or chlorine. This property is used in [[iodometric titration]]s to reduce iodine (in this reaction the endpoint is usually indicated using a [[starch]]-[[iodine]] [[Coordination complex|complex]]).<ref>G. Jander, K. Jahr, ''Maßanalyse: Theorie und Praxis der Titrationen mit chemischen und physikalischen Indikationen.'' Berlin, New York, De Gruyterm, '''2009''', 17th edition, p. 179ff. [https://doi.org/10.1515/9783110211337 https://doi.org/10.1515/9783110211337]</ref> |
==Preparation== | ==Preparation== | ||
| − | Thiosulfates can be prepared by heating a concentrated solution containing six moles of an alkali metal[[ hydroxide]] with four moles of sulfur powder, then cooling it down and let the thiosulfate fall out of solution. [[Sodium hydroxide]] is preferred for this, as [[potassium hydroxide]] forms [[potassium thiosulfate]], which is extremely [[Hygroscopy|hygroscopic]]. | + | Thiosulfates can be prepared by heating a concentrated solution containing six moles of an alkali metal[[ hydroxide]] with four moles of [[sulfur]] powder, then cooling it down and let the thiosulfate fall out of solution. [[Sodium hydroxide]] is preferred for this, as [[potassium hydroxide]] forms [[potassium thiosulfate]], which is extremely [[Hygroscopy|hygroscopic]]. |
==Projects== | ==Projects== | ||
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* [[Titration|Redox titrations]] | * [[Titration|Redox titrations]] | ||
* Removing iodine or bromine stains | * Removing iodine or bromine stains | ||
| + | |||
| + | ==Safety== | ||
| + | Thiosulfates have low toxicity. | ||
==References== | ==References== | ||
Latest revision as of 19:06, 13 August 2023
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The thiosulfate ion is a sulfur oxoanion with the formula S2O32-. It is derived from sulfate, but a sulfur atom replaces an oxygen atom within the structure.
Contents
Properties
The thiosulfate ion has a charge of -2. Sodium thiosulfate, the primary source of the ion, has an interesting hexagonal crystal structure. Because of its three oxygen molecules, the thiosulfate ion can complex, forming a thiosulfato complex. Thiosulfate isn't stable in acidic conditions, only neutral and basic. Ammonium thiosulfate and potassium thiosulfate are also sometimes encountered.
Thiosulfuric acid does not exist at temperatures obtainable by the amateur chemist. When protonated, the ion forms sulfur and sulfur dioxide. The sulfur produced forms a colloidial suspension with the water, and can be used to demonstrate Rayleigh scattering.
Thiosulfates are capable of reducing halogens to halides. When this occurs, the thiosulfate is oxidized to tetrathionate (S4O62-) in the presence of iodine, or sulfate in the presence of bromine or chlorine. This property is used in iodometric titrations to reduce iodine (in this reaction the endpoint is usually indicated using a starch-iodine complex).[1]
Preparation
Thiosulfates can be prepared by heating a concentrated solution containing six moles of an alkali metal hydroxide with four moles of sulfur powder, then cooling it down and let the thiosulfate fall out of solution. Sodium hydroxide is preferred for this, as potassium hydroxide forms potassium thiosulfate, which is extremely hygroscopic.
Projects
- Iodine clock reaction
- Redox titrations
- Removing iodine or bromine stains
Safety
Thiosulfates have low toxicity.
References
- ↑ G. Jander, K. Jahr, Maßanalyse: Theorie und Praxis der Titrationen mit chemischen und physikalischen Indikationen. Berlin, New York, De Gruyterm, 2009, 17th edition, p. 179ff. https://doi.org/10.1515/9783110211337
