Potassium thiosulfate: Difference between revisions
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| OtherNames = {{Unbulleted list |
| OtherNames = {{Unbulleted list |
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| Dipotassium monothiosulfate |
| Dipotassium monothiosulfate |
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| ChEBI = |
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| ChemSpiderID |
| ChemSpiderID = 55421 |
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| EINECS = |
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| EC_number = 233-666-8 |
| EC_number = 233-666-8 |
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| Gmelin = |
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| SMILES = [O-]S(=O)(=S)[O-].[K+].[K+] |
| SMILES = [O-]S(=O)(=S)[O-].[K+].[K+] |
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| UNII = UK1TD58L5O |
| UNII = UK1TD58L5O |
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| StdInChI=1S/2K.H2O3S2/c;;1-5(2,3)4/h;;(H2,1,2,3,4)/q2*+1;/p-2 |
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| StdInChIKey = FGRVOLIFQGXPCT-UHFFFAOYSA-L |
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| Section2 = {{Chembox Properties |
| Section2 = {{Chembox Properties |
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| pKa = |
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| Solubility = 96.1 g/100 ml (0 °C)< |
| Solubility = 96.1 g/100 ml (0 °C)<br />155.4 g/100 ml (20 °C)<br />165 g/100 ml (25 °C)<br />175.7 g/100 ml (30 °C)<br />204.7 g/100 ml (40 °C)<br />215.2 g/100 ml (50 °C)<br />238.3 g/100 ml (60 °C)<br />255.2 g/100 ml (70 °C)<br />293.1 g/100 ml (80 °C)<br />312 g/100 ml (90 °C)<ref name="soln">{{cite book |author1=Atherton Seidell |title=Solubilities of inorganic and organic compounds c. 2 |date=1919 |publisher=D. Van Nostrand Company |page=568 |url=https://books.google.com/books?id=t4LSvgY7uIEC |language=English}}</ref> |
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| Section7 = {{Chembox Hazards |
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| Section9 = {{Chembox Related |
| Section9 = {{Chembox Related |
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| OtherAnions = |
| OtherAnions = [[Potassium sulfite]]; [[Potassium sulfate]] |
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| OtherCations = [[Sodium thiosulfate]] |
| OtherCations = [[Sodium thiosulfate]] |
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'''Potassium thiosulfate''' |
'''Potassium thiosulfate''' is an [[inorganic compound]] with the formula K<sub>2</sub>S<sub>2</sub>O<sub>3</sub>. This salt can form multiple hydrates, such as the monohydrate, dihydrate, and the pentahydrate, all of which are white or colorless solids.<ref name="soln" /> It is used as a [[fertilizer]]. |
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==Formation and reactions== |
==Formation and reactions== |
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{{main| |
{{main|Thiosulfate}} |
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Thiosulfate salts are produced by the reaction of [[sulfite]] ion with elemental sulfur, and by incomplete oxidation of [[sulfide]]s. Thiosulfates are stable in neutral or [[alkali]]ne solutions, but not in acidic solutions, due to [[disproportionation]] to sulfur dioxide and sulfur:<ref name="Ullmann">{{Ullmann| |
Thiosulfate salts are produced by the reaction of [[sulfite]] ion with elemental sulfur, and by incomplete oxidation of [[sulfide]]s. For example, this salt is produced by reacting [[potassium hydroxide]] with [[ammonium hydroxide]], [[sulfur dioxide]], and elemental [[sulfur]].<ref name=":)" /> Thiosulfates are stable in neutral or [[alkali]]ne solutions, but not in acidic solutions, due to [[disproportionation]] to sulfur dioxide and sulfur:<ref name="Ullmann">{{Ullmann|author=Barbera, J. J. |author2=Metzger, A. |author3=Wolf, M. |title=Sulfites, Thiosulfates, and Dithionites|doi=10.1002/14356007.a25_477|isbn=978-3-527-30673-2|date=2012}}</ref> |
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:{{chem|S|2|O|3|2−}} + 2 H<sup>+</sup> → SO<sub>2</sub> + "S" + H<sub>2</sub>O |
:{{chem|S|2|O|3|2−}} + 2 H<sup>+</sup> → SO<sub>2</sub> + "S" + H<sub>2</sub>O |
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Due to this property, it can [[chelation|sequester]] metals, especially [[iron]].<ref name=":)" /> |
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Thiosulfate reacts with iodine to give [[tetrathionate]]: |
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Thiosulfate reacts with iodine to give [[tetrathionate]], in this case potassium thiosulfate reacts with iodine to produce potassium tetrathionate and [[potassium iodide]]: |
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Thiosulfate extensively forms diverse complexes with [[transition metal]]s. In the era of silver-based photography, thiosulfate was consumed on a large scale as a "stop" reagent. This application exploits thiosulfate's ability to dissolve silver halides. Thiosulfate is also used to extract or leach gold (sodium thiosulfate) and silver from their ores as a less toxic alternative to cyanide.<ref name="Ullmann">{{Ullmann| |
Thiosulfate extensively forms diverse complexes with [[transition metal]]s. In the era of silver-based photography, thiosulfate was consumed on a large scale as a "stop" reagent. This application exploits thiosulfate's ability to dissolve silver halides. Thiosulfate is also used to extract or leach gold (sodium thiosulfate) and silver from their ores as a less toxic alternative to cyanide.<ref name="Ullmann">{{Ullmann|author=Barbera, J. J. |author2=Metzger, A. |author3=Wolf, M. |title=Sulfites, Thiosulfates, and Dithionites|doi=10.1002/14356007.a25_477|isbn=978-3-527-30673-2|date=2012}}</ref> |
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==Uses== |
==Uses== |
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Potassium thiosulfate is commonly used as a fertilizer alone or with [[urea]] and/or [[urea ammonium nitrate]]<ref name="fert">{{cite book |author1=Western Plant Health Association |editor1-last=Barlow |editor1-first=Dave |editor2-last=Pier |editor2-first=Jerome |title=Western Fertilizer Handbook |date=2018 |publisher=Waveland Press |isbn= |
Potassium thiosulfate is commonly used as a fertilizer alone or with [[urea]] and/or [[urea ammonium nitrate]]<ref name="fert">{{cite book |author1=Western Plant Health Association |editor1-last=Barlow |editor1-first=Dave |editor2-last=Pier |editor2-first=Jerome |title=Western Fertilizer Handbook |date=2018 |publisher=Waveland Press |isbn=978-1-4786-3884-1 |page=186 |edition=Third Horticulture |url=https://books.google.com/books?id=BHeCDwAAQBAJ |access-date=6 October 2021 |language=English |format=Ebook}}</ref> due to its ability to delay [[nitrification]].<ref name=":)">{{cite book |editor1-last=Sulewski |editor1-first=Gavin |editor2-last=Thompson |editor2-first=Michael |editor3-last=Mikkelsen |editor3-first=Robert |editor4-last=Norton |editor4-first=Robert |editor5-last=Scott |editor5-first=T. |title=Improving Potassium Recommendations for Agricultural Crops |date=2020 |publisher=Springer International Publishing |isbn=978-3-030-59197-7 |page=60 |url=https://books.google.com/books?id=pfsOEAAAQBAJ |access-date=6 October 2021 |language=English |format=Ebook}}</ref> It thus has the ability to reduce the emission of [[nitrous oxide]].<ref>{{cite journal |author1=Zejiang Cai |author2=Suduan Gao |author3=Minggang Xu |author4=Bradley D Hanson |title=Evaluation of potassium thiosulfate as a nitrification inhibitor to reduce nitrous oxide emissions |journal=Science of the Total Environment |date=2017 |volume=618 |pages=243–249 |doi=10.1016/j.scitotenv.2017.10.274 |pmid=29128773 |language=English|doi-access=free }}</ref> It can also reduce the amount of [[fumigants]] being released from the soil.<ref>{{cite journal |author1=Ruijun Qin |author2=Suduan Gao |author3=Jason A McDonald |author4=Husein Ajwa |author5=Shachar Shem-Tov |author6=David A Sullivan |title=Effect of plastic tarps over raised-beds and potassium thiosulfate in furrows on chloropicrin emissions from drip fumigated fields |journal=Chemosphere |date=2008 |volume=72 |issue=4 |pages=558–563 |doi=10.1016/j.chemosphere.2008.03.023 |pmid=18440581 |bibcode=2008Chmsp..72..558Q |language=English}}</ref> If used alone it is used in very dilute solution due to its ability to cause [[phytotoxicity]] symptoms. This is caused by the elemental sulfur being oxidized to produce [[sulfuric acid]].<ref name="fert" /> |
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==References== |
==References== |
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Latest revision as of 06:16, 20 November 2024
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.030.593 |
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PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |
| K2S2O3 | |
| Molar mass | 190.32 g/mol |
| Appearance | White solid |
| Density | 2.37 g/cm3 |
| 96.1 g/100 ml (0 °C) 155.4 g/100 ml (20 °C) 165 g/100 ml (25 °C) 175.7 g/100 ml (30 °C) 204.7 g/100 ml (40 °C) 215.2 g/100 ml (50 °C) 238.3 g/100 ml (60 °C) 255.2 g/100 ml (70 °C) 293.1 g/100 ml (80 °C) 312 g/100 ml (90 °C)[1] | |
| Hazards | |
| GHS labelling: | |
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| Warning | |
| H315, H319 | |
| P264, P280, P302+P352, P305+P351+P338, P321, P332+P313, P337+P313, P362 | |
| Related compounds | |
Other anions
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Potassium sulfite; Potassium sulfate |
Other cations
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Sodium thiosulfate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Potassium thiosulfate is an inorganic compound with the formula K2S2O3. This salt can form multiple hydrates, such as the monohydrate, dihydrate, and the pentahydrate, all of which are white or colorless solids.[1] It is used as a fertilizer.
Formation and reactions
[edit]Thiosulfate salts are produced by the reaction of sulfite ion with elemental sulfur, and by incomplete oxidation of sulfides. For example, this salt is produced by reacting potassium hydroxide with ammonium hydroxide, sulfur dioxide, and elemental sulfur.[2] Thiosulfates are stable in neutral or alkaline solutions, but not in acidic solutions, due to disproportionation to sulfur dioxide and sulfur:[3]
- S
2O2−
3 + 2 H+ → SO2 + "S" + H2O
Due to this property, it can sequester metals, especially iron.[2]
Thiosulfate reacts with iodine to give tetrathionate, in this case potassium thiosulfate reacts with iodine to produce potassium tetrathionate and potassium iodide:
- 2 K2S
2O
3 + I2 → K2S
4O
6 + 2 KI
Thiosulfate extensively forms diverse complexes with transition metals. In the era of silver-based photography, thiosulfate was consumed on a large scale as a "stop" reagent. This application exploits thiosulfate's ability to dissolve silver halides. Thiosulfate is also used to extract or leach gold (sodium thiosulfate) and silver from their ores as a less toxic alternative to cyanide.[3]
Uses
[edit]Potassium thiosulfate is commonly used as a fertilizer alone or with urea and/or urea ammonium nitrate[4] due to its ability to delay nitrification.[2] It thus has the ability to reduce the emission of nitrous oxide.[5] It can also reduce the amount of fumigants being released from the soil.[6] If used alone it is used in very dilute solution due to its ability to cause phytotoxicity symptoms. This is caused by the elemental sulfur being oxidized to produce sulfuric acid.[4]
References
[edit]- ^ a b Atherton Seidell (1919). Solubilities of inorganic and organic compounds c. 2. D. Van Nostrand Company. p. 568.
- ^ a b c Sulewski, Gavin; Thompson, Michael; Mikkelsen, Robert; Norton, Robert; Scott, T., eds. (2020). Improving Potassium Recommendations for Agricultural Crops (Ebook). Springer International Publishing. p. 60. ISBN 978-3-030-59197-7. Retrieved 6 October 2021.
- ^ a b Barbera, J. J.; Metzger, A.; Wolf, M. (2012). "Sulfites, Thiosulfates, and Dithionites". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_477. ISBN 978-3-527-30673-2.
- ^ a b Western Plant Health Association (2018). Barlow, Dave; Pier, Jerome (eds.). Western Fertilizer Handbook (Ebook) (Third Horticulture ed.). Waveland Press. p. 186. ISBN 978-1-4786-3884-1. Retrieved 6 October 2021.
- ^ Zejiang Cai; Suduan Gao; Minggang Xu; Bradley D Hanson (2017). "Evaluation of potassium thiosulfate as a nitrification inhibitor to reduce nitrous oxide emissions". Science of the Total Environment. 618: 243–249. doi:10.1016/j.scitotenv.2017.10.274. PMID 29128773.
- ^ Ruijun Qin; Suduan Gao; Jason A McDonald; Husein Ajwa; Shachar Shem-Tov; David A Sullivan (2008). "Effect of plastic tarps over raised-beds and potassium thiosulfate in furrows on chloropicrin emissions from drip fumigated fields". Chemosphere. 72 (4): 558–563. Bibcode:2008Chmsp..72..558Q. doi:10.1016/j.chemosphere.2008.03.023. PMID 18440581.