Disulfur decafluoride
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| Names | |||
|---|---|---|---|
| Preferred IUPAC name
Disulfur decafluoride | |||
| Systematic IUPAC name
Decafluoro-1λ6,2λ6-disulfane | |||
| Identifiers | |||
3D model (JSmol)
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| ChemSpider | |||
| ECHA InfoCard | 100.024.732 | ||
| EC Number |
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| MeSH | Disulfur+decafluoride | ||
PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |||
| S2F10 | |||
| Appearance | colorless liquid | ||
| Melting point | -53 °C | ||
| Boiling point | 30.1 °C | ||
| Hazards | |||
| NFPA 704 (fire diamond) | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Disulfur decafluoride (S2F10) is a gas discovered in 1934 by Denbigh and Whytlaw-Gray.[1] Each S of the S2F10 molecule is octahedral, and surrounded by 5 fluorines.[2] S2F10 is highly toxic, with toxicity similar to phosgene. It was considered a potential chemical warfare pulmonary agent in World War II because it does not produce lacrimation or skin irritation, thus providing little warning of exposure. It is a possible by-product of electrically decomposed SF6 gas -- an essentially inert insulator used in high voltage systems such as transmission lines, substations and switchgear. S2F10 is also made during the production of SF6, but is distilled out.
n the +5 oxidation state.
At The analogous reaction with bromine is reversible and yields SF
5Br.[3] The reversibility of this reaction can be used to synthesize S
2F
10 from SF
5Br.[4]
Ammonia is oxidised by S
2F
10 into NSF
3.[5]
Toxicity
S
2F
10 is a colorless, odorless liquid about 4 times as poisonous as phosgene; a single breath can kill within a day. Its toxicity is thought to be caused by its disproportionation in the lungs into SF
6, which is inert, and SF
4, which reacts with moisture to form sulfurous acid and hydrofluoric acid.[6]
External links
References
- ^ Kenneth G. Denbigh and Robert Whytlaw-Gray (1934). "The preparation and properties of disulphur decafluoride". J. Chem. Soc.: 1346–1352. doi:10.1039/JR9340001346.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1021/ja01108a015, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1021/ja01108a015instead. - ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1021/ic50034a025, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1021/ic50034a025instead. - ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1016/S0022-1139(97)00096-1, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi=10.1016/S0022-1139(97)00096-1instead. - ^ Steve Mitchell (1996). Steve Mitchell (ed.). Biological interactions of sulfur compounds. CRC Press. p. 14. ISBN 0748402454.
- ^ Harold Johnston (2003). A bridge not attacked: chemical warfare civilian research during World War II. World Scientific. pp. 33–36. ISBN 9812381538.
- Loucas G. Christophorou; Isidor Sauers (1991). Gaseous Dielectrics VI. Plenum Press. ISBN 0-306-43894-1.
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