Disulfur decafluoride: Difference between revisions

Disulfur decafluoride
Ball-and-stick model of disulfur decafluoride	Space-filling model of disulfur decafluoride
Names
	Preferred IUPAC name Disulfur decafluoride
	Systematic IUPAC name Decafluoro-1λ6,2λ6-disulfane
Identifiers
CAS Number	5714-22-7;
3D model (JSmol)	Interactive image;
ChemSpider	56348 ;
ECHA InfoCard	100.024.732
EC Number	227-204-4;
MeSH	Disulfur+decafluoride
PubChem CID	62586;
CompTox Dashboard (EPA)	DTXSID2073356 ;
	InChI InChI=1S/F10S2/c1-11(2,3,4,5)12(6,7,8,9)10 Key: BPFZRKQDXVZTFD-UHFFFAOYSA-N;
	SMILES FS(F)(F)(F)(F)S(F)(F)(F)(F)F;
Properties
Chemical formula	S2F10
Appearance	colorless liquid
Melting point	-53 °C
Boiling point	30.1 °C
Hazards
NFPA 704 (fire diamond)	4 3 2
	Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Infobox references

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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

CDC page on the substance's toxicology

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/ja01108a015 instead.
^ 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/ic50034a025 instead.
^ 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)00096-1 instead.
^ 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. {{cite book}}: Unknown parameter |editors= ignored (|editor= suggested) (help)

This inorganic compound–related article is a stub. You can help Wikipedia by expanding it.

[1] Kenneth G. Denbigh and Robert Whytlaw-Gray (1934). "The preparation and properties of disulphur decafluoride". J. Chem. Soc.: 1346–1352. doi:10.1039/JR9340001346.

[2] 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/ja01108a015 instead.

[3] 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/ic50034a025 instead.

[4] 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)00096-1 instead.

[5] Steve Mitchell (1996). Steve Mitchell (ed.). Biological interactions of sulfur compounds. CRC Press. p. 14. ISBN 0748402454.

[6] Harold Johnston (2003). A bridge not attacked: chemical warfare civilian research during World War II. World Scientific. pp. 33–36. ISBN 9812381538.

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@@ Line 39: / Line 39: @@
 '''Disulfur decafluoride''' (S<sub>2</sub>F<sub>10</sub>) is a gas discovered in 1934 by Denbigh and Whytlaw-Gray.<ref>{{cite journal |journal=J. Chem. Soc. |year=1934 |pages=1346–1352 |doi=10.1039/JR9340001346 |title=The preparation and properties of disulphur decafluoride |author=Kenneth G. Denbigh and Robert Whytlaw-Gray}}</ref> Each S of the S<sub>2</sub>F<sub>10</sub> molecule is [[octahedral]], and surrounded by 5 fluorines.<ref>{{cite doi|10.1021/ja01108a015}}</ref> S<sub>2</sub>F<sub>10</sub> 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 [[Sulfur hexafluoride|SF<sub>6</sub>]] gas -- an essentially inert [[Electrical insulation|insulator]] used in high voltage systems such as [[transmission lines]], [[Electrical substation|substation]]s and [[switchgear]]. S<sub>2</sub>F<sub>10</sub> is also made during the production of  SF<sub>6</sub>, but is distilled out.
+n the +5 oxidation state.
-==Properties==
+At The analogous reaction with [[bromine]] is reversible and yields {{chem|SF|5|Br}}.<ref>{{cite doi|10.1021/ic50034a025}}</ref> The reversibility of this reaction can be used to synthesize {{chem|S|2|F|10}} from {{chem|SF|5|Br}}.<ref>{{cite doi|10.1016/S0022-1139(97)00096-1}}</ref>
-This compound contains sulfur in the +5 oxidation state.
-At temperatures above 150°C, {{chem|S|2|F|10}} decomposes slowly to {{chem|SF|6}} and {{chem|SF|4}}.
-{{chem|S|2|F|10}} reacts with {{chem|N|2|F|4}} to give {{chem|SF|5|NF|2}}. It reacts with {{chem|SO|2}} to form {{chem|SF|5|OSO|2|F}} in the presence of ultraviolet radiation.
-In the presence of excess [[chlorine]] gas, {{chem|S|2|F|10}} reacts to form [[sulfur chloride pentafluoride]] ({{chem|SF|5|Cl}}):
-: {{chem|S|2|F|10}} + {{chem|Cl|2}} → 2 {{chem|SF|5|Cl}}
-The analogous reaction with [[bromine]] is reversible and yields {{chem|SF|5|Br}}.<ref>{{cite doi|10.1021/ic50034a025}}</ref> The reversibility of this reaction can be used to synthesize {{chem|S|2|F|10}} from {{chem|SF|5|Br}}.<ref>{{cite doi|10.1016/S0022-1139(97)00096-1}}</ref>
 [[Ammonia]] is oxidised by {{chem|S|2|F|10}} into [[Thiazyl trifluoride|{{chem|NSF|3}}]].<ref>{{cite book