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Sulfur tetrafluoride

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Sulfur tetrafluoride
Structural formula of sulfur tetrafluoride, showing dimernsions
Ball-and-stick model of sulfur tetrafluoride
Ball-and-stick model of sulfur tetrafluoride
Space-filling model of sulfur tetrafluoride
Space-filling model of sulfur tetrafluoride
Names
IUPAC name
Sulfur(IV) fluoride
Other names
Sulfur tetrafluoride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.029.103 Edit this at Wikidata
RTECS number
  • WT4800000
UNII
UN number 2418
  • InChI=1S/F4S/c1-5(2,3)4 checkY
    Key: QHMQWEPBXSHHLH-UHFFFAOYSA-N checkY
  • InChI=1/F4S/c1-5(2,3)4
    Key: QHMQWEPBXSHHLH-UHFFFAOYAT
  • FS(F)(F)F
Properties
SF4
Molar mass 108.07 g/mol
Appearance colorless gas
Density 1.95 g/cm3, −78 °C
Melting point −121.0 °C
Boiling point −38 °C
reacts
Vapor pressure 10.5 atm (22 °C)[1]
Structure
Seesaw (C2v)
0.632 D[2]
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 highly toxic
corrosive
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
3
0
2
W
NIOSH (US health exposure limits):
PEL (Permissible)
 none[1]
REL (Recommended)
 C 0.1 ppm (0.4 mg/m3)[1]
IDLH (Immediate danger)
 N.D.[1]
Safety data sheet (SDS) ICSC 1456
Related compounds
Other anions
 Sulfur dichloride
Disulfur dibromide
Sulfur trifluoride
Other cations
 Oxygen difluoride
Selenium tetrafluoride
Tellurium tetrafluoride
Related sulfur fluorides
 Disulfur difluoride
Sulfur difluoride
Disulfur decafluoride
Sulfur hexafluoride
Related compounds
 Thionyl fluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Sulfur tetrafluoride is a chemical compound with the formula SF4. It is a colorless corrosive gas that releases dangerous hydrogen fluoride gas upon exposure to water or moisture. Sulfur tetrafluride is a useful reagent for the preparation of organofluorine compounds,[3] some of which are important in the pharmaceutical and specialty chemical industries.

Structure[edit]

Sulfur in SF4 is in the +4 oxidation state, with two lone pairs. The atoms in SF4 are arranged in a see-saw shape, with the sulfur atom at the center. One of the three equatorial positions is occupied by a nonbonding lone pair of electrons. Consequently, the molecule has two distinct types of F ligands, two axial and two equatorial. The relevant bond distances are S–Fax = 164.3 pm and S–Feq = 154.2 pm. It is typical for the axial ligands in hypervalent molecules to be bonded less strongly.

The 19F NMR spectrum of SF4 reveals only one signal, which indicates that the axial and equatorial F atom positions rapidly interconvert via pseudorotation.[4]

Intramolecular dynamic equilibration of SF4.

Synthesis and manufacture[edit]

At the laboratory scale, sulfur tetrafluride is prepared from elemental sulfur and cobaltic fluoride[5]

S + 4CoF3 → SF4 + 4CoF2

SF4 is industrially produced by the reaction of SCl2 and NaF with acetonitrile as a catalyst[6]

3 SCl2 + 4 NaF → SF4 + S2Cl2 + 4 NaCl

At higher temperatures (e.g. 225–450 °C), the solvent is superfluous. Moreover, sulfur dichloride may be replaced by elemental sulfur (S) and chlorine (Cl2).[7][8]

A low-temperature (e.g. 20–86 °C) alternative to the chlorinative process above uses liquid bromine (Br2) as oxidant and solvent:[9]

S(s) + 2 Br2(l; excess) + 4KF(s) → SF4↑ + 4 KBr(brom)

Use in synthesis of organofluorine compounds[edit]

Main article: Fluorination by sulfur tetrafluoride

In organic synthesis, SF4 is used to convert COH and C=O groups into CF and CF2 groups, respectively.[10] The efficiency of these conversions are highly variable.

In the laboratory, the use of SF4 has been superseded by the safer and more easily handled diethylaminosulfur trifluoride, (C2H5)2NSF3, "DAST":[11] This reagent is prepared from SF4:[12]

SF4 + (CH3)3SiN(C2H5)2 → (C2H5)2NSF3 + (CH3)3SiF

Other reactions[edit]

Sulfur chloride pentafluoride (SF
5Cl), a useful source of the SF5 group, is prepared from SF4.[13]

SF4 + Cl2 + CsF → SF5Cl + CsCl

Hydrolysis of SF4 gives sulfur dioxide:[14]

SF4 + 2 H2O → SO2 + 4 HF

This reaction proceeds via the intermediacy of thionyl fluoride, which usually does not interfere with the use of SF4 as a reagent.[6]

When amines are treated with SF4 and a base, aminosulfur difluorides result.[15]

Toxicity[edit]

SF
4 reacts inside the lungs with moisture, forming sulfur dioxide and hydrogen fluoride which forms highly toxic and corrosive hydrochloric acid [16]

References[edit]

  1. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0580". National Institute for Occupational Safety and Health (NIOSH).
  2. ^ Tolles, W. M.; W. M. Gwinn, W. D. (1962). "Structure and Dipole Moment for SF4". J. Chem. Phys. 36 (5): 1119–1121. Bibcode:1962JChPh..36.1119T. doi:10.1063/1.1732702.
  3. ^ Wang, C.-L. J. (2004). "Sulfur Tetrafluoride". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X. hdl:10261/236866. ISBN 9780471936237.
  4. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
  5. ^ Kwasnik, W. (1963). "Fluorine compounds: Sulfur tetrafluoride". In Brauer, Georg (ed.). Handbook of Preparative Inorganic Chemistry. Vol. 1. Translated by Riley, Reed F. (2nd ed.). NY, NY: Academic Press. p. 168. LCCN 63-14307 – via the Internet Archive.
  6. ^ a b Fawcett, F. S.; Tullock, C. W. (1963). "Sulfur(IV) Fluoride" (PDF). Inorganic Syntheses. Vol. 7. pp. 119–124. doi:10.1002/9780470132388.ch33. ISBN 978-0-470-13166-4.
  7. ^ Tullock, C. W.; Fawcett, F. S.; Smith, W. C.; Coffman, D. D. (1960). "The Chemistry of Sulfur Tetrafluoride. I. The Synthesis of Sulfur Tetrafluoride". J. Am. Chem. Soc. 82 (3): 539–542. doi:10.1021/ja01488a011.
  8. ^ US 2992073, Tullock, C.W., "Synthesis of Sulfur Tetrafluoride", issued 1961 
  9. ^ Winter, R.W.; Cook P.W. (2010). "A simplified and efficient bromine-facilitated SF4-preparation method". J. Fluorine Chem. 131: 780-783. doi:10.1016/j.jfluchem.2010.03.016
  10. ^ Hasek, W. R. (1961). "1,1,1-Trifluoroheptane". Organic Syntheses. 41: 104. doi:10.15227/orgsyn.041.0104.
  11. ^ Fauq, A. H. (2004). "N,N-Diethylaminosulfur Trifluoride". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X. hdl:10261/236866. ISBN 9780471936237..
  12. ^ W. J. Middleton; E. M. Bingham (1977). "Diethylaminosulfur Trifluoride". Organic Syntheses. 57: 440. doi:10.15227/orgsyn.057.0050.
  13. ^ Nyman, F.; Roberts, H. L.; Seaton, T. (1966). "Sulfur Chloride Pentafluoride" (PDF). Inorganic Syntheses. Vol. 8. McGraw-Hill. p. 160. doi:10.1002/9780470132395.ch42. ISBN 9780470132395.
  14. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  15. ^ Grunwell, John R.; Dye, Sandra L. (1975). "Novel generation of benzonitrile-N-sulfide". Tetrahedron Letters. 16 (21): 1739–1740. doi:10.1016/s0040-4039(00)72247-7. ISSN 0040-4039.
  16. ^ Johnston, H. (2003). A Bridge not Attacked: Chemical Warfare Civilian Research During World War II. World Scientific. pp. 33–36. ISBN 981-238-153-8.
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