Sulfur tetrafluoride

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

Sulfur tetrafluoride
IUPAC name
Sulfur(IV) fluoride
Other names
Sulfur tetrafluoride
3D model (JSmol)
ECHA InfoCard 100.029.103
RTECS number
  • WT4800000
UN number 2418
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
Vapor pressure 10.5 atm (22°C)[1]
Seesaw (C2v)
0.632 D[2]
Main hazards highly toxic
Safety data sheet ICSC 1456
NFPA 704 (fire diamond)
NIOSH (US health exposure limits):
PEL (Permissible)
REL (Recommended)
C 0.1 ppm (0.4 mg/m3)[1]
IDLH (Immediate danger)
Related compounds
Other anions
Sulfur dichloride
Disulfur dibromide
Sulfur trifluoride
Other cations
Oxygen difluoride
Selenium tetrafluoride
Tellurium tetrafluoride
Polonium 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).
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Infobox references


Sulfur in SF4 is in the formal +4 oxidation state. Of sulfur's total of six valence electrons, two form a lone pair. The structure of SF4 can therefore be anticipated using the principles of VSEPR theory: it is a see-saw shape, with S 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. In contrast to SF4, the related molecule SF6 has sulfur in the 6+ state, no valence electrons remain nonbonding on sulfur, hence the molecule adopts a highly symmetrical octahedral structure. Further contrasting with SF4, SF6 is extraordinarily inert chemically.

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]

Synthesis and manufacture

SF4 is produced by the reaction of SCl2 and NaF in acetonitrile:[5]

3 SCl2 + 4 NaF → SF4 + S2Cl2 + 4 NaCl

SF4 is also produced in the absence of solvent at elevated temperatures.[6][7]

Alternatively, SF4 at high yield is produced using sulfur (S), NaF and chlorine (Cl2) in the absence of reaction medium, also at less-desirable elevated reaction temperatures (e.g. 225-450 °C).[6][7]

A low temperature (e.g. ambient-86 °C) method of producing SF4 at high yield, without the requirement for reaction medium, has been demonstrated utilizing bromine (Br2) instead of chlorine (Cl2), S and KF:[8]

S + (2+x) Br2 + 4 KF → SF4↑ + x Br2 + 4 KBr

Use of SF4 for the synthesis of fluorocarbons

In organic synthesis, SF4 is used to convert COH and C=O groups into CF and CF2 groups, respectively.[9] Certain alcohols readily give the corresponding fluorocarbon. Ketones and aldehydes give geminal difluorides. The presence of protons alpha to the carbonyl leads to side reactions and diminished (30–40%) yield. Also diols can give cyclic sulfite esters, (RO)2SO. Carboxylic acids convert to trifluoromethyl derivatives. For example, treatment of heptanoic acid with SF4 at 100-130 °C produces 1,1,1-trifluoroheptane. Hexafluoro-2-butyne can be similarly produced from acetylenedicarboxylic acid. The coproducts from these fluorinations, including unreacted SF4 together with SOF2 and SO2, are toxic but can be neutralized by their treatment with aqueous KOH.

The use of SF4 is being superseded in recent years by the more conveniently handled diethylaminosulfur trifluoride, Et2NSF3, "DAST", where Et = CH3CH2.[10] This reagent is prepared from SF4:[11]

SF4 + Me3SiNEt2 → Et2NSF3 + Me3SiF

Other reactions

Sulfur chloride pentafluoride (SF
), a useful source of the SF5 group, is prepared from SF4.[12]

Hydrolysis of SF4 gives sulfur dioxide:[13]

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.[5]


reacts inside the lungs with moisture, generating sulfur dioxide and hydrogen fluoride:[14]

SF4 + 2 H2O → SO2 + 4 HF


  1. 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. doi:10.1063/1.1732702.
  3. C.-L. J. Wang, "Sulfur Tetrafluoride" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289.
  4. Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  5. F. S. Fawcett, C. W. Tullock, "Sulfur (IV) Fluoride: (Sulfur Tetrafluoride)" Inorganic Syntheses, 1963, vol. 7, pp 119–124. doi:10.1002/9780470132388.ch33
  6. 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), pp 539–542 doi:10.1021/ja01488a011
  7. Tullock C.W. (1961). "Synthesis of Sulfur Tetrafluoride". United States Patent 2992073.
  8. 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
  9. Hasek, W. R. "1,1,1-Trifluoroheptane". Organic Syntheses.; Collective Volume, 5, p. 1082
  10. A. H. Fauq, "N,N-Diethylaminosulfur Trifluoride" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289.
  11. W. J. Middleton; E. M. Bingham. "Diethylaminosulfur Trifluoride". Organic Syntheses.; Collective Volume, 6, p. 440
  12. Nyman, F., Roberts, H. L., Seaton, T. Inorganic Syntheses, 1966, Volume 8, p. 160 McGraw-Hill Book Company, Inc., 1966, doi:10.1002/9780470132395.ch42
  13. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  14. 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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