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Sulfur tetrafluoride

Sulfur tetrafluoride

Systematic name Sulfur(IV) fluoride
Other names Sulfur tetrafluoride
Molecular formula SF4
Molar mass 108.07 g/mol
Appearance colorless gas
CAS number [7783-60-0]
Density and phase 1.95 g cm−3, −78 °C
Solubility in water decomp
Other solvents benzene
Melting point −121.0 °C
Boiling point −38 °C
Dipole moment 0.632 D[1]
MSDS External MSDS
Main hazards highly toxic
NFPA 704
R/S statement R: 14-26-34-37
S: 26-36/37/39-38-45
RTECS number WT4800000
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Related compounds
Related compounds SeF4
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Infobox disclaimer and references

Sulfur tetrafluoride is the chemical compound with the formula SF4. This species exists as a gas at standard conditions. 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,[2] some of which are important in the pharmaceutical and specialty chemical industries. Aside from SF4, two other "binary" sulfur fluorides are well known, SF6 and S2F10.


Structure of the SF4 molecule

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 trigonal bipyramid, with S at the center. One of the three equatorial positions is occupied by a nonbonding lone pair of electrons. Ignoring this lone pair, SF4 resembles a "see-saw." Consequently, the molecule has two distinct types of F ligands, two axial and two equatorial. The relevant bond distances are S-Fax = 1.643 Å and S-Feq = 1.542 Å. 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."[3]  

Synthesis and manufacture

SF4 is produced by the reaction of SCl2, Cl2, and NaF:

SCl2 + Cl2 + 4NaF → SF4 + 4NaCl

Interestingly, treatment of SCl2 with NaF also affords SF4, not SF2. SF2 is unstable, it condenses with itself to form SF4 and SSF2.

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.[4] 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. 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.[5] This reagent is prepared from SF4:[6]

SF4 + Me3SiNEt2 → Et2NSF3 + Me3SiF


  1. ^ Tolles; W. M. Gwinn, W. D. “Structure and Dipole Moment for SF4” The Journal of Chemical Physics 1962, Volume 36, pp. 1119-1121. doi:10.1063/1.1732702
  2. ^ 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.
  3. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  4. ^ W. R. Hasek "1,1,1-Trifluoroheptane" Organic Syntheses, Coll. Vol. 5, p.1082; Vol. 41, p.104.
  5. ^ 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.
  6. ^ W. J. Middleton, E. M. Bingham" Diethylaminosulfur Trifluoride” Organic Syntheses, Coll. Vol. 6, p.440; Vol. 57, p.50.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Sulfur_tetrafluoride". A list of authors is available in Wikipedia.
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