|Molar mass||108.07 g/mol|
|Density||1.95 g/cm3, −78 °C|
|Melting point||−121.0 °C|
|Boiling point||−38 °C|
|Vapor pressure||10.5 atm (22°C)|
|Main hazards||highly toxic
|Safety data sheet||ICSC 1456|
|EU Index||Not listed|
|US health exposure limits (NIOSH):|
|C 0.1 ppm (0.4 mg/m3)|
IDLH (Immediate danger
Related sulfur fluorides
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is: / ?)(|
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, some of which are important in the pharmaceutical and specialty chemical industries.
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.
Synthesis and manufacture
- SCl2 + Cl2 + 4 NaF → SF4 + 4 NaCl
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. 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.
- SF4 + Me3SiNEt2 → Et2NSF3 + Me3SiF
SF4 + 2 H2O → SO2 + 4 HF
Disulfur decafluoride is a colorless gas or liquid with a sulfur-dioxide-like odor. It is about 4 times as poisonous as phosgene. Its toxicity is thought to be caused by its disproportionation in the lungs, according to the following reaction:
S2F10 → SF6 + SF4
6 is inert, and SF
4 reacts inside the lungs with moisture:
SF4 + 2 H2O → SO2 + 4 HF
- "NIOSH Pocket Guide to Chemical Hazards #0580". National Institute for Occupational Safety and Health (NIOSH).
- 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.
- 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.
- Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- F. S. Fawcett, C. W. Tullock, "Sulfur (IV) Fluoride: (Sulfur Tetrafluoride)" Inorganic Syntheses, 1963, vol. 7, pp 119–124. doi:10.1002/9780470132388.ch33
- Hasek, W. R. "1,1,1-Trifluoroheptane". Org. Synth.; Coll. Vol. 5, p. 1082
- 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.
- W. J. Middleton, E. M. Bingham. "Diethylaminosulfur Trifluoride". Org. Synth.; Coll. Vol. 6, p. 440
- 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
- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419.
- "Sulfur Pentaflu". 1988 OSHA PEL Project. CDC NIOSH.
- 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.