Sulfur tetrafluoride is the chemical compound with the formula S F₄. It is a colorless corrosive gas 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.

Structure

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–F_ax = 164.3 pm and S–F_eq = 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]

Intramolecular dynamic equilibration of SF₄.

Synthesis and manufacture

At the laboratory scale, fluorination of elemental sulfur with cobaltic fluoride suffices:^[5]

S + 4CoF₃ → SF₄ + 4CoF₂

For larger-scale syntheses, SF₄ is produced by the reaction of SCl₂ and NaF in acetonitrile:^[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 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.^[10] 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.^[11] This reagent is prepared from SF₄:^[12]

SF₄ + Me₃SiNEt₂ → Et₂NSF₃ + Me₃SiF

Other reactions

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

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]

Toxicity

SF
₄ reacts inside the lungs with moisture, generating sulfur dioxide and hydrogen fluoride:^[15]

SF₄ + 2 H₂O → SO₂ + 4 HF

References

^ ^a ^b ^c ^d 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. Bibcode:1962JChPh..36.1119T. doi:10.1063/1.1732702.
^ 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.
^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
^ 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.
^ ^a ^b Fawcett, F. S.; Tullock, C. W. (1963). Sulfur (IV) Fluoride: (Sulfur Tetrafluoride). Inorganic Syntheses. Vol. 7. pp. 119–124. doi:10.1002/9780470132388.ch33.
^ 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.
^ US 2992073, Tullock, C.W., "Synthesis of Sulfur Tetrafluoride", issued 1961
^ 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
^ Hasek, W. R. "1,1,1-Trifluoroheptane". Organic Syntheses.; Collective Volume, vol. 5, p. 1082
^ 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..
^ W. J. Middleton; E. M. Bingham. "Diethylaminosulfur Trifluoride". Organic Syntheses.; Collective Volume, vol. 6, p. 440
^ Nyman, F.; Roberts, H. L.; Seaton, T. (1966). "Sulfur Chloride Pentafluoride". Inorganic Syntheses. Inorganic Syntheses. Vol. 8. McGraw-Hill. p. 160. doi:10.1002/9780470132395.ch42. ISBN 9780470132395.
^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
^ 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.