Sulfur tetrafluoride is the chemical compound with the formula SF4. 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.
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Names | |||
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IUPAC name
Sulfur(IV) fluoride
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Other names
Sulfur tetrafluoride
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Identifiers | |||
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3D model (JSmol)
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ChEBI |
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ChemSpider |
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ECHA InfoCard | 100.029.103 | ||
PubChem CID
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RTECS number |
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UNII |
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UN number | 2418 | ||
CompTox Dashboard (EPA)
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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
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highly toxic corrosive | ||
NFPA 704 (fire diamond) | |||
NIOSH (US health exposure limits): | |||
PEL (Permissible)
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none[1] | ||
REL (Recommended)
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C 0.1 ppm (0.4 mg/m3)[1] | ||
IDLH (Immediate danger)
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N.D.[1] | ||
Safety data sheet (SDS) | ICSC 1456 | ||
Related compounds | |||
Other anions
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Sulfur dichloride Disulfur dibromide Sulfur trifluoride | ||
Other cations
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Oxygen difluoride Selenium tetrafluoride Tellurium tetrafluoride | ||
Related sulfur fluorides
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Disulfur difluoride Sulfur difluoride Disulfur decafluoride Sulfur hexafluoride | ||
Related compounds
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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
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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]
At the laboratory scale, fluorination of elemental sulfur with cobaltic fluoride suffices:[5]
For larger-scale syntheses, SF4 is produced by the reaction of SCl2 and NaF in acetonitrile:[6]
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]
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 more conveniently handled diethylaminosulfur trifluoride, Et2NSF3, "DAST":[11] This reagent is prepared from SF4:[12]
Sulfur chloride pentafluoride (SF
5Cl), a useful source of the SF5 group, is prepared from SF4.[13]
Hydrolysis of SF4 gives sulfur dioxide:[14]
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, iminosulfur difluorides result.[15]
SF
4 reacts inside the lungs with moisture, generating sulfur dioxide and hydrogen fluoride:[16]