Sulfur tetrafluoride is the chemical compound with the formula S F₄. 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–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.

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