Thiophosphoryl fluoride is an inorganic molecular gas with formula P S F₃ containing phosphorus, sulfur and fluorine. It spontaneously ignites in air and burns with a cool flame. The discoverers were able to have flames around their hands without discomfort,^[5] and called it "probably one of the coldest flames known".^[5] The gas was discovered in 1888.^[5]

It is useless for chemical warfare as it burns immediately and is not toxic enough.^[6]

Preparation

Thiophosphoryl fluoride was discovered and named by J. W. Rodger and T. E. Thorpe in 1888.^[5]^[7]

They prepared it by heating arsenic trifluoride and thiophosphoryl chloride together in a sealed glass tube to 150 °C. Also produced in this reaction was silicon tetrafluoride and phosphorus fluorides. By increasing the PSCl₃ the proportion of PSF₃ was increased. They observed the spontaneous inflammability. They also used this method:

3 PbF₂ + P₂S₅ → 3 PbS + PSF₃

at 170 °C, and also substituting a mixture of red phosphorus and sulfur, and substituting bismuth trifluoride.^[5]

Another way to prepare PSF₃ is to add fluoride to PSCl₃ using sodium fluoride in acetonitrile.^[8]

A high yield reaction can be used to produce the gas:^[9]^{[citation needed]}

P₄S₁₀ + 12 HF → 6 H₂S + 4 PSF₃

Under high pressure phosphorus trifluoride can react with hydrogen sulfide to yield:^[10]

PF₃ + H₂S → PSF₃ + H₂ (1350 bar at 200 °C)

Another high pressure production uses phosphorus trifluoride with sulfur.^[10]

Reactions

PSF₃ is decomposed by moisture and oxygen or heat. With heat, phosphorus, sulfur and phosphorus fluorides are formed:

PSF₃ → PF₃ + S

The hot gas reacts with glass producing SF₄, sulfur and elemental phosphorus. The pure gas is completely absorbed by alkali solutions. However, it does not react with ether, benzene, carbon disulfide, or pure sulfuric acid. It is stable against CaO, which can be used to remove impurities such as SiF₄ and PF₃. In air it burns spontaneously with a greyish green flame, producing solid white fumes. With dry oxygen combustion may not be spontaneous and the flame is yellow. On burning SO₂ and P₂O₅ are produced. The gas burns with one of the coldest flames known.^[5]

Reaction with water is slow:

PSF₃ + 4 H₂O → H₂S + H₃PO₄ + 3 HF

If PSF₃ is allowed to react with water in a lead glass container, the hydrofluoric acid and hydrogen sulfide combination produces a black deposit of lead sulfide on the inner surface of the glass.^[5]

It reacts with four times its volume of ammonia gas producing ammonium fluoride and a mystery product, possibly P(NH₂)₂SF.^[5]

PSF₃ is an initiator for the polymerization of tetrahydrofuran.^[11]

Sulfur removal

2 PSF₃ + SO₂ → 2 POF₃ + 3 S

This reaction indicates why PSF₃ is not formed from PF₃ and SO₂.^[10]

PSF₃ + SO₃ → POF₃ + 2 S and sulfur and sulfur sesquioxide (S2+4 polysulfate) as additional products.^[12]^{[clarification needed]}

Fluorine substitution

PSF₃ + 2 ICl → PCl₂F₃.^[13]^{[clarification needed]}

PSF₃ combines with dimethylamine in solution to produce dimethylaminothiophosphoryl difluoride (H₃C−)₂N−P(=S)F₂ and difluorophosphate and hexafluorophosphate ions.^[14]

Thiophosphoryl difluoride isocyanate can be formed by reacting PSF₃ with silicontetraisocyanate at 200 °C in an autoclave.^[9]

Cations

PSF₃ reacts with alkali solutions producing the fluoride and a thiophosphate (PSO3−3).^[5]

CsF + 2 PSF₃ → Cs[PF₆] + CsPS₂F₂^[15]

The reaction is showing the PSF₃ gas is related to the difluorodithiophosphate cation PS₂F−2.^[16]

The CsPS₂F₂ is caesium difluorodithiophosphate.

PSF₃ reacts with [SF₆]⁻ in a mass spectrometer to form [PSF₄]⁻.^[17]

PSF₃ + [SF₆]⁻ → [PSF₄]⁻ + SF₅^•

Related compounds

One fluorine can be substituted by iodine to give thiophosphoryl difluoride iodide, PSIF₂.^[18] PSIF₂ can be converted to hydrothiophosphoryldifluoride, S=PHF₂, by reducing it with hydrogen iodide.^[19] In F₂P(=S)−S−PF₂, one sulfur forms a bridge between two phosphorus atoms.^[18]

Dimethylaminothiophosphoryl difluoride ((H₃C−)₂N−P(=S)F₂) is a foul smelling liquid with a boiling point of 117 °C. It has a Trouton constant (entropy of vaporization at the boiling point of the liquid) of 24.4, and a heat of evaporation of 9530 cal/mole. Alternately it can be produced by fluorination of dimethylaminothiophosphoryl dichloride ((H₃C−)₂N−P(=S)Cl₂).

Physical properties

The thiophosphoryl trifluoride molecule shape has been determined using electron diffraction. The interatomic distances are P=S 0.187±0.003 nm, P−F 0.153±0.002 nm and bond angles of F−P−F bonding is 100.3±2°, The microwave rotational spectrum has been measured for several different isotopologues.^[20]

The critical point is at 346 K at 3.82 MPa.^[21] The liquid refractive index is 1.353.^[4]

The enthalpy of vaporisation 19.6 kJ/mol at boiling point.^[22] The enthalpy of vaporisation at other temperatures is a function of temperature T: H(T)=28.85011(346-T)^0.38 kJ/mol.^[23]

The molecule is polar. It has a non-uniform distribution of positive and negative charge which gives it a dipole moment. When an electric field is applied more energy is stored than if the molecules did not respond by rotating. This increases the dielectric constant. The dipole moment of one molecule of thiophosphoryl trifluoride is 0.640 Debye.^[24]

The infrared spectrum includes vibrations at 275, 404, 442, 698, 951 and 983 cm⁻¹.^[25] These can be used to identify the molecule.