Thiophosphoryl fluoride is an inorganic molecular gas with formula PSF3 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, and called it "probably one of the coldest flames known". The gas was discovered in 1888. It is useless for chemical warfare as it burns immediately and is not toxic enough.
Preparation Thiophosphoryl fluoride was discovered and named by J. W. Rodger and T. E. Thorpe in 1888. 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 + 2 PSF3 at 170 °C, and also substituting a mixture of red phosphorus and sulfur, and substituting bismuth trifluoride. Another way to prepare PSF3 is to add fluoride to PSCl3 using sodium fluoride in acetonitrile. A high yield reaction can be used to produce the gas:
P4S10 + 12 HF → 6 H2S + 4 PSF3 Under high pressure phosphorus trifluoride can react with hydrogen sulfide to yield:
PF3 + H2S → PSF3 + H2 (1350 bar at 200 °C) Another high pressure production uses phosphorus trifluoride with sulfur.
Reactions PSF3 is unstable against moisture or heat. The pure gas is completely absorbed by alkali solutions, producing the fluoride and a thiophosphate (PSO3−3), but stable against CaO. The latter can be used to remove SiF4 or PF3 impurities.
Hydrolysis and decomposition Reaction with neutral water is slow:
PSF3 + 4 H2O → H2S + H3PO4 + 3 HF Nevertheless, dissociation constants for related acids suggest that the phosphorus atom is at least as electrophilic as in phosphoryl fluoride. Autodecomposition from heat gives phosphorus fluorides, sulfur, and phosphorus:
PSF3 → S + PF3 → ... Hot PSF3 reacts with glass, producing SF4, sulfur and elemental phosphorus. If water is present and the glass is leaded, then the hydrofluoric acid and hydrogen sulfide combination produces a black plumbous sulfide deposit on the inner surface.
Oxidation In air, PSF3 burns spontaneously with a greyish green flame, producing solid white fumes containing SO2 and P2O5. The flame is one of the coldest known. With dry oxygen, combustion may not be spontaneous and the flame is yellow. Thiophosphoryl fluoride reduces oxygenated compounds to give phosphoryl fluoride and sulfur:
PSF3 + 3 SO3 → POF3 + 4 SO2 2 PSF3 + SO2 → 2 POF3 + 3 S The latter reaction also indicates why PSF3 is not formed from PF3 and SO2. Various oxidants can convert thiophosphoryl fluoride to phosphorus dichloride trifluoride, e.g.:
PSF3 + 2 ICl → PCl2F3 + I2 + S.
Nucleophilic substitution Thiophosphoryl difluoride isocyanate can be formed by reacting PSF3 with silicon tetraisocyanate at 200 °C in an autoclave. In general, nucleophilic substitution onto thiophosphoryl fluoride is complex, because free fluoride ions tend to induce disproportionation to hexafluorophosphate and dithiodifluorophosphate (PS2F−2). For example, with cesium fluoride:
CsF + 2 PSF3 → Cs[PF6] + CsPS2F2 Thus PSF3 combines with dimethylamine in solution to produce dimethylaminothiophosphoryl difluoride (H3C−)2N−P(=S)F2 and difluorophosphate and hexafluorophosphate ions:
4 SPF3 + 4 HNMe2 → 2 SPF2NMe2 + [H2NMe2]PF6 + [H2NMe2]S2PF2. PSF3 reacts with four times its volume of ammonia gas producing ammonium fluoride and a mystery product, possibly P(NH2)2SF.
Miscellaneous PSF3 does not react with ether, benzene, carbon disulfide, or pure sulfuric acid. It initiates tetrahydrofuran polymerization. PSF3 reacts with [SF6]− in a mass spectrometer to form [PSF4]−.
PSF3 + [SF6]−• → PSF4− + SF5•
Related compounds One fluorine can be substituted by iodine to give thiophosphoryl difluoride iodide, PSIF2. PSIF2 can be converted to hydrothiophosphoryldifluoride, S=PHF2, by reducing it with hydrogen iodide. In F2P(=S)−S−PF2, one sulfur forms a bridge between two phosphorus atoms. 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. The critical point is at 346 K at 3.82 MPa. The liquid refractive index is 1.353. The enthalpy of vaporisation 19.6 kJ/mol at boiling point. The enthalpy of vaporisation at other temperatures is a function of temperature T: H(T)=28.85011(346-T)0.38 kJ/mol. 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. The infrared spectrum includes vibrations at 275, 404, 442, 698, 951 and 983 cm−1. These can be used to identify the molecule.
See also Difluorodithiophosphate, [S2PF2]−
References
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