Phosphorus trichloride
| Names | |
|---|---|
| IUPAC name
Phosphorus trichloride
| |
| Other names
Phosphorus(III) chloride
Phosphorous chloride Trichlorophosphane Trichlorophosphine | |
| Properties | |
| PCl3 | |
| Molar mass | 137.33 g/mol |
| Appearance | Colorless fuming liquid |
| Odor | Stinging, acid-like |
| Density | 1.574 g/cm3 |
| Melting point | −93.6 °C (−136.5 °F; 179.6 K) |
| Boiling point | 76.1 °C (169.0 °F; 349.2 K) |
| Reacts | |
| Solubility | Reacts with alcohols Soluble in benzene, benzoyl chloride, bromoacetic acid, liq. carbon dioxide, carbon tetrachloride, chloroform, cyanogen chloride, liq. H2S, liq. HCl, iron pentacarbonyl, phosphoryl chloride, liq. SO2, sulfuryl chloride, thionyl chloride, tin(IV) chloride |
| Vapor pressure | 120 mm Hg (25 °C) |
| Thermochemistry | |
| Std enthalpy of
formation (ΔfH |
-319.7 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
18 mg/kg (rat, oral) 104 ppm (rat, 4 hr) 50 ppm (guinea pig, 4 hr) |
| Related compounds | |
| Related compounds
|
Phosphorus pentachloride Phosphorus tribromide Phosphorus triiodide |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Phosphorus trichloride is a chemical compound with the formula PCl3, widely used as a precursor for a variety of chemical compounds, most often organophosphorus compounds.
Contents
Properties
Chemical
Phosphorus trichloride readily hydrolyzes in water to give hydrogen chloride and phosphorus acid:
- PCl3 + 3 H2O → H3PO3 + 3 HCl
Physical
Phosphorus trichloride is a colorless or faint yellow fuming liquid, which readily hydrolyzes in moist air and water releasing a mist of hydrogen chloride, giving the compound a strong HCl odor.
Phosphorus trichloride is known to form azeotropes with sulfuryl chloride and thionyl chloride respectively.[1]
Availability
Phosphorus trichloride is sold by chemical suppliers, but due to its hazards it's out of reach for the amateur chemist.
Phosphorus trichloride is listed on Schedule 3 of the Chemical Weapons Convention, which means all transactions could be be monitored, especially for large amounts.
Preparation
Preparation of phosphorus trichloride requires dry or inert medium, as the substance is volatile, very corrosive and highly toxic. This is not something everyone can do safely and as such it's recommended that only professionals should try this.
The most convenient route involves the reaction of white or red phosphorus with chlorine, in phosphorus tichloride. The PCl3 solvent where the bulk process takes place is made by reacting chlorine with phosphorus directly, and then more phosphorus is added, followed by the addition of chlorine. The reaction takes place under reflux and the resulting phosphorus trichloride is fractionally distilled to prevent the formation of phosphorus pentachloride. Needless to say, both precursors must be perfectly dry.
Passing hot vapors of phosphoryl chloride over red hot charcoal or activated carbon at 400 °C will give phosphorus trichloride.[2]
Heating phosphorus pentoxide with sodium chloride in an iron or steel crucible at 250 °C will give phosphoryl chloride and phosphorus trichloride. The resulting POCl3 is reduced by the iron surface to PCl3. Interestingly, while some free chlorine is produced, the reaction does not give phosphorus pentachloride. Calcium chloride can also be used, but the reaction will only start above 400 °C.[3]
Heating red phosphorus with thionyl chloride at 125 °C will also give phosphorus trichloride.[4]
Mercury will reduce phosphorus pentachloride to phosphorus trichloride. While not specified directly, the reaction appears to take place in carbon tetrachloride. Mercury(II) chloride is formed as byproduct.[5] Other metals, such as tin will also work.[6]
Reaction of phosphorus triiodide with chlorine in neat will give phosphorus trichloride. Reaction of ferric chloride with phosphorus also gives PCl3.[7]
Phosphorus trichloride can also be produced by reducing phosphorus pentachloride with phosphine.[8]
Heating mercury(I) chloride with white phosphorus will give PCl3.[9]
Projects
While very useful in preparing a huge variety of phosphorus compounds and more, the great reactivity and toxicity of phosphorus trichloride means you will need a good ventilation system to safely handle the chemical.
- Make phosphorus pentachloride
- Make phosphoryl chloride
- Make alkyl and acyl chlorides
- Make phosphite esters
- Make triphenylphosphine
Handling
Safety
Phosphorus trichloride is very corrosive and toxic. In moist air, the substance hydrolyzes releasing hydrochloric acid fumes. Air concentrations above 600 ppm can cause death in just a few minutes.
Storage
Phosphorus trichloride should be stored in Schlenk flasks, as they're the best storage containers for air and water sensitive reagents.
Disposal
Phosphorus trichloride should be neutralized with a base, such as calcium hydroxide suspension outside or in a well ventilated area. Always add small amounts of PCl3 at a time to prevent splashing. Try not to use carbonates, as they will fizzle and cause some PCl3 to become airborne.
References
- ↑ Astakhova, G. V.; Bokareva, E. I.; Mironova, E. B.; Nisel'son, L. A.; Russian Journal of Inorganic Chemistry (Translation of Zhurnal Neorganicheskoi Khimii); vol. 22; (1977); p. 113 - 116
- ↑ Riban, J.; Comptes Rendus Hebdomadaires des Seances de l'Academie des Sciences; vol. 95; (1882); p. 1160 - 1163
- ↑ Tarbutton, G.; Egan, E. P.; Frary, S. G.; Journal of the American Chemical Society; vol. 63; (1941); p. 1782 - 1789
- ↑ North, H. B.; Thomson, J. C.; Journal of the American Chemical Society; vol. 40; (1918); p. 774 - 777
- ↑ Downs, J. J.; Johnson, R. E.; Journal of the American Chemical Society; vol. 77; (1955); p. 2098 - 2102
- ↑ Baudrimont, E.; Ann. Chim. Phys.; vol. 2; (1864); p. 5 - 67
- ↑ Gladstone, J. H.; Philosophical Magazine (1798-1977); vol. 35; (1849); p. 345 - 355
- ↑ Mahn, R.; Jena. Z. Med. Naturwiss.; vol. 5; (1870); p. 158 - 166
- ↑ Davy, H.; Schweiggers J. Chem. Phys.; vol. 3; (1811); p. 79 - 120
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