Phosphoryl chloride
Phosphoryl chloride (commonly called phosphorus oxychloride) is a colourless liquid with the formula POCl3. It hydrolyses in moist air releasing phosphoric acid and fumes of hydrogen chloride. It is manufactured industrially on a large scale from phosphorus trichloride and oxygen or phosphorus pentoxide.[3] It is mainly used to make phosphate esters such as tricresyl phosphate.
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| Names | |||
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| Preferred IUPAC name
Phosphoryl trichloride[1] | |||
| Other names
Phosphorus oxychloride Phosphoric trichloride Trichlorophosphate | |||
| Identifiers | |||
3D model (JSmol) |
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| ChEBI | |||
| ChemSpider | |||
| ECHA InfoCard | 100.030.030  | ||
| EC Number |
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PubChem CID |
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| RTECS number |
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| UNII | |||
| UN number | 1810 | ||
CompTox Dashboard (EPA) |
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| Properties | |||
| POCl3 | |||
| Molar mass | 153.33 g/mol | ||
| Appearance | colourless liquid, fumes in moist air | ||
| Odor | pungent and musty | ||
| Density | 1.645 g/cm3, liquid | ||
| Melting point | 1.25 °C (34.25 °F; 274.40 K) | ||
| Boiling point | 105.8 °C (222.4 °F; 378.9 K) | ||
| Reacts | |||
| Solubility | highly soluble in benzene, chloroform, CS2, CCl4 | ||
| Vapor pressure | 40 mmHg (27 °C)[2] | ||
Refractive index (nD) |
1.460 | ||
| Structure | |||
| tetrahedral | |||
| 2.54 D | |||
| Thermochemistry | |||
Heat capacity (C) |
84.35 J/mol K | ||
Std enthalpy of formation (ΔfH⦵298) |
-568.4 kJ/mol | ||
| Hazards | |||
| Safety data sheet | See: data page ICSC 0190 | ||
EU classification (DSD) (outdated) |
Very toxic (T+) Harmful (Xn) Corrosive (C) | ||
| R-phrases (outdated) | R14, R22, R26, R35, R48/23 | ||
| S-phrases (outdated) | (S1/2), S7/8, S26, S36/37/39, S45 | ||
| NFPA 704 (fire diamond) | |||
| Lethal dose or concentration (LD, LC): | |||
LD50 (median dose) |
36 mg/kg (rat, oral) | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible) |
none[2] | ||
REL (Recommended) |
TWA 0.1 ppm (0.6 mg/m3) ST 0.5 ppm (3 mg/m3)[2] | ||
IDLH (Immediate danger) |
N.D.[2] | ||
| Related compounds | |||
Related compounds |
Thiophosphoryl chloride Phosphorus oxybromide Phosphorus trichloride Phosphorus pentachloride | ||
| Supplementary data page | |||
| Refractive index (n), Dielectric constant (εr), etc. | |||
Thermodynamic data |
Phase behaviour solid–liquid–gas | ||
| UV, IR, NMR, MS | |||
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 | |||
Structure
Like phosphate, phosphoryl chloride is tetrahedral in shape.[5] It features three P−Cl bonds and one strong P=O double bond, with an estimated bond dissociation energy of 533.5 kJ/mol. On the basis of bond length and electronegativity, the Schomaker-Stevenson rule suggests that the double bond form is dominant, in contrast with the case of POF3. The P=O bond involves the donation of the lone pair electrons on oxygen p-orbitals to the antibonding combinations associated with phosphorus-chlorine bonds, thus constituting π bonding.[6]
Phosphoryl chloride exists as neutral POCl3 molecules in the solid, liquid and gas states. This is unlike phosphorus pentachloride which exists as neutral PCl5 molecules in the gas and liquid states but adopts the ionic form [PCl4+][PCl6−] in the solid state. The average bond lengths in the crystal structure of POCl3 are 1.98 Å for P–Cl and 1.46 Å for P=O.[4]
Physical properties
With a freezing point of 1 °C and boiling point of 106 °C, the liquid range of POCl3 is rather similar to water. Also like water, POCl3 autoionizes, owing to the reversible formation of POCl2+,Cl−.
Chemical properties
POCl3 reacts with water to give hydrogen chloride and phosphoric acid:
- O=PCl3 + 3 H2O → O=P(OH)3 + 3 HCl
Intermediates in the conversion have been isolated, including pyrophosphoryl chloride, P2O3Cl4.[7]
Upon treatment with excess alcohols and phenols, POCl3 gives phosphate esters:
- O=PCl3 + 3 ROH → O=P(OR)3 + 3 HCl
Such reactions are often performed in the presence of an HCl acceptor such as pyridine or an amine.
POCl3 can also act as a Lewis base, forming adducts with a variety of Lewis acids such as titanium tetrachloride:
- Cl3PO + TiCl4 → Cl3POTiCl4
The aluminium chloride adduct (POCl3·AlCl3) is quite stable, and so POCl3 can be used to remove AlCl3 from reaction mixtures, for example at the end of a Friedel-Crafts reaction.
POCl3 reacts with hydrogen bromide in the presence of Lewis-acidic catalysts to produce POBr3.
Preparation
Phosphoryl chloride can be prepared by many methods. Phosphoryl chloride was first reported in 1847 by the French chemist Adolphe Wurtz by reacting phosphorus pentachloride with water.[8]
By oxidation
The commercial method involves oxidation of phosphorus trichloride with oxygen:[9]
- 2 PCl3 + O2 → 2 POCl3
An alternative method involves the oxidation of phosphorus trichloride with potassium chlorate:[10]
- 3 PCl3 + KClO3 → 3 POCl3 + KCl
Oxygenations
The reaction of phosphorus pentachloride (PCl5) with phosphorus pentoxide (P4O10).
- 6 PCl5 + P4O10 → 10 POCl3
The reaction can be simplified by chlorinating a mixture of PCl3 and P4O10, generating the PCl5 in situ. The reaction of phosphorus pentachloride with boric acid or oxalic acid:[10]
Other methods
Reduction of tricalcium phosphate with carbon in the presence of chlorine gas:[11]
The reaction of phosphorus pentoxide with sodium chloride is also reported:[11]
- 2 P2O5 + 3 NaCl → 3 NaPO3 + POCl3.
Uses
In one commercial application, phosphoryl chloride is used in the manufacture of phosphate esters. Triarylphosphates such as triphenyl phosphate and tricresyl phosphate are used as flame retardants and plasticisers for PVC. Trialkylphosphates such as tributyl phosphate are used as liquid–liquid extraction solvents in nuclear reprocessing and elsewhere.[9]
In the semiconductor industry, POCl3 is used as a safe liquid phosphorus source in diffusion processes. The phosphorus acts as a dopant used to create n-type layers on a silicon wafer.
As a reagent
In the laboratory, POCl3 is a reagent in dehydrations. One example involves conversion of formamides to isonitriles (isocyanides);[12] primary amides to nitriles:[13]
- RC(O)NH2 + POCl3 → RCN + "PO2Cl" + 2 HCl
In a related reaction, certain aryl-substituted amides can be cyclized using the Bischler-Napieralski reaction.
Such reactions are believed to proceed via an imidoyl chloride. In certain cases, the imidoyl chloride is the final product. For example, pyridones and pyrimidones can be converted to chloro- derivatives such as 2-chloropyridines and 2-chloropyrimidines, which are intermediates in the pharmaceutical industry.[14]
In the Vilsmeier-Haack reaction, POCl3 reacts with amides to produce a "Vilsmeier reagent", a chloro-iminium salt, which subsequently reacts with electron-rich aromatic compounds to produce aromatic aldehydes upon aqueous work-up.[15]
References
- Nomenclature of Organic Chemistry: IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. p. 926. doi:10.1039/9781849733069-FP001. ISBN 978-0-85404-182-4.
- NIOSH Pocket Guide to Chemical Hazards. "#0508". National Institute for Occupational Safety and Health (NIOSH).
- Toy, Arthur D. F. (1973). The Chemistry of Phosphorus. Oxford: Pergamon Press. ISBN 9780080187808. OCLC 152398514.
- Olie, K. (1971). "The crystal structure of POCl3". Acta Crystallogr. B. 27: 1459–1460. doi:10.1107/S0567740871004138.
- Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann.
- Chesnut, D. B.; Savin, A. (1999). "The Electron Localization Function (ELF) Description of the PO Bond in Phosphine Oxide". Journal of the American Chemical Society. 121 (10): 2335–2336. doi:10.1021/ja984314m. ISSN 0002-7863.
- Grunze, Herbert (1963). "Über die Hydratationsprodukte des Phosphoroxychlorides. III. Darstellung von Pyrophosphorylchlorid aus partiell hydrolysiertem Phosphoroxychlorid (Hydration products of phosphorus oxychloride. III. Preparation of pyrophosphoryl chloride from partially hydrolyzed phosphorus oxychloride)". Zeitschrift fuer Anorganische und Allgemeine Chemie. 324: 1–14. doi:10.1002/zaac.19633240102.
- Wurtz, Adolphe (1847). "Sur l'acide sulfophosphorique et le chloroxyde de phosphore" [On monothiophosphoric acid and phosphoryl chloride]. Annales de Chimie et de Physique. 3rd series (in French). 20: 472–481.; see Chloroxyde de phosphore, pp. 477–481. (Note: Wurtz's empirical formulas are wrong because, like many chemists of his day, he used the wrong atomic mass for oxygen.)Roscoe, Henry Enfield; Schorlemmer, Carl; Cannell, John, eds. (1920). A Treatise on Chemistry. vol. 1 (5th ed.). London, England: Macmillan and Co. p. 676.
- Bettermann, Gerhard; Krause, Werner; Riess, Gerhard; Hofmann, Thomas (2000). "Phosphorus Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a19_527..
- Pradyot, Patnaik (2003). Handbook of Inorganic Chemicals. New York: McGraw-Hill. p. 709. ISBN 0070494398.
- Lerner, Leonid (2011). Small-Scale Synthesis of Laboratory Reagents with Reaction Modeling. Boca Raton, Florida: CRC Press. pp. 169–177. ISBN 9781439813126.
- Patil, Pravin; Ahmadian-Moghaddam, Maryam; Dömling, Alexander (29 September 2020). "Isocyanide 2.0". Green Chemistry. 22 (20): 6902–6911. doi:10.1039/D0GC02722G.
- March, J. (1992). Advanced Organic Chemistry (4th ed.). New York, NY: Wiley. p. 723.
- Elderfield, R. C. (ed.). Heterocyclic Compound. 6. New York, NY: John Wiley & Sons. p. 265.
- Hurd, Charles D.; Webb, Carl N. (1925). "p-Dimethylaminobenzophenone". Organic Syntheses. 7: 24. doi:10.15227/orgsyn.007.0024.
Further reading
- Handbook of Chemistry and Physics (71st ed.). Ann Arbor, MI: CRC Press. 1990.
- Stecher, Paul G. (1960). The Merck Index of Chemicals and Drugs (7th ed.). Rahway: Merck & Co. OCLC 3653550.
- Wade, L. G., Jr (2005). Organic Chemistry (6th ed.). Upper Saddle River, NJ: Pearson/Prentice Hall. p. 477.
- Walker, B. J. (1972). Organophosphorus Chemistry. Harmondsworth: Penguin. pp. 101–116.
- "CDC – NIOSH Pocket Guide to Chemical Hazards".