Phosphorus pentachloride is the chemical compound with the formula PCl5. It is one of the most important phosphorus chlorides/oxychlorides, others being PCl3 and POCl3. PCl5 finds use as a chlorinating reagent. It is a colourless, water-sensitive solid, although commercial samples can be yellowish and contaminated with hydrogen chloride.
The structures for the phosphorus chlorides are invariably consistent with VSEPR theory. The structure of PCl5 depends on its environment. Gaseous and molten PCl5 is a neutral molecule with trigonal bipyramidal geometry and (D3h) symmetry. The hypervalent nature of this species (as well as of PCl− 6, see below) can be explained with the inclusion of non-bonding molecular orbitals (molecular orbital theory) or resonance (valence bond theory). This trigonal bipyramidal structure persists in nonpolar solvents, such as CS2 and CCl4.[5] In the solid state PCl5 is an ionic compound, formulated PCl+ 4PCl− 6.[6]
Structure of solid phosphorus pentachloride, illustrating its autoionization at higher concentrations.[7]
In solutions of polar solvents, PCl5 undergoes self-ionization.[8] Dilute solutions dissociate according to the following equilibrium:
PCl5 ⇌ PCl+ 4 + Cl−
At higher concentrations, a second equilibrium becomes more prevalent:
2 PCl5 ⇌ PCl+ 4 + PCl− 6
The cation PCl+ 4 and the anion PCl− 6 are tetrahedral and octahedral, respectively. At one time, PCl5 in solution was thought to form a dimeric structure, P2Cl10, but this suggestion is not supported by Raman spectroscopic measurements.
AsCl5 and SbCl5 also adopt trigonal bipyramidal structures. The relevant bond distances are 211 pm (As−Cleq), 221 pm (As−Clax), 227 pm (Sb−Cleq), and 233.3 pm (Sb−Clax).[9] At low temperatures, SbCl5 converts to the dimer, dioctahedral Sb2Cl10, structurally related to niobium pentachloride.
PCl5 is prepared by the chlorination of PCl3.[10] This reaction is used to produce around 10,000 tonnes of PCl5 per year (as of 2000).[6]
PCl3 + Cl2 ⇌ PCl5 (ΔH = −124 kJ/mol)
PCl5 exists in equilibrium with PCl3 and chlorine, and at 180 °C the degree of dissociation is about 40%.[6] Because of this equilibrium, samples of PCl5 often contain chlorine, which imparts a greenish coloration.
Phosphorus pentachloride is a Lewis acid. This property underpins many of its characteristic reactions, autoionization, chlorinations, hydrolysis. A well studied adduct is PCl5(pyridine).[11]
In synthetic chemistry, two classes of chlorination are usually of interest: oxidative chlorinations and substitutive chlorinations. Oxidative chlorinations entail the transfer of Cl2 from the reagent to the substrate. Substitutive chlorinations entail replacement of O or OH groups with chloride. PCl5 can be used for both processes.
PCl5 reacts with a tertiary amides, such as dimethylformamide (DMF), to give dimethylchloromethyleneammonium chloride, which is called the Vilsmeier reagent, [(CH3)2N=CClH]Cl. More typically, a related salt is generated from the reaction of DMF and POCl3. Such reagents are useful in the preparation of derivatives of benzaldehyde by formylation and for the conversion of C−OH groups into C−Cl groups.[14]
Both PCl3 and PCl5 convert R3COH groups to the chloride R3CCl. The pentachloride is however a source of chlorine in many reactions. It chlorinates allylic and benzylic CH bonds. PCl5 bears a greater resemblance to SO2Cl2, also a source of Cl2. For oxidative chlorinations on the laboratory scale, sulfuryl chloride is often preferred over PCl5 since the gaseous SO2 by-product is readily separated.
As for the reactions with organic compounds, the use of PCl5 has been superseded by SO2Cl2. The reaction of phosphorus pentoxide and PCl5 produces POCl3 :[18][page needed]
Phosphorus pentachloride was first prepared in 1808 by the English chemist Humphry Davy.[20] Davy's analysis of phosphorus pentachloride was inaccurate;[21] the first accurate analysis was provided in 1816 by the French chemist Pierre Louis Dulong.[22]
^ abcPhosphorus pentachloride in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemistry WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD) (retrieved 2014-05-15)
^Corbridge, D. E. C. (1995). Phosphorus: An outline of its chemistry, biochemistry, and uses. Elsevier Science. ISBN0-444-89307-5.
^ abcHolleman, A. F.; Wiber, E.; Wiberg, N. (2001). Inorganic Chemistry. Academic Press. ISBN978-0-12-352651-9.
^Finch, A.; Fitch, A.N.; Gates, P.N. (1993). "Crystal and Molecular structure of a metastable modification of phosphorus pentachloride". Journal of the Chemical Society, Chemical Communications (11): 957–958. doi:10.1039/C39930000957.
^Burks Jr., J. E. (2004). "Phosphorus(V) chloride". In Paquette, L. (ed.). Encyclopedia of Reagents for Organic Synthesis. New York, NY: J. Wiley & Sons. doi:10.1002/047084289X.rp158. ISBN0471936235.
^Spaggiari, A.; Vaccari, D.; Davoli, P.; Torre, G.; Prati, F. (2007). "A Mild Synthesis of Vinyl Halides and gem-Dihalides Using Triphenyl Phosphite−Halogen-Based Reagents". The Journal of Organic Chemistry. 72 (6): 2216–2219. doi:10.1021/jo061346g. ISSN0022-3263. PMID17295542.
^Cotton, Frank Albert (1999). Advanced Inorganic Chemistry. Wiley-Interscience. ISBN978-0-471-19957-1.
^Bushkova, O. V.; Yaroslavtseva, T. V.; Dobrovolsky, Yu. A. (4 August 2017). "New lithium salts in electrolytes for lithium-ion batteries (Review)". Russian Journal of Electrochemistry. 53 (7): 677–699. doi:10.1134/S1023193517070035. S2CID103854243.
^Dulong (1816). "Extrait d'un mémoire sur les combinaisons du phosphore avec l'oxigène" [Extract from a memoir on the compounds of phosphorus with oxygen]. Annales de Chimie et de Physique. 2nd series (in French). 2: 141–150. On p. 148, Dulong presented the correct analysis of phosphorus pentachloride (which is 14.9% phosphorus and 85.1% chlorine by weight, vs. Dulong's values of 15.4% and 84.6%, respectively).
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