Phosphine

Philippe Gengembre (1764–1838), a student of Lavoisier, first obtained phosphine in 1783 by heating phosphorus in an aqueous solution of potash (potassium carbonate).[4][5]

Perhaps because of its strong association with elemental phosphorus, phosphine was once regarded as a gaseous form of the element, but Lavoisier (1789) recognised it as a combination of phosphorus with hydrogen and described it as phosphure d'hydrogène (phosphide of hydrogen).[6]

In 1844, Paul Thénard, son of the French chemist Louis Jacques Thénard, used a cold trap to separate diphosphine from phosphine that had been generated from calcium phosphide, thereby demonstrating that P₂H₄ is responsible for spontaneous flammability associated with PH₃, and also for the characteristic orange/brown color that can form on surfaces, which is a polymerisation product.[7] He considered diphosphine's formula to be PH₂, and thus an intermediate between elemental phosphorus, the higher polymers, and phosphine. Calcium phosphide (nominally Ca₃P₂) produces more P₂H₄ than other phosphides because of the preponderance of P-P bonds in the starting material.

The name "phosphine" first appeared in combined form in 1857.[8] The gas PH₃ was named "phosphine" by 1865 (or earlier).[9]

PH₃ is a trigonal pyramidal molecule with C_3v molecular symmetry. The length of the P-H bond is 1.42 Å, the H-P-H bond angles are 93.5°. The dipole moment is 0.58 D, which increases with substitution of methyl groups in the series: CH₃PH₂, 1.10 D; (CH₃)₂PH, 1.23 D; (CH₃)₃P, 1.19 D. In contrast, the dipole moments of amines decrease with substitution, starting with ammonia, which has a dipole moment of 1.47 D. The low dipole moment and almost orthogonal bond angles lead to the conclusion that in PH₃ the P-H bonds are almost entirely pσ(P) – sσ(H) and phosphorus 3s orbital contributes little to the bonding between phosphorus and hydrogen in this molecule. For this reason, the lone pair on phosphorus may be regarded as predominantly formed by the 3s orbital of phosphorus. The upfield chemical shift of the phosphorus atom in the ³¹P NMR spectrum accords with the conclusion that the lone pair electrons occupy the 3s orbital (Fluck, 1973). This electronic structure leads to a lack of nucleophilicity in general and lack of basicity in particular (pK_aH = –14),[10] as well as an ability to form only weak hydrogen bonds.[11]

The aqueous solubility of PH₃ is slight; 0.22 mL of gas dissolve in 1 mL of water. Phosphine dissolves more readily in non-polar solvents than in water because of the non-polar P-H bonds. It is technically amphoteric in water, but acid and base activity is poor. Proton exchange proceeds via a phosphonium (PH₄⁺) ion in acidic solutions and via PH₂⁻ at high pH, with equilibrium constants K_b = 4 × 10⁻²⁸ and K_a = 41.6 × 10⁻²⁹.

Phosphine burns producing a dense white cloud of phosphoric acid:

PH₃ + 2 O₂ → H₃PO₄

Phosphine may be prepared in a variety of ways.[12] Industrially it can be made by the reaction of white phosphorus with sodium or potassium hydroxide, producing potassium or sodium hypophosphite as a by-product.

3 KOH + P₄ + 3 H₂O → 3 KH₂PO₂ + PH₃

Alternatively the acid-catalyzed disproportioning of white phosphorus yields phosphoric acid and phosphine. Both routes have industrial significance; the acid route is preferred method if further reaction of the phosphine to substituted phosphines is needed. The acid route requires purification and pressurizing. It can also be made (as described above) by the hydrolysis of a metal phosphide, such as aluminium phosphide or calcium phosphide. Pure samples of phosphine, free from P₂H₄, may be prepared using the action of potassium hydroxide on phosphonium iodide (PH₄I).

Organophosphines are compounds with the formula PR_nH_3−n. These compounds are often classified according to the value of n: primary phosphines (n = 1), secondary phosphines (n = 2), tertiary phosphines (n = 3). All adopt pyramidal structures. Their reactivity is also similar – they can be oxidized to the phosphorus(V) level, they can be protonated and alkylated at phosphorus to give phosphonium salts, and, for primary and secondary derivatives, they can be deprotonated by strong bases to give organophosphide derivatives.

Secondary phosphines

Secondary phosphines are prepared analogously to the primary phosphines. They are also obtained by alkali-metal reductive cleavage of triarylphosphines followed by hydrolysis of the resulting phosphide salt. The latter route is employed to prepare diphenylphosphine (Ph₂PH). Diorganophosphinic acids, R₂P(O)OH, can also be reduced with diisobutylaluminium hydride. Secondary phosphines are typically protic in character. When however modified with suitable substituents as in certain (rare) diazaphospholenes (scheme 3) the polarity of the P-H bond can be inverted (see: umpolung) and the resulting phosphine hydride can reduce a carbonyl group as in the example of benzophenone in yet another way.[16]

Scheme 3. diazaphospholene phosphine hydride

Phosphine gas is denser than air and hence may collect in low-lying areas. It can form explosive mixtures with air and also self-ignite.

Phosphine can be absorbed into the body by inhalation. Direct contact with phosphine liquid – although unlikely to occur – may cause frostbite, like other cryogenic liquids. The main target organ of phosphine gas is the respiratory tract.[19] According to the 2009 U.S. National Institute for Occupational Safety and Health (NIOSH) pocket guide, and U.S. Occupational Safety and Health Administration (OSHA) regulation, the 8 hour average respiratory exposure should not exceed 0.3 ppm. NIOSH recommends that the short term respiratory exposure to phosphine gas should not exceed 1 ppm. The Immediately Dangerous to Life or Health level is 50 ppm. Overexposure to phosphine gas causes nausea, vomiting, abdominal pain, diarrhea, thirst, chest tightness, dyspnea (breathing difficulty), muscle pain, chills, stupor or syncope, and pulmonary edema.[20][21] Phosphine has been reported to have the odor of decaying fish or garlic at concentrations below 0.3 ppm. The smell is normally restricted to laboratory areas or phosphine processing since the smell comes from the way the phosphine is extracted from the environment. However, it may occur elsewhere, such as in industrial waste landfills. Exposure to higher concentrations may cause olfactory fatigue.[22]

Deaths have resulted from accidental exposure to fumigation materials containing aluminum phosphide or phosphine.[23][24][25][26] It can be absorbed either by inhalation or transdermally.[23] As a respiratory poison, it affects the transport of oxygen or interferes with the utilization of oxygen by various cells in the body.[25] Exposure results in pulmonary edema (the lungs fill with fluid).[26] Phosphine gas is heavier than air so stays nearer the floor where children are likely to play.[27]

In the pilot episode of Breaking Bad, Walter White killed Krazy-8 and Emilio Koyama by generating phosphine gas.

• Diphosphane, H₂PPH₂, simplified to H₄P₂
  • Diphosphines, R₂PPR₂, R₂P(CH₂)_nPR₂
• Diphosphene, HP=PH
  • Diphosphenes, R-P=P-R'
• Phosphine oxide, R₃P=O
• Phosphorane, PR₅, R₃P=CR₂
• Phosphinite, P(OR)R₂
• Phosphonite, P(OR)₂R
• Phosphite, P(OR)₃
• Phosphinate, R₂P(RO)O
• Phosphonate, RP(RO)₂O
• Phosphate, P(RO)₃O

• Fluck, E. (1973). "The Chemistry of Phosphine". Topics in Current Chemistry. Fortschritte der Chemischen Forschung. 35: 1–64. doi:10.1007/BFb0051358. ISBN 3-540-06080-4.
• WHO (World Health Organisation) (1988). Phosphine and Selected Metal Phosphides. Environmental Health Criteria. 73. Geneva: Joint sponsorship of UNEP, ILO and WHO.

• International Chemical Safety Card 0694
• CDC – Phosphine – NIOSH Workplace Safety and Health Topic