Chlorine trifluoride oxide

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Chlorine trifluoride oxide
Chlorine-trifluoride-oxide-3D-vdW.png
Chlorine-trifluoride-oxide-3D-balls.png
Names
IUPAC name
trifluoro(oxo)-λ5-chlorane
Identifiers
3D model (JSmol)
  • InChI=1S/ClF3O/c2-1(3,4)5
    Key: QPKQQPNQDSQNHS-UHFFFAOYSA-N
  • O=Cl(F)(F)F
Properties
ClF3O
Molar mass 108.44 g·mol−1
Density 1.865
Melting point −42 °C (−44 °F; 231 K)
Boiling point 29 °C (84 °F; 302 K)
Structure
monoclinic
C2/m
a = 9.826, b = 12.295, c = 4.901
α = 90°, β = 90.338°, γ = 90°[2]
592.1
8
Hazards
GHS labelling:
GHS03: OxidizingGHS05: CorrosiveGHS06: ToxicGHS09: Environmental hazard
Danger
Related compounds
Related compounds
BrOF3; IOF3
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Chlorine oxide trifluoride or chlorine trifluoride oxide is a corrosive liquid molecular compound with formula ClOF3. It was developed secretly as a rocket fuel oxidiser.

Production[edit]

Chlorine oxide trifluoride was originally made at Rocketdyne[3] by treating chlorine monoxide with fluorine. Other substances that could react with fluorine to make it includes sodium chlorite (NaClO2), and chlorine nitrate (ClONO2). The first published production method was a reaction of chlorine monoxide with oxygen difluoride (OF2). Yet other production methods are reactions between ClO2F or ClO3F and chlorine fluorides.[4] A safer approach is the use chlorine nitrate with fluorine.

Reactions[edit]

As a Lewis base it can lose a fluoride ion to Lewis acids, yielding the difluorooxychloronium(V) cation (ClOF2+).[5] Compounds with this include: ClOF2BF4, ClOF2PF6, ClOF2AsF6, ClOF2SbF6, ClOF2BiF6, ClOF2VF6, ClOF2NbF6, ClOF2TaF6, ClOF2UF6, ClOF2, (ClOF2)2SiF6, ClOF2MoOF5, ClOF2Mo2O4F9,[4] ClOF2PtF6.[6]

Functioning as a Lewis acid, it can gain a fluoride ion from a strong base to yield a tetrafluorooxychlorate(V) anion: ClOF4 ion.[7] These include KClOF4, RbClOF4, and CsClOF4.[8] This allows purification of ClOF3, as at room temperature a solid complex is formed, but this decomposes between 50 and 70 °C. Other likely impurities either will not react with alkali fluoride, or if they do will not easily decompose.[3]

Chlorine trifluoride oxide fluoridates various materials such as chlorine monoxide, chlorine, glass or quartz.[3] ClOF3 + Cl2O → 2ClF + ClO2F;[6] 2ClOF3 + 2Cl2 → 6ClF + O2 at 200 °C[6]

Chlorine trifluoride oxide adds to chlorine fluorosulfate, ClOF3 + 2ClOSO2F → S2O5F2 + FClO2 + 2ClF. The reaction also produces SO2F2.[3]

Chlorine trifluoride oxide can fluoridate and add oxygen in the same reaction, reacting with molybdenum pentafluoride, silicon tetrafluoride, tetrafluorohydrazine (over 100 °C), HNF2, and F2NCOF. From HNF2 the main result was NF3O. From MoF5, the results were MoF6 and MoOF4.[3]

It reacts explosively with hydrocarbons.[3] With small amounts of water, ClO2F is formed along with HF.[3]

Over 280 °C ClOF3 decomposes to oxygen and chlorine trifluoride.[3]

Properties[edit]

The boiling point of chlorine trifluoride oxide is 29 °C.[9]

The shape of the molecule is a trigonal bipyramid, with two fluorine atoms at the top and bottom (apex) (Fa) and an electron pair, oxygen and fluorine (Fe) on the equator.[7] The Cl=O bond length is 1.405 Å, Cl-Fe 1.603 Å, other Cl-Fa 1.713 Å, ∠FeClO=109° ∠FaClO=95°, ∠FaClFe=88°. The molecule is polarised, Cl has a +1.76 charge, O has −0.53, equatorial F has −0.31 and apex F has −0.46. The total dipole moment is 1.74 D.[10]

References[edit]

  1. ^ Urben, Peter (2017). Bretherick's Handbook of Reactive Chemical Hazards. Elsevier. p. 784. ISBN 9780081010594.
  2. ^ Ellern, Arkady; Boatz, Jerry A.; Christe, Karl O.; Drews, Thomas; Seppelt, Konrad (September 2002). "The Crystal Structures of ClF3O, BrF3O, and [NO]+[BrF4O]". Zeitschrift für anorganische und allgemeine Chemie. 628 (9–10): 1991–1999. doi:10.1002/1521-3749(200209)628:9/10<1991::AID-ZAAC1991>3.0.CO;2-1.
  3. ^ a b c d e f g h Advances in Inorganic Chemistry and Radiochemistry. Academic Press. 1976. pp. 331–333. ISBN 9780080578675.
  4. ^ a b Holloway, John H.; Laycock, David (1983). Advances in Inorganic Chemistry. Academic Press. pp. 178–179. ISBN 9780080578767.
  5. ^ Christe, Karl O.; Curtis, E. C.; Schack, Carl J. (September 1972). "Chlorine trifluoride oxide. VII. Difluorooxychloronium(V) cation, ClF2O+. Vibrational spectrum and force constants". Inorganic Chemistry. 11 (9): 2212–2215. doi:10.1021/ic50115a046.
  6. ^ a b c Schack, Carl J.; Lindahl, C. B.; Pilipovich, Donald.; Christe, Karl O. (September 1972). "Chlorine trifluoride oxide. IV. Reaction chemistry". Inorganic Chemistry. 11 (9): 2201–2205. doi:10.1021/ic50115a043.
  7. ^ a b Christe, K.O.; Schack, C.J. (1976). Chlorine Oxyfluorides. Advances in Inorganic Chemistry and Radiochemistry. Vol. 18. pp. 319–398. doi:10.1016/S0065-2792(08)60033-3. ISBN 9780120236183.
  8. ^ Christe, Karl O.; Schack, Carl J.; Pilipovich, Donald.; Christe, Karl O. (September 1972). "Chlorine trifluoride oxide. V. Complex formation with Lewis acids and bases". Inorganic Chemistry. 11 (9): 2205–2208. doi:10.1021/ic50115a044.
  9. ^ Pilipovich, Donald.; Lindahl, C. B.; Schack, Carl J.; Wilson, R. D.; Christe, Karl O. (September 1972). "Chlorine trifluoride oxide. I. Preparation and properties". Inorganic Chemistry. 11 (9): 2189–2192. doi:10.1021/ic50115a040.
  10. ^ Oberhammer, Heinz.; Christe, Karl O. (January 1982). "Gas-phase structure of chlorine trifluoride oxide, ClF3O". Inorganic Chemistry. 21 (1): 273–275. doi:10.1021/ic00131a050.