Chemistry:Bromine trifluoride

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Bromine trifluoride
Structural formula, showing bond lengths and angles
Bromine Trifluoride
3D model (JSmol)
EC Number
  • 232-132-1
UN number 1746
Molar mass 136.90 g/mol
Appearance straw-coloured liquid
Odor Choking, pungent[1]
Density 2.803 g/cm3 [2]
Melting point 8.77 °C (47.79 °F; 281.92 K)
Boiling point 125.72 °C (258.30 °F; 398.87 K)
Reacts with water[3]
Solubility in sulfuric acid very soluble
T-shaped (C2v)
1.19 D
Main hazards dangerously sensitive to water, source of HF
Safety data sheet
GHS pictograms GHS03: OxidizingGHS05: CorrosiveGHS06: ToxicGHS08: Health hazard
GHS Signal word Danger
H271, H330, H314, H373
P102, P103, P210, P220, P221, P260, P264, P271, P280, P283, P284, P301+310, P301+330+331, P303+361+353, P304+312, P306+360, P308+313, P370+380, P340, P363, P305+351+338+310
NFPA 704 (fire diamond)
Related compounds
Other anions
Bromine monochloride
Other cations
Chlorine trifluoride
Iodine trifluoride
Related compounds
Bromine monofluoride
Bromine pentafluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Bromine trifluoride is an interhalogen compound with the formula BrF3. It is a straw-coloured liquid with a pungent odor.[5] It is soluble in sulfuric acid but reacts violently with water and organic compounds. It is a powerful fluorinating agent and an ionizing inorganic solvent. It is used to produce uranium hexafluoride (UF6) in the processing and reprocessing of nuclear fuel.[6]


Bromine trifluoride was first described by Paul Lebeau in 1906, who obtained the material by the reaction of bromine with fluorine at 20 °C:[7]

Br2 + 3 F2 → 2 BrF3

The disproportionation of bromine monofluoride also gives bromine trifluoride:[5]

3 BrF → BrF3 + Br2


Like ClF3 and IF3, the BrF3 molecule is T-shaped and planar. In the VSEPR formalism, the bromine center is assigned two electron pairs. The distance from the bromine each axial fluorine is 1.81 Å and to the equatorial fluorine is 1.72 Å. The angle between an axial fluorine and the equatorial fluorine is slightly smaller than 90° — the 86.2° angle observed is due to the repulsion generated by the electron pairs being greater than that of the Br-F bonds.[8][9]

Chemical properties

BrF3 rapidly and exothermically reacts with water to release hydrobromic acid and hydrofluoric acid:

BrF3 + 2H2O → 3HF + HBr + O2

BrF3 is a fluorinating agent, but less reactive than ClF3[10] Already at -196 °C, it reacts with acetonitrile to give 1,1,1-trifluoroethane.[11]

BrF3 + CH3CN → CH3CF3 + 1/2 Br2 + 1/2 N2

The liquid is conducting, owing to autoionisation:[6]

2 BrF3 ⇌ BrF2+ + BrF4

Fluoride salts dissolve readily in BrF3 forming tetrafluorobromate:[6]

KF + BrF3 → KBrF4

It react as a fluoride donor:[12]

BrF3 + SbF5 → [BrF2+][SbF6]


  2. Lide, David R., ed (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3. 
  3. "Archived copy". 
  4. "Safety Data Sheet Bromine Trifluoride". Airgas. 
  5. 5.0 5.1 Simons JH (1950). "Bromine(III) Fluoride (Bromine Trifluoride)". Bromine (III) Fluoride - Bromine Trifluoride. Inorganic Syntheses. 3. pp. 184–186. doi:10.1002/9780470132340.ch48. ISBN 978-0-470-13234-0. 
  6. 6.0 6.1 6.2 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. 
  7. Lebeau P. (1906). "The effect of fluorine on chloride and on bromine". Annales de Chimie et de Physique 9: 241–263. 
  8. Gutmann V (1950). "Die Chemie in Bromitrifluorid". Angewandte Chemie 62 (13–14): 312–315. doi:10.1002/ange.19500621305. 
  9. Meinert H (1967). "Interhalogenverbindungen". Zeitschrift für Chemie 7 (2): 41–57. doi:10.1002/zfch.19670070202. 
  10. Rozen, Shlomo; Sasson, Revital (2007). "Bromine Trifluoride". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/9780470842898.rb266.pub2. ISBN 978-0471936237. 
  11. Rozen, Shlomo (2010). "Selective Reactions of Bromine Trifluoride in Organic Chemistry". Advanced Synthesis & Catalysis 352 (16): 2691–2707. doi:10.1002/adsc.201000482. 
  12. A. J. Edwards and G. R. Jones. J. Chem. Soc. A, 1467 (1969)

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