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Boron trifluoride

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Title: Boron trifluoride  
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Subject: Boron, Diborane, Lithium tetrafluoroborate, List of UN numbers 1701 to 1800, Chemical polarity
Collection: Boron Compounds, Boron Halides, Fluorides
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Boron trifluoride

Boron trifluoride
Boron trifluoride in 2D
Boron trifluoride in 3D
Other names
Boron fluoride, Trifluoroborane
(dihydrate) Y
ChemSpider  Y
EC number 231-569-5
Jmol-3D images Image
RTECS number ED2275000
UN number Compressed: 1008.
Boron trifluoride dihydrate: 2851.
Molar mass 67.82 g/mol (anhydrous)
103.837 g/mol (dihydrate)
Appearance colorless gas (anhydrous)
colorless liquid (dihydrate)
Density 0.00276 g/cm3 (anhydrous gas)
1.64 g/cm3 (dihydrate)
Melting point −126.8 °C (−196.2 °F; 146.4 K)
Boiling point −100.3 °C (−148.5 °F; 172.9 K)
exothermic decomposition [1] (anhydrous)
very soluble (dihydrate)
Solubility soluble in benzene, toluene, hexane, chloroform and methylene chloride
Vapor pressure >50 atm (20°C)[2]
0 D
50.46 J/mol K
254.3 J/mol K
-1137 kJ/mol
-1120 kJ/mol
Safety data sheet ICSC
GHS pictograms Press. GasAcute Tox. 2Skin Corr. 1A
GHS signal word DANGER
H330, H314 [note 1]
Very toxic (T+)
Corrosive (C)
R-phrases R14, R26, R35
S-phrases (S1/2), S9, S26, S28, S36/37/39, S45
NFPA 704
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
1227 ppm (mouse, 2 hr)
39 ppm (guinea pig, 4 hr)
418 ppm (rat, 4 hr)[5]
US health exposure limits (NIOSH):
PEL (Permissible)
C 1 ppm (3 mg/m3)[2]
REL (Recommended)
C 1 ppm (3 mg/m3)[2]
25 ppm[2]
Related compounds
Other anions
Boron trichloride
Boron tribromide
Boron triiodide
Other cations
Aluminium fluoride
Gallium(III) fluoride
Indium(III) fluoride
Thallium(III) fluoride
Related compounds
Boron monofluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
 Y  (: Y/N?)

Boron trifluoride is the formula BF3. This pungent colourless toxic gas forms white fumes in moist air. It is a useful Lewis acid and a versatile building block for other boron compounds.


  • Structure and bonding 1
  • Synthesis and handling 2
  • Reactions 3
    • Comparative Lewis acidity 3.1
    • Hydrolysis 3.2
  • Uses 4
    • Niche uses 4.1
  • Discovery 5
  • Notes 6
  • References 7
  • External links 8

Structure and bonding

The geometry of a molecule of BF3 is trigonal planar. Its D3h symmetry conforms with the prediction of VSEPR theory. The molecule has no dipole moment by virtue of its high symmetry. The molecule is isoelectronic with the carbonate anion, CO32−.

BF3 is commonly referred to as "electron deficient," a description that is reinforced by its exothermic reactivity toward Lewis bases.

In the boron trihalides, BX3, the length of the B-X bonds (1.30 Å) is shorter than would be expected for single bonds,[7] and this shortness may indicate stronger B-X π-bonding in the fluoride. A facile explanation invokes the symmetry-allowed overlap of a p orbital on the boron atom with the in-phase combination of the three similarly oriented p orbitals on fluorine atoms.[7]

Boron trifluoride pi bonding diagram

Synthesis and handling

BF3 is manufactured by the reaction of boron oxides with hydrogen fluoride:

B2O3 + 6 HF → 2 BF3 + 3 H2O

Typically the HF is produced in situ from sulfuric acid and fluorite (CaF2).[8] Approximately 2300-4500 tonnes of boron trifluoride are produced every year.[9]

On a laboratory scale, BF3 is produced by the thermal decomposition of diazonium salts:[10]

PhN2BF4PhF + BF3 + N2

Alternatively the chemical can be synthesized from [11]

6 NaBF4 + B2O3 + 6 H2SO4 → 8 BF3 + 6 NaHSO4 + 3 H2O

Anhydrous boron trifluoride has a boiling point of −100.3 C and a critical temperature of −12.3 C, so that it can be stored as a refrigerated liquid only between those temperatures. Storage or transport vessels should be designed to withstand internal pressure, since a refrigeration system failure could cause pressures to rise to the critical pressure of 49.85 bar (4.985 MPa).[12]

Boron trifluoride is corrosive. Suitable metals for equipment handling boron trifluoride include stainless steel, monel, and hastelloy. In presence of moisture it corrodes steel, including stainless steel. It reacts with polyamides. Polytetrafluoroethylene, polychlorotrifluoroethylene, polyvinylidene fluoride, and polypropylene show satisfactory resistance. The grease used in the equipment should be fluorocarbon based, as boron trifluoride reacts with the hydrocarbon-based ones.[13]


Unlike the aluminium and gallium trihalides, the boron trihalides are all monomeric. They undergo rapid halide exchange reactions:

BF3 + BCl3 → BF2Cl + BCl2F

Because of the facility of this exchange process, the mixed halides cannot be obtained in pure form.

Boron trifluoride is a versatile Lewis acid that forms adducts with such Lewis bases as fluoride and ethers:

CsF + BF3 → CsBF4
O(C2H5)2 + BF3 → BF3O(C2H5)2

Tetrafluoroborate salts are commonly employed as non-coordinating anions. The adduct with diethyl ether, boron trifluoride diethyl etherate or just boron trifluoride etherate (BF3 · O(Et)2) is a conveniently handled liquid and consequently is a widely encountered as a laboratory source of BF3. It is stable as a solution in ether, but not stoichiometrically. Another common adduct is the adduct with dimethyl sulfide (BF3 · S(Me)2), which can be handled as a neat liquid.

Comparative Lewis acidity

All three lighter boron trihalides, BX3 (X = F, Cl, Br) form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicities of the adduct-forming reaction. Such measurements have revealed the following sequence for the Lewis acidity:

BF3 < BCl3 < BBr3 (strongest Lewis acid)

This trend is commonly attributed to the degree of π-bonding in the planar boron trihalide that would be lost upon pyramidalization of the BX3 molecule.[14] which follows this trend:

BF3 > BCl3 > BBr3 (most easily pyramidalized)

The criteria for evaluating the relative strength of π-bonding are not clear, however.[7] One suggestion is that the F atom is small compared to the larger Cl and Br atoms, and the lone pair electron in pz of F is readily and easily donated and overlapped to empty pz orbital of boron. As a result, the pi donation of F is greater than that of Cl or Br.

In an alternative explanation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the adducts F3B-L.[15][16]


Boron trifluoride reacts with water to give boric acid and fluoroboric acid. The reaction commences with the formation of the aquo adduct, H2O-BF3, which then loses HF that gives fluoboric acid with boron trifluoride.[17]

4 BF3 + 3 H2O → 3 HBF4 + "B(OH)3"

The heavier trihalides do not undergo analogous reactions, possibly due to the lower stability of the tetrahedral ions BX4 (X = Cl, Br). Because of the high acidity of fluoroboric acid, the fluoroborate ion can be used to isolate particularly electrophilic cations, such as diazonium ions, that are otherwise difficult to isolate as solids.


Boron trifluoride is most importantly used as a reagent in

  • "Safety and Health Topics: Boron Trifluoride". OSHA. 
  • "BORON TRIFLUORIDE ICSC: 0231". International Chemical Safety Cards. CDC. 
  • "Boron & Compounds: Overview". National Pollutant Inventory. Australian Government. 
  • "Fluoride Compounds: Overview". National Pollutant Inventory. Australian Government. 
  • "Boron trifluoride". WebBook. NIST. 
  • ) Applications"3"Boron Trifluoride (BF. Honeywell. 

External links

  1. ^
  2. ^ a b c d "NIOSH Pocket Guide to Chemical Hazards #0062".  
  3. ^ Index no. 005-001-00-X of Annex VI, Part 3, to Regulation (EC) No 1272/2008 of the European Parliament and of the Council of 16 December 2008 on classification, labelling and packaging of substances and mixtures, amending and repealing Directives 67/548/EEC and 1999/45/EC, and amending Regulation (EC) No 1907/2006. OJEU L353, 31.12.2008, pp 1–1355 at p 341.
  4. ^ "Boron trifluoride", Pocket Guide to Chemical Hazards, U.S. Department of Health and Human Services (NIOSH) Publication No. 2005-149, Washington, DC: Government Printing Office, 2005,  .
  5. ^ "Boron trifluoride". Immediately Dangerous to Life and Health.  
  6. ^
  7. ^ a b c  
  8. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press.  
  9. ^ a b Brotherton, R. J.; Weber, C. J.; Guibert, C. R.; Little, J. L. (2005), "Boron Compounds",  
  10. ^ Flood, D. T. (1933). "Fluorobenzene".  
  11. ^ a b Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry Vol. 1, 2nd Ed. Newyork: Academic Press. p. 220 & 773.  
  12. ^ Yaws, C. L., ed. (1999). Chemical Properties Handbook. McGraw-Hill. p. 25. 
  13. ^ "Boron trifluoride". Gas Encyclopedia.  
  14. ^  
  15. ^ Boorman, P. M.; Potts, D. (1974). "Group V Chalcogenide Complexes of Boron Trihalides".  
  16. ^ Brinck, T.; Murray, J. S.; Politzer, P. (1993). "A Computational Analysis of the Bonding in Boron Trifluoride and Boron Trichloride and their Complexes with Ammonia".  
  17. ^ Wamser, C. A. (1951). "Equilibria in the System Boron Trifluoride–Water at 25°".  
  18. ^ Heaney, H. (2001). "Boron Trifluoride".  
  19. ^ a b ) Applications"3"Boron Trifluoride (BF.  
  20. ^ Gay-Lussac, J. L.; Thénard, L. J. (1809). "Sur l’acide fluorique". Annales de Chimie 69: 204–220. 
  21. ^ Gay-Lussac, J. L.; Thénard, L. J. (1809). "Des propriétés de l’acide fluorique et sur-tout de son action sur le métal de la potasse". Mémoires de Physique et de Chimie de la Société d’Arcueil 2: 317–331. 


  1. ^ Within the European Union, the following additional hazard statement (EUH014) must also be displayed on labelling: Reacts violently with water.


Boron trifluoride was discovered in 1808 by Joseph Louis Gay-Lussac and Louis Jacques Thénard, who were trying to isolate "fluoric acid" (i.e., hydrofluoric acid) by combining calcium fluoride with vitrified boric acid. The resulting vapours failed to etch glass, so they named it fluoboric gas.[20][21]


Niche uses


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