Boron trifluoride

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Boron trifluoride
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Molar mass 67.8062 g mol−1
Density 2.178 g dm−3
Melting point
Boiling point
Solubility in other solvents decomposes
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Boron trifluoride is the chemical compound with 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

Unlike the aluminium trihalides, the boron trihalides are all monomeric. They do undergo rapid dimerization as indicated by the high rate of the halide exchange reactions:

BF3 + BCl3 → BF2Cl + BCl2F

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

The geometry of a molecule of BF3 is described as trigonal planar. The D3h symmetry conforms with the prediction of VSEPR theory. Although featuring three polar covalent bonds, the molecule has no dipole moment by virtue of its high symmetry. Although isoelectronic with carbonate, 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-F bonds (1.30 Å) is shorter than would be expected for single bonds,[1] 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.[1]


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).[2]

On a laboratory scale, BF3 is produced by the thermal decomposition of diazonium salts:[3]

PhN2BF4PhF + BF3 + N2

Lewis acidity and related reactions

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 is a conveniently handled liquid and consequently is a widely encountered as a laboratory source of BF3.

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 commonly attributed to the degree of π-bonding in the planar boron trihalide that would be lost upon pyramidalization of the BX3 molecule.[4] which follows this trend:

BF3 > BCl3 > BBr3 (most easily pyramidalized)

The criteria for evaluating the relative strength of π-bonding are not clear, however.[1]
In an alternative explanation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the adducts F3B-L.[5][6]


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:

4 BF3 + 3 H2O → 3 HBF4 + "B(OH)3"

The heavier trihalides do not undergo analogous reactions, possibly 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 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.[7]



  1. 1.0 1.1 1.2 Greenwood, N. N.; A. Earnshaw (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  2. Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  3. Flood, D. T.. "Fluorobenzene". Org. Synth.; Coll. Vol. 2: 295. 
  4. Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. (1999). Advanced Inorganic Chemistry (6th Edn.) New York: Wiley-Interscience. ISBN 0-471-19957-5.
  5. Group V Chalcogenide Complexes of Boron Trihalides Boorman, P. M.; Potts, D. Canadian. Journal of Chemistry (Rev. can. chim.) volume 52, (1974) pp 2016-2020
  6. T. Brinck, J. S. Murray and P. Politzer (1993). "A computational analysis of the bonding in boron trifluoride and boron trichloride and their complexes with ammonia". Inorg. Chem. 32 (12): 2622–2625. doi:10.1021/ic00064a008.
  7. "Boron trifluoride". Gas Encyclopedia. Air Liquide.

External links

cs:Fluorid boritý de:Bortrifluorid nl:Boortrifluoride