Boron trifluoride is the inorganic 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.
Boron fluoride, Trifluoroborane
3D model (JSmol)
|UN number||compressed: 1008.|
boron trifluoride dihydrate: 2851.
CompTox Dashboard (EPA)
|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; 146.3 K)|
|Boiling point||−100.3 °C (−148.5 °F; 172.8 K)|
|exothermic decomposition  (anhydrous)|
very soluble (dihydrate)
|Solubility||soluble in benzene, toluene, hexane, chloroform and methylene chloride|
|Vapor pressure||>50 atm (20°C)|
Heat capacity (C)
|50.46 J/mol K|
|254.3 J/mol K|
Std enthalpy of
Gibbs free energy (ΔfG˚)
|Safety data sheet||ICSC|
|GHS signal word||DANGER|
|H280, H330, H314, H335, H373|
|P260, P280, P303+361+353, P304+340, P310, P305+351+338, P403+233|
|Lethal dose or concentration (LD, LC):|
LC50 (median concentration)
|1227 ppm (mouse, 2 hr)|
39 ppm (guinea pig, 4 hr)
418 ppm (rat, 4 hr)
|US health exposure limits (NIOSH):|
|C 1 ppm (3 mg/m3)|
|C 1 ppm (3 mg/m3)|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Structure and bondingEdit
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−.
In the boron trihalides, BX3, the length of the B-X bonds (1.30 Å) is shorter than would be expected for single bonds, 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. Others point to the ionic nature of the bonds in BF3.
Synthesis and handlingEdit
BF3 is manufactured by the reaction of boron oxides with hydrogen fluoride:
- B2O3 + 6 HF → 2 BF3 + 3 H2O
For laboratory scale reactions, BF3 is usually produced in situ using boron trifluoride etherate, which is a commercially available liquid.
Laboratory routes to the solvent-free materials are numerous. One well documented route involves the thermal decomposition of diazonium salts of BF4−:
- 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).
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.
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.
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 widely encountered as a laboratory source of BF3. It is stable as a solution in ether—an excess of ether compared to BF3—but not as a pure stoichiometric substance. Another common adduct is the adduct with dimethyl sulfide (BF3·S(Me)2), which can be handled as a neat liquid.
Comparative Lewis acidityEdit
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)
- BF3 > BCl3 > BBr3 (most easily pyramidalized)
The criteria for evaluating the relative strength of π-bonding are not clear, however. 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.
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 fluoroboric acid with boron trifluoride.
- 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.
- initiates polymerisation reactions of unsaturated compounds, such as polyethers
- as a catalyst in some isomerization, acylation, alkylation, esterification, dehydration, condensation, Mukaiyama aldol addition, and other reactions
Other, less common uses for boron trifluoride include:
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.
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