Boron trifluoride is the inorganic compound with the formulaBF3. This pungent, colourless, and 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
The geometry of a molecule of BF3 is trigonal planar. Its D3hsymmetry 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, CO2−3.
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] Others point to the ionic nature of the bonds in BF3.[8]
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).[9] Approximately 2300-4500 tonnes of boron trifluoride are produced every year.[10]
Laboratory scale
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. A well documented route involves the thermal decomposition of diazonium salts of [BF4]−:[11]
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).[13]
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 < BI3 (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.[17] which follows this trend:
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. As a concequence, the bond length between boron and the halogen increases while going from fluor to iodine hence spatial overlap between the orbitals becomes more difficult. 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.[18][19]
Yet another explanation might be found in the fact that the pz orbitals in each higher period have an extra nodal plane and opposit signs of the wave function on each side of that plane. This results in bonding and antibonding regions within the same bond, deminishing the effective overlap and so lowering the π-donating blockage of the acidity.[20]
Hydrolysis
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.[21]
4 BF3 + 3 H2O → 3 H[BF4] + B(OH)3
The heavier trihalides do not undergo analogous reactions, possibly due to the lower stability of the tetrahedral ions [BCl4]− and [BBr4]−. 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", Pocket Guide to Chemical Hazards, U.S. Department of Health and Human Services (NIOSH) Publication No. 2005-149, Washington, DC: Government Printing Office, 2005, ISBN9780160727511.
^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". Inorganic Chemistry. 32 (12): 2622–2625. doi:10.1021/ic00064a008.
^Here on Wikipedia an easy to understand table is found, which shows drawings of the several higher p orbitals.
^Sowa, F. J.; Hennion, G. F.; Nieuwland, J. A. (1935). "Organic Reactions with Boron Fluoride. IX. The Alkylation of Phenol with Alcohols". Journal of the American Chemical Society. 57 (4): 709–711. doi:10.1021/ja01307a034.
