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]

[PhN₂]⁺[BF₄]⁻ → PhF + BF₃ + N₂

Alternatively it arises from the reaction of sodium tetrafluoroborate, boron trioxide, and sulfuric acid:^[12]

6 Na[BF₄] + B₂O₃ + 6 H₂SO₄ → 8 BF₃ + 6 NaHSO₄ + 3 H₂O

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 < 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:

BF₃ > BCl₃ > BBr₃ < BI₃ (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. As a consequence, the bond length between boron and the halogen increases while going from fluorine to iodine hence spatial overlap between the orbitals becomes more difficult. The lone pair electron in p_z of F is readily and easily donated and overlapped to empty p_z 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 p_z orbitals in each higher period have an extra nodal plane and opposite signs of the wave function on each side of that plane. This results in bonding and antibonding regions within the same bond, diminishing 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.

Organic chemistry

Boron trifluoride is most importantly used as a reagent in organic synthesis, typically as a Lewis acid.^[10][22] Examples include:

• initiates polymerisation reactions of unsaturated compounds, such as polyethers
• as a catalyst in some isomerization, acylation,^[23] alkylation, esterification, dehydration,^[24] condensation, Mukaiyama aldol addition, and other reactions^{[25][citation needed]}

Niche uses

Other, less common uses for boron trifluoride include:

• applied as dopant in ion implantation
• p-type dopant for epitaxially grown silicon
• used in sensitive neutron detectors in ionization chambers and devices to monitor radiation levels in the Earth's atmosphere
• in fumigation
• as a flux for soldering magnesium
• to prepare diborane^[12]