Boron Trifluoride

Boron trifluoride is an inorganic compound with the formula BF3 . This pungent colorless poisonous gas forms white smoke in moist air. It is a versatile building block for a useful Lewis acid and other boron compounds.

Structure and binding

The geometry of a molecule BF 3 is triangular planar. Its D 3H symmetry is in line with the prediction of VSEPR theory. The molecule has no dipole moment due to its high symmetry. The molecule is isoelectronic with the carbonate anion, CO2−3.

Bf3 is commonly referred to as “electron deficient”, a description reinforced by its exothermic reactivity towards Lewis bases .

In boron trihalides, BX 3 , the length of the BX bond (1.30 k) is shorter than would be expected for a bond, [7] and this shortness may be indicative of stronger BX -bonding in fluoride. A simpler explanation invokes the symmetry-allowed overlap of the p orbital on the boron atom with the in-phase combination of three equally oriented p orbitals on the fluorine atoms. [7] Others point to the ionic nature of the bonds in BF3 .

Synthesis and Handling

BF 3 is produced by the reaction of boron oxide with hydrogen fluoride:B 2 O 3 + 6 HF → 2 BF 3 + 3 H 2 O

Typically HF is produced in situ from sulfuric acid and fluorite ( CaF2 ) . [9] About 2300–4500 tons of boron trifluoride are produced each year. [10]

Laboratory scale

For laboratory scale reactions, BF is usually produced in situ using boron trifluoride ether, which is a commercially available liquid.

Laboratory routes to solvent-free materials are numerous. In a well documented passage BF . The thermal decomposition of diazonium salts involves-
4: [11]phn 2 bf 4 → phf + bf 3 + n 2

Alternatively it is produced by the reaction of sodium tetrafluoroborate, boron trioxide and sulfuric acid: [12]6 NaBF 4 + B 2 O 3 + 6 H 2 SO 4 → 8 BF 3 + 6 NaHSO 4 + 3 H 2 O


Anhydrous boron trifluoride has a boiling point of -100.3 °C and a critical temperature of −12.3 °C, so that it can only be stored as a refrigerated liquid between those temperatures. Storage or transport vessels must be designed to withstand the internal pressure, as a refrigeration system failure can cause the pressure to rise to a critical pressure of 49.85 bar (4.985 MPa). [13]

Boron trifluoride is corrosive. Metals suitable for equipment handling boron trifluoride include stainless steel, Monel and Hastelloy. It corrosives steel, including stainless steel, in the presence of moisture. It reacts with polyamides. Polytetrafluoroethylene, polychlorotrifluoroethylene, polyvinylidene fluoride and polypropylene show satisfactory resistance. Oil based equipment should be used in fluorocarbon based, such as boron trifluoride reaction with hydrocarbon based ones. [14]


Unlike aluminum and gallium trihalides, boron trihalides are all monomeric. They rapidly undergo halide exchange reactions:Bf 3 + Bcl 3 → Bf 2 Cl + Bcl 2 F

Due to the convenience of this exchange process, mixed halides cannot be obtained in pure form.

Boron trifluoride is a Lewis versatile acid that forms adducts with such Lewis bases as fluoride and ether:Csf + Bf 3 → Csbf 4O (C 2 H 5 ) 2 + Bf 3 → Bf 3 O (C 2 H 5 ) 2

Tetrafluoroborate salts usually serve as the non-coordinating anion. Adduct with diethyl ether, boron trifluoride diethyl etherate, or simply boron trifluoride etherate, (BF 3 O(et) 2 ) is a conveniently handled liquid and is consequently widely encountered as a laboratory source of BF 3 . [15] Another common addition is with dimethyl sulfide (BF 3 S(Me) 2 ), which can be handled as a clear liquid. [16]

Comparative Lewis Acidity

All three mild boron trihalides, BX 3 (X = F, Cl, Br) form stable pairs with common Lewis bases. Their relative Lewis acidity can be evaluated in terms of the relative exothermicity of the adduct-forming reaction. Such measurements have revealed the following sequence for Lewis acidity:Bf 3 < Bcl 3 < BBr 3 (strongest Lewis acid)

This trend is normally attributed to the degree of -bonding in planar boron trihalide that would be lost upon pyramidalization of the BX 3 molecule. [17] Which follows this trend:bf 3 > bcl 3 > bbr 3 ( most easily pyramidal)

However, the criteria for evaluating the relative strength of -bonding are unclear. [7] One suggestion is that the F atom is smaller than the larger Cl and Br atoms, and the lone pair electron in the p z of F is readily and readily donated and vacated to the p z orbital of boron. goes . As a result, the pi donation of F is greater than that of Cl or Br.

In an alternative interpretation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the added F3B-L [18] [19]


Boron trifluoride reacts with water to give boric acid and fluoroboric acid. The reaction begins with the formation of the aqua adduct, HO- BF3, which then loses HF which gives fluoroboric acid with boron trifluoride [20]4Bf 3 + 3H 2 O → 3HBF 4 + B(OH) 3

The heavy trihalides do not undergo similar reactions, possibly due to the low stability of the tetrahedral anion BCl .-
4and BBR-
4, Due to the high acidity of fluoroboric acid, the fluoroborate ion can be used specifically to separate electrophilic cations, such as the diazonium ion, which are otherwise difficult to separate as solids.


Organic Chemistry

Boron trifluoride is used most significantly as a reagent in organic synthesis, usually in the form of Lewis acid. [10] [21] Examples include:

  • Initiates polymerization reactions of unsaturated compounds, such as polyethers
  • As a catalyst in some isomerization, acylation, [22] alkylation, esterification, dehydration, [23] condensation, Mukaiyama aldol addition, and other reactions [24] citation needed ]

Got used

Other, less common uses of boron trifluoride include:

  • Applied as a dopant in ion implantation
  • P-type dopant for epitaxially grown silicon
  • Used in sensitive neutron detectors in ionization chambers and instruments to monitor radiation levels in the Earth’s atmosphere
  • fumigation in
  • As a flux for soldering magnesium
  • To prepare diborane [12]


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 mixing calcium fluoride with vitrified boric acid. The resulting vapors failed to etch the glass, so they named it fluboric gas .