Boron trifluoride - Misplaced Pages

Boron trifluoride is the inorganic compound with the formula BF₃. 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 D_3h 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, CO2−3.

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–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 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). Approximately 2300-4500 tonnes of boron trifluoride are produced every year.

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]:

[PhN₂][BF₄] → PhF + BF₃ + N₂

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

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

Properties

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.

Reactions

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.

Boron trifluoride is a versatile Lewis acid that forms adducts with such Lewis bases as fluoride and ethers:

CsF + BF₃ → Cs[BF₄]

O(CH₂CH₃)₂ + BF₃ → BF₃·O(CH₂CH₃)₂

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(CH₂CH₃)₂) is a conveniently handled liquid and consequently is widely encountered as a laboratory source of BF₃. Another common adduct is the adduct with dimethyl sulfide (BF₃·S(CH₃)₂), which can be handled as a neat liquid.

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. 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. 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.

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.

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.

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.

Uses

Organic chemistry

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

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

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

Discovery

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.

References

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Index no. 005-001-00-X of Annex VI, Part 3, to Regulation (EC) No 1272/2008 of the European Parliament and of the Council of 16 December 2008 on classification, labelling and packaging of substances and mixtures, amending and repealing Directives 67/548/EEC and 1999/45/EC, and amending Regulation (EC) No 1907/2006. Official Journal of the European Union L353, 31 December 2008, pp. 1–1355 at p 341.
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Here on Misplaced Pages an easy to understand table is found, which shows drawings of the several higher p orbitals.
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External links

"Safety and Health Topics: Boron Trifluoride". OSHA.
"BORON TRIFLUORIDE ICSC: 0231". International Chemical Safety Cards. CDC. Archived from the original on 2017-11-23. Retrieved 2017-09-08.
"Boron & Compounds: Overview". National Pollutant Inventory. Australian Government.
"Fluoride Compounds: Overview". National Pollutant Inventory. Australian Government.
"Boron trifluoride". WebBook. NIST.
"Boron Trifluoride (BF₃) Applications". Honeywell. Archived from the original on 2012-01-29. Retrieved 2012-02-14.

v t e Boron compounds
Boron pnictogenides	BAs BN BP
Boron halides	BBr₃ BCl₃ BF BFO BF₃ BI₃ B₂F₄ B₂Cl₄
Acids	B(NO₃)₃ B(OH)₃ BPO₄ BH₃O
Boranes	BH₃ B₂H₄ B₂H₆ BH₃NH₃ B₄H₁₀ B₅H₉ B₅H₁₁ B₆H₁₀ B₆H₁₂ B₁₀H₁₄ B₁₈H₂₂
Boron oxides and sulfides	B₂O B₂O₃ B₂S₃ B₆O
Carbides	B₄C
Organoboron compounds	(BH₂Me)₂ BMe₃ BEt₃ Ac₄(BO₃)₂ COBH₃