Boron trifluoride
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| Boron trifluoride | |
|---|---|
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| Identifiers | |
| CAS number | [7637-07-2] |
| UN number | Compressed: 1008. Complex with acetic acid: 1742. Complex with propionic acid: 1743. Boron trifuoride diethyl etherate: 2604. Boron trifluoride dimethyl etherate: 2965. Boron trifluoride dihydrate: 2851. |
| Properties | |
| Molecular formula | BF3 |
| Molar mass | 67.8062 g mol−1 |
| Density | 2.178 g dm−3 |
| Melting point |
−126°C |
| Boiling point |
−100.3°C |
| Solubility in other solvents | decomposes |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Boron trifluoride is the chemical 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.
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[edit] Structure and bonding
Unlike the aluminium trihalides, the boron trihalides are all monomeric. They do undergo rapid reversible dimerization as indicated by the high rate of the halide exchange reactions:
- BF3 + BCl3 → BF2Cl + BCl2F
Because of the facility of this exchange process, the mixed halides cannot be obtained in pure form.
The geometry of a molecule of BF3 is described as trigonal planar. The D3h symmetry conforms with the prediction of VSEPR theory. Although featuring three polar covalent bonds, the molecule has no dipole moment by virtue of its high symmetry. Although isoelectronic with carbonate, CO32−, 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-F bonds (1.30 Å) is shorter than would be expected for single bonds,[1] 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.[1]
[edit] Synthesis
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).[2]
On a laboratory scale, BF3 is produced by the thermal decomposition of diazonium salts:[3]
[edit] Lewis acidity and related reactions
Boron trifluoride is a versatile Lewis acid that forms adducts with such Lewis bases as fluoride and ethers:
Tetrafluoroborate salts are commonly employed as non-coordinating anions. The adduct with diethyl ether is a conveniently handled liquid and consequently is a widely encountered as a laboratory source of BF3.
[edit] 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 (strongest Lewis acid)
This trend commonly attributed to the degree of π-bonding in the planar boron trihalide that would be lost upon pyramidalization of the BX3 molecule.[4] which follows this trend:
- BF3 > BCl3 > BBr3 (most easily pyramidalized)
The criteria for evaluating the relative strength of π-bonding are not clear, however.[1]
One of the suggestion is that F atom is small compared to I atom, the lone pair electron in Pz of F readily and easily donated and overlapped to empty Pz orbital of boron.
As a result, the back donation of F is greater than that of I.
In an alternative explanation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the adducts F3B-L.[5][6]
[edit] Hydrolysis
Boron trifluoride reacts with water to give boric acid,and fluoroboric acid: and [HF] The reaction commences with the formation of the aquo adduct, H2O-BF3, which then loses HF:
- 4 BF3 + 3 H2O → 3 HBF4 + "B(OH)3"
The heavier trihalides do not undergo analogous reactions, possibly 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.
[edit] Handling
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.[7]
[edit] Uses
- applied as dopant in ion implantation
- p-type dopant for epitaxially grown silicon
- initiates polymerisation reactions of unsaturated compounds. Example polyethers
- as a catalyst in some isomerization, alkylation, esterification, condensation, Mukaiyama aldol addition, and other reactions.
- used in sensitive neutron detectors
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[edit] References
- ^ a b c Greenwood, N. N.; A. Earnshaw (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Flood, D. T.. "Fluorobenzene". Org. Synth.; Coll. Vol. 2: 295.
- ^ Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. (1999). Advanced Inorganic Chemistry (6th Edn.) New York: Wiley-Interscience. ISBN 0-471-19957-5.
- ^ Group V Chalcogenide Complexes of Boron Trihalides Boorman, P. M.; Potts, D. Canadian. Journal of Chemistry (Rev. can. chim.) volume 52, (1974) pp 2016-2020
- ^ T. Brinck, J. S. Murray and P. Politzer (1993). "A computational analysis of the bonding in boron trifluoride and boron trichloride and their complexes with ammonia". Inorg. Chem. 32 (12): 2622-2625. doi:.
- ^ Boron trifluoride. Gas Encyclopedia. Air Liquide.
