Английская Википедия:Boron trifluoride
Boron trifluoride is the inorganic compound with the formula Шаблон:Chem2. 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 Шаблон:Chem2 is trigonal planar. Its D3h 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, Шаблон:Chem2.
Шаблон:Chem2 is commonly referred to as "electron deficient," a description that is reinforced by its exothermic reactivity toward Lewis bases.
In the boron trihalides, Шаблон:Chem2, the length of the B–X 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] Others point to the ionic nature of the bonds in Шаблон:Chem2.[2]
Boron trifluoride pi bonding diagram
Synthesis and handling
Шаблон:Chem2 is manufactured by the reaction of boron oxides with hydrogen fluoride:
Typically the HF is produced in situ from sulfuric acid and fluorite (Шаблон:Chem2).[3] Approximately 2300-4500 tonnes of boron trifluoride are produced every year.[4]
Laboratory scale
For laboratory scale reactions, Шаблон:Chem2 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 Шаблон:Chem2:[5]
Alternatively it arises from the reaction of sodium tetrafluoroborate, boron trioxide, and sulfuric acid:[6]
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).[7]
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.[8]
Reactions
Unlike the aluminium and gallium trihalides, the boron trihalides are all monomeric. They undergo rapid halide exchange reactions:
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:
Tetrafluoroborate salts are commonly employed as non-coordinating anions. The adduct with diethyl ether, boron trifluoride diethyl etherate, or just boron trifluoride etherate, (Шаблон:Chem2) is a conveniently handled liquid and consequently is widely encountered as a laboratory source of Шаблон:Chem2.[9] Another common adduct is the adduct with dimethyl sulfide (Шаблон:Chem2), which can be handled as a neat liquid.[10]
Comparative Lewis acidity
All three lighter boron trihalides, Шаблон:Chem2 (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:
- Шаблон:Chem2 < Шаблон:Chem2 < Шаблон:Chem2 < Шаблон:Chem2 (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 Шаблон:Chem2 molecule.[11] which follows this trend:
- Шаблон:Chem2 > Шаблон:Chem2 > Шаблон:Chem2 < Шаблон:Chem2 (most easily pyramidalized)
The criteria for evaluating the relative strength of π-bonding are not clear, however.[1] 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 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 Шаблон:Chem2 is attributed to the relative weakness of the bond in the adducts Шаблон:Chem2.[12][13]
Yet another explanation might be found in the fact that the pz 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.[14]
Hydrolysis
Boron trifluoride reacts with water to give boric acid and fluoroboric acid. The reaction commences with the formation of the aquo adduct, Шаблон:Chem2, which then loses HF that gives fluoroboric acid with boron trifluoride.[15]
The heavier trihalides do not undergo analogous reactions, possibly due to the lower stability of the tetrahedral ions Шаблон:Chem2 and Шаблон:Chem2. 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.[4][16] Examples include:
- initiates polymerisation reactions of unsaturated compounds, such as polyethers
- as a catalyst in some isomerization, acylation,[17] alkylation, esterification, dehydration,[18] condensation, Mukaiyama aldol addition, and other reactions[19]Шаблон: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[6]
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.[20][21]
See also
References
External links
Шаблон:Boron compounds Шаблон:Fluorine compounds Шаблон:Authority control
- ↑ 1,0 1,1 1,2 Шаблон:Greenwood&Earnshaw2nd
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite book
- ↑ 4,0 4,1 Шаблон:Ullmann
- ↑ Шаблон:OrgSynth
- ↑ 6,0 6,1 Шаблон:Cite book
- ↑ Шаблон:Cite book
- ↑ Шаблон:Cite encyclopedia
- ↑ Шаблон:Cite book
- ↑ Шаблон:Cite book
- ↑ Шаблон:Cotton&Wilkinson6th
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite journal
- ↑ Here on Wikipedia an easy to understand table is found, which shows drawings of the several higher p orbitals.
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite encyclopedia
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite web
- ↑ Шаблон:Cite journal
- ↑ Шаблон:Cite journal
