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| Names | |||
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| IUPAC name Boron trifluoride | |||
| Systematic IUPAC name Trifluoroborane | |||
| Other names Boron fluoride, Trifluoroborane | |||
| Identifiers | |||
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3D model (JSmol) | |||
| ChEBI | |||
| ChemSpider |
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| ECHA InfoCard | 100.028.699 | ||
| EC Number |
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| RTECS number |
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| UNII | |||
| UN number | compressed:1008. boron trifluoride dihydrate:2851. | ||
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| Properties | |||
| BF3 | |||
| Molar mass | 67.82 g/mol (anhydrous) 103.837 g/mol (dihydrate) | ||
| Appearance | colorless gas (anhydrous) colorless liquid (dihydrate) | ||
| Odor | Pungent | ||
| Density | 0.00276 g/cm3 (anhydrous gas) 1.64 g/cm3 (dihydrate) | ||
| Melting point | −126.8 °C (−196.2 °F; 146.3 K) | ||
| Boiling point | −100.3 °C (−148.5 °F; 172.8 K) | ||
| exothermic decomposition[1] (anhydrous) very soluble (dihydrate) | |||
| Solubility | soluble inbenzene,toluene,hexane,chloroform andmethylene chloride | ||
| Vapor pressure | >50 atm (20 °C)[2] | ||
| 0 D | |||
| Thermochemistry | |||
| 50.46 J/(mol·K) | |||
Std molar entropy(S⦵298) | 254.3 J/(mol·K) | ||
Std enthalpy of formation(ΔfH⦵298) | −1137 kJ/mol | ||
Gibbs free energy(ΔfG⦵) | −1120 kJ/mol | ||
| Hazards[4][5] | |||
| GHS labelling: | |||
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| Danger | |||
| H314,H330,H335,H373 | |||
| P260,P280,P303+P361+P353,P304+P340,P305+P351+P338,P310,P403+P233 | |||
| NFPA 704 (fire diamond) | |||
| Flash point | Nonflammable | ||
| Lethal dose or concentration (LD, LC): | |||
LC50 (median concentration) | 1227 ppm (mouse, 2 hr) 39 ppm (guinea pig, 4 hr) 418 ppm (rat, 4 hr)[3] | ||
| NIOSH (US health exposure limits): | |||
PEL (Permissible) | C 1 ppm (3 mg/m3)[2] | ||
REL (Recommended) | C 1 ppm (3 mg/m3)[2] | ||
IDLH (Immediate danger) | 25 ppm[2] | ||
| Safety data sheet (SDS) | ICSC0231 | ||
| Related compounds | |||
Otheranions | |||
Othercations | |||
Related compounds | Boron monofluoride | ||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Boron trifluoride is theinorganic compound with theformulaBF3. This pungent, colourless, andtoxic gas forms white fumes in moist air. It is a usefulLewis acid and a versatile building block for otherboron compounds.
The geometry of amolecule ofBF3 istrigonal planar. Its D3hsymmetry conforms with the prediction ofVSEPR theory. The molecule has no dipole moment by virtue of its high symmetry. The molecule isisoelectronic with the carbonate anion,CO2−3.
BF3 is commonly referred to as "electron deficient," a description that is reinforced by itsexothermic reactivity towardLewis bases.
In theboron trihalides,BX3, the length of the B–X bonds (1.30 Å) is shorter than would be expected for single bonds,[7] and this shortness may indicate stronger B–Xπ-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.[7] Others point to the ionic nature of the bonds inBF3.[8]
BF3 is manufactured by the reaction of boron oxides withhydrogen fluoride:
Typically the HF is producedin situ from sulfuric acid andfluorite (CaF2).[9] Approximately 2300–4500 tonnes of boron trifluoride are produced every year.[10]
For laboratory scale reactions,BF3 is usually produced in situ usingboron trifluoride etherate, which is a commercially available liquid.[how?]
Laboratory routes to the solvent-free materials are numerous. A well documented route involves the thermal decomposition ofdiazonium salts of[BF4]−:[11]
It forms by treatment of a mixtureboron trioxide andsodium tetrafluoroborate with sulfuric acid:[12]
Alternatively,boron tribromide converts variousorganofluorine compounds toorganobromines, evolving the trifluoride gas:[13]
Anhydrous boron trifluoride has aboiling point of −100.3 °C and acritical 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 thecritical pressure of 49.85 bar (4.985 MPa).[14]
Boron trifluoride is corrosive. Suitable metals for equipment handling boron trifluoride includestainless steel,monel, andhastelloy. In presence of moisture it corrodes steel, including stainless steel. It reacts withpolyamides.Polytetrafluoroethylene,polychlorotrifluoroethylene,polyvinylidene fluoride, andpolypropylene show satisfactory resistance. Thegrease used in the equipment should befluorocarbon based, as boron trifluoride reacts with the hydrocarbon-based ones.[15]
Unlike the aluminium and gallium trihalides, the boron trihalides are allmonomeric. 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 versatileLewis acid that formsadducts with suchLewis bases asfluoride andethers:
Tetrafluoroborate salts are commonly employed asnon-coordinating anions. The adduct withdiethyl ether, boron trifluoride diethyl etherate, or justboron trifluoride etherate, (BF3·O(CH2CH3)2) is a conveniently handledliquid and consequently is widely encountered as a laboratory source ofBF3.[16] Another common adduct is the adduct withdimethyl sulfide (BF3·S(CH3)2), which can be handled as a neat liquid.[17]
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:
This trend is commonly attributed to the degree ofπ-bonding in the planar boron trihalide that would be lost uponpyramidalization of theBX3 molecule.[18] which follows this trend:
The criteria for evaluating the relative strength ofπ-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 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 forBF3 is attributed to the relative weakness of the bond in the adductsF3B−L.[19][20]
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.[original research?]
Boron trifluoride reacts with water to giveboric acid andfluoroboric 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]
The heavier trihalides also hydrolyze, but toboric andhydrohalic acids, 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 asdiazonium ions, that are otherwise difficult to isolate as solids.
Boron trifluoride is most importantly used as a reagent inorganic synthesis, typically as aLewis acid.[10][22] Examples include:
Other, less common uses for boron trifluoride include:
Boron trifluoride was discovered in 1808 byJoseph Louis Gay-Lussac andLouis Jacques Thénard, who were trying to isolate "fluoric acid" (i.e.,hydrofluoric acid) by combiningcalcium fluoride with vitrifiedboric acid. The resulting vapours failed to etch glass, so they named itfluoboric gas.[26][27]
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