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Boron trifluoride – Wikipedia

Boron trifluoride
Boron trifluoride in 2D
Boron trifluoride in 3D
Names
Other names

Boron fluoride, Trifluoroborane
Identifiers
CAS Number
  • 7637-07-2 checkY
  • 13319-75-0 (dihydrate) checkY
3D model (JSmol)
  • Interactive image
ChEBI
  • CHEBI:33093 checkY
ChemSpider
  • 6116 checkY
ECHA InfoCard 100.028.699 Edit this at Wikidata
EC Number
  • 231-569-5
PubChem CID
  • 6356
RTECS number
  • ED2275000
UNII
  • 7JGD48PX8P ☒N
UN number compressed: 1008.
boron trifluoride dihydrate: 2851.
CompTox Dashboard (EPA)
  • DTXSID7041677 Edit this at Wikidata
InChI
  • InChI=1S/BF3/c2-1(3)4 checkY
    Key: WTEOIRVLGSZEPR-UHFFFAOYSA-N checkY
SMILES
  • FB(F)F
Properties
Chemical formula
BF3
Molar mass 67.82 g/mol (anhydrous)
103.837 g/mol (dihydrate)
Appearance colorless gas (anhydrous)
colorless liquid (dihydrate)
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)
Solubility in water
exothermic decomposition (anhydrous)
very soluble (dihydrate)
Solubility soluble in benzene, toluene, hexane, chloroform and methylene chloride
Vapor pressure >50 atm (20 °C)
Dipole moment
0 D
Thermochemistry
Heat capacity (C)
50.46 J/mol K
Std molar
entropy (S298)
254.3 J/mol K
Std enthalpy of
formation fH298)
-1137 kJ/mol
Gibbs free energy fG)
-1120 kJ/mol
Hazards
GHS labelling:
Pictograms
Press. Gas Acute Tox. 2 Skin Corr. 1A GHS08: Health hazard
Signal word
Danger
Hazard statements
H280, H314, H330, H335, H373
Precautionary statements
P260, P280, P303+P361+P353, P304+P340, P305+P351+P338, P310, P403+P233
NFPA 704 (fire diamond)
Health 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gas Flammability 0: Will not burn. E.g. water Instability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calcium Special hazards (white): no code

NFPA 704 four-colored diamond

3
0
1
Flash point Non-flammable
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)
NIOSH (US health exposure limits):
PEL (Permissible)
C 1 ppm (3 mg/m3)
REL (Recommended)
C 1 ppm (3 mg/m3)
IDLH (Immediate danger)
25 ppm
Safety data sheet (SDS) ICSC 0231
Other anions
boron trichloride
boron tribromide
boron triiodide
Other cations
aluminium fluoride
gallium(III) fluoride
indium(III) fluoride
thallium(III) fluoride
Related compounds
boron monofluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)
Infobox references
Chemical compound

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 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, 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.

Boron trifluoride pi bonding diagram

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:

PhN2BF4 → PhF + BF3 + N2

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

6 NaBF4 + 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 → CsBF4
O(C2H5)2 + BF3 → BF3·O(C2H5)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(Et)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(Me)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, and 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 BF3 is attributed to the relative weakness of the bond in the adducts F3B−L.

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 HBF4 + 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[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

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.

See also

  • List of highly toxic gases

References

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.
  • “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.
  • 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
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
  • v
  • t
  • e
Fluorine compounds
HF He
LiF BeF2 BF
BF3
B2F4
CF4
CxFy
NF3
N2F4
OF2
O2F2
F Ne
NaF MgF2 AlF3 SiF4 PF3
PF5
SF2
SF4
SF6
ClF
ClF3
ClF5
HArF
KF CaF2 ScF3 TiF3
TiF4
VF3
VF4
VF5
CrF2
CrF3
CrF4
CrF5
MnF2
MnF3
MnF4
FeF2
FeF3
CoF2
CoF3
NiF2
NiF3
CuF
CuF2
ZnF2 GaF3 GeF4 AsF3
AsF5
SeF4
SeF6
BrF
BrF3
BrF5
KrF2
RbF SrF2 YF3 ZrF4 NbF4
NbF5
MoF4
MoF5
MoF6
TcF6 RuF3
RuF4
RuF5
RuF6
RhF3
RhF5
RhF6
PdF2 AgF
AgF2
AgF3
Ag2F
CdF2 InF3 SnF2
SnF4
SbF3
SbF5
TeF4
TeF6
IF
IF3
IF5
IF7
XeF2
XeF4
XeF6
CsF BaF2 * HfF4 TaF5 WF4
WF6
ReF6
ReF7
OsF4
OsF5
OsF6
IrF3
IrF5
IrF6
PtF4
PtF6
AuF3
Au2F10
AuF5·F2
HgF2
Hg2F2
TlF PbF2
PbF4
BiF3
BiF5
PoF6 At RnF2
Fr RaF2 ** Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
* LaF3 CeF3 PrF3 NdF3 PmF3 SmF3
EuF2
EuF3
GdF3 Tb DyF3 HoF3 Er TmF3 YbF3 LuF3
** AcF3 ThF4 Pa UF3
UF4
UF5
UF6
NpF3
NpF4
NpF5
NpF6
PuF3
PuF4
PuF5
PuF6
AmF3
AmF4
CmF3 Bk Cf Es Fm Md No Lr
PF6, AsF6, SbF6 compounds
  • AgPF6
  • KAsF6
  • LiAsF6
  • NaAsF6
  • HPF6
  • HSbF6
  • NH4PF6
  • KPF6
  • KSbF6
  • LiPF6
  • NaPF6
  • NaSbF6
  • TlPF6
AlF6 compounds
  • Cs2AlF5
  • K3AlF6
  • Na3AlF6
chlorides, bromides, iodides
and pseudohalogenides
  • BaClF
  • CFN
  • ClFO2
SiF62-, GeF62- compounds
  • BaSiF6
  • BaGeF6
  • (NH4)2SiF6
  • Na2SiF6
  • K2[SiF6]
Oxyfluorides
  • BrOF3
  • BrO2F
  • ThOF2
Organofluorides
  • CBrF3
  • CBr2F2
  • CBr3F
  • CClF3
  • CCl2F2
  • CCl3F
  • CF2O
  • CF3I
  • CHF3
  • CH2F2
  • CH3F
  • C2Cl3F3
  • C2H3F
  • C6H5F
  • C7H5F3
  • C15F33N
with transition metal,
lanthanide, actinide, ammonium
  • CrFO4
  • CrF2O2
  • NH4F
  • Na2TiF6
  • Na2ZrF6
  • K2NbF7
  • K2TaF7
  • UO2F2
nitric acids
  • FNO
  • FNO2
  • FNO3
bifluorides
  • KHF2
  • NaHF2
  • NH4HF2
thionyl, phosphoryl,
and iodosyl
  • F2OS
  • F3OP
  • PSF3
  • IOF3
  • IO3F
Chemical formulas
Authority control Edit this at Wikidata
General
  • WorldCat (via Library of Congress)
National libraries
  • Germany
  • Israel
  • United States


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