Boron Trifluoride - Wikipedia

Boron trifluoride

Names
IUPAC name Boron trifluoride
Systematic IUPAC name Trifluoroborane
Other names Boron fluoride, Trifluoroborane
Identifiers
CAS Number	7637-07-2 Y 13319-75-0 (dihydrate) Y
3D model (JSmol)	Interactive image Interactive image
ChEBI	CHEBI:33093 Y
ChemSpider	6116 Y
ECHA InfoCard	100.028.699
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
InChI InChI=1S/BF3/c2-1(3)4 YKey: WTEOIRVLGSZEPR-UHFFFAOYSA-N Y
SMILES FB(F)F [F+]=[B-](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)
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)
Solubility in water	exothermic decomposition [1] (anhydrous)very soluble (dihydrate)
Solubility	soluble in benzene, toluene, hexane, chloroform and methylene chloride
Vapor pressure	>50 atm (20 °C)[2]
Dipole moment	0 D
Thermochemistry
Heat capacity (C)	50.46 J/(mol·K)
Std molarentropy (S⦵298)	254.3 J/(mol·K)
Std enthalpy offormation (ΔfH⦵298)	−1137 kJ/mol
Gibbs free energy (ΔfG⦵)	−1120 kJ/mol
Hazards[4][5]
GHS labelling:
Pictograms
Signal word	Danger
Hazard statements	H314, H330, H335, H373
Precautionary statements	P260, P280, P303+P361+P353, P304+P340, P305+P351+P338, P310, P403+P233
NFPA 704 (fire diamond)	3 0 1
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)	ICSC 0231
Related compounds
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 YN ?) 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

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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,[7] 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.[7] Others point to the ionic nature of the bonds in BF3.[8]

Synthesis and handling

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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).[9] Approximately 2300–4500 tonnes of boron trifluoride are produced every year.[10]

Laboratory scale

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For laboratory scale reactions, BF3 is usually produced in situ using boron 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 of diazonium salts of [BF4]−:[11]

[PhN2]+[BF4]− → PhF + BF3 + N2

It forms by treatment of a mixture boron trioxide and sodium tetrafluoroborate with sulfuric acid:[12]

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

Alternatively, boron tribromide converts various organofluorine compounds to organobromines, evolving the trifluoride gas:[13]

3 R–F + BBr3 → 3 R–Br + BF3

Properties

[edit]

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).[14]

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.[15]

Reactions

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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.[16] Another common adduct is the adduct with dimethyl sulfide (BF3·S(CH3)2), which can be handled as a neat liquid.[17]

Comparative Lewis acidity

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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.[18] 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.[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 for BF3 is attributed to the relative weakness of the bond in the adducts F3B−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?]

Hydrolysis

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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.[21]

4 BF3 + 3 H2O → 3 H[BF4] + B(OH)3

The heavier trihalides also hydrolyze, but to boric and hydrohalic 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 as diazonium ions, that are otherwise difficult to isolate as solids.

Uses

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Organic chemistry

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Boron trifluoride is most importantly used as a reagent in organic synthesis, typically as a Lewis acid.[10][22] Examples include:

• initiates polymerisation reactions of unsaturated compounds, such as polyethers
• as a catalyst in some isomerization, acylation,[23] alkylation, esterification, dehydration,[24] condensation, Mukaiyama aldol addition, and other reactions[25][citation needed]

Niche uses

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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[12]

Discovery

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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.[26][27]

See also

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• List of highly toxic gases

References

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1. ^ Prudent Practices in the Laboratory. 16 August 1995. doi:10.17226/4911. ISBN 978-0-309-05229-0. Archived from the original on 14 December 2014. Retrieved 7 May 2018. {{cite book}}: |website= ignored (help)
2. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0062". National Institute for Occupational Safety and Health (NIOSH).
3. ^ "Boron trifluoride". Immediately Dangerous to Life or Health Concentrations. National Institute for Occupational Safety and Health.
4. ^ 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.
5. ^ "Boron trifluoride". Pocket Guide to Chemical Hazards. U.S. Department of Health and Human Services (NIOSH) Publication No. 2005-149. Washington, DC: Government Printing Office. 2005. ISBN 9780160727511..
6. ^ Inc, New Environment. "New Environment Inc. - NFPA Chemicals". www.newenv.com. Archived from the original on 27 August 2016. Retrieved 7 May 2018. {{cite web}}: |last= has generic name (help)
7. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
8. ^ Gillespie, Ronald J. (1998). "Covalent and Ionic Molecules: Why Are BeF2 and AlF3 High Melting Point Solids whereas BF3 and SiF4 Are Gases?". Journal of Chemical Education. 75 (7): 923. Bibcode:1998JChEd..75..923G. doi:10.1021/ed075p923.
9. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5.
10. ^ a b Brotherton, R. J.; Weber, C. J.; Guibert, C. R.; Little, J. L. "Boron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a04_309. ISBN 978-3-527-30673-2.
11. ^ Flood, D. T. (1933). "Fluorobenzene". Organic Syntheses. 13: 46; Collected Volumes, vol. 2, p. 295.
12. ^ a b Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry. Vol. 1 (2nd ed.). New York: Academic Press. p. 220 & 773. ISBN 978-0121266011. {{cite book}}: ISBN / Date incompatibility (help)
13. ^ Hegedüs, Kristof (11 Jan 2019). "Performing a halogen exchange, a HalEx reaction on..." Pictures from an Organic Chemistry Laboratory. Tumblr. Archived from the original on 19 Jan 2019. Retrieved 12 January 2025.
14. ^ Yaws, C. L., ed. (1999). Chemical Properties Handbook. McGraw-Hill. p. 25.
15. ^ "Boron trifluoride". Gas Encyclopedia. Air Liquide. 2016-12-15. Archived from the original on 2006-12-06.
16. ^ Cornel, Veronica; Lovely, Carl J. (2007). "Boron Trifluoride Etherate". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/9780470842898.rb249.pub2. ISBN 978-0471936237. S2CID 100921225.
17. ^ Heaney, Harry (2001). "Boron Trifluoride-Dimethyl Sulfide". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rb247. ISBN 0471936235.
18. ^ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN 0-471-19957-5
19. ^ Boorman, P. M.; Potts, D. (1974). "Group V Chalcogenide Complexes of Boron Trihalides". Canadian Journal of Chemistry. 52 (11): 2016–2020. doi:10.1139/v74-291.
20. ^ Brinck, T.; Murray, J. S.; Politzer, P. (1993). "A Computational Analysis of the Bonding in Boron Trifluoride and Boron Trichloride and their Complexes with Ammonia". Inorganic Chemistry. 32 (12): 2622–2625. doi:10.1021/ic00064a008.
21. ^ Wamser, C. A. (1951). "Equilibria in the System Boron Trifluoride–Water at 25°". Journal of the American Chemical Society. 73 (1): 409–416. Bibcode:1951JAChS..73..409W. doi:10.1021/ja01145a134.
22. ^ Heaney, H. (2001). "Boron Trifluoride". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.rb250. ISBN 0-471-93623-5.
23. ^ Mani, Rama I.; Erbert, Larry H.; Manise, Daniel (1991). "Boron Trifluoride in the Synthesis of Plant Phenolics: Synthesis of Phenolic Ketones and Phenyl Stryl Ketones" (PDF). Journal of Tennessee Academy of Science. 66 (1): 1–8. Archived from the original (PDF) on 27 October 2016. Retrieved 27 October 2016.
24. ^ Sowa, F. J.; Hennion, G. F.; Nieuwland, J. A. (1935). "Organic Reactions with Boron Fluoride. IX. The Alkylation of Phenol with Alcohols". Journal of the American Chemical Society. 57 (4): 709–711. Bibcode:1935JAChS..57..709S. doi:10.1021/ja01307a034.
25. ^ "Boron Trifluoride (BF3) Applications". Honeywell. Archived from the original on 2012-01-29.
26. ^ Gay-Lussac, J. L.; Thénard, L. J. (1809). "Sur l'acide fluorique". Annales de Chimie. 69: 204–220.
27. ^ Gay-Lussac, J. L.; Thénard, L. J. (1809). "Des propriétés de l'acide fluorique et sur-tout de son action sur le métal de la potasse". Mémoires de Physique et de Chimie de la Société d'Arcueil. 2: 317–331.

External links

[edit]

• "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 B2(OH)4 BH3O
Boranes	BH3 B2H4 B2H6 BH3NH3 B4H10 B5H9 B5H11 B6H10 B6H12 B10H14 B18H22 [B12H12]2-
Boron oxides and sulfides	BO B2O3 B2S3 B6O
Carbides	B4C
Organoboron compounds	(BH2Me)2 BMe3 BEt3 Ac4(BO3)2 COBH3

v t e Fluorine compounds
v t e Salts and covalent derivatives of the fluoride ion HF ?HeF2 LiF BeF2 BFBF3B2F4+BO3 CF4CxFy+CO3 NF3FN3N2F2NFN2F4NF2?NF5+N+NO3 OF2O2F2OFO3F2O4F2?OF4 F2 Ne NaF MgF2 AlFAlF3 SiF4 P2F4PF3PF5+PO4 S2F2SF2S2F4SF3SF4S2F10SF6+SO4 ClFClF3ClF5 ?ArF2?ArF4 KF CaFCaF2 ScF3 TiF2TiF3TiF4 VF2VF3VF4VF5 CrF2CrF3CrF4CrF5?CrF6 MnF2MnF3MnF4?MnF5 FeF2FeF3FeF4 CoF2 CoF3 CoF4 NiF2NiF3NiF4 CuFCuF2?CuF3 ZnF2 GaF2GaF3 GeF2GeF4 AsF3AsF5 Se2F2SeF4SeF6+SeO3 BrFBrF3BrF5 KrF2?KrF4?KrF6 RbF SrFSrF2 YF3 ZrF2ZrF3ZrF4 NbF4NbF5 MoF4MoF5MoF6 TcF4TcF5 TcF6 RuF3RuF4RuF5RuF6 RhF3RhF4RhF5RhF6 PdF2Pd[PdF6]PdF4?PdF6 Ag2FAgFAgF2AgF3 CdF2 InFInF3 SnF2SnF4 SbF3SbF5 TeF4?Te2F10TeF6+TeO3 IFIF3IF5IF7+IO3 XeF2XeF4XeF6?XeF8 CsF BaF2   LuF3 HfF4 TaF5 WF4WF5WF6 ReF4ReF5ReF6ReF7 OsF4OsF5OsF6?OsF7?OsF8 IrF2IrF3IrF4IrF5IrF6 PtF2Pt[PtF6]PtF4PtF5PtF6 AuFAuF3Au2F10?AuF6AuF5•F2 Hg2F2HgF2?HgF4 TlFTlF3 PbF2PbF4 BiF3BiF5 PoF2PoF4PoF6 AtF?AtF3?AtF5 RnF2?RnF4?RnF6 FrF RaF2   LrF3 Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og ↓ LaF3 CeF3CeF4 PrF3 PrF4 NdF2 NdF3 NdF4 PmF3 SmF SmF2 SmF3 EuF2 EuF3 GdF3 TbF3 TbF4 DyF2 DyF3 DyF4 HoF3 ErF3 TmF2 TmF3 YbF2 YbF3 AcF3 ThF2ThF3ThF4 PaF4PaF5 UF3UF4UF5UF6 NpF3NpF4NpF5NpF6 PuF3PuF4PuF5PuF6 AmF2AmF3AmF4?AmF6 CmF3CmF4 ?CmF6 BkF3 BkF4 CfF3 CfF4 EsF3 EsF4?EsF6 Fm MdF3 No
PF−6, AsF−6, SbF−6 compounds	AgPF6 HPF6 KPF6 LiPF6 NaPF6 NH4PF6 TlPF6 KAsF6 LiAsF6 NaAsF6 HSbF6 KSbF6 LiSbF6 NaSbF6
AlF2−5, AlF3−6 compounds	Cs2AlF5 Li3AlF6 (NH4)3[AlF6] K3AlF6 Na3AlF6
chlorides, bromides, iodides and pseudohalogenides	BaClF BrSO3F Br(SO3F)3 CFN ClFO2 PbFBr PbFCl SiIBrClF SrFCl
SiF2−6, GeF2−6 compounds	BaGeF6 Li2GeF6 BaSiF6 Na2[SiF6] (NH4)2SiF6 K2[SiF6] Li2SiF6
Oxyfluorides	AcOF BrOF3 BrO2F BrO3F C2F4O C7H5FO ClOF3 ClO2F3 CrOF4 ErOF HoOF LaOF NdOF NpO2F2 OsOF5 PrOF PuOF PuO2F2 SmOF TbOF TcO3F ThOF2 UO2F2 VOF3 WOF4 YOF
Organofluorides	CBrF3 CBr2F2 CBr3F CClF3 CCl2F2 CCl3F CFNO3S CF2O CF3I CHF3 CH2F2 CH3F C2Cl3F3 C2H3F C3H5F C6H5F C6H11F C7H5F3 C15F33N
with transition metal, lanthanide, actinide, ammonium	(NH4)3CrF6 NH4F (NH4)3FeF6 (NH4)3GaF6 (NH4)2GeF6 (NH4)3InF6 NH4NbF6 (NH4)2SnF6 NH4TaF6 (NH4)3VF6 (NH4)2ZrF6 CsXeF7 Li2SnF6 Li2TiF6 LiWF6 Li2ZrF6 K2NbF7 K2TaF7 K2TiF6 K2ZrF6 Na2TiF6 Na2ZrF6 Rb2TiF6
nitric acids	FNO FNO2 FNO3
bifluorides	KHF2 NaHF2 NH4HF2
thionyl, phosphoryl, and iodosyl	FN3O2S F2OS3 F2OS F3OP PSF3 IOF3 IO3F F2O6S2 F2O5S2 FClO5S2 ISO3F IOF5 IO2F IO2F3 BrFO2S I3SO3F S3O8F2
Chemical formulas

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