Boron trifluoride

The geometry of a molecule of {{chem2|BF3}} is trigonal planar. Its D_3h 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, {{chem2|CO3(2-)}}.

{{chem2|BF3}} is commonly referred to as "electron deficient," a description that is reinforced by its exothermic reactivity toward Lewis bases.

In the boron trihalides, {{chem2|BX3}}, the length of the B–X bonds (1.30 Å) is shorter than would be expected for single bonds,{{Greenwood&Earnshaw2nd}} 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 {{chem2|BF3}}.{{cite journal | author = Gillespie, Ronald J. | title = Covalent and Ionic Molecules: Why Are BeF₂ and AlF₃ High Melting Point Solids whereas BF₃ and SiF₄ Are Gases? | journal = Journal of Chemical Education | year = 1998 | volume = 75 | issue = 7 | page = 923 | doi = 10.1021/ed075p923 | bibcode = 1998JChEd..75..923G }}

File:Boron-trifluoride-pi-bonding-2D.png

{{chem2|BF3}} is manufactured by the reaction of boron oxides with hydrogen fluoride:

:{{chem2|B2O3 + 6 HF → 2 BF3 + 3 H2O}}

Typically the HF is produced in situ from sulfuric acid and fluorite ({{chem2|CaF2}}).{{cite book |author1=Holleman, A. F. |author2=Wiberg, E. | title = Inorganic Chemistry | publisher = Academic Press | location = San Diego | year = 2001 | isbn = 0-12-352651-5 }} Approximately 2300-4500 tonnes of boron trifluoride are produced every year.{{Ullmann | title = Boron Compounds | doi = 10.1002/14356007.a04_309 | author = Brotherton, R. J. | author2 = Weber, C. J. | author3 = Guibert, C. R. | author4 = Little, J. L. }}

For laboratory scale reactions, {{chem2|BF3}} is usually produced in situ using boron trifluoride etherate, which is a commercially available liquid.{{how?|date=January 2025}}

Laboratory routes to the solvent-free materials are numerous. A well documented route involves the thermal decomposition of diazonium salts of {{chem2|[BF4]-}}:{{ OrgSynth | author = Flood, D. T. | title = Fluorobenzene | year = 1933 | volume = 13 | pages = 46 | collvol = 2 | collvolpages = 295 | prep = CV2P0295 }}

:{{chem2|[PhN2]+[BF4]− → PhF + BF3 + N2}}

It forms by treatment of a mixture boron trioxide and sodium tetrafluoroborate with sulfuric acid:{{cite book|last=Brauer|first=Georg|title=Handbook of Preparative Inorganic Chemistry|volume=1|edition=2nd|date=1963|publisher=Academic Press|location=New York|isbn=978-0121266011|page=220 & 773|url=https://books.google.com/books?id=TLYatwAACAAJ&q=Handbook+of+Preparative+Inorganic+Chemistry}}

:{{chem2|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:{{cite web|url=https://labphoto.tumblr.com/post/181931503047/performing-a-halogen-exchange-a-halex-reaction-on|title=Performing a halogen exchange, a HalEx reaction on....|date=11 Jan 2019|first=Kristof|last=Hegedüs|publisher=Tumblr|access-date=12 January 2025|archive-url=https://web.archive.org/web/20190119032227/https://labphoto.tumblr.com/post/181931503047/performing-a-halogen-exchange-a-halex-reaction-on|archive-date=19 Jan 2019|website=Pictures from an Organic Chemistry Laboratory|url-status=live}}

:3 R–F + BBr₃ → 3 R–Br + BF₃

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).{{cite book | editor = Yaws, C. L. | title = Chemical Properties Handbook | publisher = McGraw-Hill | year = 1999 | page = 25 }}

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.{{cite encyclopedia | publisher = Air Liquide | url = http://encyclopedia.airliquide.com/encyclopedia.asp?GasID=68 | encyclopedia = Gas Encyclopedia | title = Boron trifluoride | url-status = live | archive-url = https://web.archive.org/web/20061206053305/http://encyclopedia.airliquide.com/Encyclopedia.asp?GasID=68 | archive-date = 2006-12-06 |date = 2016-12-15}}

Unlike the aluminium and gallium trihalides, the boron trihalides are all monomeric. They undergo rapid halide exchange reactions:

:{{chem2|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:

:{{chem2|CsF + BF3 → Cs[BF4]}}

:{{chem2|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, ({{chem2|BF3*O(CH2CH3)2}}) is a conveniently handled liquid and consequently is widely encountered as a laboratory source of {{chem2|BF3}}.{{cite book |doi=10.1002/9780470842898.rb249.pub2|chapter=Boron Trifluoride Etherate|title=Encyclopedia of Reagents for Organic Synthesis|year=2007|last1=Cornel|first1=Veronica|last2=Lovely|first2=Carl J.|isbn=978-0471936237|s2cid=100921225 }} Another common adduct is the adduct with dimethyl sulfide ({{chem2|BF3*S(CH3)2}}), which can be handled as a neat liquid.{{cite book |doi=10.1002/047084289X.rb247|chapter=Boron Trifluoride-Dimethyl Sulfide|title=Encyclopedia of Reagents for Organic Synthesis|year=2001|last1=Heaney|first1=Harry|isbn=0471936235}}

All three lighter boron trihalides, {{chem2|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:

:{{chem2|BF3}} < {{chem2|BCl3|link=Boron trichloride}} < {{chem2|BBr3|link=Boron tribromide}} < {{chem2|BI3|link=Boron triiodide}} (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 {{chem2|BX3}} molecule.{{Cotton&Wilkinson6th }} which follows this trend:

:{{chem2|BF3}} > {{chem2|BCl3}} > {{chem2|BBr3}} < {{chem2|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. 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 p_z of F is readily and easily donated and overlapped to empty p_z 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 {{chem2|BF3}} is attributed to the relative weakness of the bond in the adducts {{chem2|F3B\sL}}.{{cite journal |author1=Boorman, P. M. |author2=Potts, D. | title = Group V Chalcogenide Complexes of Boron Trihalides | journal = Canadian Journal of Chemistry | year = 1974 | volume = 52 | issue = 11 | pages = 2016–2020 | doi = 10.1139/v74-291 }}{{cite journal |author1=Brinck, T. |author2=Murray, J. S. |author3=Politzer, P. | title = A Computational Analysis of the Bonding in Boron Trifluoride and Boron Trichloride and their Complexes with Ammonia |journal = Inorganic Chemistry | year = 1993 | volume = 32 | issue = 12 | pages = 2622–2625 | doi = 10.1021/ic00064a008 }}

Yet another explanation might be found in the fact that the p_z 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.{{OR-inline|date=May 2025}}

Boron trifluoride reacts with water to give boric acid and fluoroboric acid. The reaction commences with the formation of the aquo adduct, {{chem2|H2O\sBF3}}, which then loses HF that gives fluoroboric acid with boron trifluoride.{{cite journal | author = Wamser, C. A. | title = Equilibria in the System Boron Trifluoride–Water at 25° | journal = Journal of the American Chemical Society | year = 1951 | volume = 73 | issue = 1 | pages = 409–416 | doi = 10.1021/ja01145a134 }}

:{{chem2|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 {{chem2|[BCl4]-}} and {{chem2|[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.

Boron trifluoride is most importantly used as a reagent in organic synthesis, typically as a Lewis acid.{{cite encyclopedia | author = Heaney, H. | chapter = Boron Trifluoride | encyclopedia = Encyclopedia of Reagents for Organic Synthesis | year = 2001 | doi = 10.1002/047084289X.rb250 | isbn = 0-471-93623-5 }} Examples include:

• initiates polymerisation reactions of unsaturated compounds, such as polyethers
• as a catalyst in some isomerization, acylation,{{cite journal|last1=Mani|first1=Rama I.|last2=Erbert|first2=Larry H.|last3=Manise|first3=Daniel|title=Boron Trifluoride in the Synthesis of Plant Phenolics: Synthesis of Phenolic Ketones and Phenyl Stryl Ketones|journal=Journal of Tennessee Academy of Science|date=1991|volume=66|issue=1|pages=1–8|url=http://www.tennacadofsci.org/journal/articles/JTAS66-1-1.pdf|access-date=27 October 2016|url-status=dead|archive-url=https://web.archive.org/web/20161027124626/http://www.tennacadofsci.org/journal/articles/JTAS66-1-1.pdf|archive-date=27 October 2016}} alkylation, esterification, dehydration,{{cite journal|last1=Sowa|first1=F. J.|last2=Hennion|first2=G. F.|last3=Nieuwland|first3=J. A.|title=Organic Reactions with Boron Fluoride. IX. The Alkylation of Phenol with Alcohols|journal=Journal of the American Chemical Society|date=1935|volume=57|issue=4|pages=709–711|doi=10.1021/ja01307a034}} condensation, Mukaiyama aldol addition, and other reactions{{cite web | publisher = Honeywell | url = http://www51.honeywell.com/sm/bf3/applications.html | title = Boron Trifluoride (BF₃) Applications | url-status = dead | archive-url = https://web.archive.org/web/20120129231149/http://www51.honeywell.com/sm/bf3/applications.html | archive-date = 2012-01-29 }}{{Citation needed|reason=Honeywell citation unverifiable|date=March 2015}}

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

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.{{cite journal |author1=Gay-Lussac, J. L. |author2=Thénard, L. J. | title = Sur l'acide fluorique | journal = Annales de Chimie | volume = 69 | year = 1809 | pages = 204–220}}{{cite journal |author1=Gay-Lussac, J. L. |author2=Thénard, L. J. | title = Des propriétés de l'acide fluorique et sur-tout de son action sur le métal de la potasse | journal = Mémoires de Physique et de Chimie de la Société d'Arcueil | volume = 2 | year = 1809 | pages = 317–331 | url = https://books.google.com/books?id=Nl87AAAAcAAJ&pg=PA317}}

• List of highly toxic gases

{{Reflist}}

• {{cite web | url = http://www.osha.gov/dts/chemicalsampling/data/CH_221700.html | title = Safety and Health Topics: Boron Trifluoride | publisher = OSHA }}
• {{cite web | url = https://www.cdc.gov/niosh/ipcsneng/neng0231.html | title = BORON TRIFLUORIDE ICSC: 0231 | work = International Chemical Safety Cards | publisher = CDC | access-date = 2017-09-08 | archive-url = https://web.archive.org/web/20171123090449/https://www.cdc.gov/niosh/ipcsneng/neng0231.html | archive-date = 2017-11-23 | url-status = dead }}
• {{cite web | url = http://www.npi.gov.au/substances/boron/index.html | work = National Pollutant Inventory | title = Boron & Compounds: Overview | publisher = Australian Government }}
• {{cite web | url = http://www.npi.gov.au/substances/fluoride-compounds/index.html | work = National Pollutant Inventory | title = Fluoride Compounds: Overview | publisher = Australian Government }}
• {{cite web | url = http://webbook.nist.gov/cgi/cbook.cgi?ID=C7637072 | title = Boron trifluoride | work = WebBook | publisher = NIST }}
• {{cite web | url = http://www51.honeywell.com/sm/bf3/applications.html | title = Boron Trifluoride (BF₃) Applications | publisher = Honeywell | access-date = 2012-02-14 | archive-url = https://web.archive.org/web/20120129231149/http://www51.honeywell.com/sm/bf3/applications.html | archive-date = 2012-01-29 | url-status = dead }}

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