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Boron trifluoride
Boron trifluoride in 2D
Boron trifluoride in 3D
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number 231-569-5
RTECS number ED2275000
UN number Compressed: 1008.
Boron trifluoride dihydrate: 2851.
Properties
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, 146.4 K
Boiling point −100.3 °C, 172.9 K
Reacts
Solubility soluble in benzene, toluene, hexane, chloroform and methylene chloride
Hazards
GHS pictograms Press. GasAcute Tox. 2Skin Corr. 1A
GHS signal word DANGER
H330, H314 [note 1]
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 4: Very short exposure could cause death or major residual injury. E.g., VX gas Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
Flash point 4 °C
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
YesY verify (what is YesYN ?)
Infobox references
  • Tracking categories (test):

Boron trifluoride is the chemical compound with the formula BF3. This pungent colourless 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[edit]

The geometry of a molecule of BF3 is trigonal planar. The 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, CO32−.

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-F bonds (1.30 Å) is shorter than would be expected for single bonds,[3] 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.[3]

Boron trifluoride pi bonding diagram

Synthesis and handling[edit]

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).[4] Approximately 2300-4500 tonnes of boron trifluoride are produced every year.[5]

On a laboratory scale, BF3 is produced by the thermal decomposition of diazonium salts:[6]

PhN2BF4PhF + BF3 + N2

Anhydrous boron trifluoride has a normal boiling temperature 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[edit]

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 → BF3O(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 a widely encountered as a laboratory source of BF3. It is stable as a solution in ether, but not stoichiometrically. Another common adduct is the adduct with dimethyl sulfide (BF3 · S(Me)2), which can be handled as a neat liquid.

Comparative Lewis acidity[edit]

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 (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.[9] which follows this trend:

BF3 > BCl3 > BBr3 (most easily pyramidalized)

The criteria for evaluating the relative strength of π-bonding are not clear, however.[3] 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.[10][11]

Hydrolysis[edit]

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 fluoboric acid with boron trifluoride.[12]

4 BF3 + 3 H2O → 3 HBF4 + "B(OH)3"

The heavier trihalides do not undergo analogous reactions, possibly the lower stability of the tetrahedral ions BX4- (X = Cl, Br). 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.

Coordination Chemistry[edit]

History[edit]

The ability of Lewis acids to coordinate to transition metals as σ-acceptor ligands was recognized as early as in the 1970's, but the so-called Z-type ligands remained curiosities until the early 2000's. Over the last decade, significant progress has been made in this area, especially via the incorporation of Lewis acid moieties into multidentate, ambiphilic ligands. Our understanding of the nature and influence of TM → Z interactions has considerably improved and the scope of Lewis acids susceptible to behave as σ-acceptor ligands has been significantly extended.[13]

Z-Ligand Behavior[edit]

Owing to its vacant orbital over the Boron atom, BF3 has an incomplete octet due to which it readily accepts a pair of electrons from other atoms. Hence, the electron deficit BF3 behaves as a Z-type ligand. Z-type ligands are those that accept two electrons from the metal center as opposed to the electron donation occurring with the other two X and L ligands.[14] A Z‑function ligand interacts with a metal center via a dative covalent bond, differing from the L‑function in that both electrons are donated by the metal rather than the ligand.[15] As such, a BF3 ligand donates zero electrons to a metal center because it is a strong Lewis Acid (electron acceptor). Therefore, the coordination with the metal atom always occurs through the Boron and not the Fluorine atoms.

Molecular orbital diagram showing the dative bond character between the metal center and boron ligand



Although BF3 is a Lewis acid, it behaves as a neutral ligand in the complex without contributing to the overall charge present on the complex. But since the metal uses two of its electrons in forming the M–Z bond, the BF3 ligand raises the valence of the metal center by two units. This means that presence of the BF3 ligand changes the dn configuration of the complex without changing the total e- count.[14]

A Z ligand is usually accompanied by an L ligand, as the presence of the L ligand would add stability to the complex. As the electrons are being donated from the central metal atom to the Z ligand, the L ligand donates its pair of electrons to the metal atom. This unique type of bonding existing between two different ligands and the metal atom renders the BF3 complexes stable when present with a strong sigma donor ligand.[15] In BF3 complexes, the L and Z ligands can be written in terms of X. For example, if one Z (BF3) ligand is accompanied by one L type ligand, it can be written as a complex containing two X type ligands; i.e. MLZ type complex becomes an MLX2 type.[14]

Geometry and Bond Strength[edit]

The metal-boron bond is formed when the metal center donates two electrons to the electron-deficient boron's empty π-orbital.[14] This forms what is called a dative bond. The bond can be further stabilized by the presence of an L-type ligand (on the metal center, which donates two electrons to the metal. This flow of electrons from the L-type ligand to the metal to the boron stabilizes the normally trigonal planar BF3 into a tetrahedral M-BF3 complex.[13]

The bond geometry change for a planar BF3 molecule when bound to a metal.
*note: This is not indicative of a reaction scheme.

Similar Molecules[edit]

There are several other molecules that are similar to BF3 with similar Z-ligand functionality. These can range from the simpler BX3 molecules such as BH3, BCl3, and BR3, to the more complex boron-centered molecules such as B(C6F5)3.[14] In addition, there are many complex boron-centered molecules that act as multiple ligands on a single metal atom, forming "scaffolding" structures.[15]

Uses[edit]

Boron trifluoride is most importantly used as a reagent in organic chemistry, typically as a Lewis acid. Examples:[5][16]

Other uses:

Discovery[edit]

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.[18][19]

Notes[edit]

  1. ^ Within the European Union, the following additional hazard statement (EUH014) must also be displayed on labelling: Reacts violently with water.

References[edit]

  1. ^ 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. OJEU L353, 31.12.2008, pp 1–1355 at p 341.
  2. ^ "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 .
  3. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  4. ^ Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. San Diego: Academic Press. ISBN 0-12-352651-5. 
  5. ^ a b Brotherton, R. J.; Weber, C. J.; Guibert, C. R.; Little, J. L. (2005), "Boron Compounds", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a04_309 
  6. ^ Flood, D. T. (1933). "Fluorobenzene". Org. Synth. 13: 46. ; Coll. Vol., 2, p. 295 
  7. ^ Yaws, C. L., ed. (1999). Chemical Properties Handbook. McGraw-Hill. p. 25. 
  8. ^ "Boron trifluoride". Gas Encyclopedia. Air Liquide. 
  9. ^ 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 
  10. ^ 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. 
  11. ^ 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. 
  12. ^ Wamser, C. A. (1951). "Equilibria in the System Boron Trifluoride–Water at 25°". Journal of the American Chemical Society. 73 (1): 409–416. doi:10.1021/ja01145a134. 
  13. ^ a b Braunschweig, H; Dewhurts, R. D. (2010). "σ-Acceptor, Z-type ligands for transition metals". Dalton Transactions: 859–871. 
  14. ^ a b c d e Green, M. L. H. (1995). "A new approach to the formal classification of covalent compounds of the elements". Journal of Organometallic Chemistry: 127–148. 
  15. ^ a b c Braunschweig, H.; Dewhurst, R. D. (2011). "Transition metals as Lewis bases: "Z-type" boron ligands and metal-to-boron dative bonding". Dalton Transactions: 549–558. 
  16. ^ Heaney, H. (2001). "Boron Trifluoride". Encyclopedia of Reagents for Organic Synthesis. ISBN 0-471-93623-5. doi:10.1002/047084289X.rb250. 
  17. ^ a b "Boron Trifluoride (BF3) Applications". Honeywell. 
  18. ^ Gay-Lussac, J. L.; Thénard, L. J. (1809). "Sur l’acide fluorique". Annales de Chimie. 69: 204–220. 
  19. ^ 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]


Category:Fluorides Category:Boron compounds

Category:Boron halides