Boron monofluoride

Last updated
Boron monofluoride
Names
Other names
Boron fluoride

Boron(I) fluoride
Fluoroboronene

Fluoroborylene
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.033.970 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 237-383-0
PubChem CID
UNII
  • InChI=1S/BF/c1-2 Yes check.svgY
    Key: YFSQMOVEGCCDJL-UHFFFAOYSA-N Yes check.svgY
  • [BH0]F
  • [B-]=[F+]
  • [B-2]#[F+2]
Properties
BF
Molar mass 29.81 g·mol−1
Thermochemistry
Std molar
entropy (S⦵298)
200.48 J K−1 mol−1
115.90 kJ mol−1
Related compounds
Carbon monoxide, dinitrogen, nitrosonium, cyanide, acetylide
Related compounds
 aluminium monofluoride
aluminium monochloride
aluminium monoiodide
gallium monofluoride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Boron monofluoride or fluoroborylene is a chemical compound with the formula BF, one atom of boron and one of fluorine. It is an unstable gas, but it is a stable ligand on transition metals, in the same way as carbon monoxide. It is a subhalide, containing fewer than the normal number of fluorine atoms, compared with boron trifluoride. It can also be called a borylene, as it contains boron with two unshared electrons. BF is isoelectronic with carbon monoxide and dinitrogen; each molecule has 14 electrons. [1]

Contents

Structure

The experimental B–F bond length is 1.26267  Å. [2] [3] [4] Despite being isoelectronic to the triple-bonded species CO and N2, computational studies generally agree that the true bond order is much lower than 3. One reported computed bond order for the molecule is 1.4, compared with 2.6 for CO and 3.0 for N2. [5]

Lewis dot diagram structures show three formal alternatives for describing bonding in boron monofluoride. BF resonance lewis structures.svg
Lewis dot diagram structures show three formal alternatives for describing bonding in boron monofluoride.

BF is unusual in that the dipole moment is inverted with fluorine having a positive charge even though it is the more electronegative element. This is explained by the 2sp orbitals of boron being reoriented and having a higher electron density. Backbonding, or the transfer of π orbital electrons for the fluorine atom, is not required to explain the polarization. [6]

Preparation

Boron monofluoride can be prepared by passing boron trifluoride gas at 2000 °C over a boron rod. It can be condensed at liquid nitrogen temperatures (−196 °C). [7]

Properties

Boron monofluoride molecules have a dissociation energy of 7.8 eV or heat of formation −27.5±3 kcal/mole [1] [8] or 757±14 kJ/mol. [2] The first ionization potential is 11.115 eV. [2] The spectroscopic constants vibrational frequency ωe of BF+ (X 2Σ+) is 1765 cm−1 and for neutral BF (X 1Σ+) it is 1402.1 cm−1. [2] [9] The anharmonicity of BF is 11.84 cm−1. [9]

Reactions

BF can react with itself to form polymers of boron containing fluorine with between 10 and 14 boron atoms. BF reacts with BF3 to form B2F4. BF and B2F4 further combine to form B3F5. B3F5 is unstable above −50 °C and forms B8F12. This substance is a yellow oil. [7]

BF reacts with acetylenes to make the 1,4-diboracyclohexadiene ring system. BF can condense with 2-butyne forming 1,4-difluoro-2,3,5,6-tetramethyl-1,4-diboracyclohexadiene. Also, it reacts with acetylene to make 1,4-difluoro-1,4-diboracyclohexadiene. [7] Propene reacts to make a mix of cyclic and non-cyclic molecules which may contain BF or BF2. [2]

BF hardly reacts with C2F4 or SiF4. [2] BF does react with arsine, carbon monoxide, phosphorus trifluoride, phosphine, and phosphorus trichloride to make adducts like (BF2)3B•AsH3, (BF2)3B•CO, (BF2)3B•PF3, (BF2)3B•PH3, and (BF2)3B•PCl3. [2]

BF reacts with oxygen: BF + O2 → OBF + O; with chlorine: BF + Cl2 → ClBF + Cl; and with nitrogen dioxide BF + NO2 → OBF + NO. [10]

Ligand

A naïve analysis would suggest that BF is isoelectronic with carbon monoxide (CO) and so could form similar compounds to metal carbonyls. As discussed above (see  § Structure), BF has a much lower bond order, so that the valence shell around boron is unfilled. Consequently, BF as a ligand is much more Lewis acidic; it tends to form higher-order bonds to metal centers, and can also bridge between two or three metal atoms (μ2 and μ3). [11]

Working with BF as a ligand is difficult due to its instability in the free state. [12] Instead, most routes tend to use derivatives of BF3 that decompose once coordinated.

In a 1968 conference report, Kämpfer et al claimed to produce Fe(BF)(CO)4 via reaction of B2F4 with Fe(CO)5, but modern chemists have not reproduced the synthesis, and the original compound has no crystallographic characterization. [13] [14] The first modern demonstration of BF coordinated to a transition element is due to Vidovic and Aldrige, who produced [(C5H5)Ru(CO)2]2(μ2-BF) (with BF bridging both ruthenium atoms) in 2009. [15] To make the compound, Vidovic and Aldridge reacted NaRu(CO)2(C5H5) with (Et2O)·BF3; the boron monofluoride ligand then formed in-place. [14]

Vidovic and Aldridge also developed a substance with the formula (PF3)4FeBF by reacting iron vapour with B2F4 and PF3. [2] Hafnium, thorium, titanium, and zirconium can form a difluoride with a BF ligand at the low temperature of 6K. These come about by reacting the atomic metal with BF3. [2]

The first fully characterized molecule featuring BF as a terminal ligand was synthesized by Drance and Figueroa in 2019, by sterically hindering the formation of a dimer. In the molecule, boron is double-bonded to iron. [16]

FBScF2, FBYF2, FBLaF2, and FBCeF2 have been prepared in a solid neon matrix by reacting atomic metals with boron trifluoride. [17]

Related Research Articles

In coordination chemistry, a coordinate covalent bond, also known as a dative bond, dipolar bond, or coordinate bond is a kind of two-center, two-electron covalent bond in which the two electrons derive from the same atom. The bonding of metal ions to ligands involves this kind of interaction. This type of interaction is central to Lewis acid–base theory.

<span class="mw-page-title-main">Lewis acids and bases</span> Chemical bond theory

A Lewis acid (named for the American physical chemist Gilbert N. Lewis) is a chemical species that contains an empty orbital which is capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base, then, is any species that has a filled orbital containing an electron pair which is not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base, because it can donate its lone pair of electrons. Trimethylborane () is a Lewis acid as it is capable of accepting a lone pair. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, a lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3. The terminology refers to the contributions of Gilbert N. Lewis.

In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

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.

Phosphorus trifluoride (formula PF3), is a colorless and odorless gas. It is highly toxic and reacts slowly with water. Its main use is as a ligand in metal complexes. As a ligand, it parallels carbon monoxide in metal carbonyls, and indeed its toxicity is due to its binding with the iron in blood hemoglobin in a similar way to carbon monoxide.

<span class="mw-page-title-main">Tetrafluoroborate</span> Anion

Tetrafluoroborate is the anion BF−
4. This tetrahedral species is isoelectronic with tetrafluoroberyllate (BeF2−
4), tetrafluoromethane (CF4), and tetrafluoroammonium (NF+
4) and is valence isoelectronic with many stable and important species including the perchlorate anion, ClO−
4, which is used in similar ways in the laboratory. It arises by the reaction of fluoride salts with the Lewis acid BF3, treatment of tetrafluoroboric acid with base, or by treatment of boric acid with hydrofluoric acid.

<span class="mw-page-title-main">Selenium tetrafluoride</span> Chemical compound

Selenium tetrafluoride (SeF4) is an inorganic compound. It is a colourless liquid that reacts readily with water. It can be used as a fluorinating reagent in organic syntheses (fluorination of alcohols, carboxylic acids or carbonyl compounds) and has advantages over sulfur tetrafluoride in that milder conditions can be employed and it is a liquid rather than a gas.

<span class="mw-page-title-main">Tris(pentafluorophenyl)borane</span> Chemical compound

Tris(pentafluorophenyl)borane, sometimes referred to as "BCF", is the chemical compound (C6F5)3B. It is a white, volatile solid. The molecule consists of three pentafluorophenyl groups attached in a "paddle-wheel" manner to a central boron atom; the BC3 core is planar. It has been described as the “ideal Lewis acid” because of its high thermal stability and the relative inertness of the B-C bonds. Related fluoro-substituted boron compounds, such as those containing B−CF3 groups, decompose with formation of B-F bonds. Tris(pentafluorophenyl)borane is thermally stable at temperatures well over 200 °C, resistant to oxygen, and water-tolerant.

<span class="mw-page-title-main">Boron compounds</span>

Boron compounds are compounds containing the element boron. In the most familiar compounds, boron has the formal oxidation state +3. These include oxides, sulfides, nitrides, and halides.

<span class="mw-page-title-main">Dioxygenyl</span> Chemical compound

The dioxygenyl ion, O+
2, is a rarely-encountered oxycation in which both oxygen atoms have a formal oxidation state of +1/2. It is formally derived from oxygen by the removal of an electron:

In chemistry, an ate complex is a salt formed by the reaction of a Lewis acid with a Lewis base whereby the central atom increases its valence and gains a negative formal charge..

<span class="mw-page-title-main">Thiazyl fluoride</span> Chemical compound

Thiazyl fluoride, NSF, is a colourless, pungent gas at room temperature and condenses to a pale yellow liquid at 0.4 °C. Along with thiazyl trifluoride, NSF3, it is an important precursor to sulfur-nitrogen-fluorine compounds. It is notable for its extreme hygroscopicity.

Diboron tetrafluoride is the inorganic compound with the formula (BF2)2. A colorless gas, the compound has a halflife of days at room temperature. It is the most stable of the diboron tetrahalides, and does not appreciably decompose under standard conditions.

<span class="mw-page-title-main">Tetrafluoroammonium</span> Chemical compound

The tetrafluoroammonium cation is a positively charged polyatomic ion with chemical formula NF+
4. It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine. Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF
4, trifluoramine oxide ONF
3 and the tetrafluoroborate BF−
4 anion.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

Borane, also known as borine, is an unstable and highly reactive molecule with the chemical formula BH
3. The preparation of borane carbonyl, BH3(CO), played an important role in exploring the chemistry of boranes, as it indicated the likely existence of the borane molecule. However, the molecular species BH3 is a very strong Lewis acid. Consequently, it is highly reactive and can only be observed directly as a continuously produced, transitory, product in a flow system or from the reaction of laser ablated atomic boron with hydrogen. It normally dimerizes to diborane in the absence of other chemicals.

<span class="mw-page-title-main">Fluorine azide</span> Chemical compound

Fluorine azide or triazadienyl fluoride is a yellow green gas composed of nitrogen and fluorine with formula FN3. Its properties resemble those of ClN3, BrN3, and IN3. The bond between the fluorine atom and the nitrogen is very weak, leading to this substance being very unstable and prone to explosion. Calculations show the F–N–N angle to be around 102° with a straight line of 3 nitrogen atoms.

<span class="mw-page-title-main">Boron monofluoride monoxide</span> Chemical compound

Boron monofluoride monoxide or oxoboryl fluoride or fluoroxoborane is an unstable inorganic molecular substance with formula FBO. It is also called boron fluoride oxide, fluoro(oxo)borane or fluoro-oxoborane. The molecule is stable at high temperatures, but below 1000 °C condenses to a trimer (BOF)3 called trifluoroboroxin. FBO can be isolated as a triatomic non-metallic molecule in an inert gas matrix, and has been condensed in solid neon and argon. When an attempt is made to condense the gas to a solid in bulk, a polymeric glass is formed, which is deficient in fluoride, and when heated forms a glassy froth like popcorn. Boron fluoride oxide has been studied because of its production in high energy rocket fuels that contain boron and fluorine, and in the form of an oxyfluoride glass. BOF glass is unusual in that it can condense directly from gas.

<i>N</i>-Heterocyclic carbene boryl anion Isoelectronic structure

An N-heterocyclic carbene boryl anion is an isoelectronic structure of an N-heterocyclic carbene (NHC), where the carbene carbon is replaced with a boron atom that has a -1 charge. NHC boryl anions have a planar geometry, and the boron atom is considered to be sp2-hybridized. They serve as extremely strong bases, as they are very nucleophilic. They also have a very strong trans influence, due to the σ-donation coming from the boron atom. NHC boryl anions have stronger electron-releasing character when compared to normal NHCs. These characteristics make NHC boryl anions key ligands in many applications, such as polycyclic aromatic hydrocarbons, and more commonly low oxidation state main group element bonding.

<span class="mw-page-title-main">Tetrahalodiboranes</span> Class of diboron compounds

Tetrahalodiboranes are a class of diboron compounds with the formula . These compounds were first discovered in the 1920s, but, after some interest in the middle of the 20th century, were largely ignored in research. Compared to other diboron compounds, tetrahalodiboranes are fairly unstable and historically have been difficult to prepare; thus, their use in synthetic chemistry is largely unexplored, and research on tetrahalodiboranes has stemmed from fundamental interest in their reactivity. Recently, there has been a resurgence in interest in tetrahalodiboranes, particularly in diboron tetrafluoride as a reagent to promote doping of silicon with for use in semiconductor devices.

References

  1. 1 2 Hildenbrand, Donald L.; Murad, Edmond (1965). "Dissociation Energy of Boron Monofluoride from Mass-Spectrometric Studies". The Journal of Chemical Physics. 43 (4): 1400. Bibcode:1965JChPh..43.1400H. doi:10.1063/1.1696932.
  2. 1 2 3 4 5 6 7 8 9 Vidovic, Dragoslav; Aldridge, Simon (2011). "Coordination chemistry of group 13 monohalides". Chemical Science. 2 (4): 601. doi:10.1039/C0SC00508H.
  3. ↑ Nesbet, R. K. (1964). "Electronic Structure of N2, CO, and BF". The Journal of Chemical Physics. 40 (12): 3619–3633. Bibcode:1964JChPh..40.3619N. doi:10.1063/1.1725063.
  4. ↑ Cazzoli, G.; Cludi, L.; Degli Esposti, C.; Dore, L. (1989). "The millimeter and submillimeter-wave spectrum of boron monofluoride: Equilibrium structure". Journal of Molecular Spectroscopy. 134 (1): 159–167. Bibcode:1989JMoSp.134..159C. doi:10.1016/0022-2852(89)90138-0. ISSN   0022-2852.
  5. ↑ Martinie, R. J.; Bultema, J. J.; van der Wal, M. N.; Burkhart, B. J.; van der Griend, D. A. & de Kock, R. L. (2011). "Bond Order and Chemical Properties of BF, CO, and N2". Journal of Chemical Education. 88 (8): 1094–1097. Bibcode:2011JChEd..88.1094M. doi:10.1021/ed100758t.
  6. ↑ Fantuzzi, Felipe; Cardozo, Thiago Messias; Nascimento, Marco Antonio Chaer (28 May 2015). "Nature of the Chemical Bond and Origin of the Inverted Dipole Moment in Boron Fluoride: A Generalized Valence Bond Approach". The Journal of Physical Chemistry A. 119 (21): 5335–5343. Bibcode:2015JPCA..119.5335F. doi:10.1021/jp510085r. PMID   25531385.
  7. 1 2 3 Timms, P. L. (1972). "Low Temperature Condensation". Advances in Inorganic Chemistry and Radiochemistry. Academic Press. p. 143. ISBN   0-12-023614-1.
  8. ↑ Eyring, Leroy (1967). Advances in High Temperature Chemistry volume 1. Academic Press. p. 70. ISBN   9781483224343.
  9. 1 2 Dyke, John M.; Kirby, Colin; Morris, Alan (1983). "Study of the ionization process BF+ (X 2Σ+ ) ← BF(X 1Σ+ ) by high-temperature photoelectron spectroscopy". J. Chem. Soc., Faraday Trans. 2. 79 (3): 483–490. doi:10.1039/F29837900483.
  10. ↑ Light, G. C.; Herm, R. R.; Matsumoto, J. H. (November 1985). "Kinetics of some gas-phase elementary reactions of boron monofluoride" (PDF). The Journal of Physical Chemistry. 89 (23): 5066–5074. doi:10.1021/j100269a036. Archived (PDF) from the original on June 1, 2022.
  11. ↑ Xu, Liancai; Li, Qian-shu; Xie, Yaoming; King, R. Bruce; Schaefer, Henry F. (15 March 2010). "Major Difference between the Isoelectronic Fluoroborylene and Carbonyl Ligands: Triply Bridging Fluoroborylene Ligands in Fe3(BF)3(CO)9 Isoelectronic with Fe3(CO)12". Inorganic Chemistry. 49 (6): 2996–3001. doi:10.1021/ic902511m. PMID   20143841.
  12. ↑ Xu, Liancai; Li, Qian-shu; King, R. Bruce (May 2012). "Fluoroborylene ligands in binuclear ruthenium carbonyls: Comparison with their iron analogues". Polyhedron. 38 (1): 44–49. doi:10.1016/j.poly.2012.02.003.
  13. ↑ Drance et al. 2019: "Previously, Vidovic and Aldridge reported that two equivalents of the ruthenium-based nucleophile Na[CpRu(CO)2] (Cp, cyclopentadienyl; [C5H5]−) reacts with boron trifluoride diethyl etherate (BF·
    3    Et
    2    O    ) with the formal loss of two equivalents of sodium fluoride (NaF) to produce the bridging BF complex ((μ2-BF)[CpRu(CO)2]2) (20). The latter is the only crystallographically characterized compound in which BF functions as a ligand to a metal center."
  14. 1 2 Xu, L.; Li, Q.-S.; Xie, Y.; King, R. B.; Schaefer, H. F. III (2010). "Binuclear fluoroborylene manganese carbonyls". Inorganica Chimica Acta. 363 (13): 3538–3549. doi:10.1016/j.ica.2010.07.013.
  15. ↑ Vidovic, Dragoslav; Aldridge, Simon (4 May 2009). "Coordination and Activation of the BF Molecule". Angewandte Chemie. 121 (20): 3723–3726. Bibcode:2009AngCh.121.3723V. doi:10.1002/ange.200901022. PMID   19373822.
  16. ↑ Drance, M. J.; Sears, J. D.; Mrse, A. M.; Moore, C. E.; Rheingold, A. L.; Neidig, M. L.; Figueroa, J. S. (2019). "Terminal Coordination of Diatomic Boron Monofluoride to Iron". Science. 363 (6432): 1203–1205. Bibcode:2019Sci...363.1203D. doi: 10.1126/science.aaw6102 . PMID   30872521. S2CID   78094683.
  17. ↑ Xu, Bing; Li, Li; Pu, Zhen; Yu, Wenjie; Li, Wenjing; Wang, Xuefeng (18 February 2019). "Fluoroborylene Complexes FBMF 2 (M = Sc, Y, La, Ce): Matrix Infrared Spectra and Quantum Chemical Calculations". Inorganic Chemistry. 58 (4): 2363–2371. doi:10.1021/acs.inorgchem.8b02801.
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