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Borane

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Borane

Borane
Structural formula of borane
Ball-and-stick model of borane
Spacefill model of borane
Names
Systematic IUPAC name
borane (substitutive)
trihydridoboron (additive)
Other names
  • borine
  • boron trihydride
Identifiers
ChEBI
ChemSpider
44
Jmol-3D images Image
PubChem
Properties
H3B
Molar mass 13.83 g·mol−1
Appearance colourless gas
hydrolyses
Solubility in Ammonia 3.2 mol L−1
Thermochemistry
187.88 kJ mol−1 K−1
106.69 kJ mol−1
Structure
D3h
trigonal planar
trigonal planar
0 D
Related compounds
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Borane (also systematically named trihydridoboron), also called borine, is an chemical formula BH
3
. It is a colourless gas that only persists at elevated temperatures or in dilution. Borane is the simplest member of the boranes.

Contents

  • History 1
  • Chemical properties 2
    • Preparation of molecular BH3 2.1
    • Structure 2.2
    • Amphotericity 2.3
    • Borane as a reactive intermediate 2.4
    • Solute properties of diborane 2.5
  • Production 3
  • Uses 4
  • References 5

History

In 1937, the discovery of carbonyltrihydridoboron, the adduct of borane with carbon monoxide, among other borane adducts, played an important role in exploring the chemistry of "normal" boranes at a time when three-centre two-electron bonding was not yet known.[1] This discovery also implied the existence of borane, however, it was not until some years later that direct evidence was observed.

Chemical properties

Preparation of molecular BH3

The matrix isolated products of laser ablated boron atoms with hydrogen forms BH3 as a minor constituent along with B2H6, diborane and BH(H2) complex.[2] Studies of gas phase diborane have detected monomeric BH3.

Structure

The structure of BH3 is trigonal planar (D3h molecular symmetry) with an experimentally determined B–H bond length of 119 pm.[3] This the same as the terminal B–H bond length in diborane(6). The dominant behaviour of borane is its dimerisation to form diborane, as shown by the enthalpy of the reaction, which is predicted to be near -40 kcal/mol.[4] For such an exothermic process, the concentration of BH3 is negligible in solution.

2 BH3 → B2H6

Amphotericity

Although borane in principle acts as a Lewis acid, the many 1:1 adducts are invariably prepared from diborane or via ligand exchange of an existing adduct.

B2H6 + 2 L → 2 LBH3

The stability sequence of these estimated spectroscopically and thermochemically is:-[5]

PF3 < CO < Et2O < Me2O < C4H8O < C4H8S < Et2S < Me2S < Py < Me3N < H

BH3 has some soft acid characteristics (sulfur donors are more stable than oxygen donors).[5]

The boryl group (-BH
2
) in boranes such as borane can assimilate another hydrogen centre into the molecule by ionisation:

BH
3
+ H+
BH
2
(H
2
)+
[6]

Because of this capture of the proton (H+
), borane has basic character. Its conjugate acid is (η-dihydrogen)dihydridoboron(1+) ([BH
2
(η-H
2
)
]+). It should be noted that being a Kubas complex, [BH
2
(η-H
2
)
]+ rapidly decomposes and ejects dihydrogen. Borane does not form stable aqueous solutions due to hydrolysis.

BH
3
+ 3 H
2
O
is in a favored equilibrium with B(OH)
3
+ 3 H
2
[7]

Borane as a reactive intermediate

One example where molecular BH3 is believed to be a reaction intermediate is in the pyrolysis of diborane to produce higher boranes:[5]

B2H6 ⇌ 2BH3
BH3 +B2H6 → B3H7 +H2 (rate determining step)
BH3 + B3H7 ⇌ B4H10
B2H6 + B3H7 → BH3 + B4H10
⇌ B5H11 + H2

Further steps give rise to successively higher boranes, with B10H14 as the most stable end product contaminated with polymeric materials, and a little B20H26.

Another example is hydroboration reaction, in which "diborane" adds to an alkene. In this reaction, either the diborane dissociates forming BH3 as an intermediate or one B–H–B bridge opens to produce an electron deficient boron atom.[8] The addition to alkenes is rapid, quantitative and reversible. The addition is anti-markovnikov, that is to say that boron adds to the less substituted C atom, the attack taking place on the less hindered side of the molecule.

Solute properties of diborane

Diborane dissolves in diethyl ether and diglyme and is present as the dimer. In Tetrahydrofuran, THF, it is present as a loose 1:1 adduct, THF•BH3.[9] A diborane solution in THF is commercially available as is a solution of the DMS complex.[8] Gaseous borane, diborane(6) dissolves in polar compounds such as amines and tetrahydrofuran. This dissolution property of diborane makes it a widely used laboratory chemical, for example, as a reagent in the production of ethylborane. Borane is commercially available as a Lewis acid–base adduct with various ligands, including solutions of borane dimethylsulfide, ammonia borane (and other amines), and borane tetrahydrofuran. Borane has a solubility in liquid ammonia of 3.2 mol L−1, above which it precipitates out as the ammoniate. Any attempt to deammoniate the crystallised product thermally, only results in its decomposition.

Production

There are two main methods for producing borane. One common method is the cleaving reaction of diborane with dimethyl sulfide. The other method is the partial oxidation of a boranuide salt under a coordinating borane solvent such as trimethylamine.

Uses

Borane adducts are widely used in hydroboration, where BH3 adds across the C=C bond in alkenes to give trialkylboranes:

(THF)BH3 + 3 CH2=CHR → B(CH2CH2R)3 + THF

This reaction is regioselective, and the product trialkylboranes can be converted to useful organic derivatives. With bulky alkenes one can prepare species such as [HBR2]2, which are also useful reagents in more specialised applications. Borane dimethylsulfide which is more stable than the THF adduct of borane adduct.[10]

References

  1. ^
  2. ^
  3. ^
  4. ^ M. Page, G.F. Adams, J.S. Binkley, C.F. Melius "Dimerization energy of borane" J. Phys. Chem. 1987, vol. 91, pp 2675–2678. doi:10.1021/j100295a001
  5. ^ a b c
  6. ^
  7. ^
  8. ^ a b Hydrocarbon Chemistry, George A. Olah, Arpad Molner, 2d edition, 2003, Wiley-Blackwell ISBN 978-0471417828
  9. ^
  10. ^ Kollonitisch, J., "Reductive Ring Cleavage of Tetrahydrofurans by Diborane", J. Am. Chem. Soc. 1961, volume 83, 1515. doi:10.1021/ja01467a056
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