Borane
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A borane is an inorganic chemical compound of boron and hydrogen. The lighter boranes are notably unstable - diborane ignites in air to burn with a green flame - but higher ones are much less so. Decaborane is stable and crystalline, reacting with neither air nor water.
They are named by analogy with the alkanes, which are carbon-hydrogen compounds. The salts of boranes are called borohydrides. The bonding in boranes is not explicable by a standard covalent bonding scheme, and is best described by 3-center-2-electron bonds.
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History
German chemist Alfred Stock was the first scientist to characterize the series of boron-hydrogen compounds by analogy with hydrocarbons. The boranes remained a laboratory curiosity until World War Two, where there was some interest in using uranium borohydride as a volatile uranium compound for isotope separation. Dr Herbert C. Brown, Nobel prize winner in 1979, started working on boranes at the University of Chicago in 1942 under the auspices of this project, and never really stopped.
Borane-based reagants are now widely used in organic synthesis; sodium borohydride is the standard reactant for converting aldehydes and |ketones to alcohols.
The US and USSR both spent very substantial amounts in the fifties and early sixties researching boron-based high energy fuels (ethylboranes, for example) for very fast aircraft such as the XB-70 Valkyrie. The development of advanced surface-to-air missiles made the fast aircraft redundant, and the fuel programs were shut down: boranes were used to light the engines of the SR-71 Blackbird high-speed plane.
Chemistry
The chemistry of boranes is dominated by boron possessing only three valence electrons, but four valence orbitals. Imagining a covalently-bonded system for BH3, boron, with its three valence electrons, will bond to three hydrogen atoms, which in turn each share one electron with the boron to give it a valence electron count of six. However, the octet rule states that a top row atom such as boron must fill its valence orbitals for maximum stability, and therefore possess eight valence electrons, yet the boron in discrete BH3 has only six valence electrons, and hence a vacant p-orbital. The unsaturation of borane results in a highly-reactive species that only exists in the gas phase. It readily dimerises to form diborane and, with larger numbers of boron atoms, clusters.
Cluster formation overcomes the electron deficiency of boranes by utilising a molecular orbital bonding scheme that gives rise to 3-center-2-electron bonds. Using empirical rules developed by K. Wade and later improved by M. Mingos, known as Polyhedral skeletal electron pair theory or Wade's/Mingos' rules, the structure of a boron cluster can often be unambiguously determined from the chemical formula.
Industrial applications
Diborane is manufactured in kilotons annually; it is used as a dopant in semiconductors as well as in organic synthesis. Diborane can be prepared by reacting a hydride agent such as sodium borohydride to boron trifluoride or by adding sodium borohydride to sulfuric acid. It can be produced industrially by reducing borax with aluminium and hydrogen at high pressure with a aluminium chloride catalyst [1].
Safety information
Boranes are generally at least somewhat toxic; the exposure limit for diborane is 100 parts per billion.
References
- [1] Modern Inorganic Chemistry W.L. Jolly, ISBN 0-07-032-760-2de:Borane