Borane
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- For the chemical compound (BH3)2, see diborane.
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.
Organoboranes are compounds containing carbon as well as boron and hydrogen.
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[edit] 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, although triethylborane (TEB) was later used to light the engines of the SR-71 Blackbird high-speed plane.<ref>http://incolor.inebraska.com/hwolfe/history/sr71.pdf</ref>
[edit] Some bonding concepts relevant to boranes
Bonding for elements to the left of carbon is dominated by the fact that they possess fewer valence electrons than valence orbitals. The octet rule, which is only a simple guideline, suggests that such elements must achieve the eight electrons. One way that such electron-deficient elements achieve the octet is by forming Lewis acid-Lewis base adducts.
Discrete BH3 has only six valence electrons, which are involved with B-H bonding utilizing three of boron's four atomic orbitals. Hence one atomic orbital is un-utilized in the bonding; this is the pz orbital, where z is defined as being perpendicular to the plane defined by the BH3. The six electrons 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.
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[edit] Adducts
Consistent with its unsaturation, BH3 forms adducts with many Lewis bases. Notable examples include BH3·THF and ammonia borane (BH3·NH3).
[edit] 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 an aluminium chloride catalyst.<ref>Modern Inorganic Chemistry W.L. Jolly, ISBN 0-07-032760-2</ref>
[edit] Safety information
Boranes are generally at least somewhat toxic; the exposure limit for diborane is 100 parts per billion.
[edit] References
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