Boranes three-center bonding

The boranes are electron-deficient compounds (Section 3.8) we cannot write valid Lewis structures for them, because too few electrons are available. For instance, there are 8 atoms in diborane, so we need at least 7 bonds however, there are only 12 valence electrons, and so we can form at most 6 electron-pair bonds. In molecular orbital theory, these electron pairs are regarded as delocalized over the entire molecule, and their bonding power is shared by several atoms. In diborane, for instance, a single electron pair is delocalized over a B—H—B unit. It binds all three atoms together with bond order of 4 for each of the B—H bridging bonds. The molecule has two such bridging three-center bonds (9). 

The many higher boranes such as B5H9 and BgH 2 are similarly electron deficient and cannot be described by a single Lewis structure. They can often be described in terms of a combination of two- and three-center bonds. Alternatively, their structures can be rationalized by electron-counting schemes such as those proposed by Wade. Analysis of the electron density of these molecules by the AIM method shows that there are bond paths between all adjacent pairs of atoms. So from the point of view of the AIM theory there are bonds between each adjacent pair of atoms, but these cannot all be regarded as Lewis two-center, two-electron bonds as is the case in B2H6. 

Three-center Bonding in Boranes Lipscomb s Equations of Balance... 

Fig. 4. Localized two- and three-center bond schemes for the boranes B4H10, BbHii, and BeHio.

The boranes are an extensive series of highly reactive electron-deficient binary compounds of boron and hydrogen. The boranes have three-center bonds. 

In monoborane (BH3), monoalkylboranes RBH2, or dialkylboranes R2BII there is only an electron sextet at the boron atom. In comparison to the more stable electron octet, the boron atom thus lacks two valence electrons. It obtains them by bonding with a suitable electron pair donor. When no better donor is available, the bonding electron pair of the B—H bond of a second borane molecule acts as the donor so that a two-electron, three-center bond is produced. Under these conditions, boranes are consequently present as dimers BH3, for example, as B2H6. Still, small fractions of the monomers appear as minor components in the dissociation equilibrium of the dimer B2H6, for example, thus contains some BH3. 

B5H9 and B5Hn, Longuet-Higgins45 applied the three-center bond theory to these systems. The essence of the approach was to utilize all of the orbitals in these electron-deficient systems to generate an electronic structure involving delocalization of the electrons in the boron framework of the borane. 

The electronic structures of borane clusters were first successfully described using localized three-center- and two-center-two-electron bonds. These treatments have been replaced by the cluster electron-counting rule based on MO methods hence, why bother with the three-center bond model in a book about clusters Let s consider why there is value in a more localized approach. 

This concept of Kekule structures can be extended to the 3D deltahedral borane anionsB H " (6Kekule structures make use of three-center B-B-B bonds instead of the carbon-carbon double bonds in benzenoid Kekule structures. Lipscomb s semitopological method"" " for studying the electron and orbital balance in boron networks containing mixtures of B-B two-center and B-B-B three-center bonds is essential for extending the concept of Kekule structures from 2D benzenoid hydrocarbons to 3D deltahedral boranes. 

Eiiborane is a dimer of borane, BH3. The bonding in diborane is unusual because the hydrogen atoms bridge the two boron atoms with the two monomeric BH3 subunits being bound by two-electron, three-center bonds. Because the boron atom in borane possesses an empty p-orbital, borane is a Lewis acid, and it forms stable complexes upon reaction with tetrahydrofuran (THF) and other ethers, which function as Lewis bases, as illustrated by the formation of a borane-THF complex (Eq. 10.27). 

Hydride abstraction from IPr-borane with tris(pentafluorophenyl)borane was attempted by In s et al. It was found that only half an equivalent of the borane reacted, producing a two-electron three-center bond of the form H2B-H-BH2, which is stabilized at each boron center by an NHC ligand, giving an overall cationic species with [HB(C6Fs)3] as the anion [93]. 

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