B2H6, bonding

The compounds B2X4 are spontaneously flammable in air and react with H2 to give BHX2, B2H6 and related hydrohalides they form adducts with Lewis bases (B2CI4L2 more stable than B2F4L2) and add across C-C multiple bonds, e.g. 

What happens if there are not enough electrons to form conventional two-electron bonds Diborane (B2H6 provides a good example. Were the molecule to look like ethane, how many valence electrons would be required tc hold it together How many valence electrons does diborane possess Examine the actual structure for diborane. 

Lewis s theory also fails to account for the compound diborane, B2H6, a colorless gas that bursts into flame on contact with air. The problem is that diborane has only 12 valence electrons (three from each B atom, one from each H atom) but, for a Lewis structure, it needs at least seven bonds, and therefore 14 electrons, to bind the eight atoms together Diborane is an example of an electron-deficient compound, a compound with too few valence electrons to be assigned a valid Lewis structure. Valence-bond theory can account for the structures of electron-deficient compounds in terms of resonance, but the explanation is not straightforward. 

X 14.111 (a) Examine the structures of diborane, B2H6, and AI2CI6(g), which can be found on the Web site for this book. Compare the bonding in these two je compounds. How are they similar (b) What are the differences, if any, in the types of bonds formed (c) What is the hybridization of the Group 13/III element (d) Are the molecules planar If not, describe their shapes. 

The Tetrahedral Carbon Atom.—We have thus derived the result that an atom in which only s and p eigenfunctions contribute to bond formation and in which the quantization in polar coordinates is broken can form one, two, three, or four equivalent bonds, which are directed toward the corners of a regular tetrahedron (Fig. 4). This calculation provides the quantum mechanical justification of the chemist s tetrahedral carbon atom, present in diamond and all aliphatic carbon compounds, and for the tetrahedral quadrivalent nitrogen atom, the tetrahedral phosphorus atom, as in phosphonium compounds, the tetrahedral boron atom in B2H6 (involving single-electron bonds), and many other such atoms. 

Actually the halides all do have the formula BX3, and show no tendency to associate, while the simplest hydride, B2H6, has not begun to dissociate into 2BH3 at a temperature at which it begins to lose hydrogen. If BH3 were given an electronic formula with an electron-pair bond between boron and each H... 

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. 

Figure 14. Contour plot of the electron density of B2H6 in the plane of the bridging hydrogen. Each hydrogen is connected to the two boron atoms by a bond path to each. In contrast, the boron atoms do not share a bond path linking them to one another. (See legend to Fig. 2 for contour values.)...

Perhaps the most depressing fact associated with the consequences of the above division is the lack of consistency often found in treatments of compounds which are essentially isostructural. Take, for instance, the different descriptions of the bonding situation in B2H6 on the one hand, and the isostructural (e.g. AI2CI6) molecules on the other while the latter may be treated by the conventional bonding principles expressed in Hyps. III.l to III.5, the treatment of the former (in terms of 3-centre bonds) breaks with Hyps. III.l to III.4. A similar conclusion is in fact reached in the majority of abnormal cases. Other simple examples are provided by the alkali-metal hydrides (with NaCl-type structure), CuH (with ZnS-wurtzite type structure), etc. These examples are typical in that it is only when a scarcity of electrons and/or orbitals enforces a search for extraordinary bonding principles that Hyps. III.l to III.4 are reluctantly (partly or completely) replaced by alter-... 

In this chapter many of the basic principles related to structure and bonding in molecules have already been illustrated. However, there is another type of compound that is not satisfactorily described by the principles illustrated so far. The simplest molecule of this type is diborane, B2H6. The problem is that there are only 10 valence shell electrons available for use in describing the bonding in this molecule. 

The BH3 molecule is not stable as a separate entity. This molecule can be stabilized by combining it with another molecule that can donate a pair of electrons (indicated as ) to the boron atom to complete the octet (see Chapter 9). For example, the reaction between pyridine and B2H6 produces C5H5N BH3. Another stable adduct is carbonyl borane, OC BH3 in which a pair of electrons is donated from carbon monoxide, which stabilizes borane. In CO, the carbon atom has a negative formal charge, so it is the "electron-rich" end of the molecule. Because the stable compound is B2H6 rather than BH3, the bonding in that molecule should be explained. 

The planar framework has a bonds as just shown, which involve sp2 hybrid orbitals on the boron atoms. This leaves one unhybridized p orbital that is perpendicular to the plane. The B2H6 molecule can be considered being made by adding two H+ ions to a hypothetical B21142 ion that is isoelectronic with C2H4 because each carbon atom has one more electron than does a boron atom. In the B2I l42 ion, the two additional electrons reside in a tt bond that lies above and below the plane of the structure just shown. When two H+ ions are added, they become attached to the lobes of the n bond to produce a structure, the details of which can be shown as... 

Figure 3.92 The structure of diborane B2H6 (left) and shapes of the overlapping (sp459)b and (s)h natural hybrids contributing to the three-center tBhb bridge bond (right). (Only the outermost hybrid contours are shown, in order to reduce congestion on the diagram.)...

Figure 3.118 Three-center bond (upper row) and antibond (lower two rows) NBOs of A12H6 (left) and Ga2H6 (right) (cf. Fig. 3.93 for the corresponding orbitals of B2H6).

Hydroboration, the addition of a boron-hydrogen bond across an unsaturated moiety, was first discovered by H. C. Brown in 1956. Usually, the reaction does not require a catalyst, and the borane reagent, most commonly diborane (B2H6) or a borane adduct (BH3-THF), reacts rapidly at room temperature to afford, after oxidation, the /AMarkovnikov alkene hydration product. However, when the boron of the hydroborating agent is bonded to heteroatoms which lower the electron deficiency, as is the case in catecholborane (1,3,2-benzodioxaborole) 1 (Scheme 1), elevated temperatures are needed for hydroboration to occur.4 5... 

Based upon a halide-hydride exchange reaction, the first truly practical synthesis of B2H6 was reported by Finholt, Bond, and Schlesinger in 1947 (1 3) (Reaction (2)). 

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