Bonding orbital description

The reader can easily see that 2 is essentially an sp2 lone pair AO, 4>i essentially a un h MO, and 3 essentially a on-h MO. In short, the two approaches are equivalent and we have chosen to use the bond orbital description for conceptual simplicity. 

As a final example concerning the difference between the bond orbital and MO descriptions, let us consider how an electrophile might attack lEO to form H20 -/f. According to the bond orbital description 7.36, an electrophile would... 

Ollier than 90°, makes the actual 0 value smaller than 90°. With the bond orbital description, the gauche conformation of XChUOH is predicted by simply requiring one hybrid lone pair of oxygen to be antipcriplanar to a x- specifying various conformations of low-symmetry molecules containing OH groups, we will find it convenient to adopt a bond orbital description. 

An example of a 14-electron AH3BH system is methanol, CH3OH. Its stable conformation 10.70 has the three electron pairs of CH3 staggered with respect to those of OH in the bond orbital description. With a modified methyl group XCH2,... 

Section 2 22 Lewis structures orbital hybridization and molecular orbital descriptions of bonding are all used m organic chemistry Lewis structures are used the most MO descriptions the least All will be used m this text... 

FIGURE 8 2 Hybrid orbital description of the bonding changes that take place at carbon during nucleophilic substitution by the Sn2 mechanism... 

A molecular orbital description of benzene has three tt orbitals that are bonding and three that are antibonding Each of the bonding orbitals is fully occupied (two electrons each) and the antibonding orbitals are vacant... 

Several methods of quantitative description of molecular structure based on the concepts of valence bond theory have been developed. These methods employ orbitals similar to localized valence bond orbitals, but permitting modest delocalization. These orbitals allow many fewer structures to be considered and remove the need for incorporating many ionic structures, in agreement with chemical intuition. To date, these methods have not been as widely applied in organic chemistry as MO calculations. They have, however, been successfully applied to fundamental structural issues. For example, successful quantitative treatments of the structure and energy of benzene and its heterocyclic analogs have been developed. It remains to be seen whether computations based on DFT and modem valence bond theory will come to rival the widely used MO programs in analysis and interpretation of stmcture and reactivity. 

Fig. 4.2 Molecular orbital description of the bonding in acyclic S2N2 and (SN)x...

The molecular orbital description of the bonding in NO is similar to that in N2 or CO (p. 927) but with an extra electron in one of the tt antibonding orbitals. This effectively reduces the bond order from 3 to 2.5 and accounts for the fact that the interatomic N 0 distance (115 pm) is intermediate between that in the triple-bonded NO+ (106 pm) and values typical of double-bonded NO species ( 120 pm). It also interprets the very low ionization energy of the molecule (9.25 eV, compared with 15.6 eV for N2, 14.0 eV for CO, and 12.1 eV for O2). Similarly, the notable reluctance of NO to dimerize can be related both to the geometrical distribution of the unpaired electron over the entire molecule and to the fact that dimerization to 0=N—N=0 leaves the total bond order unchanged (2 x 2.5 = 5). When NO condenses to a liquid, partial dimerization occurs, the cis-form being more stable than the trans-. The pure liquid is colourless, not blue as sometimes stated blue samples owe their colour to traces of the intensely coloured N2O3.6O ) Crystalline nitric oxide is also colourless (not blue) when pure, ° and X-ray diffraction data are best interpreted in terms of weak association into... 

Figure 1.18 A molecular orbital description of the C=C tt bond in ethylene. The lower-energy, tt bonding MO results from a combination of p orbital lobes with the same algebraic sign and is filled. The higher-energy, -tt antibonding MO results from a combination of p orbital lobes with the opposite algebraic signs and is unfilled.

The stability order of alkenes is due to a combination of two factors. One is a stabilizing interaction between the C=C tr bond and adjacent C-H a bonds on substituents. In valence-bond language, the interaction is called hyperconjugation. In a molecular orbital description, there is a bonding MO that extends over the four-atom C=C—< -H grouping, as shown in Figure 6.6. The more substituents that are present on the double bond, the more hyperconjugation there is and the more stable the alkene. 

Having just seen a resonance description of benzene, let s now look at the alternative molecular orbital description. We can construct -tt molecular orbitals for benzene just as we did for 1,3-butadiene in Section 14.1. If six p atomic orbitals combine in a cyclic manner, six benzene molecular orbitals result, as shown in Figure 15.3. The three low-energy molecular orbitals, denoted bonding combinations, and the three high-energy orbitals are antibonding. 

Conjugated enones are more stable than nonconjugated enones for the same reason that conjugated dienes are more stable than nonconjugated dienes (Section 14.1). Interaction between the tt electrons of the C=C bond and the tt electrons of the C=0 group leads to a molecular orbital description for a conjugated enone that shows an interaction of the tt electrons over all four atomic centers (Figure 23.3). 

It may be proper at this stage to lead the reader back to the stage where we constructed the localized orbitals of a CH2 group. At that time two valence orbitals were set aside—the 2pv orbital, and the outer (2s, 2pr) hybrid. Both of these orbitals lie in the. r, y plane. Now in our description of cyclopropane, we used bond orbitals to describe the CC bonding these bond orbitals are derived from in-plane (xy y) hybrids on each carbon. The two hybrids which are required on each carbon atom—in ordet to participate in two bond orbitals—are built precisely from the 2py orbital and the (2s, 2pj.) out combination on each CH2 group. 

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