Molecular orbital description bonding

VSEPR concept, valence bond description and hybridization, molecular orbital description, bond energies, covalent and van der Waals radii of the elements, intermolecular forces... 

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... 

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... 

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). 

At first sight, the molecular orbital description of N2 looks quite different from the Lewis description ( N=N ). However, it is, in fact, very closely related. We can see their similarity by defining the bond order (b) in molecular orbital theory as the net number of bonds, allowing for the cancellation of bonds by antibonds ... 

The n molecular orbitals described so far involve two atoms, so the orbital pictures look the same for the localized bonding model applied to ethylene and the MO approach applied to molecular oxygen. In the organic molecules described in the introduction to this chapter, however, orbitals spread over three or more atoms. Such delocalized n orbitals can form when more than two p orbitals overlap in the appropriate geometry. In this section, we develop a molecular orbital description for three-atom n systems. In the following sections, we apply the results to larger molecules. 

The basic principles dealing with the molecular orbital description of the bonding in diatomic molecules have been presented in the previous section. However, somewhat different considerations are involved when second-row elements are involved in the bonding because of the differences between s and p orbitals. When the orbitals being combined are p orbitals, the lobes can combine in such a way that the overlap is symmetric around the intemuclear axis. Overlap in this way gives rise to a a bond. This type of overlap involves p orbitals for which the overlap is essentially "end on" as shown in Figure 3.5. For reasons that will become clear later, it will be assumed that the pz orbital is the one used in this type of combination. 

The bonding in the XeF2 molecule can be explained quite simply in terms of a 3-center, 4 electron bond that spans all three atoms in the molecule. The bonding in this molecular orbital description involves the filled 5pz orbital of Xe and the half-filled 2pz orbitals of the two F-atoms. The linear combination of these three atomic orbitals affords one bonding, one non-bonding and one anti-bonding orbital, as depicted below ... 

Molecular orbital description of bonding in iron oxides... 

Pyridine, symmetry group C2v, has six electrons in a system delocalized around the ring, and two lone-pair electrons in an orbital localized at the Nitrogen atom. The Is electrons, as well as the electrons in orbitals describing the a bonds, need not be considered explicitly in describing the resonance stabilization and low-lying excited states of pyridine. The simple molecular orbital description has the following characteristic assumptions ... 

To consider why the two-orbital two-electron single bond formation case can be more complex than often thought, let us consider the H2 system in more detail. In the molecular orbital description of H2, both bonding og and antibonding ou mos appear. 

This picture can qualitatively account for the g tensor anisotropy of nitrosyl complexes in which g = 2.08, gy = 2.01, and g == 2.00. However, gy is often less than 2 and is as small as 1.95 in proteins such as horseradish peroxidase. To explain the reduction in g from the free electron value along the y axis, it is necessary to postulate delocalization of the electron over the molecule. This can best be done by a complete molecular orbital description, but it is instructive to consider the formation of bonding and antibonding orbitals with dy character from the metal orbital and a p orbital from the nitrogen. The filled orbital would then contribute positively to the g value while admixture of the empty orbital would decrease the g value. Thus, the value of gy could be quite variable. The delocalization of the electron into ligand orbitals reduces the occupancy of the metal d/ orbital. This effectively reduces the coefficients of the wavefunction components which account for the g tensor anisotropy hence, the anisotropy is an order of magnitude less than might be expected for a pure ionic d complex in which the unpaired electron resides in the orbital. 

---