Molecular orbital theory bond order

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 ... [Pg.243]

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). [Pg.723]

In Molecular Orbital Theory, the greater the bond order, the more stable is the molecule or ion. Therefore, we predict ... [Pg.137]

The variations in the bond dissociation energies of H2+, H2 and He.,+ and their equilibrium bond lengths are consistent with the expectations from molecular orbital theory. As the bond order increases it would be expected that the bonds formed would be stronger and shorter. [Pg.47]

Lewis dot diagrams of nitric oxide compared to the nitrosonium ion and molecular nitrogen. Each bond contains one electron from each atom. These simple diagrams fail to properly account for the effective bond order of 2.5 predicted by molecular orbital theory and must be only considered as illustrative. The dimer of two nitric oxide molecules has five bonds, which is the same as two individual molecules. Thus, nitric oxide remains dissociated at room temperatures. [Pg.3]

Hybridization is a concept normally associated with the valence bond approach but may also be derived from molecular orbital theory by using second-order perturbation theory in the following way (78, 117, 190). [Pg.21]

In the case of tt complexes of substituted cyclopentadienones, such as the iron tricarbonyl derivatives prepared by Weiss and H libel (30), qualitative molecular-orbital theory (20) predicted a considerable reduction of the ketonic carbonyl bond order. It was observed that the ketonic carbonyl frequency dropped by as much as 65 cm-1, in agreement with theory. A similar explanation can also be provided in terms of valence bond theory (Fig. 14). It has been suggested that n complexing of arenes such as benzene results in loss of aromaticity of the ring in contrast to the dicyclopentadienyl... [Pg.25]

I. Molecules containing one to four electrons. We have already seen the H2 molecule in which there are two electrons in the ou orbital. Two bonding electrons constitute a chemical bond The molecular orbital theory does not restrict itself to even numbers of bonding electrons, and so the bond order is given as one-half the difference between the number of bonding electrons and the number of antibonding electrons ... [Pg.93]

Molecular orbital theory predicts that O2 is paramagnetic, in agreement with experiment. Note that the Lewis structure of O2 does not indicate that it has two unpaired electrons, even through it does imply the presence of a double bond. In fact, the prediction/confirmation of paramagnetism in O2 was one of the early successes of molecular orbital theory. Also, the ions 0+ (dioxygen cation), Oj (superoxide anion), and 0 (peroxide anion) have bond orders 2V2, U/2, and 1, respectively. The experimental energy levels of the molecular orbital for the O2 molecule are shown in Fig. 3.3.3(b). [Pg.95]

Earlier in this chapter, you learned the definition of bond order in the valence bond theory. In molecular orbital theory, the bond order is defined as one-half the difference between the number of electrons in bonding orbitals and the number of electrons in antibonding orbitals. Mathematically, this can be expressed as... [Pg.143]

The first chapter reviews information theoretic methods applied to determination of molecular electronic structure in the context of the Orbital Communication Theory of the chemical bond. The molecule is treated as an information network from which structural information can be obtained. The results of probing molecular charges and bond orders using this method are compared with results obtained by more commonly used means. [Pg.389]

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