Valence theory, description

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

All lone pair orbitals have a node between the two atoms and, hence, have a slightly antibonding character. This destabilizing effect of the lone pair localized molecular orbitals corresponds to the nonbonded repulsions between lone pair atomic orbitals in the valence bond theory. In the MO theory all bonding and antibonding resonance effects can be described as sums of contributions from orthogonal molecular orbitals. Hence, the nonbonded repulsions appear here as intra-orbital antibonding effects in contrast to the valence-bond description. 

The first topic has an important role in the interpretation and calculation of atomic and molecular structures and properties. It is needless to stress the importance of electronic correlation effects, a central topic of research in quantum chemistry. The relativistic formulations are of great importance not only from a formal viewpoint, but also for the increasing number of studies on atoms with high Z values in molecules and materials. Valence theory deserves special attention since it improves the electronic description of molecular systems and reactions with the point of view used by most laboratory chemists. Nuclear motion constitutes a broad research field of great importance to account for the internal molecular dynamics and spectroscopic properties. 

Vol. 30 R.D. Harcourt, Qualitative Valence-Bond Descriptions of Electron-Rich Molecules Pauling 3-Electron Bonds and Increased-Valence Theory. X, 260 pages. 1982. 

What is going to happen in the future to the idea of a bond Coulson gave two possible answers to this interesting question. The work of the next years will have to be more concerned with refining and perhaps simplifying the sort of description already worked out. [42] In a symposium commemorating the 50 years of valence theory which took place in 1970, Coulson went much further. 

The band picture of metals developed by physicists accounts very well for conduction and other electric and magnetic properties. The valence bond description of the bonds in metals related to the concepts of chemistry explains much better than the former theory such properties as lattice energies and bond distances. Today, however, the V.B. picture does not lend itself well to a priori quantitative calculations of these properties and it seems doubtful to what extent a bond in solid lithium with a bond order of o. 11 (with respect to the bond order one in a gas molecule) has any fundamental meaning. There is no doubt, however, that in less typical metals and compounds Pauling s theory is valuable as a counterpart to the band picture, just as the V.B. and the M.O. methods are both of great importance for the description of the constitution of organic molecules. 

We said in Section L6 that chemists use two models for deseribtng covalent bondfi valence bond theory and molecular orbital Uieoiy- Having now seen a valence bond description of the double bond in ethylene, let s also took at a molecular orbital descTipcion. 

The crystal structure of 864(118207)2 90, 91) has shown 864 to be square planar with an 8e-8e bond distance of 2.283(4) A, significantly less than that of 2.34(2) A found in the 8eg molecule (92), indicating some degree of multiple bonding. 8uch a result is consistent with a valence bond description of the molecule involving four structures of type VII. Alternatively the structure can be understood in terms of molecule orbital theory. The circle in structure VIII denotes a closed-shell (aromatic ) six-w-electron system. Of the four tt molecular orbitals,... 

We now outline two approaches to a description of valence in the boron hydrides. The first employs three-center bonds. This particular kind of localized molecular orbital seems most suitable for the smaller, more open hydrides. Its use in the more complex hydrides will require delocalization of the bonding electrons either by a molecular orbital modification or a resonance description. The second approach is simply that of molecular orbitals, which is particularly effective in the more condensed and symmetrical hydrides. These approaches merge as the discussion becomes more complete. It is an important result that filled orbital descriptions are obtainable for the known boron hydrides. Also some remarks about charge distribution in the boron hydrides are possible. But the incompleteness of this valence theory in this nontopological form is indicated by the lack of a large number of unknown hydrides, whose existence would be consistent with these assumptions. 

This figure is identical to Fig. 8.6, which means that we have arrived at a result similar to the valence-bond description made previously. It was, in fact, only after the introduction of equivalent molecular orbitals (which later evolved into almost localized m.o.s) by J.E. Lennard-Jones in 1949 to provide the m.o. equivalent of directed valence bonding (ref. 113) that m.o. theory found widespread application to structural problems in chemistry. 

MULTIREFERENCE PERTURBATION THEORY AND VALENCE BOND DESCRIPTION OF ELECTRONIC STRUCTURES OF MOLECULES... 

It is important to recognize that this conclusion is not a repudiation of valence bond theory. As Shaik and Hiberty pointed out, a correct valence bond theory description of the bonding in methane produces a result that is entirely consistent with the PES result." Rather, the PES result is a reminder that illustrations of C-H bonds formed by overlap of hydrogen Is and carbon sp orbitals should not be considered pictures of reality. 

As with HMO theory, the valence bond description of these resonance structures is limited to wave fimctions involving the resonance electrons (i.e., those that correspond to the n electrons in the MO method). The tr bonds are not explicitly considered. 

The electronic structure of molecules and atoms has long been at the focus of chemists interests. Soon after the electron was discovered in 1897, electron theories of valence were developed [1, 2]. However, at that time, nothing was known about the driving force behind the formation of a chemical bond or an ion. The discovery of quantum mechanics in 1925 made it possible to address the fundamental questions of chemistry. Schro-dinger s equation [3] became the central point in electronic structure theory. In 1927, Condon [4] gave a quantum mechanical explanation of the bond in H2 and initiated molecular orbital theory. In the same year, Heitler and London [5] developed the valence-bond description of H2. 

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