Orbital overlap description

An orbital overlap description of electron delocalization mil dimethylallyl cation H2C=CH—C(CH3)2 is given m Figure 10 2 Figure 10 2a shows the rr bond and the vacant p orbital as independent units Figure 10 2b shows how the units can overlap to give an extended rr orbital that encompasses all three carbons This permits the two rr electrons to be delocalized over three carbons and disperses the positive charge... 

Elements beyond the second row of the periodic table can form bonds to more than four ligands and can be associated with more than an octet of electrons. These features are possible for two reasons. First, elements with > 2 have atomic radii that are large enough to bond to 5, 6, or even more ligands. Second, elements with > 2 have d orbitals whose energies are close to the energies of the valence p orbitals. An orbital overlap description of the bonding in these species relies on the participation of d orbitals of the inner atom. 

FIGURE 1.23 Orbital overlap description of the sp -sp CT bond between the two carbon atoms of ethane. 

In this chapter, we develop a model of bonding that can be applied to molecules as simple as H2 or as complex as chlorophyll. We begin with a description of bonding based on the idea of overlapping atomic orbitals. We then extend the model to include the molecular shapes described in Chapter 9. Next we apply the model to molecules with double and triple bonds. Then we present variations on the orbital overlap model that encompass electrons distributed across three, four, or more atoms, including the extended systems of molecules such as chlorophyll. Finally, we show how to generalize the model to describe the electronic structures of metals and semiconductors. 

We begin with the Lewis structure of PH3, and apply the principles of orbital overlap. The bonding description of the molecule must be consistent with the bond angles, so we must pay particular attention to the orientation of the valence orbitals. 

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. 

Other simplified quantum treatments, such as the Lewis electron pair and orbital overlap models, have proved useful in teaching and they give qualitative predictions of the structures of molecular compounds, but they become unwieldy when applied to solids. They have not proved to be particularly helpful in the description of the complex structures found in inorganic chemistry and have therefore not been widely used in this field. 

B. Description of Homoconjugative Interactions in Terms of Orbital Overlap... 

With these definitions, useful descriptions of non-planar rc-systems and potentially homoconjugated systems have been developed. The descriptions are particularly attractive to experimentalists because orbital overlap is accepted as a major contributing factor to bonding and it is easy to visualize. Although Haddon did not formally define homoaromaticity in his work63 one can use his threshold value of S to define homoaromaticity in the following way ... 

An X-ray photoelectron spectroscopic study of Ni(DPG)2I showed no evidence of trapped valence or any appreciable change in the charge on the metal upon oxidation.97 The site of partial oxidation and hence the electron transport mechanism is still unclear but one explanation of the relatively low conductivity is that the conduction pathway is metal centred and that the M—M distances are too long for effective orbital overlap. Electron transport could be via a phonon-assisted hopping mechanism or, in the Epstein—Conwell description, involve weakly localized electronic states, a band gap (2A) and an activated carrier concentration.101... 

With a little effort, we have constructed the tools—density of states, its decompositions, the crystal orbital overlap population—that allow us to move from a complicated, completely delocalized set of crystal orbitals or Bloch functions to the localized, chemical description. There is no mystery... 

In the spin-coupled description of a molecule such as SF6, the sulfur atom contributes six equivalent, nonorthogonal sp -like hybrids which delocalize onto the fluorine atoms. Each of these two-centre orbitals overlaps with a distorted F(2p) function and the perfect-pairing spin function dominates. Of course, using only 3s, 3px, 3p and 3pz atomic orbitals, we can at most form four linearly independent hybrid orbitals localized on sulfur, with a maximum occupancy of 8 electrons, as in the octet rule. However, the six sulfur+fluorine hybrids which emerge in the spin-coupled description are not linearly dependent, precisely because each of them contains a significant amount of F(2p) character. It is thus clear that the polar nature of the bonding is crucial. 

Theoretical and spectroscopic studies have provided a more thorough and sophisticated description of the bonding, but the simple orbital overlap picture presented above remains qualitatively valid, except possibly with respect to the role of the a component of the bond, which may be less important than previously thought. For the crucial 8and 77 components, however, all theoretical and experimental results are in accord. 

---