Bonding description molecular orbital calculations

Group 2 complexes are formally electron deficient and conformationally floppy only small energies (often only 1-2 kcal mol-1) are required to alter their geometries by large amounts (e.g., bond angles by 20° or more). In such cases, the inclusion of electron-correlation effects becomes critical to an accurate description of the molecules structures. Both HF/MP2 and density functional theory (DFT) methods have been applied to organoalkaline earth compounds. DFT approaches, which implicitly incorporate electron correlation in a computationally efficient form, are generally the more widely used. Molecular orbital calculations that successfully reproduce bent... 

The nature of the bonding in heteronuclear gold cluster compounds has been investigated primarily by means of semiempirical molecular orbital calculations and the results of these studies have allowed a number of structural generalizations to be made. These theoretical developments have been discussed in detail in a series of papers by Mingos et al. (210, 256, 279-282) and only a brief description of the results of this work will be given in this section. 

Molecular orbital calculations at various levels have given reasonably satisfactory descriptions of the structure and bonding in and other related species. 

The modern theory of chemical bonding begins with the article The Atom and the Molecule published by the American chemist G. N. Lewis in 1916 [1], In this article, which is still well worth reading, Lewis for the first time associates a single chemical bond with one pair of electrons held in common by the two atoms "After a brief review of Lewis model we turn to a quantum-mechanical description of the simplest of all molecules, viz. the hydrogen molecule ion H J. Since this molecule contains only one electron, the Schrodinger equation can be solved exactly once the distance between the nuclei has been fixed. We shall not write down these solutions since they require the use of a rather exotic coordinate system. Instead we shall show how approximate wavefunctions can be written as linear combinations of atomic orbitals of the two atoms. Finally we shall discuss so-called molecular orbital calculations on the simplest two-electron atom, viz. the hydrogen molecule. 

The description of hypervalent bonding in terms of pd hybridization and two-center orbitals implies that the electron densities in the ft and dz2 orbitals are approximately equal. The description in terms of three-center orbitals. Fig. 15.10, implies that the electron density in the dz2 orbital is negligible. Accurate molecular orbital calculations show that the latter description is closer to reality. 

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