Covalent compounds molecular orbitals

We have seen that the crystal-field model provides a basis for explaining many features of transition-metal complexes. In fact, it can be used to explain many observations in addition to those we have discussed. Many lines of evidence show, however, that the bonding between transition-metal ions and ligands must have some covalent character. Molecular-orbital theory (Sections 9.7 and 9.8) can also be used to describe the bonding in complexes, although the application of molecular-orbital theory to coordination compounds is beyond the scope of our discussion. The crystal-field model, although not entirely accurate in all details, provides an adequate and useful first description of the electronic structure of complexes. 

In other words, the valence bonds approach is suitable for compounds showing purely ionic or purely covalent behaviour we require molecular orbitals for a more mature description of the bonding in such materials. So the yellow colour of silver iodide reflects the way the bonding is neither ionic nor covalent. We find, in fact, that the charge clouds of the silver and iodide ions overlap to some extent, allowing change to transfer between them. We will look at charge transfer in more detail on p. 459. 

Still another model to represent the bonding that takes place in covalent compounds is the molecular orbital theory. In the molecular orbital (MO) theory of covalent bonding, atomic orbitals (AOs) on the individual atoms combine to form orbitals that encompass the... 

The relevance of the Mdssbauer effect to chemistry is that now one can reinvestigate compounds such as the tin tetrahalides which were previously thought to be understood. There are molecular orbital calculations, quoted by Goldanskii, based on the assumption that the tin tetrahalides are primarily covalent in character. 

Although the MEG model is essentially ionic in nature, it may also be used to evaluate interactions in partially covalent compounds by appropriate choices of the wave functions representing interacting species. This has been exemplified by Tossell (1985) in a comparative study in which MEG treatment was coupled with an ab initio Self Consistent Field-Molecular Orbital procedure. In this way, Tossell (1985) evaluated the interaction of C03 with Mg in magnesite (MgC03). Representing the ion by a 4-31G wave function, holding its... 

The exceptions to the octet rule described in the previous section, the xenon compounds and the tri-iodide ion, are dealt with by the VSEPR and valence bond theories by assuming that the lowest energy available d orbitals participate in the bonding. This occurs for all main group compounds in which the central atom forms more than four formal covalent bonds, and is collectively known as hypervalence, resulting from the expansion of the valence shell This is referred to in later sections of the book, and the molecular orbital approach is compared with the valence bond theory to show that d orbital participation is unnecessary in some cases. It is essential to note that d orbital participation in bonding of the central atom is dependent upon the symmetry properties of individual compounds and the d orbitals. 

We have used the concepts of the resonance methods many times in previous chapters to explain the chemical behavior of compounds and to describe the structures of compounds that cannot be represented satisfactorily by a single valence-bond structure (e.g., benzene, Section 6-5). We shall assume, therefore, that you are familiar with the qualitative ideas of resonance theory, and that you are aware that the so-called resonance and valence-bond methods are in fact synonymous. The further treatment given here emphasizes more directly the quantum-mechanical nature of valence-bond theory. The basis of molecular-orbital theory also is described and compared with valence-bond theory. First, however, we shall discuss general characteristics of simple covalent bonds that we would expect either theory to explain. 

A striking analogy exists between localized molecular orbital, electron-domain models of organic and other covalently bonded molecules (Figs. 3—8) and ion-packing models of inorganic compounds 47). 

The primary difference between covalent and ionic bonding is that with covalent bonding, we must invoke quantum mechanics. In molecular orbital (MO) theory, molecules are most stable when the bonding MOs or, at most, bonding plus nonbonding MOs, are each filled with two electrons (of opposite spin) and all the antibonding MOs are empty. This forms the quantum mechanical basis of the octet rule for compounds of the p-block elements and the 18-electron rule for d-block elements. Similarly, in the Heider-London (valence bond) treatment... 

But how do we account for the bond angles in water (104°) and ammonia (107°) when the only atomic orbitals are at 90° to each other All the covalent compounds of elements in the row Li to Ne raise this difficulty. Water (H2O) and ammonia (NH3) have angles between their bonds that are roughly tetrahedral and methane (CH4) is exactly tetrahedral but how can the atomic orbitals combine to rationalize this shape The carbon atom has electrons only in the first and second shells, and the Is orbital is too low in energy to contribute to any molecular orbitals, which leaves only the 2s and 2p orbitals. The problem is that the 2p orbitals are at right angles to each other and methane does not have any 90° bonds. (So don t draw any either Remember Chapter 2.). Let us consider exactly where the atoms are in methane and see if we can combine the AOs in such a way as to make satisfactory molecular orbitals. 

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