Chemical bonding theoiy

In this chapter we have been looking at three types of chemical bonds covalent bond, ionic bonds and metallic bonds. The bonds are described by using different models and theoiy which introduce the molecular orbitals. These molecular orbitals are formed from atomic orbitals which we heard about in chapter 1. 

The "nature of the chemical bond", the "chemical group effects" are examples of "eoneepts" accepted by group 11 as objects of theoretical investigation. To perform these studies it is allowed to introduce other "concepts" and "quantities" which have a questionable status in the formal theoiy. 

Each atom in this structure has an octet. The carbon atom is surrounded by four bonding electron pairs (an octet), and each chlorine atom has three lone pairs and one bonding pair (an octet). We can simplify our notation by representing bonding electron pairs with dashes, reinforcing the idea that in Lewis theoiy a bonding electron pair is a chemical bond ... 

Valence bond theoiy also introduces the concept of directionality to chemical bonds. For example, we expect the bond formed by the overlap of a p orbital to coincide with the axis along which the p orbital lies. Consider the molecule H2S. Unlike the other molecules that we have encountered, H2S does not have the bond angle that Lewis theoiy and the VSEPR model would lead us to predict. (With four electron domains on the central atom, we would expect the bond angle to be on the order of 109.5 .) In fact, the H—S—H bond angle is 92°. Looking at this in terms... 

The Lewis theoiy of chemical bonding provides a relatively simple way for us to visuaUze the arrangement of electrons in molecules. It is insufficient, however, to ejqilain the differences between the covalent bonds in compounds such as H2, F2, and HF. Although Lewis theory describes the bonds in these three molecules in exactly the same way, they really are quite different from one another, as evidenced by their bond lengths and bond enthalpies listed in Table 9.3. Understanding these differences and why eovalent bonds form in the first place requires a bonding model that combines Lewis s notion of atoms sharing electron pairs and the quantum mechanical descriptions of atomic oibitals. 

Molecular Orbital Theory of the Electronic Structure of Organic Compounds V Molecular Orbital Theoiy of Bond Separation W. J. Hehre, R. Ditchfield, L. Random and J. A. Pople Journal of the American Chemical Society 92 (1970) 4796... 

Quantum mechanics provide many approaches to the description of molecular structure, namely valence bond (VB) theoiy (S-70), molecular orbital (MO) theoiy (11,12), and density functional theoiy (DFT) (13). The former two theories were developed at about the same time, but diverged as competing methods for describing the electronic structure of chemical systems (14). The MO-based methods of calculation have enjoyed great popularity, mainly due to the availability of efficient computer codes. Together with geometiy optimization routines for minima and transition states, the MO methods (DFT included) have become prevalent in applications to molecular structure and reactivity. 

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