Covalent Bonding in Molecules

After a discussion on atomic structure, we now consider the electronic structure in molecules. It is clear that, whatever theory is chosen to treat this problem, it must be able to answer Why do molecules form at all For instance, when two H atoms collide, the H2 molecule is formed. However, when two He atoms are brought together, no He2 is formed. 

One way to answer the aforementioned question is Two atoms form a molecule because the energy of the whole is lower than the sum of the energies of its parts. So the theory we need should bring about a decrease in energy when a molecule is formed by two atoms. As in the case of an atom, the energy of a molecule can also be calculated using quantum mechanical methods. In this chapter, we first use the simplest diatomic molecules Hj and H2 as examples to illustrate the basic concepts of chemical bonding. Then we will turn to other more complicated molecules. 

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The building blocks of supramolecular systems are held together by intramolecular interactions and these systems are reversible. This intrinsic property is not only a consequence of the more labile interactions within supermolecules, compared to covalent bonds in molecules, but reversibility is essential for the function expressed by supramolecular systems. Kinetics can never be inferred from thermodynamic studies. For example, the knowledge of a host-guest equilibrium constant does not... 

Double covalent bonds in molecules of oxygen, Oz, and carbon dioxide, C02, and a triple covalent bond in a molecule of nitrogen, N2. 

There are five chapters in Part I Introduction to quantum theory, The electronic structure of atoms, Covalent bonding in molecules, Chemical bonding in condensed phases and Computational chemistry. Since most of the contents of these chapters are covered in popular texts for courses in physical chemistry, quantum chemistry and structural chemistry, it can be safely assumed that readers of this book have some acquaintance with such topics. Consequently, many sections may be viewed as convenient summaries and frequently mathematical formulas are given without derivation. 

Infrared (IR) spectroscopy, this is concerned with the energy changes involved in the stretching and bending of covalent bonds in molecules. 

Soon after the development of the quantum mechanical model of the atom, physicists such as John H. van Vleck (1928) began to investigate a wave-mechanical concept of the chemical bond. The electronic theories of valency, polarity, quantum numbers, and electron distributions in atoms were described, and the valence bond approximation, which depicts covalent bonding in molecules, was built upon these principles. In 1939, Linus Pauling s Nature of the Chemical Bond offered valence bond theory (VBT) as a plausible explanation for bonding in transition metal complexes. His application of VBT to transition metal complexes was supported by Bjerrum s work on stability that suggested electrostatics alone could not account for all bonding characteristics. 

In Chapter 3, these bonding characteristics were described without explanation, because you did not yet have the tools necessary for understanding them. Now that you know more about the electron configurations of atoms, you can begin to understand why atoms form bonds as they do. To describe the formation of covalent bonds in molecules, we use a model called the valence-bond model, but before the assumptions of this model are described, let s revisit some of the important issues relating to the use of models for describing the physical world. 

In the first three solid types (metals, ionic crystals, van der Waals crystals), the forces of interaction that hold the particles together do not act in any preferred direction in covalent crystals, the bonds are formed only in special directions because of the directional character of the covalent bond. The principles governing the direction of bond formation in covalent crystals are the same as those governing the covalent bond in molecules. 

Overall, hybrid orbitals provide a convenient model for using valence-bond theory to describe covalent bonds in molecules in which the molecular geometry conforms to the electron-domain geometry predicted by the VSEPR modeb The picture of hybrid orbitals has limited predictive value. When we know the electron-domain geometry, however, we can employ hybridization to describe the atomic orbitals used by the central atom in bonding. 

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