Quantum mechanics bond theory

A second quantum mechanical bonding theory is molecular orbital theory. This theory is based on a wave description of electrons. The molecular orbital theory assumes that electrons are not associated with an individual atom but are associated with the entire molecule. Delocalized molecular electrons are not shared by two atoms as in the traditional covalent bond. For the hydrogen molecule, the molecular orbitals are formed by the addition of wave functions for each Is electron in each hydrogen atom. The addition leads to a bonding molecular... 

Inorganic chemistry draws its strength from its great practical utility, and this book presents the subject from the standpoint of applications rather than the customary one of quantum mechanical bonding theory. Since the quintessential subject matter is the properties of the 112 known chemical elements and their compounds, we begin with a consideration of the availability of the commonest elements in the Earth s crust (Table 1.1), hydrosphere (i.e., oceans, lakes, rivers, snowfields, ice caps, and glaciers), and atmosphere, along with brief summary of the production and uses of these elements and their compounds. 

In the early days of quantum mechanics, MO theory received a very bad press compared to valence bond (VB) theory because it predicted the incorrect dissociation behaviour of the hydrogen molecule. This is illustrated... 

The positive sign of the interaction constant has been interpreted in terms of the quantum mechanical resonance theory. If CO2 is represented as a combination of the structures (a) 0=C=0, (h) 0+=C—0, (c) 0 —C=0+, and if the left-hand bond is lengthened, structure (c) will be favored over structure (6). But this would imply a stiffening of the right-hand bond, and therefore a positive interaction force constant. 

Consider the behavior of the average potential and kinetic energies for large R. The forces between uncharged atoms or molecules (other than those due to bond formation) are called van der Waals forces. For two neutral atoms, at least one of which is in an S state, quantum-mechanical perturbation theory shows that the van der Waals force of attraction is proportional to l/R, and the potential energy behaves like... 

The first solution to this problem was produced phenomenologically by Mooser and Pearson. The solution for A B compounds is reproduced in Figure 9. Similar solutions apply not only to A"B semiconductors and insulators, but also to many intermetallic compounds including transition metals. This work provides the first step toward explaining structural and phase transitions in chemically homologous families of binary crystals. It has made the question of the proper treatment of chemical bonding in crystals susceptible to theoretical analysis, whereas formerly work based on mechanical models (ionic compounds) or quantum mechanical perturbation theory (nearly-free-electron metals) made the same problem appear insoluble. 

Quantum Mechanics offers the most comprehensive and most successful explanation of many chemical phenomena such as the nature of valency and bonding as well as chemical reactivity. It has also provided a fundamental explanation of the periodic system of the elements which summarizes a vast amount of empirical chemical knowledge. Quantum Mechanics has become increasingly important in the education of chemistry students. The general principles provided by the theory mean that students can now spend less time memorizing chemical facts and more time in actually thinking about chemistry. 

The main features of the chemical bonding formed by electron pairs were captured in the early days of quantum mechanics by Heitler and London. Their model, which came to be known, as the valence bond (VB) model in its later versions, will serve as our basic tool for developing potential surfaces for molecules undergoing chemical reactions. Here we will review the basic concepts of VB theory and give examples of potential surfaces for bond-breaking processes. 

What Are the Key Ideas The central ideas of this chapter are, first, that electrostatic repulsions between electron pairs determine molecular shapes and, second, that chemical bonds can be discussed in terms of two quantum mechanical theories that describe the distribution of electrons in molecules. 

Lewis s theory of the chemical bond was brilliant, but it was little more than guesswork inspired by insight. Lewis had no way of knowing why an electron pair was so important for the formation of covalent bonds. Valence-bond theory explained the importance of the electron pair in terms of spin-pairing but it could not explain the properties of some molecules. Molecular orbital theory, which is also based on quantum mechanics and was introduced in the late 1920s by Mul-liken and Hund, has proved to be the most successful theory of the chemical bond it overcomes all the deficiencies of Lewis s theory and is easier to use in calculations than valence-bond theory. 

After the discovery of quantum mechanics in 1925 it became evident that the quantum mechanical equations constitute a reliable basis for the theory of molecular structure. It also soon became evident that these equations, such as the Schrodinger wave equation, cannot be solved rigorously for any but the simplest molecules. The development of the theory of molecular structure and the nature of the chemical bond during the past twenty-five years has been in considerable part empirical — based upon the facts of chemistry — but with the interpretation of these facts greatly influenced by quantum mechanical principles and concepts. 

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