Quantum mechanics theory of bonding

This discovery was to be the beginning of the use of exchange terms in the quantum mechanics of atoms and molecules. It became the key factor that shortly afterward allowed Walter Heitler and Fritz London to obtain the first successful quantum mechanical calculation of the covalent bond in the simplest case of a diatomic hydrogen molecule. Exchange terms would also pave the way for the notion of quantum mechanical resonance and the development of the quantum mechanical theories of bonding by Linus Pauling and many others. ... 

Against this position McLaughlin maintains that the coming of quantum mechanics and the quantum mechanical theory of bonding has rendered these emergentist claims untenable. In fact he is very categorical about the prospects for modem day emergentism. 

There is another aspect of McLaughlin s above cited passage that is entirely incorrect, namely his claim that the discovery of the structure of DNA owes something to the quantum mechanical theory of bonding. As a matter of fact there is no connection whatsoever between these two developments. All I can think of to explain McLaughlin s statement is that Pauling was involved in both developments. But of course Pauling rather famously failed to find the structure of DNA and was beaten to it by Crick and Watson. 

The discovery of the structure of DNA was driven almost entirely by the X-ray diffraction evidence that became available to Crick and Watson, courtesy of Wilkins and Franklin. It did not rest on any quantum mechanical calculations or indeed any insights provided by the theoryIt involved model building and cardboard-cut outs of bases. McLaughlin does not say anything whatsoever about pre-quantum mechanical theories of bonding, except to imply that they were completely inadequate. At the same time he suggests that the quantum mechanical theory has provided a complete answer to the question of bonding. Neither of these extreme positions are correct. 

In any case, as McLaughlin himself seems to concede, the advent of a quantum mechanical theory of bonding did not in fact kill off emergentism completely since some prominent biologists and neurophysiologists such as Roger Sperry, whom he cites, continued to work in this tradition. Moreover, if one surveys the literature in science as well as philosophy of science, one cannot fail to be struck by the... 

In molecular crystals, there are two levels of bonding intra—within the molecules, and inter—between the molecules. The former is usually covalent or ionic, while the latter results from photons being exchanged between molecules (or atoms) rather than electrons, as in the case of covalent bonds. The hardnesses of these crystals is determined by the latter. The first quantum mechanical theory of these forces was developed by London so they are known as London forces (they are also called Van der Waals, dispersion, or dipole-dipole forces). 

In this section, you have used Lewis structures to represent bonding in ionic and covalent compounds, and have applied the quantum mechanical theory of the atom to enhance your understanding of bonding. All chemical bonds—whether their predominant character is ionic, covalent, or between the two—result from the atomic structure and properties of the bonding atoms. In the next section, you will learn how the positions of atoms in a compound, and the arrangement of the bonding and lone pairs of electrons, produce molecules with characteristic shapes. These shapes, and the forces that arise from them, are intimately linked to the physical properties of substances, as you will see in the final section of the chapter. 

The existing phenomenological theories of catalysis bear approximately the same relation to the electron theory as the theory of the chemical bond, which was prevalent in the last century and which made use of valence signs (and dealt only with these signs), bears to the modern quantum-mechanical theory of the chemical bond which has given the old valence signs physical content, thereby disclosing the physical nature of the chemical forces. 

After F. London, who developed a quantum-mechanical theory of the origin of these forces and also pioneered many quantum calculations of great consequence to chemistry, including bonding in H2, which will be discussed in Section 21-1. 

Recently Mulliken 5 has given a more general quantum mechanical theory of the bonding in these molecular complexes. He also considers resonance between no-bond configurations A, B with dative configurations A" B+. These latter ones are, however, equivalent to the configurations of Brack-man, which also arise from electron transfer from the N-com-pound or donor (Lewis base) to the c sextet5 -compound or acceptor (Lewis acid, p. 87). 

Van t Hoff postulated free rotation round a single bond in order to explain the lack of cis and Irons isomers in molecules of the type of di-chlorethane. In the light of the quantum mechanical theory of the chemical bond, the free rotation is explained by the axial symmetiy of the a bond between the two carbon atoms. Thus the a bond is not in itself a hindrance to free rotation, but as the rotation occurs the relative configurations of the atoms will be changed, so that the distances between the non-bonded atoms and consequently their energies of interaction will alter. 

But the point particles of physics ignore shape and size that are the axiomatic attributes of the subject of chemistry, be they atoms, molecules, proteins joined in a supposedly particular configuration by "chemical bonds", or transient lipid vesicles or micelles. And where one object ends and another begins is not so self-evident. The notion of a bond that emerges from a quantum mechanical theory of two interacting atoms is not so obvious if those objects are immersed in a sea of their neighbours, forming a solid or liquid. 

It is important to say, from the start, that covalent bonding and other molecular properties can be studied quantum mechanically without reference to molecular orbitals. We are referring to valence bond theory which, being based on the orbital concept for atoms, was, in fact, the first quantum-mechanical theory of the chemical bond. In Chapter 8, we will make a more detailed reference to this method. Our main concern in this and the coming chapters is molecular orbital theory. 

In the LCAO MO description, the H2 molecnle in its ground state has a pair of electrons in a bonding MO, and thus a single bond (that is, its bond order is 1). Later in this chapter, as we describe more complex diatomic molecules in the LCAO approximation, bond orders greater than 1 are discussed. This quantum mechanical definition of bond order generalizes the concept first developed in the Lewis theory of chemical bonding—a shared pair of electrons corresponds to a single bond, two shared pairs to a double bond, and so forth. 

Mulliken worked on valence theory and molecular structure starting in the 1920s. In 1952 he developed a quantum-mechanical theory of the behavior of electron orbitals as different atoms merge to form molecules, and in 1966 he was awarded the Nobel Prize in chemistry for his fundamental work concerning chemical bonds and the electronic structure of molecules by the molecular orbital method. ... 

Hj, is the electron orbital occupation. The BOOP, as defined here within the framework of molecular orbital theory, can be considered the quantum-mechanical equivalent of bond order, in the way it was defined by Pauling [10]. In case the periodicity of the metal molecular orbitals is explicitly considered, the BOOP is equivalent to the crystal orbital overlap population defined by Hoffmann [9]. 

It may be emphasized that the quantum-mechanical theory of homopolar valence gives a very detailed picture of this situation and in some cases makes possible a quantitative description of the bond. ... 

---