Covalent bonds valence bond theory and

How does electron sharing lead to bonding between atoms Two models have been developed to describe covalent bonding valence bond theory and molecular orbital theory. Each model has its strengths and weaknesses, and chemists tend... 

We said in Section 1.5 that chemists use two models for describing covalent bonds valence bond theory and molecular orbital theory. Having now seen the valence bond approach, which uses hybrid atomic orbitals to account for geometry and assumes the overlap of atomic orbitals to account for electron sharing, let s look briefly at the molecular orbital approach to bonding. We ll return to the topic in Chapters 14 and 15 for a more in-depth discussion. 

Throughout the book, theoretical concepts and experimental evidence are integrated An introductory chapter summarizes the principles on which the Periodic Table is established and describes the periodicity of various atomic properties which are relevant to chemical bonding. Symmetry and group theory are introduced to serve as the basis of all molecular orbital treatments of molecules. This basis is then applied to a variety of covalent molecules with discussions of bond lengths and angles and hence molecular shapes. Extensive comparisons of valence bond theory and VSEPR theory with molecular orbital theory are included Metallic bonding is related to electrical conduction and semi-conduction. 

Many solid-state physicists discuss the structure and properties of metals and alloys with use of the band theory, in its several modifications. This theory is also a quantum mechanical theory, which starts with a solution of the wave equation for a single electron, and introduces electron-electron correlation in one or another of several ways. The resonating-valence-bond theory introduces electron-electron correlation in several stages, one of which is by the formation of covalent bonds between adjacent atoms, and another the application of the electroneutrality principle to restrict the acceptable structures to those that involve only M+, M°, and M-. It should be possible to find a relationship between the band-theory calculations and the resonating-covalent-bond theory, but I have been largely unsuccessful in finding such a correlation. I have, for example, not been able to find any trace of the metallic orbital in the band-theory calculations, which thus stand in contrast to the resonating-valence-bond theory, in which the metallic orbital plays a predominant role."... 

L. Pauling and Z. S. Herman, The unsynchronized-resonating-covalent-bond theory of metals, alloys, and intermetallic compounds, in Valence Bond Theory and Chemical Structure, D. J. Klein and N. Trinastic, eds., Elsevier, Amsterdam, 1990, pp. 569-610. 

We said in Section L6 that chemists use two models for deseribtng covalent bondfi valence bond theory and molecular orbital Uieoiy- Having now seen a valence bond description of the double bond in ethylene, let s also took at a molecular orbital descTipcion. 

To systematize the structural problems of the so-called compounds of higher order he reconsidered valency. Thus, he advanced the idea of Hauptvalenz and Nebenvalenz, which could be said to contain the image of modern ionic and covalent bonding theory. Furthermore, he visualized the valence bond not as small sticks existing around atoms in a plane, as conceived by Kekul6, but as positions taken by bond arms on the surfaces of spherical atoms in three dimensions. Could this idea not be a precursor to the modern conception of bond orbitals ... 

How does electron sharing occur Two models have been developed to describe covalent bond formation valence bond theory and molecular orbital theory. Each model has its strengths and weaknesses, and chemists use them interchangeably depending on the circumstances. Valence bond theory is the more easily visualized of the two, so most of the descriptions we ll be using in this book derive from that approach. We ll take a brief look now at both theories and then return for a second look at molecular orbital theory in Section 1.9. 

At present, two quantum mechanical theories are used to describe covalent bond formation and the electronic structure of molecules. Valence bond (VB) theory assumes that the electrons in a molecule occupy atomic orbitals of the individual atoms. It permits us to retain a picture of individual atoms taking part in the bond formation. The second theory, called molecular orbital (MO) theory, assumes the formation of molecular orbitals from the atomic orbitals. Neither theory perfectly explains aU aspects of bonding, but each has contributed something to our understanding of many observed molecular properties. 

There are two quantum mechanical explanations for covalent bond formation valence bond theory and molecular orbital theory. In valence bond theory, hybridized atomic orbitals are formed by the combination and rearrangement of orbitals from the same atom. The hybridized orbitals are aU of equal energy and electron density, and the number of hybridized orbitals is equal to the number of pure atomic orbitals that combine. 

The more recent treatment of the covalent bond, based on the application of the principles of wave mechanics, has developed in two distinct forms, usually termed the valence-bond and molecular-orbital theories, respectively. Although ultimately there is no inconsistency between these two theories, they do in fact approach the problem of chemical binding from different points of view, and we shall generally find that for our purposes the valence-bond treatment is the more suitable. This theory starts from concepts already familiar to the chemist and its conclusions can usually be expressed verbally in qualitative terms the molecular-orbital theory, on the other hand, is more mathematical in its approach and lends itself less readily to such an interpretation. We shall, therefore, first discuss the valency-bond theory, and refer only briefly to the molecular-orbital treatment later in the chapter. 

The application of the (8—n) rule to elements preceding group 4 implies the availability of (8—n) electrons per atom for covalent bond formation and is to this extent artificial unless a mechanism for the provision of these electrons can be proposed. A possible mechanism in the case of zinc, based on the valence-bond treatment of metal theory, has already been outlined in 5.28, but it is difficult to feel satisfied that this is more than an ad hoc explanation designed to explain the observed crystal structure of the element if the structure of zinc were unknown there would be few grounds for treating it as other than a simple divalent element. 

This chapter walks you through the evolution of covalent bond theory, starting with the Lewis dot structures you likely covered in General Chemistry and then advancing to valence bond and molecular orbital theories that stem from quantum mechanics calculations. As with most chemistry, this chapter is purely about following the bouncing balls (electrons). 

Angyan, J. G., Loos, M. and Mayer, I. (1994) Covalent Bond Orders and Atomic Valence Indices in the Topological Theory of Atoms in Molecules. J. Phys. Chem., 98, 5244-5248. 

In the previous chapters, we discussed various models of bonding for covalent and polar covalent molecules (the VSEPR and LCP models, valence bond theory, and molecular orbital theory). We shall now turn our focus to a discussion of models describing metallic bonding. We begin with the free electron model, which assumes that the ionized electrons in a metallic solid have been completely removed from the influence of the atoms in the crystal and exist essentially as an electron gas. Freshman chemistry books typically describe this simplified version of metallic bonding as a sea of electrons that is delocalized over all the metal atoms in the crystalline solid. We shall then progress to the band theory of solids, which results from introducing the periodic potential of the crystalline lattice. 

At present, two quantum mechanical theories—valence bond theory and molecular orbital theory—are used to describe covalent bond formation and the electronic struc t ulept ce iMMi tM that the electrons in a molecule... 

With the understanding that electrons occupy regions of space called orbitals, we can now turn our attention to a deeper understanding of covalent bonds. Specifically, a covalent bond is formed from the overlap of atomic orbitals. There are two commonly used theories for describing the nature of atomic orbital overlap valence bond theory and molecular orbital (MO) theory. The valence bond approach is more simphstic in its treatment of bonds, and therefore we will begin our discussion with valence bond theory. 

The VSEPR model is usually a satisfactory method for predicting molecular geometries. To understand bonding and electronic structure, however, you must look to quantum mechanics. We will consider two theories stemming from quantum mechanics valence bond theory and molecular orbital theory. Both use the methods of quantum mechanics but make different simplifying assumptions. In this section, we will look in a qualitative way at the basic ideas involved in valence bond theory, an approximate theory to explain the electron pair or covalent bond by quantum mechanics. 

Angyan JG, Loos M, Mayer I (1994) Covalent bond orders and atomic valence indices in the topological theory of atoms in molecules. J Phys (Them 98 5244-5248... 

We 11 begin our discussion of hydrocarbons by introducing two additional theories of covalent bonding the valence bond model and the molecular orbital model... 

In valence bond theory a covalent bond is described m terms of m phase overlap of a half filled orbital of one atom with a half filled orbital of another When applied to bonding m H2 the orbitals involved are the Is orbitals of two hydrogen atoms and the bond is a ct bond... 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

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