Orbital overlap model

The Lewis stmcture of a molecule shows how its valence electrons are distributed. These stmctures present simple, yet information-filled views of the bonding in chemical species, hi the remaining sections of this chapter, we build on Lewis stmctures to predict the shapes and some of the properties of molecules. In Chapter 10. we use Lewis stmctures as the starting point to develop orbital overlap models of chemical bonding. 

In this chapter, we develop a model of bonding that can be applied to molecules as simple as H2 or as complex as chlorophyll. We begin with a description of bonding based on the idea of overlapping atomic orbitals. We then extend the model to include the molecular shapes described in Chapter 9. Next we apply the model to molecules with double and triple bonds. Then we present variations on the orbital overlap model that encompass electrons distributed across three, four, or more atoms, including the extended systems of molecules such as chlorophyll. Finally, we show how to generalize the model to describe the electronic structures of metals and semiconductors. 

Orbital overlap models proceed from the following assumptions ... 

Hydrogen sulfide is a toxic gas with the foul odor of rotten eggs. The Lewis structure of H2 S shows two bonds and two lone pairs on the S atom. Experiments show that hydrogen sulfide has a bond angle of 92.1°. We can describe the bonding of H2 S by applying the orbital overlap model. 

Example applies the orbital overlap model to phosphine, 13. ... 

Many of the Lewis structures in Chapter 9 and elsewhere in this book represent molecules that contain double bonds and triple bonds. From simple molecules such as ethylene and acetylene to complex biochemical compounds such as chlorophyll and plastoquinone, multiple bonds are abundant in chemistry. Double bonds and triple bonds can be described by extending the orbital overlap model of bonding. We begin with ethylene, a simple hydrocarbon with the formula C2 H4. 

Of all the concepts used in chemistry, that of the chemical bond is one of the most useful and, at the same time, one of the most difficult. It is useful because it helps us to understand the structures of compounds and their properties, and it is difficult because it is not easy to relate it to the physical theories, such as quantum mechanics, that underlie chemistry. This is not to say that people have not attempted to find a connection between the chemical bond and quantum mechanics. The Lewis (1923) electron pair model and the orbital overlap model (Coulson 1961) are, perhaps, among the better known attempts, but all are a posteriori rationalizations, trying to explain the properties of the empirical nineteenth-century chemical bond in terms of twentieth-century physical concepts. It is unlikely that, left to themselves, theoretical chemists in the twentieth century would have ever created the idea of a chemical bond had not the concept already been central to the language of structural chemistry. To this day the chemical bond remains largely an empirical concept. 

Other simplified quantum treatments, such as the Lewis electron pair and orbital overlap models, have proved useful in teaching and they give qualitative predictions of the structures of molecular compounds, but they become unwieldy when applied to solids. They have not proved to be particularly helpful in the description of the complex structures found in inorganic chemistry and have therefore not been widely used in this field. 

Describe the orbital overlap model of covalent bonding. 

I explain how hybridization reconciles observed molecular shapes with the orbital overlap model. 

What Is the Orbital Overlap Model of Covalent Bonding ... 

Orbital overlap models of methane, ammonia, and water. 

The orbital overlap model describes all double bonds in the same way that we have described a carbon-carbon double bond. In formaldehyde, CHgO, the simplest organic molecule containing a carbon-oxygen double bond, carbon forms sigma bonds to two hydrogens by the overlap of an sj hybrid orbital of carbon and the li atomic orbital of... 

According to the orbital overlap model, the formation of a covalent bond results from the overlap of atomic orbitals. 

B. Orbital Overlap Model of a Carbon-Carbon Double Bond... 

Restricted rotation about the carbon-carbon double bond in ethylene, (a) Orbital overlap model showing the pi bond. (b)The pi bond is broken by rotating the plane of one H — C—H group by 90° with respect to the plane of the other H — C — H group. 

According to the orbital overlap model (Section 1.6F), a triple bond is described in terms of the overlap of sp hybrid orhitals of adjacent carbons to form a sigma bond, the overlap of parallel 2py orbitals to form one pi bond, and the overlap of parallel 2p orbitals to form the second pi bond. In ethyne, each carbon forms a bond to a hydrogen by the overlap of an sp hybrid orbital of carbon with a 15 atomic orbital of hydrogen. 

According to the orbital overlap model, a carbon-carbon double bond consists of one sigma bond formed by the overlap of sp hybrid orbitals and one pi bond formed by the overlap of parallel 2p atomic orbitals. It takes approximately 264 kJ/mol (63 kcal/mol) to break the pi bond in ethylene. 

Account for the molecular geometry of allene in terms of the orbital overlap model. Specifically, explain why all four hydrogen atoms are not in the same plane. 

Orbital overlap model of the bonding in benzene, (a) The carbon, hydrogen framework. The six 2p orbitals, each with one electron, are shown uncombined, (b) The overlap of parallel 2p orbitals forms a continuous pi cloud, shown by one torus above the plane of the ring and a second below the plane of the ring. 

---