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Unformatted text preview: 423 A lthough the Lewis theory has been useful in our discussion of chemical bonding, it does have shortcomings. For exam-ple, it does not help explain why metals conduct electricity or how a semiconductor works. While we will continue to use the Lewis theory for most purposes, there are cases that require more sophisticated approaches. One such approach involves the familiar s , p , and d atomic orbitals, or mixed-orbital types called hybrid orbitals . A second approach involves the creation of a set of orbitals that belongs to a molecule as a whole. Electrons are then assigned to these molecular orbitals . Our purpose in this chapter is not to try to master theories of covalent bonding in all their details. We want simply to discover how these theories provide models that yield deeper insights into the nature of chemical bonding than do Lewis structures alone. 11-1 What a Bonding Theory Should Do 11-2 Introduction to the Valence-Bond Method 11-3 Hybridization of Atomic Orbitals 11-4 Multiple Covalent Bonds 11-5 Molecular Orbital Theory 11-6 Delocalized Electrons: Bonding in the Benzene Molecule 11-7 Bonding in Metals 11-8 Some Unresolved Issues; Can Electron Charge-Density Plots Help? ➣ F OCUS O N Photoelectron Spectroscopy C O N T E N T S C HEMICAL B ONDING II: A DDITIONAL A SPECTS Electrostatic potential maps of benzene (one solid and one transparent) showing the negative charge density due to the molecular orbitals of benzene. p Impress on the students how important hybrid orbitals are for organic and inorganic chemistry. please place this 24 points to the left PETRMC11_423-470-hr 12/20/05 3:42 PM Page 423 424 Chapter 11 Chemical Bonding II: Additional Aspects W HAT A B ONDING T HEORY S HOULD D O Imagine bringing together two atoms that are initially very far apart. Three types of interactions occur: (1) the electrons are attracted to the two nuclei; (2) the elec-trons repel each other; and (3) the two nuclei repel each other. We can plot po-tential energy—the net energy of interaction of the atoms—as a function of the distance between the atomic nuclei. In this plot, negative energies correspond to a net attractive force between the atoms; positive energies, to a net repulsion. Figure 11-1 shows the energy of interaction of two H atoms. This starts at zero when the atoms are very far apart. At very small internuclear distances, repul-sive forces exceed attractive forces and the potential energy is positive. At inter-mediate distances, attractive forces predominate and the potential energy is negative. In fact, at one particular internuclear distance (74 pm) the potential energy reaches its lowest value This is the condition in which the two H atoms combine into a molecule through a covalent bond. The nuclei continuously move back and forth; that is, the molecule vibrates, but the average internuclear distance remains constant. This internuclear distance cor-responds to the bond length . The potential energy corresponds to the negative ....
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- Spring '08