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chapter_2 - Chapter 2 Molecular Orbital Theory Everything...

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1 Chapter 2 Molecular Orbital Theory Everything said so far about bonding has been based on two very helpful simplifications. First, we have localized electron pairs, either as lone pairs on one particular atom or as bonded pairs between two particular atoms. Second, we have divided the valence electrons of a molecule into two groups, bonding and nonbonding, and assumed that they can be considered separately. Neither of these simplifications is correct. But without them, the description of the electronic structure of molecules with more than two atoms becomes extremely complex. Fortunately, such complex descriptions are usually not necessary in organic chemistry. Nevertheless, there are many aspects of the behavior of organic molecules that cannot be understood simply on the basis of Valence Bond theory alone. It is for these situations that we abandon, at least to some extent, the two simplifications of Valence Bond theory and use some of the ideas of Molecular Orbital theory. 2.1 Molecular Orbitals We have already used the idea of combining two or more atomic orbitals of a given atom to make hybrid orbitals that better describe how the atom bonds to other atoms. In a similar way, we can describe a bond by making combinations of the atomic orbitals of the two atoms that overlap to form the bond between these atoms.. The combining of atomic orbitals on different atoms must obey the same rules that we have already presented for making combinations of orbitals on the same atom to form hybrid orbitals.
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2 1. The number of combination orbitals equals the number of atomic orbitals that are combined to make them. The number of orbitals associated with any chemical species remains constant no matter how we choose to combine or blend them. 2. The total energy of the combination orbitals is equal to the total energy of the atomic orbitals that are combined to make them. The energy of a chemical species does not change just because we choose to describe it differently. When two atomic orbitals, one on each atom, overlap to form a bond between the two atoms, we can describe the process as the combining of the two orbitals. According to Rule 1, the combining of these two orbitals must produce two new orbitals. According to Rule 2, the total energy of these two new orbitals must be equal to the total energy of the atomic orbitals that combined to form them. The energy of one of the new orbitals is lower than the energy of the atomic orbitals. The energy of the other one is higher than the energy of the atomic orbitals by the same amount. These relationships are shown in Figure 2.1 Note: Figures are in a separate file called Chapter 2 Figures Figure 2.1 Energy relationships for combining two atomic orbitals on different atoms According to the Pauli Exclusion Principle, an orbital can hold a maximum of two electrons. The electron pair that is shared by the two bonded atoms is placed in the lower energy orbital. This orbital is called the bonding molecular orbital. Its lower energy helps explain why the formation of bonds is energetically favorable. The higher energy orbital
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