4 - Chapter 4 An introduction to an orbital description of...

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1 Chapter 4 – An introduction to an orbital description of bonding An energy is associated with each occupied atomic orbital The most common mistake in drawing Lewis structures is to exceed an atom’s octet of electrons. An example of an invalid Lewis structure for CH 3 NO 2 is shown below. HC H N H O O The nitrogen atom shares a total of 10 electrons and exceeds its octet The problem with this configuration is that it depicts nitrogen as sharing a total of ten electrons. In order to understand why nitrogen cannot accommodate 10 electrons, we must examine the energies of the atomic orbitals. The 2s and 2p x , 2p y , and 2p z can each hold two electrons, a total of 8 valence electrons in all. Thus, the atoms in the second row of the periodic table do not have a sufficient number of low-energy atomic orbitals to accommodate more than 8 electrons. Nitrogen cannot involve the 3 s atomic orbital in covalent bonding because the energy of the 3 s orbital is too high. This high energy is evident from Figure 1, which shows how the energy levels change for the first 18 elements. 1 Nitrogen has atomic number of 7 and according to the figure, extrapolation of the energy of the 3s atomic orbital suggests that its value is off the chart. To involve this atomic orbital in bonding – that is, to occupy these orbitals with electrons – would simply cost too much energy. 1 For a complete description, see M. Kasha and R. Latter, Phys. Rev., 1955 , 99 , 510.
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2 Figure 1. This plot shows how the energy levels for the first 18 elements in the periodic table change as a function of atomic number. The energy levels shown here are the result of computation based on modern atomic theory. Notice that the energy scale is logarithmic. Atomic orbitals describe the distribution of electron density in space Not only do occupied atomic orbitals have energy associated with them, but they also describe the spatial distribution of electron density. More details will be provided next semester, but for now recognize that the electron density of the 2s atomic orbital has spherical shape, while the electron density of the 2p x , 2p y , and 2p z orbitals is cylindrical in shape, with each cylinder (or “dumbbell”) aligning itself along one of a set of three perpendicular axes. This distribution of electron density is not ideal for rationalizing the chemical bonding that we observe in organic molecules.
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3 The figure above shows the three common bonding geometries observed in organic
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4 - Chapter 4 An introduction to an orbital description of...

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