Electronic structure Lewis approach

The Lewis approach also fails for the compound diborane, B2H6, a colorless gas that bursts into flame on contact with air. The problem is that diborane has only 12 valence electrons (3 from each B atom, 1 from each H atom), but for a Lewis structure, it needs at least 7 bonds, and therefore 14 electrons, to bind the 8 atoms together Diborane is an example of an electron-deficient compound, a compound with too few valence electrons to be assigned a valid Lewis structure. Valence-bond theory can account for the structures of electron-deficient compounds in terms of resonance, but the explanation is not straightforward. 

Our LMO approach to the electronic structure and reactivity of chemical species allowed us not only to generalize the (Lewis and Linnett) theory of valence but also to determine the mechanism of a wide variety of organic reactions. 

Thus far we have discussed the chemical bonding in polyatomic molecules in terms of the VB model, or more crudely in terms of Lewis structures. These two treatments are related in that they focus on chemical bonding in terms of the sharing of electron pairs by adjacent atoms. In many cases, however, a more sophisticated approach based on molecular orbital concepts is needed to accurately picture the electronic structure of polyatomic molecules—even on a quahtative level. 

The controversy as to how best to write Lewis structures will no doubt continue in the chemical literature, but you should not be too dismayed by this situation. Our approach to depicting the electronic structure of a molecule is based on the simplest Lewis structure and its concomitant use in determining the shape of a molecule through VSEPR theory. In order to probe more deeply into the nature of a chemical bond—for example, to understand experiment results, such as bond enthalpy values— we must analyze a computed electron density map for that molecule rather than rely just on the Lewis structure. 

There are several ways to choose the more plausible of two structures differing in their arrangement of atoms. As pointed out in Example 7.1, the fact that carbon almost always forms four bonds leads to the correct structure for ethane. Another approach involves a concept called formal charge, which can be applied to any atom within a Lewis structure. The formal charge is the difference between the number of valence electrons in the free atom and the number assigned to that atom in the Lewis structure. The assigned electrons include—... 

According to Lewis s approach and valence-bond theory, we should describe the bonding in 02 as having all the electrons paired. However, oxygen is a paramagnetic gas (Fig. 3.24 and Box 3.2), and paramagnetism is a property of unpaired electrons. The paramagnetism of 02 therefore contradicts both the Lewis structure and the valence-bond description of the molecule. 

Our approach to these molecules illustrates the general strategy for determining the electron group geometry and the molecular shape of each inner atom in a molecule. The process has four steps, beginning with the Lewis structure and ending with the molecular shape. 

Our treatment of O2 shows that the extra complexity of the molecular orbital approach explains features that a simpler description of bonding cannot explain. The Lewis structure of O2 does not reveal its two unpaired electrons, but an MO approach does. The simple (t-tt description of the double bond in O2 does not predict that the bond in 2 is stronger than that in O2, but an MO approach does. As we show in the following sections, the molecular orbital model has even greater advantages in explaining bonding when Lewis structures show the presence of resonance. 

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