Bonding localized electron model

The Lewis model of the chemical bond assumes that each bonding electron pair is located between the two bonded atoms—it is a localized electron model. However, we know from the wave-particle duality of the electron (Sections 1.5-1.7) that the location of an electron in an atom cannot be described in terms of a precise position, but only in terms of the probability of finding it somewhere in a region of... 

Carbon occurs in the allotropes (different forms) diamond, graphite, and the fullerenes. The fullerenes are molecular solids (see Section 16.6), but diamond and graphite are typically network solids. In diamond, the hardest naturally occurring substance, each carbon atom is surrounded by a tetrahedral arrangement of other carbon atoms, as shown in Fig. 16.26(a). This structure is stabilized by covalent bonds, which, in terms of the localized electron model, are formed by the overlap of sp3 hybridized atomic orbitals on each carbon atom. 

Hydrogen bridges between the beryllium atoms produce a polymeric structure for BeH2, as shown in Fig. 18.6. The localized electron model describes this bonding by assuming that only one electron pair is available to bind each Be—H—Be cluster. This is called a three-center bond, since one electron pair is shared among three atoms. Three-center bonds have also been postulated to explain the bonding in other electron-deficient compounds (compounds where there are fewer electron pairs than bonds), such as the boron hydrides (see Section 18.5). 

Compare the description of the localized electron model (Lewis structure) with that of the molecular orbital model for the bonding in NO, NO+, and NO-. Account for any discrepancies between the two models. 

Bonding in Complex Ions The Localized Electron Model... 

By this point in your study of chemistry, you no doubt recognize that the localized electron model, although very simple, is a very useful model for describing the bonding in molecules. Recall that a central feature of the model is the formation of hybrid atomic orbitals that are used for sharing electron pairs to form cr bonds between atoms. This same model can be used to account for the bonding in complex ions, but there are two important points to keep in mind. 

Although the localized electron model can account in a general way for metal-ligand bonds, it is rarely used today because it cannot predict important properties of complex ions, such as magnetism and color. Thus we will not pursue the model any further. 

The simplest member of the saturated hydrocarbons, which are also called the alkanes, is methane (CH4). As discussed in Section 14.1, methane has a tetrahedral structure and can be described in terms of a carbon atom using an sp-J hybrid set of orbitals to bond to the four hydrogen atoms (see Fig. 22.1). The next alkane, the one containing two carbon atoms, is ethane (C2H6), as shown in Fig. 22.2. Each carbon in ethane is surrounded by four atoms and thus adopts a tetrahedral arrangement and sp3 hybridization, as predicted by the localized electron model. 

A special class of cyclic unsaturated hydrocarbons is known as the aromatic hydrocarbons. The simplest of these is benzene (C6H6), which has a planar ring structure, as shown in Fig. 22.11(a). In the localized electron model of the bonding in benzene, resonance structures of the type shown in Fig. 22.11(b) are used to account for the known equivalence of all the carbon-carbon bonds. But as we discussed in Section 14.5, the best description of the benzene molecule assumes that sp2 hybrid orbitals on each carbon are used to form the C—C and C—H a bonds, while the remaining 2p orbital on each carbon is used to form 77 molecular orbitals. The delocalization of these 1r electrons is usually indicated by a circle inside the ring [Fig. 22.11(c)]. 

The concept of resonance is necessary because the localized electron model postulates that electrons are localized between a given pair of atoms. However, nature does not really operate this way. Electrons are really delocalized—they can move around the entire molecule. The valence electrons in the N03 molecule distribute themselves to provide equivalent N—0 bonds. Resonance is necessary to compensate for the defective assumption of the localized electron model. However, this model is so useful that we retain the concept of localized electrons and add resonance to allow the model to treat species such as N03. ... 

As we saw in Chapter 8, the localized electron model views a molecule as a collection of atoms bound together by sharing electrons between their atomic orbitals. The arrangement of valence electrons is represented by the Lewis structure (or structures, where resonance occurs), and the molecular geometry can be predicted from the VSEPR model. In this section we will describe the atomic orbitals used to share electrons and hence to form the bonds. 

Describe the bonding in the ammonia molecule using the localized electron model. Solution... 

We have seen that the localized electron model is of great value in interpreting the structure and bonding of molecules. However, there are some problems with this model. For example, it incorrectly assumes that electrons are localized, and so the concept of resonance must be added. Also, the model does not deal effectively with molecules containing unpaired electrons. And finally, the model gives no direct information about bond energies. 

We divide by 2 because, from the localized electron model, we are used to thinking of bonds in terms of pairs of electrons. 

It would seem that the ideal bonding model would be one with the simplicity of the localized electron model but with the delocalization characteristic of the molecular orbital model. We can achieve this by combining the two models to describe molecules that require resonance. Note that for species such as O3 and the double bond changes... 

Therefore, we conclude that the a bonds in a molecule can be described as being localized with no apparent problems. It is the tt bonding that must be treated as being delocalized. Thus, for molecules that require resonance, we will use the localized electron model to describe the a bonding and the molecular orbital model to describe the tt bonding. This allows us to keep the bonding model as simple as possible and yet give a more physically accurate description of such molecules. 

What hybridization is required for central atoms exhibiting trigonal bipyramidal geometry Octahedral geometry Describe the bonding of PF5, SF4, SFg, and IF5 using the localized electron model. 

Use the localized electron model to describe the bonding in CCI4. 

All the Group 4A elements can form four covalent bonds to nonmetals—for example, CFI4, Sip4, GeBr4, SnC, and PbC. In each of these tetrahedral molecules the central atom is described as sp hybridized by the localized electron model. 

Describe the bonding in SO2 and SO3 using the localized electron model (hybrid orbital theory). How would the molecular orbital model describe the tt bonding in these two compounds ... 

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