Hybrid orbitals the valence bond model

As useful and appealing as the concept of the shared-electron bond is, it raises a somewhat troubling question that we must sooner or later face what is the nature of the orbitals in which the shared electrons are contained Up until now, we have been tacitly assuming that each valence electron occupies the same kind of atomic orbital as it did in the isolated atom. As we shall see below, his assumption very quickly leads us into difficulties. 

Consider how we might explain the bonding in a compound of divalent beryllium, such as beryllium hydride, BeH2. The beryllium atom, with only four electrons, has a configuration of ls22s. 

This means that there are only two electrons in the outer shell of beryllium these electrons are paired up together in the 2s orbital. 

Now according to covalent bond theory, each beryllium-hydrogen bond must consist of two electrons, shared between the two different atoms. One of these electrons is contributed by the hydrogen, while the other comes from the beryllium atom. 

The question then arises how can the two outer-shell electrons of beryllium interact with the hydrogen electrons Two electrons in the same orbital have opposite spins, and constitute a stable pair that has no tendency to interact with unpaired electrons on other atoms. Electron sharing, as we have seen so far, can only take place when UNpaired electrons on adjacent atoms interact. 

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Valence-bond theory and hybrid orbitals allow us to move in a straightforward way from Lewis structures to rationalizing the observed geometries of molecules in terms of atomic orbitals. The valence-bond model, however, does not explain all aspects of bonding. It is not successful, for example, in describing the excited states of molecules, which we must understand to explain how molecules absorb light, giving them color. 

Figure 4.7 The linear geometry of BeCl2 can be explained by assuming that Be is sp hybridized. In the valence bond model, the two sp hybrid orbitals overlap with the two chlorine 3p orbitals to form two covalent bonds.

We 11 expand our picture of bonding by introducing two approaches that grew out of the idea that electrons can be described as waves—the valence bond and molecular orbital models In particular one aspect of the valence bond model called orbital hybridization, will be emphasized... 

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the valence bond model, which uses hybrid orbitals to account for the observed shapes of organic molecules. Before going on to a systematic study of organic chemistry, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the electrons in a bond are not shared equally between atoms. 

The valence bond model constructs hybrid orbitals which contain various fractions of the character of the pure component orbitals. These hybrid orbitals are constructed such that they possess the correct spatial characteristics for the formation of bonds. The bonding is treated in terms of localised two-electron two-centre interactions between atoms. As applied to first-row transition metals, the valence bond approach considers that the 45, 4p and 3d orbitals are all available for bonding. To obtain an octahedral complex, two 3d, the 45 and the three 4p metal orbitals are mixed to give six spatially-equivalent directed cfisp3 hybrid orbitals, which are oriented with electron density along the principal Cartesian axes (Fig. 1-9). 

Hybridization in an octahedral complex in the valence bond model of coordination ompound the formation of bonds to an octahedral transition metal ion would involve six d sp h /hrid orbitals. In the approach used here the diff( ring syininetry properties of the s, thii p, p and p.. and the d, and d orbitals used in bonding are explicitly taken into account. However, the 3 all of the p and two ol the d metal valence shell orbitals are otill used in o-bond tormation. 

We start with the organic series CH3, CH2, and CH. Just as there were one, two, or three nonbonding hybrid orbitals, respectively in the valence-bond model, the electronic structure in terms of... 

The valence bond model and the accompanying idea of orbital hybridization have applications beyond those of bonding in hydrocarbons. A very simple extension to inorganic chemistry describes the bonding to the nitrogen of ammonia and the oxygen of water. 

The valence bond model for the nitrogen molecule is shown in Figure 10.11. The best Lewis structure for N2 has a N=N triple bond with a lone pair on each N atom, or a linear electron geometry. Therefore, the expected hybridization from Table 10.2 is sp. One of the sp hybrids will be used to form a sigma bond between the two N atoms, while the other sp hybrid will contain a lone pair. In addition, each N atom will have two unhybridized p-orbitals. These two orbitals can overlap in a sideways manner to form two pi bonds between the N atoms. 

For molecules where the valence of the central atom violates the octet rule (the so-called hypervalent molecules), the valence bond model is forced to resort to the inclusion of d-orbitals in the hybridization. Thus, for instance, the hybridization of phosphorous in PF5 is formally dsp, as shown in Figure 10.12, even though it is doubtful that the high-lying 3d orbitals actually participate in the bonding to any significant extent. 

Q18 State the geometry of a molecule in which the bonding may be described using the valence-bond model as being made up of sp hybrid orbitals on the central atom. 

The orbital hybridization model (which is a type of valence bond model)... 

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