The Use of p Orbitals in Pi Bonding

In view of the extensive chemistry of alkenes it was only natural for organic and inorganic chemists to search for analogous Si=Si doubly bonded structures. For a long lime such attempts proved to be fruitless. The first stable C=Si and Si=Si9 compounds were synthesized about a decade ago. One synthesis involves the rearrangement of cyclottjsilane  

It is possible to add reagants across the Si=Si double bond in some ways analogous to the C=C bond in atkenes  

7 The diagonal relationship, like any other rule-or-lhumb, can be seen from several viewpoints in terms of unifying known facts (when it works), as a predictor of unknown properties (hoping it works), or in terms of the significance of its exceptions (when it does not work). See, for example, Fcinstein. H. I. J. Chan. Educ. 1984, 61. 128. Human, T. P. J. Chan. Editc. 1987. 64. 886-687. 

Compounds that are formally analogous to carbon compounds are found to have quite different structures. Thus carbon dioxide is a gaseous monomer, but silicon dioxide is an infinite single-bonded polymer. In a similar manner. gem-diols arc unstable relative to ketones  

10 The lerm silicone was coined by analogy 10 kcionc under lhe mistaken belief that monomeric R Si 0 compounds could be isolated. See Chapter 16. 

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First- and Second-Row Anomalies 858 The Use of p Orbitals in Pi Bonding 861 The Use (or Not) of d Orbitals by Nonmetals 86ft Reactivity and d Orbital Participation 875... 

This MO picture helps explain why all the atoms of ethene lie in the same plane. On each carbon the p orbital that is used to form the pi bond is perpendicular to the plane defined by the C and the two attached H s. Unless the plane of one C and its two attached H s is the same as the plane of the other C and its two attached H s, the p orbitals will not be parallel and overlap will be decreased. As we have seen, as overlap decreases, AE decreases and the energy of the electrons increases—the molecule is less stable. 

Figure 3.15c is an attempt to show how the AOs might overlap to form localized MOs in the formate anion. In this localized MO picture, a p orbital on the carbon overlaps with a p orbital on the upper oxygen to form a pi bond, corresponding to the Lewis structure of Figure 3.15a. In this structure, the lower oxygen has three unshared pairs of electrons. Whenever an atom with an unshared pair of electrons is adjacent to a pi bond, as occurs here, that atom usually assumes a hybridization that places an unshared pair in a p orbital because the overlap of this p orbital with the p orbital of the pi bond on the adjacent atom is stabilizing. It is this overlap that allows resonance to occur. In this case the p orbital with the unshared pair on the lower oxygen overlaps equally well with the p orbital on the carbon so that the pi bond could also be shown using these two orbitals with an unshared pair of electrons in the p orbital on the upper oxygen. This corresponds to the second Lewis structure (b).

In allene, the central carbon atom is sp hybridized and linear (Section 2-4), and the two outer carbon atoms are sp2 hybridized and trigonal. We might imagine that the whole molecule lies in a plane, but this is not correct. The central sp hybrid carbon atom must use different p orbitals to form the pi bonds with the two outer carbon atoms. The two unhybridized p orbitals on the sp hybrid carbon atom are perpendicular, so the two pi bonds must also be perpendicular. Figure 5-19 shows the bonding and three-dimensional structure of allene. Allene itself is achiral. If you make a model of its mirror image, you will find it identical with the original molecule. If we add some substituents to allene, however, the molecule may be chiral. 

The carbons in ethene form two bonds with each other. This is called a double bond. The two carbon-carbon bonds in the double bond are not identical. One of the bonds results from the overlap of an sp orbital of one carbon with an sp orbital of the other carbon this is a sigma (cr) bond because it is formed by end-on overlap (Figure 1.16a). Each carbon uses its other two sp orbitals to overlap the s orbital of a hydrogen to form the C—H bonds. The second carbon-carbon bond results from side-to-side overlap of the two unhybridized p orbitals. Side-to-side overlap of p orbitals forms a pi (rr) bond (Figure 1.16b). Thus, one of the bonds in a double bond is a cr bond and the other is a rr bond. All the C—H bonds are a bonds. 

In a similar manner, VBT can be used to explain why ethylene is a planar molecule. As in O2, the two C atoms are sp hybridized and lie in the xy-plane—only in H2C=CH2, each sp hybrid contains a single electron. One of these hybrids is used to form the C-C sigma bond, while the other two hybrids are used to form the C-H sigma bonds. The C-C pi bond forms as a result of the sideways overlap between the two p orbitals. In order for the two p orbitals to overlap with each other, there can be no free rotation around the C=C double bond. Therefore, both CH2 halves of the ethylene molecule must lie in the same plane, as shown in Figure 10.10. 

How can we explain the double bond between the carbon atoms In the cr bond the electron pair occupies the space between the carbon atoms. The second bond must therefore result from sharing an electron pair in the space above and below the cr bond. This type of bond can be formed using the 2p orbital perpendicular to the sp2 hybrid orbitals on each carbon atom (refer to Fig. 14.10). These parallel p orbitals can share an electron pair, which occupies the space above and below a line joining the atoms, to form a pi (zr) bond, as shown in Fig. 14.12. 

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