Ethylene orbital overlap

Figure 1.14 The structure of ethylene. Orbital overlap of two sp hybridized carbons forms a carbon-carbon double bond. One part of the double bond results from a (head-on) overlap of sp2 orbitals (green), and the other part results from (sideways) overlap of unhybridized p orbitals (red/blue). The ir bond has regions of electron density on either side of a line drawn between nuclei.

Let us now examine the Diels-Alder cycloaddition from a molecular orbital perspective Chemical experience such as the observation that the substituents that increase the reac tivity of a dienophile tend to be those that attract electrons suggests that electrons flow from the diene to the dienophile during the reaction Thus the orbitals to be considered are the HOMO of the diene and the LUMO of the dienophile As shown m Figure 10 11 for the case of ethylene and 1 3 butadiene the symmetry properties of the HOMO of the diene and the LUMO of the dienophile permit bond formation between the ends of the diene system and the two carbons of the dienophile double bond because the necessary orbitals overlap m phase with each other Cycloaddition of a diene and an alkene is said to be a symmetry allowed reaction... 

What do we know about ethylene We know from Section 1.8 that a carbon-carbon double bond results from orbital overlap of two s/Ahybridized carbon atoms. The a part of the double bond results from sp2-sp2 overlap, and the 77 part results from p-p overlap. 

Many of the Lewis structures in Chapter 9 and elsewhere in this book represent molecules that contain double bonds and triple bonds. From simple molecules such as ethylene and acetylene to complex biochemical compounds such as chlorophyll and plastoquinone, multiple bonds are abundant in chemistry. Double bonds and triple bonds can be described by extending the orbital overlap model of bonding. We begin with ethylene, a simple hydrocarbon with the formula C2 H4. 

The n molecular orbitals described so far involve two atoms, so the orbital pictures look the same for the localized bonding model applied to ethylene and the MO approach applied to molecular oxygen. In the organic molecules described in the introduction to this chapter, however, orbitals spread over three or more atoms. Such delocalized n orbitals can form when more than two p orbitals overlap in the appropriate geometry. In this section, we develop a molecular orbital description for three-atom n systems. In the following sections, we apply the results to larger molecules. 

When two p orbitals overlap in a side-by-side configuration, they form a pi bond, shown in Figure 7.7. This bond is named after the Greek letter 7t. The electron clouds in pi bonds overlap less than those in sigma bonds, and they are correspondingly weaker. Pi bonds are often found in molecules with double or triple bonds. One example is ethene, commonly known as ethylene, a simple double-bonded molecule (Figure 7.8). The two vertical p orbitals form a pi bond. The two horizontal orbitals form a sigma bond. 

Now let us return to our discussion of the conical intersection structure for the [2+2] photochemical cycloaddition of two ethylenes and photochemical di-Jt-methane rearrangement. They are both similar to the 4 orbital 4 electron model just discussed, except that we have p and p overlaps rather than Is orbital overlaps. In Figure 9.5 it is clear that the conical intersection geometry is associated with T = 0 in Eq. 9.2b. Thus (inspecting Figure 9.5) we can deduce that... 

Fig. 7.9. The possible mode of orbital overlapping of tungsten d orbitals with two ethylenic bonds...

Figure 10-4 shows the hybridization that occurs in ethylene, H2C=CH2. Each carbon has sp2 hybridization. On each carbon, two of the hybrid orbitals overlap with an s-orbital on a hydrogen atom to form a carbon-to-hydrogen covalent bond. The third sp2 hybrid orbital overlaps with the sp2 hybrid on the other carbon to form a carbon-to-carbon covalent bond. Note that each carbon has a remaining p-orbital that has not undergone hybridization. These are also overlapping above and below a line joining the carbons. 

A cr orbital can be formed either from two s atomic orbitals, or from one s and one p atomic orbital, or from two p atomic orbitals having a collinear axis of symmetry. The bond formed in this way is called a a bond. A n orbital is formed from two p atomic orbitals overlapping laterally. The resulting bond is called a n bond. For example in ethylene (CH2=CH2), the two carbon atoms are linked by one a and one n bond. Absorption of a photon of appropriate energy can promote one of the n electrons to an antibonding orbital denoted by n. The transition is then called Ti —> 7i. The promotion of a a electron requires a much higher energy (absorption in the far UV) and will not be considered here. 

It is also worthwhile to compare the ferrocenyl ethylene (vinylferrocene) anion-and cation-radicals. For the cyano vinylferrocene anion-radical, the strong delocalization of an unpaired electron was observed (see Section 1.2.2). This is accompanied with effective cis trans conversion (the barrier of rotation around the -C=C- bond is lowered). As for the cation-radicals of the vinylferrocene series, a single electron remains in the highest MO formerly occupied by two electrons. According to photoelectron spectroscopy and quantum mechanical calculations, the HOMO is mostly or even exclusively the orbital of iron (Todres et al. 1992). This orbital is formed without the participation of the ethylenic fragment. The situation is quite different from arylethylene radical cations in which all n orbitals overlap. After one-electron oxidation of ferrocenyl ethylene, an unpaired electron and a positive charge are centered on iron. The —C=C— bond does not share the n-electron cloud with the Fe center. As a result, no cis trans conversion occurs (Todres 2001). 

Thus, a (—ac) is the core energy of an electron localized to the 2p atomic orbital of a carbon atom, and / (—/ cc) is the energy associated with the interaction of two carbon 2p orbitals overlapping in (parallel) fashion at the C—C separation of benzene (or ethylene). 

This [2 + 2] cycloaddition requires the HOMO of one of the ethylenes to overlap with the LUMO of the other. Figure 15-20 shows that an antibonding interaction results from this overlap, raising the activation energy. For a cyclobutane molecule to result, one of the MOs would have to change its symmetry. Orbital symmetry would not be conserved, so the reaction is symmetry-forbidden. Such a symmetry-forbidden reaction can occasionally be made to occur, but it cannot occur in the concerted pericyclic manner shown in the figure. 

The bonding in vinylcyclopropane (3) is such that an s-trans-gauche conformational equilibrium exists to allow for maximum orbital overlap of the asymmetric component of cyclopropane orbitals with the IT- or 7T -orbitals of the ethylene unit, as shown in (3a). From thermochemical stuches it appears that conjugation of an alkene with cyclopropane stabilizes the system by 1.2 kcal mol". The conformational equilibrium for vinylcyclopropane was shown to consist of an s-trans minimum (3b) and two gauche conformers that are equ in energy and destabilized by 1.43 kcal mol with respect to the s-trans conformation. The barrier to interconversion has been determined to be 3.92 kcal mol . ... 

In this picture of ethylene, the two orbitals that make up the double bond are not equivalent. The a orbital is ellipsoidal and symmetrical about the C—C axis. The 71 orbital is in the shape of two ellipsoids, one above the plane and one below. The plane itself represents a node for the n orbital. In order for the p orbitals to maintain maximum overlap, they must be parallel. This means that free rotation is not possible about the double bond, since the two p orbitals would have to reduce their overlap to allow one H C H plane to rotate with respect to the other. The six atoms of a double bond are therefore in a plane with angles that should be 120°. Double bonds are shorter than the corresponding single bonds because maximum stability is obtained when the p orbitals overlap as much as possible. Double bonds between carbon and oxygen or nitrogen are similarly represented they consist of one a and one n orbital. 

FIGURE 7.12 Bond formation in ethylene, (a) Overlap of sp hybrid orbitals forms a cr bond between the carbon atoms, (b) Overlap of parallel 2p orbitals forms a tt bond. 

Section 2.2.5 and Figure 2.33), with the ethylene jt and n MOs (Houk, 1982). In the ortho approach, the benzene MOs and . can interact with the ethylene MOs jt and m, while in the meta approach, and 0,. can interact with Jt and Jt. Therefore, the configuration is predicted to be energetically favored along the ortho cycloaddition path, while the configuration d>s a is stabilized for the ortho and particularly the meta cycloaddition path. The magnitude of the stabilization depends on orbital overlap and relative orbital energies. The former factor favors the ortho approach and the latter the meta approach of the reactants. 

Figure 4.8 illustrates how these differences influence available orbitals. In the [100] plane, e, orbitals emerge perpendicular and parallel to the surface, orbitals at 4S Empty or partially htled Cg orbitals overlap the l5 orbital of hydrogen in two locations, with one g orbital at the on top position or with hve at the position one half of an atomic layer into the surface. Similar situations prevail for other surface planes. With this model, it is now possible to visualize interactions of molecules such as hydrogen and ethylene with the nickel surface. Figure 4.9 shows the possibilities. Currently these models are only qualitative and much more... 

---