Orbitals axial overlap

The combination of two carbon atoms, for example in ethane, results from the axial overlap of two sp3 atomic orbitals, one from each... 

Two types of interactions are often distinguished cr interactions, which concern an axial orbital overlap, and tt interactions, where the orbital overlap occurs laterally, or sideways . These two types of overlap are illustrated in 1-24 and 1-25, respectively, for two p orbitals whose axes of revolution are either co-linear (axial overlap) or parallel (sideways overlap). Notice that another way to characterize rr interactions is to observe that the orbitals involved share a common nodal plane (P, 1-25). 

In general, a interactions are stronger than tt interactions, since axial overlap is more efficient than sideways overlap. The energy separation between the resulting orbitals is therefore larger fora (bonding) and a (antibonding) MO than for the tt and jt MO. 

We shall consider the interactions between the d orbitals on two metallic centres, initially ignoring any influence of the ligands on these orbitals. Each interaction leads to the formation of a bonding and an antibonding MO, but only the first of these, written is shown in 4-45. The interaction between the orbitals (4-45a) is of ct type the axial overlap between the two orbitals has cylindrical symmetry around the intemudear axis (z). The MO and yz (4-45b) have the same characteristics as the 7T orbitals of acetylene (4-44) they are bonding, and they possess a nodal plane that passes through the nuclei. 

Sigma bonds are formed in all the examples described by the axial (head-on) overlap of orbitals (Figure 14.18). Such O bonds are relatively unreactive, but are important in determining the shape or skeleton of a molecule or ion (Chapter 4). However, there is another way in which p orbitals can overlap if they are brought together in the correct orientation. 

This bond is formed by the axial overlap of atomic orbitals. This bond is formed by the sideways overlap of atomic orbitals. 

Figure 14.66 Energy diagram for the formation of a and a orbitals from the axial overlap of a a Is and 2p orbital and b two 2p orbitals... 

With the radical 29, even though loss of an equatorial hydrogen should be sterically less hindered and is favored thermodynamically (by relief of 1,3 interactions of the axial methyl), there is an 8-fold preference for loss of the axial hydrogen (at 100 ( i. The selectivity observed in the disproportionation of this and other substituted cyclohexyl radicals led Beckwith18 to propose that disproportionation is subject to stereoelectronic control which results in preferential breaking of the C-H bond which has best overlap with the orbital bearing the unpaired spin. 

I have developed a simple theory of these potential barriers, described in the following paragraphs. According to this theory, the potential barriers are not a property of the axial bond itself, but result from the exchange interactions of electrons involved in the other bonds (adjacent bonds) formed by each of the two atoms, as determined by the overlap between the parts of the adjacent bond orbitals that extend from each of the two atoms toward the other. 

Leahy demonstrated that unsaturation at the 5-position of a 4-cyano-l,3-dioxane can lead to a reversal in selectivity [12] (Eq. 6). Alkylation of cyanohydrin acetonide 19 with benzyl bromide generated a 9 1 mixture of 20 and 21, with the flufz-isomer 20 predominating, in 57% overall yield. An alkylithium intermediate in which overlap with the methylidene tt orbital favors the axial configuration could account for this anomalous selectivity. 

As a consequence of the molecular orbital interactions, ferrocene adopts an axially symmetrical sandwich structure with two parallel Cp ligands with a distance of 3.32 A (eclipsed conformation) and ten identical Fe-C distances of 2.06 A as well as ten identical C-C distances of 1.43 A [12]. Deviation of the parallel Cp arrangement results in a loss of binding energy owing to a less efficient orbital overlap [8]. All ten C-H bonds are slightly tilted toward the Fe center, as judged from neutron-diffraction studies [13]. 

With a steric number of 5, chlorine has trigonal bipyramidal electron group geomehy. This means the inner atom requires five directional orbitals, which are provided hymsp d hybrid set. Fluorine uses its valence 2 p orbitals to form bonds by overlapping with the hybrid orbitals on the chlorine atom. Remember that the trigonal bipyramid has nonequivalent axial and equatorial sites. As we describe in Chapter 9, lone pairs always occupy equatorial positions. See the orbital overlap view on the next page. 

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