Cylindrical triple bonds

Double and triple bonds, particularly those in carbon molecules, are often described in MO terms. Thus a double bond is described as consisting of a a bond formed by the overlap of an sp hybrid orbital on each carbon atom and a tt bond formed by the sideways overlap of either the 2p or 2p orbitals (Figure 3.18). A cr orbital has cylindrical symmetry like an atomic s orbital whereas a ir orbital, like an atomic p orbital, has a planar node passing through the nucleus of each of the bonded atoms. A triple bond is similarly described as consisting of a cr orbital and two ir orbitals formed from both the 2p and 2pv orbitals on... 

The 7i electron density in the triple bond of ethyne is cylindrically symmetric. 

A triple bond is composed of one sigma bond and two pi bonds. Since the two pi bonds are perpendicular to each other, they are treated as two separate, noninteracting pi bonds. Perpendicular orbitals do not interact. The triple bond is linear, and atoms bonded to it lie in a straight line (Fig. 1.13). Organic chemists tend to draw skinny p orbitals (see Fig. 1.8) the triple bond is overall cylindrically symmetrical. 

Overlap not a problem because of the cylindrical shape of the triple bond. 

The simplest example of the triple bond is in the nitrogen molecule illustrated in Fig. 23.15. Two 7C bonds are formed in addition to the o bond. The unshared pairs are 180° from the sigma bond on each end of the molecule. The formation of two n bonds results in the charge cloud having the shape of a cylindrical sheath around the axis of the molecule. (Compare to Fig. 22.5 for m = 1.) Species that are isoelectronic with nitrogen and therefore have the same electronic structure are C=0, "C=N, and C=C. ... 

The nature of the triple bond, whether —C=C—, —C=N or ==N=N and their higher row analogues, has been reviewed several times in this series, in theoretical chapters. The nature of a chemical entity is defined by its geometric and electronic structural properties. The triple bond is characterized by short bond distances, local cylindrical symmetry, relatively high electron density prependicular to the bond axis and relatively large polarizability parallel to the bond axis. Its description in terms of a and n bonding is classic, and qualitatively explains the characteristic properties of the triple bond. 

In organic chemistry, we know that two carbon atoms can be linked by a single bond (a), a double (a - - tt), or a triple bond (a + In). The increase in bond order is accompanied by a decrease in interatomic distance, from 1.54 A (C—C), through 1.34 (C=C) to 1.20 A (C=C). From the orbital point of view, these different bond orders correspond to the occupation of one, two, or three bonding MO between the two carbon atoms (4-44), it being clear that the corresponding antibonding MO are empty. Tbe orbital that characterizes the a bond has cylindrical symmetry about the internuclear axis in each case, whereas the orbitals that characterize the tt bond(s) have a nodal plane (4-44). 

Because the 2p orbitals have a node at the atomic nucleus, the electron distribution of the VB wavefunction formed by their side-on overlap does not have cylindrical symmetry, but instead changes sign when rotated by 180° about the bond axis. Bonds that have this property are called pi (tt) bonds. Thus, the double bond in the oxygen molecule consists of a sigma bond and a pi bond, which are not equivalent. The existence of two different types of bonds in double bond formation is something that is not predicted by simple Lewis theory. A triple bond such as that in N2 is described within the VB approach as a sigma bond formed from the head-on overlap of 2p orbitals and two pi bonds formed by the overlap of the 2py and 2p orbitals on one N atom with their counterparts on the other atom (Figure 3.7). 

In acetylene, each C atom is sp hybridized, and its four valence electrons half-fiU the two sp hybrids and the two unhybridized 2p orbitals (Figure 11.11). Each C forms a C—H a bond with one sp orbital and a C—C a bond with the other. Side-to-side overlap of one pair of 2p orbitals gives one tt bond, with electron density above and below the a bond. Side-to-side overlap of the other pair of 2p orbitals gives another tt bond, 90 away from the first, with electron density in front and back of the (j bond. The result is a cylindrically symmetrical H—C=C—H molecule. Note the greater electron density between the C atoms created by the six bonding electrons. Any triple bond consists of one a and two it bonds. 

In a triple bond with an electron configuration the set of molecular orbitals has a cylindrical symmetry again so the system can rotate internally in a free way. 

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