In the Classroom edited by
Overhead Projector Demonstrations
Doris K. Kolb Bradley University Peoria, IL 61625
Visualizing Bent Bonds in Cyclopropane Thomas M. Bertolini Department of Chemistry, University of California–Irvine, Irvine, CA 92697-2025;
[email protected] In practically all sigma bonds, the region of greatest electron density is symmetrically distributed around the axis between the two nuclei. A frequently encountered exception is the carbon–carbon bonding of cyclopropane, where atomic orbitals overlap outside the internuclear axis (Figure 1) and cylindrical symmetry is not present (1). Regrettably, some introductory organic chemistry textbooks do not address bent or “banana” bonds. It is common for students to believe that the interorbital angle of a cyclopropyl carbon is 60⬚, when this is in fact the internuclear angle of cyclopropane. The interorbital angle of a cycloalkyl carbon is the angle between its two carbon-bonding sp3 orbitals. For small-ring compounds, however, the internuclear angle is smaller than the interorbital angle. To illustrate the unique bonding of cyclopropane, I employ a two-minute overhead demonstration using a molecular model kit. Molecular Visions models work particularly well, although many other model kits are suitable. To begin, I place a model of propane on the overhead, manipulating the carbon backbone to show its rotational flexibility. I then assemble a model of cyclopropane, a process that requires a small amount of practice. Most model kits, much like an sp3 hybridized carbon atom, resist forming 60⬚ bond angles. As a result, the bonds of the cyclopropane model are bent and the angles are about 102⬚ (Figure 2). This model is an excellent visual aid because it closely approximates the orientation of cyclopropane bonding orbitals (2). It is easy for students to see that the bonding in cyclopropane (relative to propane) is higher in energy. Clearly, the cyclopropane model is more difficult to construct and has warped, nonlinear bonds. Of course, it is important to emphasize to students that covalent bonds are the result of orbital overlap and greater overlap produces stronger bonds. The unique bent-bonding of cyclopropane proves this rule: since the carbon-carbon orbitals overlap only slightly, the bonds are weak. To demonstrate this point, I open the cyclopropane model (making a loud popping sound) and the carbon–carbon “bonds” of the model immediately become linear again. This physical demonstration also sets the stage for the ring-opening reactions of cyclopropane (via hydrogenation), epoxides, and halonium ions encountered in sophomore-level organic chemistry. Lastly, to prevent confusion I never refer to bent bonds as “puckered,” because “puckering” of cycloalkanes is more commonly used to describe a nonplanar conformation.
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Journal of Chemical Education
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Figure 1. Interorbital versus internuclear angles of cyclopropane. The dotted arrows show the directionality of the carbon sp3 orbitals (and hence the 102° interorbital angles) of cyclopropane; the solid lines show its 60° internuclear angles (1).
Figure 2. Model of cyclopropane. I put circular stickers at the vertices of the model so that the carbon nuclei are more obvious. The bond angles of the carbon atoms are similar to the actual 102° interorbital angles of cyclopropane (2).
Literature Cited 1. Coulson, C. A.; Moffitt, W. E. J. Chem. Phys. 1947, 15, 151. 2. March, J. Advanced Organic Chemistry, 3rd ed.; John Wiley & Sons: New York, 1985; p 131.
Vol. 81 No. 6 June 2004
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