Experiment 1 Synthesis of Geometric Isomers

Techniques to review before coming to the laboratory: Appendix C: Vacuum filtration, Refluxing
Type of Lab Report: Modified Formal, review Appendix A: Modified Formal Lab Reports, hand in at the beginning of your next lab.

Introduction: Alkenes are hydrocarbons with a C=C double bond. The double bond is stronger than a single bond, yet paradoxically, the C=C double bond is much more reactive than a C-C single bond. Historically, hydrocarbons with double bonds were known as olefins. This name comes from the Latin words oleum, an oil, and ficare, to make. It arose because derivatives of such compounds often had an oily appearance. Alkenes can exist as two different geometric isomers. This isomerism results from the fact that there is restricted rotation at the double bond. Even though geometric isomers have the same molecular formula, they usually are different in many physical properties. The heats of formation for cis and trans butene are given below. Examination of these values shows that the trans alkene is more stable than the cis isomer. (Remember that a more negative heat of formation Hof corresponds to a more stable compound). Generally trans alkenes are more stable than the isomeric cis alkenes by about 1 kcal/mole.

The distance between the adjacent methyl groups in cis-2-butene is about 3Å. Since the sum of the van der Waals radii for two methyl groups (-CH­3) is 4 Å, the hydrogens in these two groups are sufficiently close that there is a net repulsion not present in the trans compound. When atoms or groups are too close to each other, repulsion occurs between the electron clouds of the atoms or groups, this effect is known as steric hindrance.
Some Notes on the Hybridization of Carbon: sp3 hybridization (one 2s and three 2p atomic orbitals)
e.g. Methane, CH4 The electronic configuration of carbon is 1s2 2s2 2px1 2py1 2pz0. This leaves two unpaired electrons for bonding which is contrary to experience, since a carbon singly bonded to only two other atoms is very unstable. Almost all carbon compounds exhibit quadrivalency. In valence bond theory the 2s orbital and the three 2p orbitals are hybridized to give four equivalent sp3 hybrid orbitals. The four electrons are then placed, one in each orbital (Hund’s rule) by promotion. The hybridized atomic orbitals overlap with the s orbitals of H atoms to form four strong sigma () bonds which are identical and symmetrically (tetrahedrally) disposed at an angle of 109.5o to each other.

The new hybrid atomic orbitals have identical energies, intermediate between s and p, as indicated in the following energy level diagram:

2p __ __

energy forms 4 σ bonds sp3 hybrid orbital 2s unhybridized C atom hybridized C atom

sp2 hybridization (one 2s and two 2p orbitals)
e.g. Ethylene, C2H4

In sp2 hybridization the 2s and two 2p orbitals are hybridized to form 3 sp2 hybrid atomic orbitals. In ethylene, two sp2 hybrids from each carbon, bond with an s orbital of hydrogen. The third sp2 hybrid on each carbon is used to form the  C-C single bond. All of these atoms lie in the same plane and are inclined at 120o to each other (trigonal planar). This leaves a p orbital “left over” on each carbon. The p orbitals lie perpendicular to the plane of the six atoms. The two p orbitals are also parallel to each other and have regions of overlap below and above the molecular plane. This results in a bond spread over both carbon atoms called a pi () bond. The lateral overlap of a  bond is not as effective as the lateral overlap of a  bond and thus is weaker.

2p forms π bond unhybridized p orbital energy forms 3 σ bonds sp2 hybrid orbitals 2s ___ unhybridized C atom hybridized C atom

sp1 hybridization (one 2s and one 2p orbital)
e.g.