Molecular orbital theory (MO theory) can be a very challenging topic. Students come into the classroom already knowing about the octet rule, Hund’s rule, and the Pauli exclusion principle. However, the lecture component is critical for bridging the gap between their knowledge of general chemistry and organic chemistry.
As such, I start my lecture on MO theory with a concept that ties in the law of conservation of mass. MO theory requires that for every atomic orbital (AO), there must also be a molecular orbital (MO) made. It is imperative that students know the difference between AOs and MOs. For instance, s and p orbitals are AOs, whereas sp, sp2, and sp3 hybridized orbitals are MOs. I illustrate this concept using soda bottles, partly because I personally love Pepsi, and partly because I like to add fun to the classroom, which lightens the mood and makes learning a tough topic a bit easier and quicker for my students.
In this example, I tell my students that my husband goes to the grocery store and purchases his favorite soda, Sprite, while also knowing that my favorite soda is Pepsi. When he comes home, I find that he has 1 Sprite bottle and 3 Pepsi bottles. For the purposes of this analogy, I label the Sprite as “s” and the Pepsi as “p,” which helps my students visualize 4 AOs, particularly 1 s and 3 p orbitals. When these bottles are mixed together, they result in 4 bottles: 1 part s and 3 parts p. Translated back to MO theory, this equals 4 sp3 orbitals, or 4 MOs! I then rationalize to my students how this applies to the remaining MOs—sp2 and sp—and explain how the remaining p orbitals lead to the pi bonding in the alkene and alkyne. For instance, for an sp2 MO, we need 1 s and 2 p orbitals. This means that we put in 3 AOs and get 3 sp2 hybridized MOs, with one p orbital leftover that can be used for pi bonding.
After this analogy, my lecture transitions to a discussion of energy diagrams. My main goal here is to help my students visualize and review all the components of energy diagrams. For instance, we look at the following items listed below:
- How do you draw an energy diagram?
- How many MOs do I add in the diagram based on the described structures?
- How do you fill electrons in the diagram? Why do we fill them from the bottom up?
- Where do bonding vs. antibonding orbitals sit?
- Where are non-bonding orbitals? How are they designated? Are non-bonding orbitals found in all energy diagrams, like the bonding and antibonding orbitals?
- How many nodes are there? Where do you place them?
In particular, I choose to focus on allylic cation, allylic carbanion, 1,3-butadiene, 1,3-pentadienyl cation, and finally, (1E,3E)-1,3-pentadiene-1-amine, because all these examples show bonding vs. antibonding MOs. Once students begin to understand the components of MO theory and these two fundamental types of MOs (bonding and antibonding), I then bring in the non-bonding MOs.
The topic of MO theory is important because it shows us how we should view molecules in nature and in real time. In the case of 1,3-butadiene, we see it as a 4-carbon molecule with two double bonds, and between the double bonds, there is a single bond. I had mentioned earlier to my students—about 1,3-butadiene—that they can note that the most stable MO has two sets of overlapping p orbitals (representing the double bond) with one node, a place where there is no overlap or bond, between them (representing the single bond). I emphasize that the highest energy bonding orbital (also known as the highest occupied molecular orbital) is what we see in nature because it is the most stable.
A latter example is the 1,3-pentadienyl cation, which has 5 carbons, meaning that it has 5 MOs in the energy diagram. Typically, the cation is drawn on the exterior carbon—which we can call carbon #5—but this carbocation is not stable because it is a primary carbocation. So I ask my students whether the cation can be moved to make the carbocation more stable. They mention that the cation can move via resonance to the center—carbon #3—which generates the secondary carbocation (as secondary carbocations are more stable than primary carbocations). I explain to my students that this example ties back to our discussion of energy diagrams because the highest energy MO arrangement has the node at the center carbon—carbon #3—with p orbitals overlapping to generate a pi bond on each side of the cation.
Even though MO theory can be tricky for students, my favorite part of teaching it is when students experience their “lightbulb moment” in class. When I was discussing the importance of these MO diagrams and how they relate to organic molecules in my (flexible) hybrid classroom earlier, I received an emoji with a brain exploding in our chat. This made my day, not only because it meant that my students finally “got” the smaller connections between MO concepts, but also that they “got” how MO theory relates to the bigger picture in organic chemistry.
-Kerri Taylor, Columbus State University
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