Molecular-Orbital Theory
In the mid-1960's, scientists worked to find a new theory to explain the paramagnetism of oxygen and other properties of molecules that could not be accounted for by the VSEPR and hybrid orbital theories. The result of the research was the molecular-orbital theory (or "MO" for short), for which Robert Millikan, the discoverer, was awarded the Nobel Prize for chemistry in 1966. MO theory and the VSEPR/hybrid orbital theories use different mechanisms to explain atomic bonding and the orbitals involved. Neither is "more correct" than the other; they are used to explain different aspects of the molecule under study, so don't slam your head against the monitor if these two theories don't seem to fit together! MO theory is quite complex and frankly not terribly useful in general chemistry, so we will only scratch the surface of this area.
For example, MO theory reverses some of the principles of the previous pages: normally, we believe that the configuration of the atomic orbitals determines the shape of the molecule. In MO, however, the shape of the molecule is determined by experiment, and the distribution of orbitals is then studied.
MO theory reincorporates some of the wave theory of the electron, because it predicts that the interaction of two atomic orbitals will produce a bonding molecular orbital and an anti-bonding molecular orbital; this is analogous to waves interfering and either amplifying (resulting in the bonding MO) or canceling each other (producing an antibonding orbital). Each type of atomic orbital (σ and π) also has an antibonding orbital, symbolized by a sigma or pi symbol with a superscript asterisk (σ* and π*). For example, s orbitals can only produce a sigma bond, so each s orbital has one σ and one σ* bond. The three p orbitals, however, can form one sigma and two pi bonds, because they point in three directions: up, sideways, and down; for any bonding orientation, one head-on sigma bond and two sideways pi bonds can form. Therefore, each p subshell has one σ orbital, two π orbitals, one σ* orbital, and two π* orbitals.
We use a diagram like that below to allocate electrons to molecular orbitals. The example below is oxygen, since we're trying to explain its paramagnetic behavior:
This MO diagram explains quite a bit. First, the 1s and 2s bonding orbitals and antibonding orbitals cancel out, representing that these orbitals do not participate in the bonding process. In the 2p orbital, the total number of shared electron pairs not cancelled by the antibonding orbital is two, representing the double covalent bond. Finally, the unpaired electrons in the 2p π antibonding orbital finally explain O2's paramagnetism.
This page finishes our discussion of atoms and molecules. We hope you've learned a lot about how the most basic units of matter join together to form all the substances we see in the universe. You should give yourself a big pat on the back, because you've completed two of the three most difficult chapters in this web site! Put on your safety goggles, because the next chapter, "Acids and Bases" will describe some of the most important reactions in chemistry. You'll learn about subjects including acids strong enough to burn through metal, bases that can turn skin to soap, and the buffers that protect us from such dangerous compounds. Good luck!
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