At the end of lecture on Friday I reviewed some concepts pertaining to molecular orbital theory, which is covered in general chemistry classes at Ohio State in Chapter 9 of the textbook. Here are a few key points you need to consider when you are dealing with molecular orbital theory.
*It's all about the orbital overlap and the electrons residing in the orbitals.
Know the shapes of all the orbitals. If you can't draw the d orbitals like the back of your hand you will fail the 2nd midterm.
Know how to determine the oxidation state and d electron count for the transition metals in transition metal complexes (we will discuss this soon). If you can't do this by the time you get to the 2nd midterm you are in really bad shape.
*The better the orbital overlap the stronger the bond.
*Orbital overlap depends not only whether an orbital is s, p, or d, but also on how those orbitals are oriented in space.
*Before you come to class on Wed, know the shapes of the d orbitals.
*When orbitals of the same phase overlap, they form a bonding interaction which lowers the overall energy of the atomic orbitals interacting making them more stable.
*When orbitals of opposite phases overlap they form an anti-bonding interaction which raises the overall energy of the atomic orbitals interacting, making them less stable
*When you construct a molecular orbital diagram, the total number of all the molecular orbitals must be equal to the sum of all the atomic orbitals. Molecular orbitals can be bonding, anti-bonding, or non-bonding.
Song Played Before Lecture:
Bruce Springsteen - Born To Run
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