Asking to make sure I understand: Valence bond vs. Molecular Orbital theory.
Hello. I am currently a freshman in college and finding the chapter on Quantum Mechanics incredibly interesting. Whether I decide to delve into this field is still up in the air, but I am having trouble fully understanding the concepts behind Molecular Orbital Theory and how it relates to the Valence Bond Theory, and would really like to get a grasp on these theories simply for the sake of understanding.
I understand the VBT basically points out that orbitals overlap or mix to create strong bonds between molecules and they hybridize whenever necessary to make room and accommodate sigma and pi bonds. How do these normal and hybrid orbitals relate to the model shown in MO diagrams? I understand that the MO theory states that bonds become delocalized throughout the entire molecule, but it seems that the lowest state bonding orbital and anti-bonding orbital bond s orbital electrons when it isn't necessary; why is this?
For example: According to Valence Bond, O2 is made from two Oxygen atoms that overlap with their respective 2py and 2pz orbitals, leaving 2s and 2px untouched. But in the MO diagram, each oxygen atom contributes 1 electron from its own 2s orbital to the 2s molecular bonding orbital and the 2s molecular anti-bond. Then once again, each Oxygen contributes 1 electron from its 2px orbital to the 2px molecular bonding orbital. Why is this? I thought those electrons I mentioned remained separate as lone pairs? Am I just looking at this the completely wrong way?
I understand this is a full question, but I would really appreciate some help. :)