Molecular Orbital Theory

Introduction

• Just as valence bond theory was introduced as a higher level of theory (as compared to VSEPR) molecular orbital theory (MO theory) is simply another even higher level of theory that was developed to explain some experimental observations.

• The basis of MO theory is similar to hybridization. This time though we will combine all the atomic orbitals from all the atoms and create new molecular orbitals from that

• It is very difficult to predict these MO’s in general but you should know how to read and generate MO diagrams for diatomic molecules

• Another concept becomes important in MO theory and that is the idea of bonding and anti-bonding orbitals. Anti-bonding orbitals are orbitals in which there is a node between the two atoms in question as is shown below

• Once orbitals have been identified as bonding or anti-bonding we can discuss bond order. Bond order is a measure of bond strength. ie a triple bond would have a bond order of 3. To calculate the bond order from a MO diagram (we’ll see this soon) you need to use the equation shown below.

• Bond order describes the degree of bonding between the two atoms and estimates the strength of the bond. The higher the bond order, the stronger the bond and therefore the shorter the bond.

• If the bond order = 0, there is no bonding between the atoms and the molecule is unstable.

How to Draw MO Diagrams for Homonuclear Diatomics

• As an example the MO diagram for the simplest homodiatomic, hydrogen, is shown below

• Note: Now the atoms and the homodiatomic are all on the same energy axis.

• Hydrogen has one electron so each side of the diagram is filled with one electron

• H2 has two electrons so the middle column of orbitals is filled with two electrons.

• Filling the orbitals follows the same three rules that we talked about in chapter 2.

• The MO picture for H2 is fairly simple because we only have one orbital on each atom.When we move to the second row of the periodic table we have to consider the p orbitals.

• Both the π and π* orbitals are comprised of two degenerate pi orbitals.

• This orbital diagram is the same for all homodiatomics of the second row from Li to N.

• When we move to the right of nitrogen along the second row there is a small change to the MO picture. The σ2p switches places with the π2p as shown in figure 8.4 below.

• The reason for this is complex but a partial explanation is obtained by noting that the 2s-2p orbital gap in oxygen is larger (due to electronegativity) and so less 2s-2p “mixing” is observed.

• O2 is a good example of why we need MO theory. MO theory predicts that O2 is paramagnetic whereas VBT does not.