Molecular Orbital Theory
To date, we have looked at three different theories of molecular boning. They are the VSEPR Theory (with Lewis Dot Structures), the Valence Bond Theory (with hybridization) and Molecular Orbital Theory. A good theory should predict physical and chemical properties of the molecule such as shape, bond energy, bond length, and bond angles.
WHY THREE THEORIES?
One model does not describe all the properties of molecular bonds. Each model describes a set of properties better than the others. The final test for any theory is experimental data.
The Molecular Orbital Theory does a good job of predicting electronic spectra and paramagnetism, when VSEPR and the V-B Theories don't. The MO theory does not need resonance structures to describe molecules, as well as being able to predict bond length and energy. The major draw back is that we are limited to talking about diatomic molecules (molecules that have only two atoms bonded together), or the theory gets very complex.
The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike the V-B theory, which treats the electrons as localized baloons of electron density, the MO theory says that the electrons are delocalized. That means that they are spread out over the entire molecule.
Now, when two atoms come together, their two atomic orbitals react to form two possible molecular orbitals. One of the molecular orbitals is lower in energy. It is called the bonding orbital and stabilizes the molecule. The other orbital is called an anti-bonding orbital. It is higher in energy than the original atomic orbitals and destabilizes the molecule.
The MO Theory has five basic rules:
1. The number of molecular orbitals = the number of atomic orbitals combined
2. Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-bonding orbital (higher energy)
3. Electrons enter the lowest orbital available
4. The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle)
5. Electrons spread out before pairing up (Hund's Rule)
The two AO's or atomic orbitals combine to form 2 MO's - the bonding and the anti-bonding molecular orbitals. Also, notice that the five rules have been followed, the electrons having been placed in the lowest enegry orbital(rule 3) and have paired up(rule 4) and there are only two electrons in the orbitals(rule 5).
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If you notice at the very bottom of the above picture, "bond order" is mentioned. If a molecule is to be stable, it must have a bond order greater that 0. Bond order is calculated as: 1/2 ( # of electrons in bonding orbitals - # of electrons in anti-bonding orbitals). If the bond order is 0, the molecule is unstable and won't form. If the bond order is 1 a single bond is formed. If the BO (bond order) is 2 or 3 a double or triple bond will be formed respectively.
When the 2nd period atoms are bonded to one another, you have both 2s and three 2p orbitals to contend with. When this happens, you have twice as many MO's!
Finally, we can put the Molecular Orbital Theory to use!
The answer is dilithium because it has a bond order of 1 which is stable and diberylium has a BO of 0 which is unstable and therefore will not form.
Look at the following MO diagrams for some of the period two elements. Can you tell which molecules are paramagnetic? Which molecules have the highest bond energy, which has the lowest? Rank single, double, and triple bonds in order of bond energy and bond length. (hint a BO of 1 is a single bond, 2 a double...)
Finaly, it was mentioned earlier that the MO Theory did not need resonance structures to explain anything. Because the MO theory holds that electrons are not held to only one position. Instead they are spread across the entire molecule. Below is a picture of Benzene and Ozone. You can see that rather than having two resonance structures, we can picture one structure with the electrons dispersed over the entire molecule.
