Molecular orbital (MO) theory, a complex but useful way to view electron orbitals in molecules, is the bonding theory most widely used by practicing chemists. This theory assumes that pure s and p orbitals of the atoms in the molecule combine to form delocalized, or molecular orbitals. MO theory is far more useful than valence bond theory for predicting electronic properties of molecules and for predicting stability and structure in transition metal complexes.
To understand how MO theory relates to the molecules not well served by the Lewis approach, a description of the four principles of MO theory is necessary. First, the number of molecular orbitals produced is always equal to the number of atomic orbitals brought by the atoms that have combined. Second, a bonding orbital is lower in energy that the parent orbitals, whereas the antibonding orbital is higher in energy. A third principle is that the electrons of the molecule are assigned to orbitals of successively higher energy, which follows the Pauli principle and Hund's rule. Lastly, the fourth principle states that atomic orbitals combine to form molecular orbitals most effectively when the atomic orbitas are of similar energy.
This module simulates the energy level diagram for transition metal complexes of three geometries: 6-coordinate octahedral, 5-coordinate square pyramidal, and 4-coordinate square planar. Three types of ligands are possible to study: sigma donors without pi-bonding effects, sigma donor-pi acceptors, and sigma donor-pi donors. The energy level diagram changes by varying the ligand type, and subsequently varying the amount of sigma or pi interaction in the molecule.
INSTRUCTIONS
1. Select a molecular structure from the Structure menu.
2. Select the LIgand type by clicking on one of the three Ligand Type option buttons.
3. Vary the Sigma and/or Pi interaction scrollbars to view the changes in the bonding orbitals.
