Energy Level Diagram - Practice Questions & MCQ

Molecular Orbital Energy Diagrams

The relative energy levels of atomic and molecular orbitals are typically shown in a molecular orbital diagram. As given in the figure below, for a diatomic molecule, the atomic orbitals of one atom are shown on the left, and those of the other atom is shown on the right. Each horizontal line represents one orbital that can hold two electrons. The molecular orbitals formed by the combination of the atomic orbitals are shown in the center. Dashed lines show which of the atomic orbitals combine to form the molecular orbitals. For each pair of atomic orbitals that combine, one lower-energy (bonding) molecular orbital and one higher-energy (antibonding) orbital result. Thus we can see that combining the six 2p atomic orbitals results in three bonding orbitals (one σ and two π) and three antibonding orbitals (one σ* and two π*).

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molecular orbital diagram 

The molecular orbitals are filled in the same manner as atomic orbitals, using the Aufbau principle and Hund’s rule.

Bond Order

The filled molecular orbital diagram shows the number of electrons in both bonding and antibonding molecular orbitals. The net contribution of the electrons to the bond strength of a molecule is identified by determining the bond order that results from the filling of the molecular orbitals by electrons.

The MO technique is more accurate and can handle cases when the Lewis structure method fails, but both methods describe the same phenomenon.

In the molecular orbital model, an electron contributes to a bonding interaction if it occupies a bonding orbital and it contributes to an antibonding interaction if it occupies an antibonding orbital. The bond order is calculated by subtracting the destabilizing (antibonding) electrons from the stabilizing (bonding) electrons. Since a bond consists of two electrons, we divide by two to get the bond order. We can determine bond order with the following equation:

bond order = [(number of bonding electrons)−(number of antibonding electrons)]/2

The order of a covalent bond is a guide to its strength; a bond between two given atoms becomes stronger as the bond order increases. If the distribution of electrons in the molecular orbitals between two atoms is such that the resulting bond would have a bond order of zero, a stable bond does not form. 

For example, the bond order of H₂ molecule is given as follows:

The molecular orbital energy diagram predicts that H₂ will be a stable molecule with lower energy than the separated atoms.

A dihydrogen molecule contains two bonding electrons and no antibonding electrons so we have:

bond order in H₂=(2−0)/2=1

Because the bond order for the H–H bond is equal to 1, the bond is a single bond.

Magnetic Moment

The magnetic behaviour of any molecule can be determined from the number of unpaired electrons in the bonding and antibonding orbitals. The molecule is said to be diamagnetic as there is no unpaired electron present in the orbitals and not attracted by the magnet. But if any unpaired electron is present then the molecule is paramagnetic.

For example, O₂ molecule has 2 unpaired electrons can be seen from the diagram below:

Therefore, O₂ molecule is paramagnetic.