Molecular Orbital Theory and Chemical Bonding: Study Guide

Molecular Orbital Theory and Chemical Bonding

Introduction

Molecular orbital theory is a fundamental concept in general chemistry that explains the electronic structure of molecules, particularly how atomic orbitals combine to form molecular orbitals. This theory provides insight into chemical bonding, bond order, magnetism, and the properties of molecules.

Key Concepts in Molecular Orbital Theory

• Molecular Orbital (MO): A region in a molecule where electrons are likely to be found; formed by the combination of atomic orbitals.

• Bonding Orbital: An MO that is lower in energy than the atomic orbitals from which it is formed; electrons in these orbitals stabilize the molecule.

• Antibonding Orbital: An MO that is higher in energy than the atomic orbitals from which it is formed; electrons in these orbitals destabilize the molecule.

• Delocalized Electrons: Electrons that are not associated with a single atom or bond but are spread over several atoms or bonds.

• Resonance: The concept that some molecules can be represented by two or more valid Lewis structures; molecular orbital theory provides a more complete description by showing electron delocalization.

Formation of Molecular Orbitals

When two atomic orbitals combine, they form two molecular orbitals: one bonding and one antibonding. The number of molecular orbitals formed equals the number of atomic orbitals combined.

• Bonding MO: Constructive interference of atomic orbitals; increases electron density between nuclei.

• Antibonding MO: Destructive interference; creates a node between nuclei.

Bond Order

Bond order indicates the strength and stability of a bond. It is calculated as:

• Formula:

$\text{Bond Order} = \frac{(\text{Number of electrons in bonding MOs}) - (\text{Number of electrons in antibonding MOs})}{2}$

• Example: For $\mathrm{O}_2$, bond order = $\frac{8-4}{2} = 2$

Paramagnetism and Diamagnetism

• Paramagnetic: Molecules with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Molecules with all electrons paired; repelled by magnetic fields.

• Example: $\mathrm{O}_2$ is paramagnetic due to two unpaired electrons in its molecular orbital diagram.

Molecular Orbital Diagrams

Molecular orbital diagrams visually represent the arrangement of electrons in molecular orbitals. They are essential for determining bond order, magnetism, and stability.

• Steps to Draw MO Diagrams:

  1. Write the electron configuration for each atom.

  2. Combine atomic orbitals to form molecular orbitals (sigma and pi types).

  3. Fill molecular orbitals with electrons according to the Aufbau principle, Pauli exclusion principle, and Hund's rule.

  4. Calculate bond order and determine magnetic properties.

• Example: MO diagram for $\mathrm{N}_2$ shows a bond order of 3 and is diamagnetic.

Application to Heteronuclear Diatomic Molecules

When applying MO theory to molecules with different atoms (e.g., CO, NO), atomic orbitals of different energies combine, resulting in molecular orbitals that are not evenly shared between atoms.

• Example: In CO, the molecular orbitals are closer in energy to the atomic orbitals of oxygen due to its higher electronegativity.

Comparison of Bonding Theories

Molecular orbital theory, Lewis theory, and valence bond theory each provide different perspectives on chemical bonding:

• Lewis Theory: Focuses on electron pairs and resonance structures.

• Valence Bond Theory: Describes bonds as localized overlaps of atomic orbitals.

• Molecular Orbital Theory: Explains delocalization and magnetic properties.

Example: MO theory explains the paramagnetism of $\mathrm{O}_2$, which Lewis theory cannot.

Important Vocabulary

• Molecular orbital (MO)

• Molecular orbital (MO) theory

• Bonding orbital

• Antibonding orbital

• Molecular orbital diagram

• Bond order

• Paramagnetic

• Diamagnetic

Sample Molecular Orbital Table

The following table summarizes the bond order and magnetic properties of selected diatomic molecules:

Molecule	Bond Order	Magnetic Property
$\mathrm{H}_2$	1	Diamagnetic
$\mathrm{He}_2$	0	Diamagnetic
$\mathrm{O}_2$	2	Paramagnetic
$\mathrm{N}_2$	3	Diamagnetic
$\mathrm{B}_2$	1	Paramagnetic

Practice and Further Study

• Review textbook sections 6.4 - 6.5 for detailed explanations and diagrams.

• Watch molecular orbital theory videos for visual understanding.

• Practice drawing MO diagrams for period 2 diatomic molecules and calculating bond order.

Additional info:

• Questions in the file focus on molecular orbital theory, bond order, paramagnetism/diamagnetism, and comparison of bonding theories, all of which are core topics in General Chemistry Chapter 10 (Molecular Shapes & Valence Bond Theory) and Chapter 9 (Chemical Bonding I: The Lewis Model).