5 Key Concepts of Molecular Orbital Theory

Molecular Orbital Theory (MOT) is a fundamental concept in chemistry used to explain how atomic orbitals combine to form molecular orbitals and how electrons are distributed in molecules. Understanding MOT can help solve many practical problems, such as predicting molecular stability, bond order, and magnetic properties. However, students and professionals often struggle with its abstract nature and mathematical underpinnings. This guide will simplify the key concepts of Molecular Orbital Theory, providing practical examples and actionable steps to help you apply this theory effectively.

Whether you're a student trying to grasp the basics or a professional needing a refresher, this guide will break down the five key concepts of MOT in an easy-to-understand way. By the end, you'll be equipped with a solid understanding of molecular orbitals, bond order, and how to predict molecular behavior. Let’s dive in!

1. Atomic Orbitals Combine to Form Molecular Orbitals

The core idea of MOT is that atomic orbitals (AOs) on individual atoms combine to form molecular orbitals (MOs) when atoms bond. These molecular orbitals spread over the entire molecule and are classified into two types: bonding and antibonding orbitals.

Bonding Orbitals: When atomic orbitals combine constructively (in-phase), they create bonding orbitals. These orbitals have lower energy than the original atomic orbitals, making the molecule more stable.

Antibonding Orbitals: When atomic orbitals combine destructively (out-of-phase), they form antibonding orbitals. These have higher energy and destabilize the molecule.

For example, in a hydrogen molecule (H₂), the 1s orbitals of two hydrogen atoms combine to form one bonding orbital (σ1s) and one antibonding orbital (σ*1s). If both electrons occupy the bonding orbital, the molecule is stable.

Actionable Tip: When analyzing molecules, always identify how atomic orbitals combine and classify the resulting molecular orbitals.

2. Electrons Fill Molecular Orbitals Based on Energy Levels

Similar to the Aufbau principle in atomic orbitals, electrons occupy molecular orbitals starting from the lowest energy level. The order of filling depends on the specific molecule and often follows a predictable pattern, especially in diatomic molecules.

For homonuclear diatomic molecules (e.g., O₂, N₂), the energy level order typically changes depending on the number of electrons. For molecules with fewer than 14 total electrons (e.g., B₂, C₂, N₂), the order is:

• σ2s → σ*2s → π2px = π2py → σ2pz → π*2px = π*2py → σ*2pz

For molecules with more than 14 total electrons (e.g., O₂, F₂), the order shifts slightly:

• σ2s → σ*2s → σ2pz → π2px = π2py → π*2px = π*2py → σ*2pz

Pro Tip: Memorize these patterns for quick application in exams or problem-solving scenarios.

For instance, oxygen (O₂) has 16 total electrons. Following the second order, the molecular orbitals are filled, leaving unpaired electrons in the π*2px and π*2py orbitals. This explains why O₂ is paramagnetic (attracted to magnetic fields).

3. Bond Order Predicts Molecular Stability

Bond order is a key metric in MOT that predicts the stability of a molecule. It is calculated using the formula:

Bond Order = (Number of bonding electrons - Number of antibonding electrons) / 2

A higher bond order indicates a stronger and more stable bond. For example, in the hydrogen molecule (H₂), there are two bonding electrons and no antibonding electrons:

• Bond Order = (2 - 0) / 2 = 1

In contrast, for the helium molecule (He₂), there are two bonding electrons and two antibonding electrons:

• Bond Order = (2 - 2) / 2 = 0

This indicates that He₂ is not stable and does not exist as a molecule under normal conditions.

Actionable Tip: Use bond order calculations to quickly assess whether a molecule is likely to form and how strong its bonds will be.

4. Symmetry and Overlap Determine Orbital Interactions

Not all atomic orbitals combine to form molecular orbitals. For effective orbital interaction, the following conditions must be met:

• Similar Energy: Atomic orbitals must have comparable energy levels to combine.
• Proper Symmetry: Orbitals must align in a way that allows for constructive or destructive interference. For example, s orbitals combine well with other s orbitals or p orbitals aligned along the same axis (e.g., σ bonds).
• Significant Overlap: Orbitals must physically overlap to interact. Greater overlap leads to stronger bonding interactions.

For example, in carbon monoxide (CO), the 2p orbital of carbon interacts with the 2p orbital of oxygen to form bonding and antibonding orbitals. The overlap is particularly strong because of their similar energy levels.

Pro Tip: When predicting orbital combinations, always check symmetry and energy compatibility first.

5. Molecular Orbital Diagrams Provide Visual Insights

Molecular orbital diagrams are invaluable tools for visualizing how electrons are distributed in a molecule. These diagrams show the relative energy levels of molecular orbitals and the placement of electrons within them.

Here’s how to construct a molecular orbital diagram:

1. Identify the atomic orbitals: Determine the atomic orbitals for each atom in the molecule (e.g., 1s, 2s, 2p).
2. Combine orbitals: Combine atomic orbitals to create bonding and antibonding molecular orbitals.
3. Order the orbitals: Arrange the molecular orbitals based on their energy levels (use the patterns discussed earlier).
4. Fill the orbitals: Add electrons to the molecular orbitals, starting with the lowest energy level, following Hund’s rule and the Pauli exclusion principle.

For example, the molecular orbital diagram for nitrogen (N₂) shows that all bonding orbitals are filled, resulting in a bond order of 3 and a very stable triple bond.

Pro Tip: Practice drawing molecular orbital diagrams for common diatomic molecules to build confidence and intuition.