Study Guide: Chemical Bonding II – Molecular Shapes, VSEPR, and Molecular Orbital Theory

Chapter 11: Chemical Bonding II – Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

Introduction

This chapter explores advanced models of chemical bonding, focusing on the three-dimensional shapes of molecules, the Valence Shell Electron Pair Repulsion (VSEPR) theory, Valence Bond (VB) theory, and Molecular Orbital (MO) theory. Understanding these concepts is essential for predicting molecular geometry, bond properties, and the behavior of molecules in chemical reactions.

VSEPR Theory and Molecular Shapes

• VSEPR Theory: The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the geometry of molecules based on the repulsion between electron pairs (bonding and lone pairs) around a central atom.

• Electron Domains: Regions of electron density (bonds or lone pairs) around the central atom. The arrangement of these domains determines the molecular shape.

• Common Geometries:

  • Linear: 2 electron domains, 180° bond angle

  • Trigonal planar: 3 electron domains, 120° bond angle

  • Tetrahedral: 4 electron domains, 109.5° bond angle

  • Trigonal bipyramidal: 5 electron domains, 90°, 120°, and 180° bond angles

  • Octahedral: 6 electron domains, 90° and 180° bond angles

• Lone Pairs: Lone pairs occupy more space than bonding pairs, causing bond angles to decrease from ideal values.

• Predicting Shapes: Count the number of bonding and lone pairs to determine the electron geometry and molecular shape.

• Polarity: The shape and distribution of polar bonds determine if a molecule is polar or nonpolar.

Example: In water (H2O), there are two bonding pairs and two lone pairs on oxygen, resulting in a bent shape with a bond angle of about 104.5°.

Valence Bond Theory (VB Theory)

• VB Theory: Describes covalent bonding as the overlap of atomic orbitals from two atoms, each containing one unpaired electron.

• Hybridization: Atomic orbitals mix to form new, equivalent hybrid orbitals (e.g., sp, sp2, sp3).

• Types of Bonds:

  • σ (sigma) bonds: Formed by head-on overlap of orbitals; all single bonds are σ bonds.

  • π (pi) bonds: Formed by side-on overlap; present in double and triple bonds.

• Bond Strength and Length: Multiple bonds (double, triple) are shorter and stronger than single bonds.

Example: In methane (CH4), the carbon atom undergoes sp3 hybridization, forming four equivalent σ bonds with hydrogen atoms.

Molecular Orbital Theory (MO Theory)

• MO Theory: Atomic orbitals combine to form molecular orbitals that are delocalized over the entire molecule.

• Bonding and Antibonding Orbitals:

  • Bonding orbitals (lower energy): Increase electron density between nuclei, stabilizing the molecule.

  • Antibonding orbitals (higher energy): Decrease electron density between nuclei, destabilizing the molecule.

• Bond Order: Indicates the strength and stability of a bond. Calculated as: $\text{Bond Order} = \frac{1}{2}[(\text{number of electrons in bonding orbitals}) - (\text{number of electrons in antibonding orbitals})]$

• Magnetism: Molecules with unpaired electrons in MO diagrams are paramagnetic; those with all electrons paired are diamagnetic.

• HOMO-LUMO Gap: The energy difference between the Highest Occupied Molecular Orbital (HOMO) and the Lowest Unoccupied Molecular Orbital (LUMO) is important for understanding chemical reactivity and light absorption.

Example: In O2, MO theory predicts two unpaired electrons, explaining its paramagnetic behavior.

Applications and Problem Solving

• Be able to draw Lewis structures and predict molecular shapes using VSEPR theory.

• Assign hybridization states to atoms in molecules and identify the types of bonds present.

• Construct and interpret simple MO diagrams for diatomic molecules (e.g., H2, O2, N2).

• Predict bond order, bond energy, and bond length using MO theory.

• Determine if a molecule is paramagnetic or diamagnetic based on its MO diagram.

• Relate the HOMO-LUMO gap to molecular properties such as color and reactivity.

Summary Table: Common Molecular Geometries (VSEPR)

Electron Domains	Electron Geometry	Molecular Geometry	Bond Angle	Example
2	Linear	Linear	180°	CO2
3	Trigonal planar	Trigonal planar	120°	BF3
4	Tetrahedral	Tetrahedral	109.5°	CH4
4	Tetrahedral	Bent	~104.5°	H2O
5	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°, 180°	PCl5
6	Octahedral	Octahedral	90°, 180°	SF6

Additional info:

• Understanding the relationship between molecular structure and physical/chemical properties is crucial for predicting reactivity and function.

• MO theory is especially important for explaining phenomena that cannot be described by Lewis structures or VB theory alone, such as the paramagnetism of O2.