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 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 (Valence Shell Electron Pair Repulsion) predicts the shapes of molecules based on the repulsion between electron pairs (bonding and lone pairs) around a central atom.

• Electron Domains include both bonding pairs and lone pairs of electrons.

• Basic Principle: Electron pairs arrange themselves as far apart as possible to minimize repulsion.

• Common Geometries:

  • 2 electron domains: Linear (180°)

  • 3 electron domains: Trigonal planar (120°)

  • 4 electron domains: Tetrahedral (109.5°)

  • 5 electron domains: Trigonal bipyramidal (90°, 120°)

  • 6 electron domains: Octahedral (90°)

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

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

• Polarity: Molecular shape and the 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

• Valence Bond Theory describes covalent bonds as the overlap of atomic orbitals from two atoms.

• 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 ethene (C2H4), each carbon is sp2 hybridized, forming a σ bond framework and a π bond between the carbons.

Molecular Orbital (MO) Theory

• MO Theory describes electrons in molecules as occupying molecular orbitals that are spread over the entire molecule.

• Bonding and Antibonding Orbitals:

  • Bonding orbitals (σ, π): Lower energy, increase electron density between nuclei.

  • Antibonding orbitals (σ*, π*): Higher energy, decrease electron density between nuclei.

• Bond Order: Indicates bond strength and is 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 chemical reactivity and light absorption.

Example: O2 has two unpaired electrons in its π* orbitals, making it paramagnetic.

Comparison of Bonding Theories

Theory	Main Features	Applications
VSEPR	Predicts molecular shape based on electron pair repulsion	Geometry, bond angles, polarity
Valence Bond	Describes bonds as overlap of atomic orbitals; hybridization	Bond formation, hybrid orbitals, σ/π bonds
Molecular Orbital	Electrons in molecular orbitals delocalized over molecule	Bond order, magnetism, spectroscopy

Key Learning Objectives

• Predict molecular shapes using VSEPR theory.

• Assign hybridization states to atoms in molecules.

• Draw and interpret Lewis structures and resonance forms.

• Distinguish between σ and π bonds in molecules.

• Construct and analyze simple MO diagrams for diatomic molecules.

• Calculate bond order and predict magnetic properties using MO theory.

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

• Understand the basics of molecular spectroscopy and its relation to electronic structure.

Additional info:

• For more complex molecules, hybridization and MO theory can be extended to explain delocalized bonding (e.g., in benzene).

• Spectroscopic techniques such as UV-Vis and IR spectroscopy provide experimental evidence for molecular orbital transitions and bonding.