Fundamentals of Organic Chemistry

Chemical Bonding

This section introduces the foundational concepts of chemical bonding in organic chemistry, focusing on atomic structure, bonding models, and their impact on molecular properties.

Intended Learning Outcomes

• Describe the atomic structure and model of the atom.

• Explain the influence of bond polarization on a molecule's structure and reactivity.

• Describe different bonding models and appreciate how these impact the properties of a simple molecule.

• Describe the acid-base properties of organic molecules and how they impact their reactivity.

• Recognize functional groups of simple organic molecules.

• Relate the functional group of simple organic molecules with their physical properties.

The Nature of Chemical Bonds

Valence Bond Theory

Valence Bond Theory explains how covalent bonds form when two atoms approach each other closely, allowing a singly occupied orbital on one atom to overlap with a singly occupied orbital on another atom. This overlap leads to electron pairing and bond formation.

• Covalent Bond: Formed by the overlap of singly occupied orbitals from two atoms.

• Bonding Models:

  • Valence Bond Theory

  • Molecular Orbital Theory

Valence Bond Theory: Hydrogen Molecule Example

• Electrons are paired in the overlapping orbitals and attracted to both nuclei.

• The H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals.

• The H–H bond is cylindrically symmetrical, known as a sigma (σ) bond.

Example: Formation of H2 molecule:

• Each hydrogen atom has a 1s orbital with one electron.

• Overlap of these orbitals forms a σ bond.

Bond Energy

Bond energy is the energy released when a bond forms or absorbed when a bond breaks. It is a measure of bond strength.

• Reaction: releases 436 kJ/mol.

• Product (H2) has 436 kJ/mol less energy than two separate H atoms.

• Bond strength of H–H: 436 kJ/mol.

• Unit conversions: ;

Bond Length

Bond length is the distance between the nuclei of two bonded atoms that leads to maximum stability.

• If atoms are too close, they repel due to positive charge of nuclei.

• If atoms are too far apart, bonding is weak.

• Optimal bond length for H–H: 74 pm (picometers).

Example: The bond length of H–H in H2 is 74 pm, corresponding to the lowest energy and greatest stability.

Hybridization and Molecular Structure

sp3 Orbitals and the Structure of Methane

Hybridization is the process by which atomic orbitals mix to form new, equivalent hybrid orbitals suitable for bonding.

• Carbon's Electron Configuration: (4 valence electrons)

• In methane (CH4), all C–H bonds are identical and arranged tetrahedrally.

• sp3 Hybrid Orbitals: One s orbital and three p orbitals combine to form four equivalent, unsymmetrical tetrahedral orbitals.

• Each sp3 orbital overlaps with a hydrogen 1s orbital to form a C–H bond.

• Bond angle in methane: 109.5° (tetrahedral angle).

Example: Methane (CH4) has four identical C–H bonds, each formed by the overlap of a carbon sp3 orbital and a hydrogen 1s orbital.

Additional info: Hybridization theory helps explain the geometry and equivalence of bonds in organic molecules, which is essential for understanding their chemical behavior.