Chapter 9: Bonding and Molecular Structure

Transition to Chapter 9: Discussion on Lewis structures and bond lengths specific to organic molecules.

• Focus on carbon compounds (e.g., C2H4, C2H2).
• Shorter carbon-carbon bond lengths correlate with higher bond order (single < double < triple).

Recap of Bond Concepts:

• Condensed Structural Formula: Lists connectivity in organic molecules.
• Bond Orders:
• Single bond (C-C) → bond order = 1.
• Double bond (C=C) → bond order = 2.
• Triple bond (C≡C) → bond order = 3.
• Bond Length:
• Triple bond (< double bond < single bond in length).

Bond Energies:

• Comparison of bond energies: Single (368 kJ/mol), Double (682 kJ/mol), Triple (962 kJ/mol).
• Bond energies do not scale linearly due to electron repulsion effects.

Chapter 9 Focus:

• Introduction to Valence Bond Theory and Orbital Hybridization (focusing on understanding geometries and bond angles).
• Understanding the shape and angles of molecules like tetrahedral (109.5°) based on bonding theory.

Valence Bond Theory:

• Bonds form through overlap of atomic orbitals, resulting in covalent bonds.
• Types of Bonds:
• Sigma bond (σ): Created by direct overlapping of orbitals (shared electron density along bond axis).
• Example with hydrogen (H2) showing overlap of 1s orbitals forming a sigma bond with a bond length of 74 picometers, minimizing repulsion.

Hybridization of Atomic Orbitals:

• Concept: Mixing of atomic orbitals to explain bond formation and molecular geometries.
• Types of Hybridization:
• sp³: tetrahedral arrangements (e.g., methane, CH4) with 4 equivalent hybrid orbitals.
• sp²: trigonal planar arrangements (e.g., boron trifluoride, BF3).
• sp: linear arrangements (e.g., beryllium compounds, BeCl2).

Sigma Bonding in Hydrocarbons:

• Example of Carbon:
• Carbon's electron configuration: 1s² 2s² 2p².
• Hybridization mixes 1s and 3p to form 4 equivalent sp³ orbitals, each holding an unpaired electron allowing for 4 bonds.
• Importance of Bond Angles: Bond angles result from the arrangement of hybrid orbitals, which impacts molecular shape.

Example of Hybridization:

• Methane (CH4):
• Uses sp³ hybridization.
• Each sp³ orbital overlaps with one 1s orbital from hydrogen to form 4 sigma bonds at 109.5°.
• Boron Trifluoride (BF3):
• Uses sp² hybridization resulting in a trigonal planar shape with bond angles of 120°.

Electron Group Geometry: Determined by counting bonding pairs and lone pairs around the central atom.

Lone Pairs and Geometry:

• Example: Ammonia (NH3) has lone pairs influencing the bond angle.
• Hybridization for nitrogen in NH3 is sp³, which involves 3 sigma bonds and one lone pair altering the angle to approximately 107.5°.

Summary:

• Understanding hybridization of orbitals helps predict molecular shapes and bond angles.
• Each type of hybridization corresponds to different electron geometries, influenced by the numbers of sigma bonds or lone pairs present in the molecule.
• Practice recognizing hybridization types based on Lewis structures and molecular geometries to enhance understanding of molecular behavior in organic compounds.