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.