Chapter9-Orbitals

Chapter 9: Covalent Bonding and Orbitals

  • Introduction to covalent bonding and its relation to atomic orbitals.

  • Building on previous chapters (Chapter 7: quantum mechanical model, Chapter 8: covalent bonds and Lewis structures).

Hybridization and Localized Electron Model

Definition of Hybrid Orbitals

  • Hybrid orbitals explain molecular bond angles.

  • Example: Methane (CH4) involves the hybridization of one 2s and three 2p orbitals from carbon to create four sp³ hybrid orbitals.

  • Hybridization is influenced by molecular geometry, requiring the understanding of VSEPR theory.

Conservation of Orbitals

  • The number of atomic orbitals is conserved when forming molecular orbitals.

Sp³ Hybridization

Structure and Geometry

  • Hybridized orbitals are equivalent in energy (degenerate).

  • In CH4, the tetrahedral arrangement is formed with four equivalent C-H bonds created through sp³ hybridization.

  • Internal representation includes Lewis structure and the arrangement of valence orbitals.

Effective Electron Pairs

  • In ammonia (NH3), the nitrogen atom also undergoes sp³ hybridization, confirming VSEPR predictions.

Sp² and Sp Hybridization

Sp² Hybridization

  • One unhybridized 2p orbital remains for bonding with another p orbital.

  • Structure for single, double, and triple bonds:

    • Single bond: Sigma (σ)

    • Double bond: Sigma (σ) + Pi (π)

    • Triple bond: Sigma (σ) + Two Pi (π)

Sp Hybridization

  • Similar bond structure as sp² for triple bonds: sigma (σ) and two pi (π).

Sigma and Pi Bonds

  • Definitions of bonds in terms of σ- and π-bonds.

  • Example with the molecule Sevin for counting σ and π bonds.

Hybridization in Elements with d-Orbitals

  • In period 3 and above, d-orbitals can participate in bonding.

  • These can form σ-bonds, π-bonds, and δ-bonds for higher coordination numbers in rare bonds.

VSEPR Theory and Hybridization

Hybridization Summary Table

  • Effective Electron Pairs vs. Hybridization:

    • 2 pairs: Linear (sp)

    • 3 pairs: Trigonal planar (sp²)

    • 4 pairs: Tetrahedral (sp³)

    • 5 pairs: Trigonal bipyramid (dsp³)

    • 6 pairs: Octahedral (dsp³)

  • Ligand hybridization may differ from the central atom depending on Lewis and VSEPR structures.

Molecular Orbital Theory

Overview

  1. Valence electrons form molecular orbitals that are not independent.

  2. Bonding and antibonding orbitals illustrated in textbook figure.

  3. Stability occurs when bonding interactions outweigh antibonding interactions.

  4. Bond Order (BO) indicates net bonding strength; must be greater than zero for stability.

  5. Higher BO correlates with stronger bonds.

Conservation of Orbitals in MO Theory

  • Same number of bonding/antibonding orbitals formed through in-phase (bonding) and out-of-phase (antibonding) combinations.

Applications of Molecular Orbital Theory

  • Calculation examples for homonuclear and heteronuclear molecular orbitals such as O2 and CO, including bonding detections and unpaired electron counts.

  • Understanding complexities in molecular arrangements helps with predicting physical and chemical properties.