_lecture 3 CFM_ slidss -

Learning Outcomes

  • Understand sigma (σ) and pi (π) bonds

  • Describe orbital hybridization as per valence bond theory

  • Understand hybridization of atomic orbitals:

    • sp3 hybridization

    • sp2 hybridization

    • sp hybridization

  • Focus on Chemical Bonding: Orbital Hybridization (C, N, O)

Valence Bond Theory (VB Theory)

  • Valence Bond Theory considers the interaction of atomic orbitals from separate atoms forming molecules.

  • Bonds form from the pairing of unpaired electrons in atomic orbitals.

    • H atom: 1s^1 paired with another H atom's 1s^1

Sigma (σ) Bonds

  • Formed from interactions of:

    • Two s orbitals

    • Two p orbitals (no nodal plane along the bond axis)

    • One s and one p orbital

  • Important to note that σ bonds allow for free rotation about the bond axis.

Pi (π) Bonds

  • Formed from the overlap of:

    • Two p orbitals (nodal plane along the bond axis)

  • Pi bonds restrict rotation due to the presence of nodal planes.

Relationship of Sigma (σ) and Pi (π) Bonds

  • A single bond is a σ-bond.

  • A double bond consists of a σ-bond plus one π-bond.

  • A triple bond comprises a σ-bond plus two π-bonds.

Hybridisation of Atomic Orbitals

Example 1: Carbon

  • Electron configuration: 1s^2 2s^2 2px^1 2pz^1

  • Carbon forms 4 bonds but cannot do so with this electron configuration because the 2s orbital is full and stable.

sp3 Hybridization

  • Involves:

    • Promotion of an electron from 2s to 2p

  • Resulting configuration: 1s^2 2s^1 2px^1 2py^1 2pz^1

    • 4 sp3 hybrid orbitals formed from one s and three p orbitals

  • Energetic favourability leads to the formation of C-H bonds in methane.

sp2 Hybridization

  • Formed from one s orbital and two p orbitals.

  • Example: Ethylene, exhibiting σ and π bonds, restricting rotation due to double bond.

sp Hybridization

  • Involves the formation of one s orbital and one p orbital.

  • Triple bond in acetylene consists of one σ and two π bonds.

Molecular Orbital (MO) Theory

  • Valence electrons are delocalized across the entire molecule.

  • MOs arise from combining/superimposing atomic orbitals.

  • Bonding MO (σ1s) is lower in energy, stable, and arises from constructive interference.

  • Antibonding MO (σ1s*) results from destructive interference, with a node and higher energy.

Bond Order Calculation

  • Bond order = ½ (number of bonding electrons – number of antibonding electrons)

    • Example: For He2, bond order = ½(2-2)= 0 (no bond)

    • For Li2, bond order = ½(4-2)= 1 (having a bond).

Examples of Bonding in Molecular Orbital Theory

  • Hydrogen (H2): Electron configuration in bonding orbital (σ1s).

  • Di-helium (He2): No bond due to equal electrons in bonding and antibonding orbitals.

  • Beryllium (Be2): Also shows no bond with 0 bond order in total and valence electrons.

Upcoming Topics in Chemical Bonding

  • Understanding atomic structure and bonding theories, Lewis structures, molecular geometry, orbital hybridization, resonance, and aromaticity in organic compounds.

  • Discussions on thermodynamics, entropy, Gibbs energy, and kinetic experiments.