Molecular Geometry and Bonding Theories Notes

  • Molecular Shapes

    • Lewis structures provide bonding and lone pair information but not shape.

    • Used to predict molecular shapes.

  • Shape Determination

    • Bond angles and lengths define shape.

    • Electron pairs repel; they spread apart to minimize repulsion (VSEPR model).

  • Electron Domains

    • Electron domains indicate directions of electron pairs.

    • A central atom can have multiple electron domains.

  • VSEPR Model

    • Best arrangement minimizes electron domain repulsion.

  • Electron-Domain Geometry

    • Characterized by number of electron pairs, bonds, and lone pairs.

  • Molecular Geometry Steps

    1. Draw the Lewis structure.

    2. Determine electron-domain geometry.

    3. Identify molecular geometry using bonded atom arrangement.

  • Common Electron Domains

    • Linear: 2 atoms lead to linear geometry.

    • Trigonal Planar: 3 bonding domains (trigonal planar); 1 lone pair (bent).

    • Tetrahedral: 4 bonding domains (tetrahedral); 1 lone pair (trigonal pyramidal); 2 lone pairs (bent).

  • Influence of Lone Pairs and Multiple Bonds

    • Nonbonding pairs compress bond angles due to size.

    • Multiple bonds exert greater repulsion than single bonds.

  • Octet Rule

    • Some elements can expand beyond the octet, allowing 5 (trigonal bipyramidal) or 6 (octahedral) bonds.

  • Valence Bond Theory

    • Covalent bonds form through orbital overlaps.

    • Increased overlap leads to balance between repulsion and attraction.

  • Hybrid Orbitals

    • Formed from mixing atomic orbitals, resulting in degenerate orbitals.

    • Example: Be uses sp hybridization; C uses sp³ hybridization.

  • Types of Bonds

    • Sigma (σ) bonds: head-to-head overlap.

    • Pi (π) bonds: sideways overlap.

  • Molecular Orbital Theory

    • Considers wave properties and energy levels of electrons.

    • Formation of bonding and antibonding orbitals from atomic overlaps.

    • Bond order calculation indicates bond strength.

Molecular Orbital Characteristics:

-Max 2 electrons per orbital

-Electrons in the same orbital have opposite spin

  • Electrons in Bonding

    • Localized electrons are shared between two atoms; delocalized electrons are shared among multiple atoms.

      Paramagnetism and Diamagnetism

    • Paramagnetic: unpaired electrons

    • Diamagnetic: all paired.

  • MO Diagrams for Heteronuclear Molecules

    • Different energy levels affect bonding characteristics.

Overlap: Electrons of two atoms sharing the same space. This can occur in two types:

  • Sigma (σ) Bonds: Formed by head-to-head overlap of orbitals, allowing free rotation.

  • Pi (π) Bonds: Formed by sideways overlap of p orbitals, creating a bond that restricts rotation due to the presence of a nodal plane.

    • The sharing of space between 2 electrons of opposite spins creates a covalent bond

  • Double Bonds: Consist of one sigma bond and one pi bond, which enhances bond strength and affects molecular shape.

  • Triple Bonds: Comprise one sigma bond and two pi bonds, resulting in even greater bond strength and further restriction of rotational movement in the molecule.