Topic 5: Chemical Bonding II

  • Foundation in Science - Chemistry 1

Module Overview

  • Topic: Chemical Bonding II

  • Prepared by: Ms. Shasha

Learning Outcomes

  • Explain molecular shapes and bond angles using VSEPR theory with examples:

    • CO2: Linear

    • CH4: Tetrahedral

    • BF3: Trigonal Planar

    • NH3: Trigonal Pyramidal

    • H2O: Bent/V-shape

    • SF6: Octahedral

  • Predict molecule shapes and bond angles

  • Determine molecular polarity based on geometry

Important Theories

Lewis Structure

  • Represents molecule compositions & covalent bonds but not shapes.

VSEPR Theory

  • Acronym: Valence Shell Electron Pair Repulsion

  • States that electron group repulsion determines molecular shapes.

  • Types of Electron Groups: Lone pairs, bonded pairs, single electrons.

  • Strength of Repulsion Order:

    • Lone pair - Lone pair > Lone pair - Bonded pair > Bonded pair - Bonded pair

  • Optimal separation for minimum energy leads to molecular shape formation.

Predicting Molecular Geometry (Steps)

  1. Draw Lewis structure

  2. Count electron groups around the central atom (including lone pairs, single/multiple bonds)

  3. Identify bond angles based on electron pair interactions.

  4. Deduce molecular shape.

Common Molecular Geometries

  • 2 electron pairs: Linear

  • 3 electron pairs: Trigonal Planar

  • 4 electron pairs: Tetrahedral

  • 5 electron pairs: Trigonal Bipyramidal

  • 6 electron pairs: Octahedral

Molecular Examples and Shapes

  • CO2 (Linear): 2 electron groups, 180° bond angle

  • BF3 (Trigonal Planar): 3 electron groups, 120° bond angle

  • H2O (Bent): 4 electron groups, 104.5° bond angle

  • CH4 (Tetrahedral): 4 electron groups, 109.5° bond angle

  • NH3 (Trigonal Pyramidal): 4 electron groups, < 109.5° bond angle

  • SF6 (Octahedral): 6 electron groups, 90° bond angle

Polarity of Molecules

  • Factors Affecting Polarity:

    • Molecular shape

    • Electronegativities of atoms involved

  • Higher electronegativity difference correlates with increased polarity.

Dipole and Dipole Moment

  • Polar Molecules: Exhibit both high (δ-) and low (δ+) electron density areas.

  • Dipole Moment (μ): A vector quantity indicating molecular polarity.

  • Non-polar molecules have dipoles that cancel each other out.

Bond Polarity

  • Polar bonds exist in polar molecules due to unequal electron sharing.

  • Electronegativities:

    • Increases left to right across the periodic table.

    • Example trend: F > Cl > Br > I

Practical Examples of Polar and Non-Polar Molecules

  • BrCl (Linear): Polar (Cl δ-, Br δ+)

  • SO2 (Bent): Polar (O δ-, S δ+; dipole moments do not cancel)

  • NH3 (Trigonal Pyramidal): Polar due to lone pair causing asymmetry

  • CCl4 (Tetrahedral): Non-polar though C-Cl bonds are polar; symmetrical cancellation of dipole moments.

  • BF3 (Trigonal Planar): Non-polar due to symmetrical cancellation of dipole moments.

  • SF6 (Octahedral): Non-polar for the same reason as above.

Applications

  • Atorvastatin (Lipitor): Lowers cholesterol by fitting precisely into enzyme cavities; effectiveness altered by molecular shape/size modifications.