Detailed Study Notes on Lewis Structures, Molecular Shapes, and Polarity

Objectives

  • Explore molecular structures and shapes using model kits.
  • Practice predicting molecular shapes via VSEPR theory and analyze molecular polarity.

Background

  • Covalent Bonds: Formed when non-metal atoms share valence electrons to achieve filled valence orbitals akin to noble gases.
    • Octet Rule: Most atoms strive for a total of eight valence electrons; exception is hydrogen (requires 2 electrons).

Lewis Structures

  • Definition: Representation of covalent molecules/ions showing valence electrons as shared (bond pairs) or unshared (lone pairs).
  • Representation of Bonds:
    • Single bond: represented by a short line (1 pair of electrons).
    • Double bond: represented by two short lines (2 pairs).
    • Triple bond: represented by three short lines (3 pairs).
    • Lone pairs: indicated by pairs of dots.
  • Example Representation:
    • For H-N-H: lone pair represented by dots, bond pairs by lines.

Molecular Shapes

  • Determined by the arrangement of atoms and their bond angles around the central atom, influenced by electron repulsion.
  • VSEPR Theory: States that electron pairs will arrange themselves to maximize distance from each other to minimize repulsion.
Common Molecular Shapes

  1. Linear

    • Description: 2 outer atoms, no lone pairs, bond angle 180°.

  2. Trigonal Planar

    • Description: 3 outer atoms in a flat plane, no lone pairs, bond angle 120°.

  3. Bent

    • Description: 2 outer atoms and 1 lone pair, bond angle slightly less than 120°.

  4. Tetrahedral

    • Description: 4 outer atoms, no lone pairs, bond angles 109.5°.

  5. Trigonal Pyramidal

    • Description: 3 outer atoms and 1 lone pair, bond angle slightly less than 109.5°.

  6. Bent (with 2 lone pairs)

    • Description: 2 outer atoms, 2 lone pairs, bond angle slightly less than 109.5°.

Electronegativity and Bond Polarity

  • Definition: The tendency of an atom to attract shared electrons.
    • Increases across a period, decreases down a group.
  • Polar Covalent Bonds: Form when bonded atoms have different electronegativities leading to unequal sharing of electrons:
    • Atom with high electronegativity acquires a partial negative charge (-), and the other atom acquires a partial positive charge (+).
Example of Polar Covalent Bond
  • H-Cl → Electrons drawn closer to Cl leading to polarization.

Molecular Polarity

  • Definition: Unbalanced distribution of electrons across a molecule, often leading to a net dipole moment.
  • Criteria for Polarity:
    • Molecule is typically polar if it has polar bonds arranged asymmetrically around the central atom.
    • Non-polar if symmetrical distribution or all bonds are non-polar.
Examples of Molecular Polarity
  1. Non-Polar: Symmetrical distribution of electrons.
  2. Polar: Asymmetrical distribution leading to a net dipole moment.

Experimental Procedure for Drawing Lewis Structures

  1. Count Valence Electrons: Add or subtract electrons according to charge on ions.
  2. Sketch the Structure: Start with the least electronegative atom as the central atom (except for Hydrogen).
  3. Obey the Octet Rule: Distribute remaining electrons to satisfy the octet rule.
    • Use double or triple bonds if necessary to fulfill octet.

Constructing 3D Models and Determining Polarity

  1. Use a molecular model kit to build structures, following bond representation rules (short sticks for single bonds, etc.).
  2. Assess polarity based on bond characteristics and overall symmetry.

Rules for Constructing Molecules

  • Use short sticks for single bonds.
  • Use two long flexible sticks for double bonds.
  • Use three long flexible sticks for triple bonds.

Experimental Data Example

  • Molecular Formula: H₂, O₂, N₂, H₂O, NH₃, CO₂.
  • Populate Lewis structures and assess molecular polarity.