TEST REVIEW

Differences Between Ionic and Molecular Compounds

  • Ionic Compounds:
    • Formed by the transfer of electrons between atoms, resulting in cations and anions.
    • Generally solid at room temperature with high melting and boiling points.
    • Conductive when melted or dissolved in water (electrolytes).
    • Segregated structure, hard and brittle.
  • Molecular Compounds:
    • Formed by the sharing of valence electrons.
    • Can be solid, liquid, or gas at room temperature with low melting and boiling points.
    • Non-conductive in solution (non-electrolytes).
    • Soft, waxy, flexible, and can have crystalline structures.

Bond Polarity

  • Polar Bonds:
    • Occur due to a difference in electronegativity (ΔEN) between two bonded atoms.
    • A bond is considered polar if ΔEN is approximately between 0.4 and 1.7.
  • Nonpolar Bonds:
    • Very similar electronegativities (ΔEN < 0.4).
    • No significant dipole moment present.
  • Electronegativity values:
    • Fluorine (F) = 4.0, Oxygen (O) = 3.5, Nitrogen (N) = 3.0, Carbon (C) = 2.5, Hydrogen (H) = 2.1.

Molecular Polarity

  • Determine Polarity:

    1. Identify polar bonds in the molecule.
    2. Analyze the overall shape of the molecule for symmetry:
    • Asymmetrical shape = polar molecule
    • Symmetrical shape = nonpolar molecule.

  • Example: Carbon Dioxide (CO₂)

    • Lewis Structure:

    O=C=O

    • Nonpolar because the symmetry of the structure causes the polar bonds to cancel each other's dipoles.

Miscibility of Oil and Water

  • Explanation:
    • Water is a polar molecule, while oil is nonpolar.
    • “Like dissolves like” means polar substances dissolve polar; nonpolar dissolves nonpolar.
    • Due to differences in polarity, oil and water do not mix (immiscible).

Melting and Boiling Points

  • Polar vs Nonpolar:
    • Polar molecules possess stronger intermolecular forces, leading to higher melting and boiling points.
    • Polar molecules, such as water (H₂O), have higher melting points than nonpolar molecules, such as methane (CH₄).

Intermolecular Forces

  • Types of Forces (from weakest to strongest):
    • London Dispersion Forces (weak attraction between all molecules).
    • Dipole-Dipole interactions (between opposite dipoles of polar molecules).
    • Hydrogen Bonding (strong dipole interaction between H and F, O, or N).
    • Ion-Dipole Forces (attraction between ions and polar molecules).

VSEPR Theory

  • Shapes of Molecules:
    • Based on electron pair repulsion (VSEPR) to predict 3D shapes of molecules:
    • Linear, Trigonal Planar, Bent, Tetrahedral, Trigonal Pyramidal, etc.
    • Determining Geometry:
    • Identify the central atom.
    • Count bonded atoms and lone pairs.
    • Establish the molecular geometry based on these counts.

Chemical and Physical Properties

  • Impact of Intermolecular Forces:
    • Stronger intermolecular forces result in higher melting and boiling points, higher surface tension, and different solubility levels.
  • Dissociation in Water:
    • Ionic compounds like NaCl dissolve in water and break into ions due to strong polarity of water, causing dissociation.

Practice Questions

  1. Identify Polar vs. Nonpolar:
    • Examples: CF₄ (nonpolar), C₂H₆ (nonpolar), CCl₃H (polar), CH₃OH (polar).
  2. Intermolecular Forces:
    • Differences between intermolecular forces and chemical bonds
  3. Predicting Boiling Points:
    • Compare C₂H₄ and C₆H₁₄ or C₂H₄ and C₂H₅OH based on polarity and intermolecular forces involved.