Chemistry Notes: Molecular Polarity and Intermolecular Forces

1. Molecular Polarity

  • Definition: Molecular polarity refers to the distribution of electrical charge over the atoms in a molecule, which is determined by the geometry of the molecule and the electronegativity of its constituent atoms.

2. Electronegativity

  • Definition: Electronegativity is a measure of how tightly an atom can attract and hold onto electrons in a covalent bond.
  • Example: In the case of chlorine (Cl) and hydrogen (H), Cl is more electronegative than H. Therefore:
    • Cl tends to gain electrons, becoming a negative ion (Cl⁻).
    • H tends to lose an electron, becoming a positive ion (H⁺).

3. Types of Bonds

3.1 Polar Covalent Bonds

  • A polar covalent bond is formed when two atoms with different electronegativities share electrons unequally, resulting in a molecular dipole.
  • Example: In a molecule like HCl:
    • Cl has a partial negative charge (S⁻) due to its higher electronegativity, and H has a partial positive charge (S⁺).

3.2 Non-Polar Covalent Bonds

  • Non-polar covalent bonds occur when two atoms share electrons equally, such as in a hydrogen (H₂) molecule.

4. Water: An Example of Polarity

  • Water (H₂O) exhibits significant polarity due to the difference in electronegativity between H and O.
  • Hydrogen Bonding: Water molecules can form up to four hydrogen bonds per molecule, contributing to its unique properties.

5. Intermolecular Forces

5.1 Dipole-Dipole Interactions

  • Dipole-dipole interactions occur between molecules that have permanent dipoles.
  • Examples:
    • H-O-H (water) exhibits dipole-dipole interactions due to its polar nature.
    • CH₃-O-CH₃ (dimethyl ether) also exhibits dipole-dipole interactions due to the presence of oxygen.

5.2 Hydrogen Bonds

  • Hydrogen bonds are a specific type of dipole-dipole interaction, characterized by a hydrogen atom attached to a highly electronegative atom (like O, N, or F).
  • Example: H-F is known for forming strong hydrogen bonds.
    • Notation: H-F:··· H-F represents a hydrogen bond.

5.3 Dispersion Forces

  • Dispersion forces (also called London dispersion forces) arise from temporary dipoles created when electrons move around in atoms or non-polar molecules.
  • They are generally weaker than dipole-dipole and hydrogen bonds, but they become stronger with increased atom size and electron number.

6. Comparative Strength of Intermolecular Forces

  • Ranking: The strength of intermolecular forces from strongest to weakest is:
    1. Hydrogen Bonds
    2. Dipole-Dipole Interactions
    3. Dispersion Forces

6.1 Summary of Bond Strengths

  • H-bonds > Dipole-Dipole > Dispersion Forces

7. Additional Notes

  • Polarity and Solubility: Polar molecules, such as water, are soluble in solvents that can also form hydrogen bonds, whereas non-polar molecules are not soluble in polar solvents.
  • If the central atom in a molecule has lone pairs of electrons and all outer atoms are the same, the molecule can be non-polar despite having polar bonds (a notable exception to the rule).
  • Examples of Molecular Geometry Determining Polarity:
    • A molecule like CO₂ is non-polar despite having polar covalent bonds due to its linear geometry.