Atoms within a molecule are held together by intramolecular forces (ionic, covalent, metallic bonds). In contrast, intermolecular forces are attractive forces between separate molecules or ions and are generally much weaker than intramolecular bonds. These forces determine many macroscopic physical properties: boiling point, melting point, viscosity, vapor pressure, surface tension, solubility, etc.

1. London Dispersion Forces (LDF) (also called van der Waals dispersion forces):
2. Present in all molecules and atoms (including noble gases).
3. Arise from temporary fluctuations in electron density that create instantaneous dipoles, which induce dipoles in neighboring particles.
4. Magnitude depends on:
   • Polarizability: larger, more electrons, and more diffuse electron clouds → more polarizable → stronger LDF.
   • Surface area: molecules with greater surface contact (e.g., long carbon chains) exhibit stronger dispersion forces.
5. Examples:
   • Noble gases: He (liquid at –269 °C), Ne (liquid at –246 °C), Ar (liquid at –186 °C), Kr (liquid at –153 °C), Xe (liquid at –111 °C). As atomic size and polarizability increase, boiling point increases.
   • Hydrocarbons: n-octane (C₈H₁₈) has a higher boiling point than n-butane (C₄H₁₀) because of larger size and surface area.

1. Dipole–Dipole Interactions:
2. Occur between permanently polar molecules. The positive end (δ⁺) of one dipole attracts the negative end (δ⁻) of another.
3. Strength depends on:
   • Magnitude of the molecular dipole moment (larger dipole → stronger interaction).
   • Orientation: head-to-tail alignment maximizes attraction.
4. Examples:
   • Hydrogen chloride (HCl) molecules: δ⁺H—Clδ⁻ align so Hδ⁺ of one HCl is near Clδ⁻ of another.
   • Acetone (CH₃COCH₃) is polar (C=O dipole); acetone molecules exhibit dipole–dipole attractions.

1. Hydrogen Bonding (special case of dipole–dipole):
2. A particularly strong intermolecular interaction when H is covalently bonded to highly electronegative atoms (F, O, or N), and this hydrogen interacts with a lone pair on F, O, or N in a neighboring molecule.
3. Criteria for hydrogen bonding:
   • Molecule A must have H bonded to F, O, or N (A–H; A = F, O, or N).
   • Molecule B must have a lone pair on F, O, or N.
   • The H⋯B (lone-pair-bearing atom) distance is significantly shorter than the sum of their van der Waals radii (i.e., it is a true “bond”).
4. Examples:
   • Water (H₂O): each H (bound to O) can hydrogen-bond to a lone pair on O of another water molecule → extensive hydrogen-bond network → high boiling point (100 °C) relative to molecular mass.

1. Ion–Dipole Interactions:
2. Occur between an ion (cation or anion) and a polar molecule (dipole).
3. Strength depends on:
   • Charge of the ion (higher charge → stronger interaction).
   • Magnitude of dipole moment (greater δ⁺/δ⁻ separation → stronger).
   • Distance between ion and dipole (Coulomb’s law; smaller distance → stronger).
4. Examples:
   • Dissolution of NaCl in water: Na⁺ interacts with the partial negative charge on O; Cl⁻ interacts with partial positive charges on H. This ion–dipole stabilization allows NaCl to dissolve.

4.5.2 Relative Strengths of Intermolecular Forces
In order of increasing typical strength (kJ/mol):
1. London dispersion forces (≈ 0.05–40 kJ/mol depending on size and polarizability)
2. Dipole–dipole interactions (≈ 3–10 kJ/mol, can be larger in highly polar molecules)
3. Hydrogen bonds (≈ 15–40 kJ/mol, sometimes up to 60 kJ/mol in strong cases)
4. Ion–dipole interactions (≈ 50–100 kJ/mol)
5. Ionic bonds (intramolecular for ionic solids; > 400 kJ/mol in lattice energy)

4.5.3 Influence on Physical Properties
- Boiling points and melting points:
- Substances with stronger intermolecular forces require more energy (heat) to separate molecules → higher melting/boiling points.
- Examples:
- Compare boiling points of Group 17 hydrides: HF (19.5 °C), HCl (–85.0 °C), HBr (–66.7 °C), HI (–35.4 °C). HF has a high boiling point due to strong hydrogen bonding; the others rely mainly on dispersion forces and are far lower.