Intermolecular Forces

Core Concept

Intermolecular forces (IMFs) are the relatively weak attractions that act between separate molecules, ions, or atoms. They are entirely distinct from the covalent or ionic bonds that hold atoms together within a single molecule or formula unit. IMFs are what make a substance a liquid or solid at a given temperature: the stronger the IMFs, the more thermal energy is needed to pull particles apart, so the higher the melting and boiling points. The four main kinds of IMFs, from weakest to strongest, are London dispersion forces, dipole–dipole forces, hydrogen bonding, and ion–dipole forces.

Key Ideas

Intramolecular vs. intermolecular

• Intramolecular bonds (covalent, ionic, metallic) hold atoms together inside one molecule or crystal. Covalent bond strengths are typically 150–1000+ kJ/mol.
• Intermolecular forces act between separate particles and are typically 1–100 kJ/mol — about an order of magnitude weaker. Boiling a liquid breaks IMFs, not covalent bonds.

The four IMFs (weakest → strongest)

• London dispersion forces — present in all substances; the only IMF in nonpolar molecules and noble-gas atoms.
• Dipole–dipole forces — between polar molecules with permanent dipoles.
• Hydrogen bonding — a special, stronger dipole–dipole interaction where H is bonded to N, O, or F.
• Ion–dipole forces — between an ion and a polar molecule (e.g., Na⁺ dissolved in water).

What drives each force

• Dispersion: instantaneous, fluctuating electron clouds create momentary dipoles that induce dipoles in neighbors.
• Polarizability: how easily an electron cloud is distorted. Larger atoms/molecules with more electrons are more polarizable → stronger dispersion.
• Dipole–dipole: permanent partial charges (δ+, δ−) on polar molecules attract opposite ends.
• H-bond: an H atom bonded to a highly electronegative N, O, or F carries a large δ+, which is strongly attracted to a lone pair on N/O/F of a neighbor.

Equations and Variables

• Coulomb's law (governs ion–ion and ion–dipole attractions): F = k·q₁q₂/r², where F = force, q = charge magnitudes, r = separation, k = Coulomb constant. Ion–dipole strength scales with ion charge and dipole magnitude, and weakens as r².
• Dipole moment: μ = Q·r, where μ = dipole moment (debye, D), Q = magnitude of separated charge, r = distance between charges. A molecule is polar when μ ≠ 0.
• Dispersion energy (qualitative): E ∝ −α/r⁶, where α = polarizability. Bigger α (larger, more electron-rich particles) → stronger dispersion attraction.
• No single equation predicts boiling point from IMFs; rankings are made by comparing IMF types first, then size/polarizability within a type.

How It Works

1. Establish the electron distribution. A molecule may have a permanent dipole (polar) or not (nonpolar). All species, even noble gases, have moving electrons.
2. Dispersion forms everywhere. At any instant, electron density fluctuates, creating a transient dipole; this dipole distorts the electron cloud of a neighbor (induced dipole), producing a weak attraction that flickers in and out.
3. Permanent dipoles align. In polar molecules, the δ+ end of one molecule is attracted to the δ− end of another, adding dipole–dipole attraction on top of dispersion.
4. Hydrogen bonding appears when H–N/O/F is present. Because N, O, and F are very electronegative and small, an H bonded to them is nearly "bare," so its δ+ is exceptionally strong and is drawn to lone pairs on neighboring N/O/F atoms.
5. Ions interact with dipoles. In a salt solution, the ion's full charge pulls strongly on the partial charges of polar solvent molecules (ion–dipole), which is why ionic compounds dissolve in water.
6. Stronger total IMFs → higher boiling/melting point. More energy is required to overcome stronger attractions.

Worked Example

Rank ethane (C₂H₆), fluoromethane (CH₃F), and methanol (CH₃OH) by increasing boiling point.

1. Identify IMFs. C₂H₆ is nonpolar → dispersion only. CH₃F is polar → dispersion + dipole–dipole. CH₃OH has an O–H bond → dispersion + dipole–dipole + hydrogen bonding.
2. Rank. Dispersion-only (C₂H₆) is weakest; adding dipole–dipole (CH₃F) is stronger; hydrogen bonding (CH₃OH) is strongest.
3. Result. C₂H₆ (−89 °C) < CH₃F (−78 °C) < CH₃OH (65 °C). The order matches the measured boiling points. Within a single IMF type (e.g., dispersion only), the larger, more polarizable molecule boils higher.

How it works

1. Establish the electron distribution. A molecule may have a permanent dipole (polar) or not (nonpolar). All species, even noble gases, have moving electrons.
2. Dispersion forms everywhere. At any instant, electron density fluctuates, creating a transient dipole; this dipole distorts the electron cloud of a neighbor (induced dipole), producing a weak attraction that flickers in and out.
3. Permanent dipoles align. In polar molecules, the δ+ end of one molecule is attracted to the δ− end of another, adding dipole–dipole attraction on top of dispersion.
4. Hydrogen bonding appears when H–N/O/F is present. Because N, O, and F are very electronegative and small, an H bonded to them is nearly "bare," so its δ+ is exceptionally strong and is drawn to lone pairs on neighboring N/O/F atoms.
5. Ions interact with dipoles. In a salt solution, the ion's full charge pulls strongly on the partial charges of polar solvent molecules (ion–dipole), which is why ionic compounds dissolve in water.
6. Stronger total IMFs → higher boiling/melting point. More energy is required to overcome stronger attractions.

Common confusions

• "A hydrogen bond is a covalent bond to hydrogen." — Wrong. The O–H bond within a water molecule is covalent; the hydrogen bond is the attraction between the H of one molecule and the O of another.
• "Only nonpolar molecules have London dispersion forces." — Wrong. Dispersion forces exist in every substance; they are simply the only IMF in nonpolar species.
• "Dipole–dipole forces are stronger than all dispersion forces." — Wrong. A very large nonpolar molecule (e.g., I₂) can have dispersion forces stronger than the dipole–dipole forces of a tiny polar molecule.
• "IMF strength doesn't matter for solids." — Wrong. Molecular solids are held together entirely by IMFs, so IMFs set their melting points.
• "HF is a stronger H-bond donor than water because F is more electronegative." — Partly true (F is more electronegative), but water forms more H-bonds per molecule, so water's overall network is stronger. Rank IMF type, then count.

Quick review

• Intramolecular bonds are within molecules; IMFs are between particles.
• IMF types (weak→strong): London dispersion, dipole–dipole, H-bond, ion–dipole.
• Dispersion increases with size/polarizability; H-bond needs H–N/O/F.
• More electrons → more polarizable → stronger dispersion.
• Stronger IMFs → higher bp/mp, lower vapor pressure.
• Worked ranking: C₂H₆ < CH₃F < CH₃OH (boiling point).