Chemistry · Unit 3: Intermolecular Forces and Properties · 14 min read · Updated 2026-05-11
Intermolecular Forces — AP Chemistry
AP Chemistry · Unit 3: Intermolecular Forces and Properties · 14 min read
1. What Are Intermolecular Forces?★★☆☆☆⏱ 2 min
Intermolecular forces (abbreviated IMFs) are electrostatic attractive forces that act between separate discrete molecules, monatomic atoms (like noble gases), or ions. They are distinct from intramolecular bonds (covalent, ionic, metallic) that hold atoms together within a single chemical unit.
All intermolecular attractions arise from Coulombic interactions between partial or full charges, and are universally weaker than intramolecular chemical bonds: typical IMFs range from 1 to 100 kJ/mol, while covalent bonds exceed 150 kJ/mol. This difference means phase changes (melting, boiling) only overcome IMFs, not break intramolecular bonds.
2. London Dispersion Forces★★☆☆☆⏱ 3 min
The strength of LDF depends on polarizability, or how easily the electron cloud is distorted. Polarizability increases with total number of electrons (and thus molar mass). For molecules of the same molar mass, LDF strength also depends on molecular shape: linear, unbranched molecules have more surface area for intermolecular contact than spherical branched isomers, leading to stronger LDF.
V \propto -\frac{\alpha_1 \alpha_2}{r^6}
Exam tip: Always check for same-molar-mass isomers first when comparing LDF strength — if molar mass is identical, the difference comes from surface area, not another type of IMF.
3. Dipole-Dipole Interactions and Hydrogen Bonding★★★☆☆⏱ 4 min
The potential energy of dipole-dipole attraction follows the relationship below, where $\mu$ is the permanent dipole moment of each molecule, $k$ is Boltzmann’s constant, and $T$ is absolute temperature. Larger dipole moments give stronger attraction, and higher temperature disrupts dipole alignment, weakening net attraction.
V \propto -\frac{\mu_1^2 \mu_2^2}{r^6 kT}
Hydrogen bonds have strengths of 20-40 kJ/mol, compared to 1-5 kJ/mol for regular dipole-dipole interactions.
Exam tip: For hydrogen bonding between two different molecules (e.g. ethanol and water), you only need a donor H (bonded to N/O/F) on one molecule and an acceptor N/O/F on the other — both do not need a donor H.
4. Ion-Dipole Forces★★★☆☆⏱ 3 min
Ion-dipole forces range up to ~100 kJ/mol, comparable to weak intramolecular bonds. The potential energy of ion-dipole attraction follows the relationship below, where $z$ is the charge of the ion, $\mu$ is the dipole moment of the polar molecule, and $r$ is the distance between the ion and dipole center. Higher ion charge and larger dipole moment give stronger attraction.
V \propto -\frac{z \mu}{r^2}
Because attraction follows an inverse $r^2$ relationship (instead of $r^6$ for neutral IMFs), ion-dipole attraction falls off much more slowly with distance, leading to a stronger net interaction.
Exam tip: When asked to explain why an ionic compound dissolves in water, always explicitly name ion-dipole forces between ions and water molecules as the key stabilizing interaction, do not only write "like dissolves like".
5. AP-Style Worked Practice Questions★★★★☆⏱ 4 min
Common Pitfalls
Why: Students confuse branching with higher polarity, forgetting branching reduces surface area contact between molecules.
Why: Organic chemistry courses discuss intramolecular hydrogen bonding, leading students to confuse it with a full chemical bond.
Why: Students memorize that N/O/F are required for hydrogen bonding and forget the requirement that H must be covalently bonded directly to N/O/F.
Why: Students memorize the "hydrogen bonding > dipole-dipole > LDF" rule and forget it only applies to similar molar mass.
Why: Students see F and automatically assume hydrogen bonding, ignoring the bonding of H.