MCAT topics → Nature of Molecules and Intermolecular Interactions → Intermolecular Forces (London, Dipole-Dipole, H-Bonding)

Intermolecular Forces (London, Dipole-Dipole, H-Bonding)

MCAT trap: Thinks any C–H or S–H bond can participate in hydrogen bonding. Hydrogen bonding requires H bonded directly to N, O, or F; the H must also interact with a lone pair on N, O, or F of another molecule.

Intermolecular forces (IMFs) are the attractions between separate molecules — not the covalent bonds within them — and the MCAT uses this distinction constantly to test whether you know which forces control boiling point, solubility, and vapor pressure. The most common error: thinking London dispersion forces only exist in nonpolar molecules. Every molecule has them. Polarity just adds dipole-dipole or hydrogen bonding on top. The four types you need cold are London dispersion (all molecules), dipole-dipole (polar molecules), hydrogen bonding (H directly on N, O, or F), and ion-dipole (ions in polar solvents).

The MCAT tests IMFs from multiple angles. At the recall level, you need to identify which forces are present in a given molecule. At the application level, you'll rank a series of molecules by boiling point or explain why water has anomalously high surface tension. In passage-based questions, you'll see biological contexts — DNA base pairing, protein secondary structure, enzyme-substrate binding — where you're expected to recognize that hydrogen bonds are doing the work, not covalent bonds. The cross-disciplinary angle is especially high-yield: if a passage describes a protein denaturing or DNA strands separating, the MCAT expects you to immediately think 'hydrogen bonds being disrupted.'

What makes this topic tricky is that students carry in several stubborn misconceptions. Many think hydrogen bonding applies to any molecule with an H atom — wrong, C–H and S–H bonds don't qualify. Others rank boiling points by molecular weight alone and get burned when a tiny hydrogen-bonding molecule outperforms a heavier nonpolar one. A third common error is believing London forces only exist in nonpolar molecules, when in reality every molecule has them — polarity just adds additional forces on top. Nail the logic of each IMF type and these traps disappear.

Common misconceptions

Common mistake
Wrong: Any molecule containing hydrogen can form hydrogen bonds.
Right: Hydrogen bonding requires H bonded directly to N, O, or F; the H must also interact with a lone pair on N, O, or F of another molecule.
Hydrogen bonding is not just about having a hydrogen atom — the geometry and electronics have to be right. The H must be covalently bonded to N, O, or F (highly electronegative atoms that make the H strongly partial-positive), and it must then interact with a lone pair on another N, O, or F. A C–H bond doesn't polarize the hydrogen enough, and S–H doesn't qualify because sulfur is too large and not electronegative enough to create the necessary partial charge. So CH₄ and H₂S cannot hydrogen bond, even though both contain hydrogen.
Common mistake
Wrong: London dispersion forces only exist in nonpolar molecules.
Right: London dispersion forces exist in all molecules; they are simply the dominant IMF in nonpolar molecules where dipole-dipole and H-bonding are absent.
London dispersion forces arise from instantaneous fluctuations in electron density that create temporary dipoles — and every molecule has electrons, so every molecule experiences London forces. In polar molecules, dipole-dipole interactions and London forces both contribute to total IMF strength. The reason we say London forces 'dominate' in nonpolar molecules is that they're the only IMF present, not because polar molecules lack them. Forgetting this leads to underestimating total IMF strength in polar or large molecules.
Common mistake
Wrong: The molecule with the highest molecular weight always has the highest boiling point.
Right: Boiling point depends on IMF strength; a smaller molecule capable of hydrogen bonding (e.g., H₂O) can have a higher boiling point than a larger nonpolar molecule with only London forces.
Molecular weight correlates with boiling point only when you're comparing molecules of the same IMF type, because bigger molecules have more electrons and stronger London forces. The moment you introduce a molecule capable of hydrogen bonding, the IMF type changes and molecular weight becomes a secondary factor. Water (MW 18) boils at 100°C while hexane (MW 86) boils at 69°C — water wins because its hydrogen bonds are far stronger than hexane's London forces. Always identify IMF type first, then use size as a tiebreaker within that type.
Common mistake
Wrong: The base pairs in DNA are held together by covalent bonds.
Right: Complementary base pairs in DNA are held together by hydrogen bonds (2 for A-T, 3 for G-C), which allow strand separation during replication.
The bonds within each DNA strand — between nucleotides — are covalent (phosphodiester bonds), but the attractions holding the two complementary strands together are hydrogen bonds between base pairs. A-T pairs share 2 hydrogen bonds; G-C pairs share 3, which is why G-C-rich regions are harder to denature. The fact that these are hydrogen bonds (not covalent) is biologically essential: it allows helicase to unzip the strands during replication without breaking the backbone. If base pairs were covalently bonded, DNA replication as we know it couldn't happen.
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What the exam tests

  1. Identify the correct IMF type(s) present in a given molecule — London dispersion (all molecules), dipole-dipole (polar molecules), hydrogen bonding (H bonded to N, O, or F), and ion-dipole (ions near polar molecules) — and know their relative strengths.
  2. Explain mechanistically how stronger IMFs produce higher boiling points, higher melting points, and lower vapor pressures, because more energy is required to overcome intermolecular attractions.
  3. Given a set of molecules differing in size, polarity, and hydrogen-bonding capability, rank them by expected boiling point — requiring you to weigh IMF type against molecular size simultaneously.
  4. Apply hydrogen bonding to biological systems: base-pair specificity and strand separation in DNA, alpha-helix and beta-sheet stabilization in proteins, and the unique properties of water; apply ion-dipole forces to explain how ionic solutes dissolve in water.

Can you avoid these mistakes?

Rank the following molecules from lowest to highest boiling point and explain your reasoning: CH₄, NH₃, N₂, HF. What IMF types are present in each?
A student claims that because HCl is polar and Ar is nonpolar, HCl must have stronger total intermolecular forces than Ar regardless of molecular size. Is this always true? At what molecular weight might a noble gas or nonpolar molecule actually have a higher boiling point than HCl?
In a protein alpha-helix, what type of bond stabilizes the helical structure, and between which atoms does it form? What would happen to that structure if you raised the temperature significantly or added a strong denaturant?
A passage describes a newly synthesized molecule with the formula (CH₃)₂S–H. A student predicts it will hydrogen bond with water and dissolve readily. Identify the flaw in this prediction and explain what IMF the S–H group actually contributes.

Related topics

Bond and Molecular Polarity, Dipole Moments Ionic and Covalent Bonds VSEPR and Molecular Geometry Hybridization (sp, sp2, sp3) and Sigma/Pi Bonds

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