Intermolecular Forces
Imagine three ways to stick two surfaces together. Velcro is quick and easy to pull apart - it works through lots of tiny, weak hooks that grab temporarily. Glue is stronger and more directional - you have to apply it in the right spots and it holds with moderate force. Industrial-strength adhesive is the strongest of the three - it creates a bond so tight that you might tear the material before the adhesive gives way. Now imagine a fourth option: welding. That fuses the materials into one piece, and it operates on a completely different level of strength.
Intermolecular forces work the same way. They are the forces between molecules, and they come in a range of strengths. But they are all dramatically weaker than the intramolecular bonds (covalent, ionic) that hold atoms together within a molecule. The MCAT tests both the types of intermolecular forces and how they affect physical properties, so you need to understand the hierarchy.
Intramolecular vs. Intermolecular: The Critical Distinction
Before diving into the types, be absolutely clear on this distinction. The MCAT will try to confuse you.
Intramolecular forces are the bonds within a molecule - covalent bonds, ionic bonds. These hold atoms together to form molecules or formula units. They are strong (hundreds of kJ/mol).
Intermolecular forces (IMFs) are the attractions between separate molecules. These are much weaker (typically 1-40 kJ/mol). When water boils, you are breaking intermolecular forces (hydrogen bonds between water molecules), not intramolecular forces (the O-H covalent bonds within each water molecule). The water molecules themselves remain intact.
Type 1: London Dispersion Forces (LDF)
London dispersion forces (also called van der Waals forces or induced dipole forces) are present in every molecule - polar or nonpolar, large or small. They arise from temporary, instantaneous dipoles caused by the random motion of electrons.
At any given instant, the electrons in a molecule might be slightly more concentrated on one side than the other, creating a fleeting dipole. This temporary dipole can induce a complementary dipole in a neighboring molecule, and the two temporary dipoles attract each other. A fraction of a second later, the electrons shift and the dipole disappears, but a new one forms immediately. These flickering attractions add up.
Three factors increase the strength of London dispersion forces:
- Molecular weight: More electrons means larger temporary dipoles.
- Surface area: More contact area between molecules means more opportunities for induced dipoles. Long, straight-chain molecules have stronger LDFs than compact, spherical molecules of the same molecular weight.
- Polarizability: How easily the electron cloud can be distorted. Larger atoms with more loosely held electrons are more polarizable.
Type 2: Dipole-Dipole Interactions
Dipole-dipole interactions occur between polar molecules that have permanent dipole moments. The positive end (Ξ΄+) of one molecule is attracted to the negative end (Ξ΄-) of a neighboring molecule.
Unlike London dispersion forces, dipole-dipole interactions are directional and persistent. They do not flicker in and out of existence. As long as the molecules have a permanent dipole, the attraction is always there.
The strength of dipole-dipole interactions depends on the magnitude of the dipole moment. Molecules with larger electronegativity differences between bonded atoms and asymmetric geometries will have stronger dipole-dipole forces.
Examples: HCl, acetone (CH3COCH3), formaldehyde (CH2O). These molecules all have permanent dipoles and align so that opposite partial charges face each other.
Note that polar molecules also experience London dispersion forces in addition to dipole-dipole interactions. IMFs are cumulative - you add all applicable types together.
Type 3: Hydrogen Bonding
Hydrogen bonding is a special, extra-strong type of dipole-dipole interaction. It occurs when a hydrogen atom bonded directly to F, O, or N interacts with a lone pair on another F, O, or N atom.
Why is hydrogen bonding so much stronger than regular dipole-dipole forces? Two reasons. First, F, O, and N are the three most electronegative small atoms, so the dipole on the H-F, H-O, or H-N bond is very large. Second, hydrogen is tiny - it has no inner electron shells to shield its nucleus, so a nearby lone pair can get extremely close to the partially positive hydrogen, creating an unusually strong electrostatic attraction.
Hydrogen bonding has enormous biological importance:
- Waterβs properties: The high boiling point, high specific heat, high heat of vaporization, and surface tension of water are all due to extensive hydrogen bonding between water molecules.
- DNA base pairing: The two strands of the double helix are held together by hydrogen bonds between complementary bases (A-T has 2 H-bonds, G-C has 3 H-bonds).
- Protein structure: Alpha helices and beta sheets in protein secondary structure are stabilized by hydrogen bonds between backbone N-H and C=O groups.
Why Waterβs Boiling Point is So High
Water (H2O, MW = 18 g/mol) boils at 100 degrees C. Compare that to H2S (MW = 34 g/mol), which boils at -60 degrees C. H2S is almost twice as heavy, yet boils 160 degrees lower. Why?
Intermolecular forces, weighed against real boiling points
Scroll sideways to see the whole map.
Hydrogen bonding has a narrow definitionIt needs hydrogen bonded directly to nitrogen, oxygen or fluorine, and a lone pair on a neighbouring N, O or F to accept it. The HβC bonds in methane do not qualify however many of them there are, which is why methane boils at β161 Β°C.
Why dispersion still wins sometimesDispersion is the weakest force per contact, but it scales with surface area and polarisability. Enough of it beats a stronger force used once: octane boils at 126 Β°C on dispersion alone, well above polar HβS at β60 Β°C.
What this predictsStronger intermolecular forces mean higher boiling point, higher melting point, higher viscosity, higher surface tension and lower vapour pressure. They are five ways of asking the same question, so answer them all from the same ladder.
Water forms extensive hydrogen bonds - each molecule can participate in up to four hydrogen bonds simultaneously (two from its two O-H bonds donating, two from its two lone pairs accepting). H2S cannot form hydrogen bonds because sulfur is not electronegative enough. It relies only on dipole-dipole and London dispersion forces, which are far weaker.
This is why water is a liquid at room temperature while similarly sized molecules like methane (CH4, bp = -161 degrees C) and ammonia (NH3, bp = -33 degrees C) are gases. Hydrogen bonding is the key.
Comparison Table: Intermolecular Forces
| Force | Present In | Strength | Example |
|:---|:---|:---:|:---|
| London dispersion | All molecules | Weakest (0.05 - 40 kJ/mol) | Ar, CH4, I2 |
| Dipole-dipole | Polar molecules | Moderate (5 - 25 kJ/mol) | HCl, acetone |
| Hydrogen bonding | Molecules with H-F, H-O, or H-N | Strong (10 - 40 kJ/mol) | H2O, NH3, HF |
| Ion-dipole | Ions in polar solvents | Strongest IMF (50+ kJ/mol) | NaCl in water |
Note: Ion-dipole forces are covered in the next section but included here for the complete ranking. Within similar-sized molecules, the ranking above holds. However, for very large molecules, London dispersion forces can rival or exceed hydrogen bonding in magnitude.
IMFs and Physical Properties
Stronger intermolecular forces make it harder to separate molecules from each other. This has predictable effects on every physical property the MCAT might test:
- Boiling point: Stronger IMFs β higher boiling point (more energy needed to pull molecules apart into the gas phase)
- Melting point: Stronger IMFs β higher melting point
- Surface tension: Stronger IMFs β higher surface tension (molecules at the surface are pulled inward more strongly)
- Viscosity: Stronger IMFs β higher viscosity (molecules resist flowing past each other)
- Vapor pressure: Stronger IMFs β lower vapor pressure (fewer molecules have enough energy to escape into the gas phase)
Predicting Relative Boiling Points
On the MCAT, you will frequently be asked to rank molecules by boiling point. Here is the decision framework:
- Identify the strongest IMF each molecule can form. Hydrogen bonding beats dipole-dipole beats LDF (for similar-sized molecules).
- If two molecules have the same type of strongest IMF, compare molecular weight. Higher MW means stronger LDFs, so higher boiling point.
- If molecular weights are similar, compare surface area. Longer, more extended molecules have more surface contact and stronger LDFs than compact, branched molecules.
Example: Rank the boiling points of CH4, CH3OH, and CH3CH2CH2CH3 (butane).
- CH4: nonpolar, LDF only, very low MW. Lowest boiling point.
- Butane: nonpolar, LDF only, but MW = 58 (higher than CH4). Moderate boiling point.
- CH3OH (methanol): polar, can form hydrogen bonds (O-H group). Highest boiling point despite having a lower MW than butane.
Result: CH4 < butane < CH3OH.
Common MCAT Traps
- Saying nonpolar molecules have βno intermolecular forces.β Wrong. All molecules have London dispersion forces. Nonpolar molecules just lack dipole-dipole and hydrogen bonding.
- Calling any O-H interaction a hydrogen bond. The hydrogen must be bonded directly to F, O, or N. An O-H group on one molecule hydrogen bonds with a lone pair on F, O, or N of another molecule. A C-H bond next to an oxygen does not count.
- Ignoring London dispersion forces in polar molecules. Polar molecules have dipole-dipole (and possibly H-bonding) forces in addition to LDFs. All three types are cumulative.
- Confusing boiling with bond breaking. Boiling breaks intermolecular forces, not covalent bonds. When water boils, the O-H bonds remain intact. You get individual H2O molecules in the gas phase, not separated atoms.
London dispersion forces < dipole-dipole interactions < hydrogen bonding. LDFs are present in all molecules and arise from temporary dipoles. Dipole-dipole forces occur between polar molecules with permanent dipoles. Hydrogen bonding is a special, extra-strong dipole-dipole force requiring H bonded to F, O, or N interacting with a lone pair on F, O, or N.
F, O, and N. Remember βFON - Phone a Friend for H-bonds.β These are the three most electronegative small atoms. Hydrogen bonding requires H bonded directly to one of these atoms (donor), interacting with a lone pair on another F, O, or N atom (acceptor). H bonded to Cl or S does not form hydrogen bonds because those atoms are too large and not electronegative enough.
Water forms extensive hydrogen bonds; H2S cannot. Oxygen is electronegative enough (and small enough) to participate in hydrogen bonding, so each water molecule can form up to four hydrogen bonds with its neighbors. Sulfur is less electronegative and too large, so H2S relies only on weaker dipole-dipole and London dispersion forces. The far stronger hydrogen bonding network in water requires much more energy to disrupt, resulting in a boiling point over 160 degrees higher.
Stronger IMFs lead to lower vapor pressure. Vapor pressure measures how easily molecules escape from the liquid phase into the gas phase. When intermolecular forces are strong, molecules are held tightly in the liquid and fewer have enough kinetic energy to break free. The result is fewer molecules in the gas phase above the liquid, meaning lower vapor pressure. This is the inverse of the boiling point trend - substances with high boiling points have low vapor pressures.