AP Chem Intermolecular Forces: Explained

Intermolecular forces (IMFs) are the attractions between particles, and in AP Chemistry they explain physical properties like boiling point, vapor pressure, viscosity, and solubility. The main types are London dispersion forces, dipole-dipole forces, hydrogen bonding, and ion-dipole forces. To earn credit on the exam, you need to identify which forces are present and explain why one substance has stronger attractions than another, using particle-level reasoning.

What is the difference between intermolecular and intramolecular forces?

This is the most common conceptual trap in the whole topic.

When water boils, the O–H covalent bonds stay intact. The molecules simply separate from each other as hydrogen bonds between molecules are overcome. Writing "boiling breaks the bonds in water" is a classic point-losing error.

What are the types of intermolecular forces?

London dispersion forces

Dispersion forces exist between all molecules and atoms, polar or nonpolar. At any instant, the electron cloud can be unevenly distributed, creating a temporary dipole that induces a dipole in a neighbor.

Their strength depends on polarizability — how easily the electron cloud is distorted. Polarizability increases with:

• More electrons (larger molar mass usually tracks this). Br₂ is a liquid at room temperature while Cl₂ is a gas because Br₂ has more electrons and a larger, more polarizable electron cloud.
• Greater surface area / contact. Straight-chain pentane boils higher than its branched isomer neopentane because the elongated molecules have more contact area.

A key AP insight: dispersion forces are not always "weak." In large molecules, total dispersion forces can outweigh the dipole-dipole forces in a smaller polar molecule.

Dipole-dipole forces

Polar molecules have permanent dipoles, and the partially positive end of one molecule attracts the partially negative end of another. To have dipole-dipole forces, a molecule needs polar bonds and an asymmetric shape so the dipoles do not cancel.

• CO₂ has polar bonds but is linear, so dipoles cancel. Nonpolar; dispersion only.
• SO₂ is bent, so it is polar. Dipole-dipole plus dispersion.

Hydrogen bonding

Hydrogen bonding is an especially strong dipole-dipole attraction. It occurs when hydrogen is covalently bonded to N, O, or F and is attracted to a lone pair on an N, O, or F atom of a nearby molecule.

Why so strong? N, O, and F are highly electronegative and small, leaving the hydrogen with a large partial positive charge that can get very close to a lone pair.

Important precision points:

• A hydrogen bond is an intermolecular attraction, not a covalent bond.
• CH₃OCH₃ (dimethyl ether) has O but no H bonded to O, so it cannot hydrogen bond with itself. Its isomer ethanol, CH₃CH₂OH, can — which is why ethanol boils much higher.
• On the exam, draw hydrogen bonds as dashed lines between the H of one molecule and a lone pair on N, O, or F of another.

Ion-dipole forces

Ion-dipole forces occur between an ion and a polar molecule. They explain why ionic compounds like NaCl dissolve in water: the partially negative oxygen ends of water molecules surround Na⁺ ions, and the partially positive hydrogen ends surround Cl⁻ ions. Smaller ions with larger charges attract water molecules more strongly.

How do you identify which IMFs a molecule has?

Use this decision sequence:

1. Is it ionic, mixed with a polar solvent? Ion-dipole forces are present.
2. Does it have H bonded directly to N, O, or F? Hydrogen bonding is present.
3. Is the molecule polar? (Draw the Lewis structure, use VSEPR to find the shape, check whether bond dipoles cancel.) If yes, dipole-dipole forces are present.
4. Every molecule has London dispersion forces.

How do IMFs affect physical properties?

Stronger intermolecular forces mean more energy is needed to separate particles.

Vapor pressure moves in the opposite direction from boiling point. A liquid with strong IMFs has fewer molecules with enough energy to escape into the gas phase at a given temperature.

Solubility: "like dissolves like"

A substance dissolves well when solute-solvent attractions are comparable to the solute-solute and solvent-solvent attractions being disrupted. Polar and ionic substances tend to dissolve in polar solvents; nonpolar substances dissolve in nonpolar solvents. For example, I₂ is far more soluble in nonpolar hexane than in water, because water molecules are strongly attracted to each other through hydrogen bonding and gain little by interacting with nonpolar I₂.

Worked example: ranking boiling points

Rank in order of increasing boiling point: CH₄, CH₃OH, CH₃Cl.

1. CH₄: nonpolar, dispersion only, few electrons.
2. CH₃Cl: polar, dipole-dipole plus dispersion, and more electrons than CH₄.
3. CH₃OH: hydrogen bonding through its O–H group.

Order: CH₄ < CH₃Cl < CH₃OH.

A full-credit justification names the forces for each substance and compares their strengths explicitly: "CH₃OH has the highest boiling point because its molecules experience hydrogen bonding, which requires more energy to overcome than the dipole-dipole and dispersion forces in CH₃Cl."

How to write IMF answers on AP Chem FRQs

1. Name the specific forces for every substance you compare — do not just say "stronger IMFs."
2. Explain the cause (more electrons, greater polarizability, H bonded to O).
3. Connect to the property ("so more energy is needed to separate the molecules, resulting in a higher boiling point").
4. Never say bonds break during a phase change of a molecular substance.
5. Compare like with like. If two substances have the same types of forces, the deciding factor is usually polarizability from electron count or shape.

For broader exam strategy, see our AP Chem FRQ guide and the full roadmap in how to get a 5 on AP Chemistry. Periodic trends feed directly into this topic, so pair this with our guide to AP Chem periodic trends.

How Ivystar can help

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