AP Chemistry: Intermolecular Forces

Intermolecular forces (IMFs) are the attractions between separate particles, such as molecules, atoms, or ions. In AP Chemistry, they are one of the most important tools for explaining physical properties: boiling and melting points, vapor pressure, viscosity, surface tension, and many solubility trends. They are also a frequent source of mistakes, especially when students confuse them with intramolecular forces.

Intermolecular vs. Intramolecular Forces: The Distinction That Drives Everything

The first step is separating two ideas that sound similar but behave very differently.

Intramolecular forces

Intramolecular forces are the bonds within a particle: covalent bonds inside a molecule or ionic bonding inside an ionic compound’s lattice. These forces determine chemical identity and are typically strong. Breaking them is a chemical change.

Intermolecular forces

Intermolecular forces act between particles. They are weaker than chemical bonds and govern many physical changes. Boiling water, for example, does not break O–H covalent bonds; it separates water molecules from one another by overcoming intermolecular attractions.

A practical AP Chemistry checkpoint: when you are asked about a physical property (boiling point, volatility, viscosity), think IMFs first. When you are asked about bond strength or reactivity, think intramolecular forces.

Why Intermolecular Forces Control Physical Properties

Intermolecular forces influence how easily particles can move relative to each other. Stronger attractions mean particles “stick” together more, requiring more energy to separate.

Key consequences:

• Boiling point increases as IMFs get stronger (more energy needed to escape into gas).
• Vapor pressure decreases as IMFs get stronger (fewer particles escape into gas at a given temperature).
• Viscosity and surface tension increase with stronger IMFs (more resistance to flow; stronger cohesive forces at the surface).
• Solubility depends on whether solute-solvent attractions can compete with solute-solute and solvent-solvent attractions.

The Major Types of Intermolecular Forces (and How to Recognize Them)

AP Chemistry emphasizes four common categories: London dispersion, dipole-dipole, hydrogen bonding, and ion-dipole interactions. Real substances often experience more than one type simultaneously; the question is usually which one is strongest or most responsible for the trend.

London Dispersion Forces (LDF)

London dispersion forces arise from temporary fluctuations in electron density that create instantaneous dipoles, which then induce dipoles in neighboring particles. They are present in all atoms and molecules, including nonpolar substances.

What strengthens LDF?

• More electrons and larger molar mass: more polarizable electron clouds.
• Greater surface area: long, less-branched molecules have more contact and stronger dispersion than compact, highly branched isomers.

AP-style implications:

• Noble gases show increasing boiling points down the group (He < Ne < Ar < Kr < Xe) due to increasing polarizability.
• For hydrocarbons, a longer chain typically has a higher boiling point than a shorter chain because dispersion forces grow with size and surface area.

A common pitfall: saying dispersion forces are “weak so they don’t matter.” For large molecules, LDF can dominate and produce high boiling points even without polarity.

Dipole-Dipole Forces

Dipole-dipole forces occur between polar molecules. Molecules with permanent dipoles align such that the partially positive end of one attracts the partially negative end of another.

Recognizing dipole-dipole interactions:

1. Determine if the molecule is polar (bond polarity plus molecular geometry).
2. If the molecule has a net dipole, dipole-dipole forces are present.

Dipole-dipole forces generally increase boiling point relative to a similar-sized nonpolar molecule, but size still matters. A smaller polar molecule can boil at a lower temperature than a larger nonpolar molecule if the larger one has much stronger dispersion forces.

Hydrogen Bonding (A Strong Special Case of Dipole-Dipole)

Hydrogen bonding is an especially strong dipole-dipole interaction. It occurs when hydrogen is covalently bonded to a highly electronegative atom and interacts with a lone pair on another electronegative atom.

For AP Chemistry, the standard rule:

• Hydrogen bonding occurs when a molecule contains H bonded to N, O, or F (donor) and another molecule has a lone pair on N, O, or F (acceptor).

Hydrogen bonding explains why water has unusually high boiling and melting points for its molar mass. Water molecules form extensive hydrogen-bond networks, requiring significant energy to separate.

Important nuance for common test questions:

• A molecule may have electronegative atoms but still not hydrogen bond with itself if it lacks H attached to N/O/F. For example, acetone (with O) can accept hydrogen bonds from water, but it cannot donate hydrogen bonds because it has no O–H or N–H bond.

Ion-Dipole Interactions

Ion-dipole interactions occur between an ion and the partial charges of a polar molecule. They are crucial for understanding why many ionic compounds dissolve in water and why hydration is energetically favorable.

Classic example: dissolving NaCl in water. Water’s oxygen end (partial negative) stabilizes Na⁺, while water’s hydrogen end (partial positive) stabilizes Cl⁻. The resulting hydration shells lower the energy of the separated ions.

Ion-dipole forces are often stronger than hydrogen bonding and standard dipole-dipole interactions because full charges are involved on the ion side.

Ranking Intermolecular Forces: A Useful AP Chemistry Framework

A typical strength order (for comparable-sized species) is:

ion-dipole > hydrogen bonding > dipole-dipole > London dispersion

Two cautions that matter in exam settings:

• London dispersion can exceed dipole-dipole in large, heavy, highly polarizable molecules.
• Many substances have multiple IMFs at once; the strongest interaction often dictates the dominant trend, but the others still contribute.

Solubility and “Like Dissolves Like”

Solubility depends on whether forming solute-solvent attractions compensates for breaking existing attractions.

Polar with polar, nonpolar with nonpolar

• Polar solutes tend to dissolve in polar solvents (dipole-dipole and hydrogen bonding can form).
• Nonpolar solutes tend to dissolve in nonpolar solvents (dispersion-based interactions dominate).

Ionic compounds in water

Ionic compounds often dissolve in water because ion-dipole attractions between ions and water molecules can compete with the ionic lattice energy. If lattice energy is extremely high, solubility can be low, but the interaction type remains ion-dipole.

Hydrogen bonding and solubility

Compounds capable of hydrogen bonding often have higher water solubility than similar molecules that cannot hydrogen bond. For instance, alcohols (with O–H) are generally more water-soluble than comparable ethers (no O–H), especially at smaller chain lengths. As hydrocarbon chains grow, increased nonpolar character can reduce solubility even if hydrogen bonding is possible.

Connecting IMFs to Common Lab-Scale Observations

Intermolecular forces are not abstract. They show up in everyday chemical behavior:

• Volatility: Substances with weaker IMFs evaporate more easily, leading to stronger odors and faster drying.
• Viscosity: Honey-like flow often indicates strong attractions and molecular “entanglement,” both related to intermolecular interactions.
• Surface tension: Water’s high surface tension reflects strong cohesive forces, largely from hydrogen bonding.

When asked to justify a trend, tie the observation directly to energy: stronger attractions require more energy to separate particles, changing boiling points, vapor pressures, and phase behavior.

How to Avoid the Most Common AP Chemistry Errors

1. Do not confuse polarity with hydrogen bonding. Hydrogen bonding requires H bonded to N, O, or F. A polar molecule is not automatically hydrogen bonding.
2. Do not ignore dispersion forces. They exist in all substances and often dominate in large molecules.
3. Do not talk about breaking covalent bonds during boiling or melting. Phase changes overcome IMFs, not intramolecular bonds.
4. Always justify with particle-level reasoning. Mention what is attracting what (temporary dipoles, permanent dipoles, H-bond donors/acceptors, ions interacting with dipoles).

Intermolecular forces are a unifying theme in AP Chemistry. Once you can identify the dominant IMF and relate it to energy and particle motion, many “trend” questions become predictable rather than memorized.