Difficulty: Introductory to Intermediate | Prerequisites: basic molecular geometry, Lewis structures, electronegativity trends.
Intermolecular forces (IMFs) are the attractions between molecules that determine nearly every bulk physical property you will be tested on: boiling point, viscosity, vapour pressure, surface tension, and miscibility. This topic sits at the bridge between molecular-level structure (Lewis structures, polarity) and macroscopic behaviour (phase changes, solution chemistry). If you are comfortable drawing Lewis structures and identifying polar bonds, you have the prerequisites. Everything here feeds directly into the acids, bases, and buffer material that follows.
Molecules attract each other through dispersion forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole forces, listed weakest to strongest. The stronger these attractions, the higher the boiling point, the higher the surface tension, the higher the viscosity, and the lower the vapour pressure. When you dissolve an ionic compound in water, the number of dissolved particles determines colligative properties such as freezing point depression.
Dispersion forces (London dispersion forces)
Present in all molecules. These arise from temporary, instantaneous dipoles caused by the random motion of electrons. Strength increases with molar mass and surface area.
In simple terms: every molecule has them, even nonpolar ones. Bigger molecules have stronger dispersion forces.
Dipole-dipole forces
Attractions between the positive end of one polar molecule and the negative end of another. Only occur in polar molecules (those with a net dipole moment).
Think of it as: two little magnets lining up positive-to-negative.
Hydrogen bonding
A special, stronger form of dipole-dipole interaction. Occurs when H is bonded directly to N, O, or F and interacts with a lone pair on another N, O, or F atom.
In simple terms: if you see an H attached to N, O, or F, hydrogen bonding is in play. It is why water has an unusually high boiling point for its size.
Ion-dipole forces
Attractions between an ion (from a dissolved ionic compound) and the partial charge on a polar molecule such as water. These are the strongest IMF type covered here.
Think of it as: a full charge talking to a partial charge, so the pull is strong. This is the force that makes salt dissolve in water.
Vapour pressure
The pressure exerted by a vapour in equilibrium with its liquid at a given temperature. Substances with weak IMFs evaporate easily and have high vapour pressure.
In simple terms: how readily a liquid becomes a gas. Nail polish remover (high vapour pressure) evaporates quickly; honey does not.
Viscosity
A liquid's resistance to flow. Stronger IMFs and larger, more tangled molecules lead to higher viscosity.
Think of it as: how "thick" a liquid feels. Water flows easily (low viscosity); syrup does not.
Surface tension
The energy required to increase the surface area of a liquid. Molecules at the surface are pulled inward by their neighbours, creating a "skin." Stronger IMFs mean higher surface tension.
In simple terms: it is why small insects can walk on water and why water droplets bead up on a waxed car.
Surfactant
A molecule with a hydrophilic (water-loving) head and a hydrophobic (water-fearing) tail. Surfactants lower surface tension and help oil and water mix.
Think of it as: soap. The water-friendly end grabs water, the oil-friendly end grabs grease.
Colligative properties
Properties that depend on the number of dissolved solute particles, not their identity. Freezing point depression and boiling point elevation are the two tested most often.
In simple terms: more dissolved particles means a lower freezing point and a higher boiling point. This is why you salt icy roads.
Freezing point depression
The decrease in a solvent's freezing point caused by the presence of a solute. Governed by ΔTf = i · Kf · m, where i is the van 't Hoff factor (number of particles the solute produces).
Think of it as: dissolved particles get in the way of the solvent molecules trying to lock into a solid crystal.
Start with dispersion forces. Every molecule has them. Always list these first.
Check for polarity. If the molecule has a net dipole moment (asymmetric distribution of polar bonds), add dipole-dipole forces.
Check for hydrogen bonding. If H is bonded to N, O, or F, add hydrogen bonding. Note: CHCl₃ (chloroform) has C–H bonds, not O–H, N–H, or F–H, so it does not hydrogen-bond. Its IMFs are dispersion + dipole-dipole only.
Check for ions. If the system contains an ionic compound dissolved in a polar solvent, add ion-dipole forces.
Example: an aqueous NaCl solution contains dispersion, dipole-dipole, hydrogen bonding (from the water molecules), and ion-dipole forces (from Na⁺ and Cl⁻ interacting with water). All four types are present.
From weakest to strongest (in general): dispersion < dipole-dipole < hydrogen bonding < ion-dipole.
When comparing molecules of similar type, look at:
Molar mass and surface area for dispersion-dominated molecules (e.g. comparing hexane to dichloromethane).
Polarity and hydrogen-bonding ability for polar molecules.
For the trio pentane / acetonitrile / ethanol:
Pentane (C₅H₁₂): nonpolar, dispersion only. Weakest IMFs.
Acetonitrile (CH₃CN): polar, has dispersion + dipole-dipole. Moderate IMFs.
Ethanol (CH₃CH₂OH): polar with an O–H bond, so dispersion + dipole-dipole + hydrogen bonding. Strongest IMFs of the three.
Order of increasing IMF strength: pentane < acetonitrile < ethanol.
All four of these properties connect back to how tightly molecules hold onto each other.
Boiling point: stronger IMFs require more energy to pull molecules apart into the gas phase, so the boiling point is higher.
Vapour pressure: if molecules are held tightly (strong IMFs), fewer escape into the gas phase, so vapour pressure is lower. Boiling point and vapour pressure are inversely related.
Surface tension: stronger IMFs mean the surface molecules are pulled inward more strongly, so surface tension is higher.
Viscosity: stronger IMFs (and larger, more entangled molecules) resist flow more, so viscosity is higher.
The key relationship to memorise: if compound A has a higher boiling point than compound B, then compound A has a lower vapour pressure and a higher surface tension.
A cooling curve plots temperature (y-axis) against time (x-axis) as a substance loses heat.
Sloped regions: the substance is cooling within a single phase (gas, liquid, or solid). Temperature decreases over time.
Flat (plateau) regions: a phase change is occurring. Temperature stays constant because energy is being removed to change the phase, not to lower the temperature.
Reading the curve from high temperature to low temperature:
The first plateau (highest temperature flat region) is condensation (gas to liquid). The temperature at this plateau is the boiling point.
The second plateau (lower temperature flat region) is freezing (liquid to solid). The temperature at this plateau is the freezing/melting point.
For a curve starting at about 140 °C and cooling to about 25 °C with plateaus at roughly 90 °C and 50 °C: the boiling point is approximately 90 °C. The process occurring during the 50 °C plateau (e.g. from 14 to 20 minutes) is freezing.
Polar solvents dissolve polar and ionic solutes. Nonpolar solvents dissolve nonpolar solutes.
Water is polar and capable of hydrogen bonding.
Ethanol (CH₃CH₂OH) is polar with hydrogen bonding capability, so it is miscible with water.
Acetonitrile (CH₃CN) is polar, so it mixes with water.
Pentane (C₅H₁₂) is nonpolar (only dispersion forces), so it does not mix with water.
When an ionic compound dissolves, it dissociates into its constituent ions. To find the concentration of a specific ion:
Calculate moles of the compound: mass ÷ molar mass.
Use the formula's stoichiometry to find moles of the ion.
Divide by the volume in litres.
Example: dissolving 2.50 g of FeCl₂ (molar mass ≈ 126.75 g/mol) in 150 mL of water.
Moles of FeCl₂ = 2.50 / 126.75 ≈ 0.01972 mol
Each FeCl₂ produces 2 Cl⁻ ions, so moles of Cl⁻ = 0.03944 mol
[Cl⁻] = 0.03944 / 0.150 ≈ 0.263 M
ΔTf = i · Kf · m
The van 't Hoff factor (i) is the number of particles the solute produces when it dissolves.
NaCl → Na⁺ + Cl⁻, so i = 2. For 1 M NaCl: effective particle concentration = 1 × 2 = 2.
CaCl₂ → Ca²⁺ + 2 Cl⁻, so i = 3. For 0.7 M CaCl₂: effective = 0.7 × 3 = 2.1.
CH₃CH₂OH does not dissociate, so i = 1. For 1 M ethanol: effective = 1 × 1 = 1.
FeCl₃ → Fe³⁺ + 3 Cl⁻, so i = 4. For 0.4 M FeCl₃: effective = 0.4 × 4 = 1.6.
The solution with the highest effective particle concentration has the greatest freezing point depression (lowest freezing point). Here, 0.7 M CaCl₂ (effective = 2.1) wins.
A surfactant is the molecule itself: one hydrophilic head, one hydrophobic tail. This is the definition tested.
A micelle is the spherical structure surfactant molecules form in water (tails inward, heads outward).
A colloid is a mixture where particles are small enough to stay dispersed but large enough to scatter light.
A suspension is a mixture where particles are large enough to settle out over time.
(Solutions and colligative properties content is covered above in the preceding sections.)
ΔTf = i · Kf · m (freezing point depression: i = van 't Hoff factor, Kf = cryoscopic constant of the solvent, m = molality)
ΔTb = i · Kb · m (boiling point elevation: same idea, Kb = ebullioscopic constant)
Molarity (M) = moles of solute / litres of solution
IMF strength ranking: dispersion < dipole-dipole < hydrogen bonding < ion-dipole
Cooling curve: temperature on y-axis, time on x-axis. Plateaus = phase changes (condensation at the boiling point, freezing at the melting point).
Freezing point depression is why road crews spread salt on icy roads in winter: the dissolved ions lower the freezing point of water so ice melts at a lower temperature. Surfactants are the active ingredient in soap and detergents, letting grease (nonpolar) mix with water (polar) so it can be rinsed away.
Students often think CHCl₃ can hydrogen-bond because it contains an H atom. It cannot, because the H is bonded to C, not to N, O, or F.
Students sometimes confuse "surfactant" with "micelle." The surfactant is the individual molecule; the micelle is the structure many surfactant molecules form together in solution.
A common mistake is assuming a higher boiling point means higher vapour pressure. The relationship is inverse: stronger IMFs raise the boiling point and lower the vapour pressure.
When calculating freezing point depression, students forget to multiply by the van 't Hoff factor. An ionic compound that splits into three ions has three times the particle effect of a non-electrolyte at the same molality.
⚠️ Identifying all IMFs present in a substance or solution is one of the most commonly tested skills. For aqueous ionic solutions, remember to include all four types (dispersion, dipole-dipole, hydrogen bonding from the water, and ion-dipole).
⚠️ The boiling point, vapour pressure, surface tension, viscosity relationship appears frequently as a "fill in the blank" or matching question. Know the direction of each.
⚠️ Cooling/heating curve questions test whether you can identify which phase change is occurring at a plateau and read the boiling/melting point from the graph.
⚠️ Freezing point depression problems require you to calculate the effective particle concentration (i × m) for each solution and compare. Do not forget the van 't Hoff factor.
True or false: all molecules exhibit dispersion forces. (True)
Fill in the blank: hydrogen bonding occurs when H is bonded to ___, ___, or ___. (N, O, or F)
True or false: a substance with a high boiling point also has a high vapour pressure. (False, it has a low vapour pressure)
Fill in the blank: NaCl dissociates into ___ particles per formula unit, so its van 't Hoff factor is ___. (2 particles; i = 2)
True or false: pentane is miscible with water. (False, pentane is nonpolar)
Q: What intermolecular forces are present in chloroform (CHCl₃)?
A: Dispersion and dipole-dipole. CHCl₃ is polar (tetrahedral geometry with three Cl atoms pulling electron density), but its H is bonded to C, not N/O/F, so no hydrogen bonding.
Q: What intermolecular forces must be overcome to boil an aqueous NaCl solution?
A: Dispersion, dipole-dipole, hydrogen bonding, and ion-dipole. All four are present because the water molecules interact with each other (dispersion, dipole-dipole, H-bonding) and with the dissolved Na⁺ and Cl⁻ ions (ion-dipole).
Q: Rank water, hexane, and dichloromethane in order of increasing viscosity.
A: Hexane < dichloromethane < water. Hexane has only weak dispersion forces. Dichloromethane is polar (dispersion + dipole-dipole). Water has hydrogen bonding, giving it the strongest IMFs and the highest viscosity of the three.
Q: If compound A has a higher boiling point than compound B, what can you say about compound A's vapour pressure and surface tension relative to B?
A: Compound A has a lower vapour pressure and a higher surface tension.
Q: On a cooling curve, what process occurs during a temperature plateau at roughly 50 °C if the substance's boiling point is 90 °C?
A: Freezing. On a cooling curve, the higher plateau is condensation (at the boiling point) and the lower plateau is freezing (at the melting/freezing point).
Q: Which aqueous solution has the lowest freezing point: 1 M NaCl, 0.7 M CaCl₂, 1 M CH₃CH₂OH, or 0.4 M FeCl₃?
A: 0.7 M CaCl₂. Effective particle concentrations: NaCl = 2, CaCl₂ = 2.1, ethanol = 1, FeCl₃ = 1.6. The highest effective concentration produces the greatest depression.
Q: Which of pentane, ethanol, and acetonitrile will not mix with water?
A: Pentane. It is nonpolar (dispersion forces only) and does not interact favourably with polar water.
Q: A molecule with a hydrophilic head and a hydrophobic tail is best described as what?
A: A surfactant.
IMF concepts feed directly into acid-base chemistry: the ability of water to act as both an acid and a base (amphoteric behaviour) depends on its hydrogen-bonding network. Solution concentration calculations (molarity) come back constantly in buffer and titration problems. Colligative properties connect to thermodynamics and entropy, topics you will meet later in the course.
IMFs, intermolecular forces, London dispersion forces, van der Waals forces, dipole-dipole, hydrogen bonding, H-bonding, ion-dipole, vapour pressure, vapor pressure, boiling point, surface tension, viscosity, cooling curve, heating curve, phase diagram, phase change, condensation, freezing, melting, colligative properties, freezing point depression, boiling point elevation, van 't Hoff factor, molarity, molality, miscibility, like dissolves like, surfactant, micelle, colloid, suspension, CHM 11200, general chemistry, Purdue