Source: Comprehensive Chemistry Study Guide, General Chemistry (Purdue University)
Tags: intermolecular forces, IMFs, dipole-dipole, hydrogen bonding, London dispersion forces, ion-dipole, surface tension, viscosity, vapour pressure, boiling point, surfactants, van der Waals forces
Difficulty: Introductory–Intermediate Prerequisites: Basic understanding of molecular polarity, electronegativity, and Lewis structures.
Intermolecular forces (IMFs) are the attractions that hold molecules near one another in liquids and solids. They are weaker than the covalent or ionic bonds inside a molecule, but they control almost every bulk physical property you can observe: whether a substance is a gas, liquid, or solid at room temperature, how easily it evaporates, how it dissolves, and how it behaves on surfaces. This topic sits at the bridge between molecular structure (which you have already studied) and the macroscopic behaviour of matter (phase changes, solutions, colligative properties) that follows in later chapters. If you can rank IMFs correctly, much of the rest of this unit falls into place.
Molecules attract each other through dipole-dipole interactions, hydrogen bonds, London dispersion forces, and ion-dipole forces, in roughly ascending order of typical strength. These attractions determine liquid properties such as surface tension, viscosity, vapour pressure, and boiling point: stronger IMFs mean higher surface tension, higher viscosity, lower vapour pressure, and higher boiling points.
Dipole-dipole interaction
The electrostatic attraction between the positive end of one polar molecule and the negative end of another. Present whenever a molecule has a permanent dipole.
In simple terms, polar molecules line up positive-to-negative, a bit like tiny magnets attracting each other.
Hydrogen bonding (H-bonding)
A particularly strong type of dipole-dipole interaction that occurs when hydrogen is covalently bonded to a highly electronegative atom (O, N, or F) and is attracted to a lone pair on another O, N, or F atom nearby.
Think of it as the VIP version of dipole-dipole: same idea, but the small size of hydrogen and the high electronegativity of O, N, or F make the attraction unusually strong.
London dispersion forces (LDFs)
Temporary, instantaneous dipole-induced dipole attractions present in all molecules. They arise because electrons are always moving, creating fleeting asymmetries in charge distribution. Larger, more polarisable molecules have stronger LDFs.
In simple terms, even nonpolar molecules experience brief flickers of uneven charge that let them stick to their neighbours. The bigger and more spread out the electron cloud, the stickier they get.
Ion-dipole force
The attraction between an ion (cation or anion) and a polar molecule. This is the dominant force when ionic compounds dissolve in polar solvents such as water.
Think of it as what pulls NaCl apart when you drop it in water: each Na+ and Cl- gets surrounded by water molecules orienting their partial charges toward the ion.
Surface tension
The cohesive force acting at the surface of a liquid, causing it to behave as though it has an elastic skin. Molecules at the surface have fewer neighbours to attract them, so the net inward pull creates tension.
In simple terms, this is why water forms droplets and why a carefully placed paperclip can float on water.
Viscosity
A liquid's resistance to flow. Higher intermolecular forces and larger, more tangled molecules both increase viscosity.
Think of it as how "thick" a liquid feels: honey has high viscosity, water has low viscosity.
Vapour pressure
The pressure exerted by the vapour above a liquid when the liquid and its vapour are in dynamic equilibrium (evaporation rate equals condensation rate) in a closed container.
In simple terms, it measures how easily molecules escape from the liquid surface. Volatile liquids (weak IMFs) have high vapour pressures.
Boiling point
The temperature at which a liquid's vapour pressure equals the external atmospheric pressure. The "normal boiling point" is specifically the boiling point at 1 atm.
In simple terms, it is the temperature where the liquid has enough energy to push molecules into the gas phase against the weight of the atmosphere above it.
Surfactant
A substance that lowers the surface tension of a liquid, typically by sitting at the surface with a hydrophilic head in the water and a hydrophobic tail pointing away from it. Soaps and detergents are everyday surfactants.
Ion-dipole forces are the strongest IMFs listed here, relevant when ions interact with polar molecules (e.g. dissolving salts in water).
Hydrogen bonding is the next strongest among purely molecular interactions. Requires H bonded to O, N, or F, with a lone pair on a nearby O, N, or F to accept the bond.
Water (H₂O), ammonia (NH₃), and hydrogen fluoride (HF) are the classic examples.
Hydrogen bonding explains water's anomalously high boiling point compared to H₂S, H₂Se, and H₂Te.
Dipole-dipole interactions occur between any polar molecules that lack the specific H–O/N/F requirement for hydrogen bonding.
Strength depends on the magnitude of the molecular dipole moment.
London dispersion forces are present in every molecule and are the only IMF available to nonpolar molecules.
Strength increases with molar mass and molecular surface area (longer, more extended shapes are more polarisable than compact, spherical ones).
For large nonpolar molecules, LDFs can be stronger than dipole-dipole forces in small polar molecules.
Surface tension increases with stronger IMFs. Water has unusually high surface tension because of extensive hydrogen bonding.
Viscosity increases with stronger IMFs and with molecular complexity (long chains tangle).
Vapour pressure decreases with stronger IMFs. Molecules need more kinetic energy to escape a liquid held together by strong attractions.
Boiling point increases with stronger IMFs, because you must supply more energy to overcome those attractions and convert liquid to gas.
Respiratory distress in infants: Premature infants can lack pulmonary surfactant, a substance that lowers surface tension inside the lungs' alveoli. Without it, the alveoli collapse and breathing becomes extremely difficult.
Herbicides and leaf surfaces: Spray droplets need to spread across waxy leaf surfaces. Adding surfactants to herbicide formulations lowers surface tension so the liquid wets the leaf instead of beading up.
Soil wetting and water repellency: Some soils become hydrophobic after fires or organic coatings build up. Understanding IMFs helps agronomists choose treatments that restore water infiltration.
Soap and virus transmission: Soap molecules are surfactants whose hydrophobic tails disrupt the lipid bilayer of enveloped viruses, which is one reason hand-washing is effective against viruses like SARS-CoV-2.
No single formula dominates this section, but the following qualitative relationships are essential:
Stronger IMFs → higher boiling point
Stronger IMFs → higher surface tension
Stronger IMFs → higher viscosity
Stronger IMFs → lower vapour pressure
Greater molar mass / larger electron cloud → stronger London dispersion forces
The Clausius-Clapeyron equation (relating vapour pressure to temperature) appears in some courses at this level:
ln(P₂/P₁) = (ΔH_vap / R) × (1/T₁ − 1/T₂)
where ΔH_vap is the enthalpy of vaporisation, R = 8.314 J/(mol·K), and T is in kelvins.
Intermolecular forces are the reason water is a liquid at room temperature while methane (similar molar mass) is a gas. Engineers designing heat exchangers, distillation columns, or even paint formulations rely on boiling point and vapour pressure data that trace directly back to IMFs. In biology, hydrogen bonding holds the two strands of DNA together and determines how proteins fold into functional shapes.
"Hydrogen bonds are bonds inside a molecule." They are not. Hydrogen bonds are intermolecular attractions between molecules (or between different parts of a very large molecule). The covalent O–H bond inside water is not the same thing as the hydrogen bond between two water molecules.
"London dispersion forces only matter for nonpolar molecules." LDFs are present in all molecules, polar or not. In large polar molecules, LDFs can even be the dominant contribution to total IMF strength.
"Higher molar mass always means higher boiling point." This holds when comparing similar types of molecules, but a small molecule with hydrogen bonding (e.g. water, 18 g/mol) can have a higher boiling point than a larger nonpolar molecule relying only on LDFs.
"Vapour pressure increases when IMFs are stronger." The relationship is inverse. Stronger IMFs make it harder for molecules to escape into the gas phase, so vapour pressure decreases.
⚠️ Exam questions frequently ask you to rank a set of molecules by boiling point or vapour pressure. The key is to identify the strongest IMF present in each molecule, then rank accordingly.
⚠️ Be precise about hydrogen bonding criteria: H must be bonded to O, N, or F. A molecule like CH₃OCH₃ has lone pairs on oxygen but no O–H bond, so it cannot donate a hydrogen bond (though it can accept one).
⚠️ When two molecules have the same type of strongest IMF, the tiebreaker is usually molar mass (larger → stronger LDFs → higher boiling point).
⚠️ Questions on surfactants and surface tension appear in applied/contextual problems. Know that surfactants lower surface tension and why.
True or false: Hydrogen bonding can occur in any molecule that contains hydrogen atoms.
Fill in the blank: The only type of intermolecular force present between two nonpolar molecules is __________.
True or false: A liquid with strong intermolecular forces will have a high vapour pressure.
Fill in the blank: The boiling point of a liquid is the temperature at which its vapour pressure equals the __________ pressure.
True or false: Ion-dipole forces are important when an ionic compound dissolves in water.
Answers: 1. False (H must be bonded to O, N, or F). 2. London dispersion forces. 3. False (strong IMFs → low vapour pressure). 4. External (atmospheric). 5. True.
Q: Rank the following in order of increasing boiling point: CH₄, CH₃OH, CH₃CH₃. Explain your reasoning.
A: CH₄ < CH₃CH₃ < CH₃OH. Both CH₄ and CH₃CH₃ are nonpolar and interact only through London dispersion forces, with CH₃CH₃ having the higher molar mass and therefore stronger LDFs. CH₃OH can form hydrogen bonds (O–H group), giving it the highest boiling point of the three.
Q: Why does water have a higher surface tension than most other liquids?
A: Water molecules form an extensive hydrogen-bonding network. Molecules at the surface are pulled inward by these strong attractions to their neighbours, creating high surface tension.
Q: A student claims that HCl has hydrogen bonding because it contains hydrogen. What is wrong with this claim?
A: Hydrogen bonding requires H to be bonded to O, N, or F. Chlorine is not electronegative enough to support hydrogen bonding in HCl. HCl has dipole-dipole interactions (it is polar), but these are not classified as hydrogen bonds.
Q: Explain why adding soap to water allows a needle that was floating on the surface to sink.
A: Soap is a surfactant that disrupts the hydrogen-bonding network at the water's surface, lowering its surface tension. With reduced surface tension, the water can no longer support the weight of the needle.
This material connects directly to phase changes (the energy required to overcome IMFs determines heats of fusion and vaporisation), solutions and solubility (the "like dissolves like" principle is really about matching IMF types between solute and solvent), and colligative properties (the presence of solute particles disrupts solvent IMFs, altering boiling point, freezing point, and vapour pressure).
intermolecular forces, IMF, dipole-dipole, hydrogen bond, H-bond, London dispersion, van der Waals, ion-dipole, surface tension, viscosity, vapour pressure, vapor pressure, boiling point, normal boiling point, surfactant, polarisability, polarizability, like dissolves like, cohesive forces, adhesive forces, volatile, nonvolatile, Clausius-Clapeyron