Difficulty: Introductory to Intermediate | Prerequisites: Basic Lewis structures, electronegativity, molecular geometry (VSEPR)
This chapter covers the forces that act between molecules (as opposed to the bonds within them) and explains how those forces govern the physical behaviour of liquids, including boiling point, vapour pressure, viscosity, and surface tension. It then extends to phase changes and heating curves. If you are comfortable drawing Lewis structures and assigning molecular polarity, you are ready for this material. If not, review VSEPR and polarity first.
Intermolecular forces (IMFs) are attractions between molecules. Stronger IMFs mean higher boiling points, higher surface tension, higher viscosity, and lower vapour pressure. The four main types, from weakest to strongest in a general sense, are London dispersion forces, dipole-dipole interactions, hydrogen bonding, and ion-dipole forces. Phase changes absorb or release energy depending on the direction, and heating curves let you calculate the energy involved at each stage.
Intermolecular forces (IMFs)
Attractive or repulsive forces that act between neighbouring molecules, atoms, or ions. They are weaker than intramolecular bonds (covalent, ionic). Think of them as the "glue" holding molecules near each other in liquids and solids without actually bonding them together.
London dispersion forces (LDFs)
Temporary, instantaneous dipole-induced dipole attractions that arise from random fluctuations in electron distribution. Present in all molecules and atoms. In simple terms, electrons shift around briefly, creating a momentary charge imbalance that nudges a neighbour into a matching imbalance.
Dipole-dipole forces
Attractions between the positive end of one polar molecule and the negative end of another. Think of it as two small bar magnets lining up positive-to-negative.
Hydrogen bonding
A particularly strong type of dipole-dipole interaction that occurs when hydrogen is bonded to nitrogen, oxygen, or fluorine (N, O, or F) and is attracted to a lone pair on N, O, or F of a neighbouring molecule. In simple terms, the tiny size of hydrogen and the high electronegativity of N/O/F concentrate charge into a very small space, making the attraction unusually strong.
Ion-dipole forces
Attractions between an ion (cation or anion) and the partial charge on a polar molecule. These are typically the strongest IMF and are central to how ionic compounds dissolve in water.
Surface tension
The energy per unit area at the surface of a liquid, arising because surface molecules are pulled inward by IMFs with no balancing pull from above. Think of it as an invisible "skin" on the liquid surface.
Viscosity
A liquid's resistance to flow. Stronger IMFs and larger, more complex molecular shapes both increase viscosity.
Vapour pressure
The pressure exerted by the vapour above a liquid in a closed container at equilibrium. Higher IMF strength means lower vapour pressure, because fewer molecules have enough energy to escape the surface.
Boiling point
The temperature at which a liquid's vapour pressure equals the external (atmospheric) pressure. The normal boiling point is measured at 1 atm.
Normal boiling point vs. observed boiling point
The normal boiling point is defined at exactly 1 atm of external pressure. The observed boiling point changes with altitude or applied pressure: lower atmospheric pressure lowers the boiling point, higher pressure raises it.
Phase change
A physical transformation from one state of matter to another (solid, liquid, gas). Each has a specific name and an associated energy change.
Surfactant
A substance that lowers the surface tension of a liquid, typically by disrupting intermolecular forces at the surface. Soap is the classic example.
London dispersion forces (LDFs)
Present in every molecule, polar or nonpolar. They are the only IMF available to nonpolar species.
Strength increases with molar mass and surface area. A long, stretched-out molecule has stronger LDFs than a compact, spherical one of similar mass, because more electron cloud is in contact.
Example: pentane (straight chain) has a higher boiling point than neopentane (compact, spherical), despite having the same molecular formula.
Dipole-dipole forces
Occur between polar molecules (molecules with a permanent dipole moment).
The positive end of one molecule is attracted to the negative end of another.
Strength depends on the magnitude of the dipole moment.
To identify: draw the Lewis structure, determine geometry (VSEPR), then check if bond dipoles cancel. If they do not cancel, the molecule is polar and participates in dipole-dipole interactions.
Hydrogen bonding
Special, strong case of dipole-dipole interaction.
Requirements: a hydrogen atom covalently bonded to N, O, or F, interacting with a lone pair on N, O, or F of another molecule.
Hydrogen bonding explains water's unusually high boiling point, high surface tension, and many of its unique properties.
When drawing hydrogen bonds between molecules, use a dotted line from the H to the lone pair on the neighbouring N, O, or F.
Ion-dipole forces
Arise when an ionic species interacts with a polar molecule.
Central to the dissolution of salts in water: water molecules orient around cations (oxygen end toward the cation) and anions (hydrogen end toward the anion).
Typically the strongest of the four IMF types covered here.
Identifying IMFs for a given molecule
All molecules have London dispersion forces.
If the molecule is polar, add dipole-dipole.
If the molecule has H bonded to N, O, or F, add hydrogen bonding.
If the molecule interacts with an ion, add ion-dipole.
To determine the strongest IMF between two identical molecules, pick the highest-ranking force in that list.
Relationship between IMF strength and liquid properties
Stronger IMFs lead to higher boiling points, higher melting points, higher surface tension, and higher viscosity.
Stronger IMFs lead to lower vapour pressure (molecules are held more tightly in the liquid phase).
Surface tension and surfactants
Surface tension arises because molecules at the surface experience a net inward pull (they have no neighbours above them).
Surfactants reduce surface tension by inserting themselves at the surface and disrupting the cohesive IMFs among the liquid molecules.
Critical applications: lung surfactant in premature infants (reduces surface tension in alveoli so they do not collapse), herbicide formulations (surfactants help the herbicide spread across waxy leaf surfaces).
Boiling point details
A liquid boils when its vapour pressure equals the external pressure.
Normal boiling point: the temperature at which vapour pressure reaches 1 atm.
At higher altitudes (lower atmospheric pressure), liquids boil at lower temperatures.
Phase change terminology
Solid to liquid: melting (fusion). Endothermic.
Liquid to solid: freezing (solidification). Exothermic.
Liquid to gas: vaporisation (evaporation/boiling). Endothermic.
Gas to liquid: condensation. Exothermic.
Solid to gas (skipping liquid): sublimation. Endothermic.
Gas to solid (skipping liquid): deposition. Exothermic.
Energy and phase changes
Endothermic changes (melting, vaporisation, sublimation) absorb heat. The system gains energy.
Exothermic changes (freezing, condensation, deposition) release heat. The system loses energy.
During a phase change, temperature remains constant even as heat is added or removed. The energy goes into breaking or forming intermolecular interactions rather than increasing kinetic energy.
Heating curves
A heating curve plots temperature (y-axis) vs. heat added (x-axis).
Sloped regions: single phase present, temperature is rising. The slope is determined by the specific heat capacity of that phase.
Flat (plateau) regions: phase change in progress. Temperature is constant. The length of the plateau is determined by the enthalpy of the phase change (heat of fusion or heat of vaporisation).
On a cooling curve, the same logic applies in reverse.
Particle motion across phases
Solid: particles vibrate in fixed positions.
Liquid: particles move past each other but remain close together.
Gas: particles move freely with large distances between them.
Heat transfer during temperature change (single phase)
q = m \times c \times \Delta Twhere q = heat (J), m = mass (g), c = specific heat capacity (J/g·°C), and ΔT = change in temperature (°C).
Heat transfer during a phase change (constant temperature)
q = n \times \Delta Hwhere q = heat (J or kJ), n = moles, and ΔH = enthalpy of the phase change (e.g., ΔH_fus for melting, ΔH_vap for vaporisation).
Alternatively, if given mass and a per-gram value:
q = m \times \Delta Hwhere ΔH is in J/g or kJ/g.
Heating curve diagram (know how to read one)
x-axis: heat added (q). y-axis: temperature.
Rising slopes: q = mcΔT applies. The slope is 1/(mc), so phases with smaller specific heat have steeper slopes.
Flat plateaus: q = nΔH applies. Longer plateaus indicate larger enthalpies of phase change. The vaporisation plateau is always longer than the fusion plateau for a given substance (ΔH_vap > ΔH_fus).
Lewis structure diagrams for IMFs
Draw lone pairs explicitly when showing hydrogen bonds (the dotted line goes from H to the lone pair on N, O, or F).
For ion-dipole: draw the ion near the appropriate end of the polar molecule (cation near the partial negative end, anion near the partial positive end).
Respiratory distress in premature infants. The alveoli (tiny air sacs) in the lungs are lined with a thin film of water. Surface tension in that film tends to collapse the sacs. Lung surfactant, produced naturally, lowers the surface tension so the alveoli stay open. Premature infants often lack sufficient surfactant, leading to respiratory distress syndrome. Treatment involves administering synthetic or animal-derived surfactant.
Herbicide application and plant leaf surfaces. Waxy leaf surfaces are nonpolar and repel water-based herbicide solutions. Adding a surfactant to the herbicide mixture lowers the surface tension of the spray droplets, allowing them to spread across the leaf instead of beading up and rolling off.
Water repellency of soil (soil wetting). Some soils become hydrophobic after fires or due to organic coatings on soil particles. Surfactants can be added to irrigation water to improve wetting and absorption.
COVID-19 and intermolecular forces. The SARS-CoV-2 virus has a lipid (fatty) envelope held together by intermolecular forces. Soap disrupts that envelope through its surfactant action, which is why handwashing with soap is effective at inactivating the virus.
Students often confuse intermolecular forces with intramolecular bonds. IMFs act between separate molecules; covalent and ionic bonds act within a molecule or compound. Boiling breaks IMFs, not covalent bonds.
Students sometimes think hydrogen bonding can occur whenever hydrogen is present in a molecule. It requires hydrogen bonded specifically to N, O, or F, and a lone pair on N, O, or F on a neighbouring molecule.
A common error is assuming that a molecule with a higher molar mass always has a higher boiling point. Molar mass matters for London dispersion forces, but if a lighter molecule has hydrogen bonding and a heavier one does not, the lighter molecule can have the higher boiling point (e.g., water vs. a larger nonpolar molecule).
Students sometimes think temperature rises continuously as you heat a substance. During a phase change, temperature stays flat on a heating curve even though heat is still being added.
⚠️ Be prepared to identify all IMFs present for a given molecule from its Lewis structure. This is a foundational skill tested repeatedly.
⚠️ Know how to rank molecules by boiling point, vapour pressure, or surface tension using IMF type and molecular size/surface area.
⚠️ Heating curve analysis is a frequent exam question: identify the phase, the phase change, the correct formula (q = mcΔT vs. q = nΔH), and the relevant constant for each segment.
⚠️ Expect at least one question on real-world applications (surfactant in infant lungs, herbicide spreading, soap and COVID-19).
⚠️ Know the names and directions (endo- vs. exothermic) for all six phase changes.
True or False: London dispersion forces are only present in nonpolar molecules. Answer: False. LDFs are present in all molecules.
Fill in the blank: Hydrogen bonding requires H bonded to ______, ______, or ______. Answer: N, O, or F.
True or False: A substance with strong IMFs will have a high vapour pressure. Answer: False. Strong IMFs mean low vapour pressure.
Fill in the blank: During a phase change on a heating curve, the temperature ______. Answer: Remains constant.
True or False: Condensation is an endothermic process. Answer: False. Condensation is exothermic (releases heat).
Q: What intermolecular forces are present between two molecules of CH₃OH (methanol)?
A: London dispersion forces, dipole-dipole forces, and hydrogen bonding. Methanol has an O-H bond, so hydrogen bonding is present. It is also polar (dipole-dipole) and, like all molecules, exhibits LDFs.
Q: Rank the following in order of increasing boiling point: CH₄, CH₃OH, NaCl dissolved in water (ion-dipole context). Explain your reasoning.
A: CH₄ < CH₃OH < NaCl (in aqueous context, ion-dipole is strongest). CH₄ is nonpolar with only weak LDFs. CH₃OH has hydrogen bonding. NaCl's ions interact with water through ion-dipole forces, the strongest IMF type here.
Q: On a heating curve for water, which plateau is longer, the melting plateau or the boiling plateau? Why?
A: The boiling plateau is longer because the heat of vaporisation (ΔH_vap) is much larger than the heat of fusion (ΔH_fus). More energy is needed to completely separate molecules into the gas phase than to loosen them from a solid lattice into a liquid.
Q: A molecule has a high surface tension and a low vapour pressure. What can you infer about its intermolecular forces?
A: The molecule has strong intermolecular forces. Strong IMFs hold molecules tightly at the surface (high surface tension) and make it difficult for molecules to escape into the vapour phase (low vapour pressure).
Q: Why does water boil at a lower temperature on top of a mountain than at sea level?
A: Atmospheric pressure is lower at higher altitudes. A liquid boils when its vapour pressure equals the external pressure, so lower external pressure means the liquid reaches that threshold at a lower temperature.
Q: Identify the phase change and state whether it is endothermic or exothermic: frost forming on a car windscreen overnight.
A: Deposition (gas directly to solid). It is exothermic, as energy is released when water vapour transitions to ice without passing through the liquid phase.
This connects directly to Chapter 12 (Solutions): ion-dipole forces are the primary interaction that drives dissolution of ionic compounds in water, and "like dissolves like" is rooted in matching IMF types between solute and solvent. Colligative properties (Ch. 12.6-7) depend on how solute particles disrupt the solvent's IMFs at the surface and at phase boundaries.
Heating curves and enthalpy of phase changes connect to thermochemistry concepts from earlier in the course (Hess's law, enthalpy calculations).
Hydrogen bonding reappears in biochemistry: it holds DNA base pairs together and stabilises protein secondary structures.
Intermolecular forces, IMFs, London dispersion forces, van der Waals forces, induced dipole, temporary dipole, dipole-dipole interaction, hydrogen bonding, H-bonding, ion-dipole, surface tension, viscosity, vapour pressure, vapor pressure, boiling point, normal boiling point, melting point, phase change, phase transition, heating curve, cooling curve, heat of vaporisation, heat of vaporization, enthalpy of fusion, enthalpy of vaporisation, sublimation, deposition, condensation, evaporation, freezing, melting, surfactant, lung surfactant, soil wetting, CHM 11200, Purdue general chemistry, Chapter 11