Difficulty: Introductory to Intermediate | Prerequisites: General chemistry (intermolecular forces basics, electronegativity)
Big Picture
Physical properties tell you how molecules interact with each other, not just how they react. Boiling points, melting points, and solubility are all governed by intermolecular forces (IMFs). This section teaches you to predict which molecule in a pair will have the higher boiling point, which is a staple of organic chemistry exams. It also covers combustion energy as a measure of molecular stability. If you understand electronegativity and can identify functional groups (from the first set of notes), you have what you need.
Boiling points increase with stronger intermolecular forces and larger molecular surface area. The IMF hierarchy runs hydrogen bonds > dipole-dipole > London dispersion forces (LDF), but LDF can overwhelm the others when molecules are large enough. Melting points increase with molecular mass and symmetry. Combustion energy increases with more C–H bonds and, counterintuitively, higher combustion energy per carbon means lower stability.
Intermolecular forces (IMFs)
Forces of attraction between molecules (not within them). These determine physical properties like boiling point, melting point, and solubility. In simple terms, IMFs are what hold molecules near each other in a liquid or solid.
Hydrogen bond (H-bond)
A particularly strong dipole-dipole interaction that occurs when H is bonded to a highly electronegative atom (O, N, or F) and is attracted to a lone pair on another O, N, or F. The strongest common IMF in organic chemistry.
Dipole-dipole interaction
Attraction between the positive end of one polar molecule and the negative end of another. Weaker than hydrogen bonds but stronger than LDF for molecules of similar size.
London dispersion forces (LDF)
Temporary, induced dipole attractions present in all molecules. Strength increases with molecular size and surface area (more electrons = more polarisable). Also called van der Waals forces or dispersion forces. Think of these as the baseline attraction that every molecule has.
Polarizability
How easily an atom's electron cloud can be distorted by a nearby charge. Larger atoms with more electrons are more polarisable, which means stronger LDF. This is why bromine compounds have higher boiling points than chlorine compounds of similar size.
Boiling point
The temperature at which a liquid's vapour pressure equals the external pressure. Higher IMFs mean a higher boiling point, because more energy is needed to pull molecules apart.
Melting point
The temperature at which a solid becomes a liquid. Increases with molecular mass and, notably, with molecular symmetry (symmetrical molecules pack more efficiently into a crystal lattice).
Combustion energy (heat of combustion)
The energy released when a compound burns completely in O₂ to produce CO₂ and H₂O. Higher combustion energy per carbon indicates lower stability of the original molecule.
Two main factors
Boiling points increase with molecular weight (more mass = stronger dispersion forces)
Boiling points decrease with branching (more compact shape = less surface area = fewer dispersion contacts)
The shape effect
The more spherical or compact a molecule is, the lower its boiling point. A long, extended chain has more surface area for LDF contact than a highly branched isomer of the same molecular formula.
Comparing boiling points: worked examples from lecture
A straight-chain alkane vs. a branched isomer of the same formula: the straight chain has the higher boiling point because of greater surface area and stronger LDF. Example: n-pentane (−0.5°C) vs. neopentane (branched, lower).
CH₃CHO (acetaldehyde, bp 20°C) vs. CH₃CH₃ (ethane, bp −89°C): same number of electrons, but acetaldehyde has a polar carbonyl producing dipole-dipole interactions. Ethane has only LDF.
CH₃CHO (bp 20°C) vs. CH₃CH₂OH (ethanol, bp 78°C): both are polar, but ethanol can form hydrogen bonds through its –OH. Hydrogen bonding wins.
Chloroethane vs. bromoethane: bromine is larger and more polarisable, giving the bromo compound a higher boiling point despite a similar dipole.
Alcohols vs. amines of similar size: alcohols have higher boiling points because oxygen is more electronegative than nitrogen, making the O–H hydrogen bond stronger than N–H.
Melting points follow a related but slightly different logic from boiling points.
Melting point increases with molecular mass (heavier molecules need more energy to break out of the solid lattice)
Melting point increases with higher levels of molecular symmetry (symmetrical molecules pack more efficiently into a crystal, requiring more energy to disrupt the lattice)
The symmetry factor is the main wrinkle. A highly symmetrical molecule can have a surprisingly high melting point relative to its boiling point.
The general ranking of IMF strength:
H-bonds > dipole-dipole > London dispersion forces
But this hierarchy has an important caveat: LDF can overwhelm hydrogen bonds when the molecule is large enough. A very large nonpolar molecule (many carbons, high surface area) can have a higher boiling point than a small molecule capable of hydrogen bonding.
Decision framework for comparing boiling points
Identify the functional groups present in each molecule
Determine which IMFs each molecule can engage in (H-bonding? Dipole-dipole? LDF only?)
If the IMF types differ, the molecule with the stronger IMF type usually wins
If the IMF types are the same, the molecule with greater molecular weight or surface area wins
If molecular weights are similar but branching differs, the less branched (more extended) molecule wins
This framework handles the vast majority of boiling point comparison questions on exams.
The basic reaction
Combustion converts a hydrocarbon completely to CO₂ and H₂O. Example: CH₄ + 2O₂ → CO₂ + 2H₂O.
Key relationships
Combustion energy increases with the number of carbons (more C–H bonds to break and reform)
You can only fairly compare combustion energies between molecules with the same number of carbons. To compare propane with hexane, multiply propane's value by 2
Higher combustion energy per carbon = lower stability. The molecule releases more energy because its starting point was higher on the energy scale
Why this matters
Combustion energy is one way to measure relative stability of isomers. If two isomers have different heats of combustion, the one that releases less energy is more stable (it was already closer to the energy floor of CO₂ and H₂O).
Students often think hydrogen bonds are the strongest force in every comparison. They are not. A large nonpolar molecule with extensive LDF can out-boil a small molecule with H-bonding. Always consider molecular size alongside IMF type.
Confusing intramolecular bonds (covalent bonds within a molecule) with intermolecular forces (attractions between molecules). Boiling does not break covalent bonds; it overcomes IMFs.
Assuming that branching always lowers boiling point for the right reason. Students sometimes say branching "weakens bonds." The correct explanation: branching makes the molecule more compact, reducing the surface area available for LDF contacts.
Thinking higher combustion energy means higher stability. The relationship is inverted: higher combustion energy per carbon means the molecule started higher on the energy ladder, so it is less stable.
⚠️ "Which compound has the higher boiling point?" is one of the most common exam questions in this chapter. Master the five-step decision framework above.
⚠️ Know the IMF hierarchy cold: H-bonds > dipole-dipole > LDF, with the caveat that LDF scales with molecular size.
⚠️ Be prepared to explain why alcohols boil higher than amines of similar size (O is more electronegative than N, making O–H hydrogen bonds stronger).
⚠️ Combustion energy comparisons require equal carbon counts. If the exam gives you molecules with different numbers of carbons and asks you to compare stability, normalise first.
True or false: London dispersion forces are only present in nonpolar molecules.
Fill in the blank: Boiling point decreases with increased branching because of reduced ______.
True or false: Melting point is affected by molecular symmetry.
Fill in the blank: The IMF hierarchy from strongest to weakest is ______ > ______ > ______.
True or false: Higher combustion energy per carbon means higher stability.
Answers: 1. False (LDF are present in all molecules). 2. Surface area. 3. True. 4. H-bonds > dipole-dipole > LDF. 5. False (higher combustion energy per carbon means lower stability).
Q: Rank the following in order of increasing boiling point: ethane (CH₃CH₃), acetaldehyde (CH₃CHO), ethanol (CH₃CH₂OH). Explain.
A: Ethane (−89°C) < acetaldehyde (20°C) < ethanol (78°C). Ethane has only LDF. Acetaldehyde adds dipole-dipole interactions from the polar C=O. Ethanol adds hydrogen bonding from the –OH group.
Q: Two molecules have the same molecular formula. One is straight-chain, the other is highly branched. Which has the higher boiling point and why?
A: The straight-chain isomer. It has more surface area available for London dispersion force contacts, resulting in stronger overall IMFs.
Q: Chloroethane and bromoethane have similar dipole moments. Why does bromoethane have a higher boiling point?
A: Bromine is larger and more polarisable than chlorine, which gives bromoethane stronger London dispersion forces. The LDF difference outweighs the similar dipole contributions.
Q: Two isomeric hydrocarbons have heats of combustion of 5470 kJ/mol and 5500 kJ/mol, respectively. Which is more stable?
A: The isomer with the lower heat of combustion (5470 kJ/mol) is more stable. It releases less energy upon combustion, meaning it was already lower on the energy scale.
Q: Explain why an alcohol has a higher boiling point than an amine of comparable molecular weight.
A: Oxygen is more electronegative than nitrogen, so the O–H bond is more polarised than the N–H bond. This makes hydrogen bonds involving O–H stronger, raising the boiling point.
Intermolecular forces show up whenever you need to predict solubility (like dissolves like), explain separation techniques (distillation, extraction, chromatography), or reason about reaction conditions. Understanding why certain solvents work for certain reactions draws directly on this material.
Combustion energy connects to thermodynamics and enthalpy diagrams, which become central in reaction energetics (Chapters 6–7). The stability reasoning you learn here for isomers carries through to comparing the stability of intermediates and transition states.
Boiling point, melting point, intermolecular forces, IMF, hydrogen bond, H-bond, dipole-dipole, London dispersion forces, LDF, van der Waals, polarizability, surface area, branching, combustion energy, heat of combustion, molecular symmetry, electronegativity, alcohol vs amine boiling point, organic chemistry chapter 2, physical properties organic chemistry, UW organic chemistry, Prof. Forrest Michael