Intermolecular forces
The idea
Intermolecular forces (IMFs) are the attractions between molecules — far weaker than the covalent bonds within them, yet they decide boiling points, melting points, viscosity, surface tension, and solubility. The hierarchy you need: London dispersion forces act between all molecules and grow with electron count and contact area; dipole–dipole forces add on for polar molecules; hydrogen bonding — an especially strong dipole interaction requiring H bonded directly to N, O, or F — adds the most per interaction; ion–dipole forces dominate when ions dissolve in polar solvents.
The comparison strategy is sequential: first match molar masses to hold dispersion roughly constant, then ask which molecule carries the stronger additional force. Shape matters too — a long unbranched chain offers more contact area than its compact branched isomer, so it boils higher despite identical formula. Polarity comes straight from your VSEPR analysis: a molecule needs both polar bonds and a shape that fails to cancel them.
Two corrections to common errors: boiling water breaks hydrogen bonds between molecules, never the O–H covalent bonds inside them; and dispersion is not always the weakest player — in large molecules its cumulative effect can outweigh hydrogen bonding, which is why nonpolar candle wax is a solid while hydrogen-bonding ethanol is a liquid.
Worked example
Butane (CH₃CH₂CH₂CH₃, 58 g/mol), acetone (CH₃COCH₃, 58 g/mol), and 1-propanol (CH₃CH₂CH₂OH, 60 g/mol) have nearly identical molar masses. Rank their boiling points from lowest to highest and justify the order.
- Equal molar masses and similar sizes mean the dispersion contributions are comparable, so the ranking will be decided by the additional forces each molecule can form.
- Butane is a pure hydrocarbon with essentially nonpolar bonds — dispersion only, so it should boil lowest; indeed it is a gas at room temperature (bp about −1 °C).
- Acetone has a polar C=O group and a bent-at-carbonyl shape that leaves a net dipole, adding dipole–dipole attraction; but its hydrogens sit on carbon, not oxygen, so it cannot hydrogen-bond to itself. It boils in the middle (about 56 °C).
- 1-Propanol carries an O–H group: hydrogen bonding between molecules is the strongest IMF available here, so it boils highest (about 97 °C).
- Sanity-check the spread: each added force type lifts the boiling point by tens of degrees at constant mass — exactly the staircase the data shows: −1 °C, 56 °C, 97 °C.
Answer. Butane < acetone < 1-propanol (about −1 °C, 56 °C, and 97 °C), because at matched molar mass the boiling point tracks dispersion-only < dipole–dipole < hydrogen bonding.
Check your understanding
- Why must you compare molecules of similar molar mass before crediting a boiling-point difference to dipoles or hydrogen bonds?
- How can a molecule with very polar bonds, like CO₂, still have zero dipole moment, and which force is left to hold its solid together?
- What changes about hydrogen bonding when a molecule can donate H-bonds versus only accept them, and how does that show up in boiling points?
- Why does pentane boil 27 °C higher than its isomer neopentane even though their formulas are identical?
Build the foundations first
Intermolecular forces builds on these concepts. If any feel shaky, start there.