Intermolecular forces are the attractions between separate molecules, atoms, or ions in a substance. They are weaker than chemical bonds inside a molecule, but they strongly affect boiling point, melting point, viscosity, surface tension, and solubility. Understanding these forces helps explain why water is a liquid at room temperature, why oil and water do not mix well, and why some substances evaporate quickly while others do not.
Understanding Intermolecular Forces
At the particle level, attractions compete with thermal motion. Particles are always moving, rotating, and vibrating. When they move fast enough, they can pull apart despite attraction.
Cooling slows this motion, so attractions hold particles close for longer. This is why a gas can condense into a liquid when temperature falls. Distance matters greatly.
An attraction becomes much weaker as particles separate, so nearby molecules have the biggest effect. Molecular shape matters because it controls how closely particles can approach and whether charged regions can line up well.
A force is not a permanent glue between two particles. It is a shifting pull that depends on position and motion.
Boiling is a useful place to apply this idea. A liquid boils when particles throughout the liquid gain enough energy to escape into gas bubbles. The chemical bonds within each molecule usually stay intact during boiling.
What is overcome is the attraction between neighboring particles. This explains why boiling point is not simply about molecular mass. A larger molecule often has a higher boiling point because its electron cloud is easier to distort and its contact area can be greater.
Yet a smaller molecule with especially strong directional attractions may boil at a much higher temperature than expected. Evaporation happens at the surface below the boiling point. The fastest surface particles escape first, which removes energy and produces cooling.
Hydrogen bonding is especially important because its geometry is selective. The hydrogen must be attached to a strongly electron pulling atom, and a suitable lone pair on a nearby molecule must be in the right position. This lets molecules form temporary networks rather than random pairings.
Water forms an extended network, which helps explain its unusually high boiling point for such a small molecule. Shape can change the result too. Straight chain molecules usually touch over more area than compact branched molecules of similar mass.
More contact creates more opportunities for weak attractions. When comparing substances, first identify ions, polar bonds, molecular shape, and atoms with lone pairs. Then compare particle size and possible contact area.
Solubility depends on an energy balance, not on a simple rule to memorize. To dissolve a solid, particles in the solid must separate. Solvent particles must make space.
New attractions then form between solute and solvent. Dissolving is favorable when these new attractions compensate enough for the attractions that were disrupted. Water can surround many ions because its charged ends orient toward them.
This surrounding layer is called hydration. Nonpolar substances cannot form equally strong attractions with water, so they often separate into a different layer. Students often confuse polar bonds with a polar molecule.
Bond directions can cancel in a symmetrical shape. Drawing the three dimensional shape and marking charge regions makes predictions much more reliable.
Key Facts
- London dispersion forces occur in all atoms and molecules because electrons can form temporary dipoles.
- Dipole-dipole forces occur between polar molecules with permanent partial positive and partial negative ends.
- Hydrogen bonding occurs when H is bonded to N, O, or F and is attracted to a lone pair on N, O, or F nearby.
- Ion-dipole forces occur between an ion and a polar molecule, such as Na+ attracted to the O end of water.
- For substances with similar molar mass, stronger intermolecular forces usually mean higher boiling point.
- Coulomb's law helps describe attraction strength: F = kq1q2/r^2, so larger charges and shorter distances increase attraction.
Vocabulary
- Intermolecular force
- An attraction between separate particles, such as molecules, atoms, or ions.
- London dispersion force
- A weak attraction caused by temporary shifts in electron density that create temporary dipoles.
- Dipole
- A separation of electric charge in a molecule, with one region partially positive and another partially negative.
- Hydrogen bond
- A strong type of dipole-dipole attraction involving hydrogen bonded to nitrogen, oxygen, or fluorine.
- Ion-dipole force
- An attraction between an ion and the oppositely charged end of a polar molecule.
Common Mistakes to Avoid
- Calling intermolecular forces chemical bonds is wrong because they act between particles, while covalent, ionic, and metallic bonds hold atoms together within a substance.
- Assuming bigger molecules always have higher boiling points is wrong because molar mass matters, but polarity and hydrogen bonding can have an even stronger effect.
- Saying every molecule with hydrogen can hydrogen bond is wrong because hydrogen bonding requires H directly bonded to N, O, or F.
- Ignoring molecular shape is wrong because bond polarity alone does not guarantee a polar molecule if the bond dipoles cancel.
Practice Questions
- 1 Rank these substances from lowest to highest expected boiling point and explain your order: CH4, H2O, H2S. Use the main intermolecular force in each substance.
- 2 A liquid has a normal boiling point of 78 °C and another similar-sized liquid has a normal boiling point of 35 °C. Which liquid likely has stronger intermolecular forces, and by how many degrees Celsius are the boiling points different?
- 3 A student says CO2 must be polar because each C=O bond is polar. Explain why this conclusion is incorrect using molecular shape and dipole cancellation.