Intermolecular forcesEdexcel A-Level Chemistry: Revision notes

Intermolecular forces act between molecules and are much weaker than covalent bonds. Boiling or melting a simple molecular substance overcomes these forces; the covalent bonds stay intact.

• London forces (instantaneous dipole–induced dipole): present in all molecules. Electrons move constantly, so an instantaneous dipole forms and induces a dipole in a neighbour. They get stronger with more electrons and a larger surface area.
• Permanent dipole–dipole forces: between polar molecules, attraction of δ+ of one to δ− of another.
• Hydrogen bonds: the strongest, between a δ+ hydrogen bonded to N, O or F and a lone pair on N, O or F of another molecule.

Hydrogen bonding needs a hydrogen atom bonded to a very electronegative atom (N, O or F) and a lone pair on N, O or F in a neighbouring molecule. The bond is drawn as a dotted line, with the three atoms in a straight line (O–H···O).

• H₂O: two δ+ hydrogens and two lone pairs, so up to four hydrogen bonds per molecule.
• NH₃: three δ+ hydrogens but only one lone pair, so fewer hydrogen bonds per molecule.
• HF: one δ+ hydrogen and three lone pairs on F.

Prediction: molecules such as ethanol (O–H) and methylamine (N–H) form hydrogen bonds. A molecule like CH₃OCH₃ has no hydrogen on oxygen and cannot hydrogen bond to itself. CH₄ cannot hydrogen bond at all.

Water has relatively high melting and boiling temperatures for such a small molecule (compare H₂S, bp about −60 °C), because hydrogen bonds between molecules need a lot of energy to overcome.

Ice is less dense than liquid water. In ice each molecule forms four hydrogen bonds in an open, tetrahedral lattice, holding the molecules further apart. On melting, some hydrogen bonds break and the molecules pack more closely, so liquid water is denser. This is why ice floats and why lakes freeze from the top.

Alkanes, with chain length: longer chains have more electrons and a larger surface area, so stronger London forces and a higher boiling temperature.

Branching: a branched isomer is more compact, with less surface contact, so the London forces are weaker and the boiling temperature is lower (2,2-dimethylpropane boils lower than pentane).

Alcohols and alkanes with a similar number of electrons: ethanol boils far above propane because ethanol forms hydrogen bonds as well as London forces, so it is also less volatile.

Hydrogen halides: HF is much higher than HCl because of hydrogen bonding. From HCl to HBr to HI the boiling temperature rises because the number of electrons increases, so London forces are stronger.

Water dissolves some ionic compounds. Polar water molecules surround each ion (hydration): δ− oxygen atoms attract cations and δ+ hydrogen atoms attract anions. The energy released compensates for breaking the lattice.

Water dissolves simple alcohols because they form hydrogen bonds with water.

Water is a poor solvent for halogenoalkanes, which cannot hydrogen bond with water; their attractions to water are weaker than the hydrogen bonds between water molecules that would be broken.

Non-aqueous solvents such as hexane suit substances with similar intermolecular forces (London forces). Ionic compounds do not dissolve in them, since they cannot hydrate the ions.