Forces between moleculesAQA A-Level Chemistry: Revision notes

Intermolecular forces act between molecules and are much weaker than covalent bonds. The three types, from weakest to strongest, are:

• Induced dipole-dipole (van der Waals, dispersion, London) forces: present in all molecules
• Permanent dipole-dipole forces: between polar molecules
• Hydrogen bonding: the strongest, between molecules containing H bonded to N, O or F

When a molecular substance melts or boils, only these forces are overcome; the covalent bonds stay intact.

Electrons are in constant motion, so at any instant the electron density in a molecule may be unevenly spread, giving a temporary instantaneous dipole. This induces a dipole in a neighbouring molecule, and the two attract. These forces exist between all molecules, including non-polar ones such as the noble gases, alkanes and halogens.

They get stronger as the number of electrons increases (larger molecules), so boiling point rises with chain length. Branching lowers the boiling point because compact molecules have less surface contact: pentane boils at 36 °C but 2,2-dimethylpropane at 10 °C.

Molecules with a permanent dipole (polar bonds that do not cancel) attract each other through the δ+ end of one molecule and the δ− end of another, in addition to induced dipole-dipole forces. They are stronger than induced forces between molecules of similar size.

Example: propanone, (CH₃)₂CO, boils at 56 °C, compared with 36 °C for pentane, which has a higher relative molecular mass but only induced forces.

A hydrogen bond is a strong attraction between a δ+ hydrogen atom, bonded to N, O or F, and a lone pair on N, O or F of a neighbouring molecule. It requires a very electronegative atom with a lone pair and a hydrogen attached directly to it.

Examples: H₂O, NH₃, HF, alcohols (R–OH), carboxylic acids and amines. Hydrogen bonding in water is shown by a dotted line from H to the lone pair on O, with the lone pair drawn, and the angle about 180° around the hydrogen (O–H···O).

Down Group 4 the hydride boiling points rise (CH₄ to SnH₄) because induced forces increase with the number of electrons. In Groups 5, 6 and 7, NH₃, H₂O and HF have anomalously high boiling points because they form hydrogen bonds, which are stronger than the forces in PH₃, H₂S and HCl.

Within the other hydrides the boiling point rises down the group again as the induced forces increase. HF boils at 20 °C but HCl at −85 °C.

In liquid water the hydrogen bonds are continually breaking and re-forming and the molecules pack closely. In ice each molecule forms four hydrogen bonds in a fixed tetrahedral arrangement, giving an open lattice with spaces. This holds the molecules further apart, so ice is less dense than liquid water and floats. When ice melts, some hydrogen bonds break and the molecules move closer, so the volume decreases.

For any melting or boiling point question:

1. Identify each molecule as polar or non-polar and whether it has N–H, O–H or F–H.
2. List the intermolecular forces present, strongest first.
3. Compare the electrons (size) if the same type of force is present.
4. Link: stronger forces need more energy, so a higher boiling point.

For example, ethanol (78 °C) has hydrogen bonds, methoxymethane (−24 °C) permanent dipole-dipole forces and propane (−42 °C) only induced dipole-dipole forces, despite similar relative molecular masses.