Types of intermolecular forceEdexcel International A Level Chemistry: Revision notes

Intermolecular forces are the attractions between separate molecules (or atoms). They are far weaker than the covalent bonds inside a molecule.

When a simple molecular substance melts or boils, only the intermolecular forces are overcome. The covalent bonds stay intact. Three types are needed at this level: London forces, permanent dipole–permanent dipole interactions and hydrogen bonds.

All three are electrostatic attractions between regions of partial charge, written δ+ and δ−.

London forces act between all atoms and molecules, polar or non-polar. They are also called instantaneous dipole–induced dipole forces.

1. Electrons are always moving, so at any instant the electron distribution in an atom or molecule may be uneven. This gives an instantaneous dipole.
2. The instantaneous dipole induces a dipole in a neighbouring particle by repelling or attracting its electrons.
3. The opposite partial charges attract. The dipoles flicker and change constantly but the attraction is always present.

London forces are the only forces between non-polar substances such as the noble gases, alkanes and halogens. They get stronger as the number of electrons (and so the size of the molecule) increases.

A bond is polar when the atoms differ in electronegativity: the more electronegative atom attracts the bonding electrons and becomes δ−.

A molecule is polar only if the bond dipoles do not cancel. Chloromethane, CH₃Cl, is polar. Tetrachloromethane, CCl₄, has polar bonds but is symmetrical, so it is non-polar.

Polar molecules attract each other: the δ+ end of one molecule is attracted to the δ− end of its neighbour. These permanent dipole–permanent dipole interactions act in addition to London forces.

Example: iodine monochloride, ICl, has Cl δ− and I δ+, so it has both types of force. Bromine, Br₂, has London forces only.

A hydrogen bond is a strong type of permanent dipole–permanent dipole attraction. It forms between:

• a δ+ hydrogen atom bonded directly to nitrogen, oxygen or fluorine, and
• a lone pair of electrons on a nitrogen, oxygen or fluorine atom of another molecule.

N, O and F are very electronegative and small, so the X–H bond is strongly polar and the lone pair is easily approached. The hydrogen bond is drawn as a dotted line, lined up with the X–H bond (X–H···X).

• H₂O: each molecule has two δ+ hydrogens and two lone pairs, so it can form up to four hydrogen bonds.
• NH₃: three δ+ hydrogens but only one lone pair on nitrogen.
• HF: one δ+ hydrogen and three lone pairs on fluorine.

These liquids have much higher boiling temperatures than other molecules of similar size.

To predict whether molecules hydrogen-bond with each other, check both conditions:

1. Is there an H atom bonded directly to N, O or F?
2. Is there a lone pair on N, O or F to receive it?

Examples that do: ethanol (O–H), ethanoic acid, hydrazine N₂H₄ and amines such as CH₃NH₂ (N–H).

Examples that do not: H₂S and HCl (S and Cl are not electronegative enough or small enough), CH₄ and CH₃Cl (H bonded to C only), and CH₃OCH₃ or HCHO (oxygen has lone pairs but no H is bonded to it).

Hydrogen bonds are usually stronger than other intermolecular forces for molecules of a similar size, but they are still much weaker than covalent bonds.