Intermolecular forces, boiling and solubility

What you’ll learn

• Identify London forces, permanent dipole interactions and hydrogen bonding.
• Explain boiling trends and the open structure of ice.
• Relate solubility to interactions formed and disrupted.

Understand the idea

All atoms and molecules experience London forces caused by fluctuating electron distributions and induced dipoles. Larger, more polarisable electron clouds generally strengthen these attractions. Molecules with permanent dipoles have additional dipole–dipole interactions. Hydrogen bonding requires a hydrogen bonded to nitrogen, oxygen or fluorine and a suitable lone-pair acceptor.

Boiling separates molecules; it does not normally break their internal covalent bonds. Longer alkane chains generally boil at higher temperatures because London attractions are stronger. Increased branching often lowers boiling temperature among alkane isomers by reducing effective surface contact. Alcohols boil higher than similarly sized alkanes because they can form hydrogen bonds. Hydrogen fluoride is anomalously high boiling compared with the London-force trend from hydrogen chloride to hydrogen iodide.

Hydrogen bonding gives water unusually high melting and boiling temperatures. In ice it supports an open arrangement, so ice is less dense than liquid water. Dissolving requires disruption of existing attractions and formation of solute–solvent attractions. Water hydrates ions and hydrogen-bonds to short-chain alcohols; a polar halogenoalkane is not necessarily water-soluble because polarity alone does not guarantee sufficiently favourable new interactions.

Key formulae

$$\mathrm{O-H\cdots O}\quad\text{hydrogen-bond interaction}$$

$$\mathrm{H_2O(l)\rightarrow H_2O(g)}\quad\Delta H_{\mathrm{vap}}>0$$

$$\rho=\frac mV$$

$$T_b(\mathrm{HCl})<T_b(\mathrm{HBr})<T_b(\mathrm{HI})$$

Worked example

Ethanol and dimethyl ether share the molecular formula C₂H₆O. Explain why ethanol has the higher boiling temperature.

1. Both molecules have London forces and a polar oxygen-containing region.
2. Ethanol contains an O–H group, allowing hydrogen bonds between its own molecules.
3. Dimethyl ether has oxygen lone pairs but no O–H, N–H or F–H donor, so pure ether cannot form that hydrogen-bond network.

Answer: More energy is needed to overcome the additional hydrogen bonding between ethanol molecules.