London (dispersion) forces, permanent dipole-dipole forces and hydrogen bonding, how they arise, and how they explain boiling points, solubility and the anomalous properties of water.

What this topic is asking

Edexcel Topic 8 wants you to identify the three intermolecular forces, explain how each arises at the molecular level, and use them to explain physical properties such as boiling point, solubility and the unusual behaviour of water. You should be able to rank forces by strength and justify trends down a group or across a series of similar molecules.

The answer

The three intermolecular forces

All three are weak compared with covalent bonds (a few $\text{kJ mol}^{-1}$ for London forces, up to about $40\ \text{kJ mol}^{-1}$ for a hydrogen bond, versus several hundred $\text{kJ mol}^{-1}$ for a typical covalent bond). London forces are always present; polar molecules have London plus dipole-dipole; molecules that meet the N/O/F rule also have hydrogen bonding on top.

Why hydrogen bonding needs N, O or F

Hydrogen bonding requires two things together: an H atom bonded to a very electronegative atom (so the H carries a large $\delta+$), and a small, electronegative atom carrying a lone pair for the H to be attracted to. Only N, O and F are electronegative and small enough. Chlorine is electronegative but its larger, more diffuse lone pair gives only dipole-dipole attraction, which is why $\text{HCl}$ does not hydrogen bond.

Explaining boiling points

Boiling overcomes intermolecular forces; it does not break covalent bonds. Down group 4 the hydrides $\text{CH}_4, \text{SiH}_4, \text{GeH}_4, \text{SnH}_4$ show a steady rise in boiling point because each has more electrons and stronger London forces. The hydrides of groups 5, 6 and 7 follow the same trend except that $\text{NH}_3$, $\text{H}_2\text{O}$ and $\text{HF}$ are anomalously high because they hydrogen bond. Water is the most striking: its boiling point of $100\ ^\circ\text{C}$ is roughly $160\ ^\circ\text{C}$ higher than the size trend alone would predict.

The anomalies of water

Solubility

A substance dissolves when the solute-solvent attractions formed are comparable to or stronger than the solute-solute and solvent-solvent attractions broken. Polar and hydrogen-bonding solutes (such as ethanol, sugars and ammonia) dissolve well in water because they can hydrogen bond to it. Non-polar substances (such as hydrocarbons) dissolve in non-polar solvents through London forces but are nearly insoluble in water, because they cannot replace the strong hydrogen bonds between water molecules.

Examples in context

Example 1. DNA base pairing. The double helix is held together by hydrogen bonds between complementary bases: adenine pairs with thymine through two hydrogen bonds, and guanine with cytosine through three. These bonds are individually weak, which lets enzymes unzip the strands during replication, yet collectively strong enough to keep the helix intact. This is a direct biological consequence of the N/O/F hydrogen-bonding rule from Topic 8.

Example 2. Why ethanol mixes with water but hexane does not. Ethanol ($\text{CH}_3\text{CH}_2\text{OH}$) is fully miscible with water because its $\text{O-H}$ group forms hydrogen bonds with water molecules, replacing the water-water hydrogen bonds that are broken. Hexane ($\text{C}_6\text{H}_{14}$) is non-polar and can only offer weak London forces; mixing it with water would mean breaking strong water-water hydrogen bonds without compensating attractions, so hexane stays as a separate layer. This "like dissolves like" behaviour underpins solvent extraction in the laboratory.

Try this

Q1. Explain why the boiling point of $H_2O$ is much higher than that of $H_2S$. [2 marks]

• Cue. Water molecules form hydrogen bonds (O is small and very electronegative); $H_2S$ has only weaker dipole-dipole and London forces.

Q2. Explain why ice is less dense than liquid water. [2 marks]

• Cue. Hydrogen bonds hold water molecules in an open lattice with larger spacing than in the liquid.