Hydrogen Bonding: Strength and Effects Revision Notes | A Level AQA Chemistry

1. Why Hydrogen Bonding Matters

Water boils at 100 °C. By every prediction based on molecular mass, it should boil somewhere around −80 °C. That gap — roughly 180 degrees — is entirely down to hydrogen bonding. Understanding this force explains not just boiling points, but why ice floats, why DNA holds its shape, and why proteins fold correctly.

At A Level, you need to do more than recognise hydrogen bonds — you need to explain them precisely, compare their strength to other intermolecular forces, and use them to justify physical properties. This lesson builds that full picture from the ground up.

2. What a Hydrogen Bond Actually Is

Picture a water molecule. The oxygen atom is highly electronegative — it pulls electron density strongly towards itself, leaving the hydrogen atom with a significant partial positive charge ($\delta+$). That electron-poor hydrogen is now attracted to a lone pair on the oxygen of a neighbouring water molecule. That electrostatic attraction between the $\delta+$ hydrogen and the lone pair is a hydrogen bond.

Three conditions must all be met for a hydrogen bond to form: the hydrogen must be bonded to N, O, or F; the neighbouring atom must be N, O, or F; and that neighbouring atom must carry a lone pair. Remove any one of these conditions and the hydrogen bond disappears.

3. Comparing Intermolecular Forces: Strength and Context

Hydrogen bonds do not exist in isolation — they sit within a hierarchy of intermolecular forces. To explain physical properties accurately, you need to know where hydrogen bonding sits relative to London dispersion forces and permanent dipole–dipole interactions.

Notice that London dispersion forces can actually rival hydrogen bonds in very large molecules — a long-chain alkane has enormous London dispersion forces due to its large electron cloud and extensive surface area. This is why the boiling point of a large alkane can exceed that of a small molecule with hydrogen bonding.

4. How Intermolecular Forces Determine Melting and Boiling Points

Melting and boiling are both processes that require energy to overcome intermolecular forces — not to break covalent bonds within molecules. The stronger the intermolecular forces, the more energy is needed to separate the molecules, and therefore the higher the melting and boiling points.

When a molecular solid melts, the molecules gain enough kinetic energy to overcome the forces holding them in their fixed lattice positions. When a liquid boils, the molecules gain enough energy to completely overcome intermolecular attractions and escape into the gas phase. For small to medium-sized molecules, hydrogen bonds are the strongest intermolecular force and require more energy to break than London dispersion forces or permanent dipole–dipole interactions in comparable molecules — hence their dramatic effect on boiling points. Note, however, that in very large molecules, London dispersion forces can become substantial enough to rival or exceed hydrogen bonds in their overall contribution to boiling point.

This principle applies across all molecular substances: compare two molecules of similar size, and the one with stronger intermolecular forces will have the higher melting and boiling points. Hydrogen bonding is the clearest example of this, but the same logic applies when comparing London dispersion forces across a homologous series.

5. Anomalous Boiling Points of Group 16, 15, and 17 Hydrides

Here is where hydrogen bonding becomes genuinely striking. Look at the boiling points of the Group 16 hydrides: $\text{H}_2\text{S}$ (−60 °C), $\text{H}_2\text{Se}$ (−41 °C), and $\text{H}_2\text{Te}$ (−2 °C). As molecular mass increases down the group, London dispersion forces increase and boiling points rise — exactly as expected. Now add water ($\text{H}_2\text{O}$) to that trend. Based on the pattern, water should boil at around −80 °C. Instead, it boils at 100 °C. That is an anomaly of roughly 180 degrees.

The same anomaly appears in Group 15 and Group 17. Ammonia ($\text{NH}_3$) boils at −33 °C, far above the expected value from the trend set by $\text{PH}_3$ (−87 °C), $\text{AsH}_3$ (−62 °C), and $\text{SbH}_3$ (−17 °C). Hydrogen fluoride ($\text{HF}$) boils at 19 °C, far above the trend from $\text{HCl}$ (−85 °C), $\text{HBr}$ (−67 °C), and $\text{HI}$ (−35 °C).

In every case, the anomalously high boiling point belongs to the molecule that can form hydrogen bonds: $\text{H}_2\text{O}$ (O–H bonds), $\text{NH}_3$ (N–H bonds), and $\text{HF}$ (F–H bonds). The heavier hydrides in each group lack these bonds and follow the expected London dispersion trend. The extra energy needed to boil water, ammonia, or hydrogen fluoride is the energy required to break the hydrogen bonds between molecules.

6. The Low Density of Ice — A Structural Consequence

Most substances are denser as solids than as liquids — the particles pack more closely together when they freeze. Water does the opposite: ice is less dense than liquid water, which is why ice floats. This is a direct structural consequence of hydrogen bonding.

In liquid water, hydrogen bonds are constantly forming and breaking. The molecules are disordered and can pack relatively closely together, giving liquid water its maximum density of approximately 1.00 g cm$^{-3}$ at 4 °C. When water freezes, the molecules lock into a rigid, open hexagonal lattice. Each water molecule forms four hydrogen bonds directed tetrahedrally — two donated via its O–H bonds and two accepted via its lone pairs. This fixed geometry forces the molecules to sit further apart than they do in the liquid, creating a more open structure with a lower density (approximately 0.92 g cm$^{-3}$).

When ice melts, some of those hydrogen bonds break. The rigid lattice collapses partially, allowing molecules to pack more closely together — so liquid water at 4 °C is actually denser than ice. This is why ice floats on water, which has profound consequences for aquatic life: a layer of ice insulates the water below, preventing lakes and rivers from freezing solid in winter.

7. Exam Technique and Common Pitfalls

AQA questions on hydrogen bonding tend to fall into three categories: drawing hydrogen bonds correctly in diagrams, explaining anomalous boiling points using trend data, and explaining the density anomaly of ice. Each has its own mark-winning language and its own traps.

Worked example: Answering an 'Explain the anomalous boiling point' Question

The boiling point of NH₃ is −33 °C. The boiling points of PH₃, AsH₃, and SbH₃ are −87 °C, −55 °C, and −17 °C respectively. Explain why the boiling point of NH₃ is anomalously high. [4 marks]

1. PH₃ (−87 °C) → AsH₃ (−55 °C) → SbH₃ (−17 °C): boiling point increases as molecular mass increases. — [B1] Stating the expected trend earns a mark. Examiners want you to show you understand what the 'anomaly' is being compared against. The increasing boiling points down the group reflect increasing London dispersion forces as the electron cloud grows larger.
2. NH₃ contains N–H bonds. Nitrogen is highly electronegative (N, O, F are the three atoms that enable hydrogen bonding). The N–H bond is significantly polarised, giving H a δ+ charge. — [B1] Naming the specific bond type (N–H) and linking it to electronegativity is required for the mark. Do not write 'hydrogen bonding exists' without explaining why it can form.
3. The δ+ hydrogen on one NH₃ molecule is attracted to the lone pair on the nitrogen of a neighbouring NH₃ molecule, forming a hydrogen bond. — [B1] The explanation must include: δ+ H, lone pair on N, and the electrostatic attraction between them. Omitting the lone pair is a common reason for losing this mark.
4. Hydrogen bonds are stronger than the London dispersion forces present between PH₃, AsH₃, and SbH₃ molecules. More energy is required to overcome these forces in NH₃, so its boiling point is higher than expected from the trend. — [B1] The comparative language ('stronger than', 'more energy required') is what earns the final mark. A statement that hydrogen bonds exist without linking to energy and boiling point will not score.

Answer: The boiling point of NH₃ is anomalously high because it forms hydrogen bonds between molecules (via the δ+ H and the lone pair on N of a neighbouring molecule). These are stronger than the London dispersion forces present in PH₃, AsH₃, and SbH₃, so more energy is required to separate NH₃ molecules, giving a higher boiling point than the trend predicts. [4 marks]

When drawing hydrogen bonds in diagrams, always show the dotted line from the δ+ H to the lone pair on the electronegative atom of the neighbouring molecule — not from atom to atom in a vague way. The direction matters, and examiners look for it.

Across this lesson, every anomalous property of water — its unexpectedly high boiling point, its lower density as a solid, and its ability to act as both a hydrogen bond donor and acceptor — traces back to the same root cause: the highly electronegative oxygen atom polarising the O–H bonds strongly enough to create persistent, directional hydrogen bonds between molecules. Master that one idea and the rest follows logically.