Energy Changes & States of Matter Revision Notes | A Level AQA Chemistry

1. Why Energy and States of Matter Matter

Picture a block of ice sitting on a warm surface. Energy flows in, and eventually the ice melts — but the temperature stays at 0 °C the whole time. Where is that energy going? The answer lies in the intermolecular forces holding the water molecules together, and understanding this is central to AQA A Level Physical Chemistry.

Changes of state are not about breaking covalent bonds. They are about overcoming or forming the intermolecular forces between molecules. Getting this distinction right is the difference between a confident exam answer and a costly misconception.

2. The Three States and What Holds Them Together

In a solid, particles are held in fixed positions by strong intermolecular forces. They vibrate but cannot move past one another. In a liquid, those forces are partially overcome — particles can flow but remain close together. In a gas, intermolecular forces are negligible; particles move freely and are widely separated.

This distinction is critical. When water boils, the O–H covalent bonds inside each water molecule remain intact. What breaks are the hydrogen bonds between water molecules. Examiners test this distinction directly, so it is worth anchoring firmly now.

3. Energy Changes During Changes of State

Every change of state involves an enthalpy change. The direction — endothermic or exothermic — depends on whether intermolecular forces are being broken or formed. Breaking intermolecular forces requires energy input (endothermic, $\Delta H > 0$); forming them releases energy (exothermic, $\Delta H < 0$).

Think of intermolecular forces like stretched springs between molecules. Pulling molecules apart (melting, boiling, sublimation) requires you to do work — energy goes in. Letting molecules come together (condensation, freezing, deposition) releases that stored energy — energy comes out.

The chart above uses +1 for endothermic and −1 for exothermic as a visual guide. Notice the perfect symmetry: every forward process (solid → liquid → gas) is endothermic, and every reverse process (gas → liquid → solid) is exothermic by exactly the same magnitude.

Freezing is the reverse of melting, so its enthalpy change has the same magnitude but opposite sign: $\Delta_{freeze}H = -\Delta_{fus}H$. This follows directly from Hess's Law — enthalpy is a state function, so the energy released on freezing equals the energy absorbed on melting.

4. Comparing Enthalpy of Fusion and Enthalpy of Vaporisation

Here is something that surprises many students: the enthalpy of vaporisation is always much larger than the enthalpy of fusion for the same substance. For water, $\Delta_{fus}H = +6.0 \text{ kJ mol}^{-1}$ whereas $\Delta_{vap}H = +41 \text{ kJ mol}^{-1}$ — roughly seven times larger.

The reason is particle-level. When a solid melts, intermolecular forces are only partially overcome — molecules gain enough energy to move past one another, but they remain close and many interactions persist. When a liquid vaporises, all remaining intermolecular forces must be completely broken so that molecules become independent in the gas phase. That requires far more energy.

Sublimation — the direct conversion of solid to gas — has the largest enthalpy change of all, approximately equal to the sum of the enthalpy of fusion and the enthalpy of vaporisation. This makes sense: the substance must traverse both the solid-to-liquid and liquid-to-gas energy barriers in one step.

5. How Intermolecular Force Strength Affects Energy Changes

The stronger the intermolecular forces in a substance, the more energy is needed to overcome them — and therefore the larger the enthalpy of fusion and vaporisation, and the higher the melting and boiling points. This connects directly to the types of intermolecular forces present.

Hydrogen bonding is the strongest type of intermolecular force for molecules of comparable size at A Level, though London dispersion forces in very large molecules can become dominant. Substances with hydrogen bonding (such as water, ammonia, and hydrogen fluoride) have anomalously high boiling points and large enthalpies of vaporisation compared with molecules of similar molar mass that only exhibit London dispersion forces.

• London dispersion forces only (e.g. noble gases, non-polar molecules): weakest intermolecular forces, lowest $\Delta_{vap}H$, lowest boiling points
• Permanent dipole–dipole interactions (e.g. HCl, propanone): moderate intermolecular forces, intermediate $\Delta_{vap}H$
• Hydrogen bonding (e.g. $H_2O$, $NH_3$, $HF$, alcohols, carboxylic acids): strongest intermolecular forces, highest $\Delta_{vap}H$ relative to molar mass

Within a homologous series where only London dispersion forces operate — such as the alkanes — both the boiling point and the enthalpy of vaporisation increase with chain length. Longer chains have greater surface area, allowing more contact points between molecules and therefore stronger cumulative London dispersion forces.

6. Temperature During a Change of State

One of the most counterintuitive aspects of changes of state is that temperature remains constant while the change is occurring, even though energy is being supplied or removed. This is because all the energy goes into breaking (or is released by forming) intermolecular forces, rather than increasing the kinetic energy of the particles.

A heating curve illustrates this clearly. As a solid is heated, temperature rises until the melting point is reached. At the melting point, the temperature plateaus — energy input goes entirely into overcoming intermolecular forces (the enthalpy of fusion). Once all the solid has melted, temperature rises again through the liquid phase until the boiling point, where another plateau occurs (the enthalpy of vaporisation).

The two flat sections on the heating curve correspond to the melting point and boiling point respectively. The length of each plateau is proportional to the enthalpy change for that transition — so the boiling plateau is longer than the melting plateau, reflecting the larger enthalpy of vaporisation.

7. Worked Example and Exam Technique

Exam questions on this topic often use the command words 'explain' or 'compare'. The key to full marks is always linking your answer to intermolecular forces at the particle level — not stopping at 'energy is needed to melt it'. Let's walk through a typical A Level question.

Worked example: Explaining Energy Changes During Vaporisation

Explain why vaporisation of a liquid is an endothermic process, and why the enthalpy of vaporisation of water is greater than that of hydrogen sulfide ($H_2S$). [4 marks]

1. Vaporisation: liquid → gas. $\Delta H > 0$ (endothermic). — State the direction of enthalpy change first — this earns the first mark and frames the rest of the answer.
2. Energy must be supplied to break the intermolecular forces between molecules in the liquid phase. — Use the phrase 'intermolecular forces' explicitly. Saying 'bonds' without specifying intermolecular is too vague and may not earn the mark.
3. $H_2O$ molecules form hydrogen bonds with one another (O–H···O). $H_2S$ molecules only exhibit London dispersion forces and weak permanent dipole–dipole interactions — hydrogen bonding does not occur because S is not electronegative enough. — Name the specific type of intermolecular force in each substance. This is the comparative element the question demands.
4. Hydrogen bonds are stronger than London dispersion forces and dipole–dipole interactions. Therefore more energy is required to break the intermolecular forces in water than in $H_2S$. — Always close the argument by linking force strength directly to the magnitude of $\Delta_{vap}H$. This is the mark-winning step.

Answer: Vaporisation is endothermic ($\Delta H > 0$) because energy must be supplied to break intermolecular forces between liquid molecules. Water has a higher enthalpy of vaporisation than $H_2S$ because water molecules form hydrogen bonds, which are stronger than the London dispersion forces and dipole–dipole interactions present in $H_2S$. More energy is therefore required to overcome the intermolecular forces in water.

Notice the structure of that model answer: state the energy direction, name the forces being broken, compare the force strengths, and link to the enthalpy magnitude. This four-beat pattern works for almost every 'explain' question on changes of state.

8. Exam Technique and Common Mistakes

The core spec requirement for this topic is to explain the energy changes associated with changes of state. Every mark-winning answer returns to the same foundation: intermolecular forces are broken (endothermic) or formed (exothermic), and the magnitude of the energy change reflects the strength and number of those forces.