Enthalpy Changes Revision Notes | A Level AQA Chemistry

1. Why Enthalpy Changes Matter

Every chemical reaction involves an energy change. Whether you are burning fuel in an engine or dissolving a painkiller in water, energy is either released to the surroundings or absorbed from them. Understanding and quantifying that energy change is at the heart of energetics.

At A Level, the key quantity is enthalpy change, symbol \(\Delta H\). Rather than measuring total energy stored in a system — which is practically impossible — chemists measure the change in heat content at constant pressure. That is exactly what \(\Delta H\) captures.

The sign of \(\Delta H\) tells you the direction of energy flow. A negative value means energy leaves the system and warms the surroundings; a positive value means energy enters the system from the surroundings. This sign convention is non-negotiable in AQA mark schemes.

2. Exothermic and Endothermic Reactions

Picture a hand-warmer pack. When you crack it, the contents react and the pack heats up — energy is released to your hand. That is an exothermic reaction. Now picture the cold pack used for sports injuries: it draws heat away from your skin, cooling it down. That is an endothermic reaction.

In an exothermic reaction, the energy released when new bonds form in the products is greater than the energy required to break bonds in the reactants. The surplus energy is transferred to the surroundings as heat, raising the temperature of the reaction mixture.

In an endothermic reaction, the energy required to break bonds in the reactants is greater than the energy released when new bonds form. The shortfall is drawn in from the surroundings, so the reaction mixture cools down.

3. Enthalpy Change and Constant Pressure

You might wonder why chemists use enthalpy rather than measuring temperature change. The answer lies in the conditions of most laboratory reactions: they happen in open containers at atmospheric pressure, not in sealed, rigid vessels. At constant pressure, the heat exchanged between the system and the surroundings equals the enthalpy change, \(\Delta H\).

The absolute enthalpy of a substance cannot be measured directly — only the change can be determined experimentally. This is why \(\Delta H\) is always defined as a difference, and why the sign carries so much meaning in AQA questions.

4. Standard Conditions and Standard Enthalpy Changes

Enthalpy changes depend on temperature and pressure. To make values comparable across different laboratories and data sources, chemists define a set of standard conditions. Under AQA specifications, standard conditions are a pressure of 100 kPa and a stated temperature, most commonly 298 K (25 °C).

When a measurement is made under standard conditions, the enthalpy change is given the symbol \(\Delta H^{\ominus}\). The subscript 298 is added when the temperature is specifically 298 K, giving \(\Delta H^{\ominus}_{298}\). The superscript \(\ominus\) (the standard symbol, sometimes written as a circle with a horizontal line) signals that standard conditions apply.

A common point of confusion: standard conditions do not mean the reaction must be carried out at 298 K. They mean the measurement is referenced to those conditions. For example, the standard enthalpy of combustion of methane is defined at 298 K and 100 kPa, even though the actual flame is far hotter.

Notice that 100 kPa is the standard pressure used by AQA — not 101.3 kPa (1 atm), which you may encounter in older textbooks. This distinction rarely affects numerical answers at this level, but using the correct value in written explanations is good exam practice.

5. Reaction Profile Diagrams

A reaction profile diagram gives you a visual representation of how enthalpy changes as reactants convert to products. The y-axis shows enthalpy; the x-axis shows the progress of the reaction (sometimes called the reaction coordinate). The diagram is not a graph of concentration against time — it is an energy landscape.

In an exothermic profile, the reactants sit at a higher enthalpy level than the products. The curve rises to a peak — the transition state — before falling to the product level. The height of that peak above the reactants is the activation energy, \(E_a\). The overall drop from reactants to products is \(\Delta H\), which is negative.

For an endothermic reaction, the profile is inverted: products sit at a higher enthalpy level than reactants, and \(\Delta H\) is positive. The activation energy is still the height from the reactant level to the peak of the curve — not from the product level.

• Exothermic: reactants higher than products on the diagram; \(\Delta H < 0\); the gap between reactant and product levels equals \(|\Delta H|\)
• Endothermic: products higher than reactants on the diagram; \(\Delta H > 0\); the gap between reactant and product levels equals \(\Delta H\)
• Activation energy (\(E_a\)): always measured from the reactant enthalpy level to the peak of the curve, regardless of reaction type
• A catalyst lowers \(E_a\) but does not change \(\Delta H\) — the start and end levels remain the same

6. Worked Calculation: Using \(\Delta H\) in Context

Exam questions often ask you to interpret or calculate enthalpy changes from experimental data. The most common route at this stage is using temperature change data from a calorimetry experiment. The worked example below walks through the method step by step.

Worked example: Calculating Enthalpy Change from Calorimetry Data

25.0 cm\(^3\) of 1.00 mol dm\(^{-3}\) hydrochloric acid is mixed with 25.0 cm\(^3\) of 1.00 mol dm\(^{-3}\) sodium hydroxide solution. The temperature rises by 6.8 °C. Calculate the enthalpy change of neutralisation in kJ mol\(^{-1}\). (Assume the density of the solution is 1.00 g cm\(^{-3}\) and the specific heat capacity is 4.18 J g\(^{-1}\) °C\(^{-1}\).)

1. Total volume = 25.0 + 25.0 = 50.0 cm\(^3\) Mass = 50.0 cm\(^3\) × 1.00 g cm\(^{-3}\) = 50.0 g — Both solutions are mixed, so their volumes add. Using density = 1.00 g cm\(^{-3}\) converts volume directly to mass.
2. \(q = mc\Delta T\) \(q = 50.0 \text{ g} \times 4.18 \text{ J g}^{-1}\text{°C}^{-1} \times 6.8 \text{ °C}\) \(q = 50.0 \times 4.18 \times 6.8\) \(50.0 \times 4.18 = 209.0\) \(209.0 \times 6.8 = 1421.2 \text{ J}\) — \(q = mc\Delta T\) gives the heat energy in joules. Each multiplication is carried out separately to avoid errors.
3. \(q = 1421.2 \text{ J} \div 1000 = 1.4212 \text{ kJ}\) — Enthalpy changes are reported in kJ mol\(^{-1}\), so convert from J to kJ by dividing by 1000.
4. Moles of HCl = 1.00 mol dm\(^{-3}\) × 0.0250 dm\(^3\) = 0.0250 mol Moles of NaOH = 1.00 mol dm\(^{-3}\) × 0.0250 dm\(^3\) = 0.0250 mol Moles of water formed = 0.0250 mol (1:1 ratio) — Note: 25.0 cm\(^3\) = 0.0250 dm\(^3\). The neutralisation reaction is 1:1, so 0.0250 mol of water is produced.
5. \(\Delta H = -\frac{q}{n} = -\frac{1.4212 \text{ kJ}}{0.0250 \text{ mol}} = -56.8 \text{ kJ mol}^{-1}\) — The negative sign is applied because the temperature rose — the reaction is exothermic and energy was released to the surroundings. Always assign the sign based on the direction of energy transfer, not just the arithmetic.
6. The literature value for strong acid–strong base neutralisation is approximately \(-57 \text{ kJ mol}^{-1}\). Our answer: \(-56.8 \text{ kJ mol}^{-1}\) — consistent. — Checking against a known value confirms the method and sign are correct. A positive answer here would signal a sign error.

Answer: \(\Delta H_{\text{neutralisation}} = -56.8 \text{ kJ mol}^{-1}\)

7. Exam Technique and Common Mistakes

The specification points on enthalpy changes appear in almost every AQA paper, and the marks are often lost not through lack of knowledge but through imprecise language or sign errors. The guidance below targets the most frequent pitfalls seen in mark schemes.

When an exam question uses the command word explain, you must give a reason, not a description. For example, 'the reaction is exothermic' is a description; 'the reaction is exothermic because more energy is released when bonds form in the products than is required to break bonds in the reactants' is an explanation and will score the mark.

Across all three specification points — the distinction between exothermic and endothermic, the definition of enthalpy change at constant pressure, and the meaning of standard conditions — the thread connecting them is the sign and symbol of \(\Delta H^{\ominus}\). Master that notation and you have the foundation for every energetics topic that follows, from Hess's Law to Born-Haber cycles.