Energy transfer during exothermic and endothermic reactions — AQA GCSE Chemistry

Energy conservation means the total energy of a closed system stays constant during any chemical change. In a reaction, bonds in reactants break, absorbing energy, and new bonds in products form, releasing energy. The difference between these energy changes appears as energy transferred to or from the surroundings, often as heat. For example, burning methane releases energy to the surroundings, but the total energy of the universe remains unchanged; the chemical store decreases while the thermal store of the surroundings increases by the same amount. Students should track energy stores rather than treating energy as used up. This principle underpins exothermic and endothermic definitions and reaction profile diagrams.

This statement applies energy conservation to the whole universe, treated as a closed system. During a reaction, energy may move between the reacting chemicals and the surroundings, but the total energy of the universe remains constant. For example, in an exothermic reaction, the chemical store of the reactants decreases, while the thermal store of the surroundings increases by the same amount, so the total is unchanged. Students should avoid saying energy is lost; instead, they should describe energy transfer between stores. This idea supports understanding of exothermic and endothermic reactions and reaction profile diagrams, where the energy difference is shown as a transfer to or from the surroundings.

In an exothermic reaction, the products have lower energy than the reactants because energy is transferred to the surroundings. The difference in energy between reactants and products equals the energy transferred. For example, in the combustion of methane, the chemical energy stored in the reactants is greater than that in the products; the excess is released as heat, raising the temperature of the surroundings. Students should link this to reaction profile diagrams, where the products are shown at a lower energy level than the reactants, and the energy difference represents the energy transferred to the surroundings. This concept explains why exothermic reactions cause temperature increases.

An exothermic reaction transfers energy to the surroundings, causing a temperature rise. This occurs because the energy released when new bonds form in the products is greater than the energy absorbed when bonds break in the reactants. The excess energy is transferred to the surroundings, often as heat, which increases the temperature. Common examples include combustion, respiration, and neutralisation. Students should be able to identify exothermic reactions from temperature changes and explain the energy transfer in terms of bond breaking and bond making. They should also be able to represent the energy change on a reaction profile diagram, with products at a lower energy level than reactants.

Exothermic reactions transfer energy to the surroundings, so the surroundings warm up. Combustion is burning a fuel in oxygen, for example methane + oxygen → carbon dioxide + water, releasing energy as bonds form in the products. Many oxidation reactions release energy, such as reacting magnesium with oxygen or iron rusting slowly. Neutralisation is the reaction of an acid with a base or alkali to form a salt and water, for example hydrochloric acid + sodium hydroxide → sodium chloride + water, and the mixture warms. In each case, energy released when new bonds form exceeds energy taken in to break bonds, so overall energy is transferred out. Temperature rise of the surroundings is the observable evidence.

Self-heating cans and hand warmers use exothermic reactions to release useful heat. In a self-heating can, pressing a button or pulling a tab mixes reactants that then react exothermically, warming the drink or food. Hand warmers often use the oxidation of iron in air; when the sealed packet is opened, air enters and iron oxidises, releasing energy that warms the hands. The design controls the rate so heat is released safely over a useful time. These are everyday applications of exothermic change: energy transferred to the surroundings is the desired product. The reaction is not used to make a new substance for its own sake; the heat is the purpose.

An endothermic reaction takes in energy from the surroundings, so the surroundings cool down. Energy is transferred from the surroundings to the reacting system. In bond terms, the energy needed to break bonds in the reactants is greater than the energy released when new bonds form in the products, so the overall change absorbs energy. A practical example is the reaction of citric acid with sodium hydrogencarbonate, where the mixture temperature falls. Thermal decomposition of calcium carbonate also requires continuous heating. The observable evidence is a temperature decrease of the surroundings, measured with a thermometer or temperature sensor.

Thermal decomposition means breaking a compound down by heating, for example heating calcium carbonate to form calcium oxide and carbon dioxide; energy must be taken in continuously, so it is endothermic. The reaction of citric acid with sodium hydrogencarbonate is also endothermic: when the two solids or solutions mix, they fizz to produce carbon dioxide and the mixture cools. In both cases energy is taken in from the surroundings, so the temperature of the surroundings falls. These examples show that endothermic reactions are not limited to one type of chemistry; they include decomposition driven by heating and a reaction between an acid and a hydrogencarbonate.

An endothermic reaction takes in energy from its surroundings, so the surroundings cool. A sports injury pack uses this: when the pack is activated, an endothermic reaction absorbs thermal energy from the injured tissue, producing a cold pack that reduces swelling and pain. The reaction mixture and the pack feel cold because energy is transferred from the surroundings to the reacting system. For example, dissolving ammonium nitrate in water in a sealed pack absorbs energy, lowering the temperature. This is the opposite of an exothermic reaction, which releases energy and warms the surroundings. The pack must be designed so the reactants mix only when needed and the cold is safely delivered to the skin without causing ice burns.

Exothermic reactions release energy to the surroundings, so the surroundings warm; endothermic reactions absorb energy from the surroundings, so the surroundings cool. To distinguish them, measure the temperature change of the surroundings: a rise indicates exothermic, a fall indicates endothermic. For example, adding magnesium to hydrochloric acid raises the temperature, while dissolving ammonium nitrate in water lowers it. When evaluating uses, consider whether the energy transfer is useful, safe and controllable. Exothermic reactions suit hand warmers and self-heating cans; endothermic reactions suit cold packs. Evaluation should weigh benefits such as portability against risks such as burns, cost, reversibility and environmental impact.

In this topic, practical work and questions focus on measuring temperature change, not on calculating energy changes or enthalpy change, ΔH. You record the initial temperature of the surroundings or solution, carry out the reaction, and record the highest or lowest temperature reached. The difference between these temperatures shows whether the reaction is exothermic or endothermic and how large the change is. For example, measure the temperature before and after adding magnesium to copper sulfate solution, then subtract the initial reading from the final reading. You are not required to convert temperature change into energy in joules or to calculate ΔH. This keeps the focus on observing and interpreting energy transfer.

In Required practical 4, you investigate variables that affect temperature changes in reacting solutions. Suitable reactions include acid plus metals, acid plus carbonates, neutralisation and displacement of metals. You measure the temperature change of the solution before and after reaction, then vary one factor at a time, such as concentration, volume, particle size or the metal used. For example, react different concentrations of hydrochloric acid with magnesium and record the maximum temperature rise. Control other variables to make the comparison fair. The practical develops skills in measuring temperature, controlling variables, recording data and drawing conclusions about energy transfer.