Energy Changes — AQA Combined Science: Trilogy

All chemical reactions involve energy changes. This unit covers exothermic and endothermic reactions, reaction profiles and (at Higher Tier) bond energy calculations.

Exothermic and endothermic reactions

• Exothermic reactions transfer energy to the surroundings, usually as heat, so the temperature of the surroundings increases. Examples: combustion, neutralisation, oxidation, and many displacement reactions. Everyday uses: hand warmers and self-heating cans.
• Endothermic reactions take in energy from the surroundings, so the temperature decreases. Examples: thermal decomposition (e.g. heating metal carbonates) and the reaction of citric acid with sodium hydrogencarbonate. Everyday uses: sports injury cold packs.

Required practical: investigating the temperature change in reactions, e.g. neutralisation, dissolving salts, or displacement reactions, by measuring temperature before and after with a thermometer.

Conservation of energy

The total amount of energy is conserved in a chemical reaction — energy is transferred to or from the surroundings, not created or destroyed.

Reaction profiles

A reaction profile is a diagram showing the energy of the reactants and products as a reaction proceeds.

• In an exothermic reaction, the products are at a lower energy than the reactants — energy is released.
• In an endothermic reaction, the products are at a higher energy than the reactants — energy is absorbed.

Activation energy

The activation energy is the minimum energy that colliding particles must have for a reaction to occur. On a reaction profile it is the "hump" — the energy needed to get from reactants to products. A catalyst lowers the activation energy by providing an alternative pathway.

Bond energies (Higher Tier)

During a reaction:

• Breaking bonds is endothermic (energy must be supplied).
• Making bonds is exothermic (energy is released).

The overall energy change is:

$$\Delta E = (\text{energy to break bonds}) - (\text{energy released making bonds})$$

• If more energy is released making bonds than is used breaking them, the reaction is exothermic (ΔE is negative).
• If more energy is needed to break bonds than is released, the reaction is endothermic (ΔE is positive).

Worked example

For H₂ + Cl₂ → 2HCl:

• Bonds broken: 1 × (H–H) + 1 × (Cl–Cl).
• Bonds made: 2 × (H–Cl).
• Subtract energy released from energy supplied to find the overall change. A negative answer means the reaction is exothermic.

You may be given the bond energies in a table — read them carefully and count the bonds in each molecule.

Cells and batteries (overview)

Chemical reactions can be used to produce electricity in cells. A simple cell is made by connecting two different metals in an electrolyte; the bigger the difference in reactivity between the metals, the bigger the voltage. Batteries are two or more cells connected together. Rechargeable cells can have the reaction reversed by applying a current.

Exam tips

• Learn the definitions: exothermic releases energy (temperature rises), endothermic absorbs energy (temperature falls).
• For reaction profiles, label the reactants, products, activation energy and overall energy change.
• Remember: breaking bonds = endothermic, making bonds = exothermic ("break = take, make = give").
• In bond-energy calculations, a negative overall value means exothermic.
• Watch the maths: count every bond, including all the bonds in larger molecules.