Energy Changes in GCSE Chemistry — Exothermic, Endothermic, Bond Energies

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Energy changes accompany every chemical reaction, and understanding them is essential for GCSE Chemistry. This topic links directly to rates of reaction (via activation energy), to equilibrium (via exothermic and endothermic reversible reactions), and to real industrial processes. Questions range from drawing reaction profiles to calculating energy changes from bond energies — this guide covers all of it.

Exothermic and Endothermic Reactions

Every chemical reaction involves breaking bonds in reactants and forming bonds in products. Breaking bonds requires energy — it is always endothermic. Forming bonds releases energy — it is always exothermic. The overall energy change of a reaction depends on which process requires or releases more energy.

Exothermic reaction: More energy is released when bonds form in the products than is needed to break bonds in the reactants. The reaction releases energy to the surroundings — the temperature of the surroundings increases. Examples: combustion, neutralisation, oxidation of metals (rusting), many displacement reactions.
Endothermic reaction: More energy is needed to break bonds in the reactants than is released when bonds form in the products. The reaction absorbs energy from the surroundings — the temperature of the surroundings decreases. Examples: thermal decomposition, dissolving ammonium nitrate, photosynthesis.

Exothermic: temperature increases, energy released to surroundings, ΔH is negative. Endothermic: temperature decreases, energy absorbed from surroundings, ΔH is positive. The sign of ΔH is the clearest indicator — always state it in calculation answers.

Reaction Profile Diagrams

A reaction profile (energy profile diagram) shows the energy of reactants and products, and the activation energy needed to start the reaction. The x-axis shows the progress of the reaction; the y-axis shows energy.

For an exothermic reaction: the products are at a lower energy level than the reactants. The peak of the curve above the reactants is the activation energy. The difference in height between reactants and products is the overall energy change (ΔH), which is negative.

For an endothermic reaction: the products are at a higher energy level than the reactants. The peak is still above both reactants and products (activation energy is always positive). ΔH is positive.

The effect of a catalyst: it provides an alternative reaction pathway with a lower activation energy — the peak of the curve is lower. The overall energy change (ΔH) is unchanged — the start and end points remain the same. Only the peak height changes.

Drawing Reaction Profiles

When asked to draw a reaction profile: (1) draw axes labelled "Energy" (y) and "Progress of reaction" (x). (2) Draw reactants as a horizontal line. (3) Draw a smooth curve rising to a peak (the activation energy) then falling to the products line. (4) For exothermic, products are lower than reactants. For endothermic, products are higher. (5) Label reactants, products, activation energy and ΔH with a double-headed arrow. All five elements are typically needed for full marks.

Bond Energy Calculations

Bond energy values (given in kJ/mol) tell you how much energy is needed to break one mole of a particular type of bond. Using these values, you can calculate the overall energy change of a reaction.

ΔH = energy needed to break bonds − energy released forming bonds
= Σ(bond energies of reactants) − Σ(bond energies of products)

If the result is negative, the reaction is exothermic (more energy released than taken in). If positive, it is endothermic.

Worked Example

Calculate ΔH for: H₂ + Cl₂ → 2HCl

Bond energies: H−H = 436 kJ/mol, Cl−Cl = 243 kJ/mol, H−Cl = 432 kJ/mol

Energy to break bonds in reactants: H−H (436) + Cl−Cl (243) = 679 kJ/mol

Energy released forming bonds in products: 2 × H−Cl = 2 × 432 = 864 kJ/mol

ΔH = 679 − 864 = −185 kJ/mol (exothermic)

❌ Most common error: forgetting to multiply bond energies by the number of bonds broken or formed. In the example above, two H−Cl bonds form (2HCl), so the bond energy must be multiplied by 2. Always count the number of each bond from the structural formula, not just the molecular formula.

Reversible Reactions and Equilibrium

Some reactions are reversible — they can proceed in both the forward and reverse directions. This is shown with a double arrow (⇌). When the rate of the forward reaction equals the rate of the reverse reaction, the system is at dynamic equilibrium. The concentrations of reactants and products remain constant (though not necessarily equal).

Le Chatelier's Principle states: if a system at equilibrium is disturbed, it will respond to oppose the change and re-establish equilibrium.

Increasing temperature: The equilibrium shifts in the endothermic direction (to absorb the extra heat).
Increasing pressure: The equilibrium shifts toward the side with fewer moles of gas.
Increasing concentration of a reactant: The equilibrium shifts toward the products.
Adding a catalyst: The equilibrium position is unchanged — the catalyst speeds up both the forward and reverse reactions equally. Equilibrium is reached faster, but the same proportions of products and reactants are formed.

The Haber Process — A Case Study in Compromise

The Haber process produces ammonia from nitrogen and hydrogen. It is one of the most important industrial chemical processes — ammonia is the basis of all nitrogen-based fertilisers.

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ/mol (exothermic)

The industrial conditions used are: temperature 450°C, pressure 200 atmospheres, iron catalyst. Each condition is a compromise:

Temperature: The forward reaction is exothermic, so lower temperature favours ammonia formation (Le Chatelier). But lower temperature means slower rate. 450°C is a compromise — enough ammonia forms at an acceptable rate.
Pressure: There are 4 moles of gas on the left and 2 on the right, so higher pressure favours ammonia (fewer gas molecules). Very high pressure is expensive and dangerous. 200 atm is a compromise.
Catalyst: Iron catalyst increases rate without affecting equilibrium position, allowing the compromise temperature to still produce ammonia at a useful rate.

The AQA energy changes specification is at the AQA GCSE Chemistry specification page.

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