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WASSCE Chemistry · Energy Changes in Reactions

Exothermic and endothermic reactions

By the WASSCE Tutor team

Updated 12 November 2026 · 8 min read · Prepared against the WASSCE syllabus.

Charcoal glowing in a coal pot gives out heat; an instant cold pack in a first-aid kit takes heat in. Almost every chemical reaction involves an energy change, usually as heat. This lesson explains exothermic and endothermic reactions, enthalpy changes, energy profile diagrams, and how to measure heats of reaction in the laboratory.

Learning objectives

By the end of this lesson you should be able to:

  • distinguish exothermic and endothermic reactions, with examples
  • define enthalpy change (ΔH) and use its sign convention
  • draw and interpret energy profile diagrams, including activation energy
  • define standard enthalpies of combustion, formation, neutralisation and solution
  • calculate enthalpy changes from experimental data using q = mcΔT

Exothermic and endothermic reactions

ExothermicEndothermic
Gives out heat to the surroundingsTakes in heat from the surroundings
Surroundings get warmerSurroundings get colder
ΔH is negativeΔH is positive
Products have less energy than reactantsProducts have more energy than reactants
Combustion of fuels (charcoal, kerosene, cooking gas); respiration; neutralisation; rusting; dissolving NaOH or concentrated H₂SO₄ in waterPhotosynthesis; thermal decomposition of CaCO₃; dissolving ammonium nitrate or ammonium chloride in water; melting and boiling

Enthalpy change

Definition

The enthalpy change (ΔH) of a reaction is the heat energy given out or taken in at constant pressure, measured in kJ/mol (kJ mol⁻¹). It is negative for exothermic and positive for endothermic reactions.

A thermochemical equation shows ΔH with the balanced equation and state symbols:

CH4(g)+2O2(g)→CO2(g)+2H2O(l)ΔH=−890 kJ/mol\text{CH}_4\text{(g)} + 2\text{O}_2\text{(g)} \rightarrow \text{CO}_2\text{(g)} + 2\text{H}_2\text{O(l)} \qquad \Delta H = -890 \text{ kJ/mol}
CH4(g)+2O2(g)→CO2(g)+2H2O(l)ΔH=−890 kJ/mol\begin{gathered}\text{CH}_4\text{(g)} + 2\text{O}_2\text{(g)} \\[6pt] \rightarrow \text{CO}_2\text{(g)} + 2\text{H}_2\text{O(l)} \\[6pt] \begin{aligned}\Delta H &= -890 \text{ kJ/mol}\end{aligned}\end{gathered}
CaCO3(s)→CaO(s)+CO2(g)ΔH=+178 kJ/mol\text{CaCO}_3\text{(s)} \rightarrow \text{CaO(s)} + \text{CO}_2\text{(g)} \qquad \Delta H = +178 \text{ kJ/mol}
CaCO3(s)→CaO(s)+CO2(g)ΔH=+178 kJ/mol\begin{gathered}\text{CaCO}_3\text{(s)} \\[6pt] \rightarrow \text{CaO(s)} + \text{CO}_2\text{(g)} \\[6pt] \begin{aligned}\Delta H &= +178 \text{ kJ/mol}\end{aligned}\end{gathered}

Energy profile diagrams

Energy profiles: in an exothermic reaction the products are lower than the reactants (ΔH negative); in an endothermic reaction they are higher (ΔH positive); both need activation energy Ea to start

Figure 1. Energy profiles of exothermic and endothermic reactions. WASSCE Tutor diagram.

Definition

Activation energy (Ea) is the minimum energy that colliding particles must have for a reaction to take place.

Even exothermic reactions need energy to start. A match must be struck, and charcoal must be lit, to provide the activation energy. Once started, the heat given out keeps the reaction going.

Standard enthalpy changes

Standard conditions are 25 °C (298 K) and 1 atm, with substances in their normal states.

Enthalpy changeDefinitionExample
Combustion, ΔH°cHeat given out when 1 mole of a substance burns completely in oxygenC(s) + O₂ → CO₂; ΔH = −394 kJ/mol
Formation, ΔH°fHeat change when 1 mole of a compound forms from its elements in their standard states½H₂ + ½Cl₂ → HCl; ΔH = −92 kJ/mol
NeutralisationHeat given out when an acid and a base react to form 1 mole of waterH⁺ + OH⁻ → H₂O; ΔH = −57.3 kJ/mol
SolutionHeat change when 1 mole of a substance dissolves in a large amount of waterNH₄NO₃(s) → NH₄NO₃(aq); ΔH positive

The heat of neutralisation of any strong acid by any strong base is about −57.3 kJ/mol, because the reaction is always H⁺ + OH⁻ → H₂O. With a weak acid or base it is less exothermic, because some energy is used to ionise the weak acid fully.

Measuring enthalpy changes

Heat absorbed by water:

q=mcΔTq = mc\Delta T
  • m = mass of water (or solution) in g (1 cm³ of water ≈ 1 g)
  • c = specific heat capacity of water, 4.2 J g⁻¹ °C⁻¹
  • ΔT = temperature change in °C

Then ΔH=−qn\Delta H = -\dfrac{q}{n}, where n is the number of moles reacted.

Worked example

Heat of neutralisation. 50 cm³ of 1.0 mol/dm³ HCl is mixed with 50 cm³ of 1.0 mol/dm³ NaOH in a plastic cup. The temperature rises by 6.8 °C.

  1. Mass of solution =50+50=100= 50 + 50 = 100 g.
  2. q=100×4.2×6.8=2856q = 100 \times 4.2 \times 6.8 = 2856 J =2.856= 2.856 kJ.
  3. Moles of water formed =1.0×50÷1000=0.050= 1.0 \times 50 \div 1000 = 0.050 mol.
  4. ΔH=−2.8560.050=−57.1\Delta H = -\dfrac{2.856}{0.050} = -57.1 kJ/mol.
Worked example

Heat of combustion of ethanol. A spirit lamp burning 0.46 g of ethanol (M = 46) heats 200 g of water in a copper can by 15 °C.

  1. q=200×4.2×15=12 600q = 200 \times 4.2 \times 15 = 12\,600 J =12.6= 12.6 kJ.
  2. Moles of ethanol =0.46÷46=0.010= 0.46 \div 46 = 0.010 mol.
  3. ΔHc=−12.60.010=−1260\Delta H_c = -\dfrac{12.6}{0.010} = -1260 kJ/mol.

The data-book value is −1367 kJ/mol. The experimental value is lower because of heat lost to the air and the can, and incomplete combustion (soot on the can).

Improving accuracy: insulate the container (plastic cup with lid), shield the flame from draughts, put the can close to the flame, stir, and record the highest temperature reached.

In Sierra Leone

Choosing a cooking fuel is an enthalpy question. Charcoal, firewood and kerosene used in Sierra Leonean kitchens all burn exothermically. Their heats of combustion per gram, and how efficiently a stove transfers that heat to the pot, decide how much fuel a family must buy. Improved clay or metal stoves waste less heat than open three-stone fires, so they save fuel and reduce smoke.

Common mistake

A negative ΔH means heat is given out (exothermic), even though the temperature of the surroundings rises. Do not confuse the sign of ΔH with the direction of the temperature change.

Key points

  • Exothermic: heat given out, ΔH negative (combustion, neutralisation, respiration). Endothermic: heat taken in, ΔH positive (photosynthesis, decomposition of CaCO₃).
  • Activation energy is the minimum energy needed for a reaction; it appears as the peak of the energy profile.
  • Standard enthalpies: combustion, formation, neutralisation (−57.3 kJ/mol for strong acid + strong base), solution.
  • q = mcΔT with c = 4.2 J g⁻¹ °C⁻¹; ΔH = −q ÷ n.

WASSCE examination focus

WAEC tests this

Expect to classify reactions as exothermic or endothermic, to draw energy profiles labelled with Ea and ΔH, to define standard enthalpies, and to calculate ΔH from temperature-change data. Questions often ask why experimental values are lower than data-book values.

Quick check

  1. Is the thermal decomposition of calcium carbonate exothermic or endothermic?
  2. What is activation energy?
  3. Why is the heat of neutralisation of all strong acids by strong bases about the same?
  4. 25 cm³ of acid and 25 cm³ of alkali (each 2.0 mol/dm³) react and the temperature rises by 13.6 °C. Find ΔH of neutralisation.

Answers:

  1. Endothermic.
  2. The minimum energy colliding particles must have for a reaction to occur.
  3. In every case the reaction is H⁺ + OH⁻ → H₂O.
  4. q=50×4.2×13.6=2856q = 50 \times 4.2 \times 13.6 = 2856 J; moles of water =2.0×25÷1000=0.050= 2.0 \times 25 \div 1000 = 0.050; ΔH=−2.856÷0.050=−57.1\Delta H = -2.856 \div 0.050 = -57.1 kJ/mol.

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