Chemistry · Lesson 9

Energy changes & thermochemistry

JAMBWAECNECO

What the syllabus asks

JAMB: Energy changes: energy changes accompanying chemical reactions, ΔH, entropy, free energy and spontaneity

WAEC: Energy changes: exothermic and endothermic reactions, enthalpy change (ΔH), energy profile diagrams, activation energy, heats of reaction (formation, combustion, neutralisation, solution), bond energies, Hess’s law, entropy and free energy (ΔG = ΔH − TΔS), spontaneity

NECO: Energy changes: exothermic and endothermic reactions, enthalpy change (ΔH), energy profile diagrams, activation energy, heats of reaction (formation, combustion, neutralisation, solution), bond energies, Hess’s law, entropy and free energy (ΔG = ΔH − TΔS), spontaneity

What you will learn

Distinguish exothermic and endothermic reactions with everyday examples.
Draw and interpret energy profile diagrams, showing ΔH and activation energy.
Calculate heat changes using q = mcΔT and convert them to ΔH per mole.
Define and calculate heats of combustion, formation, neutralisation and solution.
Use bond energies to estimate ΔH.
Apply Hess’s law and heats of formation to find unknown enthalpy changes.
Explain entropy and use ΔG = ΔH − TΔS to predict spontaneity.

Key points

ΔH = H(products) − H(reactants); negative = exothermic, positive = endothermic.
Energy profiles show Ea (reactants to peak) and ΔH; catalysts lower Ea only.
q = mcΔT; ΔH = −q/n for a temperature rise, +q/n for a fall.
Strong acid + strong base: ΔH(neut) ≈ −57 kJ mol⁻¹; weak acids/bases give less.
ΔH = bonds broken − bonds formed; Hess’s law: ΔH = ΣΔHf(products) − ΣΔHf(reactants).
ΔG = ΔH − TΔS; spontaneous when ΔG < 0; T = ΔH/ΔS at the turning point.

Start here: Energy in chemical reactions

Every chemical reaction involves an energy change. When a woman in Onitsha lights a kerosene stove, the kerosene burns and gives out heat for cooking. When you dissolve some ammonium nitrate fertiliser in water, the container feels cold. In both cases chemical energy is being converted: in the first, stored energy is released as heat; in the second, heat is taken in from the surroundings.

Chemists call the heat content of a substance its enthalpy, H. We cannot measure H directly, but we can measure the change in enthalpy, ΔH (read "delta H"), during a reaction. ΔH is the heat given out or absorbed when the reaction occurs at constant pressure, usually quoted in kilojoules per mole (kJ mol⁻¹).

ΔH = H(products) − H(reactants)

This simple subtraction is the key to the whole topic. If the products have less energy than the reactants, ΔH is negative and energy has escaped to the surroundings as heat. If the products have more energy, ΔH is positive and energy has been taken in.

The law of conservation of energy applies: energy is not created or destroyed in a reaction, only transferred between the chemicals (the system) and everything around them (the surroundings), such as the test tube, the water and the air.

Worked example 9.1

The reactants in a reaction have a total enthalpy of 250 kJ and the products 180 kJ. Find ΔH and say whether heat is given out or taken in.

Use ΔH = H(products) − H(reactants) because ΔH is always "final minus initial".
ΔH = 180 − 250 = −70 kJ.
The negative sign tells us the system lost energy, so heat is given out.

ΔH = −70 kJ; heat given out

⚠️ Common mistake: writing ΔH = reactants − products. This reverses the sign and turns an exothermic reaction into an endothermic one.

Try it yourself. Reactants 120 kJ, products 160 kJ. Find ΔH.

Show solution

160 − 120 = +40 kJ (heat absorbed).

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Practice questions

1. A reaction that releases heat to the surroundings is

exothermic
endothermic
reversible
photochemical

Show answer

A: exothermic
Exo = out; ΔH is negative.

2. Dissolving ammonium chloride in water makes the beaker feel cold. The process is

exothermic
a combustion
neutralisation
endothermic

Show answer

D: endothermic
Heat is absorbed from surroundings.

3. The burning of charcoal is

a physical change
thermal decomposition
exothermic
endothermic

Show answer

C: exothermic
Heat is given out (ΔH negative).

4. For an endothermic reaction, ΔH is

always −57 kJ/mol
positive
negative
zero

Show answer

B: positive
Heat is absorbed.

5. Dissolving ammonium nitrate in water makes the beaker cold. The process is

endothermic
exothermic
a redox reaction
combustion

Show answer

A: endothermic
It absorbs heat.

Key words

Enthalpy (H): The heat content of a system at constant pressure.
ΔH: Enthalpy change: heat given out or taken in by a reaction, in kJ mol⁻¹.
Exothermic: Releasing heat to the surroundings; ΔH negative.
Endothermic: Absorbing heat from the surroundings; ΔH positive.
Activation energy (Ea): Minimum energy colliding particles must have for a reaction to occur.
Standard conditions: 298 K (25 °C) and 1 atm (101.325 kPa), substances in their normal states.
Heat of combustion: Heat released when one mole of a substance burns completely in oxygen.
Heat of formation: Heat change when one mole of a compound is formed from its elements in their standard states.
Heat of neutralisation: Heat released when one mole of water is formed from an acid and a base.
Heat of solution: Heat change when one mole of a substance dissolves in a large amount of water.
Bond energy: Average energy needed to break one mole of a given bond in gaseous molecules.
Hess’s law: The enthalpy change of a reaction is independent of the route taken.

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