OxfordAQA International GCSE Chemistry: Exothermic and endothermic reactions, and calculating and explaining energy change (9202) – Revision Notes

These revision notes condense sections 3.9.1 Exothermic and endothermic reactions and 3.9.2 Calculating and explaining energy change of the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. They cover statements 3.9.1 a to d and 3.9.2 a to h. The qualification is untiered, so every point is for every candidate, and section 6.2 attaches no required practical to these sections. For full teaching and longer worked examples, use the energy change study guide.

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The one-line summary

Energy goes in to break bonds and comes out when bonds form; the larger of the two decides whether the surroundings warm up or cool down.

Definitions to learn word for word

Term	Meaning
Exothermic reaction	Transfers energy to the surroundings; surroundings get hotter; ΔH negative
Endothermic reaction	Takes in energy from the surroundings; surroundings get colder; ΔH positive
ΔH	Enthalpy change of the reacting chemicals, usually in kJ/mol
Activation energy	Minimum energy particles need for a reaction to happen
Catalyst	Gives the reaction an alternative route that has a lower activation energy
Reversible reaction	Products can react to re-form the original reactants (⇌)
Bond dissociation energy	Energy needed to break one mole of a given bond

3.9.1: which reactions are which

Exothermic (3.9.1 b)

• Combustion of fuels
• Many oxidation reactions, such as iron rusting
• Neutralisation of an acid by an alkali
• Uses: self-heating cans (for example calcium oxide reacting with water to heat coffee) and hand warmers (air-activated packs where iron powder oxidises)

Endothermic (3.9.1 c)

• Thermal decompositions, such as heating calcium carbonate to make calcium oxide and carbon dioxide
• Use: sports injury packs that go cold, such as packs where ammonium nitrate dissolves in water

Reversible reactions (3.9.1 d)

                        endothermic →
hydrated copper sulfate ⇌ anhydrous copper sulfate + water
       (blue)           ← exothermic        (white)

• Heat blue crystals: they turn white and water is given off (energy taken in).
• Add water to white powder: it turns blue and gets hot (energy given out).
• The energy transferred is the same size in both directions; only the sign changes.

3.9.2 a to c: calorimetry

The equation

Symbol	Quantity	Unit
Q	energy transferred	J
m	mass of water or solution heated	g
c	specific heat capacity (4.2 for water)	J/g °C
ΔT	temperature change	°C

Q = mcΔT. Divide by 1000 to get kJ. Then divide by grams of fuel for kJ/g, or by moles for kJ/mol (3.9.2 b).

Units (3.9.2 b)

Unit	Use it for
J or kJ (1 kJ = 1000 J)	the energy Q transferred in one experiment
kJ/g	comparing fuels by the mass burned
kJ/mol	comparing reactions per mole of a named substance, and for ΔH

Method in steps: fuels (3.9.2 a)

1. Measure a fixed mass of water into a glass or metal container and record its temperature.
2. Weigh the spirit burner with its fuel.
3. Light it under the container, stir, and stop when the water has risen by a sensible amount.
4. Record the highest temperature and reweigh the burner.
5. Q = mcΔT using the water’s mass, then divide by the mass of fuel burned.
6. Repeat for each fuel with the same water mass, container and flame distance.

Small worked reminder: 0.50 g of a fuel heats 100 g of water by 15.0 °C. Q = 100 × 4.2 × 15.0 = 6300 J = 6.30 kJ. Per gram = 6.30 ÷ 0.50 = 12.6 kJ/g.

Method in steps: solutions (3.9.2 c)

1. Measure one solution into a polystyrene cup (sitting in a beaker for stability) and record its temperature.
2. Add the second solution or the weighed solid; fit a lid with a hole for the thermometer.
3. Stir and record the highest (or lowest) temperature.
4. m = total volume of solution in cm³ (1 cm³ ≈ 1 g); for a solid in water, use the mass of water.
5. Q = mcΔT, then divide by moles that reacted.
6. Temperature rose: ΔH negative. Temperature fell: ΔH positive.

Why results are lower than true values

• Energy heats the container, thermometer and air.
• Incomplete combustion of a fuel (soot forms on the can).
• Fuel evaporates from the wick before burning.
• Slow reactions lose more heat before the highest temperature is reached.

3.9.2 d and h: energy level diagrams

ENDOTHERMIC                       EXOTHERMIC
Energy                            Energy
 |        __                       |       __
 |       /  \______ products       | _____/  \
 | _____/          (higher)        | reactants\
 | reactants                       |           \______ products
 +-------------------->            +-------------------->  (lower)

• Activation energy = peak − reactant level.
• ΔH = product level − reactant level (negative if products are lower).
• Catalysed curve: lower peak, same start, same finish, same ΔH.
• The curve rises first in both cases, because bonds must start to break.

Small worked reminder: reactants 40 kJ, peak 190 kJ, products 125 kJ. Activation energy = 190 − 40 = 150 kJ. ΔH = 125 − 40 = +85 kJ, so endothermic.

3.9.2 e to g: bond energies

Process	Energy	Sign
Breaking bonds	Taken in	endothermic step
Making bonds	Given out	exothermic step
Exothermic overall (3.9.2 f)	made > broken	ΔH negative
Endothermic overall (3.9.2 g)	broken > made	ΔH positive

Method in steps

1. Write displayed formulas, or list every bond in each molecule.
2. Multiply by the coefficients in the balanced equation.
3. Total the energy for bonds broken (all reactant bonds).
4. Total the energy for bonds formed (all product bonds).
5. Energy change = broken − formed.
6. State exothermic or endothermic from the sign.

Small worked reminder: CO + H₂O → CO₂ + H₂, using C≡O 1077, O–H 464, C=O 805, H–H 436 kJ/mol.

• Broken: 1077 + (2 × 464) = 2005 kJ
• Formed: (2 × 805) + 436 = 2046 kJ
• Energy change = 2005 − 2046 = −41 kJ/mol, exothermic

Must-know distinctions

• ΔH sign vs temperature change. Temperature up means ΔH down (negative). They move in opposite directions.
• Activation energy vs ΔH. Activation energy is the hill; ΔH is the difference between start and finish.
• kJ/g vs kJ/mol. Per gram compares fuels by mass; per mole compares by number of particles. A fuel can win on one and lose on the other.
• Fuel calorimetry vs solution calorimetry. Fuels heat water in a separate container; in solution the reaction happens inside the water you measure.
• Catalyst vs heating. A catalyst lowers the activation energy; heating gives particles more energy but leaves the activation energy unchanged.

Quick self-test

1. Give the sign of ΔH for a neutralisation reaction.
2. Name the type of reaction that is endothermic in 3.9.1 c.
3. White anhydrous copper sulfate has water added. Give the colour change and whether energy is taken in or given out.
4. Calculate Q when 75 g of water rises in temperature by 12.0 °C. Use c = 4.2 J/g °C.
5. Convert 2.45 kJ into joules.
6. 0.60 g of a fuel releases 15.3 kJ. Calculate the energy released per gram.
7. On a profile, reactants are at 300 kJ and the peak is at 420 kJ. Calculate the activation energy.
8. Bonds broken total 1580 kJ and bonds formed total 1920 kJ. Calculate the energy change and state its type.
9. What effect does a catalyst have on ΔH?
10. In a neutralisation, 0.0250 mol of water forms and 1.40 kJ is released. Calculate ΔH in kJ/mol.
11. Why is a polystyrene cup used for reactions in solution?

Answers

1. Negative: neutralisation is exothermic.
2. Thermal decomposition.
3. White to blue; energy is given out (exothermic).
4. Q = 75 × 4.2 × 12.0 = 3780 J (3.78 kJ).
5. 2450 J.
6. 15.3 ÷ 0.60 = 25.5 kJ/g.
7. 420 − 300 = 120 kJ.
8. 1580 − 1920 = −340 kJ; exothermic.
9. None: ΔH is unchanged; only the activation energy is lowered.
10. 1.40 ÷ 0.0250 = 56.0, so ΔH = −56.0 kJ/mol.
11. It is a good insulator, so less energy is lost to or gained from the surroundings.

Where marks are usually lost

• Writing “exothermic reactions take in heat” by mixing up “to” and “from” the surroundings.
• Giving a positive ΔH for a reaction that warmed the solution.
• Using the mass of fuel burned as m in Q = mcΔT.
• Using 40 g instead of 80 g when 40 cm³ of each solution is mixed.
• Dividing joules by moles and labelling the answer kJ/mol.
• Drawing a catalysed curve that ends at a different energy level.
• Calling activation energy “the energy given out”.
• Missing a coefficient when counting bonds, so only one molecule’s bonds are included.
• Writing “bonds formed − bonds broken”, which gives the wrong sign.
• Describing the copper sulfate colours the wrong way round.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. These notes cover section 3.9.1 Exothermic and endothermic reactions and section 3.9.2 Calculating and explaining energy change.