Exothermic and Endothermic Reactions: Practice Questions

These are original questions written for Marlbridge, for revision and practice on this content. They are not reproduced past-paper questions, and they do not replicate the exam’s exact structure, question count or mark tariffs — examination boards hold copyright in their own papers. Use these alongside the official past papers available free from your board.

Related: Exothermic and Endothermic Reactions revision notes

Tier note: most of this page is in the Supplement column of the 0620 syllabus, so it is marked (0620 Extended, 5070 required): ΔH and its sign (syllabus 5.1 outcome 4), activation energy (5.1 outcome 5), drawing and labelling reaction pathway diagrams (5.1 outcome 6), bond breaking and bond making (5.1 outcome 7), bond-energy calculations (5.1 outcome 8) and a catalyst lowering the activation energy (6.2 outcome 7). 0620 Core covers questions 1 (apart from the sign of ΔH) and 2: exothermic and endothermic reactions in terms of energy transfer and the temperature of the surroundings (5.1 outcomes 1–2). Core candidates also need to interpret (not draw) reaction pathway diagrams (5.1 outcome 3). 5070 has no tiers, so O Level candidates need all of it. Questions 6 and 8 are background beyond both syllabuses.

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Section A

1. Define exothermic and endothermic, stating the sign of ΔH and the direction of energy transfer in each case. (The sign of ΔH is 0620 Extended, 5070 required.) [4]

2. Give two everyday or laboratory examples of an exothermic process, and two of an endothermic process. [4]

Section B

3. (0620 Extended, 5070 required) Explain, in terms of bond breaking and bond making, why a reaction is exothermic. [3]

4. (0620 Extended, 5070 required) For the reaction H₂ + Cl₂ → 2HCl, the bond energies in kJ mol⁻¹ are H–H 436, Cl–Cl 242, H–Cl 431.

(a) Calculate the energy needed to break the bonds in the reactants. [1] (b) Calculate the energy released in forming the bonds in the products. [1] (c) Calculate ΔH and state whether the reaction is exothermic or endothermic. [3]

5. (0620 Extended, 5070 required) Sketch or describe an energy profile diagram for an endothermic reaction, labelling reactants, products, activation energy and ΔH. [4]

Not on the 0620/5070 syllabus. Questions 6 and 8 use calorimetry and Q = mcΔT, which are not part of either syllabus’s chemical energetics content — the only quantitative energetics outcome required at this level is the bond-energy calculation (question 4). They’re included here only as useful background for progression to AS Chemistry (9701), and won’t be examined at IGCSE/O Level.

6. (Background — not a 0620/5070 outcome.) 50.0 cm³ of acid is neutralised by 50.0 cm³ of alkali (equal volumes); the temperature rises by 6.8 °C. Take the specific heat capacity as 4.18 J g⁻¹ °C⁻¹ and the density as 1.00 g cm⁻³.

(a) Calculate the energy released in joules. [2] (b) Give two reasons why the experimental value is lower in magnitude than the accepted value. [2]

7. (0620 Extended, 5070 required) A catalyst is added to a slow reaction.

(a) Explain, with reference to activation energy, how the catalyst increases the rate of reaction. [2] (b) State what happens to ΔH when a catalyst is used, explaining why. [2] (c) Describe how the catalysed pathway would be drawn on the same energy profile axes as the uncatalysed pathway. [2]

8. (Background — not a 0620/5070 outcome.) 100 cm³ of dilute acid (density 1.00 g cm⁻³) is neutralised by 100 cm³ of alkali (equal volumes), and the temperature rises by 5.5 °C. Take c = 4.18 J g⁻¹ °C⁻¹.

(a) Calculate the energy released, in joules. [2] (b) Give one reason for the difference between this experimental value and the accepted data-book value, other than heat loss to the surroundings. [1]

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Answers

1. Exothermic — energy is transferred to the surroundings, so the surroundings get warmer [1]; ΔH is negative [1]. Endothermic — energy is taken in from the surroundings, which get colder [1]; ΔH is positive [1].

2. Exothermic: combustion, neutralisation, respiration, oxidation [1] [1]. Endothermic: thermal decomposition, photosynthesis, the reaction of citric acid with sodium hydrogencarbonate, dissolving ammonium nitrate [1] [1].

3. Bond breaking is endothermic and bond making is exothermic [1]. In an exothermic reaction, more energy is released making the bonds in the products than is absorbed breaking the bonds in the reactants [1], so there is a net release of energy to the surroundings [1].

4. (a) 436 + 242 = 678 kJ mol⁻¹ [1]. (b) 2 × 431 = 862 kJ mol⁻¹ [1]. (c) ΔH = bonds broken − bonds made = 678 − 862 [1] = −184 kJ mol⁻¹ [1]; the negative value shows the reaction is exothermic [1].

5. Reactants drawn lower than the products on the energy axis [1]; a hump between them with activation energy measured from the reactants up to the peak [1] [1]; ΔH measured from the reactants up to the products, arrow pointing upwards and labelled as positive [1].

6. (Background.) (a) Total volume = 50.0 + 50.0 = 100.0 cm³ (acid + alkali, both warmed), so mass = 100.0 g; Q = mcΔT = 100.0 × 4.18 × 6.8 [1] = 2842 J (2.84 kJ) [1]. (b) Heat is lost to the surroundings, the thermometer and the container during the reaction, so less energy ends up warming the measured solution than the reaction actually released [1]; a simple polystyrene cup is not fully insulated, so it loses further heat in the same direction [1]. (The assumption that the solution has the specific heat capacity of pure water is a separate simplifying assumption — it does not, by itself, explain why the value is lower.)

7. (a) The catalyst provides an alternative reaction pathway with a lower activation energy [1], so a greater proportion of colliding particles have enough energy to react, increasing the frequency of successful collisions [1]. (b) ΔH is unchanged [1], because the catalyst only affects the pathway between reactants and products, not their relative energy levels, which is what ΔH actually measures [1]. (c) A second, lower curve drawn between the same two starting and ending energy levels as the uncatalysed pathway [1], with a smaller hump representing the reduced activation energy [1].

8. (Background.) (a) Total volume = 100 + 100 = 200 cm³ (acid + alkali, both warmed), so mass = 200 g; Q = mcΔT = 200 × 4.18 × 5.5 [1] = 4598 J (4.60 kJ) [1]. (b) The specific heat capacity used (4.18 J g⁻¹ °C⁻¹) is that of pure water, not the actual dilute acid/salt solution, which differs slightly [1].

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Where marks are usually lost

• (0620 Extended, 5070 required) Giving ΔH the wrong sign.
• (0620 Extended, 5070 required) Reversing “bonds broken − bonds made”.
• (0620 Extended, 5070 required) Measuring activation energy from the products instead of the reactants.
• (Background questions 6 and 8 only — not a 0620/5070 outcome) Using the mass of the solute rather than the total solution mass in Q = mcΔT.
• (0620 Extended, 5070 required) Saying a catalyst changes ΔH, or drawing the catalysed curve starting/ending at different energy levels from the uncatalysed one — only the height of the hump should change.
• (0620 Extended, 5070 required) Explaining “increases the rate” without linking it back to activation energy and the proportion of particles with enough energy to react.

Question 7 draws on the catalyst section of the Exothermic and Endothermic Reactions revision notes; questions 6 and 8 draw on the calorimetry section instead (background beyond 0620/5070), since question 8 is itself a calorimetry question and has nothing to do with catalysts.