1. 1
   A student considers four changes, all taking place at constant pressure: (1) ice melting at 5 °C; (2) the thermal decomposition 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(g) + CO₂(g); (3) the oxidation 2SO₂(g) + O₂(g) → 2SO₃(g); (4) liquid ethanol evaporating in an open dish.

   (a)

   Which change results in a decrease in the entropy of the system?

   [1 mark]

   • A(1)
   • B(3)
   • C(2)
   • D(4)

   (b)

   Which statement best explains why the standard entropy of ethanol vapour is greater than that of liquid ethanol?

   [1 mark]

   • AIn the gas, the particles are far more spread out and move freely, so there are many more ways of distributing the particles and their energy.
   • BThe intermolecular forces in the gas are stronger than in the liquid.
   • CEthanol molecules are larger in the gas phase than in the liquid phase.
   • DEntropy is calculated by dividing the enthalpy of the substance by its temperature, and the gas has more enthalpy.

   (c)

   Melting ice is endothermic. Explain, in terms of the Gibbs energy change, why change (1) is spontaneous at 5 °C.

   [2 marks]

   Total for question 1: 4 marks

2. 2
   Ammonia is manufactured by the reaction N₂(g) + 3H₂(g) → 2NH₃(g), for which ΔH⦵ = −92.2 kJ mol⁻¹. Standard entropies, S⦵, in J K⁻¹ mol⁻¹: N₂(g) 192, H₂(g) 131, NH₃(g) 193. Assume ΔH⦵ and ΔS⦵ do not change with temperature.

   (a)

   What is the standard entropy change, ΔS⦵, for the reaction in J K⁻¹ mol⁻¹?

   [1 mark]

   • A+199
   • B−130
   • C−392
   • D−199

   (b)

   What is the standard Gibbs energy change, ΔG⦵, at 298 K, in kJ mol⁻¹?

   [1 mark]

   • A−151.5
   • B−87.2
   • C−32.9
   • D+5.92 × 10⁴

   (c)

   Determine the temperature above which the reaction is no longer spontaneous under standard conditions.

   [2 marks]

   Total for question 2: 4 marks

3. 3
   Dinitrogen tetroxide dissociates reversibly: N₂O₄(g) ⇌ 2NO₂(g). For this reaction ΔH⦵ = +57.2 kJ mol⁻¹ and ΔS⦵ = +175.8 J K⁻¹ mol⁻¹. A sealed flask at 298 K is filled with a mixture that is almost pure N₂O₄, so that the reaction quotient, Q, is initially 1.0 × 10⁻³. Use R = 8.31 J K⁻¹ mol⁻¹.

   (a)

   Calculate ΔG⦵ at 298 K and hence the equilibrium constant, K, for the dissociation at 298 K.

   [3 marks]

   (b)

   Calculate ΔG for the reaction mixture in the flask at the start, deduce what happens to the composition of the mixture, and explain how ΔG changes as the system approaches equilibrium.

   [4 marks]

   Total for question 3: 7 marks

4. 4
   Quicklime, CaO, is made by heating limestone in a lime kiln: CaCO₃(s) → CaO(s) + CO₂(g). Standard enthalpies of formation, ΔfH⦵, in kJ mol⁻¹: CaCO₃(s) −1207, CaO(s) −635, CO₂(g) −394. Standard entropies, S⦵, in J K⁻¹ mol⁻¹: CaCO₃(s) 93, CaO(s) 40, CO₂(g) 214. The kiln is open to the air, and a stream of air carries the carbon dioxide away as it forms. One student states that the reaction can never be spontaneous because it is endothermic. A second student states that heating the limestone above the temperature at which ΔG⦵ becomes zero guarantees complete and rapid decomposition.

   (a)

   Determine ΔH⦵ and ΔS⦵ for the decomposition of calcium carbonate, deduce whether it is spontaneous at 298 K, and calculate the minimum temperature at which it becomes spontaneous under standard conditions.

   [6 marks]

   (b)

   Evaluate the statements made by the two students, with reference to the Gibbs energy equations and to the conditions in the kiln.

   [6 marks]

   Total for question 4: 12 marks

End of questions