Learners' AlleyDesigned by Khow Willard

Chemical Energetics II

Set AFoundations

Name

Time 1 hour

Total 34 marks

A Data Booklet may be used.

From the lesson

7 marks

CommonFind ΔH_sol(KCl), given LE(KCl) = −711, ΔH_hyd(K⁺) = −322 and ΔH_hyd(Cl⁻) = −381 kJ mol⁻¹.

[1]

Source: Learners' Alley, Chemical Energetics II slides

ChallengeWhich ion has the more exothermic ΔH_hyd, Mg²⁺ or Na⁺? Explain.

[2]

Source: Learners' Alley, Chemical Energetics II slides

CommonFor the Haber process, ΔH = −92 kJ mol⁻¹ and ΔS = −199 J mol⁻¹ K⁻¹. Is it spontaneous at 298 K?

[2]

Source: Learners' Alley, Chemical Energetics II slides

Baking powder releases CO₂ on heating: 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(g) + CO₂(g), ΔH = +130 kJ mol⁻¹, ΔS = +335 J mol⁻¹ K⁻¹. Above what temperature is this spontaneous?

[2]

Source: Learners' Alley, Chemical Energetics II slides

Practice

27 marks

CommonFor the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ mol⁻¹. Which statement about the forward reaction is correct?

AIt is spontaneous only at low temperature.

BIt is spontaneous only at high temperature.

CIt is not spontaneous at any temperature.

DIt is spontaneous at all temperatures.

Answer [1]

Source: 2024 DHS J2 Prelim P1 Q10

Use of the Data Booklet is relevant to this question. When water vaporises, 40.7 kJ mol⁻¹ of heat is required. What is the entropy change when 72 g of water boils at 100 °C?

A−436 J K⁻¹

B−109 J K⁻¹

C+109 J K⁻¹

D+436 J K⁻¹

Answer [1]

Source: IJC 2017 Promo, via EJC 2022 Chemical Energetics Topical Revision, Q20

Magnesium chloride is one of the most abundant salts in sea water.

lattice energy of MgCl₂(s)	−2521 kJ mol⁻¹
enthalpy change of formation of MgCl₂(s)	−641 kJ mol⁻¹
enthalpy change of formation of Mg²⁺(aq)	−467 kJ mol⁻¹
enthalpy change of formation of Cl⁻(aq)	−167 kJ mol⁻¹

(a)

Use relevant data above to calculate a value for the enthalpy change of solution of magnesium chloride.

[2]

(b)

Use your answer to explain why MgCl₂ is very soluble in water.

[1]

Source: EJC 2022 JC1 Chemical Energetics Tutorial, A6

CommonAssume the pressure of each system remains constant throughout each process.

(a)

Explain what is meant by the term entropy of a chemical system.

[1]

(b)(i)

Describe and explain how the entropy of each of the following systems will change during the stated process.
1 mol of Cl₂(g) at 298 K is added to 1 mol of N₂(g) at 298 K.

[1]

(ii)

1 mol of Cl₂(g) at 298 K is heated to 373 K.

[1]

(iii)

1 mol of Cl₂(g) at 298 K is reacted with iodine: Cl₂(g) + 13I₂(s) → 23ICl₃(l)

[1]

(iv)

1 mol of Cl₂(g) at 1000 K is photolysed: Cl₂(g) → 2Cl•(g)

[1]

Source: EJC 2022 JC1 Chemical Energetics Tutorial, B8

(a)

Given that the lattice energy of lithium chloride is −843 kJ mol⁻¹ and its hydration energy is −883 kJ mol⁻¹, calculate the enthalpy change of solution of lithium chloride.

[1]

(b)

50.0 cm³ of water was placed in a thin, well-lagged calorimeter of negligible heat capacity. The temperature of the water was 18.2 °C. 3.00 g of potassium chloride, KCl, was added and the mixture well stirred until all the crystals had dissolved. The minimum temperature reached was 14.9 °C. Calculate the enthalpy change of solution of potassium chloride. (Specific heat capacity of water = 4.20 J g⁻¹ K⁻¹)

[2]

(c)

Comment on the results obtained in (a) and (b).

[2]

Source: RJC H2 Chemistry Tutorial 5a (Chemical Energetics I), Q7(b) to (d)

Nitroglycerin, C₃H₅(NO₃)₃, decomposes on ignition: C₃H₅(NO₃)₃(l) → 32N₂(g) + 14O₂(g) + 3CO₂(g) + 52H₂O(g). The standard enthalpy change of this decomposition is −1423 kJ mol⁻¹.

(a)

Calculate ΔG^⦵ at 298 K for this decomposition, given that ΔS^⦵ = +208 J K⁻¹ mol⁻¹. Hence, predict the spontaneity of the reaction at 298 K.

[2]

(b)

Is the reaction spontaneous at all temperatures? Explain.

[1]

Source: 2019 CJC J1 Promo P2 Q4(c) and (d)

CommonDinitrogen pentoxide, N₂O₅, can be produced in a car engine by the following reaction sequence.

I	N₂(g) + O₂(g) → 2NO(g)	ΔH^⦵ = +180 kJ mol⁻¹, ΔS^⦵ = +24.8 J mol⁻¹ K⁻¹
II	2NO(g) + O₂(g) → 2NO₂(g)	ΔH^⦵ = −114 kJ mol⁻¹
III	2NO₂(g) + ½O₂(g) → N₂O₅(g)	ΔH^⦵ = −110 kJ mol⁻¹

(a)

Explain why reaction I does not take place at standard conditions but occurs in the car engine.

[2]

(b)

Predict the sign of ΔS^⦵ in reaction II. Explain your answer.

[2]

Source: NJC 2018 Promo, via EJC 2022 Chemical Energetics Topical Revision, Q40(i) and (ii)

ChallengeThe following data relate to the energy changes which occur when sodium hydroxide, NaOH, and magnesium hydroxide, Mg(OH)₂, dissolve in water.

	sodium	magnesium
hydration energy of metal ion / kJ mol⁻¹	−390	−1890
lattice energy of metal hydroxide / kJ mol⁻¹	−896	−2995
enthalpy change of solution of metal hydroxide / kJ mol⁻¹	−44	x
extent of hydration of metal ion (average number of attached water molecules per metal ion)	5	15

(a)

Explain why the hydration energy of Mg²⁺ is more exothermic than that of Na⁺.

[1]

(b)

Use the data above to calculate x, the enthalpy change of solution of Mg(OH)₂.

[2]

(c)

Solubility is also controlled by the entropy change which accompanies dissolution. Use the extent of hydration from the table to compare qualitatively the entropy changes of solution of Mg(OH)₂ and NaOH.

[1]

(d)

Suggest and explain one other factor which contributes to the total entropy change when Mg(OH)₂ dissolves in water.

[1]

Source: EJC 2022 JC1 Chemical Energetics Tutorial, B9