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Edexcel IAL Chemistry Topic 12 Entropy and energetics

Translate equations and physical changes into entropy direction, calculate each contribution with compatible units and use total entropy rather than system entropy alone for…

Syllabus
First assessment 2019
Course
Chemistry YCH11
Level
A2

Exam points

  • Explain entropy changes through molecular dispersal, energy distribution, state and gas moles.
  • Calculate ΔSsystem and ΔSsurroundings from tabulated entropy, enthalpy and temperature data.
  • Combine contributions to obtain ΔStotal and judge feasibility at a stated temperature.

Covered syllabus points

12.1—That, since endothermic reactions can occur spontaneously at room temperature, enthalpy changes alone do not control whether12.2—Entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between12.3—The entropy of a substance increases with temperature, that entropy increases as solid → liquid → gas and that perfect12.4—Interpret the natural direction of change as being in the direction of increasing total entropy (positive entropy change)12.5—Why entropy changes occur during: i changes of state ii dissolving of a solid ionic lattice iii reactions in which there is12.6—The total entropy change of any reaction is the sum of the entropy change of the system and the entropy change of12.7—The entropy change of the system for a reaction, ∆Ssystem, given the entropies of the reactants and products12.8—The entropy change in the surroundings, and hence ∆Stotal12.9—The feasibility of a reaction depends on: i the balance between ∆Ssystem and ∆Ssurroundings, so that even endothermic12.10—Reactions can occur as long as ∆Stotal is positive even if one of the other entropy changes is negative12.11—And distinguish between the concepts of thermodynamic stability and kinetic stability12.12—Define the terms: i standard enthalpy change of atomisation, ∆atH ii electron affinity iii lattice energy (as the exothermic12.13—Construct Born-Haber cycles and carry out related calculations12.14—A comparison of the experimental lattice energy value (from a Born-Haber cycle) with the theoretical value (obtained12.15—Polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycle12.16—Define the terms ‘enthalpy change of solution, ∆solH’ and ‘enthalpy change of hydration, ∆hydH of an ion’12.17—Energy cycles and energy level diagrams to calculate the enthalpy change of solution of an ionic compound12.18—The effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic12.19—Entropy and enthalpy changes of solution values to predict the solubility of ionic compounds and discuss trends in

Topic 12: Entropy and Energetics question 1

[Maximum number: 1]

Which of these gases would have the greatest standard molar entropy?

A

NH3\mathrm{NH}_{3}

B

H2\mathrm{H}_{2}

C

N2\mathrm{N}_{2}

D

SO2\mathrm{SO}_{2}

Topic 12: Entropy and Energetics question 2

[Maximum number: 1]

In which series does the standard molar entropy of the compounds increase?

A

CaO(s)<H2O(l)<CO2( g)<SO2( g)\mathrm{CaO}(\mathrm{s})<\mathrm{H}_{2} \mathrm{O}(\mathrm{l})<\mathrm{CO}_{2}(\mathrm{~g})<\mathrm{SO}_{2}(\mathrm{~g})

B

CaO(s)<H2O(l)<SO2( g)<CO2( g)\mathrm{CaO}(\mathrm{s})<\mathrm{H}_{2} \mathrm{O}(\mathrm{l})<\mathrm{SO}_{2}(\mathrm{~g})<\mathrm{CO}_{2}(\mathrm{~g})

C

SO2( g)<CO2( g)<H2O(l)<CaO(s)\mathrm{SO}_{2}(\mathrm{~g})<\mathrm{CO}_{2}(\mathrm{~g})<\mathrm{H}_{2} \mathrm{O}(\mathrm{l})<\mathrm{CaO}(\mathrm{s})

D

CO2( g)<SO2( g)<H2O(l)<CaO(s)\mathrm{CO}_{2}(\mathrm{~g})<\mathrm{SO}_{2}(\mathrm{~g})<\mathrm{H}_{2} \mathrm{O}(\mathrm{l})<\mathrm{CaO}(\mathrm{s})

Topic 12: Entropy and Energetics question 3

[Maximum number: 13]

Ammonia is manufactured in the Haber Process.

N2( g)+3H2( g)⇌2NH3( g)\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_{3}(\mathrm{~g})

In this process, pressures between 100 atm and 300 atm and temperatures between 675 K and 725 K are usually used.

Question (a)

(a)

Calculate ΔSsystem ⊖\Delta S_{\text {system }}^{\ominus} for this reaction at 298 K .

Include a sign and units with your answer.
[Standard molar entropy, S⊖\mathrm{S}^{\ominus}, for 1 mol of each gas, in JK−1 mol−1\mathrm{J} \mathrm{K}^{-1} \mathrm{~mol}^{-1} :
N2=191.6H2=130.6NH3=192.3]\left.\mathrm{N}_{2}=191.6 \quad \mathrm{H}_{2}=130.6 \quad \mathrm{NH}_{3}=192.3\right]

[ 2 ]

Question (b)

(b)

Explain, using ideas about disorder, whether the sign of your answer to (a)(i) is as expected.

[ 2 ]

Question (c)

(c)

At 700 K , the enthalpy change for this reaction is ΔH=−110.2 kJ mol−1\Delta H=-110.2 \mathrm{~kJ} \mathrm{~mol}^{-1}. Calculate the entropy change of the surroundings, ΔSsurroundings \Delta S_{\text {surroundings }}, at 700 K . Include a sign and units in your answer.

[ 2 ]

Question (d)

(d)

At 700 K, the total entropy change, ΔS_total = -78.7 J K⁻¹ mol⁻¹. Calculate ΔS_system for this reaction at 700 K. Include a sign and units in your answer.

[ 1 ]

Question (e)

(e)

Explain one advantage and one disadvantage of using a temperature higher than 700 K in the Haber process. Consider the effect of an increase in temperature on the rate of reaction, on the values of ΔS_surroundings and ΔS_total and on the equilibrium constant.

[ 6 ]
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