Edexcel A-Level Chemistry A2 12 Entropy and Energetics Questions

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

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 whether 12.2—Entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between 12.3—The entropy of a substance increases with temperature, that entropy increases as solid → liquid → gas and that perfect 12.4—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 is 12.6—The total entropy change of any reaction is the sum of the entropy change of the system and the entropy change of 12.7—The entropy change of the system for a reaction, ∆Ssystem, given the entropies of the reactants and products 12.8—The entropy change in the surroundings, and hence ∆Stotal 12.9—The feasibility of a reaction depends on: i the balance between ∆Ssystem and ∆Ssurroundings, so that even endothermic 12.10—Reactions can occur as long as ∆Stotal is positive even if one of the other entropy changes is negative 12.11—And distinguish between the concepts of thermodynamic stability and kinetic stability 12.12—Define the terms: i standard enthalpy change of atomisation, ∆atH ii electron affinity iii lattice energy (as the exothermic 12.13—Construct Born-Haber cycles and carry out related calculations 12.14—A comparison of the experimental lattice energy value (from a Born-Haber cycle) with the theoretical value (obtained 12.15—Polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycle 12.16—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 compound 12.18—The effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic 12.19—Entropy and enthalpy changes of solution values to predict the solubility of ionic compounds and discuss trends in

[Maximum number: 1]

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

$\mathrm{NH}_{3}$

$\mathrm{H}_{2}$

$\mathrm{N}_{2}$

$\mathrm{SO}_{2}$

[Maximum number: 1]

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

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

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

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

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

[Maximum number: 13]

Ammonia is manufactured in the Haber Process.

\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.

(a)

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

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

[ 2 ]

(b)

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

[ 2 ]

(c)

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

[ 2 ]

(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 ]

(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 ]

[Maximum number: 2]

Ammonia is manufactured in the Haber Process.

\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.

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

[Maximum number: 1]

Ammonia is manufactured in the Haber Process.

\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.

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.

[Maximum number: 2]

The equation for the formation of ammonia in the Haber Process is shown

1 / 2 \mathrm{~N}_{2}(\mathrm{~g})+1 \frac{1}{2} \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})

At 298 K the standard entropy change of the system, $\Delta S_{\text {system }}^{\ominus}=-98 \mathrm{JK}^{-1} \mathrm{~mol}^{-1}$ .

Calculate the standard entropy of one mole of ammonia.
Use the value of $\Delta S_{\text {system }}^{\ominus}$ and the data in the table.

Table for Question 6 — Edexcel A-Level Chemistry A2

states or uses equation

- calculate $S^{\circ}$ products

(1)

(1)

$\begin{aligned} \Delta S_{\text {system }}^{\ominus} & =S_{\text {products }}^{\ominus}-S_{\text {reactants }}^{\ominus} -98.0 & =S_{\text {products }}^{\ominus}-((0.5 \times 192)+(1.5 \times 131)) S_{\text {products }}^{\ominus} & =292.5-98 S_{\text {products }}^{\ominus} & =(+) 194.5 / 195\left(\mathrm{~J} \mathrm{~K}^{-1} \mathrm{~mol}^{-1}\right) \end{aligned}$

If units are given they must be correct

Allow TE on incorrect $S^{\ominus}$ reactants

Comment

Correct answer with no working scores (2)
$\begin{aligned} S_{\text {products }}^{\ominus} & =63.5 \text { scores max (1) } S_{\text {products }}^{\ominus} & =225 \text { scores max (1) } \end{aligned}$

[Maximum number: 1]

The equation for the formation of ammonia in the Haber Process is shown

1 / 2 \mathrm{~N}_{2}(\mathrm{~g})+1 \frac{1}{2} \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})

The relationship between $\Delta S_{\text {total }}$ and 1 / T can be found by combining the two equations:

\Delta S_{\text {total }}=\Delta S_{\text {surroundings }}+\Delta S_{\text {system }}

and $\Delta S_{\text {surroundings }}=-\Delta H / T$
to give

\Delta S_{\text {total }}=-\Delta H / T+\Delta S_{\text {system }}

Identify the thermodynamic quantity that can be obtained from this gradient.

An answer that makes reference to the following points:

- enthalpy change of reaction / $\Delta_{\mathrm{r}} H$ (of the Haber process)

Allow $-_{\mathrm{r}} H$

Allow enthalpy change / $\Delta H /-\Delta H$

[Maximum number: 4]

The equation for the formation of ammonia in the Haber Process is shown

1 / 2 \mathrm{~N}_{2}(\mathrm{~g})+1 \frac{1}{2} \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})

(a)

The relationship between $\Delta S_{\text {total }}$ and 1 / T can be found by combining the two equations:

\Delta S_{\text {total }}=\Delta S_{\text {surroundings }}+\Delta S_{\text {system }}

and $\Delta S_{\text {surroundings }}=-\Delta H / T$
to give

\Delta S_{\text {total }}=-\Delta H / T+\Delta S_{\text {system }}

[ 1 ]

(i)

Determine the temperature at which the reaction ceases to be thermodynamically feasible at a pressure of 100 kPa .

[ 1 ]

An answer that makes reference to the following points:

- value of T found either by reading from the graph the value of T when $\Delta S_{\text {total }}=0$

or

by calculation

460 (K)

Allow an answer between 440-480

$=\frac{\text { answer to (c)(i) }}{98}$ $=\frac{45150}{98}=460.71 / 460(\mathrm{~K})$

Or

$=\frac{- \text { answer to (b) }}{-98}$ $=\frac{-45150}{-98}=460.71 / 460(\mathrm{~K})$

ALLOW TE on graph or on answer to (c)(i)

(b)

The industrial synthesis of ammonia

1 / 2 \mathrm{~N}_{2}(\mathrm{~g})+11 / 2 \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NH}_{3}(\mathrm{~g})

is carried out at pressures of about 20000 kPa and temperatures between 700 K and 750 K . These temperatures are higher than the answer to (c)(iii).

[ 3 ]

(i)

Explain why $\Delta S_{\text {total }}$ decreases with an increase in temperature.

[ 3 ]

An answer that makes reference to the following points:

Either

- ( $\Delta S_{\text {total }}$ decreases because) $\Delta S_{\text {system }}$ (and $\Delta H$ ) do not change with temperature (significantly)

- therefore $\Delta S_{\text {surroundings }}$ must decrease (so that ( $\Delta S_{\text {total }}$ decreases)

- this is because $\Delta S_{\text {surroundings }}=-\Delta H / T$ (so as T increases $-\Delta H / T$ becomes less positive because $\Delta H$ is exothermic)

Or

- the reaction is exothermic and

so increasing temperature shifts the equilibrium to the left / towards the reactants

- the value of K decreases

- because $\Delta S_{\text {total }}$ is proportional to K /

$S_{\text {total }}=R \ln K$ the value of $\Delta S_{\text {total }}$ decreases

[Maximum number: 2]

This question is about oxides of iron.

The equation for the reduction of iron(III) oxide by carbon monoxide is shown.

\mathrm{Fe}_{2} \mathrm{O}_{3}(\mathrm{~s})+3 \mathrm{CO}(\mathrm{~g}) \rightarrow 2 \mathrm{Fe}(\mathrm{~s})+3 \mathrm{CO}_{2}(\mathrm{~g}) \quad \Delta H^{\ominus}=-24.8 \mathrm{~kJ} \mathrm{~mol}^{-1}

The entropy change, $\Delta S_{\text {system }}^{\ominus}$ , for this reaction is $+15.2 \mathrm{~J} \mathrm{~mol}^{-1} \mathrm{~K}^{-1}$ .

Explain why the reaction should be feasible at any temperature.

[Maximum number: 2]

Nitrous oxide, $\mathrm{N}_{2} \mathrm{O}$ , decomposes at high temperature to form nitrogen and oxygen.

\mathrm{N}_{2} \mathrm{O}(\mathrm{~g}) \rightleftharpoons \mathrm{N}_{2}(\mathrm{~g})+1 / 2 \mathrm{O}_{2}(\mathrm{~g})

Explain whether or not this reaction occurs at 2048 K by considering the values calculated in (a) and (b).

[Maximum number: 2]

This question is about oxides of iron.

Iron also forms iron(II) oxide. The data in the table can be used to construct a Born-Haber cycle for iron(II) oxide, FeO.

Table for Question 11 — Edexcel A-Level Chemistry A2

Suggest why the second electron affinity of oxygen is positive.
()

[Maximum number: 5]

This question is about oxides of iron.

(a)

Iron also forms iron(II) oxide. The data in the table can be used to construct a Born-Haber cycle for iron(II) oxide, FeO.

Table for Question (a) — Edexcel A-Level Chemistry A2

[ 5 ]

(i)

Complete the Born-Haber cycle by putting letters in the boxes to label the energy changes.

Figure for Question (i) — Edexcel A-Level Chemistry A2

[ 3 ]

(ii)

Calculate the value for second ionisation energy of iron, C.

[ 2 ]

correct expression

- correct calculation

Example of calculation
-759-416-272+3920-798+141-249 $=(+) 1567\left(\mathrm{~kJ} \mathrm{~mol}^{-1}\right)$
Allow 3SF
Correct answer with or without working scores 2
marks
Allow 1 mark for one mistake

[Maximum number: 3]

This question concerns calcium chloride, $\mathrm{CaCl}_{2}$ .

The theoretical lattice energy of formation of $\mathrm{CaCl}_{2}$ is $-2223 \mathrm{~kJ} \mathrm{~mol}^{-1}$ . Explain why the experimental lattice energy of $\mathrm{CaCl}_{2}$ is more exothermic than the theoretical value.

[Maximum number: 1]

Experimental lattice energy values determined using the Born-Haber cycle often differ from calculated theoretical values.
The ionic radii and charges of some ions are shown.
P, Q, R and S are not the chemical symbols of the elements.

Table for Question 14 — Edexcel A-Level Chemistry A2

Which of these compounds would be expected to show the largest percentage difference between the experimental and theoretical lattice energies?

[Maximum number: 1]

Which of these statements is true?

the enthalpy change of hydration of an ion is always negative

the enthalpy change of solution of an ionic compound is always negative

an endothermic enthalpy change of solution indicates that an ionic compound will be insoluble

the more exothermic the lattice energy of an ionic compound, the more soluble it is likely to be

[Maximum number: 2]

This question concerns calcium chloride, $\mathrm{CaCl}_{2}$ .

Calculate the enthalpy change of solution of calcium chloride, $\mathrm{CaCl}_{2}$ .

Data
$\Delta_{\text {hyd }} H^{\ominus}\left(\mathrm{Ca}^{2+}\right)=-1650 \mathrm{~kJ} \mathrm{~mol}^{-1} \quad \Delta_{\text {hyd }} H^{\ominus}\left(\mathrm{Cl}^{-}\right)=-364 \mathrm{~kJ} \mathrm{~mol}^{-1}$
Experimental lattice energy $\left(\mathrm{CaCl}_{2}\right)=-2258 \mathrm{~kJ} \mathrm{~mol}^{-1}$

[Maximum number: 3]

This question concerns calcium chloride, $\mathrm{CaCl}_{2}$ .

The table shows the theoretical lattice energies of some chlorides of Group 2.

Table for Question 17 — Edexcel A-Level Chemistry A2

Explain the trend in theoretical lattice energies on descending Group 2.

[Maximum number: 1]

The standard enthalpy change of solution of potassium chloride， KCl ，is $+17 \mathrm{~kJ} \mathrm{~mol}^{-1}$ ．
The solubility of potassium chloride in water at 298 K is $359 \mathrm{gdm}^{-3}$ ．
Which of these explains the solubility of potassium chloride in water？

the hydration enthalpy of $\mathrm{K}^{+}$ and the lattice energy of KCl are exothermic

the hydration enthalpy of $\mathrm{K}^{+}$ and the lattice energy of KCl are endothermic

the total entropy change when KCl dissolves is positive

the total entropy change when KCl dissolves is negative

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Covered syllabus points

Question 1

Question 2

Question 3

Question (a)

Question (b)

Question (c)

Question (d)

Question (e)

Question 4

Question 5

Question 6

Question 7

Question 8

Question (a)

Question (i)

Question (b)

Question (i)

Question 9

Question 10

Question 11

Question 12

Question (a)

Question (i)

Question (ii)

Question 13

Question 14

Question 15

Question 16

Question 17

Question 18