Reactivity: What Drives Chemical Reactions? — Free IB Chemistry HL Practice Questions

1FoundationMCQExothermic vs endothermic reactions1 markPaper 1~2 min

A student reacts solid sodium hydroxide with dilute hydrochloric acid and records the temperature of the mixture rising from

21.0\ ^\circ\text{C}

to

28.5\ ^\circ\text{C}

. However, the student notes that dissolving solid NaOH in water is itself exothermic. Which statement correctly identifies the overall thermal observation and its thermochemical interpretation for the neutralisation experiment?

Solutions

2MasteryMCQEnthalpy of reaction, heat capacity, and calorimetry1 markPaper 1~2 min

A student adds

0.327\ \text{g}

of zinc powder (

M_r = 65.4

) to

50.0\ \text{cm}^3

of

0.200\ \text{mol dm}^{-3}

copper(II) sulfate solution in a calorimetry experiment. The initial temperature is

21.0\ ^\circ\text{C}

and the final temperature is

26.3\ ^\circ\text{C}

. The specific heat capacity and density of the solution are

4.18\ \text{J g}^{-1}\text{K}^{-1}

and

1.00\ \text{g cm}^{-3}

respectively. What is the enthalpy change, in

\text{kJ mol}^{-1}

, for the reaction below?

\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}

Solutions

3FoundationMCQExothermic vs endothermic reactions1 markPaper 1~2 min

A student dissolves a solid salt in water at

20\ ^\circ\text{C}

. After stirring, the solution temperature falls to

14\ ^\circ\text{C}

. Which statement correctly describes thermochemistry of the dissolving process?

Solutions

4FoundationMCQExothermic vs endothermic reactions1 markPaper 1~2 min

An energy profile for a reaction shows the reactants at a higher energy level than the products, with the arrow representing

\Delta H

pointing downward from reactants to products. Which statement correctly describes this reaction?

Solutions

5FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

A student mixes

50.0 \ \text{cm}^3

of

1.00 \ \text{mol dm}^{-3}

hydrochloric acid with

50.0 \ \text{cm}^3

of

1.00 \ \text{mol dm}^{-3}

sodium hydroxide solution in a polystyrene cup. Both solutions have an initial temperature of

21.5 \ ^\circ\text{C}

, and the maximum temperature recorded after mixing is

28.0 \ ^\circ\text{C}

. The specific heat capacity of the solution is

4.18 \ \text{J g}^{-1} \text{K}^{-1}

and its density is

1.00 \ \text{g cm}^{-3}

. What is the molar enthalpy change of neutralisation, in

\text{kJ mol}^{-1}

?

Solutions

6MasteryMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

A student mixes

50.0 \ \text{cm}^3

of

1.00 \ \text{mol dm}^{-3}

hydrochloric acid with

50.0 \ \text{cm}^3

of

1.00 \ \text{mol dm}^{-3}

sodium hydroxide in a polystyrene cup of heat capacity

30 \ \text{J K}^{-1}

. The temperature rises by

6.0 \ \text{K}

. Using

q = mc\Delta T

with

c = 4.18 \ \text{J g}^{-1} \text{K}^{-1}

and solution density

1.00 \ \text{g cm}^{-3}

, but neglecting the heat absorbed by the cup, what is the error in the student's calculated enthalpy change per mole of water formed?

Solutions

7FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

The standard enthalpy of combustion of hydrogen is

-286\ \text{kJ mol}^{-1}

. An enthalpy profile is drawn for the reaction

2\text{H}_2\text{(g)} + \text{O}_2\text{(g)} \rightarrow 2\text{H}_2\text{O(l)}

, with the reactants set at the reference level of

0\ \text{kJ}

. Which statement correctly describes this reaction and its enthalpy profile?

Solutions

8FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

The Haber process reaction

\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)}

is carried out in a thermally insulated (adiabatic) reactor. Which observation would best provide experimental evidence that the forward reaction is exothermic?

Solutions

9FoundationMCQFuels and their efficiency1 markPaper 1~2 min

A student compares the heating efficiency of two fuels in a calorimetry experiment. Fuel A releases

50 \ \text{kJ g}^{-1}

when combusted, of which

20 \ \text{kJ g}^{-1}

is transferred to the water. Fuel B releases

40 \ \text{kJ g}^{-1}

when combusted, of which

30 \ \text{kJ g}^{-1}

is transferred to the water. Which fuel has the greater heating efficiency, and what is that efficiency?

Solutions

10MasteryMCQCombustion reactions and energy release1 markPaper 1~2 min

The complete combustion of methane releases

890 \text{ kJ mol}^{-1}

. When

0.0500 \text{ mol}

of methane undergoes complete combustion, all the energy released is transferred to

500 \text{ g}

of water. What is the expected temperature rise of the water? (Specific heat capacity of water

= 4.18 \text{ J g}^{-1} \text{ K}^{-1}

)

Solutions

11FoundationMCQFuels and their efficiency1 markPaper 1~2 min

A student burns equal masses of propane (

\text{C}_3\text{H}_8

) and butane (

\text{C}_4\text{H}_{10}

) in separate trials using the same calorimeter containing the same mass of water. Burning

0.50\ \text{g}

of propane raises the water temperature by

15.0\ ^\circ\text{C}

, while burning

0.50\ \text{g}

of butane raises it by

19.5\ ^\circ\text{C}

. Assuming complete combustion and no heat loss, which statement correctly compares the energy content per gram of the two fuels?

Solutions

12FoundationMCQFuels and their efficiency1 markPaper 1~2 min

A car engine burning octane (

\text{C}_8\text{H}_{18}

) produces significant quantities of nitrogen oxides (

\text{NO}_x

) in its exhaust, whereas a Bunsen burner burning methane (

\text{CH}_4

) produces negligible

\text{NO}_x

under normal operating conditions. Neither fuel contains nitrogen. What is the most likely reason for the formation of

\text{NO}_x

in the car engine?

Solutions

13FoundationMCQSpontaneous vs non-spontaneous processes1 markPaper 1~2 min

For the Haber process,

\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)}

,

\Delta H = -92.6 \ \text{kJ mol}^{-1}

. The reaction is spontaneous at 298 K but becomes non-spontaneous at sufficiently high temperatures. Which statement correctly accounts for this behaviour?

Solutions

14MasteryMCQEntropy and the second law of thermodynamics1 markPaper 1~2 min

The thermal decomposition of ammonium chloride is represented by the equation:

\text{NH}_4\text{Cl}(s) \rightarrow \text{NH}_3(g) + \text{HCl}(g) \quad \Delta H^\circ = +176 \text{ kJ mol}^{-1}

The reaction is non-spontaneous below approximately 340 °C but becomes spontaneous above this temperature. Which statement correctly explains this behaviour in terms of entropy changes?

Solutions

15FoundationMCQSpontaneous vs non-spontaneous processes1 markPaper 1~2 min

At 298 K, ammonium nitrate dissolves endothermically in water:

\text{NH}_4\text{NO}_3\text{(s)} \rightarrow \text{NH}_4^+\text{(aq)} + \text{NO}_3^-\text{(aq)}

, with

\Delta H = +25.7 \text{ kJ mol}^{-1}

and

\Delta S = +108 \text{ J K}^{-1}\text{ mol}^{-1}

. Which statement about the spontaneity of this process at 298 K is correct?

Solutions

16FoundationMCQSpontaneous vs non-spontaneous processes1 markPaper 1~2 min

At 298 K, the combustion of methane has

\Delta H = -890 \ \text{kJ mol}^{-1}

and

\Delta S = -242 \ \text{J K}^{-1} \text{mol}^{-1}

. What is the value of

\Delta G

at this temperature, and what does it indicate about spontaneity?

Solutions

17ChallengeSAQ-LExothermic vs endothermic reactions9 marksPaper 2~14 min

Data

| Bond enthalpy /

\text{kJ mol}^{-1}

|

\text{N} \equiv \text{N}

—

945

\text{H} - \text{H}

—

436

\text{N} - \text{H}

—

391

The Haber process produces ammonia:

\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)}

. The forward reaction is exothermic. Bond enthalpy

(a)

State what is meant by the term average bond enthalpy. [1 mark]

(b)

Determine the enthalpy change,

\Delta H

, for the forward reaction using the bond enthalpy data. State whether the reaction is exothermic or endothermic. [3 marks]

(c)

Explain why the value of

\Delta H

calculated using bond enthalpies differs from the value measured experimentally. [2 marks]

(d)

The Haber process is operated at approximately

450\,^\circ\text{C}

and

200\,\text{atm}

, using an iron catalyst. Evaluate the choice of

450\,^\circ\text{C}

, referring to reaction rate, equilibrium yield, and the sign of

\Delta H

. [3 marks]

Solutions

18ChallengeSAQ-LExothermic vs endothermic reactions8 marksPaper 2~12 min

A student investigates thermal decomposition of sodium hydrogencarbonate, which occurs when baking soda is heated:

2\text{NaHCO}_3\text{(s)} \rightarrow \text{Na}_2\text{CO}_3\text{(s)} + \text{H}_2\text{O(g)} + \text{CO}_2\text{(g)}

The student places

5.00\,\text{g}

of

\text{NaHCO}_3

in a test tube and heats it strongly. The mass of the test tube and contents decreases from

25.80\,\text{g}

to

22.55\,\text{g}

after complete decomposition. Molar masses /

\text{g mol}^{-1}

:

\text{NaHCO}_3 = 84.0

,

\text{Na}_2\text{CO}_3 = 106.0

,

\text{H}_2\text{O} = 18.0

,

\text{CO}_2 = 44.0

(a)

Determine the number of moles of

\text{NaHCO}_3

that decomposed, using the mass loss observed in the experiment. [2 marks]

(b)

Calculate the molar enthalpy change,

\Delta H

, for the decomposition of

1.00\,\text{mol}

of

\text{NaHCO}_3

, given that the reaction absorbs

4.50\,\text{kJ}

for every

1.00\,\text{g}

of

\text{NaHCO}_3

decomposed. [2 marks]

(c)

State whether the decomposition is exothermic or endothermic and justify your answer using your result from (b). [1 mark]

(d)

The student suggests that

\text{NaHCO}_3

could replace ammonium nitrate (

\text{NH}_4\text{NO}_3

) in a commercial cold pack.

\text{NH}_4\text{NO}_3

dissolves endothermically with

\Delta H = +25.7\,\text{kJ mol}^{-1}

(

M = 80.0\,\text{g mol}^{-1}

). Evaluate this suggestion by calculating the enthalpy absorbed per gram for each substance and considering one practical factor. [3 marks]

Solutions

19ChallengeSAQ-LExothermic vs endothermic reactions9 marksPaper 2~14 min

Data

- Density of

HCl(aq)

:

1.00 \ \text{g cm}^{-3}

- Specific heat capacity of solution:

4.18 \ \text{J g}^{-1} \text{K}^{-1}

- Molar mass of

CaO

:

56.08 \ \text{g mol}^{-1}

- Molar mass of

CaCO_3

:

100.09 \ \text{g mol}^{-1}

This question is about thermal decomposition of calcium carbonate, an important industrial process for producing calcium oxide (quicklime).

CaCO_3(s) \rightarrow CaO(s) + CO_2(g)

The following standard enthalpy of formation data are available: Substance —

\Delta H_f^\circ \ / \ \text{kJ mol}^{-1}

CaCO_3(s)

—

-1206.9

CaO(s)

—

-635.1

CO_2(g)

—

-393.5

A student attempts to verify this value experimentally by dissolving

2.80 \ \text{g}

of

CaO(s)

in excess dilute hydrochloric acid in a polystyrene cup calorimeter. The temperature of

100.0 \ \text{cm}^3

of

1.00 \ \text{mol dm}^{-3}

HCl(aq)

rises from

21.4 \ ^\circ C

to

29.6 \ ^\circ C

.

CaO(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) \quad \Delta H_1

The standard enthalpy change for the reaction of

CaCO_3(s)

with excess dilute hydrochloric acid is known to be

\Delta H_2 = -15.9 \ \text{kJ mol}^{-1}

:

CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g) \quad \Delta H_2

(a)

Calculate the standard enthalpy change for the decomposition of

CaCO_3(s)

,

\Delta H^\circ_{decomp}

, using the formation data in the table. [2 marks]

(b)

Calculate the enthalpy change

\Delta H_1

for the reaction of

CaO(s)

with excess

HCl(aq)

, using the experimental data from the calorimetry experiment. Give your answer in

\text{kJ mol}^{-1}

. [3 marks]

(c)

Using your answers to (a) and (b), and the given value of

\Delta H_2

, construct a Hess's law cycle to determine an experimental value for

\Delta H^\circ_{decomp}

. Evaluate whether the experimental value supports theoretical value, identifying one significant source of systematic error in the calorimetry experiment and explaining the direction of its effect on

\Delta H_1

. [4 marks]

Solutions

20ChallengeSAQ-LExothermic vs endothermic reactions7 marksPaper 2~11 min

Data

- Density of the solution:

1.00\,\text{g cm}^{-3}

- Specific heat capacity of the solution:

4.18\,\text{J g}^{-1}\text{K}^{-1}

- Assume the catalyst adds negligible mass and volume.

Hydrogen peroxide,

\text{H}_2\text{O}_2

, decomposes slowly at room temperature in an exothermic reaction:

2\text{H}_2\text{O}_2(l) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)

A student investigates this decomposition using a calorimeter. They place

50.0\,\text{cm}^3

of

0.500\,\text{mol dm}^{-3}

aqueous hydrogen peroxide in a polystyrene cup and measure its temperature as

20.0\,^\circ\text{C}

. They add a small spatula of manganese(IV) oxide catalyst,

\text{MnO}_2

, which causes rapid decomposition. The maximum temperature reached is

26.8\,^\circ\text{C}

.

(a)

Calculate the enthalpy change,

\Delta H

, in

\text{kJ mol}^{-1}

, for the decomposition of hydrogen peroxide per mole of

\text{H}_2\text{O}_2

. [3 marks]

(b)

The student claims that adding more

\text{MnO}_2

catalyst would increase the value of

\Delta H

measured. Explain whether this claim is correct. [2 marks]

(c)

The student repeats the experiment using a copper calorimeter instead of a polystyrene cup. The copper calorimeter has a mass of

45.0\,\text{g}

and a specific heat capacity of

0.385\,\text{J g}^{-1}\text{K}^{-1}

. Determine the corrected value of

\Delta H

in

\text{kJ mol}^{-1}

using the same temperature data, and explain why this value differs from your answer in (a). [2 marks]

Solutions

21MasterySAQ-SThermochemistry and enthalpy diagrams6 marksPaper 2~9 min

The enthalpy below (not to scale) represents the reaction pathway for the formation of hydrogen iodide:

\text{H}_2(\text{g}) + \text{I}_2(\text{g}) \rightarrow 2\text{HI}(\text{g})

The activation energy for the forward reaction is

165\,\text{kJ mol}^{-1}

and the enthalpy change for the forward reaction is

+9.4\,\text{kJ mol}^{-1}

. The reactants are assigned an enthalpy value of

0\,\text{kJ mol}^{-1}

.

(a)

State the enthalpy values, in

\text{kJ mol}^{-1}

, of the products and of the transition state. [2 marks]

(b)

Explain, using collision theory, why increasing the temperature increases the rate of the forward reaction. [2 marks]

(c)

Deduce the activation energy for the reverse reaction and explain whether the reverse reaction is more or less sensitive to a given temperature increase than the forward reaction. [2 marks]

Solutions

22MasterySAQ-SThermochemistry and enthalpy diagrams5 marksPaper 2~8 min

The thermite reaction is highly exothermic and is used in welding:

2\text{Al}(s) + \text{Fe}_2\text{O}_3(s) \rightarrow \text{Al}_2\text{O}_3(s) + 2\text{Fe}(l) \qquad \Delta H = -852\,\text{kJ mol}^{-1}

A student performs a calorimetry experiment using

0.100\,\text{mol}

of Al with excess

\text{Fe}_2\text{O}_3

. The heat released raises the temperature of

500\,\text{g}

of surrounding material (specific heat capacity

0.900\,\text{J g}^{-1}\text{K}^{-1}

) from

25.0\,°\text{C}

to

95.0\,°\text{C}

.

(a)

State what the sign of

\Delta H

indicates about the energy of the products relative to the reactants. [1 mark]

(b)

Calculate the percentage of theoretical heat released that is absorbed by the surrounding material. [2 marks]

(c)

Explain why thermite reaction, once initiated, is self-sustaining. [2 marks]

Solutions

23ChallengeSAQ-LThermochemistry and enthalpy diagrams7 marksPaper 2~11 min

Data

q = mc\Delta T

;

c_{\text{water}} = 4.18\,\text{J}\,\text{g}^{-1}\,\text{K}^{-1}

;

M(\text{C}_2\text{H}_5\text{OH}) = 46.08\,\text{g}\,\text{mol}^{-1}

;

M(\text{C}_8\text{H}_{18}) = 114.23\,\text{g}\,\text{mol}^{-1}

;

\Delta H_c(\text{C}_8\text{H}_{18}) = -5470\,\text{kJ}\,\text{mol}^{-1}

.

A student uses a spirit burner to combust ethanol,

\text{C}_2\text{H}_5\text{OH}(\text{l})

, heating

200.0\,\text{g}

of water in a copper calorimeter. The water temperature rises from

22.5\,°\text{C}

to

55.0\,°\text{C}

. The mass of ethanol burned is

0.85\,\text{g}

. The literature standard enthalpy of combustion of ethanol is

-1367\,\text{kJ}\,\text{mol}^{-1}

.

(a)

(i) Calculate the experimental enthalpy of combustion of ethanol. [2]
(ii) State one assumption made in this calculation. [1 mark]

(b)

(i) Draw an enthalpy level for the combustion of ethanol, showing both the literature and experimental values. [1]
(ii) Explain why the experimental value is less exothermic than the literature value. [1 mark]

(c)

A student claims that ethanol is a better fuel than octane because its experimental enthalpy of combustion, expressed in

\text{kJ}\,\text{g}^{-1}

, is closer to the literature value. Evaluate this claim by calculating the energy density (

\text{kJ}\,\text{g}^{-1}

) of both fuels using literature values, and assess whether the experimental data support the student's conclusion. [2 marks]

Solutions

24MasterySAQ-SThermochemistry and enthalpy diagrams5 marksPaper 2~8 min

The Haber process produces ammonia according to:

\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g}) \quad \Delta H = -92\,\text{kJ mol}^{-1}

The uncatalysed forward reaction has an activation energy of

+180\,\text{kJ mol}^{-1}

. An iron catalyst is used industrially.

(a)

State the relative energy levels of reactants and products for this reaction. [1 mark]

(b)

The iron catalyst lowers the activation energy of the forward reaction to

+130\,\text{kJ mol}^{-1}

. Determine the activation energy for the reverse (catalysed) reaction. [2 marks]

(c)

Explain why the use of a catalyst does not change the value of

\Delta H

for the reaction. [2 marks]

Solutions

25ChallengeSAQ-LCombustion reactions and energy release7 marksPaper 2~11 min

Data

\Delta H_f^\ominus(\text{CO}_2(g)) = -393.5\,\text{kJ mol}^{-1}

\Delta H_f^\ominus(\text{CO}(g)) = -110.5\,\text{kJ mol}^{-1}

\Delta H_f^\ominus(\text{H}_2\text{O}(l)) = -285.8\,\text{kJ mol}^{-1}

\Delta H_f^\ominus(\text{CH}_4(g)) = -74.8\,\text{kJ mol}^{-1}

Methane (

\text{CH}_4

) is a common fuel used in domestic heating. The complete combustion of methane is represented by:

\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \qquad \Delta H_c^\ominus = -890\,\text{kJ mol}^{-1}

In poorly ventilated gas heaters, incomplete combustion occurs, producing carbon monoxide:

\text{CH}_4(g) + \tfrac{3}{2}\text{O}_2(g) \rightarrow \text{CO}(g) + 2\text{H}_2\text{O}(l)

(a)

Calculate the enthalpy change, in

\text{kJ mol}^{-1}

, for the incomplete combustion of methane using the standard enthalpies of formation below. [3 marks]

(b)

Explain, in terms of bond breaking and bond forming, why incomplete combustion of methane releases less energy per mole than complete combustion. [2 marks]

(c)

A household uses a gas heater in a poorly ventilated room. Evaluate whether the reduced energy output and the nature of the combustion products make incomplete combustion more hazardous than simply being less efficient. [2 marks]

Solutions

26MasterySAQ-SFuels and their efficiency7 marksPaper 2~11 min

Methanol (

\text{CH}_3\text{OH}

) and hydrogen (

\text{H}_2

) are being considered as fuels for vehicles. The enthalpy of combustion of methanol is

-726\,\text{kJ mol}^{-1}

and for hydrogen is

-286\,\text{kJ mol}^{-1}

. - Density of liquid methanol

= 0.792\,\text{g cm}^{-3}

- Molar mass of methanol

= 32.04\,\text{g mol}^{-1}

- Molar mass of hydrogen

= 2.02\,\text{g mol}^{-1}

- At standard conditions,

1

mole of hydrogen gas occupies

24.5\,\text{dm}^3

(a)

State two factors, other than energy content per mole, that must be considered when comparing these two fuels for use in vehicles. [2 marks]

(b)

Calculate the energy released per gram of liquid methanol. [1 mark]

(c)

Calculate the energy released per cubic centimetre of liquid methanol. [1 mark]

(d)

The fuel tank of a vehicle holds

50.0\,\text{dm}^3

of liquid methanol. Calculate the volume of hydrogen gas, at standard conditions, that would need to be stored to provide the same total energy as a full tank of methanol. [3 marks]

Solutions

27ChallengeSAQ-LCombustion reactions and energy release7 marksPaper 2~11 min

A researcher investigates the use of biodiesel (methyl octadecanoate,

\text{C}_{19}\text{H}_{38}\text{O}_2

) as a renewable fuel. The complete combustion of biodiesel is:

\text{C}_{19}\text{H}_{38}\text{O}_2(l) + \frac{53}{2}\text{O}_2(g) \rightarrow 19\text{CO}_2(g) + 19\text{H}_2\text{O}(l)

The researcher burns

5.00\,\text{g}

of biodiesel in a bomb calorimeter with a heat capacity of

8.50\,\text{kJ K}^{-1}

. The temperature of the calorimeter increases by

12.4\,\text{K}

.

(a)

Calculate the enthalpy change of combustion of biodiesel, in

\text{kJ mol}^{-1}

. Include the sign and state whether the reaction is exothermic or endothermic. [3 marks]

(b)

Explain two reasons why the value obtained in (a) may differ from a value calculated using average bond enthalpies. [2 marks]

(c)

A transport company must choose between biodiesel and diesel (a mixture of

\text{C}_{12}

–

\text{C}_{15}

alkanes from crude oil) as fuel. Evaluate which fuel is the more sustainable choice, considering both energy density and carbon neutrality. Justify your conclusion. [2 marks]

Solutions

28MasterySAQ-SFuels and their efficiency9 marksPaper 2~14 min

Data

- Energy content of coal

= 32.5\,\text{kJ g}^{-1}

- Enthalpy of combustion of sulfur:

\text{S}(s) + \text{O}_2(g) \rightarrow \text{SO}_2(g)

,

\Delta H^\circ = -297\,\text{kJ mol}^{-1}

- Molar mass of sulfur

= 32.06\,\text{g mol}^{-1}

; molar mass of

\text{SO}_2 = 64.06\,\text{g mol}^{-1}

- The power station burns

1000

tonnes of coal per day (

1\,\text{tonne} = 10^6\,\text{g}

)

A power station burns coal to generate electricity. The coal contains

2.5\%

sulfur by mass. When burned, sulfur reacts with oxygen to form sulfur dioxide,

\text{SO}_2

, a major contributor to acid rain.

(a)

State two environmental problems caused by burning coal, other than the production of

\text{SO}_2

. [2 marks]

(b)

Calculate the mass, in grams, of

\text{SO}_2

produced per day. [2 marks]

(c)

Calculate the percentage of the total daily energy output of the coal that is contributed by the combustion of sulfur. [3 marks]

(d)

A flue gas desulfurization (FGD) system removes

95\%

of the

\text{SO}_2

produced. Evaluate whether removing

\text{SO}_2

from flue gases is worthwhile, given that the FGD system itself consumes energy. [2 marks]

Solutions

29ChallengeSAQ-LSpontaneous vs non-spontaneous processes8 marksPaper 2~12 min

This question is about thermodynamics of the industrial production of ammonia via the Haber–Bosch process.

\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) \qquad \Delta H^\circ = -92.6\,\text{kJ mol}^{-1}

(a)

State the sign of

\Delta S^\circ

for this reaction. [1 mark]

(b)

Using the standard molar entropy data below, calculate

\Delta G^\circ

for the reaction at

298\,\text{K}

. Species —

S^\circ\,/\,\text{J K}^{-1}\text{mol}^{-1}

\text{N}_2(g)

—

191.6

\text{H}_2(g)

—

130.7

\text{NH}_3(g)

—

192.8

[3 marks]

(c)

(i) Explain why the Haber–Bosch process is operated at approximately

700\,\text{K}

rather than

298\,\text{K}

, even though

\Delta G^\circ

is negative at

298\,\text{K}

. [2]
(ii) Using your answer to (b) and the expression

\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ

, estimate

\Delta G^\circ

at

700\,\text{K}

and evaluate whether the reaction remains thermodynamically feasible at this temperature. [2 marks]

Solutions

30ChallengeSAQ-LSpontaneous vs non-spontaneous processes8 marksPaper 2~12 min

Data

- Specific heat capacity of solution

= 4.18\,\text{J}\,\text{g}^{-1}\,\text{K}^{-1}

- Density of solution

= 1.00\,\text{g}\,\text{cm}^{-3}

- Molar mass of

\text{NH}_4\text{NO}_3 = 80.0\,\text{g}\,\text{mol}^{-1}

- Assume no heat is lost to the surroundings

A student investigates the spontaneity of the dissolution of ammonium nitrate in water:

\text{NH}_4\text{NO}_3(s) \rightarrow \text{NH}_4^+(aq) + \text{NO}_3^-(aq)

The student dissolves

10.0\,\text{g}

of

\text{NH}_4\text{NO}_3

in

100\,\text{cm}^3

of water at

298\,\text{K}

and records a temperature drop of

3.2\,\text{K}

.

(a)

Explain why the observed temperature decrease indicates that the dissolution is endothermic (

\Delta H > 0

), yet the process is spontaneous. [3 marks]

(b)

Calculate the enthalpy change of dissolution,

\Delta H

, per mole of

\text{NH}_4\text{NO}_3

. [3 marks]

(c)

Calculate the entropy change of the surroundings,

\Delta S_\text{surr}

, when

1.00\,\text{mol}

of

\text{NH}_4\text{NO}_3

dissolves at

298\,\text{K}

. [2 marks]

Solutions

31ChallengeSAQ-LSpontaneous vs non-spontaneous processes8 marksPaper 2~12 min

The combustion of ethanol is used as a source of energy in some fuel blends:

\text{C}_2\text{H}_5\text{OH}(l) + 3\text{O}_2(g) \rightarrow 2\text{CO}_2(g) + 3\text{H}_2\text{O}(l)

(a)

Predict the sign of the standard entropy change of the system,

\Delta S^\circ_{\text{sys}}

, for this reaction. Justify your answer in terms of the chemical equation. The following data apply at

298\,\text{K}

:

\Delta G^\circ = -1325\,\text{kJ mol}^{-1} \qquad \Delta S^\circ = -139\,\text{J K}^{-1}\text{mol}^{-1}

[1 mark]

(b)

Calculate the standard enthalpy change,

\Delta H^\circ

, for this reaction in

\text{kJ mol}^{-1}

. [2 marks]

(c)

Explain why ethanol does not spontaneously combust when left in an open container at

298\,\text{K}

, despite

\Delta G^\circ

being negative. [3 marks]

(d)

The value of

\Delta G^\circ

for this reaction becomes less negative as temperature increases. Using the relationship

\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ

, deduce and explain whethere exists a temperature at which this reaction becomes non-spontaneous. [2 marks]

Solutions

32ChallengeSAQ-LSpontaneous vs non-spontaneous processes8 marksPaper 2~12 min

Data

\Delta H^\circ = +176\,\text{kJ}\,\text{mol}^{-1}

;

\Delta S^\circ = +285\,\text{J}\,\text{K}^{-1}\,\text{mol}^{-1}

Assume

\Delta H^\circ

and

\Delta S^\circ

do not vary with temperature.

The decomposition of ammonium chloride is non-spontaneous at

298\,\text{K}

but becomes spontaneous at higher temperatures.

\text{NH}_4\text{Cl}(s) \rightarrow \text{NH}_3(g) + \text{HCl}(g)

(a)

State the sign of

\Delta S

for this reaction. Justify your answer with reference to the states of matter involved. [2 marks]

(b)

Calculate the minimum temperature, in K, at which the decomposition becomes spontaneous. [3 marks]

(c)

The decomposition is carried out at

700\,\text{K}

in a sealed container. At this temperature,

K = 1.86

. (i) Calculate

\Delta G^\circ

at

700\,\text{K}

. [1]
(ii) At a certain point during the reaction, the reaction quotient

Q = 3.50

. Determine whether the forward reaction is spontaneous at this point. Show your working. [2 marks]

Solutions

33ChallengeLAQExothermic vs endothermic reactions15 marksPaper 3~23 min

Data

- Bond enthalpies:

\text{H–H} = 436\,\text{kJ mol}^{-1}

;

\text{I–I} = 151\,\text{kJ mol}^{-1}

;

\text{H–I} = 299\,\text{kJ mol}^{-1}

- Equilibrium constant

K_c

at

500\,\text{K}

:

160

The industrial preparation of hydrogen iodide from hydrogen and iodine is an equilibrium reaction studied for its thermodynamic properties.

\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\,\text{HI}(g)

A student adds

0.500\,\text{mol}

of

\text{H}_2(g)

and

0.500\,\text{mol}

of

\text{I}_2(g)

to a constant-pressure calorimeter containing

2.00\,\text{kg}

of inert oil (specific heat capacity

= 2.50\,\text{J g}^{-1}\text{K}^{-1}

). The initial temperature of the oil is

25.0\,^\circ\text{C}

. After the reaction proceeds, the final temperature of the oil is

22.3\,^\circ\text{C}

.

(a)

State, with reference to the temperature change of the oil, whether the reaction is endothermic or exothermic. [1 mark]

(b)

Calculate the enthalpy change

\Delta H

for the reaction as written, using the calorimetry data. State the sign convention used. [3 marks]

(c)

Using the bond enthalpy data, calculate a theoretical value for

\Delta H

for the reaction as written. [3 marks]

(d)

Explain two reasons why the value obtained in (b) differs from the value obtained in (c). [4] (e) The equilibrium constant

K_c = 160

at

500\,\text{K}

. Using your answer from (c) and a consideration of

\Delta S

for this reaction, evaluate whether this reaction is better described as "enthalpy-driven" or "entropy-driven". [4 marks]

Solutions

34ChallengeLAQHess's Law and enthalpy cycles9 marksPaper 3~14 min

The industrial synthesis of hydrogen gas from methane and steam is known as steam reforming:

\text{CH}_4(g) + \text{H}_2\text{O}(g) \rightarrow \text{CO}(g) + 3\text{H}_2(g)

The following standard enthalpies of combustion are available:

\Delta H_c^\circ[\text{CH}_4(g)] = -890\,\text{kJ mol}^{-1}

\Delta H_c^\circ[\text{H}_2(g)] = -286\,\text{kJ mol}^{-1}

\Delta H_c^\circ[\text{CO}(g)] = -283\,\text{kJ mol}^{-1}

(a)

State the value of

\Delta H_c^\circ[\text{H}_2\text{O}(g)]

and justify your answer. [1 mark]

(b)

Calculate the standard enthalpy change for the steam reforming reaction using the combustion data above. [3 marks]

(c)

The reaction has a high activation energy,

E_a

. Using the Maxwell–Boltzmann energy distribution, explain why a temperature of

800

–

1000\,^\circ\text{C}

and a nickel catalyst are both required for a practical rate of reaction. [3 marks]

(d)

Using the value calculated in (b), explain why increasing temperature also improves the equilibrium yield of

\text{H}_2(g)

in this process. [2 marks]

Solutions

35ChallengeLAQThermochemistry and enthalpy diagrams15 marksPaper 3~23 min

Data

Quantity — Value Molar mass of

\text{NaHCO}_3

—

84.0\,\text{g}\,\text{mol}^{-1}

\Delta H_f^\circ\,[\text{NaHCO}_3\text{(s)}]

—

-950.8\,\text{kJ}\,\text{mol}^{-1}

\Delta H_f^\circ\,[\text{Na}_2\text{CO}_3\text{(s)}]

—

-1130.7\,\text{kJ}\,\text{mol}^{-1}

\Delta H_f^\circ\,[\text{H}_2\text{O(l)}]

—

-285.8\,\text{kJ}\,\text{mol}^{-1}

\Delta H_f^\circ\,[\text{CO}_2\text{(g)}]

—

-393.5\,\text{kJ}\,\text{mol}^{-1}

Standard enthalpy of solution of

\text{NaHCO}_3\text{(s)}

—

+16.4\,\text{kJ}\,\text{mol}^{-1}

A student investigates the enthalpy change for the decomposition of sodium hydrogencarbonate,

\text{NaHCO}_3\text{(s)}

, which is used in baking powder.

2\text{NaHCO}_3\text{(s)} \rightarrow \text{Na}_2\text{CO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}

In a separate experiment, the student dissolves

4.20\,\text{g}

of

\text{NaHCO}_3\text{(s)}

in

50.0\,\text{g}

of water initially at

25.0\,°\text{C}

. The minimum temperature recorded is

18.2\,°\text{C}

. The specific heat capacity of the solution is

4.18\,\text{J}\,\text{g}^{-1}\,\text{K}^{-1}

.

(a)

Calculate the standard enthalpy change,

\Delta H^\circ

, for the decomposition of

\text{NaHCO}_3\text{(s)}

using the standard enthalpies of formation. Give your answer in

\text{kJ}\,\text{mol}^{-1}

of

\text{NaHCO}_3

decomposed. [3 marks]

(b)

Using the calorimetry data, calculate the experimental enthalpy of solution of

\text{NaHCO}_3\text{(s)}

in

\text{kJ}\,\text{mol}^{-1}

. [3 marks]

(c)

The student claims the calorimetry experiment in (b) can be used directly to determine the enthalpy change for the decomposition reaction. Deduce, with reference to Hess's law, why this claim is incorrect. [3 marks]

(d)

The decomposition of

\text{NaHCO}_3\text{(s)}

is endothermic yet occurs spontaneously during baking at approximately

180\,°\text{C}

. Using the relationship

\Delta G = \Delta H - T\Delta S

, estimate the minimum temperature at which the decomposition becomes thermodynamically spontaneous, and evaluate whether this consistent with baking conditions. Assume

\Delta S^\circ = +0.185\,\text{kJ}\,\text{mol}^{-1}\,\text{K}^{-1}

per mole of

\text{NaHCO}_3

decomposed. [6 marks]

Solutions

36ChallengeLAQThermochemistry and enthalpy diagrams15 marksPaper 3~23 min

Data

Bond enthalpies /

\text{kJ mol}^{-1}

:

\text{C} \equiv \text{O}

:

1072

; H–H:

436

; C–H:

413

; C–O:

358

; O–H:

463

Standard enthalpies of formation /

\text{kJ mol}^{-1}

: Substance —

\Delta H_f^\circ

\text{CO}(\text{g})

—

-110.5

\text{CO}_2(\text{g})

—

-393.5

\text{H}_2\text{O}(\text{l})

—

-285.8

\text{CH}_3\text{OH}(\text{l})

— unknown Two catalysts are under consideration. Their enthalpy profile data are: Forward activation energy /

\text{kJ mol}^{-1}

—

85

—

62

Intermediate enthalpy relative to reactants /

\text{kJ mol}^{-1}

—

+45

—

+22

A chemical company is developing an industrial process to produce methanol,

\text{CH}_3\text{OH}(\text{l})

, from synthesis gas:

\text{CO}(\text{g}) + 2\text{H}_2(\text{g}) \rightarrow \text{CH}_3\text{OH}(\text{l}) \quad \Delta H = -128\,\text{kJ mol}^{-1}

(a)

Calculate the enthalpy change for

\text{CO}(\text{g}) + 2\text{H}_2(\text{g}) \rightarrow \text{CH}_3\text{OH}(\text{l})

using bond enthalpy data. [3 marks]

(b)

(i) The combustion of methanol is:

\text{CH}_3\text{OH}(\text{l}) + \frac{3}{2}\text{O}_2(\text{g}) \rightarrow \text{CO}_2(\text{g}) + 2\text{H}_2\text{O}(\text{l})

Use the standard enthalpy of formation data provided to determine

\Delta H_f^\circ

of

\text{CH}_3\text{OH}(\text{l})

, given that the standard enthalpy of combustion of methanol is

-726.0\,\text{kJ mol}^{-1}

. [3]
(ii) Explain why the bond enthalpy value calculated in (a) differs from the value

-128\,\text{kJ mol}^{-1}

. [3 marks]

(c)

Evaluate the choice of Catalyst B over Catalyst A for industrial methanol production. In your answer, calculate the reverse activation energy for each catalyst and discuss the implications for yield and rate. [6 marks]

Solutions

37ChallengeLAQEnergy density of different fuels (e.g., hydrocarbons, biofuels)13 marksPaper 3~20 min

Data

- Methane (

\text{CH}_4

):

\Delta H_c^\circ = -890\,\text{kJ mol}^{-1}

, molar mass

= 16.04\,\text{g mol}^{-1}

- Octane (

\text{C}_8\text{H}_{18}

):

\Delta H_c^\circ = -5470\,\text{kJ mol}^{-1}

, molar mass

= 114.23\,\text{g mol}^{-1}

, density

= 0.703\,\text{g mL}^{-1}

- Ethanol (

\text{C}_2\text{H}_5\text{OH}

):

\Delta H_c^\circ = -1367\,\text{kJ mol}^{-1}

, molar mass

= 46.07\,\text{g mol}^{-1}

, density

= 0.789\,\text{g mL}^{-1}

- Methanol (

\text{CH}_3\text{OH}

): energy density

= 22.7\,\text{MJ kg}^{-1}

, density

= 0.792\,\text{g mL}^{-1}

The combustion of fossil fuels and biofuels provides energy for transportation and electricity generation. Energy density, measured in

\text{MJ kg}^{-1}

, is a critical parameter in fuel selection.

(a)

Calculate the energy density in

\text{MJ kg}^{-1}

for methane, octane, and ethanol. [3 marks]

(b)

Explain why the energy density of octane is greater than that of ethanol, in terms of molecular structure and bonding. [3 marks]

(c)

A fuel cell powered by methanol operates at an energy conversion efficiency of

45\%

. Evaluate whether replacing octane in a vehicle (engine efficiency

30\%

) with methanol would require a larger fuel tank to achieve the same driving range. Support your answer with calculations of useful energy per unit volume for each fuel. [4 marks]

(d)

State three reasons why biofuels such as ethanol are often considered more sustainable than fossil fuel hydrocarbons, despite having lower energy density. [3 marks]

Solutions

38ChallengeLAQEnergy density of different fuels (e.g., hydrocarbons, biofuels)15 marksPaper 3~23 min

Data

-

\Delta H_c^\circ

for

\text{H}_2\text{(g)} = -286\,\text{kJ mol}^{-1}

- Molar mass of

\text{H}_2 = 2.02\,\text{g mol}^{-1}

- Density of liquid hydrogen at

-253\,^\circ\text{C} = 70.8\,\text{kg m}^{-3}

- Density of gasoline

= 703\,\text{kg m}^{-3}

- Energy density of gasoline

= 47.9\,\text{MJ kg}^{-1}

-

1\,\text{m}^3 = 1000\,\text{L}

Hydrogen gas (

\text{H}_2

) is being considered as a clean fuel for vehicles. The energy density of a fuel is defined as the energy released per unit mass upon complete combustion.

(a)

Calculate the energy density of hydrogen gas in

\text{MJ kg}^{-1}

. [2 marks]

(b)

Calculate the volumetric energy density of liquid hydrogen in

\text{MJ L}^{-1}

. [2 marks]

(c)

Compare the volumetric energy densities of liquid hydrogen and gasoline. Determine how many times larger a liquid hydrogen fuel tank must be to store the same total energy as a gasoline tank. [3 marks]

(d)

A hydrogen fuel cell vehicle has an overall efficiency of

60\%

from fuel to wheels; a gasoline internal combustion engine vehicle has an efficiency of

25\%

. Using your answers from (b) and (c), evaluate whether the higher efficiency of the hydrogen fuel cell fully compensates for the lower volumetric energy density of liquid hydrogen. [4] (e) Explain why the mass-based energy density of hydrogen is high, despite the H–H bond enthalpy being only

436\,\text{kJ mol}^{-1}

. [4 marks]

Solutions

39ChallengeLAQGibbs free energy and its application to chemical reactions15 marksPaper 3~23 min

The industrial synthesis of ammonia via the Haber process is represented by:

N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)

At

298\,\text{K}

, the standard enthalpy change for the forward reaction is

\Delta H^\ominus = -92.6\,\text{kJ mol}^{-1}

and the standard entropy change is

\Delta S^\ominus = -198.3\,\text{J K}^{-1}\text{mol}^{-1}

.

(a)

Calculate the standard Gibbs free energy change,

\Delta G^\ominus

, for the forward reaction at

298\,\text{K}

. State the units of your answer. [3 marks]

(b)

(i) Calculate

\Delta G^\ominus

for the forward reaction at

700\,\text{K}

and state whether the reaction is thermodynamically spontaneous under standard conditions at this temperature. [2]
(ii) The Haber process is operated at approximately

700\,\text{K}

despite the result in (b)(i). Explain this choice of temperature, referring to both kinetics and thermodynamics. [3 marks]

(c)

Evaluate the following claim: *"Because

\Delta H^\ominus

is negative, the reaction is spontaneous at all temperatures."* In your answer: - state the sign of

\Delta S^\ominus

and explain its origin terms of the reaction - use

\Delta G = \Delta H - T\Delta S

to show spontaneity depends on temperature - calculate the temperature above which the reaction becomes non-spontaneous under standard conditions [5 marks]

(d)

Predict, with a reason, how the entropy of the system changes when the reaction is carried out at

200\,\text{atm}

rather than

1\,\text{atm}

. [2 marks]

Solutions

40ChallengeLAQGibbs free energy and its application to chemical reactions13 marksPaper 3~20 min

A student investigates thermal decomposition of calcium carbonate:

\text{CaCO}_3(s) \rightarrow \text{CaO}(s) + \text{CO}_2(g)

Data at

298\,\text{K}

: -

\Delta H^\ominus = +178\,\text{kJ mol}^{-1}

-

\Delta S^\ominus = +160.5\,\text{J K}^{-1}\text{mol}^{-1}

(a)

Calculate the standard Gibbs free energy change,

\Delta G^\ominus

, for this reaction at

298\,\text{K}

, and state whether the reaction is spontaneous at this temperature. [2 marks]

(b)

Determine the minimum temperature, in

^\circ\text{C}

, at which this reaction becomes spontaneous under standard conditions. [3 marks]

(c)

Explain why this endothermic reaction becomes spontaneous at high temperatures. In your answer, refer to the relative contributions of the enthalpy and entropy terms to

\Delta G

at different temperatures. [4 marks]

(d)

A student proposes: *"Increasing the partial pressure of

\text{CO}_2

above the solid mixture would lower the temperature at which the reaction becomes spontaneous."* Using the relationship

\Delta G = \Delta G^\ominus + RT\ln Q

, evaluate this proposal. In your answer, state the direction of any shift in equilibrium position and explain the effect on the minimum temperature required for spontaneity. Useful relationships:

\Delta G^\ominus = \Delta H^\ominus - T\Delta S^\ominus \qquad \Delta G = \Delta G^\ominus + RT\ln Q \qquad T(\text{K}) = T(^\circ\text{C}) + 273

[4 marks]

Solutions

Reactivity: What Drives Chemical Reactions? — Free Chemistry HL Practice Questions