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

1FoundationMCQExothermic vs endothermic reactions1 markPaper 1~2 min

In an energy level for a chemical reaction, the reactants are positioned at a higher energy level than the products. Which statement correctly describes this reaction?

Solutions

2MasteryMCQEnthalpy of reaction, heat capacity, and calorimetry1 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. Both solutions have an initial temperature of

21.5 \ ^\circ\text{C}

. The maximum temperature recorded after mixing is

28.3 \ ^\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

3MasteryMCQEnthalpy of reaction, heat capacity, and calorimetry1 markPaper 1~2 min

In a calorimetry experiment,

0.500\ \text{g}

of magnesium is added to

100.0\ \text{cm}^3

of

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

hydrochloric acid in a polystyrene cup. The reaction is:

\text{Mg(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}

The density of the solution is

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

and its specific heat capacity is

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

. The temperature rises by

12.5\ \text{°C}

. Which statement best explains why the enthalpy change calculated from this experiment is likely to be less exothermic than the true value?

Solutions

4FoundationMCQExothermic vs endothermic reactions1 markPaper 1~2 min

A student measures the temperature of a reaction mixture and observes a rise in temperature during a chemical reaction. Which statement correctly explains why this observation indicates an exothermic reaction?

Solutions

5MasteryMCQEnthalpy of reaction, heat capacity, and calorimetry1 markPaper 1~2 min

A student burns

0.450\ \text{g}

of propan-1-ol (

M_r = 60.0

) in a spirit burner beneath a calorimeter containing

200.0\ \text{g}

of water. The temperature of the water rises from

22.0\ °\text{C}

to

35.5\ °\text{C}

. Using a specific heat capacity of

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

for water only, the student calculates the enthalpy of combustion as

-1850\ \text{kJ mol}^{-1}

. The literature value is

-2020\ \text{kJ mol}^{-1}

. Which statement best explains why the experimental value is less exothermic than the literature value?

Solutions

6FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

In an enthalpy profile for the reaction A → B, the enthalpy of A is

+50 \ \text{kJ mol}^{-1}

, the enthalpy of B is

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

, and the activation energy for the forward reaction is

+80 \ \text{kJ mol}^{-1}

. What is the activation energy for the reverse reaction B → A?

Solutions

7MasteryMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

The reaction between zinc metal and aqueous copper(II) sulfate is represented by:

\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)} \quad \Delta H = -220 \text{ kJ mol}^{-1}

Which statement correctly describes the enthalpy profile for this reaction?

Solutions

8FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

An enthalpy profile for the combustion of methane,

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

, shows the reactants at

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

, the transition state at

+265 \ \text{kJ mol}^{-1}

, and the products at

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

. Which statement about this reaction is correct?

Solutions

9FoundationMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

An enthalpy profile shows reactants at

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

, products at

-150 \ \text{kJ mol}^{-1}

, and an activated complex at

+80 \ \text{kJ mol}^{-1}

. Which statement about this reaction is correct?

Solutions

10MasteryMCQThermochemistry and enthalpy diagrams1 markPaper 1~2 min

When

0.0500\ \text{mol}

of methane is completely combusted in excess oxygen,

44.5\ \text{kJ}

of heat is released. What is the molar enthalpy of combustion of methane?

Solutions

11FoundationMCQCombustion reactions and energy release1 markPaper 1~2 min

A student burns ethanol (

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

) in a simple calorimeter and obtains an experimental enthalpy change of combustion of

-1360 \ \text{kJ mol}^{-1}

. The literature value is

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

. Which explanation best accounts for the experimental value being less exothermic than the literature value?

Solutions

12MasteryMCQCombustion reactions and energy release1 markPaper 1~2 min

When 1.00 mol of methane undergoes complete combustion under standard conditions, the enthalpy change is

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

. The heat released by burning 0.0500 mol of methane is transferred entirely to 500 g of water with a specific heat capacity of

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

. What is the expected temperature rise of the water?

Solutions

13FoundationMCQCombustion reactions and energy release1 markPaper 1~2 min

Equal masses of

1.00\ \text{g}

of methane (

\text{CH}_4

), propane (

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

), and octane (

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

) are each burned completely in separate calorimeters containing identical volumes of water. Using the standard enthalpies of combustion below, which option correctly ranks the resulting temperature increases from smallest to largest? Fuel —

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

\text{CH}_4

—

-890

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

—

-2220

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

—

-5470

Solutions

14FoundationMCQCombustion reactions and energy release1 markPaper 1~2 min

A student burns

0.50\ \text{g}

of methane,

\text{CH}_4

, completely in excess oxygen. The standard enthalpy of combustion of methane is

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

. The energy released heats

200\ \text{g}

of water with a specific heat capacity of

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

. Assuming no heat loss to the surroundings, what is the expected temperature increase of the water? (Molar mass of

\text{CH}_4 = 16.0\ \text{g mol}^{-1}

)

Solutions

15MasteryMCQCombustion reactions and energy release1 markPaper 1~2 min

A student burns

0.50\ \text{g}

of ethanol (

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

,

M = 46.08\ \text{g mol}^{-1}

) in a spirit burner beneath a copper calorimeter containing

200\ \text{g}

of water. The water temperature rises from

22.0\ ^\circ\text{C}

to

34.5\ ^\circ\text{C}

. The student's calculated enthalpy of combustion is significantly less exothermic than theoretical value of

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

. Which explanation accounts for the largest discrepancy between the experimental and theoretical values?

Solutions

16FoundationMCQGibbs free energy and its application to chemical reactions1 markPaper 1~2 min

The decomposition of calcium carbonate,

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

, is non-spontaneous at 298 K but becomes spontaneous at 1200 K. Which statement correctly explains this change in spontaneity?

Solutions

17MasteryMCQEntropy and the second law of thermodynamics1 markPaper 1~2 min

In the Haber process,

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

, the forward reaction is exothermic. Which statement correctly describes the entropy change of the system and the temperature conditions under which the forward reaction is thermodynamically spontaneous?

Solutions

18MasteryMCQEntropy and the second law of thermodynamics1 markPaper 1~2 min

When ammonium nitrate,

\text{NH}_4\text{NO}_3\text{(s)}

, dissolves in water, the solution temperature drops noticeably. The dissolution is spontaneous at room temperature. Which statement correctly identifies

\Delta H

and

\Delta S

for the dissolution and thermodynamic reason for spontaneity?

Solutions

19FoundationMCQGibbs free energy and its application to chemical reactions1 markPaper 1~2 min

For the reaction

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

, the standard enthalpy and entropy changes are

\Delta H^\circ = -92 \text{ kJ mol}^{-1}

and

\Delta S^\circ = -198 \text{ J K}^{-1}\text{mol}^{-1}

. Which statement correctly describes the spontaneity of this reaction at 298 K and 500 K?

Solutions

20MasteryMCQEntropy and the second law of thermodynamics1 markPaper 1~2 min

For the reaction

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

, the Gibbs energy change

\Delta G

is positive at 298 K but becomes negative above approximately 1100 K. Using

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

, what are the signs of

\Delta H

and

\Delta S

for this reaction?

Solutions

21ChallengeSAQ-LEnthalpy of reaction, heat capacity, and calorimetry8 marksPaper 2~12 min

A student carries out a calorimetry experiment to determine the enthalpy change for the combustion of propan-1-ol,

\text{C}_3\text{H}_7\text{OH(l)}

. A spirit burner heats

200.0\,\text{g}

of water in a copper calorimeter of mass

50.0\,\text{g}

(including stirrer). The initial temperature of the water is

21.5\,°\text{C}

and the final temperature is

35.2\,°\text{C}

. The mass of propan-1-ol burned is

0.420\,\text{g}

. The specific heat capacity of water is

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

and the specific heat capacity of copper is

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

. The student assumes all heat released by combustion is absorbed by the water and calorimeter.

(a)

Calculate the heat absorbed by the water and copper calorimeter, in J. [1 mark]

(b)

Calculate the amount, in mol, of propan-1-ol burned. [1 mark]

(c)

Determine the experimental enthalpy change of combustion of propan-1-ol, in

\text{kJ mol}^{-1}

. [2 marks]

(d)

Explain two reasons why the experimental value obtained in (c) is less exothermic than the literature value of

-2021\,\text{kJ mol}^{-1}

. [2] (e) The student notices the wick of the spirit burner continues to absorb propan-1-ol after the flame is extinguished, causing the recorded mass loss to be greater than the mass actually burned. - Identify the type of error this represents and state its effect on the calculated enthalpy change of combustion. - Suggest one specific modification to the experimental procedure that would eliminate this error. [2 marks]

Solutions

22MasterySAQ-SExothermic vs endothermic reactions5 marksPaper 2~8 min

An energy profile for a reaction is shown. The reactants have an enthalpy of

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

, the products have an enthalpy of

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

, and the activated complex has an enthalpy of

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

.

(a)

(i) State whether the overall reaction is exothermic or endothermic. [1]
(ii) Calculate the enthalpy change,

\Delta H

, for the reaction. [1 mark]

(b)

(i) Define the term activation energy. [1]
(ii) Calculate the activation energy,

E_a

, for the forward reaction. [1 mark]

(c)

Suggest how the addition of a catalyst would change the energy profile . [1 mark]

Solutions

23ChallengeSAQ-LExothermic vs endothermic reactions8 marksPaper 2~12 min

Data

Quantity — Value Mass of water in calorimeter —

250.0\,\text{g}

Initial temperature of water —

21.5\,\text{°C}

Final temperature of water —

18.9\,\text{°C}

Specific heat capacity of water —

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

Specific heat capacity of steel container —

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

Mass of steel container —

150.0\,\text{g}

Molar mass of

\text{CaCO}_3

—

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

A student investigates the enthalpy change for the decomposition of calcium carbonate:

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

The student uses a calorimeter consisting of a steel container placed in a known mass of water. The student heats

5.00\,\text{g}

of

\text{CaCO}_3

strongly until it decomposes completely, and measures the temperature change of the surrounding water.

(a)

State whether the decomposition of

\text{CaCO}_3

is exothermic or endothermic, giving one piece of evidence from the data. [1 mark]

(b)

Calculate the enthalpy change,

\Delta H

, in

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

, for the decomposition of

1.00\,\text{mol}

of

\text{CaCO}_3

. Include the heat absorbed by both the water and the steel container. [3 marks]

(c)

Identify two distinct sources of systematic error in this experimental design that could account for the difference between the student's calculated

\Delta H

and the literature value of

+178\,\text{kJ}\,\text{mol}^{-1}

. For each error, state the direction of its effect on the calculated

\Delta H

. [2 marks]

(d)

Suggest and explain one modification to the experimental design that would improve the accuracy of the measured

\Delta H

for this reaction. [2 marks]

Solutions

24ChallengeSAQ-LExothermic vs endothermic reactions8 marksPaper 2~12 min

A student uses the following reaction to investigate enthalpy changes:

\text{NH}_4\text{NO}_3\text{(s)} + \text{H}_2\text{O(l)} \rightarrow \text{NH}_4^+\text{(aq)} + \text{NO}_3^-\text{(aq)}

In an experiment,

8.00\,\text{g}

of ammonium nitrate is dissolved in

100.0\,\text{g}

of water in a polystyrene cup calorimeter. The initial temperature of the water is

22.0\,°\text{C}

. After stirring until all solid dissolves, the final temperature is

17.2\,°\text{C}

. - Specific heat capacity of the solution:

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

- Molar mass of

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

:

M = 80.06\,\text{g}\,\text{mol}^{-1}

- Density of solution:

1.00\,\text{g}\,\text{mL}^{-1}

- The polystyrene cup absorbs negligible heat.

(a)

State whether the dissolution of

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

is exothermic or endothermic. [1 mark]

(b)

Calculate the molar enthalpy change,

\Delta H

, for the dissolution of

1.00\,\text{mol}

of

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

. [3 marks]

(c)

Explain, in terms of breaking and forming of intermolecular and ionic interactions, why the dissolution has the sign of

\Delta H

that you determined in (b). [2 marks]

(d)

Evaluate the use of a polystyrene cup calorimeter for this experiment, identifying one limitation and stating how it affects the calculated value of

\Delta H

. [2 marks]

Solutions

25ChallengeSAQ-LExothermic vs endothermic reactions9 marksPaper 2~14 min

A student investigates the enthalpy change for the reaction between zinc powder and aqueous copper(II) sulfate solution.

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

The student adds

0.0500\,\text{mol}

of zinc powder to

25.0\,\text{cm}^3

of

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

copper(II) sulfate solution in a polystyrene cup. The initial temperature of both reactants is

21.5\,°\text{C}

. After stirring, the maximum temperature reached is

35.2\,°\text{C}

. The specific heat capacity of the solution is

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

and the density of the solution is

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

.

(a)

Calculate the heat energy released, in kJ, during the reaction. [1 mark]

(b)

Identify the limiting reagent and determine the enthalpy change,

\Delta H

, for the reaction in

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

. State whether the reaction is exothermic or endothermic, with justification. [3 marks]

(c)

The student repeats the experiment using

0.100\,\text{mol}

of zinc powder with the same volume and concentration of copper(II) sulfate solution. Explain why the temperature rise observed is not double the original value. [2 marks]

(d)

Evaluate the suitability of using a polystyrene cup as the calorimeter in this experiment, with reference to the accuracy of the calculated

\Delta H

value. [2] (e) Suggest one improvement to the apparatus and explain how it would reduce uncertainty in the calculated

\Delta H

value. [1 mark]

Solutions

26MasterySAQ-SThermochemistry and enthalpy diagrams5 marksPaper 2~8 min

A student investigates the exothermic reaction between zinc powder and aqueous copper(II) sulfate:

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

The student adds

0.0500\,\text{mol}

of zinc powder to

50.0\,\text{cm}^3

of

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

aqueous copper(II) sulfate in a polystyrene cup. The initial temperature of both reactants is

21.5\,°\text{C}

. After stirring, the maximum temperature reached is

38.2\,°\text{C}

. Assume the specific heat capacity of the solution is

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

and the density of the solution is

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

.

(a)

State which species — reactants or products — has the higher enthalpy in this reaction. [1 mark]

(b)

Calculate the enthalpy change,

\Delta H

, for this reaction in

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

of zinc reacted. Give your answer to three significant figures. [2 marks]

(c)

The accepted standard enthalpy change for this reaction is

-216\,\text{kJ}\,\text{mol}^{-1}

. Explain two reasons why the student's experimental value differs from this accepted value. [2 marks]

Solutions

27MasterySAQ-SThermochemistry and enthalpy diagrams5 marksPaper 2~8 min

The combustion of ethene,

\text{C}_2\text{H}_4\text{(g)}

, is exothermic. The standard enthalpy of combustion of ethene is

-1411\,\text{kJ mol}^{-1}

. Bond enthalpies: Bond — Enthalpy /

\text{kJ mol}^{-1}

C−H —

+413

C=C —

+612

O=O —

+498

C=O —

+805

O−H —

+464

(a)

State what the negative value of the standard enthalpy of combustion indicates about the relative stability of reactants and products. [1 mark]

(b)

Calculate the enthalpy change for the complete combustion of one mole of ethene using bond enthalpies. [2 marks]

(c)

Explain why the value calculated in (b) differs from the standard enthalpy of combustion of ethene. [2 marks]

Solutions

28ChallengeSAQ-LThermochemistry and enthalpy diagrams7 marksPaper 2~11 min

Consider the hydrogenation of ethene to form ethane:

\text{C}_2\text{H}_4\text{(g)} + \text{H}_2\text{(g)} \rightarrow \text{C}_2\text{H}_6\text{(g)}

Average bond enthalpies /

\text{kJ mol}^{-1}

: Bond — Enthalpy C−H —

413

C=C —

614

C−C —

348

C=O —

743

O−H —

463

H−H —

436

(a)

Calculate the enthalpy change,

\Delta H

, for the hydrogenation of ethene using the bond enthalpy data provided. [3 marks]

(b)

Explain why the value calculated in (a) differs from the experimentally determined value of

-136\,\text{kJ mol}^{-1}

, and state which value is more reliable for this specific reaction under standard conditions. [2 marks]

(c)

A student proposes that the hydrogenation of methanal (HCHO) to methanol (

\text{CH}_3\text{OH}

) is also exothermic:

\text{HCHO(g)} + \text{H}_2\text{(g)} \rightarrow \text{CH}_3\text{OH(g)}

Evaluate this prediction by calculating

\Delta H

for this reaction and comparing the pattern of bond breaking and bond forming with the hydrogenation of ethene. [2 marks]

Solutions

29MasterySAQ-SThermochemistry and enthalpy diagrams6 marksPaper 2~9 min

The dissolution of ammonium nitrate in water is endothermic. A student dissolves

8.00\,\text{g}

of

\text{NH}_4\text{NO}_3\text{(s)}

in

100.0\,\text{g}

of water in a calorimeter. The initial temperature of the water is

22.0\,°\text{C}

. After complete dissolution, the temperature falls to

18.4\,°\text{C}

. Assume the specific heat capacity of the solution is

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

. The molar mass of

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

is

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

.

(a)

State the relative enthalpy levels of reactants and products for this dissolution an enthalpy . [1 mark]

(b)

Calculate the enthalpy change,

\Delta H

, for the dissolution of ammonium nitrate, in

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

. [3 marks]

(c)

The experimental value of

\Delta H

differs from the accepted literature value of

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

. Deduce one assumption made in this experiment and explain how it leads to an underestimate of the magnitude of

\Delta H

. [2 marks]

Solutions

30MasterySAQ-SThermochemistry and enthalpy diagrams5 marksPaper 2~8 min

The Haber process synthesises ammonia:

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

(a)

On the axes below, sketch a labelled enthalpy profile for this reaction, showing the effect of a catalyst. Your must include: reactants, products, activation energy with and without catalyst, and

\Delta H

. [1 mark]

(b)

Using the bond enthalpy data below, calculate the bond enthalpy of the

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

bond. | Bond enthalpy /

\text{kJ mol}^{-1}

|

\text{H}{-}\text{H}

—

+436

\text{N}{-}\text{H}

—

+391

[2 marks]

(c)

Explain why the value calculated in (b) may differ from the experimentally determined bond dissociation energy of

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

in

\text{N}_2(g)

. [2 marks]

Solutions

31MasterySAQ-SCombustion reactions and energy release7 marksPaper 2~11 min

Methanol (

\text{CH}_3\text{OH}

) is being considered as a cleaner-burning alternative fuel for cars. In a laboratory test,

0.640\,\text{g}

of methanol is burned completely in a bomb calorimeter containing

850\,\text{g}

of water. The combustion reaction is:

2\text{CH}_3\text{OH}(l) + 3\text{O}_2(g) \rightarrow 2\text{CO}_2(g) + 4\text{H}_2\text{O}(l)

The temperature of the water rises from

22.5\,°\text{C}

to

36.8\,°\text{C}

. The specific heat capacity of water is

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

. The molar mass of methanol is

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

. The standard enthalpy of combustion of methanol is

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

.

(a)

State two conditions required for complete combustion of methanol. [2 marks]

(b)

Calculate the heat energy, in kJ, released by the combustion of the

0.640\,\text{g}

sample. [2 marks]

(c)

Using your answer to (b), determine the experimental molar enthalpy of combustion of methanol in

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

. Hence evaluate the reliability of this experiment by comparing your value with the standard value. [3 marks]

Solutions

32MasterySAQ-SCombustion reactions and energy release5 marksPaper 2~8 min

Methane (

\text{CH}_4

) undergoes incomplete combustion when the oxygen supply is limited, producing carbon monoxide (CO) and water vapour instead of carbon dioxide (

\text{CO}_2

). The equations for the two reactions are:

\text{Complete: } \text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) \quad \Delta H = -890\,\text{kJ mol}^{-1}

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

(a)

State which combustion reaction releases more energy per mole of methane. [1 mark]

(b)

Explain, using bond breaking and bond making, why complete combustion of methane releases more energy than incomplete combustion. [2 marks]

(c)

In a gas heater operating under limited oxygen, incomplete combustion of methane releases

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

. Calculate the percentage of energy lost compared to complete combustion. Give your answer to three significant figures. [2 marks]

Solutions

33ChallengeSAQ-LCombustion reactions and energy release7 marksPaper 2~11 min

A student investigates the energy released when different alcohols are used as fuels in a spirit burner. The student burns

1.00\,\text{g}

of each alcohol and measures the temperature increase of

200.0\,\text{g}

of water in a copper calorimeter. The specific heat capacity of water is

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

. The experimental results are shown: - Methanol (

\text{CH}_3\text{OH}

):

\Delta T = 17.3\,\text{K}

- Ethanol (

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

):

\Delta T = 13.0\,\text{K}

- Propan-1-ol (

\text{C}_3\text{H}_7\text{OH}

):

\Delta T = 10.4\,\text{K}

(a)

Calculate the experimental enthalpy change of combustion, in

\text{kJ mol}^{-1}

, for ethanol. [3 marks]

(b)

Explain why the experimental value is significantly less exothermic than the value calculated from bond enthalpies. [2 marks]

(c)

State whether the assumption that all heat released is absorbed by the water is valid, giving one reason. [1 mark]

(d)

Suggest one specific improvement to the experimental design that would increase the accuracy of the results. [1 mark]

Solutions

34MasterySAQ-SCombustion reactions and energy release6 marksPaper 2~9 min

Octane (

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

) is a major component of petrol. The balanced equation for complete combustion is:

2\text{C}_8\text{H}_{18}(l) + 25\text{O}_2(g) \rightarrow 16\text{CO}_2(g) + 18\text{H}_2\text{O}(g)

The standard enthalpy of combustion of octane is

-5470\,\text{kJ mol}^{-1}

. A driver uses

569\,\text{g}

of octane during a journey.

(a)

A poorly tuned engine produces black soot on the spark plugs. State the type of combustion occurring and identify the chemical substance responsible for the soot. [2 marks]

(b)

Calculate the energy released when

569\,\text{g}

of octane undergoes complete combustion. [2 marks]

(c)

The same mass of octane burned in the poorly tuned engine releases significantly less energy than calculated in (b). Using your knowledge of enthalpy and oxidation products, explain why the energy released is lower. (Do not simply restate that combustion is incomplete.) [2 marks]

Solutions

35MasterySAQ-SFuels and their efficiency10 marksPaper 2~15 min

Data

- Fuel A: ethanol,

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

, mass burned

= 0.460\,\text{g}

, temperature rise

= 15.0\,\text{°C}

- Fuel B: octane,

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

, mass burned

= 0.570\,\text{g}

, temperature rise

= 20.5\,\text{°C}

- Specific heat capacity of water

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

- Standard enthalpy of combustion of ethanol

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

- Standard enthalpy of combustion of octane

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

- Molar masses: ethanol

= 46.08\,\text{g}\,\text{mol}^{-1}

; octane

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

A student investigates the efficiency of different fuels by burning a small sample of each fuel in a spirit burner to heat

200.0\,\text{g}

of water. The temperature change of the water is measured for each fuel.

(a)

State what is meant by the standard enthalpy of combustion of a fuel. [1 mark]

(b)

Calculate the experimental energy transferred to the water when

0.460\,\text{g}

of ethanol was burned. [2 marks]

(c)

Calculate theoretical energy released by the

0.460\,\text{g}

sample of ethanol, using the standard enthalpy of combustion. [2 marks]

(d)

The student claims that octane is a more efficient fuel than ethanol because it produced a larger temperature rise. Evaluate this claim by calculating the percentage efficiency of energy transfer for both fuels and comparing the results. [5 marks]

Solutions

36ChallengeSAQ-LEntropy and the second law of thermodynamics7 marksPaper 2~11 min

Data

-

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

-

\Delta S^{\ominus}_{system} = +161\,\text{J K}^{-1}\text{mol}^{-1}

Equations:

\Delta G^{\ominus} = \Delta H^{\ominus} - T\Delta S^{\ominus}

;

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

;

\quad \Delta S_{surroundings} = -\dfrac{\Delta H^{\ominus}}{T}

This question is about the spontaneity of thermal decomposition of calcium carbonate, a key reaction in the cement industry.

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

(a)

Calculate the temperature, in K, above which the decomposition of calcium carbonate becomes spontaneous. [3 marks]

(b)

Explain why

\Delta S^{\ominus}_{system}

is positive for this reaction. [2 marks]

(c)

Calculate the total entropy change of the universe,

\Delta S_{total}

, at

298\,\text{K}

, and hence evaluate whether this consistent with the reaction being non-spontaneous at this temperature. [2 marks]

Solutions

37ChallengeSAQ-LEntropy and the second law of thermodynamics9 marksPaper 2~14 min

Data

Species —

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

NH_4NO_3(s)

—

151.1

NH_4^+(aq)

—

113.4

NO_3^-(aq)

—

146.4

-

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

Equations:

\Delta G^\ominus = \Delta H^\ominus - T\Delta S^\ominus_{system}

;

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

;

\Delta S_\text{surroundings} = -\dfrac{\Delta H^\ominus}{T}

This question explores the spontaneity of dissolving ammonium nitrate in water, a process used instant cold packs.

NH_4NO_3(s) \rightarrow NH_4^+(aq) + NO_3^-(aq)

(a)

Calculate

\Delta S^\ominus_\text{system}

for the dissolution of ammonium nitrate. [2 marks]

(b)

Determine the minimum temperature, in K, above which the dissolution of ammonium nitrate is spontaneous. [2 marks]

(c)

Explain why the entropy change of the system is positive for this dissolution. [2 marks]

(d)

Evaluate the statement: *"Since the dissolution of ammonium nitrate is spontaneous at

298\,\text{K}

, the entropy change of the universe must be positive."* Support your answer with a calculation. [3 marks]

Solutions

38MasterySAQ-SEntropy and the second law of thermodynamics5 marksPaper 2~8 min

Data

T = 298\,\text{K}

;

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

Ammonium nitrate (

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

) dissolves spontaneously in water at room temperature, even though the process is endothermic. The dissolution reaction is:

\text{NH}_4\text{NO}_3(s) \rightarrow \text{NH}_4^+(aq) + \text{NO}_3^-(aq) \quad \Delta H^\circ = +26\,\text{kJ mol}^{-1}

(a)

State the second law of thermodynamics. [1 mark]

(b)

Calculate the standard entropy change of the surroundings,

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

, at

298\,\text{K}

.

\Delta S^\circ_{\text{surr}} = -\frac{\Delta H^\circ_{\text{system}}}{T}

[2 marks]

(c)

Using your answer to (b) and the value of

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

given above, calculate

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

and deduce whether the dissolution is spontaneous. [2 marks]

Solutions

39MasterySAQ-SEntropy and the second law of thermodynamics5 marksPaper 2~8 min

Consider thermal decomposition of calcium carbonate:

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

\Delta H^\circ = +178\,\text{kJ mol}^{-1}, \quad \Delta S^\circ_{\text{sys}} = +160\,\text{J K}^{-1}\text{mol}^{-1}

(a)

State why

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

is positive for this reaction. [1 mark]

(b)

Calculate the minimum temperature, in

^\circ\text{C}

, at which this reaction becomes spontaneous. [2 marks]

(c)

Using the second law of thermodynamics and the expression

\Delta S_{\text{surr}} = -\dfrac{\Delta H}{T}

, explain why spontaneity is only achieved above the temperature calculated in (b). [2 marks]

Solutions

40MasterySAQ-SEntropy and the second law of thermodynamics6 marksPaper 2~9 min

Hydrogen gas and iodine gas react to form hydrogen iodide gas:

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

\Delta H^\circ = +9.4\,\text{kJ mol}^{-1}, \quad \Delta S^\circ_{\text{sys}} = +21.8\,\text{J K}^{-1}\text{mol}^{-1}

A student claims that because the reaction is endothermic, it cannot be spontaneous at any temperature.

(a)

State the second law of thermodynamics. [1 mark]

(b)

Calculate the minimum temperature, in °C, above which this reaction becomes spontaneous. [3 marks]

(c)

Evaluate the student's claim by considering the signs of

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

,

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

, and

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

at temperatures above and below the value calculated in (b). [2 marks]

Solutions

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