Thermodynamics (OCR A Level Chemistry A): Exam Questions

This question is about fluorine and the associated energy changes when it is reacted with magnesium to form magnesium fluoride.

i) Define the term electron affinity for fluorine.

ii) Use the data in the table below to calculate a value for the electron affinity of fluorine.

The table below gives some values of standard enthalpy changes. Use these values to answer the questions.

i) Suggest why the electron affinity of bromine is an exothermic change.

ii) Explain why the bond enthalpy of a Br–Br bond is greater than that of a I–I bond.

This question is about lithium fluoride.

Complete the Born–Haber cycle for lithium fluoride by adding the missing species on the lines.

Use the data in the table below and your completed Born–Haber cycle from part (c) to answer the questions below.

i) Calculate the enthalpy of lattice formation of lithium fluoride.

ii) Explain and justify how the enthalpy of lattice formation of NaBr compares with that of NaF. You must refer to the size of the ions in your answer.

This question is about the energy changes represented in a Born-Haber cycle when forming silver chloride.

i) Complete the Born–Haber cycle for silver chloride above.

ii) Suggest why the electron affinity of chlorine has a negative value. You should refer to electrostatic forces in your answer.

This question focuses on the difference between the theoretical enthalpy values and experimental values of ionic substances. A perfect ionic model can be used to calculate a theoretical value for the enthalpy of lattice dissociation.

i) The theoretical enthalpy of lattice dissociation for sodium fluoride is +902 kJ mol^–1. This is very different to the experimental value that can be calculated from a Born Haber cycle. Explain this difference.

ii) The theoretical enthalpy of lattice dissociation value for sodium chloride is less than the theoretical enthalpy lattice dissociation of sodium fluoride. Explain why this difference exists.

This question looks at the enthalpy steps of calcium sulfide and its Born-Haber cycle.

i) Name the enthalpy changes for Steps 1, 5 and 7 in this Born-Haber cycle for calcium sulfide.

ii) Step 6 is an endothermic process. Explain why and identify Z. 

iii) Explain why the value for Step 3 is smaller than Step 4.

This question is about enthalpy changes in solution.

i) Write the equation for the process showing the enthalpy of solution of silver fluoride. Include state symbols in your answer. 

ii) Use the data in the table below to calculate the standard enthalpy of solution of silver fluoride.

This question focuses on the enthalpy of hydration of fluoride compounds.

i) Define the term enthalpy of hydration in relation to a fluoride ion.

ii) State whether the hydration of a fluoride ion is an exothermic or endothermic process. Explain your answer.

The enthalpy of hydration of the fluoride and chloride ions are -524 kJ mol^-1 and -364 kJ mol^-1. Explain why the value for the enthalpy of hydration for the fluoride ion is more negative than that for the chloride ion. You must refer to the attractive forces involved in your answer.

Use the enthalpy values below to suggest why there is a difference between the hydration values of calcium ions and sodium ions.

The table below gives data for the ionic compound calcium bromide, CaBr₂. 

i) Define the term enthalpy of lattice formation.

ii) Using the data from the table below, calculate the enthalpy of hydration of Ca²⁺.

The table below shows some solution and hydration enthalpies.

i) Using the data from the table, calculate the enthalpy of hydration of calcium ions.

ii) When a calcium ion is hydrated, calcium ions attract water molecules and energy is released. Explain why water molecules are attracted to calcium ions. You may use a labelled diagram to illustrate your answer.

This question looks at how the entropy change of water varies with temperature.

i) The entropy of water is zero when the temperature is zero Kelvin. Explain why, with reference to the water molecules in your answer.

ii) Explain why the entropy change, ΔS, is larger at temperature T₂ than at temperature T₁. 

iii) On the graph draw the boiling point (T_b) of water on the appropriate axis.

Standard entropies can be used to calculate the entropy change of a reaction, ΔS. For example, for the reaction between nitrogen monoxide and oxygen which is shown below. 

2NO (g) + O₂ (g) → 2NO₂ (g)

i) Use the data from the table to calculate the entropy change of the reaction between nitrogen monoxide and oxygen. 

ii) Using your answer from part (i), explain what the sign of the entropy change indicates about the products (NO₂) compared to the reactants (NO and O₂) in this reaction.

The contact process is a method used by industries to form sulfur trioxide, by reacting sulfur dioxide and oxygen together over a vanadium (V) oxide catalyst.

The equation for this reaction is shown below: 

2SO₂ (g) + O₂ (g) ⇌ 2SO₃ (g)

i) Calculate the enthalpy change of the contact process reaction, using the data provided in the table. 

ii) The standard entropy change of this reaction is –189 J K^-1 mol^-1. Use this value and your enthalpy value from part (i) calculate a value for the free energy change for this reaction at 45°C.

iii) Use your answer to part (ii) to explain whether the reaction is feasible at 45°C.

This question is about the enthalpy of solution of sodium chloride.

If the value of enthalpy of solution of sodium chloride is +4 kJ mol^-1, explain why the free energy change for dissolving sodium chloride in water is negative, despite the enthalpy change being a positive value.

Calcium carbonate thermally decomposes to form calcium oxide and carbon dioxide, as shown below:

CaCO₃ (s) → CaO (s) + CO₂ (g)

The enthalpy change of the above reaction is ΔH = +178 kJ mol^-1 and the entropy change is ΔS = +161 J K^-1 mol^-1. 

Calculate the temperature at which the free-energy change, ΔG, for this process is zero.

Some ionic compounds such as potassium chloride, KCl, will dissolve in water at room temperature in an endothermic process. 

KCl (s) → K⁺ (aq) + Cl^- (aq) ΔH = +16 kJ mol^-1 

i) Using the data from the table to prove that this process is feasible at 25 ℃.

ii) Use your knowledge of structure and bonding to explain why ΔH is positive for this

Diamond and graphite are both allotropes of carbon. The table below shows the data for the conversion of graphite into diamond. Use this data to calculate values for ΔH and ΔS for the reaction. Use these values to explain why this reaction is not feasible under standard pressure at any temperature.

Carbon (graphite) → Carbon (diamond)

Magnesium hydroxide, Mg(OH)₂​, and barium hydroxide, Ba(OH)₂​, are Group 2 compounds with different solubilities in water.

i) State the trend in the solubility of the Group 2 hydroxides down the group from Mg(OH)₂​ to Ba(OH)₂​.

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ii) Explain this trend in solubility.

In your answer, refer to lattice enthalpy and hydration enthalpy.

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A student adds a sample of solid barium oxide, BaO, to a beaker containing distilled water.

The resulting solution of barium hydroxide, Ba(OH)₂​, is transferred to a volumetric flask and made up to 250.0 cm³ with distilled water.

The pH of this solution is measured as 13.12 at 298 K.

Assume the ionic product of water, K_w​=1.00 x 10⁻¹⁴ mol² dm⁻⁶ at 298 K.

i) Write the equation for the reaction of barium oxide with water.

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ii) Calculate the concentration, in mol dm⁻³, of hydroxide ions, OH⁻, in the solution.

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iii) Calculate the mass of barium oxide, BaO, that the student used to prepare this solution.

Give your answer to 3 significant figures.

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iv) Explain why using a pH meter to measure the pH in this experiment is more reliable than using universal indicator paper.

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Carbon dioxide and hydrogen can react to form methanol. 

CO₂ (g) + 3H₂ (g) ⇌ CH₃OH (g) + H₂O (g)

Data for enthalpy of formation and entropy is shown in the table.

Calculate the Gibbs free-energy change (ΔG), in kJ mol^–1, for this reaction at 790 K.

Explain whether this reaction is feasible at 790 K.

Use your values of ΔH and ΔS from part (a) to calculate the temperature below which this reaction is feasible.

Explain why the reaction in part a) has a negative entropy change. 

This question is about Born-Haber cycles. 

Draw a fully labelled Born-Haber cycle for the formation of solid potassium sulfide, K₂S, from its elements. 

Include state symbols for all species involved. 

Use the data to calculate the lattice enthalpy of K₂S.

If lithium replaced potassium in the lattice, what would you expect to happen to the value of the lattice enthalpy?

Explain your answer.   

Explain the difference in values between the first and second electron affinities of sulfur. 

This question is about the feasibility of a reaction. 

Calcium carbonate can be thermally decomposed to make calcium oxide. 

CaCO₃ (s)  →   CaO (s)   +   CO₂ (g)        ΔH= + 178 kJ mol^-1

Table 1 below shows the standard entropies of each substance. 

Table 1 

i) Use the information in the table to show that calcium carbonate is stable at room temperature (25 ^oC).

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ii) Calculate the minimum temperature needed to decompose calcium carbonate.

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Explain how the conditions can be changed for the reaction in a) to make the reaction feasible.

The evaporation of water is an endothermic reaction but also a spontaneous reaction.

Explain why.

Refer to the entropy change that occurs and the change in the arrangement of water molecules. 

Use the information in the table to calculate the temperature above which the reaction between hydrogen and oxygen to form gaseous water is not feasible.

Explain what would happen to the sample of gaseous water if it was heated to a higher temperature than that calculated in part (d). 

This question is about lattice enthalpy. 

Use the following data in Table 1 to calculate the lattice enthalpy of caesium oxide. 

Table 1 

A different Group 1 metal forms an ionic compound with chlorine.

The enthalpy of lattice dissociation for this compound is +773 kJ mol^-1 and the hydration enthalpy of a chloride ion is -363 kJ mol^-1. 

The enthalpy of solution of the Group 1 chloride is +4 kJ mol^-1. 

Using the information in the table, identify the Group 1 ion.

The same Group 1 metal from part (b) forms an ionic lattice with another halide ion. This new ionic compound has an enthalpy of lattice formation of -705 kJ mol^-1. 

Using M to represent the Group 1 metal, suggest a formula for the new ionic lattice and explain your answer.

The enthalpy of hydration becomes less exothermic as you go down Group 1. Using the values in part (b):

i) Explain why the enthalpy of hydration of Group 1 ions is negative.

ii) Explain why the enthalpies of hydration become less negative as you go down the group.

Strontium is used as a red colouring agent in fireworks as it provides a very intense red colour.

Use the data in Table 1 to calculate the atomisation energy for chlorine in strontium chloride. 

Table 1

Using information in part (a) suggest a value for the lattice enthalpy of formation of rubidium chloride.

Explain your choice of value.

Enthalpy changes can be determined using Born-Haber cycles.

i) Define the term enthalpy change of atomisation.

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ii) Define the term first electron affinity.

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The diagram below shows an incomplete Born-Haber cycle for barium iodide, BaI₂.

i) On the dotted lines, add the species and state symbols for the steps occurring.

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ii) The table below shows standard enthalpy changes.

Calculate the lattice enthalpy of barium iodide.

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Explain why a purely ionic model becomes less valid for magnesium iodide, MgI₂, than for barium iodide, BaI₂, leading to a greater difference between the theoretical and experimental lattice enthalpy values.

In your answer, you should refer to the size and charge density of the ions involved.

A student determines the enthalpy change of solution of anhydrous calcium chloride, CaCl₂, by calorimetry.

The student adds 3.00 g of anhydrous CaCl₂ to 100.0 g of water in a polystyrene cup.

The temperature increases from 20.5 ^oC to 34.0 ^oC.

The specific heat capacity of the solution is 4.18 J g^-1 K^-1.

i) Calculate the energy change, q, in Joules, for this reaction.

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ii) Calculate the enthalpy change of solution, Δ_solH, for calcium chloride in kJ mol^-1.

M_ᵣ(CaCl₂) = 111.1 g mol^-1.

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iii) The student repeats the experiment using hydrated calcium chloride, CaCl₂·6H₂O.

The enthalpy change of solution for the hydrated salt is determined to be +19.1 kJ mol^-1.

Using your answer to (a)(ii) and the value above, construct an enthalpy cycle to calculate the enthalpy change for the reaction:

CaCl₂ (s) + 6H₂O (l) → CaCl₂·6H₂O (s)

[3]

Another student investigates the reaction between magnesium and dilute hydrochloric acid.

Mg (s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g)

The student adds a sample of magnesium ribbon to excess hydrochloric acid.

The hydrogen gas produced is collected in a gas syringe at 25.0 ^oC and 105 kPa.

The volume of gas collected is 52.0 cm³.

i) Calculate the amount, in moles, of hydrogen gas produced.

The gas constant R = 8.314 J mol^-1 K^-1.

[3]

ii) The student calculates the relative atomic mass, Aᵣ, of the magnesium sample using the mass of the ribbon and the moles of gas produced.

The calculated Aᵣ is higher than the periodic table value of 24.3.

The student realises the magnesium ribbon had a layer of magnesium oxide on the surface.

Explain, using the stoichiometry of the reaction, why this contamination results in a calculated Aᵣ that is higher than expected.

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