Thermodynamics (AQA A Level Chemistry)

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 from Table 1 below to calculate a value for the electron affinity of fluorine.

Table 1

Table 2 below gives some values of standard enthalpy changes. Use these values to answer the questions.

Table 2

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.

Figure 1

Use the data in Table 3 and your completed Born–Haber cycle from part (c) to answer the questions below.

Table 3

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.

Figure 1

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.

Figure 2

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 Table 1 below to calculate the standard enthalpy of solution of silver fluoride.

Table 1

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.

This question is about the enthalpy of hydration of halide ions.

Table 1

Using the data from Table 1 above, 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 in Table 2 to suggest why there is a difference between the hydration values of calcium ions and sodium ions.

Table 2

i) Define the term enthalpy of lattice formation.

ii) Using the data from Table 3 below, calculate the enthalpy of hydration of Ca²⁺ ions.

Table 3

i) Table 1 below shows some solution and hydration enthalpies. Using the data from Table 1, calculate the enthalpy of hydration of calcium ions.

Table 1

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.

Figure 1

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 Figure 1, 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)

Table 2

i) Use the data from Table 2 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)

Table 3

i) Calculate the standard enthalpy change of the contact process reaction, using the data provided in Table 3. 

ii) The standard entropy change of this reaction is –189 JK^-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. If you did not calculate an answer for part (i), use 70 C as your value. Note that this is not the correct answer to part (i). 

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      

Table 1

i) Using the data from Table 1, show that this process is feasible at 25 C.

ii) Use your knowledge of structure and bonding to explain why ∆H is positive for this process.

Diamond and graphite are both allotropes of carbon.

Table 2 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)

Table 2

Use the following data in Table 1 to calculate the enthalpy of lattice formation 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 Table 2, identify the group one ion.

Table 2

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 Table 2 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. 

Draw a fully labelled Born-Haber cycle for the formation of sodium selenide, Na₂Se (s), from its elements. Include state symbols in all equations.

If sulfur is used as opposed to selenium in the lattice, what would you expect to happen to the value of the enthalpy of lattice dissociation. Explain your answer.

Use the data in Table 1 to calculate the value for the enthalpy of lattice formation of aluminium oxide.

Table 1

Aluminium oxide is insoluble in water, but sodium oxide is soluble. Explain why there is no enthalpy of solution data for sodium oxide.

The equations for two separate reversible reactions are as follows:

Reaction A      2SO₂ (g) + O₂ (g) ⇌ 2SO₃ (g)

Reaction B      CO (g) + H₂O (g) ⇌ CO₂ (g) + H₂ (g)

Use the information in Table 1 to calculate the temperature at which the free energy change for reaction A becomes feasible

Table 1

By using the data in Table 1, deduce if reaction B is feasible at a low temperature.

Magnesium carbonate decomposes at a relatively high temperature. Using the data in Table 2, determine if this reaction is feasible at 280 C.

Table 2

Using your answer to part (c), determine the temperature in C at which the decomposition of magnesium carbonate becomes feasible.

This question is about enthalpy changes. 

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.

A different Group 1 metal, lithium, can be used in rechargeable batteries for mobile phones, laptops, digital cameras and electric vehicles. Lithium is also used in some non-rechargeable batteries for things like heart pacemakers, toys and clocks.     

Lithium carbonate will decompose on heating.

Figure 1 shows the graph of ΔG plotted against temperature, K for the decomposition of lithium carbonate.

i) Use Figure 1 to calculate the entropy change, ΔS, for this reaction in J K^-1 mol^-1. 

ii) Determine the temperature at which this reaction becomes feasible in K. 

Figure 1

Which of the following compounds would have the greatest difference between the theoretical and experimental value for lattice enthalpy. Explain your answer.

Silver fluoride

Silver iodide

Sodium fluoride

Butane can be converted into ethene and hydrogen which can be used in hydrogenation of vegetable oil in the presence of a nickel catalyst.

i) Write a fully balanced chemical equation including state symbols for cracking of butane.

ii) Using the data in Table 1, calculate the free energy change for this reaction. 

iii) Using your answer to part (ii) explain under what conditions this reaction would be feasible.

Table 1

The free energy change for this reaction varies with temperature.

i) Use the data in Table 2 to plot a graph of free energy against temperature on the grid below.

ii) Calculate the gradient of the line on your graph and hence calculate the entropy change, ΔS in J K mol^-1 for the formation of the cracking of butane.

Table 2

Figure 1

Hydrazoic acid is a colorless, volatile explosive liquid under standard conditions. The decomposition of the acid in a gaseous state is shown in the following equation.

HN₃ (g) →  H₂ (g) + N₂ (g)

Figure 1 shows a graph of graph of free energy against temperature on the grid below.

Figure 1

i) Use Figure 1 to calculate the value for ΔS for this reaction.

ii) Determine a value for ΔH for this reaction.

Use the data give in Table 2 to explain why the reaction for the formation of ethene from its elements is never feasible at any temperature. You must show your working.

Table 2