Define the term enthalpy of lattice formation.

The lattice enthalpies of potassium fluoride, potassium chloride and calcium fluoride are shown in the table below.

Explain the differences between these lattice enthalpies.

Write an equation to represent the enthalpy change of atomisation, ΔH_at, of chlorine.

Explain why ΔH_at for chlorine is positive (endothermic).

Lattice enthalpies can be determined indirectly using Born-Haber cycles.

The table below shows the enthalpy changes that are needed to determine the lattice enthalpy of sodium chloride, NaCl.

The diagram below shows an incomplete Born-Haber cycle to determine the lattice enthalpy of sodium chloride.

The arrows labelled 1-5 correspond to the five rows of the table above.

On the five dotted lines, add the species present, including state symbols.

Using the model answer from part a), calculate the enthalpy of lattice formation, ΔH_latt, of sodium chloride.

The change that produces lattice enthalpy is spontaneous but has a negative entropy change.

Why is this change able to take place spontaneously?

Define the term first electron affinity.

Write an equation to represent the first electron affinity, ΔH_ea1, of fluorine.

The diagram below shows an incomplete Born-Haber cycle to determine the enthalpy of lattice formation of magnesium fluoride.

Complete the Born-Haber cycle diagram for magnesium fluoride.

Include the formulae and state symbols of appropriate species.

The table below contains some enthalpy data.

Use the Born-Haber cycle from part a) and data from the table above to calculate a value for the first electron affinity, ΔH_ea1, of fluorine.

Define the term enthalpy of hydration.

The table below shows enthalpy of hydration values for some group 2 ions.

Explain why the enthalpy of hydration becomes less exothermic from Mg²⁺ to Sr²⁺.

Draw a line from each ion to the enthalpy of hydration of that ion.

The diagram below is an incomplete Born-Haber cycle showing the enthalpy changes involved when lithium chloride, LiCl, dissolves in water.

On the two dotted lines, add the species present, including state symbols.

State the enthalpy changes represented by the four arrows labelled 1-4.

The table below shows some enthalpy data.

Use the tabulated data and the Born-Haber cycle from part b) to calculate the standard enthalpy of solution, ΔH_sol, of lithium chloride.

Use your answer from part c) to predict whether lithium chloride is soluble in water.

Potassium forms an ionic compound with a group 7 halide, X^-.

The enthalpy of lattice formation for this compound is -674 kJ mol^-1 and its enthalpy of solution is +6 kJ mol^-1.

The hydration enthalpies for potassium and other halide ions are shown in the table below.

Identify the halide ion, X^-. 

Bromine forms an ionic lattice with another group 1 ion. The new ionic compound has an enthalpy of lattice formation of +733 kJ mol^-1. 

Suggest a formula for the new ionic lattice. Explain your answer.

At 298 K, carbon dioxide and hydrogen gases react to produce two liquids, as shown by the equation:

CO_2(g) + 3H_2(g) ➔ CH₃OH_(l) + H₂O_(l)       ΔH^ø = -131 kJ mol^-1

Standard entropies are shown in the tabe below.

Calculate the standard entropy change, S^ø, for this reaction.

Calculate the standard Gibbs free energy change, ΔG^ø, for this reaction at 298 K.

Give your answer in units of kJ mol^-1 to 3 significant figures.

Predict the effect of decreasing the temperature on the feasability of this reaction.

The feasibility of a chemical reaction depends on the standard Gibbs free energy change, ΔG^ø. This is related to the standard enthalpy and entropy changes by the equation:

ΔG^ø = ΔH^ø - TΔS^ø

Predict, with reasoning, whether the following processes will lead to an increase or decrease in entropy:

1. solid potassium sulfate dissolving in water
2. steam condensing to water
3. the reaction of calcium carbonate with hydrochloric acid

Energy changes for the melting and boiling of H₂O are shown below.

H₂O_(s) ➔ H₂O_(l)      ΔH^ø = +6.01 kJ mol^-1

H₂O_(l) ➔ H₂O_(g)     ΔH^ø = +40.7 kJ mol^-1

Explain why ΔH^ø is positive when water melts or boils.

Standard entropies of H₂O in its three physical states are shown in the table below.

Explain why the increase in S^ø is much greater when water boils than when water melts.

Use data in the table above to calculate the change in entropy, in kJ K^-1 mol^-1, when ice melts.

Use your answer to part d) to show that ice melts at 0°C (at standard pressure).

Hydrogen can be manufactured from the reaction of steam with methane.

CH_4(g) + H₂O_(g) ⇌ 3H_2(g) + CO_(g)     ΔH^ø = +206 kJ mol^-1

The entropy change for the forward reaction, ΔS^ø, is +214.5 J K^-1 mol^-1.                          

Explain why the forward reaction has a positive ΔS^ø value.

The entropy values of some gases are shown in the table below.

Using the value of ΔS^ø and the entropy values in the table, calculate the entropy of CO_(g).

Give your answer to 1 decimal place.

Calculate whether the forward reaction is feasible at 700°C.

The temperature used for this manufacture of hydrogen is typically about 1,000°C.

Suggest one reason why a temperature of 1,000°C is used rather than 700°C.

The equation for the reaction between ammonia and oxygen is shown.

4NH_3(g) + 5O_2(g) ⇌ 4NO_(g) + 6H₂O_(g)

The table below contains some enthalpy of formation and entropy data.

Use data in the table above to calculate the enthalpy change, ΔH_r, for the reaction of ammonia with oxygen.

Use data in the table above to calculate the entropy change, ΔS, for the reaction of ammonia with oxygen.

Using your answers to parts a) and b), calculate a value for the Gibbs free energy change, ΔG, in kJ mol^-1, for the reaction between ammonia and oxygen at 800 K.

Give your answer to the nearest integer.

Explain how the feasibility of the reaction between ammonia and oxygen changes as the temperature increases.