AQA A-Level Chemistry: Energetics & Thermodynamics

Lattice enthalpies are a measure of the strength of the ionic bonding in a crystalline solid, and they vary in a predictable way from compound to compound.

Describe and explain how the lattice enthalpy of formation of an ionic compound varies with the charge and the radius of its ions. In your answer you should:

• give the definition of lattice enthalpy of formation (including the relevant equation with state symbols, using a named compound as an example);
• explain, in terms of the electrostatic forces between ions, why lattice enthalpy depends on ionic charge and on ionic radius;
• illustrate your reasoning by comparing NaCl with MgO, and NaF with NaI.

(6 marks)

A student used a polystyrene cup to measure the enthalpy change of the neutralisation reaction between hydrochloric acid and sodium hydroxide:

$\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$HCl(aq)+NaOH(aq)→NaCl(aq)+H2​O(l)

The student mixed 25.0 cm³ of 1.00 mol dm⁻³ HCl with 25.0 cm³ of 1.00 mol dm⁻³ NaOH. The temperature of the mixture rose by 7.0 °C.

Use the data below to calculate the enthalpy change of neutralisation, $\Delta_{\text{neut}}H$Δneut​H, in kJ mol⁻¹.

Quantity	Value
Total mass of mixture	50.0 g
Specific heat capacity, $c$c	4.18 J g⁻¹ K⁻¹
Temperature change, $\Delta T$ΔT	+7.0 K

Assume that all the heat released is transferred to the solution, and that the solution has the same specific heat capacity as water. Show your working, including the correct sign and units.

(6 marks)

When potassium bromide dissolves in water, the ionic lattice breaks apart and the gaseous ions become hydrated:

$\text{KBr}(s) \rightarrow \text{K}^{+}(aq) + \text{Br}^{-}(aq)$KBr(s)→K+(aq)+Br−(aq)

The enthalpy of solution can be found by constructing an enthalpy cycle that links the solid, the gaseous ions and the aqueous ions. The relevant data are given below.

Enthalpy change	Value / kJ mol⁻¹
Lattice enthalpy of formation of KBr, $\Delta_{\text{latt}}H$Δlatt​H	−670
Enthalpy of hydration of $\text{K}^{+}(g)$K+(g), $\Delta_{\text{hyd}}H$Δhyd​H	−322
Enthalpy of hydration of $\text{Br}^{-}(g)$Br−(g), $\Delta_{\text{hyd}}H$Δhyd​H	−335

Use these data to calculate the standard enthalpy of solution of potassium bromide, $\Delta_{\text{sol}}H$Δsol​H. Show the reasoning of your cycle and take care with the sign of the lattice term.

(5 marks)

A metal carbonate decomposes on heating according to the equation:

$\text{MCO}_3(s) \rightarrow \text{MO}(s) + \text{CO}_2(g)$MCO3​(s)→MO(s)+CO2​(g)

For this reaction, $\Delta H = +176 \text{ kJ mol}^{-1}$ΔH=+176 kJ mol−1 and $\Delta S = +220 \text{ J K}^{-1}\text{ mol}^{-1}$ΔS=+220 J K−1 mol−1.

(a) Using $\Delta G = \Delta H - T\Delta S$ΔG=ΔH−TΔS, calculate the minimum temperature at which this decomposition becomes thermodynamically feasible. Show your working and take care with units. (3 marks)

(b) Explain why this reaction is not feasible at room temperature (298 K) but becomes feasible above the temperature you calculated. (2 marks)

Ethene reacts with hydrogen in the gas phase to form ethane:

$\text{H}_2\text{C}{=}\text{CH}_2(g) + \text{H}_2(g) \rightarrow \text{CH}_3\text{CH}_3(g)$H2​C=CH2​(g)+H2​(g)→CH3​CH3​(g)

Use the mean bond enthalpies below to calculate the enthalpy change of this reaction, $\Delta H$ΔH, in kJ mol⁻¹. Show your working.

Bond	Mean bond enthalpy / kJ mol⁻¹
C=C	612
C–C	347
C–H	413
H–H	436

(4 marks)

(a) Define the term standard enthalpy of formation, $\Delta_{f}H^{\circ}$Δf​H∘. (2 marks)

(b) Write a balanced equation, including state symbols, that represents the standard enthalpy of formation of magnesium oxide, MgO(s). (1 mark)