Entropy, Lattice Energy and Born-Haber Cycles: Practice Questions

Which of the following processes has a positive standard entropy change of the system, \(\Delta S_{\text{system}}^{\ominus}\)?

Which equation correctly represents the standard enthalpy change of atomisation of bromine?

Consider the data given below for potassium bromide, \(\text{KBr}\):

Enthalpy of atomisation of potassium, \(\Delta H_{\text{at}}^{\ominus}[\text{K}(s)] = +89\text{ kJ mol}^{-1}\)
First ionisation energy of potassium, \(IE_1[\text{K}(g)] = +419\text{ kJ mol}^{-1}\)
Enthalpy of atomisation of bromine, \(\Delta H_{\text{at}}^{\ominus}[\frac{1}{2}\text{Br}_2(l)] = +112\text{ kJ mol}^{-1}\)
First electron affinity of bromine, \(EA_1[\text{Br}(g)] = -325\text{ kJ mol}^{-1}\)
Standard enthalpy of formation of \(\text{KBr}(s)\), \(\Delta H_f^{\ominus}[\text{KBr}(s)] = -394\text{ kJ mol}^{-1}\)

What is the lattice energy (enthalpy of formation of the crystal from gaseous ions) of \(\text{KBr}(s)\)?

Which of the following processes results in the greatest increase in the entropy of the system, \( Δ S_{\text{system}} \)?

Consider the gaseous decomposition of dinitrogen tetroxide:
\( \text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g) \)    \( \Delta H = +58 \text{ kJ mol}^{-1} \)
What is the entropy change of the surroundings, \( \Delta S_{\text{surroundings}} \), at \( 298 \text{ K} \)?

Explain why the entropy change of the system, \(\Delta S_{\text{system}}\), is positive for the following thermal decomposition reaction:
\(CaCO_{3}(s) \rightarrow CaO(s) + CO_{2}(g)\)

A chemical reaction occurs at a constant temperature of \(298 \text{ K}\) with an enthalpy change, \(\Delta H\), of \(-62.0 \text{ kJ mol}^{-1}\). Calculate the entropy change of the surroundings, \(\Delta S_{\text{surroundings}}\), in \(\text{J K}^{-1} \text{ mol}^{-1}\).

State the fundamental condition required for a chemical reaction to be thermodynamically feasible in terms of the total entropy change, \(\Delta S_{\text{total}}\).

(a) Define the term standard enthalpy change of neutralisation, \(\Delta_{\text{neut}}H^{\ominus}\).

(b) In an experiment, \(50.0 \text{ cm}^{3}\) of \(1.00 \text{ mol dm}^{-3}\) hydrochloric acid is reacted with \(50.0 \text{ cm}^{3}\) of \(1.00 \text{ mol dm}^{-3}\) sodium hydroxide solution. The initial temperature was \(19.5 \text{ }^{\circ}\text{C}\) and the maximum temperature reached was \(26.2 \text{ }^{\circ}\text{C}\).
Calculate the enthalpy change of the reaction in \(\text{kJ mol}^{-1}\).
(Assume the density of the solution is \(1.00 \text{ g cm}^{-3}\) and the specific heat capacity is \(4.18 \text{ J g}^{-1} \text{ }^{\circ}\text{C}^{-1}\)).

The decomposition of calcium carbonate is shown by the equation below:
\(\text{CaCO}_{3}(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_{2}(\text{g})\)
\(\Delta_{r}H^{\ominus} = +178 \text{ kJ mol}^{-1}\)

The standard molar entropies, \(S^{\ominus}\), at \(298 \text{ K}\) are:
\(S^{\ominus}[\text{CaCO}_{3}(\text{s})] = 92.9 \text{ J K}^{-1} \text{ mol}^{-1}\)
\(S^{\ominus}[\text{CaO}(\text{s})] = 39.7 \text{ J K}^{-1} \text{ mol}^{-1}\)
\(S^{\ominus}[\text{CO}_{2}(\text{g})] = 213.6 \text{ J K}^{-1} \text{ mol}^{-1}\)

(a) Calculate the entropy change of the system, \(\Delta S_{\text{system}}^{\ominus}\), for this reaction.

(b) Calculate the entropy change of the surroundings, \(\Delta S_{\text{surroundings}}^{\ominus}\), at \(298 \text{ K}\).

(c) Calculate the total entropy change, \(\Delta S_{\text{total}}^{\ominus}\), at \(298 \text{ K}\). Explain whether the reaction is thermodynamically feasible at this temperature.