A student analyzes the dissolution of KCl in liquid water at 25∘C and reasons as follows:
"Because hydrating the ions releases more energy (−690kJ mol−1) than is needed to separate the ions in the crystal (+657kJ mol−1), the overall free‐energy change for dissolving KCl must be negative. Therefore, KCl should be highly soluble at 25∘C."
Experimental measurements report ΔG∘=+1.7kJ mol−1 and a molar solubility of ≈4.5M under these conditions.
Answer the following questions about the student’s reasoning and the thermodynamic description of the process.
(1) Part 1: (a) Identify the two distinct energetic processes whose enthalpy values the student compared in the quoted statement.
(2) Part 2: (b) Describe how the entropy change associated with dissolving an ionic solid can offset an apparently favorable enthalpy change.
(3) Part 3: (c) Explain why comparing only the numerical magnitudes of the lattice enthalpy and the hydration enthalpy is insufficient for predicting the sign of ΔG∘.
(4) Part 4: (d) Discuss one way in which reorganization of the liquid‐water hydrogen-bond network can contribute to an unexpectedly small (or even positive) entropy term for this dissolution.
(5) Part 5: (e) Based on your analysis in parts (a)–(d), evaluate the student’s claim about the solubility of KCl and provide a corrected thermodynamic explanation.