Two ionic compounds, Solid A and Solid B, are both sparingly soluble in water at 25 °C. Calorimetric measurements show that dissolution of Solid A absorbs heat from the surroundings, whereas dissolution of Solid B releases heat. In both cases the measured standard molar entropy change for dissolution is small and positive.
Answer the following questions about the relative solubilities of the two solids at different temperatures.
(a) Explain, in molecular terms, how the signs and relative magnitudes of the enthalpy change and the entropy change together determine whether the dissolution of a solid is thermodynamically spontaneous at a given temperature.
(b) For Solid A and Solid B, predict qualitatively how heating the solution from 25 °C to 80 °C will affect their solubilities. Justify each prediction using thermodynamic reasoning.
(c) For Solid B, the dissolution is observed to become less exothermic as temperature increases. Explain how this observation affects the reliability of using the sign of the enthalpy change alone to predict temperature-dependent solubility trends.