Three crystalline ionic salts—LiF, NaCl, and CsI—are individually introduced into liquid water at 298 K. Qualitative thermodynamic information obtained from independent calorimetric and spectroscopic investigations is summarized below.
Verbal description of data table (all values are relative, not absolute):
• Column 1 – Lattice enthalpy magnitude (very large, large, moderate, small).
• Column 2 – Enthalpy of ion–water interaction (very exothermic, exothermic, mildly exothermic).
• Column 3 – Net entropy change for the dissolution process (substantially positive, slightly positive, negligible).
Data entries
LiF : very large | very exothermic | negligible
NaCl : large | exothermic | slightly positive
CsI : moderate | mildly exothermic | substantially positive
(a) Identify which salt is expected to dissolve spontaneously in water at 298 K and which salt is least likely to dissolve. Provide both names.
(b) For the salt you selected in part (a) as spontaneous, describe in detail how the relative magnitudes of the lattice enthalpy, the ion–water interaction enthalpy, and the entropy term combine to yield a negative standard free-energy change.
(c) Predict how an increase in temperature to 350 K would influence the spontaneity of dissolution for each of the three salts. Support each prediction with thermodynamic reasoning.
(d) Suppose the same three salts are instead introduced into liquid hexane (a non-polar solvent) at 298 K. Without performing calculations, rank the salts from most to least likely to dissolve in hexane and justify your ranking in terms of the enthalpic and entropic factors listed above.