12-6 (1)

12-6 (1) - The answer must lie with the enthalpy change for...

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Which is a plausible explanation for the low solubility of AgCl(s) in water? a) The lattice energy (absolute value) of AgCl is greater than the enthalpy of hydration for the ions. b) The enthalpy of hydration for Cl - is very exothermic. c) The enthalpy of hydration for Ag + is endothermic. d) The heat of solution (the delta H of the dissolving process) of AgCl is exothermic. e) Small ions, such as Cl - , have a positive enthalpy of hydration. f) The entropy change in dissolving a salt is always negative. Key: There are two driving forces which promote a process like this: making entropy more positive and lowering the energy (enthalpy). First, entropy: Clearly, when a salt like AgCl dissolves, the entropy of the system increases as the highly ordered crystal converts to the highly disordered solution. So entropy doesn’t account for the low solubility—entropy actually works in favor of dissolving. This eliminates answer f.
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Unformatted text preview: The answer must lie with the enthalpy change for the process: solid AgCl converts to dissolved AgCl. The enthalpy change must be working against this process, so must be positive (eliminating answer d). The enthalpy change is the difference in enthalpy between the undissolved crystalline state and the solution state. The lattice energy of the crystal must be really large (negative value) compared to the hydration energy (negative value). Remember that systems prefer to go to their lowest (most negative) energy state. So, in this case, the AgCl prefers to remain in the crystal rather than dissolve because its lattice energy must be more negative than its solvation energy. The enthalpies of hydration of the individual ions alone cannot provide the explanation—one must take into account all the energies in the solution process, both lattice and solvation energies. This eliminates answers b, c, and e....
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This document was uploaded on 11/04/2011 for the course CHEM 106 at BYU.

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