Chemistry
Chemistry
Problem
Silver chloride and silver bromide have experimentally determined molar enthalpies of solution that are both approximately +56 kJ mol⁻¹, yet AgCl is noticeably more soluble in water at 25 °C than AgBr. Use molecular-level reasoning to account for this difference. (a) Describe two distinct molecular processes that contribute to the entropy change when an ionic solid dissolves in liquid water. (b) Explain why the entropy change accompanying dissolution of AgCl is likely to be more positive than that for AgBr. (c) Using your explanation from part (b), account for the observed difference in solubility between AgCl and AgBr despite their similar enthalpies of solution.Intuitive explanation
The key is recognizing that solubility depends on both enthalpy and entropy—specifically, the free energy change must be negative for dissolution to be spontaneous. Since both salts have nearly identical positive values, the difference must lie in entropy. AgCl dissolves more readily because breaking apart the more rigid, highly ordered AgCl crystal lattice releases more entropy than disrupting the AgBr lattice—the smaller Cl⁻ ion creates tighter electrostatic interactions and a more constrained solid structure, so its dissolution generates a larger favorable that overcomes the unfavorable enthalpy.
Intuitive explanation
The key is recognizing that solubility depends on both enthalpy and entropy—specifically, the free energy change must be negative for dissolution to be spontaneous. Since both salts have nearly identical positive values, the difference must lie in entropy. AgCl dissolves more readily because breaking apart the more rigid, highly ordered AgCl crystal lattice releases more entropy than disrupting the AgBr lattice—the smaller Cl⁻ ion creates tighter electrostatic interactions and a more constrained solid structure, so its dissolution generates a larger favorable that overcomes the unfavorable enthalpy.