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For the solvation of ionic salts, if we look down a group, say Group II, and we keep the anion constant, we find that the solubility decreases going down the group. This is because the hydration enthalpy is more exothermic for smaller ions of the same charge. However, the smaller ions also make the lattice energy larger, yielding two counteracting terms in the equation for the full enthalpy of solvation. It seems to me that both are due to the same fundamental principle, charge density. Why is it so that the increase in exothermicity of the hydration enthalpy more than makes up for the increase in the lattice energy for smaller ions?

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    $\begingroup$ This link will answer your question... (Hopefully). Its a previous SE.chemistry post. $\endgroup$ – dval98 Oct 5 '20 at 21:19

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