Why does the reaction of dissolution stop at an equilibrium point? [closed]

Question

The formula for Gibbs free energy is $\Delta G=\Delta H-T\Delta S$. If Gibbs free energy is negative, the reaction is spontaneous. This also applies to dissolution reactions.

However, we know that every soluble element has a dissolution constant. At a certain equilibrium point, the reaction stops. Which term in the Gibbs free energy indicates that the reaction should stop?

Moreover, according to the Le Chatelier's principle, the dissolution reaction can go in either way if we change some concentrations at the equilibrium point. But only changing concentrations does not seem to affect neither the enthalpy nor the entropy. Why does the value of $\Delta G$ vary?

Finally, is there any way to predict the dissolution constant based on $\Delta H$ and $\Delta S$?

Answer 1

"Why does the reaction of dissolution stop at an equilibrium point?" It does not stop! There are two sides to a state of equilibrium:

• The forward reaction, e.g. a salt dissolving in water, $\ce{NaCl -> Na+ + Cl-}$. Simplified for illustration.
• The backwards reaction, e.g. a salt precipitating from solution in water, $\ce{Na+ + Cl- -> NaCl}$.

At the point of equilibrium, they are proceeding equally rapidly in both directions. That means salt is dissolving and precipitating constantly. If you place a crystal of salt in a just saturated solution, it may gradually change shape, as parts dissolve and precipitate elsewhere.