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Is equilibrium between dissolution and recrystallization of a salt in an aqueous solution reached only when the solution is saturated; or can this also be established in an unsaturated solution?

Will Le Chatelier's principle apply in this situation (as it is about chemical equilibria)?

I was actually studying purification of salt by common ion effect where salt crystallizes from saturated aqueous solution on addition of common ion. Will the salt still crystallize if solvent still can dissolve more of it in unsaturated solution on addition of little amount of common ion. Since this crystallization is consequence of Le Chateliers principle, will crystallization still occur if solution taken was unsaturated and the common ion added is also in very small amount?

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    $\begingroup$ If the solid salt is present together with the solution of the salt, the solution is saturated and in equilibrium with the solid (provided you have allowed time for the solid to dissolve etc.) $\endgroup$ – porphyrin May 14 '17 at 15:00
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    $\begingroup$ Technically, in an unsaturated solution, you are also at equilibrium. $\endgroup$ – Zhe May 14 '17 at 15:38
  • $\begingroup$ Le Chatelier's principle will apply $\endgroup$ – JSCoder says Reinstate Monica Oct 12 '17 at 13:04
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The solubility of a solute in a solvent is determined by the chemical potentials of the solute for the solution and the pure solute. Substances move from high chemical potential to low chemical potential. At the concentration of solubility, these two chemical potentials are equal so there is no net transfer. Below this concentration, the chemical potential of the solute in the solution is lower than that of the pure solute so molecules will transfer from the pure solute to solvent until it is all dissolved or saturation is reached.

When you invoke the common ion effect, you are increasing the chemical potential of the common ion in solution which decreases the solubility. It is possible to have a nearly saturated solution and add enough common ions to reduce the solubility below the original concentration.

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Beginning level Thermodynamics is about path independent processes near (or at) equilibrium. Almost no thermodynamic systems are left undisturbed long enough for true equilibrium to occur, so we speak about "near equilibrium" states of isolated, closed, or open systems. What exactly we mean by "near" is not obvious or easy. A solution in equilibrium is one in which the concentration of solute is constant (and stable). Visual observation of a solution to determine if it is stable is not optimum, although it is convenient. The problems with this approach are, I think, beyond the scope of your question. (They involve the actual molecular and atomic scale processes which occur during crystallization/dissolution.) So, to answer your first question: Equilibrium is reached when the concentration is constant. If sufficient solute is present to saturate the system, then it is saturated equilibrium, if less is present, then unsaturated equilibrium. The answer to your second question is a bit more complicated. If I take a saturated salt solution and add more solid salt to it, does the concentration in the solution change? No. So Le Chateliers fails here. You need to be aware of the difference between quantitative chemical relationships (for instance equilibrium constants) and qualitative ones (eg. Le Chats.Principle). I shouldn't have to tell you which one is more trustworthy. (OTOH, quantitative relationships are all specific, while qualitative ones can often be applied without knowing all of the details - they're more general because they're less specific. But along with generality comes exceptions to these "rules"....) Does Le Chatelier's Principle apply here? I suppose it could but I don't see any useful application of it. I guess the common ion effect could be looked at as an example of Le Chat's, but it seems to me easier to explain by looking at the equation for Ksp (quant wins over qual). As to your third question: It isn't coherent. "Will something crystallize out if it isn't at saturation?" is what you're asking. By definition, no. Can you reach saturation by adding a common ion rather that the solute of interest? Yes, as you suggest, this will be (generally) in solutions near saturation (and recall, saturation implies equilibrium.) But keep in mind that adding a common ion is actually adding part of the solute of interest - the [Anion] or [Cation] is increasing, and so the ratio of reactants to products also changes, moving closer to saturation.

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