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.