I should really know these things by now, but it's driving me nuts.

I'm looking at this: http://www.caton.org/images/chem/TableD.gif

Lets say I dissolve 85g of KNO3 at 50C, to make a saturated solution of it. Could I then dissolve any NaCl in it? I mean, the classic way I learned about solubility is that the ions are cars, and there are only so many parking spaces (water molecules).

So, what if I dissolved some NaNO3 in water, and some KCl. Which curve would apply? The one for the NaNO3, or the one for the NaCl, or the KCl, or the KNO3?

Help a confused guy out.

I suppose it would be whichever one is preferentially the most soluble... but remember, (using your metaphor) that once you park, the cars break in half and wind up a jumbled bag of who-knows-what. it's no longer seperate substances.

More soluble things will remain in solution, while others will precipitate first. It's the principle used when we cool down a solution to precipitate a product. The more insoluble (we hope) product precipitates, leaving the rest of the stuff in solution.

Solubility curves get more complicated when dealing with more than one thing.

Yeah, and kinda like salting out MEKP.

Well, I was trying to find an example, using actual math, but I got frustrated. Anyway, I was thinking that if you dissolved two salts each made of completely soluble ions (group 1, nitrates, and Cl-), so that both the salts you dissolved didn't saturate the solution, yet a salt made from the positive ion of one salt and the negative from the other would be supersaturated in the solution, would it precipitate out?

IE
KCl and NaNO3, both unsaturated, is like KNO3 and NaCl, but the NaCl could be supersaturated. Would it precipitate out? Once ions are in a solution, there aren't salts, just ions, right?

Ok, if you have a group of ion dissolved in a solution, you have to assume that the addition of addition ion will create new salts. So say you have an aqueous solution of K+, NO3-, Na+ and Cl-. Even though they all exist as ions, you have to assume that you are dealing with KNO3, NaNO3, NaCl, and KCl all together. In this case, according to the solubility curves you provided, for temperatures above 40C, the solubility of NaCl would be your limiting solubility. However, if you consider the reaction: KNO3 + NaCl ---> NaNO3 + KCl you should see that the Sodium ions produced by the NaCl you’re adding will displace the potassium ions in solution. In fact, at that temperature of 50C that you specified, it seems to me from the solubility curves, that the KCl produced by that reaction would precipitate out of solution.

So as long as you add different ions to a saturated solution, you can generally expect an equilibrium to form. If the solution is saturated and you add a non-reactive compound, you would expect that compound to not dissolve and simply pile up in the bottom of the beaker. However, when you have a sodium salt that you add to a solution of potassium ions, you can expect the sodium to replace the potassium because Na is more reactive than K. Yet, if you were to add a common ion to the solution, it's a whole different story. We just reviewed common ion calculations in my chem class, and they're rather complex, so I won't explain that unless you want. But basically, it sets up a different equilibrium.

So you’re right with your car analogy that a given solvent can only dissolve a certain amount of any solutes. Anything that is non-reactive with the solution will not dissolve in, and anything that does go into solution with a saturated solution will force something else from that solution to be precipitated out. All according to LeChatlier's principle.

Does that help clear things up?

Eh, kinda.

It's just that there's a contradiction here. The curve for one salt in the solution says it's completely soluble, and the curve for another salt with a mutual ion says it should precipitate out.