Let's start with a saturated solution of NaCl in water. At this point, any further addition of NaCl at the same T and P would not dissolve. However, if I then add another fully soluble salt (say AgNO3) it will also dissolve until it reaches equilibrium. If I keep adding different soluble salts (assume I don't encounter things like the common ion effect), would I ever overwhelm the solvent such that it could not dissolve anything else?

My intuition is that eventually, there would be no more solvent moleclues available for solvation and thus nothing would dissolve. Is this even valid to think about?


2 Answers 2


Your question seems to be exclusively about salts dissolving to form ions in water. The answer is fairly trivial: a saturated solution of a salt in water is not water.

At (thermodynamic) equilibrium, the solubility depends on the net free energy of the system. When the interactions in solution no longer energetically justify charge separation of the ionic solid into aqueous ions, then you've got equilibrium, right?

So, consider a solution as being (in a very crude and simplified way) Solvent-M1 interactions, Solvent-X1 interactions, solvent-solvent interactions, and M1(aq)-X1(aq) interactions as well as M1(aq)-M1(aq) and X1(aq)-X1(aq) interactions. Which of those will change when you add M2X2 ions to it? Well, at low [M2X2] none of those will change significantly, but you'll add the interactions between M2 and X2 and ALL of the species present. (as already explained, silver nitrate being added to NaCl is a very poor choice, AgCl2 is quite insoluble)

At high enough concentrations you'll begin to significantly change the energetics. In other words, adding M2X2 to a solution of M1X1 changes the interactions of all species present, so you would expect that the mixed salt solution will have different equilibrium concentrations for both M1X1 and M2X2.

If you were to carefully choose highly soluble salts, I think what would be surprising is how well their solubility was predicted by just using each independent Ksp, rather than the (expected) deviations from those predictions. Note that nothing I've said here should be taken to mean that adding a different salt to a saturated solution will decrease the solute concentration of the first salt. (Although it is the most likely scenario, just as the common ion effect would suggest. In some cases solubility will increase!) The reality is that highly concentrated solutions are far more idiosyncratic and generalizations (especially "laws") are much harder to come by.

While there probably is justification to claim that water is dissolved in sodium nitrate (at equilibrium), I've never found such foolishness useful. And I've been around for a long while. (ultra) Pure water does not freeze at 0°C, negative absolute temperatures (negative Kelvin) are attainable, and a bunch of other stuff which while true, doesn't really aid understanding (for the novice) probably should be ignored until one wants to deep dive into a subject.

  • $\begingroup$ Thanks for the responses. My take away is that you can't view each solute's interaction with the solvent independent of all the other species present so each new salt added changes the mixture and as such the mixture properties also change. Also, I'm curious about your comment for the ultra pure water not freezing at 0°C, do you mean to hint at phase meta-stability due to no nucleation sites? $\endgroup$
    – J. Ari
    Jul 15, 2016 at 21:38
  • $\begingroup$ chemistry.stackexchange.com/questions/2530/… it's not "normal" temperature that can be negative $\endgroup$
    – Mithoron
    Jul 15, 2016 at 23:06

Dont add AgNO3 to NaCl. It will form insoluble AgCl.

Yes, you can add salts M1X1, M2X2, M3X3. Eventually it will turn out that solubility of A1X3 is lover than solubility of M1X1 and M3X3 and it will precipitate.

Next by definition in physical chemistry the solvent is the one that precipitates when you cool down solution. Take a bucket of water, add a grain of salt, cool it down. Water freezes - it is a solvent. Make saturated solution of NaNO3 and cool it down. NaNO3 precipitates, so now NaNO3 is a solvent. This is not a paradox, but it does sound unusual. So now your argument doesn't hold. You add not really the "water binders", but "solvents.

Here is a different example. In organic synthesis you obtain a mixture of products. Pretty much all of them are solids. But because they are mixed the result looks like a chewing gym. Addition of yet another solid will not solidify the product (it will be more viscous, but still a liquid). Addition of a real solvent can extract small amounts of impurities and solidify your product.


Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy.

Not the answer you're looking for? Browse other questions tagged or ask your own question.