5
$\begingroup$

What concentrations of magnesium, sodium and sulfate are needed to precipitate solid MgSO4 and Na2SO4 (let's say evaporation is occurring in a system) I'm assuming probably really high because they are soluble salts? But can someone tell me in mg/L how much?

I'm asking because in this paper I read the authors are claiming magnesium and sodium in solution are coming from dissolution of MgSO4 and Na2SO4 but the concentrations are really low... like 20 mg/L of each. If those minerals were present wouldn't concentrations of Mg and Na be much higher?

$\endgroup$
1
$\begingroup$

Solubilities (taken from Wikipedia):

  • MgSO4: 35.1 g/100 mL (20 °C)
  • Na2SO4: 13.9 g/100 mL (20 °C)

It does seem that the solubilities should be an order of magnitude higher than is quoted in the paper you are reading.

Several ideas:

  1. The paper is wrong.
  2. You did not understand the paper correctly.
  3. The solubilities mentioned above are of the pure compound in pure water. If it's in an evaporite basin the the chemistry of the water might differ and the solubilities might also differ.
  4. As an example, the presence of a more soluble salt will cause the less soluble salt to be come even more soluble (a process called "salting-out"). For instance, if you have an almost saturared solution of NaCl, and you add CaCl2, a more soluble salt, you will start preciptating NaCl once enough CaCl2 is added to the system.
  5. There could be other anions in the system which make Na and Mg less soluble. Borates? Phosphates? Carbonates? Fluorides? The solubility is determined by the least soluble component of the system. Adding sulfate to magnesium silicate isn't going to do much to it in terms of solubility (unless it is sulfuric acid of course).
$\endgroup$
  • $\begingroup$ Thank you Michael. Just for clarification, this is what is said in the paper. "The higher autumn concentration of Ca, Mg and SO4 could be due to evaporation, drying of the moss and concentrating of the water during the dry summer. With Autumn rains, the Ca and Mg salts of sulfate would be dissolved and released in the water". The maximum concentration of Ca/Mg they report are 476 uEq L-1 and Mg 375 uEq L-1, respectively. $\endgroup$ – user378381 Dec 25 '17 at 18:56
  • $\begingroup$ In any kind of surface water I'd absolutely expect the carbonate concentration to be a factor. You'll get carbon dioxide from the air. $\endgroup$ – MaxW Dec 26 '17 at 20:18
1
$\begingroup$

Magnesium sulfate is considerably more soluble than calcium sulfate. What happens during evaporation will depend on the original ion concentrations but, for example with sea water, evaporation will cause gypsum to precipitate first. This will decrease the sulfate concentration until it is gone before magnesium sulfate will precipitate. The remaining calcium and magnesium will be charge-balanced by chloride and will increase in concentration past the point of halite precipitation. Only after further evaporation will magnesium and calcium chloride precipitate, but you won't get magnesium sulfate.

To get magnesium or sodium sulfate to precipitate you would have to have more sulfate in the system than calcium (on a meq basis) so that there is sulfate remaining after all the calcium is precipitated. Without any further details on the system they studied, it's hard to say if this is possible or not. Personally I would suspect that chlorides would be more likely.

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

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