Ca, Alk solubility in relation to temp

[Sleepydoc]

Somewhere in the past, probably the interweb, I read that the solubility of alkalinity goes down temperature. I was recently doing some looking, and all the information I found shows the solubility of both CaCl and Na2CO3 going up with increasing temperature. Was my original information wrong? The information I found was all for CaCl and Na2CO3 individually, not combined as with sea water, so that may be the issue as well.

 

[Mjrenz]

Interesting question, and I think is the guy to ask @Randy Holmes-Farley

 

[JimWelsh]

I think you mean the solubility of CaCO3. See section 2.3.1: http://www.iwtc.info/2003_pdf/08-5.pdf

 

[Dkeller_nc]

Nice paper, Jim - we're going to nominate you as the official documentation researcher on R2R's Chemistry Forum.

SleepyDoc - If you go to the wikipedia page on calcium carbonate, there's a comment about solubility in seawater vs. temperature - it's in the section on Carbonate Compensation Depth. This is the quote:

"Calcium carbonate is unusual in that its solubility increases with decreasing temperature. Increasing pressure also increases the solubility of calcium carbonate. The carbonate compensation depth can range from 4,000 to 6,000 meters below sea level."

This is also the case in freshwater, btw, which is why lime deposits often occur in the hot water plumbing in a house in a hard water area.

 

[Sleepydoc]

Silly me - I searched for the solubility of CaCl and Na2CO3 rather than the solubility of CaCO3, and didn't bother to check wikipedia! So 'alkalinity' in terms of Sodium Carbonate is more soluble with temperature but CaCO3 is less soluble with increasing temperatures.

It does raise a couple of other questions, though. If salt mixes use a combination of CaCl, and Na2CO3 to provide calcium and alkalinity, the solubility of both of these increases with temperature. A solution with both of these would not necessarily behave the same and I assume that there is some sort of equilibrium between CaCO3 and dissolved Ca++/CO3--. Since the solubility of CaCO3 is significantly lower than the precursors, the mechanics would seem to favor CaCO3 precipitation. I know Mg helps stabilize CO3-- and prevent CaCO3 precipitation. I've never thought about the exact mechanism, but MgCO3 is not soluble in water either, so preferential formation of MgCO3 wouldn't really explain it. Now I'm trying to figure out why I don't just have a lump of calcium carbonate at the bottom of my tank!

All this is really beyond the basic chemistry I took as an undergrad 20+ years ago, and I should probably just be happy with the fact that I can take some salt mix, throw it into my mixing tub and have saltwater for my critters, but now I'm curious.

 

[JimWelsh]

What you are wondering about is the CaCO3 saturation state, commonly referred to with the upper-case greek letter omega. Here's a little light reading on it: https://pdfs.semanticscholar.org/7192/cede9fa4e2822fec6dd282d7b11079f0220e.pdf

 

[Dkeller_nc]

Once again, Jim's come through with an excellent article - I remember reading that one some time back.

But to answer your question in shorthand, you're right - if you have a solution of simple pure water and calcium and carbonate/bicarbonate, and at neutral pH, just about all of the calcium and alkalinity will precipitate out of solution (presuming they're in equal molar ratios). The only remaining Ca & HCO3-/CO3-- will be the concentration that equates to the solubility of calcium carbonate at that temperature and pH, which is exceedingly low (about 13 milligrams per liter at 25 deg C). This, by the way, is Le Chatelier's principal at work - because the reaction product of the calcium ion and carbonate ion is largely insoluble, the precipitation of the compound continues to drive the reaction of more calcium and carbonate to form more calcium carbonate - until such point as the rate of calcium carbonate dissolution in water equals the precipitation rate of calcium carbonate, and the reaction stops.

Randy has described why this doesn't happen in saltwater in various articles; essentially, magnesium "poisons" the surface of growing calcium carbonate crystals, providing an energy barrier to further reaction of calcium and carbonate to form more calcium carbonate. In addition, organics, especially the ones that act as chelating agents, "bind up" the calcium, preventing its reaction with carbonate and subsequent precipitation.

On one level, calcium and alkalinity in seawater is fairly simple - just look up what the saturation value is for your particular temperature and pH, and there you have it. But when you start diving into the reasons for why seawater has the saturation levels of calcium and alkalinity that it does, especially in the context of the actual ocean, things get complex (and to us chemistry geeks, interesting) very quickly.