Reef Chemistry Question of the Day #162 Calcium Carbonate Precipitation

[Randy Holmes-Farley]

Reef Chemistry Question of the Day [HASHTAG]#162[/HASHTAG]

Calcium carbonate is recognized as being supersaturated in normal seawater at an alkalinity of 7 dKH, calcium at 420 mg/L, and pH of 8.2. That means there is the potential for precipitation of calcium carbonate from solution. As any of those rise, the potential for precipitation rises even more.

But the precipitation of calcium carbonate is slow in the ocean and in most reef tanks due to a variety of processes that deter the precipitation.

Which of the following is not likely to deter the precipitation of calcium carbonate in a reef tank?

A. Higher temperatures, which increase the solubility of calcium carbonate in seawater

B. Magnesium by ion pairing with carbonate ions in solution, effectively reducing the free carbonate concentration

C. Magnesium by getting onto growing calcium carbonate crystals in place of calcium ions, effectively messing up the crystal order for additional precipitation to take place on top of it

D. Sulfate by ion pairing with calcium ions in solution, effectively reducing the free calcium concentration

E. Sodium by ion pairing with carbonate ions in solution, effectively reducing the free carbonate concentration

F. Organic compounds getting onto growing calcium carbonate crystals, effectively shielding the crystal for additional precipitation to take place

Good luck!





























.

 

[beaslbob]

wild guess

f

?

 

[pfoxgrover]

B

 

[twilliard]

Tough one Randy!
I will have to dig into this one

 

[JoaoTomas]

Maybe B option

 

[zzzdiver]

Wow this will take some work to figure out

 

[nervousmonkey]

A, B, C and E

 

[Toadfish]

My chemistry is very rusty, but I want to say since the enthalpy of formation is -1433 kJ/mol for calcium sulfate and -1207 for calcium carbonate, the calcium will have a higher affinity for the carbonate than the sulfate. Mag and sodium both have a higher affinity for carbonate than calcium, hence why we dose magnesium. The sodium has interplay with other compounds too though, so I'm not sure how much it can really be said to help. It's not going to hurt at least.

As for crystallization inhibition, both C and F seem to be likely to not change the total rates of precipitation, merely the clast size of the precipitate. That being said, small clast is significantly less annoying to see in a tank, and has a higher surface area for further reaction.

Based on, again, very rusty chemistry knowledge, the only thing that won't have ANY detrimental effect to calcium carbonate formation is added sulfate, so D. There are varying degrees of effectiveness of the other methods though, and in particular I'd avoid adding free organics as that has... compounding issues.

How did I do?

 

[Randy Holmes-Farley]

How did I do?

Patience.

I usually give the answer after a few days.

 

[Toadfish]

Patience.

I usually give the answer after a few days.

Gotcha. First couple days here, didn't know how it goes. I'll keep an eye on this!

 

[Randy Holmes-Farley]

Gotcha. First couple days here, didn't know how it goes. I'll keep an eye on this!

Welcome to Reef2Reef!

 

[beaslbob]

I now thing b

Magnesium slows precipitation by physically interfering with the formation of calcium carbonate.

But I'm usually wrong on these.

 

[Whatnow]

C

 

[marke]

C

 

[marke]

Magnesium is especially important for its role of preventing the abiotic precipitation of calcium carbonate from seawater. Seawater is supersaturated with respect to calcium carbonate, but any time that it begins to precipitate, magnesium attaches to the growing crystal's surface and inhibits further precipitation. Consequently, the ocean can stay supersaturated for long periods of time from Dr Farleys artlcle

 

[Randy Holmes-Farley]

C

The question was which is NOT likely to be preventing precipitation.

 

[Taservices]

I'd like to guess A (on a technicality that higher temps would probably result in it not being a reef tank for much longer haha).

I really have no idea, it has been a long time since I've done any real non-organic chemistry. I look forward to seeing the answer.

 

[Randy Holmes-Farley]

And the answer is...A. Higher temperatures, which increase the solubility of calcium carbonate in seawater

Calcium carbonate, both in seawater and fresh water is less soluble as the temperature rises. So this cannot explain why it doesn't precipitate.

I show the temperature effects on calcium carbonate, and discuss the reasons why, here:

Chemistry and the Aquarium: Calcium ? Advanced Aquarist | Aquarist Magazine and Blog
http://www.advancedaquarist.com/2002/3/chemistry

In short, part of the reason is that the acidity of bicarbonate increases with temperature, so as the temp rises, more HCO3- ***** into CO3-- and H+, making more carbonate available for precipitation (the equilibrium below shifts to the right)

HCO3- <-- --> H+ + CO3--

The other part of the reason is that the solubility of calcium carbonate decreases with increases in temperature even if the CO3-- concentration is held steady.

The other possible answers are details in later posts.

Happy Reefing!

 

[Randy Holmes-Farley]

As to the wrong answers, here's my discussion:

C. Magnesium by getting onto growing calcium carbonate crystals in place of calcium ions, effectively messing up the crystal order for additional precipitation to take place on top of it.

That definitely happens. I discuss it here:

A Simplified Guide to the Relationship Between Calcium, Alkalinity, Magnesium and pH by Randy Holmes-Farley - Reefkeeping.com
http://reefkeeping.com/issues/2006-06/rhf/index.php

from it:

Calcium Carbonate and Magnesium
Finally, we come to magnesium's role in the calcium carbonate system. The situation for magnesium is appreciably more complex than for pH and alkalinity, but we can continue our same analysis to understand it qualitatively. When solid calcium carbonate is put into seawater, it doesn't just undergo the sorts of "on" and "off" dynamics as calcium and carbonate ions discussed above. Other ions can get into the crystal structure in place of either of these ions. In seawater, magnesium ions get into calcium carbonate crystals in place of calcium ions. Strontium ions may also do so, but their numbers are far lower than magnesium's (about 600 times lower) so they are less likely to become incorporated.

Figures 8 and 9 show how magnesium in solution gets onto and actually into a thin layer of calcium carbonate surface put into seawater. Even though magnesium carbonate itself is soluble enough that it will not precipitate from normal seawater, in a mixed calcium and magnesium carbonate structure, its solubility is lower. So solid, pure calcium carbonate (Figure 8) is rapidly converted to a material with a coating of calcium and magnesium carbonate (Figure 9).

This coating has some very important effects. The primary effect is that it makes the surface no longer look like calcium carbonate, so calcium and carbonate ions that land on it no longer find the surface as inviting as before. The magnesium ions have altered the surface in a way that does not hold calcium and carbonate as strongly, and so the "off" rate of any newly landing calcium and carbonate ions is higher (Figure 10). Consequently, even if the driving force to deposit calcium carbonate is still there, the magnesium has gotten in the way and doesn't allow it to happen (or keeps it from happening as fast).

The extent to which magnesium gets onto calcium carbonate surfaces depends strongly on the amount of magnesium in solution. The more there is, the more it gets onto the surfaces. If magnesium is lower than normal, then it may not adequately get onto growing calcium carbonate surfaces, allowing the deposition of calcium carbonate to proceed faster than it otherwise would, potentially leading to increased abiotic precipitation of calcium carbonate from seawater onto objects such as heaters and pumps. Often the inability to maintain adequate calcium and alkalinity despite extensive supplementation, and the precipitation of significant amounts of calcium carbonate on heaters and pumps, are signs that the water has inadequate magnesium.

 

[Randy Holmes-Farley]

All of the ion pairing answers are also wrong in the sense that they all happen and all reduce the likelihood of precipitation of calcium carbonate.

It is a little complicated, because these things are already figured in when determining the supersaturation of calcium carbonate. So they are happening even when calcium carbonate is just exactly saturated. These effects are big. if you had exactly saturated calcium carbonate (say, at much lower calcium or alkalinity or pH than normal seawater) and then you suddenly removed the sulfate, magnesium and sodium and chloride, the solution would suddenly be supersaturated again because this ion pairings that reduce calcium and carbonate free forms are gone.

If that doesn't make sense to someone, I'm happy to explain more.

So these answers are wrong because they do happen:


B. Magnesium by ion pairing with carbonate ions in solution, effectively reducing the free carbonate concentration

D. Sulfate by ion pairing with calcium ions in solution, effectively reducing the free calcium concentration

E. Sodium by ion pairing with carbonate ions in solution, effectively reducing the free carbonate concentration

I discuss some of these in these articles:

Aquarium Chemistry: Magnesium In Reef Aquaria ? Advanced Aquarist | Aquarist Magazine and Blog
http://www.advancedaquarist.com/2003/10/chemistry

from it:
Magnesium is present in seawater as the Mg2+ ion, meaning that it carries two positive charges, just as calcium does. Most of the magnesium is present as the free ion, with only water molecules attached to it. It is estimated that each magnesium ion has approximately eight water molecules tightly bound to it. That is, water molecules that are so tightly bound that they move with it as the magnesium ion moves through the bulk of the water. For comparison, singly charged ions like sodium have only 3-4 tightly bound water molecules. A small portion (about 10%) of the magnesium is present as a soluble ion pair with sulfate (MgSO4), and much smaller portions are paired with bicarbonate (MgHCO3+), carbonate (MgCO3), fluoride (MgF+), borate (MgB(OH)4+), and hydroxide (MgOH+).

While these ion pairs comprise only a small portion of the total magnesium concentration, they can dominate the chemistry of these other ions. An extended discussion of these facts is beyond the scope of this article, but is should be noted that these ion pairs can have huge impacts on seawater chemistry. In the case of carbonate, for example, the ion pairing to magnesium so stabilizes the carbonate that it is present in far higher concentrations than it would be present in the absence of magnesium. This effect, in turn, makes seawater a much better buffer in the pH range of 8.0-8.5 than it otherwise would be. Without this ion pairing, seawater pH might be significantly higher, and more susceptible to diurnal (daily) swings.

Chemistry and the Aquarium: Calcium ? Advanced Aquarist | Aquarist Magazine and Blog
http://www.advancedaquarist.com/2002/3/chemistry

from it:
In seawater, the situation is slightly more complicated. While the majority of calcium ions are still free, some (about 10-15%) are present as an ion pair with sulfate, forming the neutral ion pair CaSO4 (Figure 2). These types of soluble ion pairs are short lived, forming and breaking apart quite rapidly. Nevertheless, they can have significant impact on the properties of seawater. This ion pair is in turn hydrated with water molecules, as shown in Figure 2.

Calcium similarly forms ion pairs with carbonate and bicarbonate. While these comprise a small fraction of the total calcium, the calcium carbonate ion pair comprises a fairly large portion of the total carbonate (together with magnesium, about 2/3 of the carbonate). These ion pairs consequently tend to lower the free concentration of carbonate, and thereby help to inhibit precipitation of calcium carbonate, and consequently increase its solubility.

Finally, calcium forms ion pairs with fluoride, hydroxide, borate, the various forms of phosphate, and other ions to smaller extents that are unimportant to the free calcium concentration, but may impact the free concentrations of these other ions (especially phosphate, where calcium binds to more than 70% of the PO4---). In almost all cases, however, the effect of calcium is smaller than the effect of magnesium on these ions, both because the concentration of magnesium is higher, and because in some cases it actually interacts more strongly (MgF+ compared to CaF+, for example).