Reef Alchemy by Randy Holmes-Farley

What is that Precipitate in My Reef Aquarium?


As time goes by, most reef aquarists will encounter a solid precipitate of some sort in conjunction with their reef aquarium. Often these are undesirable particles that cloud the water, coat the glass or clog pumps. Others are more benign and disappear shortly after forming. This article describes a variety of different precipitates that aquarists encounter, and details what they are, how and why they form, how to tell the different types apart, whether they are a concern and how to avoid them.

The article is broken up into sections. The first ten sections discuss different precipitates, while the latter seven discuss the detailed information supporting what the precipitates are and what factors contribute to their formation (such as pH, temperature, etc.).

Contents:


Solid Residues Remaining After Preparing Artificial Seawater


Most salt mixes leave behind a solid residue when dissolved, although the extent to which this occurs varies from brand to brand. I use Instant Ocean and rarely clean out the 44-gallon trashcan that I mix it in, so a significant residue builds up over time (Figure 1). In preparation for this article I removed some of this solid material, and found that it could be almost completely dissolved in hydrochloric acid with lots of bubbling. This demonstrates that these solids were probably calcium carbonate (CaCO3), perhaps also containing magnesium. Pure magnesium carbonate is undersaturated in seawater (which is detailed in later sections of this article) 1 and should dissolve in marine systems, so it isn't likely to be the precipitated material, although there may be significant magnesium in the calcium carbonate.

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Figure 1. The residue on the bottom of the plastic trash can that I use to mix Instant Ocean. I rarely clean it out. The solid is most likely calcium carbonate.

Based on the fact that the material exists as sheets that clearly did not arrive in the mix (as opposed to a fine powder which might have), I conclude that at least a significant fraction of this residue formed in the barrel. I cannot, however, rule out the possibility that some solid calcium or magnesium carbonate may have existed in the salt mix and was cemented together by additional precipitation of calcium carbonate during dissolution or storage.

When salt mixes are dissolved, there exist local regions where the salt concentration is very high. In those local regions, the calcium and alkalinity must also be very high. In fact, as seawater is concentrated by evaporation, there is a well-established series of minerals that precipitate as the salinity increases. In this series, calcium and magnesium carbonate are the first to precipitate, appearing at a specific gravity of about 1.140, which is about a 50% solution of salt in water.1 Such conditions may well exist on the bottom of a saltwater reservoir as the salt is dissolving.

With some mixes (but not the Instant Ocean that I use), the initial pH on dissolution may be very high (pH 8.5-9 +). As shown in detail later in this article, pH can play a dominant role in determining the rate of calcium carbonate precipitation, and such a high pH would make it more likely to precipitate.

It has been suggested by some aquarists that some salt mixes may contain anti-caking agents, such as clays. I do not know if this is true, but if it is, they may form part of the residue that is left behind after dissolution.

In order to minimize the formation of insoluble carbonate salts when mixing, the following suggestions may be helpful:

1. Add the salt to a full batch of water, rather than adding water slowly to a large batch of salt. The latter allows a greater time at much higher than natural seawater salinity, which may tend to precipitate calcium and magnesium salts.

2. Stir the mixture vigorously as it is being dissolved.

3. If using a mix with a high initial pH, aerate the mixture as well as stirring it. The aeration will reduce the pH.

Precipitates on Heaters and Pumps


The precipitate that forms on heaters (Figure 2) and pumps is primarily calcium carbonate. Natural seawater and reef aquarium water is often supersaturated with calcium carbonate, but the precipitation is kept at a very slow pace, primarily by magnesium ions (for reasons described later in this article). Certain factors, however, push the supersaturation even higher, and the "pressure" to precipitate becomes too high for the magnesium to prevent it. These factors include elevated calcium, alkalinity, pH, and in the case of heaters and pumps, elevated temperature. How these factors impact supersaturation is detailed later in the article, but the following section summarizes the effects.

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Figure 2. Calcium carbonate can form thick deposits on heaters.
It is the heat itself that helps drive this precipitation.

Chief among the factors causing precipitation are alkalinity and pH. While most aquarists would be unlikely to have in their aquaria double the calcium level of natural seawater (more than 800 ppm), many folks have doubled its supersaturation level via alkalinity and pH. With all other things being similar, for example, an aquarium with a pH of 8.5 has twice the supersaturation of one with a pH of 8.2, and an aquarium with a pH of 8.2 has twice the supersaturation of one with a pH of 7.9.

Alkalinity, likewise, can be a big contributor to precipitation. Many aquarists push alkalinity to elevated levels, approaching or exceeding twice the naturally occurring levels (about 2.5 meq/L, or 7 dKH). That change can, in turn, double the supersaturation.

Elevated temperature impacts calcium carbonate precipitation in two ways: 1) by simply reducing the solubility of CaCO3 (which is more soluble at lower temperatures) and 2) by causing the formation of additional carbonate from the bicarbonate in solution. These effects are roughly similar in magnitude, and are one of the reasons that pumps and heaters can be more rapidly coated with calcium carbonate precipitate than other objects in the same aquarium.

Another possible reason that pump impellers and heaters get coated more rapidly than other surfaces is that such surfaces may be less likely to become coated with a film of bacteria that prevent calcium and carbonate from reaching the surface. Heaters may become too hot for such bacterial films to form properly, and pump impellers may have enough water motion around them to wash away bacteria as soon as they begin to attach. Exactly how large such effects are, however, is not clear to me.

Low magnesium can also contribute to precipitation because magnesium interferes with the precipitation of calcium ions. Normally, magnesium ions slip into spots where calcium would otherwise bind, altering the surface's structure so that it no longer is a good seed crystal for further calcium carbonate precipitation. Fewer magnesium ions in solution makes this process less effective.

Finally, it appears that ferric ion (Fe+++) may help initiate the precipitation of calcium carbonate. Such ions may be added directly in some products, and can be released from iron-based phosphate binders such as Rowaphos, Phosban, and Phosphate Killer.

Preventing Precipitates on Heaters and Pumps


While the calcium carbonate that precipitates onto heaters and pumps can be readily removed by soaking them in acid (undiluted vinegar or diluted muriatic acid/hydrochloric acid: 1 part acid added to 9 parts fresh water), it is often easier to prevent the buildup in the first place. The following actions can be taken if the aquarium has excessive precipitation of calcium carbonate:

Note: actions 1, 3, and 4 may also make it harder for calcifying organisms such as corals and coralline algae to calcify, so they may grow more slowly if a significant reduction is made in pH, calcium or alkalinity.

1. Reduce the overall pH. The lower the better for this purpose, but I wouldn't go below pH 7.8. Definitely target pH below pH 8.5, even at the end of the light cycle.

2. Add high pH additives (limewater/kalkwasser and high pH two-part calcium and alkalinity additive systems, especially) so that they mix quickly and do not enter pumps or pass over heaters before mixing in well. Add them at times of the day when low pH conditions exist (usually early morning). Reduce these types of additives, if necessary.

3. Reduce the alkalinity. The lower the better for this purpose, but I would not target levels below 2.5 meq/L (7 dKH). Definitely target alkalinity below 4 meq/L (11 dKH).

4. Reduce the calcium level. The lower the better for this purpose, but I would not target the calcium below 380ppm. Definitely target the calcium below 500 ppm.

5. Maintain an appropriate magnesium level of 1250 -1350 ppm. Higher levels (up to, say, 1500 ppm) will help reduce precipitation further, and some aquarists have resorted to this, but I do not know at what level magnesium becomes toxic.

6. If you use an iron-based phosphate binder, either situate it well upstream of pumps and heaters or discontinue its use. Chelated iron supplements (added to supply macroalgae with iron) are not as likely, in my opinion, to cause precipitation, but may still contribute to some degree.

7. Use pumps that do not get as hot internally (external heat is not the concern here). I do not, however, have specific advice on which pumps would be best in this regard. Plastic impellers, while wonderful from a chemical resistance standpoint, are poor heat conductors, so the heat that comes from pushing against the water is not as rapidly dissipated as it would be by impellers made from other materials.

8. Elevated phosphate and organics can reduce calcium carbonate precipitation. While I do not recommend intentionally raising these for this purpose, aquarists may find that if they reduce the levels of these materials, that precipitation may become worse.

Precipitates Where Limewater is Added


When limewater is added to seawater, a cloudiness can form almost immediately (similar to that in Figure 3, although not usually that intense). This initial cloudiness is magnesium hydroxide, Mg(OH)2, and it forms when the water's pH rises into the low to middle 10's. Theoretical and experimental reasons for believing this material to be Mg(OH)2 (and not magnesium carbonate or calcium carbonate, for example) are given later in this article. As the limewater is mixed in, the local pH around the particulates drops, and as soon as it drops below pH 10, the magnesium hydroxide dissolves.

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Figure 3. The transient cloud of magnesium hydroxide that forms when high pH additives are added. In this case, the alkalinity portion of B-ionic was added to a fairly still portion of one of my reef aquaria.

If the limewater is not allowed to disperse rapidly enough, meaning that the pH does not drop fairly quickly, additional precipitates can form, especially calcium carbonate. Additionally, if the limewater drips onto surfaces in contact with seawater (such as the sides of a sump, Figure 4), bulk calcium carbonate can form on those surfaces. This precipitation takes place primarily because the limewater has pushed the CaCO3 supersaturation very high by converting much or all of the bicarbonate into carbonate. Since the precipitation of calcium carbonate can be slow to occur, rapid dispersal of the limewater doesn't lead to much or any precipitation of calcium carbonate. But if a region maintains high pH for long enough, calcium carbonate will precipitate. How long this process takes depends on the degree of supersaturation, but can be on the order of minutes to hours.

Figure 4. When limewater drips onto surfaces, such as the sides of a sump, precipitation of calcium carbonate takes place. The off-white coloration probably comes from metals such as iron binding to the calcium carbonate surface in the place of calcium.
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Precipitates When Carbonate Solutions are Added


Many aquarists are familiar with the cloudiness that forms when high pH two-part calcium and alkalinity additive systems are added to marine aquaria. Figure 3 shows the initial cloud that forms, for example, when the alkalinity part of B-ionic is added to a relatively low flow reef aquarium. The initial cloud sinks and spreads out, eventually dissolving. A similar phenomenon is not observed when adding calcium or magnesium salts, but is observed when adding sodium carbonate solutions.

This cloudiness is, at least in part, magnesium hydroxide and is formed when hydroxide ions are added and the local pH rises. Unlike the addition of limewater, which is unlikely to form magnesium carbonate, this may, although I think it unlikely. The reason it might form here is that the addition of the carbonate ions may push the magnesium carbonate solubility product above saturation. The precipitation of magnesium carbonate can be kinetically slow, just as the precipitation of calcium carbonate can be slow, and since this cloudiness forms instantly, magnesium hydroxide is a much more likely candidate. However, if the additive is not rapidly mixed in, or worse yet, if solid globs of the initial precipitate settle out and are very slow to dissolve, then conditions may be ripe for magnesium carbonate (and calcium carbonate) to form.

In any case, any magnesium carbonate that does form will probably dissolve later as the pH returns to normal reef aquarium levels, so whether the initial cloudiness contains any magnesium carbonate or not is not a critical issue. It does not contain calcium carbonate if mixed in reasonably quickly (a couple of minutes or less), as CaCO3 would not dissolve when mixed with seawater (and this material is observed to dissolve).

Precipitates on Top of Limewater


The precipitate that forms on the top of limewater is calcium carbonate (Figure 5). Limewater is high in calcium (about 800 ppm at saturation) and is very high in pH (pH 12.54 at saturation), meaning that it contains a lot of hydroxyl ions (OH-). When carbon dioxide from the air encounters the water, it hydrates to form carbonic acid:

(1)   CO2 + H2O     H2CO3

Then, if the pH is above 11, as it is in limewater, the carbonic acid equilibrates to form mostly carbonate:

(2)   H2CO3 + 2OH-     2H2O + CO3--

It is the carbonate that we are concerned with in the formation of insoluble calcium carbonate, both on the surface of, and inside, the limewater:

(3)   Ca++ + CO3--     CaCO3 (solid)

The result of this reaction is visually obvious. The calcium carbonate can be seen as a solid crust on the limewater's surface that has been exposed to the air for a day or two (do not bother to remove this crust; it may actually be protecting the underlying limewater from further penetration by carbon dioxide). The formed solids also settle to the bottom of the container, and can, in fact, form down inside it. Since solid calcium carbonate is not an especially useful calcium or alkalinity supplement, this reaction has the effect of reducing the limewater's potency. With sufficient exposure to air, such as by aeration or vigorous agitation, this reaction can be driven to near completion, with little calcium or hydroxide remaining in solution.

Figure 5. A view of my limewater reservoir showing a thin calcium carbonate crust on part of the liquid surface, and heavy deposits of calcium carbonate, calcium hydroxide, magnesium hydroxide, and other materials on the bottom. I clean out the bottom of the reservoir only once a year or so.
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This reaction is the basis of the claims by many aquarists that limewater must be protected from the air. It is also the basis of the claim that limewater reactors (Nilsen reactors) are to be preferred over delivery from still reservoirs of limewater. Neither of these claims, however, stands up to experimental scrutiny, as I showed in a previous article.

Precipitates on Limewater Drip Tips


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Figure 6. Calcium carbonate that forms on the end of the tubing where limewater drips into my sump.

The same reaction between limewater and carbon dioxide that takes place on the surface of a limewater reservoir (Figure 5) to form calcium carbonate (described above) can also take place on a limewater drip tip (Figure 6). Such tips rapidly become clogged, and need periodic cleaning. This tip can be cleaned by soaking it in acid (straight vinegar or diluted muriatic acid/hydrochloride acid: 1 part acid added to 9 parts fresh water). One can also periodically push a pencil or similar object through the crust to ensure that the opening is large enough for limewater to exit.

I once neglected to clean this tip on my system for too long, and it completely sealed over. Without my knowledge, my Reef Filler pump, though pumping away, was not delivering any top off water. Only when my float switch shut down the main pumps did I realize that there was a problem. Unfortunately, I made a foolish mistake and poked a nail though the crust on the tip without shutting down the pump. These pumps are actually able to develop quite a bit of head pressure, and a fair amount of milky limewater sprayed out over me (fortunately I had glasses on, protecting my eyes)!

Some aquarists have suggested keeping the tip underwater to avoid a reaction with atmospheric carbon dioxide. Unfortunately, that is worse, because there is actually a lot of bicarbonate and carbonate in marine aquarium water, and as this water mixes with, and perhaps diffuses up the limewater tubing slightly, an even more rapid precipitation of calcium carbonate will occur, thus clogging the tip sooner.

Precipitates on the Bottom of Limewater Reservoirs


The solids on the bottom of a limewater reservoir (Figures 5 and 7) contain everything that did not dissolve, or that dissolved and later precipitated from solution. Such solids could contain magnesium hydroxide and carbonate, calcium hydroxide and carbonate and a variety of other impurities such as copper salts, alumina, silica, etc.

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Figure 7. A view of the side of my limewater reservoir after the water level has lowered. The precipitate is most likely calcium carbonate formed when residual limewater on the sides reacts with air, but may also include solids that settled from the initial mixing, and so may resemble the material found on the bottom of the container (Figure 5). I can't recall ever cleaning off the sides of this container.

In order to determine what is in these deposits, I tested a sample of the white solid material that had been collecting for months on the bottom of my limewater reservoir, and detailed the results in a previous article. I removed the white sludge along with some limewater. The mixture of solid and liquid was acidified to dissolve it, and it was tested for calcium, magnesium, and strontium. The results are shown in Table 1. Only relative concentrations are shown as no effort was made to dry the sample prior to analysis, making absolute concentrations meaningless.

Table 1. Alkaline earth metals in limewater sludge.
Metal
Relative Concentration (by weight)
Enrichment Relative to Solid Lime
Magnesium
0.05
13.
Calcium
1.00
1.00
Strontium
0.00019
0.5

As anticipated, based on the very low solubility of magnesium hydroxide and the high concentration of hydroxide ion in limewater, the solid material on the bottom of the limewater is enriched with magnesium. Relative to calcium, magnesium is enriched by a factor of 13 in the sludge compared to the solid starting quicklime. This magnesium may be present as both magnesium hydroxide and magnesium carbonate, but because magnesium carbonate is fairly soluble compared to calcium carbonate, it is most likely that the primary magnesium salt is magnesium hydroxide. It may also be mixed calcium and magnesium carbonates.

Interestingly, strontium is actually depleted by a factor of two relative to solid starting quicklime, indicating that it is less likely than calcium to end up on the bottom of the reservoir. While strontium carbonate is somewhat less soluble than calcium carbonate, the strontium concentration in the limewater is so low that SrCO3 may not actually be saturated, so it may not precipitate at all. The strontium that is there may simply be copreciptiated with calcium carbonate.

The solids on the bottom of a limewater reservoir or the residue left in a limewater reactor can also contain other materials. Phosphate, for example, would be insoluble in limewater, precipitating as calcium phosphate. Many toxic metals, such as copper, are also insoluble in the high pH of limewater, forming carbonates or oxides. These metals can also bind directly to undissolved lime or to calcium carbonate precipitates, as I showed in a previous article. In a sense, this precipitation can purify the limewater so that in some cases it may be even purer than the starting water or lime.

This purification is also seen in practice by many aquarists who have noticed the solids on the bottom of their limewater containers discolor, often to a bluish/green color suggesting copper. For these reasons, I recommend that lime solids not be dosed to aquaria when it is possible to avoid it. Letting the limewater settle for a few hours to overnight will permit most of the large particles to settle out, and whether it looks clear at that point or not, it is likely fine to use. In general, it is a good practice to leave residual solids on the bottom of limewater reservoirs rather than cleaning them out every time, as they may actually help purify the water by these precipitation mechanisms. Once the solids discolor, or have been collecting for 6-12 months, however, they should be discarded.

Precipitates in Limewater Delivery Tubing


Precipitates can also form inside the tubing that delivers limewater to the aquarium (Figure 8). Some of this material will have been formed in the reservoir, and would therefore have been carried into the tubing in particulate form. These solids will be similar in composition to those found on the bottom of the limewater reservoir (described in the previous section). Some of this material, however, may form in situ as carbon dioxide enters the limewater through the walls of the tubing. Different materials have different permeability toward carbon dioxide, and different thicknesses will also alter its diffusion into the tubing. The diffusion can be sufficient, however, that over the course of a year significant amounts of calcium carbonate solids may accumulate in this way, and may eventually clog the tubing. For this reason, I flush acid through the system once a year or so to dissolve this calcium carbonate buildup.

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Figure 8. The settling of precipitated calcium carbonate and possibly other materials in limewater delivery tubing.

Precipitates from Overdosing Limewater


When limewater is substantially overdosed, the transient precipitation of magnesium hydroxide from normal use may not be the only precipitate that forms. If the pH becomes elevated and stays that way long enough, calcium carbonate can precipitate throughout the water column. In such situations, the entire aquarium can become very cloudy, looking almost like skim milk (Figures 9 and 10). Such precipitation events have the beneficial effect of lowering the pH and alkalinity that were raised by the overdose, limiting the ongoing damage that takes place. In many cases, there is no apparent harm after a day or two, but in a few rare cases, when the overdose was especially extensive, a tank crash can occur, killing many organisms.

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Figure 9. A reef aquarium experiencing a limewater overdose, with calcium carbonate particles filling the water column. This is the aquarium of Reef Central members Ereefic and Klasikb.

The following important points should help in dealing with a limewater overdose:

1. Don't panic! These overdoses do not usually cause a tank to crash.

2. The primary concern is pH. If the pH is 8.6 or lower, you need not do anything. If the pH is above 8.6, then reducing the pH is the priority. Direct addition of vinegar or soda water is a good way to accomplish this goal. Either one mL of distilled white vinegar, or six mL of soda water, per gallon of tank water will give an initial pH drop of about 0.3 pH units. Add either to a high flow area that is away from organisms (e.g., a sump).

3. Do not bother to measure calcium or alkalinity while the tank is cloudy. The solid calcium carbonate particles will dissolve in an alkalinity test, and all of the carbonate in them will be counted as if it were in solution and part of "alkalinity." The same may happen to some extent with calcium tests. Wait until the water clears, and at that point, alkalinity is more likely to be low than high. Calcium will likely be mostly unchanged.

4. The particles themselves will typically settle out and disappear from view over a period of 1-4 days. They do not appear to cause long term detrimental effects to tank organisms.

5. Water changes are not necessarily beneficial or needed in response to a limewater overdose.

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Figure 10. A reef aquarium experiencing a limewater overdose (top), with calcium carbonate particles filling the water column. The same aquarium the next day was clear (similar to the bottom picture, but taken a few days later). This is the aquarium of Joe Robertson (Reef Central member Froggy).

Calcium Carbonate Precipitation: Calcium Supersaturation Theory


In seawater, calcium carbonate is supersaturated. In this context, that means that given the right circumstances, it will precipitate as solid calcium carbonate. Of course, under many other circumstances, it doesn't do so, and the answer to the question of why that happens is useful for aquarists to understand.

First, some definitions are in order. The equilibrium constant expression for the dissolution of calcium carbonate is shown below (equation 4):

(4) K = [Ca++]x[CO3--]

Where K is a value determined by multiplying the concentration of calcium by the concentration of carbonate. When K = Ksp* (the solubility product constant in seawater at any given temperature, pressure and salinity, which can be looked up in a book), then the solution is said to be exactly saturated (equation 5).

(5) Ksp* = [Ca++]x[CO3--] (at saturation)

This relationship is often quantified using the supersaturation parameter, which is symbolized as Ω:

(6)    Ω = ([Ca++]x[CO3--] )/ Ksp*

When Ω = 1, the solution is exactly saturated. When Ω exceeds one, it is supersaturated, and when Ω is less than 1, the solution is undersaturated. The higher the supersaturation, the more likely precipitation of CaCO3 will take place.

In normal seawater, Ω ~ 3 for aragonite and Ω ~ 5 for calcite, though these values have been steadily dropping as carbon dioxide has been added to the atmosphere, reducing the seawater's pH. Aragonite and calcite are just different crystalline forms of calcium carbonate. Calcite is slightly more stable, and hence slightly less soluble, than aragonite (i.e., has a lower Ksp*). Organisms can precipitate both aragonite (pteropods and corals) and calcite (foraminifera and coccoliths), but most of the precipitation in reef aquaria is aragonite (although certain organisms such as abalone form both).

Reef aquaria often have higher alkalinity and calcium levels than seawater, and hence are more supersaturated than seawater. Alkalinity is a measure of the bicarbonate and carbonate in solution. At a fixed pH, if the alkalinity is doubled, then the carbonate will also be doubled. Since many aquarists keep reef aquaria at alkalinity levels higher than natural seawater levels (2.5 meq/L; 7 dKH), the supersaturation is often higher than in the ocean.

The biggest driver of supersaturation in a reef aquarium, however, might be pH. In aquaria with a high pH (such as many aquaria using limewater) the supersaturation is much higher than in seawater. At the same alkalinity, if you raise the pH, you convert some of the bicarbonate into carbonate:

(7)     HCO3 -  +  OH-     CO3- -   +   H2O

At pH 8.2 and 25°C, only 15% of the total carbonate and bicarbonate is present as carbonate. At pH 7.8, that value drops to 7%. But as the pH is raised, that portion increases to 50% at pH 8.93 and to 75% at pH 9.4. Consequently, as the pH is raised at a fixed alkalinity, the concentration of carbonate rises, thereby increasing the supersaturation of calcium carbonate. Within the pH range of most reef tanks (up to about pH 9 or so), the amount of carbonate present is approximately linear with the pH because of the relationship seen in equation (7). So if the pH rises from 7.5 to 8.5, there is approximately a ten-fold increase in the carbonate concentration. From pH 8.0 to 8.5, the increase in carbonate is about threefold. Above pH 9, the carbonate concentration continues to rise, but more slowly, and it levels off above about pH 10 as there is very little bicarbonate left at pH 10+ to convert into carbonate.

Calcium Carbonate Precipitation: Calcium, Alkalinity, and pH


Combining the various factors described above, here are some combinations of calcium, alkalinity and pH that have equal supersaturation (that is, equal propensity to cause calcium carbonate precipitation):

Ω = 1 (dissolution of aragonite takes place at all lower values of these parameters)

pH = 7.7
pH = 8.2
Calcium = 410 ppm
Calcium = 340 ppm
Alkalinity = 2.5 meq/L
Alkalinity = 1.0 meq/L

Ω = 3 (typical of normal seawater)

pH = 8.2
pH = 8.0
pH = 8.4
Calcium = 410 ppm
Calcium = 410 ppm
Calcium = 260 ppm
Alkalinity = 2.5 meq/L
Alkalinity = 4.0 meq/L
Alkalinity = 2.5 meq/L

Ω = 6 (non-biological precipitation is more likely)
pH = 8.2
pH = 8.2
Calcium = 410 ppm
Calcium = 820 ppm
Alkalinity = 5.0 meq/L
Alkalinity = 2.5 meq/L
 
 
pH = 8.0
pH = 8.7
Calcium = 410 ppm
Calcium = 410 ppm
Alkalinity = 8.0 meq/L
Alkalinity = 2.5 meq/L
 
 
pH = 8.45
Calcium = 410 ppm
Alkalinity = 4.2 meq/L

How should we think about supersaturation? The higher it is, the more likely it is that calcium carbonate will precipitate. The reason for this is straightforward: if the "pressure" to precipitate calcium carbonate becomes too high, certain inhibiting processes (described below) will be overwhelmed, and precipitation will take place.

If Ω is not too high, some precipitation will take place before the inhibiting mechanisms take control of the crystals' surface and prevent further precipitation. This is the process that happens in normal seawater. If Ω is too high, a bigger precipitation event can take place before being halted. In the worst cases, this can lead to a snowstorm of calcium carbonate particulates throughout the tank. Such snowstorms can occur, for example, when too much limewater is added to the tank. In that case, the pH rises and converts much of the bicarbonate to carbonate. Ω is then driven to unstable levels, and a massive precipitation event takes place.

Calcium Carbonate Precipitation: Low Magnesium


Calcium carbonate is the most common precipitate that forms in reef aquaria, and often forms on pumps and heaters. As described above, calcium carbonate is supersaturated in seawater and in nearly all reef aquarium systems. In the ocean, calcium carbonate is supersaturated by a factor of about three to five.

So why doesn't it immediately precipitate? There are several reasons that calcium carbonate does not, under normal conditions, rapidly precipitate, turning the water cloudy with a white precipitate. The biggest factor is magnesium. Once a tiny crystal of calcium carbonate begins to form, magnesium pushes its way onto the calcium carbonate surface, and occupies a site that would normally have been occupied by calcium. These magnesium ions poison the surface against more accumulation of calcium carbonate as it no longer looks like a crystal of calcium carbonate to incoming calcium and carbonate ions, and precipitation stops.

Magnesium is present in natural seawater at about 1280 ppm. If magnesium is well below natural levels, then its effect on calcium carbonate precipitation may be ineffective, allowing precipitation to take place more rapidly than otherwise. Aquarists experiencing such precipitation might consider testing for magnesium, and supplementing if its level is not at least 1250 ppm. I usually recommend targeting magnesium at about 1250-1350 ppm. Exactly how low it needs to be before it begins to allow calcium carbonate precipitation is not known exactly. However, the relative impact of low magnesium in this context depends on the relative pressure for calcium carbonate precipitation to take place (which is shown in other sections of this article to be determined primarily by the calcium concentration, alkalinity, temperature and pH).

Calcium Carbonate Precipitation: Elevated Temperature


For the reasons described above, many aquarists tend to blame low magnesium levels for excessive precipitation of calcium carbonate. In fact, while it is worth testing when such precipitation is significant, low magnesium is not the only way that such precipitation takes place.

It turns out that temperature also plays a significant role in precipitation of calcium carbonate. This effect is noticeable not so much via the aquarium's water temperature, although that may have some impact, but most apparently, on unusually warm objects in the aquarium, such as pumps and heaters. There are two primary ways that warmer temperatures drive precipitation of calcium carbonate. The first is simply by reducing the calcium carbonate's solubility.

At S=35 (that is, normal seawater salinity) and 1 atmosphere pressure, the Ksp* decreases slightly as the temperature rises. Millero has published a series of long equations for calculating Ksp* for both aragonite and calcite.1 For aragonite, the log Ksp* drops from -6.19 at 25°C, to -6.23 at 40°C to -6.44 at 80°C. In relative terms, the Ksp* has fallen from 1 to 0.91 to 0.55 over this temperature range. Likewise for calcite, the relative Ksp* has changed from 1 to 0.96 to 0.73 over this range.

Consequently, if a reef aquarium has a supersaturation of about 3 for aragonite and 5 for calcite at 25°C (typical for seawater), then at 40°C the supersaturation has increased to about 3.3 and 5.2, respectively. At 80°C this supersaturation has increased to 5.4 and 6.8, respectively. Since the supersaturation has increased, the likelihood of precipitation has increased, and this increase is part of the explanation of why precipitation takes place on the surface of heaters, where temperatures can greatly exceed the bulk water temperature.

A second factor is the acidity of bicarbonate. Bicarbonate's acidity increases with temperature, so the equilibrium for equation 7 is pushed to the right as the temperature rises.

(7)     HCO3 -  +  OH-     CO3- -   +   H2O

In a previous article I worked through a variety of calculations based on literary values for the acidity of HCO3- as a function of temperature in seawater, and showed that as temperature rose from 25°C to 40°C, the relative concentration of carbonate increased by a factor of 1.45 as equation 7 is pushed to the right. We also find, of course, that the supersaturation of calcium carbonate (equation 6) has increased by a factor of 1.45 over this temperature range. Running the same calculations for 80°C, we see that carbonate concentration increases by a factor of 2.4x compared to 25°C.

Combining these two ways that temperature impacts supersaturation between 25 and 40°C for aragonite, we find a change from Ω = 3.0 to 3.3 due to solubility, and from 3.0 to 4.4 due to the bicarbonate pKa shift. Together these effects yield a supersaturation of 4.8 for aragonite.

Combining these two ways that temperature impacts supersaturation between 25 and 80°C, we find a change from Ω = 3.0 to 5.4 due to solubility, and from 3.0 to 7.2 due to the bicarbonate pKa shift. Together these effects yield a supersaturation of 13 for aragonite.

How do we put these supersaturation changes into perspective? The following combinations of calcium and alkalinity have the same supersaturation in seawater:

1. Normal seawater at 80°C

2. Seawater at 25°C with calcium raised to 1300 ppm

3. Seawater at 25°C with the alkalinity raised to 8.2 meq/L

It makes sense that all three of the situations above could lead to precipitation of calcium carbonate, and that is exactly what happens on the surfaces of hot objects in reef aquaria.

Magnesium Hydroxide


Magnesium hydroxide, Mg(OH)2, can form in a variety of different ways in reef aquaria, but it most often forms when the aquarium water's pH rises. At higher pH there are more hydroxide ions (OH-) , with about 10 times more OH- for each 1 pH unit increase. The Mg(OH)2 that forms is quite insoluble, and it comes out of solution as an amorphous white powder:

(8)  Mg++  +  2OH-     Mg(OH)2 (precipitate)

In fact, this reaction is a well-known industrial process for forming magnesium hydroxide that is later heated, driving off water, to form magnesium oxide that has many industrial uses. Most nonchemists know magnesium hydroxide as a suspension of particles in water. Yes, it is Milk of Magnesia, sold at the drugstore as a laxative. In that application, the solid particles partially or completely dissolve, and the released magnesium ions are poorly absorbed from the human intestines. To maintain osmotic balance, water then enters the intestines, giving a laxative effect as the stool is softened.

At normal reef aquarium pH, there is not enough hydroxide to cause magnesium hydroxide to precipitate. At what pH does it begin to precipitate? In a previous article on metals in limewater, I showed how to calculate the solubility of a variety of metals as a function of pH in FRESH WATER. Figure 11 shows such data, and focusing on magnesium, it can be seen that with a magnesium concentration of 1280 ppm (the concentration in seawater), magnesium becomes insoluble as magnesium hydroxide at pH values above about pH 9.2 (that is, the solubility above pH 9.2 is below 1280 ppm, or below 1.3 x 103 ppm).

Click here for larger image
Figure 11. The solubility of magnesium in FRESH water as a function of pH. The solubility of magnesium drops below 1280 ppm (natural seawater levels) at pH values above 9.2 (red line in graph), indicating that magnesium hydroxide precipitates above this pH. The much higher solubility of calcium/calcium hydroxide is shown for comparison.

But is that also true in seawater? We can answer the question both theoretically and experimentally. First, the theory. The solubility of magnesium hydroxide is governed by a standard solubility product equation:

(9)  K = [Mg++]x[OH-]x[OH-]

where K is called the solubility product, Mg++ stands for the concentration of free magnesium ions, and OH- is the concentration of free hydroxyl ions. So it is readily seen that as either the concentration of magnesium or the pH rises, K increases.

The specific value for K, called the Ksp or the solubility product constant, can be found in a book. When the aquarium's K value exactly equals the Ksp, the water is exactly saturated with magnesium hydroxide. When K exceeds the Ksp, magnesium hydroxide, it is likely to precipitate from solution. So, we can calculate when K exceeds Ksp in seawater since we know the pH (from a measurement), the magnesium concentration (from a measurement or known values for natural seawater) and the Ksp (from a book).

It turns out that such a calculation is a little more complicated than that, since not all of the magnesium ion present in seawater is "free" magnesium. Rather, some is tied up interacting with other ions. In seawater, it is known that magnesium concentrations need to be adjusted downward in such calculations by a factor of about 0.255 (with 0.255 being called the activity coefficient). Consequently, a normal magnesium concentration (1280 ppm = 0.053 M) needs to be corrected to 0.0135 M.

Hydroxide ion is rarely measured directly, but is calculated from the pH. We can calculate it from the known relationship:

(10)  [OH-] = Kw*/[H+]

Where Kw* is a constant (for seawater at a given temperature) that can be looked up in a book, and

(11)  [H+] = 10-pH

Like magnesium, the hydroxide and H+ ion concentrations must also be corrected with an activity coefficient in seawater, and those values are known to be about 0.236 and 0.688 respectively (in practice, however, the value coming out of equation 11 for H+ is already corrected by the activity coefficient correction for H+, so we need not do that correction twice). So knowing the magnesium concentration, the constants Ksp (= 1.2 x 10-17 when the concentrations are in molar units) and Kw* (= 6.45 x 10-14 in molar units at 25°C) and the activity coefficients, we can determine K as a function of pH (Table 2). As can be seen, K exceeds Ksp when the pH is above about pH 9.4 (which is not especially different from fresh water).

Table 2. K for magnesium hydroxide as a function of pH in seawater at 25°C.
pH
K
8.2
7.7 x 10-14
8.5
3.0 x 10-13
8.8
1.2 x 10-12
9.1
4.8 x 10-12
9.4
1.9 x 10-11
9.7
7.7 x 10-11
10.0
3.0 x 10-10
10.3
1.2 x 10-9
10.6
4.8 x 10-9
10.9
1.9 x 10-8

Experimental Measurements of Magnesium Hydroxide Formation


In order to substantiate the theoretical calculations described above for magnesium hydroxide, I also tested the precipitation of magnesium hydroxide at elevated pH. Starting with fresh Instant Ocean artificial seawater, I added dropwise a concentrated solution of sodium hydroxide in fresh water. As each drop entered, a cloudy/white glob formed that settled to the bottom. At first, each drop could be fully dissolved with sufficient swirling. By the time the pH of the solution reached the upper 9's, the glob was only very slowly dissolving. By pH 10.4, the solution was very cloudy, even after extensive swirling.

Taking this cloudy solution at pH 10.4, I added hydrochloric acid to reduce the pH to 10.0. All of the cloudiness disappeared within a few minutes. When set aside for 90 minutes, however, a solid precipitate formed. Upon reducing the pH to 8.2 with hydrochloric acid, this precipitate remained. It remained undissolved for many days.

My conclusion is that the initial cloudiness is magnesium hydroxide that rapidly forms and dissolves (not magnesium carbonate, as described below). It appears somewhat more soluble than the theoretical calculations above would suggest. That is, it does not precipitate out until the pH is above 10, and seems fully soluble at pH 10, while the calculations suggest that the initial point of precipitation should be lower (mid 9's). This result appears to be because the magnesium hydroxide that precipitates first is not the most stable form, and it is the most stable form (the crystalline form brucite) that is typically reported in Ksp values. The initial precipitate is faster to form, but is more soluble.

The notion that there are different forms of magnesium hydroxide that can form is supported in the chemical literature. One group has reported, for example, that the initial precipitate that forms when raising seawater's pH is not the most stable form, that only more slowly is made as the initial precipitate slowly converts into a less soluble crystal. They note: "In the case of Mg(OH)2, a considerable difference between the solubility of freshly precipitated hydroxide and that of aged hydroxide was revealed."2 A second group reported similar findings: "Precipitated Mg(OH)2 is at first more soluble than brucite, but in time its solubility decreases to that of brucite."3

When sitting at high pH for more than a few minutes, calcium carbonate can precipitate from these high pH solutions, as was observed. Since seawater is supersaturated with respect to calcium carbonate, even at pH 8.2, this material does not dissolve on reducing the pH to 8.2, even after many days of exposure.

Magnesium Carbonate Solubility


According to Millero in "Chemical Oceanography,"1 natural seawater is undersaturated in magnesium carbonate1 (pure magnesium carbonate, not mixed calcium/magnesium carbonates) by a factor of 27. When the pH of seawater is raised with hydroxide ion (as by the addition of limewater or sodium hydroxide solution), the only thing that could lead to magnesium carbonate precipitation is the conversion of bicarbonate to carbonate, raising the solubility product:

(7)     HCO3 -  +  OH-     CO3- -   +   H2O

(12)    K = [Mg++]x[CO3--]

This process, however, is limited by how much bicarbonate can be converted into carbonate. At pH 8.2 in seawater, 15% of the total carbonate plus bicarbonate is present as carbonate. Even if all of the remaining bicarbonate were converted into carbonate, the carbonate concentration would rise by only 5.7-fold. Consequently, the increase in pH is not enough to cause magnesium carbonate to reach saturation and permit solid magnesium carbonate precipitation. This fact is consistent with the commercial process of forming magnesium hydroxide (and not magnesium carbonate) by adding hydroxide ion to seawater.

Even artificial seawater with elevated carbonate alkalinity does not contain enough carbonate to exceed saturation. For example, an alkalinity of 5 meq/L (14 dKH) at pH 8.2 has double the total bicarbonate and carbonate compared to natural seawater. This doubling and the 5.7-fold increase from converting all the bicarbonate to carbonate (2 x 5.7 = 11.4) still does not attain the 27-fold increase in K that would be required to reach the point of precipitate formation.

Consequently, I conclude that magnesium carbonate does not form when hydroxide-containing additives (such as limewater) are added to reef aquaria.

Can it form when carbonate salts are added? For example, when sodium carbonate solutions are added as part of a two-part additive system? Theoretically, yes. However, calcium carbonate is kinetically slow to precipitate even when it can do so (a well established fact that is experimentally reproduced in the preceding section), and it is anticipated that the kinetic issues that slow calcium carbonate precipitation (the need to fit ions carefully into a crystalline structure) are similar for magnesium carbonate (but not magnesium hydroxide, which is amorphous) and will slow calcium carbonate precipitation as well.

So I think it likely that magnesium carbonate may form if carbonate salts are maintained at high concentration for extended periods, but not likely when carbonate-containing additives are well mixed into a reef aquarium within a couple of minutes of their addition.

Summary


Many precipitates are formed in the various processes that aquarists use to maintain reef aquaria. Some of these may be beneficial, such as the solid material that settles out of limewater, reducing the load of impurities it delivers to the aquarium. Others are neither beneficial nor detrimental, such as the initial cloudiness that forms when high pH additives are added to reef aquaria. Others are downright detrimental, such as solid calcium carbonate that can clog pumps. The information provided in this article should help aquarists understand what these are and why they form. Using that information, aquarists may be able to better maintain their reef aquaria without excessive worry over insignificant issues.

Happy Reefing!



If you have any questions about this article, please visit my author forum on Reef Central.

References:

1. Chemical Oceanography, Second Edition. Millero, Frank J.; Editor. USA. (1996), 496 pp. Publisher: (CRC, Boca Raton, Fla.).

2. Solubility of the constituents of seawater under the conditions of the PWR secondary circuit. Lambert, Irma; Lefevre, Antoinette; Montel, J. Cent. Etudes Nucl. Saclay, CEA, Gif-sur-Yvette, Fr. Water Chem. Corros. Probl. Nucl. Power Plants, Proc. Int. Symp. (1983), Meeting Date 1982, 349-63.

3. The solubility of natural and artificial magnesium hydroxides. Quartaroli, Alfredo; Belfiori, Ofelia. Annali di Chimica Applicata (1941), 31 56-61.




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What is that Precipitate in My Reef Aquarium? by Randy Holmes-Farley - Reefkeeping.com