[saltjunkie]

What is the major use, fdactor , reason, and anything else about mag. What is its relation ship to salinity, calcium, alk. etc...
Need to know as much as possible as soon as I can.
Thanks for any help.
Rick

[Elmo18]

Ok, I'm no expert here, but I did find several great Bookmarkable articles.

This one's from the Aquarium Frontiers Library:

Magnesium is an interesting atom that has tremendous biological and chemical relevance to reef tanks. Fortunately for reefkeepers, it is present in abundance in seawater and is depleted only slowly. Consequently, maintenance of magnesium levels is not typically a big issue if using an appropriate salt mix. Nevertheless, magnesium is a very important ion and engenders much discussion among hobbyists. In this article I'll try to add to the extensive writings that Craig Bingman has published in the past.

First, a little background on magnesium. In seawater, magnesium is invariably present in the form of a divalent cation, Mg++. It is present in seawater at a concentration of about 1300 ppm, and that concentration does not vary appreciably with depth or location in the world (besides estuaries and other places where all ions are distorted). In seawater, Mg ions outnumber Ca ions by a factor of 5. Most magnesium in seawater is present as the free ion hydrated with tightly bound water molecules. Some of it, however, forms tight ion pairs (i.e., soluble complexes) with negatively charged ions, such as sulfate, bicarbonate, carbonate, borate, fluoride and hydroxide.

Interestingly, the average residence time for a magnesium ion in seawater is tens of millions of years, substantially longer than calcium (a few million years) and aluminum (100 years), but less than sodium (about 250 million years). In a sense, this is an indication of how reactive magnesium is: it stays in seawater a long time because it's fairly unreactive, but it does get taken out of solution more readily than does sodium.

Another interesting characteristic of ions is whether they are excluded from organisms, actively taken up or just “allowed” to be present. Like sodium and sulfate, the relative concentration of magnesium in organisms is approximately the same as in seawater. More than anything else, this probably results from the fact that there is plenty of magnesium present, and that it is used by organisms for many purposes. Chloride, on the other hand, is actively rejected by organisms, and most other ions are substantially concentrated.

Why do we talk so much about magnesium?

So what is it about magnesium that gets reefkeepers talking about it so much? Is it the fact that it is important for many enzymes? No. Is it the fact that it's important for the skeletons of many organisms in reef tanks? No. It's because reefkeepers have become infatuated with measuring and “optimizing” the levels of calcium and alkalinity in their tanks, and the fact that magnesium can play a role in these levels.

What does magnesium have to do with calcium and alkalinity? Good question. In order to answer that, one has to have a basic understanding of the calcium and carbonate systems in seawater. This system is detailed in many reef-oriented publications, so I won't go into it in great detail. Suffice to say that calcium carbonate is supersaturated in seawater, meaning that given enough time calcium ions will interact with carbonate ions and precipitate as calcium carbonate. If you run the concentration of either too high, CaCO3 will start to precipitate. Magnesium interferes with this process, permitting both calcium and carbonate to be elevated above where they would be in the absence of magnesium.

How does magnesium interfere with precipitation of CaCO3? Two ways have been suggested, and these are detailed below. The first involves magnesium poisoning the surface of growing CaCO3 crystals, slowing the precipitation. The second involves the interaction between soluble magnesium ions and soluble carbonate ions, forming an ion pair and effectively lowering the free concentration of carbonate that is available to precipitate with calcium.

Measurement of the Impact of Magnesium on the Calcium/Carbonate System

In Stephen Spotte's book Captive Seawater Fishes there is an extensive discussion of the impact of magnesium on the calcium/carbonate system. Buried in that discussion is a set of data that indicates the magnitude of the impact that magnesium can have. In this experiment, batches of artificial seawater were made up with varying magnesium and carbonate levels. The scientists then measured how long it took for calcium carbonate to precipitate from each solution. Not surprisingly, the higher the carbonate was raised, the more rapid was the precipitation of calcium carbonate.

More interestingly, the magnesium levels were found to have a very large impact on the rate of precipitation. In batches with no magnesium, and at natural calcium and elevated carbonate levels, calcium carbonate was found to precipitate in minutes. With a natural seawater level of magnesium added to that mix, the precipitation was delayed to 13 to 20 hours. With double the natural magnesium concentration, the precipitation was delayed to 22 to 29 hours.

Even more strikingly, at a lower level of carbonate (closer to that of natural seawater and probably similar to that in many reef tanks), precipitation was delayed from a few minutes in the absence of magnesium to 750 hours in the presence of natural levels of magnesium. Consequently, we conclude that magnesium has a big impact on the rate of precipitation of calcium carbonate (a fact that has been confirmed by many researchers).

Poisoning of Growing CaCO3 Surfaces

In Captive Seawater Fishes, poisoning of growing CaCO3 surfaces is presented as the only explanation involving magnesium for the delay in precipitation of CaCO3, and there is an extensive discussion there about how this takes place. Like many real life problems, this one is not easy to fully explain with simple chemistry, but I'll try! To get an idea of the complications involved, however, Spotte says “The study of carbonate minerals involves nuances of solubility that pose some of the most difficult problems in chemical oceanography and geochemistry.”

In short, while magnesium carbonate is not supersaturated in seawater (or in reef tanks), and will not precipitate on its own, magnesium is attracted to calcium carbonate surfaces where the carbonate ions are already held in place by the calcium ions. With the carbonate ions held in place, magnesium finds this an attractive place to precipitate. A similar effect happens for phosphate and many organics, where they precipitate onto the calcium ions that are held in place by the carbonate ions — but that's another story for another day.

When calcium carbonate (as calcite or aragonite) is placed into seawater, calcium carbonate rapidly begins to precipitate onto the fresh surface. After a short time, however, a thin coating of Mg/CaCO3 (magnesian calcite) begins to form as magnesium pushes its way into the growing surface. Eventually, the surface contains a substantial amount of magnesium. The extent to which this happens depends on the underlying mineral, and is apparently much more extensive on calcite than aragonite (perhaps another reason that aragonite dissolves more readily than calcite). It also depends upon the relative amounts of calcium and magnesium in the water. Regardless, a new type of material is formed that contains both calcium and magnesium.

This new mineral surface containing both calcium and magnesium is not a good nucleating site for precipitation of additional calcium carbonate (as aragonite or calcite), and precipitation of additional CaCO3 slows down substantially. Consequently, seawater and reef tanks are readily supersaturated with respect to calcium carbonate.

Ion Pairing of Magnesium to Carbonate

A second factor in the impact of magnesium on calcium and alkalinity levels involves the ion pairing of magnesium to carbonate in solution. The magnitude of this effect is likely smaller than the surface poisoning effect, but it may still be significant to the calcium and alkalinity levels attainable in reef tanks. To understand how this works, let's start back at the equation governing the equilibrium solubility of calcium carbonate in water. From the solubility product we have

Ksp = [Ca++] x [CO3--]

(This means the concentration of the calcium times the concentration of the carbonate equals a constant, Ksp, the solubility product constant. This is easily thought of as follows: it takes one Ca++ and one CO3-- to come together to form CaCO3, so as the concentration of either rises you are more likely to get them to come together, and so when you reach a certain amount of both, precipitation will begin).

But that is the freshman chemistry simplification. It really should read:

Ksp = ?c x [Ca++] x ?o x [CO3--]

where ?c and ?o are the activity coefficients of the calcium and carbonate respectively. These factors take into account the fact that some of the calcium and carbonate may be tied up or otherwise prevented from easily interacting with each other. In seawater, there are many things that contribute to the activity coefficient, but magnesium is one, and Millero quantifies this effect in his book Chemical Oceanography. In seawater, more than half of the carbonate ions present at any given point in time are ion-paired to magnesium, and this substantially reduces the free concentration of carbonate ions available to precipitate with calcium.

Before suggesting the magnitude of this effect, let me turn the solubility equation around a bit.

Saturation state = (?c x [Ca++] x ?o x [CO3--])/Ksp

The saturation state is a measure of how much calcium and carbonate are in solution. If the saturation state = 1, then the solution is exactly saturated. If it is above 1, the solution is supersaturated. If it is below 1, the solution is undersaturated. Natural seawater has a saturation state of 3 to 5, so it is substantially supersaturated. The higher the supersaturation, the more "pressure" there is for CaCO3 to precipitate from solution.

The interesting data presented by Millero is as follows:

If you put CO3-- into pure sodium chloride solution at the salinity of natural seawater, you get a certain activity coefficient, ?o = 0.164. Alternatively, if you put it into a mixture of sodium chloride and magnesium sulfate at natural seawater salinity and in the appropriate ratios for natural seawater, you get a lower activity coefficient, ?o = 0.134. The magnesium has lowered the activity coefficient substantially (being an anion, the sulfate would have a small impact on the activity coefficient of carbonate, also an anion).

What does this mean? Well, if we look at the saturation state equation using these values, we find that for a given concentration of calcium and carbonate, the supersaturation is lower when magnesium is present than when it is not.

Alternatively, for a given saturation state, the concentration of calcium and carbonate will be higher in the presence of magnesium. How big is the effect?

Let's say we have a certain saturation state with 400 ppm calcium in the presence of magnesium. With everything else unchanged, removing all of the magnesium (and replacing it with sodium) will result in the same saturation state at a calcium level of 327 ppm. Thus, if reef tank seems to be maxing out Ca++ at 400 ppm, then without any Mg++ it might max out at 327 ppm (through this effect alone and assuming that the calcium and alkalinity is being limited by precipitation, not by the amount being supplemented to the tank).

Note, however, that this isn't the perfect answer. It is examining the difference between having and not having Mg++ in pure sodium chloride solution, not in natural seawater. The answer in the latter case may be different. Nevertheless, it gives an idea of what magnitude of effect magnesium might be expected to have through mechanisms involving reduction of the activity of the carbonate ions in solution.

Conclusion

I hope this article provides reefkeepers with an understanding of how magnesium is playing a role in the calcium/carbonate system in their tanks. As is usually the case, these kinds of discussions do not tell us what levels of magnesium are optimal or how far above or below natural seawater one can safely keep a reef tank. Nevertheless, understanding what is happening in our tanks can help lead us to solutions to vexing problems, as well as to provide us with an appreciation for the great complexities that nature provides in our tanks.


This next one is by Dr. Craig Bingham, PhD.

This month I will discuss possible ways in which magnesium could be depleted from reef aquaria and the dangers of excessively high or low magnesium concentrations.

Depletion phenomena, real and imagined
Let’s start with the imagined depletion phenomena. These stories often start with someone buying a magnesium test kit, measuring the concentration in his or her aquarium for the first time, and discovering that it is lower than expected. The expected value for 35 parts per thousand salinity seawater is 52.82 millimoles per kilogram or 1284 parts per million (ppm;mass/mass).

Low salinity
The first “mysterious” cause of lower than expected magnesium concentrations in a reef aquarium is low overall salinity. Although I don’t advise running a reef system at lower than 35 ppt salinity, it isn’t uncommon for people to recommend running saltwater tanks at 30 ppt salinity or even less. If you use a synthetic seawater mix that perfectly emulates the ionic ratios in natural seawater, and only use enough of it to attain 30 ppt salinity, the magnesium concentration in your system will be only 0.0453 moles per kilogram or 1100 parts per million (ppm; w/w.) That is significantly lower than you might be expecting. The easiest way to correct this is simply to increase the salinity of the system to 35 ppt salinity.
Obviously, this doesn’t just affect magnesium. If your system is operating at lower than 35 ppt salinity, then it is likely that the concentrations of all major and minor ions is systematically low in your system. Aquarists should also be aware that all of the synthetic seawater mixes on the market have substantial water or hydration in them. In many cases, a “50-gallon bag” of synthetic seawater mix will make only 42 gallons of 35 ppt salinity synthetic seawater. (The mean salinity of the eight brands of synthetic seawater mixed to 35 grams per kilogram was 29.68 ppt, 50 x 29.68/35 = 42.4 gallons at 35 ppt salinity.)


Variations in magnesium content of synthetic seawater mixes
A second leading cause of mysteriously low magnesium concentrations in reef tanks is batch-to-batch or systematic variations in the amount of magnesium present in a given brand of synthetic seawater. The assay of several brands of synthetic seawater presented last month ( ”The Composition Of Several Synthetic Seawater Mixes”) doesn't discriminate between these two sources of variability, but it is clear that some brands of salt can deviate substantially from the ionic ratios found in natural seawater. So it is possible that an anomalous magnesium concentration in the aquarium might be due to the salt you are using.
Looking through the plot of the magnesium concentrations observed in that study, most salts actually do a reasonably good job of providing enough magnesium. There were really only two outliers, one high and one low, of the eight brands of salt assayed. If you are using those two brands of salt, you might want to pay particular attention to the magnesium concentration both in your tank and in fresh batches of water you are adding to your system. These are the two leading causes of imagined depletion of magnesium in reef aquaria. Now let’s examine some effects that could cause real loss of soluble magnesium from the system.


Real depletion phenomena
Having examined cases where the magnesium concentration was wrong in the first place, we should examine mechanisms by which magnesium might be selectively depleted from the water in the tank. There are several possibilities: incorporation into living tissues and precipitation as magnesian calcites as skeletons or tests of marine organisms, and hydrochemical formation of magnesium-containing precipitates.
There are a number of types of calcium carbonate. There are three crystal forms: calcite, aragonite and vaterite. Of those, the calcite and aragonite lattices are commonly found in nature. In many cases, the calcites and aragonites formed are relatively “dirty” because other ions are co-crystallized in the skeletal matrix. Some organisms, most notably some types of coralline algae, secrete a high magnesian calcite of up to 30 mole percent magnesium content (Morse and Mackenzie 1990).

Octocorals have internal skeletal components, called sclerites, that are composed of magnesian calcite. Weinbauer and Velimirov (1995) have shown that the sclerites of gorgonians have 6.4 to 9.8 mole percent magnesium. While calcites produced by coralline algae and octocorals are probably the major biogenic sinks for magnesium in the reef aquarium, other minor organisms also secrete calcitic skeletons. For example, echinoderms produce magnesian calcite spikes and skeletal plates (e.g., see Tsipursky and Brusek 1993). Agegian and Mackenzie (1989) found that the magnesium content of the calcite produced by a tropical benthic community off Hawaii was 6 to 16 mole percent magnesium. Calcified arthropod cuticle also contains magnesian calcite (e.g., see Giraud-Guille and Quintana 1982).

The topic of the precipitation of magnesium from reef aquaria as magnesium hydroxide has been covered in the past (Bingman 1997) The major conclusion of that column was that gradual limewater additions are unlikely to cause the stable formation of magnesium hydroxide in an aquarium because brucite (the name of magnesium hydroxide found in mineral form) and hydromagnesite (which contains magnesium ion, carbonate ions, hydroxide ions and waters of hydration) form only at carbon dioxide concentrations, orders of magnitude lower than found in reef aquaria, except at the point where limewater meets aquarium water. We expect that any formation of these minerals will be minimized by adding limewater at a very turbulent point in the aquarium.

Recently, Morse et al. (1997) have shown that at 25 degrees Celsius (77 degrees Fahrenheit), formation of calcite precipitates from seawater requires that the magnesium to calcium ratio be less than 1:4 (on a molar basis). Robbins et al. (1996) have recently presented data consistent with the famous carbonate whitings on the Great Bahama Bank being driven by photosynthesis from picoplankton. These whitings are mainly aragonite with small quantities of magnesian calcite. This is consistent with my unpublished observations that the hydrochemical (abiogenic) carbonates that form around the point of limewater addition have a low magnesium content. One can tenatively conclude that hydrochemical precipitation of mangesium in a reef system is probably less important than the formation of biogenic magnesian calcites by coralline algae and various invertebrates.


Dangers of excessively high or low magnesium
Some aquarists are apparently adding magnesium supplements to their aquariums without performing any testing for magnesium levels. This is a potentially dangerous practice. I can point to no evidence that higher than natural seawater magnesium concentration has a beneficial effect on marine organisms. However, there is evidence that higher than natural seawater concentrations of magnesium can harm marine organisms.
Magnesium salts have been used as a “chemical shucking aid” on oysters and other shellfish for many years (Whyte and Carswell 1983). Magnesium chloride is also used as a relaxant or anaesthetic in pearl oyster culture (Norton et al. 1996). Namba et al. (1995) report persistent relaxation of the adductor muscles of the oyster, Crassostrea gigas, when exposed to solutions of magnesium chloride. The degree of relaxation was dose and time dependent. Higher concentrations of magnesium chloride and longer periods of exposure increased the number of individuals showing persistent relaxation. Significant relaxation was observed in groups of animals exposed to concentrations of magnesium only six times higher than seawater for periods as brief as 30 minutes.

Culloty and Mulcahy (1990) indicate that magnesium salts make an inexpensive anaesthetic agent for Ostrea edulis. In studies of anaesthetic agents for the scallop, Pecten fumatus, Heasman et al. (1995) found that magnesium chloride at 0.31 molar was a good anaesthetic. However, magnesium sulfate at the same dose caused excessive post-anaethesia mortality, a finding that may be interesting and relevant to individuals who use magnesium sulfate heptahydrate (Epsom salts) as a magnesium additive.

It is also known that magnesium concentrations significantly lower than natural seawater can negatively influence the development of many marine organisms. For example, Lee and Krishnan (1986) have shown that synthetic culture media lacking magnesium or with half seawater concentrations of magnesium impair the development of larval dolphin fish (Coryphaena hippurus), although it is not universally required.

The same author also demonstrated that magnesium was not especially important in the development of gray mullet (Lee and Hu 1983). The juxtaposition of these two studies highlights the importance of making conservative judgements about the importance of major and minor ions in seawater.

Relatively few studies have investigated the effects of much higher and lower than natural seawater concentrations of magnesium. Based on the relatively few scientific studies where this has been examined, it seems prudent to maintain our aquaria at as close to natural seawater magnesium concentrations as feasible. Next month I’ll cover expected depletion rates, using the simulation program ION, and describe a “homebrew” magnesium supplement.

Also by Dr. Bingham, this one discusses more chemistry issues.

INTERACTIVE
Discuss issues with aquarists around the world in our Interactive Aquarist Forum.
It has been stated that dosing limewater causes the loss of significant amounts of magnesium ions from reef aquaria. The supposed mechanism for this loss is through the formation of magnesium hydroxides. In my view, it is extremely unlikely that limewater causes the significant loss of magnesium from reef aquaria. This view is based on well-established data from the formation and transformation (diagenesis) of recent carbonate sediments in the ocean, on a failure to observe significant magnesium ion depletion in my aquaria, and on other recently published research from Aquarium Systems in SeaScope (1996, 1997).

FIGURE 1

Before addressing this issue directly, it is important to introduce the vocabulary of this field, as well as the various carbonate minerals found in marine sediments and the reactions through which magnesium loss is supposed to occur. Many of you are familiar with the reactions through which limewater maintains calcium and alkalinity in reef aquaria. Limewater is a solution or suspension of calcium hydroxide in water. Calcium oxide can also be used to make limewater because it is immediately hydrated to calcium hydroxide in a highly exothermic reaction (generating much heat).


CaO + H2O ---> Ca(OH)2
(calcium oxide) (calcium hydroxide)


This reaction is also called "slaking." Calcium oxide is also called "quick lime" because of the great heat (and often steam) generated by this reaction. Although the hydration of a lot of calcium oxide with a little water can quickly boil much of the added water, aquarists usually add small quantities of calcium oxide to a lot of water and might not even notice the strongly exothermic nature of this reaction.


When calcium hydroxide is slowly dripped into the aquarium, it captures carbon dioxide and converts it to bicarbonate ions.


Ca++ + 2 OH- + 2CO2 ---> Ca++ + 2(HCO3-)


If insufficient carbon dioxide is present in the aquarium, bicarbonate ions will be converted to carbonate ions.


Ca++ + 2 OH- + 2(HCO3-) ---> Ca++ + 2(CO3--) + 2 H2O


The carbonate ions formed may cause the precipitation of calcium carbonate. Rapid addition of limewater may cause the loss of alkalinity from the aquarium through the following reaction, in which I take a pre-existing calcium ion and two bicarbonate ions from the aquarium, add calcium hydroxide, and form solid calcium carbonate.


Ca++ + 2(HCO3-) + Ca++ + 2(OH-) ---> 2 CaCO3 + 2 H2O
(from aquarium) (from limewater)


Therefore, it is possible to reduce the carbonate alkalinity and the calcium concentration of the aquarium by too rapid dosing of calcium hydroxide into the tank. In practice, it is difficult to distinguish the loss of calcium and alkalinity by this mechanism from rapid depletion of calcium and alkalinity from the system by calcifying organisms. One of my recent columns in Aquarium Frontiers (January/February 1997) gave some estimated calcification rates for reef aquaria, and information was given on how much limewater various types of aquaria are expected to consume. The rates given in that article are approximate -- your mileage will vary based on how intensely your aquarium is illuminated, the water temperature and several other variables. However, these approximate rates have utility, and match the actual rates observed in my aquarium, in those of my friends and the results from two unpublished surveys that were used to verify the applicability of these rates for captive ecosystems (Bingman, unpublished data).


How Might Magnesium Be Lost From Reef Aquaria?
The supposed magnesium-depleting reactions are closely related to the previous two reactions. In place of calcium carbonate, some form of magnesium carbonate or hydroxide is supposed to form in the system.

Ca++ + 2(OH-) + Mg++ ---> Ca++ + Mg(OH)2
(from limewater) (from the aquarium) (solid)


or

Ca++ + 2(OH-) + 2 (HCO3-) + Mg ---> CaMg(CO3)2
(from limwater) (from the aquarium) (solid)


Our job is to see if these reactions are realistic, and to what extent they occur in functioning reef tanks.


The Carbonate Minerals
The properties of the carbonates that form in marine sediments have been studied intensively for several decades. Photographs of the various types of calcium carbonate can be found at http://mineral.galleries.com/mineral.../aragonit.htm.

These compounds are well characterized, and their stability under a wide range of solution conditions is known. There are two common forms of calcium carbonate: aragonite and calcite. The structure of calcite is depicted at http://darkwing.uoregon.edu/~jrice/g...1/fig14-2.html The structure of aragonite is depicted at http://darkwing.uoregon.edu/~jrice/g.../fig14-6.html. There is a third, less common form, vaterite, that does not seem to be produced in appreciable quantities biologically, and is rarely formed from marine sediments. Aragonite has a density of 2.83 kilograms per liter (kg/L), and calcite has a density of 2.711 kg/L.


The two major allomorphs of calcium carbonate are related to each other as the two familiar allomorphs of elemental carbon: diamond and graphite. Diamond is more dense, but is thermodynamically unstable at 1 atmosphere of pressure. (Don't be too concerned about that diamond in your ring turning into a lump of coal -- the reaction to form graphite is very very slow at room temperature.) Aragonite, being more dense, is the energetically most favored phase of calcium carbonate at high pressures. A more complete discussion of the pressure dependence of the stability of calcite and aragonite can be found at http://darkwing.uoregon.edu/~jrice/g.../fig14-5.html.


However, the universe is not ruled exclusively by the absolute stability of compounds. The relative ease of their formation (kinetics) is also important. As we will see later, although aragonite is unstable with respect to calcite in aquarium conditions, aragonite is the form of calcium carbonate preferentially formed by modern corals, and the form that precipitates most easily from near-seawater conditions.


An overview of various carbonate minerals, with pictures of specimens, can be found at http://mineral.galleries.com/mineral...nat/class.htm.


A similar page with pictures of hydroxides can be found at http://mineral.galleries.com/mineral...e/brucite.htm.


An overview of carbonate and hydroxide minerals from a geological standpoint is available at http://darkwing.uoregon.edu/~jrice/g...carbonate.html and http://darkwing.uoregon.edu/~jrice/g...ydroxide.html.


A discussion of the chemistry of Group II metals is available at http://www.nidlink.com/~jfromm/elements/alkaline.htm.


Carbonate Mineral Composition
Aragonite CaCO3
Calcite CaCO3
Vaterite CaCO3
Magnesite MgCO3
Hydromagnesite Mg4(CO3)3(OH)2:3H2O
Dolomite CaMg(CO3)2
Brucite Mg(OH)2
At this point, I'll introduce the two figures that form the backbone of this column, which were adapted from Morse and Mackensie (1990). Figure 1 shows the relationships between the energetically favorable phases of calcium carbonate in aquarium conditions (temperature and pressure). The x-axis indicates the pressure of carbon dioxide in atmospheres. Lower pressures of CO2 correspond to higher pH values. So, the highest pH values are on the right side of the figure.


The y-axis is the ratio of the activities of calcium and magnesium. The activity (a in the figures) of an ion is related to the concentration of that ion by a weighting or intensity factor called the activity coefficient. In terms of parameters that you can measure, low calcium to magnesium concentrations are lower on the y-axis, and high calcium to magnesium ratios are higher. Natural seawater conditions are indicated with a box. As you can see, this box sits in the dolomite domain. Therefore, dolomite is the energetically most favored form of carbonate mineral in near-seawater conditions. The mineral brucite is magnesium hydroxide, FIGURE 2

which is what is supposed to be accumulating in our systems when calcium hydroxide is added. As you can see from the phase diagram, brucite is not a stable form at anywhere near aquarium conditions.


Figure 2 shows the most kinetically favorable phases of carbonate minerals in near-seawater conditions. Natural seawater conditions are again indicated with a box. The box barely sits in the magnesian calcite phase, and is very near the aragonite phase. Only a small decrease in magnesium concentration will shift the point into the aragonite domain. So, if precipitation of magnesian calcites happened under aquarium conditions, it would be a self limiting phenomenon -- it would deplete the magnesium concentration in the tank and shift the situation higher in the figure into the aragonite domain. Inorganically precipitated aragonite typically has 1 percent or less magnesium content, so magnesium loss through the formation of aragonite in the reef would be very slow.


Most importantly, note that the magnesium-rich hydroxides and hydroxycarbonates brucite, hydromagnesite, magnesite and huntite are not stable at seawater conditions. REFERENCES
Bingman, C. 1997. Defining the limits of limewater. Aquarium Frontiers Jan/Feb:8-11.

Frakes T. and J. Studt. 1996. Notes on the use of kalkwasser, Part I. SeaScope Fall:1.

Frakes T. and J. Studt. 1997. Notes on the use of kalkwasser, Part II. SeaScope Spring:1.

Morse, J. W. and F. T. Mackensie. 1990. Geochemistry Of Sedimentary Carbonates. Elsevier Science Publ. Co., New York. Pp. 707.

The magnesium-rich phase closest to stablity is huntite, and the stability of this phase is influenced more by the magnesium to calcium activity than pH. In a diluting plume of limewater, the pCO2 (partial pressure of CO2) will be lower than natural seawater, but so will be the magnesium concentration, so diluting limewater plumes are up and to the right of the natural seawater (NSW) point in these figures. Calcium test kits, which increase the pH of the sample to extremely high values (greater than 12 in some cases) are capable of moving the system into the hydromagnesite and brucite domain.


I agree with a recent SeasScope article (1997) that states that magnesium loss in reef tanks is not as significant an issue as was thought in the hobby a couple of years ago. I've always held the opinion that the formation of magnesian calcites (both biologically and chemically) is the major source of magnesium loss. The reason that many people observe much lower than natural seawater values for the magnesium concentration in their reef aquaria is that the sea salt they are using starts out deficient in magnesium. As people in the hobby have become aware of this issue, I've noticed that the magnesium concentrations in certain salts have increased substantially.

------

Hope you find what you're looking for. All in all magnesium levels is very tightly woven with levels of calcium in our system.

- Elmo

[saltjunkie]

Thanks Elmo!!!!!!


That was alot of detailed info, Im glad somebody out there can find stuff like this!


So assuming this data, am i correct in saying that if I/we dont attain NSW level of mag. (1300ppm, i think) that we wont be able to maintain the correct levels of calcium or alk?
Also before hand, If the salinity in my system isnt at 35 ppt, then my mag level and reading will be incorrect?


If this is so, if my water is not at nsw levels which is???? How do i affectivly maintain and balance all these different chemicals?

[saltjunkie]

anybody else bother dealing with MAG?

[Sueet]

Originally Posted by saltjunkie

anybody else bother dealing with MAG?

Nope... I figure whatever I need comes in the salt and gets added with water changes...lazy me

[saltjunkie]

Originally Posted by Sueet

Nope... I figure whatever I need comes in the salt and gets added with water changes...lazy me

thats pretty much what i go by.. but mag, is a major componet of the nsw make up and keeping all levels balanced correctly..
i am still just curious if all of our water is right.. salt.. and even test kits are calibrated to this chemical and its neccisity to keep all other chemical levels stable????

[Ninong]

Originally Posted by saltjunkie

anybody else bother dealing with MAG?

Yes, I test for magnesium about once every two months or so and add a little magnesium chloride if necessary to keep my levels near NSW levels (~1290 ppm). I would not want my magnesium levels to fall below 1000 ppm or rise above 1600 ppm. I have never tested high on magnesium but I do regularly test around 1150 ppm, in which case I add about a tablespoon of magnesium chloride daily until I get it up around 1250 ppm.

This is something that you do not want to fall too low or rise too high.