Marine Chemistry 101

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geologeek
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Marine Chemistry 101

Postby geologeek » Fri Nov 13, 2009 3:42 pm

Blatantly taken from the user contributed article - for easier reference :)

There are quite a few questions being asked about the chemistry within reef aquaria. There is a wealth of information out there but i thought i would do a series of threads that try to address these from a hobbiest perspective.

I am a simple geologist by trade but have had geochemistry drummed into me to the point that it hurt and within my career i have had to deal with the legacy our predecessors have left us in the form of contaminated ground which ultimately affects controlled waters.

So i am not a reef expert by any means but i have a fair understanding of the processes that occur in our aquariums and have experienced many of the same problems as you all in my 20 years of reefing.

So.............

Calcium and alkalinity are amongst the most important parameters in the marine aquarium. The relationship between these and how one interacts with the other is often confusing or poorly understood by marine and reef keepers.

Reference to a chemistry book or a search on the internet will show just how well such relationships are understood by the scientific community but these interactions are often portrayed in chemical equations which unfortunately can often be quite baffling to the average aquarist.

Such clinical descriptions are dissatisfactory for many aquarists without a scientific background who can become quickly bored or confused by such terminology. The interactions between calcium and alkalinity are further complicated when pH and magnesium are added to the mixing pot and the relationships become more convoluted.

The following aims to set out the relationship between calcium, alkalinity, pH and magnesium in an informative way that will outline the reasonably complex chemical processes that are continually occurring unseen in our aquaria until a problem arises and we tend to panic and over rectify any deficiencies indicated by our test kits.

The eventual goal of this article is to prevent any misunderstandings that could ultimately confuse the aquarist.

What is calcium and alkalinity?

Whilst not naturally found in its elemental state, calcium is an alkaline earth metal that is essential for living organisms.

Within the Earth’s crust calcium is the fifth most abundant element, and accordingly it is the fifth most abundant dissolved ion in seawater. Its concentration in natural seawater is approximately 420 ppm and any variation from this concentration is largely as a result of changes in salinity or due to the fact that it has a variable residence time and tends to precipitate out of solution quickly.

Calcium is extremely important in a reef aquarium as many of the organisms we keep, including corals and coralline algae use it to form and deposit calcium carbonate skeletons as well as proper function within their cells. If calcium is not maintained at adequate levels, such organisms can be distressed and if not rectified they can ultimately perish. As such the general consensus is to maintain calcium at 380-450 ppm within the reef aquarium.

With regards to alkalinity we are not looking for a single determinand within the water, but more accurately the summation of many determinands. Anything that absorbs a proton (Hydrogen ion) when the pH falls will contribute to alkalinity and a discussion of all the potential contributors would only complicate matters further and has been left outside the scope of this article.

For our purposes as aquarists, the reason that we measure alkalinity is that a large proportion of normal seawater consists of both bicarbonate (HCO3-) and carbonate (CO3-2). As previously mentioned corals and other organisms deposit calcium carbonate within their skeletons and other body parts. In order to do this they must generate calcium and carbonate at the surface of the growing calcium carbonate crystal (outlined below).

While it is far beyond the scope of this article to fully describe the exact process, it is readily apparent that if corals deposit these chemicals, they are using them up from the water that they inhabit. So, if that is the case then why not just measure carbonate levels as we do with calcium?

Well, there are two real reasons; the first is that there is no simple way for the average aquarist to measure carbonate without undertaking a more controlled titration than our hobbyist test kits can allow. Secondly, corals may actually use bicarbonate instead of carbonate as their ultimate source of carbonate (which they split into H+ and CO3-2).

So what we are doing when we test for alkalinity is using a very simple test as a surrogate measure for the sum of all bicarbonate and carbonate.

Normal to high alkalinity implies adequate levels of bicarbonate, whilst a low alkalinity implies that it may be depleted. In the absence of any method of supplementing alkalinity in a reef aquarium, the water can rapidly become exhausted of bicarbonate.
This depletion of alkalinity from normal to unacceptable levels can take only a day or two in reef aquaria with a high biological demand for carbonate.

As with calcium, when the levels of bicarbonate are depleted, corals that deposit calcium carbonate can become stressed and may ultimately perish.

It is generally regarded that reef aquarists maintain an alkalinity of 2.5-4 meq/L (7-11dKH) within the aquarium.

Why do we need to understand Calcium Carbonate?

Due to the aforementioned requirement that calcifying organisms, including corals and coralline algae, have for both calcium and alkalinity, there is a requirement to continually ensure that appropriate levels are present and maintained within the aquarium.

To complicate matters, there is a natural tendency for nonbiological (abiotic) precipitation of insoluble calcium carbonate from the water. This tends to occur as calcium ions and carbonate ions combine.

Understandably this tendency for precipitation of calcium carbonate plays a big role in the relationship between calcium and alkalinity in reef aquaria.

Here is where the first of the complexities in reef chemistry comes to the fore! As mentioned calcium is the fifth most abundant ion in seawater and is found in much larger quantities than bicarbonate or carbonate.

If we discounted natural replenishing of the components of alkalinity and the sum of these were removed by precipitating calcium carbonate, the calcium levels in the ocean would theoretically drop by only about 50 ppm.

It is for this very reason that alkalinity varies much more rapidly and extensively, on a percentage basis, than calcium when both are over or under dosed, relative to their demand.

So how does this all work?

Forgetting stoichiometric calculations to determine solubility we all know that if we pour sugar into our coffee in the morning it will reach a point where no more sugar can dissolve into solution.

As far as the salt we use in our tanks are concerned this dissolution arises because of the attraction between the positive charges of some of the salt and the partially negative oxygens in the water or the negative charges of some of the salts being attracted to the partially positive hydrogens in the water.

The same is true for salt as it would be for the sugar in our coffee, in that there is a limit to how much salt can be dissolved in a given volume of water. For example, sodium and chloride ions combine to make solid sodium chloride crystal (table salt). Imagine a small crystal of sodium chloride is placed into a glass of freshwater and you will see that it dissolves as ions leave its surface and go into solution within the water.

Inevitably as this occurs the number of ions in the water increases as the salt dissolves. Whilst the crystal of salt becomes smaller there is actually a movement of ions on a molecular level. As ions leave the surface of the salt crystal to go into solution, ions are also landing on the surface of the salt and coming out of solution.

This is where the solubility limit comes into play. If the salt crystal was small enough it would fully dissolve, resulting in all the ions going into solution (as in our aquariums) but if the crystal was larger, the solubility limit would eventually be reached whereby no more salt can be taken into solution.

In other words the rate at which ions leave the surface of salt crystal and going into solution equally matches the rate at which they come from solution and become part of the salt crystal.
Now if we complicate matters further and take our sodium chloride solution and add more sodium to the solution we would increase the rate at which sodium ions land on the surface of our salt crystal. In doing so these ions can then trap the chloride ions beneath them preventing them from escaping. Therefore by increasing the concentration of one the ions we can decrease the concentration of the other ion in solution by forcing it into its solid phase.

Supersaturation is the term used to describe a solution that has more dissolved material than could be dissolved by the solvent, which in our case is water. If for example we took 20g of sodium chloride and added enough water to fully dissolve the salt we would have a solution at saturation. By adding a further 5g of sodium the solution would then have 15g of sodium and 10g of chloride and could be termed supersaturated. The effect is there are more ions in solution than is stable.

If we then add our salt crystal to this solution the rate at which sodium and chloride ions land on the surface of the crystal exceeds the rate at which they can leave. This will result in precipitation of sodium chloride, increasing the size of the salt crystal until the point at which the amount in solution returns to saturation.

Thus it becomes clear that when the water is supersaturated some precipitation will inevitably take place and result in the amount of ions in solution decreasing to the point of saturation, whereby the rate of dissolution and precipitation becomes equal.

What has this got to do with Calcium Carbonate?

To take this analogy one step further, let us now discuss the solubility of calcium carbonate. The process crudely outlined above is also applicable for calcium carbonate in that the solubility is determined by the rate at which calcium and carbonate ions land on and leave a solid surface. As with the supersaturation of the sodium chloride solution, if we were to elevate either the levels of calcium or carbonate ions artificially with an additive, this can result in forcing the other ion onto the surface of the solid.

This however is an over simplistic view as some of the effects of calcium carbonate solubility are much more complicated than that of sodium chloride and will be discussed further below.

Many references state that the surface waters of the ocean typically have levels of calcium at 420 ppm, an average pH of 8.2 and an alkalinity of 2.5 meq/l (7 dKH). Normal seawater is significantly supersaturated with calcium carbonate meaning that there is more in solution than is stable. Discounting the effects of magnesium at this stage, the rate at which the calcium and carbonate ions land on a solid surface in the ocean far exceeds the rate at which they can leave.

Within our aquaria, should we have very high calcium or alkalinity then precipitation of calcium carbonate can reduce the level of both of these in solution. If we increase the levels of these ions by dosing our two part additives to the point where the ions are in excess of the saturation point then the rate of calcium carbonate precipitation is sped up and would result in accelerated coral growth.

Conversely if the levels of calcium and/or carbonate ions in solution are below the point of saturation there would be no net gain in precipitation of calcium carbonate.

What is the relationship between Calcium Carbonate and pH?

pH as you will be aware is a measure of the acidity or basicity of a solution. Pure water is said to be neutral. The pH for pure water at 25 °C (77 °F) is close to 7.0. Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are said to be basic or alkaline.

The pH scale is logarithmic; for example a pH of 8 is ten times more alkaline than a pH of 7 and a pH of 9 is 100 times more alkaline than a pH of 7.The solubility of calcium carbonate depends strongly on pH in that the lower the pH, the more soluble calcium carbonate becomes.

Both bicarbonate and carbonate are forms of the same ion. At a lower pH, the bicarbonate form (HCO3-) predominates but at a higher pH, more of the carbonate form (CO3-2) exists. The effect of varying pH can be very acute in that each drop of 0.3 pH units below a pH of 9 causes a two-fold drop in the carbonate concentration. A full pH unit drop would correspond to a ten-fold decrease in carbonate concentration.

Accordingly it can be seen that by varying the pH of a solution we can also change the amount of carbonate ion in solution. As previously discussed it is the concentration of these ions that determines the rate at which carbonate ions land on the surface of the solid. So the higher the pH, the faster the rate at which these ions land on the surface and thus the solubility of calcium is lower at higher pH’s.

This lower solubility implies that as the pH increases the amount of calcium and alkalinity that can be kept in solution without precipitation occurring decreases. To put this into context, if we were to add a large amount of kalkwasser (limewater) solution to the aquarium, the result would be an increase in pH. This would rapidly permit the precipitation of calcium carbonate. The precipitation is not necessarily as a result of increasing the levels of calcium or alkalinity, although these can play a role, but is also due to the fact that as we increase the pH, a lot of the existing bicarbonate within the water converts to carbonate, with a resultant spike in carbonate concentration.

The opposite is true with a falling pH in that the amount of calcium and alkalinity that can be kept in solution without precipitation occurring increases. This is why adding carbon dioxide to a reactor dissolves the media.

You may therefore think that it would be better to run your aquarium with a lower pH as you can maintain calcium and alkalinity levels better and abiotic precipitation of calcium carbonate will not occur. However, doing so would be detrimental to your corals due to the fact they have to work much harder converting bicarbonate to carbonate to allow calcification.

So what does Alkalinity have to do with Calcium Carbonate?

To muddy the waters further we must now consider the role of alkalinity on the solubility of calcium carbonate.

If we take our body of water at a fixed pH of 8.2 and increase the alkalinity then the amount of carbonate would be directly proportional. That is for example if we raised the alkalinity to 14 dKH there would be twice as much carbonate within the body of water than that of natural seawater at 7 dKH.

The same mechanisms discussed above have a role in the solubility of calcium carbonate with changes in alkalinity. The rate at which the calcium and carbonate ions land on, and leave the surface of a solid can affect the level of alkalinity by changes in the concentration of carbonate within the solution.

Lower calcium carbonate solubility at higher alkalinity implies an increased precipitation of calcium carbonate. In other words, as the level of alkalinity rises, the amount of calcium that can be kept in solution without precipitation decreases.

If we again try to put this into context; those aquariums that are maintained at a very high alkalinity will tend to observe abiotic precipitation of calcium carbonate on circulation pumps. Should we subsequently reduce the alkalinity of the said aquarium, then the amount of calcium that can remain in solution increases.

And Magnesium?

The role of magnesium is far more complex than that of both pH and alkalinity. If we take the molal composition of standard seawater it can be seen that magnesium is the third most abundant ion after chloride and sodium.

So its presence complicates our simplistic view of calcium and carbonate landing on and leaving the surface of solid calcium carbonate. Magnesium ions have the ability to be incorporated into the crystal structure of calcium carbonate and in doing so replace the role of the calcium ions.

Over time as more and more magnesium is incorporated onto the surface of the calcium carbonate a layer of calcium and magnesium carbonate is formed. This coating results in the surface of the calcium carbonate no longer resembling that of calcium carbonate and prevents a firm bond of calcium and carbonate to the surface. As a result further precipitation of calcium carbonate is largely reduced.

The extent to which magnesium is incorporated within calcium carbonate surfaces depends largely on the amount of magnesium ions in solution. The greater the amount of magnesium within the solution, the more it is incorporated. Inversely if levels of magnesium are 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. This could potentially lead to an increase in abiotic precipitation of calcium carbonate.

Our inability as aquarists to maintain adequate calcium and alkalinity levels, despite extensive supplementation, or significant evidence of abiotic precipitation of calcium carbonate on heaters and pumps, are signs that levels of magnesium are inadequate within the aquarium.

So what has this got to do with my corals?

Now that we have a rudimentary understanding of the abiotic process of calcium carbonate formation we can apply this to biological deposition by corals, coralline algae and other organisms that deposit calcium carbonate.

Whilst the process is not strictly the same as abiotic deposition as these organisms have a certain level of control over the process, many of the interactions between pH, calcium and alkalinity as well as magnesium holds true.

These organisms utilise calcium and alkalinity almost exclusively (with the exception of magnesium and other trace elements) to deposit calcium carbonate. The chemical composition of calcium carbonate (CaCO3 ⇋ Ca2+ + CO3-2) indicates that these organisms use a 1:1 ratio of calcium and alkalinity. However as previously touched upon, the consumption of calcium can vary from species to species due to incorporation of magnesium within calcium carbonate.

Within the aquarium, where an aquarist can modify the water conditions, the deposition of calcium carbonate can be expedited by increasing the alkalinity which will result in the reduction of both calcium and alkalinity as the calcium carbonate is deposited. The opposite is true in that if levels of calcium are low with an average alkalinity the rate of deposition of calcium carbonate by the organism will be retarded or prevented until levels of calcium are increased. At which point the level of both calcium and alkalinity will begin to fall due to calcium carbonate deposition.

Accordingly, if the concentration of calcium or carbonate becomes too low within the aquarium, corals will have a far more difficult time depositing their skeletons.

pH levels can also have dramatic effects on the deposition of calcium carbonate. Even if the levels of calcium and alkalinity are in line with those of natural seawater, a low pH could ultimately result in the dissolution of coral skeletons (just as within our calcium reactors).

It can therefore be seen that there are several complex chemical reactions and interactions that occur within our aquaria. Whilst the above outlines a slightly simplistic view of these it should hopefully give you a better understanding of your aquarium and help you keep the parameters in check, but more importantly help you to appreciate how and why this should be achieved in a closed environment.

If you want to discuss this further just say - or if there is another aspect you want discussed feel free to ask and ill do my best to give it to you straight!

Remember! Go slow and let it grow.............
Go Slow And Let It Grow! - But remember Detritus Happens!

weasley
Posts: 354
Joined: Mon Dec 29, 2014 9:23 am
Location: Cheshire

Postby weasley » Sat Sep 05, 2015 10:27 am

Fascinating, thank you...


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