CO2 Calculator

Enter the measured pH and KH of the aquarium water. The CO2 concentration will appear in the lowest box.


pH:


KH:


CO2:

ppm

(JavaScript version, for a temperature of 25C.  See below at right for a downloadable Excel version)

Planted Aquaria

Control of Algae in Planted Aquaria

Paul Sears and Kevin Conlin wrote an article with this title in 1996. It has been reproduced widely in several languages. The original URL is no longer active. The same article may be found along with lots of other material at The Krib, a site mainly about planted aquaria. For a mailing list about plants, go to The Aquatic Plants Mailing List.

More Recent Impressions

The method described in the original article has demonstrated its effectiveness. The idea is to supply the plants in the aquarium with all the necessary nutrients, but to restrict phosphorous. The major nutrients required, other than phosphorous, are nitrogen and potassium. These may most easily be supplied by the addition of potassium nitrate (KNO3), which provides both, and potassium sulphate, K2SO4, which supplies potassium alone. Trace elements are most easily provided by using a commercial trace element mix intended for hydroponic use. Magnesium may be supplied as Epsom salts, magnesium sulphate, MgSO4,7H2O. Phosphorous must be restricted, but not eliminated. If and when it is necessary to add it, potassium dihydrogen phosphate (KH2PO4) or dipotassium hydrogen phosphate (K2HPO4) may be used. Supplementary carbon dioxide may be added, but it is not necessary.

The original article contained a recipe for "Poor Man's Dupla Drops", (PMDD), later renamed "Poor Man's Dosing Drops", which contained all the required nutrients except phosphorous in fixed proportions. In retrospect, I think that publishing this recipe was a mistake. There was a tendency to add this mixture to aquaria without paying sufficient attention to what was going on. It is much better to measure and control the concentrations of some at least of the individual nutrients.

Up to a point, the appearance and growth of the plants (and algae!) in the aquarium may be used to decide what nutrients are required at any given time. It is very easy, though, to measure the concentrations of nitrate and of iron using commonly available test kits. Phosphate test kits have also become reasonably common. Potassium kits do exist, but I can't say I've seen one yet. When I get around to it, I'll buy some sodium tetraphenyl boron and make one up myself. This test relies on the turbidity (cloudiness) of the test solution, so it is somewhat less satisfactory than the more usual ones based on colour. Whatever test one uses, and from whatever source it comes, the test itself should be checked from time to time. A standard solution (even a rough one) should be made up and tested to check that the test gives a reasonably correct result. Equally important is to run a blank test, on a solution containing none of the substance in question. One test to forget about is carbon dioxide. The reagent used to measure this is very dilute sodium hydroxide containing an indicator. It is virtually impossible to prevent this reagent being neutralised by atmospheric carbon dioxide. The reagent steadily becomes weaker, and the measured results keep climbing, sometimes to completely ridiculous values. I'll come to a better method for measuring carbon dioxide concentration a bit later.

Armed with at least a nitrate test and an iron test, one should attempt to maintain the concentration of nitrate ion (NO3-) at 5 to 10 ppm. Iron should be just detectable, a fraction of a ppm, but definitely there. If I could measure it, I would try to maintain potassium also at 5 to 10 ppm. As it is, when I add potassium nitrate, I just add rather less potassium sulphate than nitrate. Phosphate should also be in the "just detectable" range. I have deliberately put about 5 ppm into an aquarium, and algae rapidly became a problem.

If one is starting with an aquarium in which algae are rampant, it will take a while to get things under control. It is essential to keep the plants supplied with all the required nutrients, so that they can grow. The algae will grow as well, but eventually the phosphate concentration will fall sufficiently that algae will cease to grow. In my experience to date, I have the impression that the long green hair algae requires more phosphate than does green powder algae or black brush algae. Green spot algae is the last to remain.

If the aquarium is starved of phosphate beyond this point, the plants will also begin to die out. In general, the "easy", fast-growing plants die out first. If you have the algae under control but the plants are starting to suffer, check light, nitrate and iron concentrations (and potassium if possible). If all those are adequate, it is time to add phosphate. Use a small amount, and see what happens. If the plants begin to grow and the nitrate concentration begins to fall, shortage of phosphate was the problem.

pH Buffers

The pH of aquarium water is decided by the material dissolved in it.  In theory, at any rate, the pH of pure water would be 7.  That is, the concentration of hydrogen ions would be 10-7 molar, as would the concentration of hydroxide ions.  The pH is defined as minus the log of the molar hydrogen ion concentration.  In practice, water is very rarely so pure that the concentration of the impurities is less than 10-7 molar.  What is more, the conductivity of water that pure is so low that making measurements of its properties is very difficult.

One frequently sees comments in material on aquaria that water with little carbonate hardness is poorly buffered, and that its pH is not very stable.  This is literally half the story.  In a buffered solution, an acid and its salt with a (usually) metal ion, are both present and between them set the pH.  For an acid, which we will call for now HA, there is an equilibrium:

HA  <-->  H+  +  A-

The double-ended arrow indicates that the reaction goes in both directions all the time, and is in a state of dynamic equilibrium.  There is an equilibrium constant, Ka, defined as (square brackets mean "molar concentration of"):

Ka = [H+][A-]/[HA]

Rearranging this equation shows how a buffer system sets the pH:

[H+] = Ka[HA]/[A-]

The effectiveness of the buffer in resisting changes to the pH depends on the concentrations of both the acid and the anion.  The buffer best resists the changes caused by addition of another acid or base when the two concentrations are equal.  At that point, the hydrogen ion concentration is equal to the Ka for the equilibrium, or the pH is equal to the pKa.

The acid of interest in most aquaria is carbonic, H2CO3, and the negative ion (anion) is the bicarbonate ion, HCO3-.  The positive ion that goes along with bicarbonate ion is usually calcium, Ca++.  As luck would have it, understanding how this system works is considerably complicated by the fact that this is a most unusual acid.  It is formed when carbon dioxide dissolves in water, but only a very small fraction of the carbon dioxide dissoved is actually in the form of carbonic acid.  Nevertheless, since this fraction is more or less constant, we can get an effective Ka for:

[CO2](aq)  <-->  [H+]  +  [HCO3-]

The pH of the solution is set by the CO2 concentration and the bicarbonate concentration, which is usually referred to as the KH.  Note that it is always this way round.  You cannot set the CO2 concentration by setting the KH and then adjusting the pH.  The result of that is an alteration of the KH and a temporary alteration of the carbon dioxide concentration.  The reason that the CO2 concentration change is temporary is that there is another equilibrium involving the CO2 - it is free to come and go.  Plants absorb it, fish excrete it and there is continuous exchange with the air.  That is what sets the CO2 concentration.  It is possible to change it by using a CO2 addition system, but not by altering the pH by adding another acid.

It is nevertheless possible to find out what the carbon dioxide concentration is by measuring the pH and the KH and doing a calculation.

Carbon Dioxide Concentration Calculator

As discussed above, the test kits sold for measuring the carbon dioxide concentration in the aquarium are of somewhat dubious value.  When they are first made, they will probably give reasonable results, but the very dilute reagent will readily absorb atmospheric carbon dioxide, becoming still more dilute.  As time passes, more and more of it will be required to react with the carbon dioxide in a sample of the same aquarium water.  The Aquatic Plants Mailing List has more than once had a posting in which someone comments on their extremely high CO2 concentration.  In most of these cases, I suspect that a defective test kit was being used.  It is much better to measure the pH and KH of the water, and to infer the CO2 concentration from those.  Even when this is done, care must be used to get a good pH reading.  A sample of water removed from the aquarium will lose CO2 rapidly to the atmosphere, giving a falsely high pH, so speed is required if one is to get a meaningful result, particularly in aquaria to which supplemental CO2 is being added.  KH tests should be reasonably reliable, provided that the "alkalinity" of the water is due to bicarbonate ions (almost always the case).  Tables provided in aquarium books or on the web may then be used to calculate the CO2 concentration.  Alternatively, the result may be calculated.

An Excel spreadsheet that calculates the concentration of carbon dioxide in the aquarium water can be downloaded here.  The pH, KH and temperature of the water are entered in the appropriate cells, and the calculated CO2 concentration appears.  The pKa used for the first ionisation of carbonic acid is calculated as a function of the temperature.  For a simpler calculation, use the JavaScript version at the top of the page.

Stay tuned for more....

 

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