# [Wet Thumb Forum]-Ionic iron scavenges ammonia.



## wetmanNY (Feb 1, 2003)

_Here's a further iron reaction, which I don't remember hearing mentioned in the context of daily iron dosing in planted aquariums:Though ionic iron reacts with many species in the aquarium, I hadn't realized there could be a reaction with NH3 that removes ammonia from circulation. I discovered it in Prof JP Birk's chemistry pages at the Arizona State U. website http://www.public.asu.edu/~jpbirk/qual/qualanal/iron.html In ordinary oxygenated water, both Fe(II) and Fe(III) react with dissolved ammonia (NH4) to produce a rusty-colored iron hydroxide Fe(OH)3. The Fe(II) reaction involves an intermediate that's rapidly oxidized to the same end product. The equation that makes it run is there at Prof. Birk's pages: even color pix of the products in test tubes.

How do I escape this conclusion: that if you are dosing with iron daily, the ionic iron is in direct competition for NH4 with your plants? Plants need such minute traces of iron, but they require lots of nitrogen. Why isn't this competition considered counter-productive in planted-tank circles?_

--cut n pasted from www.skepticalaquarist.com "Water" folder, "Nutrient Cycles" pages, for your criticisms.


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## wetmanNY (Feb 1, 2003)

_Here's a further iron reaction, which I don't remember hearing mentioned in the context of daily iron dosing in planted aquariums:Though ionic iron reacts with many species in the aquarium, I hadn't realized there could be a reaction with NH3 that removes ammonia from circulation. I discovered it in Prof JP Birk's chemistry pages at the Arizona State U. website http://www.public.asu.edu/~jpbirk/qual/qualanal/iron.html In ordinary oxygenated water, both Fe(II) and Fe(III) react with dissolved ammonia (NH4) to produce a rusty-colored iron hydroxide Fe(OH)3. The Fe(II) reaction involves an intermediate that's rapidly oxidized to the same end product. The equation that makes it run is there at Prof. Birk's pages: even color pix of the products in test tubes.

How do I escape this conclusion: that if you are dosing with iron daily, the ionic iron is in direct competition for NH4 with your plants? Plants need such minute traces of iron, but they require lots of nitrogen. Why isn't this competition considered counter-productive in planted-tank circles?_

--cut n pasted from www.skepticalaquarist.com "Water" folder, "Nutrient Cycles" pages, for your criticisms.


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## Roger Miller (Jun 19, 2004)

wetmanNY,

In the reference you site iron is not "scavenging" ammonia. Ammonia only participates in the reaction as a buffer. It's in the left side of both reactions as NH3 and on the right side as an equivalent amount NH4+, both of which are in solution.

It looks like Dr. Birk's web page supports a class in Qualitative Analysis. Aquarium conditions are not comparable to Prof. Birk's test tubes. 


Roger Miller


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## wetmanNY (Feb 1, 2003)

I understand now. No gain or loss of NH3/NH4, even in the test tube. ...Well, no wonder this "reaction" isn't "considered counter-productive in planted-tank circles>"

I better get that paragraph right out of that page.

(Learning in public has always required a good deal of courage...)


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## anonapersona (Mar 11, 2004)

So, what is the difference between NH3 and NH4+? Isn't that ammonia and ammonium? I thought that the presence of ammonia vs ammonium was controlled by pH. Will this reaction change in my alkaline water vs wetman's acidic water?


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## Roger Miller (Jun 19, 2004)

The reaction was in a test tube in a fairly concentrated ammonia/ammonium solution. It isn't significant under aquarium conditions.

One thing to keep in mind is that the precipitation of iron hydroxides produces quite a bit of H+. It would be interesting to work out how much H+ is produced and whether there might be a measurable effect on alkalinity.


Roger Miller


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## Roger Miller (Jun 19, 2004)

I found myself near the end of my work day, waiting for a program to finish a hideously long run. So I figured this out myself.

Precipitating Fe(III) from aqueous solution works like this:

Fe(III) + 3H2O <-> Fe(OH)3 + 3H+

the buffering reaction we are concerned with is:

3HCO3- + 3H+ -> 3H2O + 3CO2

So the net reaction is:

Fe(III) + 3HCO3- -> Fe(OH)3 + 3CO2

Precipitating one mole of Fe(III) as its hydroxide uses 3 moles of bicarbonate. The ratio in weight units works out to... (drum roll please):

0.15 dKH per mg/l Fe(III).

Let's take the hypothetical case where someone doses their tank to a measurable 0.1 mg/l every day and by the end of the day they have unmeasurable iron. They repeat the next day. They continue this routine every day. Forever.

The plants don't use anywhere near that much iron and they never will. Almost all of that iron is precipitated as an hydroxide. Suppose that it's Fe(OH)3. Over a period of 30 days the aquarist added 3 mg/l of Fe(III). Most all of that iron precipitated, taking with it 3*0.15 dKH of alkalinity, or 0.45 dKH.

Ok, so if the hypothetical aquarist never changes water the effect will eventually build up to something to get concerned about. In a more realistic assessment, the acidifying effect of iron precipitation is one factor that tends to break down the alkalinity. There are others and iron precipitation is certainly not the biggest factor.

Roger Miller

P.S. Fe(III) hydroxide probably doesn't really precipitate as Fe(OH)3. Reality is far more complex. This simplification is common and probably close enough for almost any purpose.


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