Iron transformations

Iron, another essential nutrient for plants and animals,  can become toxic in excess, partly because there's no natural mechanism for eliminating it from the organism.
 
Most of the iron in fish is represented in the hemoglobin of their blood. The fish prefer to get their iron ready bound in an organic "heme group" that is derived from their carnivorous diet of blood and tissue. Ultimately, down the food web, the iron has been derived from photosynthesizers. In plants on the other hand, iron is used, in molecular traces only, as part of metabolic pathways, not to build structure.
 
Aquarists who are trying to achieve "optimal" fertilizer concentrations may become fixated on iron. You'll often get the mis-impression that iron is a macronutrient of aquatic plants. In fact, iron retards plant growth at levels higher than 2.0 mg/l, and it's toxic to plants at levels greater than 5.0 mg/l, several agricultural extension websites will tell you. I'm not sure any more whether most of the iron in my aquaria occurs in the substrate porewater or in the water column above, but I think now it's mostly an academic curiosity, because iron is highly reactive in water and transforms from phase to phase.
 
Iron is not so unlike phosphate in several ways:  iron has no gaseous reservoir in the atmosphere, passes between soluble and precipitated forms, co-precipitates with other minerals, and forms hydroxides. In the heme groups of hemoglobin, iron even has an organic state, as does phosphate. Years ago I quit adding chelated iron to water as a fertilizer, though. There's already a lifetime's supply of iron in the aquarium, partly in the reddish Flourite that enriches my substrates. Instead I keep a supply of natural chelators constantly leaching into my waters from humic substances. And my acidic anoxic porewater keeps that deep substrate iron soluble. How does this work?
 
Ionic iron. In water, iron can take two ionic forms, soluble ferrous forms — Fe++ or Fe(II) — and the more oxidized insoluble ferric forms — Fe+++ or Fe(III). In general, the more soluble the form of iron, the more bio-available it is, both for animals and for plants.
 
In the aquarium, at normal pH values and the usual levels of oxygen, virtually all the iron in the water is insoluble Fe(III). Deep in the porewater of anoxic layers of your substrate, though, bacterial action has stripped oxygen from the iron and reduced it to colorless, soluble Fe(II). In oxygenated water, the reactive Fe(II) ions don't last long. They are quickly chelated to organic acids (or artificial chelators), or they are lightly bound in various ways to insoluble humus, or else the Fe is swiftly further oxidized to insoluble amorphous precipitates called ferric hydroxides (Fe(OH₂)) or oxyhydroxides (FeOOH), which coat the grains and colloids in the upper levels of your substrate and aren't available to plants. These processes happen within minutes to hours. 
 
Chelated iron. Which route Fe takes depends partly on the pH and alkalinity of the water. In acidic water such as mine (pH <7.0), if there were no added artificial chelator to bind it, the iron would rapidly bind to polyphenols (including humic substances such as tannins) in the water. It binds equally to long-chain insoluble polyphenols, like those in peat, or to short-chain ones, such as organic acids dissolved in the water. These form complexes with iron, and prevent further polymerization. And, since in general the longer the polymeric chain, the less soluble the molecule, chelation keeps the iron more soluble and so more bio-available.
 
The chelating bond isn't like a mineral binding. It binds and loosens many times a second, so that a percentage of free ions of iron is constantly being made available to plants and algae. When chelated with dissolved humins, iron of both ionic species remains potentially mobile enough for plants to take it up when the chelating bond eventually degrades.
 
Thus chelation works a bit like a time-release capsule. Though we think of chelated iron as "bio-available," the iron only becomes available to plants as the chelating bond breaks down and the ionic iron is momentarily freed. Roger Miller contributed a post "Re: adding iron" to the Aquatic-Plants Digest in July 2001 that adds some science to "balanced aquarium" lore. One point he made is especially useful: terrestrial plants don't show iron deficiency until soluble iron in the soil porewater drops to a few parts per trillion, he reported. Only then does plant metabolism actively start to break down chelated iron.
 
The chelated iron fertilizer often used in aquaria is NaFe(III)EDTA ("sodium iron ethylene diamine tetra-acetic acid"); that EDTA component is the widely used artificial chelating molecule). It's been thoroughly investigated because it's also used in dietary supplements to counter human dietary iron deficiency — anemia. The water-soluble sodium-iron complex (NaFe) is stabilized and kept soluble by chelating it with the man-made amino acid EDTA, a common chelator found in many water "conditioners."
 
Investigators working with anemic lab rats found that dissolved tannins did not affect the bio-availability of NaFe(III)EDTA, resolving a possible concern for aquarists with soft acidic waters laden with humic acids, such as mine. In other words,   organic acids such as tannins supplement the chelating effect of EDTA, rather than compete with it.
 
Ferric precipitates. Ionic iron that isn't chelated will precipitate in various ways. For instance, in alkaline waters, where calcium is plentiful, unchelated iron would rapidly complex with the calcium or the carbonate instead, and co-precipitate out.
 
Such ferric precipitates can cause milky-white cloudiness after iron is dosed as a plant fertilizer. A  thread at Aquatic-Plants Digest begun 15 Jan 2001 records an aquarist who was adding Seachem Flourish weekly and Flourish Iron daily, and who eventually found that the water was turning milky seconds after dosing, apparently precipitating with something. After Flourish Iron was added, iron levels were as high as 0.5 mg/l, but within two to three hours iron was undetectable again.
 
Dr. Greg Morin of Seachem responded to the post, suggesting that the precipitate might be iron carbonate, which usually is reported only when KH levels are high, around 12. Yet the aquarist reported low KH around 2, and the absence of household scale on his tea kettle etc. seemed to corroborate his test results. "Or maybe your iron level is already quite high?" Dr Morin politely suggested. No one at the mailing-list seemed eager to take that up. Or it could be another co-precipitate, iron phosphate, Dr Morin thought. (It seems odd to me that fertilizer dosers remain convinced that their plants are usingthis iron.)
 
In any of these ferric precipitates, the iron is insoluble— unavailable— and evades your iron test the following day. Farmers with troublesome levels of iron in well water may often add hydrated lime to a holding tank, in order to raise the pH to about 8.0 and precipitate out the iron in reactions very much like these.
 
That is why iron supplements for plants figure prominently in dosing highly-buffered waters but are less urgent in "soft" waters. Each approach is correct in its appropriate water. There's no all-purpose iron fertilization schema for plants in every water: it depends on your own pH and GH and alkalinity too.
 
Ferric hydroxides and oxyhydroxides(e.g. Fe(OH) ₂, FeOOH, Fe₂ O₃) comprise other common stable and highly insoluble forms of unchelated Fe, which precipitate out in aerobic environments at normal aquarium pH levels, especially above pH 7.0. Ferrous Fe(II) is the soluble phase, but in ordinary oxygenated aquarium water, at aquarium pH values, it is rapidly oxidized to ferric Fe(III). Ferric oxides are generally coated with negatively-charged reactive -OH functional groups that adsorb organic and inorganic molecules in the water column and at the surface of the substrate. Lower in the substrate, where oxygen is sufficiently low, the Fe(III) dissolves again, freeing adsorbed and co-precipitated species. There in the anoxic pore water, iron hydroxides are redissolved to iron's colorless ferrous form Fe(II), which is the only form in which iron is available to plants— or algae. If this ferrous iron isn't taken up by plant roots, it may diffuse towards the surface levels of the substrate. When it reaches pore water with some dissolved oxygen, this ferrous iron is reoxidized to its ferric Fe(III) form, which precipitates again as ferric hydroxides.
 
You can see that in a substrate that's not artificially disturbed, these opposing gradients—  diminishing oxygen supplies diffusing from above meeting diminishing supplies of soluble Fe(III) below, tend to lay down a narrow, well-defined dark layer of precipitated iron. Such a black layer could be mistaken for the presence of hydrogen sulfide. Iron bacteria take advantage of this "redox gradient" to make a living in microzones in your substrate.
 
Pumping iron: alkaline precipitation  is the most common industrial-scale technology for removing heavy metals from wastewater. Like many modern wastewater technologies, it's quite relevant to the aquarium, where the very same chemistry harnessed in the industrial process is at work. Unocal Corp. developed an alkaline precipitation system they call "Unipure" manipulating pH and ferric hydroxides to strip heavy metals from effluent. At their website they lay open the chemistry so clearly that even I understand the basics now
 
The essence of alkaline precipitation is this: iron will attract and adsorb other heavy metals, among other substances. The two ions will co-precipitate and could be caught in the floc that accumulates in the aquarium's mechanical filter. This co-adsorption is reversible. One of the substances that iron will precipitate, you'll be interested to hear, is phosphate (PO₄-).
 
Alkaline precipitation is based on a chemical equation
 
Fe⁺⁺ + 2(OH^-) <-> Fe(OH)₂
 
that I will translate for you: "A divalent ferric ion plus two hydroxide ions (which carry two negative charges) gives insoluble ferric hydroxide." 
 
Any divalent metal ion (i.e., one with a double positive charge) would do the same: copper (Cu++) for instance. The reaction is reversible (hence the double-headed arrow), for it is pH-dependent. Adding OH¯(hydroxide ions) to the water moves the reaction to the right, producing more insoluble hydroxide at higher pH. As the reaction moves to the left (with more H+ ions, i.e. with lower pH) the metal hydroxide comes back into solution. Copper medications are well known to be more toxic at lower pH. Now I see why. As the pH drops, the Cu++ is dissolved from its hydroxide and becomes bio-available and toxic once more. You'll probably sense that there's some kind of buffering situation here, analagous to the carbonate/bicarbonate buffer. You'd be right. Unipure offers a graph that shows the solubility of metal hydroxides at pH values above 7.0. Solubility drops precipitously: at pH 7.5 hydroxides are only a tenth as soluble as at 7.0, and by pH 8.0 they are only a hundredth as soluble. This process doesn't continue, because as available hydroxide ions increase, the metal hydroxide takes on another 0H¯, to form a metal complex that's soluble. That reaction only takes over in very high pH. The two curves intersect about pH 8.5— above that the metals are increasingly soluble once more.
 
Iron bacteria. These purely chemical sources of iron hydroxides I've been noting so far aren't the only sources. Bacteria power all the aquarium's nutrient cycles, and the constant restless transformations of iron aren't an exception. In the microaerobic gradients of diminishing porewater oxygen below the surface of the substrate, there are bacteria that can compete successfully with the abiotic chemical oxidation I was describing a few moments ago. A community of bacteria that are collectively called "iron bacteria," though they aren't all genetically related, make a living oxidizing iron, independent of light, wherever they can get as little as 0.3 ppm dissolved oxygen. That's pretty much down to the limits of oxygen diffusion into the substrate. They deposit the resulting ferric hydroxides in their cell sheaths and in the slimy biofilm they generate. Plants (and algae) are in direct competition with iron bacteria for scarce supplies of ferrous iron. Iron bacteria aren't toxic to us or to fish, but they can cause esthetic problems in drinking water systems with their slimy, reddish-brownish biofilm with funky odors described as a blend of fuel oil, overripe cucumbers and low-tide mudflats. In unplanted tanks, iron bacteria could be responsible for some untraceable odors and mysterious staining. Are these some of the odors we mistake for hydrogen sulfide in the bands of precipitated iron? (I've recently got the overall picture from an excellent technical paper by Eric E. Roden et al., U. of Alabama, "New insights into the biogeochemical cycling of iron in circumneutral sediments."
 
Elusive iron. In brief, when an aquarist adds iron to a planted tank, then tests the following day and finds iron levels that are once again untraceably low, there are many routes the iron may have followed, aside from being taken up by the plants or algae: it may have been scavenged by iron bacteria, or abiotically oxidized by dissolved oxygen, it may have precipitated out or co-precipitated, or it may have got bound to an organic chelating compound dissolved in the water, such as one of the extremely various humic substances, tannins or humic or fulvic acids and the like. In more alkaline waters than mine, these humins would be chelating calcium and magnesium in competition with iron, which remains more vulnerable to oxidation. So plants in alkaline water can suffer Fe deficiency, even when the element is plentiful in the system—a less troublesome problem for plants in waters with pH values below 7.0. Iron toxicity in over-fertilized planted aquaria is an unexplored issue, I feel. I would generally prefer to add humic substances, in order to chelate the iron that is already there, rather than add iron to my aquarium systems.