Iron is another essential nutrient for plants
and animals, one that can become toxic in
excess, partly because there's no natural
mechanism for eliminating it. 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 and transforms
from phase to phase. Not so unlike phosphate,
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.
I've quit adding chelated iron to water as
a fertilizer, though. There's 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 the aquarium, at normal aquarium
pH values, virtually all the iron in the
water is 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(III) 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. These processes
happen within
minutes to hours.
In general, the more soluble the form of
iron, the more bio-available it is, both
for animals and for plants.
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 (humic substances such as tannins) in the
water, both long-chain insoluble ones, like
peat etc.-- or to short-chain ones, organic
acids dissolved in the water. These would
form complexes with iron, and prevent further
polymerization. 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 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"--
the EDTA is the 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 is
stabilized and kept soluble by chelating
it with the man-made amino acid EDTA, a common
chelator found in many water "conditioners."
I didn't just know this stuff, such as it
is, though I hope I've sketched it clearly:
I've been reading Z. Cheng and W. Oldewage-Theron,
"Food fortification to prevent and control
iron deficiency" in the African Journal of Food & Nutritional
Sciences,-- I eagerly await every issue you can be
sure! Skim the details concerned with human
anemia, read the intro and conclusion and
concentrate on the solubility, bio-availability
and stability of NaFe(III)EDTA in solution.
That's the aspect that directly relates to
aquarium water.
Other investigators working with anemic lab
rats found that dissolved tannins did not
affect the bio-availability of NaFe(III)EDTA,
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 using this 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
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, Fe2O3) 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., 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."
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.
