# Phophate inhibits Fe uptake??



## Laith (Sep 4, 2004)

*Phosphate inhibits Fe uptake??*

When I was looking for more info on Proper pH to respond to a question on another thread, I found the following FAQ (http://www.aquariumpharm.com/articles/propPH.asp) about it:

"Can I use PROPER pH with aquarium plants?

No. Phosphate buffers interfere with the metabolism of trace elements such as zinc and iron.* This causes aquatic plants to yellow and die from trace element deficiencies.

*Bennet, W. F. 1993. Nutrient Deficiencies and Toxicities in Crop Plants. APS Press, Minnesota."

I think this is the first I've heard of phosphate inhibiting Fe uptake... Anyone else know about this?


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## NE (Dec 10, 2004)

I have heard that Fe and phosphate cold create strong bindings to each other, but this is just rumours and i dont know the real fact of it.


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## Laith (Sep 4, 2004)

I know about Fe and Phosphate binding together (if in high enough concentrations) and precipitating out of solution.

But here it says "interferes with the metabolism". The use of the word metabolism makes it sound like the phosphate is actually inhibiting plants' capacity to uptake Fe and Zinc.


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## Edward (May 25, 2004)

Hi Laith
Shouldn't we rename this thread from *Phosphate inhibits Fe uptake?* to *Phosphate Buffer inhibits Fe uptake?* 
This is about a pH buffer product that detoxifies metals. This detoxification is the reason why plants turn white and die. It does make the trace elements useless to plant. 
Also, I would like to point out that we should not be using pH buffers in our aquariums. It is a quick fix that may temporarily bandage the problem but it will not cure it.

Thank you
Edward


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## DonaldmBoyer (Aug 18, 2005)

Phosphate oxidizes the iron into an unusuable form of Fe3+ which plants can't use in their metabolism.


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## Laith (Sep 4, 2004)

So I'm just misinterpreting the word "metabolism" in the original link. What they're actually talking about is a chemical reaction outside of the plant that creates a form of Fe that is unavailable or less available to plants. Which is the standard reason of not mixing phosphate and Fe together...

Let me try again to explain: I interpreted the link to mean that the phosphate was having an effect on the plants biological capability to uptake Fe and Zinc, not that the phosphate was binding with the Fe outside of the plant. I don't know if I can explain the issue any better!

By the way, I'm not too sure that the phosphate buffer is what actually detoxifies metals. As far as I understand these products the phosphate buffer is only used as a method to buffer pH. The detoxifying of metals in water conditioners is done by something else since most conditioners do not contain phosphate.

And yes, phosphate buffers are not necessary and, in general, should be avoided. All this came up because someone had a question and was using a phosphate buffer, Proper pH. In trying to find out what exactly this was, I found the FAQ.


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## plantbrain (Jan 23, 2004)

Yep

I think the issue has to do with the FePO4 precipitate, not metabolism.

But this precipitate goes somwhere, the substrate. Then it can be reduce either by the redox/bacteria, or by the roots of plants which add H+'s to reduc things like FePO4 nto Fe2+ and PO4-3.

Regards, 
Tom Barr


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## defdac (May 10, 2004)

In my case it seems the FePO4-precipitate get stuck in the filter where bacteria seems to be dead and the plants can't reach it. My filter media have turned black or very dark brown. The Ehfimech is very dark brown and the Efhisubstrate is completely black. My blue sponge is completely black also. The filter is rust-brown on all surfaces inside.

I have dosed large amounts of a CSM+B-equivalent since I have had problems with pale new shoots, and I usually dose 2,5 ppm PO4 at the same time as the micro.

This has never been an issue before, and my guess is it has something to do with higher pH:s because of misting instead of dissolving CO2. The tank has been really jumpy algaewise (dead nitrificiation?) and I can't get GSA if I try to limit KH2PO4-dosages like I did before.

Most of the plants is covered with thread algae and new shoots are small. I mist until my fish gasps at the surface via a spraybar that runs along the back side of the tank by the substrate (3-6 bubbles/sec in a 300-litre tank) and the growth has been terrible.

Something is very wrong.


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## JLudwig (Feb 16, 2004)

Heh, this subject keeps coming up again and again. I think I have a way to cover all bases, at least from a pseudo science standpoint... 

First, to prevent excess precipitation, you can dose the same molar quantities of PO4(2-) and Fe(2+). This ensures that in the worse case of everything precipitating, you aren't building an excess of one or the other. If I dose 2-to-1 PO4 and the reaction happens fast, I'll always have an excess of PO4 waiting around to react the next time I dump iron in the tank.

Space the dosing out time-wise, I do macros on change day, iron daily. This way on change day there's an excess of PO4 which the plants can soak up, and assuming they use some PO4 during the week, there's an excess of iron by the end of the week. The reaction rate between the two effectively controls the level in the tank, you can up both ammounts but that's shooting yourself in the foot to a certain extent as that will speed up the reaction.

Now, what about the precipitate. As Tom points out, I'm not sure the plants can't use it, you just need the correct conditions - a deep substrate with peat or some kind of high-CEC material at the bottom. Then the plants can get at the precipitate. I'm also starting to favor just topping stem plants to allow them to develop a more extensive root systems which can suck this stuff up.

Cheers,
Jeff


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## averater (Dec 14, 2004)

i've counted a bit on solubility of ironphosphates (both fe+++ and fe++)
the solubility for iron (2+) in an 1ppm phosphate solution is at:
pH5 ~300ppm
pH6 ~10ppm
pH7 ~1ppm
pH8 ~0.1ppm

in an 100ppm PO4 solution you'll get these solubilities:
pH5 ~10ppm
pH6 ~0.5ppm
pH7 ~0.05ppm
pH8 ~0,005ppm

the calculations are a bit complicated so i dont include then here but i can present them if anyone want me to.

iron (3+) will at ph 6-8 and PO4<10ppm form Fe(OH)3 instead of FePO4, so there PO4 isn't a problem.


(all numbers ar heavily rounded, if anyone can please check for miscalculations, constants from "SI chemical data", Aylward, Findlay (2002).)


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## defdac (May 10, 2004)

*Fe(OH)3 instead of FePO4*

Hej averater, kul att du hittat hit =)



> on (3+) will at ph 6-8 and PO4<10ppm form Fe(OH)3 instead of FePO4, so there PO4 isn't a problem.


Interesting. So do you think my brown-black Ehfisubstrate is due to some kind of Fe(OH)3?
A hint I guess would be that bleach didn't work to get it to normal color, but citric acid did:
http://www.defblog.se/picture/1648.html
(tried to include it with a img-tag but it didn't work)

With my extremely limited chemical knowledge anything OH would be ripe to be "nuked" with some acid? The process goes fairly fast, within an hour the color have gone from dark to white.


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## BryceM (Nov 6, 2005)

Nice info averater! So if this is true, then Fe solubility at around pH 6.0 is somewhere between 1 and 10 ppm for around 2 ppm of PO4. This seems to indicate to me 2 things:

1) Worrying about precipitation probably isn't all that big of a deal in the relatively dilute environment of the aquarium.

2) All of this worry about Fe & PO4 probably started when people tried to mix up a fertilizer solution with both macros & micros.

Maybe someone could set up 2 identical tanks and dose micros & macros on the same day vs alternating days in the other tank. My guess is that there would be no apparent difference, as long as the ferts were added a few minutes apart to allow for even mixing.


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## dennis (Mar 1, 2004)

Sort of...it depends on the valence of the Fe. Fe3+(OH)3 is stable up to pH 3-4. Fe2(OH)2 is stable up to about pH6. This means that at pH lower than the a for mentioned levels, the Fe does not become available as easily obtained Fe2+ and Fe3+ cations. In our aquariums, this is OK as the substrate should have pH levels low enough to help us. In nature, this is bad as high levels of free Fe2+/3+ are extreemly toxic (think red coal mine run off in the Appalachians).

The percipitate we see when adding Fe is due to an enthusiastic attraction for bicarbonate ions. Once again though, the percipitate falls to the substrate and will become available to the roots via ReDox(I think).

All right, theres the current limit to my knowledge...


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## JLudwig (Feb 16, 2004)

guaiac_boy said:


> 1) Worrying about precipitation probably isn't all that big of a deal in the relatively dilute environment of the aquarium.


I disagree with your two points. There is a noticable clouding of the water at least in my tanks, even when the two are dosed a day apart. It could be the gluconate fueling a bacteria bloom but something reacting/decomposing, it does happen outside of the fert bottle.

Second point I would like to make is that we, in general, add way too much iron to our tanks. We don't understand this chemistry and our ability to deliver iron to plants is not effective. There's no reason we should need to add anywhere near the level of iron to our tanks as phosphate for example, since one is a macro and the other is a micro. We have a lot to learn in this area.

averater, there are enough scientists on the board, I would love a detailed chemistry model for that reaction. I threw together a quick matlab model assuming some linear kinetics and precip irreversible but eq constants would be much better. Perhaps what I'm seeing in my tank is the breakdown of the chelating agent. Even so I still add too much iron 

Jeff


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## JLudwig (Feb 16, 2004)

dennis said:


> The percipitate we see when adding Fe is due to an enthusiastic attraction for bicarbonate ions. Once again though, the percipitate falls to the substrate and will become available to the roots via ReDox(I think).


Dennis I have a book on groundwater remediation at home, I'll get your a citation, but you're basically right - several bacteria use iron as a reducing agent in anaerobic conditions. I just question whether we actually attain this in an aquarium. Our substrate is relatively coarse (not like a SW deep sand bed with 1mm sand) and I would imagine that we might have a difficult time depleting the substrate of oxygen. If there's O2 around (and plant roots will release O2 IIRC) that reduction is a no-go also you can form H2S quite easily which will kill fish. With a soil substrate I'd believe it's possible to get these conditions, but with Flourite? If it is a reducing environment we can save a lot of money and skip the chelating agent.

Jeff


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## averater (Dec 14, 2004)

here's the detailed description for my calculations:

first constants:
H3PO4
Ka1=10^-2,15
Ka2=10^-7,20
Ka3=10^-12,3

Kw=10^-14

solubility products:
Fe(OH)2:4,1*10^-15
Fe3(PO4)2:1*10^-36
Fe(OH)3:2*10^-39
FePO4:4*10^-27

[P tot]=[H3PO4]+[H2PO4-]+[HPO4--]+[PO4---]
Ka1=[H2PO4-]*[H+]/[H3PO4]
Ka2=[HPO4--]*[H+]/[H2PO4-]
Ka3=[PO4---]*[H+]/[HPO4--]

reararanging gives:
[H3PO4]=[P tot]/(1+Ka1/[H+]+Ka1*Ka2/[H+]^2+Ka1*Ka2*Ka3/[H+])
inserting into the equillibrum equations gives:
[PO4---]=(insert h3po4 here)*Ka1*Ka2*Ka3/[H+]^3

i know this equation easily can be simplifyed but since i made these calculations in a simple spreadsheet i made everything in these kind of small steps...

[OH-]=Kw/[H+]

Fe(OH)2:
4,1*10^-15=[Fe++]*[OH-]^2
Fe3(PO4)2:
1*10^-36=[Fe++]^3*[PO4---]^2
Fe(OH)3:
2*10^-39=[Fe+++]*[OH-]^3
FePO4:
4*10^-27=[Fe+++]*[PO4---]

the calculated values for [Fe] is the maximum solubility, so of course the lovest value (from FeOH or FePO4) would be the maximum solubility.

the values for Fe++ and Fe+++ must be separated sinc they are different kind of ions.

i understand a lot of the parts in these equations are really basic stuff but i rather include to much than too little.

i love the chemistry part of all this and thats what i'm currently studing (or should be studing instead of writing here) but when it comes to biology i'm pretty much a newbee.


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## BryceM (Nov 6, 2005)

I agree that we probably have a poor understanding of the biochemistry involved with iron uptake. We add tons more iron than the plants really need, but only a small portion is available for uptake. I know if I cut back though that my plants quickly show signs of chlorosis. A little more Flourish Fe and they recover in a couple of days.

Is the cloudiness a result of a precipitate, a micro algae bloom, or something else? I'm sure I don't know for certain. It may very well be coming out of solution. Something tells me that uptake in the root zone is more important for plants in the wild than in our tanks.


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## jude_uc (Feb 7, 2006)

The reason you see a precipitate when you use flourish iron, but not when you use CSM+B is because CSM+B uses a chelating agent which keeps the iron in solution. My understanding is that at standard pH values, binding with the chelator is favored over being unbound (and available for precipitation). Flourish iron uses gluconate which forms a weak complex. A large part of Seachem's marketing for gluconate is that although the binding is much weaker, the iron is in the proper form. Of course, that has to be balanced with whether or not you have so much phosphate that it will all precipitate out. When I had high phosphates ( > 5ppm), all flourish iron would precipitate. The CSM+B wouldn't.

-Adam


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## dennis (Mar 1, 2004)

averater,

That chemistry is way over my head at the moment. At what levels of PO4 will we have issues with Fe precititation?


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## averater (Dec 14, 2004)

first there are two kinds of iron that must be separated Fe++ and Fe+++, i've heard but never found any published document to confirm it that plants only wants Fe++ and not Fe+++.

about Fe+++: it doesnt matter how much phosphate you add the iron solubility wont be affected.

about Fe++: here phosphtes might be a problem and stuff get a lot more complicated.
first, if you have youre iron chelated, it cant bind with the phosphates and that iron can be left out of the equations (kind of... unless you have very high pH).
for unchelated iron the effects of po4 is much more drastic and you can only have 0.00001ppm po4 if you want 1ppm fe++ thats unchelated.
if you want 1ppm of po4 you'll only be able to have 0,0005ppm fe++.

this shows the importanc of chelates and that this talk about 1ppm of iron is a lot of iron.


to completely eliminate the problem: use chelated iron.

the only place these solubilites is an issue i think is in fertilizer solutions. those should be acidic and not a mix including both po4 and iron.


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## dennis (Mar 1, 2004)

That's what I though, although I was unaware of the PO4 differences between the di- and trivalent Fe. So in the end, the PO4 amounts in our tanks have no effect on Fe precipitation? That means the cloudyness isuues are strictly bicarbonate based, right?


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## plantbrain (Jan 23, 2004)

averater said:


> first there are two kinds of iron that must be separated Fe++ and Fe+++, i've heard but never found any published document to confirm it that plants only wants Fe++ and not Fe+++.
> 
> about Fe+++: it doesnt matter how much phosphate you add the iron solubility wont be affected.
> 
> ...


This concurs with everything I know.
The plants can use either Fe2+ or Fe3+, they reduce it on the cell's surface to Fe2+ (or Fe2+ as a free ion-> as we can see, this is a very rare situation, thus plants need remove the chelator on Fe or Fe3+ and then reduces it at the cell 's surface.), or the plant can remove the chelater at the cell's surface and take in the Fe and then the plant has it's own internal chelator once inside the plant.

Regards, 
Tom Barr


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## JLudwig (Feb 16, 2004)

If plants can use Fe 3+ readily, and it has a relatively high solubility, why can't I just use rust as a substrate and establish equilibrium levels of Fe 3+ and be done with things? Why bother with a chelating agent at all if Fe+++ is all we need? And if Fe 2+ needs a very strong chelating agent to prevent possible side reactions, is it still bioavailable? Enzymes don't change the thermodynamics, they only change the kinetics.

Averater, I'll get to the math sometime hopefully soon, thanks for the eq constants. My view is that we're interested in free Fe++. The chelating agent and Fe++ are in a pH dependent equilibrium, so if Fe++ and PO4-- is a very fast reaction with a high eq. constant, its essentially irreversible:

A<->B->C

Where A is the chelated iron, B is free Fe++, and C is a precip.

Jeff


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## averater (Dec 14, 2004)

JLudwig said:


> If plants can use Fe 3+ readily, and it has a relatively high solubility, why can't I just use rust as a substrate and establish equilibrium levels of Fe 3+ and be done with things?


the solubility of Fe+++ is very low compared to Fe++



JLudwig said:


> Averater, I'll get to the math sometime hopefully soon, thanks for the eq constants. My view is that we're interested in free Fe++. The chelating agent and Fe++ are in a pH dependent equilibrium, so if Fe++ and PO4-- is a very fast reaction with a high eq. constant, its essentially irreversible:
> 
> A<->B->C
> 
> ...


i think i got eq. constants for edta with iron to so it should be possible to make an eq.equation that includes edta of some concentration. ill see.

M + nL <-> ML
K = [ML]/[M][L]^n

log(K) (Fe++ + EDTA----) = 14,30
log(K) (Fe+++ + EDTA----) = 25,1

i dont have pka values for edta here, but from an unreliable source (internet  ):
pka1:1,70
pka2:2,60
pka3:6,30
pka4:10,60

i dont have time now but i probably calculate solubilites for fe++ and fe+++ with edta tomorrow.


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## plantbrain (Jan 23, 2004)

JLudwig said:


> If plants can use Fe 3+ readily, and it has a relatively high solubility, why can't I just use rust as a substrate and establish equilibrium levels of Fe 3+ and be done with things?
> 
> Jeff


Because plants will do better at high light and growth rates with foliar application. We can supply all the needed Fe from rust in non CO2 tanks and tanks that have slow growth rates.

You can make this arguement for most nutrients.
As the growth rates increase, the hydrophytes can have the option of the foliar uptake and the root uptake.

The difference between the energy required to remove the chelator is less and pH dependent than reducing straight rust.

I'm not sure it's really that big of an issue for roots, but for the foliar uptake, I'd say it is.

Different plant organs have different abilities and sources and the reality is that there is a little of both no matter what unless you do a lab set up for this specific question.

Regards, 
Tom Barr


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## averater (Dec 14, 2004)

i've calculated some more on iron solubilites with edta.

the solubilites are still low and gets higher with decreasing ph.

at ph 6(or lower) it's possible to have 1ppm of both iron(II) and PO4

to make some kind of conclusion: po4-levels of more than 5ppm or so does affect the solubilite of iron.


offtopic: i made a test today where i tried to dissolve 100mg FeCl2 in my 3l nano tank. I had to take up the undissolved remains after a while since the water turned wery brown (so aparantly my po4-levels cant be very high).
(i had some kind of deficiency)


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## dennis (Mar 1, 2004)

I thought Fe(II) was toxic to plants and animals, especially at lower pH's (<4?)? Ironically, isn't low pH's the only enviroment Fe(II) is available before it starts to oxidize?


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## plantbrain (Jan 23, 2004)

Thanks averator for the stats, useful stuff.



dennis said:


> I thought Fe(II) was toxic to plants and animals, especially at lower pH's (<4?)? Ironically, isn't low pH's the only enviroment Fe(II) is available before it starts to oxidize?


Fe is most certainly toxic as are many metals if you have enough of them, same for Mn, Zn, Cu.

While the forms of Chelators varies, from ETDA(Various brands, most noteably, CMS plantex), DTPH(Tropica Master grow) and Fe-Gluconate complex (Flourish), the end result is fairly similar and difficult to tell the differences in most cases.

Some plants have a nicer sheen using TMG for example, likely due to the higher copper content.

Some would like to argue that gluconate is less tightly bounded and therefore easier for the plant to use than say ETDA.

I think it matters little in the larger scope of things.
Testing for Fe is a waste of time simply put.

we can define the CO2 tank pH's to 5.8-7.2 or so for a range, the Flourish and the TMG have the best pH ranges for optimal use for the additives.

The hard part is teasing apart if the Fe alone is responsible and the chelators/complex is the key, or is it Mn, Cu etc ratios, more, less etc.

Hard to say if things are limiting, but adding more and doing a ratio based of the totals can help give some insight.

Regards, 
Tom Barr


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## dennis (Mar 1, 2004)

plantbrain said:


> T
> 
> Fe is most certainly toxic as are many metals if you have enough of them, same for Mn, Zn, Cu.
> 
> While the forms of Chelators varies, from ETDA(Various brands, most noteably, CMS plantex), DTPH(Tropica Master grow) and Fe-Gluconate complex (Flourish), the end result is fairly similar and difficult to tell the differences in most cases.


In our aquarium conditions Fe(II) would never exhist, right. I assume that the typical substrate would never get a pH that low either? So, the only way Fe(II) could exhist in our aquariums is if it is chelated and if it is chelated, then the solubility of the Fe(II) becomes a mute point.

Sorry if I am being redundant but al this is boarderline at teh limit of my understanding


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## defdac (May 10, 2004)

dennis said:


> In our aquarium conditions Fe(II) would never exhist, right. I assume that the typical substrate would never get a pH that low either?


I thought "good" substrates actually maintained a sufficiently low pH below 6.0 e g. ADA-substrates, and one of the reasons a dusting of peat is good in more inert substrates?


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## dennis (Mar 1, 2004)

Do they get down to >4.0 pH though? I would think that the plants transport of O2 down to the roots would prevent pH that low..... I don't know though.


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## plantbrain (Jan 23, 2004)

dennis said:


> Do they get down to >4.0 pH though? I would think that the plants transport of O2 down to the roots would prevent pH that low..... I don't know though.


Dennis, lets get something very clear, aquatic hydrosoils tend, **all of them** after flooding, toward neutrality, where they have plants or not.

The redox is lowered(more negative) when you you flood a soil, but the pH can either go down if alkaline or up if acid.

We can observe and do observe highly organic and also low organic soils doing this, if there is a fair amount of CaCO3 etc, then the pH will be buffered, but the Redox is the process that controls everything in hydrosoils rather than pH directly. Think redox, not pH.

Regards, 
Tom Barr


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## plantbrain (Jan 23, 2004)

dennis said:


> In our aquarium conditions Fe(II) would never exhist, right.


It most certainly does exist and the plants allow it to exist regardless of pH etc at their cell surface(Both foliar and root mediated uptake mechanisms). Bioavailability vs what is in the water column are two very different mechanism and since ultimately we want a nutrient that is easyt to deliver to the plant, should be what you focus on when thinking about Fe in aquatic systems.

We do not add Fe etc to play organic and inorganic chemistry.
We add it for the plants.



> I assume that the typical substrate would never get a pH that low either? So, the only way Fe(II) could exhist in our aquariums is if it is chelated and if it is chelated, then the solubility of the Fe(II) becomes a mute point.
> 
> Sorry if I am being redundant but al this is boarderline at teh limit of my understanding


I think you have this point down.
It can exists in other form than chelated, but only for a brief time frame at the cell surface, then the plant chelates it intenally.

Regards, 
Tom Barr


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## JLudwig (Feb 16, 2004)

Dennis as Tom pointed out I think you're confusing redox potential and pH a little bit. In an oxygen poor environment, a deep fine aquatic soil for example, bacteria will reduce Fe(3+) to Fe(2+), thats a redox reaction.

It's also very important to distinguish between thermodynamic and kinetic limits in a reacting system. The pKa values establish the equilibrium concentration of ions in a system - given infinite time and a closed system....you know the usual frictionless surface in a vacuum that scientists like to theorize about  The solubility curves tell you nothing about the speed of the reaction. You may be able to supersaturate a solution with ions and if the reaction has a high activation barrier it could take a very long time to establish equilibrium. A lot of reactions are very fast, some are not, and I don't have an idea about time scales in aquatic systems, not my area. Just keep in mind the thermodynamic limit may never be actually attained.

Cheers,
Jeff


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