Iron gluconate
Contents:
- Iron gluconate - Aquatic Plants Digest V3 #647
by George Slusarczuk <yurko/warwick.net> (Sun, 15 Nov 1998)
- iron gluconate
by Beverly Erlebacher <bae/cs.toronto.edu> (Mon, 16 Nov 1998)
- Ferrous Gluconate
by "Christopher Coleman" <christopher.coleman/worldnet.att.net> (Tue, 17 Nov 1998)
- EDTA/gluconate
by eworobe/cc.UManitoba.CA (Wed, 18 Nov 1998)
- reducing agents
by Paul Sears <psears/nrn1.NRCan.gc.ca> (Thu, 19 Nov 1998)
- Ferrous Gluconate
by Greg Morin <greg/seachem.com> (Fri, 20 Nov 1998)
- Aquatic Plants Digest V3 #658
by Roger Miller <rgrmill/rt66.com> (Sat, 21 Nov 1998)
- EDTA
by eworobe/cc.UManitoba.CA (Sat, 21 Nov 1998)
- iron gluconate
by Greg Morin <greg/seachem.com> (Mon, 23 Nov 1998)
- iron gluconate
by "Roger S. Miller" <rgrmill/rt66.com> (Mon, 23 Nov 1998)
- EDTA chelation
by Paul Sears <psears/nrn1.NRCan.gc.ca> (Tue, 24 Nov 1998)
- Iron Gluconate
by "Christopher Coleman" <christopher.coleman/worldnet.att.net> (Tue, 24 Nov 1998)
- EDTA vs. Gluconate
by krandall/world.std.com (Tue, 24 Nov 1998)
by George Slusarczuk <yurko/warwick.net>
Date: Sun, 15 Nov 1998
Hello Steve,
Iron gluconate is the salt of gluconic acid, which is a product of
glucose oxidation. Gluconic acid contains one carboxylic acid group and
5 hydroxy groups, so it tends to chelate. Because it is a sugar, I would
imagine that "bugs" love it!
Gluconic acid does exist in nature, as a byproduct of fungal oxidation
of corn syrup and from other sources. If it "finds" an available iron
ion nearby, it may well form a chelate with it. I don't think that in
nature it lasts for very long.
Best,
George
>
> What is iron gluconate(C12H22FeO14, 2H2O)? Does it occur naturally?
>
> What are its merits and demerits as compared to EDTA and DTPA chelated
> Fe?
>
> Steve
>
by Beverly Erlebacher <bae/cs.toronto.edu>
Date: Mon, 16 Nov 1998
> Date: Sun, 15 Nov 1998 06:12:21 -0800
> From: Steve Pushak <teban@powersonic.bc.ca>
>
> What is iron gluconate(C12H22FeO14, 2H2O)? Does it occur naturally?
It's the ferrous salt of gluconic acid... I think gluconate is one
of those intermediary metabolism thingies that is a step in metabolising
glucose to CO2, water and energy. If so, it would be found in all aerobically
metabolising critters, with the probable exception of some weird bacteria.
I could be wrong about the above, but at any rate, it's an innocuous compound
of a type called a sugar acid.
> What are its merits and demerits as compared to EDTA and DTPA chelated Fe?
I don't know if it is a chelator. It is easily consumed by bacteria, et al.
EDTA and DTPA both contain nitrogen.
Perhaps the makers of Flourish can be persuaded to disgorge their reasoning
in using Fe-gluconate in their products.
by "Christopher Coleman" <christopher.coleman/worldnet.att.net>
Date: Tue, 17 Nov 1998
I received the following today from Seachem in response to my
question regarding ferrous gluconate, which is used in their Flourish
and Flourish Iron fertilizers.
>>Christopher Coleman (christopher.coleman@worldnet.att.net) wrote:
>>
>> 1) why should I use gluconate as a 'natural' chelator over
>> EDTA.
>> 2) will ferrous gloconate work as well as an EDTA based
>> chelator in a tank with plain gravel substrate?
>Gluconate actually "complexes" the iron vs the "chelation" found
>with EDTA. The distinction between a complex and chelate is that
>there is no formal bonding in a complex which means that the association
>is not as strong as that found in a chelate. The problem with EDTA
chelation
>is that it is too strong and the plants have a very difficult time
"cracking the
>nut" to get the iron out. Gluconate complexation is not as strong so it is
much
>easier for the plants to extract the needed iron. Both chelates and
complexes
>give an overall charge neutral species. The gluconate also (like EDTA)
helps to
>keep the iron in solution longer than if the iron were free. However
because
>gluconate is not as strong as EDTA with respect to iron association you
will
>see some precipitation with Flourish Iron when dosing the tank... however
the
>key here is that more of the iron in Flourish Iron will be utilized by the
plants
>than will be utilized if EDTA-iron is used. The amount of EDTA-iron that
the
>plants are able to use is so small as to not really be useable. EDTA-iron
complexes
>look better to the consumer because it appears that after adding such a
>product to the tank the iron levels stay up for quite some time whereas
>with Flourish Iron the levels drop off more quickly (looks like the
EDTA-iron
>is more economical doesn't it?). However the Flourish Iron is being
utilized much
>more rapidly (and some of it is precipitating). I think the key here is to
see which
>product works best in your system. I'm confident that you would find that
Flourish
>Iron gives the best response. A third advantage to the gluconate is that it
is a
>reducing agent and so helps to keep the Fe+2 from being oxidized Fe+3 (the
>product also contains other reducing agents to aid in this process as
well). It is
>also my understanding that EDTA iron is actually in the +3 oxidation state.
>Greg Morin
>Gregory Morin, Ph.D. ~Research Director~~~~~~~~~~~~~~~~~~~~
>Seachem Laboratories, Inc. www.seachem.com 888-SEACHEM
>~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~
As an aside, I recently inserted a potassium gluconate tablet consisting of
90mg
equivelent potassium in the root zone of a giant hygro I suspected of having
K
defficiency. Actually, there are two gian hygros with more or less the same
level
of symptom. Only one of them got the supplement. I'll post result if it is
encouraging.
by eworobe/cc.UManitoba.CA
Date: Wed, 18 Nov 1998
The iron in the Fe-EDTA complex is definitely available to plants. To
suggest otherwise is incorrect.
The reason that Fe-gluconate drops off is that the iron precipitates as
Fe+3 rather quickly and the gluconate (being a labile sugar) is digested by
bacteria and algae.
In an oxic environment "reducing agents" will not prevent the oxidation
of ferrous iron.
The charge on chelated iron is irrelevant to iron uptake.
For K deficiency it is preferable to add the K to the water column since
this is one of the elements that aquatic plants require in the water.
dave.
by Paul Sears <psears/nrn1.NRCan.gc.ca>
Date: Thu, 19 Nov 1998
> From: "James Purchase" <jpurch@interlog.com>
> Subject: RE: Iron supplementation
>
> > A third advantage to the gluconate is that it is a reducing agent
> > and so helps to keep the Fe+2 from being oxidized Fe+3 (the product
> > also contains other reducing agents to aid in this process as well).
>
> To which, Dave (eworobe@cc.UManitoba.CA) commented:
>
> > In an oxic environment "reducing agents" will not prevent the oxidation
> > of ferrous iron.
What is more, gluconate will not be an effective reducing agent
for metal ions. Glucose itself can be, but the aldehyde group on
glucose has already been oxidised when we get to the acid. The OH
groups won't reduce metal ions.
HOCH2(HCOH)4COOH Five OH groups and one COOH
>
> Is this difference (E.D.T.A. vs. gluconate) merely one of two equally
> effective alternatives or is one product to be preferred over the other?
I would (and do) go with the EDTA or other chelating agent. The
idea of the chelating agent is to keep the iron in solution, not to
maintain a II oxidation state. The main advantage of gluconic acid is,
I suspect, that it is very cheap. It is used in cleaning solutions
because (under alkaline conditions, at least) it will hang on to metal
ions. EDTA will be a lot more effective.
- --
Paul Sears Ottawa, Canada
by Greg Morin <greg/seachem.com>
Date: Fri, 20 Nov 1998
Hello,
I recently responded to some e-mail correspondences from James Purchase
and Christopher Coleman concerning a discussion on this list about our
Flourish Iron product. I thought it would be helpful to join the list and
keep an eye out for any other questions concerning our products that might
pop up. I realize there is a strong anti-commercial posting clause in the
list charter so let me just say from the outset I intend to abide by that
in the strictest sense.
With that said, I would like to respond to the responses to my posted
letter to Christopher:
(1) EDTA-Iron is in the Fe+3 state, not Fe+2. Every source I have
investigated shows this to be the case (Merck Index, 12th ed, pg 684#4076,
Aldrich Catalog, 1998-99 pg#764 Item # 35,961-0 and others). It is actually
a Fe+3 and Na+ chelate which gives an overall charge neutral species. I
realize one could argue that you could have 2 Fe+2 to one EDTA to give the
same net charge, but apparently this compound does not form (or does not
form readily). There are a lot of compounds that I can draw that are
chemically correct, but that doesn't mean they can be made. Actually there
is an experiment you can do to prove this yourself. I'll even provide
anyone with some samples of these chemicals if you can't get access to
them. To understand this little experiment you need to know that ferrous
(Fe+2) compounds give greenish compounds and that ferrous iron in solution
gives solutions with a greenish cast to it. Ferric (Fe+3) compounds are
reddish to brown (like rust) and give brownish solutions. Okay, in a glass
container, add some ferrous sulfate... as it dissolves the solution turns
from greenish to more of a light brown. It is being oxidized to Fe+3 in the
water. Then add some Prime (or another reducing agent, such as ascorbic
acid (Vitamin C)), about 10 drops or so. The solution then turns back from
brown to a light green. Now add some tetrasodium EDTA, the solution turns
back to brown! This is because the EDTA binds up the iron and at some point
in the process the iron is oxidized to Fe+3. BTW, the tan/brown color of
Flourish Iron is from the gluconic acid, so, no, Flourish Iron is not all
oxidized in the bottle.
(2) With regard to this statement: "In an oxic environment "reducing
agents" will not prevent the oxidation of ferrous iron." I think the answer
in (1) above with the little experiment shows that reducing agents not only
prevent the oxidation of ferrous iron, but actually will reduce ferric back
to ferrous, even with a relatively mild reducing agent such as ascorbic
acid. Also, the complexation of ferrous iron will change the oxidation
potential of the iron. From Shriver, Inorganic Chemistry 2ed, 1994, WH
Freeman & Co., pg 308 "The formation of a more stable complex when the
metal has the lower oxidation number favors reduction and the reduction
potential becomes more positive." In this case the comparison is with the
hexahydrate of Fe+2 vs the gluconate of Fe+2. The gluconate complex is more
stable and therefore favors remaining in the reduced state. This does not
mean that the Fe+2 will _never_ be oxidized while complexed to gluconate,
only that it is going to take longer than if the Fe+2 were not complexed.
The presence of additional reducing agents (in Flourish Iron) makes that
process take just that much longer. Keeping it in the reduced form longer
gives the plants that much more opportunity to absorb it.
(3) With respect to this statement: "The charge on chelated iron is
irrelevant to iron uptake." It may be irrelevant in that both _can_ be
adsorbed, however it is relevant in that Fe+2 is much more easily taken up.
The evidence I offer for this is mainly empirical. When humans are iron
deficient the form the iron supplements that they are given is the ferrous
state. You will never find such supplements in the ferric state. One could
argue that plants are different from humans. Actually in this respect they
are not, but, Ok. To this I cite the Sigma Bio Sciences 1996 Plant Culture
Catalogue pg 68-73 which lists all of the chemical parameters of their
Plant Culture Basal Salt Mixtures and Media. This is basically a list of
different culture media for the hydroponic growth of plants in research. So
it is assumed here the researchers _want_ the plants to grow well :-). Of
the 50 different mixtures listed 86% listed exclusively a ferrous source of
iron (ferrous sulfateĀ7H2O), only 12% listed exclusively a ferric source of
iron (2% had no iron source). Again, I want to reiterate, this does not
mean that plants _can't_ use ferric iron, only that they can use ferrous
much more easily. When a researcher is not interested in evaluating the
iron parameter he/she uses the most readily accessible source, ferrous. The
12% and 2% cited are most likely special mixtures where they do want to
evaluate ferric uptake and utilization. Also, here I will finally cite
something that most people on this list do have available to them and that
is the Aquarium Atlas Volume 2 by Hans Baensch, 1ed, 1993, pg 150. Here it
is stated that "... as a rule, only ferrous (Fe+2) is available for aquatic
plants." This leads into my next response...
(4) The biological relevance of EDTA-Iron vs Gluconate-Iron. From the above
Baensch reference on page 151 is where the idea must be coming from that
EDTA complexed iron is available to plants (he says exactly that). Here I
must disagree with Mr. Baensch. The best way to look at this is to simply
back up and look at what EDTA actually is and does. EDTA is a non-natural
compound that is used to _sequester_ metal ions in living organisms. For
example EDTA is used in severe cases of lead or heavy metal poisoning
[Merck Index 12th ed, pg 593#3559]... i.e. it is so strong it can chelate
these toxic metals and render them _inert_ . Here I would like to quote
from a message I received from James Purchase which is relevant: "On the
topic of E.D.T.A and its dangers, I came across just such a reference in
the literature - but it was dealing with E.D.T.A. being added to an
aquarium by itself (in attempt to bind nutrient ions already present).
Caution was urged as overdosing E.D.T.A could lead to the nutrients
actually being sucked out of the plants, leading to collapse. Most of us,
when dealing with E.D.T.A., use it in a form where it has already been
associated with the nutrient ions." This is the crux of the matter. If EDTA
sucks nutrients (cations) out of plants, how likely is it that if one adds
EDTA with a nutrient (Iron) that the plant is going to pluck it right back?
If the EDTA can pull the nutrients out, then clearly the EDTA is the
stronger of the two in this tug of war. Now on the other side of the
coin... Gluconate is a naturally occurring, thus biological systems are
going to be equipped with the tools necessary to metabolize/interact with
it; they are not equipped to metabolize/interact with EDTA.The natural
result of such interactions will be the release of its complexed iron. And
again, I come back to the medical approach: What do physicians give
patients that are low in iron, do they give them EDTA-Iron? No. They give
them ferrous gluconate (which I believe Beverly Erlebacher was eluding to
in her message regarding ferrous gluconate as being a cheap substitute for
aquarium chemicals).
(5) With respect to the statement that Flourish Iron simply "oxidizes and
precipitates as rust": although I can't rule out the possibility that some
of it is doing just that, the only precipitates we have observed are
ferrous compounds (probably carbonate), this is based on their green
coloring (all ferrous compounds have a green coloring). Even though
Flourish Iron is targeted as a supplement for the stems and leaves,
whatever material that does fall out before being adsorbed by the stems and
leaves can still be adsorbed by the root system of the plants.
(6) With respect to Mortimer Snerd's query about massive algae outbreaks
after using Flourish and Flourish Iron with each water change: The Flourish
Iron definitely is not the source of the problem here. It is formulated
such that constant dosing will not lead to unwanted organic build up that
can contribute to algae growth. This product was actually an outgrowth of a
difficulty many people had with the initial release of Flourish. Flourish
has the same gluconate iron and to keep their iron up, people would dose
quite often. The problem was that they were adding too much of all of the
other stuff in Flourish that isn't consumed as rapidly as the iron. This
led to a build up of too much organics. Flourish still has the iron in it,
but you wouldn't want to use Flourish exclusively to maintain an iron
level. We then spun off Flourish Iron so it could be dosed on a separate
schedule without leading to organic buildup. The dosing for Flourish is
based on an assumed average plant load, and it may be that you should cut
back the amount of Flourish you are using because your plant load is either
lower than average or the variety of plants you have don't utilize the
vitamins and amino acids as rapidly as other plants.
(7)
>I think there is this opinion on Seachems part that EDTA is
>two strong ( actually they commented that "The amount of EDTA-iron
>that the plants are able to use is so small as to not really be useable" )
>This is a pretty strong statement,
All I can say is this statement is in line with the known chemical and
biological facts concerning EDTA as I've outlined above. If anyone has any
evidence otherwise, I would be more than happy to take a look at it.
>And I believe there are pH dependencies here clouding any
>blanket statements about either product.
Actually, pH plays no role (in the aquarium) with either EDTA or gluconate,
The pKa's of the 4 carboxylic acids of EDTA are 0.26, 0.96, 2.00 and 2.67
and for gluconic acid it is 3.86 (source Handbook of Biochemistry and
Molecular Biology, 3ed, Vol 1, p. 310 Fasman, Gerald D. Ed, CRC Press 1976)
Since plant tank pH is going to be well above pH 4 all the acids are fully
ionized.
>It also seems to me that if the gluconate form does get into the water
>more readily, that it has the potential to preferentily benefit algae.
I'm afraid I don't understand the reasoning behind this statement. Are you
saying plants are better able to "extract" iron from EDTA-iron when
compared to algae? And that if the iron is more easily accessible then the
algae has the upper hand?
(8)
>While it may be true that only Fe+2 is taken up by plants, the charge on
>the chelated Fe is irrelevant to its uptake.
I'm sorry, I don't follow here. If as you're saying plants only uptake Fe+2
and the chelated iron is Fe+3, then how is the plant going to uptake the
Fe+3?
(9)
>What is more, gluconate will not be an effective reducing agent
>for metal ions. Glucose itself can be, but the aldehyde group on
>glucose has already been oxidised when we get to the acid. The OH
>groups won't reduce metal ions.
Well, actually the gluconate is a reducing agent in the same manner as
ascorbic acid is a reducing agent (although not as strong as ascorbate).
This question is also answer in question (2) above with respect to the
change in oxidation potential of complexed iron. The additional reducing
agents we employ in Flourish Iron also help to keep the iron in the reduced
state. But as I said above, this only prolongs the inevitable, giving the
plants more opportunity to to adsorb the iron.
> The idea of the chelating agent is to keep the iron in solution, not to
>maintain a II oxidation state.
But if you can do both, then wouldn't that be better? Also, keeping it in
the water doesn't matter if the plants can't sufficiently utilize it. Of
course that depends on one's definition of "sufficient."
>The main advantage of gluconic acid is,
>I suspect, that it is very cheap. It is used in cleaning solutions
>because (under alkaline conditions, at least) it will hang on to metal
>ions. EDTA will be a lot more effective.
Actually gluconic acid is more expensive that EDTA. I'm curious, what
cleaning solutions employ gluconic acid? Why would they use it for the
metal chelating ability... EDTA would be a much better choice? ;-)
I'm sorry for the length of this message, but I wanted to be complete and
thorough and to try and answer everyone's questions regarding gluconate
iron as best I could. And if anyone has any cited references that
contradicts anything I've said, I would very much like to take a look at
them. If I find any of my points are in error based on such references I
will of course retract or clarify such points. If anyone can not gain
access to a reference I've cited, I'd be happy to photocopy and mail it to
you. I can't post this kind of stuff on the web due to copyright issues
(with the respective publisher of each source).
Regards,
- -Greg Morin
Gregory Morin, Ph.D. ~Research Director~~~~~~~~~~~~~~~~~~~~
Seachem Laboratories, Inc. www.seachem.com 888-SEACHEM
~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~
by Roger Miller <rgrmill/rt66.com>
Date: Sat, 21 Nov 1998
~
On Sat, 21 Nov 1998, Greg Morin wrote:
> I realize there is a strong anti-commercial posting clause in the
> list charter so let me just say from the outset I intend to abide by that
> in the strictest sense.
The anticommercial clause in the charter is aimed (I think) at avoiding
commerical promotion. Personally, I welcome non-promotional information
and discussion from commercial sources as long as they are clearly
identified as such. Of course the line between information and promotion
is very hard to draw and some people may simply take exception to the
whole idea.
>
> (2) With regard to this statement: "In an oxic environment "reducing
> agents" will not prevent the oxidation of ferrous iron." I think the answer
> in (1) above with the little experiment shows that reducing agents not only
> prevent the oxidation of ferrous iron, but actually will reduce ferric back
> to ferrous, even with a relatively mild reducing agent such as ascorbic
> acid.
I think that the original statement refers to oxidation in the aquarium,
and much of your discussion seems more relevant to oxidation in the
bottle. Certainly reducing agents and complex stability can stabilize
ferrous compounds in sufficiently concentrated solution, but I doubt that
can be extended to the dilute, oxidized and biologically active
environment of a planted aquarium.
> (3) With respect to this statement: "The charge on chelated iron is
> irrelevant to iron uptake." It may be irrelevant in that both _can_ be
> adsorbed, however it is relevant in that Fe+2 is much more easily taken up.
My understanding is that the Fe+2 ion is readily taken up. Ferrous
gluconate is a salt containing Fe+2. In order for the plant to access the
Fe+2 ion in the ferrous gluconate either the plant must take in the entire
salt molecule or the salt must be broken down before the Fe+2 is adsorbed.
The first mechanism isn't specific to the oxidation state of the iron; it
works with ferric chelates as well. The second, unless it happens in the
rhizosphere, will probably just lead to oxidation and precipitation of the
iron, not to plant uptake of the FE+2 ion.
>
> (4) The biological relevance of EDTA-Iron vs Gluconate-Iron. From the above
> Baensch reference on page 151 is where the idea must be coming from that
> EDTA complexed iron is available to plants (he says exactly that).
EDTA-chelated iron (as well as DTPA chelates and a few others) have been
in horticultural use for decades with a record of apparent effectiveness.
Either the chemical industry has foisted a huge hoax on the agricultural,
horticultural and hobby industries and we've all been naively duped or
EDTA chelated iron is biologically useful. I tend to think that EDTA-iron
is useful.
I have no reason at all to believe that ferrous glauconate isn't similarly
useful.
>
> (5) With respect to the statement that Flourish Iron simply "oxidizes and
> precipitates as rust": although I can't rule out the possibility that some
> of it is doing just that, the only precipitates we have observed are
> ferrous compounds (probably carbonate), this is based on their green
> coloring (all ferrous compounds have a green coloring). Even though
> Flourish Iron is targeted as a supplement for the stems and leaves,
> whatever material that does fall out before being adsorbed by the stems and
> leaves can still be adsorbed by the root system of the plants.
It sounds to me like the precipitate you're describing might be found in
the bottle. Ferrous carbonate is the natural mineral siderite, which is
usually brown, not green; maybe a fresh precipitate would be green. I
doubt that you would see iron carbonate in aquarium use, where the iron
would be too dilute and the pH too low to precipitate the carbonate.
I've read that gluconates have become the favored form of complexed
metals in the European industrial community specifically because it does
break down quickly. EDTA in particular breaks down so slowly that its
presence in industrial effluents might effect the availability of metals
in the sediments of receiving waters; gluconates don't last long enough
to do that.
How is ferrous gluconate "targeted" to the leaves and stems?
> (7)
> >I think there is this opinion on Seachems part that EDTA is
> >two strong ( actually they commented that "The amount of EDTA-iron
> >that the plants are able to use is so small as to not really be useable" )
> >This is a pretty strong statement,
>
> All I can say is this statement is in line with the known chemical and
> biological facts concerning EDTA as I've outlined above. If anyone has any
> evidence otherwise, I would be more than happy to take a look at it.
EDTA-iron has been in common use for decades. That looks to me like
pretty strong evidence.
>
>
> >And I believe there are pH dependencies here clouding any
> >blanket statements about either product.
>
> Actually, pH plays no role (in the aquarium) with either EDTA or gluconate,
> The pKa's of the 4 carboxylic acids of EDTA are 0.26, 0.96, 2.00 and 2.67
> and for gluconic acid it is 3.86 (source Handbook of Biochemistry and
> Molecular Biology, 3ed, Vol 1, p. 310 Fasman, Gerald D. Ed, CRC Press 1976)
> Since plant tank pH is going to be well above pH 4 all the acids are fully
> ionized.
My reference puts the pKa of EDTA at 2.008, 2.683, 6.098 and 10.181. These
values would indicate considerable importance for pH on EDTA chemistry.
The values I sited are listed by James Plambeck on the University of
Alberta web site at:
http://c.chem.ualberta.ca/~plambeck/che/data/p0040x.htm
I haven't tried to discuss the differences with Dr. Plambeck.
Virtually every source I've read on iron chelates indicates that the
stability of the complex is pH dependent. For EDTA, in particular the
complex starts to break down at pH over about 6 and it's generally not
useful at pH well over 7. Other chelates are more stable at higher pH
values.
>
> (8)
> >While it may be true that only Fe+2 is taken up by plants, the charge on
> >the chelated Fe is irrelevant to its uptake.
>
> I'm sorry, I don't follow here. If as you're saying plants only uptake Fe+2
> and the chelated iron is Fe+3, then how is the plant going to uptake the
> Fe+3?
I found the original statement a little confusing, too. I interpreted it
to mean that iron in an ionic form is taken up only in its Fe+2 state.
Uncomplexed Fe+3 iron is virtually non-existent under aquarium conditions.
Iron in organic complexes may be either Fe+2 and Fe+3 and the oxidation
state of the iron in the complex isn't important to plant uptake.
There's an excellent British web site that details iron uptake mechanisms
in plants. I liked the document so much that I downloaded it locally, and
I don't have the URL anymore. I posted the URL on this list a few weeks
back and more recently Steve Pushak announced that he provided a link to
that document on his web site.
Very briefly, plants have several mechanisms available for using Fe+3 as
long as the iron is in an organic complex. Information on the site is
also very relevant to the occasional arguments on this list over
availability of iron in laterite.
>
> > The idea of the chelating agent is to keep the iron in solution, not to
> >maintain a II oxidation state.
>
> But if you can do both, then wouldn't that be better?
Probably yes, but certainly that isn't clear-cut.
>
> I'm sorry for the length of this message, but I wanted to be complete and
> thorough and to try and answer everyone's questions regarding gluconate
> iron as best I could.
I can't accept your arguments that EDTA-chelated iron is an ineffective
source of iron. It does appear to me that ferrous glauconate might be a
better choice at pH > 7; other chelated iron products might also be good
alternatives.
My tanks generally run at pH > 7, so maybe next time I see ferrous
gluconate for sale I'll give it a try.
Roger Miller
in pleasantly cool, dry and sunny Albuquerque.
by eworobe/cc.UManitoba.CA
Date: Sat, 21 Nov 1998
> (1) EDTA-Iron is in the Fe+3 state, not Fe+2.
You can complex either ferrous or ferric iron with EDTA since this
chelator is a sexadentate ligand (six charged sites). Chelated iron
will not precipitate out of solution but yes, the Fe:chelator is a CLEAR
yellow-brown color. "One of the valuable properties of EDTA is that it
combines with metal ions in a 1/1 ratio regardless of the charge of the
cation(Skoog and West. Fundamentals of Analytical Chemistry)."
>
> (2) With regard to this statement: "In an oxic environment "reducing
> agents" will not prevent the oxidation of ferrous iron." I think the answer
> in (1) above with the little experiment shows that reducing agents not only
> prevent the oxidation of ferrous iron, but actually will reduce ferric back
> to ferrous, even with a relatively mild reducing agent such as ascorbic
In a biologically active medium saturated with oxygen, ferrous iron is
simply not stable for any length of time (minutes or less).
> (3) With respect to this statement: "The charge on chelated iron is
> irrelevant to iron uptake." It may be irrelevant in that both _can_ be
> adsorbed, however it is relevant in that Fe+2 is much more easily taken up.
No, in countless experiments over many years, research has clearly shown
that chelated Fe:EDTA can deliver adequate amounts of iron for maximal
growth rates. In my own experiments I have found that 0.2 to 0.9 uM
FE:EDTA (0.1 to 0.5 ppm) will suppport relative growth rates of between
75 to 300 (1 to 4 day doubling time).
> This is the crux of the matter. If EDTA
> sucks nutrients (cations) out of plants, how likely is it that if one adds
> EDTA with a nutrient (Iron) that the plant is going to pluck it right back?
This is an interesting paradox that I spent many sleepless nights
thinking about during my PhD years. The paradox, simply stated, is that
EDTA can remove nutrients from solution and cause deficiency, while at
the same time it can provide Fe to the plants. The answer lies in the
relative solubilities of the ions in question. Soluble cations such as Ca
or Mg are taken up as cations and EDTA will prevent this uptake. Fe, on
the other hand, is insoluble in water so the EDTA will prevent its
precipitation. Uptake of Ca:EDTA or Fe:EDTA is about the same, but for Ca
it is much too slow relative to Ca+2 uptake, while for Fe it is much
faster than uptake of the free Fe ion.
Since Fe never occurs unchelated in natural, oxygenated systems, plants are
perfectly capable of absorbing chelated Iron.
> (all ferrous compounds have a green coloring).
Ferrous hydroxide is green and quickly turns to Ferric oxyhydroxide in an
oxic environment.
> All I can say is this statement is in line with the known chemical and
> biological facts concerning EDTA as I've outlined above. If anyone has any
> evidence otherwise, I would be more than happy to take a look at it.
Well, theres about 50 years of hydroponics and aquatic plant research
that clearly shows that Fe:EDTA delivers adequate amounts of Fe for
sustained maximal growth.
> Actually, pH plays no role (in the aquarium) with either EDTA or gluconate,
> The pKa's of the 4 carboxylic acids of EDTA are 0.26, 0.96, 2.00 and 2.67
> and for gluconic acid it is 3.86 (source Handbook of Biochemistry and
> Molecular Biology, 3ed, Vol 1, p. 310 Fasman, Gerald D. Ed, CRC Press 1976)
Pka values for EDTA are approximately 2.0, 2.7, 6.2, and 10.3.
> (8)
> >While it may be true that only Fe+2 is taken up by plants, the charge on
> >the chelated Fe is irrelevant to its uptake.
>
> I'm sorry, I don't follow here. If as you're saying plants only uptake Fe+2
> and the chelated iron is Fe+3, then how is the plant going to uptake the
> Fe+3?
Current thinking suggests that plants attach the Fe:EDTA complex onto
the membrane and remove the Fe ion. How this occurs is not fully
understood but it is certainly an effective uptake system.
dave.
by Greg Morin <greg/seachem.com>
Date: Mon, 23 Nov 1998
- --============_-1300269047==_ma============
Content-Type: text/plain; charset="us-ascii" ; format="flowed"
>EDTA-chelated iron (as well as DTPA chelates and a few others) have been
>in horticultural use for decades with a record of apparent effectiveness.
Does anyone know where I might find the documentation that supports the
record of apparent effectiveness? I just want to be able to take a look at
the studies first hand...
>It sounds to me like the precipitate you're describing might be found in
>the bottle. Ferrous carbonate is the natural mineral siderite, which is
>usually brown, not green; maybe a fresh precipitate would be green. I
>doubt that you would see iron carbonate in aquarium use, where the iron
>would be too dilute and the pH too low to precipitate the carbonate.
Well we haven't observed any precipitation in the bottle. The precipt I was
referring to was seen in a container of fresh water from a planted tank
here, after a day or so there was a _very_ light dusting of light green
powder on the bottom which we presumed might be the carbonate (since the
other counter anions in the water would produce only soluble Fe+2 salts.
The Merck Index (12ed, pg 4087#4089) states that ferrous carbonate is
practically insoluble in water. Since we know that Fe+2 is rapidly oxidized
in the water by itself (within seconds) this would suggest that the
gluconate is keeping the iron in the Fe+2 state long enough for an anion
exchange (carbonate for gluconate) process to yield the precipitate we're
seeing; if it is around long enough that the precipt is not noticeable
until a day later, then this would suggest the ferrous gluconate is around
long enough to be utilized in the Fe+2 state (i.e. at least a day).
>How is ferrous gluconate "targeted" to the leaves and stems?
By keeping the iron solubilized the leaves and stems have a greater
opportunity to absorb the iron than do the roots. This "opportunity" is
based on two factors: (1) The leaves and stems represent a much higher
percentage of the plant's surface area than do the roots and (2) over a
given unit of time a greater volume of water passes by the leaves and stem
(unless an UGF is employed, in which case the water flow discrepancy is not
as great).
At this point I would like to publicly retract/modify my initial statement
concerning the near inability of plants to utilize EDTA-Iron. Based on the
responses I've received to my posting and further reading, I now suggest
the following mechanisms.
The success the horticulture industry has seen with EDTA-Iron has been in
its capacity as a fertilizer/root supplement. Based on what we know about
plant roots and their ability to release chelates (phytosiderophores) and
extract iron from their chelates it would be reasonable to assume the
plants can extract the iron from the EDTA-Iron chelate as well. Also, since
we know that the iron in the soil is by far mainly in the Fe+3 state, the
fact that EDTA-Iron is in the Fe+3 state as well has little impact on the
benefit of EDTA-Iron vs normal soil extraction. Although I think we all
agree that the plant would much "prefer" to use Fe+2 if available.
In an aquatic environment the same rules apply, the benefit that people
see from using EDTA-Iron supplements is from the capacity of those plant's
roots to extract the iron from the EDTA. The leaves and stems do not have
the chemical machinery to extract iron from a chelate. This last statement
is based on conjecture on my part in that (a) I have seen no evidence to
the contrary and (b) there would be no evolutionary advantage to leaves and
stems producing siderophores to extract iron as those siderophores would
simply be washed away. If they don't produce siderophores then they have no
need for chemical machinery to readsorb chelates and extract the metal.
The advantage to Flourish Iron is that it supplies stabilized ferrous iron.
The leaves and stems _can_ use the ferrous iron and I suppose they can use
ferric iron as well, just not nearly as well as ferrous. So if you have
plants that have a great capacity to adsorb iron from their stems and
leaves, Flourish Iron is what you would want to use. If that capacity is
not so great or you're satisfied with the level of iron absorbed through
the roots in relation to the plants growth and appearance, then EDTA-Iron
will be suitable (although gluconate iron is still going to be _better_
than EDTA-Iron with respect to charge and ease of absorption). So the short
summary is: EDTA-Iron suitable for roots, gluconate iron good for leaves,
stems and roots.
>> Actually, pH plays no role (in the aquarium) with either EDTA or gluconate,
>> The pKa's of the 4 carboxylic acids of EDTA are 0.26, 0.96, 2.00 and 2.67
>> and for gluconic acid it is 3.86 (source Handbook of Biochemistry and
>> Molecular Biology, 3ed, Vol 1, p. 310 Fasman, Gerald D. Ed, CRC Press 1976)
>> Since plant tank pH is going to be well above pH 4 all the acids are fully
>> ionized.
>
>My reference puts the pKa of EDTA at 2.008, 2.683, 6.098 and 10.181. These
>values would indicate considerable importance for pH on EDTA chemistry.
>The values I sited are listed by James Plambeck on the University of
>Alberta web site at:
>
>http://c.chem.ualberta.ca/~plambeck/che/data/p0040x.htm
> Virtually every source I've read on iron chelates indicates that the
> stability of the complex is pH dependent. For EDTA, in particular the
> complex starts to break down at pH over about 6 and it's generally not
> useful at pH well over 7. Other chelates are more stable at higher pH
> values.
Well, we're both right. I listed the pKa's for the carboxylic acids only.
The 6.0 and 10.1 pKa values are for the amines. However, I think you got
the pH importance backwards ;-) As the pH goes below 6 the interaction of
one of the amines becomes less and less important (as it becomes
increasingly ionized). Going above 6 decreases the ionization of the first
amine resulting in a stronger interaction of the amine with whatever metal
is chelated. You would need to go above a pH of 10 for maximum chelation
(so that both amines are unionized) but there is really no need to do that.
(So I guess I just talked you out of trying a ferrous gluconate product!
;-)
>Uncomplexed Fe+3 iron is virtually non-existent under aquarium conditions.
Why? If there is no EDTA present in the aquarium then what happens to the
Fe+3 if it is non-existant? I initially had thought this too... but I see
in the Merck Index that there are several soluble ferric compounds, so that
would imply the Fe+3 remains solubulized.
>> (1) EDTA-Iron is in the Fe+3 state, not Fe+2.
>
>You can complex either ferrous or ferric iron with EDTA since this
>chelator is a sexadentate ligand (six charged sites). Chelated iron
>will not precipitate out of solution but yes, the Fe:chelator is a CLEAR
>yellow-brown color. "One of the valuable properties of EDTA is that it
>combines with metal ions in a 1/1 ratio regardless of the charge of the
>cation(Skoog and West. Fundamentals of Analytical Chemistry)."
Based on the known properties of EDTA, it's not unreasonable to have
_assumed_ EDTA can complex Fe+2... however, based on everything I've read
this complex simply does NOT form. I don't know _WHY_ it doesn't form, but
the fact that remains is that it just doesn't. If anyone has any reference
that specifically refers to an Iron(II) EDTA complex I would be very
interested to see it
>In a biologically active medium saturated with oxygen, ferrous iron is
>simply not stable for any length of time (minutes or less).
Even if that ferrous iron is complexed by a reducing agent and there is a
suitable amount of reducing compounds that have been added to the water?
Are you saying that these have NO EFFECT at all on the average length of
time that the Fe+2 will remain in the ferrous state before being oxidized?
Based on what we've seen, Flourish Iron will keep the iron in the ferrous
state for at least a day.
>No, in countless experiments over many years, research has clearly shown
>that chelated Fe:EDTA can deliver adequate amounts of iron for maximal
>growth rates. In my own experiments I have found that 0.2 to 0.9 uM
>FE:EDTA (0.1 to 0.5 ppm) will suppport relative growth rates of between
>75 to 300 (1 to 4 day doubling time).
This sounds very interesting... could you provide references to the
research you cite? Please tell me this was a root supplement only? ;-) ;-)
>Well, theres about 50 years of hydroponics and aquatic plant research
>that clearly shows that Fe:EDTA delivers adequate amounts of Fe for
>sustained maximal growth.
Again, I would be very interested in being pointed to some of the relevant
literature here...
Regards,
- -Greg Morin
Gregory Morin, Ph.D. ~Research Director~~~~~~~~~~~~~~~~~~~~
Seachem Laboratories, Inc. www.seachem.com 888-SEACHEM
~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~
- --============_-1300269047==_ma============
Content-Type: text/enriched; charset="us-ascii"
>EDTA-chelated iron (as well as DTPA chelates and a few others) have
been
>in horticultural use for decades with a record of apparent
effectiveness.
Does anyone know where I might find the documentation that supports the
record of apparent effectiveness? I just want to be able to take a look
at the studies first hand...
>It sounds to me like the precipitate you're describing might be found
in
>the bottle. Ferrous carbonate is the natural mineral siderite, which
is
>usually brown, not green; maybe a fresh precipitate would be green.
I
>doubt that you would see iron carbonate in aquarium use, where the
iron
>would be too dilute and the pH too low to precipitate the carbonate.
Well we haven't observed any precipitation in the bottle. The precipt I
was referring to was seen in a container of fresh water from a planted
tank here, after a day or so there was a _very_ light dusting of light
green powder on the bottom which we presumed might be the carbonate
(since the other counter anions in the water would produce only soluble
Fe+2 salts. The Merck Index (12ed, pg 4087#4089) states that ferrous
carbonate is practically insoluble in water. Since we know that Fe+2 is
rapidly oxidized in the water by itself (within seconds) this would
suggest that the gluconate is keeping the iron in the Fe+2 state long
enough for an anion exchange (carbonate for gluconate) process to yield
the precipitate we're seeing; if it is around long enough that the
precipt is not noticeable until a day later, then this would suggest
the ferrous gluconate is around long enough to be utilized in the Fe+2
state (i.e. at least a day).
>How is ferrous gluconate "targeted" to the leaves and stems?
By keeping the iron solubilized the leaves and stems have a greater
opportunity to absorb the iron than do the roots. This "opportunity" is
based on two factors: (1) The leaves and stems represent a much higher
percentage of the plant's surface area than do the roots and (2) over a
given unit of time a greater volume of water passes by the leaves and
stem (unless an UGF is employed, in which case the water flow
discrepancy is not as great).
At this point I would like to publicly retract/modify my initial
statement concerning the near inability of plants to utilize EDTA-Iron.
Based on the responses I've received to my posting and further reading,
I now suggest the following mechanisms.
The success the horticulture industry has seen with EDTA-Iron has been
in its capacity as a fertilizer/root supplement. Based on what we know
about plant roots and their ability to release chelates
(phytosiderophores) and extract iron from their chelates it would be
reasonable to assume the plants can extract the iron from the EDTA-Iron
chelate as well. Also, since we know that the iron in the soil is by
far mainly in the Fe+3 state, the fact that EDTA-Iron is in the Fe+3
state as well has little impact on the benefit of EDTA-Iron vs normal
soil extraction. Although I think we all agree that the plant would
much "prefer" to use Fe+2 if available.
In an aquatic environment the same rules apply, the benefit that
people see from using EDTA-Iron supplements is from the capacity of
those plant's roots to extract the iron from the EDTA. The leaves and
stems do not have the chemical machinery to extract iron from a
chelate. This last statement is based on conjecture on my part in that
(a) I have seen no evidence to the contrary and (b) there would be no
evolutionary advantage to leaves and stems producing siderophores to
extract iron as those siderophores would simply be washed away. If they
don't produce siderophores then they have no need for chemical
machinery to readsorb chelates and extract the metal.
The advantage to Flourish Iron is that it supplies stabilized ferrous
iron. The leaves and stems _can_ use the ferrous iron and I suppose
they can use ferric iron as well, just not nearly as well as ferrous.
So if you have plants that have a great capacity to adsorb iron from
their stems and leaves, Flourish Iron is what you would want to use. If
that capacity is not so great or you're satisfied with the level of
iron absorbed through the roots in relation to the plants growth and
appearance, then EDTA-Iron will be suitable (although gluconate iron is
still going to be _better_ than EDTA-Iron with respect to charge and
ease of absorption). So the short summary is: EDTA-Iron suitable for
roots, gluconate iron good for leaves, stems and roots.
>> Actually, pH plays no role (in the aquarium) with either EDTA or
gluconate,
>> The pKa's of the 4 carboxylic acids of EDTA are 0.26, 0.96, 2.00 and
2.67
>> and for gluconic acid it is 3.86 (source Handbook of Biochemistry
and
>> Molecular Biology, 3ed, Vol 1, p. 310 Fasman, Gerald D. Ed, CRC
Press 1976)
>> Since plant tank pH is going to be well above pH 4 all the acids are
fully
>> ionized.
>
>My reference puts the pKa of EDTA at 2.008, 2.683, 6.098 and 10.181.
These
>values would indicate considerable importance for pH on EDTA
chemistry.
>The values I sited are listed by James Plambeck on the University of
>Alberta web site at:
>
>http://c.chem.ualberta.ca/~plambeck/che/data/p0040x.htm
<excerpt>Virtually every source I've read on iron chelates indicates
that the
stability of the complex is pH dependent. For EDTA, in particular the
complex starts to break down at pH over about 6 and it's generally not
useful at pH well over 7. Other chelates are more stable at higher pH
values.
</excerpt>
Well, we're both right. I listed the pKa's for the carboxylic acids
only. The 6.0 and 10.1 pKa values are for the amines. However, I think
you got the pH importance backwards ;-) As the pH goes below 6 the
interaction of one of the amines becomes less and less important (as it
becomes increasingly ionized). Going above 6 decreases the ionization
of the first amine resulting in a stronger interaction of the amine
with whatever metal is chelated. You would need to go above a pH of 10
for maximum chelation (so that both amines are unionized) but there is
really no need to do that. (So I guess I just talked you out of trying
a ferrous gluconate product! ;-)
>Uncomplexed Fe+3 iron is virtually non-existent under aquarium
conditions.
Why? If there is no EDTA present in the aquarium then what happens to
the Fe+3 if it is non-existant? I initially had thought this too... but
I see in the Merck Index that there are several soluble ferric
compounds, so that would imply the Fe+3 remains solubulized.
>> (1) EDTA-Iron is in the Fe+3 state, not Fe+2.
>
>You can complex either ferrous or ferric iron with EDTA since this
>chelator is a sexadentate ligand (six charged sites). Chelated iron
>will not precipitate out of solution but yes, the Fe:chelator is a
CLEAR
>yellow-brown color. "One of the valuable properties of EDTA is that it
>combines with metal ions in a 1/1 ratio regardless of the charge of
the
>cation(Skoog and West. Fundamentals of Analytical Chemistry)."
Based on the known properties of EDTA, it's not unreasonable to have
_assumed_ EDTA can complex Fe+2... however, based on everything I've
read this complex simply does NOT form. I don't know _WHY_ it doesn't
form, but the fact that remains is that it just doesn't. If anyone has
any reference that specifically refers to an Iron(II) EDTA complex I
would be very interested to see it
>In a biologically active medium saturated with oxygen, ferrous iron is
>simply not stable for any length of time (minutes or less).
Even if that ferrous iron is complexed by a reducing agent and there is
a suitable amount of reducing compounds that have been added to the
water? Are you saying that these have NO EFFECT at all on the average
length of time that the Fe+2 will remain in the ferrous state before
being oxidized? Based on what we've seen, Flourish Iron will keep the
iron in the ferrous state for at least a day.
>No, in countless experiments over many years, research has clearly
shown
>that chelated Fe:EDTA can deliver adequate amounts of iron for maximal
>growth rates. In my own experiments I have found that 0.2 to 0.9 uM
>FE:EDTA (0.1 to 0.5 ppm) will suppport relative growth rates of
between
>75 to 300 (1 to 4 day doubling time).
This sounds very interesting... could you provide references to the
research you cite? Please tell me this was a root supplement only? ;-)
;-)
>Well, theres about 50 years of hydroponics and aquatic plant research
>that clearly shows that Fe:EDTA delivers adequate amounts of Fe for
>sustained maximal growth.
Again, I would be very interested in being pointed to some of the
relevant literature here...
Regards,
- -Greg Morin
Gregory Morin, Ph.D. ~Research Director~~~~~~~~~~~~~~~~~~~~
Seachem Laboratories, Inc. www.seachem.com 888-SEACHEM
~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~~
- --============_-1300269047==_ma============--
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by "Roger S. Miller" <rgrmill/rt66.com>
Date: Mon, 23 Nov 1998
On Monday, November 23, 1998, Greg Morin wrote (among other things):
>
> >EDTA-chelated iron (as well as DTPA chelates and a few others) have been
> >in horticultural use for decades with a record of apparent effectiveness.
>
> Does anyone know where I might find the documentation that supports the
> record of apparent effectiveness? I just want to be able to take a look at
> the studies first hand...
Gee, I learned at my mother's knee that when the roses get chorosis, you
add chelated iron. You mean I need a reference?
My current favorite on-line reference is a US Forest Service site:
http://willow.ncfes.umn.edu/fnn_7-97/cult_per.htm#fe
This is one part of a rather large document. I may eventually download
the whole thing (with graphics) and keep it locally so I can study it off
line. There is a small references list at the end of the document if you
want to follow up.
> >How is ferrous gluconate "targeted" to the leaves and stems?
>
> By keeping the iron solubilized the leaves and stems have a greater
> opportunity to absorb the iron than do the roots. This "opportunity" is
> based on two factors: (1) The leaves and stems represent a much higher
> percentage of the plant's surface area than do the roots and (2) over a
> given unit of time a greater volume of water passes by the leaves and stem
> (unless an UGF is employed, in which case the water flow discrepancy is not
> as great).
I'm not sure that (1) is generally correct. Certainly it is for some
plants (many in the frogbit family - hydrocharitaceae - for instance)
but probably not for most plants with fully developed roots (and root
hairs).
All of the information I've read addresses root uptake of Fe. I haven't
read anything about foliar uptake of iron. Anybody have any good
references (I'm really not limited to on-line resources)?
> > Virtually every source I've read on iron chelates indicates that the
> > stability of the complex is pH dependent. For EDTA, in particular the
> > complex starts to break down at pH over about 6 and it's generally not
> > useful at pH well over 7. Other chelates are more stable at higher pH
> > values.
>
>
> Well, we're both right. I listed the pKa's for the carboxylic acids only.
> The 6.0 and 10.1 pKa values are for the amines. However, I think you got
> the pH importance backwards ;-).
I don't know whether or not the pH sensitivity of the iron chelate is
related directly to the behavior of the acid groups. It may be for some
other reason; e.g. at elevated pH the hydroxide complexes may be relatively
more stable than the chelate.
>
> >Uncomplexed Fe+3 iron is virtually non-existent under aquarium conditions.
>
> Why? If there is no EDTA present in the aquarium then what happens to the
> Fe+3 if it is non-existant?
Fe+3 precipitates as the hydroxide and oxyhydroxide (FeO(OH)), which have
very low solubility. Even at concentrations allowed by the solubility of
those compounds, Fe+3 forms complexes with hydroxides and even polymeric
hydroxide chains. Fe+3 would be found as an uncomplexed ion only at
rather low pH. Plants often acidify their rhizosphere and that tends to
break down the hydroxides and hydroxide complexes and make the Fe+3 in the
soil more readily available.
There are a number of excellent references on the behavior of iron and
other metals at trace concentrations under natural conditions - certainly
more good references than I've read. If you email me off the list I
might be able to find one or two around my office.
Roger Miller
by Paul Sears <psears/nrn1.NRCan.gc.ca>
Date: Tue, 24 Nov 1998
> From: Greg Morin <greg@seachem.com>
> Subject: Re: iron gluconate
>
> Well, we're both right. I listed the pKa's for the carboxylic acids only.
> The 6.0 and 10.1 pKa values are for the amines. However, I think you got
> the pH importance backwards ;-) As the pH goes below 6 the interaction of
> one of the amines becomes less and less important (as it becomes
> increasingly ionized). Going above 6 decreases the ionization of the first
> amine resulting in a stronger interaction of the amine with whatever metal
> is chelated. You would need to go above a pH of 10 for maximum chelation
> (so that both amines are unionized) but there is really no need to do that.
> (So I guess I just talked you out of trying a ferrous gluconate product!
> ;-)
EDTA complexes generally have both amine nitrogens and three of the
carboxylates attached to the metal atom. The sixth co-ordination position
on the metal atom is associated with a water molecule. The pH does _not_
have to be very high for that. Quite strongly basic amines co-ordinate
well with transition metal ions at moderate pH, because the bond formed
between the nitrogen lone pair and the metal atom is strong.
- --
Paul Sears Ottawa, Canada
by "Christopher Coleman" <christopher.coleman/worldnet.att.net>
Date: Tue, 24 Nov 1998
Greg Morin wrote:
>That actually falls in nicely with what I said earlier... the
>plants you described lack roots, therefore the leaves and stems
>have taken on the critical characteristics of roots (i.e. iron chelate
>adsorbption). Of course there are always going to be exceptions... I'm
>sure there probably are plants out there that can utilize EDTA-Iron
>through a foliar route... and probably just as many that cannot... so the
>trick is, figuring out which of your plants can and cannot do this. As
>stated elsewhere the best way really is going to be trying both types
>of products (EDTA vs gluconate) and see which one performs best
>in your system.
A few points:
1) Between leaves, stems, and roots, the basic physiology of plants is is
to
allow the roots to serve as _the_ major nutrient uptake mechanism. I
am
even given pause to consider why it is often suggested to dose
potassium
and magnesium into the water column. And I am challanged to find any
physiological structure in the leaves which aid int the uptake of
nutrients
( exception: stomata take in CO2 ).
2) Much attention is paid to CEC values of substates to underscore
the desirability of getting nutrients _into_ the substrate where they
than
become available to roots. Seachem itself has developed Flourite
(calcined
clay) product in part to achieve this.
3) It is central to algae control rational to keep nutrients out of the
water
column where they will be available to algae more than if some of them
were in the substate. I think the goal is get as much nutrient into
the
substrate as possible. Again I am given pause to consider why potassium
and magnesium are said to be more effectively taken up in the water
column.
These points above and gluconate's weaker bonding capacity of fe++
provided the rational behind my earlier statement that ferrous gluconate has
the potential to preferentially benefit algae.
Christopher Coleman
christopher.coleman@worldnet.att.net
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by krandall/world.std.com
Date: Tue, 24 Nov 1998
>If one is using EDTA-Iron on
>a weekly basis and getting the growth and response they desire, then that's
>great. I'm just saying that if a gluconate-iron supplement had been used
>instead of the EDTA-Iron (and depending on the foliar requirements of your
>plants) you would have seen an even better growth/response in your plants
I'm not sure that is true. If you add iron out of balance to other
nutrients, you are going to quickly push your plants into a deficiency of
what ever is in next least supply.
For my tanks personally, the limiting factor for growth at most times seems
to be either N or K. I have to really stay on top of those or growth is
rapidly affected. But of course, I run high light/strong growth systems.
In systems with slower growth and without adequate supplementation, I would
suspect iron deficiencies to show up before N and/or K deficiencies
>I guess the only real way for each individual
>to determine which is best for them is to set up two systems that are
>nearly identical and use only one product in each and after a certain
>period of time look at the results.
That's probably overkill.<g> And I doubt you'll have many people switch to
your product if they have to do that to decide if they like it.
I'd suggest that people do what I usually advocate anyway. Watch your
plants. If they look like they are suffering from an iron deficiency,
(which, BTW, won't show up as slower growth at first, but as chlorotic new
leaves. This will be apparent in fast growing plant species first) maybe
it's worth trying an iron gluconate supplement. Or for that matter, if you
just want to experiment with a new fertilizer, give it a try, but give it a
try for at least 6 months unless you see clear signs of a problem earlier.
Watch your plants and let _them_ tell you how they're doing.
One additional caveat would be that if you are using an iron-only
supplement as opposed to a balanced trace element + K supplement, keep an
eye out for signs of other nutritional deficiencies.
Karen Randall
Aquatic Gardeners Association
The Krib
Plants
Plants