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Re: iron gluconate

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:


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

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