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Re: FeIII solubility

> From: "Peter Hughes (X)" <peterh at pican_pi.csiro.au>
> Date: Mon, 25 Mar 1996 13:23:55 +1000 (EST)
> Subject: Iron solubility
> 
> There have been some comments on Iron solubility that I thaught I would 
> post some information on.
> 
> FeCl3 has a solubility of 74g/100ml (ie Fe3+)
> FeCl2                     64.4g/100ml (ie Fe2+)
> 
> The solubility of the phosphate forms of these ions is very limited but 
> at the common concentrations in aquariums should not precipitate as 
> phoshates. The sulphates are listed as sparingly soluble in water, but 
> soluble in acids. This data is from the CRC Handbook of chemistry and 
> physics. It is written for people trying to make solutions considerably 
> more concentrated than those that should be encountered in the aquarium. 
> Another thing is that it is a long way from being a handbook, it has a 
> great deal of very interesting data in it.
> 
> So to round off the discussion on Iron I would like to put my two cents 
> worth in. The solubility of both 2+ and 3+ Iron should not be a problem 
> under most situations, especially if in a well planted tank that the 
> phosphate concentration is kept low relataive to the iron and also 
> chelated with EDTA. If peopel would direct me to some papers that 
> actually detail the unavailability/unusability of Fe2+ over 3+ I would be 
> grateful. The chemist in me thinks that this would be highly unlikely, 
> the cytosol of eukaryotic cells is reductive and so the Fe3+ would be 
> reduced there to Fe2+. I have a great deal of difficulty in accepting 
> that the valence state of the ions makes that much difference. Perhaps 
> what is really being talked about are the variety of Iron oxides that 
> really are very insoluble and as such unavailable for use ? A word of 
> caution here to, something that has not appeared in an externally 
> reviewed scientific paper is often pretty dodgy in terms of its 
> reproducibility and rigor. I know that it seems pedantic, but it is 
> usually the unrefereed journals that we have problems with.
> 
> Hoping that this stirs some comments
> 
> Peter Hughes     ANGFA(ACT)
>
	I'll comment!  I would also very much like to see
some papers that explain why (if) Fe2+ is preferable to Fe3+ as far as
plants are concerned. 

	The problem with Fe3+ in aquatic solution
is its tendency to hydrolyse.  The ion Fe(H20)6 3+ is quite a strong acid:

*    (Fe(H20)6)3+     <->   (Fe(H20)5OH)2+    +    H+       K= 10^-3.05

	Other reactions follow, and the result is that, unless the pH is 
very low, oxohydroxides precipitate out.  The iron _can_ be kept in solution
as the EDTA complex Fe(EDTA)(H20)-.  This is still iron(III), but its 
three positive charges have been offset by four from the EDTA.  Paradoxically,
even phosphates may _help_ keep FeIII in solution by complex formation and
prevention of the hydrolysis reactions.  One should also bear in mind that 
the complex iron(III) oxohydroxides may take a long time to actually
leave the solution if the iron concentration is low, as it is in our tanks.


	The question is:  Is the problem the precipitation of iron in the III
oxidation state, or an inability of plants to use it in this oxidation state?

	Bear in mind too that our test kits measure iron(III).


Paul Sears (another chemist)   Ottawa, Canada

*  Cotton and Wilkinson, Advanced Inorganic Chemistry, second ed.