Bedlasky
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Aerial oxidation of ferrous ions
Hi.
I some time ago examined aerial oxidation of ferrous ions. It's generally known fact that ferrous ions (in solutions and also in solid form) are
slowly oxidized by atmospheric oxygen to ferric ions. But this reaction is faster then one could expected. I wrote about this short article on my
website in czech and also in english.
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woelen
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I have bad experience with this. I purchased very nice and pure mint-green crystalline FeSO4.7H2O. Some time later, when I opened the jar, the
material was covered by a brown crust of basic ferric sulfate. I also once had reagent grade FeCl2.4H2O. This completely turned brown into an
insoluble basic ferric chloride.
I now only keep ferrous ammonium sulfate as iron(II) salt. That compound is quite air-stable, much more so than the plain sulfate or chloride.
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Bedlasky
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I had in home ferrous sulfate for a few years. Surface was brown as you described. So I made from him some ferrous ammonium sulfate and a bunch of
ferric ammonium sulfate. I still have some of this old ferrous sulfate, maybe in future I will convert it in more ferrous and ferric ammonium
sulfates.
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Boffis
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I use ferrous chloride as a highly selective nitro-group reducing agent in the preparation of aromatic amino acids. I have given up trying to preserve
it in any form and now prepare it immedately prior to use by passing a solution of ferric chloride in hydrochloric acid slowly through steel wool and
filtering off the excess wool and black residue. Ferrous sulphate is definitely more stable but the sulphate ions left in solution after the reduction
complicate the work up.
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Amos
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Quote: Originally posted by Boffis  | | I use ferrous chloride as a highly selective nitro-group reducing agent in the preparation of aromatic amino acids. I have given up trying to preserve
it in any form and now prepare it immedately prior to use by passing a solution of ferric chloride in hydrochloric acid slowly through steel wool and
filtering off the excess wool and black residue. Ferrous sulphate is definitely more stable but the sulphate ions left in solution after the reduction
complicate the work up. |
Do you have a paper on this or general outline of the procedure?
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morganbw
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| Quote: Originally posted by Boffis | | I use ferrous chloride as a highly selective nitro-group reducing agent in the preparation of aromatic amino acids. I have given up trying to preserve
it in any form and now prepare it immedately prior to use by passing a solution of ferric chloride in hydrochloric acid slowly through steel wool and
filtering off the excess wool and black residue. Ferrous sulphate is definitely more stable but the sulphate ions left in solution after the reduction
complicate the work up. |
That sparked my interest. A little more if you have it.
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Boffis
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Yes I have several papers, I am away from home at present but I'll dig out the digital ones at least when I get home at the weekend. The problem with
the sulphate is that sodium sulphate crystallises out with the product while sodium chloride says in solution and salts out the amino acids. You can
get around this by using a barium chloride precipitation to remove the sulphate but this is messy and difficult to filter.
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Boffis
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Right, the original paper on the use of ferrous hydroxide reduction of nitro compounds was:
Benda; Berichte der Deutschen Chemischen Gesellschaft v47; p1316 [1914]
but this paper give specific details and explains the difference in methods from ferrous sulphate v ferrous chloride in the reduction of
nitrophenylarsonic acids:
Jacobs et al., J. Amer. Chem. Soc.; v40, p1580, is10 [1918]
Hope these help, the first is in German though.
[Edited on 22-2-2020 by Boffis]
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AJKOER
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From Chemosphere, 2007 Aug;68(11):2080-4. Epub 2007 Mar 21.
"The effect of pH on the kinetics of spontaneous Fe(II) oxidation by O2 in aqueous solution--basic principles and a simple heuristic description" by
Morgan B1, Lahav O.
Abstract
The spontaneous chemical oxidation of Fe(II) to Fe(III) by O(2) is a complex process involving meta-stable partially oxidized intermediate species
such as green rusts, which ultimately transform into a variety of stable iron oxide end-products such as hematite, magnetite, goethite and
lepidocrocite. Although in many practical situations the nature of the end-products is of less interest than the oxidation kinetics, it is difficult
to find in the literature a description of all the basic steps and principles governing the kinetics of these reactions. This paper uses basic
aquatic-chemistry equilibrium theory as the framework upon which to present a heuristic model of the oxidation kinetics of Fe(II) species to ferric
iron by O(2). The oxidation rate can be described by the equation (in units of mol Fe(II)/(l min)): -d[Fe(2+)]/dt = 6 x
10(-5)[Fe(2+)]+1.7[Fe(OH)(+)]+4.3 x 10(5)[Fe(OH)(2)(0)]. This rate equation yields a sigmoid-shaped curve as a function of pH; at pH values below
approximately 4, the Fe(2+) concentration dominates and the rate is independent of pH. At pH> approximately 5, [Fe(OH)(2)(0)] determines the rate
because it is far more readily oxidized than both Fe(2+) and FeOH(+). Between pH 5 and 8 the Fe(OH)(2)(0) concentration rises steeply with pH and the
overall oxidation rate increases accordingly. At pH values> approximately 8 [Fe(OH)(2)(0)] no longer varies with pH and the oxidation rate is again
independent of pH. The paper presents a heuristic overview of the pH dependent kinetics of aqueous ferrous oxidation by O(2(aq)) which we believe will
be useful to professionals at both research and technical levels.
[Edited on 29-2-2020 by AJKOER]
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