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Author: Subject: Aerial oxidation of ferrous ions
Bedlasky
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[*] posted on 17-2-2020 at 15:23
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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[*] posted on 17-2-2020 at 23:26


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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[*] posted on 18-2-2020 at 00:51


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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[*] posted on 18-2-2020 at 05:52


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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[*] posted on 18-2-2020 at 16:23


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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[*] posted on 19-2-2020 at 12:35


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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[*] posted on 20-2-2020 at 03:56


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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[*] posted on 22-2-2020 at 11:28


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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[*] posted on 28-2-2020 at 19:00


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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