What are the characteristics of a solution containing ferrate(VI) ions?

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Offline Bill.D.Katt.

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I recently heard of the ferrate (VI) ion. I attempted a synthesis with NaClO and Fe(OH)2, and believe I attained partial success when I got a pink solution. It looked similar to a dilute MnO4 solution. Would heating this solution concentrate the FeO4 2- ions, or would it break them down?
« Last Edit: 16/06/2010 21:38:47 by chris »

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Offline lightarrow

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I recently heard of the ferrate (VI) ion. I attempted a synthesis with NaClO and Fe(OH)2, and believe I attained partial success when I got a pink solution. It looked similar to a dilute MnO4 solution. Would heating this solution concentrate the FeO4 2- ions, or would it break them down?
Ferrate(VI) is quite unstable and lasts a few hours even at room T in solution.

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Offline Bill.D.Katt.

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If it only lasts a few hours wouldn't I see iron oxide or hydroxide slowly precipitating out? I've observed the solution for several days and there has been no precipitate.

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Offline lightarrow

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I don't have idea of the FeO42- concentrations you got, maybe it was so small that you couldn't observe any precipitate (of course the fact your solution had a significant colour intensity doesn't say anything).

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Offline Bored chemist

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How sure are you that there was no Mn present?
Please disregard all previous signatures.

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Offline Bill.D.Katt.

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I am very sure there was no Mn present. Here is what I did (I have repeated since): old BBs (with copper coating removed) + HCl +H2O2 (to speed up reaction) = FeCl2. FeCl2 +NaHCO3 (to remove potential unreacted HCl) =NaCl +CO2 +Fe(OH)2. Fe(OH)2 +NaClO. A lot of the Fe(OH)2 was unreacted so I filtered it. I added bicarb to remove potential Cl2 formation. After bicarb was added I didn't see any distinct color change, it just turned to a grey sludge. With the NaClO it immediately turned into the brown-red usually associated with iron. Also the BBs are steel, so there might be some reaction with the C that I am not aware of.

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Offline Bill.D.Katt.

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The solution looked like less than a gram of KMnO4 was added to 250 mls of water. Light pink.

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Offline lightarrow

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How sure are you that there was no Mn present?
Don't know about him, but I got a pink, MnO4- - like solution too. My process was different: I molted NaOH with pure Fe2O3 (I'll write further on how much pure) and then added Na2S2O8. After have cooled the melted, I dissolved it with cold NaOH water solution and got the dark pink solution, which however got much less coloured after ~ a day, forming a brown, Fe(OH)3 - like precipitate.

Mn in the Fe2O3 can is, according to the label, less than 0.25% so it could even be that. But why the solution was darker before? No trace of black, MnO2 - like precipitate.
« Last Edit: 17/06/2010 12:56:45 by lightarrow »

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Offline Bill.D.Katt.

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I'm not sure, I've never been able to get it past a sharp pink color, which I got by just letting it sit and slowly evaporate. There might be a definite equilibrium in both of our solutions keeping it from reaching a higher concentration. From what I've seen in experiments the MnO4 ion is fairly stable unless you have organic molecules near it, so it shouldn't decompose in solution. I also tried the Fe +KNO3 procedure, but I can't claim any success.

The NaOH should take care of the acidic solution-decomp problems, but if it is an equilibrium then there will have to be another solution that can be concentrated because neither solutions seem to be working.