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Author: Subject: Potassium ferrate
neutrino
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[*] posted on 19-11-2004 at 14:34


Quote:
Originally posted by S.C. Wack
You can convert your sulfate to carbonate the same way you would with PbSO4, by making a paste with baking soda and water, then heating with an ordinary gas flame.


How does this work? Where does the sulfate go?
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S.C. Wack
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[*] posted on 19-11-2004 at 16:24


Sodium sulfate, to be washed away with water later.

BTW, looking at Qualitative Chemical Analysis, it says that not just the nitrate and iodate of Ba give BaO on heating, but all of the organic salts as well. Well, OK, maybe they mean 1400C. Maybe not. This quote also: "Boiling BaSO4 with at least 15 times its equivalent weight of 2-4 N Na2CO3 will convert 99% of the BaSO4 to Na2SO4 in one hour, in the case of a fresh precipitate. About double the time is required for native barite."
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BromicAcid
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[*] posted on 29-11-2004 at 17:33


I found some evidence today that barium peroxide would not work. In the book I checked out from the library "Peroxides, Superoxides, and Ozonides of Alkali and Alkaline Eath Metals"

Quote:
Barium peroxide reacts with molybdate according to the equation

Mo + 3BaO2 ---> BaMoO4 + 2BaO

The reaction with iron is as follows

2Fe + 3BaO2 ---> Fe2O3 + 3BaO

It has been established that Fe2O3 accelerates the liberation of oxygen from barium peroxide. In the process, Fe2O3 acts as a catalyst up to 500C, and over 600C it reacts with BaO2 forming BaO*6Fe2O3


Maybe it acts catalytically by producing ferrate which readily decomposes at these temperatures?




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[*] posted on 29-11-2004 at 21:28
two quick questions


how is it possible to chemically oxidize iron to a +6 state??
(FeO4-2) is the ferrate anion, right?
i have never seen iron oxidized beyond +3...

as well, could one do the same reaction with aluminum instead of iron and achieve a similar result?




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S.C. Wack
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[*] posted on 29-11-2004 at 22:01


Well, it is in the literature but who knows. One of my books mentions it, it is in Gmelin's, and in Z. anorg. - BaFeO4 from Fe2O3 and BaO2. But my German sucks. I wonder about using Fe2O3, Ba(OH)2, and a nitrate or chlorate. Kind of like the Fe or Fe2O3/KNO3, with the Ba ferrate formed in situ.

I see that FeCl3 was mentioned earlier - there is a method using it, from J. Chem. Phys. Low yield though, and it uses insane amounts of NaOH and KOH relative to the FeCl3. They say that deviation from the procedure decreases the yield even less than their 10-15%, though the authors of the JACS article that Polverone wrote down improved the yield a little, but claim a product of much lower purity than the original. Touchy.

As for Al, it is trivalent - only.

[Edited on 30-11-2004 by S.C. Wack]
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[*] posted on 1-12-2004 at 13:10


Budullewraagh said: "how is it possible to chemically oxidize iron to a +6 state??
(FeO4-2) is the ferrate anion, right?
i have never seen iron oxidized beyond +3..."

Of course it is possible. I have done it myself, using virtually kitchen reagents, by adding an excess of household bleach, which is alkaline sodium hypochlorite solution, NaOCl, to rouge, which is powdered ferric oxide, Fe2O3, or indeed to any soluble iron compound. The Fe(III) dissolves (or redissolves) as sodium ferrite(III), which is then oxidized by excess hypochlorite to the intensely purple-magenta colored ferrate(VI), Na2FeO4, which is similar to permanganate(VII) in color although slightly more reddish. It can also be obtained by electrolysis of an alkaline solution of sodium ferrite. In this, the Fe has two unutilized 3d electrons, which can be confirmed by its paramagnetism. Ferrate could be used as a cheap substitute for permanganate in water treatment.

Because Fe is potentially octovalent, like Ru and Os, it just might be possible to obtain perferrate(VII), FeO4-, and the tetroxide(VIII), FeO4, by electrolysis of a supercooled alkaline solution of ferrate at higher voltages. These would be very liable to decomposition on warming. However, no serious attempt to make them seems to have been been reported.
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[*] posted on 1-12-2004 at 13:33


i think im gonna go find some rust now...



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[*] posted on 10-7-2005 at 06:13


Quote:
Originally posted by S.C. Wack
I've made most all of the simple inorganic Ba cpds. starting from the carbonate or BaSO4. You can convert your sulfate to carbonate the same way you would with PbSO4, by making a paste with baking soda and water, then heating with an ordinary gas flame.

The peroxide is the easy part, the temperature is 500C. Of course it is preferred to do this in O2 or at least CO2-free air, but...homemade H2O2 (the purpose of my peroxide experiments) came out fairly well for me.

An alternative way to BaO, and to BaO2 from there, by heating BaNO3, is in Inorganic Laboratory Preparations.

I recall reading somewhere that superheated steam is an industrial route to the hydroxide from the carbonate, and the hydroxide gives the oxide on strong heating.

The carbon reduction is at 1100, not a big deal for small amounts.

...Just took a look at Thorpe, heating the iodate is mentioned but that is a little much. He also mentions the steam, but that isn't where I saw it. Is Thorpe (A Dictionary of Applied Chemistry vols 1-7 except for 3, because Gallica doesn't provide it) not on the FTP somewhere? a_bab?

EDIT: Have cropped the Thorpe pdfs, am converting to djvu, Thorpe will be up as soon as the djvu virtual printer will convert it. So this might lead one to think that I highly recommend that everyone should download it once up. Although of an industrial bent, there is enough lab work, refs, and just general DIY knowledge to make it A Good Thing. Just a coincidence that the missing volume is the one covering explosives.

[Edited on 19-11-2004 by S.C. Wack]
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[*] posted on 11-2-2007 at 20:55


Here's a picture of a ferrate solution I prepared recently (apologies for the picture quality).



I prepared it by reacting FeSO4 with a very large excess of NaOCl, and a fairly large quantity of NaOH. This reaction is a very good example of the effect that temperature can have. Upon the reaction of the FeSO4 with the NaOCl, the solution immediately turned a dirty-brown as a result of the precipitation of Fe2O3. It took maybe half an hour before a very pale purple colour could be noticed, but the solution was going berserk producing bubbles, indicating that the ferrate was decomposing very soon after formation, despite the very high pH. Placing the solution in the fridge seemed to stabilise the ferrate considerably, with very few bubbles being produced, to give the much deeper purple solution shown in the picture.

[Edited on 12-2-2007 by Pyrovus]




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[*] posted on 11-2-2007 at 22:42


Did anyone here try to fuse Fe2O3 with NaOH/KNO3 yet?



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[*] posted on 11-2-2007 at 22:55


I think Bromic did, even with electrolysis.

I can take a stab at it, when I get my induction heater back together..




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[*] posted on 12-2-2007 at 05:21


I tried to fuse Fe<sub>2</sub>O<sub>3</sub> with KOH/NaNO<sub>3</sub> and got just a hint of ferrate, I don't think I was able to get the melt hot enough. With electrolysis though of a highly concentrated solution of KOH with iron electrodes I did manage to get a good ferrate color but nothing significantly recoverable. Best method I found was just to add bleach to FeCl<sub>3</sub> solution with heating but I never got around to perfecting it.



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[*] posted on 11-9-2011 at 11:59


Why the Fe2O3 didn't work? Does anybody else tried that?

I want to react it with KNO3, using NaOH as an alkaline source.
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[*] posted on 11-9-2011 at 17:07


Well the missing explosives-kyrofin volume of Thorpe has been available for a long time now from google.
These days there's an ACS symposium series book called Ferrates that looks like it goes into preparation some: http://dx.doi.org/10.1021/bk-2008-0985
- it isn't in my archive.
Someone registers as dont kill bill, posts right away, and never logs in again, ok. What do they do -
...weirdness...This thread still freaks me out, see -
google led to the first article below today. It was read over some time and the pdf was closed, then straight here to see if it had been mentioned in an organic thread, but here is this old thread on the top of recent posts. It's stalking me, insurance if I hadn't gone all the way to page 4 of google search results. Hi bdbstone. This is awkward. This is all about me, not you.

Unglaze.....now.

This is an "improved" preparation from potassium hypochlorite, the ubiquitous excess KOH, ferric nitrate, and some experimentation with their product and organics like benzyl alcohol and toluene, and some clays, silica gel, etc. They insist on high grade KOH. Ferrate is used in several other articles where benzaldehyde is obtained in various yields.
A Novel Oxidizing Reagent Based on Potassium Ferrate(VI) J. Org. Chem. 1996, 61, 6360-6370
http://dx.doi.org/10.1021/jo960633p

Earlier articles along those lines are mentioned there, Polverone coming up with the first earlier -

Preparation and Purification of Potassium Ferrate. J. Am. Chem. Soc., 1951, 73 (3), pp 1379–1381
http://dx.doi.org/10.1021/ja01147a536

Preparation and Alcohol Oxidation Studies of the Ferrate(VI) Ion
Inorganica Chimica Acta Volume 8, 1974, Pages 177-183
http://dx.doi.org/10.1016/S0020-1693(00)92612-4

Attachment: K ferrate.zip (468kB)
This file has been downloaded 340 times





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[*] posted on 23-9-2011 at 00:40


I added bleach to FeSO4 solution and got a brown precipitate. A added bleach in excess, and heated the mixture for about 7 mins. What i got was a pink non turbid solution, which i think is Na2FeO4
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[*] posted on 23-9-2011 at 01:42
Iron(VIII), Fe+8 oxidation state


After reading one of the posts in this forum suggesting that ferrate(VIII) might exist, I did some online research. Here is what I found:

translated from russian: "K2FeO5 green, easily decomposed, especially in the light."

"Mellor’s Treatise on Inorganic Chemistry devotes only a couple of paragraphs to the preparation of potassium perferrate K2FeO5 and iron tetraoxide FeO4.
Mellor stated that K2FeO5 was obtained by heating Fe2O3, KOH and a large excess of KNO3. A green melt is obtained, which becomes a green solid on cooling. When I tried it, I certainly obtained a green melt, but it became white when cooled. Sadly, I can't remember anything else" (quote by "ferrocene" from scienceforums.net)

http://www.intechopen.com/source/pdfs/14558/InTech-Ferrate_v...
this link, which mainly discusses ferrate(IV) has a table of different iron oxidation states. It clearly shows ferrate(VIII) right below ferrate (VI).

ferrate (IV) Na2FeO3
ferrate (V) K2FeO4
ferrate (VI) Na2FeO4, K2FeO4
ferrate (VIII) Na2FeO5

"in addition to the stable oxidation states of iron, 0, +2, +3, the strong oxidizing envirorment caused the occurrence of higher oxidation states of iron, +4, +5, +6, +8. These higher oxidation states of iron are commonly known as ferrates. Among the ferrates the +6 is the most stable and easy to synthesize..."

Ferrate(VI) in the Treatment of Wastewaters: A New Generation Green Chemical
Diwakar Twaria and Seung-Mok Lee (from India and Korea)

I do not think that ferrate(VIII) is very stable, it likely decomposes shortly after it is formed. Ferrate(VI) is also unstable in aqueous solutions if the concentration is too high or the pH is not high enough. Even the more stable aqueous solutions typically decompose after 30 minutes.

[Edited on 23-9-2011 by AndersHoveland]
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