Sciencemadness Discussion Board - Potassium ferrate

Potassium ferrate

vulture - 5-7-2002 at 12:39

This should be an even stronger oxidizer than potassium permanganate, IIRC the formula is K2FeO4. It can be produced by adding finely divided iron powder to molten KNO3.
A violent reaction ( I assume a deflagration) should take place and the potassium ferrate can be isolated by dissolving the mess ( what else is going to be left?

) into cold water.

Anybody ever tried this?

Polverone - 5-7-2002 at 22:46

What a coincidence! I recently tried to prepare this myself. Whatever reaction occurs certainly is not violent - or maybe I just needed finer iron powder. With 100 mesh iron powder and KNO3 heated over a gas burner, I obtained no potassium ferrate (did not observe characteristic color on dissolving in water) but some quantity of potassium nitrite (fumes evolved on addition of acid). I will have to try again in the future. I'll hunt some articles on ferrates down also.

Nils - 10-7-2002 at 09:11

Jup, KNO2 is formed when KNO3 is been reduced with an easily-to-oxidize metal(Pb, Fe etc.).

Has anyone more information about K2FeO4? it sounds interesting.
What about K2BrO4?

Ferrate information now online

Polverone - 10-7-2002 at 13:29

I have summarized and condensed a couple of journal articles on ferrate production and properties. The summary is available in the budding sciencemadness library; http://www.sciencemadness.org and then click on "Library."

vulture - 6-8-2002 at 01:28

Nils, what do you mean with K2BrO4? Only KBrO3 exists, that is potassium bromate, a powerful oxidizer, much like KClO3.

bonemachine - 28-10-2002 at 09:57

I just want to ask if kno3 could be used in a solution electrolised with Fe electrodes to form potassium ferrate and if any dangerous fumes is given of.

Also anyone have tested how good oxidiser it is?

madscientist - 28-10-2002 at 14:26

Electrolyzing a solution of potassium nitrate with iron electrodes will initially yield a solution of iron nitrates and potassium hydroxide, which will soon reform potassium nitrate and precipitate iron hydroxides.

bonemachine - 29-10-2002 at 00:19

Is there any other method than melting kno3 with iron oxide to prepare ferrate? Actualy the only material close to it's synthesys i got is kno3.

thanks.

bonemachine - 29-10-2002 at 11:56

When iron nitrate decomposes in a deflagrating reaction, in what would it decompose? I mean something like FeO + N + O2 ?

bonemachine - 29-10-2002 at 13:25

Damn not a good thing that posts can not be edited. I just was curious if melting kno3 and Fe2O3 is safe (fumes giving of and stuff like that) and if it yields pottasium ferrate. I want to try that actually but i want to be sure first. Also has anyone tested it as an oxidiser and if yes is it good enough at low temperatures?

bonemachine - 1-11-2002 at 06:30

What color suposed to be potassium ferrate?

Theoretic - 31-7-2003 at 09:35

It's red/purple.
Could Na2FeO4 be made by electrolysing a NaOH solution with a steel or iron cathode?
The oxidizing potential of FeO4-- (converted to Fe(OH3)) is 2.20 - a stronger oxidizer than ozone!

vulture - 31-7-2003 at 11:28

IIRC, the reduction potential of ozone is -2.7V, so still more than FeO4-

Theoretic - 1-8-2003 at 07:09

Mistake! Not 2.7, but 2.07!
Well, 2.07 according to some, 2.08 according to another, and 2.75 according to yet another source.

vulture - 1-8-2003 at 11:18

Argh. Not even mentioning the fact that it should all be negative voltages, since it's always indicated as reduction potential.

Why oh why is the world (and my psyche, but on second thought that is none of your bussiness) so confusing?

Old topic worth reconsidering....

chemoleo - 18-5-2004 at 16:34

Well I checked... but I didnt find a detailed method on the ferrate production - via the Fe/KNO3 route. Again it's taken from Jander&Blasius, preparative inorg. chemistry.

Here it is:

10 g of Fe powder and 20 g of KNO3 are mixed. The latter has to be molten first, then pulverised, to remove any water present.
The mixture is placed ca 1 cm thick (1/2 inch) onto an iron plate. At the edge a 1:1 mixture is added, this is needed for igniting the rest of the mix.
Using a Bunsen burner, the 1:1 mix is ignited, and the heat is sufficient to ignite the rest of the mix. I am not quite sure why sparklers wouldnt do.
Once the reaction starts, white fumes develop, and the reaction proceeds through the whole mix (therefore it it slow, and rel. safe).
Now and here comes the important bit, potentially explaining why people before had trouble making it:
After cooling down, the product is dissolved in 50 ml ice-cold water, and filtered rapidly.
The red-violet filtrate is immediately mixed with an ice-cold BaCl2 solution. This is allowed to settle (BaFeO4 is not well soluble), and filtrated.
This is then washed with aldehyde-free EtOH/Acetone, and then dried in a desiccator.

Acidified solutions of BaFeO4 produce immediately O2, and are reduced from Fe VI to Fe III.

What I found noteworthy about this prep is that it stresses the need for ice-cold conditions, plus precipitation with barium - maybe that increases stability of the salt, and the ease of isolation of course.
Anyway - it also means, in warm solutions, the ferrate solution (i.e, K2FeO4) is by itself not stable for a long time - possibly explaining why it's hard to isolate this.

Apparently ferrates are stronger oxidisers than permanganates, so beware!
Although i think this is precisely what's making ferrates interesting!!!

[Edited on 19-5-2004 by chemoleo]

Theoretic - 20-5-2004 at 06:32

"Apparently ferrates are stronger oxidisers than permanganates, so beware!"
True, they oxidize ammonia to N2 at room temperature (not the actual potential but kinetic factor I'm talking about).
Ferrates are stable in Alkaline solution, maybe dissolving them in ice cold alkali-metal hydroxide solution will stabilise the product.

chemoleo - 20-5-2004 at 08:20

Oh dear, what happened to muyos attachments? They were extremely useful

- and I failed to save them!
Whoever has saved those ferrate pdfs, could you upload them onto the FTP please?

vulture - 20-5-2004 at 08:58

There are quite a few PDFs about ferrates on the FTP. Atleast one about bariumferrate.

Esplosivo - 20-5-2004 at 10:38

I've got no access to the ftp. Guess I don't have books to scan lol. Well, is this the pdf you were looking for, describing the synthesis of barium ferrate? I'm not sure this is what you attached.

Attachment: viewthread2.pdf (107kB)
This file has been downloaded 2351 times

Ferrates

Alchemist - 20-5-2004 at 13:35

Hello all,

Just uploaded "The Ferrates" to the FTP upload file.
Have fun!

BromicAcid - 22-5-2004 at 18:05

How about instead of using iron powder as the starting material Fe2O3 is used instead. In this case it could be added directly to the molten nitrate without the worry of deflagration mentioned in Vulture's first post. But for me it is simply a matter of convince, I've got 2 kg of Fe2O3 from my pyro days for making thermite. I think I will give this a try next weekend:

Large excess of KNO3 (with KOH added) is heated to melting point and slowly Fe2O3 is sifted in with stirring. The reaction is allowed to proceed for several minutes at melting point then the resultant mass will be dissolved in basified ice water. The solution will be filtered then a saturated cold BaCl2 solution will be added and any precipitate collected.

[Edited on 5/23/2004 by BromicAcid]

BromicAcid - 23-5-2004 at 09:16

I tried my above reaction outline today, but had to stop due to threat of hail and tornado's. I took a large excess of NaNO3 in prill form and added Fe2O3 till coated, I then heated the mix in a test tube over a hotplate then after it melted and I saw the reaction was going no further I took the massive heat to it with a propane torch. The contents began to bubble vigorusly and NO2 started to come off slightly. But I was hurried so I had to stop here. I let the test tube cool and NaNO3 crystalized almost clear at the top whereas when hot it was a homogenous sludge. There were areas at the bottom of the test tube that were yellow/green but that does not match the color of the ferrate anion mentioned further upthread. However when I took the mass and added to dil. HCl a vigorous bubbling was observed. Due to time constraints this was the only test I could preform regardless of how inconclusive it may have been.

[Tried again later, glass test tube cracked from the heat so used iron crucible, heated to bubbling for 30 minutes, dissolved mass in water, filtered, very faintly purple/red solution, added sat. BaCl2, solution became cloudy, too little BaFeO4 to even attempt isolation so in retrospect Fe2O3 is not the best starting material, as a matter of fact I would recomend against it.]

Quote:

Could Na2FeO4 be made by electrolysing a NaOH solution with a steel or iron cathode?

About 3 weeks ago I tried the electrolysis of KOH using an iron pipe as my vessel. I took the KOH and packed it into the pipe and added somewhat significant quantities of water till I had free water at the bottom. Two nickel electrodes were lowered into the solution and electrolysis commenced, this is what it looked like after 20 minutes:

However I don't remember which is the cathode and which is the anode in the picture. However the solution to the right does have a red/purple color that could be K2FeO4 from the KOH reacting with the vessel. I was really wondering what it could be so now I have a possibilitiy.

I guess I'll have to give this a try again except this time with iron cathode and anode doing electrolysis in a KOH paste, the reaction was fast last time and a very distinct purple, sounds much easier then the waste of time I had with the Fe2O3/NaNO3 method that I tried eariler today. I'll keep everyone posted.

[Edited on 5/24/2004 by BromicAcid]

chemoleo - 25-5-2004 at 17:35

Regarding your Fe2O3/NaNO3 thing, the problem may be that the reaction is simply too endothermic for Na2FeO4 to form (i.e. the oxidation potential is not enough at the temperatures employed to oxidise the iron III+ (Fe2O3) to Fe VI+. Quite conversely, oxidation of iron to iron III + happens quite easily, with a large reaction enthalpy. I just thought that this enthalpy is needed to further oxidise it to Fe VI+. How knows, if you heated it up enough it would form Fe VI+ with NaNO3. But then, it would probably decompose due to the high heat....

As to electrolysis - wouldnt you just get precipitating iron hydroxide? Still it's an interesting experiment!

The_Davster - 26-5-2004 at 15:37

I found the folowing line in "basic inorganic chemistry" by cotton and wilkenson.
"Na2O2 vigourously oxidizes some metals. For example, Fe reacts with Na2O2 violently to give FeO42-"
Thats all thats mentioned about ferrates in the entire book. Unfortunatly sodium peroxide is expensive (9$ US a ounce), but if anyone has any lying around it would be worth a try.

BTW: Is sodium peroxide a more powerfull oxidizer than potassium ferrate?

BromicAcid - 26-5-2004 at 19:31

From "The Chemistry of the Elements:

"But the best known oxo-anion of iron is the ferrate (VI) prepared by oxidizing a suspension of hydrous Fe2O3 in conc alkali with chlorine, or by the anodic oxidation of iron in conc alakli. The tetrahedral [FeO4]-2 ion is red-purple and is an extremely strong oxidizing agent. It oxidizes NH3 to N2 even at room temperature..."

So there is even more possibility that I generated the ferrate anion in my electrolysis experiment above. And Theoretic, where did you get the oxidizing potential for ferrates to Fe2O3? I was reading about strong oxidizing agents and it listed bismuthates as one of the strongest in aqueous solutions and their potential is 2.04 I believe if ferrate ---> Fe2O3 is 2.20 then that would be incredible!

Theoretic - 27-5-2004 at 03:39

Yes, 2.20 V for the three electrons taken when iron goes down from +6 to +3 oxidation state in acid solution - FeO4-- => Fe+++. Which is why I said it's stronger than ozone. I got my info from a 6-cm thick book blandly entitled "Inorganic chemistry". I have also found this information onthe web.

Why, oh why do all the oxidation potential tables give potential for ACID solution? Wouldn't it be more sensible to give NEUTRAL potentials? Most potential tables, even when they give both the acid and the alkaline potentials, they almost never give the neutral potential...
A nice link here...
http://env.snu.ac.kr/research.html

Edit: and reportedly, iron 5+ is a "much more powerful oxidant then Fe 6+"

How on Earth...

[Edited on 27-5-2004 by Theoretic]

unionised - 27-5-2004 at 14:27

Because they give them with respect to a standard hydrogen electrode which is acidic. This way they can avoid having a liquid junction potential to try to estimate.

BromicAcid - 27-5-2004 at 16:19

Tried more work with ferrates today. Decided on the anodic oxidation method. Took an iron crucible and filled it with about 10 g of KOH and about 15 ml of H2O. Iron electrodes were inserted and electrolysis commenced. Initially the electrode solution turned purple but after 10 - 15 minutes the current had jumped and it was black. I poured the boiling hot solution in ice water. Initially it was black but the darkness settled and it was dark red but bubbles started continuily rising from it. And the black precipitate settled to the bottom. I let it dissolve a little more but when I came back 5 min or so later the solution was even more decolorised and bubbles kept rising, I filtered the solution but what was left hardly looked concentrated enough to attempt isolation if anything was there.

Second attempt same deal but more KOH and the electrolysis ran for less time. This time when I took the electrodes out the cathode was covered in a flaky mass, figured out my mystery precipitate from the first try, the iron was reacting with the hydroxide and re-plating on the cathode. The molten mix this time was green with streaks of red. I poured into iced water and stirred, knowing that the solution decomposed quickly the time before I immediately filtered it then seeing that there was a considerable color left to it I added saturated BaCl2 solution. A precipitate did form initially some very tiny red specks but overall it eneded up being a fluffy light chocolate colored precipitate. Considerably lighter then I was thinking it would be. Filtered, put a tiny amount in the middle of a watch glass and added HCl from the side and let it run down into the precipitate, it made a yellow solution and ate away at it somewhat rapidly with gas evolution. No Cl2 was evolved that I could smell. Tried oxidizing Xylene with it in acid but all it did was decompose my precipitate, the xylene became cloudy but apperaed vastly unchanged. I believe my precipitate was mostly barium hydroxide

So now I'm sick of doing this the round about way, Friday is fusing KOH with Fe2O3 and if that doesn't work steel wool :mad:

I took pictures of the above if anyone is interested.

Saerynide - 30-5-2004 at 07:46

Me is interested

BromicAcid - 30-5-2004 at 08:29

Concentrated solution of potassium hydroxide being electrolyzed in a steel crucible with iron electrodes at 12V and 10A. When the electrolysis was first commenced a definate purple color came from the anode but as the solution heated up and it progressed the whole solution became black.
Added the mess still boiling hot to iced water, the water was distilled but the ice was not Regardless, some solid iron fell to the bottom of the beaker and from there I filtered it twice while still ice cold.
Added conc BaCl2 solution and after 10 minutes this is the precipitate that settled.
Picture of precipitate after decanting the water, looks darker almost chocolate.
Put some of the precipitate into the middle of a watch glass and put a drop of HCl on the edge. It bubbled like neutralizing a carbonate but nothing of the sort was present in the solution. After adding a few drops more I was left with a green solution with a precipitate of BaCl2 on the bottom. I do not believe I had Ba(OH)2 because that would not have given an efforvesence however I believe if I had indeed made the ferrate it would have caused chlorine evolution from oxidizing the chlorine anion in the HCl. So I'm not really sure what I eneded up with.

Another Ferrate Prep!

chemoleo - 5-7-2004 at 19:00

The reaction is
2Fe(NO3)3 + 3NaOCl + 10 NaOH ---> 2 Na2FeO4 + 6 NaNO3 + 3 NaCl + 5 H2O

Make a solution of 1 g of iron III nitrate enneahydrate in 5 ml of water, and drip this into a solution of 2g NaOH in 50 ml of 2M NaOCl.
Make sure the drops come about at 1 per 15 seconds ( so very slow addition) to the stirred NaOH/NaOCl solutionl.
Once all the iron nitrate has been added, boil the solution for 1-2 minutes.

Then filter the hot solution through a sintered glass, straight into 1g of barium nitrate (or chloride I guess) in 20 ml of water.
Filter off the red barium ferrate precipate, wash with water and dry.

Notes: I wouldnt see why FeCl3 wouldnt work, which is VERY easy to obtain from electronics suppliers. Else... it couldnt be easier to do, could it? Someone try... please!

BromicAcid - 5-7-2004 at 19:09

Sounds simple enough, I will make me some ferric nitrate tomorrow! Maybe the nitrate is necessary, acting in some way in the reaction not apparent in the overall equation?

I did some experiments to make ferrates the other day but nothing turned out. Tried anodic oxidation in highly basic solution, an iron anode simply covered in rust and hardly reacted further. Steel wool as an anode turned the solution purple like permanganate but the color disappeared in 20 seconds or less. Precipitation with barium chloride gave the same orange precipitate as my previous experiments.

But this sounds to be the easiest method yet, no molten hydroxides, pyrotechnic mixtures, bubbling of boiling solutions with chlorine, good work Chemoleo!

But 2M NaOCl, doesn't that work out to be around 15% w/v ? If so that is not something to be accomplished lightly, such solutions would have to be made and used quickly. That is the principle behind bubbling Cl2 though an alkaline suspension of ferric oxide.

[Edited on 7/6/2004 by BromicAcid]

Theoretic - 8-7-2004 at 04:15

Why not use CaCl2 istead of BaCl2?
It's MUCH more easier to obtain, and I believe CaFeO4 will be slightly soluble, so that it could be precipitated at low temperatures. Bring out the ice bath (also with CaCl2)

Also... what about evaporating an ice-cold (even a little bit below zero) solution of Na2FeO4 under vacuum? The latter being obtained by a converted bycicle pump (reverse the valves and you have a vacuum pump)!

BromicAcid - 12-7-2004 at 16:58

Quote:

Make a solution of 1 g of iron III nitrate enneahydrate in 5 ml of water

Check, my Fe(NO3)3*9H2O turned out decent from reacting Fe2O3 with boiling HNO3, upon dissolving slightly more then 1 gram in 5 ml of H2O, there was a slight amount of Fe2O3 at the bottom undissolved. I jammed a cotton wad into a pipette and sucked up the liquid and put it into a different test tube, sucking up the liquid, taking the bulb off the pipette and putting it on the other side then putting it in the new container, forcing it all to be strained. What was left was a stable red solution.

Quote:

and drip this into a solution of 2g NaOH in 50 ml of 2M NaOCl.

2M NaOCl is close to 15%, very concentrated stuff, the highest concentration I could come up with was 10% although I could have basified it and bubbled Cl2 thought it I decided to use what I had, I put 50ml into a 150 ml Pyrex beaker with 2g NaOH, I inserted a medium magnetic stirring bar and set it to a medium stirring speed (4.5 out of 7ish), the NaOH dissolved quickly.

Quote:

Make sure the drops come about at 1 per 15 seconds ( so very slow addition) to the stirred NaOH/NaOCl solutionl.

It was a mistake to think that adding only 5 ml, one drop at a time by hand would be an easy task. Even this small amount seemed to take forever and I was counting out the seconds in my mind, I think I skipped several of them numerous times, about 6 pipettes full

At first there was almost no effect on the solution (first 5 or less drops) then it started to turn from orange to amber, then later red, when it was red a very very fine colloidal solution of Fe(OH)3 had developed that would not precipitate on standing. The solution continued to darken to almost black and individual particles were no longer visible, magnetic stirring is a must, and do not let the solution drip down the side of the beaker, it makes a solid cake where it contacts the solution that does not break up.

Quote:

Once all the iron nitrate has been added, boil the solution for 1-2 minutes.

I heated quickly to get it boiling as fast as possible. It remained black for the first 30 seconds or so but a purple foam started appearing on the top rapidly after that, ferrate?

Quote:

Then filter the hot solution through a sintered glass

Glass is a must! I used a cotton ball in the bottom of a funnel with a vacuum pump attached to the flask it was filtering into. The mixture rapidly attacked the cotton and in addition the fine Fe2O3 and Fe(OH)3 still suspended very rapidly clogged up the cotton. This resulted in not one, but two swaps of filtration apparatuses in the middle of filtration, I believe my yield was significantly dropped at this point as bubbles continuously rose to the surface, both from the hot conditions where it was oxidizing the water, and from attacking the cotton.

Quote:

straight into 1g of barium nitrate (or chloride I guess) in 20 ml of water.

I used about 3 g of BaCl2 in 100 ml of H2O, because I didn't remember the amount of water called for.

Quote:

Filter off the red barium ferrate precipate, wash with water and dry.

Yes, it worked, I got a red precipitate and was left with an off yellow solution, without the smell of NaOCl unless you literally stuck it right under your nose. Very different then the original starting solution.

Quote:

Notes: I wouldnt see why FeCl3 wouldnt work, which is VERY easy to obtain from electronics suppliers.

I am totally in agreement, FeCl3 added to the solution should make the same super fine Fe(OH)3 precipitate that is oxidized as any other soluble Fe3+ salt.

Some things I did with my supposed ferrate. Took 5 ml Xylene and added a few drops of HCl(aq) to it, to this under the influence of magnetic stirring I added two drops of the ferrate slurry at the bottom of the test tube I suctioned it into. Immediately the solution turned yellow (Oxidized Cl-) and within a short amount of time there was a white precipitate bumping around in the aqueous phase.

Took 'ferrate' and added one drop to one drop HCl, the two drops bubbled very vigorously and the solid precipitate was quickly consumed. The same experiment with NaOCl straight from the bottle showed no where near the aggressiveness of the ferrate mixture, in addition one drop of the liquid that the ferrate had precipitated from with HCl showed almost no reaction.

Adding ferrate to methanol with a few drops of HCl resulted in vigorous reaction that ended up with a green solution and was hot, 50C+.

I tried the reaction again later, this time using Fe2O3 in place of the slow precipitation of Fe(OH)3, I added it all at once and heated to a boil where I held it for 3 minutes. Filtered into BaCl2 and got no visible precipitate, the solution was OVERPOWERINGLY NaOCl scented, no reaction.

My conclusion, this is ferrate, the prep was simple but I need to get something better to filter with, the hot solution is somewhat menacing.

[Edited on 7/13/2004 by BromicAcid]

seen this?

Hjalmar_Poelzig - 14-7-2004 at 02:28

Abstract of US5746994

A method of producing ferrate is disclosed, in which Fe3+ is oxidized with monoperoxosulfate (HSO5-) to form K2FeO4/K2SO4. The isolation of the potassium ferrate (K2FeO4) product in a sulfate matrix (K2SO4) stabilizes the ferrate against decomposition and inhibits clumping of the solid product by inhibiting moisture adsorption. The method is a safe, simple process for the production of ferrate that is reliable, fast, and inexpensive, and that avoids the use of chlorine or chlorinated products, thus avoiding their harmful side effects. The improved ferrate product of this method is particularly useful for water and wastewater treatment, especially in the treatment of sulfides and hydrazines, and in other applications.

other patents

Hjalmar_Poelzig - 14-7-2004 at 03:11

Other processes for preparation of ferrates are known and used, many of them also involving the reactions with hypochlorite. U.S. Pat. No. 5,202,108 to Deininger discloses a process for making stable, high-purity ferrate (VI) using beta-ferric oxide (beta-Fe2 O3) and preferably monohydrated beta-ferric oxide (beta-Fe2 O3 --H2 O), where the unused product stream can be recycled to the ferrate reactor for production of additional ferrate.

U.S. Pat. Nos. 4,385,045 and 4,551,326 to Thompson disclose a method for direct preparation of iron an alkali metal or alkaline earth metal ferrates from inexpensive, readily available starting materials, where the iron in the product has a valence of +4 or +6. The method involves reacting iron oxide with an alkali metal oxide or peroxide in an oxygen free atmosphere or by reacting elemental iron with an alkali metal peroxide in an oxygen free atmosphere.

U.S. Pat. No. 4,405,573 to Deininger et al. discloses a process for making potassium ferrate (K2 FeO4) in large-scale quantities (designed to be a commercial process) by reacting potassium hydroxide, chlorine, and a ferric salt in the presence of a ferrate stabilizing compound.

U.S. Pat. No. 4,500,499 to Kaczur et al. discloses a method for obtaining a highly purified alkali metal or alkaline earth metal ferrate salts from a crude ferrate reaction mixture, using both batch and continuous modes of operation.

U.S. Pat. No. 4,304,760 to Mein et al. discloses a method for selectively removing potassium hydroxide from crystallized potassium ferrate by washing it with an aqueous solution of a potassium salt (preferably a phosphate salt to promote the stability of the ferrate in the solid phase as well as in aqueous solution) and an inorganic acid at an alkaline pH.

U.S. Pat. No. 2,758,090 to Mills et al. discloses a method of making ferrate, involving a reaction with hypochlorite, as well as a method of stabilizing the ferrate product so that it can be used as an oxidizing agent.

U.S. Pat. No. 2,835,553 to Harrison et al. discloses a method, using a heating step, where novel alkali metal ferrates with a valence of +4 are prepared by reacting the ferrate (III) of an alkali metal with the oxide (or peroxide) of the same, or a different, alkali metal to yield the corresponding ferrate (IV).

U.S. Pat. No. 5,284,642 to Evrard et al. discloses the preparation of alkali or alkaline earth metal ferrates that are stable and industrially usable as oxidizers, and the use of these ferrates for water treatment by oxidation. Sulfate stabilization is also disclosed.

FeCl3 may be a bad alternative to Fe(NO3)3

BromicAcid - 22-7-2004 at 16:52

The other day I retried the oxidation of a Fe+3 cation in a NaOCl solution using the drip method discussed in my last post, except this time I used OTC FeCl3 purchased from Radio Shack.

The results were not encouraging. In complete contrast to my Fe(NO3)3 solution the FeCl3 solution sizzled and popped when I added it to the stirring hypochlorite. Also it made coagulated blobs that only broke up under magnetic stirring. The colloid that eventually formed though was similar to what I had with the first runs of Fe(NO3)3 solution. However the suction filtered liquid that I ended up with before precipitation was green/yellow, whereas my other runs were purple/red. I got no precipitate.

My conclusion is that either the FeCl3 solution was way to concentrated which may have had some detrimental effect, or perhaps there was some organic material in the solution (e.g., ethanol) that the NaOCl solution oxidized in preference to the ferric cation. Or maybe FeCl3 just can't be used for this reaction, there was a reaction taking place, maybe the NaOCl was oxidizing the Cl- anion, there was an awful odor that burned my eyes but I couldn't place it exactly.

Regardless, other methods will be tested eventually, such as those in the above post.

Ferrates & Other High-Valent Fe Compounds

JohnWW - 27-7-2004 at 01:49

I once obtained a reddish-purple solution (slightly more red than permanganate) of sodium ferrrate (VI) by simply further alkalizing a precipitate of Fe(OH)3, and then adding ordinary household sodium hypochlorite solution. That even hypochlorite (far from being the strongest oxidant) can produce it suggests that even higher-valent Fe is possible.

Noting that Fe2O3 is amphoteric, I am fairly sure than an easier and much safer way (than many of the methods suggested here) of obtaining ferrate(VI), FeO4--, would be by the electrolysis using an appropiate voltage of an alkaline solution of sodium ferrite, NaFeO2, obtained by alkalizing a Fe(OH)3 precipitate with NaOH as before. Lower voltages would result in Fe(IV) and (V) anions.

Like its homologs ruthenium and osmium, Fe is theoretically able to form compounds with a valence of up to 8, although of course with much greater difficulty. Os forms OsO4 simply by heating in air. I have in my possession a PDF (if anyone wants it) of a theoretical study of the stability of Fe(VIII) oxide, FeO4, isoelectronic with MnO4-, which concluded that it would be marginally stable and that its production would be highly endothermic. It, and Fe(VII) as FeO4-, might just be obtainable by the further electrolysis, at an increased voltage and just above the freezing point of the solution, of an alkaline concentrated solution of Na2FeO4. Anyone like to try it?

JohnWW

Theoretic - 27-7-2004 at 13:07

Fe (IV) and Fe (V) is very unstable, so I don't think it's worth preparing them.
Electrolytic oxidation of NaFeO2 would work with some NaCl added, because this would make chlorine, this would make hypochlorite, this would oxidize FeO2- to FeO4--. Or you can directly oxidize NaFeO2 with NaClO. NaFeO2 could be prepared by dissolving iron (not steel) in molten NaOH or Na2CO3 or a eutectic of both.
If FeCl3 could be oxidized by NaClO in a non-pleasing way (to produce Cl2), and Cl- wouldn't work, then Fe2(SO4)3 could be used. On second thoughts... you can make Fe(NO3)3 by dissolving Fe in molten AN - can you get NH4HSO4 that easily to melt and dissolve Fe in it? I think that's the main reason for Fe(NO3)3 being suggested.

BromicAcid - 9-10-2004 at 14:41

Just messing around today I mixed together about 1.2 g of Fe(NO3)3*9H2O and 7 g of KOH in a test tube, enough water was added to cover the mixture and the solution had solid KOH at the bottom (white) and a dark red collodial solution above it.

In a seperate test tube I had mixed 5 g of KBrO3 and 10 ml of HBr (aq) and at the bottom about 1.9 ml of Br2 had collected. I drew this off with a pipette and added it to the first test tube a few drops at a time. It made hissing sounds and rumbled physically as the Br2 contacted the solid KOH at the bottom and the mixture gradually darkended. It turned black but right at the interface where the KOH (s) was touching the solution there was a layre of purple crystals (potassium ferrate), I decanted the solution and sure enough the crystals were one of the last things to slide out.

It was pretty.

tokat - 9-10-2004 at 23:40

Question

If you have a mix of KOH + H2O with a iron electrode, will the K2FeO5 decompose because of the OH- and H+?

BromicAcid - 10-10-2004 at 19:19

The potassium ferrate K2FeO4 will decompose in water, slower in basified water. When you electrolyze a solution of concentrated KOH in water with iron electrodes the electrodes are eaten away and the solution changes color. If your amps are high enough though you can form the K2FeO4 faster then it is decomposed. But once you are done you immediately toss the solution into iced water.

tokat - 12-10-2004 at 20:12

Thanks Bormic.

[Edited on 13-10-2004 by tokat]

Hummm... AC electrolysis gives better yeilds....

BromicAcid - 11-11-2004 at 09:55

Journal of Applied Electrochemistry
29 (5): 569-576, May 1999
Copyright © 1999 Kluwer Academic Publishers
All rights reserved
Electrochemical production of ferrate(vi) using sinusoidal alternating current superimposed on direct current. Pure iron electrode

K. Bouzek
Institute of Chemical Technology, Department of Inorganic Technology, Technická 5, 166 28 Prague 6, Czech Republic

L. Flower
University of Exeter, School of Engineering, North Park Road, Exeter, EX4 4QF, UK

I. Roušar
Institute of Chemical Technology, Department of Inorganic Technology, Technická 5, 166 28 Prague 6, Czech Republic

A. A. Wragg
University of Exeter, School of Engineering, North Park Road, Exeter, EX4 4QF, UK

Abstract
The current yield for the anodic oxidation of a pure iron (99.95%) electrode to ferrate(VI) ions in 14 M NaOH between 30 and 60 °C using a sinusoidal alternating current (a.c.) at amplitudes in the range 38–88 mA cm-2 and frequencies in the range 0.5 mHz to 5 kHz superimposed on direct current (d.c.) of 16 mAcm-2 was measured under conditions of bubble induced convection in a batch cell. The current yield for ferrate(VI) synthesis exhibited a complex dependence on temperature and a.c. frequency, but generally a maximum was observed in a frequency range 2–50Hz depending on the a.c. amplitude. A global maximum current yield after 180 min of electrolysis of 33% was reached at the following conditions: a.c. amplitude of 88 mA cm-2, a.c. frequency of 50 Hz and temperature of 40 °C. At the optimum conditions the highest d.c. electrolysis yield was 23%. Thus, operation with the a.c. component leads to an increase in the yield by 43% with respect to d.c. electrolysis alone.

Keywords
alternating current, batch electrolysis, current yield, ferrate(VI), pure iron

Article ID: 197640

From Kluwer Online

BromicAcid - 18-11-2004 at 14:26

Like I've said before, the reaction between sodium peroxide and ferric oxide yields ferrate reasonably well, but sodium peroxide is somewhat expensive, I think this would work out decently with barium peroxide, aside from density sodium and barium are very similar, more so then any other pair of alkali metals with alkali earths. But the barium would be preferred simply because of its wider availability (two pyrotechnics sites sell barium peroxide by the pound for $20 or less). If anyone can see any major reason why this won't work speak up, otherwise I might make a small initial investment soon.

S.C. Wack - 18-11-2004 at 19:01

If you don't mind CO from BaCO3 + C, BaCO3 is $1 a pound. Then the peroxide with gentler heating. Less convenient but not difficult, at least on a small scale. Everyone should have some Ba. Very useful stuff, and easy to recycle.

chemoleo - 18-11-2004 at 19:14

SC Wack, have you got more on this (although it's beside the topic)? I.e. what temps are required? Easily recycled... hmm, my BaSO4 is still waiting for it...and recycling requires high prolonged temps afaik.
How easily does the BaO convert to BaO2? Temp? Duration? I bet it's more difficult than it sounds

BromicAcid - 18-11-2004 at 20:10

Barium peroxide can be made from barium oxide by heating barium oxide in air free from water vapor at about 500C, it is fairly stable, not decomposing till about 800C which is well beyond its melting point. I said that it was similar to sodium above because of this aspect, that they are both the first metals of their period to form a large percentage of peroxide when burned in air. As for the conversion to barium oxide from carbonate.... not sure about those temps, barium carbonate is stable to its melting point (811 C) and probably beyond, and carbon reductions usually seems to run a little high for my tastes.

S.C. Wack - 18-11-2004 at 20:33

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]

neutrino - 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?

S.C. Wack - 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."

BromicAcid - 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?

two quick questions

budullewraagh - 29-11-2004 at 21:28

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?

S.C. Wack - 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]

JohnWW - 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.

budullewraagh - 1-12-2004 at 13:33

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

dont_kill_bill - 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]

Pyrovus - 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]

guy - 11-2-2007 at 22:42

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

12AX7 - 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..

BromicAcid - 12-2-2007 at 05:21

I tried to fuse Fe2O3 with KOH/NaNO3 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 FeCl3 solution with heating but I never got around to perfecting it.

bdbstone - 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.

S.C. Wack - 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 545 times

rstar - 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

Iron(VIII), Fe+8 oxidation state

AndersHoveland - 23-9-2011 at 01:42

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]