Sciencemadness Discussion Board

Overlooked Ferrates

ave369 - 9-7-2015 at 00:17

I have an interest in the ferrate (VI) salts, and I've found it very hard to find useful information about ferrate salts other than potassium, sodium and barium.

There is almost nothing to be found in the Internet about calcium ferrate, lithium ferrate, magnesium ferrate. Are they unable to be synthesized? Is there any ferrate solubility table?

hyfalcon - 9-7-2015 at 03:56

Read, read, read.

https://www.google.com/?tbm=pts&gws_rd=ssl#tbm=pts&q...

BromicAcid - 9-7-2015 at 12:58

Sometimes it's amazing how little information is on the internet for a given subject. While on the other hand it would not shock me in the slightest if there where whole books written previously on ferrates.

Oh, and a quick amazon search says there is at least one:

Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (ACS Symposium Series) Hardcover – July 25, 2008

Did you check the literature in the SM library? There is likely entries in Mellor's for at least some of these. But as far as I know ferrates do not need a large cation to stabilize them so any metal cation (in a high oxidation state so it itself is not subject to the oxidation) could be used with the ferrate. So there is nothing on in the way (at least on the surface) to making calcium, lithium, magnesium ferrate etc. If you could make the barium then you would be able to easily do a displacement with the sulfate salt of some of those to make your target.

As for the ferrate solubility table, it's hard to find solubility tables for some common salts, let alone something as exotic as ferrates, you may be out of luck.

aga - 9-7-2015 at 13:23

Are you looking to achieve something specific with Ferrate (VI) salts or is it for general interest ?

Edit:

Ah. I saw your post in an old thread relating to water purification.

From a Survivalist's point of view, surely the biological contaminants can be easily removed by heat.

In extreme conditions, perhaps the removal of heavy metals, radioactivity etc would be more useful ?

[Edited on 9-7-2015 by aga]

blogfast25 - 9-7-2015 at 13:35

It's very difficult (virtually impossible) to prepare solid ferrates, they're far too unstable. See this preparation where the solid obtained contained only barely 42 w% of the target Na2FeO4:

http://pubs.acs.org/doi/abs/10.1021/ja01155a545

One could look into preparing ferrates by anhydrous route but with water the resulting ferrate invariably oxidises water! Yield is therefore by definition small.

[Edited on 9-7-2015 by blogfast25]

chemister2015 - 9-7-2015 at 15:24

http://pubs.acs.org/action/doSearch?text1=ferrates%28VI%29&a...

ave369 - 9-7-2015 at 22:38

Quote: Originally posted by blogfast25  

One could look into preparing ferrates by anhydrous route but with water the resulting ferrate invariably oxidises water! Yield is therefore by definition small.

[Edited on 9-7-2015 by blogfast25]


My idee fixe is using solvents other than water. Specifically, solvents that can be cooled to very low temperatures and remain liquid. The idea is that low temperatures will make the ferrates stable and more easily isolable.

My current idea is electrolyzing a very cold KOH-EtOH solution. As far as I know, ferrates do not react with very cold EtOH and aren't soluble in it at all. And if you replace KOH with NaOH, you will get an easy way to isolate solid sodium ferrate.

[Edited on 10-7-2015 by ave369]

AJKOER - 10-7-2015 at 04:22

Per this source, "Hydrogen Generation from Alkaline Solutions of Methanol and Ethanol by Electrolysis" by Islay Ulusoy, I suspect you will convert your KOH into K2CO3 and generate H2 at unusually low electric energy consumption as compared to the standard electrolysis of water.

Link: http://www.academia.edu/534836/Hydrogen_Generation_from_Alka... and see also http://www.evworld.com/news.cfm?newsid=14290 and http://www.sciencedirect.com/science/article/pii/S0378775306...

My speculation as to the underlying radical chemistry with methanol containing water (and similarly with ethanol) at low voltage:

CH3OH ---electrolysis→ CH3 • + HO•

(this source acknowledges the formation of the Methyl radical under electrolysis: http://pubs.acs.org/doi/abs/10.1021/ja01529a021?journalCode=... and see also http://www.ncbi.nlm.nih.gov/pubmed/20580488 and http://journals.plos.org/plosone/article?id=10.1371/journal.... with respect to possible hydroxyl radical and H2O2 formation via electrolysis):

CH3 • + H2O → CH3OH + 1/2 H2 (g)

where the active Methyl radical attacks water (see https://en.m.wikipedia.org/wiki/Methyl_radical ), which may explain the low energy formation of hydrogen as compared to the higher voltage required to liberate H2 directly from H2O.

HO• + HO• → H2O2

and finally:

CH3OH + 3 H2O2 → CO2 + 5 H2O
2 KOH + CO2 → K2CO3 + H2O

Note, if any O2 was directly formed, I would expect:

CH3• + O2 → CH3O2•
CH3O2• + O2 → HCHO + HO2•
HO2• + HO2• → H2O2 +O2 (reference: http://www2.nau.edu/~doetqp-p/courses/env440/env440_2/lectur... )

that is, a distinct odor of Formaldehyde would be detected. Also, the hydrogen gas created may not be pure due to the formation of a small amount of Ethane via:

CH3• + CH3• → C2H6

The overall number of products formed is undoubtedly more extensive, but if my take on the chemistry approaches correct, perhaps I have contributed to a better understanding of the underlying processes.
---------------------------------------------

Note, with respect to the electroIysis, I would also expect, per Wikipedia on the Methyl radical:

"Methyl is very corrosive to metals, forming methylated metal compounds:

M + n CH3• → M(CH3)n "

that an Aluminum electrode in CH3OH, especially containing water, would quickly be subject to corrosion.

Caution: the net reaction upon the electrolysis of say just 404 cc of pure methanol contained in an additional body of water could potentially produce 10 moles of CO2 (or some 224 liters) and 5 moles of hydrogen (112 liters) presenting both a suffocation and a fire/explosion hazard. Fumes from other reaction products are likely toxic and carcogenic as well.

[Edited on 10-7-2015 by AJKOER]

DFliyerz - 10-7-2015 at 09:14

I have been very interested in ferrates this past week, and have been doing some experiments in preparing barium ferrate, which is a stable insoluble ferrate salt. I'm linking two videos, one on methods of preparing ferrate, and one on preparing barium ferrate/the properties of ferrates.

https://www.youtube.com/watch?v=5u8LOfTrO6w
https://www.youtube.com/watch?v=2L906yG0-14

The Volatile Chemist - 10-7-2015 at 09:30

Interesting, as I have a small quantity of soluble Barium salt. Could be of interest. Aren't there a lot of interesting mixed-oxidation oxides of iron that are interesting, at least to a degree?

DFliyerz - 12-7-2015 at 16:00

I have some observations on preparing sodium ferrate by the first method outlined in the first video I linked, and couldn't think of any other place to report them, so here it goes:
When filtering the solution through fiberglass, I've found that the filtrate starts green and eventually transitions to the violet/amethyst of sodium ferrate, but the previous filtrate stays green. I'm not sure what's causing this, but an old paper I found suggests the formation of sodium manganate.

ChemPlayer_ - 12-7-2015 at 16:59

DFliyers, thanks for sharing; great videos!

I'd read that using stronger solutions of iron (III) chloride and freshly made hypochlorite it was actually possible to obtain dark purple crystals of the alkali ferrate salts on cooling the reaction mixture. Did you try this or have any luck with obtaining solid salts?

They're interesting compounds and definitely going on my experimentation list.

AJKOER - 16-7-2015 at 12:46

Actually, just found this by accident, "A New Way to Prepare Ferrate Solutions by Microwave Assisted Hypochlorite-Oxidation Method". Per the abstract, to quote:

"In this article, a new way to prepare ferrate solution was proposed by combining microwave irradiation and hypochlorite-oxidation method. It was demonstrated that the reactivity speed can be improved as high as 30 times compared to the single hypochlorite -oxidation method."

To repeat, a 30 fold reaction speed bump!

My take is that apparently, free radicals, which can be formed via microwave irradition, are likely very instrumental in ferrate creation.

Link: http://www.scientific.net/AMR.554-556.632

The Volatile Chemist - 16-7-2015 at 12:54

Where did you read this?

aga - 16-7-2015 at 14:25

Quote: Originally posted by ave369  
As far as I know, ferrates do not react with very cold EtOH and aren't soluble in it at all.

Solubility at various temperatures is something you can determine experimentally, although there is probably a mathematical route working that out.

Thermodynamics, Gibs's/Higgs's and Kinetics are involved in working out if stuff will react, and at what rate in a given temperature/situation.

It would be of IMMENSE interest (at least to me) if someone could give the workings out of an example.

E.g. NaOH and HCl readily react at room temperature at fairly low concentrations.

At what temperature/relative concentrations will they Not react (as in an imperceptible rate) ?

(0 of everything obviously would be a case, yet significantly less interesting)

AJKOER - 16-7-2015 at 14:33

Quote: Originally posted by The Volatile Chemist  
Where did you read this?


More details available for a fee at D. S. Li, X. L. Chen, "A New Way to Prepare Ferrate Solutions by Microwave Assisted Hypochlorite-Oxidation Method", Advanced Materials Research, Vols. 554-556, pp. 632-635, Jul. 2012.

And no, I have not purchased the article.

On free radical formation, if that your question, speaking generally, they can be created via irradition including a radioactive source, plasma pulses, microwaves, uv light (atmospheric chemistry is mosty radical based), high temperatures (like in the science of combustion of gases), electrolysis (pollution abatement technique to breakdown nasty organics via methyl, actyl radicals,..)....

The augmenting role of radicals in the formation of ferrates is my speculation as a microwave heating of aqueous mixtures differs not significantly otherwise, in my opinion, from normal heating by contact with a burning fossil fuel with the noted exceptions of rapid heating, more uniform heating and bulk heating effects. As such, assuming a non-thermal effect, a likely guess on my part is a radical based increase in reactivity, which is characteristic of some radicals themselves (and how they can be used to boost chlorate formation, for example, in one of my more obtuse SM threads, see http://www.sciencemadness.org/talk/viewthread.php?tid=34429#...).

I should also note with respect to ferrous and ferric ions there is a body of work referred to as Fenton chemistry, Electro-Fenton and Photo-Fenton chemistry involving the formation of the highly reactive hydroxyl radical (as some background, see for example, "Fundamental Mechanistic Studies of the Photo-Fenton Reaction for the Degradation of Organic Pollutants" at https://www.google.com/url?sa=t&source=web&rct=j&... ). In the case of Photo-Fenton chemistry, the precise pH at which the OH. radical is efficiently produced appears to be a function of the absorbance ability of the particular Iron species existing at that pH. In the case of ferrate formation, different Ferric species are likely present at higher pHs which may contribute to a similar ferrate promoting radical based reaction.

The cited article above may only state factually the performance jump, and be clueless (or perhaps unwiling to speculate), as to the underlying (I would say possibly radical based) chemistry involved.


[Edited on 17-7-2015 by AJKOER]

blogfast25 - 16-7-2015 at 16:05

Quote: Originally posted by AJKOER  
Actually, just found this by accident, "A New Way to Prepare Ferrate Solutions by Microwave Assisted Hypochlorite-Oxidation Method". Per the abstract, to quote:

"In this article, a new way to prepare ferrate solution was proposed by combining microwave irradiation and hypochlorite-oxidation method. It was demonstrated that the reactivity speed can be improved as high as 30 times compared to the single hypochlorite -oxidation method."

To repeat, a 30 fold reaction speed bump!

My take is that apparently, free radicals, which can be formed via microwave irradition, are likely very instrumental in ferrate creation.

Link: http://www.scientific.net/AMR.554-556.632


All you've got is an abstract, 1 page and a nonsensical sentence: "compared to the single hypochlorite -oxidation method".

But we all know you're pleased with a dead mouse as long as it's got the word 'hypochlorite' in it.

Pfffff...

AJKOER - 16-7-2015 at 16:35

Here is a Patent (please see "Method for preparing alkali metal ferrates and novel alkali metal ferrate granules", US 20050053543 A1), to quote the abstract:

"The invention concerns a method for preparing an alkali metal ferrate, which comprises the following steps: (i) preparing a reaction mixture comprising at least an iron salt, at least an alkali or alkaline-earth metal hypochlorite and an alkali metal hydroxide; (ii) heating the mixture prepared at step (i) to a temperature ranging between 45 and 75° C., so as to form the alkali metal ferrate; (iii) recuperating the alkali metal ferrate formed at step (ii). The invention also concerns novel alkali metal ferrates."

Link: http://www.google.com/patents/US20050053543 that discusses the hypochlorite method among other things.

So, replace "heat" with "heat in a microwave", and yield may improve. In fact, the author of this Patent cites said use, to quote:

"8. Process according to claim 1, characterized in that said mixture is heated using infrared."

[Edited on 17-7-2015 by AJKOER]

blogfast25 - 16-7-2015 at 16:54

Sigh. The cations here are spectator ions. I can make any solution of any alkali metal ferrate in a beaker on a hot plate.

Re. how microwaves help in the process, please put up or shut up.

You're chasing another mirage, one of your favourite pass times...

AJKOER - 16-7-2015 at 18:46

Blogfast, a reference "Microwave-Assisted Chemistry: Synthetic Applications for Rapid Assembly of Nanomaterials and Organics", to quote:

"Microwave-assisted chemical reactions are now well-established practices in the laboratory setting although some controversy lingers as to how MW irradiation is able to enhance or influence the outcome of chemical reactions. Much of the discussion has focused on whether the observed effects can, in all instances, be rationalized by purely thermal Arrhenius-based phenomena (thermal microwave effects), that is, the importance of the rapid heating and high bulk reaction temperatures that are achievable using MW dielectric heating in sealed reaction vessels, or whether these observations can be explained by so-called “nonthermal” or “specific microwave” effects."

Link: http://pubs.acs.org/doi/abs/10.1021/ar400309b

blogfast25 - 17-7-2015 at 08:29

I.o.w whether microwaves have an effect other than purely thermal remains up for grabs. Good luck to anyone trying to separate the two effects (assuming there really is a non-thermal effect).

Try and be a little less credulous.

AJKOER - 17-7-2015 at 10:41

Actually Blogfast, here is a tenative explanation as to how yield could be augmented with microwaving (MW):

1. We start by adding NaOH to say aqueous Fe2(SO4)3, forming a nano-suspension of Fe(OH)3. MW is say to be benefical in creating more uniform particle sizes,...

2. Apparently, Fe(OH)3 is capable of creating hydroxyl radicals on exposure to UV in a manner similar to TiO2 (source: http://pubs.acs.org/doi/abs/10.1021/j100354a021 ).

3. Apparently, TiO2 forms hydroxyl radicals both via UV and MW. Please see http://www.researchgate.net/publication/223007525_Hydroxyl_r... So, let's assume that Fe(OH)3 behaves likewise in this regard.

4. The equilibrium reaction between hydroxyl radicals and hydrogen peroxide (see, for example, https://books.google.com/books?id=rjEgq4Qf69MC&pg=PA296&...

OH. + OH. + M = H2O2 (where M is a mass)

5. The reaction of H2O2 in the presence of a basic hypochlorite forms nascent oxygen. Please see one of my prior SM's thread and references therein at http://www.sciencemadness.org/talk/viewthread.php?tid=31729#...

6. The action of any created nascent oxygen could augment the power of the hypochlorite in forming ferrates.

Interestingly, here is an actual historical journal report of the formation of both ferrates and nascent oxygen from the action of Fe(OH)3 on chlorine bleach over time which likely encompasses uv exposure (see https://books.google.com/books?id=tjIAAAAAMAAJ&pg=PA731&... ).

[Edited on 18-7-2015 by AJKOER]

blogfast25 - 17-7-2015 at 13:14

Quote: Originally posted by AJKOER  

3. Apparently, TiO2 forms hydroxyl radicals both via UV and MW. Please see http://www.researchgate.net/publication/223007525_Hydroxyl_r... So, let's assume that Fe(OH)3 behaves likewise in this regard.



And why would we 'assume' that, huh?

As so often with you, lengthy chains of assumptions designed to make you 'sound interesting'.

Sadly you remain as crashing a bore as ever. Not to mention a prime example of how a little knowledge is a dangerous thing...

But hey, you're in a hole, so keep digging!

[Edited on 17-7-2015 by blogfast25]

aga - 17-7-2015 at 13:49

Personally i like AJ's posts.

Lots of things to click on.

I like yours too, but will not say which i like best ;)

[Edited on 17-7-2015 by aga]

blogfast25 - 17-7-2015 at 13:59

Quote: Originally posted by aga  

Lots of things to click on.



How about a random tinkerwebs surf? Infinite links!

Start here:

https://en.wikipedia.org/wiki/Molten_salt_reactor

... and let the Click-a-Thon begin! :D

[Edited on 17-7-2015 by blogfast25]

aga - 17-7-2015 at 14:03

Context m'lud. Context.

aga - 17-7-2015 at 14:06

Quote: Originally posted by AJKOER  
formation of both ferrates and nascent oxgygen

What is 'nascent' oxygen ?

Oxygen that is about to be born or has just been made ?

AJKOER - 17-7-2015 at 15:11

It is basically oxygen in an excited state.

See my previous provided link to a thread I have on it. Wikipedia also has some commentary as well.

It is actually fun stuff. It glows in the dark, doesn't like water (a quenching reaction forming H2O2), and exists for less than an hour!

[Edited on 18-7-2015 by AJKOER]

AJKOER - 19-7-2015 at 09:23

I can up with some very strong work on how MW radiation may be augmenting the formation of ferrates based on my research on Fe(OH)3. See, for example, "Nanostructures: Synthesis, Functional Properties and Application" edited by Thomas Tsakalakos, ..., pages 56 and 57 at https://books.google.com/books?id=z2ryCAAAQBAJ&pg=PA57&a...

To paraphrase the authors, apparently, Ferric hydroxide is very good at absorbing electromagnetic radiation and converting it into a homogeneous and effective heating source. In a nano-suspension of Fe(OH)3, this can increase nucleation sites and limit average particle growth while narrowing the particle size distribution. This is attributed to increase kinetics and more homogeneous heating resulting in an increase in yield and reduced processing time. In general, magnetic materials under MW heating are superior as a consequence of a ferromagnetic resonance effect.

Also, I found some research supportive of the possible presence/influence of radicals. Please see http://www1.lsbu.ac.uk/water/magnetic_electric_effects.html ),
but not as a strong case, in my opinion, over the specific case of Fe(OH)3 as detailed above. To quote:

"The solubility properties of the water will change in the presence of such fields and may result in the concentration of dissolved gases and hydrophobic molecules at surfaces followed by reaction (for example, due to reactive singlet oxygen (1O2) or free radical formation such as OH·) or phase changes (for example, formation of flattish surface nanocavities, termed nanobubbles [506]). It is also possible that these processes may result in the production of low concentrations of hydrogen peroxide in a similar manner to mechanical vibrations [1066, see equations]. Such changes can clearly result in effects lasting for a considerable time, giving rise to claims for 'memory' effects. One of the curious facts, concerning reports of the effects of magnets and electromagnetic radiation on the properties of water, is the long lifetime these effects seem to have (for example, [757]..."

And also:

"In addition to the breakage of hydrogen bonds electromagnetic fields may perturb in the gas/liquid interface and produce reactive oxygen species [110]. "

From the above, I would suspect that increasing solid or solution contact with air (or better O2) would promote radical formation and a possible increase in yield.

As a safety note, the authors note that MW radiation is applied in a few seconds burst followed by powering off. This obviates rapid extreme temperature changes and the risk of explosion.

[Edited on 19-7-2015 by AJKOER]

The Volatile Chemist - 21-7-2015 at 13:06

Quote: Originally posted by ChemPlayer_  
DFliyers, thanks for sharing; great videos!

I'd read that using stronger solutions of iron (III) chloride and freshly made hypochlorite it was actually possible to obtain dark purple crystals of the alkali ferrate salts on cooling the reaction mixture. Did you try this or have any luck with obtaining solid salts?

They're interesting compounds and definitely going on my experimentation list.

My where did you read this was directed here, not to AJOKER.

ave369 - 4-8-2015 at 15:14

I've found a very interesting material about rare and unusual ferrates, by the Russian chemist under the pseudonym "Odin". Here I shall provide the one and only translation into English. A Sciencemadness exclusive!

An analysis of double displacement of ferrates

1. Description of the experiment

The displacement reactions were performed with potassium ferrate (VI), which is soluble in water. The goal of the experiment is checking if ferrate compounds of various metals exist. Three groups of reactions were performed by me:

Group I: with alkali metal chloride salts and ammonium chloride

Group II: with alkaline earth metal chloride salts

Group III: with miscellaneous metal salts including rare earths, silver and lead

2. Analysis of displacement reactions with chlorides of alkali metals and ammonium

Test tube 11: NH4Cl
Test tube 12: LiCl
Test tube 13: NaCl
Test tube 14: RbCl
Test tube 15: CsCl

To each test tube, crystals of potassium ferrate (PF) were added.
In test tube 11 PF instantly changed color to light brown, an insoluble brown precipitation formed. The liquid in the test tube became yellow (supposedly a colloidal solution of Fe(OH)3). I suppose that ammonium ferrate does not exist.
In test tube 12 PF was fully dissolved. There are all reasons to suppose that Li2FeO4 exists, is stable and water-soluble.
In test tube 13 PF was fully dissolved. There are all reasons to suppose that Na2FeO4 exists, is stable and water-soluble.
In test tube 14 PF formed a cherry-pinkish suspension. After a while, the suspension precipitated to the bottom of the tube, the upper part of the liquid became clear. A pink precipitate formed. There are all reasons to suppose that Rb2FeO4 exists, is stable and has low solubility in water.
In test tube 15 PF was fully dissolved. There are all reasons to suppose that Cs2FeO4 exists, is stable and water-soluble.

3. Analysis of displacement reactions with chlorides of alkaline earth metals

The salts were dissolved in the following test tubes:

Test tube 21: MgCl2
Test tube 22: CaCl2
Test tube 23: SrCl2
Test tube 24: BaCl2

To each test tube, PF was added.
In test tube 21, PF crystals slowly sank to the bottom of the tube, leaving purple trails. After shaking, a cherry-colored suspension formed. It started to slowly precipitate, but after an hour it fully decomposed. I suppose that MgFeO4 exists, has low solubility in water and is unstable in aqueous solution.
In test tube 22, PF crystals slowly sank to the bottom of the tube, leaving brownish cherry-red trails. After shaking, a suspension of this color formed. One hour later, the suspension precipitated, its color did not change. I suppose that CaFeO4 exists, has low solubility in water and is stable.
In test tube 23, PF crystals slowly sank to the bottom of the tube, leaving cherry-red trails. After shaking, a suspension of this color formed. One hour later, the suspension precipitated, its color did not change. I suppose that SrFeO4 exists, has low solubility in water and is stable.
In test tube 24, PF crystals sank to the bottom of the tube, leaving no trails, which hints at the very low solubility of this ferrate. After shaking, a suspension of cherry-red color formed. One hour later, the suspension precipitated, faster than the other AEMs, its color did not change. I suppose that BaFeO4 exists, is stable and almost insoluble in water.

4. Analysis of displacement reactions with miscellaneous metal salts

The salts were dissolved in the following test tubes:

Test tube 31: AgNO3
Test tube 32: Pb(NO3)2
Test tube 33: Y2(SO4)3
Test tube 34: Ce2(SO4)3
Test tube 35: Gd2(SO4)3

To each test tube PF was added.

In test tube 31, a thick black flakey precipitate was instantly formed. Shaking resulted in more precipitate, the liquid became grayish. Soon a white precipitate formed in addition to the black one. I suppose that PF and the silver nitrate reacted with the formation of Fe(OH)2, Ag2O2 and KNO3. Since there was no brown precipitate, no Fe(OH)3 was formed. The black precipitate has to be Ag2O2, because Ag2O is dark brown.

In test tube 32, a light brown precipitate was immediately formed. Shaking did not change the color. Supposed products of the reaction: Fe(OH)3, Pb3O4, KNO3

In test tube 33, PF was fully dissolved and a gas was evolved. After adding another portion of PF, a precipitate of Fe(OH)3 was formed. I could not identify the products of the reaction.

In test tube 34, PF formed a pink suspension, a gas was evolved. Soon the suspension precipitated to the bottom of the tube, and after 2 hours the precipitate partially decomposed. Probably, Ce(m)(FeO4)(n) exists, is unstable and insoluble in water.

In test tube 35, PF partially dissolved, evolving a gas. A yellowish-pink suspension formed and precipitated. Its color soon changed to brown. 2 hours later, no more changes in the color of the precipitate were found. Probably, Gd(m)(FeO4)(n) exists but is unstable and insoluble. The solution contained a low amount of Gd ions, because gadolinium sulfate has low solubility in cold water.

Later, the author comments that he had discovered a new method of separating rubidium from other alkali metals, with ferrate a convenient anion with which rubidium can be precipitated and other alkali metals cannot.

[Edited on 4-8-2015 by ave369]

blogfast25 - 4-8-2015 at 16:03

@ave369:

Please provide a reference to 'Odin's' text. Who made the translation?

ave369 - 4-8-2015 at 22:20

I did. The original is here:

http://mirror.pirotehnika-ruhelp.com/forum/303ca03332b7dbc26...

If anyone here can read Russian, of course.

deltaH - 5-8-2015 at 00:24

Some nice photos on the original link provided by ave369. Nice work translating, thank you!

ave369 - 5-8-2015 at 00:52

Odin is also the author of a new method of ferrate synthesis, the "ferrite-hypochlorite" method. The difference from the standard hypochlorite method is that a ferrite with a formula of KFeO2, synthesized by melting KOH and Fe2O3 together, is used as the starting compound which is oxidized by hypochlorite. Odin claims that this improves the yield and stability of resulting potassium ferrate.

I'm going to repeat all his experiments with ferrates when my Buchner funnel arrives. I'll write all results here on the forum.

deltaH - 5-8-2015 at 01:56

Although hypothetical, I've shown by calculation that the fusion of potassium bromate and iron oxide in the presence of excess base could also yield potassium ferrate.

If you're interested see http://www.sciencemadness.org/talk/viewthread.php?tid=26760#...

Please note there is a mistake in my final equation, I forgot to add the KBr, it should read:

3Fe3O4(s) + 5KBrO3(s) + 18KOH(l) => 9K2FeO4(s) + 5KBr(s) + 9H2O(g)

[Edited on 5-8-2015 by deltaH]

The Volatile Chemist - 9-8-2015 at 14:28

Interesting. Nice translation ave. I don't suppose ferric ferrite exists? :P

ave369 - 10-8-2015 at 10:02

Quote: Originally posted by The Volatile Chemist  
Interesting. Nice translation ave. I don't suppose ferric ferrite exists? :P


Ferric ferrite, which means iron (III) ferrate (III), doesn't exist. Amphoteric hydroxides do not work this way.

Ferrous ferrite, that is, iron (II) ferrate (III), exists. It is better known as Fe3O4.

Ferric ferrate, that is, iron (III) ferrate (VI), as far as I know, doesn't exist. I don't think any amphoteric metal ferrates exist at all, at least in aqueous form: ferric acid appears to be at least of middle strength, so its salts with weak bases will be acidic. And "acidic" and "ferrate" aren't words that like to be in the same sentence.

Update: found the calculated pKa of ferric acid. It's 3.5. It's middle strength, weaker than phosphoric, stronger than formic. Amphoteric metal ferrates will be acidic, and, therefore, will decompose themselves.

My little addendum to Odin's text above. There is a probability that some of the precipitates in the tubes are, in fact, double salts (I've read something in some paper about a Na, Ca double ferrate, so maybe others exist). When my Buchner funnel arrives (and when my tartric acid arrives, too), I'm going to perform a modified form of Odin's experiments, by filtering, washing and then decomposing all precipitated ferrates with acid and testing the resulting solutions with tartric acid to detect potassium.

If there will be no potassium in the magnesium ferrate, then bingo! - I've found my holy grail of water purificants. As far as Odin's results can tell, it has a low but noticeable solubility in water, small amounts of it dissolve in water and then decompose, making room for more magnesium ferrate to dissolve, and the products are harmless to ingest: magnesium and iron hydroxides. Then it is possible to take this ferrate to a camping trip and purify and disinfect any water in the forest with it.


[Edited on 10-8-2015 by ave369]

Mabus - 14-6-2016 at 07:55

Wow this is amazing. It's always nice to see less known chemistry being experimented.

zwt - 12-8-2016 at 10:20

How does one determine the concentration of ferrate in a sample? It seems like it decomposes too quickly in solution for an accurate titration. For the purposes of this question, assume you have a solid purporting to be or contain "sodium ferrate".

Thanks.

clearly_not_atara - 12-8-2016 at 11:57

To me the most interesting potential application of ferrate is either as:

* a precursor to permanganate (VII), in which case there is no issue with accidentally producing manganate (VI) instead, since Fe6+ oxidizes Mn6+.

* a co-oxidant recycling MnO2 to MnO4-

Permanganate, after all, is dangerous, (relatively) expensive, difficult to produce, and explosive.

Quote:
How does one determine the concentration of ferrate in a sample? It seems like it decomposes too quickly in solution for an accurate titration. For the purposes of this question, assume you have a solid purporting to be or contain "sodium ferrate".


It might be possible to titrate with ammonia, which I assume undergoes a relatively predictable oxidation which does not generate any acids (which are what usually decomposes FeO4(2-)):

2FeO4(2-) + 2NH3 (aq) >> N2 + 4OH- + Fe2O3 + H2O

I just made that up, though, so first you should check that ferrate does not react too enthusiastically with ammonia. Anyway one possibility is to simply add a little ferrate solution to excess ammonia, decompose it all, and weigh the iron oxide.

This only makes sense if the ferrate is dissolved. Ferrate is, however, quite soluble, so it should be fine to first take it up in alkaline water (pH > 12).

zwt - 13-8-2016 at 06:35

Quote: Originally posted by clearly_not_atara  
Anyway one possibility is to simply add a little ferrate solution to excess ammonia, decompose it all, and weigh the iron oxide.

This only makes sense if the ferrate is dissolved. Ferrate is, however, quite soluble, so it should be fine to first take it up in alkaline water (pH > 12).


Good point. Strong hydroxide solution should slow the decomposition, as hydroxide is one of the products of the decomposition.

Is there any risk of the iron(III) oxide/hydroxide produced by the decomposition redissolving in a concentrated hydroxide solution? If not, then it should be as simple as dissolving the sample in hydroxide solution, filtering (no paper, of course), allowing to decompose (no ammonia required), filtering, and drying and weighing the precipitate. I know pottery-grade ferric oxide won't really dissolve in hydroxide solution, but I've also read (in other threads here) that pottery-grade oxides can be extremely inert compared to their freshly-precipitated counterparts.

byko3y - 13-8-2016 at 09:05

What is application of ferrates? Is there any task that can't be done with permanganate, but can be done with ferrate?
By the way, just as the ferrate is prepared by oxidation of iron salt with bleach, the permanganate can prepared via the same route.

zwt - 13-8-2016 at 09:28

There are applications where ferrates may be superior to permanganates, though mainly from an industrial perspective. Because ferrates have iron instead of manganese, they may be slightly cheaper, the products of decomposition less toxic, and more "environmentally friendly" in industrial applications and water treatment. It is also a more powerful oxidizer in solution than permanganate, though I haven't seen an application where only ferrate would work.

The liquid bleach method gives some purple color with either iron or manganese oxides, but in both cases, the solutions are dilute and prone to decompose (the ferrate much more so). I'm exploring water-free methods of ferrate synthesis.
____________________________
Quote: Originally posted by zwt  
Is there any risk of the iron(III) oxide/hydroxide produced by the decomposition redissolving in a concentrated hydroxide solution?
Anybody?

[Edited on 13-8-2016 by zwt]

byko3y - 13-8-2016 at 11:52

Bleach methods yields a solid product with approx 40% amount of ferrate (Preparation of Sodium Ferrate(VI)), the rest are hydrous ferric oxides

zwt - 13-8-2016 at 12:16

Ah, that's with specially-prepared super-concentrated sodium hypochlorite/hydroxide solution. When you said "bleach", I thought you meant <10% hypochlorite laundry bleach. The latter will produce a small amount of ferrate or permanganate when ferric oxide or manganese dioxide is added, but is not a practical synthesis.

byko3y - 13-8-2016 at 13:06

In my region bleach concentration is usually 3-5% because it's stored for a long time at r.t, despite the fact a freshly manufactured bleach is 30% strong. In the article researchers used 55g/100ml bleach, but it has such a short shelf time it should be used immediately. Same applies to the regular 30%+ commercial bleach. You can easily reconcentrate the bleach and get any concentration you want, and it's a preferred way to obtain highly concentrated bleach. And if you look for a long shelf life, then ferrates are not the compounds you should mess with.

Magpie - 13-8-2016 at 16:09

Quote: Originally posted by byko3y  
In my region bleach concentration is usually 3-5% because it's stored for a long time at r.t, despite the fact a freshly manufactured bleach is 30% strong. In the article researchers used 55g/100ml bleach, but it has such a short shelf time it should be used immediately. Same applies to the regular 30%+ commercial bleach. You can easily reconcentrate the bleach and get any concentration you want, and it's a preferred way to obtain highly concentrated bleach. And if you look for a long shelf life, then ferrates are not the compounds you should mess with.

Magpie - 13-8-2016 at 16:10

How do you reconstitute bleach?

byko3y - 13-8-2016 at 18:14

Acid-base extraction :P Kidding. Reconstitute - is a correct word. Evaporate the chlorine from the bleach with acid, then absorb it with alkali. The device can be as simple, as a flask with bleach-acid covered with beaker, standing in the center of a large closed but non-sealed vessel with a NaOH solution and some ice on a bottom of it. Strong mineral acids react too virgously.
I prefer to use two separated flasks, but that's just because I really want to see the process and not just a result.

ave369 - 19-8-2016 at 12:43

Quote: Originally posted by byko3y  

I prefer to use two separated flasks, but that's just because I really want to see the process and not just a result.


Do you need a funnel and beaker trap to dissolve chlorine in alkali? Is there any risk of suckback?

byko3y - 19-8-2016 at 14:38

Tendence to suck back is observable, but it's not as strong as with H2O-HCl case. Anyway, I prefere not to place chlorine inlet under surface of NaOH solution.