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aga
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Context m'lud. Context.
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aga
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What is 'nascent' oxygen ?
Oxygen that is about to be born or has just been made ?
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AJKOER
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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]
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AJKOER
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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]
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The Volatile Chemist
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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.
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ave369
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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]
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blogfast25
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@ave369:
Please provide a reference to 'Odin's' text. Who made the translation?
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ave369
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I did. The original is here:
http://mirror.pirotehnika-ruhelp.com/forum/303ca03332b7dbc26...
If anyone here can read Russian, of course.
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deltaH
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Some nice photos on the original link provided by ave369. Nice work translating, thank you!
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ave369
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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.
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deltaH
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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]
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The Volatile Chemist
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Interesting. Nice translation ave. I don't suppose ferric ferrite exists? 
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ave369
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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]
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Mabus
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Wow this is amazing. It's always nice to see less known chemistry being experimented.
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zwt
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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.
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clearly_not_atara
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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).
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zwt
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| 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.
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byko3y
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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.
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zwt
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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?
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Anybody?
[Edited on 13-8-2016 by zwt]
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byko3y
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Bleach methods yields a solid product with approx 40% amount of ferrate (Preparation of Sodium Ferrate(VI)), the rest are hydrous ferric oxides
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zwt
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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.
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byko3y
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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.
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Magpie
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| 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. |
The single most important condition for a successful synthesis is good mixing - Nicodem
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Magpie
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How do you reconstitute bleach?
The single most important condition for a successful synthesis is good mixing - Nicodem
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byko3y
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Acid-base extraction 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.
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