| Pages:
1
2
3 |
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
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
International Hazard
Posts: 5128
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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 
I took pictures of the above if anyone is interested.
|
|
|
Saerynide
National Hazard
Posts: 954
Registered: 17-11-2003
Location: The Void
Member Is Offline
Mood: Ionic
|
|
Me is interested 
"Microsoft reserves the right at all times to monitor communications on the Service and disclose any information Microsoft deems necessary to...
satisfy any applicable law, regulation or legal process"
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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.
|
|
|
chemoleo
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
Another Ferrate Prep!
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!
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
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
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
| 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]
|
|
|
Hjalmar_Poelzig
Harmless
Posts: 20
Registered: 4-8-2003
Member Is Offline
Mood: No Mood
|
|
seen this?
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.
<a href=http://rbenda.bei.t-online.de/ret/ret33.jpg>Der Mann mit den Röntgenaugen </a>
|
|
|
Hjalmar_Poelzig
Harmless
Posts: 20
Registered: 4-8-2003
Member Is Offline
Mood: No Mood
|
|
other patents
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.
<a href=http://rbenda.bei.t-online.de/ret/ret33.jpg>Der Mann mit den Röntgenaugen </a>
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
FeCl3 may be a bad alternative to Fe(NO3)3
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.
|
|
|
JohnWW
International Hazard
Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline
Mood: No Mood
|
|
Ferrates & Other High-Valent Fe Compounds
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
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
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
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
Hazard to Self
Posts: 64
Registered: 1-6-2004
Location: 6feet south
Member Is Offline
Mood: No Mood
|
|
Question
If you have a mix of KOH + H2O with a iron electrode, will the K2FeO5 decompose because of the OH- and H+?
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
Hazard to Self
Posts: 64
Registered: 1-6-2004
Location: 6feet south
Member Is Offline
Mood: No Mood
|
|
Thanks Bormic.
[Edited on 13-10-2004 by tokat]
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
Hummm... AC electrolysis gives better yeilds....
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
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
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
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
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
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
BromicAcid
International Hazard
Posts: 3253
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
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]
|
|
|
| Pages:
1
2
3 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
