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Butterflywings
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Reduction of cerium oxide to metal
One of the method found was by electrochemical reduction. I am searching for other ways to reduce cerium oxide to metal.
http://ac.els-cdn.com/S0013468610016737/1-s2.0-S001346861001...
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AndersHoveland
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cerium fluoride can be reduced with calcium
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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elementcollector1
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I've been discussing and thinking about this for a while, and the best way we came up with on another forum would be to have a crucible with your
mixture surrounded by a steel can of lit coals, and a blanket of argon being slowly pumped through. The mixture would be CeO2 and Ca,
stochiometrically mixed.
Not an easy setup for the amateur, but it will give you pure cerium metal.
I'd like to see if anyone can find a 'wet' method for cerium, like how lithium can be electrowon from its salt in pyridine.
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bfesser
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Couldn't you have posted this in your <a href="http://www.sciencemadness.org/talk/viewthread.php?tid=22194">other cerium oxide thread</a>?
[Edited on 11/10/12 by bfesser]
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Butterflywings
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elementcollector1:
Do you have the procedures to the method that you explained?
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elementcollector1
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What do you mean? This isn't an official patent or something, it's a DIY thing we came up with. If you want stochiometry or something, the ratio is
3.28 CeF3 to 1 Ca, in grams. (I think CeF3 would be better than CeO2, as the rare earth fluorides generally tend to be more volatile than their
oxides.)
The crucible itself would be a large soup can full of coals, with a smaller, sealed soup can with the reactants inside, with steel tubes leading
through the inner crucible and pumping argon.
Would someone mind calculating the reaction enthalpy of CeO2-Ca mix and CeF3-Ca mix to see which one actually is better? I haven't gotten quite that
far in my chemistry textbook yet. 
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AJKOER
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OK, here is an interesting alternative and speculative idea although more expensive. First, prepare Cerium(III) oxalate, Ce2(C2O4)3, by the reaction
of oxalic acid with cerium(III) chloride. Then, like many oxalates, it decomposes on heating.
Possible products are:
Ce2(C2O4)3 --Heat-?-> 2 Ce + 6 CO2
in a manner simliar to the decomposition of Iron(II) oxalate (see http://www.youtube.com/watch?v=adhE1m2vX38&feature=relat... )
FeC2O4 --> Fe + 2 CO2
or:
Ce2(C2O4)3 --Heat-?-> 2 CeO2 + 4 CO + 2 CO2
which would be a path similar to Magnesium oxalate (see http://pubs.acs.org/doi/abs/10.1021/j100886a013 ):
MgC2O4 --> MgO + CO + CO2
or other combination paths, but I would not be surprised if the first reaction is what is observed.
[Edited on 15-11-2012 by AJKOER]
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woelen
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I hardly can believe that ceric oxalate will give metallic cerium on heating. Cerium simply is too reactive. Some oxalates do give metal on
decomposition, e.g. silver oxalate, but more reactive metals give oxide. Many people say that ferrous oxalate gives pyrophoric iron, but as far as I
know this is not iron, but iron(II) oxide, which also is very reactive. With cerium I hence expect ceric oxide or maybe cerous oxide, but not cerium
metal.
If things indeed turn out otherwise, then I certainly am open for correction, but I would like to see a literature reference for that, or a
confirmation, based on personal experimental results and not on speculation.
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AJKOER
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Have found a reference citing that the decomposition of Cerium Oxalate proceeds in two steps with the formation of Cerium oxide and not metallic Ce.
See http://www.springerlink.com/content/h038736p2617479g/
To quote from the Abstract;
"Cerium oxalate and mixed cerium-gadolinium oxalates containing 20 and 50 mol% gadolinium were subjected to thermal decomposition. Thermal analysis
showed that cerous oxalate is transformed to cerium oxide in two steps. The first step involves the endothermic removal of 10 mol of water, with a
calculated activation energy of 78.2 kJ/mol. The second step involves the exothermic decomposition of the anhydrous oxalate, with an activation energy
of 112.6 kJ/mol. The water content in the mixed cerium-gadolinium oxalates decreases with increasing gadolinium content, while the temperature of
exothermic decomposition of the anhydrous oxalate increases with it."
My original speculation was fueled by another source (http://onlinelibrary.wiley.com/doi/10.1002/zaac.19633240113/... ) whose abstracts states "anomalous behaviour differing from that of other
rare-earth metal oxalates". To quote in its entirety:
"The thermal decomposition of the hydrated oxalates of cerium(III) and thorium has been studied by means of a Stanton thermobalance. The thermolysis
pattern of cerium oxalate shows an anomalous behaviour differing from that of other rare-earth metal oxalates. Constant weight levels corresponding to
lower hydrates have not been obtained. Several points of resemblances are noticed in the decomposition patterns of cerium oxalate and thorium oxalate.
Formation of thorium oxalate monohydrate at 140° C, as reported, has not been observed during the present studies."
Again another study, which also avoids giving results in its abstract, may confirm or deny the above in the absence of oxygen. See "THE DECOMPOSITION
OF CEROUS OXALATE IN A REDUCING OR INERT ATMOSPHERE AND A NEW PROPERTY OF THE HIGHER OXIDES OF CERIUM" at http://pubs.acs.org/doi/abs/10.1021/ja02253a005
Now that I surveyed the for profit literature, I probably still have to speculate.
[Edited on 15-11-2012 by AJKOER]
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elementcollector1
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What is the temperature at which these oxalates typically decompose? IIRC, it was aroudn 1000 C, or calcining temperatures, but I could never find an
actual reference for the temperature of decomposition.
The reason I'm asking is that my ceria is reddish brown, which I know is nowhere near pure. So, I want to use the oxalate route (or other routes, if
applicable) to remove the iron and get the pure, white ceria. And then turn *that* to cerium fluoride. Or chloride, considering chloride would be so
much easier to get.
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AJKOER
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Quote: Originally posted by elementcollector1  | What is the temperature at which these oxalates typically decompose? IIRC, it was aroudn 1000 C, or calcining temperatures, but I could never find an
actual reference for the temperature of decomposition.
The reason I'm asking is that my ceria is reddish brown, which I know is nowhere near pure. So, I want to use the oxalate route (or other routes, if
applicable) to remove the iron and get the pure, white ceria. And then turn *that* to cerium fluoride. Or chloride, considering chloride would be so
much easier to get. |
My Youtube video reference on Iron oxalate shows no extra-ordinary heating involved although it did take 10 minutes for completion.
Please note, many Oxalates on decomposition seem to be a path to nano-sized metal or metal oxides.
On the separation out of the Iron, it is my understanding that an excess of H2C2O4 will produce a soluble Fe complex (this may or may not be true for
the CeC2O4 as well).
[Edited on 15-11-2012 by AJKOER]
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elementcollector1
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| Quote: Originally posted by AJKOER |
My Youtube video reference on Iron oxalate shows no extra-ordinary heating involved although it did take 10 minutes for completion.
Please note, many Oxalates on decomposition seem to be a path to nano-sized metal or metal oxides.
On the separation out of the Iron, it is my understanding that an excess of H2C2O4 will produce a soluble Fe complex (this may or may not be true for
the CeC2O4 as well).
[Edited on 15-11-2012 by AJKOER] |
Yes, and I did use that separation method for my original isolation of neodymium oxalate from a magnet (suggested by blogfast25). Rare earth oxalates
do not form such a complex.
Nano-sized oxide means faster reactions, right? 
So, what temperature did the iron oxalate require? I'm guessing somewhere like 200-400 C?
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blogfast25
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Forget getting REs from RE oxalates: you get the oxides (even in inert atmosphere). This is generally true except in a few rare cases where the oxide
itself is unstable, for instance silver. RE oxides, by contrast, have very high Heats of Formation, so much so that reduction with Al doesn't even
work.
EC1: altough I have no figures, a lot of experience with similar calculations tells me that CeF3 + Ca is to be preferred over Ce2O3 + Ca. But neither
are simple to do in home lab setting...
[Edited on 15-11-2012 by blogfast25]
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blogfast25
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Just a few calcs on ceric oxide (CeO2). Because it contains more oxygen it’s more likely that alumina or magnesiothermic reduction
is more exothermic, compared to the cerous oxide (Ce2O3) reaction.
Acc. Wolfram Alpha, the Standard Heat of Formation (SHoF) for CeO2 is -1089 kJ/mol. NIST Webbook gives SHoFs for Al2O3 and CaO of respectively –
1676 kJ/mol and – 635 kJ/mol.
Using those data, for CeO2 + 4/3 Al === > Ce + 2/3 Al2O3 we get a Standard Enthalpy of Reaction of – 28 kJ/mol (negative but negligibly small!)
and for CeO2 + 2 Ca === > Ce + 2 CaO, we get – 181 kJ/mol. That’s not very much either and almost certainly not enough to heat the reaction
product mix to the MP of calcia (a whopping 2572 C!), needed to obtain slag-metal separation.
As shown somewhere in the neodymium thread, NdF3 + 2/3 Ca appears to be strongly exothermic and enough so that separation between the metal and the
CaF2 slag would be possible. Presumably CeF3 + 2/3 Ca should work too…
[Edited on 15-11-2012 by blogfast25]
[Edited on 15-11-2012 by blogfast25]
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AJKOER
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Apparently the decomposition of Ferrous oxalate per my recent research search is apparently best described as 'controversial'. As such Woelen opinion
is also correct as per the extract below from the attached pdf suggests varying results depending on the type of the equipment arrangement (especially
to exposure to decomposition gases) , temperature and atmosphere (air, vacuum, Hydrogen, Nitrogen or Argon). To quote from "Thermal behaviour of
iron(II) oxalate dihydrate in the atmosphere of its conversion gases" by Martin Hermanek, Radek Zboril, Miroslav Mashlan, Libor Machala and Oldrich
Schneeweiss:
"The conversion process in an inert atmosphere (N2, Ar) or
vacuum proceeds in two steps too but, unlike in oxidative
conditions, the individual steps are much better separated and
the composition of the reaction products seems to be more
controversial. Most authors identify FeO as the primary conversion
phase that subsequently decomposes to Fe3O4 and a-Fe
due to its instability below 570 C.17,24,25 The uncertainties
however concern the following conversion steps and the composition
of the final transformation products. Thus, a-Fe2O3
along with FeO and Fe3O4 were identified in the XRD pattern
of a sample isothermally treated in dry nitrogen at 440 C.26
Primary creation of FeO in dry nitrogen was also observed by
Rane et al.,18 who suggest the subsequent reaction of FeO with
FeC2O4 yielding Fe3O4 as the final decomposition product. In
the case of a nitrogen atmosphere containing water vapour,
Fe3O4 as a single phase was detected,16 similarly as in argon
atmosphere.26,27 Despite the same reaction atmosphere (Ar), a
completely different phase composition of the samples including
Fe4C, Fe3O4 and Fe is reported,28 evidently due to the
participation of the conversion gas in the reaction system.
Thermal decomposition of FeC2O4.2H2O carried out in a
reducing atmosphere of H2 proceeds again in two steps taking
place at quite separate temperatures, attributed to primary
dehydration followed by reductive decomposition. Among the
solid reaction products, FeO, Fe3O4, a-Fe and Fe3C were
detected depending on the temperature conditions. As expected,
FeO as the primary product consequently decomposed into
Fe3O4 and a-Fe.27,29 In the next reaction steps magnetite and
iron carbide were transformed to a-Fe, although the intermediate
formation of Fe3C is marginally discussed. In addition
to pure hydrogen, its mixture with NH3 in a ratio of 1 : 1 was
used as a reduction atmosphere to monitor the decomposition
mechanism of FeC2O4.2H2O. Mossbauer spectra revealed a
mixture of FeO with superparamagnetic Fe3O4 at the very
beginning of the conversion. At higher temperatures reduction
of Fe3O4 to a-Fe and its nitridation towards various iron
nitrides took place simultaneously.30"
Also:
"Being aware of the literature discrepancies concerning the
mechanism of the decomposition process, especially in inert
and reducing atmospheres, it seems to be highly probable that
gaseous conversion products more or less participated in the
formation of the solid phases through solid–gas reactions. This
can be related to the experimental arrangement, especially to
insufficient draining of the decomposition gases by passing
through the used atmosphere. Surprisingly, there are no
literature data on the thermal decomposition of ferrous
oxalate dihydrate in the atmosphere of the conversion gases."
And as to the final results:
"Following the experimental results presented in this study, we
suggest a unique five-step decomposition mechanism of
FeC2O4.2H2O in the atmosphere of its conversion gases:
I. Release of crystal water proceeding within the temperature
range of 170–230 C:
FeC2O4.2H2O --> FeC2O4 + 2H2O
II. Thermally induced conversion of FeC2O4 to Fe3O4 and
carbon oxides at temperatures above 230 C:
3FeC2O4 --> Fe3O4 + 4CO + 2CO2
III. Reduction of ferrous oxalate by carbon monoxide to
Fe3C (above 360 C):
3FeC2O4 + 2CO --> Fe3C + 7CO2
IV. Thermal conversion of Fe3C (415–535 C):
Fe3C --> 3Fe + C
V. Thermally induced reduction of magnetite to FeO by
carbon monoxide (above 535 C):
Fe3O4 + CO --> 3FeO + CO2 "
------------------------l--------------------
With respect to Ce2(C2O4)3, I would speculate that its thermal decomposition in a stream of dry NH3 might be interesting.
Attachment: Thermal_behaviour_of_iron__II__oxalate (1).pdf (347kB)
This file has been downloaded 690 times
[Edited on 16-11-2012 by AJKOER]
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elementcollector1
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Topic bump!
I plan to find a source of fluoride so I can make the CeF3 to find out if this works - but where do I find a good source of sodium fluoride? I don't
want to work with HF, can't use my CaF2 without running into the aforementioned risk (unless I mix with sodium metal, and get calcium or something?
Not likely.), and can't exactly extract it from toothpaste without scaling up into the kilogram quantity to get gram amounts of products. No natural
sources, either.
I read on another thread that some insecticides may contain high amounts of soluble fluoride, but the link that listed brands was shut down.
Toothpaste is a no-go as well, because it contains so little fluoride.
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franklyn
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Pay attention to what woelen said. Cerium metal is more than just reactive ,
it is pryrophoric. The reason it's an ingredient of lighter flints. Reduction at high
temperature will require an inert atmosphere. http://periodic.lanl.gov/58.shtml
The Preparation & Properties of Metallic Cerium ( a thesis from 1911 )
http://books.google.com/books/download/The_preparation_and_p...
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elementcollector1
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Ah. I forgot to mention that I would be conducting this under argon, pumped through from a tank. My bad.
The book was excellent - exactly the stuff I needed to know. I'll be switching back to chlorides, then - I didn't really want to deal with fluorides
anyway (although it seems like I change my mind every hour these days...), back to chlorides! Only this time, I think electrolysis of a molten
eutectic may better suit my needs.
Using this source (http://www.crct.polymtl.ca/fact/documentation/FTsalt/FTsalt_...), I found that a 0.7 mole(of NaCl) eutectic of CeCl3/NaCl melts at 500 C, a 0.625
mole-KCl eutectic melts at ~540 C, and a 0.75 mole-LiCl eutectic melts at ~500 C.
I assume these refer to the anhydrates, so dehydrating the chloride without decomposing to the oxide and oxychloride is going to be a challenge (I
could use NH4Cl, but would have to work out a setup for that in my head).
NaCl is the obvious choice, but at 0.7 moles NaCl (and presumably 0.3 moles CeCl3), that's a lot of molten salt!
As for heating, I think I will go with an air-assisted coal furnace (no idea what they're actually called). I think plante199 posted pictures of one
in his 'calcium oxide from seashells' Prepublication. As I remember, it consisted of a brick 'tower' filled with coal, with a hole near the bottom. A
steel tube would lead out of that hole, connected to a hairdryer to blow air through and increase the temperature. Mine will be much smaller - perhaps
a large soup can full of coal, with a hole drilled or punched at the bottom - I have used one of these without the hairdryer addition to melt aluminum
cans (MP of 660 C) with hardly a problem.
[Edited on 5-6-2013 by elementcollector1]
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blogfast25
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EC1:
Thanks for posting that eutectic database, very useful!
Any idea how you're going to obtain the CeCl3 hydrate (never mind the anhydrous form for a minute)? I'm guessing you'll be starting from CeO2?
If you change your mind back to CeF3, NH4HF2 (ammonium bifluoride) is fairly easy to get. NaF isn't very soluble (about 1 M/L)
[Edited on 5-6-2013 by blogfast25]
[Edited on 5-6-2013 by blogfast25]
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phlogiston
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elementcollector, thanks for the link. Very interesting and useful. The links to phase diagrams of mixtures containing Th, U or Pu don't work (but
none of us will every use those I guess).
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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elementcollector1
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Well, I can imagine a few uses...
Probably CeO2 boiled and stirred in the azeotropic (purified by distillation) of HCl at around 20%?
I'm wondering if I should keep it stoichiometric or just use an excess, because I feel some HCl is going to be given off by heating. Plus, the oxide
will take forever to dissolve (being most likely calcined).
Interestingly, the oxide (which is a light brown in color) is readily attacked by dilute H2SO4, with no visible reaction but for a change in color
from the usual pale brown/yellow to a much stronger golden-yellow (about the color of cerium sulfate). Speaking of cerium sulfate, it is listed as
soluble in dilute acids, so this checks out.
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blogfast25
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No, even 36 % HCl is unlikely to put a small dent in CeO2, even with prolonged boiling. Calcinated oxides are quite resistant to even conc. hot HCl.
You will need 95 % (or better) H2SO4. Alternatively it could be worth trying to fuse the ceria with anhydrous NaHSO4, then dilute with 20 % H2SO4,
after digestion. When NaHSO4 melts you get a much higher temperature than with boiling H2SO4, which will promote digestion of the oxide.
Then you'd need to chill to promote solubility of the Ce(SO4)2, so it can be reduced to Ce(III) (H2O2 or methanol/ethanol) and precipitated as
Ce(OH)3.nH2O with ammonia. Ce2(SO4)3 is poorly soluble in water too but using strong ammonia that should be convertible to Ce(OH)3.nH2O.
Another possibility is to try and dissolve the ceria in a mixture of conc. H2SO4 and (NH4)2SO4, that should yield ceric ammonium sulphate which is
reported as 'moderately soluble'. Then reduce and precipitate as cerous hydroxide.
The freshly precipitated cerous hydroxide hydrate should be soluble in conc. HCl, yielding CeCl3.
[Edited on 5-6-2013 by blogfast25]
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elementcollector1
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Well, I'll be sure to do that when I get home. Ce(OH)3 is bright yellow, correct?
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blogfast25
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I'm not sure. In my opinion it's only the Ce (IV) hydrated ion that is orange/yellow. Ce(III) halides are all white, except CeI3.
There's a picture of Ce2(SO4)3 [if I can believe the author] on one of the two threads and that product is white.
What grade of CeO2 are you planning to use?
[Edited on 6-6-2013 by blogfast25]
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MrHomeScientist
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Interesting stuff. What is your source of cerium?
If you do want to go the fluoride route, I have quite a bit of ammonium bifluoride. I can give you a bit, as long as you promise to be OBSCENELY
CAREFUL with it! (and release me from all liability  )
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