Picric-A
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Cerium (III) Oxide
I am looking to aquire some Cerium (III) oxide.
Ebay sells 'Cerium Oxide' for use in rock tumbling however i am not sure which oxide this is; IV or III.
Judging by the pictures this looks like the III oxide as wiki describes it as 'gold-yellow in color' whilst the IV is a 'pale yellow-white powder' and
the ebay pictures certainly make it look orange.
Now i know looks can be decieving and this is backed up by the fact cerium III is less stable than cerium IV oxide and so it is doubtful this is
actually the III oxide.
Ok, so lets say for the sake of argument this IS ceirum III oxide, it is likely this stuff will be very unreactive, like calcined oxides normally are
for pottery ect... Is there any way i can make it more reactive, ie not inert?
For the other argument, if this stuff on ebay turns out to be the IV oxide, is there any way of reducing it on lab scale which doesnt involve reducing
it at '1400degC(sic) in hydrogen' as this is WAY out of my league...
I know if temps below this are used the resultant oxide will be pyrophoric but it doesnt matter as i will handle it in a glove box and will use it
straight away.
If some of my questions can be answered it will be a great help.
Particually Woeelen who seems to have a lot of experience in this area!
A few refs;
http://en.wikipedia.org/wiki/Cerium(III)_oxide
http://en.wikipedia.org/wiki/Cerium(IV)_oxide
http://cgi.ebay.co.uk/100g-Cerium-Oxide-Regipol-FREE-POSTAGE...
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DJF90
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Ceria, CeO2, is used for polishing, particularly in the glass industry. The orangy-brown colour is likely due to Iron contamination and also any other
additives they put in to make it better for its intended application.
The M2O3 oxides of the Lanthanide series are best produced by the thermal decomposition of the trivalent nitrates:
4Ln(NO3)3 => 2Ln2O3 + 12NO2 + 3O2
However, the metals with +2 and +4 oxidation states ( that includes cerium) can afford other stoichiometries. The trivalent oxides can be synthesised
from such oxides under a reducing atmosphere, e.g. H2, CO.
Whats the interest in Cerium III oxide anyway?
[Edited on 16-5-2010 by DJF90]
[Edited on 16-5-2010 by DJF90]
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JohnWW
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Transitions of that unpaired 4f electron in Ce2O3 would make it fairly strongly absorb UV light, and the absorption band is likely to stretch into the
visible.
However, why use Ce2O3 or CeO2 (which have better uses) as an abrasive, if Al2O3 (microcrystalline corundum), which can polish siliceous minerals,
will do the job at least as well?
[Edited on 17-5-10 by JohnWW]
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not_important
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Because alumina doesn't do the job as well, it is too hard for softer materials. Cerium oxide is commonly used on quartz, opal, and softer stones, and
for polishing glass optics. Alumina and SiC may be used to cut such materials, but cerium and tin oxides are used to polish them.
There are several grades of cerium oxide used for polishing, and increasingly there are cerium oxide based formulations. We chemist types don't always
want the highest grade stuff, for reasons explained later.
The use of CeO2 for polishing goes back to the 1930s or perhaps a few years earlier, but it really took off after WW-II. Back then you commonly found
polishing grades that were only about half CeO2, with a lot of other rare earth oxides making up the rest. Those oxides contributed little to the
polishing action, but the low grade was much cheaper to produce. Colours of the low grade stuff ranged from mocha to chocolate, as you went to purer
CeO2 the colour lightened and became orange-tinted, then tan, beige, buff, cream, and ivory as the purity increased.
The highest concentration contaminates were/are lanthanum, which doesn't really affect the colour, praseodymium, whose mixed valence oxides are
black-brown, and neodymium, contributing red-violet. Neodynmium is more common than praseodymium, so Nd is a significant contaminate although it is
not immediately adjacent to cerium.
Now I've a pamphlet on the use of Ce(IV) salts in redox titrations, from back in the 1940s or so, when you just aabout had to make your on cerium
reagents. And it said that the not-very-pure stuff was the best starting because it was easier to get it to dissolve. In a fashion a bit similar to
CrCl3, you need to get some Ce(III) or other reducing agent in solution to reduce the fairly inert CeO2 so acids will attack it. And the impure
storngly coloured cerium oxide had both some Ce2O3 and those mixed III or IV praseodymium oxides that greatly facilitate getting the bulk of the CeO2
into solution.
I've purchased polishing grade CeO2 a number of times over the years, getting differing grades. The 80% stuff was always easy to into solution, the
90% could be difficult, and the 97% on up grades generally were a struggle. I had one small batch of supposedly 95% CeO2 that when I started tapping
some into hot HCl, foamed away bubbling off waves of Cl2. I've had off-white high purity batches that just would not dissolve.
HCl is a good starting point, as it can function both as acid and reducing agent. Sometimes I had to drip ethanol or isopropanol into the acid-CeO2
mix; H2O2 has worked on several occasions. You end up with a solution of the REE as their trichlorides. From that you can oxidise the Ce back to the
+4 state and drop it out of solution as the hydroxide/hydrated oxide which is easily soluble in acids.
The pamphlet was mainly about redox stuff, ceric ammonium nitrate was the main target production. So it focused on that route - dissolve the crude
cerium oxide in HNO3, do a little magic, add NH4NO3 and then concentrated HNO3 to crystalise out CANwhile leaving the other REE and whatever in
solution as their +3 nitrates.
Now as currently there is a trend to add suspending agents to cerium oxide to produce formulations for the big users and the polishing machines they
use, as those pump suspensions of the polishing agent over the working surfaces. These suspending agents may contribute the the shade of the oxide,
and possibly cause other problems.
So my advice is to run a few tests on a small test tube scale. Heat a bit of the oxide dry to it's quite hot, just to see if it shows signs of
containing organics. A a bit of the oxide to a few ml of concentrated hydrochloric acid and heat it to see if you get a reaction. Repeat, but when
hot add a few drops of EtOH. Try H2O2 with acid, experiment to see what does bring about reaction.
U2U me in a couple of days if I haven't posted more detail from that old pamphlet, I'm going to try to find it.
[Edited on 17-5-2010 by not_important]
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woelen
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I have some experience with CeO2 and these are not good. I have very pure CeO2 (which is almost white, just a very pale color) and this stuff is
remarkably inert, just like many other calcined oxides.
I expect Ce2O3 to be much less inert than CeO2. I have several other La2O3 oxides (e.g. see my experiments with Nd2O3) of very high purity and all of
these seem to be soluble fairly easily, just by adding them to a hot solution of dilute H2SO4. The oxides I have, however, all are white, except the
Nd2O3 (grey with a bluish hue) and the Pr2O3 (light green). I also expect Ce2O3 to be purely white.
This 'regipol' stuff does not look to me like a good source of cerium. A simple search on google reveals that it is cerium oxide (probably CeO2) with
other stuff mixed in to make its polishing properties better. There also are different grades of 'regipol', such as 'regipol 80', 'regipol 95' and
other numbers I have seen as well. I think that these are percentages CeO2 and the remaining part being some proprietary mix of compounds and the
different grades certainly will have different polishing properties and the customer can choose from these different grades. I have no experience
though with dissolving 'regipol'. I myself certainly would not buy such a compound, I am afraid that it will be a lot of a hassle to get the cerium
free from all the other crap in that material.
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not_important
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| Quote: | | In your letter, you indicate that "Regipol" is a rare earth oxide based glass polishing powder that is produced from a cerium rich bastnasite
concentrate. You also indicate that grades 800 CA and 790 CA typically contain approximately 69% and 58% rare earth oxides respectively, with the
remainder made up of clays and other minerals. |
So the question becomes, is it just bastnasite concentrate, a concentrate that has been roasted and acid leeched, or one of those plus additional
additives. A simple bastnasite concentrate would be OK, cerium is pretty easily purified from such material by changing is oxidation state. A
roasted material will be less soluble and thus more trouble, the bother vs the price may be the deciding factor.
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Picric-A
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Quote: Originally posted by DJF90  | Ceria, CeO2, is used for polishing, particularly in the glass industry. The orangy-brown colour is likely due to Iron contamination and also any other
additives they put in to make it better for its intended application.
The M2O3 oxides of the Lanthanide series are best produced by the thermal decomposition of the trivalent nitrates:
4Ln(NO3)3 => 2Ln2O3 + 12NO2 + 3O2
However, the metals with +2 and +4 oxidation states ( that includes cerium) can afford other stoichiometries. The trivalent oxides can be synthesised
from such oxides under a reducing atmosphere, e.g. H2, CO.
Whats the interest in Cerium III oxide anyway?
[Edited on 16-5-2010 by DJF90]
[Edited on 16-5-2010 by DJF90] |
My interest in Cerium III oxide is in the making of Ceric ammonium nitrate, a very usefull oxidiser in organic chemicstry and inorganic chemistry.
thanks for all the great replies, this stuff will take me a while to read so keep it comming!
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Magpie
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I presume you have seen this:
http://www.sciencemadness.org/talk/viewthread.php?tid=9505#p...
The single most important condition for a successful synthesis is good mixing - Nicodem
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Picric-A
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Magpie- Yep i found that before i posted dont worry however it did not contain as much information based soley on the production of the oxide which i
am looking for, ie conversion of IV to III and purification ect..
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Butterflywings
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How can you produce cerium (III) oxide from a solution of concentrated cerium (III) sulfate?
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woelen
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Precipitate the cerium as the hydroxide, rinse very well and then heat the precipitate. I'm not sure how pure it will be then. Maybe better purity is
achieved with ammonia instead of a soluble hydroxide, because then you won't have other metal ions from the hydroxide in the precipitate. Still,
sulfate may be a contaminant as well.
The heating also must be done in an atmosphere, free of oxygen. I'm quite sure that heating a cerium(III) compound in air will lead at least partially
to formation of cerium(IV).
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Butterflywings
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| Quote: Originally posted by woelen | Precipitate the cerium as the hydroxide, rinse very well and then heat the precipitate. I'm not sure how pure it will be then. Maybe better purity is
achieved with ammonia instead of a soluble hydroxide, because then you won't have other metal ions from the hydroxide in the precipitate. Still,
sulfate may be a contaminant as well.
The heating also must be done in an atmosphere, free of oxygen. I'm quite sure that heating a cerium(III) compound in air will lead at least partially
to formation of cerium(IV). |
But what I'm trying to achieve is cerium (III) instead. Is there a way?
[Edited on 5-12-2012 by Butterflywings]
[Edited on 5-12-2012 by Butterflywings]
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woelen
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What I wrote is about making cerium(III). The precipitated hydroxide will be the species in oxidation state +3. You only have to take certain
precautions to assure that no cerium(IV) is made.
It is not without reason that cerium(III) oxide is not a common article of commerce, while cerium(IV) oxide is. I think that making pure cerium(III)
oxide is very hard.
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IrC
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Not the expert chemist but from the wiki it mentions carbon monoxide as a route: "When there is a shortage of oxygen, cerium(IV) oxide is reduced by
carbon monoxide to cerium(III) oxide:
2 CeO2 + CO → Ce2O3 + CO2 ".
If the oxygen present is in too high a percentage then reversal occurs.
"When there is an oxygen surplus, the process is reversed and cerium(III) oxide is oxidized to cerium(IV) oxide:
2 Ce2O3 + O2 → 4 CeO2 ".
Is this not a possible route: 2 CeO2 + CO → Ce2O3 + CO2?
"Science is the belief in the ignorance of the experts" Richard Feynman
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12AX7
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If that's the case, then it should be easy to pack charcoal and CeO2 in a steel can and toss it on the fire.
Doesn't say what temperature it's done at, but red heat or above would be a good bet.
You could also lead a thin steel tube into the powdered mass, so that, once it's heated up some, a little air blown in will heat up the rest of the
mass rapidly (allowing much higher temperatures than passive heating).
The powder on top will be CeO2, while the stuff deep inside will be a mixture of Ce2O3, a little CeO2 (how much depends on the reaction rate,
temperature, time, actual diffusion rate of the stuff, etc.), and leftover charcoal and ashes.
Most likely, the whole mass will dissolve in acid much easier after treatment than the relatively inert oxide was before. Filter out charcoal and
residues, and precipitate the hydroxide to remove alkali and calcium (but not magnesium). The precipitate could be washed with excess alkali to
remove aluminum and other amphoterics.
Tim
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WGTR
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Cerium salt fluorescence under UV light
I purchased some 99.99% (!) purity CeO2 from Paul Brown Studios on eBay a year or two back. It's a sort of off-white color...not quite white, but not yellow by any means.
This stuff is practically inert. I tried dissolving a bit of it in molten sodium bisulfate, and only got a bit of it to dissolve after several
minutes. Most of it sat in the bottom of the test tube, unaffected, even with me blasting the tube with a propane torch. I tried concentrated acids,
HCl, HNO3, H2SO4, HCl mixed with H2O2, and none of those really worked. I eventually tried a
1:1 mix of 30% H2O2 and 68% HNO3, and that is doing the trick. Once the powder is added it immediately turns
yellow/orange. Under heavy stirring the powder dissolves in an hour or so. As it dissolves the solution turns lighter in color; once it dissolves
completely the solution becomes colorless and clear.
The interesting thing, though, is that I also tried dissolving the powder in molten sodium sulfite, since I have about 100lbs of this sitting around
in large jars (don't ask). It's probably oxidized a bit from age, so there is likely some sulfate in there as well. I placed a few 10's of
milligrams of CeO2 in a soda-lime test tube, and then covered it with my impure sodium sulfite. Upon heating with a torch, the
CeO2 began turning black as it reduced. The sulfite began melting a bit at the point that the test tube began softening, and some of the
reduced powder started dissolving into this molten salt.
The melt wasn't working too well, so I decided to quit on that attempt. When I went to wash out the test tube, however, I smelled a faint
H2S odor. Out of curiosity I looked up cerium sulfide, and noticed some references related to its fluorescent abilities.
Since I have a UV mercury vapor lamp, I decided to fire it up and see if there was anything interesting left in the test tube. I scraped out a few
white, undissolved, bits of the sulfite salt and the residual cerium oxide. The cerium oxide did not appear to be fluorescent. The sulfite salt,
however, with bits of reduced cerium oxide dissolved in it:
..appears to glow red/orange under UV light.
Once I placed this material in a small, sealed, vial, I illuminated it again:
..much brighter this time. Since the material was damp, I used dry nitrogen to dry it, and then illuminated the sample again. It still fluoresced,
but the brightness was comparable to the first picture instead of the second. Perhaps the H2S environment in the second picture made a
difference in glow intensity.
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WGTR
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Also, if one wants to reduce CeO2, this is not too difficult. I used a small borosilicate glass tube. This was heated with a propane
torch towards one end, so that I could bend an "L"-shape in the tube. About 1mL of alcohol was added into the tube, and the tube was held in a stand
at a slight angle, so that the alcohol would stay in the bottom of the "L".
Dry nitrogen was bubbled slowly through the alcohol, and a sample of CeO2 was placed into the other end of the tube. With the nitrogen
flowing, heat was applied to the powder sample with a propane torch, and a color change from off-white to black was quickly noted.
Heat was then removed, and the sample allowed to cool under flowing nitrogen. It seems to be a bit stable if it's being moved rapidly from the tube
to an oxygen-free vial, but will soon re-oxidize in air to a light gray color.
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elementcollector1
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I may have managed reduction of CeO2 to Ce2O3 a few years ago in high school. Reagents used were pure, white CeO2 and granular, 99.99% Mg, and the
stoichiometry was set up so that (in theory) the Mg would reduce the oxide to elemental Ce. However, I was met with a strange green-black substance
instead, and was unable to quite explain what this was. Given the purity of the reagents involved, I would suspect a mixture of MgO and Ce2O3, as
there was little in the setup that could have contaminated the mix with much of anything else.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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WGTR
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I should clarify that I don't claim that I reduced it all the way to metallic cerium, but rather just a lower oxide. I would be surprised if I had
zero valent cerium from what I did. The black color is interesting, though. I thought Ce2O3 was more of an orange color, from
the limited information I've seen online.
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AJKOER
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| Quote: Originally posted by IrC | Not the expert chemist but from the wiki it mentions carbon monoxide as a route: "When there is a shortage of oxygen, cerium(IV) oxide is reduced by
carbon monoxide to cerium(III) oxide:
2 CeO2 + CO → Ce2O3 + CO2 ".
If the oxygen present is in too high a percentage then reversal occurs.
"When there is an oxygen surplus, the process is reversed and cerium(III) oxide is oxidized to cerium(IV) oxide:
2 Ce2O3 + O2 → 4 CeO2 ".
Is this not a possible route: 2 CeO2 + CO → Ce2O3 + CO2?
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For those working with Ce salts, you may want to look at my related comments at http://www.sciencemadness.org/talk/viewthread.php?tid=94078#... .
[Edited on 29-9-2018 by AJKOER]
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