j_sum1
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Dissolution of cerium (IV) oxide
I have just received 200g of CeO2 to play with. I thought I might try preparing some solutions of Ce salts but obviously oxides of this nature do not
play well with acids. I have come across a couple of ideas:
- Reaction with sulfuric acid at elevated temperatures (takes days with almost zero happening)
- Reaction with HCl
- Reaction with H3PO4
- Reaction with HBr (which I don't have)
- Reaction with HNO3 / H2O2
- Reaction with HCl with KI
Sources:
https://www.researchgate.net/post/How_can_I_dissolve_CeO2_ce...
https://www.jstage.jst.go.jp/article/matertrans/53/11/53_M-M... (Which I have not fully read yet)
Does anyone have any experience with playing with CeO2?
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woelen
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Dissolving this kind of inert oxides is amazingly difficult. You could try it in overheated H2SO4 (at a red heat, the acid confined in a pressure
chamber) or in molten Na2S2O7 (can be made by heating NaHSO4 to a red heat). HCl, HBr, HNO3 are useless. I only give you some chance of success with
H2SO4 and H3PO4. The other acids are too volatile.
Industrially I think that processing of such oxides (in ores) is done in the molten state.
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j_sum1
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I have some Na2S2O7. Looks like I need to build myself a furnace then.
I might run some tests in the meantime -- based on the ideas suggested. It can't hurt.
[edit]
No I don't. I have some Na2S2O8.
[Edited on 8-4-2016 by j_sum1]
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blogfast25
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My exploits dissolving reagent grade CeO2 in conc. H2SO4.
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j_sum1
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Wow! Thanks bloggers.
I did do a search but obviously not thoroughly enough.
That little experiment series is a mine of information. I should be able to use some of that.
I think I will pursue hot sulfuric acid first and target the ceric sulfate. Then go for the hydroxide and see where it takes me. It seems to me that
the hydroxide could be easily made into all sorts of salts.
The molten sodium pyorsulfate route sounds interesting too but I don't have anything with enough heat output that I can do safely in my little shed at
present. And my temperature reading ability tops out at 300°C. I will think that one through a bit more before trying.
I don't really see how it could work but it will cost me almost nothing in time and resources – I am going to try out the HCl with added KI
suggested in the link I found. If it does something I will report it.
Unfortunately it will be at least a week before I can get into shedworld to have a play.  I guess that gives me a chance to do some reading first.
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blogfast25
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There were a few more experiments that for some reason I didn't report, like the reduction of Ce4+(aq) with H2O2(aq)
to Ce(+3), in acid conitions. Handy if you just want some stable Ce(+3) salt.
Ce(OH)3 is ironically easily oxidised to CeO2 (Ce(OH)4 ?) with that same peroxide but in alkaline conditions.
And I still don't understand why Ce(+4) is coloured and Ce(+3) colourless. Basic QC suggests the other way around!
[Edited on 8-4-2016 by blogfast25]
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j_sum1
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Yeah, that one has me puzzled too. But I am far from knowledgeable on that one. Is this a phenomenon that is repeated elsewhere among the rare
earths?
I did find this beast – the hexanitratocerrate ion which is formed in solutions of cerium ammonium nitrate. If similar structures are formed with other anions it might
account for some colour. Or, more to the point, with my limited understanding I wouldn't be surprised to find that such an object absorbed visible
light.
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elementcollector1
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Strangely, my CeO2 from Elemental Scientific dissolved just fine in HCl and H2SO4. Maybe the thermal history of yours is different? I wouldn't know if
it changes between suppliers.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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Quote: Originally posted by j_sum1  |
Yeah, that one has me puzzled too. But I am far from knowledgeable on that one. Is this a phenomenon that is repeated elsewhere among the rare
earths?
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The colours of the RE(+3) compounds are due to part-filled 4f orbitals.
Ce(+3) should have a 4f1 orbital and should exhibit colour. But Ce(+4) should be 4f0, expected colourless. Ligand
Field Theory predicts the otherwise degenerate 4f energy levels are split into groups, due to the electrical field of the ligands. 4f electrons can
then toggle between the groups by VIS absorption. But here that theory doesn't work so well...
| Quote: Originally posted by elementcollector1 | | Strangely, my CeO2 from Elemental Scientific dissolved just fine in HCl and H2SO4. Maybe the thermal history of yours is different? I wouldn't know if
it changes between suppliers. |
Thermal history does play a part, that's known.
Did chlorine evolve when you dissolved the CeO2 in HCl/H2SO4? Off the top of my head, Ce(+4) oxidises
chloride to chlorine. If so, that could help explain the dissolution.
Edit:
SRPs confirm that:
Ce(+4)/Ce(+3) = + 1.44 V
2Cl-/Cl2 = - 1.36 V
(Source)
[Edited on 8-4-2016 by blogfast25]
[Edited on 8-4-2016 by blogfast25]
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