Sciencemadness Discussion Board

CeCl3 from CeO2

elementcollector1 - 24-9-2012 at 17:35

How would I go about preparing this compound?
I understand it involves HCl to some extent, but my CeO2 appears to contain iron (very reddish-brown color) and my HCl contains the same. Also, I have heard that CeO2 reacts slowly, if at all, with HCl to produce CeCl3 and H2O.
As for why I need this rare-earth compound, it's for synthesis of cerium metal by a reduction with calcium metal and iodine (for additional heat added to the reaction).

madcedar - 24-9-2012 at 18:29

These two links will help you greatly :P
http://www.sciencemadness.org/talk/viewthread.php?tid=9505
http://www.sciencemadness.org/talk/viewthread.php?tid=13856

They will help you with ceriium III and IV too.

elementcollector1 - 24-9-2012 at 19:00

Many thanks! I thought H2O2 might help.
Evidently, iron can be gotten rid of by adding sodium sulfate; but how would sulfuric acid do? I'd imagine it would have the same effect, being technically the "hydrogen salt" of the sulfate ion and all. Alternatively, I could mix precious sodium hydroxide and sulfuric acid to form sodium sulfate and water.
So, then I would isolate Ce(OH)4 by adding the ceric sulfate (in solution) to a hot solution of NaOH and add NaOCl. This creates an emulsion of yellow Ce(OH)4, which can then be redissolved in pure HCl to yield the desired salt.
Which can then be reacted in a 200:103:56 ratio of CeCl3:I2:Ca under anhydrous conditions at 400 C (under argon?) to isolate cerium yield of 93%.
(But seriously, shouldn't that last part be under argon? It mentions using a crucible of sintered CaO or dolomite (beyond my capabilities) inside an iron tube with a screw-on lid and welded bottom (pipe bomb! Yay!) You'd think this could be somewhat explosive, but maybe the iron would hold out.)

blogfast25 - 25-9-2012 at 04:00

EC1:

What makes you think sulphates will get rid of iron?? Both ferrous and ferric sulphate are quite water soluble, although the ferric form does have strong tendency to hydrolyse. Separate the iron and cerium using the potassium cerium (III) double sulphate in acid solution (to prevent hydrolysis of the iron sulphate)

Re. CeCl3/Ca/I2: iodine is indeed used as a heat booster in some calciothermic reactions (Ca + I2 == > CaI2 + heat) but this can only be carried out in closed reactor, preferably under vacuum or low pressure argon. And your CeCl3 needs to be anhydrous.

What’s your reference for this method?


[Edited on 25-9-2012 by blogfast25]

elementcollector1 - 25-9-2012 at 05:21

"Preparation of Metallic Lanthanides".

Upon reading the other threads, it seems that cerium sulfate was precipitated out by adding concentrated sodium sulfate to the chloride solution, leaving ferrous sulfate behind in solution.

The literature said "an iron pipe with a welded-shut bottom, and a screw-on lid". This is probably as close as one can get to a closed reactor. As for flooding with argon, this could be difficult. Maybe drill two holes in the reactor, place steel fittings on here (weld them?) and cycle argon carefully through these? This could be a problem once the actual reaction starts, as something might reach the tubes and block them.

For anhydrous CeCl3, then I need to get some thionyl chloride or ammonium chloride.
For thionyl chloride, I suppose I could prepare it through the chlorination of elemental S to make SCl2, and then react this with sulfur dioxide and additional Cl2. This would be complicated, though, and I would have no good idea of how to make sure the thionyl produced is anhydrous itself.
For ammonium chloride, I would need dry HCl and dry ammonium gas to produce an anhydrous result.

Attachment: Metallic lanthanides.doc (467kB)
This file has been downloaded 474 times

blogfast25 - 25-9-2012 at 05:35

Quote: Originally posted by elementcollector1  
"Preparation of Metallic Lanthanides".

Upon reading the other threads, it seems that cerium sulfate was precipitated out by adding concentrated sodium sulfate to the chloride solution, leaving ferrous sulfate behind in solution.

The literature said "an iron pipe with a welded-shut bottom, and a screw-on lid". This is probably as close as one can get to a closed reactor. As for flooding with argon, this could be difficult. Maybe drill two holes in the reactor, place steel fittings on here (weld them?) and cycle argon carefully through these? This could be a problem once the actual reaction starts, as something might reach the tubes and block them.

For anhydrous CeCl3, then I need to get some thionyl chloride or ammonium chloride.
For thionyl chloride, I suppose I could prepare it through the chlorination of elemental S to make SCl2, and then react this with sulfur dioxide and additional Cl2. This would be complicated, though, and I would have no good idea of how to make sure the thionyl produced is anhydrous itself.
For ammonium chloride, I would need dry HCl and dry ammonium gas to produce an anhydrous result.


Fluxing with argon would drive off the iodine, once the reaction gets started, trust me: been there, done that. With I2 you really need a closed system, pressure resistant.

Your reference also mention KClO3 as a heat booster but I'm not keen on mixing chlorates with a chloride based reduction reaction. Iodine really is probably best here...

Heating the CeCl3 with an excess of NH4Cl should give dry CeCl3 but my first attempt with NdCl3 hydrate did not go very well.

[Edited on 25-9-2012 by blogfast25]

elementcollector1 - 25-9-2012 at 08:42

How did it go?
(And furthermore, how does this dehydration work? Mixing the powders would make for difficult separation, unless the NH4Cl is heated to the point of decomposition...)

So, no argon, just a closed container. This actually makes my job easier! :D

Now, because I don't have a sintered CaO crucible, what are adequate replacements? The casserole crucible I ordered is far too big to fit in an iron pipe of ordinary diameter, but I have a stainless steel cup that *might* stand a chance. They specifically say in the lit. reference not to let the reactants touch the iron, but I don't know how stainless would hold up. Probably not too well. In that case, I'll visit the pottery shop; they're bound to have something.

How well would a method like this work for other elements, such as neodymium, lanthanum, etc.?

blogfast25 - 26-9-2012 at 04:00

For it to work, I think your hydrate should be fairly dry (my NdCl3 wasn’t because it’s hygroscopic) and fine and then mixed VERY INTIMATELY with dry, pure NH4Cl. For MnCl2 I used 2 mol salmiac per mol of MnCl2 hydrate. I mixed them by grinding them into each other in a granite mortar and pestle. That dehydration worked really well.

On heating to well above the sublimation point of NH4Cl it sublimes and temporarily decomposes to NH3(g) and HCl(g). It’s the dry HCl gas that protects the hydrate from decomposing (while it dehydrates), or so the theory goes.

Another way is to dehydrate in a stream of dry HCl(g).

Re. other REs, each case is somewhat different (see also you r own reference). IIRW I calculated the reduction of NdF3 by Ca to be possible (see ‘the trouble with neodymium’, near end of thread). Dunno about the others.
http://www.sciencemadness.org/talk/viewthread.php?tid=14145&...

Crucibles? Steel isn’t recommended. Lab ceramic might just do the job. Ordinary ceramic has a tendency to crack due to thermal shock, unless you treat it gently (heat up and cool down fairly slowly) A cracked crucible would be a small price to pay for decent metal though…


[Edited on 26-9-2012 by blogfast25]

jam_kemist - 26-10-2024 at 12:10

About the reaction of CeO2 with 23%HCl,
I've tryed it myself twice by adding small amounts of 30% H2O2 as well, all under reflux for multiple hours.
But I never managed to dissolve all the CeO2 I used, only a partial amount.

I suspect the reaction is extremely slow, today I refluxed it for ~3h and got no CeCl3 at all (after filtration of unreacted CeO2 and boilling the filtrate down)
The time it worked I think I refluxed it for 30-40 hours

How much of a difference would 36% HCl make ? It's hard for me to get so I usually try to avoid it but if it speeds up the dissolution significantly... I'll have to consider again.

j_sum1 - 26-10-2024 at 15:57

I have had the same experience -- leftover CeO2. However, I have simply added it to my next batch. It looks really white and clean with no apparent iron contamination. In none of my batches have I dissolved all the CeO2. I am not sure why this is.