woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Very low solubility of erbium oxide in mineral acids
I just received some erbium oxide (Er2O3), which I purchased from eBay. This is a fine compact pink powder.
To my bad surprise it is very hard (nearly impossible) to dissolve this material in mineral acids. I added some of the powder to 20% H2SO4 and boiled
the solution for several minutes. Only a very small amount of the oxide dissolves and a very pale solution is obtained. I tried the same with 30% HCl
and the result is similar. I just boiled off a lot of HCl, but the amount of Er2O3, which goes in solution is small, very small.
Are there other people over here who have experience with erbium oxide? How can this oxide be dissolved, such that a decently concentrated solution of
Er(3+) ions can be obtained?
I also have Nd2O3, Pr6O11 and La2O3 and all of these dissolve easily in moderately concentrated mineral acids, especially when the solution is heated
gently. I expected the Er2O3 to behave similarly, but this appears not to be so.
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
Rare earth sulfates are less soluble in hot water. This may be one reason? Are you sure, Er2O3 is left and not Erbium sulfate? But if the other RE
dissolve well, this should not be the reason.
You should get the chloride by heating Er2O3 with NH4Cl:
http://de.wikipedia.org/wiki/Erbium%28III%29-chlorid
I also have some Er2O3 and always wanted to dissolve it in mineral acids. I will try it today with various acids and tell you about the success.
[Edited on 6-10-2012 by Pok]
|
|
|
haroldramis
Harmless
Posts: 21
Registered: 6-1-2012
Member Is Offline
Mood: No Mood
|
|
http://books.google.com/books?id=zNJDAAAAYAAJ&pg=PA140&a...
Manual of qualitative chemical analysis
By C. Remigius Fresenius
I bought Eu,Er,La,Nd,Sm,Pr and Tb oxides from an ebay seller and Erbium was the most stubborn so far. Praseodymium right flew out of the flask and
quite near murdered me with the chlorine smelling fumes that escaped.
Here is what I did: 1g Erbium oxide to a minimum of 2mL conc.HCl. Stir at a temp just below boiling and replace any HCl that evaporates off.
You will have a lovely looking pink solution within an hour or so, probably sooner.
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
Got it!
0,29 grams Er2O3 + 1ml HCl (30%) ---heating 2 minutes---> pure ErCl3 * xH2O
Erbium oxide (more pink in reality)
Cold crystallized Erbium chloride
After addition of some water. Some particles still remaining (much less than 5% of added oxide). Probably not heated long enough.
Edit: the HCl was boiled for 2 minutes, not only heated.
BTW: an interesting property of Er2O3 is that emits a greenish light, wenn heated directly and very strong with a bunsen burner flame. This has
something to do with complementary colours (oxide red, glow green). The same is visible with thulium oxide (oxide slightly greenish, glow red). Almost
impossible to get by a camera.
[Edited on 6-10-2012 by Pok]
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Pok  | | an interesting property of Er2O3 is that emits a greenish light, wenn heated directly and very strong with a bunsen burner flame. This has something
to do with complementary colours (oxide red, glow green). |
I knew erbium oxide has been used as a phosphor
for decades. What I didn't know was what was reported in this paper, an up-conversion emission in the green band. This sounds a lot like what you're
reporting.
Preparation and characterization of erbium oxalate and erbium oxide nanoparticles by microemulsion technique
Wenxiu Que, et al.
Materials Science and Engineering: C, Volume 16, Issues 1–2, 20 October 2001, Pages 51–54
http://dx.doi.org/10.1016/S0928-4931(01)00297-1
Here's the abstract: | Quote: | | This paper reports the preparation and characterization of erbium oxalate and erbium oxide nanoparticles. The erbium oxalate and erbium oxide phosphor
materials were synthesized through the microemulsion technique. Transmission electron microscopy (TEM) was used to characterize the phosphor materials
and reveal a nanocrystal structure for the studied phosphors. A strong green up-conversion emission at 543 nm
(4S3/2→4I15/2) and another weak emission at 528 nm
(2H11/2→4I15/2) have been measured for the studied phosphors upon excitation at 993 nm
(4I11/2→4I15/2). The up-conversion emission mechanism has been explained by means of an energy level
diagram. The lifetimes of these emissions have been measured. The green up-conversion emission equally strong in both the phosphors and for these
phosphors, color coordinates (X, Y) have been obtained in order to assess their color emitting efficiency. |
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Is there any reason to believe that what holds for many metal oxides, that their solubility in mineral acids depends on temperature and duration of
calcination, isn’t true for RE oxides? I don’t thinks there is.
Woelen’s Er2O3 may have been calcined harder than Pok’s. My homemade Nd2O3 was also surprisingly resistant to 36 % HCl, even though it had been
calcined ‘softly’. Only hot conc. H2SO4 converted it slowly to insoluble Nd sulphate (the ‘sandy’ version).
|
|
|
strontiumred
Harmless
Posts: 36
Registered: 6-7-2007
Location: Southern England
Member Is Offline
Mood: Delocalised
|
|
Hi woelen,
I was able to dissolve 50mg of erbium oxide in 1.5ml of 10% HCl with some heating. Maybe you are trying to form too concentrated a solution?
[Edited on 7-10-2012 by strontiumred]
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Only 50 mg of erbium oxide in 1.5 ml of 10% HCl? I'll have to check out solubility of Er-chloride and Er-sulfate.
I now have been able to dissolve some of my Er2O3 in dilute H2SO4. But still, the amount is very small. The solution is very pale pink, nearly
colorless and a lot of fine powder remains undissolved in the liquid. It now is standing here for more than 2 days and still hardly any material
dissolved. The same is true for my brew in 10% HCl. The HCl-stuff is even more dilute.
I am somewhat disappointed about this behavior. I want reasonable concentrated solutions of Er(3+) salts, which I can use for further experimenting.
Maybe I should make very dilute solutions and then boil them down.
I personally am inclined to think that blogfast25 is right with his remark about calcining of oxides. Apparently mine is calcined more strongly than
Pok's.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Woelen:
Try fusing with NaHSO4? That tends to digest even the hardest of oxides.
Then convert back to the hydroxide, the fresh hydroxide should dissolve much more easily.
Re. calcination. My homemade Nd2O3 (magnet Nd) was prepared simply by precipitating as hydated hydroxide with strong ammonia, very scrupulous washing
and then drying on a full heat electical hot plate. It hardly constitutes 'calcining' in the narrow sense of the word. Yet the product proved almost
completely impervious to hot 36 % HCl. And NdCl3 is of course highly soluble (I can't actually get it to crystallise!)
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I just allowed the material to stand for a long time. I added the Er2O3 to 10% HCl 5 days ago and it did not dissolve. Now, however, 5 days after
adding it to the dilute HCl, all of the Er2O3 has dissolved and the solution is pale pink. Two days ago, the solution also was pale pink, but at that
time, there still was some fine powder at the bottom of the test tube. I swirled it to mix the powder through the acid. So, my Er2O3 does dissolve in
dilute HCl, but it takes a long time.
I also have some of the Er2O3 still standing in 20% H2SO4. This material still did not dissolve. However, its appearance has changed. It is not a fine
powder anymore, but a thin crust, sticking to the glass. The color of this crust is more intense. The solution is pink. It seems as if the erbium
oxide is converted to crystalline erbium sulfate, which does not dissolve further in the 20% H2SO4. Apparently Er2(SO4)3 is only sparingly soluble in
water, especially in the presence of a high concentration of sulfate ions from another source.
But I am happy with these results. I now put a somewhat larger amount of Er2O3 in a little bottle, add some HCl and put it aside for a week or so. I
then expect to have a decent amount, suitable for experimenting.
I did not try the suggested method of melting with NaHSO4. Now I have a method of dissolving it, although this method requires some patience and
planning.
[Edited on 11-10-12 by woelen]
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
It's a non-standard usage, but has anybody tried dissolving these calcined rare earths with a Soxhlet apparatus and azeotropic HCl? If the problem is
that there are low-solubility phases on the exterior, washing with fresh solvent should help. I'll freely acknowledge this might not be the problem. A
fritted thimble might be required for solvent compatibility. Possibly no thimble is needed, if decomposition to powder get you rapidly enough to the
end goal.
|
|
|
annaandherdad
Hazard to Others
Posts: 387
Registered: 17-9-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by woelen | | I just received some erbium oxide (Er2O3), which I purchased from eBay. This is a fine compact pink powder. |
I recently bought some, too. I hope we weren't bidding against each other. All I've done with mine so far is admire its pretty color.
BTW, I like your work on Faraday's law, I may try to build something following your ideas. My daughter is studying electricity in school.
Any other SF Bay chemists?
|
|
|
annaandherdad
Hazard to Others
Posts: 387
Registered: 17-9-2011
Member Is Offline
Mood: No Mood
|
|
I hope this is not a naive comment. A while back I tried dissolving some magnetite, Fe3O4, in HCl. I obtained the magnetite by magnetic separation
out of common sand. I found that some would dissolve in the HCl, coloring the solution as I would expect for a mixture of FeCl2 and FeCl3. But even
after waiting weeks the magnetite was only partially dissolved. I tried grinding the magnetite into a fine powder before applying the acid, with the
same result. I did not try boiling it, nor did I try H2SO4.
Any other SF Bay chemists?
|
|
|
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I recently purchased a lot of Nd2O3 again, because of its very low price (only EUR 53 per kilo!!). The material is a light grey very fine powder with
a bluish hue under fluorescent light, it has a pinkish hue in daylight.
This Nd2O3 is surprisingly hard to dissolve. The other Nd2O3 I have (from 2012 or so) dissolves easily when added to dilute mineral acid. My new batch
does not dissolve at all, not even in boiling hot acid. It does not dissolve at all, not even the faintest pink/lavender color can be observed after
prolonged boiling in 10% HNO3.
The other older sample of Nd2O3 gives a faint hissing noise when added to the acid, leads to slight warming of the liquid and quickly dissolves.
Getting the last little amount of solid particles in solution only requires gentle heating for a few minutes.
So, different samples of the same oxide can have very different behavior.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is something that may be worth a try based on this paper ("SOLUBILITIES OF SOME CALCIUM MINERALS AND PREPARED CALCIUM COMPOUNDS IN EDTA ... ")
noting for some calcium calcined salts EDTA, via complexing, was able to improve solubility.
Link: https://www.google.com/url?sa=t&source=web&rct=j&...
With respect to acids, try Oxalic acid which is noted for forming complexes as well.
[Edited on 14-8-2015 by AJKOER]
|
|
|
annaandherdad
Hazard to Others
Posts: 387
Registered: 17-9-2011
Member Is Offline
Mood: No Mood
|
|
Can anyone tell me what calcining does to the material (its crystal structure or whatever) to reduce the effectiveness of the reaction with acid?
What goes on at the atomic level?
Any other SF Bay chemists?
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by annaandherdad | | Can anyone tell me what calcining does to the material (its crystal structure or whatever) to reduce the effectiveness of the reaction with acid?
What goes on at the atomic level? |
I don't know exactly either and it probably depends also on material, temperature and duration of the calcination.
One effect of drying/calcining is reduction of water content to the point of complete elimination and water is obviously a vector for water soluble
acids.
After that, high temperature calcining must somehow 'close' the crystalline structure of the material, making it less and less 'accessible' to acids
(or H<sub>3</sub>O<sup>+</sup> ions).
It's important to note that the solubility limit of the material in acid is probably not affected, but the dissolution rate
is (very negatively).
[Edited on 14-8-2015 by blogfast25]
|
|
|
annaandherdad
Hazard to Others
Posts: 387
Registered: 17-9-2011
Member Is Offline
Mood: No Mood
|
|
Thanks, my Fe3O4 was probably calcined very well in the molten magma out of which it crystallized. I get a slight solution in HCl after several days,
but nowhere near complete. I may try fusing it with NaHSO4 to see what happens.
Any other SF Bay chemists?
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Sodium bisulphate is great for that kind of thing. Needless to say, the finer you grind the target substance, the higher the dissolution rate is
likely to be.
I'd appreciate it if you would post any experimental results with that particular oxide here or in the 'Rock molester's club' thread. Thanks.
[Edited on 15-8-2015 by blogfast25]
|
|
|
Lion850
National Hazard
Posts: 514
Registered: 7-10-2019
Location: Australia
Member Is Offline
Mood: Great
|
|
I purchased pink Er2O3 and light blue Nd2O3 from a pottery supplier. Both dissolve reasonably easy in either strong sulphuric acid or hydrochloric
acid.
When dissolving these in sulphuric acid remember the sulphate salt is not very soluble, I assume 5g per 100ml at my room temp of about 35 C. With
sulphuric acid I perform the reaction in approx 50% acid and then once it seems complete resulting in sulphate ppt I add water to see that it all
dissolves. Earlier today I made 25g of erbium sulphate and that completely dissolve to clear pinkish solution in 550ml water. Tomorrow I will boil
that down to around 200, at 100C the sulphate solubility is very low, meaning the boiled down boiling solution can be filtered hot to recover more
than 80% of the sulphate. This also get rid of the majority of the excess acid.
Cerium oxide, CeO2, is a different beast to dissolve in acid: hot (120 to 140c) concentrated sulphuric acid is the way to go. Once all turned yellow
decant off the excess sulphuric acid. Wash with glacial acetic acid, filter and dry in a desiccator to get the bright yellow anhydrous cerium
sulphate. This is not the only way but worked well for me.
|
|
|
B(a)P
International Hazard
Posts: 1110
Registered: 29-9-2019
Member Is Offline
Mood: Festive
|
|
| Quote: Originally posted by Lion850 | I purchased pink Er2O3 and light blue Nd2O3 from a pottery supplier. Both dissolve reasonably easy in either strong sulphuric acid or hydrochloric
acid.
When dissolving these in sulphuric acid remember the sulphate salt is not very soluble, I assume 5g per 100ml at my room temp of about 35 C. With
sulphuric acid I perform the reaction in approx 50% acid and then once it seems complete resulting in sulphate ppt I add water to see that it all
dissolves. Earlier today I made 25g of erbium sulphate and that completely dissolve to clear pinkish solution in 550ml water. Tomorrow I will boil
that down to around 200, at 100C the sulphate solubility is very low, meaning the boiled down boiling solution can be filtered hot to recover more
than 80% of the sulphate. This also get rid of the majority of the excess acid.
Cerium oxide, CeO2, is a different beast to dissolve in acid: hot (120 to 140c) concentrated sulphuric acid is the way to go. Once all turned yellow
decant off the excess sulphuric acid. Wash with glacial acetic acid, filter and dry in a desiccator to get the bright yellow anhydrous cerium
sulphate. This is not the only way but worked well for me. |
35C room temperature! You are clearly well north of me! I long for those temperatures again and we are only halfway though autumn.....
|
|
|
Bezaleel
Hazard to Others
Posts: 444
Registered: 28-2-2009
Member Is Offline
Mood: transitional
|
|
| Quote: Originally posted by Lion850 | I purchased pink Er2O3 and light blue Nd2O3 from a pottery supplier. Both dissolve reasonably easy in either strong sulphuric acid or hydrochloric
acid.
|
Lucky you. Mine requires heating in H2SO4 to far over 200C, and only then water comes over in the amounts of 1 drop every 10 seconds (using ~100ml of
H2SO4 and a spoonful of Nd2O3).
| Quote: Originally posted by Lion850 | | Cerium oxide, CeO2, is a different beast to dissolve in acid: hot (120 to 140c) concentrated sulphuric acid is the way to go. Once all turned yellow
decant off the excess sulphuric acid. Wash with glacial acetic acid, filter and dry in a desiccator to get the bright yellow anhydrous cerium
sulphate. This is not the only way but worked well for me. |
Sounds doable to me, if it already reacts at 140C.
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
Cerium oxide (CeO2) is different to the others mentioned here in that it has cerium in oxidation state of +4 as opposed to +3 which would be Ce2O3.
CeO2 should be soluble in sulphuric acid to give Ceric sulphate which is yellow and an oxidising agent. Ce(IV) solutions are not very stable and tend
to be reduced to Ce(III) easily which is colourless in solution.
If you're not part of the solution, you're part of the precipitate.
|
|
|
Fantasma4500
International Hazard
Posts: 1677
Registered: 12-12-2012
Location: Dysrope (aka europe)
Member Is Offline
Mood: dangerously practical
|
|
googling a bit, i dont anything on ultrasonic cleaner being used to dissolve calcined metal oxides- has it been done before? i believe if you hit the
perfect frequency you can make anything happen. and dissolving some damn metal oxide with warm acid shouldnt require super god like precision, the
problem is the crystalline structure. but how much does it take to alter the structure of an acid, or make it penetrate deeper or maybe the metal
oxides structure?
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
