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bfesser
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<strong>adamsium</strong>, are you suggesting a <a href="http://en.wikipedia.org/wiki/Double_salt" target="_blank">double
salt</a> <img src="../scipics/_wiki.png" />? Something like <a href="http://en.wikipedia.org/wiki/Didymium"
target="_blank">didymium</a> <img src="../scipics/_wiki.png" /> sulfate, perhaps?
[Edited on 7/9/13 by bfesser]
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adamsium
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Quote: Originally posted by bfesser  | | <strong>adamsium</strong>, are you suggesting a <a href="http://en.wikipedia.org/wiki/Double_salt" target="_blank">double
salt</a> <img src="../scipics/_wiki.png" />? Something like <a href="http://en.wikipedia.org/wiki/Didymium"
target="_blank">didymium</a> <img src="../scipics/_wiki.png" /> sulfate, perhaps? |
Possibly, but I really don't know, to be honest. I was just wondering if the neodymium giving the orange-coloured compound might not even be
neodymium, but one of, or a mixture of, the lanthanides (quite possibly including neodymium). I suppose this might result in a double salt and some
quick searches reveal a number of papers on lanthanide double salts (but they seem to be with alkali metals). I'm confident that woelen wouldn't have
been likely to make such an error, but it's certainly not unheard of for sellers to mislabel things (presumably - and hopefully - accidentally).
[Edited on 2-7-2013 by adamsium]
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[Edited on 7/9/13 by bfesser]
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woelen
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@Blogfast25: The tests with iron were as follows:
- Solution of thiocyanate was nearly saturated, i.e. of very high concentration
- I used a test tube for the tests and mixed appr. 1 : 1 by volume the test solution and the concentrated thiocyanate.
- The test tube is a standard 16 mm test tube, so I decided positive testing on the basis of viewing through appr. 15 mm of liquid. Maybe if I worked
on a larger scale I could achieve even better sensitivity, but I did not want to spend too much of my thiocyanate.
Your red Nd-sulfate looks quite much like mine (in the small vial), but mine is somewhat more brown/orange. The material in the large vial is more
pink/purplish.
@adamsium: The red/orange material is made from Nd2O3, which was sold to me as 99.9% pure material by eBay seller Jarmond Brinkley. Of course, the
99.9% number must be taken with a grain of salt, but it would be really bad if this is so-called "didymium". The color of the oxide is as it should
be. Pure Nd2O3 has a very pale blue/gray color and my sample has exactly that color. When dissolved in dilute acids, it gives a nice lavender solution
under TL-light and a pinkish solution under daylight.
Another indication that it is not "didymium" is that on strong heating of the oxide, it does not turn dark. Didymium is a mix of praseodymium and
neodymium and Pr2O3 turns dark on strong heating in contact with air, due to formation of Pr6O11 (which contains Pr in oxidation state +4 besides Pr
in oxidation state +3). Samples of "didymium oxide" turn dark on heating, but my sample does not, I tried that some time ago and I heated until the
oxide was red hot and glowing.
[Edited on 2-7-13 by woelen]
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blogfast25
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Thanks woelen. Acidified, that means a detection limit of just under 1 ppm (expressed as Fe2O3).
My sulphate looks a little darker to the naked eye than the photo shows. I think your reddish sulphate, the reddish sulphate obtained by wizzard (and
a few others including me on the 'trouble with neodymium' thread) are one and the same compound. But the pinkish substance must differ a bit. This is
not a double salt and the colour isn't due to contamination, I'd put money on that.
I'd like to offer help by means of Na2EDTA titration, if I can find a pret-a-porter method on the Tinkerwebs, to determine actual Nd content
of both substances.
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turd
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| Quote: Originally posted by bfesser | | <strong>adamsium</strong>, are you suggesting a <a href="http://en.wikipedia.org/wiki/Double_salt" target="_blank">double
salt</a> <img src="../scipics/_wiki.png" />? Something like <a href="http://en.wikipedia.org/wiki/Didymium"
target="_blank">didymium</a> <img src="../scipics/_wiki.png" /> sulfate, perhaps? |
"Didymium sulfate" certainly is a (Pr,Nd)2(SO4)3 solid solution not a double salt, don't you think so?
<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: fixed
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[Edited on 7/9/13 by bfesser]
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bfesser
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Yes, <strong>turd</strong>, I think you are correct. Thank you for pointing that out.
<a href="http://en.wikipedia.org/wiki/Solid_solution" target="_blank">Solid Solution</a> <img src="../scipics/_wiki.png" />
<strong>woelen</strong>, perhaps I missed it, but have you tried desiccating the crystals?
[edit]
Did anyone else notice the "neodymium metal reacts <em>enthusiastically</em> with all the halogens" on the Wikipedia page?
Enthusiastically‽
[Edited on 7/9/13 by bfesser]
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blogfast25
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For what it's worth, the hexahydrate of NdCl3, looks quite pinkish, at least going by the Wiki photo:
http://en.wikipedia.org/wiki/Neodymium(III)_chloride
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adamsium
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| Quote: Originally posted by blogfast25 | For what it's worth, the hexahydrate of NdCl3, looks quite pinkish, at least going by the Wiki photo:
http://en.wikipedia.org/wiki/Neodymium(III)_chloride |
So does, the sulfate, though. http://en.wikipedia.org/wiki/File:Neodym(III)sulfat.JPG
I checked, and it's not woelen's picture this time 
Also, I googled for images of neodymium oxide and, while many of them were the pale blue/gray colour as described by woelen and seen in his picture
(http://en.wikipedia.org/wiki/File:Neodymium_oxide_170g.jpg), there were also a couple of pink coloured ones.
The plot thickens.....
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woelen
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The color of the material in that picture of Nd-sulfate matches mine quite well.
I actually have two batches of Nd2O3, purchased at very different times and from different sellers. Both batches, however, are very pale blue/gray.
One of the batches was sold to me as 99.9% and the other was sold to me as being free of iron. I tested that and indeed it is free of iron (up to PPM
or so level). The other "99.9%" sample has clearly visible presence of iron, but it is at the 10...100 ppm level, not more. That sample is slightly
less bluish than the iron-free sample, but the difference is VERY small and I could not make convincing pictures, showing the difference.
Are you sure the pink coloured oxide is Nd2O3? I also have some Er2O3 and that material has a pale pink color, but the color is much stronger than the
color of Nd2O3.
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adamsium
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No, I can't be sure. Just that it showed up in a few google image results, even though most of them were exactly like yours, so it may just be
mislabelling.
Here's one example (there were only really a few): http://www.made-in-china.com/showroom/gonare/product-detailA... Note the description of "Lavender powder"
[edit:] I followed the link to the erbium oxide on the same page linked above: http://www.made-in-china.com/showroom/gonare/product-detailL... and it looks quite different, much paler. This could be lighting, but at least it
isn't exactly the same picture or anything like that.
[Edited on 4-7-2013 by adamsium]
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blogfast25
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Which of your two samples, though?
Edit: my homemade Nd2O3 was off-white but I'm nor sure about Fe content. Low but how low?
I'm pushing ahead with determining the MW of my Nd sulphate sample, with Na2EDTA titration. Just collecting a few bits and bobs still needed. The
offer of analysing your samples still stands (and you'd get them back but as wet hydroxides). I'd need about 4 g of material.
It all reminds me a bit of ferric ammonium alum: I've obtained that (own preparation) as beautiful amethyst-like crystals but also as slightly beige
crystals. Never got to the bottom of that either.
[Edited on 4-7-2013 by blogfast25]
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woelen
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The sample in the larger vial, the other sample is much more orange/brown.
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phlogiston
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If you want I would be happy to record UV/Vis absorbtion spectra of solutions of both samples if we can exchange a small sample on short notice
(before the end of next week). I'll be changing employers and not sure I will have access to UV/Vis equipment after that. A few small crystals should
be more than enough and I will return the solution/leftover crystals to you ofcourse.
This may resolve any questions regarding the relative abundance of rare earths in your samples, as they tend to have different, specific and sharp
absorbtion bands in the UV/Vis region.
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woelen
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That would be very nice if you can try that. I can send you both complete samples in a small bubble envelope. if you provide me with your address
through U2U, then I'll post the samples next weekend, and you will have them at the start of next week.
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Bezaleel
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Phlogiston's offer is nice, but I bet the outcome will be of little help.
About half a year back, I had the impression that my neodymium sulphate sample was impure due to the presence of praseodimium. This would not be
strange, as my Nd-sulphate was made from magnets. The colour of the Nd-sulphate was reddish, as in Woelen's small vial, and as on Blogfast25's picture
on the 1st page. My sulphate consisted of larger crystals, a few mm in size.
I dissolved these crystals in order to perform tests for praseodymium, and they dissolved slowly, forming a reddish solution. Now, time passed by, and
I never took the time to get to analysis or separation. Because the beaker with the solution was not completely covered, it evaporated from 80 ml to
40 ml in about 3 months' time. Crystals started to appear on the bottom and walls of the beaker. Funny to say that the newly formed crystals are pink,
as in the large vial on Woelen's picture. So without any addition of chemicals (except for deionised water), the crystals have gone from reddish to
pink.
What I suspect is that there exist two different crystal types, one with a reddish colour, the other with a pink colour. What determines the choice of
the crystal type, I don't know. It could even be a distinct region in a multi parameter space (concentration, pH, temperature, evaporation speed, type
and concentration of impurities, presence of (undissolved) nano crystalline nuclei, etc.).
Note that polymorphism is known for other substances, e.g. telluric acid, H6TeO6, which "crystallises both cubic (in octahedra) and monoclinic (in
pseudo-trigonal doublets and triplets). From hot nitric acid of a certain concentration both modifications are formed side by side; a gradual
transformation of the cubic form into the monoclinic takes place only with dilution of the nitric acid. In the dry state both forms are very stable,
and no transformation of the one into the other takes place." (Gossner, Zeits. f. Kryst. 1903, 38, 501 as cited in An Introduction to
Crystallography, P. Groth, 1906, pp. 28-29)
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adamsium
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You may well be right in regard to the polymorphism, and it will be very interesting to find out what is going on here, but the test will still be
helpful. Even just getting a sort of 'non-result' is helpful and, in this case, will lend credence to your theory.
A common example of running tests that really don't tell us anything new but are still important is when we are doing synthesis. We already know what
we have (hopefully) made, but we still run it through NMR and perhaps some sort of MS, etc, to be sure. It's not telling us anything groundbreaking
(again, hopefully!), but it's still very important and can prevent a lot of wild-goose chases. Confirmatory tests are always helpful.
Regardless, I look forward to seeing phlogiston's results, whatever they may be.
[edit]:
I agree, and with Bezaleel's description, this seems even more likely.
It will be interesting for woelen to attempt recrystallising under different conditions when he receives his samples back.
[Edited on 5-7-2013 by adamsium]
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blogfast25
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| Quote: Originally posted by Bezaleel | | What I suspect is that there exist two different crystal types, one with a reddish colour, the other with a pink colour. |
That's certainly a plausible hypothesis.
UV/VIS would exclude contaminants. My money is on not finding any.
[Edited on 5-7-2013 by blogfast25]
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phlogiston
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Allright, so Woelen's send me his beautiful samples last weekend, and here are the results of UV-VIS absorbance spectrometry.
I will refer to the purple sample as 'purple' and the other one as 'brown'.
<b>Methods and results</b>
1. I took 5 crystals of each sample, weighed them (both about 50 mg) and added 1 ml of very pure water.
2. The samples did not completely dissolve at room temperature. I gave it a few hours with occasional agitation.
3. I noticed the solution from the brown sample was turbid, while the one from the purple sample appeared clear (brown on the right):
4. Recorded spectra 200-1100 nm of both solutions (quartz cuvettes, specord 205 analytik Jena double beam spectrophotometer. Blanked the instrument
with water)
5. In an attempt to remove the insoluble material, I centrifuged the samples for 2 minutes at 20,000*g. The brown sample yielded an off-white pellet
and a seemingly clear solution. The purple solution did not yield a visible pellet and appears unchanged.The two faintly purple solutions now look
indistinguishable to me, both in sunlight and in fluorescent light (picture taken in sunlight):
6. Recorded new spectra. (Legend: <b>brown</b>=before centrifugation / <b>centrifugation</b>= supernatant of brown sample
after centrifugation, <b>purple</b>=purple before centrifugation):
7. Because the baseline improved but a small difference remained, I vacuum-filtered the solution through a 0.2 micron filter and recorded new spectra.
I will post the spectrum later (forgot to take the data with me), but the baseline at low wavelengths remains somewhat increased compared to the
purple sample.
<b>Discussion</b>
1. The brown crystals contain a small amount of very fine insoluble particles which presumably cause light scattering at low wavelengths. Supporting
this, after centrifugation and filtration, the light scattering is reduced. If true, then some of the suspended particles must be extremely small,
because the are still present after 0.2 um filtration. They will be very difficult to remove by the means of filtration available to the amateur
chemists. The off-white insoluble material was not identified, but I think one possibility is residual
Nd<sub>2</sub>O<sub>3</sub>.
2. The major peaks (355nm, 523 nm, 577 nm, 742 nm, 796 nm, and 866 nm) ) correspond with those of other neodymium salts, so I think they can be
ascribed to Nd<sup>3+</sup>
3. Any peaks that differ between the two spectra are of special interest, since they may represent impurities different between the two preparations,
perhaps of other lanthanides. I can find only a single small peak in the brown (centrifuged) sample and not in the purple sample, at 446 nm.
Praseomymium is reported to have a peak at this wavelength, but I don't know if it is large at all, and there should be other peaks if the sample
contains significant amounts of Pr. Without a reference spectrum I don't feel confident attributing this peak to Pr<sup>3+</sup>. I do
have a sample of praseodymium oxide, but I will not have an opportunity to retrieve it from storage before monday, which is the last occasion that I
can access this UV-VIS machine in the near future.
I have a UV-Vis machine at home too, but it needs repairs before I can use it for making spectra, so when I get around to do that I will try to record
a Pr<sup>3+</sup> spectrum. I will see if I can obtain a suitable reference spectrum from literature.
I will post (edit) the raw spectral data as excel files later.
[Edited on 14-7-2013 by phlogiston]
[Edited on 14-7-2013 by phlogiston]
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blogfast25
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Excellent work, phlogiston.
My own 'brown' Nd sulphate was also very slightly turbid when I prepared a 0.05 N solution of it (for complexometric MW determination).
I think your results conclusively show there is no significant (or significantly different) contamination between the samples.
This seems to leave only the possibilities of:
1. both are different hydrates
2. they are different modifications (crystalline structures) of the same hydrate
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Bezaleel
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Great work, Phlogiston!
A nice surprise is that the brown sample contains very fine undissolved matter. It might explain the recurrence of a different crystal
structure when it is brought in solution and allowed to crystallise again - a structure different from the other sample that dissolves completely....
I'm looking forward to see the spectra!
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woelen
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Phlogiston, thanks very much for doing this test! Interesting to see that the red/brown sample has very fine solid particles. This may of course
explain the difference in color. Even a very faint opaqueness in the crystals may cause a significant difference in perceived color. I indeed made the
brown sample from Nd2O3 (which I purchased from eBay seller Jarmond Brinkley). The pink sample was made from metallic Nd (heavily oxidized, due to me
storing it improperly), purchased from an eBay seller, called Metallium.
I will make a webpage about all of this and then I'll include the spectrum-images as well and of course a word of thank to you 
----------------------------------------------------------------
In the meantime I also did some new experiments. I dissolved a few grams of Jarmond Brinkleys Nd2O3 in dilute H2SO4 (appr. 1 M). Initially, I had some
fizzling (apparently the oxide contains some carbonate as well) and there was slight warming up of the liquid. Not all of the oxide dissolved. I
needed to heat the liquid to get all of it dissolved.
At this point there were quite a few interesting observations:
1) Initially, the liquid I obtained is lavender in TL-light and pink in daylight.
2) On stronger heating, the liquid turns yellowish in TL-light, still pink in daylight, albeit somewhat duller/brownish pink.
3) On boiling, a lot of pink fine crystalline solid drops out of the liquid. The remaining liquid becomes yellowish brown in TL-light.
Next, I decanted the yellowish brown liquid and allowed this to evaporate slowly in a petri dish.
I rinsed the pink crystalline solid two times with boiling hot distilled water and after this treatment I dissolved the pink solid in water. I needed
quite a lot of water to get all of it dissolved. For appr. 5 grams I needed well over 100 ml of water. The solution is pink/lavender under TL-light,
pink under daylight. I have the impression that it is slighlty opalescent, but only very slightly so. Just to be sure, I added a single drop of 1 M
H2SO4 to the solution and tranferred this to an evaporation disk. I allowed it to evaporate slowly as well.
The solution in the petri dish evaporated in nearly three days, giving red/brown crystals again, very much like the ones in the small vial. I allowed
appr. 90% of liquid to evaporate, the rest I transferred to a test tube and I rinsed the crystals with boiling hot distilled water.
The solution in the eavporation dish was allowed to evaporate for nearly 95%, I tranferred appr. 5 ml to a test tube and rinsed the crystalline solid
with boiling hot distilled water. The crystalline solid from the evaporation disk was pink, but I did not obtain nice large crystals, they were much
finer. The pink color is the same as the pink color of the fine crystalline solid which crashed out of the boiling hot solution.
From the above, I get the impression that from a strongly acidic solution, you get nice large crystals, while from a (nearly) neutral solution you get
very small crystals.
Finally, I tested for the presence of iron with thiocyanate:
- The small amount of liquid, remaining from the petri dish was positive on iron. Actually, the color was quite strong.
- The 5 ml of liquid, remaining after crystallization from the evaporation disk also was positive on iron quite strongyl, albeit not as strong as the
other one. Apparently the crystals, crashed out of solution by boiling it contained quite some iron, which after recrystallization remains in the
solution instead of in the pink crystals.
- The red crystals and pink crystals, obtained after recrystallization both have negative test on iron.
So, the presence of a (small) quantity of iron apparently does not significantly change the perceived color. The color of the crystalline solid,
obtained initially after boiling the solution is the same as the color after recrystallization.
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Bezaleel
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I wonder how the Jarmond Brinkley oxide would behave if you dissolve it in hydrochloric acid, and reflux that for 10 minutes or so. The rare earth
chlorides have a much higher solubility than the sulphate, also at high temperatures.
My guess is that in hydrochloric, all of the oxide will dissolve, and I wonder what happens if you exchange the chloride for sulphate and crystallise
then. I expect (but am not confident) that then the sulphate will have a pink colour too.
(Exchange of chloride for sulphate: precipitate with NaOH, suction filter and rinse a few times; then dissolve in H2SO4.)
Edit: Oh yes, in order to get rid of the iron, the sulphate should be cleaned as you did in the experiments in your last post, in order to make a
sensible comparison.
[Edited on 15-7-2013 by Bezaleel]
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blogfast25
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| Quote: Originally posted by Bezaleel | I wonder how the Jarmond Brinkley oxide would behave if you dissolve it in hydrochloric acid, and reflux that for 10 minutes or so. The rare earth
chlorides have a much higher solubility than the sulphate, also at high temperatures.
|
The problem is that with HCl you're severely limited in concentration (max. 37 w%). Concentrated HCl often doesn't even make a dent in oxides that
have been calcined. That's the primary reason for using H2SO4 95 % (or better).
@woelen:
I made an interesting observation. Preparing my Nd sulphate to make a solution of exactly known concentration (about 0.05 M), I decided to dry it
overnight in a CaCl2 desiccator, to get rid of possible surface moisture. In that period, it seemed to me to have undergone a bit of efflorescence. I
weighed it and put in in the desiccator for another 24 H and it lost another 2.5 w%. That's equivalent to about 1 mole of water based on the
octahydrate. It appeared (subjectively) more pinkish).
I really need to repeat this experiment to be sure of what I'm writing.
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woelen
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The Nd2O3 I have has no problems of inertness at all. It very easily dissolves in acid, also in dilute 10% HCl. In dilute sulphuric acid some heating
is required to get all of it dissolved, in dilute hydrochloric acid or dilute nitric acid it simply dissolves in a few minutes. This is in stark
contrast with my Er2O3 I also have from Jarmond Brinkley. That material must be boiled for hours to get only a small part dissolved, or you have to
let it stand at room temperature for many days. But finally, with sufficient patience and with sufficient excess amount of acid that also dissolves.
But back on the topic of Nd-ions. I also noticed a peculiar effect when a solution of NdCl3 in excess dilute HCl is heated. On heating, the solution
turns from lavender through grey to yellowish grey and on cooling down it slowly (takes a few hours) turns back to lavender. Apparently, with chloride
it gives some coordination complex in the hot solution, which slowly decomposes again when it cools down.
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blogfast25
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Well, mine was homemade, not calcined at all (just dried very thoroughly) and didn't dissolve in 37 % appreciably!
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