Sciencemadness Discussion Board - two times neodymium sulfate; different colors

This crystals are having a bad habit if so then!

<a href="http://en.wikipedia.org/wiki/Word_salad" target="_blank">"Has anyone really been far even as decided to use even go want to do look more like?"</a> <img src="../scipics/_wiki.png" />

[Edited on 7/9/13 by bfesser]

Wizzard - 30-6-2013 at 08:03

I think the color is a product of thickness - Thin cross-sections make the samples rather red, and a Amethyst like purple... Thicker crystals are more rosy-purple to lavender.

But your crystals look downright orange!!!! Perhaps it's some other Lanthanide?

I'm going to second that it may be different levels of hydration - Like copper sulfate

Lesser hydration (unknown, exactly... I have a thread somewhere about this)
http://www.theodoregray.com/periodictable/Samples/CuSO4.c/s1...
Normal hydration:
http://www.theodoregray.com/periodictable/Samples/CuSO4.a/s1...

Poppy - 30-6-2013 at 08:32

This crystals are having a bad habit if so then!

lol, I don't know the purpose of that, but I'll try to clarify it some more:
from Crystals and Crystal Growing, by Alan Holden and Phylis Morrison
pag. 310 ref, picked at line 17, referred pag 123 Crystallographers call the characteristic shape of a crystal its "habit". They might say, "The habit of alum is depicted j]in that figure", or they might speak of alum as "having an octahedral habit."
Moreover, for proof there isn't any kind of single path idea of determinism in chemistry,
Pag. 270 "...Notice that both of the substances, sugar and rochelle salt, which have this effect on light are derived from living thigs (?)" ... "Unless some "life process" takes part, either in making the starting material for the chemical operations or as one of the process deliberately used in manufacture, the product will be as inert as water or glass between two sheets of Polaroid" ... "You may protest that your crystals of sodium chlorate have this effect on light, whereas their solutions does not, and no life process intervenes between the solution and the crystals growing from it" ... "In other words, the "life process" which intervenes is you - the choice you make among your collection of seeds"
Then, finally, on retaliation:
pag. 273 "Almost always, when two kinds of molecules are mixed together, they form crystals in which the two kinds are still mixed together, and the crystal has a plane or a center of symmetry. This mixture is like that of potassium sulfate and aluminium sulfate in alum; equal numbers of each kind of molecules cooperate to form the orderly structure. In rare cases (perhaps not that rare as in the case of Nd compounds) the two kinds of moelcules will crystallize out separately. Then a Crystal made from either kind of molecule alone will lack a plane of symmetry."
Which is more than enough to start suspecting the apparent unreasonable color differences of both samples.

Perhaps the method of assay or contaminants may trigger one or another of the crystal habits, as it may have many, and so many are the sorts of shapes relating to its symmetry, which implies color changes sometimes. sorry if thats not the case.
Hope I'm not convincedly playing jibber-jabber around, thank you.

[Edited on 6-30-2013 by Poppy]

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: fixed broken image(s)]

[Edited on 7/9/13 by bfesser]

blogfast25 - 30-6-2013 at 08:50

woelen:

So much for '99.9 %' of course. The ferric thiocyanate test is sensitive but is it that sensitive? That would be worth testing or lit. researching. In any case, the crystal samples are free of Fe.

I think I may have obtained both colours at once in one of my Nd experiments. I obtained fairly large crystals of Nd sulphate very much the same colour as the 'small vial' and a much lighter, pinkish form of smaller crystals. The 'small vial' colour is what dominates the 'trouble with neodymium' thread, in terms of Nd sulphate colour.

I would do this. Redissolve some of both substances in iced water, in identical conditions of mass, volume of solvent and temperature.

Test one. Heat some of both solutions and see what happens.

Test two. Add some conc. H2SO4 to some of both solutions and see what happens.

Test three. Allow some of both solutions to evaporate at RT and see what kind of crystals you get.

I believe the REs can be titrated with Na2EDTA. That would open the way to determine how much water each substance contains and thus which hydrate it is. Or: gravimetrically by precipitating as oxalate. Higher hydrates tend to form at lower temperature and vice versa, as a rule of thumb.

[Edited on 30-6-2013 by blogfast25]

Poppy - 30-6-2013 at 09:00

Or just for god sake take LARGER PICTURES OF ALL THE CRYSTALLSS!! and let us analyse their shapes...
Surely we can't decide which building blocks they belong to but we can see if they differ from one another quite easily.
Also, on addition to Blogfast25's suggested tests, repeat the one in which you dissolved them first with HCl.

I just redissolved the pink sample of Nd2(SO4)3 I had. I don't know if that was hydrate or not. I think I messed up a bit by not acidifying properly the water before dissolving it, but then added like 5mL H2SO4 to a solution comprising 400mL filter water containing 21g of the Nd2(SO4)3, which didn't dissolve completly until now. The solution is turbid. Anyway, lets see what crap I get after this refresh.

[Edited on 6-30-2013 by Poppy]

kmno4 - 30-6-2013 at 10:02

- Dissolve some of the crystals in cold water and add thiocyanate. Not even a faint red coloration can be observed, so the iron has not contaminated the crystalline solid. This is true for both samples of crystals.

SCN test for Fe(III) is pH sensitive when concentration of Fe is small. I would repeat it, but first add some (pure) HCl to the Nd solution. Next, you can add some insoluble in water alcohol (butanol etc) and extract water phase by shaking (in test tube).
Colourless organic phase means "Fe-free" solution.

[Edited on 30-6-2013 by kmno4]

bfesser - 30-6-2013 at 10:37

woelen, how do the colors compare outside of those glass vials. I don't expect to see much difference, but perhaps the glass of the different vials has slightly different absorbance/transmission. I'd still like to see what happens when a crystal of each is gently heated.

blogfast25 - 30-6-2013 at 12:14

Quote: Originally posted by kmno4

SCN test for Fe(III) is pH sensitive when concentration of Fe is small. I would repeat it, but first add some (pure) HCl to the Nd solution.
[Edited on 30-6-2013 by kmno4]

I used this test professionally and quantitatively some 30 years ago and can't remember the actual detection limit. It would be worth making a 100 ppm Fe³⁺ solution, and progressively dilute it by factors of 2, up to 'no signal', just to see...

woelen, how do the colors compare outside of those glass vials. I don't expect to see much difference, but perhaps the glass of the different vials has slightly different absorbance/transmission. I'd still like to see what happens when a crystal of each is gently heated.

I appreciate the genuine concern about perception of these colours but the difference is FAR too great to be attributed to a trick of the light or the glass of the vials. This difference is genuine.

Different forms of Nd sulphate (maybe different allotropes [metaphorical use] even?) have been discussed here many times before.

[Edited on 30-6-2013 by blogfast25]

bfesser - 1-7-2013 at 05:50

Do you have any cation exchange resin available for column chromatography? I don't see people from ScienceMadness doing much <a href="https://en.wikipedia.org/wiki/Ion_chromatography" target="_blank">ion chromatography</a> <img src="../scipics/_wiki.png" />, and it's a little sad. It's such an easy and effective means of separation for so many inorganic (and some organic) components. And the resin is usually able to be regenerated and reused. I see many members pursuing HPLC, GC, and even centrifugal radial TLC, yet I don't recall seeing a single member performing ion-exchange chromatography. With how easy and cheap it is, it's a wonder nobody here has continually advocated its use. It also seems to me, that it would lend itself well to the style of homebrew 'technochemistry' that people here like to tinker with. You can start with nothing more than a glass tube, resin beads, and a couple dozen test tubes. But members could build their own micro-controller based automated control and detection systems and get as fancy as they like.

[/end off-topic rant]
Sorry.

[edit]
Sadly, a <a href="search.php?token=&srchtxt=exchange+chromatography&srchfield=body&srchuname=&f%5B%5D=all&srchfrom=0&filter_distinct=yes&a mp;searchsubmit=Search">search for "exchange chromatography"</a> yields less than one page full of results—many quite old, most unrelated.

[second edit]

Different forms of Nd sulphate (maybe different allotropes [metaphorical use] even?) have been discussed here many times before.

I think they're called crystal <a href="http://en.wikipedia.org/wiki/Crystal#Polymorphism_and_allotropy" target="_blank">polymorphs</a> <img src="../scipics/_wiki.png" />, but I'm not positive. I'm just beginning to study crystallography (under mineralogy as I pursue geology). Or perhaps you're refering to something like <a href="http://en.wikipedia.org/wiki/Crystal_habit" target="_blank">crystal habit</a> <img src="../scipics/_wiki.png" />?

[Edited on 7/9/13 by bfesser]

watson.fawkes - 1-7-2013 at 06:27

That would open the way to determine how much water each substance contains and thus which hydrate it is.

Calcination would be another way. You can't always take a hydrated salt to an anhydrous salt, but it's typical to be able to take a more-highly hydrated salt to one with lower hydration.

You'd need a balance with good repeatability for it. Even without being able to measure any mass difference quantitatively, you can get a qualitative result with a simple balance beam. First balance the crucible and charge with a nonce weight, such as a dish with sand. Second perform the calcination. Third put the calcined dish back and see if it still balances.

adamsium - 1-7-2013 at 08:41

Perhaps you're seeing some NdCl_n^(3-n)+ complexes, Woelen.

These papers seem to indicate that chloro complexes of lanthanides are increasingly stable at higher temperatures and one (the second one) discusses different results based on the use of neodymium chloride vs neodymium oxide. I've not had time to read them properly and I'm (obviously) far from being an expert on this, anyway, but they would certainly be worth checking out:

Spectrophotometric study of Nd, Sm, and Ho complexation in chloride solutions at 100-250°C, S. A. Stepanchikova and G. R. Kolonin; Russian Journal of Coordination Chemistry, Vol. 31, No. 3, 2005, pp. 193–202. Translated from Koordinatsionnaya Khimiya, Vol. 31, No. 3, 2005, pp. 207–217

The aqueous geochemistry of the rare earth elements and yttrium: VI. Stability of neodymium chloride complexes from 25 to 300°C, C. H. GAMMONS, S. A. WOOD, and A. E. WILLIAMS-JONES; Geochimicaet CosmochimicaActa,Vol.60, No. 23, pp. 4615-4630, 1996

What colour were the solutions? You will also notice that the colour of NdCl₃ (far right) is very much like your product where HCl was used, while the sulfate (far left) is more like the one where HCl was not used. Interestingly, according to the neodymium page on wikipedia, some compounds of neodymium have varying colour under different lighting conditions - see https://en.wikipedia.org/wiki/Neodymium#Compounds and the images of crystalline Nd₂(SO₄)₃, confusingly, look like your product where you used HCl. Perhaps a small amount of a chloro complex was stable enough in solution that the silver nitrate test didn't detect it. So, I really don't know what's going on, but I'm quite curious to find out.

Image reference: https://en.wikipedia.org/wiki/File:Neodymium_tl1.jpg

I'd love to see:

if the appearance changes under different lighting conditions
the effect of adding a large excess of sulfuric acid

bfesser - 1-7-2013 at 09:43

I've used both Amberlite® and Dowex® ion exchange resins in the past. Both are easily obtainable, even for the home chemist on a limited budget. With the ingenuity of the members of this board, I wouldn't be surprised if someone managed to make their own, as well.

Quote: Originally posted by adamsium

Spectrophotometric study of Nd, Sm, and Ho complexation in chloride solutions at 100-250°C, S. A. Stepanchikova and G. R. Kolonin; Russian Journal of Coordination Chemistry, Vol. 31, No. 3, 2005, pp. 193–202. Translated from Koordinatsionnaya Khimiya, Vol. 31, No. 3, 2005, pp. 207–217

The aqueous geochemistry of the rare earth elements and yttrium: VI. Stability of neodymium chloride complexes from 25 to 300°C, C. H. GAMMONS, S. A. WOOD, and A. E. WILLIAMS-JONES; Geochimicaet CosmochimicaActa,Vol.60, No. 23, pp. 4615-4630, 1996

I've <a href="viewthread.php?tid=19098&page=19#pid290268">requested</a> these articles.

[Edited on 7/1/13 by bfesser]

woelen - 1-7-2013 at 10:12

Just to be sure that the effect of the presence of iron can be ruled out definitely, I assessed the sensitivity of my test with thiocyanate.

I dissolved 55 mg of FeNH4(SO4)2.12H2O in 1000 ml of water. This 55 ppm by weight, the iron ppm of course is much lower.

With 55 mg per liter, a bright red/rose color is obtained with concentrated solution of NH4SCN.
When diluted 10 times, the color cannot be observed anymore, but amazingly, when a few drops of reagent grade HCl (30%) are added, then the color becomes pale pink again and the iron/thiocyanate complex is clearly visible again. I of course tested the HCl also without the iron and then it does not give a positive with the thiocyanate.
The liquid can be diluted two to three times before the iron is not visible to me anymore. So, I can go down to let's say 2 mg per liter of dissolved FeNH4(SO4)2.12H2O.

Next, I did the test with thiocyanate again, but now with added drop of 30% HCl. No positive reaction occurred.

I also did another very sensitive test on iron. I prepared a mix of K3Fe(CN)6 and K4Fe(CN)6 and added this solution to another sample of solution of my Nd-sulfate samples. This only results in formation of a pale yellow precipitate (Nd-ferrocyanide), not the faintest indication of a blue color.

So, I a quite sure that no serious amounts of iron are present in the material, at least not so much that it clearly affects the color of the crystals.

@adamsium: The picture you show is from my website

I donated it to Wikipedia. I made it myself some years ago. I still have these ampoules, they are very nice demo objects and sometimes I show them to friends and show the differences under different types of light.

crazy complexing sulphate

Poppy - 1-7-2013 at 12:25

Quote: Originally posted by adamsium

I'd love to see:

if the appearance changes under different lighting conditions
the effect of adding a large excess of sulfuric acid

This sample I have was prepared by rapid hot addition of Nd2O3 to conc. sulfuric acid. Note the picture with a precipitate at the bottom is white, while the solution is pink. Comparing to the undissolved salt, which was pink in the same conditions, this ppt. must be Nd oxide or hydroxides. Their origin is not certain (as from what caused them to precipitate, if they were a component from the first crystalization or if they were formed by dissolution of the salt in neutral conditions which ultimately caused hydrolysis to occur. As the salt was neutralized with aq. NH4OH probably its a residue of bad operating of the ingredients, but that can be ruled out once I recrystalize this from the water - too bad its filtered water

!!! - and dissolve again). Anyway, the hydroxides seems stable as I lately added about 5mL conc. sulfuric acid to this solution and that didn't dissolve the white ppt. seen.

Redissolving to rule out color differences is useless: note the solution after shaking - I know its the ppt which is making it turbid all the way but salts generally don't fully ionize once dissolved (not fully solvated) so ridiculus small seed might stil be wandering around and this experiment can be considered "dead" as it won't give other crystal forms through with simple recrystalizations. Philosophically, maybe not even renautralization of an oxide obtained from that solution may nerf this crystal eed problem: one ought to go into the metallic state and watch out the paths he chooses to dissolve it in first place.

PS.: my samples were INITIALLY dissolved Nd from magnets in HCl, neutralized to hydroxide, calcined and added to hot conc. H2SO4. As Alan Holden's book says, this might influentiate, but not rule out completely, the formation of this specific habit (that is, neutralization gave rise to a habit on the hydroxyde that survived all the way to salt).

[Edited on 7-1-2013 by Poppy]

blogfast25 - 1-7-2013 at 12:33

Just to be sure that the effect of the presence of iron can be ruled out definitely, I assessed the sensitivity of my test with thiocyanate.

I dissolved 55 mg of FeNH4(SO4)2.12H2O in 1000 ml of water. This 55 ppm by weight, the iron ppm of course is much lower.

With 55 mg per liter, a bright red/rose color is obtained with concentrated solution of NH4SCN.
When diluted 10 times, the color cannot be observed anymore, but amazingly, when a few drops of reagent grade HCl (30%) are added, then the color becomes pale pink again and the iron/thiocyanate complex is clearly visible again.

Excellent. What was the testing protocol? In plain English, how much ferric solution, how much SCN- solution and at what concentration [ooops: it does say 'concentrated']?

Quote: Originally posted by watson.fawkes

[Calcination would be another way. You can't always take a hydrated salt to an anhydrous salt, but it's typical to be able to take a more-highly hydrated salt to one with lower hydration.

I fear hydrolysis, Watson. RE salts are quite prone to it. Especially as it is difficult to demonstrate that it occurred (or not). Ideally you’d need a thermo-gravimetrical scan or possibly DSC to determine when the hydrate loses water and when that stops.

@bfesser: chromatography enabled with detection of various REs is outside the capability envelope/budget of 99.9 % here, sadly...

My own neodymium sulphate (it appears slightly darker to the naked eye - cr*pshoot camera):

... but I did also obtain a much more pink version once.

[Edited on 1-7-2013 by blogfast25]

[Edited on 1-7-2013 by blogfast25]

watson.fawkes - 1-7-2013 at 13:03

Quote: Originally posted by watson.fawkes

Calcination would be another way. You can't always take a hydrated salt to an anhydrous salt, but it's typical to be able to take a more-highly hydrated salt to one with lower hydration.

I'll certainly acknowledge that the quantitative versions are more difficult. The qualitative one, though, seems easy to check. Just heat them up and see if they lose mass. If it does, it's either water of hydration (my initial thought) or perhaps driving off HCl if there's a chloride complex about. Either way it's a null result or some non-null result, even if a non-null result wouldn't yield much further information.

One possible result is that one color loses mass and the other doesn't. That would be new information.

Separately, I might mention that a student spectroscope would allow you to see how different the emission spectra actually are, even without being able to take a calibrated spectrum. In addition, if the two samples have different crystal structures, they would likely also have different indices of refraction. I don't know what your crystallographic test equipment is, but given your interests, you might think about some.

P.S. Re-hydration shouldn't be a problem if you weigh the samples quickly after they come out of the oven and if you use a lid.

Edit: P.S.

[Edited on 2013-7-1 by watson.fawkes]

blogfast25 - 1-7-2013 at 13:11

@Watson:

My money would be on crystal water. In the instance where I obtained the material picture above I also obtained a pink, much less crystalline looking material. The latter had formed first, from higher temperature, the former, more reddish compound developed the next day from cold (perhaps saturated) solution. Higher hydrates tend to form in colder areas of the phase diagram.

I might subject a few samples of the reddish salt (believed to be the octahydrate) to a few temperatures and see what happens. Unfortunately I don't have any of the pinkish version left for comparison.

[Edited on 1-7-2013 by blogfast25]

Random - 1-7-2013 at 15:20

NdCl3 is a mauve colored hygroscopic solid whose color changes to purple upon absorption of atmospheric water. The resulting hydrate, like many other neodymium salts, has the interesting property that it appears different colors under flourencent light- In the chloride's case, light yellow (see picture).[1] The color difference is likely due to the charge transfer converting the Nd3+ ions into the Nd2+ state; the orange absorption band of Nd3+ decreases and the green Nd2+ absorption increases producing the yellowish color.[2][3]

From wikipedia. Maybe you somehow made Nd3+ and Nd2+ in separate vials?

Maybe some double salt formed?

Poppy - 1-7-2013 at 16:13

Okay, like gold, after each precipitation it changes color
see: http://philosophicalmercury.page.tl/Blue-Gold.htm

Statistics anyone? seemingly the first crop always looks a little oranger...

bfesser - 1-7-2013 at 20:07

Thanks to solo, <a href="viewthread.php?tid=19098&page=19#pid290291">here</a> are the papers adamsium referred to.

Poppy, as of this post, your link is broken. Here's a <a href="http://webcache.googleusercontent.com/search?q=cache:wdAEMhVPJYwJ:philosophicalmercury.page.tl/Blue-Gold.htm+&cd=1&hl=en&ct=clnk& ;gl=us" target="_blank">cached copy</a> <img src="../scipics/_ext.png" /> of the page on Google—which I can't be arsed to read, because it's on a website entitled Alchemy Illuminated.

[Edited on 7/9/13 by bfesser]

adamsium - 1-7-2013 at 20:12

@adamsium: The picture you show is from my website

Haha, well that's funny. I didn't even see that it was yours and apparently haven't looked at that page on your site. I normally recognise your pictures on wikipedia; many of the ones with compounds in small bottles on a plain background.

Anyway, another interesting thing to try would be to try saturating a solution with chloride, especially with strong heating.

Is there any reasonable possibility that your neodymium was mislabelled? From what I have seen, I know you're much more organised than that, but perhaps your source of the neodymium was not. It seems that most neodymium compounds will be pink, but other lanthanides give other colours, including yellowy orange-ish.

From http://www.aukcjoner.pl/gallery/013246618-.html#I2

bfesser - 1-7-2013 at 20:30

adamsium, 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]

adamsium - 1-7-2013 at 20:50

adamsium, 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]

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: fixed broken image(s)]

[Edited on 7/9/13 by bfesser]

woelen - 1-7-2013 at 22:41

@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]

blogfast25 - 2-7-2013 at 04:40

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.

turd - 2-7-2013 at 04:47

"Didymium sulfate" certainly is a (Pr,Nd)₂(SO₄)₃ solid solution not a double salt, don't you think so?

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: fixed broken image(s)]

[Edited on 7/9/13 by bfesser]

bfesser - 2-7-2013 at 06:31

Yes, turd, 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" />

woelen, perhaps I missed it, but have you tried desiccating the crystals?

[edit]
Did anyone else notice the "neodymium metal reacts enthusiastically with all the halogens" on the Wikipedia page? Enthusiastically‽

[Edited on 7/9/13 by bfesser]

blogfast25 - 4-7-2013 at 05:04

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

adamsium - 4-7-2013 at 05:41

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.....

woelen - 4-7-2013 at 06:32

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.

adamsium - 4-7-2013 at 06:49

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]

blogfast25 - 4-7-2013 at 07:11

The color of the material in that picture of Nd-sulfate matches mine quite well.

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]

woelen - 4-7-2013 at 23:04

The sample in the larger vial, the other sample is much more orange/brown.

phlogiston - 5-7-2013 at 00:42

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.

woelen - 5-7-2013 at 00:47

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.

Bezaleel - 5-7-2013 at 06:37

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)

adamsium - 5-7-2013 at 08:10

Quote: Originally posted by Bezaleel

Phlogiston's offer is nice, but I bet the outcome will be of little help.

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]:

UV/VIS would exclude contaminants. My money is on not finding any.

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]

blogfast25 - 5-7-2013 at 08:14

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]

phlogiston - 13-7-2013 at 18:17

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'.

Methods and results
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: brown=before centrifugation / centrifugation= supernatant of brown sample after centrifugation, purple=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.

Discussion
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 Nd2O3.

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 Nd3+

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 Pr3+. 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 Pr3+ 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]

blogfast25 - 14-7-2013 at 04:28

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

Bezaleel - 14-7-2013 at 13:44

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!

woelen - 14-7-2013 at 23:27

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.

Bezaleel - 15-7-2013 at 08:18

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]

blogfast25 - 15-7-2013 at 11:13

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.

woelen - 15-7-2013 at 23:24

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.

blogfast25 - 16-7-2013 at 05:26

Well, mine was homemade, not calcined at all (just dried very thoroughly) and didn't dissolve in 37 % appreciably!

watson.fawkes - 16-7-2013 at 08:48

Have you ever tried a Soxhelet apparatus for this? Freshly distilled HCl wouldn't change the ultimate solubility, but it might increase the reaction rate.

blogfast25 - 16-7-2013 at 12:05

Quote: Originally posted by watson.fawkes

Have you ever tried a Soxhelet apparatus for this? Freshly distilled HCl wouldn't change the ultimate solubility, but it might increase the reaction rate.

I can see where you're coming from this: a solvent free of NdCl3 would marginally improve rate of dissolution. I doubt if the effect is worth deploying the Soxhlet extractor for it though...

watson.fawkes - 17-7-2013 at 02:35

I doubt if the effect is worth deploying the Soxhlet extractor for it though...

It might not be. Another thing that might be tried with one, however, would be to introduce gaseous HCl into an adapter at the lower end of the condenser. This would allow working with HCl of higher concentration. If the condenser is operating at sufficiently low temperature, it should keep fuming down to a minimum.

blogfast25 - 17-7-2013 at 03:53

Hmmm... I think I'll stick to:

* treat oxide with hot, conc. H2SO4
* treat cooled sulphate with strong NH3, filter and wash filter cake
* dissolve Nd hydroxide in HCl, filter off any non-dissolved matter.

woelen - 17-7-2013 at 04:33

This is another interesting observation. It really is strange that you need hot conc. H2SO4 to get your non-calcined oxide dissolved, while mine dissolves with just gentle heating in dilute 10% HCl or 2 M H2SO4. I even note formation of some heat when the Nd2O3 dissolves.

blogfast25 - 17-7-2013 at 06:52

Yes, it's strange. I only dried it in a glass beaker on a hot plate, full heat, while stirring. It was completely insoluble in HCl.

What do you think of the water loss in CaCl2 desiccator? I need to repeat that one.

Poppy - 30-8-2013 at 14:47

Your results are incompatibly outrageous, my crystal samples, made by evaporation, showed color differences normally not seen with liquid or powdered samples.
Here follows the link of a similar event, for those not familiar with.
http://www.sciencemadness.org/talk/viewthread.php?tid=25501
Pic are going to come when the camera is ready.
You have bigger of a problem than you can handle apparently.

Poppy - 1-9-2013 at 18:55

EDITED:
Sorry mister, for our delay, here comes therefore the mentioned images. The difference is unoticeable, and it comes to being the crystals are very hard, and perhaps what else one would have expected for having crystalizinng skills like that?

Grating you on showing, good moods to follow!

C ya Poppy

[Edited on 9-3-2013 by Poppy]

Bezaleel - 2-9-2013 at 04:54

Poppy, in the poicture on the left I notice that you have various crystals that are _very_ light red, whereas you also have cystals in there that are normal red, i.e. normal with respect to what you would expect from Nd.

What's the size of these crystals? Are you sure you don't have any lanthanum mixed in your Nd?

blogfast25 - 2-9-2013 at 05:32

That difference is not significant, IMHO. One photo is out of focus and the materials seem to be lit differently. Even the background is a slightly different hue in both photos.

To compare the colour of these materials much tighter control of the photographing conditions is needed, otherwise one is comparing apples and pears.

To conclude that this 'difference' may be due to an impurity at this point is simply not possible.

Sorry, for mister weith your deklqay, [...]

Huh?

[Edited on 2-9-2013 by blogfast25]

Poppy - 2-9-2013 at 13:54

Ops sorry, I used a translator which apparently showed some typing erros, and as I was not at home I had to go and forgot to check the spellings - what the hell is that ?!

Referring to the images, one was picked useing TL bulbs and the other was taken on the shine of a flurescent tube. there is no great difference, and sunlight doesn't do different on them. They are 5mm - 9mm wide, mostly, and all of them showed this cube like shape, with stritiations visible in all sides. Cool, but I really think Woelen's original samples are fairly contaminated.
I used extra sulfuric acid in the motherliquor whereby the crystal where recrystalized, so now the crystal should be pretty pure.
Sorry, my camera can't handle zooming closer than that.

phlogiston - 2-9-2013 at 16:47

It would be best if you could take a picture of both of your samples of crystals in a single picture (and therefore under one set of lighting conditions).

PS. please double-check you spelling, Poppy. Your English is good enough. Just get the spelling right, otherwise some words become ambiguous.

Poppy - 2-9-2013 at 17:27

Quote: Originally posted by phlogiston

It would be best if you could take a picture of both of your samples of crystals in a single picture (and therefore under one set of lighting conditions).

The crystals are the very same exact sample!
Woelen's vials are showing an unacceptable color pattern and most pictures of the dimmie's sulphate posted here in the forums seems to appear as a different sort of shape. Also, maybe they didn't grow large enough perhaps.

woelen - 2-9-2013 at 22:42

If these crystals are 5 to 9 mm size, then I think that they only contain Nd-sulfate at low concentration (especially the crystals in the left picture). They almost look like crystals of a colorless salt, in which some Nd-ions are mixed. I would expect crystals of such a large size to be MUCH less transparent. My reddish crystals are as transparent, but they only have sizes of 0.5 to 1.0 mm.

Poppy - 3-9-2013 at 14:58

This can't be! Check the thread on neodimium and you 'll see my yiled was already cut down by the sake of purification, and I steped through the oxide first, so iron contamination could go above 5% in the first crop. Its the third time I'm crystallizing the thing, and only sulfuric was added as a booster.

blogfast25 - 4-9-2013 at 10:26