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Author: Subject: The trouble with neodymium...
Poppy
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[*] posted on 12-11-2011 at 08:51


Hello, so has your method really succeded? good job!!
Thank you!

As in the case of the autoprotolysis I meant some Fe(OH)3 changes in situ into Fe2O3 because the autoprotolysys mechanism, as evidenced by the Schikorr reaction
Sorry for the misleading with the equation for O(2-).

Schikorr reaction starts at about 100ºC and increases as temperature goes up. At temperatures <100ºC the conversion of Fe(OH)3 into Fe2O3 is not even likely to occur, but my advice is to beware of high temperatures!
------------------------
Blogfast:
I get it now, Fe(OH)3 = Fe2O3.n H2O = HFeO2 = FeO(OH) in a different notation.
But powdered dry Fe2O3 would not turn into Fe2O3.n H2O by damping in water correct?
Heating 2 HFeO2 (the mud from reaction of FeIII ions with strong alkali) to 200ºC dehydrated it into Fe2O3.
Would this dry Fe2O3 work too???


For those who wants Fe(OH)3 but are not dealing with H2O2, you can prepare it without H2O2, following this long road production, starting with Schikorr reaction and then drying.
------------------------------------------
2 (Fe2+ → Fe3+ + e–) (oxidation of 2 iron(II) ions)
2 (H2O + e– → ½ H2 + OH–) (reduction of 2 water protons)

To give:
2 Fe2+ + 2 H2O → 2 Fe3+ + H2 + 2 OH–
Adding to this reaction one intact iron(II) ion for each two oxidized iron(II) ions leads to:
3 Fe2+ + 2 H2O → Fe2+ + 2 Fe3+ + H2 + 2 OH–

For electrical neutrality, OH(-) are shown:
3 Fe2+ + 6 OH– + 2 H2O → Fe2+ + 2 Fe3+ + H2 + 8 OH–
3 Fe(OH)2 + 2 H2O → Fe(OH)2 + 2 Fe(OH)3 + H2
Autoprotolysys mechanism: Its essential in the formation of the final product and occurs but only at temperatures about 100ºC or above.
OH– + OH– → O2– + H2O
acid 1 + base 2 → base 1 + acid 2, or also,
2 OH– → O2– + H2O
Then that all progresses as
3 Fe(OH)2 + 2 H2O → (FeO + H2O) + (Fe2O3 + 3 H2O) + H2
3 Fe(OH)2 + 2 H2O → FeO + Fe2O3 + 4 H2O + H2
3 Fe(OH)2 → FeO + Fe2O3 + 2 H2O + H2
The former Schikorr reaction being
3 Fe(OH)2 → Fe3O4 + 2 H2O + H2

I will make a test with Fe(OH)3 as soon as providing some HCl and H2O2 to see if it protolyses into pure Fe2O3!!
For that I'll just heat the ferric hydroxide mud until dryness and somehow check for pure Fe2O3 for its beatyful red color.
Another note, when preparing conventional Schikorr reaction for the production of magnetite, leaving the heating vessel with any opening causes oxygen to enter like crazy and get absorbed by the magnetite turning it into orange Fe2O3, which is a mixture consisting of mostly Fe2O3 and a little bit FeO.
2 Fe3O4 + 1/2 O2 --> 3 Fe2O3
The reaction with atmospheric oxygen was so strong that at 100ºC finely ground magnetite was actually burning!
The test for this orange Fe2O3 was made diluting it into 30% HCl and observing that some residue of FeO (black) was not atacked even with overweighted HCl.
Further burning of the oxide mix seemed not to oxydise anymore FeO.
The Fe2O3 can be dissolved in sulphuric acid and then put to react with strong alkali to produce Fe(OH)3.



[Edited on 11-12-2011 by Poppy]

[Edited on 11-12-2011 by Poppy]
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[*] posted on 12-11-2011 at 10:51


Alright, I'm first going to admit I have absolutely no formal training in chemistry- I'm strictly a math major :)

But I noticed over a weeklong evaporation operation- The nd sulfate fell out first, in both cases, so much so that it was able to crystallize, and appears nearly Fe free.

Iron sulfate is VERY soluable at ~95*C, the Nd sulfate is not... Am I the only person evaporating the water out slowly at just below boiling, causing the Nd to drop out selectively?

I can post some pictures, I'm working on narrowing down the time spend evaporating- So far, I manage to drop out all the ND sulfates, and then the iron sulfates right down on top of it (but the layering is very distinct).
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[*] posted on 12-11-2011 at 13:39


Quote: Originally posted by Poppy  

Would this dry Fe2O3 work too???




No, it needs to be sufficiently hydrated (‘freshly’ precipitated) to still react with low concentrations of H3O+: Fe(OH)3 + 3 H3O+ === > Fe3+ + 2H2O.

Quote: Originally posted by Wizzard  
But I noticed over a weeklong evaporation operation- The nd sulfate fell out first, in both cases, so much so that it was able to crystallize, and appears nearly Fe free.


Welcome to the murky world of hobby chemistry, mathman!

Do NOT let appearances fool you: at least with the double salt (presumed Nd2(SO4).K2SO4.2H2O) I’ve obtained a nice, pink product which on further scrutiny contained… substantial amounts of iron! There seems to be co-precipitation going on. Whether that is also the case with ‘naked’ Nd2(SO4)4 I cannot vouch for. Test by converting your sulphate to hydroxide by treating with strong alkali. Filter off the precipitate and wash it, then dissolve it in clean (iron free) HCl. Add a bit of 9 % H2O2 to oxidise any Fe2+ to Fe3+. Test visually or with KSCN.

I wonder also about the iron content of Nd oxalate precipitated from iron rich solutions, see Mr Home Scientist higher up and kmno4 in a separate Nd thread...

£$£$£$£$£$£$

Well, here’s my ‘magnet chloride’ after heating and standing overnight:



Plenty of precipitated Fe(H)3 but the dark red-brownish supernatant liquid is the tell tale that plenty of Fe didn’t drop out. pH was less than 1 at that point, so no surprise there.

After that the work became more of a rescue operation than a chemical separation!

The slurry filtered only with great difficulty and adding some freshly precipitated Fe(OH)3 did not bring the pH up enough to precipitate the remaining dissolved iron. I had to resort to manually adjusting the pH to pH >= 4, first with small amounts of 5 M NaOH, the final adjustment with household ammonia. At that point the slurry became much thicker and filtered better. The filtrate contained all the Nd, but still tested slightly positive for Fe3+, mainly due to a bit of colloidal Fe(OH)3 running through the filter I believe. The raw NdCl3 has now been precipitated one last time and will be redissolved in HCl tomorrow.

In conclusion I’d say that separating Nd3+anf Fe3+ by playing to the Fe(OH)3 equilibrium works best for removing relatively small amounts of iron, not to remove iron in concentrations 7 times or larger than that of Nd3+.


[Edited on 13-11-2011 by blogfast25]




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[*] posted on 13-11-2011 at 09:35


Will test once I get home :) It all dissolved in 15cc of very cold water (1-2*C), about .5g of the physically extracted Nd salt (scraped from the bottom- It's much softer than the very hard iron sulfate. Very faint purple color. I'm working on getting the rest out, by watching the evaporative process carefully.

I only was seeing purity because the Nd salt crystallized- I know some purity is needed for this. There were small 2-3mm crystals among the mess on the bottom of the container.
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[*] posted on 13-11-2011 at 09:47


Ok wizzard, what was your 'mother liquor' made from? 'magnet sulphate'? Did you simply dissolve a magnet in H2SO4? Then what do you do next?

Also, when you refer to 'iron sulphate', do you mean ferric or ferrous sulphate?




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[*] posted on 13-11-2011 at 19:32


Mother aqueous solution was 'magnet sulfide', with the nickel COMPLETELY removed manually.

Then left to dissolve, in full, the remaining slush of unknown material and boron removed. (pic 1)

From this, total evaporation with small addition of distilled solution, but still acidic (I did not test pH).

From this, the layered substance in photo 2.

Scraping the purple Neodymium Sulfate into a glass (a small sample amount), dissolved and nearly frozen, then dried on a hot plate over the course of 6 hours, the resulting pure crystals formed- Picture 3 in 4 parts :)

My method seems low yield, but very easy! I prefer manual means- Not so much measurement (but I would like to learn). I'm a scientist for sure, but a chemist I am not!

The largest (centered) crystal cluster is about 7-8mm across.

@Blogfast- Ferrous, Iron(II) of course :)
IMG_20111108_185543.jpg - 214kB IMG_20111111_181507.jpg - 78kB IMG_20111113_221601.jpg - 126kB

[Edited on 11-14-2011 by Wizzard]
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[*] posted on 14-11-2011 at 05:12


So, if I understand well you simply slowly evaporated the ,magnet sulphate' until crystals of Nd2(SO4)3 started to appear? Certainly the Nd sulphate is much less soluble than ferrous sulphate but it's interesting you managed to isolate the minority constituent that way and of seemingly good quality too. AFAIK no one else has applied that approach here yet. Personally I had trouble dissolving the magnet in 50 % H2SO4, not sure why...

I would still urge you to test one crystal for presence of iron. :) If contaminated you could recrystallise for instance, or did you already do that?


[Edited on 14-11-2011 by blogfast25]




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[*] posted on 14-11-2011 at 05:52


Yes, after filtering out the insoluables, the Nd fell out of solution before the Fe did with slow evaporation at 90-95*c.

I will test for iron once I get my hands on some good hydrogen peroxide... My lab chemical supply is still expanding.

I have not recrystallized a second time, only a first. I could do it a second, but I fear there is little enough already! I'll wait until the larger batch I have going right now finishes.
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[*] posted on 14-11-2011 at 07:03


Quote: Originally posted by Wizzard  
Yes, after filtering out the insoluables, the Nd fell out of solution before the Fe did with slow evaporation at 90-95*c.

I will test for iron once I get my hands on some good hydrogen peroxide... My lab chemical supply is still expanding.



It's quite amazing you got crystals rather than microcrystals (powder). You must have gotten a bit lucky there, I'm guessing; right water-to-salt ratio and all that... It's also surprising your that your ferrous sulphate wasn't more oxidised by air and heat: it then has a tendency to drop out.

3 or 9 % H2O2 from a pharmacy will do. It's dirt cheap.

[Edited on 14-11-2011 by blogfast25]




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[*] posted on 14-11-2011 at 07:59


Well I know from experience that ferrous sulfate oxidizes with contact to air in solution, unless kept mildly acidic... I've grown 1cm crystals in 2cm tubes :) I'm no chemist, but I do know a thing or two about keeping one's ingredients pure- I've been growing crystals of various salts since I was in high school.

And I was unaware that pharmacy grade H2O2 is good enough! Off I go!

Attached good macro pic, with scale.

macro.png - 773kB

[Edited on 11-14-2011 by Wizzard]
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[*] posted on 14-11-2011 at 08:32


Very nice crystals indeed: they look good enough to eat! ;)

Do you recall:

* weight of the magnet
* amount and concentration of the sulphuric acid used?

Did you get any early precipitation (of Nd2(SO4)3) during dissolving?

Acidity slows the oxidation of Fe (II) down, that's very true. But you should have seen my acidic stockpile of FeCl2 crystals after a couple of weeks: it's a brown mass of Fe(OH)3, ferric oxychlorides and what not: I bet it would test negative for Fe (II)! Just for standing in open air...




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[*] posted on 14-11-2011 at 09:43


The approximate weight of the magnets, before breaking up and removing the plating, was about 200g- I'd say 125-150g of magnet material went in.

I then added about equal parts distilled water and 98% sulfuric acid, maybe 50ccs total. When the boiling would stop, I would add more conc. sulfuric acid or water (4ccs at a time), depending on what the solution wanted, and I did this many times. It sometimes would run out of acid (and a light tinge of yellow iron hydr/oxide would form on the meniscus) or it would cease when it ran out of water. When it stopped and only small bubbles would come out of the insoluable slush at the bottom for some hours (and sometimes only when shaken slightly), then it was time to filter.

No visible precipitation occurred as the solution was made, but there could be some in the insoluable mess. It did sometimes have a tinge of purple. I am working on extracting what remaining material I can from it.

The solution was kept at room temperature while dissolving- Only in the end did I heat it (to wrap up any action of the sulfuric acid, at least 90*C) before I froze it (about 5*C), then filtered it cold (maximizing output of the Nd sulfate).
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[*] posted on 14-11-2011 at 10:07


Quote: Originally posted by blogfast25  
Fenton's reagent doesn't involve Fe (IV).

Really? I thought it did, after reading this:
http://woelen.homescience.net/science/chem/solutions/fe.html
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[*] posted on 14-11-2011 at 10:35


Quote: Originally posted by barley81  

Really? I thought it did, after reading this:
http://woelen.homescience.net/science/chem/solutions/fe.html


I think it's a matter of debate, TBH. If Fe IV really does exist it must be very unstable.

@Wizzard:

I see. Well, maybe next time I'll try again with fully peeled magnets and a stoichiometrically determined amount of H2SO4...



[Edited on 14-11-2011 by blogfast25]




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[*] posted on 14-11-2011 at 11:06


I'll run another batch, fully documented :)

I have stacks of old HDDs and all the motivation I need for beautiful, hexagonal air semi-stable (left one on the paper overnight- no water loss!!) color changing crystals of an element I have very little of in my collection.
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[*] posted on 14-11-2011 at 12:37


Lucky mathman! I asked a computer repair shop to keep any duff HDDs for me (for a token price) but it turns out they 'forgot'. Barstools.

The crystals, being of a compound that's not particularly highly soluble nor insoluble, should be neither particularly deliquiescent nor efflorescent, IMHO... My NdCl3 on the other hand deliquiesced like mad: very hygroscopic.




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[*] posted on 14-11-2011 at 12:50


Wizzard:
I would assume you just filtered the liquor out of the muddy magnet solution: this little ammount of liquor justifies the nicely dissolved Nd2(SO4)3 in there, and very small yield for 125g?
Also we must point out your evaporation method is a very good way for 1 stepped refining of neodimium sulphate., good job.
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[*] posted on 14-11-2011 at 13:04


Thanks! Very small yield was due to my physical method of separating the Nd salt from the Fe salt- I merely scraped the easiest-to-get-at soft/wet and purple neodymium sulfate crystal from the hard, green ferrous sulfate.

Low yield, yes, but I'm still extracting the salts from that aqueous solution- This was just a quick and dirty trial. I'd also still say there's 50% more in my evaporation dish- There was a good quantity of smaller crystals, but I do need to remove the film which grew on the edges before I reattempt this (and recycle the dried and extracted but not removed Nd sulfate).

Partial evaporation yielded the following crystals- Such a lovely color! Recrystallizing tonite- It doesn't take long. Note the green tinge in the center mass, and along the edges- This is the start of when the iron sulfate crystallizes- Time to stop evaporating!

[Edited on 11-14-2011 by Wizzard]

IMG_20111114_161104.jpg - 123kB
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[*] posted on 15-11-2011 at 06:27


Wizzard:

Although I had come to believe that your crystals must be high purity and essentially free of iron, now I’m not so sure anymore or at least a little befuddled. A spammer inadvertently dragged up this old post on neodymium sulphate:

http://www.sciencemadness.org/talk/viewthread.php?tid=12934#...

(go to the last post by kmno4)

The colour obtained by you seems highly unusual.

I will now cross-post your results into that thread.




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[*] posted on 15-11-2011 at 07:36


My lighting conditions are a normal 60W incandescent, and big-bulb commercial flourescent lighting. Perhaps I'll get a sunlight shot today- That will be nice for comparison.

My are certainly a bit more pink/red than those purple/pink ones you link to- But I will say I cant wait to grow a single crystal that size :)

[Edited on 11-15-2011 by Wizzard]
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[*] posted on 15-11-2011 at 08:19


On the the leftover liquor:
After separating in part the neodimium from iron from a magnet solution via the sulphate method, a leftover supernatant was obteined which seemed to contain, by its color, initially at least a big portion of Fe III sulphate, Fe II sulphate, unnoticeable by eye, and some neodimium as well.
Theoretical obtention of Nd2(SO4)3.8H2O would be slightly round 50g, but as there was obteined just 44g (oops, thats really close to the goal, and because some 5g were lost propositally =P) I am inclined to believe some neodimium was left in the supernatant solution (at least 6g) as well as some iron is present in my final powder.
The main reason is that the leftover supernatant gradually, and very slowly, formed crystals which took much longer than conventional ferrous sulphate crystals to form (6 hours for a supersaturated solution of FeSO4 and one week to badly precipitate some of the strange liquor).
Here you can see the leftover after about 2 days:


filtrate liquid crystalization.JPG - 27kB filtrate liquid crys closeup.JPG - 25kB

Here you can see the leftover solution after about 1 week:



filtrate liquid after.JPG - 25kB
Those crystals were filtered and dissolved again to see what happens: the result was a slightly pink solution from those green crystals!!!!





green dissolved as pink.JPG - 10kB green dissolved as pink big.JPG - 11kB

That after decanting (seeing pink depends on your mood on that day...)
green dissolved as pink decanted.JPG - 12kB
The green crystals initially become decoloured when washed
then gives the solution above.

So there must be double salts of Nd, Fe in the solution which modifies how the precipitation of pure ferrous sulphate crystals would do, and my obteined Nd salt powder is highly contamined with Pr or Fe in order of having its yellow color, even after it was thorughly washed with hot water when it was prepared.

And this is the Nd sulphate under flash and daylight, theres no real difference.
Nd salt daylight.JPG - 11kB Nd salt powder flash.JPG - 15kB

[Edited on 11-15-2011 by Poppy]
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[*] posted on 15-11-2011 at 08:24


Quote: Originally posted by Poppy  
On the the leftover liquor:
After separating in part the neodimium from iron from a magnet solution via the sulphate method, a leftover supernatant was obteined which seemed to contain, by its color, initially at least a big portion of Fe III sulphate, Fe II sulphate, unnoticeable by eye, and some neodimium as well.
Theoretical obtention of Nd2(SO4)3.9H2O would be slightly round 50g, but as there was obteined just 44g (oops, thats really close to the goal, and because some 5g were lost propositally =P)


Can you describe what you did slightly more in detail?
Also, literature describes an octahydrate (.8H2O), not a nonahydrate (.9H2O)...

Edit:

Dang!! Quite a bit of confusion on the colour of neodymium sulphate: higher up in this thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=14145&...

Quote: Originally posted by DerAlte  
All references I have seen say Nd sulphate is red (my emph.) and Sm is light yellow, Ce white (colourless); and these were the colors I got. All these sulphates are very difficult to dissolve at 0C, the most soluble point. Once dissolved, heating to 100C usually will precipitate them. Better crystals can be obtained by slow evaporation.

Regards, Der Alte


This photobucket entry which shows crystals very much like Wizzard’s:

http://media.photobucket.com/image/neodymium%20sulphate%20cr...

And here, slightly more pinkish:

http://www.metall.com.cn/ndso.htm


[Edited on 15-11-2011 by blogfast25]

[Edited on 15-11-2011 by blogfast25]




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[*] posted on 15-11-2011 at 11:59


It's quite possible there's some other Lanthanide impurity involved there- Neodymium magnets dont need absolute purity, as far as the Nd is concerned, I think (and was this also mentioned previously in this thread?).
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[*] posted on 15-11-2011 at 12:27


Quote: Originally posted by Wizzard  
It's quite possible there's some other Lanthanide impurity involved there- Neodymium magnets dont need absolute purity, as far as the Nd is concerned, I think (and was this also mentioned previously in this thread?).


Actually I believe it was praseodymium that was mentioned by kmno4. Absolute purity of Nd for magnets is probably not a prerequisite but to affect colour you'd need substantial contamination (I think).

Other things that might affect colour could be:

* light (UV content): already accounted for
* powder (microcrystaline) v. macro crystals
* variability in hydration (.xH2O). It's well known the lower hydrates (generally speaking - not spec. to Nd) tend to form at higher temps. and vice versa.

Looking at the body of evidence, yours seem to have the 'right' colour though...




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[*] posted on 15-11-2011 at 13:12


What I'll experiment with next will be how much alcohol I can add to the water solution of high-Nd/low-Fe and plot the soluability tables with different mixtures of alcohol/water down to -10*C (the limits of my scientific "Wal-Mart" mini-fridge).

Maybe with some alcohol, I can lower temperatures and dissolve more Nd sulfate, and drop out more Fe sulfate.
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