Sciencemadness Discussion Board

The trouble with neodymium...

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blogfast25 - 27-10-2014 at 14:09

MrHS:

I see now I've missed part of the action (your posts of 26/10).

Looks like a failure alright. Most likely explanation must indeed be oxygen.

I'm looking at that 'nested design' with argon blanket, with a small steel or iron crucible sitting inside a sealed off (bottom) copper piece of pipe (1 1/2" to 2"). I don't have a furnace, so that means relying heavily on the anticipated exotherm and full blast propane from beneath.
Quote: Originally posted by MrHomeScientist  

Come to think of it, now that I have this furnace setup, is magnesium reduction a possibility again? We scrapped that because it didn't produce enough heat of reaction, but if I heat the whole thing externally that won't be a problem. Then I wouldn't need inert atmosphere either, which would simplify greatly. In fact I think I still have the Mg + NdF3 mixture that failed from my first attempt...


Remember that the low value of ΔG is related to the equilibrium constant of the reaction, via Nernst. For ΔG = 0 basically K = 1.

That means conversion will be very poor and you'll obtain a mix of metals, possibly an alloy.

It could be worth revisiting my reference to Brauer's methods for REEs. I had referenced them in the La thread before it got deleted by the spam robot.

Brauer's part II suggests a method using anh. RECl<sub>3</sub> by reduction with calcium, boosted with sulphur, iodine or chlorate.

Magnesium seems generally useless for reduction of REE halides.

[Edited on 27-10-2014 by blogfast25]

Brain&Force - 27-10-2014 at 21:10

The reference to Brauer is in the new La thread. Download it if you haven't; it's a must read.

I'm wondering, if you flush out O2 with propane, will that work to prevent the lithium from reacting or will it just introduce side reactions?

blogfast25 - 28-10-2014 at 05:28

Quote: Originally posted by Brain&Force  
I'm wondering, if you flush out O2 with propane, will that work to prevent the lithium from reacting or will it just introduce side reactions?


KA-BOOOOMMM!

gdflp - 28-10-2014 at 07:38

How about using strontium metal for the reduction of RE fluorides? It's a lot cheaper than calcium or lithium($80.75/5.5lbs, Free shipping from GalliumSource), and it's more electronegative than calcium. According to WolframAlpha, it's exothermic, -335kJ/mol. In addition, it shouldn't have the problem of reacting with the crucible. It's melting point is kind of high though at 777°C, that might cause problems.

[Edited on 10-28-2014 by gdflp]

blogfast25 - 28-10-2014 at 09:32

Quote: Originally posted by gdflp  
According to WolframAlpha, it's exothermic, -335kJ/mol.


Acc. Wolfram Alpha, what is exothermic? Nice number but what does it represent?

Texium - 28-10-2014 at 10:03

Perhaps he's referring to the reduction of neodymium fluoride with strontium?

blogfast25 - 28-10-2014 at 10:09

Quote: Originally posted by zts16  
Perhaps he's referring to the reduction of neodymium fluoride with strontium?


If so, he's miscalculated, it's only about - 167 kJ/mol.

Brain&Force - 28-10-2014 at 10:10

http://www.sciencemadness.org/talk/viewthread.php?tid=32521&...

Strontium is definitely a cheaper choice, and it will probably work to reduce the Nd but it probably won't work without flushing out O2.

In Brauer there are references to reduction of anhydrous LnCl3 in alcoholic solution. It appears to be possible to make the anhydrous chloride by heating with ammonium chloride, but the product will be contaminated with ammonium chloride afterwards. Will this affect the reduction in any way?

gdflp - 28-10-2014 at 11:29

Sorry I wasn't more specific. Yes, I was talking about the reduction of neodymium fluoride with strontium, and my number was off because I calculated using the equation 2NdF3 + 3Sr --> 2Nd + 3SrF2 and I didn't divide by 2.

Why would the product be contaminated with ammonium chloride, I thought the whole point of it was to allow it to decompose into NH3 and HCl, effectively heating the RE chloride in a stream of HCl. Anhydrous RE chlorides should be stable to strong heating, so after the water has been removed the mixture could be heated to sublime away any excess ammonium chloride.

blogfast25 - 28-10-2014 at 12:29

Quote: Originally posted by Brain&Force  
http://www.sciencemadness.org/talk/viewthread.php?tid=32521&...


In Brauer there are references to reduction of anhydrous LnCl3 in alcoholic solution. It appears to be possible to make the anhydrous chloride by heating with ammonium chloride, but the product will be contaminated with ammonium chloride afterwards. Will this affect the reduction in any way?


I see no reason why it should. It should blow off nicely.

My own singular experience with NdCl3 hydrate/NH4Cl wasn't very positive even though previously I got it to work well for MnCl2 hydrate.

The mixture need to be very intimate (very finely ground) and starting from quite dry hydrate. I think partial vacuum probably helps things too.

Removing those very last traces of NH4Cl may be tricky, as indicated in the Brauer entry.

Brain&Force - 29-10-2014 at 15:06

Quote: Originally posted by blogfast25  
Quote: Originally posted by Brain&Force  
I'm wondering, if you flush out O2 with propane, will that work to prevent the lithium from reacting or will it just introduce side reactions?


KA-BOOOOMMM!


I'm not entirely sure what would cause the explosion after re-reading this. Would the lithium react with the propane?

I specified that the oxygen would be flushed out, which should prevent an explosion from occurring (in theory).

Dan Vizine - 29-10-2014 at 15:17

Molten lithium will react with it at higher temperatures.

Brain&Force - 29-10-2014 at 15:22

Yeah, that's what I thought. Molten lithium doesn't particularly like concrete, either.

Dan Vizine - 29-10-2014 at 15:35

And yet it's stable in the molten state in pure dry oxygen.

I love the seeming contradictions in chemistry. Like this one:

A sample of Na at 200 C will spontaneously ignite in the air. A sample of K at 280 C won't. Yet we all know that K is more "reactive" than Na. So..??

[Edited on 29-10-2014 by Dan Vizine]

blogfast25 - 30-10-2014 at 05:38

Quote: Originally posted by Brain&Force  

I'm not entirely sure what would cause the explosion after re-reading this. Would the lithium react with the propane?

I specified that the oxygen would be flushed out, which should prevent an explosion from occurring (in theory).


I was specifically referring to the flammability of propane as cause for great danger with open flames, nothing else.

blogfast25 - 30-10-2014 at 05:41

Quote: Originally posted by Dan Vizine  
A sample of Na at 200 C will spontaneously ignite in the air. A sample of K at 280 C won't. Yet we all know that K is more "reactive" than Na. So..??



Assuming these data points are factually correct then it may be related to volatility or activation energy.

You're right to use quote marks: "reactivity" is a poorly defined and elastic term.

Nature defies our simple models quite a lot.

Praxichys - 30-10-2014 at 05:47

Quote: Originally posted by Dan Vizine  
A sample of Na at 200 C will spontaneously ignite in the air. A sample of K at 280 C won't. Yet we all know that K is more "reactive" than Na. So..??


I wonder if this has something to do with passivation?

Aluminum is "more reactive" than iron, but iron reacts with the air faster. Perhaps the oxide properties of K are better suited to form a film on the liquid than that of Na?

blogfast25 - 30-10-2014 at 10:46

Quote: Originally posted by Praxichys  
Quote: Originally posted by Dan Vizine  
A sample of Na at 200 C will spontaneously ignite in the air. A sample of K at 280 C won't. Yet we all know that K is more "reactive" than Na. So..??


I wonder if this has something to do with passivation?

Aluminum is "more reactive" than iron, but iron reacts with the air faster. Perhaps the oxide properties of K are better suited to form a film on the liquid than that of Na?


It's possible but the comparison with Al/Fe is flawed because at these temperatures both Na and K are liquids, with constant surface renewal.

Brain&Force - 31-10-2014 at 18:10

Quote: Originally posted by blogfast25  
Brauer's part II suggests a method using anh. RECl<sub>3</sub> by reduction with calcium, boosted with sulphur, iodine or chlorate.


Sulfur may not be a good idea for thermiting. The main reason is that many lanthanide monochalcogenides tend to be stable. With neodymium this may not be a significant factor, as it's difficult to reduce to a 2+ state, but it will likely be signifiant if you intend to repeat the work with samarium, ytterbium, or europium-containing mixtures.

blogfast25 - 1-11-2014 at 06:45

Quote: Originally posted by Brain&Force  
Sulfur may not be a good idea for thermiting. The main reason is that many lanthanide monochalcogenides tend to be stable. With neodymium this may not be a significant factor, as it's difficult to reduce to a 2+ state, but it will likely be signifiant if you intend to repeat the work with samarium, ytterbium, or europium-containing mixtures.


CaS is likely to have a much more negative enthalpy of formation than REE sulphides. Properly formulated mixtures would then not lead to the latter forming, only CaS.

Sulphur would not be my first choice of booster system though.

[Edited on 1-11-2014 by blogfast25]

MrHomeScientist - 5-11-2014 at 09:07

Yesterday I purchased a 40 cubic foot cylinder of pure argon, along with a regulator and some fittings, as step 1 in creating an inert atmosphere for this reaction. Progress!

The graphic I posted on the previous page was intended as a 'nested crucible' which would be lowered into my small furnace and be heated externally via propane torch. Instead of that, what about simply blowing a stream of argon over the mouth of the crucible containing the reaction? That would provide a blanket of Ar over the reaction without needing to make any complex (for me) apparatus. I'd make the gas outlet into something like a wing top attachment for bunsen burners, like this one: http://www.enasco.com/product/SB06947M

The Volatile Chemist - 5-11-2014 at 13:23

Quote: Originally posted by MrHomeScientist  
Yesterday I purchased a 40 cubic foot cylinder of pure argon, along with a regulator and some fittings, as step 1 in creating an inert atmosphere for this reaction. Progress!

The graphic I posted on the previous page was intended as a 'nested crucible' which would be lowered into my small furnace and be heated externally via propane torch. Instead of that, what about simply blowing a stream of argon over the mouth of the crucible containing the reaction? That would provide a blanket of Ar over the reaction without needing to make any complex (for me) apparatus. I'd make the gas outlet into something like a wing top attachment for Bunsen burners, like this one: http://www.enasco.com/product/SB06947M

In my own uneducated opinion, it sounds good; It doesn't seem like you would be able to run the reaction for long, though, without running out of the stuff (Argon)? also, wouldn't there be a pocket of air underneath the stream that wouldn't necessarily be replaced by the argon?

blogfast25 - 5-11-2014 at 13:45

Mr HS:

My plan (if I ever get round to it) would be to hermetically seal the bottom of the outer crucible, then have a copper or steel pipe reach down into it (to near the bottom), to pipe the Ar into the outer crucible. At first give it all a good flux, then switch to lower flow rate. Load the reactants into the inner crucible while it's submersed into about half the depth of the outer crucible, then lower it to the bottom and start heating.

This should give a reasonably O2/N2 free atmosphere w/o the need to seal everything off perfectly.

Ar is denser than air, so that should help (a bit).

[Edited on 5-11-2014 by blogfast25]

MrHomeScientist - 1-3-2015 at 11:58

My progress on this experiment continues at glacial pace! I finally got around to ordering and fabricating my inert atmosphere crucible, made of black iron pipe and fittings. Here it is, and I've dubbed it "The Gavel". Or if you prefer, "The Croquet Mallet".

gavel 1.jpg - 352kB

Argon enters from the 1/4" pipe on the right. The 2" pipe on the left is the reaction chamber, inside which a small crucible which contains the reactants will rest. A hole in the top of the chamber allows gas to exit, and the whole assembly will be heated from below (probably). Here's another view:

gavel 2.jpg - 283kB

You can see the exhaust hole in the lid, and the space inside where the small inner crucible will rest. I may need some sort of heat transfer media inside to help heat the small crucible, something like stainless steel ball bearings.

The latest problem is that this whole thing just barely will not fit into my mini propane furnace. So I need to find a way to heat this thing up to the requisite several hundred degrees. Maybe wrap the whole thing in insulating kaowool?

blogfast25 - 1-3-2015 at 12:12

Looking good. Insulating will help, of course.

Remind me again which REE you're gunning for? REF<sub>3</sub> + 3 Li ?


[Edited on 1-3-2015 by blogfast25]

MrHomeScientist - 1-3-2015 at 13:35

Correct, it's neodymium:

NdF<sub>3</sub> + 3 Li --> Nd + 3LiF

DalisAndy - 8-5-2015 at 13:11

Whats Nd reactivity compared to Cu? If not, could someone try something for me?

[Edited on 8-5-2015 by DalisAndy]

DalisAndy - 8-5-2015 at 13:14

Quote: Originally posted by MrHomeScientist  
Correct, it's neodymium:

NdF<sub>3</sub> + 3 Li --> Nd + 3LiF

The only problem I see is that NdF3 is a ionic compound. And would need to be dissolved in something, I order to react. If I'm correct

gdflp - 8-5-2015 at 13:25

Quote: Originally posted by DalisAndy  
Quote: Originally posted by MrHomeScientist  
Correct, it's neodymium:

NdF<sub>3</sub> + 3 Li --> Nd + 3LiF

The only problem I see is that NdF3 is a ionic compound. And would need to be dissolved in something, I order to react. If I'm correct

Not necessarily. At the temperatures being discussed, the lithium is molten, and the liquid-solid reaction interface is fine for the reaction.
Quote: Originally posted by DalisAndy  
Whats Nd reactivity compared to Cu? If not, could someone try something for me?

Neodymium is significantly more reactive than copper, I would say that it's much closer to magnesium in reactivity.

DalisAndy - 8-5-2015 at 13:32

Quote: Originally posted by gdflp  

Neodymium is significantly more reactive than copper, I would say that it's much closer to magnesium in reactivity.

So Mg could be used to precipitate Nd after decomposing the FeSO4 and separated the Boron?

gdflp - 8-5-2015 at 13:36

I highly doubt it. You would need a solvent that wouldn't react with the magnesium or neodymium, would dissolve and ionize both of their salts, and even then I don't think that the reaction is favorable, it would also leave a metal heavily contaminated with magnesium. You would also need the neodymium salt to be anhydrous, which is a difficult task in it's own regard.

[Edited on 5-8-2015 by gdflp]

blogfast25 - 8-5-2015 at 13:43

Quote: Originally posted by DalisAndy  
Whats Nd reactivity compared to Cu? If not, could someone try something for me?

[Edited on 8-5-2015 by DalisAndy]


DalisAndy:

ALL RE elements are very electropositive, more or less on a par with Al.

Some here (like me) believe reduction of NdF3 with Li might just about be possible. MrHomeScientist is setting things up to try that.

Otherwise it's electrolysis of anhydrous RE trichlorides (industrially).

FORGET about aqueous reductions with WHATEVER: these elements react immediately with water. They cannot be electroplated from water.

Possibly electrodeposition from deep eutectic solvents or aprotic solvents but that's firmly in the future, it at all possible even.

You need to read this long thread: no gain without pain! And people get very tired of answering the same questions over and over.


[Edited on 8-5-2015 by blogfast25]

DalisAndy - 8-5-2015 at 13:48

Quote: Originally posted by blogfast25  
Quote: Originally posted by DalisAndy  
Whats Nd reactivity compared to Cu? If not, could someone try something for me?

[Edited on 8-5-2015 by DalisAndy]


DalisAndy:

ALL RE elements are very electropositive, more or less on a par with Al.

Some here (like me) believe reduction of NdF3 with Li might just about be possible. MrHomeScientist is setting things up to try that.

Otherwise it's electrolysis of anhydrous RE trichlorides (industrially).

FORGET about aqueous reductions with WHATEVER: these elements react immediately with water. They cannot be electroplated from water.

Possibly electrodeposition from deep eutectic solvents but that's firmly in the future, it at all possible even.

How whould make NdF3? Short of using F2 gas or roach poisin

gdflp - 8-5-2015 at 13:58

Quote: Originally posted by DalisAndy  

How whould make NdF3? Short of using F2 gas or roach poisin

As blogfast said, read the thread. All of the information is here. You need to do some of your own homework, many methods have been discussed in this thread already.

[Edited on 5-8-2015 by gdflp]

blogfast25 - 8-5-2015 at 13:58

Quote: Originally posted by DalisAndy  
How whould make NdF3? Short of using F2 gas or roach poisin


Full preparation is in the thread.

2 NdCl3(aq) + 3 NH4HF2(aq) === > 2 NdF3(s) + 3 NH4Cl(aq) + 3 HCl(aq)

NdF3 is insoluble. Filter, dry and calcine.

DalisAndy - 8-5-2015 at 14:30

Ok thank you. A weekend of fun planned I sense. Oh I have discovered an odd occurrence. When I brought one of my hard drive magnets near the crushed motor magnet I have, when I accidently dropped it in to the motor magnet, it made the hard drive one stronger. Sticking it to its sister magnet , I can't separate them. Are there any Nd super alloys?

[Edited on 8-5-2015 by DalisAndy]

blogfast25 - 8-5-2015 at 14:54

If the magnets are 'large', separating two of them can be hard. There do exist several grades though. Composition is the same though, I think...

DalisAndy - 10-5-2015 at 11:07

Ok I just want conformation my math is correct for these reactions

Nd2Fe14B + 17CuSO4 -> 17Cu + 14FeSO4 + Nd2(SO4)3 + B

14 FeSO4 + heat -> 14FeO + 7O2 + 14SO2

Nd2(SO4)3 + 6NH3F -> 2NdF3 + 3(NH3)2SO4

then proceed with lithium replacement

blogfast25 - 10-5-2015 at 12:16

Quote: Originally posted by DalisAndy  
Ok I just want conformation my math is correct for these reactions

Nd2Fe14B + 17CuSO4 -> 17Cu + 14FeSO4 + Nd2(SO4)3 + B

14 FeSO4 + heat -> 14FeO + 7O2 + 14SO2

Nd2(SO4)3 + 6NH3F -> 2NdF3 + 3(NH3)2SO4

then proceed with lithium replacement


Copper sulphate as solvent? Hmmm... good luck with that! (It'll probably work but why waste CuSO4 on dissolving Nd magnets?)

Ferrous sulphate will oxidise to ferric oxide.

[Edited on 10-5-2015 by blogfast25]

DalisAndy - 10-5-2015 at 12:20

Quote: Originally posted by blogfast25  
Quote: Originally posted by DalisAndy  
Ok I just want conformation my math is correct for these reactions

Nd2Fe14B + 17CuSO4 -> 17Cu + 14FeSO4 + Nd2(SO4)3 + B

14 FeSO4 + heat -> 14FeO + 7O2 + 14SO2

Nd2(SO4)3 + 6NH3F -> 2NdF3 + 3(NH3)2SO4

then proceed with lithium replacement


Copper sulphate as solvent? Hmmm... good luck with that! (It'll probably work but why waste CuSO4 on dissolving Nd magnets?)

Ferrous sulphate will oxidise to ferric oxide.

[Edited on 10-5-2015 by blogfast25]

I was going to use copper sulfate because of easy of access (I'm a student and not 18 yet). But I can get sulfuric acid. also can ferric oxide be decomposed to FeO?

[Edited on 10-5-2015 by DalisAndy]

blogfast25 - 10-5-2015 at 12:24

Quote: Originally posted by DalisAndy  

I was going to use copper sulfate because of easy of acess (I'm a student and not 18 yet). But I can get sulfuric acid. Why wont copper sulfate work?


I said it would 'probably work'. HCl or H2SO4 are simply easier/cheaper. With CuSO4 you'll definitely need to grind it quite fine. Not a big problem.

Use 10 % CuSO4 than needed: this way you won't have any free Cu<sup>2+</sup> in solution, which would then have to be separated out again.

[Edited on 10-5-2015 by blogfast25]

DalisAndy - 10-5-2015 at 13:36

Quote: Originally posted by blogfast25  
Quote: Originally posted by DalisAndy  

I was going to use copper sulfate because of easy of acess (I'm a student and not 18 yet). But I can get sulfuric acid. Why wont copper sulfate work?


I said it would 'probably work'. HCl or H2SO4 are simply easier/cheaper. With CuSO4 you'll definitely need to grind it quite fine. Not a big problem.

Use 10 % CuSO4 than needed: this way you won't have any free Cu<sup>2+</sup> in solution, which would then have to be separated out again.

[Edited on 10-5-2015 by blogfast25]


And use water for the other 90%? I know Nd will react with water

blogfast25 - 10-5-2015 at 13:59

Quote: Originally posted by DalisAndy  

And use water for the other 90%? I know Nd will react with water


I meant: calculate the amount of CuSO4 needed by your equation. Then use 10 % less than the calculated value. A small amount of magnet will not react but all your CuSO4 will. This way there's no free CuSO4 in the resulting solution.

elementcollector1 - 10-5-2015 at 14:49

Just thought I'd share a quick note, as I restarted my own Nd experiments: I just successfully prepared pure neodymium sulfate solution using the oxalate method.

I started with some freshly made 'magnet chloride', and added concentrated oxalic acid solution. This gave me a bright yellow precipitate, indicating iron was still present. In an attempt to start over, I ran fresh HCl through some of the precipitate while it was still in the funnel, only to encounter a curious effect: The iron oxalate dissolved away, giving an emerald green solution, but the neodymium didn't. A single wash with water gave me a pale lavender solid, which easily dissolved in a 50% mix of freshly prepared (hot) H2SO4 to yield a pinkish solution (though it looks orange in incandescent light). Right now, it's in the freezer so that the oxalic acid crystallizes out.

blogfast25 - 10-5-2015 at 15:26

ec1:


Looks like you've created a variant of the oxalate method. In my version the iron is oxidised with peroxide to ferric. With oxalate the ferric forms green and soluble trisoxalato ferrate(III), while the Nd forms insoluble Nd oxalale.

In your case you formed Fe(II)Ox and Nd oxalate. The former dissolves into dilute HCl (well known, that), the latter does not. But I'm quite surprised that Nd oxalate dissolved in hot 50 % H2SO4. RE oxalates are extremely insoluble. Interesting but requiring corroboration, I feel.

[Edited on 10-5-2015 by blogfast25]

elementcollector1 - 10-5-2015 at 19:15

Quote: Originally posted by blogfast25  
ec1:


Looks like you've created a variant of the oxalate method. In my version the iron is oxidised with peroxide to ferric. With oxalate the ferric forms green and soluble trisoxalato ferrate(III), while the Nd forms insoluble Nd oxalale.

In your case you formed Fe(II)Ox and Nd oxalate. The former dissolves into dilute HCl (well known, that), the latter does not. But I'm quite surprised that Nd oxalate dissolved in hot 50 % H2SO4. RE oxalates are extremely insoluble. Interesting but requiring corroboration, I feel.

[Edited on 10-5-2015 by blogfast25]


The oxalate was freshly prepared, and had never been dried - that may have had something to do with it. I remember the first time I tried this, after calcining - nothing in Hell, Heaven or Earth could have dissolved that powder.

Fe(III), not Fe(II) - so that's what I was forgetting! I was wondering why it was so much more difficult to do now. Ah well, nothing for it but to finish off this batch and convert to NdCl3 hydrate.

What to do with all this Fe(II)-contaminated oxalic acid? Currently I have 1.5 gallons of the stuff just laying around... I suppose I should recrystallize it and put it back into its container.

Does a solution of neodymium sulfate normally look orange in artificial light and pink in natural light?

blogfast25 - 11-5-2015 at 05:16

Quote: Originally posted by elementcollector1  
[What to do with all this Fe(II)-contaminated oxalic acid? Currently I have 1.5 gallons of the stuff just laying around... I suppose I should recrystallize it and put it back into its container.

Does a solution of neodymium sulfate normally look orange in artificial light and pink in natural light?


Neutralise, so Fe(OH)2 drops out, filter. Evaporate water to get sodium or potassium oxalate (depending on base used). Quite a job on 1.5 gallons though!

Edit: it might even be easier to simply boil it down. The Fe(II)Ox should crystallise out. Remove it, then boil to dryness to obtain crude oxalic acid. It's easy to recrystallize, BTW.

Nd(III) looks different under different lighting:

http://en.wikipedia.org/wiki/Neodymium(III)_chloride

Those three 4f electrons are the cause of these colour variations.

[Edited on 11-5-2015 by blogfast25]

elementcollector1 - 12-5-2015 at 10:42

Yep, some oxalic acid crystals crashed out upon cooling overnight. They look neat - sort of acicular with an odd square shape.

Anyway, as it turns out, the oxalate didn't dissolve in sulfuric acid the second time around, and something precipitated out of the first one. I think the first one was just a suspension, and had me fooled. Anyway, I now have a bright pink, wet solid - what can I use to dissolve this? Nitric acid?

So close... yet so far.

MrHomeScientist - 15-5-2015 at 05:43

Quote: Originally posted by DalisAndy  
Ok I just want conformation my math is correct for these reactions

Nd2Fe14B + 17CuSO4 -> 17Cu + 14FeSO4 + Nd2(SO4)3 + B

14 FeSO4 + heat -> 14FeO + 7O2 + 14SO2

Nd2(SO4)3 + 6NH3F -> 2NdF3 + 3(NH3)2SO4

then proceed with lithium replacement


Copper sulfate is an interesting idea and like blogfast said it may work, but an acid is your best bet.

I also wanted to point out / remind you that any soluble fluoride salts, like NH3F, are horrifically dangerous. They can and will kill you with a very small spill on your skin. That's not an exaggeration. Seriously not something for a beginner / young person to handle. I highly advise against that path unless you have a professional chemist to help you or a professional lab space. Have calcium gluconate on hand as a first response measure in case of spills (but still immediately go to the hospital).

An interesting alternative I investigated (briefly) is electroplating from a deep eutectic solvent (DES). Supposedly the choline chloride - urea system is capable of dissolving a wide variety of metal oxides, and being anhydrous makes it a promising candidate for plating out reactive metals. I made the DES but never tried dissolving anything in it. It's much safer, since you only need Nd-oxide, but also more speculative.


Edit:
EC1: You could calcine to decompose it to the oxide, but I'm not sure that that leaves you any better off than you are now. Possible use in the DES route?

[Edited on 5-15-2015 by MrHomeScientist]

blogfast25 - 15-5-2015 at 06:04

Quote: Originally posted by elementcollector1  
Anyway, I now have a bright pink, wet solid - what can I use to dissolve this? Nitric acid?


This has to be one of the two forms of neodymium sulphate.

Nitric acid is useless here.

Convert it to Nd(OH)3 with ammonia solution, simmer a bit to complete conversion. Then filter off and wash carefully. The pure hydroxide can then be dissolved in the acid of your choice.

DalisAndy - 26-5-2015 at 12:38

Could one use permagnetic acid, and/or carbonatic acid? And why would Oxalic acid easier? I'm just curious

blogfast25 - 26-5-2015 at 13:02

Quote: Originally posted by DalisAndy  
Could one use permagnetic acid, and/or carbonatic acid? And why would Oxalic acid easier? I'm just curious


'permagnetic acid'? That's either an Error 101 or a Googlewhack. :D

You mean permanganetic acid. "HMnO<sub>4</sub>". Doesn't REALLY exist. Potassium permanganate + sulphuric acid would work but NO ONE would do it like that.

You mean carbonic acid. That's soda water, basically. Do YOU think that would work? (Hint: NO!)

The oxalates of Nd and Fe are basically insoluble. It's a recipe for stalled reaction/big mess.

Go down the hardware store and get some 'muriatic acid' (HCl) or toilet drainer (H2SO4).

DalisAndy - 26-5-2015 at 13:36

Quote: Originally posted by blogfast25  
Quote: Originally posted by DalisAndy  
Could one use permagnetic acid, and/or carbonatic acid? And why would Oxalic acid easier? I'm just curious


'permagnetic acid'? That's either an Error 101 or a Googlewhack. :D

You mean permanganetic acid. "HMnO<sub>4</sub>". Doesn't REALLY exist. Potassium permanganate + sulphuric acid would work but NO ONE would do it like that.

You mean carbonic acid. That's soda water, basically. Do YOU think that would work? (Hint: NO!)

The oxalates of Nd and Fe are basically insoluble. It's a recipe for stalled reaction/big mess.

Go down the hardware store and get some 'muriatic acid' (HCl) or toilet drainer (H2SO4).


http://i.word.com/imedical/permanganic%20acid
It does by the way

Brain&Force - 26-5-2015 at 13:51

Making permanganic acid is an extremely difficult endeavor, and it's not even stable for long in water. The use of oxalic acid works by complexing the iron away as the soluble tris(oxalato)ferrate(III) complex as the neodymium precipitates as the normal oxalate.

blogfast25 - 26-5-2015 at 14:11

Quote: Originally posted by DalisAndy  


http://i.word.com/imedical/permanganic%20acid
It does by the way


A word of advice: medical dictionaries aren't the beez knees for looking up chemical stuff, you know? :cool: <=== that's a picture of a doctor in cool shades, nodding approvingly, BTW...


[Edited on 26-5-2015 by blogfast25]

blogfast25 - 26-5-2015 at 14:13

Quote: Originally posted by Brain&Force  
Making permanganic acid is an extremely difficult endeavor, and it's not even stable for long in water. The use of oxalic acid works by complexing the iron away as the soluble tris(oxalato)ferrate(III) complex as the neodymium precipitates as the normal oxalate.


Errrmmm... except that oxalic acid won't oxidise Fe to Fe(III). So that separation method I proposed (and which works) only works AFTER the Fe has been oxidised to Fe(III).

DalisAndy - 27-5-2015 at 09:32

Quote: Originally posted by blogfast25  
Quote: Originally posted by Brain&Force  
Making permanganic acid is an extremely difficult endeavor, and it's not even stable for long in water. The use of oxalic acid works by complexing the iron away as the soluble tris(oxalato)ferrate(III) complex as the neodymium precipitates as the normal oxalate.


Errrmmm... except that oxalic acid won't oxidise Fe to Fe(III). So that separation method I proposed (and which works) only works AFTER the Fe has been oxidised to Fe(III).


So using sulfric then oxalic would separate the Neodynium? Since Fe is in the +2 oxidation state

elementcollector1 - 27-5-2015 at 09:45

It would, as I can attest to, but now you have to deal with the neodymium oxalate. So far, mine's resisted H2SO4, HCl, and even a mix of H2SO4 and NaNO3. Good luck to ya.

MrHomeScientist - 27-5-2015 at 09:48

Once again I'd recommend you read this thread in its entirety. Many of your questions have been answered right here. I know it's long; It's a complicated process. It's worth it to learn as much about it as possible if you're serious about pursuing this project.

DalisAndy - 27-5-2015 at 12:14

Does neodymium and other rare earths have black light florescent properties?

blogfast25 - 27-5-2015 at 12:35

Quote: Originally posted by DalisAndy  
Does neodymium and other rare earths have black light florescent properties?


Also in the thread.

Also here:

http://en.wikipedia.org/wiki/Neodymium%28III%29_chloride

aga - 27-5-2015 at 12:38

Has anyone successfully isolated Nd yet ?

There is a tiny vial of Neodymium Sulphate somewhere in the lab as i recall ...

blogfast25 - 27-5-2015 at 13:02

Quote: Originally posted by aga  
Has anyone successfully isolated Nd yet ?

There is a tiny vial of Neodymium Sulphate somewhere in the lab as i recall ...


Compound yes (sulphate), metal no. Re. latter, MrHomeScientist is the 'Amundsen' of that race here on SM.

aga - 27-5-2015 at 13:47

Huh ?

I followed some instructions in this thread around page 4000 or so and have had some nice lilac powder in a vial for many months.

That sample even survived the last Cull (to get the vials back) as the colour is rather nice.

blogfast25 - 27-5-2015 at 13:51

Quote: Originally posted by aga  
Huh ?



Listen, I've had 3 Chimay's and I'm not drunk yet. Try and keep up! ;)

aga - 27-5-2015 at 14:15

In a universe far, far a way, a long, long time ago, ISTR that it began with some magnets in an acid ...

One appears to have made NdSO4 simply following something (i forgot what) in this thread.

Lilac powder awaiting further instructions.

blogfast25 - 27-5-2015 at 14:24

Quote: Originally posted by aga  


Lilac powder awaiting further instructions.


There are two Nd sulphates. One's nice ruby red crystals, the other a sandy, lilac powder. We don't know what causes the difference. Be a trailblazer! :D

Hint: I think it's hydration. Not sure though... The easiest way would be to determine sulphate content of both 'varieties'.


[Edited on 27-5-2015 by blogfast25]

j_sum1 - 27-5-2015 at 14:27

Quote: Originally posted by aga  
Lilac powder awaiting further instructions.

1. Leave in vial to look pretty.
2. Start work on Lanthanum from pool phosphate remover.
3. Obtain another vial with pretty crystals.
4. Copy mrhomescientist's beautiful and highly effective apparatus
5. Come in second with a better yield on the Nd race
6. Win the La race.

aga - 27-5-2015 at 14:41

7. Get some more beers and just post nonsense.

Ok. sounds like a Plan.

blogfast25 - 27-5-2015 at 14:53

8. Look up protocol for BaSO4 gravimetry

elementcollector1 - 27-5-2015 at 16:18

I would also like to add that the pink oxalate seems completely resistant to a mixture of elemental bromine and sulfuric acid as well.
Running out of ideas to dissolve this thing...

blogfast25 - 27-5-2015 at 16:38

Quote: Originally posted by elementcollector1  
I would also like to add that the pink oxalate seems completely resistant to a mixture of elemental bromine and sulfuric acid as well.
Running out of ideas to dissolve this thing...


Random chemical mixing won't get you anywhere. There's nothing bromine can do here.

Calcine the oxalate to oxide. Then treat with boiling 95 % (at least) H2SO4. Or fuse with NaHSO<sub>4</sub>.

[Edited on 28-5-2015 by blogfast25]

DalisAndy - 27-5-2015 at 17:28

Quote: Originally posted by blogfast25  
Quote: Originally posted by aga  


Lilac powder awaiting further instructions.


There are two Nd sulphates. One's nice ruby red crystals, the other a sandy, lilac powder. We don't know what causes the difference. Be a trailblazer! :D

Hint: I think it's hydration. Not sure though... The easiest way would be to determine sulphate content of both 'varieties'.


[Edited on 27-5-2015 by blogfast25]

Ummm if they are crystals, shape can be determined by dissolving them and letting them solidify on a sting. Like borax ornaments

DalisAndy - 27-5-2015 at 17:36

Quote: Originally posted by blogfast25  
Quote: Originally posted by elementcollector1  
I would also like to add that the pink oxalate seems completely resistant to a mixture of elemental bromine and sulfuric acid as well.
Running out of ideas to dissolve this thing...


Random chemical mixing won't get you anywhere. There's nothing bromine can do here.

Calcine the oxalate to oxide. Then treat with boiling 95 % (at least) H2SO4. Or fuse with NaHSO<sub>4</sub>.

[Edited on 28-5-2015 by blogfast25]


Dude, oxalates decompose easily. Go look it up, like I did today. or you could use Lithium in ethanol (or what ever non-polar solvent you have). or when in doubt electrolysis, of it's sulphate form.

DalisAndy - 27-5-2015 at 17:45

I don't get why we haven't created a location to consolidate knowledge for projects? I mean like log all expirements done, so others can test it

blogfast25 - 27-5-2015 at 17:46

Quote: Originally posted by DalisAndy  


Dude, oxalates decompose easily. Go look it up, like I did today. or you could use Lithium in ethanol (or what ever non-polar solvent you have). or when in doubt electrolysis, of it's sulphate form.


Why, what did you find, huh? Lithium in ethanol??? That's lithium ethoxide. How is that supposed to decompose oxalates?

Oxalates in acid conditions are oxidised to CO<sub>2</sub> by KMnO<sub>4</sub> but RE oxalates are extremely insoluble, making that process very slow, I would think...

The sandy pink form of Nd sulphate would only crystallise into the other, ruby form, so that solves nothing. Direct experience, BTW. The two forms have been encountered by several experimenters, myself included, on this forum.

[Edited on 28-5-2015 by blogfast25]

DalisAndy - 27-5-2015 at 17:56

Quote: Originally posted by blogfast25  
Quote: Originally posted by DalisAndy  


Dude, oxalates decompose easily. Go look it up, like I did today. or you could use Lithium in ethanol (or what ever non-polar solvent you have). or when in doubt electrolysis, of it's sulphate form.


Why, what did you find, huh? Lithium in ethanol??? That's lithium ethoxide. How is that supposed to decompose oxalates?

Oxalates in acid conditions are oxidised to CO<sub>2</sub> by KMnO<sub>4</sub> but RE oxalates are extremely insoluble, making that process very slow, I would think...

The sandy pink form of Nd sulphate would only crystallise into the other, ruby form, so that solves nothing. Direct experience, BTW. The two forms have been encountered by several experimenters, myself included, on this forum.

[Edited on 28-5-2015 by blogfast25]


First, they were to separate ideas. Second, couldn't you ignite it with Carbon, since it's similar in chemistry to magnesium. Which it's known to react with Carbon in sulphate form

blogfast25 - 27-5-2015 at 18:00

You don't need to: you need to pyrolyse the Nd oxalate in the presence of plenty of air oxygen. Probably 700 C or higher. That affords the oxide Nb<sub>2</sub>O<sub>3</sub>.

Treat that as I suggested to convert to sulphate. Eliminate any sodium as explained in the thread, if needed.

j_sum1 - 27-5-2015 at 21:27

Quote: Originally posted by DalisAndy  
I don't get why we haven't created a location to consolidate knowledge for projects? I mean like log all experiments done, so others can test it
Well, we do have a prepublications forum (that is somewhat underused.) we do try to encourage posting in existing threads rather than starting new ones -- so all the information is together in one place. And I think that some of the larger more significant threads have been summarised, but that is a large job for someone. It also puts a metaphorical lid on the thread.
What often happens however is that thread contributors are engaged in similar but slightly different experiments that benefit from discussion but the thread does not lend itself to a single tidy description of a procedure.
Besides, half the fun is in reading and encountering ideas and solutions that you had not considered before. That is where the learning happens. So what if you have to read a bit.

blogfast25 - 28-5-2015 at 05:44

There's also a lot of stuff in the open literature. The double sulphates method of separation I got from there. It's should not be hard to find it again via the Googly Interwebs.

That method is used for the industrial recovery of Nd from Nd magnets, acc. one paper.

[Edited on 28-5-2015 by blogfast25]

DalisAndy - 29-5-2015 at 09:07

When using the oxalic acid on the magnets. Which oxidation state is iron move to? Also there is a +2 iron oxalate

blogfast25 - 29-5-2015 at 09:21

Quote: Originally posted by DalisAndy  
When using the oxalic acid on the magnets. Which oxidation state is iron move to? Also there is a +2 iron oxalate


Non-oxidising acids oxidise iron to Fe(II).

Oxidising acids like HNO3 oxidise it to Fe(III).

DalisAndy - 29-5-2015 at 11:09

I think I fix/finished the oxalic acid method.

1. Treat with sulfuric acid to remove boron
2. Treat with oxalic acid, filter out neodymium oxalate
3. Treat neodymium oxalate with a base
4. Decompose hydroxide

blogfast25 - 29-5-2015 at 11:19

Quote: Originally posted by DalisAndy  
I think I fix/finished the oxalic acid method.

1. Treat with sulfuric acid to remove boron
2. Treat with oxalic acid, filter out neodymium oxalate
3. Treat neodymium oxalate with a base
4. Decompose hydroxide


Fe(II) oxalate is very poorly soluble.

Nd oxalate is more insoluble than Nd hydroxide. Converting the Nd oxalate with base to Nd hydroxide therefore cannot work.

What do you mean by 'decompose hydroxide'?

[Edited on 29-5-2015 by blogfast25]

DalisAndy - 29-5-2015 at 11:23

The base treatment is to allow isolation of an oxide, thus a pure sample. If Fe(II) is not soluabe. Heat the Nd-Fe sulfate solution untill the Fe sulfate decomposes.

blogfast25 - 29-5-2015 at 11:34

Quote: Originally posted by DalisAndy  
The base treatment is to allow isolation of an oxide, thus a pure sample. If Fe(II) is not soluabe. Heat the Nd-Fe sulfate solution untill the Fe sulfate decomposes.


You really haven't got the foggiest, have you?

READ THE THREAD.

aga - 29-5-2015 at 12:16

What ?

Reading ?! Eeek !

Nobody ever said anything about having to read things before i signed up.

I want my money back !

blogfast25 - 29-5-2015 at 12:47

Quote: Originally posted by aga  

I want my money back !


Done. You want a tax deductible receipt for that refund?

DalisAndy - 2-6-2015 at 09:34

At what temperature does neodymium sulfatd decompose? Just curious, couldn't find it online. Also could someone explain the double salt method more? Mg has a high affinity for Nd. I read in one paper that you could melt down Mg and the magnets, and separate the Mg and Nd along with a Fe-B sludge. But that would require a blast furnace or a klim.

blogfast25 - 2-6-2015 at 09:54

Nd sulphate requires seriously high temperatures to fully decompose (to oxide). At a guess, several hours at near 1,000 C.

The double salt method works as follows.

The magnets are dissolved in an excess of 33 % HCl and filtered after dissolution and cooling.

Nd is now present as NdCl3, Fe as FeCl2, in a dark but transparent solution (no solids).

To the filtrate, somewhat diluted, is added powdered K2SO4. The amount is calculated so that the solution ends up saturated in K2SO4. The slurry is heated and simmered for about 1/2 hour. A sandy precipitate of Nd2(SO4)3.K2SO4.2H2O (could be .3H2O, can't remember right now) forms. The slurry is cooled and chilled (to maximise the insolubility of the double salt) and filtered off.

The filter cake is washed with cold, acidified saturated K2SO4 solution several times to remove any remaining iron(II). It's important to use an acidic wash to avoid oxidation of the Fe(II) to Fe(III) (which can precipitate as Fe(OH)3).

The washed filter cake doesn't have to be dried. It is treated with warm, strong ammonia which sets the K2SO4 free and converts the Nd to Nd(OH)3.

Filtering and washing extensively of the filter cake then affords relatively pure (RE > 99 % or better) Nd(OH)3, a starting point for Nd chemistry in general.


[Edited on 2-6-2015 by blogfast25]

DalisAndy - 2-6-2015 at 10:10

Ok thanks. I'm think I'm going to to cryastalize off neodymium sulfate after decomposing off the iron sulfate. Before it goes Fe(III)

blogfast25 - 2-6-2015 at 10:15

Quote: Originally posted by DalisAndy  
Ok thanks. I'm think I'm going to to cryastalize off neodymium sulfate after decomposing off the iron sulfate. Before it goes Fe(III)


It WON'T work.

Decomposing the iron sulphate will lead to partial hydrolysis of the Nd sulphate. Crystallising Nd sulphate is in itself not so easy: it's a poorly soluble compound with (quite exceptionally) higher solubility in the cold than hot.

[Edited on 2-6-2015 by blogfast25]

elementcollector1 - 2-6-2015 at 14:21

Just calcined the oxalate to a gray oxide, then placed this in a mix of sulfuric acid and water. Left this overnight, and checked again this morning to find a sandy pink powder with traces of gray still mixed in, and a completely clear solution. Would this be pure, dilute neodymium sulfate? There's no way the oxalate could have survived in such quantity - it melted during calcination...

blogfast25 - 3-6-2015 at 05:09

Quote: Originally posted by elementcollector1  
Just calcined the oxalate to a gray oxide, then placed this in a mix of sulfuric acid and water. Left this overnight, and checked again this morning to find a sandy pink powder with traces of gray still mixed in, and a completely clear solution. Would this be pure, dilute neodymium sulfate? There's no way the oxalate could have survived in such quantity - it melted during calcination...


The only way to find out is to try and dissolve it in iced water in which it slowly dissolves.

DalisAndy - 3-6-2015 at 20:34

Would using a modified Alkaline extraction method work? Sulphates -> hydroxides?

DalisAndy - 3-6-2015 at 20:50

Oh gentleman I found a way to extract rare earth elements with sulfric acid, using a solvent extraction method.
http://etheses.whiterose.ac.uk/638/1/uk_bl_ethos_432307.pdf

blogfast25 - 4-6-2015 at 06:08

Quote: Originally posted by DalisAndy  
Would using a modified Alkaline extraction method work? Sulphates -> hydroxides?


REEs don’t respond to alkalis, they’re not amphoteric.
Quote: Originally posted by DalisAndy  
Oh gentleman I found a way to extract rare earth elements with sulfric acid, using a solvent extraction method.
http://etheses.whiterose.ac.uk/638/1/uk_bl_ethos_432307.pdf


Solvent extraction of lanthanides/actinides is nothing new. We (amateurs) tend to avoid it because of the required chemicals.

elementcollector1 - 4-6-2015 at 12:17

Quote: Originally posted by blogfast25  

The only way to find out is to try and dissolve it in iced water in which it slowly dissolves.


Is it soluble in cold, diluted sulfuric acid? I tried putting the original mix in the freezer, just to see what happened - nothing did.

blogfast25 - 4-6-2015 at 12:33

Quote: Originally posted by elementcollector1  


Is it soluble in cold, diluted sulfuric acid? I tried putting the original mix in the freezer, just to see what happened - nothing did.


And for your next experiment you will put some sugar at the bottom of a glass of cold water and observe.

x days later: your sugar is still at the bottom of your glass. Conclusion: sugar is insoluble in cold water (W-R-O-N-G ! ! !)

At 0 C (or near) the rate of dissolution of Nd sulphate (or sugar) is very low. And without prolonged and intensive stirring it will NOT dissolve. Same as your sugar. And relying solely on diffusive dissolution is very, very slow.

Dilute H2SO4 won't have much effect but remember the Common Ion effect.

[Edited on 4-6-2015 by blogfast25]

MrHomeScientist - 8-6-2015 at 09:43

Quote: Originally posted by j_sum1  
...
4. Copy mrhomescientist's beautiful and highly effective apparatus
...

I'll grant it looks neat, but 'highly effective' is still yet to be seen! I just got back from a week-long vacation so I haven't thought about this in a while. (I hadn't done much of anything with it before that week either, but still...)

The current plan of attack is to do a dry run just heating a crucible of lithium metal under argon, to see if I can at least melt it without it catching fire. That will tell me if my inert blanket is sufficient. Then hopefully the heat won't have welded the pipe caps onto the body, and I can actually open the thing up after it cools. Then, if those checks pass, we'll move on to the main event.

If that works (fingers crossed!) I will post a nicely formatted article with photos and video going over the entire process in the Prepublication section (among other places).


Oh and Andy, I found that magnesium melt method a while back and posted it in the thread some pages ago. It's very interesting, but again you'd need an inert atmosphere, high temperature, sealed crucible setup which is very difficult for the amateur to achieve. It involves distilling Mg metal, so that'll give you some idea of the temperatures involved. I feel like this will be ignored again, but let me repeat once more: Read the thread. People have been remarkably patient in repeating themselves here so far.

[Edited on 6-8-2015 by MrHomeScientist]
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