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Author: Subject: The trouble with neodymium...
Atrum
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[*] posted on 12-7-2016 at 15:02


@blogfast25

Thanks, I will give that a go.

Here is a photo of the precipitate that I centrifuged to wash it. Filtering was impossible. Very Very fine precipitate. It goes through filter paper like water through a sieve.



insol precipitate.jpg - 816kB




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[*] posted on 12-7-2016 at 15:17


Quote: Originally posted by Atrum  
@blogfast25

Thanks, I will give that a go.

Here is a photo of the precipitate that I centrifuged to wash it. Filtering was impossible. Very Very fine precipitate. It goes through filter paper like water through a sieve.



Well, you you've got enough for that test.




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[*] posted on 12-7-2016 at 15:21


Quote: Originally posted by blogfast25  
Quote: Originally posted by Atrum  
@blogfast25

Thanks, I will give that a go.

Here is a photo of the precipitate that I centrifuged to wash it. Filtering was impossible. Very Very fine precipitate. It goes through filter paper like water through a sieve.



Well, you you've got enough for that test.


Yeah, that is just one tube out of 4 that have the same amount. Hopefully I will be able to extract the rest of the nd from it. I am hoping for some large crystals of the sulfate eventually.




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[*] posted on 13-7-2016 at 00:13


This was probably obvious to everyone else, but I felt it was worth noting:

-I've left a solution of oxalic acid and precipitate of iron/neodymium oxalate to chelate the iron out for over a year now. Not much happened, so I thought to just let time do the work.

-Put in some concentrated hydrogen peroxide yesterday, after seeing other users' experience with oxalates. Within 24 hours, most of the iron had leached out of the precipitate, leaving the supernatant solution a vivid green and the precipitate an increasingly lighter shade of whitish-pink.

Given that I started with about 8 hard drive magnets, this leaves me with quite a bit of oxalate to work with. Hopefully I can calcine it at a low enough temperature that it won't become inert like the previous stuff...




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[*] posted on 13-7-2016 at 10:06


Quote: Originally posted by elementcollector1  
This was probably obvious to everyone else, but I felt it was worth noting:

-I've left a solution of oxalic acid and precipitate of iron/neodymium oxalate to chelate the iron out for over a year now. Not much happened, so I thought to just let time do the work.

-Put in some concentrated hydrogen peroxide yesterday, after seeing other users' experience with oxalates. Within 24 hours, most of the iron had leached out of the precipitate, leaving the supernatant solution a vivid green and the precipitate an increasingly lighter shade of whitish-pink.

Given that I started with about 8 hard drive magnets, this leaves me with quite a bit of oxalate to work with. Hopefully I can calcine it at a low enough temperature that it won't become inert like the previous stuff...




So, I am curious. How many different solvents have you tried when trying to dissolve Nd oxalate?




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[*] posted on 13-7-2016 at 10:35


Quote: Originally posted by Atrum  

So, I am curious. How many different solvents have you tried when trying to dissolve Nd oxalate?


None but water, actually. I didn't figure oxalate would be soluble in much. At any rate, my pink neodymium oxalate is currently filtering out, and looking very nice while doing so.

I do still have the results of my attempt to dissolve the previous, calcined stuff - the precipitate's turned a sandy pink, and the solution itself is quite clear. Could this be neodymium sulfate, or just 'wet', unreacted neodymium oxalate? Hard to tell.




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[*] posted on 13-7-2016 at 11:12


Quote: Originally posted by Atrum  


So, I am curious. How many different solvents have you tried when trying to dissolve Nd oxalate?


It's not really a question of 'How many different solvents have you tried'.

1. Nd oxalate is supremely insoluble, that makes it hard to dissolve in anything including strong acids.

2. Oxalic acid is not volatile: no matter what acid you throw at it, the oxalate ions are always there.

One's 'best bet' is to transform the oxalate, e.g. by oxidising it. Permanganate oxidises H2Ox to CO2 (leaving you also with Mn(II), of course).

Hypochlorite (bleach) might also do it but would probably slow because thin bleach contains only a few w% of ClO-.

[Edited on 13-7-2016 by blogfast25]




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[*] posted on 13-7-2016 at 16:46


Quote: Originally posted by blogfast25  
Quote: Originally posted by Atrum  


So, I am curious. How many different solvents have you tried when trying to dissolve Nd oxalate?


It's not really a question of 'How many different solvents have you tried'.

1. Nd oxalate is supremely insoluble, that makes it hard to dissolve in anything including strong acids.

2. Oxalic acid is not volatile: no matter what acid you throw at it, the oxalate ions are always there.

One's 'best bet' is to transform the oxalate, e.g. by oxidising it. Permanganate oxidises H2Ox to CO2 (leaving you also with Mn(II), of course).

Hypochlorite (bleach) might also do it but would probably slow because thin bleach contains only a few w% of ClO-.

[Edited on 13-7-2016 by blogfast25]


That sounds like it could be interesting to try out. I appreciate the information blogfast.




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[*] posted on 13-7-2016 at 17:30


This source:

http://www.nrcresearchpress.com/doi/pdf/10.1139/v61-211

... indicates oxidation of oxalate with HClO (acidic conditions) is possible, with high activation energy (heating required, basically):


Quote:
There is some indication that, if the reaction is relatively exothermic, the activation energy is lower. This is perhaps to be expected, since if the new bond to oxygen is strong, this would tend to lower the energy of the activated complex as it is formed. However, this trend is not universal; for instance, the very exothermic oxidation of oxalate [by HClO] has quite a high activation energy. This particular reaction may be a little different from the others, since it is apparently not an oxygen transfer but an electron transfer; although it could be formulated either way.


A combination of bleach and a weak acid (like vinegar) + some heat and time, might just slowly dissolve an insoluble oxalate to chloride/acetate.

[Edited on 14-7-2016 by blogfast25]




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[*] posted on 19-7-2016 at 09:39


A picture of the neodymium oxalate:



Still has oxalic acid crystals in it, but I don't particularly mind - those should simply decompose upon heating, right?

I think I'll take this to the university this fall and calcine it there, with their temperature-controlled furnace - what's the lowest temperature I can do this at? The furnaces typically run for 24 to 48 hours, so time is not imperative.




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[*] posted on 22-7-2016 at 11:55


That's quite a lot...is the neodymium oxalate itself very crystalline, or just the oxalic acid mixed in with it - it's hard to tell in the picture.



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[*] posted on 22-7-2016 at 12:18


I am slowly building up a good amount of Nd Oxalate. It seems to me that it is the best method to get it free of iron contamination. I am still playing around though.

On the side I am trying to get a crop of nice green Iron(III)Oxalate crystals. I find the thermochromism of iron oxalate to be interesting. Bright Green when cold, brownish red when hot.




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[*] posted on 28-7-2016 at 21:32


I'm considering using sodium metal to reduce my neodymium fluoride. This would presumably be rammed into a graphite crucible, followed by the neodymium fluoride, followed by external heating from a propane blowtorch. Does anyone see any problems with this method?



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[*] posted on 1-8-2016 at 06:16


I did some calculations a while back that suggested sodium won't work as well. It's probably back a few pages in the thread somewhere, so here they are again:

∆H values:
NdF3 = -1657 kJ/mol
MgF2 = -1124..2 kJ/mol
NaF = -576.6 kJ/mol
LiF = -616.0 kJ/mol

For Mg
2NdF3 + 3Mg = 3MgF2 + 2Nd
∆Hrxn = 3(-1124.2) – 2(-1657) = -58.6 kJ/mol

For Na
NdF3 + 3Na = 3NaF + Nd
∆Hrxn = 3(-576.6) – (-1657) = -72.8 kJ/mol

For Li
NdF3 + 3Li = 3LiF + Nd
∆Hrxn = 3(-616) – (-1657) = -191 kJ/mol


So lithium produces quite a bit more heat than the others, though I’m not sure how that number corresponds with actual reaction temperature. It at least suggests that Li should favor the easiest reaction conditions.

Also I built my setup to deliver inert atmosphere to the crucible, mostly because I'll be using lithium but also to protect the newly formed Nd from the air. It's possible it will sink to the bottom of the melt and be protected anyway, but it's something to consider.

God I need to finish this project. I finally have all the pieces in place and can proceed at any time. I just need to get off my ass and do it.
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[*] posted on 1-8-2016 at 09:11


Quote: Originally posted by MrHomeScientist  

So lithium produces quite a bit more heat than the others, though I’m not sure how that number corresponds with actual reaction temperature. It at least suggests that Li should favor the easiest reaction conditions.

Also I built my setup to deliver inert atmosphere to the crucible, mostly because I'll be using lithium but also to protect the newly formed Nd from the air. It's possible it will sink to the bottom of the melt and be protected anyway, but it's something to consider.

God I need to finish this project. I finally have all the pieces in place and can proceed at any time. I just need to get off my ass and do it.


Li is the go to reducing agent in this case.

Na might work but will need stronger heating.

Both will need strong heat + inert blanket (Ar), to obtain the Nd as a regulus.

Powdered Nd in a alkali metal fluoride matrix will be inseparable, unless you heat to to well above 1000 C: the liquid Nd will then separate out.

No matter how you play it, you'll need to heat to close to 1000 C, with adequate Ar blanket.

It might be tempting to remove heat after reaction has taken place but it's probably not a good idea, if you want adequate metal/slag separation.



[Edited on 1-8-2016 by blogfast25]




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[*] posted on 1-8-2016 at 09:36


The last problem I had to solve was connecting the Ar tank to "The Gavel" (my inert atmosphere crucible). For some reason the fittings for argon tanks don't connect with regular pipe, so I had to order a special adapter.

The next problem I foresee is being able to heat to the proper temperature. The Gavel won't fit in my mini propane furnace, so I bought some kaowool to wrap it in. If that won't hold in heat well enough, I'll have to spring for a larger, perhaps custom-made, furnace setup.
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[*] posted on 2-8-2016 at 12:27


I've been constructing furnaces that operate to 1200 C with ceramic fiber board, Alumel wire, some high-temp. ceramic adhesives or (even furnace cement) and ceramic coil forms for many years. I'd estimate you could buy everything needed to bring a 200 to 300 cubic inch space to 1200 C for $150. Control is a whole different issue. I bought a $25 PID controller on eBay and a $10 high current solid state relay for $10. I had a K thermocouple just laying around. This will give you a nice furnace with years of service. Of course, gas-heated rigs are cheaper but control is limited and of course, there are exhaust issues...

If you seriously consider the electric option and want some pointers, I've made all the mistakes that you can make, I've got it down to a reliable formula.





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[*] posted on 10-8-2016 at 09:57


I'd love to learn from your experiences. An electric furnace would be a very valuable thing to have, not just for this particular experiment.
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[*] posted on 20-8-2016 at 21:31


I think I may have accidentally come across a much better separation method for neodymium from magnets.

I've been experimenting with gadolinium, terbium, and holmium acetates, which are conveniently non-hygroscopic but very soluble in water. They are also quite poorly soluble in ethanol, though crystals of the acetates will fall apart in it.

Contrast this behavior with boric acid, iron(III) acetate and nickel acetate, which are highly soluble in ethanol.

I recommend you test this method (I'm about to test this myself). Dissolve a neodymium magnet in acetic acid and hydrogen peroxide. Remove the supernatant and evaporate it to dryness. Redissolve the material in ethanol. The neodymium acetate should remain as a powder. Filter with ethanol rinses and you should get quite pure neodymium acetate (of course there could be other lanthanides).

I haven't tried this myself, but theoretically it should get rid of nickel, iron, and boron in the magnets.




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[*] posted on 20-8-2016 at 23:03


That's some good thinking B&F. I like where that is going.
I have been pondering recently how easy it is too overlook acetic acid for reacting with metals. Sure, it tends to be slow but it has a lot of plus sides.




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[*] posted on 23-8-2016 at 09:50


Yeah, the non-hygroscopic salts but very water-soluble salts are the biggest advantage IMO. Just a quick note - keep the concentration of acetic acid low. I've found that highly concentrated acetic acid will not solubilize the acetates, and that glacial acetic acid does not appreciably attack holmium at room temperature for thirty minutes. (I suggested hydrogen peroxide to keep iron as iron(III) acetate which retains extremely high solubility in ethanol). I think iron(II) acetate will also be soluble in ethanol, so that should be another bonus.

Also, I think you can form a soluble neodymium acetate ethanol complex with strong heating, so I'd avoid heating it while it's in ethanol. That is definitely the case in methanol for dysprosium acetate.




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[*] posted on 26-11-2016 at 15:27


Okay, time to try this again. Me and Brain&Force have been discussing possible alternative amateur routes to pure neodymium salts for a while now, and while I ran an earlier attempt this year, it culminated just a week ago in utter failure. So, this time, I'm going to post my method(s) beforehand, in the hopes that somebody will be able to spot problems before they occur.

Also, this might turn into my senior thesis.

The goal of this series of tests is to identify which chemical separation route provides the most promising results for the amateur when trying to isolate neodymium compounds from magnets. Several separation methods will be evaluated, consisting of the potassium oxalate route, the ethanol route, and the selective precipitation route. From there, the oxalate will be further tested for its solubility in ionic neodymium(III) salts, as according to this source.

The reaction idea is as follows:

First, a 2"-by-2"-by-1" neodymium magnet (approximate formula Nd2Fe14B) is heated to 500 degrees Fahrenheit (approx. 260 degrees Celsius) in a standard convection oven for 1 hour to at least partially demagnetize it (for ease of processing). Once done, the nickel casing is stripped away as best as possible, leaving only the sintered material underneath. The approximate weight of this material, assuming none has been lost so far, should be 478.5 to 491.6 grams, assuming the reported density of 7.3-7.5 grams/centimeter cubed is correct. This results in an estimated neodymium content of 127.7 to 131.2 grams.

This material is then dissolved in an excess of 25% acetic acid (diluted from glacial acetic acid) to form green iron(II) acetate tetrahydrate, Fe(C2H3O2)2(H2O)4, pale mauve neodymium(III) acetate tetrahydrate, Nd(C2H3O2)3(H2O)x, and an unknown boron product (possibly elemental boron, 'boron acetate,' or boric acid, or a mixture). Assuming the boron product is mostly insoluble, it is filtered off, along with any plastic or organic binding material that may have been used in the manufacture of the magnet.

Nd2Fe14B (s) + HC2H3O2 (aq) -> Fe(C2H3O2)2(H2O)4 (aq) + Nd(C2H3O2)3(H2O)4 (aq) + (unknown boron compound, likely boric acid?)

Finally, this mixture is treated with an excess of approx. 11-14% hydrogen peroxide solution, which has been freeze-distilled from topical grade 3%. This oxidizes the iron(II) acetate to iron(III) acetate, also known as basic iron acetate, Fe3O(C2H3O2)6(H2O)3]C2H3O3. The hydrogen peroxide does not affect the neodymium(III), because it is a more stable ion.

Fe(C2H3O2)2(H2O)4 (aq) + Nd(C2H3O2)3(H2O)4 (aq) + H2O2 (aq) -> Fe3O(C2H3O2)6(H2O)3]C2H3O3 (aq) + Nd(C2H3O2)3(H2O)4 (aq)

The solution thus far is split into three equal parts.

The first third of the solution is mixed with an excess of aqueous potassium oxalate (K2C2O4) to precipitate neodymium oxalate (Nd2(C2O4)3), chelating the iron(III) into soluble, vivid green potassium ferrioxalate (K3(Fe(C2O4)3)3). The pale pink neodymium oxalate is filtered off, washed, and dried to remove any excess water.

3 Fe(C2H3O2)2(H2O)4 (aq) + 2 Nd(C2H3O2)3(H2O)4 (aq) + 6 K2C2O4 (aq)-> Nd2(C2O4)3 (s) + K3(Fe(C2O4)3)3 (aq) + 3 KC2H3O2 (aq)

(This reaction is not properly balanced.)

This oxalate is then split once again into three equal parts by mass. The first part is calcined with a torch, forming pale blue neodymium oxide (Nd2O3) and carbon dioxide according to the following reaction:

2 Nd2(C2O4)3 + 3 O2 -> 12 CO2 + Nd2O3

This is weighed to determine the yield of Nd, assuming a perfect 1/6 of the Nd from the magnet was in the starting oxalate. Theoretical perfect yield, therefore, is approximately 24.85 to 25.5 grams of neodymium oxide, representing 21.3 to 21.9 grams of neodymium.

The second portion of the oxalate is weighed and placed in a pre-prepared solution of aqueous neodymium chloride, and left to dissolve. Because this test is more qualitative than quantitative, results will depend on how much, if any, of the oxalate remains.

The third portion, as with the second, is weighed and placed in a pre-prepared aqueous solution of pure neodymium acetate, to see if a similar complexing reaction occurs.

Returning to the second portion of the mixed iron and neodymium acetates, it is boiled to dryness and mixed with ethanol. According to Brain&Force's suggestion, this should result in the iron(III) acetate dissolving in the ethanol, leaving behind insoluble neodymium acetate tetrahydrate. This acetate is then dried, weighed and compared to the theoretical yield of 128.4 to 132.0 grams. This is then dissolved in water, and tested with thiocyanate for iron content.

The third and final portion of the initial mix of iron and neodymium acetates is subjected to NurdRage's method of selective precipitation, to see if it works between iron and neodymium. This would be done by taking about half of the solution and precipitating out the hydroxides with NaOH, followed by adding the washed hydroxide mixture back into the other half of the solution. The reactions are as follows:

Fe3O(C2H3O2)6(H2O)3]C2H3O3 (aq) + Nd(C2H3O2)3(H2O)4 (aq) + NaOH (aq) -> Nd(OH)3 (s) + Fe(OH)3 (s)+ NaC2H3O2 (aq)

Nd(OH)3 (s) + Fe3O(C2H3O2)6(H2O)3]C2H3O3 (aq) -> Fe(OH)3 (s) + Nd(C2H3O2)3(H2O)4 (aq)

While this has never been tested to my knowledge, the difference in electropositivity between iron and manganese (+0.28) is less than the difference in electropositivity between iron and neodymium (+0.69), which from my admittedly ignorant interpretation means that this reaction will go forward with greater ease than its manganese counterpart. If the reaction does indeed work, then this would be an easy alternative to both the oxalate and the ethanol route for separating iron from neodymium. The resultant solution would be filtered, and tested with thiocyanate for presence of iron.

The eventual goal, once all has been said and done, is to convert any soluble neodymium salts produced into NdF3, which will be then converted into Nd metal.

Questions so far:

-Why does basic iron acetate have such a ridiculously complex formula?
-Is there anything in this post that strikes you as partially or completely incorrect?


[Edited on 11-26-2016 by elementcollector1]




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[*] posted on 28-11-2016 at 12:08


Some comments:

To demagnetize the magnets, all that's needed is to blast them with a propane torch for a minute or two until they separate from the metal tongs you're holding them with. But perhaps you wanted a more repeatable process for your procedure.

I don't believe 25% acetic acid would do anything to elemental boron, so your first step is probably just B as a product.

Once you've made Nd<sub>2</sub>O<sub>3</sub>, it will likely be difficult to transform it into anything else. I seem to remember Woelen saying his commercial sample was impervious to just about everything.

This is a great project for a thesis. So many things to experiment with and test. The ultimate extraction method would be from the paper I've attached: "Direct Extraction and Recovery of Neodymium Metal from Magnet Scrap". The apparatus is a little tough to build though!

Attachment: Direct Extraction and Recovery of Neodymium Metal from Magnet Scrap.pdf (257kB)
This file has been downloaded 303 times


Every time I see a post in this thread I get a little kick to finish my part of the experiment. I need to beat everyone else to Nd! :P
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[*] posted on 28-11-2016 at 16:49


Sadly, I lack a propane torch. One of the many drawbacks to trying to do chemistry while at university...

In that case, the boron can be easily filtered off and saved for later.

I actually have a commercial sample on the way as we speak, so that I can give comparisons between what I isolate from the magnets and the 'pure' substances. From what I've studied, the trend seems to be that the reactivity largely depends on the thermal history - the lower the decomposition temperature of the oxalate, the better.

That paper's fascinating, but you're right, the setup's nearly impossible! It requires molten magnesium, for one thing, which I'm almost certain is pyrophoric, and for another, it requires essentially that the magnesium/neodymium mix be distilled in a stainless steel container. I bet the apparatus could be simplified, and if I ever get my furnace working it's definitely something to try, but I doubt this is the most practical route for an amateur. On a related note, though, how soluble are the other lanthanides in molten magnesium?





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[*] posted on 28-11-2016 at 23:41


Quote: Originally posted by blogfast25  
Due to beaker breakage the investigation into the presence of Pr had to be abandoned (evidence lost!). Shame.


Same thing here, I think I need some stronger beakers. Did you give it another try?

[Edited on 29-11-2016 by Tomo]

praseodymiummaybe.jpg - 37kBs1600.jpg - 29kBs1600.jpg - 29kBNdCl3.jpg - 21kB




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