Sciencemadness Discussion Board - The trouble with neodymium...

The trouble with neodymium...

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blogfast25 - 2-4-2012 at 12:08

Have you added oxalic acid yet?

elementcollector1 - 2-4-2012 at 13:22

Nope. I'm waiting for the iron hydroxide to precipitate out (or should I just proceed even if it hasn't?)

blogfast25 - 3-4-2012 at 06:15

Yes, you can add oxalic acid at this point.

What should happen is that the solution goes green (slightly fluorescent too - this is the trisoxalato ferric complex) and a precipitate of neodymium oxalate will form. Allow the neodymium oxalate to settle, then filter off the supernatant liquid. Wash the filter cake with ample water, this is your relatively pure neodymium oxalate. It should look distinctly different under tungsten light and TL light.

CAUTION: oxalate bearing liquors are toxic. Good gloves and goggles are in order. Work calmly and know what you're doing. Be prepared to clean up any spills quickly and completely!

The reasons why no Fe(OH)3 precipitated can be:

* pH was too low
* iron concentration was too low
* both

But this should not impede anything.

[Edited on 3-4-2012 by blogfast25]

elementcollector1 - 3-4-2012 at 17:09

Still with heat, right?

blogfast25 - 4-4-2012 at 04:26

No. The heat would have been needed to help dissolve the precipitate of Fe(OH)3. Now, RT is fine.

elementcollector1 - 6-4-2012 at 18:41

Alright, I now have a relatively small amount of white precipitate (which probably contains leftover oxalic acid) and a pale-ish green solution. How do I test this for neodymium content? (HCl? Calcination?)

blogfast25 - 7-4-2012 at 05:18

Filter off and wash the Nd oxalate copiously with clean warm water to remove traces of oxalic acid and other solutes. Dry it, it should now look distinctly different in TL light and incandescent light.

Calcining the oxalate in air gives Nd2O3 which is reportedly soluble in strong (hot?) HCl, to give NdCl3 which can be crystallised (albeit with difficulty because it's very soluble and hygroscopic).

Nathaniel - 2-5-2012 at 04:44

I also finally managed to obtain the magnet from an old hard drive, but instead of only getting neodymium out, I'd also like to obtain the nickel. I found out that the brackets for the magnet are made of permalloy (80% Ni, 20% Fe), so that's quite some nickel. But after reading these posts I have doubts about that, because I got the impression that the casing itself is mostly iron, which is just electroplated with nickel... Does anyone know if the nickel content is really high enough and would be worth extracting? The wikipedia also has some info on permalloy and I think it would be possible...
Thanks for help

blogfast25 - 2-5-2012 at 05:07

Quote: Originally posted by Nathaniel

I also finally managed to obtain the magnet from an old hard drive, but instead of only getting neodymium out, I'd also like to obtain the nickel. I found out that the brackets for the magnet are made of permalloy (80% Ni, 20% Fe), so that's quite some nickel. But after reading these posts I have doubts about that, because I got the impression that the casing itself is mostly iron, which is just electroplated with nickel... Does anyone know if the nickel content is really high enough and would be worth extracting? The wikipedia also has some info on permalloy and I think it would be possible...
Thanks for help

That’s quite interesting, Nathaniel, do you have some references for that claim? Personally I’m doubtful about it.

Of course it’s relatively easy to find out yourself. Dissolve the suspected Permalloy in nitric acid (nickel dissolves only with great difficulty in HCl or sulphuric acid, it resembles copper in that respect) and precipitate the metals as hydroxides with strong ammonia. Excess ammonia causes the nickel to re-enter solution as the nickel hexammonium complex (blue), so consider it a colour test, because smaller amounts of nickel would go unnoticed.

I’ve a load of these magnets, still in their brackets, so I might have a go myself…

Nathaniel - 2-5-2012 at 06:35

Thanks for reply, I had the same plan for separation in mind as the one you suggested

I'll dissolve small amount of the casing in HCl/H2O2 (that should dissolve it). I can also test for the nickel with homemade DMG just to be sure

You can try yourself, but I'll deffinately do it on friday and post the results

http://www.scrapmetaljunkie.com/269/how-to-scrap-hard-drives...

http://en.wikipedia.org/wiki/Permalloy

The data in the first link seems very probable, because it says the alloy doesn't "let through" much magnetic field, which could damage other parts of computer so it's very useful for these purposes.

blogfast25 - 2-5-2012 at 08:23

Quote: Originally posted by Nathaniel

Thanks for reply, I had the same plan for separation in mind as the one you suggested

I'll dissolve small amount of the casing in HCl/H2O2 (that should dissolve it). I can also test for the nickel with homemade DMG just to be sure

You can try yourself, but I'll deffinately do it on friday and post the results

Oh, but I'll gladly wait for your results first.

HCl/H2O2 will probably do it but you'll need topping up with H2O2, as much of it tends to decay away w/o doing much. You could also try HCl/Na(or K)NO3 ("poor man's Aqua Regia")

[Edited on 2-5-2012 by blogfast25]

Nathaniel - 4-5-2012 at 02:00

Ok, so I did a few reactions and here are the results:

First I cut a piece of the suspected permalloy with a saw, with a piece of paper below to trap the filings.
I actually wanted to dissolve the piece that I cut, but the amount of filings was surprisingly large (looked large

) so I decided to test those instead. The scale however did not show anything (so it was below 0,1g)

I put the filings in a 50ml erlenm.flask and added 3,5ml 30% HCl. I waited for the bubbling to cease and then heated the solution to dissolve everything and added 3ml water. The solution was strongly yellow. I added about 0,5ml H2O2 (30%) - the solution turned dark orange (Fe²⁺ --> Fe³⁺) I boiled a solution for a bit to destroy excess peroxide

Then I did the ammonia test: To 1ml of the solution I added 25% ammonia dropwise. Fe(OH)3 started forming after about 3 drops; I added 3,5ml total. After the percipitate settled, the solution above was clear and just a bit yellowish in colour.

Being quite dissapointed I took a few drops of this yellow solution and added some very diluted dimethylglyoxime solution. The colour didn't appear immediately, but after some shaking, the distinctive pink colour appeared.

So what should I assume from that? Everything being so diluted I don't know if (compared to iron) there's just a small amount of Ni present (and the DMG test just managed to detect it) or is the complex just too diluted for the solution to be blue? It would probably be best to dissolve the big piece too and make a more conc. sample, but I'd like to know if it's even worth the effort....

Thanks for helping out

blogfast25 - 4-5-2012 at 03:38

I think it's worth the effort for corroboration purposes. But I think the Ni you detected is coating, not alloy.

[Edited on 4-5-2012 by blogfast25]

Update:

My own quick test didn’t show up any Ni either. I dissolved a piece (about 0.1 g) of hard drive neomagnet bracket in a few ml of 70 % HNO3. In the mean time I prepared a solution of NiCl2.6H2O (0.4 g) in a few ml of DIW, as a control. I have no DMG so had to rely on (relatively weak) ammonia solution.

To the control was added 1 M NaOH till precipitation (Ni(OH)2 hydrate) occurred, then NH3 solution was added. The precipitate dissolved and a pale blue-lavender Ni ammonia complex solution was obtained. It’s much less intensely blue than its copper homologue.

After cooling (it is then pale yellow - not a hint of green) to the sample solution 1 M NaOH was added dropwise till a permanent Fe(OH)3 hydrate precipitation resulted. To this ammonia solution was added but the supernatant solution remained colourless. By now all ferric hydroxide has collected south and no blue colouration of the clear liquid is seen. This is not the behaviour of a 70 % Ni alloy, even though the NH3 test isn’t near as sensitive as DMG. Unless any nickel irreversibly co-precipitated with the ferric hydroxide, it strongly suggest the sample contains no or very little nickel.

Well known object lesson to all: don’t believe everything you read on the Tinkerwebs.

[Edited on 4-5-2012 by blogfast25]

Nathaniel - 6-5-2012 at 22:43

I dissolved a 2g piece in HCl/HNO3 and the filtrate, after adding ammonia, was colourless. Test with DMG positive (strong pink), but no blue or green so I guess it really isn't a nickel alloy...
I thought the info on the web was ok, but I guess I was wrong...I was really happy to finally get some nickel, but I guess I'll have to look for an anode.. Thanks for help again

blogfast25 - 7-5-2012 at 05:26

Quote: Originally posted by Nathaniel

I thought the info on the web was ok, but I guess I was wrong...I was really happy to finally get some nickel, but I guess I'll have to look for an anode.. Thanks for help again

Get some scrap Nichrome wire (hair driers etc): 80 % Ni/20 % Cr. Try separation based on Cr's amphoterism.

Neodymium sulphate: another ‘funny turn’…

blogfast25 - 7-7-2012 at 07:45

I got 165 g of neomagnets from a scrap merchant in return for some scrap metal and decided to try and turn it into neodymium sulphate once again. This time I chose to try and precipitate the sulphate by direct addition of 50 % sulphuric acid to a solution of ‘neomagnet chloride’.

So the magnets were crushed and reacted with about 500 ml of 36 % HCl, filtered to clarity and then to the filtrate a calculated amount of sulphuric acid was added. At first nothing happened but upon heating to about 90 C the sulphate dropped out as flaky crystals. And then I made a mistake. After decanting off the supernatant liquid I added cold water for rinsing and low and behold a good dollop of it dissolved immediately! Nd2(SO4)3 is of course much more soluble in cold than hot water but it’s only supposed to dissolve slowly in cold water. In short, I lost about 75 % of the crop to the wash water and filtrates. But that’s not the end of it…

The total volume had by then swollen to a hefty 1.2 L and I tried to precipitate the Nd as Nd2(SO4)3.K2SO4.3H2O, a very sparsely soluble double salt (discussed above), by taking advantage of the acid reserve of the solution. A calculated amount of 50 % KOH was added but nothing happened. Frustrated I decided to precipitate the lot, Fe included, with NaOH 50 % and start again. But low and behold, on slowly adding NaOH an off white sandy precipitate started dropping out. Adding the NaOH slowly while checking pH was low enough, I precipitated all the Nd, while keeping Fe3+ in solution, presumably as Nd2(SO4)3.Na2SO4.3H2O (the ammonium double salt is also known, all three are poorly soluble in water).

Filtered and washed with 0.1 M NaHSO4 (to prevent any Fe3+ from precipitating), the double salt was then treated with NaOH to convert it to Nd(OH)3 and Na2SO4. The Nd(OH)3 was then converted to Nd2(SO4)3, which tomorrow will be washed with small amounts of boiling 0.1 M H2SO4 and finally with boiling water.

[Edited on 8-7-2012 by blogfast25]

Poppy - 7-7-2012 at 15:05

I would not recommend the proposal of existence of the double salt you mentioned. As myself I tried putting some ammonium sulphate within the freshly purified Nd sulfate and when it just precipitated by evaporation it seemed not double salts formed as the Nd sulfate yield was just about the predicted within the calculations.

All of the (NH4)2SO4 was recovered from the supernatant solution later.
I dont believe you are right about the ammoniacal double salt, but

Can you please support that with stoichometric weighing?

I wan't to put that claim into my own post on double salt possibilities

Thanks!!

[Edited on 7-7-2012 by Poppy]

blogfast25 - 8-7-2012 at 04:30

Quote: Originally posted by Poppy

I would not recommend the proposal of existence of the double salt you mentioned. As myself I tried putting some ammonium sulphate within the freshly purified Nd sulfate and when it just precipitated by evaporation it seemed not double salts formed as the Nd sulfate yield was just about the predicted within the calculations.
[Edited on 7-7-2012 by Poppy]

I got the information from a book but it did not state solubility of the salt. There also exists apparently a poorly soluble Cs2SO4.Nd2(SO4)3.3H2O.

Any next batch of Nd salts will be precipitated as K2SO4.Nd2(SO4)3.3H2O (the more insoluble one, apparently), because the single sulphate itself is too capricious for my liking. Then convert to Nd2O3.

Yesterday I had another mishap with a small left over (estimated less than 5 g worth of Nd). I converted it to Nd(OH)3, washed carefully, then dissolved in HCl 36 % (visual estimate of amount needed). I heated that and added what I felt sure was the right amount of H2SO4 50 % and… NOTHING happened. No precipitate at all. Not on cooling either.

Nd prices keep coming down. Reported FOB ex China about $150/kg for metal now.

Poppy - 8-7-2012 at 07:56

Yea I got an article discussing how Nd prices changed over time. To you think China will ever ship such an ammount of a single kilo of the product?
$150 is really cheap, such a good deal. Stock it and wait for inflation. Voila, you gain nothing, lol, but gets big money

Look at this:
"In general, alums are easier formed when the alkali metal atom is larger. This rule was first stated by Locke in 1902:[7] The failure of the sulphate of a given trivalent metal to unite with caesium sulphate to a compound of the type CsMIII(SO4)2·12H2O may therefore betaken as an almost positive indication that such element has no alum-forming power whatsoever."

Attachment: 740798.pdf (72kB)
This file has been downloaded 761 times

[Edited on 7-8-2012 by Poppy]

blogfast25 - 8-7-2012 at 09:26

At least one pleasant surprise: the neodymium sulphate that was hot precipitated yesterday had partly crystallised as the red octahydrate:

(the crystalline matter is about 5 cm across)

But there’s also some of the ‘sandy’ variety. I guess with much patience that could be dissolved in ice cold water and then by gradually taking up the temperature the octahydate might form.

Poppy:

Note that these Nd/alkali metal double sulphates aren’t really alums: they lack crystal water for that.

I’ve seen Chinese suppliers offering MOQ of 1 kg for Nd. But the price of REs is likely to go down further before they go back up again, IMHO… Large 'Chinky' finds, you see. Lucky barstools!

For spot FOB prices see also this:

http://www.metal-pages.com/metalprices/neodymium/

[Edited on 8-7-2012 by blogfast25]

Poppy - 8-7-2012 at 09:46

Why you growing fat crystals full of water and other double salt forming elements at all?
You making us jealous !!!

Just a question: Can you change color of those larger crystals too???

blogfast25 - 8-7-2012 at 10:30

Quote: Originally posted by Poppy

Why you growing fat crystals full of water and other double salt forming elements at all?
You making us jealous !!!

Just a question: Can you change color of those larger crystals too???

The real expert on crystalline neosulphate hydrate is 'Wizzard', he's grown some real whoppers, a couple of pages back!

The hydrate is much more colour stable than anhydrous Nd salts. The trifluoride for instance changes dramatically from incandescent to save bulb light.

Wizzard - 8-7-2012 at 20:55

I've never seen that bright red color! All I have is purple/pink (depending on lighting)... But the pink is beautiful. What lights do you use? Maybe particulate size affects color slightly? I know the smallest crystals of the substance are blue instead of purple, I have not bothered to gauge them under 'pink' lighting conditions.

blogfast25 - 9-7-2012 at 05:15

I don't worry too much about the colour: too subjective; lighting, camera, mass effects etc.

Wizzard - 9-7-2012 at 10:14

I tried taking pictures of my crystals in as many light sources as possible- The neodymium sulfate is so pretty under so many different lights!

blogfast25 - 9-7-2012 at 12:10

Well, look forward to the pix.

Wizzard - 9-7-2012 at 12:21

Have you tried growing any larger-than-normal sized crystals, or other Neodymium salts? I've been disturbingly sidetracked by my K project, and radio controlled airplanes... Summer fun! I still have a box filled with hard drives, waiting for disassembly to get magnets and platters, and aluminum for recycling.

blogfast25 - 9-7-2012 at 13:12

I've grown some large alum crystals some twenty years ago. Nd salts remain a little too expensive for crystal growing projects, IMHO... But the Nd price is coming down fast. It'll be a while before that trickles down to our level, though.

Gibberator - 2-8-2012 at 09:38

I would appreciate your input if you guys have any. I tried Isolating Neodymium Sulfate by directing dissolving the magnets into sulfuric acid, and then precipitating it out by boiling the solution. The only problem I have is that when I went to precipitate it out it didn't. When I heat the solution a very small amount of very fine precipitate does form, but it is nearly impossible to filter out. The only method that has proven any worth is to cool it down and crystallize out the Iron Sulfate except the Neodymium Sulfate is proving very hard to get into solution. After reading this thread and woelen's experiment on Neodymium Sulfate everything I've seen with mine are starting to look very different from the rest. Does anyone have any insight into this?

MrHomeScientist - 2-8-2012 at 09:46

I used the same method you did and it worked out for me. The only issue was I had to do a LOT of boiling to get crystals out. For your problem, your solution may be too dilute - try boiling the solution volume down considerably and see what happens.

It's interesting because you don't want to boil away too much water, because then iron sulfate would start to precipitate as well. My procedure for recovering went like this:
1) Heat solution near boiling until an appreciable amount of Nd-sulfate precipitated
2) Filter the precipitate and let the solution cool
3) Fe-sulfate crystals usually form, decant the solution from these
4) Repeat until you can't get any more Nd, or your solution volume is nearly gone

blogfast25 - 2-8-2012 at 10:20

Personally I much prefer the potassium neodymium sulphate double salt method, which is used industrially and is described in the thread above.

In short:

• Dissolve magnets into strong, hot (> 20 %) HCl. Filter off residue.
• Add fine K2SO4 to filtrate, an amount calculated so that the volume of solution is saturated, at RT, stir well. Insoluble Nd2(SO4)3.K2SO4.3H2O drops out as a sandy, lightly pink precipitate. Leave to stand overnight to maximise yield.
• Filter and wash filter cake (Nd2(SO4)3.K2SO4.3H2O) with small amounts of acidified saturated K2SO4, until all iron has been removed. Then wash with small amount of cold tap water.
• Convert the Nd2(SO4)3.K2SO4.3H2O by treating with strong NaOH or NH3 to Nd(OH)3 and soluble sulphates.
• Filter off Nd(OH)3, wash filter cake with plenty hot, then cold, then dionised water to remove soluble sulphates.

Convert the Nd hydroxide to the compound of your choice. It dissolves readily in HCl or H2SO4.

Wizzard - 2-8-2012 at 11:31

I'm working on a 1L batch of magnet soup as we speak!

Although my exact method isn't know to work in batches larger than 1L (about 200g of magnets), it's fantastically simple:

1. Add crushed magnets to 1:1 mix water and Sulfuric acid, enough to dissolve while hot (at least 60*C).
2. While hot, filter the boron and chromium out. No fancy filter needed.
3. Heat and concentrate the solution slowly- Boiling off the water JUST until the solution is EXACTLY saturated for any one present salt- When crystals start to form, normally at the surface of the water, you've gone too far

4. Take the hot, concentrated solution, and throw it in the freezer (but to not let freeze).
5. Once cooled to just before freezing (if ice has just started to form, you should be OK), pour and filter the solution, leaving behind a large amount of ferrous sulfate.
6. Repeat evaporation till saturation (steps 3-5), and freezer crash one more time.

At this point, you should have a VERY pure Nd salt solution. If you went too far in any previous step, loosly crush the ferrous sulfate crystals while they are moist (they feather and dry out easily)- Under flourescent light, you will see little blue dots of neodymium sulfate- You can pick them out if you're so inclined.

I've made some very large crystals of Nd sulfate this way- Straight out of the boiling pot. Recrystallization is also easy, and recommended... Massive single crystals are very possible, the solution likes large crystals more than making crystal masses.

Gibberator - 2-8-2012 at 12:34

The only problem with that is I have tried putting it in the freezer and adding in lots of extra water but my Neodymium Sulfate does not want to dissolve even with persistent stirring and constant cooling.

blogfast25 - 2-8-2012 at 12:41

Quote: Originally posted by Gibberator

The only problem with that is I have tried putting it in the freezer and adding in lots of extra water but my Neodymium Sulfate does not want to dissolve even with persistent stirring and constant cooling.

It's well known to take a long time to dissolve in iced water. What does your sulphate look like?

Gibberator - 2-8-2012 at 13:46

While it sits in the water, kind of a red-orange colour (in daylight).

blogfast25 - 3-8-2012 at 04:49

Quote: Originally posted by Gibberator

While it sits in the water, kind of a red-orange colour (in daylight).

That sounds like the regular octahydrate.

blogfast25 - 3-8-2012 at 06:06

And talking about neodymium, the past few days I’ve been trying to crystallise some NdCl3.6H2O w/o success. It was a left over sample, only a few gram and as Fe-free as I can get it.

I started off with about 150 ml of yellowish solution and reduced that by gentle boiling to about 15 ml. On fridging no crystals formed. I reduced further on low heat until a thick liquid resulted but still no crystals resulted, not even on icing: only a viscous, yellowish semi-solid. Probably amorphous.

The chloride is of course highly soluble in water, hygroscopic and with a flat temperature-solubility dependence. Frustrating...

I think I might mix it with an excess NH4Cl, then fume off the NH4Cl, in an attempt to get the anhydrous chloride...

[Edited on 3-8-2012 by blogfast25]

MrHomeScientist - 3-8-2012 at 06:08

Blogfast,

Quick question on your potassium double salt method: I dissolved my magnets in sulfuric rather than hydrochloric acid, and still have about 1.5L of 'magnet sulfate' solution. Would the double salt procedure change at all when starting with magnet sulfate rather than chloride? I'd think it would require less K2SO4 to reach saturation, since there are already a lot of sulfate ions in solution.

blogfast25 - 3-8-2012 at 06:13

Quote: Originally posted by MrHomeScientist

Blogfast,

Quick question on your potassium double salt method: I dissolved my magnets in sulfuric rather than hydrochloric acid, and still have about 1.5L of 'magnet sulfate' solution. Would the double salt procedure change at all when starting with magnet sulfate rather than chloride? I'd think it would require less K2SO4 to reach saturation, since there are already a lot of sulfate ions in solution.

I don't think it would matter much: the double salt is much more insoluble than the straight sulphate, so the equilibria predict the double salt should form in the presence of sufficient K2SO4. You could try adding a stoichiometric amount of K2SO4.

The equivalent sodium double salt also exists and is also insoluble. I had a go with NaHSO4 (very OTC) and it seemed to work also. But the K double salt is reportedly the least soluble.

elementcollector1 - 28-8-2012 at 21:49

It's been a long time since I've read this darn thing...
Anyway, I have some neodymium oxalate, partially decomposed to oxide sitting around. Can this dissolve in HCl?
And furthermore (here's the important question!), can I reduce this to elemental Nd using calcium metal? Lithium metal? Unobtainium metal? (Wait, what?)
Your thoughts?

blogfast25 - 29-8-2012 at 03:53

I don't really like 'partially decomposed': the oxalate is very, very insoluble.

Chemical reduction of REs is very problematic. Mostly electrolysis is used to obtain the elements.

elementcollector1 - 29-8-2012 at 10:12

Well, cherry-red heat isn't easy where I live, especially not in experimental settings (like a sterile lab sort of thing). The best I could hope for is to place the mix in a small soup can, and place that inside a larger can with lit coals inside it.
How does solubility affect concentrated HCl? I know the oxide is soluble in hot oxide, but wouldn't the oxalate dissolve too?

Wizzard - 29-8-2012 at 11:23

You may be able to reduce it with Lithium, but the metal melts at 1064*c... It would have to be in a sealed environment of Ar, or He.

blogfast25 - 29-8-2012 at 12:04

Quote: Originally posted by Wizzard

You may be able to reduce it with Lithium, but the metal melts at 1064*c... It would have to be in a sealed environment of Ar, or He.

For chemical reduction of a Nd compound, you need at least to start from the trifluoride (synthesis reported higher up by Mr HomeScientist). Apparently Ca could just about carry it off (Delta G < 0 for the fluoride reduction reaction) but Li does not. Mg falls short too.

blogfast25 - 29-8-2012 at 12:10

Quote: Originally posted by elementcollector1

Well, cherry-red heat isn't easy where I live, especially not in experimental settings (like a sterile lab sort of thing). The best I could hope for is to place the mix in a small soup can, and place that inside a larger can with lit coals inside it.
How does solubility affect concentrated HCl? I know the oxide is soluble in hot oxide, but wouldn't the oxalate dissolve too?

Whether an 'insoluble' compound dissolves in a strong acid depends on some factors, not the least the so-called 'solubility product', Ks. RE oxalates have very low solubility products and thus take a lot more to solubilise.

But an excess of concentrated sulphuric acid will convert even partially decomposed oxalate. Poorly soluble Nd sulphate then forms (give it time and heat to form).

Then wash off the excess acid with hot water and convert the clean sulphate with strong alkali to Nd(OH)3.

elementcollector1 - 29-8-2012 at 14:28

No, I actually think I'll keep them as chlorides for this reaction (chlorides are more volatile than oxides).
Although I have no idea of the reaction enthalpy for either, so I could very well be wrong.
I guess my plan would be to melt the Ca (842 C) or Li (181 C) with the NdCl3 mixed in, and wait for the thermite-type reaction to occur. Is there any way to protect the neodymium produced from oxidation?

blogfast25 - 30-8-2012 at 04:04

It's unlikely to work: even if you get reaction (very doubtful) the chlorides are way too volatile for open crucible conditions.

In a decent metallothermic reduction, the slag would protect the formed metal. Trust me, that ain't going to happen here: the Heats of Formation of most RE binary compounds are too high for reduction to be possible. Only NdF3 + Ca stands a fighting chance.

But if you're gonna mess with fluorides read up on the dangers: soluble fluorides are highl;y toxic.

For example for NdF3 + 3/2 Ca → Nd + 3/2 CaF2

We have: Standard HoF for NdF3 = - 1657 kJ/mol (Wolfram Alpha)

And Standard HoF for CaF2 = - 1226 kJ/mol (NIST webbook)

So for the reaction the Enthalpy of Reaction = 3/2 (- 1226) + 1657 = - 182 kJ/mol (of Nd reduced).

Since as this is negative that indicates the reaction should proceed (ignoring entropic effects) with heat generation. But the heat generated would probably not be high enough to obtain liquid Nd and liquid CaF2, unless the reactor assembly was strongly heated from the outside.

This problem is pretty universal for the REs, although I expect the HoF for the higher atomic numbers to be even more negative.

[Edited on 30-8-2012 by blogfast25]

elementcollector1 - 30-8-2012 at 12:55

What about a coal fire? Those can reach 700-1000 C fairly easily, and immersing the crucible in that kind of heat could lead to auto-ignition, as well as melting the neodymium if hot enough (MP 1010 C).

blogfast25 - 30-8-2012 at 13:14

Quote: Originally posted by elementcollector1

What about a coal fire? Those can reach 700-1000 C fairly easily, and immersing the crucible in that kind of heat could lead to auto-ignition, as well as melting the neodymium if hot enough (MP 1010 C).

Heating to auto-ignition is common practice in some cases of metallothermy. Getting suitable (fine!) calcium is another problem though. It's usually pellets or shavings...

elementcollector1 - 30-8-2012 at 16:59

It needs to be powdered? I...don't think a ball mill would work for this one.
Wouldn't the calcium just melt and start the reaction by itself?

blogfast25 - 31-8-2012 at 05:03

Quote: Originally posted by elementcollector1

It needs to be powdered? I...don't think a ball mill would work for this one.
Wouldn't the calcium just melt and start the reaction by itself?

Yes, but it's always preferred to start from something more homogeneous. An Ar blanket would also help here, to prevent Ca from reacting with air.

So, in short: NdF3 (possibly REF3?) + Ca pellets or shavings + external heating + argon blanket = worth trying.

[Edited on 31-8-2012 by blogfast25]

elementcollector1 - 31-8-2012 at 13:42

Well, time to get some argon, then.
An argon blanket doesn't have to be refilled, right? It just sits there?

EDIT: Neodymium fluoride is insoluble, correct?

[Edited on 31-8-2012 by elementcollector1]

blogfast25 - 1-9-2012 at 05:15

Quote: Originally posted by elementcollector1

An argon blanket doesn't have to be refilled, right? It just sits there?

EDIT: Neodymium fluoride is insoluble, correct?

No, no: an argon blanket needs to be maintained with a constant flow. But a smallish flow should suffice.

NdF3 is water insoluble, thus quite easy to make. For synth. see a few pages upwards.

elementcollector1 - 1-9-2012 at 21:57

So, setup would be something like this:
-A crucible with the reactant material of NdF3 and Ca metal is placed inside a steel can with a coal fire burning in it. A tube (steel? Ceramic?) directs the flow of argon into the reactant crucible, and the argon is removed by (I don't know. It can't reach the coal bucket, otherwise the coal will stop burning).

blogfast25 - 2-9-2012 at 05:24

In your set up the Ar would extinguish the coal fire, because Ar is denser than air. The heat source would have to be OUTSIDE of the steel container.

Lambda-Eyde - 2-9-2012 at 07:29

Wouldn't it be possible to use a flux, like calcium chloride?

elementcollector1 - 2-9-2012 at 19:40

So, how would I lead the argon away from the coal? A vacuum?
Also, the CaCl2 flux is probably a good idea, because no reactions can take place between it and the components, and it melts at 772 C or less (772 C being for anhydrate).

blogfast25 - 3-9-2012 at 06:09

Heat crucible inside a steel box with heat source acting on the box. Argon is the lead into the box, keeping it relatively air free.

Poppy - 3-9-2012 at 19:51

I heard no calcium would suffice to withsand the reducting powder of my gadjet! POWDER OF MY GADJET some1 whos already by means wasted a little effort on putting that to words. Ok.

elementcollector1 - 3-9-2012 at 21:30

Er...what?

Poppy - 3-9-2012 at 22:08

On what grounds you conceive calcium would reduce most other chemicals?

[Edited on 9-4-2012 by Poppy]

blogfast25 - 4-9-2012 at 05:33

Quote: Originally posted by Poppy

On what grounds you conceive calcium would reduce most other chemicals?

[Edited on 9-4-2012 by Poppy]

Specifically in the case of NdF3, see here:

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

But no one is claiming that "calcium would reduce most other chemicals". It's case-specific.

[Edited on 4-9-2012 by blogfast25]

Gibberator - 16-9-2012 at 21:29

Alright, so I have had an interesting development with my most recent attempt. I dissolved the magnets in Sulfuric acid and filtered, you know how it goes. Anyways, after I precipitated the Neodymium as NdSO4*NaSO4. Then I treated it with a fairly strong solution of NaOH, upon which the precipitate turned into the classic light blue of Neodymium Hydroxide. Unfortunately I couldn't get around to filtering it right away, partly to busy, partly just laziness (heh), but when I finally got around to filtering I noticed that the precipitate had begun reverting its colour back to the same as the double salt. I just assumed it had formed some more of the double salt so I added some more NaOH, except the colour remained the same. No matter how many times I wash with NaOH, the precipitate remains the same. I'm not sure if what I have is Neodymium Hydroxide that doesn't resemble any Neodymium hydroxide I have ever seen, or some weirdly nonreactive Double salt. Any insight to what this might be?

blogfast25 - 17-9-2012 at 09:30

I have no idea and would simply start again, this time filtering immediately.

Try also dissolving what you're got now in strong HCl: appearances, especially with neodmium, can be deceptive. If it dissolves it's at least part hydroxide.

The sulphate of neodymium is a strange compound, there seem at least to exist two forms of it, see also the neodymium thread by 'woelen' for nice details and write up.

elementcollector1 - 17-9-2012 at 12:06

So, I have that run of partially decomposed oxalate sitting pretty in a beaker of HCl. Do I need to add H2O2, or will this stuff dissolve on its own?
Also, how do I get rid of the newly formed oxalic acid?

blogfast25 - 21-9-2012 at 11:28

I'm not sure it will dissolve. If it does, add NaoH to precipitate as Nd(OH)3.

If it doesn't, try conc. H2SO4, that should convert it to Nd sulphate. Decant off excess acid and wash a few times with hot water.

elementcollector1 - 21-9-2012 at 12:32

The supernatant solution turned as black as if it had a fresh Nd magnet in there, but there is still plenty of precipitate.

blogfast25 - 23-9-2012 at 11:54

Black???

Decant (or filter) it off and set aside (but don't throw).

Try attacking the rest with conc. sulphuric acid. It should then turn pinkish/red.

elementcollector1 - 24-9-2012 at 19:03

Apparently the black color is due to the extremely dark purple color of the Nd (III) ion in this case, when a flashlight is shined upon the solution it shows purple. Akin to permanganate, maybe.

The acid hasn't finished attacking the magnet, and now the magnet itself is covered in gray sludge (presumably undissolved metal).
If I add NaOH to this purple solution, I should get pale lavender Nd(OH)3, correct?

blogfast25 - 25-9-2012 at 02:28

Check the colour of the 'black' stuff under incandescent light and TL light: it should look quite different under these sources if it is neodymium chloride.

Wizzard - 16-6-2013 at 17:02

How's everybody's Neodymium isolation going?

A brief update of my own - I gathered all of my now well-dried small crystals, close to 30g of them (although I did not weigh it), and decided to purify and recrystallize them into as large of crystals as I can.

On adding them to 300mL of 1*C pure water ("Infant" grade: steam purified, deionized, no additives etc), and waiting nearly 3 days for it all to dissolve, my first stem I knew was I had to clean out the present iron sulfate impurity...

I first tried temperature swing, which had worked so well for me in the past... But there's so little of it, it was not effective at all. Humbug.

So I thought I would let the partially saturated liquid heat and slowly evaporate, leaving a more saturated liquid for me to try to chill to drop out the iron sulfate.

After 2 days of this (I chose a very slow evaporation), liquid being held at approx. 55*C, I was reminded of my early difficulty growing Iron Sulfate crystals from powder... Without the addition of sulphuric acid to lower the pH, it reacts with the air to form insoluble (and not particularly flocculent) iron hydroxide! It sinks and sticks to the bottom of the vessel (for the most part) and my solution will be easily decanted and filtered. The neodymium compounds remains in solution, but I will be testing the filtrate for neodymium content.

I'll be posting pictures of the crystals as soon as the batch is done... However, I'll be growing them rather slowly.

I can take a picture of the vessel as it is now - Lovely lavender on top, pukey yellow-brown on bottom. And I will add, the lavender color of the Nd(III) is enhanced with purity.

[Edited on 6-17-2013 by Wizzard]

blogfast25 - 17-6-2013 at 04:46

Hey Wizzard!

Pictures would always be appreciated.

Wizzard - 24-6-2013 at 08:53

Here's one

The cluster of 2 crystals is more than 2" across, and still growing. Today, the junk crystals started growing... I think I need slower growth. But evaporation is certainly my method of choice.

http://imgur.com/FBn637q

Also, of note - ANY agitation causes nucleation and sporatic crystal growth. Do not rotate the vessel for pictures

That happened the other day - The mess of small crystals are the result.

[Edited on 6-24-2013 by Wizzard]

Poppy - 25-6-2013 at 11:09

For the growth of large neodimmy crystal, I can't believe you guys have gone this far thats an absurd, do you just heat it up or carefully control the temperatures the whole time?
I developed a method for growing large ferric ammonium sulfate crystal by adding elemental aluminium, slowly, into a saturated solution of the later, bringing strange shift in the equilibrium which ultimately cause slow temperature-constant precipitation onto a single crystal. That might work with neodymium, as changing the temperature might cause different lattices to overlap over the crystal as it grows, leading to imperfections.
Too bad my mon throwed the heavy load of dimmy's magnets I had so imma kind not inclined to use my 20g of material in this task as purification proves a real pain when dealing with Al contamiantions!!

Wizzard - 25-6-2013 at 13:39

Actually, it's all in the preparation. I work out of a storage room, I don't even really have a countertop. My "hot plate" is a 17W candle heater. I have a vent immediately behind my work surface, and I keep the area as dust and cat hair free as humanly possible.

Using the absolutely cleanest recycled labware, I recrystallize and purify the magnet soup, removing most of the iron sulfate in a few cycles, the boron within 2. From this, I crystallize the mixture, disposing of the remaining dissolved Nd and Fe sulfates once I cannot stop the Iron Sulfate from crystallizing at high temperature.

From the resulting solids, I manually extract the Nd sulfate crystals. They were small, but easily picked up and collected.

From this, I dissolved them over the course of 3 days in alternating freezing/cold steam-purified water, after a quick hot rinse to remove any clingy Iron sulfate. I used only lab-grade Coffee filters, but I am meticulous about keeping EVERYTHING clean. No fingerprints, breathing, etc.

Once the crystals are all dissolved, the mixture is filtered again, into it's final container. This container is topped with a 1-ply paper towel, to keep ANYTHING out, but slowly let the hot water vapor out.

The container is put on the hot plate, which will heat the liquid to around 60*C. I insulate the container with a sheet of insulation, to keep the condensing liquid on the walls to a minimum. The dripping can be problematic.

Over the course of nearly a week, the crystals will grow. Without agitation, and with as few controlled nucleation sites as humanly possible, you will get very large, single crystals on the bottom of your container.

Mine grew for a few days, with only 2 crystals, before I attempted to rotate the container to snap a pic... This action is not recommended, I soon had a city of crystals, but at least it was not a blanket

As of today, I am re-dissolving the crystals into a splash more water, to try single crystal growth again. Wish me luck!

Renewed Interest

MrHomeScientist - 19-1-2014 at 16:07

After a very long hiatus, I uncovered some information that is extremely interesting and makes me want to have a go at this project once more. I'll get to that in a later post. In the meantime, I tried out the neodymium double-sulfate separation method and it looks like it works great. I read through the whole thread again and I guess I never tried this route myself before.

I started with 15mL of "magnet soup" leftover from the last time I pursued this experiment. It's been sitting on the shelf for a year or so (!), in a flask covered with parafilm. As you recall from earlier posts, this soup consists of iron(II), iron(III), and neodymium sulfate from dissolving magnets in sulfuric acid. Not having any potassium sulfate handy, I tried out sodium sulfate instead.

I made up a near-saturated solution of Na₂SO₄ by dissolving 7.4g of it into 25mL distilled water. Some very fine powder was left over that refused to dissolve, so I let that settle and pipetted the liquid above it for the following procedure.

The sodium sulfate solution is on the left, and "magnet soup" on the right.

I added about half of the Na₂SO₄ solution to the magnet liquor. Nothing happened immediately, but after a few minutes a light pink sandy precipitate slowly appeared and settled fairly quickly. I left this to settle/react overnight.

The next day, a layer of brown gel had formed on top of the Nd precipitate!

The liquid tested as a pH of 5 to pH paper, as I suspected it might. I added a few drops of ~50% H₂SO₄ to bring pH back down to 1, stirred vigorously for a few minutes, and allowed to settle. This effectively dissolved the brown goo (iron hydrolysis, most likely).

Here it is after adding the acid.

This was filtered off and rinsed with several small portions of acidified (with sulfuric acid) Na₂SO₄ solution, then a final rinse with cold distilled water. I collected 0.6g of the neodymium double salt, Nd₂(SO₄)₃ * Na₂SO₄ * 3H₂O . Unfortunately, I don't have any thiocyanate to test for iron in the precipitate, but it looks quite nice.

Next, I treated this powder with sodium hydroxide to convert the neodymium to the hydroxide. I believe this goes via:

Nd₂(SO₄)₃*Na₂SO₄*3H₂O + 6NaOH --> 2Nd(OH)₃ + 4Na₂SO₄ + 3H₂O

Accordingly, 0.6g of the double salt should require 0.2g of NaOH, and produce 0.32g Nd(OH)₃ and 0.47g Na₂SO₄. Based on the solubility of the sodium sulfate and required amount of sodium hydroxide, I made up a solution of 0.8g NaOH in 12mL of distilled water. This is a large excess of hydroxide, to ensure full conversion of the neodymium.

I poured this solution directly onto the dry double-sulfate powder, and stirred vigorously for ~1 hour. I then allowed this to settle for 2 hours. The precipitate had lightened into a white color with a hint of pink. I filtered this off and rinsed with a small portion of hot water, followed by a few rinses with cold water. After drying, I recovered 0.3g of Nd(OH)₃ - a near-stoichiometric yield! (My scale is only accurate to 0.1g, though)

Here's a photo of the hydroxide:

So in conclusion, I concur with blogfast: this appears to be an excellent way to separate the iron from the neodymium! I should be getting some thiocyanate in the near future, and when I do I'll make a more quantitative determination.

Brain&Force - 19-1-2014 at 19:20

I'm going to try making some neodymium salts soon from some neodymium magnets I was careless with and shattered. I also have access to calcium, magnesium, and sodium fluoride so I may be able to do a Nd thermite (which someone has already done) - that is, if I can find a place to do it.

Why not make some unusual Nd salt? I haven't seen much on the acetate, citrate, benzoate, perchlorate, bromide, iodide, or carbonate.

Other than surface area reasons, is removing the nickel coating really necessary? NiSO₄ is soluble in water, so it shouldn't interfere with Nd precipitation.

MrHomeScientist - 19-1-2014 at 22:07

Wow that's a godawful video. No description, terrible focus, no idea what's going on. He could have filmed just about anything and claimed it was an Nd thermite...

The problem with the fluorides you listed is they are basically insoluble in water, so you can't really use them to arrive at NdF₃. You need a soluble fluoride salt; in my case, ammonium bifluoride. (EXTREMELY DANGEROUS!)

You need to remove at least some of the nickel coating because nickel doesn't react with sulfuric acid in these conditions (maybe in hot concentrated acid). So you need to expose the magnet material itself to allow the reaction to proceed. If some did dissolve, then you're correct in that it shouldn't interfere.

Brain&Force - 19-1-2014 at 22:54

By calcium and magnesium I meant the metals, not the fluorides. And I thought sodium fluoride was soluble in water.
All my magnets are chipped and corroding at the minimum, so I should be OK with just leaving them in acid.

blogfast25 - 20-1-2014 at 06:02

Quote: Originally posted by Brain&Force

I also have access to calcium, magnesium, and sodium fluoride so I may be able to do a Nd thermite

That won't work, trust me.

Reduction of the fluoride NdF3 with lithium might just about work at elevated temperatures, going by the Heats of Formation.

MrHS already tried Nd trifluoride + Mg w/o much success.

The RE oxides are almost impossible to reduce chemically. Aluminum certainly does NOT do it. The guy in the video is a plain old LIAR: the heats of formation of Nd2O3 and Al2O3 simply do NOT allow such a reduction to take place.

NaF: a fairly useless chemical, max. solubility at RT about 1 M. You really need ammonium bifluoride. Quite a dangerous chemical, read up and suit up before use.

@Mr HS: that Nd hydroxide look very much the real deal! Lovely pink. Never got mine quite that good...

[Edited on 20-1-2014 by blogfast25]

Brain&Force - 20-1-2014 at 10:10

Quote: Originally posted by MrHomeScientist

This same exact thing happened to me with the terbium extraction (but with potassium sulfate). It seems the double sulfates have a tendency to supersaturate. Conveniently, most of the K-Tb sulfate came in a single flat bed of crystals. Did this happen with the Na-Nd sulfate?

This thread seriously needs to be stickied.

blogfast25 - then I guess thermite is out of the question. I'll just try to make some of the salts mentioned previously.

blogfast25 - 20-1-2014 at 12:48

A few words on the metallothermic reduction of NdF3 with Li.

The Standard Heat of Formation of NdF3 is – 1657 kJ/mol (Wolfram alpha), a real whopper!

For LiF NIST gives – 616.93 kJ/mol.

So for NdF3 + 3 Li === > Nd + 3 LiF3 the Standard Heat of Reaction would be – 194 kJ/mol. Ignoring any entropic effects, it’s safe to say this should proceed, if you can get it to light.

But – 194 kJ/mol is almost certainly not enough to heat the reaction products to above the MP of Nd (1024 C, the MP of LiF is 845 C), needed to allow metal and slag to separate out, thermite style. So external heating would be needed to ensure the end temperature is, say, 1100 C or higher.

NdF3 is easy to prepare, both Mr HS and me have done so: it’s water insoluble.
Lithium is easy to obtain, but expensive and hard to handle.

A graphite crucible loaded with Li metal on the bottom and finely powdered, dry NdF3 on top, perhaps rammed a bit to reduce air, then heated quickly to about 600 C, may well give the desired result. Heavy, lump Nd metal covered with LiF slag is what you’d want here, post cooling.

MrHomeScientist - 21-1-2014 at 09:43

I think it's still worthwhile to pursue the magnesium thermite route, just with a larger charge than I did previously. I now have a charcoal furnace capable of melting aluminum, so I could also get to a higher temperature than before. Metallothermic reaction with lithium metal is a little scary to me - molten lithium is just about one of the most aggressive things there is, so your choice of crucible would need to be pretty careful.

Another really interesting method I came across uses a setup that refluxes magnesium metal onto magnet chips directly - the molten metal being able to dissolve Nd but not Fe or B! It collects in a crucible at the bottom of the apparatus, where Mg boils off and Nd remains. Really cool, but quite likely out of our league.

Attachment: Direct Extraction and Recovery of Neodymium Metal from Magnet Scrap.pdf (257kB)
This file has been downloaded 897 times

=====================================
=====================================

Now, for the idea that renewed my interest in this project. Completely unrelated to neodymium, I was browsing around for new things to try. I found a thread here on deep eutectic solvents, and that sent me on a quest of discovery that I'll sum up for you.

Deep Eutectic Solvent (DES) thread here on SM:
http://www.sciencemadness.org/talk/viewthread.php?tid=10529

I thought it'd be neat to make one of these room temperature "ionic liquids," specifically choline chloride & urea. In the reference link in the first post of that thread, it states that many metal oxides are soluble in these ionic liquids. Scanning through other links in that thread, I found the attached PDF that shows the electrodeposition of aluminum from AlCl₃ dissolved in such a liquid.

Attachment: Conductivities of AlCl3 in Ionic Liquid Systems and Their Application in Electrodeposition of Aluminum.pdf (191kB)
This file has been downloaded 1708 times

Further, ShadowWarrior4444 (the author of the DES thread) claimed in this thread that "as for Ionic Liquids, one of them should be able to dissolve" NdF₃.

All of these bits of information got me very excited. I thought it just might be possible to electrodeposit neodymium metal from a solution of NdF₃ in a choline chloride/urea DES.

Now, I don't have much experience with electrolysis. Looking at the standard electrode potentials, here's what I think would happen using Al electrodes. (starting with Al because that's what they used in the PDF above)

Nd³⁺ + 3e^- <---> Nd ........... -2.323V
Al <---> Al³⁺ + 3e^- ............... 1.676V

Overall Cell Potential: E_cell = E_cathode - E_anode = -2.323 - 1.676 = -3.999V

So the cell would need to be run at 4V, using Al electrodes. This metal is usually unsuitable for electrolysis, though, so I'd prefer to use graphite rods as I've done in the past. I'm not sure what the anode reaction would be in that case, though.

My concern is that I don't really know the reactivity of this solvent system, so I'm not sure if Nd will react with it in some way. I already received my choline chloride (smells awful, reminds me of triethylamine), and the urea is on the way. I can't wait to start experimenting with this.

DraconicAcid - 21-1-2014 at 09:59

Quote: Originally posted by MrHomeScientist

Now, I don't have much experience with electrolysis. Looking at the standard electrode potentials, here's what I think would happen using Al electrodes. (starting with Al because that's what they used in the PDF above)

Nd³⁺ + 3e^- <---> Nd ........... -2.323V
Al <---> Al³⁺ + 3e^- ............... 1.676V

Overall Cell Potential: E_cell = E_cathode - E_anode = -2.323 - 1.676 = -3.999V

So the cell would need to be run at 4V, using Al electrodes. This metal is usually unsuitable for electrolysis, though, so I'd prefer to use graphite rods as I've done in the past. I'm not sure what the anode reaction would be in that case, though.

My concern is that I don't really know the reactivity of this solvent system, so I'm not sure if Nd will react with it in some way. I already received my choline chloride (smells awful, reminds me of triethylamine), and the urea is on the way. I can't wait to start experimenting with this.

You're mixing conventions here. If you're going to subtract one half-reaction potential from the other, then they either both have to be reductions, or both have to be oxidations. If you've written one as an oxidation and the other as a reduction, then their signs are different, and you have to add them, giving something like -0.6 V.

As you've written the half-reaction, the Nd one does not want to go, but the Al one does. The overall reaction will be somewhere in the middle.

Those standard electrode potentials- are those the ones from aqueous solution? They won't apply at all in a different solvent.

[Edited on 21-1-2014 by DraconicAcid]

MrHomeScientist - 21-1-2014 at 10:56

Bah you're right - I got them from a table of standard potentials here, which is probably aqueous. I guess that means I'll have to slowly raise the voltage until I see something start to happen?

As for the equations, they were both listed in the table as reductions. I flipped the aluminum equation and changed the sign on the potential before inserting them into the equation. I was under the impression you needed to do this to form a complete chemical equation for the cell, before you could plug numbers into the potential equation. At the cathode Nd is plated out, while the anode dissolves into solution.

I appreciate the help. Like I said, I'm a real beginner at electrochemistry. I haven't been able to find a decent tutorial on the subject online either, so if anyone knows of a good one I'd appreciate it.

[Edited on 1-21-2014 by MrHomeScientist]

DraconicAcid - 21-1-2014 at 11:11

Quote:

Bah you're right - I got them from a table of standard potentials here, which is probably aqueous. I guess that means I'll have to slowly raise the voltage until I see something start to happen?

Yes.

Quote:

As for the equations, they were both listed in the table as reductions. I flipped the aluminum equation and changed the sign on the potential before inserting them into the equation. I was under the impression you needed to do this to form a complete chemical equation for the cell, before you could plug numbers into the potential equation. At the cathode Nd is plated out, while the anode dissolves into solution.

That's the trouble- there's two ways of looking at it. The electrochemist's view is that reduction potentials are the only thing anyone would ever look at, so you subtract one from the other. The thermodynamicist's viewpoint is that you flip one half-reaction, add the equations, and *add* the half-potentials (just like you would add enthalpy changes, Gibb's Free Energy changes, etc.). Both will give you exactly the same answer (and exactly the same "you're doing it wrong" sneer from the other group), but if you change the sign *and* subtract, you'll get the wrong answer.

Consider this:

A -> A⁺ + e E = +1.0 V
A⁺ + e --> A E = -1.0 V

If you add these up, you get a null reaction (A -> A, nothing happens), which must have E = 0 V, right? You can get this by a) adding the reduction potential to the oxidation potential (+1.0 V + -1.0 V), or b) subtracting one reduction potential from the other (+1.0 V - +1.0 V). Either way will give you the correct result. BUT changing the sign of one reduction potential (to make it negative) and then subtracting will give you 2.0 V, which is wrong.

[Edited on 21-1-2014 by DraconicAcid]

[Edited on 21-1-2014 by DraconicAcid]

blogfast25 - 21-1-2014 at 13:12

Mr HS:

I don't think lithiothermic reactions are any scarier than magnesiothermic ones.

Remember though that I made a mistake calculating the thermochemistry for NdF3 + Mg. On correction we found that basically ΔH = 0 for that reduction. There's no getting around that.

A graphite crucible (see gold smelting) should resist molten Li well, bearing in mind also that there won't be any Li for very long. It's quite a 'common' reduction agent for higher fluorides (tri or tetra).

blogfast25 - 22-1-2014 at 10:27

Well, well, well. I ran the numbers for the adiabatic end temperature of NdF3 + 3 Li === > Nd + 3 LiF and got a fairly crude estimate of 1100 to 1200 C. For a thermite that would be abysmal but here it's fairly close to target...

MrHomeScientist - 22-1-2014 at 13:36

Hmmm...
I do have some small graphite crucibles, and a strip of Li foil from a lithium battery. Do you think it would be sufficient to cram the foil into the bottom of the crucible, pour powdered (dry) NdF₃ on top, and heat with a propane torch until some kind of ignition? My Li is stored under oil, though, and surely I'd want to get rid of that somehow.

The crucible should be totally dry and free of any contaminants for safety. I'd read elsewhere that molten lithium can rip the oxygen right out of glass!

DraconicAcid - 22-1-2014 at 13:48

I want to hear about what happens in the choline-urea melt....

MrHomeScientist - 22-1-2014 at 14:10

Me too!

I received my urea in the mail yesterday, so I'm set to start down that road as well. First I'd like to dry my choline chloride, though - it came as pretty damp white crystals. I'll put some in a dessicator bag with calcium chloride (Damp Rid) tonight.

One thing that bothers me, though: numerous sources list choline chloride as completely nontoxic (indeed, it's an additive in animal feed), yet the MSDS from Sigma gives it a 2 health hazard. My sample has a pretty 'scary' smell, which reminds me of tiethylamine. Might there be some contamination in my sample? I hadn't seen anything referencing it having a smell.

blogfast25 - 22-1-2014 at 14:19

Quote: Originally posted by MrHomeScientist

Hmmm...
I do have some small graphite crucibles, and a strip of Li foil from a lithium battery. Do you think it would be sufficient to cram the foil into the bottom of the crucible, pour powdered (dry) NdF₃ on top, and heat with a propane torch until some kind of ignition? My Li is stored under oil, though, and surely I'd want to get rid of that somehow.

The crucible should be totally dry and free of any contaminants for safety. I'd read elsewhere that molten lithium can rip the oxygen right out of glass!

My guess is that that should suffice. Ram the NdF3 powder down onto the Li to reduce oxygen availability to the Li. And propane should be enough: I doubt if ignition doesn't start from about 500 C or so.

Re. liquid lithium and glass, yeah, ok. A bit extreme though, isn't it? And I'd still like to see proof of it: lithium isn't used often in oxidic reductions...

The deep-eutectic idea is of course wonderful but unchartered territory (as far as I know). Be a trail blazer!

By electrolysis, which anion will you expect to be decharged?

[Edited on 22-1-2014 by blogfast25]

DraconicAcid - 22-1-2014 at 14:28

My sample also has a bit of a smell. I didn't try drying it.

blogfast25 - 22-1-2014 at 14:30

Quote: Originally posted by DraconicAcid

My sample also has a bit of a smell. I didn't try drying it.

Have you prepared a deep eutectic with it?

DraconicAcid - 22-1-2014 at 14:43

I tried a choline chloride-copper(II) chloride eutectic, which didn't work properly.

Brain&Force - 22-1-2014 at 15:22

Why not try electrolysis of a soluble neodymium salt in pyridine? Apparently lithium can be made with that method, so Nd should be possible in theory. However, the salt would have to be perfectly dry for it to work properly, and lanthanide salts are pretty much impossible to dry without thionyl chloride or heating in a HCl atmosphere.

Interesting 1919 reference that states NdCl₃ is soluble in pyridine: http://books.google.com/books?id=W2zxN_FLQm8C&pg=PA116&a... There's a lot of useful info in these older references.

blogfast25 - 23-1-2014 at 05:56

@B&F:

I think you need anhydrous NdCl3 for that. The anhydrous form is tricky to prepare from the hexahydrate. Once I tried with NH4Cl and only got a mess.

Edit: oops, you said so.

[Edited on 23-1-2014 by blogfast25]

MrHomeScientist - 23-1-2014 at 14:59

I started my experiments with the deep eutectic choline chloride / urea mixture last night. I managed to make the liquid, but haven't done anything else so far. Tough to do stuff during the work week, especially when its been so cold outside lately!

Until I have something more relevant to neodymium come out of that route, I'll be posting my eutectic experiments in that thread instead of this one. Here's the link: http://www.sciencemadness.org/talk/viewthread.php?tid=10529&...

Brain&Force - 23-1-2014 at 15:15

If I dissolve hydrous NdCl₃ in pyridine, will I be able to dry it with calcium oxide? I can't find any info on calcium hydroxide's solubility in pyridine though.
Unfortunately I cannot acquire pyridine; even if I am able to, I can't use it in my school - the substance is banned. Very odd, because I've heard that there was this one teacher somewhere who would denature alcohol with a drop of pyridine.

NdCl₃ is soluble in alcohol (44.5g/dL), but I don't see how one could produce Nd metal in such solution.

MrHomeScientist - 30-1-2014 at 21:43

Finally got my potassium sulfate in the mail, but it's highly impure. Off white color with black specks throughout; turns out it's fertilizer grade. I guess that's what I get for ordering "organic sulfate of potash." At least it was cheap - I'm recrystallizing some now and then I'll be able to proceed.

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