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blogfast25
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Well, look forward to the pix.
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Wizzard
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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.
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blogfast25
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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.
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Gibberator
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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?
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MrHomeScientist
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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
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blogfast25
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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.
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Wizzard
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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.
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Gibberator
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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.
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blogfast25
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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?
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Gibberator
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While it sits in the water, kind of a red-orange colour (in daylight).
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blogfast25
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That sounds like the regular octahydrate.
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blogfast25
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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]
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MrHomeScientist
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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.
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blogfast25
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| 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.
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elementcollector1
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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?
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blogfast25
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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.
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elementcollector1
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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?
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Wizzard
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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.
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blogfast25
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| 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.
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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.
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blogfast25
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| 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.
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elementcollector1
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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?
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blogfast25
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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]
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elementcollector1
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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).
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blogfast25
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| 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...
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elementcollector1
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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?
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