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blogfast25
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Object lesson: never calcine anything you might later want to redissolve. 
What colour is your oxide?
[Edited on 8-5-2014 by blogfast25]
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Brain&Force
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If it's already seperated, then why not dissolve it in hydrochloric acid now?
When you add hydrochloric acid, advise me on whether excess hydrochloric acid causes the solution to become more lavender in color. I saw this in
woelen's site, and I want to confirm its occurence.
http://woelen.homescience.net/science/chem/exps/neodymium/in...
At the end of the day, simulating atoms doesn't beat working with the real things...
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elementcollector1
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Color: White, white, and more white.
Given that sulfuric acid is doing precisely nothing to it, I would assume hydrochloric will be equally ineffective.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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aga
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Sulphuric Acid !
So, if i electrolyse CuSO4 back into elemental copper and sulphuric acid, would that not work with
Nd2(SO4)3, precipitating Nd out of the solution ?
One way to find out i guess.
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elementcollector1
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Heh. Nope. Look up the reactivity of neodymium - it's not friendly with water, to put it one way.
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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blogfast25
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Quote: Originally posted by aga  | Sulphuric Acid !
So, if i electrolyse CuSO4 back into elemental copper and sulphuric acid, would that not work with
Nd2(SO4)3, precipitating Nd out of the solution ?
One way to find out i guess.
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Relatively few metals can be successfully plated out of watery solution, Nd is DEFINITELY not one of them: it's very electropositive and reacts with
water.
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MrHomeScientist
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Yeah if it were that easy this entire thread wouldn't exist...
I plan on processing more of my magnet soup via the double sulfate method this weekend. Re-reading this thread has got me excited again. I want to try
both the lithium reduction and deep eutectic solvent electrolysis. For lithium, I'm thinking harvest the foil fresh from a new battery and use it
immediately. That way I avoid having to deal with getting rid of the oil on my current sample.
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aga
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Well, i tried it, but couldn't even dissolve the Nd2(SO4)3 even using an ice bath.
The aim is to recover elemental Neodymium right ?
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MrHomeScientist
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That's my aim, at least.
Dissolving Nd-sulfate takes forever, even in cold solution where it is most soluble. Just crush it up as finely as possible and leave it stirring for
an hour or so, if I remember right.
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blogfast25
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It's much more soluble on ice but the rate of dissolution is likely to be slower as well.
Aga: recovering elemental Nd is very difficult. So far we know from literature that electrolysis of the molten chloride works. Little else seems to be
possible, going by what's been discussed in the various Nd threads.
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Töilet Plünger
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I still attempting reduction of a dehydrated neodymium salt in something like pyridine, if at all possible. The problem is that a lot of neodymium
salts are hydrated or hygroscopic. (And how would you produce the molten chloride if it tends to decompose when heated? Would it be possible to
electrolyze the oxide?)
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blogfast25
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If it wasn't clear, I meant of course the electrolysis of anh. NdCl3, in a eutectic mixture.
Acc. wiki that's how mischmetal is (was?) produced, from the mixed anh. LnCl3.
There's little reason to believe NdCl3 isn't stable enough: if it wasn't we'd be able to reduce it by chemical means.
Electrolysis in non-watery electrolytes remains on the table. Think deep eutectics?
Oxides: find a suitable solvent. Cryolite won't work here, I think
[Edited on 9-5-2014 by blogfast25]
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Brain&Force
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The problem with cryolite is that if you reduce it, you would probably get an Al-Nd alloy. (Which may be suitable if I study yttrium-scandium-aluminum
alloys at some point.)
What about dissolving neodymium oxide in neodymium fluoride? I'm sure that's beyond the scope of our equipment, but it may be possible (I don't know
if a sodium fluoroneodymate [that is an AWESOME chemical name] could exist).
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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I don't think the Cryolite would be reduced, if you use the right voltage. But do we know it forms a eutectic with it, like alumina does? I don't
think so.
NdF3/Nd2O3? Well, firstly look for eutectics of that system. The MP of NdF3 is about 1400 C (acc. http://www.phelly.com/ndf3/) Don't they use NdF3 in special glass? There may be some clues there...
The Ln don't seem very prone to forming complex anions.
[Edited on 10-5-2014 by blogfast25]
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MrHomeScientist
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Sort of thinking out loud here, so forgive me if this sounds rambling.
I'm processing more of my "magnet soup" neodymium and iron sulfate solution via the potassium double salt method, and initially got similar results as
my test at the top of page 12. This time I combined 180mL of magnet soup with about 200mL of saturated potassium sulfate solution - after several
minutes, a sandy pink precipitate slowly formed, and after letting it sit overnight a thin layer of tan gunk settled on top. I again attempted to
dissolve this by adding sulfuric acid (thinking it must be iron hydrolysis products), but the tan precipitate remains. Maybe I just need more acid?
The solution is quite acidic already, though.
I also recently got some NaSCN, so I tried my hand at testing for iron in the filtrate from rinsing the precipitate with acidified
K2SO4 solution. Initially this showed no color change at all, but I added some hydrogen peroxide and the test tube immediately
turned deep red. So amazingly enough it looks like iron(II) survived in solution for over a year and a half (!) and in fact there is barely any
iron(III) present at all. Acidity doesn't interfere with the test for iron, does it? Maybe adding hydrogen peroxide to the precipitate will make the
tan stuff soluble? This is going somewhat less smoothly than the sodium sulfate test I did, unfortunately.
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blogfast25
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| Quote: Originally posted by MrHomeScientist | This time I combined 180mL of magnet soup with about 200mL of saturated potassium sulfate solution - after several minutes, a sandy pink precipitate
slowly formed, and after letting it sit overnight a thin layer of tan gunk settled on top. I again attempted to dissolve this by adding sulfuric acid
(thinking it must be iron hydrolysis products), but the tan precipitate remains. Maybe I just need more acid? The solution is quite acidic already,
though.
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It's better (and also mentioned in the lit) to add the potassium sulphate directly to the solution, as a fine powder. That's what I do. Calculate the
amount to achieve saturation at RT. Then heat and stir for about 1/2 hour and allow to cool. This way you're not affecting the pH of the solution,
which risks precipitating ferric thingymejibs.
You may have to start over: alkalise with strong ammonia to convert double salt to Nd(OH)3 + K sulphate, filter off and dissolve solids in HCl, add
new K suphate.
Remember that the concentration of iron in your soup is quite high: that makes it even more prone to hydrolysis, very simply put:
Fe3+ + 3 OH- < === > Fe(OH)3
... is pushed to the right by high [Fe3+].
[Edited on 23-5-2014 by blogfast25]
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Brain&Force
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In my work with terbium I add potassium sulfate directly to the solution as crystals. It works fine and a needle-like coating of the double sulfate
grows on the crystals.
When I did this with terbium I had a clear coating of the gunk on top of the solution. It may be the neodymium that is forming this layer.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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I doubt it: the Nd double salt is white to pink too.
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aga
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Neodymium from magnets is simple.
All you need is a some Spin.
Magnet, acid ... Nd Sulphate.
You got the Nd didn't you ?
The Fe got lost yeah ?
The SO4 is a Free Bonus !
Vote Aga at the next election. Vote Spin.
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MrHomeScientist
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I'm still working on this, but in the meantime I wanted to say that wow, adding solid potassium sulfate to the magnet solution is
definitely the way to go. I figured people that added the solid were only concerned with keeping the volume as low as possible, but blogfast is right
about the need to not change the pH!
I took 50mL of magnet soup (Nd and Fe sulfates) and added in just over 5g of solid K2SO4 with stirring. Within just a few
minutes, the solution had gone totally opaque with pink precipitate of the neodymium double salt. I continued stirring for 40 minutes, then filtered.
I washed the precipitate with 3 portions of acidified, saturated K2SO4 solution, followed by 2 portions of cold distilled water.
I captured the last water wash and tested for iron by adding NaSCN and a small amount of H2O2 to oxidize the iron(II) to
iron(III). This test came back positive, giving an orange color. I did several more water washes and tested again, with the same result. I'm now
letting the precipitate dry on the filter paper, after which I'll scrape it off and vigorously stir in some distilled water to hopefully dissolve the
remaining iron. I might need to acidify this a bit if I end up with iron hydroxides, etc. after drying.
TLDR: The precipitate from adding solid K2SO4 to the magnet solution appears much, much cleaner than my previous trial of adding
a saturated solution instead.
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blogfast25
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MrHS:
Strange how hard you find it to get rid of the last bits of Fe. I would try washing the dry precipitate with small amounts of hot 20 % H2SO4.
Distilled water won't get rid of ferric iron: it'll be highly insoluble Fe(OH)3 by then...
"followed by 2 portions of cold distilled water" is pointless: all you do is precipitate the last bit of iron as Fe(OH)3. Think about it: at
the stage where the Nd is still as double sulphate a bit of free potassium sulphate doesn't hurt at all. So there's no need to wash with pure water.
[Edited on 27-5-2014 by blogfast25]
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MrHomeScientist
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blogfast: I see, thanks for the input. I thought I was following the procedure you used way back earlier in the thread, but I could
be mistaken. I could also be mistaken on how I'm doing the thiocyanate test. I take my sample and add a few drops of sodium thiocyanate solution
(conc. not measured, just dissolved a few crystals in several mL of water). No reaction. I then add a small amount of 3% hydrogen peroxide to oxidize
iron(II) to iron(III), and the solution turns a dark-ish orange. I believe this is a positive result, but everywhere I've read states the color of
this complex should be very dark, blood red.
I found a procedure (attached) that uses potassium permangante to oxidize the iron instead. I imagine they use this so the end of oxidiation is clear
("Stop adding potassium permanganate drops when the purple colour persists for several seconds after addition"). Contrary to what I just said above,
there's also a picture (Fig. 4) of several standard solutions of the complex - mine is usually about the color of the tube second from the right. Even
they use "blood-red" as a descriptor, which isn't at all how I would label those colors.
Attachment: Determination of Iron by Thiocyanate Colorimetry.pdf (168kB)
This file has been downloaded 934 times
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blogfast25
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It's slightly unusual you're getting orange. Try this. Add peroxide to the iron but no thiocyanate yet. Heat a bit, allow to cool again or tap cool,
then add thiocyanate.
When extracting Nd from magnet chloride with potassium sulphate, it helps to work kind of fast: don't allow the FeCl2 to oxidise to Fe3+. The Fe2+ is
easier to get out because FeCl2 doesn't hydrolyse so much.
So, dissolve magnets in HCl, filter while still hot, add potassium sulphate in the right amount to filtrate, filter off double sulphate and take it
from there.
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MrHomeScientist
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My magnet solution has been sitting on the shelf for over a year now, so I think I'm beyond the point of working fast 
I took your advice from earlier and started over. Unfortunately I realized after doing all of the following that I forgot to take pictures!
Recovered 20.2g of iron-contaminated double sulfate. To break the double salt, I used 18.2g KOH in 200mL water. Stoichiometrically, only 8.6g is
required, but I wanted a large excess to ensure full conversion to the hydroxides. The amount of water used was calculated to ensure all potassium
sulfate was able to dissolve. This solution was added directly to the double sulfate powder, and stirred for ~1 hour.
Immediately, the solution turned coffee brown, indicating severe iron contamination. After stirring, this was filtered off to yield a chocolate brown,
sandy solid of iron and neodymium hydroxides. The solution after filtering was water clear.
I then convered these hydroxides back into sulfates using 6.5mL conc. H2SO4 in 280mL H2O. This was calculated
starting from the 20.2g of double salt, rather than drying and weighing the mixed hydroxides themselves. After ~15 minutes of stirring, there was a
small amount of solid that stubbornly refused to dissolve. This was filtered off to obtain 283mL of clear, tan colored (under lab fluorescents)
solution.
To saturate this again with potassium sulfate, I weighed out 29.0g K2SO4 and added the solid directly to the solution with
stirring. Stirring continued for ~1 hour. The precipitate that formed was light pink and very clean-looking (compared to the last contaminated batch).
This was filtered and washed with several portions of acidified K2SO4, and the filtrate captured for testing.
I used a different strategy for testing this time. I still tried NaCSN to test for iron(III), but also took a second test tube and used potassium
ferricyanide to test for iron(II). The iron(III) test came back negative, and over time this solution went milky from colloidal sulfur. I'd guess the
acidity decomposed the NaCSN. The iron(II) test came back weakly positive, with the yellow solution slowly turning somewhat green. I washed twice more
with acidified potassium sulfate, then tested this again. It was an even weaker positive, taking all night to eventually turn a light green. I think
this is an acceptable level of iron contamination, since I believe this test is supposed to be pretty sensitive.
It looks like I've finally been able to achieve effective separation of Fe and Nd!
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The Volatile Chemist
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Congradulations! Are you going to 'blog' about it? Or are there no pictures to show?
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