The trouble with neodymium...

Alright, I'm first going to admit I have absolutely no formal training in chemistry- I'm strictly a math major

But I noticed over a weeklong evaporation operation- The nd sulfate fell out first, in both cases, so much so that it was able to crystallize, and appears nearly Fe free.

Iron sulfate is VERY soluable at ~95*C, the Nd sulfate is not... Am I the only person evaporating the water out slowly at just below boiling, causing the Nd to drop out selectively?

I can post some pictures, I'm working on narrowing down the time spend evaporating- So far, I manage to drop out all the ND sulfates, and then the iron sulfates right down on top of it (but the layering is very distinct).

Quote: Originally posted by Poppy

Would this dry Fe2O3 work too???

Quote: Originally posted by Wizzard

But I noticed over a weeklong evaporation operation- The nd sulfate fell out first, in both cases, so much so that it was able to crystallize, and appears nearly Fe free.

Welcome to the murky world of hobby chemistry, mathman!

Do NOT let appearances fool you: at least with the double salt (presumed Nd2(SO4).K2SO4.2H2O) I’ve obtained a nice, pink product which on further scrutiny contained… substantial amounts of iron! There seems to be co-precipitation going on. Whether that is also the case with ‘naked’ Nd2(SO4)4 I cannot vouch for. Test by converting your sulphate to hydroxide by treating with strong alkali. Filter off the precipitate and wash it, then dissolve it in clean (iron free) HCl. Add a bit of 9 % H2O2 to oxidise any Fe2+ to Fe3+. Test visually or with KSCN.

I wonder also about the iron content of Nd oxalate precipitated from iron rich solutions, see Mr Home Scientist higher up and kmno4 in a separate Nd thread...

£$£$£$£$£$£$

Well, here’s my ‘magnet chloride’ after heating and standing overnight:



Plenty of precipitated Fe(H)3 but the dark red-brownish supernatant liquid is the tell tale that plenty of Fe didn’t drop out. pH was less than 1 at that point, so no surprise there.

After that the work became more of a rescue operation than a chemical separation!

The slurry filtered only with great difficulty and adding some freshly precipitated Fe(OH)3 did not bring the pH up enough to precipitate the remaining dissolved iron. I had to resort to manually adjusting the pH to pH >= 4, first with small amounts of 5 M NaOH, the final adjustment with household ammonia. At that point the slurry became much thicker and filtered better. The filtrate contained all the Nd, but still tested slightly positive for Fe3+, mainly due to a bit of colloidal Fe(OH)3 running through the filter I believe. The raw NdCl3 has now been precipitated one last time and will be redissolved in HCl tomorrow.

In conclusion I’d say that separating Nd3+anf Fe3+ by playing to the Fe(OH)3 equilibrium works best for removing relatively small amounts of iron, not to remove iron in concentrations 7 times or larger than that of Nd3+.


[Edited on 13-11-2011 by blogfast25]

Will test once I get home It all dissolved in 15cc of very cold water (1-2*C), about .5g of the physically extracted Nd salt (scraped from the bottom- It's much softer than the very hard iron sulfate. Very faint purple color. I'm working on getting the rest out, by watching the evaporative process carefully.

I only was seeing purity because the Nd salt crystallized- I know some purity is needed for this. There were small 2-3mm crystals among the mess on the bottom of the container.

Mother aqueous solution was 'magnet sulfide', with the nickel COMPLETELY removed manually.

Then left to dissolve, in full, the remaining slush of unknown material and boron removed. (pic 1)

From this, total evaporation with small addition of distilled solution, but still acidic (I did not test pH).

From this, the layered substance in photo 2.

Scraping the purple Neodymium Sulfate into a glass (a small sample amount), dissolved and nearly frozen, then dried on a hot plate over the course of 6 hours, the resulting pure crystals formed- Picture 3 in 4 parts

My method seems low yield, but very easy! I prefer manual means- Not so much measurement (but I would like to learn). I'm a scientist for sure, but a chemist I am not!

The largest (centered) crystal cluster is about 7-8mm across.

@Blogfast- Ferrous, Iron(II) of course


[Edited on 11-14-2011 by Wizzard]

So, if I understand well you simply slowly evaporated the ,magnet sulphate' until crystals of Nd2(SO4)3 started to appear? Certainly the Nd sulphate is much less soluble than ferrous sulphate but it's interesting you managed to isolate the minority constituent that way and of seemingly good quality too. AFAIK no one else has applied that approach here yet. Personally I had trouble dissolving the magnet in 50 % H2SO4, not sure why...

I would still urge you to test one crystal for presence of iron. If contaminated you could recrystallise for instance, or did you already do that?


[Edited on 14-11-2011 by blogfast25]

Quote: Originally posted by Wizzard

Yes, after filtering out the insoluables, the Nd fell out of solution before the Fe did with slow evaporation at 90-95*c. I will test for iron once I get my hands on some good hydrogen peroxide... My lab chemical supply is still expanding.

Well I know from experience that ferrous sulfate oxidizes with contact to air in solution, unless kept mildly acidic... I've grown 1cm crystals in 2cm tubes I'm no chemist, but I do know a thing or two about keeping one's ingredients pure- I've been growing crystals of various salts since I was in high school.

And I was unaware that pharmacy grade H2O2 is good enough! Off I go!

Attached good macro pic, with scale.



[Edited on 11-14-2011 by Wizzard]

Very nice crystals indeed: they look good enough to eat!

Do you recall:

* weight of the magnet
* amount and concentration of the sulphuric acid used?

Did you get any early precipitation (of Nd2(SO4)3) during dissolving?

Acidity slows the oxidation of Fe (II) down, that's very true. But you should have seen my acidic stockpile of FeCl2 crystals after a couple of weeks: it's a brown mass of Fe(OH)3, ferric oxychlorides and what not: I bet it would test negative for Fe (II)! Just for standing in open air...

Quote: Originally posted by blogfast25

Fenton's reagent doesn't involve Fe (IV).

Quote: Originally posted by barley81

Really? I thought it did, after reading this: http://woelen.homescience.net/science/chem/solutions/fe.html

I'll run another batch, fully documented

I have stacks of old HDDs and all the motivation I need for beautiful, hexagonal air semi-stable (left one on the paper overnight- no water loss!!) color changing crystals of an element I have very little of in my collection.

Thanks! Very small yield was due to my physical method of separating the Nd salt from the Fe salt- I merely scraped the easiest-to-get-at soft/wet and purple neodymium sulfate crystal from the hard, green ferrous sulfate.

Low yield, yes, but I'm still extracting the salts from that aqueous solution- This was just a quick and dirty trial. I'd also still say there's 50% more in my evaporation dish- There was a good quantity of smaller crystals, but I do need to remove the film which grew on the edges before I reattempt this (and recycle the dried and extracted but not removed Nd sulfate).

Partial evaporation yielded the following crystals- Such a lovely color! Recrystallizing tonite- It doesn't take long. Note the green tinge in the center mass, and along the edges- This is the start of when the iron sulfate crystallizes- Time to stop evaporating!

[Edited on 11-14-2011 by Wizzard]

My lighting conditions are a normal 60W incandescent, and big-bulb commercial flourescent lighting. Perhaps I'll get a sunlight shot today- That will be nice for comparison.

My are certainly a bit more pink/red than those purple/pink ones you link to- But I will say I cant wait to grow a single crystal that size

[Edited on 11-15-2011 by Wizzard]

On the the leftover liquor:
After separating in part the neodimium from iron from a magnet solution via the sulphate method, a leftover supernatant was obteined which seemed to contain, by its color, initially at least a big portion of Fe III sulphate, Fe II sulphate, unnoticeable by eye, and some neodimium as well.
Theoretical obtention of Nd2(SO4)3.8H2O would be slightly round 50g, but as there was obteined just 44g (oops, thats really close to the goal, and because some 5g were lost propositally =P) I am inclined to believe some neodimium was left in the supernatant solution (at least 6g) as well as some iron is present in my final powder.
The main reason is that the leftover supernatant gradually, and very slowly, formed crystals which took much longer than conventional ferrous sulphate crystals to form (6 hours for a supersaturated solution of FeSO4 and one week to badly precipitate some of the strange liquor).
Here you can see the leftover after about 2 days:




Here you can see the leftover solution after about 1 week:




Those crystals were filtered and dissolved again to see what happens: the result was a slightly pink solution from those green crystals!!!!







That after decanting (seeing pink depends on your mood on that day...)

The green crystals initially become decoloured when washed
then gives the solution above.

So there must be double salts of Nd, Fe in the solution which modifies how the precipitation of pure ferrous sulphate crystals would do, and my obteined Nd salt powder is highly contamined with Pr or Fe in order of having its yellow color, even after it was thorughly washed with hot water when it was prepared.

And this is the Nd sulphate under flash and daylight, theres no real difference.


[Edited on 11-15-2011 by Poppy]

Quote: Originally posted by Poppy

On the the leftover liquor: After separating in part the neodimium from iron from a magnet solution via the sulphate method, a leftover supernatant was obteined which seemed to contain, by its color, initially at least a big portion of Fe III sulphate, Fe II sulphate, unnoticeable by eye, and some neodimium as well. Theoretical obtention of Nd2(SO4)3.9H2O would be slightly round 50g, but as there was obteined just 44g (oops, thats really close to the goal, and because some 5g were lost propositally =P)

Quote: Originally posted by DerAlte

All references I have seen say Nd sulphate is red (my emph.) and Sm is light yellow, Ce white (colourless); and these were the colors I got. All these sulphates are very difficult to dissolve at 0C, the most soluble point. Once dissolved, heating to 100C usually will precipitate them. Better crystals can be obtained by slow evaporation. Regards, Der Alte

Quote: Originally posted by Wizzard

It's quite possible there's some other Lanthanide impurity involved there- Neodymium magnets dont need absolute purity, as far as the Nd is concerned, I think (and was this also mentioned previously in this thread?).

Quote: Originally posted by Wizzard

Just reading through some other threads, came across a gas chemistry link, this may be of interest to this thread- I think it may be applied to Nd2O3 to possibly generate Nd metal! http://mattson.creighton.edu/PipetteRxn/index.html

Unfortunately, the reduction of Nd2O3 with hydrogen is far from possible. It is possible with the oxides of Fe, Co, Ni and a few others but not with the REs.

Allow me to explain. Take a generic reduction by hydrogen:

MO_n + n H₂ === > M + n H₂O

This can only proceed if the change in Gibbs Free Energy, ΔG = ΔH - T ΔS is negative (< 0): systems seek a state of minimal Free Energy. Assuming for argument’s sake that the entropic effects for this reaction are small (so T ΔS ≈ 0), then ΔG ≈ ΔH < 0.

ΔH can be found by adding the Heats of Formation of the reaction products and subtracting from this the sum of the Heats of Formation of the reagents. In the cases of Fe2O3, CoO and NiO this ΔH is indeed negative and these reactions can proceed.

But in the case of Nd2O3 its HoF is too large and ΔH > 0 (in reality we should take the entropic effects into account). It’s also for this reason that Nd2O3 cannot be reduced by aluminium or magnesium, both far stronger reducing agents than hydrogen.

REs are generally produced by electrolysis. But reduction of the trifluorides with magnesium is also possible (for the same reason: ΔH < 0). It’s this latter reaction Mr Home Scientist and me will endeavour to produce Nd metal. If we ever get round to making enough NdF₃, that is!


[Edited on 16-11-2011 by blogfast25]

In the midst of all this talk about separation and purification, I wanted to get back on track with the final step of this process to obtain Nd metal. I have a few crops of NdF3 already made, so I wanted to try one out to see how it went.

These are my samples of NdF3 (under fluorescent lighting), each vial from a different batch of magnets. The one on the left is the one I showed in a previous post (distinctly green), the middle one is a very nice white, and the righthand sample is a bit more brown. The rightmost one is what I used in this experiment - since it's a bit more contaminated it was a good candidate for initial experimentation.


That righthand vial contained 7.1g of NdF3. I mixed this with 1.4g of Mg powder - somewhat more than the stoichiometric amount, but I wanted it in excess. I placed this in a fused silica crucible in which the reaction was to take place.


Since the total mass of the reactants was only 8.5g, I decided that normal thermite procedures wouldn't be sufficient (as per an earlier suggestion). So, I placed this crucible on a ring stand and heated it from below with a propane torch. I was hoping to get to the initiation heat just from this, but after about 15 minutes of blasting the bottom of the crucible with the torch there was no change. So, to get things started I grabbed a small strip of Mg ribbon with tongs, lit this with the torch, and dropped it into the hot crucible. This did the trick! The whole interior lit up with a nice orange glow and a good amount of smoke as the reaction took place. It reminded me of chromium thermite I've done in the past - lots of light and smoke but no sparks. I was very excited to see something happen Here's what it looked like shortly after the smoking stopped. You can see the ring stand and blue propane bottle beneath it.


I waited about an hour to let this cool, then took it inside to see what I could find. Disappointingly, I didn't find much at all. The vast majority was just a dark grey powder.


Clockwise from top left, here's what each pile is.
Top left (TL): crust from top of the reaction. It's mostly white MgO.
TR: Bulk of the material below that. It's pretty uniformly grey.
Bottom right (BR): Much blacker material from the bottom of the reaction products.
BL: A small amount of scrapings from stuff stuck to the bottom of the crucible.


Now here's where things get interesting.

In scraping away the black material from the bottom, I revealed an extremely thin layer of a very shiny metal!


I scraped some of this out and picked out some of the larger pieces to put in a vial to see them better. "Larger" of course meaning extremely tiny. This can't be more than 50mg. There were more shiny dots in the bottom left pile of the previous picture.


What's encouraging is that these flakes are not attracted to a strong magnet! That means they can't be iron, so I tentatively think that these are actual (miniscule) chips of Nd metal!
Some of the rest of the reaction material did seem to be affected by a magnet, so as I suspected there's certainly some iron contamination in this batch. Magnetite could explain some of the black coloration.

Now to test this metal, I placed a small amount of the black powder with shiny chips in a test tube and added a few drops of sulfuric (battery) acid. It immediately produced a good amount of gas and a large amount of heat. Some of the black material didn't dissolve. I was hoping to get a pink solution to confirm Nd, but it still looks colorless to me. It was such a tiny amount though I bet it'd be impossible to see if it was there anyways. I'd think testing with oxalate should show me something, even at low concentrations. I have yet to try this. It's hard to say if I actually made Nd metal because it's nearly impossible to separate the shiny bits for testing alone since they are so tiny. What's in the vial is by far the best of what I could collect.

So I suppose this confirms the difficulty of getting this reaction to work at small scales. I imagine the whole mess didn't get hot enough, despite my additional propane torch heating, to react all the way or produce molten metals. Something obviously happened to the reaction mixture, but it just didn't make an appreciable amount of metal. Just a lot of grey and black powdery garbage, plus that promising tiny amount of shiny metal.

Second-round extraction and recrystallization of my magnet sulfate soup produced 5.45g of neodymum sulfate, plus .50g in 2 large single crystals- pics attached

The small one is 7x4x3mm, the large one is 9x6x6mm.

Quote: Originally posted by Poppy

I'm jealous of your crystals

Quite possibly. I dont take advantage of other people ignorance, however Plus I dont know if this sulfate is toxic or not.

Quote: Originally posted by Wizzard

Quite possibly. I dont take advantage of other people ignorance, however Plus I dont know if this sulfate is toxic or not.

Quote: Originally posted by blogfast25

You claim to see MgO but the main reaction product should be MgF2 of course. I suggest in any case to grind down the slag and treat it with strong HCl. If the reaction took place then the dark material is likely caused by fine Nd metal dispersed through sintered MgF2. It should still react with strong acid though. With an MP of 1263 C your orange glow may have been just a little short of melting the MgF2 down.

Quote: Originally posted by blogfast25

One possibility is that the metal you found is unreacted Mg (Fe seems unlikely, as it was only a contaminant) and the only way to prove it is neodymium would be to try and dissolve it in as little H2SO4 as possible and see if you get a precipitate at BP, or by saturating the solution with K2SO4. If you go Nd2(SO4)3, it can be tested with oxalic acid (redissolve it in iced water first). MgSO4 is of course highly soluble over a wide range of temps. But it’s unlikely to be Mg, IMHO…

Quote: Originally posted by blogfast25

The other point is that I found two literature references claiming NdF3 is purple/violet. I looked at your procedure again and believe your product(s) are relatively pure samples of neodymium trifluoride and that your colour is the right one, bar maybe slight distortion by FeF3. But that should have stayed in solution because it’s soluble though, bar a little occlusion/co-precipitation…

Quote: Originally posted by blogfast25

And here’s the strange news that I was going to U2U about. ... It really means that if the reaction proceeds AT ALL, strong external heating would be needed to get the metal and slag to melt and separate (whereas various thermite-like reactions generate enough heat to achieve this) But it seems to me you’ve provided invaluable evidence that the reaction does proceed, albeit w/o generating much extra heat, otherwise you’d have obtained a neat slag puddle with the metal at the bottom.

Argon may work well. Just don't try nitrogen or CO2 Speaking of which, anybody know anything about lanthanide nitrides?

Quote: Originally posted by Endimion17

If you're about to try that melting scenario, consider argon, because if you don't use it, you will fail miserably, meaning the yield will be dreadfuly small, and you'll end up with your precious starting batch heavily contaminated.

Nope. Failure is by no means guaranteed w/o argon. See aluminothermic reactions of which I carried out more than a dozen on different metal oxides (Fe, Cu, Si, Ti, Cr, Mn, V, Nb and combined ones as well): despite end temperatures of around 3000 K, you usually get good quality metal (if you get everything right! - they’re not as easy as you might think) because the once molten slag (alumina) forms a perfect protective barrier for the metal regulus. No need for argon here neither, assuming you can get the MgF2 to actually melt (1263 C, not even that hard).

@Mr Home Scientist:

Yes, I suggest trying to recover as much Nd by extracting it from the slag. Nd and Mg should be easily separable with the saturated K2SO4 method. Careful not to use too strong acid (32 % HCl will be OK) or your MgF2 will start acting like fluorite: a HF generator!

Way to go with the furnace. I’m thinking of carrying out my reaction with my dinky toy charcoal furnace: that melts copper but only just…

There is however another ‘theoretical’ aspect that I haven’t touched upon.

Assume we heat a stoichiometric mix of NdF3 and Mg to above the MP of MgF2 and reaction occurs. Now we hold the crucible there at that temperature: in the melt the equilibrium 2 NdF3 + 3 Mg < === > 2 Nd + 3 MgF2 rules.

The equilibrium constant is given by Nernst:

ΔG = - RT lnK with ΔG the change in Free Energy (at temperature T) and K the equilibrium constant (at T):

K = ( a_Nd² x a _MgF2³ / ( a_NdF3² x a _Mg³

With each a the activity of each species (we cannot use concentrations here)

We would have to calculate ΔG (at T) from the Standard change of Free Energy, ΔG^{298 K} but examples with similar reactions show that the temperature correction, even for T > 1000 K, is usually small.

Although we don’t know the exact ΔG, we know from the above that it is likely to be quite small. Well, from Nernst we know that if ΔG ≈ 0, then K ≈ 1 or in plain English: there should be considerable free Mg in the metal phase. I thus predict that obtaining high purity Nd with this method is probably impossible.

In practical news, I’ve finally finished precipitating, filtering and washing the first batch of NdF3. Tomorrow it will be dried. So far it is very clearly lavender coloured.

And another 25 g of magnet material has been dissolved in HCl. I’m going the ‘sat. K2SO4’ route again, just to explore that a bit more. The double salt has been precipitated and will be further processed, hopefully tomorrow…


[Edited on 18-11-2011 by blogfast25]

After filtering and drying (on a hot plate), I obtained a disappointing 6.3 g NdF3. TY should have been about 20, I was hoping for about 10 g and got 6. On the bright side, the product is clearly pink/lanender in tungsten light and switches quite dramatically to this eerie green under TL. With what I’ve made today I should have about 10 g or more, enough for an ‘all in’ attempt.

So with the last magnet (28 g) I revisited the K2SO4/Nd sulphate double salt method, to try and clear a few things up.

Everything was uneventful at first. I set the first filtrate (after precipitation of the Nd double salt) but before any washing, aside as ‘Filtrate A’, then washed with cold, saturated K2SO4 and obtained this ‘pretty in pink’: the NdK double sulphate (hydrate):



But it was during washing I stumbled on a mistake I made in the previous runs and which explains I found a lot of iron in the washed double sulphate: without controlling the pH in the washing solution iron becomes oxidised to Fe3+ and at the higher pH (4 - 5 or above) starts dropping out. That then get caught in the filter cake and there’s your contamination! Remedy: simply acidify the washing solution of sat. K2SO4 slightly.

The liquid obtained from washing was labelled ‘Filtrate B’.

I then proceeded to test both A and B for presence of Nd with oxalate. Obviously this can’t be done in the presence of ferrous material because FeOx is poorly soluble (and there’s a lot of Fe). Therefore 50 ml of both solutions were oxidised on ice bath with 9 % H2O2, simmered to destroy any excess H2O2, then calculated amounts of oxalic acid dehydrate were added. This complexes the Fe3+ to soluble trioxaloferrate (III) (FeOx₃^3-. This for instance is B after that treatment: the trioxaloferrate complex is emerald green:



What is very interesting is that in neither the cases of A nor B after treatment does any insoluble neodymium oxalate precipitate, none whatsoever, indicating the separation is complete.

The double sulphate was then converted to Nd(OH)3 with KOH, filtered and washed and dissolved in 10 ml of glacial acetic acid and 50 ml water, chosen because my HCl contains traces of Fe3+. The hydroxide dissolved well in that solvent but needed a little heat. And that caused precipitation! First I thought Fe(OH)3 but it’s whitish. Remnants of Nd sulphate? A hydrolysis product of Nd acetate? To be figured out tomorrow…

A slightly better video:
http://youtu.be/hY51xRO7SjA

Anybody else still working on this? I've accumulated almost 10g of Nd2(SO4)3

Quote: Originally posted by Wizzard

A slightly better video: http://youtu.be/hY51xRO7SjA Anybody else still working on this? I've accumulated almost 10g of Nd2(SO4)3

Quote: Originally posted by blogfast25

Nope. Failure is by no means guaranteed w/o argon. See aluminothermic reactions of which I carried out more than a dozen on different metal oxides (Fe, Cu, Si, Ti, Cr, Mn, V, Nb and combined ones as well): despite end temperatures of around 3000 K, you usually get good quality metal (if you get everything right! - they’re not as easy as you might think) because the once molten slag (alumina) forms a perfect protective barrier for the metal regulus. No need for argon here neither, assuming you can get the MgF2 to actually melt (1263 C, not even that hard).

That's true for larger batches, but the amounts we're dealing with here are laughably small. If it was measured in kilos and heated in cylinders (like uranium was initially made) in furnaces, it could end up well because layers would form, assisted by the weight of the mixture.

Things I see here, though very interesting and commendable, produce dispersed, impure metals poorly embedded in the crust, prone to oxidation unless argon is applied... Not that there's anything wrong with it ...

Neat thing I tried- I reduced the Nd2(SO4)3 with an amount of KI of equal molar, in about 1cc of water, and I believe K2(SO4) (clear, soluble), NdI2(green) and Nd2O3 (pale blue purple precipitate, turns blue under flourescent light) was produced by the colors of the precipitates and solution.

Evaporating now I believe all of the Nd2(SO4)3 was consumed.

Quote: Originally posted by kmno4

To make it clear, for reaction: 3Mg + 2NdF3 -> 3MgF2 + 2Nd enthaly = +14 kJ gibbs = -7 kJ You can forget about thermite-like reaction ( but the reaction can give Nd-Mg alloy)

Just a quick update, I haven't forgotten about this! It's getting colder here and I tend to go into hibernation, so posting's probably going to slow down a bit.

I built a crude furnace out of fire bricks and a propane torch that I hope will be able to hold the heat in and get me to the temperature required for this reaction. The lid is on in the picture, but of course below that is the rectangular heating chamber. I drilled a slanted hole in the bottom of the chamber that the torch nozzle sticks through to point directly at the crucible.

I'll try melting aluminum with it this weekend and see if it does anything for me. Nothing's mortared together or anything, so I'll certainly lose a good bit of heat through the cracks but we'll see how it does.

Brand news:
The neodimium sulfate obteined from the sulphate method was relatively impure because of its slight yellowish color but not purple ~ dissolving it in water was improductive as it would take a week stirring. So to check for coloration differences a little hydrochloric acid was poured just to see if it would dissolve the sulphate and reveal its propoerties under flashtube and tungsten lamps: it did. The solution is yellow green under FL light and olive brown under tungsten lamp.
The interesting part is: different eyes are differently opened for different views. While showing the substance to my grandma, she noticed the precipitate (because just a little HCl was added to the Nd sulphate contaminated salt) was from white (under FL light) to purple (under Tungsten lamp light). The changed runned unnoticed to me until she told, so I must assume her eyes were much better prepared than mine..
As a consequence I supposed the following reactions took place unpredictedly:
Nd2(SO4)3 + HCl --> NdCl3
As a result the predicted Nd/ Fe double salt contaminants reacted with HCl but, slowly, the Nd went back along the sulphate anions and re-precipitated, while all the iron seemed to stay in solution in the form of iron II chloride together with some Nd chloride.
Very luckly, this way a purer Nd2(SO4) was believely obteined. I recall it luck because the dissolution of insoluble salts with acids is not supposed to be obvious, but follows some pH rules to happen, so I was lucky to add enough HCl into it to happen the so said reactions.
The complete analisys of the accomplishments can be described with analitical chemistry formulae. I was studying it recently so I dispose myself to publish it here as soon as possible, along with pictures of the product I have.

Till next c

Quote: Originally posted by LanthanumK

Upon dissolution of a neodymium magnet in hydrochloric acid, precipitation with sodium bicarbonate, oxidation with hydrogen peroxide, and redissolving the precipitate, a blood red solution is formed. Is this a coordination complex with iron?

Well, some updates from me!

Grew out my second batch- 60g of magnets, 500mL water and 100mL 98% sulfuric yeilds about 24g of neodymium sulfate, and an unknown amount of waste material!

I have extracted 18.5 grams into LARGE crystals, the largest of which I have photographed here. I have about 4 g of material being reprocessed and purified. They are a very, very beautiful rose red when large and transparent (in sunlight).

Notes:
1. In the initial purification by crystallization, small Nd sulfate crystals grew, which were then overtaken by pink crystals, which did not change color- I do not know what they were. I think I saved some, but the rest went to reprocessing.

2. While the large batch was purifying (slow, high-temp evap), there was a growth of very light, flakey crystals, like Mica, when gathered, looked silver/white. I have saved a small vial of these crystals, what did not fly away when I gathered them from the top and sides of the vessel, outside of the water line (where water may have collected as steam and evaporated).

3. Also noticed while evaporating, the mixture took a dark brown appearance, but the Nd sulfate crystals grew just the same.

4. Crystal growth over 5 days stopped on the 5th- At this point, the solution was nearly devoid of the Nd sulfate, but was still not oversaturated with Fe sulfate.

I've also made Nd iodide, if anybody is interested Only a very small amount, but the shift in color of the crystals from green to red (in different lighting) is quite lovely.

1cm grid in photo- Small bagged crystals are the largest from the previous batch. All crystals seem to split on their own- Unknown cause- maybe they are naturally twinned?

To extract like I do, I first take the soup, add some hot water, heat to 'warm' (so as little as possible solubles are left out of solution), and filter with a fine filter paper.

Then take as much Fe sulfate out as possible- After first heating, seal the vessel (not airtight, just enough to little vapor travels in r out) and stick it in the freezer- LARGE crystals of Fe sulfate will grow, and then pour out the rest, add a bit of water, filter, and start the next step.

This filtered liquid is then heated to about 90-95*C, and a thick paper towel put over the top- The heat of the liquid pushes Fe sulfate solubility way up, and Nd sulfate way down- As the liquid SLOWLY loses water through the paper, the Nd sulfate is all pushed out of solution as evaporation occurs, and the Fe sulfate is left in, not quite saturated

Quote: Originally posted by Wizzard

@Poppy- I've found the heating and evaporating the solution will not separate the two sulfates at any temperature if the mixture is too acidic.

@Blogfast- Yes, like paint I have a small vial of the material, perhaps I should test it's composition... But I would not know where to start.

@Blogfast- It doesn't need to be filtered, it can be isolated in the evaporation condensate on the sides and top of the vessel, free from water. I'm very sure it's 100@ 'shiny' Picture in a bit.

New crystal pictures- Evaporative growth over 2 weeks (in 2 stages to remove iron) lead to these crystals from the smaller, 'scrap' Nd2(204)3 stock I have prepared.

The largest single crystal at the bottom, double-terminated (but flat on one side) is 1.8cm in length!

All large crystals of this Nd sulfate show the same fracturing seen in previous large crystals- I do not know what causes this, but the previous largest crystal (I would guess 2.2cm!!) broke on removal from the growing glass.

And I can get some pics of the Nd iodide soon, also- I'll be making another batch! The oxide is also a by-product, which is just as pretty.

Poppy, can you describe your ammon. Nd solution? I'm building a beautiful color-changing display for my RE salts

Photos of 99.9999%-iron free Nd sulfate. The ammoniacal salts remained in the crystalization liquor.
Neodimium sulfate under large fluorescent tubes:


[img]http://www.sciencemadness.org/talk/files.php?pid=240121&aid=17879[/img]

" under small fluorescent tubes:

[img]http://www.sciencemadness.org/talk/files.php?pid=240121&aid=17882[/img]

"under tungsten lamp lighting:

[img]http://www.sciencemadness.org/talk/files.php?pid=240121&aid=17884[/img]



[Edited on 3-8-2012 by Poppy]

[Edited on 3-8-2012 by Poppy]

Quote: Originally posted by Poppy

Photos of 99.9999%-iron free Nd sulfate. The ammoniacal salts remained in the crystalization liquor.

Blogfast,

Ammonia comes into nothing, apparenttly, it was used to neutralise a fraction of the excess sulfuric acid

The purity is a relative 99.9999% iron free (which means no iron at all), not 99.9999% pure Nd sulfate. The acidic washing bath with sulfuric acid should have reached this, or maybe its just wrong?

I reallty guess it was too straight an assumption, but considering that by througly sinkings the products allowed for a high quality trustworthy final product.

99.9999% was simply too much way out the strathosphere.