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Author: Subject: The trouble with neodymium...
Endimion17
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[*] posted on 22-11-2011 at 15:17


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 ...




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Wizzard
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[*] posted on 22-11-2011 at 17:16


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.
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blogfast25
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[*] posted on 23-11-2011 at 06:03


Endi:

I doubt if you've run many a thermite reaction: I've done reactions with as little as 20 g of batch size and obtained metal of very good quality. Bigger is definitely better and gives higher yield but kilos just aren't necessary at all.

AFAIK, nuclear grade uranium metal is still produced by UF4 + 2 Mg. And yes, for optimum results vacuuming the reactor and back filling with low pressure argon is recommended. But it's not essential.

The REAL problem here is not argon/no argon but the low reaction enthalpy. I had wrongly banked on about (-) 300 to 400 kJ/mol.




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kmno4
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[*] posted on 24-11-2011 at 10:13


To make it clear, for reaction:
3Mg + 2NdF3 -> 3MgF2 + 2Nd
enthalpy = +14 kJ
gibbs = -7 kJ
You can forget about thermite-like reaction ( but the reaction can give Nd-Mg alloy)
Values taken from experimetal papers (by DOI: 10.1021/j100792a041 and 10.1063/1.438183), not from somebody's assumptions or "Tungsten Gamma" similar sites.

[Edited on 24-11-2011 by kmno4]
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blogfast25
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[*] posted on 24-11-2011 at 10:58


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)


Which is what I've been saying: I computed a reaction enthalpy of - 29 kJ/mol Nd.

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

[Edited on 24-11-2011 by blogfast25]




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MrHomeScientist
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[*] posted on 8-12-2011 at 20:23


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.

furnace.jpg - 88kB
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blogfast25
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[*] posted on 9-12-2011 at 05:49


Testing with aluminium is a good idea.



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Poppy
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[*] posted on 11-12-2011 at 15:13


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
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LanthanumK
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[*] posted on 12-12-2011 at 09:39


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?



hibernating...
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blogfast25
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[*] posted on 12-12-2011 at 10:09


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?


Not really. It's simply concentrated Fe<sup>+3</sup> which tends to hydrolyse to FeOH<sup>2+</sup> (which gives it the colour) in high concentrations. Add a lot of oxalic acid to that solution, heat and the iron will complex to FeOx<sub>3</sub><sup>3-</sup> which is green and soluble and the neodymium will precipitate out as the insoluble oxalate. That's one separation method...




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Wizzard
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[*] posted on 12-12-2011 at 12:17


Noticed while processing my second batch- The nickel coating has much copper in it. I reiterate my suggestion to remove it as fully as possible.

My 60g batch of magnets (latest) had only about 1/2 the coating or less removed, and the magnet sulfate soup is a different shade- I hope the copper sulfate doesn't become a problem :C
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[*] posted on 13-12-2011 at 05:41


Copper isn't very soluble in H2SO4 or HCl without an oxidiser present. Small amounts of CuSO4 or CuCl2 should remain in the mother liquor, as both salts are highly soluble.

Judging a solution of coloured salts is very subjective: concentration is a big factor.

[Edited on 13-12-2011 by blogfast25]




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Wizzard
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[*] posted on 23-12-2011 at 08:35


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?

DSCN0311.JPG - 244kB
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MrHomeScientist
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[*] posted on 23-12-2011 at 09:55


Wizzard,

Nice crystals! I'm currently recrystallizing some of my Nd-sulfate as well, and its forming just as yours look. I have also noticed the color change in the solution when hot - when heated to near boiling, my 'magnet soup' became much darker, nearly black, and lightened again after cooling.

I'm curious - what separation method did you use to get your Nd-sulfate from your magnet solution? A number of ways have been discussed. Sounds like you heated it below boiling to speed evaporation, but if you have any more details I'd be interested. I'm still searching for the easiest way to separate good purity Nd-sulfate from the rest of the magnet crud.
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Wizzard
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[*] posted on 23-12-2011 at 10:47


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 :)

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[*] posted on 23-12-2011 at 13:00


The darkening of the solution on heating and lightening upon cooling is very typical of Fe<sup>3+</sup>, which hydrolyses to FeOH<sup>2+</sup> (simply put) and that is strongly favoured by temperature: ferric solutions are thermochromic. Now I know that the Fe is supposed to be there as ferrous sulphate but it's possible you've oxidised it (with air) to ferric sulphate which at low pH (definitely less than 3, depending on concentration) will stay in solution (it's very soluble, as long as the pH is really low).

Without calculation, 500 ml water + 100 ml conc. H2SO4 would sound like an excess of acid and the Fe concentration would be roughly 2 M, based on 60 g of magnet. So that kind of fit the bill. Of course concentrated solutions of Fe<sup>2+</sup> are quite dark too but not thermochromic...

Very nice crystals, wizz!




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[*] posted on 24-12-2011 at 18:06


Interesting note I made today- Nd2(SO4)3 is paramagnetic, just as ferrous sulfate is!
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[*] posted on 25-12-2011 at 14:15


Might bother telling us how did you pick the measure of ferrous sulphate paramagnetism? Can ferrous sulphate just be attracted by a magnet?
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[*] posted on 25-12-2011 at 14:18


No measure, but holding a large Neo magnet up to the crystal causes medium crystals (5mm or so) to stick- smaller ones aren't strong enough, larger ones too massive. You can also put a magnet to the crystals when in a bag- the bag will sway towards the magnet, strongly sometimes!
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[*] posted on 14-1-2012 at 13:34


As pomissed I've reached expertise in metalic ion separation via pH control. I've purchased a pH meter which will fit best for the purpose of determining the correct scales involved. I just need to resuply my chemicals this month, becuse a car crash took my mney away last month :p . Data shall be uploaded very soon. C you guys

Wizzard: Could you please post pictures of your neodimium iodide, it should seem pretty!
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[*] posted on 8-2-2012 at 19:34


Okay here follows the calculations involving the required pH to operate these chemical purification processes.
From a source entitled "The Thermodynamic Properties of Neodymium Hydroxide Nd(OH)3, in Acid, Neutral and Alkaline Solutions at 25°; the Hydrolysis of the Neodymium and Praseodymium Ions, Nd3+, Pr3+" the following values goes for neodimium hydroxide:
solubility: 4,8 x 10^(-3) mol per 1000g water.
Nd(OH3) (s) <--> (Nd3+) (aq) + 3(OH-) (aq) K = (8,7 +/- 4,47)x10^(-21)
Let us use the average 8,7x10^(-24) value for K, just for comodity.
The value of the K constant agrees well with the formula invonving the solubility of Nd(OH)3:

Nd(OH3) (s) <--> (Nd3+) (aq) + 3(OH-) (aq)

K = [Nd3+].[OH-]^3
as for [OH-], there will be a kind of common ion equilibrum with water molecules: Kw = [H+].[OH-]; Kw/[H+] = [OH-]
Thus follows that
K = [Nd3+].(Kw/[H+] )^3
K = [Nd3+].(Kw^3).[H+]^(-3)
considering the solubility of the compound is so low it wont move the kW value away from 1,0x10^(-14), and also considering the concentration of (H3O+) cations will be compensated by the acid behaviour of the Nd(OH2)6 aquacomplex formed by the dissolution of the compound, we can state for the sake of aproximation that [H+] = 1,0x10^(-7) too, the equation becomes:
K = 4,8x10^(-3) . 1,0x10^(-42) . 1,0x10^21 = 4,8x10^-24, which is close to the result the high grade guys obtained. They must have complicated a bit more this equation but I don't know how.
They also give another number if you guys find it useful:
Nd3+ (aq) + Nd(OH)3 (s) + (H+) (aq) --> (Nd3+) (aq) + 3 H2O (l)
K = 8.7 x 10^13 (thats about it because the article is a hard-to-visualise free version)
I had an doub't if K was in the order of 10^(-21) or 10^(-24) because of the bad visualisation allowed for a free article, but that equations above provided insurance.

So, considering the value of K, neodymium hydroxide, as well as iron III hydroxide, phydrolises too easily in solution depending on pH, it must be adjusted to prevent that. The pH at which neodymium hydroxide starts precipitating out of the solution can be calculated by switching [Nd3+] by some hypothesized value of Nd3+ concentration.
8,7x10^(-24) = [Nd3+].[OH-]^3, you find [OH-] and calculate kW=[H+][OH-], then use -log [H+] to know the pH maximum pH the solution of your Nd3+ ions can go before hydrolising and falling out of the solution, later being mistaken as "salts".
The K value for the dissolution of iron II hydroxide:
K = 2,79x10^(-39) = [Fe3+][OH-]^3 giving a somewhat lower pH of precipitation to start.
The different pH at which iron and neodymium precipitate could be used to separate the two via a pH separation, but for a series of reasons I wont do the calculations x)
So, first of all, when dissolving the magnets, keep in mind the solution must be very acidic, or otherwise very dilute for the final products, remember oxydising the Fe II into Fe III to evade problems and if you trying the sulfate separation method keep the pH high enough just to prevent hydrolisys. Nd(OH)3(s) will not add to crystal formation or color enhancing of the solution, because they neither get into solution.

Sorry for taking this long, :p
Now I gotta recycle my shitty Neodimium powder to get rid of the double salts and hydroxides left mistaken by Nd3(SO4)2 and do it all over again.
SO i hope the crystals so formed will get pretty much like wizard's ones

Muahahah!


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[Edited on 2-9-2012 by Poppy]
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[*] posted on 9-2-2012 at 07:23


Poppy:

I think differential (by pH) precipitation of Nd and Fe hydroxides will be very difficult to do. Concentration effects need also be taken into account, although the high molar ratio Fe/Nd = 7 would be an advantage here.

In an accidental sense I've already tried this, see above, w/o success. Filtering problems with large amounts of Fe(OH)3 (which peptises easily) would also need to be taken into account. Co-precipitation problems might also show up.

This is why the use of poor solubility of hot Nd sulphate or cold (Nd, K) double sulphates remains preferred industrially, or so I believe.

[Edited on 9-2-2012 by blogfast25]




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[*] posted on 9-2-2012 at 16:12


Agreed, see it would be quite difficult.

The sulfate was recycled back into hydroxide form, dryed in an oven, weighed and put to react directly with 98% sulfuric acid. The reaction proceeded very well, releasing heat and with the formation of a strong pink solution. the yield was 28g. This was then dissolved in a 1L beaker filled with 530mL water and sulfuric acid at pH 1.7. Initially there was added twice the necessary ammount of anhydrous sulfuric acid to dissolve the neodymium/ iron hydroxide because a gel was forming which turned to complicate further dissolving. So after dilution the pH was even lower than 1.7: that to prevent iron III species to fall as a precipitate. The 530mL were measured out to dilute the 28g of sulfate at 30ºC, then heating to 100ºC would bring 20g of the neodymium sulfate dissolved to come as crystals. But they didn't. Maybe the very excessive sulfuric acid is playin' a role on holding the solublity of the Nd sulfate up.
The solution at the moment is slowly evaporating and hopefully will give crystals.
How many surprises are we yet to see!
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[*] posted on 10-2-2012 at 05:36


@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.
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[*] posted on 10-2-2012 at 07:31


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.


Explain a bit more, if you please?




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