Sciencemadness Discussion Board
Not logged in [Login - Register]
Go To Bottom

Printable Version  
 Pages:  1  2  
Author: Subject: Rare earth metals, mischmetal, etc
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 6-7-2007 at 01:11


Quote:
Originally posted by ciscosdad
Der Alte,...
Is there any varieties of activated alumina that has differential absorption properties for the rare earths? (charcoal?)
Perhaps a patent search would suggest a suitable resin, or something as simple as alumina may have the desired properties (not likely I guess)....


I think the 'not likely' part is correct. Anionic exchange resins are used, alumina just doesn't interact that much with fairly simple aqueous cations.

Even with the proper resins it may not be simple. The photos I've seen of lab separations of REE showed columns of maybe 1 cm diameter and several meters length.
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 13-7-2007 at 22:07


Have been doing a bit on this. Interim report:

I took a so-called Fe14Nd2B hard drive magnet. Not currently having strong H2SO4 I used ~ 30% HCl to 'dissolve' it. The German method (Loc. cit. above, garage_chemist) used sulphuric, but all the chlorides involved are highly soluble, so I intend to get a solution of the chlorides and then do a double decomp with a sulphate later.

The magnet peeled easily from its Ni(? - it's magnetic) coat, revealing a black sintered material, dense and hard. Weight, 15.43g.

Using stoichiometric amount of HCl (but ignoring the boron) a very exothermic reactcion ensued. Temp rose rapidly. However, the reaction slowed as concentration reduced, as usual. The magnet material was finally about 9/10 consumed in 7 hours. Maybe my HCl wasn't as strong as I thought. Added a bit more next morning and now almost the whole is 'dissolved'.

The resulting solution has a lavender color plus a black powder which may be elemental boron. During the gas evolution a peculiar smell was noted - perhaps that was borane. The books say it is malodorous and this was, but not highly penetrating, and I have a sensitive nose. It did not seem like the acetylene stench that impure iron often gives off due to carbide.

Simultaneously I also used HCl to react with my 5g 'mischmetall' flints. The reacion with a few drops HCl was so fierce I decided to dilute! Here the estimated quantity to 'dissolve' proved correct ( I assumed a RE component X of AW 146 as a weighted average, a valency of 3 - ignoring the possibility that Ce might be 4). The solution does not seem to be colored. A white flocculent precipitate has appeared. No smell to the gas evolved except for inevitable HCl.

As a control, I reacted iron filings of unknown purity with the same HCl in excess. This reacted with about the same energy as the magnet. The smell of the gas indicated acetylene or some unpleasant hydrocarbon. What did surprise me was that the solution is colorless - i expected green, like the sulphate, but it was white, with precipitation of white solid. Apparently FeCl2 is white. You live and learn!

I suspect that the flints contain iron and Cerium and little RE metals.

Any comments from you real chemists lurking out there? Will filter the solution tomorrow and attempt a sulphate precipitation.

Regards, DerAlte.
View user's profile View All Posts By User
tito-o-mac
Hazard to Others
***




Posts: 117
Registered: 30-6-2007
Member Is Offline

Mood: No Mood

[*] posted on 14-7-2007 at 00:31


Best thing you wan't to get is a clump of meteriorite containing some unknown metal that landed in your backyard...:P
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 17-7-2007 at 22:27


The flints contain iron by ferrocyanide test. And, I suspect, little else but cerium.

The casing of the magnet is pure nickel as far as I can tell. Electolytically deposited, obviously. The magnet material is highly conductive electrically.

The filtered solution from the reacted magnet at first showed purple in thin layers and green through thicker. It has now gone deep green. Haven't got a black light - any fluorescence is not obvious. The green color is probably due to gradual oxidation of the ferrous chloride. Splashes turn yellow on the side of the container. The solution is still acidic.

The filtered solution from the flints was clear and now shows signs of light green - iron again. Pity it isn't praseodymium!

Sulphation is next...

Regards,

DerAlte
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 5-8-2007 at 18:03


DerAlte, I think this might be of interest to you, for the qualitative looks of the various rare earths in solution. I always thought Ce(III) solutions were colorless, but they can be coloured when the slightly oxidized metal is dissolved. My pure cerium nitrate is pure white. Unfortunatly I do not own all the rare earths, this is what I have, except Er, which I cannot seem to find...
All pics are from dissolving a piece of the metal(in the case of La, Ce, Pr, Nd, all heavily oxidized because they are powders) in half mL of conc. HCl diluted by half.

EDIT: The reaction of powdered rare earths with only 15%ish HCl is very violent, Nd and La, despite being heavily oxidized, frothed over the testtubes.

[Edited on 5-8-2007 by The_Davster]

RARE EARTHS DISSOLVED.JPG - 26kB




View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 5-8-2007 at 21:16


Very interesting to me, The_Davster.

I agree 100+ % that the reaction with HCl conc. is very violent. I was surprised. Like calcium, and I was only using lighter flints (suspected Ce/Fe) or the magnet material (Suspected Fe/Nd).

I am especially interested in the colors. I understood that Ce salts are all white - the stuff I have is a pleasant clear green due to Fe. In contrast the Nd stuff was very interesting. When the reaction started (rapidly!) the first color seen was lavender, like the German posting, but Fe took over and the resultant was green like the Ce stuff, but much deeper... But it had a peculiar appearance when poured which I can only liken to motor oil - it looked lavender again in thin layers. Wish I had a UV light to test fluorescence. Side illumination with halogen lamp showed a streak but there may have been particles present.

I have not made progress with separation of the RE components (due to a fixation with permanganates as shown in another thread!) But the Ce solution has remained perfectly clear while the ND stuff has deposited a lot of iron hydroxide and looks rusty. Haven't added HCl to it - I think that solution was neutral whereas the Ce was still acidic.

Question for anyone with an answer - when Ce is dissolved in HCl in excess, does the Ce become Ce+++ or Ce++++? Nd I know is trivalent.

Regards,

Der Alte
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 5-8-2007 at 22:39


On the separation of a rare earth from a aqueous solution of it and iron, like Ce or Nd, I wonder that if since the rare earths are water reactive(thus redox potential more negative than water) wheras iron's potential is -0.45V, if electrolysis would selectivly reduce the iron to the metal, and leave the rare earth in solution?
It is something I would like to try...

I myself came into a absolute pile of ancient hard drives, I probably have a dozen, I am thinking of trying to extract Nd myself. I have probably 50g of the pure metal already, but isolating it from a consumer source would be quite nifty.
Is it almost guaranteed that HD magnets are NdFeB? or are older ones a different composition?

Problem is I don't have the screwdriver type to open them:mad:

[Edited on 6-8-2007 by The_Davster]




View user's profile View All Posts By User
12AX7
Post Harlot
*****




Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline

Mood: informative

[*] posted on 5-8-2007 at 22:39


Ce(IV) is strong enough to oxidize Cl-.

For that matter, I've noticed when dissolving powdered Fe2O3.nH2O in HCl, some Cl2 is given off. I'm not sure if there are any other oxidizers present.

Tim




Seven Transistor Labs LLC http://seventransistorlabs.com/
Electronic Design, from Concept to Layout.
Need engineering assistance? Drop me a message!
View user's profile Visit user's homepage View All Posts By User This user has MSN Messenger
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 5-8-2007 at 23:09


Hmm I apologize for the picture of the cerium in the last picture, it appears my cerium powder/filings are contaminated with Pr. The filings were collected all in one day, off the bandsaw I was using to saw ingots of each, so contamination is possible. I just dissolved a piece of Ce I cut off my lump of it in HCl, and it is pure white, so the last picture is incorrect, the cerium should be colorless after dissolving.

I am inclined to say it dissolves into trivalent, the below picture shows, from left to right, a Ce4+ salt(ceric ammonium nitrate), the reaction product of pure cerium(slight oxide coating) dissolved in HCl, followed by a tube of pure Ce3+ nitrate.
The reaction product if cerium and HCl is actually colorless, the picture makes it look purple for some reason, when really it is a colorless liquid with some suspended black powder(does Ce form a nitride with air?)

ceric and cerrous.JPG - 9kB




View user's profile View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline

Mood: No Mood

[*] posted on 6-8-2007 at 00:54


I am fairly sure that Ce (and other rare earths, or at least those early in the series which are the most electropositive ones) would form a nitride with air; - Mg does (and in fact burns in N2 as well as in O2), and Ce is fairly similar to Mg in the electrochemical series.
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 9-8-2007 at 17:53


@ The_Davster:-

The color of Ce++++ noted - and I've seen it since so described else where - so the stuff I have is definitely colorless Ce+++.

The electrolysis idea seems sound - the half reaction Nd+++ --> Nd has Eo about 2.3 volts vs Fe++ --> Fe, 0.47 or so. Good wide range. Damn reactive these REs! Should be able to avoid chlorine production at the anode , too, if you keep the volts down and solution fairly dilute. No separator, I think. Nd(OH)3 is insoluble, like all RE hydroxides. The HCl hanging around should see that it doesn't form.

Am currently gearing up for a series of electrolyses. Getting in some anodes etc., and considering whate to use for diaphrams where necessary... But I'll try the sulphate precipitations tomorrow. I've put them off too long.

Regards,

DerAlte
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 13-8-2007 at 22:11


SULPHATES FINALLY EXTRACTED AND PURIFIED

Finally got around to it.

The Ce2(SO4)3 went off in text book fashion, no problem at all. Result (after 1 stage of purification) Type orthorhombic or monoclinic under magnification (IIRC - not up on crystallography) - pure whitre transparent crystals, well formed and fairly large. LIke calcite, even to twinning and clustering together. Yield, about 4g (about 65% depending upon %Ce in the flints and degree of hydration of the sulphate, assumed octahydrate). The initial batch was stained yellow from Fe ions, which disappeared on recrystallization. Once formed the crystals are quite difficult to redisssolve (same is true of the Nd salt).

Conclusion: Manufacurers today use Ce/Fe only, not mischmetall - or the sulphate would not be pure white.

Neodymium sulphate proved a bit more difficult to extract. The original liquid was dark and cloudy due to iron oxide/hydroxide formation, having deteriorated while I left it for several weeks (the cerium did not suffer). I filtered it first to remove this, and the solution cleared. Heating to near 100C in a water bath made the liquid cloudy again, making it difficult to see whether any precipitation had taken place.

But finally a reddish powder settled. Unlike the Ce2(SO4)3, which I decanted the liquid off, this stuff was too fine so had to be filtered. Among the powder were some well formed small crystals, again monoclinic, almost needles. Under magnification they appeared transparent and slightly reddish and just like the Ce sulphate..

After some trial and error I managed to collect about 5g of this reddish powder, which on drying became pink. Recrystallization was trickly - the stuff redissolved with difficulty at even at 0C, it's most soluble point, alleged to be at about 13g/100g aq. Howver, it reprecipitated fairly well at near 100C, wher the solubility is said to be about !g/100g. Net yield, about 60% (assuming octahydrate and magnet material of Nd2Fe12B).

The German site referenced by Garage_chemist apparently got something they called 'violettefarben' i.e. violet color. My sulphate is a definite pink. The chloride first produced by HCl (it must react before the iron) was lavender colored in solution, and the chloride (+6H2O) is alleged (CRC) to be violet and with hexagonal crystals.The bromide is violet, iodide green, nitate purple, oxide blue, sulphate (Anh?} pink.



One more thing. Tried to oxidize some Ce+++ to Ce++++ with NaClO and adding dilute HCl to neutralize the NaOH stabililizer buffer. Produced an yellowish red color, so I assume it worked.


QUESTION: does anyone have an easily obtained refeerence to the Absorbtion Spectra for visible wavelengths for the RE meatlas, or even some of them?


I have another RE magnet. Samarium cobalt, sia to be Co5Sm. Might try that some day too!

Regards,

DerAlte.
View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 13-8-2007 at 22:44


Ce will go from 3+ to 4+ under alkaline conditions. Make Ce(OH)3, leave it exposed to air and moist and it oxidises up to Ce(IV). H2O2 dives it nicely, adding ammonia and H2O2 to a Ce(III) salt will give at first mixed Ce(III)-Ce(IV) salts, which are dark red-orange, and the yellow Ce(OH)4. (air oxidation takes the same path)
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 6773
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 13-8-2007 at 22:48


Quote:
Originally posted by The_Davster
Unfortunatly I do not own all the rare earths, this is what I have, except Er, which I cannot seem to find...

The following may be interesting for you.

http://woelen.scheikunde.net/science/chem/solutions/

Click on the 4 elements in the lanthanide series, which are highlighted. Erbium is between these metals. It has a light pink color, much lighter than Nd.

I have been able to isolate PrCl3.xH2O. Initially, my solid was nice pale green, but now it has turned yellow/brown.

http://woelen.scheikunde.net/science/chem/compounds/praseody...

I am wondering what caused the change of color. The solid still dissolves fairly easily, especially if a drop of acid is added to the water. These solutions then are pale green.

I also have done some experiments with cerium(III) and I find it remarkably hard to oxidize this to cerium(IV).




The art of wondering makes life worth living...
Want to wonder? Look at http://www.oelen.net/science
View user's profile Visit user's homepage View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline

Mood: No Mood

[*] posted on 14-8-2007 at 00:40


The change of color of PrCl3 is probably due to hydrolysis to Pr(III) hydroxide, and subsequent (partial) oxidation by atmospheric oxygen of that to hydrous Pr(IV) oxide.

[Edited on 15-8-07 by JohnWW]
View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 14-8-2007 at 06:35


I'll agree with JohWW, most likely a surface coating containing some Pr(IV). Pr tends to go to the mixed valency oxide Pr6O11, which is quite dark.
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 17-8-2007 at 13:05


Just a few final notes. I used magnesium sulphate heptahydrate finely ground as a source of SO4-- ions. I avoided the sodium salt because apparently it produces double sulphates with many RE metals.

The colors of the RE ions must depend heavily on the hydration levels, i.e. complexes formed with H20, as for instance the copper green / blue ot the deeper blue of the tetrammine salt. Otherwise one ought to get a uniform color due to the Nd+++ ion. Perhaps the German violet sulphate is with 6H2O or maybe 2H2O. The chloride is lavender in solution, the sulphate I made pink.



Regards,

Der Alte
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 15-5-2008 at 22:52


Samarium

Long ago I mentioned above that I had some small samarium cobalt RE magnets. A couple of months ago I sacrificed one. Again I used hydrochloric acid to ‘dissolve’ the magnet. My aim was to produce cobalt chloride to add to some I had, and a samarium salt. I could find little information on the solubility of the sulphate except at a couple of temps in the solubility table in Wiki, near RT. However, all the other RE metal sulphates decrease their solubility with temp. so I crossed my fingers and hoped samarium did the same. It does.

Another uncertainty was the reaction rate of cobalt metal with acids. Nickel is not easy to dissolve in HCl – emission of H2 is very slow. Fe is very reactive. In the usual gradual transition of properties in the transition metal series, Co should be faster to react than Ni but less than Fe. So it proved. The Sm, of course, is very reactive.

On adding the magnet to excess of HCl, the first thing noticed was a yellow coating appearing. I took this to be SmCl3. Then a blue coloration appeared. I was expecting red, since CoCl2.6H2O is red, but apparently the strong acidic environment (30% HCl) and possibly the temperature rise cause this intense blue color. The magnet was far more slow to react than the Nd magnet mentioned earlier in the thread. I left it 24 Hrs. By then it had been consumed and some blue crystals ( probably CoCl2.2H2O) were left, plus an intense blue solution. This was evaporated to dryness to remove the excess HCl, producing a dark blue solid. I then redissolved this in water to produce a hot saturated solution. On letting this stand for another 24 hrs. fine large ruby red crystals of CoCl2.6H2O separated. On cooling to below freezing, a further crop was obtained. The weight suggested that these contained 45% of the expected cobalt chloride yield. The remaining liquid was assumed to contain the samarium (the chloride is very soluble, about 94g/100g aq at RT).

As before, magnesium sulfate was used to precipitate the Sm2(SO4)3. At room temp no such precipitate was obtained. Since calculations showed that this should have definitely happened (solubility being about 3g/100g, at RT (Wiki)). However, after heating to just below boiling for some time, a powder with a whitish color, difficult to see in the intense blue of hot CoCl2, was thrown down. I concluded that the samarium salt supersaturates . I separated this powder by filtration and dried it. It had a beige color and was verylight green when hot. This was apparently due to contamination by CoCl2. Later I redissolved and re-precipitated it at near 95 C and it became a light canary yellow. It was quite difficult to dissolve and almost equally difficult to re-precipitate easily. Once dissolved it did not want to recrystallize, and unlike cerium and to lesser extend, neodymium, the crystals were extremely small and powsery. It obviously supersaturates, which caused my earlier trouble.

On weighing I found my yield of the Sm2(SO4)3 was over 100% of expected. This puzzled me for some time. I wondered if I had the hydration level right (8H2O) but all references said I had. Finally I found out from Arnold Magnetics site (well worth reading if you are interested in magnets) that the original formulation of samarium cobalt magnets was about SmCo5 but more recently Sm2Co7 is commonly used. This allowed my yield to be a respectable 87% and explained why my CoCl2 yield was so poor! (The latter was also due to an error in technique – I poured away a solution still containing significant CoCl2, which if cool is a not very intense red).

The remaining liquid containing CoCl2 was treated with NaOCl to precipitate the cobalt as Co2O3, and then redissolved in a stoichiometric amount of HCl to reproduce the chloride as follows (I think):

2CoCl2 +3NaClo ---> Co2O3 + 3NaCl + Cl2 (smell it!)
2Co2O3 + 8HCl ---> 4CoCl2 + 4H2O + O2 (glowing splint test)

Co2O3 is very reminiscent of MnO2, being brown to black and possibly hydrated.

Regards

Der Alte
View user's profile View All Posts By User
IrC
International Hazard
*****




Posts: 2710
Registered: 7-3-2005
Location: Eureka
Member Is Offline

Mood: Discovering

[*] posted on 24-5-2008 at 18:25


"except Er, which I cannot seem to find..."

Rods about 30 grams are 65 dollars at the following site. He has just about anything you may wish for.

http://www.elementsales.com
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 24-5-2008 at 19:26


Thanks for the link, but I did not word my original post well. I physically could not find it at the time on my metals shelf, :P I already have some of the metal.
(It was behind the bases for some reason)




View user's profile View All Posts By User
kmno4
International Hazard
*****




Posts: 1365
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline

Mood: No Mood

[*] posted on 30-6-2008 at 01:16


A few days ago I decided to try to recover Nd from neodymium magnes (~25% of Nd) - it turns out quite simple.
Magnets (two pieces) were heated (burner) untill became no more magnets. Protective Ni coating has been mechanically removed (pain in the ass :mad: ). Next, grey Fe/Nb/B has been broken into smaller pieces (hammer)
2,0g of these pieces has been disslolved in HCl(aq), H2O2 added, heated to boiling and filtered. Clear, yellowish solution is obtained.
Do not give too much HCl. It is better to wait 24h (or heat it for dissolving) than add 10-fold excess of acid.
Acidic solution is diluted to 50-100 ml with demi water.
Next, ~50 ml 10% solution of oxalic acid is prepared.
Both solutions should be warm ~50 C.
A half of this solution is added to "magnets solution" - at once large amouts of gummy violet-wite precipitate is formed. In several seconds, this gummy mass turns into violet-wite powder (with aid of glass rod), which is filtered out.
Very important - do not wait longer than 2-3 minutes from the moment of adding of oxalic acid, to complete precipitation.
Almost whole Nd precipitates as oxalate at once and prolongated waiting (as I did before :mad: ) causes precipitation of yellow crystaline compound - some Fe oxalates.
Violet-white powder is washed, in several portions, with remaining part of oxalic acid solution and demi water in the end. Wet precipitate is dried at ~100 C.
Yield: about 0,90 g of neodymium oxalate as white powder with violet shade. By heating it at ~700C in the air you can convert it in Nd2O3.
On the picture Nd<sub>2</sub>[(COO)<sub>2</sub>]<sub>3</sub> (propably with some small amount of bounded water)





[Edited on 30-6-2008 by kmno4]

DSCN9843.JPG - 54kB
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 30-6-2008 at 16:02


Nice, kmno4! The color of the oxalate (lavender) is far more pleasing and characteristic of Nd+++ ion than the sulphate I made as above. Most Nd salts are purple/lilac but the sulphate is light reddish. CRC says the iodide is green. A variable ion.

Why did you add the H2O2? Just for luck, or to convert Fe++ to Fe+++? I considered the oxalates ( nearly all the RE ones are insoluble, IIRC) but rejected the idea because Fe(II) oxalate is also rather insoluble. But Fe(III) maybe is not. In which case the H2O2 sounds like a very good idea.

Also why demagnetize? And why bother to smash the thing up? The nickel coat can be peeled off rather easily, I found. Incidentally one of the failings of the NdFeB magnets is a very low Curie temp (~150C, IIRC) which makes them useless for military applications – hence SmCo is still used for high temp equipment to MIL spec.

What happens to the magnetic energy when you separate the matrix under reaction? It must go to heat – the entropy certainly increases. How much energy is released? Not much. I don’t have a figure to hand, but I worked it out some time ago (from the energy of magnetization, BH/2) and it was a few joules or so. Insignificant compared with the energy released by the reaction.

Did you notice the repulsive smell I attributed to borane, kmno4?

If anyone repeats this, do use diluted HCl as kmno4 suggests. The reaction is quite aggressive.

Regards

Der Alte

[Edited on 30-6-2008 by DerAlte]
View user's profile View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline

Mood: No Mood

[*] posted on 30-6-2008 at 20:07


You say you heated the Nd(III) oxalate to obtain the oxide, Nd2O3. I wonder if you conducted some sort of assay on it, perhaps based on the amount and composition of the Nd-Fe-B magnet you started with, to try to confirm that composition? There is some tendency of some rare-earth metals to form higher oxides, often non-stoichiometric, on heating in air, e.g. Pr6O11 and Tb4O7. I have read somewhere that such an oxide of Nd also can be made.

As for the low (only 150ºC) Curie point of Nd-Fe-B magnets, which I understand are used in the motors inside hard drives (which sometimes get fairly close to that temperature), rare-earth magnets made with those metals later in the series, especially around Gd, should be better with regard to both magnetic moment and Curie point because of their greater numbers of unpaired 4f electrons. Also, I wonder if depleted U-238, with a good number of unpaired 5f (and 6d) electrons, could be used in ferromagnetic alloys, rather than wasted in ammunition used in Iraq. (It is known that Pu is strongly ferromagnetic, but of course its radioactivity and cost rules it out).
View user's profile View All Posts By User
DerAlte
International Hazard
*****




Posts: 779
Registered: 14-5-2007
Location: Erehwon
Member Is Offline

Mood: Disgusted

[*] posted on 30-6-2008 at 21:11


Correction

@JohnWW and readers:

My IIRC re Curie temp of NdFeB magnets was faulty. Research turned up 310C, still very low with possible demagnetization at 150-200C.SmCo is in region 700-800C. Your hard drive is probably safe @ 150C, but I wouldn't trust the CPU!

Gd has been considered for RE magnets. PtCo also, but I've never heard of U being considered. All RE metals produce fairly good results with Fe,Co, for the reason JohnWW has indicated, I guess. Certainly the d,f electron band connection is there.

The RE metals are no longer rare, nor that expensive. Co is about as rare and getting rather expensive. I believe I read that Sm2Co7 is less expensive raw material-wise than SmCo5...

Regards

Der Alte
View user's profile View All Posts By User
kmno4
International Hazard
*****




Posts: 1365
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline

Mood: No Mood

[*] posted on 14-7-2008 at 15:48


Ha, why H2O2 ?
I just wanted to have "clear" composition: Nd(III)/Fe(III).
Fe(II) salts quickly are oxidized by air to Fe(III) but I wanted to have it at once. I also wanted to make all procedure quickly, so I had to smash demagnetized magnet into pices, to enlarge reacting surface. Oxalic acis (but not its salts !!) is using for precipitating of Ln(III) as oxalates, not soluble in diluted acids and water (solubility of oxalates in water is about single mg/dm3). As far as I know in the case of Nd, oxalate decomposes to oxide (via oxycarbonate) when heating in the air.
During dissolution of magnet in HCl indeed I noticed not too pleasant smell of (propably) BxHy compounds, but most of B remained as black slury.
I will try similar procedure for Sm/Co magnet (soon), but without H2O2.
View user's profile View All Posts By User
 Pages:  1  2  

  Go To Top