Sciencemadness Discussion Board

Rare earth metals, mischmetal, etc

DerAlte - 1-7-2007 at 21:16

As a new menber I feel a bit guiltly about opening threads. I did a search but found little. I have been accumulating odd ends of lighter flints due to incessant pipe smoking, and also have a few high gauss permanent magnets, cerium and possibly samarium, that I ordered years ago when designing & making circulators for millimeter wave applications.

The lighter flint end collection has reached critical mass of 5g. I am also prepared to sacrifice the magnets.

Has anyone, especially Woelen, king of the rare metals here, attempted separation of the rare earth metals, which I understand is a heroic task?



pantone159 - 1-7-2007 at 23:03

I have not tried any separation myself, but I can report what one of my books says... (Concise Inorganic Chemistry, 5th ed, J D Lee):

In general the separation is very difficult, but a few lanthanides have valence states other than +3 and this can be used.

If you have a solution of Ln+3 ions, Ce can be oxidized to +4, book says NaOCl in alkaline solution. The Ce+4 is separated by carefully controlled prcipitation of CeO2 or Ce(IO3)4, leaving the Ln+3 ions in solution. (Book has no details on the careful control...) Alternatively, Ce+4 can be readily extracted by solvent extraction in HNO3 solution using tributyl phosphate. 99% Ce can be obtained in one stage from a mixture containing 40% Ce. (This probably approximates your misch metal.)

Similarly, Eu+2 exists, and your Ln+3 solution can be reduced electrolytically with a Hg cathode or Zn amalgam. If sulfate ions are present, Eu(II)SO4 precipitates as it is insoluble. Sm and Yb can also be put into the +2 state ths way, but these ions are oxidized slowly by water.

Besides that, some sort of fractional separation is needed. Assuming you don't want to do hundreds of fractional crystallizations... Ion exchange chromatography is I think the best method. A solution of Ln+3 ions is run down a Dowex-50 or similar column, Ln is absorbed onto the column. The H+ ions resulting are washed through the column. The Ln ions are then eluted with a complexing agent, i.e. 'a buffered solution of citric acid/ammonium citrate, or a dilute solution of (NH4)3(H-EDTA) at pH 8'. Separation can be monitored spectroscopically with atomic fluorescence. (Maybe not at home!) The metals are precipitated as oxalates, then heated to give the oxides.

not_important - 2-7-2007 at 00:45

As a slight extension to pantone159's post, see

which describes the current manufacturing technique as well, and gives a few references.

Andding hydrogen peroxide and then aqueous ammonia to a solution of the mixed salts will precipitate the 3+ hydroxides first, then Ce(OH)4. Usually the L(III) hydroxides would be redissolved and precipitation to fully remove Ce, and the cerium fraction would be treated similarly.

The 2+ ions reduce to form amalgams more quickly that the 3+, using a mercury electrode a solution of the 3+ ions can have the Eu, Sm, and Yb removed as an amalgam.

I believe that one of the Inorganic Synthesis volumes had a procedure for separating the REE.

garage chemist - 2-7-2007 at 00:53

If you are going to use neodymium magnets as well, you should work them up separately as they only contain neodymium, sparing you its separation from the other rare earths.

In the german forum, there is the documented production of neodymium sulfate from magnets:,9178,0,-Neodym%28III%29-s...
Although there are some concerns about the purity of the product thus obtained.

[Edited on 2-7-2007 by garage chemist]

woelen - 2-7-2007 at 06:19

I have never tried to separate rare earths from mischmetall. As already is pointed out by others, this is very difficult, certainly for a home chemist with limited equipment and resources. One rare earth (cerium) can be relatively easily be separated, because it can go in the +4 oxidation state and as such can be handled in a special way, but separating it from mischmetall hardly is of any use. I have seen all rare earths being sold on eBay for very decent prices, the cheapest ones being cerium, praseodymium, erbium, samarium and gadolinium.

not_important - 2-7-2007 at 07:30

Ceramilcs suppliers are another source of reasonably pure oxides for cerium and the coloured REE - erbium, europium (as II), neodymium, praseodymium. For example

But then there's not much challange, which here may think is an important part of doing home science :-)

DerAlte - 2-7-2007 at 22:50

Thanks for the replies, all! I knew Ce was fairly easy to separate but it is also the least interesting. What makes the rare earths interesting is that unfilled shell, like the transition metals (one of my favorites in all inorganic chem..) It manifests itself as fluorescence, colors, sharp absorption bands – and hence the use in CRT phosphors and lasers (Nd, e.g). I’m a sucker for nice colors. Magnetic properties too.

The rare earths are not as rare today as in my distant youth. They now have considerable commercial uses. At present I am too busy with other projects to contemplate hairy separations so I will shelve the project and accumulate some more bits of flints, do some more research. Maybe even get a gram or two on ebay.

I wanted to see if anyone ever took on the challenge. I’m as ready as the next guy for a challenging fractional crystallization, I revel in it, but I currently draw the line at half a dozen attempts max., about what’s needed to really clean potassium chlorate from sodium contamination and get rid of that yellow flame - always use distilled water, never ‘de-ionized’… Perchlorate is a bit easier.



JohnWW - 3-7-2007 at 02:15

The characteristic colors of those rare earth cations which have unpaired 4f electrons, typically pastel shades of pink, green, yellow, blue, are due to transitions between energy levels of those electrons with energies which correspond to visible light wavelengths, and between the 4f and 5d orbitals. This, and paramagnetism and ferromagnetism, are most evident around the middle of the rare earth series, where the metals or their cations have up to 7 unpaired 4f electrons. See my recent post on the "High oxidation states of rare earth metals" in this section, for further information. Some of their oxides have recently found uses, in crystalline mixtures with Ba and Cu and some other oxides, as superconductors with transition temperatures higher than the boiling point of liquid nitrogen (77ºK).

As for separating them, I once had a final-year laboratory experiment on separating an unrefined mixture of rare earths (called "didymium", a crystalline mixture with a pale blue-green color), as the chlorides I think, by means of elution through a column packed with an ion-exchange resin, with aliquots of the eluted solution being collected separately and analysed by UV/visible spectroscopy and ESR spectroscopy.

As well as La and the following 14 metallic elements (the electronegativities of which steadily increase with atomic number due to increasing nuclear charge resulting in valence electrons being progressively more tightly held), "rare earth" elements also include scandium and yttrium, and also the actinide series, comprising actinium and the succeeding 14 elements of which only Th and U occur naturally significantly. They are found mostly as a mixture of the silicates in "pegmatites", huge crystals which occur at (originally) very deep levels in continental granite areas, having been the last part of granite masses to crystallize from the liquid state; rare-earth oxides are apparently very soluble in molten granite. Because of the depth at which these deposits occur, mining of them is possible only where a considerable thickness of granite has been removed by weathering and erosion, e.g. in Scandinavia. Some geochemical differential mineral formation occurs in places, resulting in some separation of Sc and Yt e.g. as the rare mineral thortveitite, Sc2Si2O7, found in Greenland, and of Th and U e.g. as pitchblende, U3O8.

chemkid - 3-7-2007 at 13:24

some of tit-o-matic's posts have been useful....or am i thinking of a different user?

Anyway, does misch metal react with HCl to form some rare earth chlorides.


12AX7 - 3-7-2007 at 13:25

I imagine it does, though they may be prone to hydrolysis, especially if you try to crystallize. I would imagine CeCl3 is hard to produce, and CeCl4 impossible (Ce oxidizing Cl-).


woelen - 3-7-2007 at 13:48

Yes, mischmetal makes chlorides with hydrochloric acid. The reaction is very vigorous and a lot of heat and hydrogen is produced.

The solutions are quite stable and hardly hydrolyze, but when an attempt is made to obtain the dry chemical, then things become different. I have made some solid PrCl3.xH2O from Pr-metal, but I needed several attempts before I obtained the nice light green material. Just slightly overheating the material results in loss of HCl and formation of oxide or oxychloride. And even then, I still think that my product contains a small percentage of oxychloride. It does not dissolve completely as a clear liquid in water. A single drop of dilute HCl, however, makes the solution clear.

Finally I succeeded, by VERY careful heating and slowly blowing away any water vapor, released from the solid:

The solid apparently looses water on storage. It still dissolves fairly easily.

[Edited on 3-7-07 by woelen]

chemkid - 3-7-2007 at 13:53

A vigorous reaction like aluminum or something more violent?

[Edited on 3-7-2007 by chemkid]

The_Davster - 3-7-2007 at 19:03

I think if you choose to attempt to do a good separation of these compounds from michmetal, you may want to start with more michmetal as depending on the source of the michmetal the composition can have varrying ammounts of the various rare earths, and some make up a miniscule fraction of the michmetal.

You can buy lumps of it for cheap from camping stores, like these here

"I feel a bit guiltly about opening threads"
no need to:)

One of the things I regret most is not making a video of the time I was instructed to cut a 1kg ingot of cerium into wafers, using a bandsaw:o:D. And followed it up with similar ingots of Nd, La, Pr.

The reaction of the rare earths with acids is along the lines of Al and HCl but milder, similar to calcium in cool water.

Crystalization of some are possible, I have not tried with many, only cerium and Nd, Ce did not work well at all, but NdCl3 I was able to obtain as light purple crystals from a slow room temp evaporation of the aqueous solution. I also did something with lanthanum chloride, but completly forget everything about it, I do not think the crystals were nice.

not_important - 3-7-2007 at 20:05

One method for getting anhydrous chlorides of metals, when those chlorides are not volatile - not like AlCl3, is to mix the hydrated chloride with ammonium chloride, slowing heat the mixture to several hundred C, and then pull vacuum on it to removed excess NH4Cl. It's in one of the early Inorganic Synthesis collections, specifically for the lanthanides but it does work with other chlorides as well.

DerAlte - 3-7-2007 at 20:47

The response has been far greater than I expected. Thanks!

@Pantone 159: chromatography is a bit beyond my capabilities, but emphasizes the difficulty of separation.

@ not_important: I read the Oxford article; again, shows it’s a tough proposition not to undertaken lightly. I believe Brauer has a bit on separation.

@ garage chemist: I read (as best I could) that fascinating German post. One thing I have plenty of is old hard drives – my son is always discarding them from old computers. My understanding of German is piss poor, unfortunately, just enough to get by with, but I shall be having a further look at that forum…

@Woelen: I am not an element collector, I guess it’s just the challenge that appeals to me. But as I said, not 100 fractional crystallizations!

@JohnWWW; Nice succinct summary of the lanthanides. I think I’ll frame it and hang it on the wall.

@tito-o-mac: You are a bit of a tit. Get a life.

@The_Davster: I feel I may not have searched the archives well enough, that’s why I am a bit bothered about opening new threads.

Have also been reading the contributions of Chemophiliac’s thread re Oxidation state of REE’s.



chemkid - 4-7-2007 at 04:36

The_davaster: What sort of bandsaw are yolu cuting theese with?

The_Davster - 4-7-2007 at 05:28

I am not a tool specialist, it was what was in the uni machine shop. it was over 6' tall, and had a blade about 1cm deep and 1mm wide. It was a medium toothed blade. I have no idea if this is what you mean?

chemkid - 4-7-2007 at 06:26

Sadly i am not much of a tool specialist either. However i would think, unless neodynium is a very soft metal, or you have access to a really nice bandsaw, that cutting neodynium would ruin most any blade. The bandsaw I use would have a tough time cutting aluminium (i'm not going to try either metal becuase it is a six hundred dollar bandsaw and it's not mine).


Fleaker - 4-7-2007 at 08:17

Yeah some of those rare earths are really hard to cut, almost like to fracture rather than cut easily. I only have very small (<15g) samples of them though.

Still, with a high carbon blade, I've cut through 1" thick steel bolts on a horizontal band saw. Aluminum doesn't like to cut with a finer blade: ""

The_Davster - 4-7-2007 at 16:45

The rare earths I cut were decently soft, I'd say similar to magnesium? You can shave off pieces of all of them with a pocketknife. We eventually gave up cutting the cerium, because of all the fires that started. Then I was given a hacksaw and did that for a few hours. I did not notice any blade issues.

unionised - 5-7-2007 at 10:18

I have obtained fairly pure Nd salts from NIB magnets. The first thing I did was heat them until they stoped being magnets- I was annoyed every time I tried to do anything with them and they stuck together. Then I peeled off the coating (Ni I think) because that was going to be easier than removing it chemically.
The next step was to dissolve them in dilute HCl, I'm not sure what happens to the boron. I guess it oxidises to boric acid because I think I would have noticed BH3.
This gave a mixture of Fe and Nd as chlorides. Then I oxidised them with bleach to get the Fe (II) to Fe (III).
The separation relied on the fact that adding a base to a mixture of Fe(III) and Nd(III) causes the Fe to precipitate first. Once you can't see the colour of the Fe(III) chloride complex you can decant the solution and ppt the Nd as a hydroxide. Dissolving this in dilute H2SO4 and leaving it for the water to evaporate gave nice crystals.

Another possibillity would be to dissolve the mixed chlorides in fairly conc HCl and extract the "HFeCl4" complex with ether. Fine if you have lots of ether- I didn't.

DerAlte - 5-7-2007 at 15:48

@ unionised

Now that I can do! In fact, I've already stripped a magnet. Any idea of the approx composition of the magnets and the Curie temp. for reference? I guess it's not really necessary to de-magnetize, though.

I'm curious - Is that "un-ionised" or "union-ised"? suspect the former, somehow...



JohnWW - 5-7-2007 at 15:51

See my recent posts in the parallel thread "High oxidation states of rare-earth metals", and, for links for downloads of ebooks on the rare earths, in the Inorganic Chemistry thread in the References section.

ciscosdad - 5-7-2007 at 16:23

Der Alte,

I think JohnWW's suggestion to use an ion exchange resin would be the most feasible for the amateur. Given that you seem to be interested in only gram quantities, the relatively dilute solutions should not be too big a problem.

Maybe an approach closer to chromatography would work.
Years ago I worked in an oil refinery laboratory, and one of the tests was for aromatics content in power kerosene (and carrot spray for gods sake!). The medium was alumina and the identity of the eluent escapes my failing memory. Propanol I think. We cheated and simply used a UV light to show the relative lengths of the various components of the developed sample. Given a constant bore column, that was not a big problem.
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).
Please continue starting new threads. I've loved every one!

DerAlte - 6-7-2007 at 00:03

Re the magnets, the initiator of the German thread cited by garage chemist states the composition as Nd2Fe12B. Further down the thread there is a suggestion (if my quick perusal and poor German are correct) that other RE metals may be involved, a sort of 'mischmetall/iron' magnet. In which case the Nd won't be as easy to separate. The process, as stated, depends upon the inverted solubility/temp curve of Nd sulphate - but, as one might expect, ALL the RE metals exhibit the same phenomenon (even La, Ce). At a given temp there is some difference in solubility, however. If it's a mixture, you will have the same separation poblem all over again...

Nitrates and chlorides are v. soluble, selenates like the sulphates, inverted curve. No surprises here - these elements are chemically too close, almost as close as recent presidential elections...



not_important - 6-7-2007 at 01:11

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.

DerAlte - 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.

tito-o-mac - 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

DerAlte - 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...



The_Davster - 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]


DerAlte - 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.


Der Alte

The_Davster - 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]

12AX7 - 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.


The_Davster - 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

JohnWW - 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.

DerAlte - 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.



DerAlte - 13-8-2007 at 22:11


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!



not_important - 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)

woelen - 13-8-2007 at 22:48

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.

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.

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).

JohnWW - 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]

not_important - 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.

DerAlte - 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.


Der Alte

DerAlte - 15-5-2008 at 22:52


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.


Der Alte

IrC - 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.

The_Davster - 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)

kmno4 - 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

DerAlte - 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.


Der Alte

[Edited on 30-6-2008 by DerAlte]

JohnWW - 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).

DerAlte - 30-6-2008 at 21:11


@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...


Der Alte

kmno4 - 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.

bahamuth - 1-2-2010 at 17:01

Tried your, knmo4, method but i couldn't get it right, yellow precipitate formed right away. Was very quick to filter the solution, and the solution may have been too acidic (became bored of waiting for the magnets to dissolve..). Probably that's my errors right there.

But, upon standing, a white precipitate formed in the mother liq..

Added my yellowish precipitate from first filtrate and let it all settle, and decanted it to recover the precipitate.

To this I added conc. HCl so to dissolve it, turned green, proving it contained a lot of iron, diluted this with water to again get white precipitate, decanted on added water and brought to a boil to dissolve the iron salts. Decanted and repeated boil and decanting. Filtered the perfect white crystals and washed them several times with water and twice with acetone to dry them. None seemed to dissolve in the acetone.

The total yield is ~20% of theoretical, 6.37g Neodymium oxalate from 22.73g magnets.


Tested it under a regular lamp, since we do not have much daylight here quite yet, (live far north in Norway) and it was faintly pinkish in appearance.

[Edited on 2-2-2010 by bahamuth]

densest - 4-2-2010 at 14:37

Would Fe(CN)6 be useful for precipitating iron from a mixed solution of Fe & lanthanides? I haven't googled it yet; inorganic chemistry doesn't seem to be well documented on the web.

Or would there be some other ion/complex which characteristically reacts with Fe++ or Fe+++? SCN? The oxalate precipitation of lanthanides set off this thought....

bahamuth - 5-2-2010 at 10:19

Are you thinking of adding potassium ferricyanide and precipitating prussian blue, filter it and have a solution with only Nd ions?

Never tried that, might work but you have to be careful not to have an acidic media when you add the ferricyanide because of possible evolution of hydrogen cyanide gas.

Also, think I've read somewhere that prussian blue is unfilterable, because of the small small particles it precipitates as.

Perhaps you must centrifuge it:)