Sciencemadness Discussion Board - Lanthanum from pool cleaner

If I wanted Lanthanum, I would get it from soil, rocks, or flint from lighters! Flint (not mineral with the same name) in lighters is mixture of La, Ce, Pr, Nd. Soil contains La in miligram amounts.

[Edited on 13-10-2014 by AsocialSurvival]

Or, you know, you could just buy a perfectly OTC and relatively pure source of lanthanum. But by all means, attempt to extract significant quantities from soil in an amateur setting. Have fun!
The lighter flint is something I've attempted before, and separating the 50% Ce, 25% La, 1% other RE's and balance of Fe and Mg is no easy task.
I wondered why I couldn't find the relevant thread here for a while... Never did see if anyone responded to my question of getting a smaller (but presumably more concentrated) amount of phosphate remover or a larger (but specified as diluted) different brand.

Texium - 13-10-2014 at 14:51

Good luck purifying visible amounts of it from those sources. Lanthanides, in case you weren't aware, are extremely hard to separate chemically. Also, the tiny amounts found in most soil would be so miniscule and hard to isolate that you might as well try and find gold in your soil.

AsocialSurvival - 13-10-2014 at 15:05

OK, anybody interested in these procedures, feel free to subscribe to my yt channel: http://youtube.com/user/AsocialSurvival

Also, I will never buy anything for Chemistry, I will only sell something. Chemistry's purpose, in my case is to become rich, not poor.

[Edited on 14-10-2014 by AsocialSurvival]

Texium - 13-10-2014 at 15:10

I'm eager to see these procedures, but unfortunately it appears that your YouTube channel doesn't exist.
Anyway though, how is it that you expect to isolate lanthanum from your soil?

Bert - 13-10-2014 at 15:16

Chemistry's purpose, in my case is to become rich, not poor.

Hey, it worked for Alfred Nobel. But Nicolas Leblanc, not so much. Let us know how this pans out for you...

BTW, YouTube link is not working.

AsocialSurvival - 13-10-2014 at 15:31

First dissolve soil or rock in Nitric Acid (which I have in unlimited quantity). That would dissolve almost all elements present there, except Si, Ti and few more. The rock will immediately disintegrate into insoluble powder. Then, by boiling solution of soluble elements, I could perfectly seperate them all. It's long explanation why would only Nitric acid in small quantity would be needed, and boiling that all, but once again, video will come one day! Fixed: http://youtube.com/user/AsocialSurvival

[Edited on 14-10-2014 by AsocialSurvival]

Bert - 13-10-2014 at 16:52

Oh, the YouTube channel is there. Just no content of any type.

"Unlimited (free?) nitric acid"... Perhaps means someone else is paying for nitric acid they are not getting? Did you surreptitiously tap into an explosives factory pipeline or what?

Tdep - 13-10-2014 at 17:01

Surely you could just sell your unlimited nitric acid and get rich that way?

But we're getting distracted. The only method people said was viable was precipitating the fluoride and then reducing it with lithium?

This can all be done OTC in Australia then, yielding a few grams of lanthanum metal for about $50. Not cheap but an interesting project

j_sum1 - 13-10-2014 at 17:25

@ AsocialSurvival
You have a lot more confidence in your ability to separate compounds than I do. One of the characteristics of the REs is their similarity in chemical and physical properties which is why, historically, they have been so hard to isolate. (Hence the challenge and interest that they present.)
I am envious of your supply of nitric acid. Wish I had the same kind of resource. I have to make mine.

One of the appeals of using OTC LaCl3 is it enables some experimentation without going through the difficult process of isolating. And indications are that reduction of LaCl3 to La is by no means a trivial process in itself. If posters' experience with Nd is anything to go by, this is a difficult task in a home lab situation.

Given that the original thread disappeared into the ether, let me summarise what I remember of the findings.

Lanthanum Chloride is available OTC – sold as a swimming pool additive to precipitate phosphates that can then be filtered out. Interesting use of the chemical.
A likely route for refining is to convert to LaF3 and reduce with lithium. Potassium has also been suggested.
The thermodynamics for reduction directly from LaCl3 are questionable. Blogfast reported a paper that had some success with CeCl3. No hard data presented for La.
LaCl3 exists as a heptahydrate and is difficult to crystallise. It is extremely hygroscopic. LaF3 is insoluble which is another reason for suggesting this route.
Precipitating LaF3 could conceivably be done using a fluoride salt. NaF possible but NH4F2H is suggested. This is available OTC as glass etchant and tile roughener. In Aus, Home Hardware sells a litre at 48g/L for around $27. I found a chemical supplier who will sell for $71 plus freight for 500g. It is available in bulk bags of 25kg each. I don't think i need that much.
Blogfast has begun work with 2.5L of LaCl3 solution that he bought. I will let him post the details. Looked very interesting though.
Some side discussion on the logistics of storing rare earths was also had. Standard techniques of ampouling under argon or oil seemed to be the way to go.
I am sure I have missed something. I am fascinated by the whole project and while my participation in practical experimentation will probably have to wait for a few months, I am interested in the ideas offered. It seems like the original thread was of significant interest to a number of people on this board.

Now. Let me not die again and take my posts with me.

In the interests of keeping this topic alive does someone want to quote this post so that if another glitch happens the information will not disappear.
Blogfast, if you remember what you posted before, I would love to take a look at some of those links.

J.

AsocialSurvival - 13-10-2014 at 20:39

Supply of nitric acid is simple, air contains nitrogen, wood gives energy.
Why are you all complicating this? So simple! Rare earth metals are easy to seperate just like any other element. I am confident in that that I can separate them all using only Nitric acid!

What fluorides? What Lithium? That is too complicated. OK, let's see how would I separate all rare earth metals?

Just add enough acid for one to dissolve (if all others are carbonates at the beginning). For reduction use K or Na or Li or Rb or Cs or electricity or something else. There're many ways.

Just like scientists complicate nuclear transmutation and want to force us to believe it is hard or impossible or complicated or a big deal. Not really! Some people just want to put me down, to force me to believe in negativity, and to waste my life trying and not succeeding. It's not about complications, not really - it's easy to do many things much more than you all think. But everybody will tell you "impossible", "complicated", "expensive".

Here's my promise for everybody: I won't talk with you until I prove to you that many things are possible. Next time I come here, I will show you how to separate all elements from soil or rocks using nitric acid only, and to make gold with nuclear fusion to become rich.

Bye until than!

j_sum1 - 13-10-2014 at 21:13

Thanks for your thoughts AsocialSurvival.
I rather suspect you are missing a few critical things. It is not like we are making things complicated just for the hell of it. Things become complicated out of necessity because they are difficult. Chemists consider energy changes that are required for chemical reactions. And that is why energetic materials like potassium, lithium and fluorine are needed. If you can obtain nitric acid from atmospheric nitrogen and energy from a wood fire then you really need to patent that. Reaction mechanics suggest that there are quite a few obstacles in that path.

And your promise... That sounds like a reasonable challenge you have set yourself. Good luck.

Bert - 13-10-2014 at 21:52

Quote:

Just like scientists complicate nuclear transmutation and want to force us to believe it is hard or impossible or complicated or a big deal. Not really! Some people just want to put me down, to force me to believe in negativity, and to waste my life trying and not succeeding. It's not about complications, not really - it's easy to do many things much more than you all think. But everybody will tell you "impossible", "complicated", "expensive".

There is chaos under Heaven, and the situation is excellent...

You go, son.

blogfast25 - 14-10-2014 at 05:02

Before I resume my actual contribution here, this little gem below merits a short answer:

What fluorides? What Lithium? That is too complicated. OK, let's see how would I separate all rare earth metals?

Just add enough acid for one to dissolve (if all others are carbonates at the beginning). For reduction use K or Na or Li or Rb or Cs or electricity or something else. There're many ways.

Just like scientists complicate nuclear transmutation and want to force us to believe it is hard or impossible or complicated or a big deal. Not really! Some people just want to put me down, to force me to believe in negativity, and to waste my life trying and not succeeding.

You are a nut and possibly need help. Possibly a troll.

But do by all means come back when you've actually got something... I won't be holding my breath though.

[Edited on 14-10-2014 by blogfast25]

blogfast25 - 14-10-2014 at 05:43

I’ll summarise my work on the lanthanum bearing (10.3 % La glycolate, acc. MSDS) pool phosphate remover called ‘Lo-Chlor Starver’ (from Lo-Chlor Chemicals, AU) as follows.

2.5 L of this solution was neutralised with 33 % NH3 and the snow-white precipitate, presumed lanthanum hydroxide, quantitatively Buchnered and washed. This yielded about 1 L of product.

It was dissolved gradually and quantitatively in 280 ml HCl 37 % with mild heating, followed by simmering. It dissolved completely to a clear solution that was… slightly green! On further boiling that colour changed to yellow. I have very strong reasons to believe this colour is due to MORE than a bit of ferric chloride, more about that later. Possibly there is Pr present.

The solution (about 1.25 L) was then gradually boiled down to about 290 g, by which time it was urine yellow. On refrigeration a mass of small, white crystals formed (photos later) although much of the presumed lanthanum chloride heptahydrate remained in solution, as expected. The rest of the LaCl3 will now be recovered using the double sulphate method (with K2SO4), to try and separate it from the yellow contamination. If it cannot be separated I’m inclined to believe the original phosphate remover contained also Praseodymium. If it can, that would point to ferric chloride.

As regards possible methods of preparing La (and other REE) metals, I proposed the reduction of LaF3 with Li, which appears to be a thermodynamically favourable reaction, possibly with enough reaction heat to ensure clean separation between metal and slag. LaF3 is water insoluble, so fairly easy to prepare.

I earlier dismissed the use of reduction of anh. RECl3 but Brauer’s ‘Handbook of Preparative Inorganic Chemistry’ proved me wrong on that. I suggest strongly that anyone interested in contributing to this discussion go to the Library of this site, download Brauer and print off pages 1135 to 1150 from Volume 2.

There is an absolute wealth of information there that seems to have been overlooked by the ‘RE nutters in residence’ so far, including myself.

Reduction with K, with Ca, alcoholic electrolysis of RECl3 solutions with Hg cathode to REE amalgam and details on electrolysis of anh. RECl3/KCl melts are but a few topics covered. Details on preparing anh. RECl3 is also covered. And for the very curious (Brain&Force?): preparation of some REE(II) salts too, like EuSO4.

[Edited on 14-10-2014 by blogfast25]

Amos - 14-10-2014 at 06:15

Does anybody else think we have a new resident troll here on SM?

blogfast25 - 14-10-2014 at 06:46

This was the LaCl3 solution at the start of the boiling process, light green:

Here after concentrating and chilling, with crystals of presumed LaCl3.7H2O on the bottom. The solution is over 50 w% LaCl3:

[Edited on 14-10-2014 by blogfast25]

Little_Ghost_again - 14-10-2014 at 07:09

Quote: Originally posted by No Tears Only Dreams Now

Does anybody else think we have a new resident troll here on SM?

I am not that new FFS

Little_Ghost_again - 14-10-2014 at 07:11

I rather suspect you are missing a few critical things.

Would they be just above the neck?

hyfalcon - 14-10-2014 at 07:36

Maybe if he could manipulate matter with will alone he might get somewhere. Otherwise, he doesn't stand a chance.

Bert - 14-10-2014 at 07:54

Please let us now restrict further remarks to the science and engineering aspects of the OP's topic.

If being trolled, DO NOT FEED THE TROLL. It will go away or die.

If site rules are violated, user that little blue "report" button or PM a moderator.

And I JUST LOVE the swimming pool chemical supply section at various stores! Right up there with the ceramic glaze supply, photographic supply and auto supply stores.

Amos - 14-10-2014 at 07:58

Quote: Originally posted by Bert

Please let us now restrict further remarks to the science and engineering aspects of the OP's topic.

If being trolled, DO NOT FEED THE TROLL. It will go away or die.

If site rules are violated, user that little blue "report" button or PM a moderator.

And I JUST LOVE the swimming pool chemical supply section at various stores! Right up there with the ceramic glaze supply, photographic supply and auto supply stores.

I wish I had the 2nd and 3rd ones

Brain&Force - 14-10-2014 at 09:25

Quote: Originally posted by blogfast25

I’ll summarise my work on the lanthanum bearing (10.3 % La glycolate, acc. MSDS) pool phosphate remover called ‘Lo-Chlor Starver’ (from Lo-Chlor Chemicals, AU) as follows.

2.5 L of this solution was neutralised with 33 % NH3 and the snow-white precipitate, presumed lanthanum hydroxide, quantitatively Buchnered and washed. This yielded about 1 L of product.

It was dissolved gradually and quantitatively in 280 ml HCl 37 % with mild heating, followed by simmering. It dissolved completely to a clear solution that was… slightly green! On further boiling that colour changed to yellow. I have very strong reasons to believe this colour is due to MORE than a bit of ferric chloride, more about that later. Possibly there is Pr present.

The solution (about 1.25 L) was then gradually boiled down to about 290 g, by which time it was urine yellow. On refrigeration a mass of small, white crystals formed (photos later) although much of the presumed lanthanum chloride heptahydrate remained in solution, as expected. The rest of the LaCl3 will now be recovered using the double sulphate method (with K2SO4), to try and separate it from the yellow contamination. If it cannot be separated I’m inclined to believe the original phosphate remover contained also Praseodymium. If it can, that would point to ferric chloride.

As regards possible methods of preparing La (and other REE) metals, I proposed the reduction of LaF3 with Li, which appears to be a thermodynamically favourable reaction, possibly with enough reaction heat to ensure clean separation between metal and slag. LaF3 is water insoluble, so fairly easy to prepare.

I earlier dismissed the use of reduction of anh. RECl3 but Brauer’s ‘Handbook of Preparative Inorganic Chemistry’ proved me wrong on that. I suggest strongly that anyone interested in contributing to this discussion go to the Library of this site, download Brauer and print off pages 1135 to 1150 from Volume 2.

There is an absolute wealth of information there that seems to have been overlooked by the ‘RE nutters in residence’ so far, including myself.

Reduction with K, with Ca, alcoholic electrolysis of RECl3 solutions with Hg cathode to REE amalgam and details on electrolysis of anh. RECl3/KCl melts are but a few topics covered. Details on preparing anh. RECl3 is also covered. And for the very curious (Brain&Force?): preparation of some REE(II) salts too, like EuSO4.

[Edited on 14-10-2014 by blogfast25]

I'm actually downloading some references from the library for the first time. Frankly, this is embarrassing. I don't have europium (yet) but I will be making some EuSO₄ shortly after I have some.

blogfast25, potassium sulfate may not be the ideal reagent for precipitating out the ferric ions. The sulfates of very light REEs appear to be relatively soluble, per the data I linked earlier:

The oxalate method may be preferable, as all lanthanide oxalates are ridiculously insoluble.

blogfast25 - 14-10-2014 at 10:44

B&F:

The double sulphates method has been prescribed many times for La to Pr. The REE sulphate solubilities may not reflect the double sulphate solubilities very well.

I don't like the oxalates: the only thing you can do with them is to calcine to oxide. And if it is iron (?) iron oxalates aren't that soluble either, unless you complex to trisoxalato ferrate (III).

[Edited on 14-10-2014 by blogfast25]

blogfast25 - 14-10-2014 at 13:13

For laughs:

Here's a twit who claims he's 'thermited' Y2O3:

http://www.sciencemadness.org/talk/viewthread.php?tid=10249&...

Quote: Originally posted by siegfried

After a semi-failure with zinc thermite (the entire pile of thermite did not burn), I tried using 5 grams of Y2O3 and 2 grams of powdered Al. The mixture was placed on a piece of tile flooring outside with an air temp of 40 degrees F.

The results were very gratifying. The mixture burned much quicker than conventional iron thermite. The temperature was high because the slab of tile cracked and all that was left were chunks of fused Yttrium metal and Al2O3.
I have no way of measuring the actual temperature of the reaction but the melting point of Y2O3 is 2,410C, Yttrium metal melts at 1,522 and boils at 3,338.

The 99.99% pure Y2O3 was purchased on Ebay and, for a "rare earth" metal oxide, was relatively cheap.

It burned even better than conventional iron thermite!

Bert - 14-10-2014 at 14:16

Quote: Originally posted by No Tears Only Dreams Now

Quote: Originally posted by Bert

And I JUST LOVE the swimming pool chemical supply section at various stores! Right up there with the ceramic glaze supply, photographic supply and auto supply stores.

I wish I had the 2nd and 3rd ones

For the price of shipping, you can...

The Photographers Formulary

Duda Diesel (click on "all chemicals")

The Big Clay Store

[Edited on 14-10-2014 by Bert]

Metacelsus - 14-10-2014 at 14:45

I highly recommend Duda Diesel. The prices are good (on most things), and I've never had a problem with them.

Texium - 14-10-2014 at 19:22

Their shipping is a bit steep though. You can buy a lot of their products for a slightly higher price but a much lower shipping cost from their Amazon store.

MrHomeScientist - 15-10-2014 at 06:43

Quote: Originally posted by No Tears Only Dreams Now

Does anybody else think we have a new resident troll here on SM?

Actually I think we have an old resident troll - ThePHDChemist returns.

Anyways on topic, I'm always interested in reading about rare earth experiments. I'm pretty surprised a lanthanum compound can be found at a pool store!

Back off topic: As for my own RE exploits, I now should have enough NdF₃ and Chinese lithium to try that route to Nd metal. If it works, that's a good sign for lanthanum as well. The thing that concerns me, though, is molten lithium's reactivity with my crucible (graphite or fused silica) and it's "explosive" reaction with concrete floors that I mentioned in another thread. Any recommendations for a better crucible material?

blogfast25 - 15-10-2014 at 10:36

Anyways on topic, I'm always interested in reading about rare earth experiments. I'm pretty surprised a lanthanum compound can be found at a pool store!

Back off topic: As for my own RE exploits, I now should have enough NdF₃ and Chinese lithium to try that route to Nd metal. If it works, that's a good sign for lanthanum as well. The thing that concerns me, though, is molten lithium's reactivity with my crucible (graphite or fused silica) and it's "explosive" reaction with concrete floors that I mentioned in another thread. Any recommendations for a better crucible material?

Mr HomeScientist:

Graphite would be my choice. I actually have a small graphite crucible and if I go ahead with LaF3 + 3 Li that's the one I would use.

I can't really see any significant interaction between C and Li or Nd but I could be horribly wrong on that.

Not sure what you mean by the '"explosive" reaction with concrete floors', I must have missed that bit. I assume you'll be working on small scale, so that reduces any risk enormously anyway.

Also interesting: Wiki's entry on Yttrium mentions reduction of YF3 with an alloy of Ca and Mg.

Other Tidbit:

The rather unusual co-precipitation behaviour of various La precipitates was used in the development of the Bomb, to 'decontaminate' Pu239. These precipitates would concentrate Pu239 (also some Np isotopes) from the liquors obtained from dissolving irradiated uranium.

[Edited on 15-10-2014 by blogfast25]

MrHomeScientist - 15-10-2014 at 12:00

Here's the post where I mentioned the concrete, with sources attached: http://www.sciencemadness.org/talk/viewthread.php?tid=32946&...

I guess the best course of action would be to initially melt a few slugs of Li in my graphite crucible and see what happens. From a safe distance away!

Brain&Force - 15-10-2014 at 12:09

Yeah, NurdRage's lithium extraction video mentions that pretty much any non-metal surface will cause an explosion on contact with molten lithium. I'm torn on the reaction of graphite though - maybe a lithium carbide will form?

blogfast25 - 15-10-2014 at 12:41

Yeah, NurdRage's lithium extraction video mentions that pretty much any non-metal surface will cause an explosion on contact with molten lithium.

Which video is that, B&F?

Brain&Force - 15-10-2014 at 22:18

Get Lithium Metal From an Energizer Battery: http://youtu.be/BliWUHSOalU

It's in the annotations. My own experience agrees with this.

blogfast25 - 16-10-2014 at 04:28

Get Lithium Metal From an Energizer Battery: http://youtu.be/BliWUHSOalU

It's in the annotations. My own experience agrees with this.

Well, describe your own experience because I can't see any annotations that state that.

However aggressive liquid lithium may be, it still needs something to react with for any 'explosion' to occur.

Do we know how it behaves towards steel, for instance? Does it alloy iron easily?

With a heat of formation of Li2O of about - 600 kJ/mol (NIST) Li would reduce both alumina and silica, thereby excluding ceramics as crucible materials. And glass, of course.

[Edited on 16-10-2014 by blogfast25]

MrHomeScientist - 16-10-2014 at 05:55

According to the sources in my link above, "Molten lithium reacts explosively with concrete flooring, and any area
wherein a liquid lithium spill may occur must have welded steel flooring."

So steel appears to be fine. Which suggests another potential crucible that I have a few of - those stainless steel condiment cups.

blogfast25 - 16-10-2014 at 11:39

So steel appears to be fine. Which suggests another potential crucible that I have a few of - those stainless steel condiment cups.

I was afraid the REE might stick to steel on solidifying, hence my anticipated choice of graphite.

I’m now almost 100 % certain that the yellow contamination in my lanthanum based phosphate remover is our old friend Fe(III) and not partly Pr (III) as I previously thought. Experiments with precipitating the La as double potassium sulphate shows clearly that the contamination does not precipitate but stays in the supernatant, while reacting positively with thiocyanate.

I also isolated some of the early crystals with difficulty because the syrupy supernatant (it is about 50 w% LaCl3!) makes filtering very difficult. So decanting off the supernatant as best as possible was the only short term option.

I tested the crystals (with some supernatant clinging to them) with acetone and there appears to be no dissolution of the crystals but the yellow colour does concentrate in the acetone phase (ferric chloride is acetone soluble). The slurry is difficult to filter. Should have my first dry and iron free LaCl3 hydrate crystals tomorrow.

This also opens up the possibility of using acetone as an anti-solvent to precipitate fairly pure LaCl3, iron free.

Brain&Force - 16-10-2014 at 12:22

Yes, it'll work acc. to this reference:

http://books.google.com/books?id=uTmc-BeVbZoC&pg=PA439&a...

blogfast25 - 16-10-2014 at 13:14

Yes, it'll work acc. to this reference:

You mean acc. that ref. acetone is an anti-solvent for LaCl3?

Brain&Force - 16-10-2014 at 14:33

It appears that lanthanum chloride is insoluble in acetone, so it'll just drop out.

You mean by antisolvent, "less soluble?"

j_sum1 - 16-10-2014 at 16:47

Blogfast
Thought you might be interested -- your work here is being used to inspire some young chemistry students.
I was teaching my chemistry class this morning. Lesson two in learning about equilibrium. We had dome a whole bunch of demonstrations of equilibrium systems involving transition ions and their complexes. Discussion of the various things we might change to shift the equilibrium point, selectively precipitate, concentrate products etc. Returned to the textbook to reinforce the concept and fill in the gaps and the example was the Fe3+, SCN- = FeSCN2+ system.
Next was a quick trip to this thread. Here you are adding ammonia, selectively precipitating, heating, cooling, filtering, adding HCl, boiling down to increase concentration and finish up with using SCN- to test for Fe3+ ions.

Thanks for writing my lesson for me. Seriously though. It is great to be able to give students something fresh and practical and keep interest levels high.

Edited to add...
The pictures posted and the write up of what you had done was also really beneficial in reinforcing principles of record-keeping and documentation that we have also been discussing this year.

[Edited on 17-10-2014 by j_sum1]

Brain&Force - 16-10-2014 at 17:46

It would be particularly interesting to do some experiments with lanthanides in a school chemistry lab. There's lots of color changing and fluorescence. If only they weren't so expensive.

On the topic of electrolysis of lanthanum compounds, I may have access to a bunch of gadolinium chloride soon and I may attempt reducing it back to the metal, per material in Brauer.

[edit] I found another rare earths reference ("A Text-book of Inorganic Chemistry Volume IV") but it is dated 1917 and it seems to be a blank pdf, save for page 31, I believe.

[Edited on 17.10.2014 by Brain&Force]

blogfast25 - 17-10-2014 at 04:41

@B&F:

'antisolvent': clumsy term for 'not a solvent'. Adding it to a watery solution causes the solute to drop out. Seems to be the case here... Interesting as FeCl3 will not precipitate. Could certainly beat hours of boiling down too! I will experiment with this for the rest of the batch.

@j_sum1: thanks and glad this work serves at least one purpose. RE "selectively precipitating", selective precipitation of the REE hydroxides with ammonia has actually been done: the lighter REE are slightly more basic than the heavier ones due to the ionic radius contraction, so they can be made to precipitate at slightly higher pH. It was von Welshbach (IIRW but don't quote me on that) who used this to separate some REEs.

[Edited on 17-10-2014 by blogfast25]

blogfast25 - 17-10-2014 at 10:49

The first batch of lanthanum potassium double sulphate was Buchnered off and washed with about 100 ml cold 1 M H2SO4 saturated with K2SO4. The La/K double sulphate is very reminiscent of the Nd/K double sulphate: both are sandy, non fluffy precipitates that filter and wash very well.

The filter cake was suspended in 400 ml water and then 30 ml of 33 w% NH3 was added. The texture of the slurry changed immediately as the K sulphate is ‘released’ and the lanthanum sulphate converts to lanthanum hydroxide, right in the photo:

The slurry was digested at about 90 C with magnetic stirring for about 30 minutes and will be filtered and washed tomorrow. Then it will be dissolved in strong HCl to obtain the iron free lanthanum chloride.

To the left is the rest of the batch, already converted to double sulphate. The yellow iron can just about be seen in the supernatant.

===================

I also played around with the first teaspoon of lanthanum chloride crystals, still coated in thick syrup containing acetone and water. I can certainly conform the lanthanum chloride hydrate appears to be quite insoluble in acetone (99.5 %). But there’s some strange goings on too. Adding more acetone to the slurry thins it but it seems only temporarily. Each time I decanted off the liquid phase, a new syrupy liquid formed. I’m not sure how to explain that.

In the end I added 20 ml of water in which it dissolved quickly. Then I boiled off that little bit of acetone, followed by most of the water, adjusting power to avoid bad bumping. From this hydrate, expect no pretty crystals: the material gradually turned into a clear, colourless semi-solid mass (this is in line with my more limited experience with NdCl3, minus the 'colourless' of course). After about a half hour I could not see any solvent coming off anymore. This is it a bit before the very end:

[Edited on 17-10-2014 by blogfast25]

blogfast25 - 19-10-2014 at 08:19

Now it looks like praseodymium is back on the agenda.

After re-precipitating the hydroxide, careful washing with hot DIW to remove sulphates and dissolving it again in hot 37 % HCl, the obtained solution was identical in colour to previously: greenish when cold and pure yellow when hot.

Despite the effort for removing all Fe3+ the solution still tested weakly positive for it. Presumably some co-precipitation of ferric sulphate with the La/K double sulphate took place. Grr. It also tested very, very weakly positive for sulphate (with barium nitrate).

Then 25 ml of the solution (about 1 M of La) was precipitated as hydroxide with ammonia and the precipitate filtered and washed. It looked perfectly white. It was thoroughly dried on a hot plate and ground into a powder.

I then took about half a teaspoon of this powder and calcined it in a ceramic crucible at max. propane Bunsen heat. Almost immediately the edges of the powder started to darken. After about 15 minutes the whole mass had turned dark grey to black. That could be indicative of Pr(IV), of course. Calcined Pr oxide is also dark due to Pr(IV). Without eliminating the iron completely it’s not the strongest evidence possible. But ferric iron is thermochromic in a different way: solutions turn slightly darker when heated due to Fe(OH)+ formation, not greenish to yellow.

The calcined oxide dissolved into hot, 37 % HCl but only slowly and incompletely, even after 30’ of refluxing there was some blackness left.

I’m tempted to try pH selective precipitation with NH3: Fe3+ and Pr3+ should gather in the first permanent precipitate if this works.

[Edited on 19-10-2014 by blogfast25]

j_sum1 - 19-10-2014 at 19:31

This is looking interesting.

A couple of questions.
You began by precipitating using NH3. Is this just because it is what you had on hand? Couldn't the same La(OH)3 precipitate be formed using NaOH. Any advantage to the ammonia? Or is it related to the La glycolate that you began with? (I am not really familiar with glycolates.)

You are quick to narrow down the impurity to Fe3+ and Pr(3+?). Is this just based on the colour? Aren't Yttrium, Cerium and Neodymium also possibilities? Is the possible formation of complexes a hindrance to positive identification? And while we are at it, what happens to Pr3+ in the presence of SCN-. Could there be a false positive for Fe3+?

For me, given that my product is stated as being LaCl3, I am probably not going to go through all the hoops of trying to purify it too much. If I manage to get a reduction to a metal that contains at least a decent proportion of La then I will be more than happy.

With that in mind I am looking forward to hearing from MrHomeScientist and his NdF3/Li adventures with the steel crucible.

[edited spelling]

[Edited on 20-10-2014 by j_sum1]

Brain&Force - 19-10-2014 at 22:16

From what woelen has done work on praseodymium(III) chloride is much like manganese(II) chloride in its color - it only becomes visible in nearly saturated solutions.

Are you sure a tetrachloroferrate(III) complex isn't responsible for the color change? If all else fails acetone seems the way to go.

blogfast25 - 20-10-2014 at 08:47

You are quick to narrow down the impurity to Fe3+ and Pr(3+?). Is this just based on the colour? Aren't Yttrium, Cerium and Neodymium also possibilities? Is the possible formation of complexes a hindrance to positive identification? And while we are at it, what happens to Pr3+ in the presence of SCN-. Could there be a false positive for Fe3+?

I’m not quick AT ALL: having no access to UV/VIS I’m taking into account circumstantial pieces of evidence and it’s very slow and tentative. Y, Ce and Nd are all possibilities but even harder to see without UV/VIS and what I can’t see I can’t worry about. Pr3+ doesn’t not complex with SCN-, that’s unique to Fe3+, as far as I know.

The evidence for Pr(III), not in order of importance:

• Strange yellow/green colour of solutions and slight thermochromism.
• Coloured impurity cannot be removed by double sulphate method. Pr(III) fits the bill because it also forms a poorly soluble K/Pr double sulphate: it stays with the La i.o.w.
• Hydroxide does not appear tinged with Fe(OH)3 which surely would be the case at high levels of Fe. It’s snow white. Pr(OH)3 is white, I think.
• Calcining the hydroxide causes it to turn black/dark grey: that doesn’t fit Fe(OH)3 very well either. But it does Pr(IV), formed on calcining (Pr6O11 mixed oxide)…

With that in mind I am looking forward to hearing from MrHomeScientist and his NdF3/Li adventures with the steel crucible.

Yes, ditto here.

From what woelen has done work on praseodymium(III) chloride is much like manganese(II) chloride in its color - it only becomes visible in nearly saturated solutions.

Are you sure a tetrachloroferrate(III) complex isn't responsible for the color change? If all else fails acetone seems the way to go.

I hate it when people invoke all kind of complexes to imperfectly try and explain small mysteries (‘Don’t know what’s going on? Blame an unspecified complex!’). Hidden parameter theories.

I’ve boiled in dilute ferric chloride in 20 % HCl to almost dryness in the past. NEVER did I see any green. The solutions go from yellow to darker to eventually almost black. Yellow, cold solutions darken visibly when heated due to Fe(OH)+, reversibly so.

Do you have any evidence/references of a tetrachloroferrate (III) complex and its colour?

[Edited on 20-10-2014 by blogfast25]

Brain&Force - 20-10-2014 at 10:25

Yellowish orange, it seems. You're right, as I think I'm confusing it with some iron(II) complex. After seeing this post it's evident that there is plenty of praseodymium.

Do you see any color change in different lighting?

blogfast25 - 20-10-2014 at 11:19

After seeing this post it's evident that there is plenty of praseodymium.

Do you see any color change in different lighting?

I'm still not 100 % convinced but find it really difficult to otherwise explain what I've seen.

Unfortunately we've changed over to saver bulbs completely in our house. But I do have a powerful incandescent torch that needs batteries. Will definitely have a look...

[Edited on 20-10-2014 by blogfast25]

MrHomeScientist - 20-10-2014 at 12:20

Quote: Originally posted by blogfast25

With that in mind I am looking forward to hearing from MrHomeScientist and his NdF3/Li adventures with the steel crucible.

Yes, ditto here.

I guess I need to quit wasting time playing Destiny and get to it then!

Lithium makes me nervous though, to be honest. Dan Vizine in another thread cited that lithium's reactivity can change drastically depending on temperature. Just melting Li in my crucible as a 'dry run' doesn't necessarily mean I'll get the same results when NdF₃ is around. The goal of producing molten Nd will take me to 1024 C, far above Li's melting point of 180 C (but within my furnace's capability, as I've melted copper before). I'd like to try it this week if at all possible. I'll post results in the neodymium thread to avoid further derailment but link to it from here for continuity.

blogfast25 - 21-10-2014 at 04:15

Let me try and put some of your fears to rest.

The reaction between NdF3 and Li, assuming it proceeds (thermodynamics say nothing about kinetics, remember), is likely to start at 400 to 600 C and the hope is that the calculated ΔH is enough to take the reaction products to above 1024 C (the MP of LiF is lower). This week I'll try and calculate the actual exotherm ΔT from NIST Shomate data.

Melting Li without argon blanket will set fire to it, so be prepared for that during your dry run. It burns extremely brightly apparently.

What physical form is your Li metal?

Edit:

Adiabatic exotherm for that reduction, based on - 194 kJ/mol reaction heat comes out as ΔT about 830 C. So even if auto-ignition would start as low as 200 C, end temperature should still surpass both MPs.

[Edited on 21-10-2014 by blogfast25]

MrHomeScientist - 21-10-2014 at 07:45

Here is my Li metal, stored under argon:

Small slugs, around 3/8" long. The plan will be to add NdF₃ powder to the crucible first then a few of these on top. The Li will melt and cover the fluoride before reaction starts. I think my biggest concern at this point is the hot, molten lithium reacting with whatever crucible material I use. I currently have fused silica (no good), graphite (questionable), and stainless steel (slightly less questionable). I think it would also be a good idea to have a cover for whatever material I end up using, to help exclude air.

Edit: Your delta-H calculation is very promising! Even if this did not produce enough heat on its own, I could reach that temperature with my furnace so I think it's looking good.

[Edited on 10-21-2014 by MrHomeScientist]

blogfast25 - 21-10-2014 at 08:17

Nice metal (you couldn't teleport some over here, could you?)

But I was thinking of doing it the other way around: load lithium first, then fluoride, so the latter protects the metal somewhat, at least at first. Then when it melts reaction will start at some point. Argon blanket would of course be better. I'm thinking of a fairly simple design: crucible within a crucible for instance. A wide copper tube (shut off at one end), pump argon into it, load crucible into it and then load reagents into inner crucible. Lower assembly into furnace, maintaining low Ar flux into outer crucible...

For what it's worth on Li2C2 Wiki mentions this:

"To prepare pure samples in the laboratory molten lithium + graphite are reacted at high temperature."

One of the entry's references mentions 800 - 900 C, so graphite may be more problematic than I previously thought.

No oxidic materials are really recommendable so that kind leaves various metals, like SS.

What scale are you attempting this? One slug? Less?

[Edited on 21-10-2014 by blogfast25]

MrHomeScientist - 21-10-2014 at 12:24

I'd be happy to send you a sample, assuming international shipping isn't a big deal. My source is the Chinese eBay seller woelen found not too long ago - he's still got some available.
What I received isn't nearly as clean as his picture, but apparently woelen got great quality. Chinese (lack of) quality control I guess, but definitely worth the price in my opinion.

My only source of argon right now is bloxygen canisters (little aerosol cans of Ar), so I can't really set up a continuous blanket. The 'nested crucible' idea is pretty good; I might have to check out the hardware store. Maybe some black iron fittings and caps, after degreasing with acetone, would work for both crucible and outer casing. Do you think it would be particularly disastrous with no inert cover?

Going by 3Li + NdF₃ == Nd + 3LiF , 0.5g Li would react with 4.79g NdF₃ to produce 3.5g Nd metal. That seems a good size to work but minimize danger. I'll have to weigh the slugs to see how many I'd need. I'd guess 2-3. The issue with small amounts is that my steel crucible is fairly wide, so I'd need to tilt it a little so the Nd is fully covered post reaction.

This has gone fairly off-topic for this thread, for which I apologize. But, if it works, the lithium reduction should be applicable to lanthanum as well!

blogfast25 - 22-10-2014 at 04:15

Ah yes, I saw that seller too. Good price.

Here argon for welding is pretty cheap, proper canister and expander included.

I don't know how deleterious oxygen will be. The worst case scenario is that your 0.5 g of Li ends up burnt to Li2O, I guess.

Dan Vizine - 22-10-2014 at 08:24

Here is my Li metal, stored under argon:

Small slugs, around 3/8" long. The plan will be to add NdF₃ powder to the crucible first then a few of these on top. The Li will melt and cover the fluoride before reaction starts. I think my biggest concern at this point is the hot, molten lithium reacting with whatever crucible material I use. I currently have fused silica (no good), graphite (questionable), and stainless steel (slightly less questionable). I think it would also be a good idea to have a cover for whatever material I end up using, to help exclude air.

Edit: Your delta-H calculation is very promising! Even if this did not produce enough heat on its own, I could reach that temperature with my furnace so I think it's looking good

[Edited on 10-21-2014 by MrHomeScientist]

Even though molten Li is fearsome to many materials, unless you want to pay big money for the "right" materials (a tantalum crucible liner slipped into a larger crucible), or at minimum, a ferritic SS like something in 400 series, the default material is likely to be SS304 or pure iron. It is a compromise, to be sure.

Here's why...as the nickel content in an alloy increases, it's resistance to heat usually increases while its resistance to leaching by hot, molten lithium or calcium decreases. This results in some leaching of Ni and Fe but not to destructive failure. For SS 304, leached material may be several tenths of a per cent in your final product. You can make the reactor and the material contacting the melt out of just one thing. This greatly simplifies construction, usually.

Pure iron is available, relatively cheap and resists hot, molten Li and Ca quite well. The problem is that you have to rely on whatever shapes you can find for sale (depending on your mechanical skills and equipment).
It is best used as a liner in a more oxidation resistant larger container.

[Edited on 23-10-2014 by Dan Vizine]

blogfast25 - 22-10-2014 at 09:31

Thanks Dan.

Finally, even stronger proof there is a considerable amount of praseodymium in this lanthanum.

Remember that I had calcined about a quarter of a teaspoon of the lanthanum hydroxide and that it had turned dark grey to black. That oxide was then subjected to an excess hot HCl 37 %. It turned out that nearly all of it had dissolved minus a black residue which I suspect is the mixed Pr(IV,III) oxide.

Today I isolated it with some mini decantations and then dried it by boiling off the supernatant in a 40 ml ceramic crucible. Black as the ace of spades:

About 2 ml of 98 % sulphuric acid was added and strongly heated to nearly 400 C (the maximum of my heater/stirrer). Despite strong fumes of SO3 coming off I could not see the acid making a dent in the black stuff (no wonder it didn’t dissolve in conc. HCl!) So after about 20 minutes I cooled it down to reasonable levels and added about 5 g of NaHSO4. This was heated to medium high propane Bunsen heat and it did the trick. Bar a few black bits that had crept up to near the top of the crucible the colour disappeared in about 10 minutes. Here’s what it looked like after cooling:

The mass was soaked in about 10 ml of hot water until it released from the crucible and then transferred to a glass beaker. With stirring/heating it dissolved to a thin slurry with white insoluble particles, presumed to be poorly soluble Na/Pr double sulphate.

To it was added 5 g of NaOH dissolved in 10 ml of DIW and that mixture simmered/stirred for about 20 minutes. This was then filtered with vacuum on a ceramic frit filter and washed with minimal amounts of hot water. The presumed hydroxide was scraped off the filter, amounting to about 0.05 ml of green material which was dissolved in a few drop of conc. HCl forming a green solution. Compared to a tube of DIW:

The photo doesn’t do it justice at all, making even the white kitchen tissue look yellow! It’s quite an intense, slightly ‘dirty’ green. It may look darker than it is due to a few particles of black residue that had carried over.

[Edited on 22-10-2014 by blogfast25]

Dan Vizine - 22-10-2014 at 10:21

Oh Oh...I read more.

I also apologize for the off-topic comments, but Mr.Home Scientist, I hate to spoil the party, but when you asked "Do you think it would be particularly disastrous with no inert cover?", I'm sorry to say that you don't have a prayer. You need to protect Li at 1000 C from oxygen, moisture and nitrogen completely. If you want good results, you'll need good quality argon, you need to sweep all of the atmospheric gas out and maintain a quality cover. The last part is the hardest.

The good thing is that burning lithium is kind of quiet, not too violent. But, oh those fumes...intensely irritating.

blogfast25 - 22-10-2014 at 11:46

I also apologize for the off-topic comments, but Mr.Home Scientist, I hate to spoil the party, but when you asked "Do you think it would be particularly disastrous with no inert cover?", I'm sorry to say that you don't have a prayer. You need to protect Li at 1000 C from oxygen, moisture and nitrogen completely.

If I may interject here. We are hoping that this reaction will start much, much lower than 1000 C.

But argon would certainly be better, no contest.

MrHomeScientist - 22-10-2014 at 12:00

Right, hopefully it will start not too far past the melting point of Li. But the desired end result is at least 1024C, for complete melting of all the products. Will most of the lithium be gone by then, and not pose a large fire risk / problem to the crucible? Unsure at this point. I hope by starting with 0.5g of Li risks will be minimized.

I should probably just say damn it all and go for it, and see what happens. If there is severe attack of the crucible I'd like to try lining it with some tantalum foil I have, if for no other reason than to find a use for it.

Edit:
Great progress on the lantanum, blogfast. About how much black material did you have? i.e. what percent Pr would you estimate this material has? How would bisulfate dissolve it when concentrated sulfuric would not?

[Edited on 10-22-2014 by MrHomeScientist]

blogfast25 - 22-10-2014 at 12:40

Great progress on the lantanum, blogfast. About how much black material did you have? i.e. what percent Pr would you estimate this material has? How would bisulfate dissolve it when concentrated sulfuric would not?

The amount of black material was small because I'd only calcined about half a teaspoon of the contaminated lanthanum hydroxide. Probably about 50 - 100 mg.

I'm not quite sure why the bisulphate made it work. It certainly increases the BP of the mixture. It also converts to pyrosulphate (S2O72-

on melting. I've used it a few times on stubborn residues/oxides, with success usually.

I don't know how much Pr is in that pool phosphate remover but going by everything that I've seen, I'd say 20 mol % easily of La + Pr. I'll definitely have a stab at getting it out.

[Edited on 23-10-2014 by blogfast25]

j_sum1 - 22-10-2014 at 14:44

@MrHomeScientist
In case you missed it, I posted in the lithium thread details of a molybdenum vessel at low price that you might find useful.
http://www.sciencemadness.org/talk/viewthread.php?tid=32946&...

Dan Vizine - 22-10-2014 at 20:38

Bisulfate and heat rips into many things that even the strongest, most aggressive acids fail to touch.

Nice detective work blogfast25 !

[Edited on 23-10-2014 by Dan Vizine]

blogfast25 - 23-10-2014 at 04:26

Nice detective work blogfast25 !

Thanks Dan. It really felt like a whodunit. Could have been over in 1/2 hour with VIS spectrometry but not as much fun as with a good old looking glass, I guess.

What do you think about molten Li + copper crucible?

j_sum1 - 23-10-2014 at 04:42

Well if it gets up to temp you might get molten copprr as well.
I wonder what the Cu La phase diagram looks like. I bet that's never been done.

plante1999 - 23-10-2014 at 04:46

Very interesting work being done here, after the experimenting is done, it might be nice to make a short summary of the work done.

blogfast25 - 23-10-2014 at 05:14

Well if it gets up to temp you might get molten copprr as well.
I wonder what the Cu La phase diagram looks like. I bet that's never been done.

Yep. Off the top of my head I thought Cu MP was 1400 K but that's K. Bummer.

Quote: Originally posted by plante1999

Very interesting work being done here, after the experimenting is done, it might be nice to make a short summary of the work done.

Will do that on my blog. This weekend I will attempt pH selective precipitation/fractionation.

Here's a slightly better pic of the green REE chloride. Slightly more dilute, in natural sunlight and with a few grans of the black oxide at the bottom:

[Edited on 23-10-2014 by blogfast25]

Dan Vizine - 23-10-2014 at 08:15

Quote: Originally posted by blogfast25

Nice detective work blogfast25 !

Well, I see that the Cu question was addressed. The suggested materials for molten Li are TMZ Molybdenum and pure iron. But the reality is that manufacturers use tantalum or tungsten liners or crucibles inside SS reaction bombs to reduce LaF3. The most frequent reductants are Li and Ca. The first successful (I2 boosted) reactions were done in dolomite-lined steel bombs, but the purity wasn't good enough.

I'd say that the choice of the best crucible material for your application depends largely on how you plan to exclude oxygen. If a gas-shielded open crucible is planned, I'd go with pure iron. Pure iron crucibles are on eBay for very low cost.

[Edited on 23-10-2014 by Dan Vizine]

Brain&Force - 23-10-2014 at 09:48

I've got to ask, have you excluded the possibility of any other lanthanides in the mix? I wouldn't be surprised, though, if all of the neodymium and samarium were removed to make magnets, and the cerium was oxidized away in the beginning. I doubt any other lanthanides would be present, but yttrium could potentially be present.

I'd like to feature this thread on my radio show. If you're interested in calling in to talk about this, U2U me so we can iron out the details. Also check the relevant links in my sig.

blogfast25 - 23-10-2014 at 11:36

The suggested materials for molten Li are TMZ Molybdenum and pure iron. But the reality is that manufacturers use tantalum or tungsten liners or crucibles inside SS reaction bombs to reduce LaF3. The most frequent reductants are Li and Ca. The first successful (I2 boosted) reactions were done in dolomite-lined steel bombs, but the purity wasn't good enough.

"TMZ"?

You have references for the reduction of LaF3? We thought we were being trailblazers here because all I can find are methods like the ones in Brauer: molten chlorides electrolysis, chlorides reduction with Ca boosted with I2 or chlorate, mercury amalgam (electrolysis) and even oxides reduction of higher REE with lanthanum at high T.