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blogfast25
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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]
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blogfast25
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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! 
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Bert
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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]
Rapopart’s Rules for critical commentary:
1. Attempt to re-express your target’s position so clearly, vividly and fairly that your target says: “Thanks, I wish I’d thought of putting it
that way.”
2. List any points of agreement (especially if they are not matters of general or widespread agreement).
3. Mention anything you have learned from your target.
4. Only then are you permitted to say so much as a word of rebuttal or criticism.
Anatol Rapoport was a Russian-born American mathematical psychologist (1911-2007).
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Metacelsus
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I highly recommend Duda Diesel. The prices are good (on most things), and I've never had a problem with them.
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Texium
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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.
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MrHomeScientist
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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?
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blogfast25
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| Quote: Originally posted by MrHomeScientist | 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?
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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 YF<sub>3</sub> 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]
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MrHomeScientist
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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!
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Brain&Force
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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?
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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| Quote: Originally posted by Brain&Force | | Yeah, NurdRage's lithium extraction video mentions that pretty much any non-metal surface will cause an explosion on contact with molten lithium.
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Which video is that, B&F?
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Brain&Force
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Get Lithium Metal From an Energizer Battery: http://youtu.be/BliWUHSOalU
It's in the annotations. My own experience agrees with this.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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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 Li<sub>2</sub>O 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]
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MrHomeScientist
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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.
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blogfast25
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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% LaCl<sub>3</sub>!) 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 LaCl<sub>3</sub> hydrate crystals tomorrow.
This also opens up the possibility of using acetone as an anti-solvent to precipitate fairly pure LaCl<sub>3</sub>, iron free.
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Brain&Force
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Yes, it'll work acc. to this reference:
http://books.google.com/books?id=uTmc-BeVbZoC&pg=PA439&a...
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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You mean acc. that ref. acetone is an anti-solvent for LaCl<sub>3</sub>?
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Brain&Force
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It appears that lanthanum chloride is insoluble in acetone, so it'll just drop out.
You mean by antisolvent, "less soluble?"
At the end of the day, simulating atoms doesn't beat working with the real things...
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j_sum1
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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]
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Brain&Force
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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]
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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@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]
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blogfast25
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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]
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blogfast25
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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 Fe<sup>3+</sup> 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)<sup>+</sup> 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: Fe<sup>3+</sup> and Pr<sup>3+</sup> should gather in the first
permanent precipitate if this works.
[Edited on 19-10-2014 by blogfast25]
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j_sum1
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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]
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Brain&Force
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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.
At the end of the day, simulating atoms doesn't beat working with the real things...
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blogfast25
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| Quote: Originally posted by j_sum1 | | 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)…
| Quote: Originally posted by j_sum1 |
With that in mind I am looking forward to hearing from MrHomeScientist and his NdF3/Li adventures with the steel crucible. |
Yes, ditto here.
| Quote: Originally posted by Brain&Force | 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)<sub>+</sub>, reversibly so.
Do you have any evidence/references of a tetrachloroferrate (III) complex and its colour?
[Edited on 20-10-2014 by blogfast25]
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