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Eddygp
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This should really be in bfesser's Topical Compendium, under terbium. Very interesting topic. In fact, I like the lanthanides' chemistry as it is
fascinating. I'm going to try a similar synthesis soon, with ytterbium (hopefully).
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blogfast25
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That would be very interesting. Don't forget to include some good quality photos!
We should take advantage of the increasing availability of the REs: when I was a student the lanthanide block was like a beach you could only dream
of, only the rich and powerful could visit it.
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Brain&Force
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Speaking of bfesser, can a mod change this thread's title to something like "General discussion of terbium compounds?" This is a whole lot more than
just the synthesis of terbium iodide.
Eddygp, if you are going to work with ytterbium, tell us what color the pure metal is. There was an edit war on Wikipedia on whether ytterbium is silvery or slightly golden, and it's still not clear as to what color it is. As one of the users
states, it's like the debate over the color of cesium - almost all references before 1984 list it as silver.
And don't forget to try a pyro mixture with it - ytterbium burns very bright green, purer than copper (bluish-green) or barium (slightly
yellowish-green). If you have equipment for it you can also see the strong IR emission of burning Yb metal. Militaries are investigating its use in
decoy flares because it can emit more IR than a Mg based flare (and the formed Yb2O3 is more emissive in the IR spectrum than
MgO).
Ytterbium has a lot of really interesting applications.
At the end of the day, simulating atoms doesn't beat working with the real things...
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siegfried
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All RE elements can be ordered from Metallium, Inc A 20 gm lump of Tb is $135 and a 5 gm piece is $32. Most oxides can be ordered at Smart Elements
is Austria. Some RE compounds can be ordered on Ebay.
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Eddygp
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Quote: Originally posted by Brain&Force  | Speaking of bfesser, can a mod change this thread's title to something like "General discussion of terbium compounds?" This is a whole lot more than
just the synthesis of terbium iodide.
Eddygp, if you are going to work with ytterbium, tell us what color the pure metal is. There was an edit war on Wikipedia on whether ytterbium is silvery or slightly golden, and it's still not clear as to what color it is. As one of the users
states, it's like the debate over the color of cesium - almost all references before 1984 list it as silver.
And don't forget to try a pyro mixture with it - ytterbium burns very bright green, purer than copper (bluish-green) or barium (slightly
yellowish-green). If you have equipment for it you can also see the strong IR emission of burning Yb metal. Militaries are investigating its use in
decoy flares because it can emit more IR than a Mg based flare (and the formed Yb2O3 is more emissive in the IR spectrum than
MgO).
Ytterbium has a lot of really interesting applications. |
Ytterbium is silvery with a tinge of cream. Definitely not golden or brass.
[Edited on 30-12-2013 by Eddygp]
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Brain&Force
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Comparing thulium to ytterbium shows ytterbium is definitely yellowish. My friend calls it a champagne color. It reminds me of the color of that new
iPhone that came out.
Back on topic...
Exposing the yellow terbium compound to air has led to no change - it's still as yellow as it was before. Perhaps I need to reduce the
ICl4- to prove its existence.
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blogfast25
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| Quote: Originally posted by Brain&Force |
Exposing the yellow terbium compound to air has led to no change - it's still as yellow as it was before. Perhaps I need to reduce the
ICl4- to prove its existence. |
Try treating a bit of it with strong sulphite solution?
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Eddygp
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A bit on topic, are any lanthanides amphoteric?
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blogfast25
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No, not remotely.
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Brain&Force
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I have a better idea regarding removal of ICl4- - just precipitate terbium as Tb(OH)3 and any
ICl4- will be removed. The remaining Tb(OH)3 will then be reused to make pure TbCl3.
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 | | I have a better idea regarding removal of ICl4- - just precipitate terbium as Tb(OH)3 and any
ICl4- will be removed. The remaining Tb(OH)3 will then be reused to make pure TbCl3.
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Sure but I was only referring to testing the ICl<sub>4</sub><sup>-</sup> hypothesis, not recovering the Tb(III).
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Brain&Force
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So break ended. I began to search for sulfite or thiosulfate, but nothing turned up. But I did find a vial of an unusual chemical - ammonium
heptamolybdate. I had previously looked upon it as a curiosity, but then I realized I could prepare terbium molybdate, a compound that is poorly
characterized.
I attempted to prepare Tb-molybdate through metathesis. Some solid ammonium molybdate was added to a solution of terbium chloride (a different
solution that did not have any noticeable ICl4- contamination) and added the ammonium molybdate. The solution became milky
white, but quickly became perfectly clear. I'm wondering if I made a complex between Tb and either molybdate or ammonia.
Also, I'm trying to make Tb-citrate to hopefully make sodium or potassium tris(citrato)terbate (is that even correct? I'm dead sure I'm wrong here).
More to come soon.
At the end of the day, simulating atoms doesn't beat working with the real things...
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Brain&Force
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blogfast25, you were right to begin with.
There's iron in the terbium - what seems to be a large amount. I knew that could be the only cause because the terbium citrate I had made was as
yellow as a school bus. I tried a thiocyanate test, this time with a different source (our school has some very old, degraded chemicals, and I'm sure
I got an old batch), and the solution turned reddish-black. I did this test with all terbium sources and got the exact same result. How wonderful.
I'm attempting a very roundabout way to recover the remaining terbium in solution. Manganese pieces are reacted with the solution, replacing Fe, which
precipitates as either Fe metal or Fe2O3. In solution are Tb3+ and Mn2+. Mn is removed by oxidation by
acidified H2O2, and Tb is precipitated as hydroxide.
I'll try to get an accurate measurement on the amount of iron in solution soon.
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 | I did this test with all terbium sources and got the exact same result. How wonderful.
I'm attempting a very roundabout way to recover the remaining terbium in solution. |
No, no. Use the sulphate method! RE sulphates are very poorly soluble but Fe (III) sulphate is very soluble. So add quite a bit of sulphuric acid to
your contaminated Tb solution, make sure pH < 3. Tb (and other REs) sulphate precipitates, ferric sulphate remains in solution. Then filter and
rinse filter cake with cold H2SO4 (say about 25 %).
Works: been there, done that, have T-shirt...
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MrHomeScientist
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Well, in the case of neodymium it's been a bit more complex than that. Regarding our work in processing neodymium magnets, the Nd-sulfate remains
quite soluble along with Fe(III) sulfate. The way I've been separating is by heating to boiling, whereupon some Nd-sulfate precipitates out while the
iron remains very soluble. Filter the pink crystals and wash with boiling water. After cooling down to room temp, usually lots of iron(III) has
crystallized. This is removed, and the process repeated several times. I know you did a lot of work in separating the two, certainly beyond just
adding a lot of sulfuric acid. So I don't think the general statement of "RE sulphates are very poorly soluble" is accurate. For this particular case
with terbium, it may well be.
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blogfast25
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| Quote: Originally posted by MrHomeScientist | | So I don't think the general statement of "RE sulphates are very poorly soluble" is accurate. For this particular case with terbium, it may well be.
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Mr HS, you're right, I'm not being very accurate (haste makes waste). The low solubility of Re sulphates is at 100 C. For Ce(III), Tb, Ce, La (and
others) Wiki's solubility table shows that clearly.
I actually prefer to use the double sulphates with potassium, which are almost insoluble, also at low temperature.
Here a strongly acidic (to prevent Fe(OH)3 precipitating) RE solution is saturated with K2SO4. The double sulphates precipitate, the iron remains in
solution. Wash filter cake with acidic K2SO4 solution. This is one industrial route.
Potassium is released by treatment of filter cake with NH3 solution, which then yields RE(OH)3.
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Brain&Force
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At this point, it's too late to continue experimentation - my project is supposed to be complete by the 24th. I already have a lot of new information
(although some of it is incorrect now that I know that the Tb was contaminated).
However, I am attempting to determine the amount of iron in the terbium using the terbium potassium alum method. The reason I had decided to use my
convoluted method was because I already had added KSCN to some of the mixtures, and from what you had stated earlier, blogfast25, I thought terbium
potassium alum was highly insoluble and could not be precipitated as hydroxide with NH3. These recent posts say otherwise.
I read somewhere that NH3 should be used to precipitate Ln3+ and not NaOH/KOH. Am I correct? And if so why?
I may have stumbled upon another method of removing iron from terbium. More on that later.
Do you have a link to this table? I could use it for my project.
[edit] Correction: I will continue experimentation, it just won't be included in the writeup and other materials for the project. I still have a bit
of Tb left.
[Edited on 12-1-2014 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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| Quote: Originally posted by Brain&Force | I read somewhere that NH3 should be used to precipitate Ln3+ and not NaOH/KOH. Am I correct? And if so why?
I may have stumbled upon another method of removing iron from terbium. More on that later.
Do you have a link to this table? I could use it for my project.
|
The REs don't from complexes with NH3, so I can't see an immediate impediment to using it. Plenty references cite its use with REs. Ammonium is easier
to wash out of precipitates, or so I've read. It also leaves more easily on calcining, if the actual oxide is what one wants.
The double sulphates convert to hydroxides on treatment with NH3 easily because the Ksp of the hydroxides is much lower than that of the double
sulphates.
Solubility table:
http://en.wikipedia.org/wiki/Solubility_table
Glad to hear you'll carry on experimenting with your terbium/iron!
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Brain&Force
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The solubility table only lists the solubility of terbium at 20 C. By solubility table I thought you had a graphic, I had used that page before.
I just realized that the manganese extraction route has a flaw. The solution currently contains these ions: H+, Mn(II), Tb(III),
K+, likely a small amount of Fe(III), complemented with chloride, thiocyanate, and a bit of triiodide. By adding
H2O2, I may generate chlorine gas (through the formation of MnO2), oxidize triiodide to iodine or iodate, or have a
whole host of other side reactions take place. However, it doesn't seem too late to just precipitate the terbium as terbium potassium alum, then
produce terbium hydroxide with ammonia.
I figured out what was causing the smell of oranges. Several of the vials at my school smell like assorted fruits, and appear to have been used for
some sort of demo. It has nothing to do with the Tb salt.
And about the new extraction method: From reacting terbium (actually ferroterbium) with citric acid, I discovered that a yellow precipitate formed. I
can't find much information on ferric citrate (regarding solubility and color), but that's what i think the precipitate is. Terbium citrate seems to
remain in solution. I'll continue looking into this. I'm surprised that terbium citrate did not precipitate.
And this paper indicates that terbium metal can be precipitated out of solution in the presence of cobalt and boric acid. If I can find a suitable
power source I'll try to reform terbium metal. This is very unusual for an element of such electropositivity.
At the end of the day, simulating atoms doesn't beat working with the real things...
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Brain&Force
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Wait, what? Strange result
I've gotten an odd result out of the separation attempts. After allowing terbium to dissolve in HCl, a yellowish solution formed (of course, the Fe
contamination). I had used too much terbium and the HCl I have isn't concentrated enough, so I decanted some of the solution into another test tube.
However, after leaving the solutions to sit for a few hours, ferric hydroxide precipitated out - but not in the form I expected. There was powder at
the bottom, but fluffy clumps of red powder were also suspended in the solution.
I think the terbium ion acted as a flocculating agent similarly to aluminum sulfate, and that caused the ferric hydroxide to precipitate rather than
remain suspended. I'll re-acidify the solution and continue separation. At what concentration does Fe hydroxide become visible in suspension? That may
explain the yellow stuff and why it clarifies so clearly when only a few drops of HCl are added. There may not be as much Fe contamination as I
thought.
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 | I've gotten an odd result out of the separation attempts. After allowing terbium to dissolve in HCl, a yellowish solution formed (of course, the Fe
contamination). I had used too much terbium and the HCl I have isn't concentrated enough, so I decanted some of the solution into another test tube.
However, after leaving the solutions to sit for a few hours, ferric hydroxide precipitated out - but not in the form I expected. There was powder at
the bottom, but fluffy clumps of red powder were also suspended in the solution.
|
Most likely explanation: your acid became depleted and wasn't capable of keeping the ferric chloride in solution. The latter would need a pH of less
than 4, less than 3 even if there's quite a lot of Fe<sup>3+</sup>.
The K<sub>sp</sub> of Fe(OH)<sub>3</sub> is very, very low, which causes ferric ions to precipitate from about said pH values.
Ferric hydroxide has a tendency to peptise: go colloidal when ionic strength of the supernatant solution is low. I've seen such precipitates run right
through a filter as if they were proper solutes!
As I wrote above, if you use the double sulphates method for separation of iron and terbium, make sure the solution and any filter cake washing
solution is pH < 3, to avoid the iron precipitating...
[Edited on 14-1-2014 by blogfast25]
| Quote: Originally posted by Brain&Force |
And this paper indicates that terbium metal can be precipitated out of solution in the presence of cobalt and boric acid. If I can find a suitable
power source I'll try to reform terbium metal. This is very unusual for an element of such electropositivity. |
It talks about Co/Tb alloys mainly, if I'm not mistaken.
[Edited on 14-1-2014 by blogfast25]
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Brain&Force
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No precipitate?
I reacidified the solution, the precipitate dissolved, and the yellow color disappeared, as expected. Then I added in K2SO4.
That's when things started to get weird. Nothing precipitated out of solution. I thought I didn't add in enough potassium sulfate, so I dropped in
several crystals, and allowed them to dissolve. They dissolved slowly (is it just me, or are all sulfates slow to dissolve?), but nothing precipitated
out. I heated the solution to boiling and only a few crystals dropped out of solution. Considering I only added half a gram of terbium to the
solution, and about a whole gram of potassium sulfate, I'm surprised only that tiny amount precipitated.
Perhaps the solution is not concentrated enough? It's probably only .3 molar terbium ion in solution. I couldn't find any 6M HCl. I'm afraid to boil
the solution in case hydrolysis occurs.
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 | Considering I only added half a gram of terbium to the solution, and about a whole gram of potassium sulfate, I'm surprised only that tiny amount
precipitated.
Perhaps the solution is not concentrated enough? It's probably only .3 molar terbium ion in solution. I couldn't find any 6M HCl. I'm afraid to boil
the solution in case hydrolysis occurs. |
0.3 M is quite low. What was the volume of solution? Hydrolysis won't occur if there's enough acid present.
You really need more concentrated HCl to carry on. Try to get 'Patio cleaner'? That's 15 - 20 %, I think.
But while searching for the original reference for the separation of iron from REs using double sulphates, I came across this reference:
http://www.kth.se/en/che/divisions/transport_phenomena/resea...
“One of the most widely used precipitation method to separate rare earth elements (REE) from acidic solutions is by precipitation of sodium
double sulphate hydrates (NaRE(SO4)2.xH2O) through the addition of sodium sulphate (Gupta 1992). These salts are only slightly soluble in acidic
solutions. Double sulphate precipitation result in separations in two fractions one enriched in the light-REE group rare earths and the other enriched
in the heavy-REE group rare earths (Gupta, 1992). The different REE can then be separated from each other by converting the double sulphate into a
highly soluble compound, such as RE-hydroxide. Although monazite leach liquor is usually purified by precipitation of REE as sodium double sulphate
hydrates the influence of the process variables on the double sulphate precipitation and on its conversion into RE-hydroxide is rarely mentioned in
literature (Abreu and Morais, 2010).”
… “in two fractions one enriched in the light-REE group rare earths and the other enriched in the heavy-REE group rare earths”, this
maybe the snake in the grass here: perhaps Tb doesn’t precipitate as eagerly or fully as e.g. Nd (for which I used the method)?
One side point: my original reference advised K2SO4 because that double sulphate is even less soluble, so the use of K2SO4 shouldn’t be the cause of
the problem.
Oxalates: I described elsewhere on this forum practical experimentation using oxalates to separate Fe and RE (all RE oxalates are exceedingly
insoluble, with very low Ksp). It involves adding an excess oxalic acid, than potassium hydroxide or potassium oxalate. The RE drops out as RE2Ox3,
the iron remains soluble as K3FeOx3 (potassium trisoxalato ferrate (III), a grass green, slightly fluorescent complex). Drawback: the RE oxalate is
harder to convert back into something soluble. I'll try and find that work now.
Here's the oxalate method at work (by MrHS) on neodymium/iron from some large magnets:
http://www.sciencemadness.org/talk/viewthread.php?tid=14145&...
[Edited on 15-1-2014 by blogfast25]
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Brain&Force
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I finally got my hands on 6M HCl (iron free and perfectly colorless) and added about 10 mL to the solution. The solution turned greenish-yellow,
almost the color of highlighter ink. Evaporating the solution to 50 mL (was about 75 mL initially) resulted in no precipitation or hydrolysis.
I can't get my hands on oxalates (someone is taking all the chemicals out of our classroom), and they're not a good idea with terbium. Tb oxalate
converts to a higher oxide when it is heated, and the product cannot be dissolved easily - it takes a week and generates large amounts of chlorine
gas.
Right now I think the best course of action is to boil the solution dry and redissolve in 6 molar HCl. That should keep the Fe(III) salts in solution
and leave the Tb sulfate or terbium potassium alum out.
Here is the acid after adding extra HCl. Note that the color just suddenly appeared after a few seconds after adding the acid.
[edit] Terbium sulfate is only slightly soluble in water: http://www.youtube.com/watch?v=DJcObFauFIc
[Edited on 16-1-2014 by Brain&Force]
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blogfast25
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3.56 g / 100 g at 20 C, acc. Wiki. This value should further decrease at higher temperature (going by the other REs), so boiling the solution with
H2SO4 should precipitate most of the terbium sulphate. If pH is low enough iron should remain in solution.
If you dissolve a fresh batch of your terbium in the 6 M HCl, the iron will be present as Fe<sup>2+</sup>, which is much less prone to
precipitating as hydroxide than Fe<sup>3+</sup>.
Oxalic acid is very, very OTC. The higher Tb oxide could probably be 'cracked' by fusion with NaHSO<sub>4</sub>.
[Edited on 16-1-2014 by blogfast25]
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