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blogfast25
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The chloride is hygroscopic, remember? Besides, high hydrates are rarely used in gravimetry because they're not very stable.
Oxalate based gravimetrical determinations of REs are mainstream.
[Edited on 21-3-2014 by blogfast25]
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Brain&Force
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Oh wow...it's been so long since I've been experimenting.
Say, what is terbium again? 
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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Here's the Nd oxalate / potassium trisoxalatoferrate (III) separation method as applied by MrHomeScientist on 'magnet soup':
http://www.sciencemadness.org/talk/viewthread.php?tid=14145&...
Needless to say the amount of Fe in your terbium is much smaller than in a Nd magnet, so it's even easier to convert that bit of Fe to highly soluble
K<sub>3</sub>FeOx<sub>3</sub>.
[Edited on 22-3-2014 by blogfast25]
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Brain&Force
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This is really, really interesting.
http://nopr.niscair.res.in/bitstream/123456789/7190/1/IJCA%2...
First of all, the compound fluoresces in water. It states that planar ligands have a strong antenna effect. Maybe this is the breakthrough I need?
Second of all, its fluorescence increases in low pH enviroments.
Third of all, it has potential applications in molecular logic gates.
And it gives me the quantum transitions characteristic of each emission line, which is something I've been looking for for a while.
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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I added some solid sodium oxalate to the solution, but didn't have time check on it after I added it, so it's just sitting there undissolved in the
beaker now. I'll report back tomorrow.
I've also discovered that the sulfate may still be soluble enough to cause significant losses - it appears to have passed through the filter paper and
deposited on the Buchner funnel.
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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I made a bit of a mistake when attempting to seperate the terbium using the oxalate method. I accidentally added the sulfate instead. Nothing appeared
to have precipitated, so I added oxalate, which caused the yellow color to disappear from solution and a white powder to precipitate. I may need to
add more oxalate because the tris(oxalato)ferrate complex doesn't seem to have formed. It takes a lot of sodium oxalate just to get anything to
precipitate!
Will the sulfate affect anything? I doubt it, the Ksp for terbium oxalate is much lower than that of the double sulfate, from what I understand.
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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Big update
I haven't been here for a long time...anyway, I have several updates to make. I likely won't be continuing any of this research over the summer, but
I'll see what I can do.
First of all, I determined that all of my HCl sources had some iron contamination in them. The problem is that iron(II) will not test positive in a
thiocyanate test, and hydrogen peroxide must be used to oxidize it to iron(III).
Second, it appears that using a large nugget of terbium reduced the amount of contaminants in the solution after dissolving it in HCl. So I think it's
safe to say that parts of the chisel may have broken off and contaminated the solution. Extremely tiny amounts of iron(III) appear to completely kill
the fluorescence of terbium, and IrC and I were discussing the possibility of using this as a sensitive test for iron.
(Quick note - the solutions were never anywhere near as dark as the commercial ferric chloride etching solutions. They were just pale yellow. So the
concentration of Fe3+ must have been pretty low.)
Third, I got terbium and iodine to react by adding a drop of water to the terbium in an iodine atmosphere, as shown in several YouTube videos
involving sodium in a chlorine atmosphere. It wasn't very violent or even noticeable, but a dark triiodide complex did form, implying that a reduction
occured. The reaction seems to only work at elevated temperatures.
I have some leftover terbium metal, but I also have some leftover terbium sulfate and oxalate. I know how to convert terbium sulfate to the carbonate,
but what can I do about the oxalate - in other words, how can I convert it into a soluble form? If I heat it, it'll ignite to form the higher
terbium(III,IV) oxide, which is not as handy as the oxide/hydroxide/carbonate.
[edit] Maybe I should do this to the terbium?
<iframe sandbox width="420" height="315" src="//www.youtube.com/embed/Noftcq8g7p8" frameborder="0" allowfullscreen></iframe>
I don't know if this has been faked, nor do I know if such an explosion will occur with pure terbium metal. This is the best video I could find, even
though it's rotated, and there are several others - search "exploding terfenol-D."
[Edited on 4.6.2014 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:
Ferric chloride etching solutions are very concentrated: several M. Almost anything is weak compared to that.
Oxalate: I see no other option, due to the insane insolubility of the Ln oxalates, to calcine to the oxide.
[Edited on 5-6-2014 by blogfast25]
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Brain&Force
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Well, research for me is pretty much done.
I still have 2.1 grams of terbium metal, as well as some leftover terbium compounds (terbium oxalate and potassium terbium sulfate). Filtering out the
compounds is a very lossy process, as I still have some leftover terbium compounds on the filter paper, as well as in the filtration flask. I wouldn't
have known this if I hadn't added sodium oxalate to one of the solutions. So I'll have to go for another round of filtration.
<a href="http://imgur.com/gTCd9HC"><img src="http://i.imgur.com/gTCd9HC.jpg" title="That's the brightest Buchner funnel I've ever seen."
width=800 /></a>
Here are the collected terbium compounds. As you can see, there are some dead crystals, likely due to leftover iron. These crystals are very slightly
greenish under tube lighting, the powders are the truest white I've ever seen. There's not a hint of tint to them. Quick note: they won't fluoresce at
all in longwave UV - you need shortwave. Any good ideas for a shortwave lamp?
<a href="http://imgur.com/VNvW08N"><img src="http://i.imgur.com/VNvW08N.jpg" title="SHINY" width=800/></a>
At the end of the day, simulating atoms doesn't beat working with the real things...
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Polverone
Now celebrating 21 years of madness
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I have a nice little handheld UV lamp that does shortwave and longwave. A simple sliding shutter system lets you do long, short, or both wavelengths
at once. It was a garage sale find -- made in the 1960s -- so I don't know what the current crop of devices is like. You need something small and
convenient, but not necessarily battery powered, because you aren't doing field work. I wouldn't recommend trying to repurpose germicidal lamps or
anything like that because they are a) excessively hazardous for your needs and b) emit too much visible light for easy observation of fluorescence;
the rock hound lamps are purpose-built for making fluorescence easy to see.
PGP Key and corresponding e-mail address
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Brain&Force
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The lamp I have is a children's invisible ink toy. It does the job for longwave UV. I also have a blue LED diode from a different toy. I've used both
in experiment with GFP. As far as I know, europium will fluoresce with blue and longwave UV light - providing another simple way to tell which species
are present in a mixture.
I don't experiment in my house, but I am curious, do you know what make/model your lamp is? I was considering using a geological UV lamp, but I can't
access one, and the germicidal ones are FAR too powerful, as you said.
At the end of the day, simulating atoms doesn't beat working with the real things...
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Polverone
Now celebrating 21 years of madness
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I have a Raytech Industries LS-4. Wow, it's still in production: http://www.raytechultraviolet.com/product-model4.php
But mine is from 1965 and I paid maybe $20 for it, not $235.
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