Sciencemadness Discussion Board - General Discussion of Terbium Compounds

General Discussion of Terbium Compounds

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Brain&Force - 17-1-2014 at 16:52

I checked on my beaker this morning and a mass precipitation occured! I think the solution was supersaturated and needed something to set it off. Some of the precipitate was collected, dried, and stored in a vial. A test with a UV source shows that it fluoresces! The sample exhibits a mint green color, likely due to fluorescence. As expected, the color is enhanced in sunlight. Some of the crystals were cubic, so likely KCl precipitated out as well.

A test was done, which involved adding ammonia to a few crystals of the potassium terbium sulfate. The crystals turned to a white powder. On acidification with HCl the powder dissolved (and a cloud of ammonium chloride was formed). This shows that terbium hydroxide precipitated in the ammonia.

I've only got one problem now: the crystals I collected weren't completely dry, so there's a bit of HCl in them. What is the best way to remove it?

I took a picture of the solution under a black light (with the precipitate) and it appears to glow green. I don't know whether it is due to the solution or the precipitate. I'll try to figure that out soon.

I added some terbium to 6 M HCl. This was apocalyptic compared to 1 M acid - the terbium is completely destroyed within 5 seconds of addition. Adding several pieces created a clear solution which became yellow after a day. I'll try make some video of the reaction (but I don't know if I can post it).

[Edited on 18-1-2014 by Brain&Force]

blogfast25 - 18-1-2014 at 05:32

Quote: Originally posted by Brain&Force

I've only got one problem now: the crystals I collected weren't completely dry, so there's a bit of HCl in them. What is the best way to remove it?

Good news! So the Tb/K double sulphate too is poorly soluble.

It should be sufficient to wash them with several small aliquots of cold, acidified (remember the Fe(III)!) sat. K2SO4 solution to get rid of the iron and other soluble contaminants.

But if you want to make Tb2(SO4

3 do as follows. Suspend the mass in 3 - 5 times the amount of water and add strong ammonia. Simmer gently for 20 minutes or so to fully convert the Tb double sulphate to Tb hydroxide + K2SO4, filter and wash filter cake with plenty hot water, then some cold deionised water.

Again, suspend the Tb hydroxide in a bit of water and add a good dollop of H2SO4 and bring to the boil, simmer for a while to ensure full conversion to the Tb2(SO4

3. Filter it off hot and wash with small aliquots of hot 20 % H2SO4, then with a bit of boiling deionised water. Let the crystals dry on a low setting electrical hot plate or equivalent. These should UV fluoresce in the dark like mad. They may also look slightly different under incandescent light (low UV) and saver bulb light (quite a bit of UV), see also Neodymium.

[Edited on 18-1-2014 by blogfast25]

Brain&Force - 18-1-2014 at 11:39

Quote: Originally posted by blogfast25

These should UV fluoresce in the dark like mad. They may also look slightly different under incandescent light (low UV) and saver bulb light (quite a bit of UV), see also Neodymium.

[Edited on 18-1-2014 by blogfast25]

Ok then, something isn't entirely right.
I believe a true alum may not have formed. The sample is only very weakly fluorescent. However, I did get a similar result with the Fe-contaminated TbCl₃ (almost zero fluorescence). It appears the presence of other metal cations (and possibly even water) interferes with the fluorescence of terbium (likely due to UV absorption). I don't have incandescent lights in my house, but CFL lights and sunlight show the same mint green.

I left the vial open overnight and the compound dried up. It doesn't seem to be very hygroscopic.

blogfast25 - 18-1-2014 at 13:07

@B&F:

I didn't claim the double salt would fluoresce (I don't know that). I would expect the pure double sulphate to do so but maybe weaker than the pure sulphate. It's possible 'foreign' cations influence the degree of fluorescence but I can't see by what mechanism.

These double salts aren't really alums. Their formula and structure of K2SO4.RE2(SO4)3.nH2O (n appears to be 2 or 3) is quite different from MK(SO4)2.12H2O (M is a trivalent metal), for actual alums. They used to be written as K2SO4.M2(SO4)3.24H2O (for potassium alums).

Brain&Force - 20-1-2014 at 10:36

From my experience and MrHomeScientist's thread on terbium tetrakis(dibenzoylmethide)triethylammonium, it appears that terbium is a difficult to get to fluoresce properly without the proper organic chromophore ligand. These sites have a lot of information about lanthanide fluorescence.

I also need to figure out the spectrum of the UV lamp I used. The sites I looked at show that the lanthanides fluoresce at wavelenghts from 340 nm 254 nm. I'll try making terbium chloride and hexakis(antipyrine)terbium chloride to see if there's a significant difference in the fluorescence, especially from different sources. This may explain why the double sulfate fluoresces under only the UV tube lamp. (I used 3 sources for my experiments: a UV LED, a tube lamp, and a UV compact fluorescent lamp.)

In the compounds europium (tetrakis)dibenzoylmethide triethylammonium and hexakis(antipyrine)terbium iodide, the dibenzoylmethane and antipyrine ligands appear to be the chromophores that the lanthanide luminescence site discusses.

Eddygp - 20-1-2014 at 11:27

Lanthanides are so interesting. Actually, their chemistry differs more than I had expected. It only turns out to be too similar when you try to separate them. :mad:

Hahahah anyway, I'll see whether I can check a weird mineral I found (that I presume is pure maghemite) in case I find something interesting. Lanthanides in geological strata are not that common, unfortunately.

Brain&Force - 20-1-2014 at 12:55

According to Wikipedia's solubility table, lanthanum iodate is insoluble. So forming terbium iodide through reaction with iodine water just might be possible as I had thought previously! Also, the corresponding iron salts are also insoluble, so getting pure terbium iodide just might be possible (and it could be a good extraction method for ferrolanthanide alloys).

Eddygp - The europium(II) and scandium cations tend to blend in with calcium-bearing minerals (notably fluorite, the Eu(II) causes blue fluorescence), which reduces the amount of those metals in lanthanide-bearing ores; see europium anomaly. Similarly, cerium(IV) is often found in zirconium-containing minerals like zircon.

Woo hoo!

Brain&Force - 22-1-2014 at 15:54

I decided to check the mother liquor for fluorescence.

Now that's what I'm talking about! But now I'm reluctant to destroy the crystals for Tb(OH)₃...

The failure of my other batch of double sulfate to fluoresce was due to the UV compact fluorescent not actually emitting any UV radiation (test was done with my highly fluorescent T-shirt).

On a side note, my terbium crystals match my T-shirt now.

I made some terbium hydroxide by adding the first batch of double sulfate to ammonia. A white, sandy precipitate formed at the bottom of the beaker. It does not fluoresce with the good lamp. I'll try to convert it to TbCl₃ and then hexakis(antipyrine)terbium chloride to see if there's increased fluorescence, as antipyrine is expected to act as a chromophore. Hopefully I can make a fluorescent solution soon.

Here's the terbium hydroxide precipitate:

[edit] Is it necessary to separate the glowing flakes from the dead ones?

[Edited on 23-1-2014 by Brain&Force]

blogfast25 - 23-1-2014 at 05:42

The green fluorescing stuff is the double sulphate?

Quote: Originally posted by Brain&Force

Here's the terbium hydroxide precipitate:

[edit] Is it necessary to separate the glowing flakes from the dead ones?

[Edited on 23-1-2014 by Brain&Force]

Not sure what you mean. I see some finer and some coarser material but nothing that glows?

[Edited on 23-1-2014 by blogfast25]

Brain&Force - 23-1-2014 at 16:28

Sorry about that edit, I was talking about whether or not I should separate the dead crystals from the fluorescing double sulfate.

I've come across a very strange and unusual result. When the crystals are dry, they fluoresce. But submerged under the original solution (an acidic solution consisting of mostly iron and potassium salts with chloride and sulfate), the crystals do not exhibit any fluorescence at all. This is very strange. The original solution is yellow with a tinge of green, so I don't know if it's just some absorption effect or fluorescence quenching. I made a video demonstrating this effect, but I can't upload it. A Google search brought up nothing relating to what I've observed - I'm surprised no one seems to have noticed this before.

I was unable to present this as a science fair project. I've been too caught up attempting to synthesize and purify the chemicals that I was unable to group all of it under a single question and put it in the "proper" problem, hypothesis, procedure, trials, conclusion format, which is absolutely necessary for this fair. On the other hand, the only reason I said I'd enter the fair is because I wanted to do research, and I can't build a home lab at this time. I'm doing this for science, not for the fair.

I could try to publish this here; I would just need a little bit more data.

Brain&Force - 26-1-2014 at 13:13

I've made some big discoveries:

The first one is that the water itself does not stop the fluorescence of the potassium terbium sulfate. The ferric ions do - Fe³⁺ really kills the fluorescence of Tb³⁺. It's probably another good (though expensive!) way of determining whether Fe³⁺ is in solution. I wonder if this works with other transition metal ions.
The second one is that terbium hydroxide is not fluorescent at all. This is really useful when converting the double sulfate to the hydroxide - just wait until the powder stops fluorescing.
Unfortunately, we suffered a spill and lost about half of our product. Thankfully, the sulfate and hydroxide are both magnetic, so collecting them as powders is no challenge.

Here are some photos from the lab:

blogfast25 - 27-1-2014 at 06:15

That IS interesting, B&F.

You are positively CERTAIN that ferric ions suppress the UV fluorescence of Tb sulphate?

[Edited on 27-1-2014 by blogfast25]

Brain&Force - 27-1-2014 at 15:46

The fluorescence stops when the terbium crystals enter the acidified solution containing ferric ions, potassium ions, chloride, and a bit of sulfate. I don't know exactly what component is suppressing the fluorescence, but I'll try soaking some of the crystals in ferric sulfate solution to see what happens. It may be due to the tetrachloroferrate (FeCl₄^-) ion, so I'll also try ferric sulfate acidified with HCl.

Happy 4444th post blogfast25!

[edit] I overlooked the fact that the Fe contaminated Tb-chloride did not fluoresce. So it must be either the ferric or tetrachloroferrate ion!

[Edited on 28-1-2014 by Brain&Force]

Brain&Force - 30-1-2014 at 19:54

Ferric ions most definitely kill the Tb fluorescence. I made another batch of double sulfate and when it came out, they were yellowish. Rinsing with water removed the Fe and the powder immediately became fluorescent. I wonder if this is true with other transition metals as well.

Also, I looked at the emission spectrum of Tb fluorescence, and, as expected, it matches that of tube and CFL lighting. I'll try to get wavelength measurements, and hopefully images with the available diffraction grating.

blogfast25 - 31-1-2014 at 06:25

B&F:

Good to see confirmation of the ferric ion influence.

Brain&Force - 3-2-2014 at 20:21

I finally was able to make some terbium hydroxide. Oh, wait, let me correct myself: terbium carbonate. I made terbium hydroxide by adding the terbium sulfate to ammonia, but it ended up absorbing carbon dioxide to form the carbonate. This was confirmed when the sample was added to hydrochloric acid - it fizzed somewhat.

Contrary to my above posts, terbium hydroxide is fluorescent. It's just a lot weaker. The sulfate is far brighter. I suspect large ligands enhance the fluorescence of Tb - from photos I've seen, the sulfate and nitrate are brighter than the hydroxide and chloride.

Cool link: http://perso.bretagne.ens-cachan.fr/~mwerts/lanthanides/ln_d...

blogfast25 - 4-2-2014 at 11:27

Quote: Originally posted by Brain&Force

I finally was able to make some terbium hydroxide. Oh, wait, let me correct myself: terbium carbonate. I made terbium hydroxide by adding the terbium sulfate to ammonia, but it ended up absorbing carbon dioxide to form the carbonate. This was confirmed when the sample was added to hydrochloric acid - it fizzed somewhat.

Contrary to my above posts, terbium hydroxide is fluorescent. It's just a lot weaker. The sulfate is far brighter. I suspect large ligands enhance the fluorescence of Tb - from photos I've seen, the sulfate and nitrate are brighter than the hydroxide and chloride.

Cool link: http://perso.bretagne.ens-cachan.fr/~mwerts/lanthanides/ln_d...

Yes, several insoluble hydroxides absorb CO2 readily from the air, Zr(OH)4 does it too.

But I doubt if it went 'all the way', in your case...

Interesting how the hydroxide/carbonate also fluoresces.

[Edited on 4-2-2014 by blogfast25]

Brain&Force - 11-2-2014 at 18:42

I was able to dry the terbium chloride, but somehow, more iron got into solution! I'm going to redissolve the crystals in HCl and reprecipitate it as terbium potassium sulfate.
The iron impurity was concentrated at the edges of the beaker in which I dried the crystals. The edges were filled with terbium chloride powder, and the center was filled with glassy crystals. The center crystals did not fluoresce, and the edges only barely.
I'm surprised zirconium hydroxide absorbs carbon dioxide. It doesn't seem the hydroxide would be basic enough to do that!

[Edited on 12-2-2014 by Brain&Force]

blogfast25 - 12-2-2014 at 09:17

Quote: Originally posted by Brain&Force

I'm surprised zirconium hydroxide absorbs carbon dioxide. It doesn't seem the hydroxide would be basic enough to do that!

[Edited on 12-2-2014 by Brain&Force]

Venable's famous monography on Zr and its compounds describes it and I saw it first hand when I extracted zirconyl chloride from Zircon powder.

I'm baaaaaaaaaack!

Brain&Force - 18-3-2014 at 16:47

I thought I figured it out, but no.

I contacted Metallium about the issue with purity, and the owner reported that the 99.5% expected purity had been confirmed by two analytical labs. I don't dispute this result, but I'm trying to determine the purity myself, just to see how contaminated the sample is. At the same time I tried figuring out if any other factors may have contaminated the terbium.

It may be the chisel.

Looking back to one of my previous posts, I mentioned that the terbium was extremely hard to break and only a chisel worked. I had been using the finest pieces, and bits of chisel may have been mixed in. So I decided that a larger piece would result in a pure terbium(III) solution being produced.

But guess what? That didn't happen! There was still a significant amount of iron in the sample, enough to produce a yellow solution. Back to square one.

Also, I happened upon some more interesting results: iron(II) contamination doesn't mask the fluorescence of terbium(III) salts! It must be inherent to the iron(III) salt's absorption bands.

For future reference, 0.4723 grams of terbium metal are being used for the analysis (measured with a high-precision analytical balance that we were very lucky to happen upon).

blogfast25 - 19-3-2014 at 05:57

B&F:

If iron is still the problem, the thiocyanate method can be turned into a quantative method w/o much fuss and w/o instrumentation, if you like. It boils down to making a few standard solutions of Fe3+ of known Fe ppm, with an excess KSCN in them and comparing their colour to samples obtained with the terbium. You will of course need strictly Fe free reagents.

I have difficulty believing the chisel is the cause, although the reasoning is correct of course.

[Edited on 19-3-2014 by blogfast25]

Brain&Force - 20-3-2014 at 18:25

blogfast25:

I'm just going to calculate the moles of terbium hydroxide produced from the sample through the sulfate method. I would have asked about colorimetry, but the balance find was very lucky, and I'll put it to good use.

blogfast25 - 21-3-2014 at 05:29

Gravimetry to determine small amounts of impurities? Good luck with that! Terbium hydroxide isn't thermally stable and the double sulphates have some limited solubility.

Separating the Fe and Tb as oxalates would be much better: Fe(III) forms a highly water soluble trisoxalatoferrate (III) complex, but RE oxalates are truly insoluble. There's a thread on it somewhere here by MrHS.

[Edited on 21-3-2014 by blogfast25]

Brain&Force - 21-3-2014 at 10:01

I forgot about the hydroxide's tendency to absorb all sorts of acidic vapors. I might just reconvert it to the chloride and either mass the chloride or precipitate it as the oxalate.

I have reason to believe there's a pretty significant Fe impurity (ca. 20%), enough that gravimetry can determine the % impurities. The impurities have a very positive test for Fe.

IrC - 21-3-2014 at 10:52

Quote: Originally posted by Brain&Force

From my experience and MrHomeScientist's thread on terbium tetrakis(dibenzoylmethide)triethylammonium, it appears that terbium is a difficult to get to fluoresce properly without the proper organic chromophore ligand. These sites have a lot of information about lanthanide fluorescence.

I also need to figure out the spectrum of the UV lamp I used. The sites I looked at show that the lanthanides fluoresce at wavelenghts from 340 nm 254 nm. I'll try making terbium chloride and hexakis(antipyrine)terbium chloride to see if there's a significant difference in the fluorescence, especially from different sources. This may explain why the double sulfate fluoresces under only the UV tube lamp. (I used 3 sources for my experiments: a UV LED, a tube lamp, and a UV compact fluorescent lamp.)

In the compounds europium (tetrakis)dibenzoylmethide triethylammonium and hexakis(antipyrine)terbium iodide, the dibenzoylmethane and antipyrine ligands appear to be the chromophores that the lanthanide luminescence site discusses.

I have done hours of searching on the subject of antenna chromophores and information is almost exclusively contained within many paid sites which I have no access to. Other than that most sources are discussing biological systems. Do you have any sources (preferably with in depth explanation) pertaining to synthetic organic chemistry on this which are not related to biological organisms? I did find some (non pay per view) information here:

http://parc.wustl.edu/search/node/antenna%20chromophores

such as:

http://parc.wustl.edu/research/themes/biohybrid

however they have a mainly biological focus. What interested me was an organic complex acting as an antenna to absorb light which it then transfers to a Lanthanide atom. I was wondering about purely synthetic crystal structures with a Lanthanide atom trapped in a crystal lattice defect site. This sounds very similar to the theory behind long persistence glow powders. I can even see similarities to room temperature superconducting Perovskites.

blogfast25 - 21-3-2014 at 13:28

Quote: Originally posted by Brain&Force

I might just reconvert it to the chloride and either mass the chloride or precipitate it as the oxalate.

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]

Brain&Force - 21-3-2014 at 14:17

Oh wow...it's been so long since I've been experimenting.

Say, what is terbium again?

blogfast25 - 22-3-2014 at 05:39

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 K3FeOx3.

[Edited on 22-3-2014 by blogfast25]

Brain&Force - 22-3-2014 at 19:36

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.

Brain&Force - 25-3-2014 at 15:38

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.

Brain&Force - 8-4-2014 at 17:36

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.

Big update

Brain&Force - 4-6-2014 at 15:25

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]

blogfast25 - 5-6-2014 at 12:43

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]

Brain&Force - 10-6-2014 at 14:05

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>

Polverone - 10-6-2014 at 17:25

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.

Brain&Force - 10-6-2014 at 20:05

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.

Polverone - 10-6-2014 at 20:22

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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