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Brain&Force
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A miracle has occured.
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]
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 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?
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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 Tb<sub>2</sub>(SO<sub>4</sub> <sub>3</sub> 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 Tb<sub>2</sub>(SO<sub>4</sub><sub>3</sub>. 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]
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Brain&Force
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| 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.
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blogfast25
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@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).
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Brain&Force
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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.
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Eddygp
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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.  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.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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Brain&Force
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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.
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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Woo hoo!
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]
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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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]
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Brain&Force
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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.
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'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:
 
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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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]
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Brain&Force
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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]
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Brain&Force
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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.
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blogfast25
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B&F:
Good to see confirmation of the ferric ion influence.
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Brain&Force
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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...
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blogfast25
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| 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 CO<sub>2</sub> 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]
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Brain&Force
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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]
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blogfast25
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| 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.
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Brain&Force
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I'm baaaaaaaaaack!
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).
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blogfast25
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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 Fe<sup>3+</sup> 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]
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Brain&Force
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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.
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blogfast25
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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]
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Brain&Force
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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.
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IrC
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| 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.
"Science is the belief in the ignorance of the experts" Richard Feynman
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