| Pages:
1
2
3
4 |
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
So, I just made Thulium Chloride, Bromide, and Nitrate solutions. I have now used up about half of my Tm sample.
This time, I was much more careful to use only slightly excess acid (HCl, HBr, and HNO3). Because of this, the reactions are proceeding much more
slowly and are not complete yet - each test tube still has a tiny piece of Tm at the bottom with a trail of bubbles rising from it.
However, I can definitely tell that the TmCl3 and Tm(NO3)3 solutions are very pale green. The TmBr3 is more ambiguous.
Maybe the problem with the H2SO4 was way too much acid. Still - you would think that the crystals I would get at the end would be Tm2(SO4)3 still.
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
IrC - thanks for the photos - I would say that the comparison to my sample is unclear, but that is probably due to your sample being formed by a
different process, resulting in a different surface appearance.
|
|
|
IrC
International Hazard
Posts: 2710
Registered: 7-3-2005
Location: Eureka
Member Is Offline
Mood: Discovering
|
|
Quote: Originally posted by sbreheny  | | IrC - thanks for the photos - I would say that the comparison to my sample is unclear, but that is probably due to your sample being formed by a
different process, resulting in a different surface appearance. |
More likely due to mine sitting in a plastic envelope for 10 years in a hot storage room. The grain structure is what you should be closely comparing.
More important than minor color differences from various levels of surface oxidation over years. Your sample may be closer to it's date of creation
and stored better.
"Science is the belief in the ignorance of the experts" Richard Feynman
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
No, but a hydrate. Probably the octahydrate. Your photo shows rather a crystal mass and not well formed crystals. Thulium sulfate often forms a
supersaturated solution and crystallizes out very fast after a while. It's not so easy to get well formed crystals.
Here is a comparison of Tm(III) in water and in ca. 25 % HCl. There is only a very little difference. But this difference is not easy to identify with
the naked eye. I would still call it "very pale green":
left: water, right: HCl
The solid and "neutral" TmCl3 * x H2O - pale green like the sulfate
If the chloride is heated with sulfuric acid and further heated until white fumes of sulfuric acid appear, white crystals form. They don't look the
slightest bit green. This is probably the stuff which you got in your first try. Thulium sulfate with less than 8 moles water per mole:
Do you have a source for this claim? Which complex do you mean?
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
http://www.researchgate.net/publication/231348013_Crystal_st...
I don't have access, but this paper references a complex, hexaaquadichloroterbium(III) chloride, which appears to be highly triboluminescent.
Those are interesting results Pok, but they don't seem to explain the black stuff in sbreheny's crystals. I'm guessing any complexation that's
occurring due to the presence of the chloride ion is minimal.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
= 6 H2O + 2 Cl + Tb + Cl
= TbCl3 * 6 H2O
This is ordinary terbium chloride (hexahydrate). The complicated name is just a more precise description of the structure of the molecule. The article
says that it's triboluminescence (TL) wasn't observed until 1989. Therefore, it's very improbable that the TL is bright. The article doesn't specify
the brightness of terbium trichloride hexahydrate (not even "bright" or so). This info should be specified in the literature elsewhere. But I think it
is very very weak and not comparable to the other 2 named in the paper. I think the authors are a bit imprecise when talking about "3
brilliant triboluminescent complexes". They also call it "terbium hexahydrate" in one passage which is absolute nonsense.
But yes. Here we see that nearly every hydrated salt is a complex in the solid state, also the lanthanide chloride hexahydrates. This can explain the
difference of the colour of hydrated vs. water-free thulium sulfate. IrC cited a wikipedia article which claims that. Although the "reference" is a
simple web page, it sounds reasonable.
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
Hmm...I knew that most salts were solid state complexes. It's good to know that terbium chloride isn't as triboluminescent as the paper states. I
guess these people didn't bother to revise their paper...
I've heard that neodymium can form acid and basic sulfates, so that may be possible with thulium as well. But the loss of water is the more likely
cause, as water molecules do most of the coordinating and produce the colors we see in transition metal salts, and the same is seen with many
dehydrated sulfates (like copper sulfate, for example).
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
sbreheny, I'm curious, is your thulium salt magnetic? I have some terbium sulfate which is definitely paramagnetic.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
My Tm oxide and sulfate are magnetic (using a very strong magnet). Also the oxides of: Er, Ho, Eu (weak)
Not the oxides of: Pr, Nd, Y, Sm, Yb, Sc, Lu
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
Yes, the salt (presumably Thulium Sulfate) is slightly attracted to a magnet. It is not strong enough to be affected through the wall of a glass
bottle but if I lower the NdFeB magnet down on top of it, at about 1mm away the powder jumps to the magnet and sticks to it.
Now, I have some interesting results to show. I tried the thiocyanate test and I need some help interpreting it. I have five test tubes. The first one
(on the left) had a sample of Tm added. The next one had 1mL of my technical grade H2SO4 added. Next one had 1mL of ACS reagent grade H2SO4 added. The
next one had a tiny amount of iron powder added, and the last one had only distilled water added.
For the iron one, I actually put in 100mg of iron powder and dissolved it in 2mL water and 1mL 70% HNO3 (note to self once again - this produced a
nice little cloud of NO2 which I had to avoid - I really do need to get a hood!). I then poured out most of the resulting solution and diluted it
about 10:1 with distilled water.
Anyway, I then added to each one (except the iron) 1mL of 70% HNO3, my idea being that it would convert any iron into Fe3+ ions. I then added 3mL more
distilled water to each and finally approximately 250mg of KSCN.
At first, this happened:
Start
Then, after about 20 seconds, it looked like this:
Settled
I tried agitating the two H2SO4 samples to re-mix the red and clear layers, but then they released a small amount of NO2 and the solution became a
very pale blue:
Almost Final
Finally, even the Tm sample underwent the same change (released a small amount of NO2 and turned light blue). The distilled water sample remained
light pink and the Fe remained dark blood red.
Final
Now, of course, the Fe behaved as expected. The fact that the distilled water showed some iron could be real or it could be due to some accidental
cross-contamination when I was using a dropper to add to each test tube in succession. However, I really do not understand what happened with the
first three. Any ideas?
Thanks,
Sean
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
Another interesting data point - I purchased a 1g sample of Tm from Metallium. When I compare this to my previous sample, the appearance is similar
BUT I cannot get the Tm from Metallium to show visible attraction to a magnet whereas the earlier sample does show very slight attraction. I suspect
that there is some iron contamination in my original Tm. Note that the Tm sample I used in the thiocyanate test was the earlier one (from
RareEarthMetalsLLC).
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
If you look at the magnetism video that I posted, you'll see thulium is attracted to a neodymium magnet. What kind of magnet are you using? Thulium
should be slightly paramagnetic.
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
IrC
International Hazard
Posts: 2710
Registered: 7-3-2005
Location: Eureka
Member Is Offline
Mood: Discovering
|
|
| Quote: Originally posted by sbreheny | | Another interesting data point - I purchased a 1g sample of Tm from Metallium. When I compare this to my previous sample, the appearance is similar
BUT I cannot get the Tm from Metallium to show visible attraction to a magnet whereas the earlier sample does show very slight attraction. I suspect
that there is some iron contamination in my original Tm. Note that the Tm sample I used in the thiocyanate test was the earlier one (from
RareEarthMetalsLLC). |
I ran into similar troubles buying different rare earths in 20 to 150 gm amounts from various ebay sellers. In my case of using the elements for
doping glow powders even microscopic Fe impurities caused large problems. I surmised this was caused by so many sellers using hacksaw blades (or
motorized saws) to cut the pieces from a larger one. Of course this was after making the same mistake myself in my first year conducting glow powder
experiments and learning the hard way never cut rare earths with metal tools if purity is important.
Also, have you tried putting them on styrofoam floating in water as I have seen done in a few videos on the subject? He was able to observe even tiny
effects in this way. Here is a link to the first of 3 videos he did on this subject.
http://www.youtube.com/watch?annotation_id=annotation_828992...
"Science is the belief in the ignorance of the experts" Richard Feynman
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
sbreheny:
Re. your thiocyanate tests.
Firstly, check all reagents used for iron.
Secondly, use hydrogen peroxide (check it too for iron) instead of nitric acid. With the latter the oxidation is often delayed somewhat (my experience
also with Fe(II)), not to mention the NO2 fumes that stink. Heat the tubes to destroy any excess H2O2, then cool and add thiocyanate.
[Edited on 7-7-2014 by blogfast25]
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
Thanks for the advice, I will try that. However, do you have any possible explanation of why the first three tubes would turn red initially but then
become clear? Do you think that the red color was not the Fe-Thiocyanate complex but rather NO2 and as the solution cooled, it released the NO2?
|
|
|
Brain&Force
Hazard to Lanthanides
Posts: 1302
Registered: 13-11-2013
Location: UW-Madison
Member Is Offline
Mood: Incommensurately modulated
|
|
I've had that occur sometimes with extremely small amounts of ferric ion and extremely large amounts of thiocyanate. The same thing occurs with small
amounts of starch and large amounts of triiodide. Not sure what causes this but it's likely an equilibrium thing.
It's sad to see your distilled water contains iron. Try deionized water, if available.
[Edited on 7.7.2014 by Brain&Force]
At the end of the day, simulating atoms doesn't beat working with the real things...
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
I am using an NdFeB magnet which is N35 or N42 (I don't know why the package does not say for sure which it is but that is the range). I just tried it
again and I see that the difference I observed was just due to the shape of the bottom of the bottles the two samples are in. The older sample is in a
smooth-bottom bottle. The newer one (Metallium) is in a concave-shaped bottom bottle. When I put the magnet directly near the samples, both are weakly
attracted to the magnet. If I put the new sample in the old one's container, I can indeed slide it around the bottom by applying the magnet from the
outside.
| Quote: Originally posted by Brain&Force | | If you look at the magnetism video that I posted, you'll see thulium is attracted to a neodymium magnet. What kind of magnet are you using? Thulium
should be slightly paramagnetic. |
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by sbreheny | | However, do you have any possible explanation of why the first three tubes would turn red initially but then become clear? |
My own guess is destruction of the thiocyanate by the nitric acid.
With peroxide, try using sparing, reasonable amounts (calculate, if you can).
A little known fact is that when you add nitric to Fe(II) initially the solution darkens due to formation of an Fe(NO)<sup>2+</sup>
complex. This can create the illusion that the oxidation (Fe II == > Fe III) is over, when it isn't.
|
|
|
kmno4
International Hazard
Posts: 1495
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
Fe/SCN test for small (expected) amounts of Fe must be done in strongly acidic conditions, 1 M total concentration of HCl would be ok. High
(relatively) amount of thiocyanate should be used (because of not very large stability of Fe(III) complex).
And the most important: extraction of formed red complex into polar organic phase. Extracted complex (in form of "ferrithiocyanate acid") is more
stable in this medium than in water.
For testing in tubes, 0,5 cm3 (or similar amount) of organic solvent should be used, with shaking and allowing to separate.
Good extractants are higher alcohols, for example (n-)butanol (interesting, n-butyl acetate does not extract this complex from water)
PS. instead of 1 ml 70% HNO3, one or two drops would be better, but I prefer H2O2 instead of HNO3.
Слава Україні !
Героям слава !
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
Hmmm. OK, so I have one recommendation to use H2O2 and another one to use H2SO4 and then separate the red complex using a polar organic solvent. I
don't really need the result to be stable, just as long as it persists long enough for me to see it and I can be sure that I didn't screw up the test
and do something wrong (I was not expecting the color to disappear so I assumed I did something wrong and perhaps the initial color change was not
indicative of iron).
If I use H2SO4, won't I have Fe(II) ions instead of Fe(III)? Will they still react with the SCN to form the red complex?
As for using Butanol to separate out the red complex from water - the wiki article on Butanol says that it is fairly soluble in water, so I don't
understand why it would form a distinct layer.
Sean
| Quote: Originally posted by kmno4 | Fe/SCN test for small (expected) amounts of Fe must be done in strongly acidic conditions, 1 M total concentration of HCl would be ok. High
(relatively) amount of thiocyanate should be used (because of not very large stability of Fe(III) complex).
And the most important: extraction of formed red complex into polar organic phase. Extracted complex (in form of "ferrithiocyanate acid") is more
stable in this medium than in water.
For testing in tubes, 0,5 cm3 (or similar amount) of organic solvent should be used, with shaking and allowing to separate.
Good extractants are higher alcohols, for example (n-)butanol (interesting, n-butyl acetate does not extract this complex from water)
PS. instead of 1 ml 70% HNO3, one or two drops would be better, but I prefer H2O2 instead of HNO3.
|
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by sbreheny | If I use H2SO4, won't I have Fe(II) ions instead of Fe(III)? Will they still react with the SCN to form the red complex?
|
Both HCl and H2SO4 are non-oxidising acids (in this context) and rely on the oxidative power of H<sub>3</sub>O<sup>+</sup>.
This ion can oxidise iron only to the ferrous form, so both HCl and H2SO4 lead to Fe<sup>2+</sup>, which does NOT form the blood red
coloured complex with thiocyanate.
That is why an oxidiser like H2O2 needs to be used prior to adding the thiocyanate, so that the peroxide oxidises all ferrous promptly to ferric iron,
which does yield the coloured complex.
I've never heard of extracting the complex with organic solvents. It may present some improvement but in my experience simple testing for
Fe<sup>3+</sup> does not require this extra complication.
In strongly oxidising conditions the thiocyanate ions probably won't last very long but as you pointed out it is sufficient to see the complex for a
brief but definite time to confirm the presence of iron.
As kmno4 suggested, keep the amount of oxisiser low: there really isn't that much iron in a couple of ml of your solutions. And you don't even need to
oxidise ALL iron to Fe (III) to actually get a positive result, even a partial oxidation will give a positive if Fe is present at all.
Re. n-butanol, with 73 g/L water solubility (Wiki) a two phase system will form if you use more than 73 g of n-butanol per L of water. 1 ml n-butanol
per ml of solution would definitely do it. But this is more useful if one wanted to actually isolate the complex than for a simple test, IMHO.
[Edited on 9-7-2014 by blogfast25]
|
|
|
Pok
potassium Prometheus
Posts: 176
Registered: 5-12-2010
Member Is Offline
|
|
I don't see a reason for you to test for iron. Thulium is very slightly magnetic as your sample is. This cannot be explained by the very low iron
content! Even if it is contaminated with iron, it is very very low as you can see from your test. This low contamination won't lead to such a rather
strong attraction to a magnet!
Your 100 mg of iron are far to much! Do you really believe that much iron is in your other solutions? This would mean about 0,7 grams of iron nitrate
hydrate!
You can test for iron with potassium ferrocyanide (for Fe3+) or potassium ferricyanide (for Fe2+). In the presence of iron you
will get a blue solution or precipitate ("prussian blue").
Both tests (cyanide-complex and thiocyanate) are very sensitive. Your thulium metal is not 99.999 % pure, so there will be iron in it! This
also isn't a usefull method to decide wether you have thulium or something else. There is no reason to believe that it is something else but thulium.
You have to find a real test for thulium (in the way it is done for other lanthanides like here).
[Edited on 9-7-2014 by Pok]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Pok | | This also isn't a usefull method to decide wether you have thulium or something else. There is no reason to believe that it is something else but
thulium. |
Wow. Calm down already (what's with the barrage of exclamation marks?)
As far as I can see it's been established it is Thulium. But there are appears to be a small contamination of iron. That appears to be quite common
with some Ln elements.
Nothing wrong with investigating the possible source of that contamination. My bet is on the cutting tool.
|
|
|
sbreheny
Hazard to Others
Posts: 145
Registered: 30-1-2014
Member Is Offline
Mood: No Mood
|
|
As you can see from some of my mistakes, I am still very much learning. I am an electrical engineer, professionally, and I tinker in lots of stuff and
have broad general science knowledge but I have a lot to learn. I wanted to test for iron in order to learn how its done.
I do agree that it is most likely Thulium. There was some doubt for a while because the Thulium sulfate crystals I got do not appear light greenish
(at least, I do not perceive it) but then I made the nitrate and chloride salt solutions and they DID have a noticeable green tint so I am satisfied
that it is Tm. In any event, this is all about learning and fun.
As for the 100mg iron - yes, I do understand that is a huge amount for this test. If you read my posts again, you'll see that I diluted it 10:1 after
reacting it with the nitric acid to make sure all of it was in solution. This is in part because I have a limited ability to measure much less than
100mg accurately. In retrospect, I probably should have diluted it 100:1, but it still served the purpose of showing that the test could detect iron.
I do appreciate your help and the help of everyone here. I am learning a lot.
| Quote: Originally posted by Pok | I don't see a reason for you to test for iron. Thulium is very slightly magnetic as your sample is. This cannot be explained by the very low iron
content! Even if it is contaminated with iron, it is very very low as you can see from your test. This low contamination won't lead to such a rather
strong attraction to a magnet!
Your 100 mg of iron are far to much! Do you really believe that much iron is in your other solutions? This would mean about 0,7 grams of iron nitrate
hydrate!
You can test for iron with potassium ferrocyanide (for Fe3+) or potassium ferricyanide (for Fe2+). In the presence of iron you
will get a blue solution or precipitate ("prussian blue").
Both tests (cyanide-complex and thiocyanate) are very sensitive. Your thulium metal is not 99.999 % pure, so there will be iron in it! This
also isn't a usefull method to decide wether you have thulium or something else. There is no reason to believe that it is something else but thulium.
You have to find a real test for thulium (in the way it is done for other lanthanides like here).
[Edited on 9-7-2014 by Pok] |
[Edited on 9-7-2014 by sbreheny]
|
|
|
kmno4
International Hazard
Posts: 1495
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
Mentioned extraction is useful only for comparation of very small amounts of Fe. When concentration is relatively high, no extraction is needed. Very
small I mean < 1 μg/dm3 of Fe.
On the picture, solution of SCN complex in water and the same sample + butanol.
Back to topic - without VIS spectrum there is no simple way to check purity of your thulium.
Mentioned paper (DOI: 10.1021/ja02221a007) gives interesting example - even small erbium contamination may give colourless thulium salts solutions (I
think it is also valid for Nd contamination).
Слава Україні !
Героям слава !
|
|
|
| Pages:
1
2
3
4 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
