Sciencemadness Discussion Board - Thulium (III) color

Hmmm. Thanks for all the replies. I purchased 1.125g for $14.50. This is comparable to the price which Metallium charges.

Another possible clue is that this sample is very slightly attracted to a magnet. I can just barely make it slide along a glass surface with a strong NeFeB magnet applied from below. This is consistent with the Wikipedia article's statement that it is paramagnetic at 300K, although I have not yet found a reference to how strongly paramagnetic it is.

[Edited on 29-6-2014 by sbreheny]

Yup, that's kinda what thulium does. (skip to 7:03)

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It'll help if you do two things:

a) look at it under fluorescent (tube or CFL) lighting

b) take pictures! Thulium oxide tends to be a pretty faint green.

Best image I could find. It's not that green. Not even mint green.

Pok - 29-6-2014 at 08:09

c) Thulium sulfate is not very soluble, but it is soluble. Keep that in mind, because you're working with a rare element.

Thulium sulfate is very soluble! Rare earth sulfates show very different solubilities:

http://link.springer.com/article/10.1007%2FBF00651459#page-1

Quote: Originally posted by blogfast25

No green colour at all, not even the hydroxide, is rather suspicious.

The rare earth hydroxides are not that deep in colour as the salts in my experience. Sm(OH)3 for example is nearly white, whereas the sulfate is yellow in the crystalline form.

Here you can see 60 grams of Thulium sulfate octahydrate:

This is the true colour which you will see with your naked eyes in daylight. Thulium oxide is weaker in green colour, but still very slightly greenish. At 405 nm none of the two compounds show a fluorescence.

[Edited on 29-6-2014 by Pok]

IrC - 29-6-2014 at 09:14

"Another possible clue is that this sample is very slightly attracted to a magnet."

Along with a few other elements. However the lack of green tells something about electronic structure which I don't think I could overlook. Remember the seller could be OK but they themselves could have been taken by who they buy from. Especially with little experience as a feedback of 181 implies. This seller possibly has no way to test the elements he is selling, and equally possible is the thought that no one who bought Thulium from him did the tests you discuss in this thread. Woelen said it, contact the person with your questions and concerns. I wouldn't rush to leave him negative feedback as his score is a perfect 100 and one has to work hard to maintain this. As I said possibly he was taken as well but is willing to make it right. Try making the oxide out of some and compare it to what is known concerning properties.

Just had another thought. What is the source of your acid? How pure is it? I ask as many use sources from OTC products which contain many impurities, which could possibly radically alter the reactions you were trying to conduct. Something to consider. His Dy looks correct if the pictures are of the actual item he is selling. I have a lot of Dy and it looks identical to his images in the listing. All in all the seller looks legit yet anyone can be taken including him. A good seller will correct the issue if there is one so your next step should be talking to him.

nezza - 29-6-2014 at 10:26

This post piqued my interest. I have a set of Lanthanides bough a few years ago when they were a bit cheaper including 1 gram of Thulium. I sacrificed 0.3 of this sample to investigate. It dissolved easily in 2.5ml of 1M sulphuric acid (Slight excess of acid used) to give a very pale green solution. I did a visible spectrum of this which shows a typical lanthanide absorbance spectrum with some sharp peaks especially at about 680nm. I have attached pictures of the solution and the spectrum.

IrC - 29-6-2014 at 10:46

Quote: Originally posted by nezza

very pale green solution

This reinforces my thought about his source of acid. I had been wondering if some other impurity/reaction could produce something which would tend to visibly mask the green color sbreheny had been looking for. Especially after carefully looking at all the items for sale, all of which appear to be legit. Is it possible some side reaction has 'covered up' the green?

Brain&Force - 29-6-2014 at 11:17

Iron, iron, iron.

In order to work with the lanthanides, you MUST have an iron-free source of acid. This is absolutely important, especially if you're going after the fluorescence, or any color changes.

Here's the kicker: terbium sulfate is supposed to be white, but a tiny amount of iron (< 3%, I think) turned it slightly green. Not as green as Pok's thulium sulfate, but still noticeably green, similar to the color of the thulium oxide I posted here.

And that's interesting, Pok. As expected, terbium sulfate was quite insoluble when I was working with it. I thought thulium sulfate would be the same.

Here is the spectrum of fluorescent lighting:

sbreheny - 29-6-2014 at 11:57

Thanks for all the help. This forum is great! The photo of the Thulium sulfate solution looks so pale green to me that I would be in doubt as to whether I was actually perceiving green or just thinking it by auto-suggestion.

In my previous attempt, I used technical grade sulfuric acid. I also have some ACS reagent grade (began as 98% but I diluted it with distilled water). I re-tried the reaction with the purer acid and I don't see much difference. But, I boiled off the water and I am now drying the precipitate to see what it looks like when dry. It looks slightly brownish-yellowish right now, but mostly white.

Additional data points - the metal does react with water (I can see bubbles forming under a microscope), but very slowly. Even if I heat the water, I didn't get much of the metal to dissolve. This could be due to the metal being a solid chunk and not porous enough to present adequate surface area.

Also, reaction with iodine is supposed to produce a yellow salt. I tried that, with enough heat to sublime the iodine, and I did get a yellow residue around the edges of the chunk. Again, I think that it might need to be made into a powder to react fully with the iodine.

I tried filing a small piece to get a powder but it is actually quite hard and doesn't file easily.

When I get the precipitate dried I will take photos of it as well as the bulk material and post them.

Brain&Force - 29-6-2014 at 12:50

Brownish-yellow? Try a thiocyanate test on your acid.

I was never really able to get bulk terbium to react with iodine with simple, but powder should work. If it still doesn't work, add a few drops of water to the mix.

Also, try burning a small piece and figure out what color it burns. It'll probably be greenish:

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[Edited on 29.6.2014 by Brain&Force]

Boffis - 29-6-2014 at 13:20

some time ago I read an article about thermoluminescense and it gave thulium oxide as one of the examples, it apparently glows red well below "red heat" c200-250 C. As far as I can ascertain this is unique amongst the RE's. Europium oxide itself and europium or samarium produce red and orange fluorescence in some compounds ( eg calcium tungstate) but as far as I can see from the information I have found their oxides are not thermoluminescent. However, thermoluminescense is usually a response to past irradiation so it may need to be irradiated first (with shortwave UV?).

I don't recall where the article was from but it shouldn't be too difficult to tract it down. Lots of papers about thulium and dysprosium induced thermoluminescense in other substances come up on a google search. Particularly in calcium sulphate which should be fairly easy to produce.

IrC - 29-6-2014 at 13:27

"I tried filing a small piece to get a powder but it is actually quite hard and doesn't file easily."

Reminds me of a thread around here somewhere and the troubles I had making glow powders. I had to give up hack sawing my Lanthanides since Fe in incredibly small amounts poisoned my work killing the glow. The statement by Brain&Force "Iron, iron, iron" is something to think about. I was mostly working with Eu, La, Nd as dopants for both Sr and Ca Aluminates. I resorted to using abrasive cutting wheels. This trouble was created by Fe in such tiny amounts I could not even determine it was present. Hack sawing a 2 pound block of Eu and the putting the unused portions back in the paint can of oil left microscopic Fe particles in the oil which contaminated the Lanthanides. As I removed a piece the Fe particles stuck to it, carried in the oil. I washed my pieces with Acetone many times putting them back into new cans with new oil. Only then did my contamination troubles cease. A file will also cause this trouble. I think you should use a Dremel with an abrasive cutting wheel to cut it. Not sure what would work to yield Fe free powder for you but a Ball Mill with ceramic media might do it.

Brain&Force - 29-6-2014 at 14:01

Quote: Originally posted by IrC

...sawing a 2 pound block of Eu...

Where can I get one? Can I get one for free? Please?

[Edited on 29.6.2014 by Brain&Force]

Pok - 29-6-2014 at 15:44

Quote: Originally posted by Boffis

some time ago I read an article about thermoluminescense and it gave thulium oxide as one of the examples, it apparently glows red well below "red heat" c200-250 C. As far as I can ascertain this is unique amongst the RE's.

I also read about this property and can confirm it. But I think the temperature is much higher than 200-250 °C. Tm2O3 glows "carmine red" when heated directly (!) with a bunsen burner flame. Unfortunately, the colour can't be catched with a camera. If heated to higher temperatures, the colour turns to yellow. Both colours are clearly different from a usual thermal radiation.

"The oxalate was of a very pale green tint when moist. Artificial light increased this tint so that it became a delicate pale green color. The oxide, obtained by igniting this oxalate, also showed this green tint and yielded a carmine glow, when heated carefully by a Bunsen burner."
http://pubs.acs.org/doi/abs/10.1021/ja02221a007

Or here.

[Edited on 29-6-2014 by Pok]

IrC - 29-6-2014 at 15:50

Where can I get one?

China but not cheap. In fact had I not bought my Lanthanides about a decade ago no way could I do it today. Some like Nd (2 lb) and La (1.5 lb) I bought from Hamric around 2004 when he had a killer deal for these two large ingots. Today under his Metallium name he has La 50 gm for $38 and Nd 50 gm for $45. Meaning he is paying a hell of a lot more for everything today as opposed to a decade ago since I doubt in experimenter quantities any sellers are more reasonable and fair in pricing. Emovendo is similar but I think he is more into selling magnets than elements today.

I've seen prices go insane in the last 5 years. To be honest I do not see why either. Eu is going up to $40/gm on ebay in sample quantities looking at several sellers. No way could I start out today and it does not look like this trend is ever going to stop. Look at Gold. $271/oz in 2001 and ~$1315/oz today.

Brain&Force - 29-6-2014 at 15:58

Considering that europium is more common than iodine in Earth's crust, I don't see how it can be so expensive. And demand is pretty low anyway. Samarium, however, is much easier to obtain. Emovendo doesn't sell any rare earths now. I know of no other suppliers.

Once the US gets its rare earth mines back online prices will (hopefully) drop dramatically.

Texium - 29-6-2014 at 17:16

That demand is low is another reason why the prices are high, rather contradictorily to supply and demand theory. Since there is no reason to develop more efficient means of mass production to meet a high demand, most rare earths are highly priced partly due to the effort it takes to process them.

sbreheny - 29-6-2014 at 17:57

Mostly finished drying the precipitate. I think that the slightly yellowish-brown color was due to concentrated sulfuric acid - the drying removed the water but not the acid until it reached the azeotropic concentration. Then, I increased the temperature of the hotplate to boil off the azeotropic acid (only a few drops). Note to self: it only takes a few drops of nearly pure h2so4 to produce a lot of irritant vapor! I had a fair amount of ventilation but not enough - I had to put on my respirator to go back into the room to open the window wider and put a fan in the window. It's only slightly detectable now.

sbreheny - 29-6-2014 at 18:21

OK, here are the photos.

Crystals on Watchglass

Crystals magnified

Raw Metal

sbreheny - 29-6-2014 at 18:24

The crystals (which should be Thulium sulfate) look white to me. There is an intensely green tiny crystal in the magnified image but that crystal is actually iridescent and fairly transparent. For scale, the piece of metal is about 0.75 inch long. The crystals unmagnified cover an area of about 1 inch by 0.75 inch. The magnification is roughly 80x in the magnified image.

IrC - 29-6-2014 at 18:57

Your image of the raw metal looks like the piece of Thulium metal I'm holding right now. One question is, you keep mentioning concentrated acid (and boiling off the solution) yet wiki clearly mentioned using dilute acid? Could this be why you are having different results?

sbreheny - 29-6-2014 at 20:04

The first time I dissolved a piece in acid, I used some acid from a bottle which I had pre-diluted a while ago to 1 M. I used technical grade 98% H2SO4 and diluted it with distilled water.

Because someone mentioned the possibility of contamination of the acid, I remembered that I also had a bottle of 98% ACS reagent-grade H2SO4. So, for the second try, I put a small piece of the metal in a test tube and added about 3mL of distilled water. I then added a few drops of the ACS reagent-grade 98% H2SO4. I would say that I ended up with probably a 20% acid concentration - a bit stronger than the first time but still not terribly strong.

I also mentioned strong acid because, as I was removing the water to recover the crystals of the supposed Thulium sulfate, of course, the remaining acid became more and more concentrated until it would no longer evaporate at the temperature it was at (maybe 80 or 90C), so I increased the temperature considerably (to maybe 250 or 300C) and it dried almost completely in short order, but produced a small cloud of very irritating vapor, which I believe was h2so4 vapor.

Pok - 30-6-2014 at 00:25

Many rare earth metals can look like your sample. The fact that IrC's thulium looks the same, doesn't prove anything. It's just a hint, that it could be thulium.

The crystals don't look like thulium sulfate octahydrate at all! They should be clear (!) and greenish. If you prepared them by boiling off the concentrated acid, you will propably get anhydrous thulium sulfate. And I doubt that this compound is greenish (many anhydrous salts look totally different compared to the hydrates)! Just dissolve the crystals in a minimum amount of water and let the water evaporate slowly at room temperature. The crystals obtained in this way will be the octahydrate which looks greenish. In general, faint colours of crystals are only easy to observe, if you have a larger amount of them. If you only have tiny amount, the characteristic colour is difficult to see!

BTW: using metallic Tm as a starting material is not the best way. First, high purity RE metals are much more difficult to produce than the oxides (very high purities available). Second, the metals are usually more expensive than the oxides (molar mass difference included in this calculation).

sbreheny - 30-6-2014 at 07:39

Thanks, Pok. Unfortunately, the whole point of this exercise is to see if my sample is Thulium and I don't want to sacrifice the entire thing to get a lot of the sulfate. I did re-dissolve the crystals in a small amount of water and left it sitting in air - I'll take another photo of the crystals when they form again. I had to find a place to put this where my inquisitive and acrobatic cat would not be able to get to it

I hope I'll be able to see the subtle color with magnification and proper light.

I checked the pH of the new solution (crystals re-dissolved in a small amount of water) and it was about 3 - so there is still a tiny amount of acid left but not that much - looks like I did get rid of the vast majority of it.

sbreheny - 30-6-2014 at 07:43

I also did try to burn a small amount of the raw metal. I used a MAP-Pro (propylene gas) torch because it was handy. It glowed a bright orange but no flame. Maybe once again the surface area is not high enough? Examination under magnification after trying to burn it showed a black layer (carbon?) had formed on the surface but there were a few spots where white crystals emerged from the edges - possibly some of the oxide?

Brain&Force - 30-6-2014 at 13:01

No flame? That's really weird. Rare earths tend to burn really easily.

I'm an international hazard now! (just noticed that)

[Edited on 30.6.2014 by Brain&Force]

sbreheny - 30-6-2014 at 14:27

Well, this link says that it is "not flammable in ingot form", only in dust or thin foils:

http://www.espimetals.com/index.php/msds/288-thulium

I would hardly call mine an ingot but it is definitely thicker than a foil.

Brain&Force - 30-6-2014 at 16:07

Well, neither is magnesium, in that sense: you'll need minutes of continuous direct torching to light them in ingot form. They need tremendous amounts of energy to ignite.

I've had success in taking tiny shards of terbium (3 - 5mm long × 1 mm across) and dropping them directly into a Bunsen burner flame. That gave the expected golden flame. Note that the metal will not ignite through heating alone - you must expose it to the flame.

IrC - 30-6-2014 at 17:55

Many rare earth metals can look like your sample

Can you name the ones that do? I have them all except for Pm, some in multi pound quantities. Honestly I spent quite a while comparing them all yesterday and for the life of me I cannot confuse a single one with Tm. Would be helpful if he would post another pic of his sample but a side view of the piece in more natural lighting, just to be positive it's not La, Lu, or Dy. Even with the image he took however I still see a difference between them and the similarity to my Tm.

Brain&Force - 30-6-2014 at 18:13

Why do you have to make me jealous, IrC?

Yes, the rare earths are all very different, contradictory to what many chemistry books say about them. Terbium is quite dark, from my experience (probably due to the formation of dark nonstoichiometric terbium oxides on the surface) but I've never seen any other rare earths. Ytterbium appears to be slightly brassy compared to most other rare earths.

I just saw your pictures, sbreheny, and I don't know why your crystals have that black stuff in them. They look almost like marble. Did you take the photos of the metal under soft white incandescent lighting?

IrC - 30-6-2014 at 19:34

"I've never seen any other rare earths"

You can look at David's 4 year chart.

http://www.elementsales.com/re_exp/index.htm

Only thing a mystery to me is My Sm came from David out of the same batch AFAIK he used a chunk of in his time test. His still exists. Going through mine one day after 4 years I noticed the once sealed bag (fairly thick plastic vacuum sealed) had expanded to bursting. A big huge amount of powder. Now I figured out the O2 inlet came from moving in storage and the stringy nature, a string had rubbed through making a hole. Yet his only turned color. All I had left was a nice large amount of Samarium Oxide. Mind you I had planned on making some but damn not all of it. My replacement quantity I now keep under oil as several are stored. All I can figure is even though the time was the same 4 years for both of us, it must be the much greater humidity I live in. Not being very good at Chemistry I cannot provide a better reason. I only study enough for what I am doing at the time since electronics has always been my main area. Mad Science is my hobby. I just cannot explain why mine is all dust in the wind (actually in a big bottle now) when David's sat on a shelf on a card in the air for the same 4 years. Unless anyone has a better theory my 55% or greater humidity all year has to be the reason Oxidation accelerated as it did.

Look at his Sm after 4 years:

Brain&Force - 30-6-2014 at 19:35

Yeah, I've referenced that chart many times. I just haven't seen them up close and personal.

If the Sm was stringy, it may have been the increased surface area that allowed for accelerated corrosion.

IrC - 30-6-2014 at 19:57

If his piece in the pic was stood up a closeup would reveal it's exactly as stringy as mine. Meaning lengthwise crevasses grooves, whatever you want to call it exist all through down the length. If you grabbed the piece off his card in the pic above with pliers you could pull it apart several strings at a time. Implying O2 can go far inside the chunk easily between the strings. Of all my Lanthanides this makes Sm by far the simplest to get small pieces for experiments. You should try taking a small sample off a chunk of Gadolinium. Even with a good hacksaw it's work. That along with the Fe contamination troubles I had is why I use a Dremel cutting wheel at high RPM now.

Look closely at this pic from wiki is easy to see what I mean:

sbreheny - 30-6-2014 at 21:52

I will take more photos of the raw metal when I get a chance - maybe tomorrow. As I said earlier, I re-dissolved the crystals in pure water and let that evaporate today. It is almost completely dry but I want to give it another day. The crystals do look very different now - I will get a photo of them as soon as I can. I took the unmagnified photos of the raw metal and crystals under a compact fluorescent light. The magnified one is using my microscope illuminator, which is white-LED based.

IrC - 30-6-2014 at 22:44

I think POK has the answer when he was talking about the state of hydration concerning your difficulties.

Pok - 1-7-2014 at 01:17

Quote: Originally posted by IrC

http://woelen.homescience.net/science/chem/exps/neodymium/in...

Many rare earth metals can look like your sample

Can you name the ones that do?

Samarium, Europium and Scandium for example can show such a golden oxide layer. Of course, you can exclude at least Eu (too reactive) and Sc (too expensive).

Quote: Originally posted by IrC

If his piece in the pic was stood up a closeup would reveal it's exactly as stringy as mine. Meaning lengthwise crevasses grooves, whatever you want to call it exist all through down the length.

The similarities you talk about can be explained by the preparation method of the metals and their degree of surface oxidation. The "stringy" looking metal is simply a result of destillation. All RE metals show this structure when prepared by destillation afaik. And the golden colour is a result of a more or less thick oxide layer. Remember steel with it's colour variation which can result from heating/oxidation. I think all RE metals can show this structure and many can show a golden tint, depending on the degree of oxidation.

In general, the attempt to identify thulium only by some visible properties (colour of metal and salts, magnetic properties, etc.) will not be sufficient. There are good reasons why analyses are made with expensive equipment in the case of RE metals (very similar chemical behaviour). Especially the purity isn't easy to determine. But I think that business sellers usually don't cheat you (except for purities).

Brain&Force - 1-7-2014 at 09:06

Can you make some thulium chloride? It should be pretty intensely green compared to the sulfate.

Pok - 1-7-2014 at 09:54

Why? Only the cation makes the colour.

This is how the chloride looks like. Very faint green, comparable to the sulfate. But there the crystals are smaller than my sulfate crystals, which may explain the weaker colouration.

[Edited on 1-7-2014 by Pok]

Brain&Force - 1-7-2014 at 10:15

I'm not sure what it is - might be some sort of a complex forming between chloride and thulium (probably a cationic one).

The anion does have an effect on the color of salts, especially colorful rare earth salts. See woelen's work on neodymium salts - the nitrate is less pinkish and more of a lavender color.

The one on the left is also neodymium chloride, but dissolved in excess HCl. Notice the difference in color:

Pok - 1-7-2014 at 11:16

These are examples of dissolved ions with excess of acid in solution. In this case complex formation can result. But I don't think that this can happen in the solid crystal. Maybe parts of water of crystallization can be replaced by acid molecules and form a solid complex in this way. I didn't notice a strong dependence of thulium(III) colour from the light source as in the case of neodymium ions. I will make some tests with thulium oxide and different HCl concentrations tomorrow.

Brain&Force - 1-7-2014 at 17:34

Terbium does form a hexaaquadichloro complex (which is highly triboluminescent). Unfortunately, terbium compounds are all white, so there's no significant color change.

Wow, this is my 20th post today.

sbreheny - 1-7-2014 at 17:44

OK, here are the additional photos.

Naturally-dried crystals

Crystal Habit (magnified)

Raw metal, side A

Raw metal, side A, magnified

Raw metal, side B

Raw metal, side B magnified

Raw metal, edge, magnified

I will probably try making other salts of this to see if they have a different appearance. I can readily make the chloride, bromide, and nitrate.

I also bought a small Thulium sample from Metallium to compare. It has not yet arrived.

Finally, as suggested, I will try testing the acid for iron. I am waiting for some KSCN to arrive.

Thanks for the help and interest!

sbreheny - 1-7-2014 at 17:52

By the way - when I do the Thiocyanate test on the acid, I will dilute it first to prevent HCN formation. I believe that it is not likely given that this is the SCN ion and not CN, but I have read that if you combine KSCN with a strong acid and heat, you may get HCN.

IrC - 1-7-2014 at 18:30

Here is one of my 5 gm pieces of Thulium.

Best I can do, I sent my entire set of 58 mm closeup and macro lenses to a family member who is studying photography. I knew I should have kept at least one around for just such a need as this thread. Lighting is typical dual 4 foot florescent tubes with camera flash turned off.

sbreheny - 1-7-2014 at 20:08

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 - 1-7-2014 at 20:10

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 - 1-7-2014 at 21:00

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.

Pok - 2-7-2014 at 04:33

you would think that the crystals I would get at the end would be Tm2(SO4)3 still.

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:

Terbium does form a hexaaquadichloro complex (which is highly triboluminescent).

Do you have a source for this claim? Which complex do you mean?

Brain&Force - 2-7-2014 at 08:50

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.

Pok - 2-7-2014 at 10:11

hexaaquadichloroterbium(III) chloride

= 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 - 2-7-2014 at 13:38

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).

Brain&Force - 4-7-2014 at 09:51

sbreheny, I'm curious, is your thulium salt magnetic? I have some terbium sulfate which is definitely paramagnetic.

Pok - 4-7-2014 at 12:25

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 - 6-7-2014 at 21:12

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 - 6-7-2014 at 21:15

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 - 6-7-2014 at 21:50

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.

IrC - 7-7-2014 at 00:54

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...

blogfast25 - 7-7-2014 at 04:56

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 - 7-7-2014 at 13:14

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 - 7-7-2014 at 13:22

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]

sbreheny - 7-7-2014 at 16:31

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.

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 - 8-7-2014 at 04:35

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 - 8-7-2014 at 15:14

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.

sbreheny - 8-7-2014 at 18:16

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 - 9-7-2014 at 04:32

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 - 9-7-2014 at 08:00

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 - 9-7-2014 at 08:57

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 - 9-7-2014 at 09:07

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.

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 - 9-7-2014 at 10:03

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).

but.bmp - 502kB

Pok - 9-7-2014 at 10:29

Quote: Originally posted by kmno4

even small erbium contamination may give colourless thulium salts solutions

I also found this claim. But how should it work? It's substractive colours, not additive. This means that the absorption of Er and Tm is added. So erbium can't make the greenish Tm colour weaker or even "colourless". I think it is a mistake in the article or an imprecise description.

@blogfast: I don't see a barrage of exclamation marks. And no: it's not established that it is thulium! There are 2 weak hints (magnetic and pale green) but no positive proof! The metal could also be 50 % Tm and 50 % xxx (with colourless ions) or even 0 % Tm, 1 % x (with green ions) and 99 % y (with colourless ions).

[Edited on 9-7-2014 by Pok]

IrC - 9-7-2014 at 10:59

it's not established that it is thulium!

He said he bought a 1 gm sample from Metallium. They have 99% Thulium. I have done business with David Hamric for over a decade. A more honest conscientious seller I have never found. If he says it's 99% Thulium then that is what it is. He can also provide assay information upon request.

Brain&Force - 9-7-2014 at 11:16

I can definitely corroborate IrC's statement - Metallium even offered to help me out when I suspected contamination issues with my terbium sample.

Have you tried a density test? That can help us weed out any issues. (Just make sure to use oil rather than water - thulium will slowly react with the water.)

Pok - 9-7-2014 at 11:21

He bought the first sample from another supplier. I also believe that it is thulium. But "belief" is not a postive proof. And in the case of metallium, neither a CoA nor the reputation of a seller is a proof.

kmno4 - 9-7-2014 at 11:44

I think it works in similar way like in case of Co(II) + Ni solutions: at some ratio solution becomes nearly colourless.
Er absorbs "green ligth", Tm absorbs "pink light", total effect depends on Er/Tm ratio (and source of ligt).
It is very similar situation to Nd-Pr pair - colour of n electron on 4f orbital is similar to 14-n configuration (Pr, Tm are green, Nd, Er are pink).
Some time ago I was playing with Nd-Pr mixtures: Nd/Pr oxalate was pale green under fluorescent lamp and pale pink under daylight.
The same situation was in solutions and it seems that at some Nd/Pr ratio solution would be coloureless.
However, Nd has larger molar absorbance coefficient in visible spectrum (let's say it is 10) than Tm does (let's say it is 2,5)

But this "colourless" is oversimplification, as Pok noticed: when some bands dissapear, the rest cannot be strictly colourless !
Human eyes interpret it as "darker", without any particular colours.
Without spectroscopy it is hard to say anything concrete....

sbreheny - 9-7-2014 at 12:18

Pok is correct - my original sample is from a different seller. I bought an additional sample from Metalium to compare against but I have not tried doing any chemical reactions on the second sample - only appearance and magnetic tests.