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[*] posted on 1-3-2006 at 01:32
Iridium (IV) chloride


One of my sources for chemicals recently added a compound, called iridium (IV) chloride, formula IrCl4, to its list of chemicals. It can be purchased at a very good price of appr. $20 per gram, which is really cheap for iridium compounds (e.g. compare with Alfa Aesar or Chemetal). I intend to buy a few grams of this, but I have some questions about this.

Two different books (Chemistry of the elements, second edition by Greenwood and Earnshaw, 1999) and an old 1960 book about inorganic chemistry both tell me that no iridium (IV) chloride was characterized, it only is characterized in its +4 oxidation state in compounds like K2IrCl6 and H2IrCl6. These are two independent sources of information, which I think are quite reliable.

Now, to my surprise I see IrCl4 on the list of a few suppliers (e.g. Alfa Aesar, Chemetal). I asked my supplier about the material and how it looks like. He tells me that it is a very dark (almost black) glassy amorphous solid, soluble in water with an Ir-content of at least 56.5%, the rest being the chlorine and some minor impurities. Theoretical pure anhydrous IrCl4 has an Ir-content of 57.5% by weight. I also looked at the MSDS of Alfa Aesar and indeed its description matches that of my supplier quite well.

Now my question is, what can be the source of such contradicting information? I strongly doubt that this "IrCl4" really is iridium (IV) chloride, but I have no clue what it really is. If any one of you has experience with iridium or happens to know something more about this, I would be very pleased. A Google search did not clarify anything, I really need more information from good textbooks, but I unfortunately do not have access to more info on this (at least not what I know of).




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[*] posted on 1-3-2006 at 03:23


I would assume that this is the product of heating of hexachloroiridic acid, H2IrCl6*6H2O (there's a procedure for this in Brauer!), until a constant weight is reached, hence the glassy (fused?) structure. Maybe this has a more definite iridium content that the hydrated H2IrCl6 (hydrates not always have the exact stochiometric water content).

H2IrCl6*6H2O should give IrCl4 on heating.
If the supplier sells this as IrCl4, then I am sure it is. No chemical supplier is going to sell something that doesn't exist!
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[*] posted on 1-3-2006 at 12:04


"No chemical supplier is going to sell something that doesn't exist! "
Ammonium hydroxide anyone?
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[*] posted on 1-3-2006 at 13:12


Quote:
Originally posted by garage chemist
I would assume that this is the product of heating of hexachloroiridic acid, H2IrCl6*6H2O (there's a procedure for this in Brauer!), until a constant weight is reached, hence the glassy (fused?) structure. Maybe this has a more definite iridium content that the hydrated H2IrCl6 (hydrates not always have the exact stochiometric water content).

H2IrCl6*6H2O should give IrCl4 on heating.
If the supplier sells this as IrCl4, then I am sure it is. No chemical supplier is going to sell something that doesn't exist!

But still I think this is a strange situation. I looked up a third book and it tells that iridum (IV) is stable, but only in complexes, such as IrCl6(2-). Is this "IrCl4" then some special, ill-defined compound? I know of other such ill-defined compounds, e.g. "ferric ammonium citrate" which exists in brown and green variations. The structure of that is very variable and only empirical formulas can be given without exactly knowing the structure of these compounds. Is a similar thing true for "iridium (IV) chloride"?

Anyway, I ordered 3 grams of the stuff :). I'm eager to experiment with iridium compounds, one of the metal, with which I do not have any experience up to now.




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[*] posted on 1-3-2006 at 17:46


Can you give me the supplier so I can order some? I find it hard to locate Ir compounds, and Ir is my favorite element to play with.
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[*] posted on 1-3-2006 at 18:13


You can order it from a variety of companies if you have some affiliation with an institution. As a private individual it will prove difficult. I know Fisher, Sigma, and Alfa sell it. I can not think of a single "cheap" source for iridium chloride.

I know that Sigma Aldrich sells lrCl4 · xH2O (which does exist, Woelen, much like auric (III) chloride as a hygroscopic compound that is difficult to dehydrate). I think I might have seen a bottle of it somewhere; next time I'm in the lab I'll look around for it. IrCl3 also exists, and has a melting point of about 763*C whereup it decomposes to the elements. That decomposition temperature is twice that of the IrCl4. Note that tri- and tetravalent iridum cations exist as ionic solids, and in solution (typically complexed).

Now Woelen, when he said "with an Ir-content of at least 56.5%" that was more or less a guarantee that he's not selling you lrCl4 · 10H2O, but rather a less hydrated compound with a closer to theoretical percentage of iridium. More or less, he's clarifying that you will not be getting the anhydrous compound (if it's anything like its platinum analogue, it would be a feat to dry it). I think the price is quite fair considering the difficulty of preparation. Woelen, I would, again, be much obliged to you if you could send me the source for this compound.

IrC, if you ever have need of the pure metal, I have a friend who can get the pure metal to you for $16/gram.
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[*] posted on 1-3-2006 at 18:51


"As a private individual it will prove difficult."

The story of my life. So many things I have not been able to get due to this, not the least of which is P. Will get back on Ir metal after I get $ back up where it belongs. Been buying a lot of laser rods lately, ruby and yag, no idea why. Some kind of madsci dream I had for something or other.
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[*] posted on 9-7-2006 at 22:41


Iridium -- goddess of the rainbow

I now have done some experimenting with this "iridium chloride" and I must say WOW, what a colorful element this is. It really is great. I made a webpage about this:

http://woelen.scheikunde.net/science/chem/exps/iridium/index...

All colors from the rainbow can be made with it, and the chemicals (besides the "IrCl4") needed for that are really simple! I'm quite sure, that with some effort, and throwing in some more fancy ligands, that many more colors can be obtained.




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[*] posted on 10-7-2006 at 03:13


See if you can make any compounds of it in the (VII), (VIII) and (IX) oxidation states, which are theoretically possible but which have so far eluded synthesis. But for that, you would need powerful fluorinating agents, or access to special electrolytic cells.
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[*] posted on 10-7-2006 at 23:18


I do not think that would be possible in a simple home lab. If powerful fluorinating oxidizing agents are needed for that, then most likely the solvent also must be special. I'm quite sure, that water is not OK for this purpose and probably a solvent of HF or some other high-oxidation state fluoride is needed. Not something you would do in the average home lab.



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[*] posted on 11-7-2006 at 16:17


Very nice, woelen.

All colours of the spectrum - incredible! I am not aware of any other element that can do that.

Quote:
Complex formation with ammonia, and subsequent oxidation

When the green solution of "IrCl4", mentioned above, is mixed with dilute ammonia (5% NH3), then a bluish/green/gray solution is formed. This is shown in the left picture. When some sodium persulfate is added to this bluish/green/gray solution, then the liquid becomes bright purple. On gentle heating, without boiling, the purple liquid becomes beautifully golden yellow. That is shown in the rightmost picture.


This one I can't quite follow chemically. What's going on? Is your starting material IrCl6(2-), or "IrCl4"? Why the quotation marks?




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[*] posted on 11-7-2006 at 18:43


Quote:
Originally posted by chemoleo
Very nice, woelen.

All colours of the spectrum - incredible! I am not aware of any other element that can do that.

Quote:
Complex formation with ammonia, and subsequent oxidation

When the green solution of "IrCl4", mentioned above, is mixed with dilute ammonia (5% NH3), then a bluish/green/gray solution is formed. This is shown in the left picture. When some sodium persulfate is added to this bluish/green/gray solution, then the liquid becomes bright purple. On gentle heating, without boiling, the purple liquid becomes beautifully golden yellow. That is shown in the rightmost picture.


This one I can't quite follow chemically. What's going on? Is your starting material IrCl6(2-), or "IrCl4"? Why the quotation marks?


I'm pretty certain that plutonium can also form a great deal of colored compounds, but that's not something a home chemist would really want to work worth. ;):D




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[*] posted on 11-7-2006 at 20:29


if coloured compounds are what you are looking for, there is a whole lot of elements which you can choose from



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[*] posted on 12-7-2006 at 00:24


Quote:
if coloured compounds are what you are looking for, there is a whole lot of elements which you can choose from

Praseodym, I agree with you that colors can be found with many other elements, and also with many organic compounds. What is striking me, however, is the fact that a SINGLE element can produce so many colors, using only the very simple reagents dilute ammonia, dilute hydrochloric acid and sodium sulfite.

Quote:
This one I can't quite follow chemically. What's going on? Is your starting material IrCl6(2-), or "IrCl4"? Why the quotation marks?

My starting material is "IrCl4" in this experiment. Just a solution of "IrCl4", which was allowed to stand for some time.
I use "IrCl4", instead of IrCl4, because I do not think that IrCl4 exists. It is sold to me as IrCl4, but the seller only knows that this compound contains Ir and Cl in a 1 : 4 molar ratio and no other elements are involved. I have read multiple textbooks and they all write that IrCl4 does not exist or at best exists as an ill-defined, very elusive, badly specified compound, whose existence is doubtful. I think that my "IrCl4" is some complicated mix of IrCl6(2-), IrCl3, Cl2, and positively charged Ir(3+) complexes. It has a smell of Cl2. So, one only can say that its average empirical formula is close to IrCl4, but the compound definitely is not a well-defined iridium (IV) chloride, otherwise the textbooks would not state that IrCl4 does not exist or is very doubtful.




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[*] posted on 12-7-2006 at 09:53


Quote:
Originally posted by woelen
$20 per gram, which is really cheap for iridium compounds


Unfortunately, it's still far more than it should cost. At the time of posting, iridium was trading for $384 per troy ounce (i.e. $12.35 per gram). Therefore, a gram of pure IrCl4 will contain $7.10 worth of iridium.

Of course synthesis will add to the cost, but I very much doubt that it justifies a $12.90 per gram markup. I don't know how IrCl4 is prepared, but IrCl3 is prepared by passing chlorine over the metal at 300°C to 400°C – a very simple process in industrial terms.

Unfortunately, this goes for all precious-metal salts. Most suppliers sell silver nitrate for over twice the price it should be.




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[*] posted on 12-7-2006 at 14:15


I know that this still is a high price, but it is MUCH lower than the price, asked by the real chemical supply houses (which do not sell to individuals anyways).

The same supplier sells AgNO3 for $160 or something like that per pound. Seems a fair price to me, and that also is MUCH cheaper than the price asked by chemical supply houses.

Of course, such a seller also needs to have a living, so he also may have some $$$ per gram ;). Otherwise this seller will be no-more after a few months.

[Edited on 12-7-06 by woelen]




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[*] posted on 14-7-2006 at 14:45


Quote:
Originally posted by Jdurg
I'm pretty certain that plutonium can also form a great deal of colored compounds, but that's not something a home chemist would really want to work worth. ;):D


My impression was that plutonium can form a number of compounds and complexes that are blues, greens, and purples in very nice, pastel shades.

Come now, with a very dilute solution, how harmful could it be, eh? :P
Now the problem is acquiring the principal reagent…
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[*] posted on 14-7-2006 at 15:12


Quote:
Originally posted by ethan_c
My impression was that plutonium can form a number of compounds and complexes that are blues, greens, and purples in very nice, pastel shades.


I don't know about blue, but I found some purty pictures of Pu solutions in red, green, purple, pink, yellow.

See
http://gotexassoccer.com/elements/094Pu/Pu.htm

Unfortunately (well, maybe not :)) these aren't from my home lab.
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[*] posted on 14-7-2006 at 15:50


Just as a word of warning I've read about Ir salts that they should be considered extremely toxic Woelen. Your site doesn't mention it do I thought it might be usefull info.

Could you explain a little more about the chemistry of Ir perhaps? Is it chemically comparable to any other element?




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[*] posted on 14-7-2006 at 16:18


Quote:
Originally posted by pantone159
Quote:
Originally posted by ethan_c
My impression was that plutonium can form a number of compounds and complexes that are blues, greens, and purples in very nice, pastel shades.


I don't know about blue, but I found some purty pictures of Pu solutions in red, green, purple, pink, yellow.

See
http://gotexassoccer.com/elements/094Pu/Pu.htm

Unfortunately (well, maybe not :)) these aren't from my home lab.


Oh wow, those are some beautiful solutions! I would pay iridium prices to pick up some of that.
Actually, I've got a nice Be hemisphere (old missile part, you know the deal), a strong alpha source, and a nice-sized puck of a particular depleted metal (the latter two safely in lots and lots of lead foil, as has been said). Nothing is actually STOPPING me from making my own…I would first make a lead pig-type container for the alpha source, and pick up a whole bunch of borated paraffin to make a little enclosure with (and maybe smear all over myself, couldn't hurt)………it'd be like that radioactive boyscout fiasco all over again! Woohoo!
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[*] posted on 14-7-2006 at 19:24


The problem is the refining of the produced plutonium. In addition to having to wait for a very long time for it to be made, you then have to separate it from the other products of the reactions and avoid any problems with doing that. Too much of a pain. It would be neat to have a tiny little button of Pu metal though. :D



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[*] posted on 14-7-2006 at 19:45


For that to be anything like safe to possess, it would have to be Pu-244, the longest-lived isotope (½ life 82 million years), rather than the Pu-239 most commonly obtained from spent enriched U fuel rods, which has a much shorter half-life.
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[*] posted on 14-7-2006 at 19:50


Quote:
Originally posted by ethan_c
Actually, I've got a nice Be hemisphere (old missile part, you know the deal), a strong alpha source, and a nice-sized puck of a particular depleted metal (the latter two safely in lots and lots of lead foil, as has been said). Nothing is actually STOPPING me from making my own…


I hope you aren't in a hurry :)

Let's say you wanted to get 0.1 g of Pu, that might be enough to make some purty solutions. That's 2.52e+20 atoms. You didn't describe your alpha source, so let's say you have 1 kg of Ra-226. :o Ra-226 emits 3.7e+10 alphas/sec/gram. These alphas may release neutrons when they hit Be, but not all of them do, in fact, just 1.2e-4 of them do, according to one source I have.

Combine all those factors, and it will take you 1800 years to get your 0.1 g. That is assuming perfect efficiency - ALL the alphas hit Be nuclei, and ALL the emitted neutrons are absorbed by your U-238 and turn into Pu-239. Since you won't have perfect efficiency, you'll have to wait even longer than that.

And if you have to settle for something less 'hot' than a kilo of pure radium, even longer still.

So, patience is required. :D
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[*] posted on 14-7-2006 at 19:55


Quote:
Originally posted by JohnWW
For that to be anything like safe to possess, it would have to be Pu-244, the longest-lived isotope (½ life 82 million years), rather than the Pu-239 most commonly obtained from spent enriched U fuel rods, which has a much shorter half-life.


My main concern would be neutron emission. If I had such a button, it would be sealed in a glass ampule, for sure. Pu-239 is an alpha emitter, and those have such a short range, that none of them would escape the ampule. There would be some gammas, but I don't think Pu-239 is a strong gamma emitter, so that probably isn't a problem. Most of the neutrons come from spontaneous fission of Pu-240 which is present as an impurity, so the safest stuff would be Pu with as little Pu-240 present as possible.

That is the description of super-weapons-grade plutonium. :)
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[*] posted on 15-7-2006 at 05:44


Quote:
Originally posted by pantone159
Quote:
Originally posted by JohnWW
For that to be anything like safe to possess, it would have to be Pu-244, the longest-lived isotope (½ life 82 million years), rather than the Pu-239 most commonly obtained from spent enriched U fuel rods, which has a much shorter half-life.


My main concern would be neutron emission. If I had such a button, it would be sealed in a glass ampule, for sure. Pu-239 is an alpha emitter, and those have such a short range, that none of them would escape the ampule. There would be some gammas, but I don't think Pu-239 is a strong gamma emitter, so that probably isn't a problem. Most of the neutrons come from spontaneous fission of Pu-240 which is present as an impurity, so the safest stuff would be Pu with as little Pu-240 present as possible.

That is the description of super-weapons-grade plutonium. :)


One only really has to be worried about alpha emitters if they are inside you. However, one speck of something like the Pu we're talking about in a lung, and you might as well start saving up for the chemotherapy you're going to begin in a few years.

EDIT: I apologize to Mr. Oelen. I have catastrophically derailed his relevant and interesting thread. Woops.

[Edited on 15-7-2006 by ethan_c]
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