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Author: Subject: Exotic Primaries - Complex Salts
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[*] posted on 13-8-2008 at 02:19


Here's a funny dream I had last night; dunno if it has any connection to reality tough:

I dreamt I found a way to make anhydrous TACN which dets. readily from a hammerblow. The instructions I dreamed up were: Precipate deep-blue/violett crystals of TACN from a conc. solution by addition of equal volume of ethanol. Wash the crystals with ethanol twice, then press dry between filter paper. Wash them two more times with dry acetone. Volume will shrink considerably due to loss of water of crystallization (in my dream I had the impression that just 1/5 of the initial volume was left). Remove acetone by carefully pressing between filter paper. Finally dry in a steam of warm air. At this point, most of the water has been removed and the crystals will no longer decompose on contact with air. Now its ready for the final drying: Put them in a decissator over CaCl2 for a few days.
I dreamt that by following these instructions, I obtained a light blue powder. In my dream I even took a hammer and it went BOOM readily. The hydrated deep-blue/violett crystals never did, plus they turned into green junk within a few days.

If this dream was true, TACN was not useless at all :cool:
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[*] posted on 13-11-2008 at 07:55


I recently came across a rather interesting compund: [Ag(DMSO)<sub>2</sub>] [ ClO<sub>4</sub>].
Method:
Excess DMSO is added to a solution of AgClO<sub>4</sub> in Acetone, solvent removal by vacuum.
Interestingly, the method uses Me<sub>2</sub>CO as a short form for acetone (for IUPAC fans: propanone).

Unfortunately, the only thing I can say about this compound is that it has explosive properties.
I assume a similar procedure would probably work for Cu, too. Which, of course, would be good in terms of money ^^


Next point, I am a bit confused with the fact that perchlorate complexes seem to be relatively easy to prepare. After all, ClO<sub>4</sub>- is quite a weak ligand, definitely weaker than water. Which seems to be a slight contradiction, considering most (or all) syntheses of perchlorate complexes were carried out in aqueous medium. EDIT: Mistake corrected, ClO<sub>4</sub>- not as a ligand.

Finally, I wondered how the persulfate ion would act as a ligand.. my (rather uneducated) guess would be bridging chelating, possibly creating rather interesting macromolecular structures -> oligomer (?). I am somewhat surprised with the stability of tetraaminecopper(II)persulfate, of course it does decompose/ react under heat, however, considering we are talking about a (macromolecular) compound containing peroxide bridges and plenty of oxygen... I would assume that a number of different compounds can be formed depending on reaction conditions. Or does anyone actually have crystallographic (or any other sorts of proper analytical) data of tetraaminecopper(II)persulfate?

EDIT:
Oh, and here is an idea to madmen and potential terr.. err.. don't wanna wake a sleeping dog there. Anyways, idea:
Complex Ni not only to it's tetracarbonyl (know as "liquid of death"), but add some nice groups with oxidating power. No idea if it works, should that be the case, whoever makes it will probably also succeed in killing himself or herself by preparing the compound. Like the guy who first synthesised Ni(0)CO<sub>4</sub> and almost killed himself.

[Edited on 13-11-2008 by natriumperoxid]

[Edited on 13-11-2008 by natriumperoxid]




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[*] posted on 13-11-2008 at 08:31


Perchlorate isn't a ligand. It is an (almost "perfect") anion. You should recognize the compound isn't even written to show ClO4- as a ligand but rather as the anion it is.

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[*] posted on 13-11-2008 at 09:16


Bugger >.< That would mean another: So obvious I don't realise it. Cheers for the information.
However, ClO<sub>4</sub>- can act as a ligand, although described as non-coordinating. An example are pentaaminecobalt complexes in which the sixth coordination position can be occupied by chlorate and perchlorate (amongst others). (Duval, Ann. Chim., 1932).
Ok, this question probably sounds rather... uninteligent to every chemist, think it's worth a try anyway:
The position of the molecules/ ions involved in complexes should have an effect on the reactivity or rather stability. What I am trying to get at: Would ClO<sub>4</sub>- as a ligand be closer to adjacent ligands such as NH<sub>3</sub>, and would the overall compund therefore be more unstable (less energy needed to make the oxidising and reducing parts react)? Same for the NO<sub>3</sub> - anion...

[Edited on 13-11-2008 by natriumperoxid]




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[*] posted on 13-11-2008 at 11:12


Ligand formation is about electron donation, although not usually as strongly as in chemical bonding. As such, more highly charged ions, like carboxylate (one electron shared between two oxygens) and to some extent carbonate (2/3) make reasonable ligands. Nitrate, chlorate and perchlorate, on the other hand, have much more electronegative components (N, Cl, O) and less charge per atom. So they have less charge to donate. I shouldn't be so arrogant as to say perchlorate isn't a ligand, I should correct that by saying, it simply has a very small formation constant in most cases.

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[*] posted on 24-1-2009 at 05:48


I once formed TANN under ethanol absolute.
By adding a stream of dry NH3 gass trough a solution of nickel nitrate (dehydrated as far as possible).
Even though I used everything within my reach to make it under as dry conditions as possible it was still wet.
IMO quite useless stuff.

Has anyone tried making chromate of iodate complexes of tertaaminenickel or other metals?
And any ideas of forming these substances?

I am very skeptical about the potentials of these compounds but still they seem very attractive and interesting.
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[*] posted on 24-1-2009 at 11:20


There seems to be a polymerized form of Ni chromate which is insoluble; however Ni2+ gives no precipate with chromate, so the normal chromate is soluble. You can make it from CrO3 and Ni(OH)2. Maybe it gives a precipate with NH3.

The iodate ammines have no or poor energetic properties. The chromate ammines however mightbe worth a try. According to Urbanski, ammonium chromate is explosive and more sensitive to impact than the dichromate.
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[*] posted on 11-2-2009 at 14:25


Are you sure you added a chromate solution to the Ni(II) solution, and not a dichromate solution? Many chromates are insoluble, while dichromates are not. For example, silver and barium chromate hardly dissolve, while silver dichromate and barium dichromate are much more soluble.
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[*] posted on 1-3-2009 at 14:02


Quote:
Originally posted by natriumperoxid I am somewhat surprised with the stability of tetraaminecopper(II)persulfate, of course it does decompose/ react under heat, however, considering we are talking about a (macromolecular) compound containing peroxide bridges and plenty of oxygen...


Hello. Do you have any data about the stability and power (perhaps VoD) of the persulfate?

It looks quite interesting.....

PS: I think it will decompose after a while, but can you increase the chemical half-life by keeping it completely dry?

[Edited on 1-3-2009 by Jetto]
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[*] posted on 14-4-2009 at 06:05


Btw, is it possible to use cartridges as blasting cap? I got some ammunition containing cordite and a percussion cap.

That's basically quite similar to a commercial detonator.
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[*] posted on 16-4-2009 at 16:01


I just had a couple of drinks , this makes me think better/ more free (which ever one you prefer) :)
Would it be possible to ppt the persulfate salt of hexamethylenetetramine, or would it be destroyed due to oxidation.

Writing this makes me wonder if there would be other organics which could get persulfate groups attached to them ( excluding TACpersulfate compounds because i know it can be done)
It is a strong oxidizer so a lot of organic molecules would be attacked by it and probably not survive.
but still some products of the oxidation might be capable of getting a persulfate group attached to them.
Iam just speculating, maybe this gives anyone inspiration.


Anyone any ideas, is this going anywhere or I am being naive/drunk :P.

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[Edited on 17-4-2009 by User]




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[*] posted on 16-4-2009 at 22:52


[Cu(NH3)4]S2O8 decomposes @RT within a couple of weeks, no matter how well it is dried. You could probably store it for month in a freezer tough. But what for? Its a nice curiosity but of no practical interest. Btw there are various decomposition products which show weak energetic properties as well. Initially the stuff is dark violett but as it decomposes it takes on a much lighter blue color. At this stage it still burns, comparable to ammonium dichromate. It will decomposes further to the point where it can be hardly ignited anymore. A funny thing happens when you smack this mostly decomposed complex: It turns green where you hit it. Its quite fascinating to see a stuff change color from impact.

Also if you put the dry crystals in a glass dish and gently warm in air, they start jumping around like fleas. Probably due to the crystal structure being broken as it decomposes.

Theres also the zinc persulphate which is said to be explosive.
This might be an interesting target. Copper doesnt seem to be terribly compatible with O-O groups - copper peroxide for example decomposes quickly. Zinc peroxide however is perfectly stable even @RT. ZnO2 can be made from boiling ZnO in H2O2 which shows that its very stable towards thermal decomposition.

I had no luck with nickel and cobalt persulphates however. These metals are simply oxidized and some insoluble oxide/hydroxide/peroxide crap precipates.

If barium persulphate could be made somehow then one could form a lot of other interesting persulphates via double decomposition from the corresponding sulphates.

If you want a really nice energetic compound containing lots of peroxide, then go for the tetraperoxochromates :) Woelen has a nice writeup of its synthesis.
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[*] posted on 17-4-2009 at 00:45


Thanks for your reply.
I am fully aware that al lot of these substances have no practical use what so ever, call it interest.

Zinc persulfate might be a nice try indeed, any special procedures or would a simply fall out of solution? When for example zinc nitrate is combined with sodium persulfate.

Bariumpersulfate solute able ?
(would it form hydrated salts ?? )
So using for example bariumnitrate/oxide might do the trick.

But would there be organics that can be combined with persulfate groups?

[Edited on 17-4-2009 by User]




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[*] posted on 17-4-2009 at 01:03


[Cu(NH3)4]S2O8 loses NH3 in air and a part of its active oxygen - Barbieri, Calzolari (Z. anorg. Ch. 71 [1911] 347-55, 351). That ref should also have something pertaining to other persulfate complexes (zinc persulfate complexes) since it's about divalent metals.
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[*] posted on 17-4-2009 at 02:26


Quote: Originally posted by User  

Zinc persulfate might be a nice try indeed, any special procedures or would a simply fall out of solution? When for example zinc nitrate is combined with sodium persulfate.


No we're not talking about zinc persulphate but its ammine complex. Add ammonia to a solution of any soluble zinc salt until the precipated hydroxide redissolves due to complexation. Then add a solution of sodium/ammonium persulphate. Cool/add ethanol to precipate.

I havent tried this but its the "standard procedure" to make such compounds.

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[*] posted on 17-4-2009 at 04:04


Yes i know, but that is the thing amine complexes are often very hygroscopic or hydrated salts, i've made a few, know how the system works.
As i said in the first poist, obviously i wasn't clear.
I mean other salts,anything but the amine complexes.
This tread goes a lot about them , i know but there is a lot more than that.




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[*] posted on 17-4-2009 at 04:23


Well here's the general formula of such a complex: [MLx]Cy
You have the metal M, the ligand L and the cation C. Each of them you can vary. So what exactly do you mean? Do you want to try other ligands together with persulphate as the cation?

The only other ligands that gave interesting compounds in my experiments were ethylenediamine and hydrazine. The copper complex [Cu(en)2](ClO4)2 gives very beautiful crystal needles, also quite impact sensitive. Nickel hydrazine dinitrate is a relatively harmless substance, not friction sensitive and quite insensitive to impact too. When heated in open air it just deflagrates, however very little confinement is needed to make it DDT.

Hydroxylamine also looks intriguing, altough the complexes might be very unstable.
Hydrazine gives some exceedingly dangerous complexes, nickel-hydrazine-perchlorate has mutilated the poor guy who "discovered" it.

There is said to be an acetato-complex with Cr(III) which can be precipated as the perchlorate. This one surely is very explosive. If you have acess to perchloric acid and any soluble Cr(III) salt, this one would be an interesting experiment.

[Edited on 17-4-2009 by Taoiseach]
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[*] posted on 7-8-2009 at 15:19


Quote: Originally posted by Formatik  
[Cu(NH3)4]S2O8 loses NH3 in air and a part of its active oxygen - Barbieri, Calzolari (Z. anorg. Ch. 71 [1911] 347-55, 351). That ref should also have something pertaining to other persulfate complexes (zinc persulfate complexes) since it's about divalent metals.


I've seen this reference and they have prepared the ammines: ZnS2O8.4NH3, CdS2O8.6NH3, NiS2O8.6NH3 and CuS2O8.4NH3. They also prepared some compounds of pyridine and hexamethylenetetramine with metal persulfates, e.g. CuS2O8.4C5H5N (blue-violet cryst.), there are no energetic properties described of those. But they say all of the ammine salts explode on strong heating or by impact. They also decompose losing NH3 more or less rapidly. Barbieri and Calzolari also say the Cu ammine salt is more air stable than Ni or Zn ammine salts.
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[*] posted on 15-8-2009 at 00:22
Persulfate Electrolysis


I am aware that this is a little off topic, so I shall keep it as brief as possible:

It's not that Persulfates are impossible to come buy, but in keeping with the ethos of this home-brew is best.

I became intrigued whilst reading 'A Course in Inorganic Preparations' by Henderson and Fernelius' pages 97/8 from the SciMad library.

The basis seems to be electrolysing a cold saturated solution of Potassium Hydrogen Sulfate with a high current and little allowance for mixing. With Pt electrodes.

Does anyone have personal insight into this electrolysis? Specifically, has anyone achieved a yield with graphite and stainless steel?

Tr

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[*] posted on 17-8-2009 at 05:33
Diazidodiamminocopper(II)


was prepared according to

CuSO4*5H2O + 2NH3 + 2NaN3 ---> [(NH3)2(N3)2Cu] + Na2SO4 + 5H2O

NH3 was added to a solution of CuSO4 until the precipate of Cu(OH)2 redissolved. A saturated warm solution of NaN3 was added. Upon cooling and addition of an equal volume of ethanol, a nice crop of [(NH3)2(N3)2Cu] precipated.

Green-blue glistering crystals. Unlike copper azide these are not friction sensitive. Explodes upon flame contact.

Obviously the compound is oxygene-deficient, so addition of a strong oxidizer should increase its power.
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[*] posted on 17-8-2009 at 11:41


I tried the same experiment but I failed. I did the following:

200 mg of CuSO4.5H2O was dissolved in a little amount of warm water
12% ammonia was added dropwise until the precipitate of Cu(OH)2 completely redissolved. One extra drop of 12% ammonia was added to be sure that really no precipitate remained.
104 mg of NaN3 was dissolved in some water and added to the deep blue solution.

The result of the final step was that the liquid became more greenish blue, but on shaking and mixing all of the liquid, it became deep blue again.
10 mg more of sodium azide were added with a drop of water. Now there is slight excess of sodium azide.

My experiment is 0.8 mmol of CuSO4.5H2O, appr. 1.7 mmol of sodium azide (slight excess) and quite a large excess of NH3.

To the deep blue somewhat cooled down solution, I added an equal volume of ethanol. This resulted in formation of a deep blue/purple precipitate. Formation of the precipitate was not at once, it took a minute or so before all material had formed as a precipitate. I let the precipitate settle at the bottom and a green/cyan solution was above the precipitate.

The solution was decanted and the precipitate was rinsed with more 96% ethanol (denatured, but distilled once such that it is purely volatile, no oily compounds in it) until the liquid, running from the solid, was colorless.

The solid now is dark blue/purple. This was spread out on a watch glass and allowed to dry at a dry place of appr. 40 C. The final result is a beautiful blue/indigo solid, consisting of small granules (max. size less than 1 mm).

I took a small amount of the solid and put it on the tip of a small screw driver, which I kept in the flame of an alcohol burner. The material does not explode, but it seems to 'evaporate' without any noise, leaving behind a small globule of brown material. The 'evaporation' looks as if almost all of the material is converted to gaseous products, but there is no bang and no visible flame.

Could you tell something about the relative quantities you used? Do I need to use a large excess of NaN3? I just used a little bit more than the stoichiometrically needed amount.

If you provide a little more details about your experiment, then I'll try again and might make a write-up on this. It is a nice little and interesting experiment (provided it succeeds and can be reproduced by others, such that it is worth to write about).

[Edited on 17-8-09 by woelen]




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[*] posted on 17-8-2009 at 21:55
hmmm


Some indications for preparation (below) say to add 1/2 vol. methanol or so much dilute acetic acid until the soln turns from blue to green. Hot H2O dec. the compound to basic azides (which are a lot less brisant than Cu(N3)2), this is also what Cu(N3)2 on standing in air converts to after 2 months. Strange nothing of air stability, etc. is mentioned.

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[*] posted on 17-8-2009 at 23:12


Thx for the paper. I wish I had it before my experiments :)

I didnt find a prep so I just tried it 4 times until "it looked right" :) First thing I did was I made some Cu(N3)2 and added NH3 to it. The dried material still pretty much looked like plain Cu(N3)2, most of it was dark brown with a few specks of green/blue in it so I guess the insoluble Cu(N3)2 doesnt fix the NH3 easily. I didnt keep any of the stuff as it seemed quite sensitive. Dry copper azide is quite scary IMHO especially when u have to use gram quanitites because the lab scale is broken and all u have is a digital bowl kitchen scale precise to 1g at most :mad:

When I added NaN3 to a solution of CuSO4 in ammonia I also observed that some violet stuff precipates first. It didn't burn after drying so I figured it must be [Cu(NH3)4]SO4 and discarded of it. The clear filtrate then deposited the green crystals after further standing.

I wasted quite some NaN3 because the solution was too dilute. You must use very concentrated solutions, otherwise little or no crystals precipate even after adding ethanol.

Attached is a photo of the green crystals. I kept a small sample of it. For an explosive azide it is amazingly insensitive.

I wonder if there is other nice azide complexes to be discovered, preferably a perchlorate with NH3 and N3 ligands :)

Btw woelen there is a stable modification of [Cr(NH3)3(O2)2] which is said to keep well for month. Its described in "Zur Kenntniss der Perchromate" by K.A.Hofmann and J.Hiendlmaier here's the preparation:

20% ammonia is saturated with amminium bichromate, after cooling it is saturated with gaseous NH3. After filtering, 100cc of the filtrate are added slowly to 11cc 30% H2O2 @0°C. The mixture is kept in the cold for 12hours. The crystals are then washed with 20% ammonia of 15°C to remove impurities (ammonium perchromate?).

This "chromium tetroxide triammine" is a modification of the brown stuff you described on your webpage and is said to be much more stable.

Now the interesting part: It fixed CN- quite easily, replacing all NH3 ligands.

I wonder if it would also fix N3- and form azido-peroxo-chromate :cool:

CuN3.jpg - 77kB
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[*] posted on 3-9-2009 at 07:15
more ammines


According to T. Klobb (Compt.rend. 103, 384) various metal permanganates is said to give relatively stable compounds with ammonia (I don't know how far I'll believe this). Like the silver salt described below. Copper, cadmium, nickel, zinc and magnesium treated in a similar fashion as below yield analogous compounds. A lot of the complexes are sol. in water, decomposing. Klobb, Bull. soc. (3) 3, 509 should have more,

Diamminesilver permanganate [Ag(NH3)2]MnO4: made from sat. KMnO4 soln. with NH3 aq., then add aq. AgNO3 (1:10 in water). The resulting ppt. is collected on guncotton, filtered off, washed with ice water, dried around CaO which is mixed with some NH4Cl. It's a violet powder, rhombic plates under microscope, little sol. in cold, more in hot water, which dec. gradually, liberating NH3, and changes into an insoluble powder. It explodes under a hammer blow. PATR says it dec. slowly on standing. It was also briefly mentioned on page 3 of this thread.
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[*] posted on 4-9-2009 at 17:52
lead chlorate complexes


There is a patent which describes mixing conc. aq. basic lead chlorate with various compounds like mannite, sugar, glucose, dextrin, tannin which produces compounds that explode violently by heat or shock. But said ref. says dissolving Pb(ClO3)2 in hot glycerin, the then formed compound explodes the strongest of all those, the patent claims it detonates similar to Hg(ONC)2 and diazobenzol salts.

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