Sciencemadness Discussion Board - Exotic Primaries - Complex Salts

Exotic Primaries - Complex Salts

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jlithen - 15-11-2006 at 13:03

I just have to say one thing related to my previous post on this thread (long time ago now).
LNP (Lead nitrato phosphite) is actually very sensitive to friction...when I noticed it I got very frightened as I have handled it rather uncarefully. It say on megas site that it is "stable" of course not necessarily meaning insensitive.
How stupid of me to not immediately check the sensitivity of a substance that I have been using for years. Anyway the detos seem to work relatively well, even after months of storage, but I would like to change the LNP to something more powerful and maybe even less sensitive. Although senitivity has not been a problem in the caps.

DeAdFX - 23-12-2006 at 11:12

I tried making a complex Triethanolamine(TEA) of my own. I first made copper (II) nitrate by reacting copper sulfate and potassium nitrate in 30mL of water(1/10 mol scale). I fucked up though and added half as much KNO3 as I should have(christ this is the 2nd time I have neglected the fact that SO4 is a -2 ion). Well anywho I added in 20mL of TEA into the CuSO4/Cu(NO3)2 solution. The solution turned a dark turquiose(SP) color. After swirling the solution for a few seconds the color intensity went down a bit. I added 30mL of ethanol and a blue/turquiose perciptate formed. The perciptate is really fine. I assume its about >200 mesh easy. Right now the perciptate is drying off in the air.

I wonder if Copper (II) Nitrate would form a complex with nitrosoamines like NitrosoEthanolamine or Dinitrosoethanol amine. These complexes would probably be more energetic due better OB.

Anywho Copper(II) Nitrate TEA complex should be more stable* than the tetramine variant. By stable I mean the complex won't fume off ammonia and become useless after a few days. TEA has a much higher boiling point.

[Edited on 23-12-2006 by DeAdFX]

DeAdFX - 24-12-2006 at 13:13

UPDATE:

The complex has dried off and still has the same turquiose color. I burned a small sample of the complex. The sample turned from a turquiose color to a black color indicating the complex decomposed to copper(II) oxide. The decompisition was not violent nor was there any flame. To verify that I had a copper complex I burned the sample with some Ammonium Perchlorate in a paper cup. The resulting flame was a very pretty blue.

quicksilver - 25-12-2006 at 08:34

Copper (II) nitrate is SO hygroscopic that there would be some water in the crystals unless special efforts were taken and that would effect the outcome to a degree that it would deflagrate. Under other circumstances it may well detonate but the water would present a problem. Copper (II) nitrate has often had this as a stumbeling block. The issue has arisen with many primaries. Even lead-based materials, if water still existed within the final product, they did not respond well. Really -=DRY=- material often tests far differently. Determine if any water existed before making a final assesment.

Zinc - 15-1-2007 at 07:13

Quote:

Originally posted by quicksilver
Copper (II) nitrate is SO hygroscopic that there would be some water in the crystals unless special efforts were taken and that would effect the outcome to a degree that it would deflagrate. Under other circumstances it may well detonate but the water would present a problem.

You mean that copper (II) nitrate can detonate?
If yes do you know anything about its properties(stability,VOD and how to dehidrate it)?

quicksilver - 17-1-2007 at 06:19

No it does not. It is a nitrate salt that is very hygroscopic but that in use, it brings with it so much water that unless steps are taken to deal with the water - the compound / resultant will have that issue. Some primaries are desensitized by water to a greater degree than others., etc.

chemoleo - 17-1-2007 at 17:26

DeadFX what I really wonder is why you think TEA is an appropriate complexing agent. I mean, for one Cu ion there are two molecules of TEA, and 2x NO3. If you calculate the oxygen balance, you'll note it is appalling. No way this will ever make it into some sort of energetic material.
Try hydrazine, hydroxylamine (although I think that I tried that before, and it doesn't work-> decomposition in solution), guanidine, and other small nitrogen containing bases, these are feasible targets! Or even tetrazole, I think Nick F may have tested that.

halogen - 14-2-2007 at 12:01

Would it not be one molecule of TEA per copper atom?

DeAdFX - 15-2-2007 at 00:50

Quote:

Originally posted by chemoleo
DeadFX what I really wonder is why you think TEA is an appropriate complexing agent. I mean, for one Cu ion there are two molecules of TEA, and 2x NO3. If you calculate the oxygen balance, you'll note it is appalling. No way this will ever make it into some sort of energetic material.
Try hydrazine, hydroxylamine (although I think that I tried that before, and it doesn't work-> decomposition in solution), guanidine, and other small nitrogen containing bases, these are feasible targets! Or even tetrazole, I think Nick F may have tested that.

The reason I thought TEA would be a more suitable complexing agent is because it has a way higher boiling point than ammonia[liquid vs gas]. The TEA complex should be more stable in the sense of having a somewhat longer shelf life with a major sacrifice in power. However my experiment was a failure because I had a miscalculation so the power is yet to be determined. The perchlorate variant would probably show more promise.

The nitrogen bases that I posses are urea and hexamine[yeah yeah big molecule..]. Technically I could also posess ammonia, guanidine and methyl amine if I weren't so god damn lazy/bogged with work. I also have azide however my time is limited in my home lab so I can't synthesize tetrazole and even then my knowledge in high explosives is rather limited.

Even if the compound is only capable of a simple burn it will still be of practical use for me atleast. Somewhat off topic but I am working on a blue star composition involving Ammonium Nitrate( I hear it burns a lot cooler than Perchlorates] as the primary oxidizer. So far I have very little success.

quicksilver - 15-2-2007 at 06:46

OT remark:

Blue is a tough color to get right. However I have had some success. The problem is that metal fuels (other than copper) will drown it out very fast. Ammonium nitrate could work just fine if you used something like: ammonium nirate 33% nickle or cupric nitrate 20% copper powder (325 mesh) 13% chlorine doner 20%, burn rate enhanser 7% binder 6% - The burn rate enhancer is needed due to ammonium nirate's slower burn and less O than the perchlorate. It's a problem and a good thing as well.

[Edited on 15-2-2007 by quicksilver]

nitro-genes - 13-3-2007 at 17:26

Haven't seen any mention of nitrato/nitrito-metal complexes for energetic complex salts in this thread (or I have overlooked them

). Some would be pretty promissing, like the hexanitritocobaltate(III) ion. The sodium salt is one of the few well soluble salts, while it readily precipitates with K+ or even NH4+...

Forming a salt with methylamine, urea, hexamine, etc could give interesting results, no idea what it's properties, or how stable they would be, as most NO or NO2 ligand metal complexes are not very stable...

Are there even any known nitrito/nitrato complexes that have been reported to have primary explosive properties? I seem to remember this nitrato-Hg(II) aldehyde complex, though it was a rather feeble explosive IIRC...

[Edited on by nitro-genes]

Rosco Bodine - 13-3-2007 at 20:17

Joeychemist got this file from solo ,
and hasn't posted it here so I will .

Attachment: Explosive Properties of Metal Ammines.pdf (273kB)
This file has been downloaded 3374 times

nitro-genes - 13-3-2007 at 22:12

Thanks for posting that article!

Ammonium hexanitrocobaltate(III) goes out the window as it seems concerning it's potential as a primary. Hydrazine hexanitrocobalate still seems pretty spicy though, if it would exist that is...

Most of the used oxidizers seem to stable to accellerate fast enough to behave like a primary, like the perchlorates or nitrates. On the other hand the chlorates are capable of doing so, but are not very storage stable. For that matter would the peroxo chromates Woelen posted about be a good starting point, since it was pretty obvious that it is a very powerfull oxizer, but seemingly stable in storage...

[Edited on by nitro-genes]

Ballermatz - 6-8-2007 at 11:59

[Edited on 24-12-2007 by Ballermatz]

12AX7 - 6-8-2007 at 13:21

Nitrito-?

Ballermatz - 6-8-2007 at 13:46

[Edited on 24-12-2007 by Ballermatz]

Nerro - 6-8-2007 at 14:45

Quote:

Originally posted by 12AX7
Nitrito-?

nitrite as a ligand

12AX7 - 6-8-2007 at 15:50

Yes; he spelled "nitrato" is all. I wasn't aware NO3- formed ligands, but I've heard of the NO2 cobalt complex.

Tim

Nerro - 6-8-2007 at 17:03

Why shouldn't nitrate form ligands? It does, it even forms reasonably stable eta-2 ligands (bridging ligands with 2 hapticity).

JohnWW - 7-8-2007 at 01:27

Getting the anions of strong oxy-acids like NO3-, SO4--, ClO4-, BrO4-, IO4-, and similar, and of strong fluoro-acids like PF6-, SbF6-, BF4-, to act as ligands is very difficult, if not impossible, because of their high symmetry and steric crowding reasons, and, in the case of the oxy-acids, their resonance-stabilization which results in the negative charge being evenly distributed over the oxygens.

Nerro - 7-8-2007 at 04:33

Which makes it impractical for an amateur chemist but not impossible on the whole.

woelen - 7-8-2007 at 10:45

I'm almost 100% sure that the yellow material, mentioned here is ammonium hexanitritocobaltate (III). So, nitrite as ligand and not nitrate. I have the sodium salt of this complex ion and indeed it forms beautul precipitates with ammonium salts and even better results are obtained with potassium salts. I know of nitrato complexes, e.g. UO2(NO3)2 and also some cerium compounds, such as (NH4)2[Ce(NO3)6].

Nerro - 7-8-2007 at 12:04

Are those the nitrito complexes (-O-N=O) or the nitro complexes (with the bond being made by the lone pair on the N)?

And how feasible is it to make a complex like ferrocene but with N5-? (the structure of which being similar to that of the cyclopentadienyl ion). Is anything known of such complexes and their stability?

-Edit-
The complex would be called bis-pentazolido iron(II).

Quote:

source
Abstract : The pentazole anion has been generated from para-hydroxyphenylpentazole and identified by electrospray ionization mass spectrometry. Whereas at low collision voltages the para-phenoxypentazole anion undergoes stepwise N2 elimination generating the corresponding azide and nitrene, at high collision voltages the N5(-) anion is formed. Fragmentation of the pentazole anion produces the N3(-) anion as the principal negative ion. These experiments provide the first experimental proof for the existence of the pentazole anion. They also demonstrate that under suitable reaction conditions the C-N bond in a phenylpentazole can selectively be broken with conservation of the pentazole ring, thus providing a potential synthetic route to the pentazole anion.

Doesn't sound like a very usefull synthesis... I remember synthesizing the cyclopentadienyl ion by letting a saturated solution of KOH with excess KOH at the bottom react with C5H6, why shouldn't something similar work on HN5? A stronger base will probably be needed...

Could an oxidizing agent be applied to synthesize N10 (+2H+ and two electrons)? Sounds like a pretty dangerous substance to me but perhaps at -80°C it might be stable enough?

And I also read that phenylpentazole is much more stable than just pentazole. Could para-dipentazolebenzene be synthesized somehow? And how is phenylpentazole synthesized anyway?

[Edited on 7-8-2007 by Nerro]

[Edited on 7-8-2007 by Nerro]

quicksilver - 8-8-2007 at 16:51

Hey Nerro: if find out more stuff about ferrocenes please post 'em! I'm interested also and have had little luck......

The_Davster - 8-8-2007 at 17:38

I have done minor looking into of pentazoles
Phenylpentazole is formed by the reaction of benzenediazonium chloride with sodium azide, this is analogous to the making of tetrazoles by the reaction of cyanides, 5-phenyl -1H-tetrazole being formed from phenylacetonitrile and sodium azide.

I have seen a paper somewhere about the free pentazoles, it was posted on this site years ago. I forget one very important thing, whether it was theoretical, or had been done before. It was on a 'pentazolium pentazolate' compound which could cyclize into a nitrogen flavour of buckyballs.

1H,5H-tetrazole can be formed from anhydrous HCN and HN3, problem is finding a nitrogen analogue of HCN that would be substitutable here. Acetylides/acetylene(forget which exactly) also react with HN3/NaN3 making the triazoles, but no such easy lab analogues exist for N. I wonder if under extreme conditions nitrogen gas would react with HN3? Some sort of exotic excitation would likely be necessary,and under pressure.

EDIT:

Quote:

And how feasible is it to make a complex like ferrocene but with N5-? (the structure of which being similar to that of the cyclopentadienyl ion). Is anything known of such complexes and their stability?

There is something close, nitrotetrazolate CN4NO2-, ion would seem to be similar to a substituted cyclopentadiene, but its complexes are not eta5, all I have seen is eta1. A tetrazolate might make a ferrocene type compound, but I have seen nothing on it as a ligand.

[Edited on 8-8-2007 by The_Davster]

The_Davster - 9-8-2007 at 17:56

This is going to be a multi-post. I used scifinder to look up pentazoles. So attached is some toilet reading

A few papers.

[Edited on 9-8-2007 by The_Davster]

Attachment: pentazole.pdf (110kB)
This file has been downloaded 2018 times

The_Davster - 9-8-2007 at 17:59

2

Attachment: pentazole2.pdf (137kB)
This file has been downloaded 2941 times

The_Davster - 9-8-2007 at 18:02

3

Attachment: pentazole3.pdf (456kB)
This file has been downloaded 4829 times

Nerro - 10-8-2007 at 03:51

One question micht be relevant in all this, if a phenyl-ring adds stability to the pentazolyl-group, why remove it? How good would Ph-N5 be as a ligand? This appears not to have been researched...

Also, it seems like a molecule such as antracene might be able to stabilize it even better.

Thanks for the papers btw, very nice.

[Edited on 10-8-2007 by Nerro]

ssdd - 9-9-2007 at 14:18

OK so I was going to attempt to make some Nickel Hydrazine Nitrate.

But before I do so I am trying to gather as much info as possible, and I was wondering if anyone knows the structure of NHN.

Or better yet, if anyone has a program that can generate these structures.

Any info would be great.

-ssdd

quicksilver - 10-9-2007 at 05:52

Get most any version of ChemOffice. You can have hours of fun with it and it also has 3D modeling effects, etc, etc. One thing that I like to do is import the structure of items I found in the Merck Index (v13+) program (it allows for this specifically) and model what I am working with. then you can add to it from what your project is at the time. You save the basic structure for later, etc. - It's a lot of fun. You DON'T have to get the very latest one. I did and it's like getting the latest MS product.....bell & whistles. Most recent versions are just fine.

Ballermatz - 3-10-2007 at 02:50

[Edited on 24-12-2007 by Ballermatz]

Ballermatz - 21-10-2007 at 13:04

[Edited on 24-12-2007 by Ballermatz]

woelen - 21-10-2007 at 22:34

Quote:

woelen has described this compound in detail on his homepage already, including the fact that it explodes upon heating - however one important detail that has not been mentioned so far is that it explodes violently upon impact too! I'd say it is even more sensitive than TACP. Since there is nothing in this complex to be oxidized, the explosion must be caused by a sudden loss of its oxygene. Thus mixing with anything combustible will certainly increase its explosive powder.

I have 15 grams or so of this compound around for approximately 18 years already

. Do you think it is wise to destroy it and not keep it around any longer? I did not know it is impact sensitive too. I'll do some tests myself and see how sensitive it actually is. After all those years of storage, the compound still is very energetic, so this at least is very stable, when stored in a dry and air-tight container.

[Edited on 22-10-07 by woelen]

Ballermatz - 24-10-2007 at 11:13

[Edited on 24-12-2007 by Ballermatz]

12AX7 - 24-10-2007 at 18:11

Nah, it could just be unusually sensitive to thermal anisotropy and vacuum fluctuation.

Tim

Nixie - 20-12-2007 at 05:57

Quote:

Originally posted by Ballermatz
50mg give such a loud report that it is paineful in the ears.

Man, I hope you realize that permanent, cumulative hearing damage occurs from levels even below the threshold of pain.

chemoleo - 21-12-2007 at 17:06

Well, except it's going to be too unstable probably ... and OB wise...is there any?

Very nice work above ... do you have pictures of the above compounds?
With your resources, perhaps it's worth testing other transition metals, such as Ni, Cd, Cu etc?
For instance, copper hydrazine, guanidine, imidazole, pyridine, dinitrophenylhydrazine perchlorate, and the nitrates etc etc
We need a comprehensive review of all these things!

Ballermatz - 22-12-2007 at 04:17

[Edited on 24-12-2007 by Ballermatz]

Ballermatz - 22-12-2007 at 04:25

Merry christmas to all!

[Edited on 24-12-2007 by Ballermatz]

Bert - 22-12-2007 at 08:17

Of course I admire the photos- But feel some misgivings about storing energetics known or suspected of being primary explosives in glass!

jlithen - 16-1-2008 at 06:00

Hi!

I must comment the sensitivity of NHN.
I have prepared it many times in about 10-20g batches and it is extremely difficult to filter. After drying and at some point grinding it (do this in small ammounts if it is already dry) it is a very flammable powder that does not seem to detonate easily unless confined.
I have tested friction sensitivity by means of a mortal and pestle. It is very well possible to make it detonate. It is deffinitely much more sensitive than TNT that I think someone compared it with. Anyway I have found it to me much less sensitive than e.g. PbN3. Maybe in the order of ETN or MHN.
I have also stored it for about 2 years at some point and haven't noticed any change. When i tried to make the same salt of Cobalt it was a bit weaker but needed less confinement to detonate in 10g quantities. Anyway it degraded in 2 weeks to something useless (water + something I guess) When burned the H from NH3 and O from -NO3 at least forms H2O, the metal forms an oxide and N2 is also formed.

Sorry for al typos et.c. I'm really in a hurry right now

Engager - 21-2-2008 at 10:41

Somebody asked for structure of NHN, it is attched below. Made in ChemDraw.

sm1.jpg - 101kB

Nerro - 21-2-2008 at 12:48

Ni2+ only needs ten more electrons to fill up all of it's orbitals. You have too many ligands.

12AX7 - 21-2-2008 at 12:51

It's flat planar? That doesn't make sense, shouldn't it be an octahedral coordinate like trisoxalatoferrate (and for that matter, Ni(en)3 and more)?

Tim

Engager - 21-2-2008 at 15:47

Quote:

Originally posted by Nerro
Ni2+ only needs ten more electrons to fill up all of it's orbitals. You have too many ligands.

Bingo! You just got point, why it is unstable and explosive.

Engager - 21-2-2008 at 15:49

Quote:

Originally posted by 12AX7
It's flat planar? That doesn't make sense, shouldn't it be an octahedral coordinate like trisoxalatoferrate (and for that matter, Ni(en)3 and more)?

Tim

It should be, but my reference in russian chemistry book points at this structure. However all this subject is doubtfull, correct answer may be acchived only by dirrect physical measurements.

[Edited on 22-2-2008 by Engager]

Zinc - 30-3-2008 at 11:19

Tetraamine zinc peroxide. [Zn(NH3)4]O2 I searched but couldn't find any information about it. Does anyone know its properties or how to make it?

Axt - 1-4-2008 at 10:25

What makes you think Zn(NH3)4O2 even exists? ZnO2 is covalently bound unlike the Zn salts.

Zinc - 1-4-2008 at 11:17

Quote:

Originally posted by Axt
What makes you think Zn(NH3)4O2 even exists?

I didn't know exactly if it exists, I only heard once about it and saw the formula.

Taoiseach - 5-4-2008 at 03:27

Where did you hear about it? I also doubt it exists.

There seems to be an "zinc acetate-peroxide complex" tough:

http://www.freepatentsonline.com/4172841.html
http://www.freepatentsonline.com/5152996.html

Zelot - 7-4-2008 at 17:20

Yesterday I did electrolysis of an ammonium nitrate solution with copper electrodes in hopes of getting TACN. After an hour or two at ~5 volts and 1 amp, I filtered the solution and ended up with CuOH in the filter and a very deep blue-violet solution in the container. I assume this is either a tetraamine copper or tetraamine copper nitrate solution.

You add acetone to precipitate it, right? Flame tests will follow after precipitation.

Zinc - 8-4-2008 at 01:11

Quote:

Originally posted by Taoiseach
Where did you hear about it? I also doubt it exists.

There seems to be an "zinc acetate-peroxide complex" tough:

http://www.freepatentsonline.com/4172841.html
http://www.freepatentsonline.com/5152996.html

I don't remember where I heard about it.

The acetate complex seems interesting and easy to make. As far as I understand dissolve some zinc in acetic acid, remove the excess piece of Zn, and then add some H2O2 and heat it so the complex precipitates. Then just filter it, wash and let dry. I will try today and post the results.

Is there perhaps some similar copper complex?

[Edited on 8-4-2008 by Zinc]

Zinc - 9-4-2008 at 04:16

Unfortunately I wont be able to make the complex as I have trouble to make zinc acetate. My Zn doesn't want to dissolve in 80% acetic acid.

Microtek - 9-4-2008 at 23:47

Just make some Zn(OH)2 and react that with acetic acid to get your acetate.

Zelot - 10-4-2008 at 05:45

Don't you make Zn(OH)2 by reacting a zinc salt with ammonia?

not_important - 10-4-2008 at 06:27

Check the solubility of zinc acetate (dihydrate), you need enough water to get all of it into solution. I think it's about 1 g/2 cc I think you are going to want to dilute the acetic acid a bit.

Also try touching a bit of platinum, gold, silver, or even nickel, iron, or copper to the zinc while it is in the acid.

Another thing you could try is to treat the zinc with a little HCl, to get it reacting, then quickly rinse and drop into warm acetic acid.

Zinc - 12-4-2008 at 13:16

Today I added some more acetic acid (quite a lot has evaporated) and diluted it with some warm water.

Zelot - 16-4-2008 at 20:11

@zinc
Were you successful in making the ZAP yet?

Zinc - 17-4-2008 at 00:57

Not yet. As my attempt to make Zn acetate failed. Today I am going to dissolve Zn in HCl and later precipitate the hydroxide with NH4OH (first have to buy that). Then dissolve the hydroxide in acetic acid.

Taoiseach - 17-4-2008 at 02:03

Be careful when you precipate with NaOH because zinc is ampother! If you add too much, the precipate will redissolve:

[Zn(H2O)6]2+(aq) + 2OH-(aq) ---> Zn(OH)2(s) + 6H2O(l)
Zn(OH)2(s) + 2OH-(aq) ---> Zn(OH)42-(aq)

Na2[Zn(OH)4] is well soluble.

If you have sodium carbonate, use that! You will get basic zinc carbonate plus quite some CO2 bubbling. Zinc carbonate does not redissolve in excess sodium carbonate and it will be easier to neutralise your acetic acid with it. Just add zinc carbonate until no more is dissolve and no more CO2 is formed, then filter & evaporate the solution.

Zinc - 17-4-2008 at 12:05

Very good. I will then use that method (which is better for me as then I don't have to buy NH4OH and I already have sodium carbonate).

Today I put some Zn in HCl to dissolve and tomorrow I will precipitate the the basic carbonate.

Zinc - 18-4-2008 at 12:59

Right now I am filtering the precipitated Zn carbonate. When it filters I will dissolve it in acetic acid and add 30% H2O2 to precipitate the peroxide complex.

Zinc - 18-4-2008 at 13:24

When it filtered I dissolved it in an excess of acetic acid and added some 30% H2O2. Nothing precipitated. Only the solution turned to a brown (but transparent) color. Does anyone know why did it happen?

And one more question, is that complex explosive? As the patent doesn't say that it is.

Zelot - 26-4-2008 at 09:52

I may have made some TACN by adding a large excess of NH4OH to a CuSO4+NH4NO3 mixture. It became deep blue, almost violet. Some copper metal was precipitated, but was then formed into CuOH. I filtered to get just the solution. Can you precipitate TACN with rubbing alcohol?

Zinc - 4-5-2008 at 05:30

I think that it can but I never tried.

Taoiseach - 5-5-2008 at 00:49

TACN should be insoluble in ethanol but copper(II)nitrate is soluble. I doubt you can precipate TACN from a solution containing [Cu(NH3)4](2+), SO4(2-) and NO3(-) ions. You might be able to get a mix of the tetrammine sulfate and nitrate tough.

Jor - 29-5-2008 at 02:17

I just made some tetraamminecopper(II) persulphate.

When I put some of the dry material in a test tube, and heat the test tube, it explodes with a small 'Poof' sound.
White smoke evolves and a light green residue remains.
This residue dissolves in dilute hydrochloric acid. Small amounts of gas (bubbles) are evolved while it dissolves.
Now who knows what reaction is taking place when this compound explodes/decomposes? I have no idea.

woelen - 29-5-2008 at 04:06

The persulphate ion oxidizes the ammonia in the cuprammine complex, but the compound is highly oxygen deficient, so the explosion is not powerful.

Jor - 29-5-2008 at 05:21

Yes, I know the persulphate is oxidising the ammonia, but I can't figure out what's really happening... What is the equation?

woelen - 29-5-2008 at 05:42

You can't give a precise equation of this. There are so many reactions.

For the oxidation part, you could write: 2NH3 + 3[O] --> N2 + 3H2O

Here, the [O] comes from the persulfate ion: S2O8(2-) --> S2O7(2-) + [O]

The persulfate ion then with the water formed builds two bisulfate ions, free pyrosulfate certainly will not exist under these conditions with ammonia and water present. These bisulfates consume other ammonia molecules and the heat of decomposition also certainly will expell some ammonia from the complex as well. This explains why the solid becomes light blue/green, the tetrammine complex is destroyed.

Most likely, there also will be some unreacted persulfate as well. This explains the bubbling when it is added to hydrochloric acid.

[Edited on 29-5-08 by woelen]

Zombiekitten - 22-6-2008 at 14:17

is there tetramminecopper(II) picrate?

Axt - 22-6-2008 at 23:35

Well, yes there is. However I don't see any reason to pop my eyes out over it.

The Combustion of the Salts of Tetraamine Copper (II).

franklyn - 26-7-2008 at 06:52

Not a free download yet, but thought I would pass along this citation

http://www.ntis.gov/search/product.aspx?ABBR=AD779649

http://stinet.dtic.mil/oai/oai?verb=getRecord&metadataPr...

http://stinet.dtic.mil/stinet/jsp/docread.jsp?K2DocKey=http%...

Another related thread on TACN
http://www.sciencemadness.org/talk/viewthread.php?tid=2187&a...
and this post
http://www.roguesci.org/theforum/showpost.php?s=a13507bed910...

______________

Just a thought,
It occurs to me that formaldehyde will condense with a coordimated metal amine complex.
Assuming coordination is not completely destroyed releasing the metal oxide, this would add
fuel to the oxygen rich compound.

A Textbook of Inorganic Chemistry Ed. J. Newton Friend
Vol X The Metal Amines M. M..J. Sutherland 1928
This I must have obtained on this forum but the location eludes me so download here
http://www.keepmyfile.com/download/df912e2298535

Explosive Properties of Metal Ammines
Previuosly provided by Roscoe Bodine, here
http://www.sciencemadness.org/talk/viewthread.php?action=att...

Related to my conjecture above

Attachment: Reaction between Formaldehyde and Ammonia.pdf (763kB)
This file has been downloaded 1816 times

Taoiseach - 27-7-2008 at 02:53

Interesting stuff, thx 4 the upload!

I wonder if an explosive complex might be obtained by heating Ca(NO3)2 in a steam of dry NH3. Bis-en-copper(II)bromate might also be obtained this way.

I believe there is some kind of water of hydration in TACN and TACP. Both salts dont burn readily even when completely dry. However, when heated carefully in a steam of dry air for some time their blue color shifts towards a deep violet color and these violet salts DO burn/explode readily. In moist air I observed that my putative anhydrous TACP turns deep blue again after a short time. Also I once left a small bottle of TACP in a very hot spot in the sun (not a wise thing I know :cool

and droplets of water deposited inside and the salt turned violet. When put in a cold place again the droplets disappeared and the color gradually shifted towards the familiar blue again. Freshly precipated tetammine-Cu-persulphate also forms blue needles. These dont burn very well. After washing with ethanol and heating in a stream of warm air however a violet powder is obtained which explodes from flame and impact readily.

TACN will decompose in open air into green inert junk when its not completely dry so dont try heating it until its carefully washed with anhydrous ethanol.

Btw Bisethylenediamine-copper(II)-perchlorate easily precipates from solution when ethanol is added. Its insoluble in ethanol and thus can be obtained easily in a perfectly dry state. Forms violet needle-shaped crystals and explodes readily from impact however it is slightly hygroscopic.

Slightly off-topic but might be of interest anyways: There is a potassium-chlorate-tartrate double salt which is said to be explosive. I made some of this by mixing stoichiometric hot solutions of potassium chlorate and potassium tartrate. First little cubes started to crystalize which did not show any energetic properties - I suppose this was the tartrate alone. Then plate-like crystals appeared which when ignited with a match burned instantaneously with a little flash. I did not test impact sensitivity tough.

Here's some more double salts you might want to try:

Calcium-acetato-perchlorate: 2/50 mol Ca-acetate are dissolved in 50ccm H2O and 1/50Mol 50% HClO4 is added. The sol is evaporated over H2SO4.

Barium-sodium-acetato-nitrate: 54g Ba-acetate and 34g sodium nitrate are dissolved in 120cc hot H2O and left to evaporate @RT. First Ba(NO3)2 will crystalize; the double salt forms long prisms and explodes upon heating.

Theres also a lead-sodium-acetato-nitrate double salt which should be easy to prepare: 15,2g lead acetate and 3,4g sodium nitrate are disolved in 80ccm H2O and left to evaporate at RT. The double salt forms small leaflets and cannot be recrystallized from water w/o decomposition.

Taoiseach - 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:

natriumperoxid - 13-11-2008 at 07:55

I recently came across a rather interesting compund: [Ag(DMSO)2] [ ClO4].
Method:
Excess DMSO is added to a solution of AgClO4 in Acetone, solvent removal by vacuum.
Interestingly, the method uses Me2CO 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, ClO4- 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, ClO4- 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)CO4 and almost killed himself.

[Edited on 13-11-2008 by natriumperoxid]

[Edited on 13-11-2008 by natriumperoxid]

12AX7 - 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.

Tim

natriumperoxid - 13-11-2008 at 09:16

Bugger >.< That would mean another: So obvious I don't realise it. Cheers for the information.
However, ClO4- 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 ClO4- as a ligand be closer to adjacent ligands such as NH3, and would the overall compund therefore be more unstable (less energy needed to make the oxidising and reducing parts react)? Same for the NO3 - anion...

[Edited on 13-11-2008 by natriumperoxid]

12AX7 - 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.

Tim

User - 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.

Taoiseach - 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.

Jor - 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.

Jetto - 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]

Jetto - 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.

User - 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

.

User

[Edited on 17-4-2009 by User]

Taoiseach - 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.

User - 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]

Formatik - 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.

Taoiseach - 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.

User - 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.

Taoiseach - 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]

Formatik - 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.

Persulfate Electrolysis

TechnologicallyRetarded - 15-8-2009 at 00:22

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

Diazidodiamminocopper(II)

Taoiseach - 17-8-2009 at 05:33

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.

woelen - 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]

hmmm

Formatik - 17-8-2009 at 21:55

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.

Attachment: [Cu(NH3)2](N3)2.pdf (1022kB)
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Taoiseach - 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:

more ammines

Formatik - 3-9-2009 at 07:15

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.

lead chlorate complexes

Formatik - 4-9-2009 at 17:52

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.

Attachment: US1206456.pdf (146kB)
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