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Author: Subject: Exotic Primaries - Complex Salts
chemoleo
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thumbup.gif posted on 30-3-2004 at 17:44
Exotic Primaries - Complex Salts


A while ago I had a conversation with Ph Z. on nickel complex salts. With his permission, I post the preparation of the extremely potent
Nickel Hydrazine Nitrate primary, which is known to have high VODs (betw 6-7 km/sec).
This primary has the advantage of being relatively insensitve to friction and shock, yet sensitive to heat (flame).

The preparation is very simple, once you have hydrazin or hydrazin sulphate. For the preparation of hydrazin and hydrazin sulphate, see
http://www.sciencemadness.org/talk/viewthread.php?tid=1128
http://www.sciencemadness.org/talk/viewthread.php?tid=376
http://www.sciencemadness.org/talk/viewthread.php?tid=757
for details.


PREPARATION:

1. Take Nickel nitrate (green emerald salt) saturate hot water with it...then add for 1 volume of this 1 volume of ethanol (80-99%) is fine.You have now a green ethanolic solution of Ni(NO3)2.

2. Perform the same mix but with concentrated hydrazine...use i.e. 80% NH2-NH2 what equals NH2-NH2.H2O the monohydrate....take one volume of this and add one volume ethanol. Alternatively, react hydrazine sulphate with Ca/Ba hydroxide, and filter Ca/BaSO4 off to be left with a solution of hydrazine.

3. Take the whole hydrazine batch and add dropwise under stirring the green alcoholic Ni(NO3)2 solution. A deep blue (saphire) precipitate appears which upon agitation turns lila pinkish. Continue adding the nickel nitrate solution until no more precipitate occurs. It is not a bad idea to filtrate the precipitate from time to time to ascertain whether a green coloration persists (if it does, then stop adding the nitrate solution). While mixing and precipitating, heat the mix a bit so as to adjust the addition speed in a way you can hold the mix (<50°C) (with this phrase i am not sure what he meant - surely the precipitation is works better at lower temperatures?)
The complex salt filtrate is insoluble in ethanol/water which helps fast and smooth precipitation.You get plenty of precipitate and yield is close to 100%.
Make sure the nickel nitrate solution is added in excess so as to completely use up the hydrazine.

4. Once filtered (and possibly washed with 50% EtOH/H2O0, dry on a radiator (60°C max), and after a few hours/days one gets a solid clay like cake. This can be broken ground into small pieces/dust (no risks of explosion as long as you don't hit it with a strong hammer blow...grinding with a spoon is riskless). After grinding allow to dry one more day. After this the complex salt has turned into a pale white dust that burn fiercely and bright white when lit (like nitrocellulose). When wrapped in Al foil, and thrown into a fire, one gets a powerful detonation (whereby the Al foil acts as a weak confinement)


Side Notes:
Make sure the hydrazine is free of ammonia, best way to assure this is to simply use the hydrazin sulphate, which doesnt dissolve well, while ammonium sulphate does.
NH3 makes a blue complex with Ni(NO3)2 that is unsoluble in pure ethanol but soluble in water. So if you have only a blue precipitate and no pink, this means that you have a majority of NH3 in your NH2-NH2.
The pink hydrazine complex is weakly soluble in cold water while the blue ammonia complex is very soluble!
Upon shock the nickel hydrazine nitrate comple has the sensitivity of TNT (which is very safe) but still don't hit a hammer on it and grind it mildly.

This primary can be used in combination with Ag2C2.HNO3x (1/2-1/3), which increases the sensitivity a bit, while reducing the VOD/Power only very little...

To confirm you have the right stuff, the NHN complex salt should dissolve in excess ammonia, i.e. 6NH3 + Ni(NH2-NH2)3(NO3)2 <---> NH2-NH2 + Ni(NH3)6(NO3)2



Other hydrazine nitrate complexes:
Of course there are many to be tried.
Cu(NH2-NH2)3or2(NO3)2 decomposes immediately... so the corresponding Co, Ag, etc complexes are definitely worth trying.
In fact, I just realised, he actually tried the Co complex too, it is pale orange-brown and when dry it displays identical properties to Ni salt (same burning rate) .



Some more notes on energetic complex salts:

Copper tetraamine nitrates can be seemingly prepared the same way. Partially alchoholic solutions of each, and precipitation of the complex salt upon mixing (which by itself suggests a new compound formed). To achieve better precipitation, one can add ether, which precipitates ionic compounds.
Copper tetraamine perchlorates/chlorates exist too, they have a higher VOD than the nitrates (4-6 km/s, as opposed to 3-4 km/s for the nitrates). The perchlorates go to up to 95% of that of TNT, however they are more sensitve than the nitrates.

On the note of nitrate/chlorate/perchlorate complexs, consider this: Ni(NH2-NH2)3(ClO4)2 is explosive in solution while its nitrate is stable (! see above preparation); Cu(NH2-NH2)2(ClO4)2 doesn't exist but its nitrate is on the edge of existance.
Tendency-wise, such Co and Ni complexes are more likely to exist than their Cu brother.


At last

What I'd like to see in this thread is the general discussion of more or less energetic complex salts, as long as they haven't been discussed before (such as acetylides, TACC etc).
I am thinking of, i.e. ethylenediamine perchlorate complexes, hydroxylamine nitrate/perchlorate complexes, and so on!
Plus, all the different types of ammonia-nitrate/perchlorate complexes, whatever you can think of!
Get your imagination going!

[Edited on 31-3-2004 by chemoleo]




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[*] posted on 30-3-2004 at 18:24


I have done some work with tetraaminecopper(II) nitrate, and I hope to try tetraaminecopper(II) acetylide soon. I had started a thread on these types of copper compounds already, do you think I should repost my experiments here in this thread for completeness?
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[*] posted on 30-3-2004 at 18:41


Yes, I realise now it's a bit tricky.

In that thread ( http://www.sciencemadness.org/talk/viewthread.php?tid=1732 ), copper complex ions are discussed in general, while here, I'd hope to focus on the energetic ones, with any metal ions.
Might make sense to separate the two.
The reason why I posted this as a general topic (i.e. complex salt primaries) was because it doesn't just mention the nickel hydrazine complexes, but a few more, plus the tendencies between different cations.

I don't know how well it will work out in reality, but it may make sense to keep these things separate...plus it will keep such ideas combined in a single thread, which is easier to follow.




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[*] posted on 31-3-2004 at 04:44


I've made nitroguanidine and aminotetrazole complexes of silver nitrate, I can't remember if I tried copper... it was a while ago so my memory is a bit hazy, but I can make some more and post a bit of info.
They seemed quite good from what I can remember, but they're probably less effective, and no more OTC, than azides...




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[*] posted on 1-4-2004 at 07:13
energetic complexes


How about trying some of these! I have many more I’ll try finding this weekend. The Alchemist.

hexamminenickel perchlorate (HANP)
tetramminecopper perchlorate (TeACP)
tetramminecopper nitrate (TeACN)
tetramminecopper perchlorate-nitrate (TeACPN)
tetramminezinc perchlorate (TeAZP)
tetramminezinc permanganate (TeAZM)
tetramminecopper permanganate (TeACM)
hexamethylenetetramine sulfate-perchlorate (HSP)
tetramminezinc chlorate (TeAZC)
hexamethylenetetraminemagnesium chlorate (HMC)
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[*] posted on 15-4-2004 at 15:27


Found this in a patent :US 6,241,281

Example 7

The nitrate salt of copper tetraammine was prepared by dissolving 116.3 g of copper(II) nitrate hemipentahydrate in 230 mL of concentrated ammonium hydroxide and 50 mL of water. Once the resulting warm mixture had cooled to 40.degree. C., one liter of ethanol was added with stirring to precipitate the tetraammine nitrate product. The dark purple-blue solid was collected by filtration, washed with ethanol, and air dried. The product was confirmed to be Cu(NH.sub.3).sub.4 (No.sub.3).sub.2 by elemental analysis. The burning rate of this material as determined from pressed 1/2-inch diameter pellets was 0.18 inches per second at 1,000 psig.
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[*] posted on 15-4-2004 at 16:43
Interesting!


This surely confirms that non - aqueous conditions are not needed, ultimately. Essentially, precipitation by EtOH/aceton should be sufficient. I wonder whether that holds for copper tetraamine chlorate, too - and still I don't see why not, despite, the fact that the synth. is recommened to be done by bubbling NH3 gas through the solution.

Anyway... keep checking, there are a number of interesting complexes to test - i.e. ethylenediamine nitrates, hyrdroxylamine nitrates, and hydrazine nitrates (and possibly chlorates, perchlorates, with various metals).
It's quite a task I know.. but an interesting one!

[Edited on 16-4-2004 by chemoleo]




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[*] posted on 16-4-2004 at 16:22
Other Alcohols?


I wonder if it has to be ethanol? Would isopropyl alcohol work. I take it the only purpose of the alcohol is to produce a solution that the desired product has a lower soluability in than the original aqueus solution?
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[*] posted on 25-4-2004 at 17:01
Guanidine complexes?!?


Recently I did a few trials with guanidinium chloride.
I prepared the guanidine base by mixing GdnHCl with NaOH, specifically by dissolving GdnHCl in hot ethanol, and adding a 1.2 fold stoichiometric excess of NaOH. The NaCl precipitates (mostly), and I am left with a reasonably clean alcoholic guanidine-base extract.
This was diluted 50% with H2O.

I mixed this with 50% alcoholic CuCl2, NiCl2, and CoCl2 (as I dont really know the stoichiometry, I added an excess of guanidine base).
Interestingly, for the first two no colour change was observed, and no precipitation occurred. However, the cobalt solution changed to a very dark beautiful blue ink-like solution... clearly indicating that complexation occurred!
Unfortunately, adding more alcohol did not lead to precipitation. However, this will be tested with ether, and various other solvents.
Anyway... these experiments where done to see whether guanidine forms complex salts, and it was found indeed they do!
I shall, when I get the chance again, try the same story with cobalt nitrate, or possibly perchlorate. Hoping I can purify this!

Let's see where this adventure shall lead!!!


PS hodges, isopropylalchohol should work, too, if not better. Its hydrophobic character is more pronounced. In fact, I think I should try my future experiments with isopropylalcohol, maybe I could have precipitated the cobalt-guanidine complex with that!

[Edited on 26-4-2004 by chemoleo]




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[*] posted on 27-6-2004 at 17:03
Trinitrotriamine cobalt (III) Co[(NO2)3(NH3)3]


Here's a prep I recently found. I suspect it has some interesting properties :)

Procedure
Dissolve 6.2 g (0.025 mole) of Co(II) acetate tetrahydrate in 50 ml of hot water, then cool in an ice bath and add a cooled solution of 5.2 g NaNO2 in 25 ml 0.88 molar (?? it's not clear on that) ammonia.
Cool to 10 deg C.
Add 7 ml of 20 volume (2 M ) H2O2 slowly, with sitrring, and then add 1 g of activated charcoal. Keep the mix cool for 10 minutes. Then bring to the boil, and keep boiling for 30 minutes.
Filter the solution while it is still hot, then cool and filtrate in an ice bath/refrigerator. Filter off the crystals which form, wash with ethanol and leave to dry.
A further crop of crystals can be obtained by adding to the filtrate about 0.5 g of charcoal and evaporating to half the volume, filtering and cooling as before.
The mustard-yellow crystals are slightly soluble in H2O.

The H2O2 is used to oxidise the Co from II+ to III+ .
I wonder whether one can do the same with the NO3- ion?!? or how about the ClO4- ion?




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[*] posted on 29-6-2004 at 11:16
Like nitrogen?


Diaminotetrazole silver nitrotetrazole, [Ag(CHN4NH2)2]CN4NO2 :D

Now that is one that would scare me... I'll have to try making it!! I have aminotetrazole, silver nitrate and sodium nitrotetrazole all in stock...




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smile.gif posted on 2-7-2004 at 12:57
confused by Co[(NO2)3(NH3)3]


is it a complex ion or a neutral coordination compound without coupling ions? if the former how the nitro ligands could form without being oxidized? if the later then why is it named Trinitrotriamine cobalt (III) ?
are you sure the ammonia solution is 0.88 M? maybe it's been 8.8 M ?  25 ml 0.88 M means 0.025*0.88=0.022 mol NH3 which means less than 1 mol NH3 per 1 mol Co!:o.
anyway nitro and nitoso complexes are indeed cool especially that most energetic complexes we have discussed had a "fuel" cation and an "oxidizer" anion. but these cations can play the roll of an "oxidizer" or both oxidizer and fuel together (aka an energetic cation!) which makes use of energetic anions with moderate to perfect OB like azide, picrate, styphnate ... (as their coupling ion I mean)




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[*] posted on 2-7-2004 at 15:09


"0.88 ammonia" is an aqueous ammonia solution that has a density of 0.88g/mL, not 0.88 M.
I'm not sure what it's concentration is, but I assume it is saturated, because otherwise it would be a rather strange concentration to have as the standard. So just look up ammonia's solubility...




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[*] posted on 5-7-2004 at 18:48


Yes, the info where I got it from says '0.880' ammonia, without a specification as to whether it's M or % or whatever. I guess Nick F is correct on this.
Kaboom, as to your question, I have no clear answer. I guess it is a complex ion of the Co[(NH3)3]3+ type. So the same thing shold work with KNO3 instead of KNO2, too. But this is guesswork :(
The information I posted is all I have, with no omissions.




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smile.gif posted on 6-7-2004 at 13:35


True that there is an ambivalence with nitro nitrito complexes; sometimes with cobalt you have nitro complexes but nitrito too...this would mean there is a solid Co-N bond in the first and a Co-O bond in the second.

I just wonder why they feel like using Co II and H2O2 instead of Co III imediately?
Also activated charcoal?

Many nitro complexes and nitroso are obtained by bubbling NO, NO2 or N2O4 though solutions of complexating metal salts.




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smile.gif posted on 6-7-2004 at 14:00


In principle the field of investigation is infinite....
oxoAnions:
ONO2(-)
ONO(-)
-NO2
-NO
OCl(-)
OClO(-)
OClO2(-)
OClO3(-)
(idem with bromine and iodine core)
-MnO4

reducing coordinators or ligands:
-C=O
-CN
CN(-)
-NC
N3(-)
CS3(2-)
Sx(2-)
S(2-)
P(3-)
-N=S
-S-CN
SCN (-)
-N=C=S
-O-N=C
-S-N=C
NH3
NH2OH
NH2NH2
NH2-CH2-CH2-NH2
orthoC6H4(NH2)2 (eventually nitro or polynitro)
NH2-CO-NH2
NH2-CO-CO-NH2
NH2-CO-NH-CO-NH2
O=CH-CH=O
O=CH2
NH2-CH2-CO2H and AA
NH2-CH=NH
NH2-CH=O
(NH2)2C=NH
(NH2)2C=S
NH2-CN
NH2-C(=NH)-NH-CN
NxSx
NC-CN
....

on metal core catalyst
with various valence/oxydation states:

Zn II
Fe II Fe III
Ni II Ni III
Co II Co III
Mn II Mn III Mn IV Mn V
Cu I Cu II
....
Unusual metals...
Ru III
Pt II
Au III
...

Also many complexes have ... chirality and isomers...thus different cristallinity, sensitivities, stabilities, colours, optical activity...and then explosive properties.

PH Z




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sad.gif posted on 6-7-2004 at 14:13


Here is a synthesis I have but I don't know form where since it is a long time ago and it is not obvious for the second one if it was Ni(NH3)6(NO3)3 or Co(NH3)3(NO3)3 because the writter mixed up Co and Ni; maybe is it acceptable for the two and here too HNO3 or H2O2 is used as an oxydiser of the Co II or Ni II into Co III or Ni III...

***********************
1. Synthesys of {Ni(NH3)6}(NO3)2.
Ni(NO3)2 + 6NH3 = {Ni(NH3)6}(NO3)2.
Calculated quantities of Ni(NO3)2 are mixed with the calculated quantity of 25% NH3. The mixture is stirred until the Ni(NO3)2 dissolves and then cooled 20-30 minutes with ice-water mix. The suspension of {Ni(NH3)6}(NO3)2 is then filtered, washed with alcohol and dried at 40-50 C.
(pretty much the same procedure I used thus but I didn't knew I had this)

2. Synthesys of {Ni(NH3)6}(NO3)3.
The synthesys is done by the reaction:
4Co(NO3)2 + 4NH4NO3 + 20 NH3 + O2 = 4{Ni(NH3)6}(NO3)3 + 2H2O
7g Co(NO3)2 is dissolved in 10 cm3 water.This is mixed with 8g NH4NO3 and 18g 25% NH3 and 1 g active charcoal. After this for 2,3 hours air is bubbled through the solution. Then the solution is tranfered into 250 cm beaker glass and 150 cm water with 2,3 ml conc HNO3 is added to it. This solution is heated on water bath for 30 minutes.Filter this to discard the carbon, and add 20 cm3 conc HNO3. Left it some days until orange chrystals({Ni(NH3)6}(NO3)3 ) form, filter and dry in 100 degrres C.
The oxidizing from air can be substituted with adding 3%H2O2. 1mol H2O2 to 2 mols Co(NO3)2 . H2O2 is added to the solution after the
NH4NO3.
**********************************
I guess we have to try both to be sure ;)




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smile.gif posted on 6-7-2004 at 14:19


***********************************
2-(5-cyanotetrazolato)pentaaminecobalt (III) perchlorate
 
The compound 2-(5-cyanotetrazolato)pentaaminecobalt(III) perchlorate (CP) is a Primary explosive compound, and has been utilized in
low-voltage detonators because it reliably undergoes
deflagration-to-detonation transition (DDT). One of the precursers in its preperation is cyanogen.
 
Another compound similar is:
 
cis-bis (5-nitrotetrazolato) tetraaminecobalt (III) perchlorate (or "BNCP";)
*************************************




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[*] posted on 23-9-2004 at 20:05
Complex Persulphates


I found a reference to "Amminated cupric Persulphates", reference in full:

G. Morgan & F. Burstall, Cupric salts stabalised by ethylenediamine, Journal of the Chemical Society, (1926) Pg. 2026, London

It mentions that the ethylenediamine complex has explosive properties, but only mentions that the tetramine has been prepared and gives a reference, in full:

Cupric persulphate has not been isolated, but its tetramino- and tetrapyridine derivatives, [Cu,4 base]S2O8, have been prepared (Barbieri and Calzolari, Z. anorg. Chem., 1911, 71, 347).

One would think the tetraamine woud have better explosive qualities then the ethylenediamine, so I tried it.

Synthesis I Tried

10g copper sulphate pentahydrate was dissolved into 50ml warm water, to this was added 22g 25% ammonia, solution turned dark bluey purple on formation of the complex sulphate. Another solution was made of 10g potassium persulphate in 200ml warm water. The solutions were poured together, put into the freezer and brought down to 0°C.

On pouring off the cooled solution, a good yeild of purple "shaved-hair-like" crystals were left, they were scooped out and placed on absorbent paper in the sun to dry.

When dry the crystals of tetraminocopper (II) persulphate responded very energetically, a pinch of it explodes in a thump when exposed to a flame. Even though very sensitive to flame, in the drop test it wont fire at a height that will reliably fire PETN.

Ive added a movie of its ignition here - http://geocities.com/roguemovies7/

Looks promising, good yield, not very soluble, not hygroscopic, not impact sensitive and quite energetic. Hopefully it doesnt decompose.

EDIT: Igniting 0.2g resulted in a crack. Dont think you will have to much higher to get a true detonation.

[Edited on 24-9-2004 by Axt]
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[*] posted on 24-9-2004 at 19:31


Disregard what I said about impact sensitivity before, as it was based on a one off test, today it's firing at a drop of 20-25cm (PETN was 40-45cm) leaving a green molten plastic-like residue.






[Edited on 9-12-2005 by Axt]
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[*] posted on 27-9-2004 at 17:05


Wow, who'd have thought that! Complex persulphates.... I'd have assumed that the low OB would make them pretty useless!
It might be interesting to test this on additional metal ions (Ni in particular).

Did you try to see whether it is friction sensitive? How about larger amounts, do they leave a dent on a surface at all (trying to find out whether it's a true det.)

Edit: Merged this with existing thread to related complex salts.




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[*] posted on 27-9-2004 at 17:56


I have tried with nickel ammonium sulphate (only nickel salt I have), which yielded a very small amount of precipitate looking just like the copper salt, but it decomposed on drying to inert black gunk. The next day the purple solution itself had gone black and left a black film stuck fast to the glass. H2SO4 removed it.

The copper salt also looks to be decomposing, turning light blue, and lighting some up today resulted in a smoldering green blob :( It has been sitting in the hot sun this whole time, so that cant help, but its probably too unstable to have any uses.

The ethylenediamine salt is said to decompose explosively, selected quotes from the reference I gave above "Extremely unstable ..... decomposed, leaving a blue residue ..... exploded when contained in cooled specimen tube ..... detonated on percussion" So that aint real useful either :(

Since Im now in this thread, I've found Ag(NH3)2MnO4 to be pretty crap. Will burn simular to "good" KNO3/sugar, and ignites within a few minutes exposure to sunlight. Maybe its only possible practical use is as a sun sensitive ignitor.

[Edited on 30-9-2004 by Axt]
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[*] posted on 29-9-2004 at 07:46


Regarding the copper amine persulphate, how does it behave if it is stored in absolute ethanol? I am sure you could increase stability that way, and use it when desired.
Interesting, about the ethylenediamine. As I have some, might be interesting to test as the nitrate complex!




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[*] posted on 29-9-2004 at 13:29


Umm, run that theory by us once again chemleo, you want to store a peroxydisulphate in a liquid reducing agent to make it more stable?
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[*] posted on 30-9-2004 at 02:19


Yes, I would.
Just like I'd store HMTD in absolute ethanol.
As long as not acid/base traces are present, I wouldn't see much of a problem. Admittedly one never knows, so I'd test it with a small amount first... You could also try acetone, if worried about oxidation.

PS who is 'chemleo'? :P




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Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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