Sciencemadness Discussion Board

Exotic Primaries - Complex Salts

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User - 5-9-2009 at 02:27

This is most interesting.
Too bad that no exact procedures or tests are described.
Anyone willing to test this ?

Also this makes me wonder if a perchlorate form would be of any interest.

[Edited on 5-9-2009 by User]

Taoiseach - 5-9-2009 at 11:48

Fascinating stuff. The patent is quite vague tough. Here's the relevant info I scraped together: "Usually, one molecule of neutral chlorate of lead, one molecule of glycerine, or mannite or of any other polyatomic alcohol, and one molecule of plumbic oxide or hydrate respectively are caused to react [...] Likewise, a solution of bibasic or tribasic acetate of lead gives instantaneously in presence of a readily soluble chlorate, such as chlorate of soda, of calcium, of magnesium or of ammonium, or by addition of alcohol, with mannite, mannitan or even wih glycerin, fluminating compounds".


Due to its insolubility (0.037/100g H2O @ 20°C) Pb(ClO3)2 should precipate from hot solutions of KClO3/NaClO3 and any soluble lead salt.

A somehow related tidbit of info from the E&W board:

"Lead glycerolate can be done by boiling excess glycerol with lead hydroxyde (1/2hour)
CH2OH-CHOH-CH2OH + Pb(OH)2 --> CH2OH-CH(O-)-CHO-Pb + 2 H2O
Filtrate the white precipitate (Pb(OH)2 is also white so it appaers nothing has happened but stil) wash with water several times and allow to dry at 40-60C.
The dry compound deflagrates when heated confined."

This must be related somehow, maybe the chlorate complexes actually being double salts of lead glycerolate and lead chlorate.

A similar double salt forms from lead acetate and lead perchlorate.

There's plenty of info on metal glycerolate complexes:

BRYLANTS, J and PHILLIPPE, A N (1980). IR and moessbauer study of
iron glycerolates. J. Inorg. Nucl. Chem., 42(11): 1603-1611.

GADD, K F (1981). Complexes of copper (II) with polyhydric alcohols.
Educ. Chem., 18(6): 176-178.

HAZIMAH, A H (1998). Characterization and reactions of copper (II)
glycerol complex. Ph. D. thesis. Universiti Putra Malaysia.

HAZIMAH, A H; KAREN ANNE CROUSE; BADRI, M; CHONG, F K and
ABDUL RAHMAN MANAS (2000). Isolation and characterization of
copper (II) glycerol complex. Malaysian Journal of Chemistry No.1:
008-0011.

KOYANO, H; KATO, M and TAKONOUCHI, H (1992). Electrodeless
copper plating from copper-glycerin complex solution. J. Electrochem.
Soc., 139(11): 3112-3116.

NAGY, L; BURGER, K; KURTI, J; MOSTAFA, M A; KORECZ, L and
KIRICSI, I (1986). Iron (III) complexes of sugar-type ligands. Inorg.
Chim. Acta., 124(1): 55-59.

RADOSLOVICH, E W; RAUPACHI, M; SLADE, P G and TAYLOR, R M
(1970). Crystalline cobalt, zinc, manganese and iron alkoxides of
glycerol. Aust. J. Chem., 23: 1963-1971.

TAYLOR, R M and BROCK, A J (1989). Zn glycerolate complex. US
Patent 4 87 378.

I wonder if similar complexes of copper, iron or nickel exist.

woelen - 6-9-2009 at 09:41

I retried the experiment of Taoiseach and now I had partial success. I took some copper sulfate and dissolved this in 12% ammonia, I added just enough ammonia to get all of the copper sulfate dissolved. This gives a deep blue solution. To this I added a few drops of highly concentrated solution of sodium azide. When this is done, the color remains deep blue.

Then I added 96% ethanol, approximately the same volume as the deep blue aqueous solution. Again, this results in formation of a purple/blue precipitate. This time, however, did not keep the precipitate, but the liquid. Once the precipitate had settled at the bottom, the liquid had a green/cyan color and it was fairly dark. This liquid was allowed to stand for many days. After several days, a crop of nice green glittering crystals was at the bottom of the test tube. These crystals were put on a piece of filter paper, in order to absorb most of the liquid and the crystals were allowed to dry in a warm place in a petri dish for a few hours.

The dry material consists of many small needle-shaped green crystals, some of them glittering very nicely. The material looks nice. Some of these crystals were put on the tip of a screw driver and this was put in a flame. The result is a soft crackling noise and many little sparks are created. Some larger crystals give a somewhat louder PENG sound, but not an impressive bang. Most beautiful, however, is the color which is given to the flame. When a crystal of the material crackles, it produced many fine particles, which give a beautiful green color to the flame of the alcohol burner. This is due to the copper content of the particles.

So, again, I did not get the heavy bang explosive, but the material I obtained was cute and had funny properties. My conclusion, however, is that making the ammonia/azide complex of copper(II) is not that easy, one easily ends up with other compounds.

lead oxohalogenates

Formatik - 6-9-2009 at 11:24

Note that Pb(ClO3)2.H2O is actually very soluble in water (see Mellor, et al.), the compd. is deliquescent. There is conflicting data on alcohol solubility; whereas in Mellor it says alc. precipitates it, Wächter (J.pr. 1843, 30, 329) says it is "very easily" sol. in both water and alc.

The lit. methods of preparation go off from dissolving PbO in calculated amounts of aq. chloric acid. If Wächter is right then one might be able to make it by combining aq. Pb(NO3)2 and NaClO3 and then separating by ppt. the NaNO3 with alcohol (at 70 wt.%, rest water not more than 7.81g NaNO3 dissolves per 100g), cryst. impure chlorate.

The monohydrate itself can explode if heated quickly to about 235º (Wächter just described vigorous dec. under sizzling), at 150º there is 4.59% water loss. Pb(ClO2)2 is also explosive and was briefly mentioned in this thread.

[Edited on 6-9-2009 by Formatik]

PHILOU Zrealone - 8-9-2009 at 05:29

Quote: Originally posted by Taoiseach  

A somehow related tidbit of info from the E&W board:

"Lead glycerolate can be done by boiling excess glycerol with lead hydroxyde (1/2hour)
CH2OH-CHOH-CH2OH + Pb(OH)2 --> CH2OH-CH(O-)-CHO-Pb + 2 H2O
Filtrate the white precipitate (Pb(OH)2 is also white so it appaers nothing has happened but stil) wash with water several times and allow to dry at 40-60C.
The dry compound deflagrates when heated confined."

Looks familiar to me... as I was the writer :D
As mentionned the white Pb(OH)2 is hard to distinguish from the glycerolate as its displays the same colour.
The compound is not very impressive: when dry and in contact with a flame in the open, it puffs like does the Ag oxalate, but under confinement of an aluminium foil and if warmed, it explodes mildly.
Maybe with glycol in place of glycerol, one would add a bit more to the explosive power.

Lead chlorate will add oxygen to the stuff en renders it more energetic. Although I wonder if a simple mix of Pb(ClO3)2 and polyol co-cristallised will not do as good...because Pb(ClO3)2 being detonating stuff on its own in a chlorate/sugar like composition must be quite powerfull.

Taoiseach - 8-9-2009 at 22:44

Wikipedia once again proved guilty of spread of bullshit and disinformation :mad: Yes Pb(ClO3)2 is highly soluble. I suppose there's no easy way to obtain it from soluble chlorates.

Dilute solutions of HClO3 can be made from oxalic or tartaric acid and KClO3. Potassium hydrogen tartrate/oxalate is highly insoluble.

I found a procedure on a German forum - basically 14,5 grams of KClO3 are dissolved in 70ml H2O @60°C. Another solution is prepared from 17,8grams L+ tartaric acid and 30ml H2O. These solutions are cooled to 10°C and poured together to yield ~10% HClO3 and insoluble K tartrate which is then removed by filtration. It obviously works. Of course this says nothing about the purity of product . Then again why give a shit about purity when the aim is to make some esoteric stuff that we don't even know the composition of.


lead bromate

Formatik - 10-9-2009 at 01:03

Pb(BrO3)2.H2O is much less soluble than the Pb(ClO3)2 salt according to Mellor, so it shouldn't be too hard to collect from NaBrO3 and Pb(NO3)2 aq. It might build similar complexes, but who knows I can't determine it at present.

chloric1 - 10-9-2009 at 04:51

Quote: Originally posted by Taoiseach  
Wikipedia once again proved guilty of spread of bullshit and disinformation :mad: Yes Pb(ClO3)2 is highly soluble. I suppose there's no easy way to obtain it from soluble chlorates.

Dilute solutions of HClO3 can be made from oxalic or tartaric acid and KClO3. Potassium hydrogen tartrate/oxalate is highly insoluble.

I found a procedure on a German forum - basically 14,5 grams of KClO3 are dissolved in 70ml H2O @60°C. Another solution is prepared from 17,8grams L+ tartaric acid and 30ml H2O. These solutions are cooled to 10°C and poured together to yield ~10% HClO3 and insoluble K tartrate which is then removed by filtration. It obviously works. Of course this says nothing about the purity of product . Then again why give a shit about purity when the aim is to make some esoteric stuff that we don't even know the composition of.



You will need to filter with mineral wool, glass wool etc. IIRC filter paper soaked in chloric acid ignites. It would be adviable to recrystalize the lead chlorate you would form. Maybe a water/alcohol mix 1:1 would suffice.

Taoiseach - 10-9-2009 at 22:26

I never heard that filter paper soaked in 10% HClO3 ignites and I highly doubt it would. The concentrated acid is very dangerous yes, but don't shit your pants about something as dilute as 10%.

Of course Ba(ClO3)2 + H2SO4 will give a product of higher purity and concentration. The advantage of the above procedure is that the precursors are easier to come be. KClO3 can be made by electrolysis and tartaric acid is a common food additive and doesnt raise any suspicion when ordered in a drugstore.


Rosco Bodine - 11-9-2009 at 08:19

Quote: Originally posted by Taoiseach  

Dilute solutions of HClO3 can be made from oxalic or tartaric acid and KClO3. Potassium hydrogen tartrate/oxalate is highly insoluble.

I found a procedure on a German forum - basically 14,5 grams of KClO3 are dissolved in 70ml H2O @60°C. Another solution is prepared from 17,8grams L+ tartaric acid and 30ml H2O. These solutions are cooled to 10°C and poured together to yield ~10% HClO3 and insoluble K tartrate which is then removed by filtration. It obviously works. Of course this says nothing about the purity of product . Then again why give a shit about purity when the aim is to make some esoteric stuff that we don't even know the composition of.


Neutralizing the so obtained 10% HClO3 with glycine or betaine should provide the corresponding organic chlorate salt, and these salts should be highly energetic materials
of possible interest gotten via a relatively mild and mundane synthetic route.

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[Edited on 11-9-2009 by Rosco Bodine]

Lead Salt Solubility Info.

TechnologicallyRetarded - 15-9-2009 at 04:17

I'll take this opportunity to post this link to the 'The lead salts preparation thread!'....

http://www.sciencemadness.org/talk/viewthread.php?tid=5490

This may all be known to you, I haven't read through this entire thread yet. Please don't flame me (again), I'm still getting used to the skin-grafts from last time.

Basically:

Pb(ClO3)2.0H2O - 144 @ 18C (vs in EtOH)
Pb(BrO3)2.1H2O - 1.33 @ 20C
Pb(IO3)2.0H2O - 0.0025 @ 25C

Pb(ClO4)2.0H2O - 441 @ 25C (s in EtOH)

All the above info was nicked from the other thread.

Tr


TechnologicallyRetarded - 15-9-2009 at 04:23

... Just edited the solubility of the Chlorate on wiki.

And Rosco, that's some trippy music... not sure if I like it or not. ; )

Tr

Rosco Bodine - 15-9-2009 at 06:06

The lead chlorate can form by double decomposition from lead nitrate and potassium chlorate when the lead chlorate is precipitated as an inclusion in a low solubility salt such as a clathrate type quasi-multiple salt. There is always the possible unpredictable formation of possible double salts or multiple salts which can complicate matters, or can produce different results depending upon temperature and concentrations. Wife swapping, threesomes....who knows
for certain what may happen in an untested mix of reactive components where the unexpected anomaly may be lurking :P

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Hexamine complexes?

TechnologicallyRetarded - 20-9-2009 at 05:30

I've done some research into Polyol complexes, I don't know if we'll win this battle. The complexes consist of deprotonated polyol species, usually adjacent groups. For instance, when glycerol form's complexes, they consist of -OCH2CH(O-)CH2OH groups. This leaves no room for other anions... does it?

But Hexamine complexes however...
Quote:

The 3CuCl2·4hmta·2HCl·2H2O complex (hmta=hexamethylenetetramine) has been obtained from acid solution. Spectroscopic and magnetic investigations indicate that, in the crystal structure of this compound, the polymeric chains, 3CuCl2·2hmta·2H2O, and hmta·HCl groups occur.

... may be promising.

more info on some lead salts

Formatik - 30-9-2009 at 13:50

After some more digging, I found out Pb(BrO3)2.H2O has some energetic properties itself also (easily explodes if it attempting to dry it by heat). It also forms complexes, e.g. preparation of it using lead acetate isn't recommended since lead(II) bromate-acetate (highly expl.) can form. The lead acetate-bromate complex already has its own thread here.

Some other complexes are basic lead (II) chlorate-acetates: Pb2(OH)(ClO3)2CH3CO2.2.5H2O detonates with loud bang if dry heated. Pb3(OH)2(ClO3)2(CH3CO2)2.3H2O on dry heating, explodes with a large brisance. Various lead(II) acetate-perchlorates which all explode on impact or by heating. Lead (II) perchlorate-oxalates, e.g. [Pb2(C2O4)](ClO4)2.3H2O is said to only deflagrate weakly when heated.

Lead chlorite-formate Pb(ClO2)(HCO2), unstable at regular temps., yellow compd. made by digesting a cold and quickly precipitated Pb(HCO2)2 precipitate with a NaClO2-solution (in excess of up to 50%), then washed with alcohol and ether. It discolors brown in a few hours. No energetic props. mentioned of this compd, but it could have some.

Attachment: Pb(BrO3)2.pdf (1.1MB)
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Attachment: Pb(ClO3)2 no.1.pdf (1.7MB)
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Attachment: Pb(ClO3)2 no.2.pdf (882kB)
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Attachment: Pb complexes.pdf (1.7MB)
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Attachment: Pb comp2.pdf (511kB)
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[Edited on 30-9-2009 by Formatik]

chloric1 - 30-9-2009 at 15:07

My understanding of German is a bit sketchy. But I caught a route to synthesis by PbSO4 and KClO3:cool:! Most lead sulfate I would have would be waste from lead(II) reactions! I cannot recall to what extent but I do know that lead sulfate is considerably more soluble in water than the carbonate. Of coarse pretty dilute solutions(lesss than 10%) would be needed. Come to think of it, the reaction could be easily reversed so it might be better to work with 80% alcohol solution in water instead of pure water. I also seed that the crystline habit is monoclinic prisms. A beautiful sight to behold for sure!:D:D:D

Taoiseach - 30-9-2009 at 22:48

@Formatik

thx for the interesting upload. From which book are these excerpts?

I feel some misgivings now about the possible tartaric acid + KClO3 route to lead chlorate. Quite possible that a dangerous tartrate-chlorate double salt would form.

What remains to be investigated is the insoluble clathrates of lead chlorate, e.g. a solution of sodium chlorate, lead acetate and glycerine should yield such a compound after addition of ethanol. That'd be an interesting novelty to fiddle with - no fancy chemicals required. Anyone dare to try? :cool:

Gmelin

Formatik - 1-10-2009 at 10:26

Quote: Originally posted by Taoiseach  
@Formatik

thx for the interesting upload. From which book are these excerpts?


They're from Gmelin's Handbuch der anorganischen Chemie. Basically it's a huge set of books found at decent libraries. I believe Beilstein Crossfire also gives access, but I've never set out to use it.

Btw, calcium oxalate is many times less soluble than potassium bitartrate. It's said though its sol. increases in strong acids, when I prepd. HClO3 using impure Ca(ClO3)2 (HClO4 thread), most to near all oxalate seemed to ppt. Using PbSO4 to prepare the chlorate looks like a pain, Gmelin mentions formation from boiling PbSO4 in KClO3 soln. citing the Ann. Chim. Phys. ref.

And a quick comment about HMTA complexes above. Although nitrogen containing, I doubt they show much interest because the energetic aspect may be muffled by the high oxygen imbalance. Like with the divalent metal hmta.persulfates, no energetic properties were noted in the journal cited earlier.

Pope - 16-10-2009 at 13:29

Is there a metathesis process I can use with KNO3 and Copper ******? to get copper nitrate?

User - 16-10-2009 at 18:50

Wrong topic, dumb question...?

Pope - 18-10-2009 at 02:55

Yea it is a dumb question cause it most probably has been answered god knows how many times, but it still has relevance that comes back to this initial thread and I thought that by posting here and not necro posting or making a new thread just to ask a simple question so I posted here expecting someone to understand, I guess not sorry.

Hydrazine hydrate copper perchlorate complex

Evilblaze - 12-11-2009 at 11:55

Yesterday I made a complex salt. On the forum it was previously mentioned, that this complex exist, but there was nothing else wrote from it. I found some information in: Encyclopedia of Explosives and Related Items - Fedoroff - Vol 7 of 10, but there was also just written, that this compound exist, but no other information.

The complex was prepared by the following way: a freshly prepared solution of copper perchlorate (from copper sulfate and sodium perchlorate) and hydrazine hydrate solution was reacted with each other and a brown precipitate formed. It was filtered and washed with water.

The brown precipitate when heated decomposes slowly and thus can be easily dried. It detonate when contact with flame. It can not be stored it decomposes to nitrogen, water and copper perchlorate. After 4 hour storing the compound, the copper-perchlorate-s blue color can be seen on the top of the complex.

Anyone got idea to a stable hydrazine-metal-perchlorate complex? The nickel salt detonates immediately. Maybe I will try the complex with Cd, Zn perchlorates.

I upload a video from this complex when it is ignited by flame.

Attachment: hydrazine-hydrate-copper-perchlorate complex 1.wmv (1.5MB)
This file has been downloaded 1080 times


Taoiseach - 12-11-2009 at 13:10

Interesting... nice video thx.

I have never heard of any other stable complexes of this kind. You should be VERY careful when you experiment with hydrazine perchlorate complexes as the stability seems to depend on crystal size and/or modification. The poor guy who made nickel hydrazine perchlorate first isolated a small amount of the dry crystals and didn't find them to be very sensitive. The next batch however exploded when he stirred the solution. Scary stuff :o




[Edited on 12-11-2009 by Taoiseach]

Taoiseach - 15-1-2010 at 03:30

Nitroaminoguanidine was mentioned here:

https://www.sciencemadness.org/whisper/viewthread.php?tid=81...

It is a condensation product of hydrazine and nitroguanidine. Due to its low O content its probably not very energetic. However its ease of preparation and its many nitrogens with unpaired electron pairs gave rise to the question if it could find use as an energetic fuel kind of ligand in coordination complexes.

I prepared a small batch following patent US2617826. This procedure was originally posted by Sobrero:

14g hydrazine sulfate and 8g NaOH were dissolved in 100ml warm water. This was added to a suspension of 11g nitroguanidine in 150ml boiling water. The solution was then hold at 80-90°C and stirred vigorously until all nitroguanidine had gone into solution. The liquid was colorless at first, then slowly turned yellow and finally deep orange.

The solution was then neutralized with conc. HCl while still hot, and then cooled in an ice bath to around 5°C. About 7.5g of pale yellow crystals were obtained. Note: Using less dilute solutions might give higher yields.

The lead and copper salts were prepared as follows:

Pb(NAG)2:
1g NAG was dissolved in 30ml of hot water (85°C).
1g Pb(OH)2 (prepared by adding 0,5g NaOH to a hot solution of 1,4g Pb(NO3)2) was added and the mixture stirred well. Upon cooling, yellow crystals formed slowly; these were pressed dry on a filter paper and dried with warm air from a heat gun.

Cu(NAG)2:
1g NAG was dissolved in 25ml of hot water (85°C).
1g CuSO4*5H2O in 10ml hot water was added, followed by 1g NaOH. The mixture was stirred well, filtered and the crystals dried with warm air.

NAG just flares upon ignition whereas the lead salt gives a pop sound similar to Pb(C6H2N3O7)2 in small amounts. The Cu salt is less energetic.

I then tried to make [Cu(NAG)2](ClO4)2. Papers on the nickel nitrate/perchlorate complexes are available in references section and suggest NAG acts as a bidentate ligand.
No measurements were made this time. A few mililiters of perchloric acid were neutralized with CuCO3. The liquid was filtered and a solution of NAG in 80°C water was added dropwise. The solution turned a deep blue color, not unlike an ammoniacal complex. Upon cooling in the freezer glistering blue crystals were obtained. These were carefully pressed dry on filter paper and dried with warm air, then dried over NaOH (to remove excess HClO4).
Trying to obtain a further crop by adding an equal volume of EtOH did not yield the complex but pale blue crystals of pure Cu(ClO4)2.

Upon ignition, the complex deflagrates violently. Its ability to make the DDT in very small amounts is comparable to NHN (nickel hydrazine nitrate) if not even better. Tightly wrapping 50mg in a few layers of aluminium foil and heating on a bunsen burner caused detonation. A quantity amounting to 1/4 of a matchhead was placed on concrete floor and given a good smack with a hammer and the resulting pain in the ears suggested its very energetic :) The crystals are perfectly dry, nonhygroscopic and air stable. Quite encouraging so far.

I also attempted to make [Cu(NAG)2](BrO3)2. A solution of copper bromate was prepared from barium bromate and copper sulphate. This was evaporated until a saturated solution of about 80°C was obtained. A solution of NAG in 80°C water was added with stirring. Complexation was evidenced by a deep blue coloration but the color quickly faded away. Effervescence was noted (O2 or CO2?) and the solution turned a pale green, also a faint smell of bromine was noted. I figure the bromate oxidized the NAG in solution. At a lower temperature the complex might persist tough.

(Left to right:) NAG, NAG-Cu, NAG-Pb, [Cu(NAG)2](ClO4)2

NAGs.jpg - 24kB

[Edited on 15-1-2010 by Taoiseach]

chemoleo - 15-1-2010 at 19:02

Very nice!
I don't like the method of drying the sample with NaOH (to remove HClO4), at least use KOH!
Why do you need 80 deg C for reaction (complexation) of NAG with copper perchlorate?
Impressive that the Cu-NAG-ClO4 crystals are stable at air! That is very unusual!
Pb(C6H2N3O7)2 - please clarify for the uninformed what it is.


[Cu(en)2]3[Co(NO2)6]2

Bear_with_vodka - 16-1-2010 at 08:25

Hi everyone. First, please forgive my rusty english :(
I can’t find anything about this odd complex in literature, any advice?
It was made by the reaction of somewhat concentrated warm solutions of Na3[Co(NO2)6] and [Cu(en)2](NO3)2, using stechiometric amounts. The mixture was brown and… nothing happened. But after 5-10min glistening small needles started to separate. After 2 hours the solution was light-brown with heavy crystals on the bottom. It was filtered, the crystals washed with cold water and dried at room temperature.
Dark-brown small needles, very sparingly soluble in water, non-hygroscopic, decomposes fiercely at ~150C, does not decompose with applied friction or strike, looks stable in air.
Small 300mg pile was ignited. It burned fast, but smooth and quiet with ~15cm high yellow-orange flame with a little light-green at the end without smoke and smell. There was a lot of black residue.
I know nothing about corresponding NH3 complex as it does not precipitate.

[Cu(en)2](NO3)2 was made by adding ethylenediamine into warm conc. solution of Cu(NO3)2 – from copper wire and 70% nitric acid (neutralized with a little NaHCO3 in the end). Beautiful big dark-violet plates. They melt in a spoon on the torch, then burn with a light-blue flash with loud fizzing.
The hydrazine complex of Cu(NO3)2 is very unstable. It’s an amorphous light green powder, which turns black after a few minutes, so it’s impossible to isolate it…


it's better to take cold, not warm solutions - that increases yield

one more edition:)
just finished another reaction... total failure:(
Both solutions were room temp. The mixture turned dark-green and a little of colourless gas evolved... 10min after - it started to smell intensively, very strange unpleasant smell
It was filtered in a gasmask (as I'm afraid of nitrosoamines, who knows what's in there...) The yield this time was 10% AND the compound burned much slower:(
Why so? But more interesting - what's the smell??

[Edited on 16-1-2010 by Bear_with_vodka]

Bear_with_vodka - 17-1-2010 at 03:42

I'll just keep talking to myself then:)
This time two solutions were cold, below zero. They were mixed and left outside for the night (-12 or so). The yield was 50% of nice quite big heavy crystals, cubic this time.
Now then, I'm not so sure about the right formula of complex as it is essencial to use excess of cobaltinitrite otherwise the stinky green solution is formed.




влажный комплекс2.jpg - 252kB

slightly wet

[Edited on 17-1-2010 by Bear_with_vodka]

Evilblaze - 17-1-2010 at 05:33

Quote: Originally posted by Bear_with_vodka  

[Cu(en)2](NO3)2 was made by adding ethylenediamine into warm conc. solution of Cu(NO3)2 – from copper wire and 70% nitric acid (neutralized with a little NaHCO3 in the end). Beautiful big dark-violet plates. They melt in a spoon on the torch, then burn with a light-blue flash with loud fizzing.
The hydrazine complex of Cu(NO3)2 is very unstable. It’s an amorphous light green powder, which turns black after a few minutes, so it’s impossible to isolate it…


[Edited on 16-1-2010 by Bear_with_vodka]


Wow, I'm happy to read this. I tried similar things in last month. At me Cu(ClO4)2 was used instead of the Cu(NO3)2, but I had similar results. The Cu-perchlorate solution with ethylenediamine gave a purple/pink solution. I just didn't had time to wait until it crystallizes so I thrown it out (but I will make it again soon).

But with the hydrazine it gave a really good complex. It is unstable until it is not dry. I made the complex quick, I waited for 10 minutes to participate (I added some alcohol, it made this quicker). I just filtered it fast, I washed it twice with water and twice with IPA. It was dried at 40 Celsius and it was used. I just have a small sample from this. The sample is 1 month old and there is no problem with it. If I ignite a small piece it just work's correctly. I think it is a really good complex salt :D

I will also try with the Cu(NO3)2. I think your solution was a bit acidic because of the nitric acid and maybe this caused the "color change from green to black".

Attachment: hydrazine-hydrate-copper-perchlorate complex test 2.wmv (1.5MB)
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Bear_with_vodka - 17-1-2010 at 05:52

hydrazine-hydrate-copper-perchlorate
Great thing! It does not burn but detonate.
I always wanted to make it, but I don't have HCLO4 :(
And there is no way of making that complex without HCLO4, am I right?
And yes, you have to wait after adding ethylendiamine. Just put it in the refregirator for a night. Thats about [Cu(en)2](NO3)2, maybe CLO4 complex is more soluble, I don't know

The only CLO4 complex I had made is [Ni(NH3)6](ClO4)2. It is insoluble, so the preparation is very easy. NiSO4 solution + NH3aq = deep-blue solution + NaClO4 solution = immediate light-blue-violet amorphous precipitate. It burnes violently with flashes.
[Ni(N2H4)2(3??)](NO3)2 is also insoluble (light green), but I think it's not interesting at all. Hard to burn it.

[Edited on 17-1-2010 by Bear_with_vodka]

Evilblaze - 17-1-2010 at 06:03

Quote: Originally posted by Bear_with_vodka  

I always wanted to make it, but I don't have HCLO4 :(
And there is no way of making that complex without HCLO4, am I right?


No! I made it from NaClO4 and CuSO4. NaClO4 or KClO4 can be obtained easier than HClO4. KClO4 = pyrotechnic supply shop;)
The only thing that you have to do is to wash it with water twice to get out the NaCl/KCl, because if the salt is in the mix, it won't detonate, it will just burn really quick, but no bang:D

Long ago I made a complex of hydrazine perchlorate (hyrdrazine hydrate and HClO4) and CuCl2. It was also a really good explosive, but it was not so stable as the hydrazyne hydrate -copper perchlorate complex (what is actually the BEST!).

I got some methylamine solution on the shelf.... Maybe I will try a complex with that too.....:D

[Edited on 17-1-2010 by Evilblaze]

Taoiseach - 17-1-2010 at 06:23

Evilblaze I had a virus warning when I DL'ed your video (Conficker worm).

AFAIK Copper(II)nitrate dissolves quite well in EtOH, so you could try to make the complex in almost water-free enviroment (Cu(NO3)2 comes in hydrated form however).

Copper(II)perchlorate is insoluble in EtOH tough.

Interesting to see that the hydrazine complex is reasonably stable. I tought they would be way too dangerous to make. From looking at your video I'd say its similar to the nitroaminoguanidine complex I described previously. It deflagrates *very* quickly, sort of on the verge to making DDT even in smallest amounts. It gives a dull report but not the high pitched bang you get when it is smacked with a hammer. Try wrapping in a few layers of aluminium foil, it should DDT even in minute amounts.

en-copper(II)-perchlorate crystallizes in *beautiful* dark purple needles. I made these by adding ethylenediamine to a cold solution of copper perchlorate, then boiled down to half its volume, added an equal volume of EtOH and put in the freezer. Even when quickly cooled down the solution deposits crystals of impressive size. It can be made anhydrous by carefully heating with warm air; the dark purple crystals then desintegrate into a pale light pink powder which is quite sensitive to impact.

[Edited on 17-1-2010 by Taoiseach]

Taoiseach - 17-1-2010 at 06:31

@Bear_with_vodka

>[Ni(N2H4)2(3??)](NO3)2 is also insoluble (light green), but I think it's not interesting at all. Hard to burn it.

Something must be wrong with your hydrazine. An ethanolic solution of nickel(II)nitrate upon addition of hydrazine *immediately* precipates light pink microcrystalline powder of [Ni(N2H4)3](NO3)2. This compound deflagrates violently upon ignition, also giving a dull report. It makes DDT *very* easily. Just the confinement from a few layers of aluminium foil is sufficient to make it detonate. I dont know what the green stuff is you described, possibly just hydrated nickel nitrate.

Evilblaze - 17-1-2010 at 06:34

Quote: Originally posted by Taoiseach  
Evilblaze I had a virus warning when I DL'ed your video (Conficker worm).


Virus? What the....? The video was recorded half month ago, I just made 1 video from 4 with windows movie maker and that's all. Who packed a virus in it?

Quote: Originally posted by Taoiseach  
Try wrapping in a few layers of aluminium foil, it should DDT even in minute amounts.


In aluminum foil with a hammer it ignites from a 8cm drop. The hammer weighted 140g.

It is not really stable to use it in larger amounts, but it was interesting to make it and test it.

Bear_with_vodka - 17-1-2010 at 06:47

[Ni(N2H4)3](NO3)2
*looking at notes*
aarrr... yes it's light-pink, sorry
I had only about 200mg, it burned with very small flashes-strikes and I didn't like it))


[Edited on 17-1-2010 by Bear_with_vodka]

Taoiseach - 17-1-2010 at 07:14

@chemolo

>Very nice!
>I don't like the method of drying the sample with NaOH (to remove HClO4), at least use KOH!
>Why do you need 80 deg C for reaction (complexation) of NAG with copper perchlorate?

The complex seems to be quite soluble in water. Thats why I wanted to minimized amount of water and used hot saturated solutions.

>Impressive that the Cu-NAG-ClO4 crystals are stable at air! That is very unusual!

The reasoning was that if hydrazine complexes of perchlorate are unstable, maybe the condensation product of hydrazine with nitroguanidine would be more stable. Also I figure that the bromate complex would DDT even easier - comparing en-copper-perchlorate (deflagrates) with en-copper-bromate (detonates in smallest amounts). Unfortunately it is quite temperature sensitive and cannot be made from hot saturated solutions.

>Pb(C6H2N3O7)2 - please clarify for the uninformed what it is.

Well think about it: C6 suggests a benzene ring, the N3O6 amount for 3 nitro groups. Its lead picrate :D

[Cu(en)2]3[Co(NO2)6]2 or something...

Bear_with_vodka - 17-1-2010 at 07:22

when completly dry, it decomposes without flame:(
just somewhat stupid decomposition, although there a lot of gas
200mg


Attachment: 9090908.avi (441kB)
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[Cu(en)2](NO3)2

Bear_with_vodka - 17-1-2010 at 07:52

forgive the dirty bowl:)
100mg overall

Attachment: copper en no3.avi (1.2MB)
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Burning of _finely_powdered_ [Cu(en)2]3[Co(NO2)6]2 is much more fun:D
250mg



Attachment: Cu-Co.avi (666kB)
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666kB:D


[Cu(en)2](ClO4)2
if I'm bothering with those small videos, feel free to tell me that:)
200mg, finely powdered



[Edited on 18-1-2010 by Bear_with_vodka]

Attachment: Cu en ClO4.avi (702kB)
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Taoiseach - 30-1-2010 at 12:45

Concerning the lead chlorate complexes, I have to report that a mixture of lead(II)acetate, sodium chlorate and glycerine doesn't precipate shit upon addition of ethanol :( The patent maybe isn't complete bullshit but likely you have to start with lead(II)chlorate to form these complexes :(

woelen - 31-1-2010 at 06:46

I tried making Ni(NH3)6(ClO4)2 as mentioned further above in this thread. I found that making this compound is more succesful if the solutions are not too strong.

I did two experiments.

In the first experiment I added excess 5% ammonia to a solution of nickel sulfate. This gives a blue/purple solution. To this I added a fairly dilute solution of ammonium perchlorate, but not enough to precipitate all of the nickel complex. When this is done, then at first nothing seems to happen, but after a few seconds the solution turns opaque and then quickly a glittering crystalline precipitate settles at the bottom. The crystalline precipitate is compact and can easily be separated from the liquid, simply by decanting the liquid from the precipitate. The liquid above the precipitate still is pale blue/purple. I also could easily rinse with some cold distilled water and I obtained a light blue/purple solid. The rinse-water also was pale blue.

In the second experiment I added a slight excess of 5% ammonia to a solution of nickel sulfate. This gives a blue/purple solution. To this I added a fairly large excess amount of a moderately concentrated solution of ammonium perchlorate. This results in quick formation of a slimy precipitate which only settles very slowly. The liquid above the precipitate is completely colorless. Even after a hour the precipitate still did not settle (still appr. 60% of liquid contained the precipitate). I did not attempt to isolate this slimy precipitate, it simply is too much of a hassle.

The other crystalline batch was dried and tested in a flame. The dry material easily can be ignited and it burns very fast with only moderately strong light output and a hissing sound. Only a little amount of smoke is produced.

Personally I find this nickel-complex moderately interesting. It is energetic but it is not really impressive.

Anders Hoveland - 22-6-2010 at 21:25

Unsymetric dimethylhydrazine perchlorate mixed with Au(ClO4)4- anion dissolved in an inert ionic liquid, might form a coordination complex. Au(ClO4)3 would precipitate, the extra ClO4- makes it soluble. The protonation of the amine group and the two methyl groups on the other side should protect the hydrazine from oxidation by the gold compound. Although Au(ClO4)3 is a sensitive explosive, the complex with the DMhydrazine should sterically hinder any interaction of the Au+3 with the perchlorates, greatly stabilizing the perchlorates. This would be an extremely powerful salt.

mewrox99 - 24-7-2010 at 17:08

How do you make TACPC (Tetraminecopper(II) Perchlorate

I have CuO, CuSO4, NH3, and NH4ClO4

Thanks

Taoiseach - 26-7-2010 at 01:05

TACP is badly soluble in EtOH so you could try

CuSO4 + 2NH4ClO4 + 4NH3 ---(ethanol)---> [Cu(NH3)4](ClO4)2 + (NH4)2SO4

Any other route would involve making copper(II)perchlorate first. Unfortunately I dont see how you can make this compound starting with NH4ClO4. Ammonium carbonate decomposes at 60-70°C thus

2NH4ClO4 + CuCO3 ---> (NH4)2CO3 + Cu(ClO4)2 ---> NH3 + CO2 + H2O + Cu(ClO4)2

seems possible at first glance. However I'm afraid that boiling copper carbonate with ammonium perchlorate will simply ecompose the copper carbonate into black copper(II)oxide.

mewrox99 - 26-7-2010 at 01:40

What about tetraminecopper(II) persulfate

Saw a video once bout it wasn't in english.

I have ammonium persulfate but can make sodium persulfate if needed

Taoiseach - 26-7-2010 at 03:11

Doesnt matter if u use sodium or ammonium persulphate. You could easily make this compound from the chemicals u have but I'd advise against it. The complex is very unstable and decomposes within a few weeks. Storing in a glass bottle is potentially hazardous. The compound dehydrates easily in warm air, and then it is becomes a very sensitive and unstable low-powered explosive. To sum it all up: its useless dangerous crap.

Try precipating TACP with ethanol, I guess it'd work. This complex is stable, stores indefinetly. It comes in hydrated form and is absolutely safe to handle. When dehydrated it is quite powerful and sensitive tough; according to literature it has initiating power.

mewrox99 - 26-7-2010 at 14:11

TACP is tetraminecopper(II) perchlorate not persulfate right?

I'll try making it when I come home.

Engager - 27-7-2010 at 15:52

Quote: Originally posted by mewrox99  
How do you make TACPC (Tetraminecopper(II) Perchlorate

I have CuO, CuSO4, NH3, and NH4ClO4

Thanks


This compound is pretty simple to make. First react copper sulfate with excess of ammonia to get solution of [Cu(NH3)4]SO4, product can be separated by adding several volumes of ethanol.

CuSO4(aq.) + 4NH3(aq.) => [Cu(NH3)4]SO4(aq.)

Sulfate can be easily replaced by perchlorate by action of Ca/Sr/Ba perchlorates, witch can be easily produced by dissolving corresponding oxide/hydroxide/carbonate in commercially available perchloric acid and evaporating solution to dryness. Dissolve some pure tetramminocopper sulphate in water and add solution of Ca/Sr/Ba perchlorate, corresponding sulphate will be precipitated leaving pure solution of tetraaminocopper perchlorate:

[Cu(NH3)4]SO4(aq.) + Ba(ClO4)2 => [Cu(NH3)4](ClO4)2(aq.) + BaSO4 (precipitate)

Evaporation of this solution will give you pure tetraaminocopper perchlorate.

DACP

Taoiseach - 1-8-2010 at 10:28

DACP supposedly stands for "double azido cobalt perchlorate", the correct formula is trans-tetraamminediazido-cobalt(III) perchlorate [(NH3)4(N3-)2Co(3+)](ClO4-). Whats interesting about this complex is the fact that it contains azide ligands, fuel and an oxidizer in a single compound.

It is briefly investigated in a freely available paper "Lead Azide Replacement Program NDIA Fuze Conference" by Bichay and Hirlinger. It is also discussed in more depth in several Chinese publications. I cannot download nor understand these papers but the google translation of the abstracts gives some valuable hints. The compound is a powerful initiating explosive with properties comparable to BNCP (bis-nitrotetrazolate-cobalt-perchlorate) and has been proposed as replacement for both BNCP and lead azide.

A German paper "light absorption of azido-cobalt-ammines" is available in reference section and sheds more light on this family of complexes.

Experimental:

2,6ml 70% perchloric acid was neutralized with basic cobalt carbonate and filtered. Another 1ml of perchloric acid was neutralized with ammonia and the solutions combined. The resultung solution was a light magenta color .

In another beaker, 4g sodium azide was dissolved in 50ml water and 9ml 25% ammonia was added with stirring. The solutions where then combined. Complexation occured, as evidenced by deep blue coloration of the liquid.

A stream of air was then drawn through the solution for 3 hours. During this time the color changed from a deep blue to a lavender color. After standing for several hours, a microcrystalline powder had settled which was filtered and air dried.




dacp2.jpg - 7kB

woelen - 1-8-2010 at 11:35

Did you make the compound yourself? Is this picture from a sample you made yourself? It sounds quite interesting. I however am missing a test on its sensitivity and explosive properties (such as igniting the compound or trying to hit with a hammer).

I'll consider doing the experiments at a small scale (I have all required chemicals and could do the experiment at a 10 times smaller size for safety), but if you have done the experiment yourself then some extra tips from you are welcomed.

[Edited on 1-8-10 by woelen]

Taoiseach - 1-8-2010 at 12:04

Yes I made this very small sample myself. The compound is almost insoluble in water so preparation was a breeze. It is annoying to filter tough; probably it is a good idea to heat the solution and cool down slowly to grow larger crystals.

The first time I tried the reaction I tried to use H2O2 to oxidize the Co(II). This works nicely to make other cobalt(III) complexes. In this case however it turned the solution dark-brown and the resulting precipate was not explosive (tough it did burn). I figure H2O2 destroys the complex and air must be used for its preparation.

As for explosive properties: It can be handled safely IMHO. I tested friction sensitivity and it is very low. It crackled a bit when rubbed with the head of a hammer on concrete ground but even picric acid does that. Flares off upon flame contact. The wee bit of confinement given by a few layers of aluminium foil make it DDT in smallest amounts.

I believe this is a very interesting class of energetic coordination compounds. DACP is very badly soluble and thus there is a good chance it can be prepared from a soluble cobalt(II) salt, any soluble perchlorate and sodium azide. I believe the bad solubility is partly due to the perchlorate ion. For example, hexamminecobalt(III)perchlorate, hexamminecobalt(II)perchlorate, hexamminenickel(II)perchlorate, pyridinecopper(II)perchlorate are all insoluble in water, and tetramminecopper(II)perchlorate is badly soluble in cold water.

It might be possible to form similar azido-cobalt(III)-perchlorate complexes with more energetic fuel kind of ligands. Hydrazine, hydroxylamine, aminoguanidine, nitroguanidine, nitraminoguanidine pop to mind...

Hoveland - 1-8-2010 at 23:55

Quote: Originally posted by Taoiseach  
Unfortunately I dont see how you can make this compound starting with NH4ClO4. Ammonium carbonate decomposes at 60-70°C thus

2NH4ClO4 + CuCO3 ---> (NH4)2CO3 + Cu(ClO4)2 ---> NH3 + CO2 + H2O + Cu(ClO4)2

seems possible at first glance. However I'm afraid that boiling copper carbonate with ammonium perchlorate will simply <d>ecompose the copper carbonate into black copper(II)oxide.


A solution of Copper Perchlorate can easily be boiled without any decomposition, it will not even be oxidizing. Ammonium Carbonate begins to decompose when mixed with water, let alone being heated in water, so it is fairly safe to say that the solution need not even be boiled, just heated. Your reaction seems very reasonable. I have made Cu(NO3)2 this way.

"Ammonium bicarbonate decomposes at 36 to 60 °C into ammonia, carbon dioxide and water vapor in an endothermic process (as it is with many ammonium salts) and so causes a drop in the temperature of the water." (wiki)

[Edited on 2-8-2010 by Hoveland]

Rosco Bodine - 2-8-2010 at 07:26

Seeing the DACP, it would definitely be interesting to look also at nickel complexes and possible nickel styphnate multiple salts as a possibility.

woelen - 2-8-2010 at 11:59

I also tried to make the cobalt complex. I proceeded as follows:

Take a big spatula of cobalt carbonate
Add a small amount of 60% perchloric acid and heat a little. All of the cobalt carbonate dissolves, a pink/rose solution is obtained and also quite a lot of a crystalline mass of solid cobaltous perchlorate. On cooling down a reddish brown slurry is obtained, hardly any liquid remains.
To this I added a large excess amount of 10% NH3. When this is done, a blue precipitate is formed, which is air-sensitive. Also a white crystal meal is obtained, the latter most likely is NH4ClO4. After a while, the blue color changes to pink/brown. This is not due to aerial oxidation, the precipitate just changes color, but this is a known property of many basic cobaltous compounds.
Then I added a concentrated solution of NaN3. When this is done part of the precipitate redissolves and the liquid becomes turbid and pink/purple.
Then I added some more water and some more ammonia and heated for several minutes. Most of the precipitate dissolves again, the liquid becomes beautiful deep purple/red.
On cooling down a dark purple precipitate is formed, consisting of many fine crystals. The color also shifts from deep purple/red to a much duller pink color.

I filtered this precipitate and obtained something which very much resembles what is shown in the picture of Taioseach. Next, I rinses three times with distilled water. When this is done, then the color of the precipitate becomes much darker, it becomes dark purple. Only a very faint smell of ammonia is left after the three rinses.

Now the material is drying. Even the wet material already is capable of burning somewhat, so its energetic properties look promising.

Quantities used in the experiment are not critical at all. No need to do precise measurements of chemicals, just combine them as described above. I assured that there is excess ammonia and excess azide relative to the cobalt. Enough water must be added to assure that all ammonium perchlorate formed in the process is dissolved.

----------------------------------------------------

I however have doubts about that this is a cobalt(III) compound. I hardly had the material in contact with air. I also find the purple color of the compound quite a lot like a cobalt(II) compound.

woelen - 3-8-2010 at 22:37

The sample I now have has changed color considerably. It now is a green/grey solid. When the solid is ignited it quickly burns without visible flame, giving a black smoke. It burns fast with a soft hissing noise and is very easily ignited, also in tiny amounts. No need to keep it over a flame for a long time, even when it comes near a flame it already ignites. When larger pieces of the solid are ignited, then the piece of material jumps around while it is burning, leaving a trail of a very fine black powder.

Taoiseach - 3-8-2010 at 22:54

Doesn't sound like the stuff I have. No guarantee tough I have the correct stuff - the synthesis was simply an adoption of known procedure to make bis-ethylenediamine-diazido-cobalt(III)nitrate.

Whats interesting is that your solution turned purple upon addition of NaN3. In my case it retained a deep-prussian-blue like color. It only turned purple slowly while blowing in air.

I wonder how one can make sure that two azido ligands are fixed. The mono-azido-diperchlorate could form just as well. And it surely would be less energetic.

woelen - 4-8-2010 at 11:29

I carefully reread what you have done. You bubbled air through the mix for 3 hours, that is what I did not do. I just boiled for a while to get all of the material dissolved and then allowed to cool down and form a precipitate. I think that is the main difference.

The compound I made also can easily be brought to explosion. I took a piece of 20 mg and folded this in a small piece of household aluminium foil. When this is put in a flame, then a fairly loud high pitched report is obtained and the aluminium foil is punched.

Taoiseach - 4-8-2010 at 12:12

>When this is put in a flame, then a fairly loud high pitched report is obtained and the aluminium foil is punched.

That means it has fixed at least 1 azido ligand I think. The ammine cobalt perchlorates dont DDT that easily.

woelen - 4-8-2010 at 23:24

Yes, I also think it must have azide in it. The explosion is obtained very easily, just keeping the piece of foil with the solid in it above a flame of an alcohol burner for a few seconds, that's all what is needed.

Finally, I added the last small amount of my green/grey solid to 2M H2SO4. When this is done, the solid quickly dissolves, giving a lot of a colorless and odorless gas and the liquid becomes pink, just like plain cobalt(II). I think that the green solid contains cobalt(III) and that on addition to the sulphuric acid the azide ligands are stripped off and the free cobalt(III) probably oxidizes water to make cobalt(II) and oxygen. On heating, the liquid had the smell of HN3, which to me is exceedingly unpleasant (it gives a very weird intense feeling of fear, just as if I have witnessed something really horrible).

franklyn - 8-8-2010 at 10:06

Characterization and Output Testing of the Novel Primary Explosive, Bis(furoxano)nitrophenol, Potassium Salt

http://www.users.qwest.net/~m-williams/Articles_PDF/AIAA-200...

.

Rosco Bodine - 15-8-2010 at 07:27

Treating a hot nickel perchlorate solution with three equivalents of semicarbazide
solution might be interesting.

It might also be interesting to see if anything interesting may further result from treatment with sodium nitrite.

Mildronate - 15-8-2010 at 07:43

How about Hg and Fe complexes?

[Edited on 16-8-2010 by Mildronate]

Chainhit222 - 27-8-2010 at 21:14

http://pubs.acs.org/cen/science/88/8834sci1.html


oh oh..

but then again, he did put what looks to be a primary explosive in a mortar and pestle.... I would not have tried to break up a large crystal of primary... maybe dissolve it in something and try to precipitate it out.... what the fuck was he thinking...

[Edited on 28-8-2010 by Chainhit222]

Formatik - 29-8-2010 at 10:29

Quote: Originally posted by Chainhit222  
http://pubs.acs.org/cen/science/88/8834sci1.html


oh oh..

but then again, he did put what looks to be a primary explosive in a mortar and pestle.... I would not have tried to break up a large crystal of primary... maybe dissolve it in something and try to precipitate it out.... what the fuck was he thinking...


An hour in the library sometimes saves hours and days of work in the lab, but it can also save extremities. Here was used far too large of an amount of a material which they knew very little about. In Gmelin a citation for a nickel hydrazine perchlorate, [Ni(N2H4)2(ClO4)2], pale-blue precipitate made from standing of aq. Ni(ClO4)2 with N2H4, describes it as extremely explosive and that even in aqueous dilute suspensions it explodes. There is also a reference (H. Ellern, D.E. Olander, J. Chem. Educ. 32 [1955] 24) which claims that the nitrate (which has been investigated as a primary), Ni(NO3)2.3N2H4, as a moist compound occasionally deflagrates spontaneously, but that the dry salt can explode spontaneously.

Attachment: Gmelin, Ni, Tl. C, 110.pdf (1.9MB)
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bbartlog - 30-8-2010 at 04:30

Quote:
what the fuck was he thinking


It's pretty clear if you read the whole article that the guy was pretty much advice-proof. Adviser tells him to make no more than a few hundred mg, he scales it up to 10g. He had numerous unlabeled bottles. He used all of 14 pages of his notebook to describe 3 months of lab work (!!!). Lots of bottles and stuff out in the open, in the photo it looks like his workspace was a mess even before the explosion. And I would bet that his claim of just taking off his safety goggles right at the end was probably BS too.
Still much of the blame has to go to the uni... the existence of sloppy and inexplicably self-destructive people is well-known and there are (or should be) rules and enforcement measures that would have stopped this guy a long time ago.

basic lead guanyl azide styphnate , (exist ?) , theoretical compound

Rosco Bodine - 17-9-2010 at 06:25

A hypothetical compound which may or may not exist has occurred to me, basic lead - guanyl azide styphnate, a possible compound which may be of interest by itself and may further be of interest as a possible substrate for a series of clathrate compounds having higher energy than the already known series of clathrates which are possible using basic lead picrate as the substrate. There has been from time to time mention made of what other possible substrates for clathrates may exist, and may be an improvement over the basic lead picrate, and this purely hypothetical compound, basic lead - guanyl azide styphnate seems a possibility, so I thought I would share this idea, for the entirely speculated compound, which at this point I truly have no idea if it even exists beyond the idea that it seems possible. I offer here my pure speculation for the hypothetical compound, for whatever it may be worth as a possibility or interest for experimentation. :D

[Edited on 17-9-2010 by Rosco Bodine]

holmes1880 - 18-12-2010 at 22:26

Rosco,

To your best knowledge, which is the most friction/impact/heat stable primary that is practical to make? I've read a thread on azo-clathrates, lead azide and its properties, and heard plenty good about silver acetylide. Btw, silver acetylide has only hobbyist-derived information, so annoying.

Which one has the best overall stability?



[Edited on 19-12-2010 by holmes1880]

Rosco Bodine - 19-12-2010 at 03:19

There are several pages in Fedoroff PATR regarding silver acetylide and its double salts, and there are patents and other references, so you are incorrect that there is only hobbyist derived information.

Your question I can only answer in part. I haven't studied or tested to identify one which one would have an average of different properties that represents highest overall stability. Generally with primaries there is a tradeoff of one good property for another and you can't have everything great at the same time for what properties are most desirable. That being true generally leads to binary compositions that are mixtures where the properties of one compound complement another. My experiments are incomplete even concerning my own ideas of things to test. I spend a hundred times as much time reading as actually doing any experimenting. So I have a substantial backlog of ideas to test, most of them having already been posted in case anyone else wants to followup before I get to it. Earlier in this thread I mentioned a binary primary that is unusual in its reported property of low sensitivity to impact, a lead azide - barium styphnate mixture. There are probably several other binary compositions.
http://www.sciencemadness.org/talk/viewthread.php?tid=1778&a...

holmes1880 - 19-12-2010 at 21:42

You're right, DS have been discussed in scientific literature, but they seem to overlook specifics on sensitivity. "Very sensitive to shock and friction" is about all they state. 1kp can be considered very sensitive, but 0.1kp is a whole lot more sensitive.

Powerlabs has shock sensitivity of 2kg--3.4cm, more than TATP, which doesn't seem right.

[Edited on 20-12-2010 by holmes1880]

otonel - 9-1-2011 at 05:20

Can make primary explosive from caseine nitrate. I search information about acsein nitrate from internet but I don`t fiind more information.
I remeber from book from an primary explosive made from milk.

hissingnoise - 9-1-2011 at 07:26

Casein nitrate is energetic but I don't think it's a primary explosive.
I might be wrong though!


Lord Emrone - 18-1-2011 at 03:54

@otonel : That book was probably KIBC ? Remove it from your pc.

DNA - 18-1-2011 at 05:21

I didn't want to start a new topic about this single question so I thought it would be best suited here.
I made lead azide according to the method described on the megalomania website.
I made a blasting cap with it as follows:
Drinking straw which is closed on one end with hot glue.
Pb(N3)2 is pressed in the drinking straw and a little bit of blackpowder is placed on top with a fuse in it, and a little bit of tissue paper to hold it in place.
Then some tape was put over the top (over the tissuepaper).
When the fuse was lit it only burned with a bit of sound, like pressed blackpowder would burn....

Anyone got an idea what is wrong with it?

[Edited on 18-1-2011 by DNA]

nitro-genes - 9-10-2011 at 12:01

Dextrinated LA under slightly alkaline conditions alters crystal size and shape to produce a 92-97% pure material that is less sensitive to static, friction and impact, but unfortunately also less prone to make DDT.

There are 3 solutions:

1. Provide better confinement. It is inadvisable to use a drinking straw anyway, since it can easily deform while manipulating it with a sensitive primary inside.
2. Use a hotter burning pyrotechnic composition to cook it off.
3. Use a very small amount of an initiator like the silver acetylide/nitrate double salt. (Commercially lead styphnate has been used.)

A fourth option is to use collodial LA, by instantly precipitating LA from solutions of a lead salt and NaN3, but this is rather sensitive material. Better improve your cap design...

Formatik - 27-10-2011 at 17:02

Recently made tris(hydrazine)zinc nitrate (ZHN). It is a very shiny, glittering compound (every crystal glitters) with a beige tint. I made it by mixing Zn(NO3)2 aqueous from zinc and nitric acid, with conc. aq. NH3 which was added until the Zn(OH)2 precipitate redissolved and the liquid took on a light yellow color to form a clear solution, then added about the same volume of 55% N2H5OH and boiled it on low boil until crystals started appearing where the boiling was stopped. Then let it crystallize, washed it several times with alcohol. The compound is also mentioned in Gmelin.

It is an insensitive initiator, it hasn't exploded in any heating tests done on small amounts, like over a flame, in flame directly, etc. It doesn't even flash like tris(hydrazine)nickel nitrate (NHN) does. But it does burn very fiercely, very nicely actually even though it leaves behind a lot of white ZnO. It burns splendid with an interesting bluish-white-orange flame. Sprinkling it onto aluminum foil and wrapping this up into a fuse caused it to burn completely to the end when one end was lit. It does have some energetic properties, but needs to basically be blasted itself.

Also tried out tris(hydrazine) nickel nitrate: when exposed to open flame it flashes like mercury fulminate, but this compound didn't explode if heated over flame in small amounts, the best detonators like silver azide and tetrazolates detonate and blow holes through aluminium sheets. Not even under confinement could I get a detonation in small amounts despite how it flashes.

Tetrazole detonators can easily initiate both of these insensitive materials. 0.02g mercury diazoaminotetrazolate (HgDTZ) initiated 0.08g both NHN and ZHN when underneath them and unconfined, in both cases a small amount of complex powders spread since it was unconfined. But the tests showed the holes these two made against aluminium sheets were far larger than the 0.02g HgDTZ made alone. The two holes were close in size to one another.

Both hydrazinate compounds would be good for demonstrations, the nickel salt for its unique color and flashes, and the zinc salt for its glitter and fierce burning and interesting flame.

Rosco Bodine - 27-10-2011 at 17:19

Interesting. Did you try conversion of the tris(hydrazine)zinc nitrate (ZHN) to a perchlorate derivative by mixing with 2 mole equivalents of ammonium perchlorate. I am thinking that a double decomposition may result leaving ammonium nitrate in solution while precipitating tris(hydrazine)zinc perchlorate (ZHP) but I have no solubility or other data to confirm the validity of this idea.

Primary resistant to impact

AndersHoveland - 27-10-2011 at 17:39

Nickel hydrazine nitrate is a potential primary explosive that shows good resistance to impact, while still being easily initiated by flame. This means that the complex salt may likely be safer to handle than other primaries. The sensitivity to impact (the hammer drop height value at which the sample has a 50% chance of detonation) is 84cm. It is resistant to friction up to 10N, resistant against electrostatic discharge, but is sensitive to flame and will explode in contact with a red hot wire.

Another advantage, small amounts of the compound do not need confinement to detonate. Nickel hydrazine nitrate shows excellent initiating power for a primary. The detonation velocity 7km/s.

One note of warning, nickel hydrazine perchlorate is dangerously sensitive to friction. An accident involving only 5 grams cracked a fragment off a laboratory table and resulted in severe injury.

Here is a video:
http://wn.com/Nickel_Hydrazine_Nitrate

Procedure:
An aqueous solution of nickel nitrate was prepared, containing 8% Ni(NO3)2 by weight. 50mL of the solution was poured into a steel container, which was then heated to 65degC. Separately, 100mL of distilled water was warmed and maintained at around 60degC. Gradually over the period of 30 minutes, 7cm3 of hydrazine hydrate was added into the steel container, simultaneously together with 50mL of the water that had been separately prepared, the remaining water was discarded. The hydrazine hydrate used was somewhat wet to begin with. The color of the reactants in the steel container changed from a bluish tint to purple over the course of the reaction. The reaction was stirred for an addition 10 minutes, maintaining the temperature at 60degC. After cooling to 20degC, the purple colored product was filtered out over two layers of coffee filter paper, washed once with 50cm3 distilled water. The moist caked solid was then partially dissolved in >98% alcohol (50mL ethanol was used), then the alcohol was allowed to evaporate out on an electric hot plate set to only 60degC. The evaporation should be carried out in the dark, but with plenty of ventilation. About 5 hours are required for complete evaporation. From this procedure, about 11 grams of nickel hydrazinium nitrate is obtained, which is a 90% yield. Heating of the reactants/reaction is not in any way necessary, as similar yields were obtained at room temperature, but the product obtained from heating shows better physical properties, as the salt is of a more crystalline form. The crystalline form has a density of about 0.89 g/cm3. The nickel hydrazinium nitrate thus obtained, when gradually heated, explodes at 219degC. The compound appears thermally stable even up to 200degC. Sensitivity (50% probability of explosion using 2kg drop hammer from variable heights) value of 84cm. Velocity of detonation: about 7km/sec.

The co-crystallization of NHN with silver azide, such that the resulting clathrate contains 10% by weight of AgN3, increases the drop height sensitivity to a value of 66cm. Even such a clathrate containing only 2% silver azide is not much less sensitive, having a drop height value of 68cm. As a side note, cobalt hydrazinium nitrate, which can be similarly prepared, is even more sensitive, having a sensitivity drop height value of 59cm. The cobalt salt also explodes at a lower temperature, 188degC. Interestingly, however, the cobalt salt is actually somewhat less sensitive to friction than NHN.

Formatik - 27-10-2011 at 17:46

Quote: Originally posted by Rosco Bodine  
Interesting. Did you try conversion of the tris(hydrazine)zinc nitrate (ZHN) to a perchlorate derivative by mixing with 2 mole equivalents of ammonium perchlorate. I am thinking that a double decomposition may result leaving ammonium nitrate in solution while precipitating tris(hydrazine)zinc perchlorate (ZHP) but I have no solubility or other data to confirm the validity of this idea.


No, I haven't. It is insoluble in water, and hot water decomposes it (hot water also slowly decomposes NHN). It is soluble in conc. aq. NH3 though. But I would be weary of forming any kind of perchlorate here considering instability of the copper perchlorate complex and dangerous sensitivity described above on the nickel perchlorate complex. I've seen no data on it.

Rosco Bodine - 27-10-2011 at 17:57

@AndersHoveland
What is the basis for NHN - AgN3 to be characterized as a clathrate?
There are a lot of coprecipitates that are definitely not clathrates, and
that would seem to be the case here .



[Edited on 28-10-2011 by Rosco Bodine]

Formatik - 27-10-2011 at 18:03

Quote: Originally posted by AndersHoveland  
Another advantage, small amounts of the compound do not need confinement to detonate.


Try that for yourself and see if it's true. I couldn't even get it to detonate under confinement.

Note: yes it can explode when heated but these are likely in what I would call at least larger amounts. Probably gram amounts. It has likely a higher critical mass than many other primaries.

[Edited on 28-10-2011 by Formatik]

Rosco Bodine - 27-10-2011 at 18:10

Quote: Originally posted by Formatik  
Recently made tris(hydrazine)zinc nitrate (ZHN). It is a very shiny, glittering compound (every crystal glitters) with a beige tint. I made it by mixing Zn(NO3)2 aqueous from zinc and nitric acid, with conc. aq. NH3 which was added until the Zn(OH)2 precipitate redissolved and the liquid took on a light yellow color to form a clear solution, then added about the same volume of 55% N2H5OH and boiled it on low boil until crystals started appearing where the boiling was stopped. Then let it crystallize, washed it several times with alcohol. The compound is also mentioned in Gmelin.

It is an insensitive initiator, it hasn't exploded in any heating tests done on small amounts, like over a flame, in flame directly, etc. It doesn't even flash like tris(hydrazine)nickel nitrate (NHN) does. But it does burn very fiercely, very nicely actually even though it leaves behind a lot of white ZnO. It burns splendid with an interesting bluish-white-orange flame. Sprinkling it onto aluminum foil and wrapping this up into a fuse caused it to burn completely to the end when one end was lit. It does have some energetic properties, but needs to basically be blasted itself.


Please attach that Gmelin excerpt if you have it.

You describe boiling the solution on a low boil until crystals begin to appear. What I was thinking was to add a boiling hot solution of ammonium perchlorate before the point where a low boil would make crystals begin to appear, seeking possible precipitation of a less soluble perchlorate derivative via double decomposition of the nitrate.

The reaction mixture you describe is already a solution of ammonium nitrate and probably diamminezinc nitrate or diamminezinc hydroxide, so the presence of ammonium perchlorate may cause the perchlorate salt to precipitate.
I know this works for tetramminecopper perchlorate, but haven't tried the hydrazine complex for any of these.
Generally though not always, the perchlorate salt will be less soluble so that will be what precipitates.

[Edited on 28-10-2011 by Rosco Bodine]

AndersHoveland - 27-10-2011 at 18:18

Perchlorate complex salts of hydrazine are dangerously friction sensitive. Hydroxylamine perchlorate is very sensitive, having a drop height value of only 2cm. Hydrazine perchlorate is also much more sensitive than hydrazine nitrate, although I do not know exactly how much. Woelen has demonstrated that it can at be heated over a flame in a metal spoon without detonation, but I would not want to risk trying this. It seems that compounds that contain perchlorate together with either hydrazine or hydroxylamine (which are reducing agents) are dangerously friction sensitive. I would wonder about nitroformate salts instead of perchlorate, since hydrazinium nitroformate is apparently stable enough for use as a rocket oxidizer. Nitroformates are more powerful than nitrates, and almost as powerful as perchlorate. However, the sensitivity and stability of complex transition metal salts (especially Cu, Ag, Hg, Pb) of nitroformate may be completely different, due to the unique chemistry of the nitroformate anion.

Whereas trinitromethane, even and especially when dissolved in an organic solvent is dangerously sensitive, its nitroformate salts, including even those of hydrazine, are much more stable. One difficulty of synthesis may arise in the fact that trinitromethane itself will oxidize hydrazine (or its salts). The trinitromethane must be added in the form of its nitroformate salt.

[Edited on 28-10-2011 by AndersHoveland]

Rosco Bodine - 27-10-2011 at 18:33

@AndersHoveland
Zinc diammine Perchlorate is actually used as a phase stabilizer at 1.8% content in NH4NO3 oxidizer based solid rocket propellants, according to US5071630

You often make generalizations or even specific statements about things which are simply inaccurate, the dubious clathrate is a recent example, and you continue opining here about the hydrazinium complexes, making generalizations which are somewhat or altogether dubious generalizations.

You go through these threads with a broad brush pseudointellectualizing about things you really don't know, and I for one wish you would stop it with the pretended knowledge, and qualify what you think by just saying, "I think" ...but it may not be so, as a kind of disclaimer :D

Attachment: US5071630 Zinc diammine Nitrate.pdf (489kB)
This file has been downloaded 602 times


AndersHoveland - 27-10-2011 at 18:40

Quote: Originally posted by Rosco Bodine  
Zinc diammine Perchlorate is actually used as a phase stabilizer at 1.8% content in NH4NO3 oxidizer based solid rocket propellants,


Zinc diammine Perchlorate is an ammonium complex, not a hydrazine complex. There are no compatibility problems between perchlorate and ammonia. Indeed, ammonium perchlorate is very stable, virtually impossible to detonate by itself.

Rosco Bodine - 27-10-2011 at 18:46

Do you have specific information on the hydrazinium complex of zinc perchlorate?

It is possible it could be unacceptably sensitive and it is possible it may not be unacceptably sensitive. Unless there is a specific reference for it, then why make such generalizations ? The hydrazinium zinc perchlorate is described in the following references, which may provide some reliable information.

W. Friederich and P. Vervoorst, SS 21, 49 (1926)

SS = Z. ges. Schiess-Sprengstoffw.

CA 21, 1184, 1927

Semicarbazide perchlorate complexes could also be interesting ....but there is nothing I can find about them.

[Edited on 28-10-2011 by Rosco Bodine]

Formatik - 27-10-2011 at 19:55

Quote: Originally posted by Rosco Bodine  
Please attach that Gmelin excerpt if you have it.


It's attached below. The original description goes: a clear solution of Zn(NO3)2 in NH3 and N2H5OH is heated on the water bath, causing a brown coloration and separation of a skin of black metal particles (I didn't see that). Upon cooling the compound separates as a white crystalline powder (it's whitish). The Franzen and Mayer reference also mentions it burns with a green flame, I suppose there is some green in its flame also. It burns brilliantly once ignited.

Quote:
You describe boiling the solution on a low boil until crystals begin to appear. What I was thinking was to add a boiling hot solution of ammonium perchlorate before the point where a low boil would make crystals begin to appear, seeking possible precipitation of a less soluble perchlorate derivative via double decomposition of the nitrate.


Maybe it could work, but I wouldn't be willing to risk the glassware on it.

Attachment: Gmelin, Zn, 156.pdf (386kB)
This file has been downloaded 626 times

Rosco Bodine - 27-10-2011 at 20:19

Thanks for the Gmelin page.

You could put your hot NH4ClO4 solution in a foam coffee cup and dribble in the other hot component solution .....no glass at risk.

Another possible way of doing this would be to react Calcium Hydroxide in slight excess with Ammonium Perchlorate and heat to boiling to drive off the ammonia, filtering the residual solution of Calcium Perchlorate. React that solution of Calcium Perchlorate with Hydrazine Sulfate and filter out the Calcium Sulfate to leave a solution of Hydrazine Perchlorate. Then you could add your metal complex candidate as an ammonia solution of the metal nitrate.

Formatik - 9-11-2011 at 21:14

Here are a few deflagrations of the aforementioned hydrazine compounds:

Tris(hydrazine)nickel nitrate (NHN):


nhnpic.png - 495kB


Deflagration to heat and flame:


nhn.png - 173kB
Videos:

http://www.megaupload.com/?d=65O1W3J9
http://www.megaupload.com/?d=YHCILN1Q


Tris(hydrazine)zinc nitrate (ZHN):

zhnpic.png - 600kB


Deflagration to heat and flame:


zhndef.png - 358kB

Videos:

http://www.megaupload.com/?d=1GNTZZVG
http://www.megaupload.com/?d=65O1W3J9


The same aforementioned fuse of ZHN is also in one of the videos above. It burned to the end. I'm not sure how reliable it is yet though. And around 30mg of the ZHN intermixed with about 30mg Ag2C2.AgNO3 into a small pile on an iron plate gave a very loud report on contact with a long wooden match flame (originally Philou's idea to mix the nickel salt with acetylide complex).

Concerning hydrazine toxicity, this works by inhibiting vitamin B6. Hydrazine drugs also inhibit vitamin B6. Vitamin B6 ingestion can offset some hydrazine toxicity, however too much vitamin B6 ingestion can cause permanent nerve damage so it has to be very carefully dosed. Respirators work pretty good against vapor. But they do nothing to protect the eyes, which hydrazine damages.

Below is the aforementioned reference claiming spontaneous deflagration and detonation of the nickel complex. It sounds to me like they had some kind of nasty contaminant, perhaps in their nickel source (copper maybe) in their lab. I had it sit around for about three weeks and it never misbehaved. Copper would not be readily discernible in their original nickel but would impart a green color to the lilac salt (their impurity maybe more bogus than copper). I've ignited 1.0g in the open and it only flashed, yet they had a violent spontaneous detonation with only 1.5g of the unconfined salt. That's not the NHN I know.

Attachment: J. Chem. Educ. 32 [1955] 24.pdf (1MB)
This file has been downloaded 694 times

[Edited on 10-11-2011 by Formatik]

franklyn - 19-12-2011 at 21:01

Studies of Complex Perchlorates
http://www.dtic.mil/dtic/tr/fulltext/u2/607947.pdf

Studies of Complex Perchlorates (final)
http://www.dtic.mil/dtic/tr/fulltext/u2/634105.pdf


Chlorates & Perchlorates their Characteristics & Uses
http://www.dtic.mil/dtic/tr/fulltext/u2/318741.pdf

Chlorates & Perchlorates Their Manufacture Properties & Uses
http://www.dtic.mil/dtic/tr/fulltext/u2/242192.pdf


Related Post _
http://www.sciencemadness.org/talk/viewthread.php?tid=1081&a...


[Edited on 20-12-2011 by franklyn]

Formatik - 27-12-2011 at 21:51

Some more thoughts on the previous nickel hydrazine nitrate:

The spontaneous explosion of the previous nickel hydrazine nitrate is a complete mystery. Rather than admit the possibility of having impure material in the facility, they speculate the presence of some kind of higher complex, noting an excess of hydrazine, or that residual N2H5OH did some kind of funky stuff. I think I may have read Franzen and Mayer explicitly doubting any higher complex possible. Mine also smelled of hydrazine for a few days, but I doubt the presence of a higher complex. It's just excess hydrazine, hydrazine being hydrazine and lingers around for a few days.

The previous crude test of initiation using HgDTZ may not have worked well at all, because I didn't think of the action of a powder on top of another simulating the effect of confinement, but it does. Initiation of ZHN and NHN was done with 0.3g PbN6 in the Journal of Hazardous Materials 171 (2009) 1175–1177, which showed maximum pressure generation of 86.9% and 105.0% of TNT for the zinc and nickel salts respectively.

franklyn - 28-12-2011 at 14:32

Random Idea

Hypofluorous Acid (HOF) CAS 14034-79-8
Formed much the same as the Chlorination of water H2O + Cl2 => HOCl + HCl
except it only forms on ice and is only barely stable at cryogenic temperature.
If mixed with Nitrogen Trifluoride may perhaps reform as
TetraFluoroAmmoniumoxyhydride (NF4OH)
The Lewis structure is perfectly satisfied manifesting a TetrafluoroAmmoniumoxy
cation and hydride anion.

Tetrafluoroammoniumoxyhydride.GIF - 2kB

http://en.wikipedia.org/wiki/Hypofluorous_acid
http://accessscience.com/content/Hypohalous-acid/334300
http://pubs.acs.org/doi/abs/10.1021/ja00327a016
http://www.sciencedirect.com/science/article/pii/00404020778...
http://www.ncbi.nlm.nih.gov/pubmed/18355062
Properties
http://www.chemeo.com/cid/26-601-5

.

Adas - 29-12-2011 at 03:34

Formatik, please, upload your videos on MediaFire next time if you can. It is free and I must not wait for download and all that crap. Thanks :)

Nice videos, btw.

[Edited on 29-12-2011 by Adas]

Formatik - 29-12-2011 at 09:05

Thanks. Ah yes, mediafire is much better.

double salt lead bromate acetate

AndersHoveland - 3-1-2012 at 17:56

Here is an obscure explosive double salt. It is doubtful that it has any practical applications, probably having poor chemical stability in storage because of the bromate.

http://www.youtube.com/watch?v=_Xx-zhALHjM

Quote:

lead acetato-bromate Pb2(C2H3O2)2(BrO3)2 precipitate deflagrates with a yellowish-white puff of smoke. hammer drop height sensitivity: 15-20cm. 175 g of KBrO3 are dissolved in 1.5 L hot water, and mixed with a solution of 175 g acetic acid (100%) and 260 g lead acetate hydrate in 2 L water. The solution remains clear initially. It is then filtered, cooled down and seeded with a few crystals of lead bromate while rubbing the side of the glass vessel with a glass rod. Soon after, the solution becomes turbid, and yields a heavy crystalline precipitate, which does not increase anymore after 12 hours in the cold. The solution above the precipitate is decanted off or extracted out, and discarded (possibly add some Na2CO3 to salvage the remaining lead in the waste solution, in the form of PbCO3). The crystalline precipitate is washed with cold water until it is free of acetic acid and Na/K acetate/bromate, then is hot dry air is passed over until free from moisture. Yield is only 123 g


[Edited on 4-1-2012 by AndersHoveland]

Formatik - 14-1-2012 at 16:19

Making that large of an amount of a primary is never a good idea. This lead salt complex also has a thread on here.

Below is the original reference of Franzen and Mayer mentioned above and some references of similar salts.

Über die Hydrazinate einiger Metallsalze. Franzen, Mayer. Z. Anorg. Chem. 60, 247 (1908).

Verbindungen von Hydrazin mit Quecksilbersalzen. K.A. Hofmann, E.C. Marburg. Ber. 30, 2019 (1897).

Zur Kenntniss der Stickstoffquecksilberverbindungen. K.A.Hofmann, E.C. Marburg. Annalen, 305, 214 (1899).

NatashaJurievna - 5-2-2012 at 09:11

Co-ordination Compounds as Sensitizers for Percussion Cap Compositions
http://www.dtic.mil/dtic/tr/fulltext/u2/a492392.pdf

Most notable, less known compounds:

Thiourea cuprous perchlorate - cuprous cyanamide complex
Behavior on ignition: Deflagrates fiercely with a dark red flame and little smoke.

Basic triethanolamine lead perchlorate
On ignition deflagrated vigorously.
:cool:

NatashaJurievna - 6-2-2012 at 14:36

BTW, German patent DE922216 says about the heavily oxygen deficient basic triethanolamine lead perchlorate that "unter Einschluß detonierende Substanz" - detonating substance under confinement.

Tetramine Zinc Permanganate

AndersHoveland - 12-2-2012 at 17:44

Do not know if this complex salt has been mentioned before, here is a translation from http://www.pirotek.info/Vv2/TeAZM.htm

Quote:

Tetramine Zinc Permanganate - [Zn (NH3) 4] (MnO4) 2 – is an ammonia complex salt, which is a relatively weak but also sensitive initiating explosive. It has an oxygen balance of -4.3 - 8 6% (since the reaction products contain a mixture of both MnO and Mn2O3), and the volume of detonation gases about 483 l / kg.


Tetramine Zinc Permanganate is unstable in storage, gradually decomposing with the loss of ammonia, resulting in a characteristic odor, accompanied by self oxidation-reduction. The latter reaction is accompanied by the release of nitrogen, water and ammonia and, because this reaction is autocatalytic, if not properly stored it can lead to an explosion, both because of the ignition, and just because of the increase of pressure in a sealed vessel. Keep substance preferably at low temperature in a dry atmosphere, as in the presence of moisture it is less stable than ammonium permanganate.

The solubility of Tetramine Zinc Permanganate in water is very small - the reaction of its formation is used in analytical chemistry (microcrystalline microscopy) for the detection of zinc (with a detection limit of 1:7000 ). As a consequence, from a solution of salt precipitated in the form of fine black powder, shiny needle crystals are formed only with careful crystallization from relatively dilute solutions. In general, due to low solubility, synthesis of Tetramine Zinc Permanganate is simple.

Required equipment: chemical glasses (or better disposable plastic) , refridgerator (or freezer), and desirable but optional, an electric mixer and filter.

Materials needed: potassium permanganate, zinc oxide and zinc chloride (or nitrate), ammonium chloride (or nitrate), 10 - 30% aqueous ammonia (ammonia), distilled water.


Stages of Preparation:

1) In 100 ml of distilled water dissolve 4g KMnO4; to accelerate the dissolution it is recommended to use an electric mixer;

2) to 8 ml 20% aqueous ammonia is added 1.1 g ZnO, and 1.4 g NH4Cl (or NH4NO3); zinc oxide should dissolve completely, instead of zinc oxide and ammonium salts can take an equimolar amount of the corresponding zinc salts;

3) The prepared solutions were mixed and rapidly cooled before freezing;

Note: The precipitation of Tetramine Zinc Permanganate takes place during mixing of solutions, cooling slows the recovery of permanganate ion and free ammonia yields a relatively pure product.

4) The bulk of the cold solution is poured, and the residue washed precipitate on the filter, the solution to further remove the filter with the precipitate on the stack of filter paper sheets, after which the product is loosened as soon as possible and dried at low temperature (optionally under reduced pressure).

The result is a 4 – 4.5 g of fine black powder (yield of permanganate 85 - 95%), the presence in it a brown powder of manganese oxides indicates that the product is partially decomposed by high temperature or duration of precipitation and drying.

In principle, the use of dilute (and not contaminated colloidal particles), the starting solution can be precipitated more pure Tetramine Zinc Permanganate in the form of small brilliant needles, but in this case the output drops sharply.

Formatik - 13-2-2012 at 00:25

There are some salts of salts of azotetrazole likely ammines which hold more potential I would think. Adding tetraamminecopper(II) nitrate powder to aq. sodium azotetrazole under stirring affords a bright green solid which is uncomparably mechanically far less sensitive than the extremely sensitive copper azotetrazole, but also less brisant.

Quote: Originally posted by NatashaJurievna  
BTW, German patent DE922216 says about the heavily oxygen deficient basic triethanolamine lead perchlorate that "unter Einschluß detonierende Substanz" - detonating substance under confinement.


The copper salt appears to hold energetic properties also. It was obtained from addition of perchloric acid onto aq. triethanolamine-copper complex and drying the liquid over H2SO4, left a deep blue solid which flashed violently with a wonderful blue light when heated over a flame. The experiment was done several years ago and perchloric acid may have been present in excess.

AndersHoveland - 14-2-2012 at 00:42

Quote: Originally posted by Taoiseach  
DACP supposedly stands for "double azido cobalt perchlorate", the correct formula is trans-tetraamminediazido-cobalt(III) perchlorate [(NH3)4(N3-)2Co(3+)](ClO4-). Whats interesting about this complex is the fact that it contains azide ligands, fuel and an oxidizer in a single compound.

It is briefly investigated in a freely available paper "Lead Azide Replacement Program NDIA Fuze Conference" by Bichay and Hirlinger. It is also discussed in more depth in several Chinese publications. I cannot download nor understand these papers but the google translation of the abstracts gives some valuable hints. The compound is a powerful initiating explosive with properties comparable to BNCP (bis-nitrotetrazolate-cobalt-perchlorate) and has been proposed as replacement for both BNCP and lead azide.

A German paper "light absorption of azido-cobalt-ammines" is available in reference section and sheds more light on this family of complexes.

tetraminecobaltazideperchlorate.jpg - 12kB

Experimental:

2,6ml 70% perchloric acid was neutralized with basic cobalt carbonate and filtered. Another 1ml of perchloric acid was neutralized with ammonia and the solutions combined. The resultung solution was a light magenta color .

In another beaker, 4g sodium azide was dissolved in 50ml water and 9ml 25% ammonia was added with stirring. The solutions where then combined. Complexation occured, as evidenced by deep blue coloration of the liquid.

A stream of air was then drawn through the solution for 3 hours. During this time the color changed from a deep blue to a lavender color. After standing for several hours, a microcrystalline powder had settled which was filtered and air dried.


This has to be the most interesting complex salt that has been mentioned in this thread.

There must be variation of this compound that is even better.
Perhaps hydrazine instead of ammonia? I would be hesitant to make a perchlorate complex hydrazine salt after reading about the accident with cobalt hydrazine perchlorate, which suggests dangerous friction sensitivity.
http://www.sciencemadness.org/talk/viewthread.php?tid=14587
It is not just complex salts, because hydroxylamine perchlorate is very sensitive (drop height 2cm). But hydrazinium nitroformate apparently is not very sensitive, as it has been researched as a solid rocket fuel oxidizer. But nitroformate probably has different compatibility issues than perchlorate. Ammonium nitroformate appears to be the most stable nitroformate salt.
Quote:

Silver nitroformate slowly decomposes at room temperature, and is a very sensitive explosive. The potassium salt, KC(NO2)3 is a lemon yellow crystalline solid that decomposes slowly at room temperatures and explodes above 95 °C. Ammonium nitroformate deflagrates or explodes above 200 °C. Hydrazinium nitroformate is thermally stable to above 125 °C. Most salts of trinitromethane derive from the aci-form. However, the silver and mercuric salts exist in two forms: colourless and yellow. This may indicate that two forms of these salts - nitro and aci - can exist.

So it would seem nitroformate would be stable in the presence of hydrazine, but not with certain metals (probably including copper).

Quote: Originally posted by Formatik  
Adding tetraamminecopper(II) nitrate powder to aq. sodium azotetrazole under stirring affords a bright green solid which is uncomparably mechanically far less sensitive than the extremely sensitive copper azotetrazole, but also less brisant.

This is good to know. I would have been a little concerned about whether nitroformate would be ligand towards Co+3, as this could result in dangerous sensitivity (nitroformate is stabilized by being an anion). But apparently complexed ammonia molecules make the transition metal ion less "noble" in character. (difficult to explain, hopefully you understand what I mean)

Nitroformate salts are usually bright yellow in color, but might not contribute a any color if the nitroformate acts as a ligand rather than anion, which would also presumably make it more sensitive.

Perhaps...
Hydrazine-diazido-cobalt nitroformate ?
(N2H4)2(N3)2Co[+] C(NO2)3[-]

Not sure what color it would be, but probably would have a grayish tone from being a mix of yellow and light purple. But difficult to speculate, as complex cobalt salts can take a range of colors.

Complex cobalt nitroformate salts have already been prepared:

Journal of Structural Chemistry
Volume 32, Number 3, 339-344, DOI: 10.1007/BF00745741
Crystal structure of (4-amino-1,2,4-triazole)pentaamminecobalt(III) nitroformate [Co(NH3)5 (C2H4N4)][C(NO2)3]3
N. V. Podberezskaya, N. V. Pervukhina and V. P. Doronina
http://www.springerlink.com/content/q50434k7334m1wt6/

ethylenediamine copper(II) nitroformate
[C2H8N2]2Cu[C(NO2)3]2
Li Yang, Jin Zhang, Tonglai Zhang, Jianguo Zhang, Yan Cui
Beijing Institute of Technology, China


A quick note about the chemistry of cobalt, solutions of Cobalt(II) salts in water are not oxidized chlorine, but the presence of ammonia renders the cobalt much more vulnerable to oxidation, even by air.

It might actually be fairly complicated to prepare a hydrazine cobalt(III) salt, since non-complexed cobalt(III) salts generally do not exist. Oxidizing the cobalt(II) with hydrazine already in the solution would also likely lead to oxidation of the hydrazine. Perhaps Co2O3 reacted with a solution of hydrazinium nitroformate, hydrazinium azide, and some additional hydrazoic acid would work.


Here is what the decomposition formula for the compound would be:

(N2H4)2[N3]2Co[C(NO2)3] --> CoO + 4 H2O + CO2 + 6½ N2

[Edited on 14-2-2012 by AndersHoveland]

Formatik - 15-2-2012 at 20:39

Yes, the ammine does desensitize the original energetic compound, though it will also take away from its power unless an oxidizer or high nitrogen is present. Another example of large sensitivity reduction is the complex of tetraamminecopper(II) nitrotetrazolate, a fairly soluble blue solid with impact sensitivity of manyfold less than copper (II) nitrotetrazolate. Copper nitrotetrazolate is a frightening material to work with anywhere near the dry state.

Speaking of cobalt salts, cobalt diazoaminotetrazole has very low sensitivity to mechanical action (in strong contrast to Hg, Cu, Ag, Pb salts which are highly sensitive), cobalt azotetrazole is also of relatively low sensitivity (again in contrast to the same metals). Though both salts are not that brisant, the latter probably more than the former. There is good potential of forming low sensitive salts with cobalt.

Thinking on the perchlorate above, how about a polyamminecobalt-azido-azotetrazolate, or nickel hydrazininium azotetrazolate. :P Maybe a question too general, what is better for an energetic then, to have an oxidizing moiety or high nitrogen content?

AndersHoveland - 16-2-2012 at 19:01

Yes, I was thinking about the compound with nitrotetrazolate ligands instead of azide.

(N2H4)2[N3]2Co[C(NO2)3]

(N2H4)2[N4CNO2]2Co[C(NO2)3]

But I am not sure if nitrotetrazolate can act as a suitable ligand to replace the azide. This is important because if the two anions have near the same ability to act as a ligand, a double salt would be unlikely to form, and the ions would just precipitate out as separate crystals of two different compounds. There is also nitrotetrazole-2-oxide that exists. So I suppose this would be possible,

(N2H4)2Co[CN5O3]3

One of the advantages of cobalt is that it would bind to the hydrazine very strongly. For example, the ammonia complexed in tetramine cobalt chloride do not come off even if the compound is dissolved in hydrochloric acid, and indeed the complex salt can be crystallized out of a solution of hydrochloric acid. The Co+3 ion is very acidic, one of the reasons that Co(II) salts will not be oxidized in the absence of ammonia. So no worries about hydrazine decomposing out from the complex salt.

Hydrazine is much more energetic than ammonia, I was just thinking about how to combine several strategies together into one compound: a complex double salt with hydrazine, azide, and nitroformate.


Quote: Originally posted by Formatik  

cobalt diazoaminotetrazole has very low sensitivity to mechanical action (in strong contrast to Hg, Cu, Ag, Pb salts which are highly sensitive), cobalt azotetrazole is also of relatively low sensitivity. There is good potential of forming low sensitive salts with cobalt.


Is the cobalt in cobalt diazoaminotetrazole in the +2 or +3 oxidation state? formula?
Also that just plain cobalt azotetrazole, or the tetramine complex with cobalt?
That would be amazing if it were just plain azotetrazole that only was not especially sensitive. Would that not mean that the ammonia complex of that compound would be even less sensitive? Or would there be no room for the ammonia to complex?


Quote: Originally posted by Formatik  

what is better for an energetic then, to have an oxidizing moiety or high nitrogen content?


That is a complex answer. My opinion is that it is best to combine both.

[Edited on 17-2-2012 by AndersHoveland]

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