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woelen
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Right now, I have the clear liquid, spread out in a petri dish in a dry and fairly warm place (40 C or so). I first start with a few ml. Is there any
chance that the material explodes or ignites, when it becomes dry? Or could it be that I make something insanely sensitive like ammoniakal silver,
which explodes on the slightest provocation?
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PHILOU Zrealone
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Quote: Originally posted by woelen  | | Right now, I have the clear liquid, spread out in a petri dish in a dry and fairly warm place (40 C or so). I first start with a few ml. Is there any
chance that the material explodes or ignites, when it becomes dry? Or could it be that I make something insanely sensitive like ammoniakal silver,
which explodes on the slightest provocation? |
Anything is possible!
-Isolation without troubles
-Decomposition
a)smooth via disproportionation of chlorite into chloride and chlorate (40°C might be too high (or not) and favors (or not) this reaction)
b)more vigorous if acidity of EDA chloride(*), dichloride, chlorite(*), dichlorite, chlorate(*), dichlorate is high enough resulting into Cl2, Cl2O,
ClO2 generation
--> self heating and vapourization by autoxydation and/or flame and/or explosion
--> generation of amine oxydation products (hydroxylamine(R-NHOH), nitroso/oxime compound (R-N=O <==> R'=N-OH), nitro compound (R-NO2),
halogenated nitroso/nitro compound, haloamines (R-NHCl or R-NCl2)
Cl2N-CH2-CH2-NCl2 and solid Br brother are relatively stable (vs NCl3 and NBr3) as experienced by Axt...but quite explosive.
Safety notes
-possible light sensitivity (especially to solar UV).
-possible heat sensitivity (40°C is maybe overkill for a first attempt)
-possible friction sensitivity
-possible shock sensitivity
(*)
Mono salts will remain on the safe side since the counterpart of the molecule stil contains a basic group (less efficient than the first because of
first protonation); disalts are more of a concern because of the second protonation and no more base to play(**)
with...
You know the song "salts of a weak base and a strong acid are weak acids" ...
--> what are:
- salts of a strong base (EDA first basic group) and a medium strength acid (HClO2)?
- or salts of a medium strenght base (EDA second basic group) and a medium strenght acid (HClO2)?
- or of stronger acids like HCl and HClO3 in the case of disproportionation
--> TERRA INCOGNITA    = Real research
and discovery(***)
(**)Reminds me the famous song of the ABBA group...with into the lyrics "no more ace to play"
(***)The winner takes it all...
[Edited on 16-12-2016 by PHILOU Zrealone]
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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woelen
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I have left the material to evaporate and dry. A white solid remains behind, but this solid does not become really dry. It remains a little sticky,
more like a paste than a crystalline solid.
I scraped some of the paste from the petri dish and kept it in a flame. It does not show any energetic properties at all. It just boils away when
strongly heated. It also chars a little.
I also did 3 experiments with the remaining material:
- Add 20% H2SO4: A colorless and odorless gas is produced, no smell of Cl2 or ClO2. The liquid becomes colorless and slightly turbid (probably due to
a small amount of remaining lead).
- Add 36% HCl: Again, a colorless and odorless gas is produced. The liquid becomes totally clear and pale yellow. Probably a tiny amount of chlorite
is left, leading to the pale yellow color.
- Add 40% HBr: A colorless and odorless gas is produced. A faint odor of bromine is produced, but this smell only is faint. The liquid turns
completely clear and orange/yellow. This is in stark contrast with a pure chlorite to which some HBr is added. If that is done, then there is a
near-explosion, the chlorite reacts extremely violently with the bromide ions at low pH.
Conclusion: Nearly all chlorite is gone. The colorless gas mostly is CO2. I am afraid that the liquid picked up CO2 from the air, is turned into a
carbonate and that the chlorite simply decomposed to chloride and oxygen. Or maybe part of the EDA is oxidized to some more complicated organic
compound. Anyway, making EDA chlorites in this way probably is not possible. I might try this again, but next time with a slight excess of Pb(ClO2)2
instead of some excess EDA.
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PHILOU Zrealone
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| Quote: Originally posted by woelen | I have left the material to evaporate and dry. A white solid remains behind, but this solid does not become really dry. It remains a little sticky,
more like a paste than a crystalline solid.
I scraped some of the paste from the petri dish and kept it in a flame. It does not show any energetic properties at all. It just boils away when
strongly heated. It also chars a little.
I also did 3 experiments with the remaining material:
- Add 20% H2SO4: A colorless and odorless gas is produced, no smell of Cl2 or ClO2. The liquid becomes colorless and slightly turbid (probably due to
a small amount of remaining lead).
- Add 36% HCl: Again, a colorless and odorless gas is produced. The liquid becomes totally clear and pale yellow. Probably a tiny amount of chlorite
is left, leading to the pale yellow color.
- Add 40% HBr: A colorless and odorless gas is produced. A faint odor of bromine is produced, but this smell only is faint. The liquid turns
completely clear and orange/yellow. This is in stark contrast with a pure chlorite to which some HBr is added. If that is done, then there is a
near-explosion, the chlorite reacts extremely violently with the bromide ions at low pH.
Conclusion: Nearly all chlorite is gone. The colorless gas mostly is CO2. I am afraid that the liquid picked up CO2 from the air, is turned into a
carbonate and that the chlorite simply decomposed to chloride and oxygen. Or maybe part of the EDA is oxidized to some more complicated organic
compound. Anyway, making EDA chlorites in this way probably is not possible. I might try this again, but next time with a slight excess of Pb(ClO2)2
instead of some excess EDA. |
Sad to read those first negative results 
Maybe were 40°C for the evaporation a little too hot for the sensitive compound?
Maybe crashing the salt out of solution with isopropanol or ethanol would speed up things by sparing the evaporation step and the risk of
decomposition by heating?
[Edited on 18-12-2016 by PHILOU Zrealone]
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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woelen
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I decided to revisit this line of experimenting. I now am in the process of making a new batch of Pb(ClO2)2 and I also have an experiment with Co(2+)
and Ni(2+) to which chlorite is used. With cobalt(II) a chocolate-brown precipitate is formed. I am allowing this to settle and hope to isolate some
of the brown material and with nickel(II) I get a pale green precipitate. I also have some of this settling. In a third bottle I have some yellow
Pb(ClO2)2 settling.
Chlorite remains one of my favorites to experiment with. So easy to obtain and at the same time so reactive.
More will follow.
[Edited on 17-2-18 by woelen]
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woelen
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Isolating the cobalt precipitate is very difficult. It is a very very fine dark brown solid, dispersed in the liquid. When it settles, then it does
not form a compact layer of precipitate, but a nearly liquid dark brown layer, which mostly is water with just a small amount of solid matter in it.
This is in strong contrast with the pale yellow precipitate of lead(II) chlorite. The latter can easily be separated. I have it rinsed one more time
to dissolve as much of PbCl2 and then I'll let it dry.
The nickel precipitate is flocculent, much like a precipitate of nickel hydroxide or copper(II)hydroxide. The amount of precipitate only is small and
the supernatant liquid is fairly strongly green, while I did use quite a large amount of sodium chlorite. I have the impression that only a small
amount of nickel(II) is precipitated. Maybe this is just nickel(II) carbonate or hydroxide. The NaClO2, according to its label is 80% by weight, the
balance being NaCl, but it is somewhat alkaline. I think that it also contains a little Na2CO3, although it is not mentioned on the package.
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DraconicAcid
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Hmmmm....nitrite can act as a ligand....have you tried adding a large xs of chlorite to the cobalt or nickel solutions to see if you get coordination?
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Laboratory of Liptakov
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I looking for on S-M how energetics properties has Pb (ClO2)2. During one hour I don´t fond almost nothing. Google: Over 100 C explode. Thats all.
VoD ? Initiation properties? Sorry my question.....
Development of primarily - secondary substances CHP (2015) Lithex (2022) Brightelite (2023) Nitrocelite (2024)
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woelen
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Yes, I tried adding nickel(II) salts to a large excess amount, but this does not show formation of a complex with a new color. With cobalt(II) I just
get the brown precipitate, the solution becomes very pale mustard/brown, maybe the brown compound is slightly soluble in water, giving a pale
mustard/brown solution.
Another thing I have to add, with the cobalt(II) sulfate added to a solution of sodium chlorite, you also get a small amount of ClO2. The air above
the mix becomes very pale yellow and there is a clearly noticeable smell of ClO2. I know that smell quite well, it differs a lot from the smell of
Cl2, more spicy, less choking. This formation of ClO2 indicates hydrolysis.
I also have done experiments with nitrite and chlorite (I described that in another thread on sciencemadness), but a mix of these is very dangerous.
On addition of a small amount of acid the mix explodes.
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woelen
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The nickel(II) solution has changed a lot while standing.
At the surface of the liquid, a black solid has formed, at the bottom of the liquid, the precipitate has turned dark brown. The liquid itself still is
green, due to nickel(II)-ions. After shaking vigorously, the liquid looks dark gray and turbid. I'll let it settle again and see what is formed.
I did another explosion test with Pb(ClO2)2 with the goal of looking at the light output of the explosion. I put a little amount in a test tube and
then heated the test tube above a flame. At a certain point, there is a POP sound and a very weird somewhat pale orange/gray flash of light. The light
is not particularly strong, but it is a most peculiar color. It looks as if you take a picture of a bright orange object or bright orange flame and
then tone down saturation of the image until it is almost black and white and just a pale orange hue is left.
Another special thing is that the reaction occurs without any fuel added to the mix. I know of other compounds which can explode violently on heating,
but these have fuel and oxidizer in the same compound (e.g. ammonium bromate, organic amine perchlorates).
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Laboratory of Liptakov
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Many great discoveries were made by mere observation. Without complicated measurement. Pb(ClO2)2 and his examination can be this case.
Development of primarily - secondary substances CHP (2015) Lithex (2022) Brightelite (2023) Nitrocelite (2024)
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woelen
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With cobalt(II) it seems to be impossible to make a chlorite by simply mixing solutions of the chemicals. There is precipitation of a dark brown
solid, but there also is formation of some ClO2. I have the impression that there is hydrolysis, leading to formation of some cobalt oxide and
lowering of pH, which leads to loss of some ClO2.
I rinsed some of the dark brown precipitate and allowed it to dry in a petri dish. This leads to formation of a nearly black solid. This solid has no
energetic properties at all. When it is added to concentrated hydrochloric acid, then it dissolves, giving a blue solution of the common well known
tetrachloro complex of cobalt(II). Some gas is produced, but no ClO2. Not the faintest appearance of the typical yellow of ClO2. The color of ClO2 is
very intense, even at low concentrations it is clearly visible. No ClO2 can be observed at all. There is a smell of chlorine though.
I think the brown solid is a mixed oxidation state material of cobalt(II) and cobalt(III), which on acidification with HCl gives a little Cl2 and
cobalt(II), which is coordinated by the high concentration of chloride ion in the conc. HCl.
Up to now, only the lead(II) salt of chlorite could be made succesfully.
I still have to try the silver salt.
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Rhodanide
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| Quote: Originally posted by woelen |
Up to now, only the lead(II) salt of chlorite could be made succesfully.
I still have to try the silver salt. |
Good to hear that you'll be working with the silver salt, as it's pretty cool.
It's easily made by addition of a NaClO2 solution to an AgNO3 solution. Silver Chlorite precipitates as a heavy, yellow
precipitate. Upon drying in a room temp, very dark cupboard, a dry powder is produced. When heated, it seems to melt, then explode. When mixed with
Sulfur powder, it can be set off by merely tapping it with a stick. I learned this the hard way, silver vapor is painful. Here's a photo of it soon
after precipitating. I have probably around 50-70 g of it in a brown glass bottle sealed with teflon, kept in a dark cupboard. Seems pretty stable so
far!
[Edited on 3-5-2018 by Rhodanide]
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woelen
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What NaClO2 did you use? I only can get the 80% grade, the remainder being NaCl and some Na2CO3. These also produce a precipitate with silver, so you
get an impure precipitate.
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Rhodanide
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| Quote: Originally posted by woelen | | What NaClO2 did you use? I only can get the 80% grade, the remainder being NaCl and some Na2CO3. These also produce a precipitate with silver, so you
get an impure precipitate. |
Yes, I used 80%. I don't think there's a way to separate the pure Chlorite, but oh well :/
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Rhodanide
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| Quote: Originally posted by woelen |
At a certain point, there is a POP sound and a very weird somewhat pale orange/gray flash of light. The light is not particularly strong, but it is a
most peculiar color. It looks as if you take a picture of a bright orange object or bright orange flame and then tone down saturation of the image
until it is almost black and white and just a pale orange hue is left.
. |
I'd expect that to happen, lol.
Pb's emission color is awesome! If you've ever set fire to Nitromethane, then it's like that, with a more of a blue hint. And since you were burning a
Pb compound, I'd expect it to be grey.  The orange was likely something in the
test tube already, or maybe dust, or maybe even the Si in the glass being excited and releasing light. Who knows?
Here's an old photo of when I wet a Tungsten rod with water and dipped into Pb Acetate crystals, then put them under a flame.
v v v v
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woelen
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Interesting info about lead flame colors. I did not know that. The pale orange color then also can be explained. I made the Pb(ClO2)2 from NaClO2 and
it almost certainly will contain a little amount of sodium ions. These introduce an orange color to flames.
Also good to know that the silver-salt can be prepared from 80% NaClO2. Did you use excess AgNO3 or excess NaClO2? Or did you use precisely computed
and weighed stoichiometric amounts of both chemicals? I can imagine that by playing with the relative amounts of these chemicals that you can find an
optimum in terms of purity of the final product (depending on which of AgCl and AgClO2 has lowest solubility in water).
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Rhodanide
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| Quote: Originally posted by woelen | Interesting info about lead flame colors. I did not know that. The pale orange color then also can be explained. I made the Pb(ClO2)2 from NaClO2 and
it almost certainly will contain a little amount of sodium ions. These introduce an orange color to flames.
Also good to know that the silver-salt can be prepared from 80% NaClO2. Did you use excess AgNO3 or excess NaClO2? Or did you use precisely computed
and weighed stoichiometric amounts of both chemicals? I can imagine that by playing with the relative amounts of these chemicals that you can find an
optimum in terms of purity of the final product (depending on which of AgCl and AgClO2 has lowest solubility in water). |
Carefully measured compounds? Pah. A Turtle would be better at math than me. I used excess Chlorite, got to make sure I don't waste soluble Ag!
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woelen
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Good, easy to do. I'll try next weeked with a gram or so of AgNO3 and excess NaClO2.
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Rhodanide
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You know, I wonder if there's any way that we could possibly use purification through solubility to purify the Chlorite...
Is it soluble in something that NaCl isn't? Or alternatively, does one have a property which would allow one to be more soluble in water than the
other?
It's obvious that the AgCl present in the Chlorite formation dulls down the Ag Chlorite's energetic properties. :/
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woelen
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The reason that NaClO2 is sold at 80-85% concentration is that the pure solid is not stable.
Pure NaClO2 slowly decomposes, giving NaCl and NaClO3. This process slows down at increasing concentration of NaCl, until appr. 20% is converted and
from that point it remains stable.
NaClO2 as it is available for private individuals frequently is intended for human consumption (used for water sterilization, also used as so-called
MMS solution) and for that purpose it must be absolutely free of NaClO3. So, they prepare a solution with 80-85% of NaClO2, the remainder being mostly
NaCl. This is evaporated to dryness and a very intimate mix (solid solution?) is formed, which is stable and has indefinite shelf life and does not
form NaClO3 when stored.
So, do not expect to find anything else than 80-85% pure NaClO2.
Purifying NaClO2 is not easy. Maybe you can recrystallize it and keep NaCl behind in solution. If you succeed in making it more pure, you will have to
use it up quicky. I do not think it is worth the effort.
It also may be that when excess NaClO2 is used (as you did) that your precipitate is nearly 100% AgClO2. This depends on the solubility of AgClO2,
relative to the solubility of AgCl. I do not know which of these two is less soluble. That determines whether one can best use excess NaClO2 or excess
AgNO3 for getting the most energetic product. I could not find solubility info on AgClO2. It is not a common compound and is not tabulated at many
places.
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Rhodanide
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| Quote: Originally posted by woelen | The reason that NaClO2 is sold at 80-85% concentration is that the pure solid is not stable.
Pure NaClO2 slowly decomposes, giving NaCl and NaClO3. This process slows down at increasing concentration of NaCl, until appr. 20% is converted and
from that point it remains stable.
NaClO2 as it is available for private individuals frequently is intended for human consumption (used for water sterilization, also used as so-called
MMS solution) and for that purpose it must be absolutely free of NaClO3. So, they prepare a solution with 80-85% of NaClO2, the remainder being mostly
NaCl. This is evaporated to dryness and a very intimate mix (solid solution?) is formed, which is stable and has indefinite shelf life and does not
form NaClO3 when stored.
So, do not expect to find anything else than 80-85% pure NaClO2.
Purifying NaClO2 is not easy. Maybe you can recrystallize it and keep NaCl behind in solution. If you succeed in making it more pure, you will have to
use it up quicky. I do not think it is worth the effort.
It also may be that when excess NaClO2 is used (as you did) that your precipitate is nearly 100% AgClO2. This depends on the solubility of AgClO2,
relative to the solubility of AgCl. I do not know which of these two is less soluble. That determines whether one can best use excess NaClO2 or excess
AgNO3 for getting the most energetic product. I could not find solubility info on AgClO2. It is not a common compound and is not tabulated at many
places. |
I've found out the same thing, it seems that there's little to no information about them online. Another thing, if Methyl Hypochlorite is stable
enough to exist at STP, then why not Methyl Chlorite? If I'm correct, the formula would be CH3ClO2... now I'm just
wondering how I'd make this, or anyone else for that matter. I think that Methyl Hypochlorite is made by rxn of MeOH with Hypochlorous acid in
solution made by acidification of NaClO/Bleach. BUT - Chlorous acid is probably wildly unstable and we risk running into our old friend Chlorine
Dioxide. That brings me to the next thing. Seeing as Chlorine Dioxide is a pretty good oxidizer, I'm not too keen on seeing if it'll make
MeClO2 by bubbling it through Methanol. I don't want to die, not yet!
Seeing as how Methyl Perchlorate is an absolutely INSANE chemical, and how I've never even heard of Methyl Chlorate, who knows what Methyl
Chlorite would be like!! MAYBE, it could be made in small amounts by slow, careful addition of NaClO2 to (liquid, might I add!!) Methyl
Chloride/Chloromethane? In the possible reaction: CH3Cl (l) + NaClO2 (s) -> CH3ClO2 (???) + NaCl (s)
It'd probably be a good idea to perform this reaction at or below -30 C, for not only keeping the MeCl liquid, but also minimizing the risk of
exotherm (and decomposition).
Maybe it would need a catalyst... hrm... Let me know what you think! Sorry for replying so late.
-R
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woelen
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Methyl hypochlorite does exist, actually, it is very easy to make it. Mix 0.5 ml of methanol with 1 ml of conc. acetic acid (80% or better). Mix this
with a few ml of 10% bleach, slowly. You get a colorless gas, it bubbles nicely. If you keep a flame near the gas, then you get a deafening report! Be
careful with igniting the gas, its explosions are powerful and can cause a test tube to shatter. It is best if you do this in a petri dish or
hourglass.
I also tried to make methyl chlorite in the same way with 10% NaClO2, but that does not work. You get slow production of ClO2, just a little bit is
formed, it makes the mix pale yellow. No colored gas can be observed above the liquid.
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Bert
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Such a stupid simple, OTC reagent sourced gas generant, yet perhaps offering some interesting possibilities. If possible to reliably (AND remotely!)
generate a few litres volume when desired, at any rate.
Woelen, do you know what is the stoichiometry of the methyl hypochlorite reaction? Quite OB negative I would assume- first H is burned to water, then
as much C as possible to CO, and on to CO2 if sufficient O2 were available. Which it does not appear to be, excepting by use of O2 from surrounding
air. Is the decomposition of the gas itself energetic, as with acetylene?
There is a thing we have done for flame effects on stage or as SFX for video/movies. We take a small sprinkler as shown in attached image, place it in
the bottom of a bucket and connect to a low pressure source of a flammable gas, propane, butane, etc. The bucket is then filled with a few inches of
water and a bit of dish washing soap, perhaps a small ammount of glycerine is added to extend foam life in case of dry weather.
Gas is turned on, you make FIRE FOAM. One may scoop up the foam and fill a prop desired to emit a good puff of flame when lit, such as a BBQ grill
when wanting to safely emulate a big flare up on lighting the grill. There is no continuous supply of flammable gas to prop, it burns and is done, no
concerns of burning on after the desired effect djration, no pipes need be run to a prop, valves, pressure regulators, etc. are not required- it is
easy to control, one just scoops up and emplaces as much volume of foam as needed for desired volume of flame. Fire foam self deactivates after a few
minutes in the open air as the soap bubbles break and the gas diffuses away, the area around the foam goes below minimum % to sustain burning in air
quite quickly.
Methyl hypochlorite behaves as a single compound explosive? Then I would expect a rather faster reaction. Perhaps an explosive foam could be made.
Initiating with some kind of HE booster rather than a flame might even provoke a detonation.
Of course, just going ahead and making a few cubic litres testing might be the last thing one ever did.
[Edited on 16-3-2018 by Bert]
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woelen
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Methyl hypochlorite indeed is a single compound explosive. No air needed for explosion. The pure gas gives very powerful explosions when ignited. The
gas can be condensed to a liquid if you provide sufficient cooling (I don't know its boiling point though). Igniting the liquid must be really
impressive.
In my experiments with small amounts of gas, I noticed that the explosion leads to formation of another gas, which burns with a pale blue flame, in
bright light this is nearly invisible.
I am inclined to think that the gas decomposes to HCl and CH2O and that the latter simply burns in contact with air, after the explosion. I could
smell the same smell as when I open a bottle of conc. HCl, so that certainly will be in the mix after explosion. Otherwise I could not smell anything.
I did not smell CH2O (formaldehyde). The problem is that I did my tests in contact with air.
If you really want to test the properties of the gas, then you should fill a small bottle with the pure gas (which is not that dificult), have a thin
nichrome wire in the bottle (through the cap or something like that) and then ignite while the bottle is closed. Next, under water you should open the
bottle and allow some water to suck in and then dissolve the gas mix in the bottle. Finding a bottle, which does not shatter in the explosion may be a
challenge. One could try with a small 100 ml bottle and bury this under a layer of soil before igniting it and having long leads to the bottle so that
on ignition you are at least a few meters away from the bottle. A lot of hassle.
I personally think that the methyl hypochlorite ester is nothing more than a funny curiousity. Nice for a little demo, the more so because it can be
done with very common chemicals.
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