Sciencemadness Discussion Board - Heavy metal chlorites

Heavy metal chlorites

woelen - 28-11-2016 at 00:42

Last weekend I investigated the properties of lead chlorite. This chemical is easy to make by precipitation. Simply mix a solution of NaClO2 and a solution of Pb(NO3)2. Both solutions should not be too strongly concentrated.

When this is done, a yellow precipitate is formed, which has a color, very much resembling the color of powdered sulphur. This precipitate can be washed with distilled water and then can be dried.

It has energetic properties. When a small amount of the dry powder is heated in a test tube, then it explodes with a loud POP noise and white smoke is produced. When it is kept in an open flame, it very quickly burns (like good quality black powder) with a pale flame, producing white smoke. The funny thing is that this material does not require any fuel for quick burning or explosion. This is in strong contrast with NaClO2 which only burns with fuels like sugar, sulphur or red phosphorus.

I do not know about its long term stability. I keep appr. 500 mg of this material in a small vial and after one day of storage no deterioration can be observed.

Have any of you experience with other heavy metal chlorites? Next one I want to try is the silver salt.

Please, if you do experiments with this, be careful. Chlorites are very reactive and at low pH you get the explosive gas ClO2. Besides that, lead is toxic and the white smoke from burning/decomposing Pb(ClO2)2 certainly should not be inhaled!

DraconicAcid - 28-11-2016 at 16:53

I'm impressed by your courage in working with this material.

Bert - 28-11-2016 at 19:33

New material to me-

Perhaps you could try hitting a small pinch with a hammer on a steel anvil, or place a few mg on a porcelain tile and rub hard with a similarly hard tool- Just a quick and rough check for impact and friction sensitivity. If you use a mortar and pestle, be aware that hot Lead oxide tends to melt itself into glassy substances, so don't use your favorite tools.

The reaction products from that burning, Lead oxide(s)? Lead chloride???

And from this earlier thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=3314

Per Axt i this thread, PATR says the mixture of Lead chlorite and sugar is explosive, to the degree it may be used as a primary?!

Lots of old references to Lead chlorite either spontaneously burning on contact with Sulfur, or firing on slight friction.

Further descriptions of Silver and Mercury chlorites as being explosive- Not going to try Mercury, myself.

[Edited on 29-11-2016 by Bert]

woelen - 29-11-2016 at 02:19

I am willing to try the silver salt and maybe copper and other row one transition metals. I do not try the mercury salt. Producing smoke of a mercury salt is not good at all. Mercury is much more toxic than lead and even lead is borderline already for me (hence I only tried ultra small amounts of 10 mg or so in open air and tried a little bit more like 50 mg in a test tube).

The chlorites are interesting, however, to experiment with and if you use very small quantities (appr. 50 mg) in the experiments, then the risk in terms of explosion or fire at least are acceptable. With such small quantities you do not run the risk of blowing up your hand or getting big things on fire.

PHILOU Zrealone - 30-11-2016 at 08:04

@Woelen,
I know from a long time that Pb, Ag and Hg chlorite are interesting alone for energetic material (research), but I have never had easy acces to pure chlorite so I never made those.

Thanks to you now I know for sure what is written in old books is right

.

Now one step further:
-Test mixes (carefully since probably more sensitive than chlorate)
-Test complexes (ethylene diamine, ...) ammonia will probably be too dangerous
--> Axt already tested AgClO2 and N2H4...decomposed into fumes upon contact.

[Edited on 30-11-2016 by PHILOU Zrealone]

Tdep - 2-12-2016 at 05:46

Quote: Originally posted by PHILOU Zrealone

never had easy access to pure chlorite

I think this thread is convincing me to buy some sodium chlorite, but when you say 'pure chlorite', it only ever comes as an 80% grade yeah? There's a source roughly nearby me on eBay for $21/200g including shipping and I think that's an alright-ish price

woelen - 3-12-2016 at 08:53

It can be made at higher purity, but this is not stable. If you have pure NaClO2, then it decomposes to NaCl, NaClO3 and possibly some O2 as well. With increasing concentration of NaCl the decomposition slows down until at appr. 20% concentration of NaCl the decomposition comes to a halt completely. In the commercial product NaCl is added intentionally to prevent further decomposition and formation of NaClO3, especially when the material is used for human consumption (e.g. in decontamination of water, or when used as MMS). Hence, the product you can buy has 20% NaCl and 80% NaClO2. The 20% NaCl is no issue though in most experiments.

PHILOU Zrealone - 3-12-2016 at 10:16

Quote: Originally posted by woelen

It can be made at higher purity, but this is not stable. If you have pure NaClO2, then it decomposes to NaCl, NaClO3 and possibly some O2 as well. With increasing concentration of NaCl the decomposition slows down until at appr. 20% concentration of NaCl the decomposition comes to a halt completely. In the commercial product NaCl is added intentionally to prevent further decomposition and formation of NaClO3, especially when the material is used for human consumption (e.g. in decontamination of water, or when used as MMS). Hence, the product you can buy has 20% NaCl and 80% NaClO2. The 20% NaCl is no issue though in most experiments.

No issue, except the fact that chlorides of the heavy metals of interest are also precipitated...

--> Hg2Cl2 (but not HgCl2 what is soluble to some extend)
--> PbCl2
--> AgCl

So if you used that brand of NaClO2/NaCl (80/20) the Pb(ClO2)2 you made also contains PbCl2...

and knowing that the solubility increases from PbCl2 to Pb(ClO4)2, one may conclude that PbCl2 < Pb(OCl)2 < Pb(OClO)2 < Pb(OClO2)2 < Pb(OClO3)2 ("<" = less soluble than)...
so the energetic properties of the Pb(ClO2) are far from what it should be (

) because most of the precipitate is tempered by a vast amount of PbCl2 (probably larger than 20%)

Tdep - 3-12-2016 at 17:45

Interesting about the decomposition, how does the NaCl stabilize the chlorite like that? It all happens in the solid state, I wouldn't have thought such a thing happened.

Ok so to see the best the heavy metal chlorites can do, the chloride must really be removed. How are we going to do that? I suppose you could fractionally crystallise it out, given that the chlorite has over three times the solubility?

Actually no, a better option might be to first add a ~0.3Mol equivalent of Ag(NO₃) to the technical grade sodium chlorite, give it a bit of time to equilibriate and then filter. That should remove the chloride and leave just the chlorite, or at least heavily enrich the chlorite/chloride ratio?

PHILOU Zrealone - 4-12-2016 at 11:04

Quote: Originally posted by Tdep

Interesting about the decomposition, how does the NaCl stabilize the chlorite like that? It all happens in the solid state, I wouldn't have thought such a thing happened.

Ok so to see the best the heavy metal chlorites can do, the chloride must really be removed. How are we going to do that? I suppose you could fractionally crystallise it out, given that the chlorite has over three times the solubility?

Actually no, a better option might be to first add a ~0.3Mol equivalent of Ag(NO₃) to the technical grade sodium chlorite, give it a bit of time to equilibriate and then filter. That should remove the chloride and leave just the chlorite, or at least heavily enrich the chlorite/chloride ratio?

Following what Woelen wrote,
One may assume that there is an equilibrium disproportionation reaction at work (into the dissolved state).
3NaOClO <===> NaCl + 2NaOClO2
or
2NaOClO <===> NaCl + NaOClO2 + 1/2 O2

Fractional crystallisation yes like you wrote.
or
Fractional dissolution/solvatation (but one needs to know better the properties of halides vs chlorites into specific solvents and relation to temperature).
or
Fractional precipitation like you wrote.

Assuming no chlorate formation during purification proces...purity may be assessed by O2 content (probably by heating moderately with a minute MnO2 amount (like it is done for chlorates)).

[Edited on 4-12-2016 by PHILOU Zrealone]

woelen - 4-12-2016 at 11:34

I did my experiments with Pb(ClO2)2 with the 80% NaClO2 which can be obtained from eBay. I have no access to 100% pure NaClO2 and that would not remain pure anyway.

The 80% material works fine and the Pb(ClO2)2 obtained with this is highly energetic.

The simple order, given by PHILOU Zrealone, does not hold.

Pb(ClO)2 does not exist, it immediately has an internal oxidation reaction in which the lead(II) is oxidized to lead(IV). If you add some hypochlorite to a solution of a lead(II) salt, then you get a brown precipitate of PbO2.

Pb(ClO2)2 probably is less soluble than PbCl2, otherwise indeed I would have obtained a lot of PbCl2 with just a little Pb(ClO2)2 in it. I do not think that my Pb(ClO2)2 is 100% pure, but I am quite sure that it is does not contain more than 20% PbCl2.

I think that the order of solubility for the lead salt is something like Pb(ClO2)2 < PbCl2 << Pb(ClO3)2 ~ Pb(ClO4)2.

For the silver salt, things may indeed be somewhat more problematic. AgCl is really insoluble, while PbCl2 is somewhat soluble. For the copper salt, CuCl2 is soluble, so with that there is no issue.

This week I hope to find some time for the experiments suggested by PHILOU Zrealone. I do have ethylene diamine and I can try with hydrazine hydrate. With the latter I expect immediate decomposition though. If you add solid NaClO2 to hydrazine hydrate then you get an immediate violent reaction (I once tried that) and I do not expect otherwise with lead chlorite.

PHILOU Zrealone - 5-12-2016 at 12:11

Quote: Originally posted by woelen

I did my experiments with Pb(ClO2)2 with the 80% NaClO2 which can be obtained from eBay. I have no access to 100% pure NaClO2 and that would not remain pure anyway.

The 80% material works fine and the Pb(ClO2)2 obtained with this is highly energetic.

The simple order, given by PHILOU Zrealone, does not hold.

Pb(ClO)2 does not exist, it immediately has an internal oxidation reaction in which the lead(II) is oxidized to lead(IV). If you add some hypochlorite to a solution of a lead(II) salt, then you get a brown precipitate of PbO2.

Pb(ClO2)2 probably is less soluble than PbCl2, otherwise indeed I would have obtained a lot of PbCl2 with just a little Pb(ClO2)2 in it. I do not think that my Pb(ClO2)2 is 100% pure, but I am quite sure that it is does not contain more than 20% PbCl2.

I think that the order of solubility for the lead salt is something like Pb(ClO2)2 < PbCl2 << Pb(ClO3)2 ~ Pb(ClO4)2.

For the silver salt, things may indeed be somewhat more problematic. AgCl is really insoluble, while PbCl2 is somewhat soluble. For the copper salt, CuCl2 is soluble, so with that there is no issue.

This week I hope to find some time for the experiments suggested by PHILOU Zrealone. I do have ethylene diamine and I can try with hydrazine hydrate. With the latter I expect immediate decomposition though. If you add solid NaClO2 to hydrazine hydrate then you get an immediate violent reaction (I once tried that) and I do not expect otherwise with lead chlorite.

Thank you for the clarification.
I just want to say that PbO2 may result from Pb(2+) with ClO(-) because commercial hypochlorite contains a lot of OH(-).
--> Oxydo redox potential generally express into acidic or into basic media...but does this reaction stil happens when in neutral aqueous media?

About hydrazine chlorite...I have never found infos about it but wel over hydrazine chlorate, perchlorate and diperchlorate.

Axt did test AgClO2 with hydrazine, but then the decomposition must be due to the Ag(+) with what hydrazine is uncompatible (see AgNO2 and AgNO3 decompose into a silver mirror...although in books the nitrate is supposed to deliver AgN3)...with Pb(ClO2)2 it is terra incognita.
--> Be very cautious --> soft plastic beaker and minute amount

woelen - 12-12-2016 at 07:41

Right now, I have the Pb(ClO2)2 standing for two weeks and it still is as good as when I made it. This stuff appears to be stable on storage. Quite a few chlorites are said to be unstable, but the lead-salt is not.
It is unstable in contact with hydrazine. Making a complex with hydrazine is not possible. It reacts immediately with hydrazine. I think that this is due to incompatibility of chlorite with hydrazine, regardless of the cation in the compound.

PHILOU Zrealone - 12-12-2016 at 11:53

Quote: Originally posted by woelen

Right now, I have the Pb(ClO2)2 standing for two weeks and it still is as good as when I made it. This stuff appears to be stable on storage. Quite a few chlorites are said to be unstable, but the lead-salt is not.
It is unstable in contact with hydrazine. Making a complex with hydrazine is not possible. It reacts immediately with hydrazine. I think that this is due to incompatibility of chlorite with hydrazine, regardless of the cation in the compound.

It seems right conclusion...Pb(2+) is unreactive towards hydrazine...so the chlorite must be responsible. How did the reaction proceed? Massive gas evolution, puff, flame?
Where both compounds into solution (cooled or not) or the solid Pb(ClO2)2 into the liquid hydrazine (concentrated or not)?

Maybe a hydrazine complex may form with another cation...Hg do form nice complexes with amines...Zn, Co, Fe, Cd, Ca, Mg remains to test and Cu which will most likely be very unstable.
1°)I did made some truquoise blue Cu(N2H4)2(NO3)2 complex from conc Cu(NO3)2, ethanol and N2H4.H2O (80% hydrazine) while most did turn into N2, Cu and Cu2O brown sponge; the remaining turquoise blue complex precipitate decomposed in big part by itself when drying (self inflamation) and a tiny part remained as very inflamable and energetic compound (burning with a nice blue-green flame).
2°)I did put NaNO2 cristals into hydrazine hydrate (N2H4.H2O / 80%N2H4) with phenol phtaleine solution. The red-violet colour dissapeared without gas bubbles nor heating while the NaNO2 cristals dissolved...so something happened.
I was hoping that since into books AgNO2 (or AgNO3) turns into AgN3 upon exposure to N2H4; maybe some NaN3 would form...but the study was not pursued further.

To confirm unstability of hydrazine with chlorites, could you just mix a little of your NaClO2/NaCl with hydrazine?

woelen - 13-12-2016 at 01:33

The reaction with the hydrazine is not really violent, it just is fast. The yellow solid, when added to a solution, containing 30% or so of hydrzine, turns white at once and falls apart and makes a cloudy white liquid.

I also did other experiments with the Pb(ClO2)2:
- Grind it with a metal spatula while pressing it hard and putting strong force on it. This does not ignite the material. The final result is a yellow powder, looking like sulphur.
- Mix the yellow powder with sulphur (appr. same volume of powdered sulphur mixed carefully with the lead chlorite). When this mix is ignited it burns very fast with a peculiar grey/blue flame color. The burning is not explosive though. Ignition is remarkably easy. You do not have to keep it in a flame, just keeping the mix near a flame is enough to ignite it.
- Mix the yellow powder very carefully with red phosphorus. Tap the mix with a glass rod. This leads to a small explosion. The mix gives a high pitched report (PANG) and produces a wonderful smoke ring which rises up slowly and becomes wider while rising up.
- Mix the yellow solid with a solution of ethylene diamine in water. The material slowly turns white. It does not dissolve, the liquid becomes cloudy. I think that no complex is formed, the material simply hydrolyses, forming Pb(OH)2, free ClO2(-) ions and protonation of some of the ethylene diamine. Nothing interesting.

I will try adding NaClO2 to a solution of hydrazine. In the past I have done the experiment of adding NaClO2 to a solution of N2H4.2HCl. That leads to a very violent and dangerous reaction. Using plain hydrazine and not the hydrazinium salt may lead to different results, however. It will come back on that.

Tsjerk - 13-12-2016 at 03:22

Maybe you could add a couple of milliliters of an organic slightly polar solvent (ethyl-acetate e.g.), to the reaction mixture of the ClO^2- with the ethelyne diamine. If you are partly oxidizing the amine groups the white stuff should dissolve in the organic layer, if it is Pb(OH)₂ it should not.

It would be interesting to see such a aquous oxidation of an organic compound.

Edit: If it is an oxidation it should also give an white percipitate with NaClO₂

[Edited on 13-12-2016 by Tsjerk]

PHILOU Zrealone - 13-12-2016 at 04:18

Quote: Originally posted by Tsjerk

Maybe you could add a couple of milliliters of an organic slightly polar solvent (ethyl-acetate e.g.), to the reaction mixture of the ClO^2- with the ethelyne diamine. If you are partly oxidizing the amine groups the white stuff should dissolve in the organic layer, if it is Pb(OH)₂ it should not.

It would be interesting to see such a aquous oxidation of an organic compound.

Edit: If it is an oxidation it should also give an white percipitate with NaClO₂

[Edited on 13-12-2016 by Tsjerk]

The oxydation of the amines may lead to:
-N-hydroxylamine (R-NHOH about as oxydable as hydrazine)
-chloramine (R-NHCl)
-nitroso compound or resonance form aldehyde oxime (R-CH2-N=O <==> R-CH=N-OH)(*)
-dichloramine (R-NCl2)
-nitro group (R-NO2)
-cyano group (R-C#N)

And further oxydation products by O or Cl.

(*) Note that nitrosocompound can be seen as a dehydrated from of a putative N,N-dihydroxylamine
R-CH2-N(OH)2 <==> R-CH2-N=O + H2O

[Edited on 13-12-2016 by PHILOU Zrealone]

PHILOU Zrealone - 13-12-2016 at 04:25

Quote: Originally posted by woelen

The reaction with the hydrazine is not really violent, it just is fast. The yellow solid, when added to a solution, containing 30% or so of hydrzine, turns white at once and falls apart and makes a cloudy white liquid.

I also did other experiments with the Pb(ClO2)2:
- Grind it with a metal spatula while pressing it hard and putting strong force on it. This does not ignite the material. The final result is a yellow powder, looking like sulphur.
- Mix the yellow powder with sulphur (appr. same volume of powdered sulphur mixed carefully with the lead chlorite). When this mix is ignited it burns very fast with a peculiar grey/blue flame color. The burning is not explosive though. Ignition is remarkably easy. You do not have to keep it in a flame, just keeping the mix near a flame is enough to ignite it.
- Mix the yellow powder very carefully with red phosphorus. Tap the mix with a glass rod. This leads to a small explosion. The mix gives a high pitched report (PANG) and produces a wonderful smoke ring which rises up slowly and becomes wider while rising up.
- Mix the yellow solid with a solution of ethylene diamine in water. The material slowly turns white. It does not dissolve, the liquid becomes cloudy. I think that no complex is formed, the material simply hydrolyses, forming Pb(OH)2, free ClO2(-) ions and protonation of some of the ethylene diamine. Nothing interesting.

I will try adding NaClO2 to a solution of hydrazine. In the past I have done the experiment of adding NaClO2 to a solution of N2H4.2HCl. That leads to a very violent and dangerous reaction. Using plain hydrazine and not the hydrazinium salt may lead to different results, however. It will come back on that.

Thank you for those valuable experiments and reports.

About cloudy EDA and "nothing interesting", I think that it is actually a very interesting fact...maybe a way to safely make (a probably unmade to date) EDA dichlorite!

[Edited on 13-12-2016 by PHILOU Zrealone]

woelen - 13-12-2016 at 10:05

EDA dichlorite would be very interesting. How would you make this compound? Adding a chlorite salt to a solution of EDA does not work. EDA is quite basic and chlorite is the conjugate of a fairly weak acid (HClO2). One should be able to isolate HClO2 (dissolved in water) and mix this with a solution of EDA.

One possible route to EDA dichlorite could be making EDA perchlorate (which is not difficult, I have done this myself from EDA and perchloric acid) and mixing a solution of that with a solution of potassium chlorite. KClO4 will precipitate, the protonated EDA ions and the chlorite ions remain behind in solution.
Unfortunately I know of no source for KClO2. On eBay I only see sellers of NaClO2 and this trick does not work for sodium salts, due to the high solubility of NaClO4. Making KClO2 yourself in a safe way is not going to be easy. One possible route is to bubble ClO2 through a solution of KOH to which H2O2 is added as well, but this reaction is quite dangerous, due to the explosive nature of ClO2.

DraconicAcid - 13-12-2016 at 11:26

Quote: Originally posted by woelen

EDA dichlorite would be very interesting. How would you make this compound? Adding a chlorite salt to a solution of EDA does not work.

It was just suggested that Pb(ClO2)2 + EDA + 2 H2O -> Pb(OH)2(s) + H2EDA(ClO2)2(aq) might work.

PHILOU Zrealone - 13-12-2016 at 13:57

Quote: Originally posted by DraconicAcid

Quote: Originally posted by woelen

EDA dichlorite would be very interesting. How would you make this compound? Adding a chlorite salt to a solution of EDA does not work.

It was just suggested that Pb(ClO2)2 + EDA + 2 H2O -> Pb(OH)2(s) + H2EDA(ClO2)2(aq) might work.

Yes DraconicAcid you got it

Woelen didn't realise that what he wrote was actually the solution..."the material simply hydrolyses, forming Pb(OH)2, free ClO2(-) ions and protonation of some of the ethylene diamine."

If Pb(OH)2 precipitates and doesn't complexate then only H(+), EDA and ClO2(-) remains into solution --> filtration and cristallization --> EDA dichlorite.

Although his new proposal of KClO4 precipitation is also a valid one.

ClO2 can be made safely from oxalic acid and NaClO3 but it is then dilluted with CO2 what may be anoying to get pure KClO2 since KHCO3 and K2CO3 will enter into account.

[Edited on 13-12-2016 by PHILOU Zrealone]

woelen - 14-12-2016 at 03:00

Yeah, you're right. I must have been sleeping while writing the previous post

I will do this experiment with a small amount of Pb(ClO2)2. In a practical experiment I think that you should add a slight excess of EDA to be sure that all chlorite is used and allow the water and excess EDA to evaporate while drying the material.

I am afraid that solid H2EDA(ClO2)2 (if it exists and does not decompose already before it becomes really dry) is a very sensitive chemical. In my experience, chlorite, combined with fuel, is more sensitive/reactive than chlorate, combined with the same fuel. And you all know how sensitive chlorate based mixes are, so what would be the chlorite based mix?

But nevertheless, it is worth a try. I have appr. 300 mg of my Pb(ClO2)2 left and I will use it up in the EDA experiment.

DraconicAcid - 14-12-2016 at 09:24

If you use excess EDA, then surely you will get HEDAClO2.

woelen - 14-12-2016 at 11:06

I added some EDA to the remaining amount of my Pb(ClO2)2. I first added the solid to some water and I had to add 3 big drops of EDA to have all of it converted to a white solid. Two drops was not enough, three drops was an excess amount. So, there is some excess amount.
Now I have a cloudy white liquid and I allow the white solid to settle at the bottom. Tomorrow I will pipette away the clear liquid and allow this to evaporate. I hope that solid material remains behind. This may become HEDA(ClO2), or a mix of H2EDA(ClO2)2 and HEDA(ClO2). Both are interesting. First let's see whetther it is possible at all to get an interesting energetic compound. Further optimizations always can be done in the future with a new batch of Pb(ClO2)2.

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

As another experiment, I added solid NaClO2 to 30% N2H4 in water. This does not react, not even after 15 minutes of waiting. Next, I added a small amount of dilute H2SO4. When this is done, then after a short delay, suddenly a violent reaction starts in which a colorless gas is formed (most likely N2) and after a while the liquid suddenly turns yellow (formation of ClO2). So, first, when there still is N2H4, the chlorite is reduced to colorless chloride, but once the hydrazine is used up, the remaining chlorite reacts with formation of ClO2.

PHILOU Zrealone - 14-12-2016 at 13:14

Quote: Originally posted by woelen

I added some EDA to the remaining amount of my Pb(ClO2)2. I first added the solid to some water and I had to add 3 big drops of EDA to have all of it converted to a white solid. Two drops was not enough, three drops was an excess amount. So, there is some excess amount.
Now I have a cloudy white liquid and I allow the white solid to settle at the bottom. Tomorrow I will pipette away the clear liquid and allow this to evaporate. I hope that solid material remains behind. This may become HEDA(ClO2), or a mix of H2EDA(ClO2)2 and HEDA(ClO2). Both are interesting. First let's see whetther it is possible at all to get an interesting energetic compound. Further optimizations always can be done in the future with a new batch of Pb(ClO2)2.

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

As another experiment, I added solid NaClO2 to 30% N2H4 in water. This does not react, not even after 15 minutes of waiting. Next, I added a small amount of dilute H2SO4. When this is done, then after a short delay, suddenly a violent reaction starts in which a colorless gas is formed (most likely N2) and after a while the liquid suddenly turns yellow (formation of ClO2). So, first, when there still is N2H4, the chlorite is reduced to colorless chloride, but once the hydrazine is used up, the remaining chlorite reacts with formation of ClO2.

Thank you again for the experiments and reports...with such a fast follow up.

The EDA chlorite or EDA dichlorite are indeed very interesting.
They may be hygroscopic at least EDA monochlorite (I doubt EDA dichlorite will be).

Thanks to the high enough pKa of chlorous acid, evaporation into plain air is no troubles (CO2 can't displace it); so the salt(s) must remain integer.

Since the salt(s) are unknown or little studied better be cautious...if perchlorates are sensitive to shock, heat and friction (I have made N2H5ClO4 and EDA diperchlorate into the 50g range and those are not to worry so much about)...following books chlorates of organic amines are even worst (I have never made any organic amine chlorates so I am of no help and I have no feeling to share about it)...
Chlorites may (or not) reveal to be even less stable...hard to predict...so better avoid glass.

Your experiments with chlorites and amines are very interesting!
They prove that:
1°)chlorite anion is more reactive/oxydizing into acidic media which can be direct acid or discrete acid like into an acidic salt (hydrazine hydrogenosulfate or hydrazinium chloride) or a salt of a weak base (mild acid).
2°)chlorite seems stabler into basic media like organic amines.
You potentially made hydrazinium chlorite and EDA chlorite/dichlorite.

woelen - 16-12-2016 at 10:09

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?

PHILOU Zrealone - 16-12-2016 at 12:52

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]

woelen - 18-12-2016 at 05:43

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.

PHILOU Zrealone - 18-12-2016 at 13:53

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]

woelen - 17-2-2018 at 13:20

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]

woelen - 19-2-2018 at 03:56

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.

DraconicAcid - 19-2-2018 at 08:13

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?

Laboratory of Liptakov - 19-2-2018 at 08:37

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..... :cool:

woelen - 19-2-2018 at 23:35

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.

woelen - 27-2-2018 at 12:20

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).

Laboratory of Liptakov - 27-2-2018 at 13:55

Many great discoveries were made by mere observation. Without complicated measurement. Pb(ClO2)2 and his examination can be this case.

woelen - 5-3-2018 at 04:25

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.

Rhodanide - 5-3-2018 at 11:57

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]

woelen - 5-3-2018 at 13:24

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.

Rhodanide - 6-3-2018 at 06:13

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 :/

Rhodanide - 6-3-2018 at 06:29

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

woelen - 6-3-2018 at 23:55

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).

Rhodanide - 7-3-2018 at 06:12

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!

woelen - 7-3-2018 at 07:55

Good, easy to do. I'll try next weeked with a gram or so of AgNO3 and excess NaClO2.

Rhodanide - 7-3-2018 at 10:16

Quote: Originally posted by woelen

Good, easy to do. I'll try next weeked with a gram or so of AgNO3 and excess NaClO2.

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. :/

woelen - 7-3-2018 at 13:29

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.

Rhodanide - 15-3-2018 at 11:15

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

woelen - 16-3-2018 at 10:31

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.

Bert - 16-3-2018 at 11:58

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.

plastic-ring-sprinkler-garden-irrigation-water-sprayer-nozzle-watering-the-flowers-tools_1626988.jpg - 93kB

[Edited on 16-3-2018 by Bert]

woelen - 18-3-2018 at 23:52

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.

Rosco Bodine - 19-3-2018 at 00:29

Quote: Originally posted by woelen

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.

Hazards for the organic hypochlorites include light !
They are photo sensitive.

Wonder if methyl hypochlorite would react with urea solution, possibly with some MEK to form the azine.
I'm thinking hydrazine or methyl hydrazine.

There is an old thread here linked
http://www.sciencemadness.org/talk/viewthread.php?tid=1896

Also this reports isopropyl hypochlorite by same method as tert-butyl hypochlorite.
http://www.orgsyn.org/demo.aspx?prep=CV5P0184

Potentially there would be chloroform as a byproduct from bleach reacting with ethanol, isopropanol, or acetone. This a nasty reaction in terms of toxic and explosion hazards.

[Edited on 3/19/2018 by Rosco Bodine]

woelen - 19-3-2018 at 04:14

I also did the experiment with ethanol and propanol. With ethanol you get a yellow liquid, but this liquid is VERY volatile. You see it evaporate, a big drop of this liquid on a piece of glass evaporates in half a minute or so. The propanol ester is much less volatile.

The ethanol and propanol esters are not explosive like the methanol ester. The vapor burns very quickly, with an orange flame and production of some soot. The liquids burn like nitrated cotton. They produce a WHOOSH sound and a very fast burn, but there is no report. The methyl ester really explodes with a loud report, so there is quite some difference.

Production of chloroform is virtually none, as long as you use sufficient acid. For this reason, you must mix your alcohol with twice its volume of acetic acid of high concentration and this mix must be added to the bleach. If you don't use lots of acid, then with ethanol and isopropanol you may end up with chloroform. If you first add the acid to the bleach, then you get a lot of chlorine and no ester is formed at all.

This reaction indeed is not very safe, that's why I told to use very small amounts, do not scale up.

Tdep - 19-3-2018 at 04:27

I too managed to make some (what I believed to be) ethyl hypochlorite, present as a green liquid that would float to the surface of the water solution it was made from transiently, before evaporating away into nothing, escaping capture and analysis. I'm amazed these compounds exist, there's very little information online about them

Rhodanide - 19-3-2018 at 04:50

Quote: Originally posted by Tdep

I too managed to make some (what I believed to be) ethyl hypochlorite, present as a green liquid that would float to the surface of the water solution it was made from transiently, before evaporating away into nothing, escaping capture and analysis. I'm amazed these compounds exist, there's very little information online about them

You didn't try and light it?

Tdep - 19-3-2018 at 05:14

Ya I mean what good is a synthesis if you don't end with trying to light it on fire, but it was in a water ethanol mix and there was a lot of chlorine and a lot going on... so none of it made it to the flames (that's really what I mean by '''further analysis''')

Bert - 19-3-2018 at 07:01

Having seen what tiny drops of the somewhat related ethyl perchlorate do, and given the description of the violent explosion of small bubbles? I am quite willing to limit the ammounts of methyl hypochlorite to a few ml...

I certainly agree that anything we make which includes a potential for explosion MUST be tested in a flame, for science! Also for entertainment value, if you're anything like me.

Also ran across this bit about trimethyl Perchlorate while searching for Blasters spoon photo, attached below.

Quote:

Another unusual ester was made in 1930 - trichloromethyl perchlorate from anhydrous carbon tetrachloride and AgClO4. It was so explosive that it could only be made in minute quantities and couldn't be completely isolated as it detonated with extreme violence when the CCl4 was evaporated off.

http://www.sciencemadness.org/talk/viewthread.php?tid=1081

I note Blaster has been absent from the forum since February of 2013. Hope he is still above the ground-

Not sure I am feeling lucky today, and carbon tetrachloride is depressingly hard to come by OTC in USA these days.

[Edited on 19-3-2018 by Bert]

woelen - 19-3-2018 at 08:33

I know of the ethyl perchlorate. It is very powerful.
Ethyl hypochlorite is much less powerful, it burns like nitrocotton when liquid. Maybe liquid methyl hypochlorite is as powerful as ethyl perchlorate, but I do not have the equipment to make this in liquid form. When you mix the chemicals it simply bubbles away as a gas. The gas, however, is very powerful already.

Rhodanide - 19-3-2018 at 10:56

Quote: Originally posted by Bert

Having seen what tiny drops of the somewhat related ethyl perchlorate do, and given the description of the violent explosion of small bubbles? I am quite willing to limit the ammounts of methyl hypochlorite to a few ml...

I certainly agree that anything we make which includes a potential for explosion MUST be tested in a flame, for science! Also for entertainment value, if you're anything like me.

Also ran across this bit about trimethyl Perchlorate while searching for Blasters spoon photo, attached below.

Quote:

Yeah, that's EXACTLY the thread I was thinking of when I said that Alkyl Perchlorates are absolutely INSANE chemicals.

Bert - 19-3-2018 at 11:01

I noted the description of Silver hypochlorite a few posts up.

Could Carbon tetrachloride and Silver hypochlorite react as Carbon tet and Silver perchlorate do?

[Edited on 19-3-2018 by Bert]

woelen - 19-3-2018 at 23:46

A few posts up there was writing about silver chlorite, not silver hypochlorite.
Silver chlorite is stable, a yellow compound: AgClO2
Silver hypochlorite does not exist it seems. I once tried making that, but if you add a silver salt to bleach you get a black precipitate, most likely a higher oxide of silver. I myself have done experiments with peroxodisulfate and silver(I) salts and these lead to formation of a mixed silver(I)/silver(III) oxide, Ag2O2. Hypochlorite most likely also produces this, but the result is much more messy than with peroxodisulfate, due to the high pH involved and due to the presence of chloride. No clean simple product.