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[*] posted on 22-3-2008 at 15:18


That is a very interesting and relatively easy way to make isopropyl hypochlorite. As I have read there the isopropyl hypochlorite must be extracted from the reaction mixture with methanol. I thought that it separates as a top (or bottom) layer as it is as far as I know insoluble in water.



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[*] posted on 22-3-2008 at 19:37


What I am wondering about is the difference between pickling salt and table salt.



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[*] posted on 23-3-2008 at 04:42


I think there is no difference.



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[*] posted on 25-3-2008 at 23:00


Quote:
Originally posted by Norrys
lol, maybe, but what would be able to make the EtOCl explode?
Has anyone data about sensitivity and uncompability with other materials?
I am afraid, that this could happen again.



For some saftey start with Bretherick's Handbook or PATR2700. Beilstein's Handbuch also has some saftey information sometimes and goes into reactivity. Try also looking at similar compounds to get an idea about the nature of that family of compounds.

E.g. for EtOCl:

Quote:
from Bretherick's Handbook of Reactive Chemical Hazards

Hazard: Though distillable slowly (at 36°C), ignition or rapid heating of the vapour causes explosion, as does contact of copper powder with the cold liquid.
Reference(s): Sandmeyer, T., Ber., 1885, 18, 1768.


So above 36 deg.C . it can explode (Beilstein says "overheating" the vapor causes explosion), so from that we can gather not good news on a very hot day. Copper powder is also enough to make it explode, so who knows how other metallic particles and impurties can and likley do act similar upon it.

MeClO is of course even worse:

Quote:
Bretherick's Handbook
Hazard: The liquid could be gently distilled (12°C) but the superheated vapour readily and violently explodes, as does the liquid on ignition.
Reference(s): Sandmeyer, T., Ber., 1886, 19, 859.


PATR (pgs H261-62) says: that lower member organic hypochlorites explode on contact with flame or bright light! And that in absence of light all of them decompose on standing except tertiary compounds. They are usually yellow oily liquids [Beilstein Syst. No. 21, pg. 325 also says EtOCl is colorless but that light turns it quickly yellow with the liberation of chlorine, that means as we will see moisture makes it especially unstable if light is present] and they are much more stable if any formed HCl is immediatley neutralized with NaHCO3 (they mention if not removed quickly with NaHCO3 solution residual HCl can make EtOCl decompose (even explode) spontaneously within a couple minutes: HCl + C2H5ClO -> Cl2 + C2H5OH). Direct sunlight or heating also makes EtClO explode. Also notes they generally explode on contact with copper powder. All propyl hypochlorites especially isopropyl are said to be instable. Further mention that n-Propyl- and isopropyl hypochlorite explode when exposed to light. Though one reference contends despite the literature, action of sunlight is "inappreciable" and that if HCl is removed EtClO can be stored at room temperature for several hours. Beilstein mentions that in storage EtClO decomposes under formation of ethyl acetate and other products.

Nothing is mentioned of shock-sensitivity. However, they can't be that friendly to a hammer blow. The perchlorate brothers are most probably worse in that aspect! Beilstein says the oil ethyl perchlorate (C2H5.O.ClO3, 110% of PA in trauzl test) explodes with severity when dry simply by pouring it from one container into another.
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[*] posted on 29-3-2008 at 17:27


Ethyl hypochlorite can also be formed from dichlorourea. For this see: Chattaway in The Chemical News and Journal of Industrial Science, V. 98, p. 166. The reference mentions that dichlorourea is made when chlorine is passed into a cooled saturated solution of urea. The equation should be: (NH2)2C=O + 2 Cl2 -> OC(NHCl)2 + 2 HCl [can neutralize with calculated mild NaHCO3 solution]. The reaction occurs without any significant heat development and the dichlorourea crystallizes out as a white powder (transparent plates). Chattaway thus expects it to be highly explosive, but it isn't. Heating does not cause it to explode itself, it does decompose around 83 deg., liberating nitrogen chloride, which can detonate if it is not allowed to escape or the temperature raised a few degrees, like when heating dichlorourea on a test tube over a water bath. It is preservable in a dry atmosphere for some time, although it is not very stable. But pretty safe to handle, easily sol. in water, alcohol and ether, and is very reactive.

It readily hydrolyzes to nitrogen chloride, CO2, N2, and NH4Cl. If it is dissolved in water or kept in a moist atmosphere, the hydrolysis slowly occurs at normal temperatures, but occurs very rapidly at 30 deg.C. Both acids and bases accelerate the hydrolysis. Dichlorourea gives iodine from hydroiodic acid, chlorine from hydrochloric acid, and reacts with ethyl alcohol to give ethyl hypochlorite and reforming urea in each case.

Some more reactions: Diurea: excess ammonia added to aq. soln. of dichlorourea hydrolyzes and liberates N2 and carbonate, also forming diurea, CO(NH.NH)2CO, which separates in considerable amounts as a sparingly soluble crystl. powder. Heating diurea with excess strong H2SO4 at a little over 100 deg.C., hydrolyzes it, CO2 is evolved and hydrazine sulfate is formed which crystallizes out perfectly pure in almost theoretical quantity by cooling and addition of small amount of H2O. Reaction between dichlorourea and solution of caustic potash is very energetic, releasing N2 with effervesence where excess ammonia and alkaline carbonate remains in the liquid: 3 CO(NHCl)2 + 12 KOH -> 3 K2CO3 + 2 NH3 + 6 KCl + 2 N2 + 6 H2O. For more see reference.

From the dry compound, although risky itself (potential NCl3 when wet with warm, though according to Chattaway this formation shouldn't be a problem in a frozen icebath and at low temperatures) this could be "safer" than just from the chlorine directly. A two-part process like this could give less hazards to worry about at a time (toxic chlorine and ethyl hypochlorite and acidity). What do you guys think?

And just after they talked about it decomposing in storage and decomposability in aqueous solutions, Beilstein also says ethyl hypochlorite solutions in carbon tetrachloride are "pretty stable" (Reference they give for this is: T., Macm., Ga., Go., E., Di., Berichte der Deutschen Chemischen Gesellschaft, 58, p. 572). I would suppose that means decomposing.
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[*] posted on 26-5-2008 at 22:01


Do you think the NaOH/EtOH/Cl method would work with Drano crystals? Or are they too impure? They have a blueish tint to them. Could this be copper?



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[*] posted on 8-11-2008 at 12:36


U.S. patent 2,694,722 "preparation of alkyl hypochlorites"
http://www.google.com/patents?id=PQ5mAAAAEBAJ

the method outlined by Axt and also succeeded by davster seems the simplest one to me. ;)
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[*] posted on 8-11-2008 at 16:45


They are just using CO2 as the acid there. I don’t trust organic hypochlorites, they seem extremely unpredictable, much less so than NCl3, though this is likely in the case of trace amounts of acidity. But they are not violent chlorinators of fats like NCl3, i.e. isoPrClO doesn't detonate on contact with olive oil.
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[*] posted on 11-11-2008 at 05:36


Hypochlorous Acid and the Alkyl Hypochlorites
M. C. Taylor, R. B. MacMullin, and C. A. Gammal
J. Am. Chem. Soc.; 1925; 47(2) pp 395 - 403;

;)

Attachment: HYPOCHLOROUS ACID AND THE ALKYL HYPOCHLORITES.pdf (618kB)
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[*] posted on 27-6-2009 at 18:24


Some comments on alkyl hypochlorites from Sandmeyer:

Speaking of EtClO, Very surprising is the action of sunlight. If the ester is subjected to it, so after a few minutes a turbulent boiling results which ends in explosion. ... illuminating in ice water does not change the course of decomposition in the slightest. But even with very weakly diffused light it can be kept only for a few hours, after which time it heats itself and suddenly boils, but without explosion, with the largest part volatilizing leaving behind a liquid which smells strongly acidic like acetic ether.

He describes what happens when the hypochlorite is confined: It may thus only be kept in loosely closed vessels. Two decigrams [so, only 0.2g] for the purpose of analysis was melt-closed into a small glass bulb, which was then placed in a middle-large beaker. Through the resulting decomposition, it was able to completley shatter the bulb as well as the beaker due to the explosion that followed.

Some reactions: In its further reactions, the ester showed greatest similarity to hypochlorous anhydride, for example, it acts on ammonia, and several organic compounds such as phenol or aniline, chlorinating and oxidizing them nearly explosively.

Decomposition described from Cl2 during its preparation from the NaOH and Cl2 method: ... One needs to make sure that the bubbling of chlorine is interrupted before the bubbles pass into the liquid, otherwise its decomposition begins several minutes after its preparation.

The methyl ester is also described, and it is just as ever bit highly explosive in the condensed liquid phase as the gas phase (contrast to the higher esters).

These observations are also a nice testament to the inherent unstable nature of primary alkyl hypochlorites which will crackle and boil away on standing even without acidity, under very weak lighting and in an unconfined state.

Attachment: Ber. 19, 857.pdf (316kB)
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[Edited on 28-6-2009 by Formatik]
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[*] posted on 28-6-2009 at 16:16


So far my attempts at synthesizing this organic hypochlorite appear to be in vain. I'm hypothesizing that I'm not adding enough chlorine to my solution, as .5mol of Cl2 occupies 11.2 liters, if I recall correctly.
I prepared an EtOH-NaOH solution from 37.897ml 70% EtOH(~23g EtOH, or approximately .5mols) and 20g NaOH(approximately .5mols)
The NaOH was in the form of very light granules and didn't completely dissolve at 20degC.
This mixture was placed in a 100ml pyrex beaker.
As for my chlorine generator, I used PE pipe with an interior diameter of 3/8" to fit snugly around the nipple of a 250ml vacuum filtration flask. To this I affixed a length of 1/8th" PE tubing and firmly adhered the two using a heat gun. I added a gross excess(100g) of crushed TCCA, a common pool chlorinator to the bottom of the flask. I added 50ml 35% HCl and chlorine was immediately produced in a fast-paced reaction. The top of the filter flask was capped and the chlorine gas was led into the solution.
After chlorine gas was done being produced, a small amount of moderately viscous, yellowish fluid appeared on top of the cloudy NaCl/NaOH laden layer below. The yellow fluid was completely immiscible with the fluid below, had a strong, unpleasant, and sickly sweet odor, and burned unimpressively. The yellowish fluid was more dense than H2O but completely miscible with tap water.
Can anyone identify the product of the reaction? It doesn't sound like EtOCl.


[Edited on 2009.6.29 by DyD]
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[*] posted on 28-6-2009 at 20:49


That's it. The esters - except the methyl one - burn when ignited unconfined. They also smell horrible. They should be pretty poisonous since they are quite nauseating.
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[*] posted on 4-5-2012 at 21:30


Ethyl hypochlorite can be formed by passing chlorine into a solution of water, ethyl alcohol, a limited quantity of sodium hydroxide. Methyl hypochlorite is a dangerously sensitive explosive.
Journal of the Chemical Society, Volume 50, p607

Ethyl hypochlorite reacts with sulfur dioxide to form ethyl chlorosulfonate, CH3-CH2-O-SO2Cl, which is similar to ethyl sulfate in reactivity, although I am not sure if it is as hazardously toxic.

I suspect that the alkyl hypochlorites are not very chemically stable, especially when heated, because the chlorination of ethanol under different conditions can result in the production of chloral hydrate instead.

[Edited on 5-5-2012 by AndersHoveland]
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[*] posted on 24-5-2012 at 18:37


I had an idea for making a hypochlorite ester; not as an explosive, but rather as a chemical regent for other reactions.

Chloroform can be condensed with acetone to form chlorobutanol. Chlorine gas could then be passed into a solution of the chlorobutanol in more chloroform to form the organic hypochlorite ester, 1,1,1-trichloro-tert-butyl hypochlorite. Do not know whether this would be explosive, but the advantage would be chemical stability.

Quote:

Very recently Chattaway and Backeberg (2) described other hypochlorite including the various isomeric propyl, butyl, and amyl hypochlorites. They found that the hypochlorites of tertiary alcohols are much more stable then those of the corresponding primary and secondary alcohols.

"Hypochlorous Acid and the Alkyl Hypochlorites",
M. C. Taylor, R. B. MacMullin, C. A. Gammal



1,1,1-trichloro-tert-butyl hypochlorite
Cl-O-C(CH3)2CCl3

Hypochlorite esters of this type may be useful regents for preparing hypochlorite salts free from chloride ions. The chemistry of hypochlorites when they are acidified is significantly different when there is no HCl around to reduce the HOCl. Acidifying normal hypochlorite bleach solutions (NaOCl, NaOH, NaCl) generally just makes Cl2 gas. The reaction of hypochlorite esters with NaOH can make a solution of pure NaOCl (without any NaCl), in addition to the respective alcohol of course.

Quote:

Ethyl hypochlorite reacts with alkali [a base] to form alkali hypochlorite free form chlorides.

The resulting ethanol, however, is likely more vulnerable to oxidation by hypochlorites (haloform reaction), whereas tertiary alcohols (such as tert-butanol) are much less vulnerable to oxidation.
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thumbup.gif posted on 22-3-2013 at 12:13


mmm hypochlorinated olive oil :P
I suppose it should have been given just a bit more light and time :D

W/regards to ethyl chlorosulfate it's been used in WW1 as an irritant (as opposed to a poison gas), so it doesn't seem to be very toxic.
I didn't see any acute toxicity warnings in the msds's.

the acetic acid is there seemingly just to protonate the hypochlorite to promote the forward reaction (which with the HOCl is over in two minutes) and be a backstop to prevent significant OH- presence that would decompose the product. The hypochlorite solutions that everyone uses are ridiculously dilute - five or thirteen or whatever percent - and of course they would not store otherwise. Therefore it seems that using 80% and especially glacial acetic is a waste as the amount of water in your reaction would be huge no matter what acetic you use. stoichiometrically the amount of acetic is the same as the amount of OCl-, so using >30% seems pointless.
if you desire 'glacial' acetic for any purpose, the amateur's best friend here seems to be cold rather than heat - freeze it out. the freezer, or if you're lucky enough to live in a place where the nightfrosts outmatch the freezer, the great outdoors (put yer acid far from cover and especially dwellings - the colder the better)
it occured to me that straight TCCA might be a better option than making chlorine gas from it, piping that goodness, then disproportionating it with alkali, then again protonating it with acetic acid... TCCA has a labile hydrolysis equilibrium, so the main consideration here is not kinetic but thermodynamic. i thought using sodium carbonate (or bicarbonate?) would deprotonate it as it forms, forcing the equilibrium towards completion. The presence of cyanuric anion would also have a 'salting-out' effect and force the product from the solution. Reversely if you're using bleach, sodium bicarbonate would serve as an acid, to force the reaction to completion and serving as a backstop against product decomposition (without this you have NaOH left where you previously had NaOCl)
Another thought that occured to me... what if you own sodium bromide? ;) The hypobromites are an interesting possibility, especially with regard to their flavors (it seems i was wrong to assume the hypochlorites would simply smell of 'chlorine', as no one actually mentioned it even once) - to compare the smells of ROCl's and ROBr's seems interesting :o:cool: These customers seem like they would be more unstable because the O-Br bond is weaker, but i wonder if tert-butyl hypobromite would not be stable still.
Last, but not least, the discussion about glacial acetic put me onto another thought... what about *acetyl* hypohalites? Freeze-concentrated or glacial acetic acid might promote its own reacton (going via autoprotonation), or a little bit of sulfuric acid would help otherwise (there is also bulk consumption of acid to protonate your hypochlorite, so using sulfuric to do this seems more efficient). The different (from R-O-H) properties of the acid oxygens, the effects of hydrogen bonding and the possibility of acetate radical decarboxylation all make it an interesting thought.




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[*] posted on 3-4-2013 at 08:21


I have reviewed some of the literature and the patent and have a different take and recommendation on the best preparation.

First observation is that the action of CO2 on NaOCl forms HOCl, but does so only slowly (this is due perhaps to the fact that H2CO3 is only a slightly stronger acid than HOCl). Other mentioned routes also involve the preparation of Hypochlorous acid, for example, by the action of Cl2 on a slurry of CaCO3 (I recall this being discussing in some patent, where the logic probably is that the Cl2 reacts with water, but not readily, to form HCl and HOCl, whereupon the HCl is readily neutralized by the Calcium carbonate). Also, possible addition of CCl4 to extract the HOCl is another route covered, per my recollection, in a patent for the preparation of HOCl in an organic solvent. The patent cited above also discusses the use of acetic acid on NaOCl, another path to HOCl, but dismisses the procedure apparently due to acidity in the presence of the alcohol.

All of the above, however, feel obligated to immediately encompass the alcohol in the preparation of the HOCl. Not the best idea in my opinion, assuming the patent is correct, as the best conditions occur reputedly in a pH range of 7.0 to 8.0. Also, the organic hypochlorite appears to readily undergo decomposition (even violently/explosively, in the presence of acid/low pH) resulting in loss of yield to say the least. So why is it necessary to place the alcohol in an acidic environment (with Cl2, HCl,..) on the route to nearly neutral HOCl? Also, a bad idea would be to mix the alcohol with NaOCl in place of HOCl. My reason is that the action of NaOCl on the alcohol, ROH (while we are waiting for the H2CO3, for example, to form HOCl) could proceed as follows:

ROH + NaOCl = ROCl + NaOH (very basic)

and the patent notes when pH is over 8, a loss in yield for the organic hypochlorite due to hydrolysis.

So my recommendations:

Step 1. Prepare separately your dilute and nearly neutral solution of HOCl (which need not be free of chlorides). For example, add H2CO3 to Bleach (a mix of NaOCl, Na2CO3, NaCl and a little NaOH), and then wait a day. Or, much more rapidly, add Acetic acid to NaOCl, and then neutralize with NaHCO3. As the patent even mentions one should add water, preparing concentrated HOCl (less stable, liberating HCl) is certainly not a good idea.

Step 2. Add a neutral salt (like NaCl) to the dilute HOCl to foster the salting-out of the organic hypochlorite. This will also increase the 'activity level' of the Hypochlorous acid.

Step 3. Finally, and only now, add alcohol while controling pH (by adding NaHCO3) and temperature in dim light. Note, per Bretherick, Vol 1, page 438, on isopropyl hypochlorite to quote:

"Of extremely low stability; explosions occurred during its preparation if cooling
was inadequate [1], or on exposure to light [2]."

Step 4. Separate out the organic hypochlorite per the usual procedures (to quote from the patent, whatever this means precisely).

It may be a while, however, before I get to try out this procedure on a very small scale on isopropyl alcohol.

[EDIT] There is a possible issue using the hydrolysis of TCCA as the source of the HOCl. As I have elsewhere given a reference, apparently Cyanuric acid (a by product of the hydrolysis of TCCA) is attacked by hypochlorite (in this case, this could include the organic hypochlorite).


[Edited on 4-4-2013 by AJKOER]
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[*] posted on 10-6-2013 at 12:46


I found this old thread and decided to try out a few things. I tried it the easy way, not wanting to bubble chlorine through a NaOH/ethanol solution. I used the method of Axe with acetic acid and bleach.

Preparation of ethyl hypochlorite:
- Take 4 ml of bleach with appr. 13% active chlorine. Put this in test tube A.
- Take 1 ml of acetic acid (80%) and put this in test tube B.
- Take 0.5 ml of ethanol (96%, denatured with 1% MEK) and add this to test tube B. Swirl the contents of test tube B.
- Pour the contents of test tube B in test tube A. Do this not at once, but add in 4 steps. After each time swirl the test tube to mix the solutions.
A yellow oil separates and remains floating on the aqueous layer. The layer only has a thckness of 3 mm or so. The yellow oil is very volatile and quickly disappears (in 10 minutes or so, all of it is gone). The oil has a strong and pungent smell, not very pleasant.
Using a glass pasteur pipette, take some of the oil and put on a watch glass. Immediately keep a small flame near the oil. The oil catches fire very easily and burns explosively violently with a WHOOSH sound. If you wait too long with igniting the oil, then nothing is left, it evaporates very easily.

The experiment was carried out at 15 C, the test tube A cooled under a running tap with water of appr. 15 C as well. No excessive heat was observed.


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

The same experiment was repeated with tert-butyl alcohol instead of ethyl alcohol. Now appr. 1 ml of t-butyl alcohol was mixed with 1 ml of acetic acid (80%) and this mix was added to 7 ml of 13% bleach. Now, a layer of more than 5 mm of oily yellow liquid was obtained. This liquid is much less volatile. Some of this liquid was put on the watch glass and ignited. It can be lighted very easily and burns fiercely with a sooty flame. The smell of the soot is the same as that when chlorinated alkanes are burnt.

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

I did a final experiment, I took appr. 1 gram of KBr and dissolved this in 7 ml of water. To this I added appr. 0.25 ml of the t-butyl hypochlorite. As soon as that liquid comes in contact with the solution of KBr, it turns orange. The liquid becomes more dense and starts to float around in the liquid. On shaking the test tube, the liquid turns red. Red oily drops are formed, which partially sink to the bottom while others move up to the surface. The red drops most likely are either t-butyl hypobromite, or t-butyl hypochlorite, contaminated with Br2. However, no Br2-vapor, nor any smell of Br2 could be observed.

Some of the red drops were collected in a pasteur pipette and transferred to a watch glass and ignited. These drops burn much more violently than t-butyl hypochlorite, the burn is nearly explosive. There is no soot and the flame is not orange, but a weak flash occurs with a pale color (hard to observe, the effect only lasts for a very short time). I think that I made some (impure) t-butyl hypobromite.

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

A final remark: It is important to mix acetic acid with an alcohol and add this mix to the solution of NaOCl. If you add the 80% acetic acid to the NaOCl first, then you get a cloud of Cl2 and the liquid which remains has lost most of its active chlorine and is not useful anymore.

[Edited on 10-6-13 by woelen]




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[*] posted on 10-6-2013 at 14:22


Very interesting woelen, as usual. Didn't you try to prepare these based on TCCA rather than hypochlorite ? Would be interesting if that works. Also - what will be the most stable organic hypochlorite and how stability changes in the row of: from primery to ternary alcohols?
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[*] posted on 11-6-2013 at 04:31


Quote: Originally posted by papaya  
Very interesting woelen, as usual. Didn't you try to prepare these based on TCCA rather than hypochlorite ? Would be interesting if that works. Also - what will be the most stable organic hypochlorite and how stability changes in the row of: from primery to ternary alcohols?


Tertiary hypochlorites are the most stable, example: t-isobutyl hypochlorite




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[*] posted on 11-6-2013 at 06:04


Quote: Originally posted by Adas  
Tertiary hypochlorites are the most stable, example: t-isobutyl hypochlorite

Hmm, isn't this in contrast with organic nitrates?
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[*] posted on 11-6-2013 at 11:49


I also did the experiment with methanol. This is a very spectacular experiment, but you need to use a thick-walled test tube, wrapped in a thick layer of towel!!!

- In a thick-walled test tube, put 3 to 4 ml of bleach with 13% active chlorine.
- In a separate test tube put a little under 1 ml of acetic acid (0.75 ml or so).
- To the acetic acid, add half a ml of methanol and swirl such that the liquids mix well.
- Add the acetic acid/methanol mix to the bleach in two steps and swirl the test tube after each addition.
When this is done, then on swirling a colorless gas is produced. The formation of the gas is accompanied by a fairly loud bubbling noise. Big bubbles of gas escape from the liquid. There is no strong formation of heat.
When swirling does not result in formation of gas anymore, then wrap the test tube in a thick layer of towel. Keep the test tube at an angle of 45 degrees and then keep a flame in front of the open end of the test tube. When this is done, a remarkably powerful explosion occurs. Only the gas explodes, the liquid is not altered. The sound of the explosion is thundering loud! Given the fact that only appr. 20 ml of gas explodes in the test tube, this is quite remarkable. Methyl hypochlorite has no practical application at all, but this is a very nice and impressive demo!

Do not scale up this experiment for added effect

[Edited on 11-6-13 by woelen]




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[*] posted on 25-3-2018 at 03:44
t-butyl hypochlorite


Since I couldn't find too much about the synthesis of t-butyl hypochlorite, I thought I'am just gonna post my attempt, which actually worked.
I used about 230g of a 14% sodium hypochlorite solution, which I just cooled with ice to maybe 5 to 10°C, and then I poured a mix of 30g t-butanol and 35g of acetic acid to it, and that's it. It immediately became cloudy and a yellow green phase seperated on the top.
After seperating that, washing it with NaHCO3 and drying it with Na2SO4, I got about 21g, which is a 50% yield. I had a lot of mechanical losses dure to drying and filtering, so you might be able to improve the yield there.
One important thing is, working with this stuff is no fun, unlike most esters, it doesn't smell nice, it smells more like a war gas. It smells quite like chlorine, but a little different, but it stings extremly in when you inhale a bit and makes you cough, sooo, just be aware of that if you wanna try it yourself, have fun! Oh, and here is a pic of what the ester looks like:



2018-03-25 13.32.41.jpg - 1.9MB
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CobaltChloride
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[*] posted on 25-3-2018 at 06:55


Have you tried any tests to see how explosive it is? I know that methyl hypochlorite is pretty explosive, so this might be as well.
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CharlieA
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[*] posted on 25-3-2018 at 07:02


You obviously got a product. How did you identify it? How do/did you assess its purity? Just curious...
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theAngryLittleBunny
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[*] posted on 25-3-2018 at 07:49


Quote: Originally posted by CobaltChloride  
Have you tried any tests to see how explosive it is? I know that methyl hypochlorite is pretty explosive, so this might be as well.


It isn't explosive, there is just too little oxidizer in the molecule for this to be the case.

Quote: Originally posted by CharlieA  
You obviously got a product. How did you identify it? How do/did you assess its purity? Just curious...


I followed and orgsyn synthesis, I really don't have anything I could analyse it with, but the liquid looks like discribed (yellow green) and floats on water, which is sensible, since the density is supposed to be 0.96g/mL.
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