Sciencemadness Discussion Board - Organic Hypochlorites

Organic Hypochlorites

halogen - 20-4-2004 at 07:37

-So i dont know what to call it-
How about methyl hypochlorite?
CH3OH+HClO-->CH3ClO+H2O
or
(CH3)2O+Cl2O-->2CH3ClO

Edited by Chemoleo: Changed title to something more comprehensible

[Edited on 27-10-2004 by chemoleo]

[Edited on 25-6-2007 by halogen]

KABOOOM(pyrojustforfun) - 21-4-2004 at 19:28

you'll just oxidize the methanol to formaldehyde and formic acid. other primary and secondary alcohols can be oxidized this way. only tertiary alcohols can form hypochlorous esters eg t-butyl hypochloite is a lachrymator and unstable liquide which is formed by passing chlorine in an alkaline solution of t-butyl alcohol (taken from penguin dictionary of chemistry) I don't see why dripping t-butanol into bleach wouldn't work.

halogen - 26-4-2004 at 02:24

1st, bleach is sodium hypochlorite. (I dont know how that would affect it but I think it would)
2nd, There would be a lot of water so wouldnt hydrolysis set in?
Its only speculation, unfortunately.

[Edited on 25-6-2007 by halogen]

DDTea - 26-4-2004 at 10:50

Methyl Hypochlorite can easily be formed by the reaction between Methanol and NaOCl...that's why NaOCl is listed as an incompatible chemical on Methanol's MSDS. It is unstable and explosive, so keep this in mind if you work with it.

Theoretic - 8-5-2004 at 09:51

A good way to prepare methyl hypochlorite would be to react methyl chloride with sodium hypochlorite. This would give a higher yield than methanol + NaOCl, because NaOH formed would partially decompose your product.
I bet that the compound is a liquid, is explosive, stinks of chlorine bigtime, and isn't good for your fingers.

The_Davster - 8-5-2004 at 13:05

I gave methyl hypochlorite a try. A testtube was filled 3/4 with 5.25% bleach and about 5 mL of methanol was added. Small bubbles were released soon and after a day the bleach had lost all its yellow color. I had been assuming that methyl hypochlorite would be an insoluble liquid that would settle on the bottom of the testtube. After 3 days there is no separate layer of methyl hypochlorite. So is it in solution? Or were the reaction conditions not right?

halogen - 9-5-2004 at 09:01

1.Youre assuming I did this before...
2. Could it be that methyl hypochlorite is a gas?
3. Where can you get chlorine heptoxide?

[Edited on 11-1-2008 by halogen]

BromicAcid - 9-5-2004 at 10:54

Quote:

3. Where can you get chlorine heptoxide?

I'm fairly certain that it is not shipped by any means of transportation and is, as a concequence, only produced where it is needed on site, regardless of it being the most stable chlorine oxide. For preparation though the only method that I know of is adding P2O5 to highly concentrated HClO4 until the Cl2O7 is evolved. Me though, I've got better things to do with my P2O5 and 72%+ HClO4 (Which would take a lot of P2O5 to get to release any Cl2O7 anyways). I'm sad to say for the most part Cl2O7 is very much inaccessable as far as I know.

I was wrong

KABOOOM(pyrojustforfun) - 14-5-2004 at 19:41

only tertiary alcohols can form hypochlorous esters

other alcohols can form hypochlorous esters too. but they are toooo unstable
from <a href="http://www.chemistry.mcmaster.ca/~chem2o6/labmanual/expt7/2o6exp7.html" target="_blank">http://www.chemistry.mcmaster.ca/~chem2o6/labmanual/expt7/2o6exp7.html</a>:
<blockquote>quote:<hr>Aqueous sodium hypochlorite (NaOCl), or common household bleach, can be used to oxidize secondary alcohols to ketones. The reaction occurs more rapidly under acidic conditions, so it is thought (we're not actually sure) that the actual oxidizing agent is hypochlorous acid (HOCl), generated by the acid base reaction between sodium hypochlorite and acetic acid...
...In alcohol oxidations, the reaction probably proceeds via E2 elimination of the alkyl hypochlorite produced by reaction of the initial alcohol with HOCl<hr></blockquote>
also take a look at:
<a href="http://www.wellesley.edu/Chemistry/chem211lab/Orgo_Lab_Manual/LabManual/week9.html" target="_blank">http://www.wellesley.edu/Chemistry/chem211lab/Orgo_Lab_Manual/LabManual/week9.html</a>
------------------
<a href="http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv4p0125" target="_blank">tert-Butyl Hypochlorite synthesis</a>
<a href="http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv5p0184" target="_blank">another (better) tert-Butyl Hypochlorite synthesis</a>

I still think t-alkyl hypochlorites are more stable 'cause they can't be oxidized=dehydrochlorated. in other words there's no hydrogen attached to the carbon which has the -OCl group to be removed with the Cl as HCl

I should apology to halogen

KABOOOM(pyrojustforfun) - 30-6-2004 at 13:23

because of my post in his "oxyazides?".
now that is in Detritus and I couldn't reply to that thread...I see he hasn't been online since 23-5-2004 so I thought he might have subscribed the thread he has made... sorry for my being off topic anyway.

[Edited on 30-6-2004 by KABOOOM(pyrojustforfun)]

BromicAcid - 26-10-2004 at 17:57

Often I'll just randomly open up a chemistry book and browse through the index looking for anything that catches my eye, today I was browsing the ACS archives, specifically from Feb, 1925 when I ran across an article titled:

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

Huzzah, I rejoiced, I remembered this thread and made some photocopies. No mention of any explosive properties of methyl or ethyl hypochlorite although preparations and properties are described.

Here are some highlights:

Quote:

Our knowledge of the alkyl hypochlorites begins with the work of Sandmeyer (1) in 1885 who prepared methyl hypochlorite by chlorinating a mixture of alcohol and sodium hydroxide. The esters are relatively insoluble in water, and separate as yellow oils, unstable towards heat and light and possessing very irritating odors. 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.

.......

The unstable nature of the pure material made is unsuitable for use in a technical process, but it was found that a solution of the ester in carbon tetrachloride, or any other solvent immisible with water, was stable for several hours below 25C.

.......

Twenty to 25 g of precipitated chalk or ground limestone is added to a liter of water and the mixture chlorinated with stirring until about 25 g of chlorine has been absorbed. With this excess chalk, the available chlorine is present almost entirely as hypochlorous acid. The excess of carbonated is filtered off and the filtrate shaken immediately with and equal volume of carbon tetrachloride containing about 2% or more of ethyl alcohol. About two minutes shaking is sufficient to cause reaction between hypochlorous acid and alcohol to such an extent that the greater portion of chlorine is removed as ethyl hypochlorite from the water to the solvent layer which can then be separated.

.......

A pure solution of ethyl hypochlorite in carbon tetrachloride made in this way lost only 10% of the available chlorine on standing for two days at 20C in the diffused light of the laboratory.

.......

Such a solution of ethyl hypochlorite liberates iodine quantitatively form potassium iodide and acetic acid. It reacts rapidly with phenol, aniline, methylaniline, sodium picrate and similar easily oxidized substances and in general shows most of the reactive properties of hypochlorous acid. When ethylene is passed into a mixture of ethyl hypochlorite, slightly acidified water and carbon tetrachloride, absorption occurs with the probable formation of ethylene chlorohydrin, which dissolves in the carbon tetrachloride. We did not, however, carry out this experiment to the point of isolating the chlorohydrin.

.......

Ethyl hypochlorite decomposes on standing, giving high yields of ethyl acetate, probably by the following steps:

C2H5OCl ----> CH3CHO + HCl
CH3CHO + C2H5OCl ----> CH3COOC2H5 + HCl

.......

Ethyl hypochlorite reacts with alkali to form alkali hypochlorite free form chlorides.

.......

The constitution of the ethyl hypochlorite has been shown by our work to be that of a true ester. The oil was prepared by the method of Sandmeyer, washed with sodium bicarbonate solution and dried with calcium chloride. It was kept in the dark at 0C during analysis. Six different determinations of chlorine content were made, of which the average value is 44.13% as compared with the calculated value of 44.05.

.......

If the substance had been a double compound of one molecule of alcohol to one of acid, EtOH.HOCl, the molecular weight would have been 98.5. This shows that the compound which we have assumed to be ethyl hypochlorite is a true ester. Although, as we will show below, hypochlorite is largely hydrolyzed in water, its insolubility in water permits esterfication by removing the reaction product from the water phase. Solvents aid in this removal of ester from water, giving even greater yields of ester then when no solvent is used.

.......

We investigated the formation of hypochlorous acid esters of various alcohols by the method which Sandmeyer used for preparing ethyl and methyl hypochlorite. Yellow, insoluble, unstable oils were obtained with propyl, isopropyl, sec.butyl, tert.butyl, isoamyl, sec.amyl, and tert.amyl alcohols. Benzyl alcohol is oxidized to benzaldehyde by hypochlorous acid. Glycerol, glycol and ethylene chlorohydrin gave no appreciable amounts of insoluble oil.

(1) Sandmeyer, Ber., 18, 1767 (1885); 19, 857 (1886)

(2) Chattaway and Backeberg, J. Chem. Soc., 123, 2999 (1923).

Hope someone found this interesting. So as of now I've read up on alkyl hypochlorites, alkyl perchlorates, and understandably not alkyl chlorites, however what about alkyl chlorates?

Blaster - 27-10-2004 at 11:35

I just thought I'd add a bit to Bromic Acid's post above.

To my knowledge the alkyl chlorites and chlorates have not been isolated. They are just too unstable. The extreme instability of the perchlorates has of course been discussed!

Hypochlorous esters are best prepared by the action of Chlorine gas on a cooled solution of the appropriate alcohol in 10% aqueous NaOH. However, they can also be prep'd by the action of the alcohol on a conc. solution of HOCl or by passing Cl2O into the alcohol and precipitating the liquid ester with water. (Sandmeyer's ref as quoted by Bromic Acid).

The boiling points are 12'C and 36'C for the methyl and ethyl esters respectively. They explode (far less violently than the perchlorate esters) when ignited or exposed to a bright light and decompose on standing to the appropriate aldehyde.

Axt - 28-10-2004 at 14:13

I made EtOCl a while ago, by the method Blaster has just mentioned. Using HCl/MnO2 to generate the chlorine, with the EtOH/NaOH in a measuring cylinder in a jar of cold water containing prilled AN. The yellow EtOCl seperates quickly and is easily extracted from the top of the liquid.

Acts simular to MEKP, though less violent and accompanied by a large fireball. Below is the ignition of ~5 drops if I remember right.

<center><img src="http://www.sciencemadness.org/scipics/axt/etoclapp.jpg"> -----> <img src="http://www.sciencemadness.org/scipics/axt/etocl.jpg"></center>

[Edited on 9-12-2005 by Axt]

chemoleo - 29-10-2004 at 08:19

I wonder if there are similar hypochlorous esters for i.e. amino, or nitro ethanol? This should make the compound more interesting, but probably also less stable.
Although I have my doubts this could be prepared by direct chlorination.
Anyway, good work, who'd have thought it's this easy!

Axt, could you comment on the shock sensitivity, and flammability? How did you ignite it?

Axt - 29-10-2004 at 18:58

I cant give any more info, only did it once. I chucked a match at it. It takes fire very easily.

If its extracted straight off the top and ignited it always leaves burney bits left behind, like MEKP does if it isnt dried. Salting it out helps somewhat, but I expect to get remnant EtOH out you would have to distill it, not something im going to do.

Question...

kazaa81 - 1-11-2004 at 15:15

Hallo to all,
I post a question that is in my mind: won't be common blach too dilute to react with ethyl or methyl alchohol?
I've not yet tested this experiment because if 5% NaClO react indeed, I will be unprepared...
Any other suitable organic hypochlorites?

Thanks for collaboration.

The_Davster - 17-12-2004 at 22:31

I gave ethyl hypochlorite a go this evening. 25mL of 10% NaOH was mixed with 25mL of denatured ethanol. These 50mL of liquid was placed in a graduated cylinder and placed in a snow/water bath. Chlorine produced via HCl +NaOCl was bubbled through for around 1h. The ammount of chlorine bubbled through was produced by the complete reaction of 60mL 31.45% HCl with excess bleach. After around the half hour mark I added around 1g more of NaOH to the cylinder because I was worried that HCl gas carry-over from the gas production setup would react with the NaOH instead of the chlorine reacting with the NaOH. The solution was a light green upon completion of bubbling the chlorine through. Smelled unusual, not bad kind of pleasant actually although the smell was different from that of ethanol. I recieved no separate layer of liquid on top in the graduated cylinder. I know that some ethyl hypochlorite was produced as when I lit samples of the solution on fire it burned orange instead of the charastic light blue of alcohol. There was still some blue in the flame indicating excess ethanol left over.

Thoughts...(especially on why I got no separate layer formed-I am thinking that I should have bubbled more chlorine through, but I am not completly sure)

EDIT: I need a gas mask *cough*

[Edited on 18-12-2004 by rogue chemist]

Axt - 17-12-2004 at 23:09

Water!

12g NaOH, 12g EtOH, 100ml water, bubble chlorine through that. It only takes a few minutes but its not very interesting. The methanol derivative may be more interesting, though you would have to keep it cool to prevent evaporation, or better, catch the pure stuff in a condenser.

I expect NaOH/EtOH will give the orange flame, as is.

[Edited on 18-12-2004 by Axt]

The_Davster - 17-12-2004 at 23:29

Thanks Axt

Just looking at the ratios that you just posted I realize that I used way to much ethanol and not enough water. "25mL of 10% NaOH" was all the water I had in there

. The ethyl hypochlorite must be semi-soluble in ethanol solutions as well which would explain the green colour but not a separate layer.

I'll try this again tomorow.

I may be able to try the methyl ester soon as well seeing as it is a Canadian winter here so I should have low enough temperatures to keep the methyl ester liquid

The_Davster - 19-12-2004 at 00:06

I tried it again, with the ammounts Axt posted. It worked quite well, I got nearly 8mL of ethyl hypochlorite. It smells quite terrible actually. Just from washing out the glassware inside the house the room I was in reeked of it afterwards. After I removed the majority of the EH from the surface with an eyedropper I left about 1mL of EH on the surface of the graduated cylinder. I bubbled chlorine through for quite some time after that and about 30min later the EH was on the bottom of the graduated cylinder. Chloroethyl hypochlorite anyone?

Does anyone know anything about the toxicity of ethyl hypochlorite? I spilled a single drop on my wrist( just missed the glove) and I am feeling quite lightheaded right now, I feel quite strange and am not thinking to clearly right now. I even had difficulty remembering what the compound I had just made was when I started posting this reply.

mark - 13-1-2005 at 05:14

Could you be sensitive to chloroform? I guess this reaction may be possible. Chlorine on sodium hydroxide to produce sodium hypochlorite reacting with ethanol to produce chloroform and getting suspended in the ethyl hypochlorite?

Why I am at I once tried it, except I used no water. After four tries I gave up but some gave me a lot yellow liquid that didn’t burn different to Metho except using a yellow flame. The solution was prepared by bubbling chlorine through EtOH and NaOH solution. The solution was prepared by dissolving NaOH in warm EtOH. On one I did actually get a crystalline suspension that dissolved after more chlorine was dissolved in it.

Esplosivo - 13-2-2005 at 05:26

I was rereading this thread and it struck me that on passing chlorine through ethanol will cause the ethanol to be oxidized to ethanal, which will then be chorinated to produce chloral as I've read. The presence of NaOH is known to 'convert' chloral into chloroform. Why doesn't this occur here also? I mean, what was produced by Axt is surely not chloroform (which is relatively non-combustible). It can't be chloral either due to the presence of a hydroxide.

In short, what is/are the differences in the synthesis of chloral and the organic hypochlorites? I'm confused.

BromicAcid - 14-2-2005 at 11:07

I didn't know it till today, but orgsyn.org has two preparations for tert-butylhypochlorite, it is of course more stable then the methyl and ethyl versions, apparent by the fact that the creators of one of the procedures made batches in upwards of 700 ml at a time, it is interesting to note that the compound decomposes slowly to methyl chloride and acetone, and that irradiation of the product with UV light leads to evolution of the gasses so fast and such a marked increase in temperature that the ester begins to boil.

Axt - 15-2-2005 at 01:04

I've no answer Esplosivo, but its made by the reaction of hypochlorite rather then free Cl2, Ethyl hypochlorite also seperates readily from a solution of NaOCl/ethanol/acetic acid. Its not much better then a very flammable liquid, though explodes when confined.

If you have PATR, check Vol 7, H 262 for a decent description.

mark - 28-3-2005 at 04:12

Can anyone specify the conditions under which the reaction takes place? With the chlorine method all that seems to happening, Is Cl2 on NaOH forms NaOCl than the NaOCl reacts with EtOH to form C2H5OCl (right?). Therefore there is no need for an acid?

However, in Axt’s second method he uses an acid? Is this just to catalyze the reaction or is there something I am missing? If it just a catalyst is their another acid that can replace it?

S.C. Wack - 21-7-2005 at 14:09

And once you do have the hypochlorite, what can you do with it, other than chlorinating amines as in OS? Well, you can chlorinate alkanes.

[Edited on 21-7-2005 by S.C. Wack]

Attachment: jacs_82_6108_1960.pdf (673kB)
This file has been downloaded 1868 times

S-Bursic - 19-6-2006 at 04:23

Me and my friend tried to make C2H5OCl and it didn't work. We didn't measure anything so is that the reason that it didn't work? We put in about 50 ml of H2O and then then about 8 ml of conc. NaOH solution. After that we added the same amount of C2H5OH. At that point a white precipitate was formed. We think that was sodim ethoxide. Then we bubbled chlorine though the solution for 5 to 10 min. Nothing was formed. The only thing was that the sodim ethoxide dissolved. Can you please tell me what went wrong except the measurment?

Zinc - 18-10-2006 at 11:51

Quote:

Originally posted by Axt
Ethyl hypochlorite also seperates readily from a solution of NaOCl/ethanol/acetic acid.

Axt Did you try it? If yes what ratios did you use?

Axt - 22-10-2006 at 21:42

Quote:

Originally posted by Zinc
Axt Did you try it? If yes what ratios did you use?

Yeh I tried it, it worked but I dont remember anymore about it. Other then it seperates quickly as the yellow top layer, and decomposes pretty quickly as well. Forget the ratios. try 20ml glacial acetic acid, 80ml 12.5% NaOCl and less then stoichiometric EtOH.

Zinc - 12-1-2007 at 06:52

Quote:

Originally posted by Axt
try 20ml glacial acetic acid, 80ml 12.5% NaOCl and less then stoichiometric EtOH.

Can I use 80% acetic acid instead of glacial acetic acid?

Sauron - 12-1-2007 at 07:06

t-Butyl hypochlorite is in Org.Syn, which see.

Maybe they will have comments and references pertinent to the others.

First of two Org.Syn procedures attached below

[Edited on 13-1-2007 by Sauron]

Attachment: CV4P0125[1].pdf (140kB)
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Sauron - 13-1-2007 at 06:51

Second Org.Syn procedure:

In both cases let me call your special attention to the safety precautions

Good hood is mandatory.

Dim lighting is required

Total avoidance of contact with rubber is required.

Do not heat above boiling point.

Etc etc.

[Edited on 13-1-2007 by Sauron]

[Edited on 13-1-2007 by Sauron]

Attachment: CV5P0184[1].pdf (144kB)
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Zinc - 5-1-2008 at 17:20

Does anyone know how to prepare amyl hypochlorite and what are its properties?

tr41414 - 10-1-2008 at 01:02

GAA obivously is diluted when mixed with hypochlorite solution...

Is it possible to use another/maybe inorganic acid, or does acetic acid work as some kind of solvent/protectant?

Norrys - 14-1-2008 at 10:57

I tested Axt`s Method, with 80% acetic acid instead, and it worked very well.
I also tried to ignite some EtOCl. In my first test I put some drops on a piece of paper and touched it with a match. It made a big orange fireball, as it schould do. But when I droped a little bit on a stone, I observed a selfignition, and the flames burned the EtOCl, which was left in the pipet.

Because it was confined, it exploded and destroyed the glass-part of the pipet completely. Luckily I wasnt hit by a schrapnell.
But how could it ignite by only beeing droped on a stone?

[Bearbeitet am 14-1-2008 von Norrys]

halogen - 14-1-2008 at 12:18

Perhaps something was on the stone?

Norrys - 14-1-2008 at 12:23

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.

Norrys - 14-1-2008 at 13:29

Quote:

Ursprünglich verfasst von Zinc
What ratios did you use?

Look at Axt's Synthesis:

80ml NaOCl to 20ml Acetic acid and less than stoichiometric ethanol.
But I made a smaller batch, ~20ml NaOCl and ~5-6ml Acid, usw. Although that was a very small batch, it worked well and was enough to make my ears ring, when it exploded

Bromine - 15-1-2008 at 08:14

Today try synthesis of EtOCl two times by bubbling chlorine trough NaOH H2O EtOH solution. Chlorine was generated with TCCA + HCl.

In first try all EtOCl vapourized before I could extract it with droper. In second atempt I provided better cooling and got few small self detontions when bubbling chlorine troug soultion.

Norrys - 15-1-2008 at 11:38

Quote:

Ursprünglich verfasst von Zinc

What conc. of NaClO did you use?

Roughly 13%. Nearly the same concentration, as the bleach, which was used in Axt's synthesis. Please dont ask, how much ethanol I used, I didn't measure it, because it was a pretty litle amount, just some ml, maybe two or three.

tr41414 - 15-1-2008 at 12:12

I have tried a modified synth today... I tried using oxalic and formic acid, with bad (no)results.
Oxalic is not a good choice as it isn't realy soluble in the mix, so it is a bit problematic to add and forms insoluble oxalate which makes the rxn "dirty". Even if there was some product it couldn't be isolated.
The reaction with the formic acid foamed quite a lot... After addition the testtube was shaken a bit which resulted in a geyser of hypochlorite/water/chlorine/co2(?) going into air

The tube got very warm... Maybe the formic is oxidised to co2 or sth like that...

@Norrys: Does the AcOH rxn foam? Did you use cooling and at what rate did you mix the reactants?

Norrys - 15-1-2008 at 14:31

Quote:

Ursprünglich verfasst von tr41414
@Norrys: Does the AcOH rxn foam? Did you use cooling and at what rate did you mix the reactants?

No I didn't use any cooling, but ist frosty cold (~3°C) in my lab. It didnt foam much, but it would have done so, if I had put all the acid at once into the cylinder.
I firstly mixed a little amount of ethanol with 20ml 13% NaOCl and added then ~5-6ml of 80% acetic acid dropwise with a pipet. When I aded it a bit too fast, gas evolved, but no foam formed.
Formic acid is said to be an reducing agent. So you can't mix it with an oxidiser, such as sodium hypochlorite solution.

[Bearbeitet am 15-1-2008 von Norrys]

tr41414 - 19-1-2008 at 13:50

I made ethyl and isopropyl hypochlorite today using that metod

There is absolutely no foam made and the hypochlorite readily separates as yellowish oil, although i didn't get much of the ethyl- product (isopropyl went much better)...

Zelot - 27-1-2008 at 21:50

Do the reactants have to be pure? It would be a real money-saver if I could buy bleach, rubbing alcohol, and white vinegar to make this. Otherwise I would have to shell out for lab-grade chems.

tr41414 - 28-1-2008 at 01:49

Acetic acid should be bought I think (lab/photo chemical), I used lab grade GAA... You could also concentrate it by yourself using solvent extraction or make acetate, dry it and then free the acid using H2SO4...

Zelot - 28-1-2008 at 16:27

*sigh*

It's just that a gallon of 5% vinegar only costs $4. Maybe I'll try it with low concentrations first. If it doesn't, then I guess I'll boil some down to 30%. Wait, how do you salt the water out? I just don't like to try experiments that no one else has done before, because it might blow up in my face.

[edit]

In one of the links previously provided, it said to bubble chlorine through a mixture of alcohol and lye. This would produce sodium hypochlorite. So, is the acetic acid really needed? Also, doesn't sodium hypochlorite react with acetic acid to give chlorine gas?

[Edited on 1/28/2008 by Zelot]

tr41414 - 29-1-2008 at 12:59

You can't distil acetic acid without a very good column, also solvent extraction is way quicker... Add IPA (say 2ml per 10ml of 5% vinegar), salt to your acid, then separate IPA+acid layer, after removal of IPA, you should be left with quite high conc. acid...

Zelot - 29-1-2008 at 14:30

How would you remove the IPA? By heating? Would it be about 80%? Wait, I wouldn't need to remove the IPA, because I'm making isopropyl hypochlorite. D'oh!

Also, is it able to DDT, or do you have to use a primary?

What temperature does the reaction have to take place to avoid it being in a gaseous state? I thought I would do it outside, where it never goes above 50 degrees F, but no oily liquid was recovered.

[Edited on 1/29/2008 by Zelot]

[Edited on 1/30/2008 by Zelot]

PHILOU Zrealone - 30-1-2008 at 08:55

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.

With sensitive unstable materials it is not unusual to have surface catalysed decompositions...so what need to be present is a finely divided solid compound and...it goes booom.
One of the main example that comes to my mind is diazomethane...glas rod may allow this to self decompose catalytically to D2D.

It is also the case with sensitive radicalar reactions like Cl2/H2 and H2/O2, under certain circumstances on unpolished glas reaction is catalysed, it heats up, then speed of reaction becomes higher what heats further and faster up into runnaway, flame and sometimes explosion.
Hypochlorite ester are prone to radicalar reactions...

:cool:

[Edited on 30-1-2008 by PHILOU Zrealone]

Zelot - 20-3-2008 at 05:53

Sorry to bring this thread up again, but I thought this might be a bit interesting:

http://books.google.com/books?id=y5mZrW1KB_AC&pg=PA239&a...

On page 242, It has a diagram showing the electrochemical preparation of isopropyl hypochlorite.

Zinc - 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.

Zelot - 22-3-2008 at 19:37

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

Zinc - 23-3-2008 at 04:42

I think there is no difference.

Formatik - 25-3-2008 at 23:00

Quote:

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.

Formatik - 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.

Zelot - 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?

kazaa81 - 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.

Formatik - 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.

kazaa81 - 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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Formatik - 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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Attachment: Ber. 18, 1767.pdf (190kB)
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[Edited on 28-6-2009 by Formatik]

DyD - 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]

Formatik - 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.

AndersHoveland - 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]

AndersHoveland - 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.

Theoretic - 22-3-2013 at 12:13

mmm hypochlorinated olive oil

I suppose it should have been given just a bit more light and time

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

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.

AJKOER - 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]

woelen - 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]

papaya - 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?

Adas - 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

papaya - 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?

woelen - 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]

t-butyl hypochlorite

theAngryLittleBunny - 25-3-2018 at 03:44

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:

CobaltChloride - 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.

CharlieA - 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...

theAngryLittleBunny - 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.

woelen - 25-3-2018 at 11:30

This stuff is not explosive, but if you ignite it, it will burn very fast. Try this with half a ml or so in a petri dish. Keep the rest of the stuff far away from flame!

theAngryLittleBunny - 26-3-2018 at 05:36

Quote: Originally posted by woelen

This stuff is not explosive, but if you ignite it, it will burn very fast. Try this with half a ml or so in a petri dish. Keep the rest of the stuff far away from flame!

Thanks for that information, I'am on to making t-butyl hypochlorite next, because I heard that burns even faster.

theAngryLittleBunny - 26-3-2018 at 12:07

Okay, I tried something and I feel like it's worth posting. I just made a little isopropyl hypochlorite by mixing 50g of 14% NaOCl with isopropanol and acetic acid, and a green layer immediately seperated on top. But as I went to get a pipette, it suddenly violently decomposed with a crackling sound O.o pretty scary, but also kinda fun. So for anyone playing with hypochlorite esters, just be a bit careful, they might blow up in your face.

[Edited on 26-3-2018 by theAngryLittleBunny]

woelen - 18-2-2020 at 14:46

I just tried making the isopropyl ester, but apparently this is much more unstable than the other esters I tried (methyl, ethyl and tert-butyl). I prepared approximately 0.5 ml, but within one minute of its preparation, it made a very loud crackling noise. Then it became silent again, then it made crackling noise again, and so on. I decided not to wait and see what happens further. I took a towel, wrapped it around my hand (thick layer of a few cm) and then grabbed the test tube and dumped its content in the sink.

The reason I made this is for making a nice video for my website, and I decided to take isopropyl, because that is cheap and very common where I live. I'll refrain from that, however. I'll try the tert-butyl variation again. That one is more stable.

[Edited on 18-2-20 by woelen]

DraconicAcid - 18-2-2020 at 15:05

Hypochlorites are commonly used to oxidize alcohols to ketones....why wouldn't you just get acetone from this reaction?

woelen - 19-2-2020 at 08:50

Under certain conditions, you can make hypochlorite esters, instead of ketones (or aldehydes).

If you add a mix of an alcohol and a weak acid (acetic acid or propionic acid is fine) to a solution of sodium hypochlorite, then you get the hypochlorite ester and not the oxidized alcohol, e.g. ethanol yields CH3CH2OCl.

These hypochlorite esters are quite interesting and reactive compounds. E.g. CH3OCl is a gas, which on ignition explodes, CH3CH2OCl is a volatile yellow liquid, which burns very fiercely when ignited. C(CH3)3OCl also is a yellow liquid, less volatile, but also very flammable. The isopropyl ester apparently is more unstable than the others and decomposes spontaneously.

It is quite interesting to perform the experiments of making CH3OCl or CH3CH23OCl, by adding a mix of acetic acid and the alcohol with 4 to 5 times its volume of concentrated bleach (10+ % of active chlorine). Only try this with small quantities (e.g. 0.5 ml of alcohol and 0.8 ml of acetic acid). Scaling up may lead to very violent reactions and explosions.