Sciencemadness Discussion Board - How to kick start the haloform reaction

How to kick start the haloform reaction

mycotheologist - 21-4-2012 at 14:30

I attempted a haloform reaction today by adding 100g Ca(ClO)₂ to 250mL of water then adding it to a flask. I equipped the flask with a reflux condenser, thermometer and an addition funnel that I loaded with 44mL of acetone and clamped the over an ice bath and stirrer. I slowly added acetone while watching the thermometer but the temperature didn't go up a single degree. I ended up adding half the acetone and no rise in temperature occured at all. I can only conclude that the reaction hasn't started. What can I do to start up the reaction? I added a few mils of boiling water and raised the temperature from 10C to 20C but it just gradually dropped back down to around 10C. I tried manually mixing the flasks contents with a glass rod to help the magnetic stir bar but still no indication of a reaction occuring. Are there any tricks for safely kick starting a reaction in a situation like this. I was thinking of a catalyst but from what I've read, this reaction is difficult enough to control as it is so enhancing the rate might not be the best idea.

[Edited on 22-4-2012 by mycotheologist]

mnick12 - 21-4-2012 at 15:52

Solid sodium hypochlorite? Hmmm considering the instability of hypochlorites I doubt you have pure sodium hypochlorite.

Also the haloform reaction is extremely exothermic and does not require any sort of initiator. But it is hard to say what went wrong without knowing what reagents you used, my guess is your hypochlorite is not hypochlorite at all. Do you mind telling us where you got this hypochlorite?

mycotheologist - 21-4-2012 at 16:23

Sorry, I meant Ca(ClO)₂. I'll edit that. I got a 4kg tub of CaClO from a swimming pool supplier. I added some dilute HCl to a bit of it and it reacted violently, releasing Cl2 but thats the only test I did on it so far. You can smell Cl2 when you open the tub. I'm fairly sure the acetone is pure.

When I started the reaction, I could hear the stirbar moving but it wasn't creating a whirlpool because the liquid was so thick and viscous due to undissolved hypochlorite. I left it for 5 hours and when I came back, there was now a whirlpool in the center of the flask and the liquid there was far less viscous. The temperature had risen by 5C but thats only because the ice bath had melted. I added the rest of the acetone but again, no temperature rise occured. I didn't bother to crush the Ca(ClO)2 granules into powder because I assumed it would just dissolve rapidly but I was clearly wrong. Maybe the reaction is just occuring extremely slowly due to the poor surface area of the granules.

UPDATE: Its been a few hours since I added the 2nd half of the acetone and now all the liquid has thinned out. Temperature still hasn't risen.

[Edited on 22-4-2012 by mycotheologist]

BromicAcid - 21-4-2012 at 18:06

Are you sure it's calcium hypochlorite? Trichloroisocyanuric acid will give the same results with hydrochloric acid and also will not give chloroform (I only mention this because you did test the material so you are already looking into this). The old reaction with calcium hypochlorite used to be the industrial method to make chloroform. I even scanned in the paper somewhere on the site on how to make it on the industrial scale using this method. In my experience there is no 'kick start' to get this reaction going, it should be spontaneous. Do you have some undissolved solids still that might be masking the formation of chloroform?

mycotheologist - 21-4-2012 at 19:27

Yeah I'm positive. The supplier told me he no longer stocks calcium hypochlorite but he had a few tubs left over. He even gave me a label with Calcium Hypochlorite on it. Maybe he made a mistake and accidentally gave me one of those cyanuric chlorides. The tubs weren't labeled. As for undissolved solids, yes. Loads of it. I added the hypochlorite granules to ice cold water (to minimize the amount of Cl2 release) and stirred it manually for a bit. I left to dissolve for at least 30 minutes but I still ended up with a sludge. At first it was more like a paste than a liquid, its been gradually becoming more liquid like. Last time I checked it was more like soup. I'm a bit worried now. I've added all the acetone. Is there a possibility that the reaction hasn't started at all and that it will startup once everything is fully dissolved? If thats the case I suppose all I can do is keep it in the ice bath.

UPDATE: I removed a stopper and smelled it and all I could smell is acetone. I suppose the only logical explanation is that the supplier gave me DCCA or TCCA. I'll find out for sure tomorrow.

[Edited on 22-4-2012 by mycotheologist]

mycotheologist - 22-4-2012 at 05:28

I don't know what to do with this beaker now. I could use the TCCA (or whatever it is) to bubble Cl2 through something but I don't know what side reactions could occur with the acetone in there. On another thread I read the following:

Quote:

I know from experience that heating TCCA with an organic and water can lead to the vapor phase blowing up

So distilling the acetone is out of the question. I suppose I'll just filter out the insoluble material, pour the solution into a pyrex baking dish and leave it outside until the acetone evaporates.

bbartlog - 22-4-2012 at 06:57

I think it's still entirely possible that you have Ca(OCl)2. The procedure you describe would lead to a slow reaction; I'm curious, where did you get the amounts and instructions you're using? Youtube? A patent? They look dubious to me; I note the following:

- if the reaction were to proceed rapidly, the exotherm would overwhelm the reflux condenser. With only 250ml of water I figure you have the potential to release about 300 calories per gram of water, i.e. you'd be boiling off chloroform and acetone.
- the water is not sufficient to dissolve the Ca(OCl)2. On the one hand, maybe this is intended to prevent a runaway; on the other, if it leads to the reaction taking days, I hardly think you're coming out ahead over using sufficient water to dissolve the hypochlorite and just using a larger container.
- In fact there is scarcely enough water to dissolve the calcium acetate that would theoretically form...
- the stoichiometry for your attempt seems weird to me. You have 700mmol of hypochlorite (maximum; actually less as the commercial product always has Ca(OH)2 as impurities, plus it's likely the hydrate). Then you have 600mmol of acetone (35g). But the correct molar ratio of (Ca) hypochlorite to acetone is 1.5:1, not ~1:1. This doesn't just result in unreacted acetone: there is a followup reaction that can take place, because acetone+chloroform+base react to form chloretone/chlorobutanol. It proceeds rapidly enough even at 0C that I would actually expect it to consume any chloroform that was produced in your setup.

If you want to test your compound to see whether it is in fact Ca(OCl)2 or TCCA, I would suggest measuring out a sub-gram quantity into a test tube, adding HCl dropwise with mild warming until evolution of Cl2 ceases (do this outside!), and then seeing whether you have a significant precipitate remaining. Ca(OCl)2 (plus any Ca(OH)2, CaCO3 or other likely calcium impurities) should result only in highly soluble CaCl2, whereas TCCA should leave relatively insoluble isocyanuric acid behind.
Alternatively you could try using 10g of your putative hypochlorite, using 100g of water and 3ml of acetone; dissolve as much of the solid as possible first (at room temperature), then add the acetone and see what happens.

garage chemist - 22-4-2012 at 07:17

You have way too little water in there, mycotheologist!
What I think is happening is that the acetone isn't even mixing into the thick calcium hypochlorite slurry (the presence of large amounts of salts in aqueous solution often strongly reduces the solubility of organic liquids- it's called salting out!). If it did, the reaction would start immediately and erupt as a geyser of steam and boiling solution out of the flask! Even with 10% aqueous sodium hypochlorite, the reaction is so exothermic that it immediately boils off all the formed chloroform when no cooling is employed!
You should use about 0,75 L of water for the amount of hypochlorite you are using, and about 0,5kg of ice, and most importantly, vigorous stirring!

[Edited on 22-4-2012 by garage chemist]

BromicAcid - 22-4-2012 at 07:39

I gave this same quote in a different thread but it is worth repeating (feel free to scale down if necessary):

From 'Thorpe's Dictionary of Applied Chemistry'

Quote:

Manufacture of Chloroform from Acetone and Bleaching-powder.

-This is the process most generally employed. The method differs in minor detail with the various manufactures, but the following may be taken as representatives. The reaction is carried out in a cast-iron still of about 800 gallons capacity, which is provided with stirring gear, steam-coils, and cooling-coils, and is connected with a condenser; 300 gallons of water are run into the still, and 800 lbs of bleaching powder are added through a manhole, which is then securely bolted down. During addition of the bleaching powder the mixture is very thoroughly stirred. (In some processes the mixing is carried out in a separate vessel, and the suspension is strained from the larger unbroken lumps of bleaching powder before being allowed to run into the still.) The container (A in the diagram shown on p. 78) is charged with 70 lb of acetone, which is then slowly run into the bottom of the still by means of a valve B. The introduction of the acetone is accompanied by a rise in the temperature which is not allowed to exceed 110 F., cooling being effected if necessary by stopping the flow of acetone and circulating cold water though the cooling coil in the still. When all the acetone has been introduced the contents of the still are raised to 134 F. At this temperature chloroform begins to distill over. The temperature is then very gradually raised to 150 F., so as to keep the chloroform readily distilling. Towards the end of the reaction the mixture is stirred and the temperature raised until no more chloroform distills over.

The crude chloroform obtained is separated and purified first by agitation with concentrated sulfuric acid. This operation is carried out in the vessel shown in the diagram ; 1,500 lb. of crude chloroform are introduced into the vessel and thoroughly stirred, by means of the agitation gear shown, with 600 lb. of sulfuric acid. The stirring is continued until a sample of the chloroform when thoroughly shaken with pure concentrated sulfuric acid does not impart the slightest color on the latter. The time required for complete purification is usually about 3 hours. The chloroform is next separated from the sulfuric acid and finally distilled over lime. The yield obtained from the above quantities averaged from over 2,000 batches was 124 lb., the highest yield in any one case being 131 lb. Variation in yield is attributed to the varying composition of bleaching powder, though doubtless other factors influence the result. Bleaching powder containing less then 33% of available chlorine gives unsatisfactory results, while samples containing more then 35% of chlorine are also unsatisfactory. The best results appear to be obtained with bleaching powder containing 34% of available chlorine.

garage chemist - 22-4-2012 at 08:17

300 gallons of water = 1137 liters, assuming US gallons are meant
800 lbs of bleaching powder = 400 kg
So this is 400g of bleaching powder with 34% active chlorine in 1,14L water, which is almost the same ratio that the thread starter has used. Now, how much active chlorine does his calcium hypochlorite have?
And did he stir the reaction well enough?

[Edited on 22-4-2012 by garage chemist]

bbartlog - 22-4-2012 at 11:33

400g of bleaching powder in 1.14 liters of water along with 32g of acetone. So with a further fourfold scaling down we would end up using 8g of acetone, less than a fourth of what was used here. It looks like they use an excess of hypochlorite.

I wouldn't advise aggressive stirring once the acetone is added unless A) you know the exotherm is manageable and B) you are removing the chloroform by boiling it off. I realize that stirring will also improve cooling, but in cases where the reaction is already running at the boiling point of acetone it will likely speed the reaction even more. In cases where the reaction is being run ice cold, it is desirable to have the chloroform sink to the bottom and separate itself so as to avoid the base-catalyzed reaction of chloroform with acetone.

Magpie - 22-4-2012 at 13:58

Brewster's procedure calls for the following:

100g calcium hypochlorite, HTH (68% available chlorine)
300 mL water
37 mL acetone

The acetone is to be added slowly through a reflux condenser, and the flask swirled to provide mixing. An ice-bath is to be used as necessary.

IIRC if you are not careful you can have a runaway reaction.

mycotheologist - 23-4-2012 at 05:34

Last night I filtered the contents of the flask and first thing I noticed was that I could not smell chlorine off the white solid material. I poured some dilute HCl onto the solid white material and no reaction occured. I must conclude that this white material is not TCCA/DCCA/hypochlorite and that a reaction actually did occur. I have about 300 mL of liquid but I don't see any layers. I can't really remember what chloroform smells like, but the liquid smells like acetone to me. If there was leftover acetone, would it cause the chloroform and water layers to mix? I don't know if testing this solutions flammability will tell me anything because there was only 44 mL of acetone for 250 mL of water so I'm not sure if the acetone would be flammable at that concentration. Assuming that chloroform has been produced, would adding salt force it to separate into its own layer? I'm going to try that now anyway.

bbartlog - 23-4-2012 at 06:22

Quote:

If there was leftover acetone, would it cause the chloroform and water layers to mix?

No, but it would react with the chloroform, given that there is plenty of Ca(OH)2 in your mix as well.

Quote:

Assuming that chloroform has been produced, would adding salt force it to separate into its own layer?

It wouldn't be miscible with water and even if there were a small amount of acetone remaining I do not think it would be enough to render it miscible. If you had gram quantities of chloroform, it would be in a blob at the bottom.

I gather that your white solid does not dissolve in HCl? That is interesting in its own right since that means it isn't Ca(OH)2 or calcium acetate.
Try heating the solid under a bit of water and see whether it will melt in the range of 80-100C without dissolving (forming a separate liquid layer). If so, then you have chlorobutanol and you know where your chloroform went. Chlorobutanol also has a distinctive medicinal sort of smell. Given the conditions of your reaction - excess of acetone added in two separate shots with five hours in between, together with magnetic stirring sufficient to create a whirlpool - I would actually be puzzled if all your chloroform *hadn't* reacted to form chlorobutanol.

bbartlog - 23-4-2012 at 06:24

Quote:

If there was leftover acetone, would it cause the chloroform and water layers to mix?

No, but it would react with the chloroform, given that there is plenty of Ca(OH)2 in your mix as well.

Quote:

Assuming that chloroform has been produced, would adding salt force it to separate into its own layer?

mycotheologist - 23-4-2012 at 09:05

Quote: Originally posted by bbartlog

I gather that your white solid does not dissolve in HCl? That is interesting in its own right since that means it isn't Ca(OH)2 or calcium acetate.[/rquote]
I don't know, I'm allowing some of it to dry in the sun right now and will test its solubility when its dry. This white solid seems to be more soluble than the hypochlorite starting material. I'll find out of theres any Ca(OH)2 in there now with a pH strip.

[rquote]
Try heating the solid under a bit of water and see whether it will melt in the range of 80-100C without dissolving (forming a separate liquid layer). If so, then you have chlorobutanol and you know where your chloroform went. Chlorobutanol also has a distinctive medicinal sort of smell. Given the conditions of your reaction - excess of acetone added in two separate shots with five hours in between, together with magnetic stirring sufficient to create a whirlpool - I would actually be puzzled if all your chloroform *hadn't* reacted to form chlorobutanol.

I pretty much followed this guide here:
https://www.erowid.org/archive/rhodium/chemistry/chloroform....
So from what you've said, I don't think there is any chloroform in the solution I have. Whatever is in there must either be acetone or chlorobutanol then. I'm still wondering why there was no rise in temperature at all though. I'm going to test the reaction again on a much smaller scale, and see what happens.

mycotheologist - 23-4-2012 at 09:06

Quote: Originally posted by bbartlog

Alternatively you could try using 10g of your putative hypochlorite, using 100g of water and 3ml of acetone; dissolve as much of the solid as possible first (at room temperature), then add the acetone and see what happens.

I'm going to give this a try now.

Quote: Originally posted by bbartlog

I gather that your white solid does not dissolve in HCl? That is interesting in its own right since that means it isn't Ca(OH)2 or calcium acetate.

I don't know, I'm allowing some of it to dry in the sun right now and will test its solubility when its dry. This white solid seems to be more soluble than the hypochlorite starting material. I'll find out of theres any Ca(OH)2 in there now with a pH strip.

Quote: Originally posted by bbartlog

Try heating the solid under a bit of water and see whether it will melt in the range of 80-100C without dissolving (forming a separate liquid layer). If so, then you have chlorobutanol and you know where your chloroform went. Chlorobutanol also has a distinctive medicinal sort of smell. Given the conditions of your reaction - excess of acetone added in two separate shots with five hours in between, together with magnetic stirring sufficient to create a whirlpool - I would actually be puzzled if all your chloroform *hadn't* reacted to form chlorobutanol.

mycotheologist - 23-4-2012 at 09:47

Quote: Originally posted by bbartlog

I gather that your white solid does not dissolve in HCl? That is interesting in its own right since that means it isn't Ca(OH)2 or calcium acetate.

I added about a gram of the white solid to around 40mL of water and dissolved as much as I could (its actually relatively soluble). The pH was around 11. Is this evidence that I really do have hypochlorite, rather than TCCA?

I also tested out what you suggested. I dissolved 10g of the hypochlorite in 100mL of water at room temperature. It was still a bit foamy at the top and a bit murky so maybe I should have gave it more time to dissolve but anyhow, I added the 3mL of acetone all at once and stirred manually. In the space of about a minute, the temperature rose by 10C so there definitely is an exothermic reaction going on in there but nothing like I was expecting, from what I've read. I'm starting to suspect that my problem yesterday was insufficient amount of water to dissolve the hypochlorite. This smaller scale reaction actually is rising in temperature but then again, I don't have this one in an ice bath.

UPDATE: Its now at 30C which is 20 degrees higher than what it started at. Its definitely working this time. So I suppose I can safely assume I actually do have hypochlorite. I notice that a load of white precipitate has formed inside in the beaker. Is that calcium acetate? Anyhow, I'm glad I got the reaction working, thanks for the help. I didn't know chlorobutanol could be formed like this too so that knowledge may come in useful in the future. I learn way more from my little backyard experiments than I do in the labs at college. Then again, in college I get to use expensive instruments like IR spectrometers and HPLC machines.

EDIT: BTW I notice that Ca(ClO)2 has a solubility of 21g/100mL. Why did you recommend using 10g of hypochlorite for that test reaction? Would using saturated hypochlorite solution cause any problems?

[Edited on 23-4-2012 by mycotheologist]

garage chemist - 23-4-2012 at 10:08

The white precipitate could be calcium hydroxide, since the haloform reaction produces that as a byproduct as well.
Also, what sort of calcium hypochlorite product do you have?
There is bleaching powder with 34% active chlorine (not very common today anymore) and there is "high test hypochlorite" (HTH) with up to 70% active chlorine. Pure calcium hypochlorite would have 99% active chlorine, this is not an article of commerce.
Can't you just take a picture of the container of the product you are using, or provide a link to the MSDS?

mycotheologist - 23-4-2012 at 10:18

On the label it says "HTH Chlorine". Heres an MSDS for HTH granular chlorine:
http://www.pollardwater.com/pdf/MSDS_Sheets/HTH%20Granular%2...
According to this site:
http://www.hth.co.uk/wt_cal_hypochlorite.shtml
it has 68% available chlorine content.

[Edited on 23-4-2012 by mycotheologist]

garage chemist - 23-4-2012 at 10:38

The reaction should work very well with this.
Try stirring magnetically until all granules of calcium hypochlorite have dissolved into a uniform suspension and add acetone to this. Also, try using up to 20g HTH per 100ml water.

I don't think that chlorobutanol will form in any appreciable amount in this reaction. Its formation requires very strong bases like NaOH and nonaqueous conditions (liquid chloroform and acetone with powdered NaOH). Ca(OH)2 is a much weaker base.

To get out the chloroform, I would simply distill it out of the reaction mixture. There is a lot of insoluble stuff in calcium hypochlorite products and you won't get a good phase separation unless you acidify with HCl.

What kind of stirrer are you using? I recommend magnetic stirring throughout the reaction.

mycotheologist - 23-4-2012 at 11:45

garage chemist: Thanks a lot! You answered every question I had on my mind. I was beginning to get pessimistic about this reaction but from that info, I can see how to make it viable now. Yeah I will just distill because filtering was a lot of hassle and lots of product probably gets lost during the filtration process. I'm using a cylindrical magnetic stir rod.

I'm curious about what you said about the insoluble stuff in Ca(ClO)₂ products. When they say 67% chlorine availability do they mean the hypochlorite is 67% pure? Could you not filter out that insoluble material beforehand or would that be dangerous (i.e. some of the insoluble stuff may be there to stabilise the hypochlorite).

Magpie - 23-4-2012 at 11:49

I'm strongly suspecting that you don't have HTH calcium hypochlorite, ie, it's mislabeled. Either that or it is so old it has lost its chlorine via decomposition. With fresh HTH this reaction will proceed like gangbusters.

When you make chloroform you will know it. It has a very distinctive and characteristic smell.

bbartlog - 23-4-2012 at 12:03

Quote:

I don't think that chlorobutanol will form in any appreciable amount in this reaction.

Maybe not... but then how do we account for the missing chloroform (or hypochlorite)? Given that the white precipitate that myco had turned out *not* to be unreacted hypochlorite, I'm inclined to think that he did perform the haloform reaction. The lack of visible temperature increase could be accounted for by the cooling and a relatively slow reaction. I suppose all the hypochlorite could be in solution, at which point the question would be how he could have avoided the haloform reaction...

Quote:

Why did you recommend using 10g of hypochlorite for that test reaction?

Because actually dissolving something to the point of achieving a saturated solution is a pain in the ass. 50% saturation is normally pretty easy to achieve, so if it serves the purpose I aim for something in that ballpark instead. Also, 10g/100ml seemed like a good mark for a noticeable exotherm without the risk of making the container too hot to hold, boiling off acetone etc.

bbartlog - 23-4-2012 at 18:17

Quote: Originally posted by garage chemist

I don't think that chlorobutanol will form in any appreciable amount in this reaction. Its formation requires very strong bases like NaOH and nonaqueous conditions (liquid chloroform and acetone with powdered NaOH). Ca(OH)2 is a much weaker base.

Now I'm curious about this... I always thought of Ca(OH)2 as a strong base with solubility issues, not a weak base. And I had thought that the nonaqueous conditions used for the chlorobutanol synthesis were just to maximize yield.
Anyway, since I happen to have the necessary chemicals handy I decided to do the following test:
I put 13.3g of chloroform (110mmol) in a 250ml RBF, then added 50ml of water, 9.6g of acetone (165mmol), and finally 10g of slaked lime, Ca(OH)2 (135mmol). Temperature of everything was around 5C (ambient in my lab). I stoppered this and shook it for about half an hour, then left it to sit for three hours, shaking briefly every hour or so. Finally I neutralized the base with a slight excess of 31% HCl (using pH indicator rather than a calculated amount).
Anyway, it still smells slightly of chloroform, so it clearly has not been quantitatively destroyed. On the other hand, the layer of chloroform that was separate at the bottom when the reagents were first mixed is no longer there. Oddly, neutralizing the lime with HCl does not result in a clear solution - it remains turbid. Silica contamination maybe? It is agricultural lime so surely contaminated with other things, but I had assumed CaCO3/Mg(OH)2/MgCO3.
I'll see whether it's separated/settled tomorrow, and distill if necessary. Of course small losses of chloroform would not be indicative of anything much, but if I can't recover more than a couple of grams then I'd say it's pretty strong evidence that it can be destroyed by these conditions.

bbartlog - 25-4-2012 at 18:46

So, the attempt described in my previous post was an embarassing failure: I mistook a bag of diatomaceous earth for my slaked lime. /facepalm. After dumping the results, I tried again with a little more care. Can chloroform react significantly with acetone, given aqueous conditions and a Ca(OH)2 slurry?

To a 500ml Erlenmeyer flask, I added
12.0g chloroform (100mmol)
11.6g acetone (200mmol)
50g ice and 50g water
13g slaked lime (Ca(OH)2) (175mmol)

I added a stir bar and set it to stir at medium speed, leaving it for two hours. At the end of this time all the ice had melted and the temperature had risen to about 10C. I then left it to sit at room temperature (15C in my barn) for a couple of hours before returning to it.

I then began adding 38g of 31% HCl in order to neutralize the base, but ended up diluting the acid with 25g of additional water in order to reduce the exotherm. There was considerable fizzing, suggesting substantial CaCO3 contamination in my slaked lime (pretty much as expected, it is an agricultural product and not reagent grade).
Much of the chloroform was still visible at the bottom of the flask. I transferred the contents to a separatory funnel and separated the chloroform; it weighed 7.8g.
There is also a flocculent orange-brown precipitate in the aqueous phase. Condensation products of acetone? Hard to say.

Unfortunately this result is not particularly decisive. Some chloroform may have been destroyed, but it's not beyond the realm of possibility that the losses here were mostly mechanical or due to some dissolution of the chloroform in the aqueous phase. I will probably try this one more time, stirring for 24 hours and then using distillation to remove the remaining chloroform; while this is less relevant to the normal course of the haloform reaction it should at least give me a better idea of whether any reaction of the chloroform is taking place at all under these conditions.

barley81 - 25-4-2012 at 18:55

I'm not completely sure of the rate of this reaction, but the chloroform might have hydrolyzed into formate in the basic mixture. Correct me if I'm wrong.

mycotheologist - 26-4-2012 at 10:35

bbartlog: Good work. 7.8g of chloroform is a very high recovery rate, would that not be an indication that chloroform did not participate (as a reactant) in any of the reactions that occured during the 2 hours? I wonder what reaction caused that 10 degree rise in temperature.

EDIT: I'm trying this reaction again today, I tried to dissolve 100g of Ca(ClO)2 in 500mL of water but it wasn't enough water. My lab is outdoors so the temperature is around 10C. I added another 50mL or so of water then added the contents to a 1000 mL flask and am adding 30mL of acetone in very small increments. This time, I'm going to distill directly out of the reaction flask but whats bothering me is the tiny amount of chloroform I'll end up with. 30mL of acetone makes around 50mL of chloroform, but thats theoretical yield so I'll probably have less than that and have to distill a tiny volume out of this 1000 mL flask. I'm planning on using the chloroform as a solvent so ideally, I'd need to make at least 100mL at a time to make it worth the hassle. Is there a way I can make this more viable? For example, when the reaction has finished and all acetone is consumed, could I add additional reactants to flask?

One problem I see is all that insoluble Ca(OH)2. I could neutralise it and turn it into a more soluble salt. It won't take me very long to rearrange the glassware for simple distillation, whats time consuming is disassembling the apparatus and cleaning out the flask etc.

[Edited on 26-4-2012 by mycotheologist]

bbartlog - 26-4-2012 at 13:03

7.8g is less than 2/3 of what I put in, so I don't think it's *that* high a recovery rate; I agree however that it raises the suspicion that not much at all happened, hence my notion of trying again with a much longer reaction time (and better recovery via distillation) in order to try to get a more definitive answer.

As for the tiny volumes of chloroform for large volumes of reagent... yes. That's how it is. In order to obtain the chloroform for the experiment described above (the previous, failed one having used up the last of what I had stored), I put two liters of bleach in my largest Erlenmeyer flask and added 33g of acetone, yielding a comparatively very small blob of chloroform (49g) at the bottom of the flask (which was actually a very good yield). I generally decant the vast majority of the aqueous phase before putting the rest in a separatory funnel. Distilling in such cases could improve recovery, but it's just not worth it.

mycotheologist - 26-4-2012 at 13:55

Magpie must be right. What I have either isn't HTH hypochlorite, or its so old that its lost most of its chlorine content. I think its the latter because the supplier said he no longer sells Ca(ClO)2 products so god knows how long its been there for. I did the reaction again, this time using adequate solvent but once again, it was barely exothermic at all. This stuff is useless for the haloform reaction. I'm just going to use it as a Cl2 source. I know it works well for that, I added a few mLs of very dilute HCl to a bit of it, and it hissed and bubbled. I'll use it to make some NaClO solution, that way at least I'll its exact concentration.

mycotheologist - 26-4-2012 at 14:28

Quote: Originally posted by bbartlog

As for the tiny volumes of chloroform for large volumes of reagent... yes. That's how it is. In order to obtain the chloroform for the experiment described above (the previous, failed one having used up the last of what I had stored), I put two liters of bleach in my largest Erlenmeyer flask and added 33g of acetone, yielding a comparatively very small blob of chloroform (49g) at the bottom of the flask (which was actually a very good yield).

Was the bleach 5% NaClO solution? I'm gonna make my own NaClO solution and was wondering if higher concentrations such as 40% could be safely used in this reaction. I don't see why not. As long as one carefully controls the drip rate of acetone, the temperature can be controlled.

bbartlog - 26-4-2012 at 17:24

What I had was 6%. Higher concentrations can definitely be used, using your method (active cooling and gradual acetone addition). If you want to do it like I did, adding the acetone in one fell swoop, then 10-12% is probably max, and even for that you need to prechill the reagents to 0C.

mycotheologist - 27-4-2012 at 08:38

After distilling, I got about 10mL of a clear, colourless liquid. I got a mild spicy odour from it, not the odour I remember of chloroform but I added some water and it separated into 2 layers, the aqueous layer being the top one. Its definitely chloroform. Any left over acetone would have distilled over with it so since I couldn't smell any, I assume it was all consumed. The still head wasn't fitted properly and I noticed mild leakage so thats probably where most of the yield went. Either that or the CCl3 particles couldn't make it out of the enormous mass of Ca(OH)2 precipitate. Since there are such massive quantities of this crude Ca(OH)2 formed during these reactions, I'm gonna keep it and try and come up with a use for it. What I love about gas bubbling, is it usually doesn't matter how filthy and contaminated the reagents you use to generate the gas are. Its a brilliant way to recycle chemicals that would ordinarily be disposed of, due to their high degree of contamination.

Separation of acetone from chloroform

mycotheologist - 12-5-2012 at 10:37

What would be the easiest way to purify chloroform that is contaminated with a small amount of acetone? First thing that comes to mind is to use a series of liquid liquid extraction to gradually extract the acetone into the aqueous layer until there is a negligible amount of acetone left in the chloroform. A more effective approach though would be to add some kind of drying agent that is completely insoluble in chloroform but will absorb the acetone. Anyone know of a drying agent like this?

Pyro - 12-5-2012 at 10:58

best would be to just add some more bleach, and in the future use bleach in excess, thats a lot easier to seperate

sargent1015 - 12-5-2012 at 11:29

Well, I did some research on it (Woelen made some good points on this in another Chemistry forum) and lab extraction is NOT going to work. Acetone will dissolve in both water and chloroform, therefore making it incredibly difficult to separate the two entirely.

I did however find this from Wiki:

Quote:

Extractive distillation is similar to azeotropic distillation, except in this case the entrainer is less volatile than any of the azeotrope's constituents. For example, the azeotrope of 20% acetone with 80% chloroform can be broken by adding water and distilling the result. The water forms a separate layer in which the acetone preferentially dissolves. The result is that the distillate is richer in chloroform than the original azeotrope.

I'm not sure how reliable it is, but if you do try it and attain good results, be sure to let us know!

mycotheologist - 12-5-2012 at 12:45

Only problem is I'm not sure if there actually is an acetone in there. Its invisible and impossible to smell in such small quantities. I want to use chloroform as a solvent and acetone would alter its solubility properties. I suppose I can always just add hypochlorite when I'm done, then distill.

BTW a single liquid liquid extraction wouldn't get rid of much acetone, but a series of them would. The partition coefficient determines what percentage of the acetone will go into the water phase and what percentage goes into the chloroform phase. So lets say I have 100 mL of chloroform with 20 mL of acetone in it. Lets say the partition coefficient is 50:50. If I add 100 mL of water, then 10 mL of acetone goes into the water phase. I dump the water phase and add another 100 mL of water. Now 5 mL goes into the water, leaving only 5 mL in the chloroform. If I do this enough times, I can bring the amount of acetone in the chloroform down to microlitres. I think it works even better if I use more water. For example, if I add 500mL of water to the 100mL of acetone, I think even more of the chloroform will go into the water phase.

I just got a good idea for producing large volumes (i.e. 500mL at a time) of chloroform very easily. At first I was thinking about using one of those glass demijohns they use for brewing beer:

for making CCl3 from NaClO but decanting is a pain in the ass and surely tilting the vessel to decant causes some of the chloroform to mix with the water and thus, increases the amount lost. Heres a way better idea. Use one of the 10 litre HDPE water containers:

and make a small hole at the bottom and make it into a stockcock type device. That way you simply make a big batch of CCl3, then open the stopcock and slowly drain the chloroform out of the container. To make the end point easier to see, you could add food colouring to the aqueous layer (assuming it doesn't dissolve in chloroform). Well, theres nothing revolutionary here, I'm just talking about making a big, opaque seperatory funnel.

[Edited on 12-5-2012 by mycotheologist]

[Edited on 12-5-2012 by mycotheologist]

mycotheologist - 13-5-2012 at 09:21

I've been making chloroform using this 5% NaClO solution my neighbour gave me and I'm starting to get the hang of it. Whats weird though is the chloroform at the bottom of the flask (a 1L erlenmeyer) is white and goopy, it looks almost like vaseline. It smells like chloroform but why is it so viscous and white? Is it some kind of emulsion? I'm guessing I can just distill it and get pure chloroform (unless theres acetone in there) but I'm just wondering why the CCl3 looks like this.

watson.fawkes - 13-5-2012 at 10:01

Quote: Originally posted by mycotheologist

I just got a good idea for producing large volumes (i.e. 500mL at a time) of chloroform very easily.

You very much should learn about thermal runaway right now. Scaling up haloform reactions is notorious, even here, for being problematic.

mycotheologist - 22-5-2012 at 05:41

Quote: Originally posted by watson.fawkes

Quote: Originally posted by mycotheologist

I just got a good idea for producing large volumes (i.e. 500mL at a time) of chloroform very easily.

You very much should learn about thermal runaway right now. Scaling up haloform reactions is notorious, even here, for being problematic.

What I have in mind is increasing the quantity, not the rate. The temperature can kept low by adding the acetone slowly.

Help!!! I guarantee this will have you baffled!

NiCoLi_BrLiNTe - 19-7-2013 at 05:58

Right firstly, thanks in advance for any help or advice you can offer. Basically i'm having some serious issues pulling off a haloform reaction with both sodium hypochlorite and calcium hypochlorite. I first tried to do it using 4.5% household bleach. Domestos extended kill brand to be precise. I wasn't having much luck with this, I assumed because of all the thickening agents and surfactants. I then tried to make my own both through electrolysis and gassing NaOH with chlorine, but still nothing. finally I decided to switch to calcium hypochlorite, technical grade from a actual chemical supplier, still nothing, the reaction vessel didn't even rise by a single degree in temperature! Assuming it must have been poor quality, I decided to pay the excess and get some from one of the international chemical supply houses. As such I purchased 10 litres of sodium hypochlorite from VWR. This is graded as having less that 10 ppm of contaminating factors and is clearly extremely strong bleach. I also thought it may be something to do with my acetone (although its fairly easy to tell acetone) so bought some 99.9 percent assay lab grade acetone as well. I brought the temperature down to about -5C and mixed near stoichiometric amounts (with the NaClO in slight excess) in a dark, super clean vessel. Much to my absolute disgust still nothing. I've tried mixing them with out cooling to see I can tell if there reacting from the heat generation, and again not even a single degree rise.
This has made me start to question the fabric of reality, I'm going to end-up drinking the acetone and hypochlorite to see if I can regurgitate chloroform.
Also, I couldn't upload any pictures so I've attached them as files instead, just so you might be able to tell me if I'm doing somthing wrong, or if i've simply gone insane!
Thanks again,

NiK

[Edited on 19-7-2013 by NiCoLi_BrLiNTe]

VWR - analytical grade sodium hypochlorite.jpg - 19kB

Bot0nist - 19-7-2013 at 06:55

In my experiance, 10ml of cold acetone is added to 300ml cold (from the freezer, near 0°C) 8% commercial bleach with mag stirring in an icebath. After 5 minutes or so, the solution becomes turbid and cloudy, then from yellow to clear, and a blob of about 8ml of chloroform can be sep funneled off the bottom. Good luck, and yes, UTFSE for more data. Its been covered to death.

[Edited on 19-7-2013 by Bot0nist]

NiCoLi_BrLiNTe - 19-7-2013 at 07:10

Thanks, I've been using either 100mls of a saturated calcium hypochlorite solution, ice cold - to 40mls of pure ice cold acetone. Alternatively I've been using 1litre of 7% sodium hypochlorite to 33mls of acetone. How hot should the reaction get? Using these ratios on a test tube scale but without cooling, I not noticed any temperature increase or any other sign of a reaction for that matter.
Also I have been reading through the other threads on this topic for over a week now in quite some depth, as well as using other sources. I can't find any reason why it shouldn't work. Not only is there no mention of this type of problem in conventional textbooks, no one else seems to behaving quite the same problem. However, it didn't occur to me to simply add this post to an existing thread, so my apologies for that and I'll be sure to bare that in mind next time I go to post something!

Thanks,
Nik

Bot0nist - 19-7-2013 at 07:37

It is exothermic, and extensive cooling is needed. Pre cool your reagents and use an icebath, then report back. With just 10ml of acetone, the outside of the flask get noticably warm when the solution gets cloudy. Also be patient. It takes several minutes of stirring and cooling before the reaction takes off.

bbartlog - 19-7-2013 at 07:53

So 33 ml of acetone = 26g = about 450mmol.
1 liter of 7% bleach = (1.1kg * 0.7) = 77g sodium hypochlorite = about 1030mmol.

Your amounts here are not near stoichiometric and your sodium hypochlorite is not in excess (rather your acetone is). The correct molar ratio is 3:1 bleach to acetone. It matters for two reasons: first, excess acetone may result in some 1-chloroacetone being produced that then doesn't get completely chlorinated (thus consuming the limiting reagent, bleach, to no purpose). Second, if there is acetone remaining at the end, it can react with the chloroform under sufficiently basic conditions, further damaging yield.
However, this doesn't explain why you didn't see a rise in temperature, and it surprises me that you got no chloroform at all. It is true that if you start with ice cold reagents and mix them immediately on addition, it can take quite some time for the reaction to get going (tens of minutes), which is why people sometimes just leave it overnight.
I suggest you proceed as follows: take one liter of your 7% bleach (chilled to 0C) and gently pour 22ml of acetone (room temperature) on top of it (so that it sits in a layer atop the bleach). Swirl gently every minute or so, just to mix a little at the boundary.

(edited to correct relatively irrelevant error in stoicho calculation)

[Edited on 19-7-2013 by bbartlog]

Sublimatus - 19-7-2013 at 08:07

NiCoLi_BrLiNTe, at least to my mind, your results suggest to me that one of your materials is not what it's labelled as. This reaction is spontaneous, and quite exothermic. In my experience, I've always been more concerned with slowing this reaction down, rather than trying to get it started.

Have you done anything to confirm the identity of your materials? I know it seems a silly suggestion, but it shouldn't be too much work to achieve. If you are indeed mixing acetone and NaOCl solution (with NaOCl in excess), and no reaction takes place, then as the title of your previous thread said, I am baffled.

Fantasma4500 - 19-7-2013 at 08:31

Quote: Originally posted by mnick12

Solid sodium hypochlorite? Hmmm considering the instability of hypochlorites I doubt you have pure sodium hypochlorite.

actually, it is possible to make solid NaClO
im very certain there will be some way to seperate the NaCl and NaClO3 from the NaClO
the reason i know this is possible is because i have evaporated water off to get a powder, in which was reacted with HCl.. the reaction i got was not by NaClO3
also when i added water to it all again, i got the classic bleach smell
starting product was approx. 5% NaClO
also for KClO3 synth. i have needed to boil down the water and completely away several times, still at this point a clear bleach smell was noticed, several times dried by boiling the water off
it is possible for sure, the question would be how to purify it

NiCoLi_BrLiNTe - 19-7-2013 at 09:30

I'll try it again using the above mentioned ratio's. also I haven't had it an a stir-plate so maybe that will help. As far as testing the reagents, the acetone is definitely acetone, it smells like it, is highly flammable and tests positive as a ketone with Brady's reagent. The sodium hypochlorite smells like bleach, and bleaches litmus paper; additionally you would hope coming from a supplier like VWR it's what it says it is. the rest of the stuff i've tried could well be flour for all I know. However, I try again with the information you guys have already provided, then if it's still not working I'll see if I can go into Uni at some point and use an IR spec or something.
Anyway, I'm off to fill the freezer with chemicals! I'll post images of what I'm doing; maybe that will help identify if I'm making a mistake somewhere.

Thanks,

NiK

Magpie - 19-7-2013 at 11:01

I agree that this synthesis should not need kickstarting - just the opposite. On my first try I had a runaway. I used HTH calcium hypochlorite pool chlorinator and acetone, no NaOCl was needed. Are you following an established procedure? Surely there is one in the forum library books.

Edit: see Norris (p. 105) and Vogel

[Edited on 20-7-2013 by Magpie]

NiCoLi_BrLiNTe - 20-7-2013 at 02:08

Yeah, I've recently modified it according to the 1:3. Ratio surggested above! The sodium hypochlorite and calcium hypochlorite aren't used in the same reaction. After failing with calcium hypochlorite, I moved to the sodium! Although, last night after I posted on here, I tried again. This time instead I used the new ratio, and I hadn't had a chance to get the reagents as cold as I would have liked, about 0*C. Finally I used a mulithead adapter and set up for reflux, After the reaction had spent half an hour mixing an high speed I re-set up for simple distillation. From the first batch I got about 5mls. But then I put more calcium hypochlorite in to the hot flask and added the acetone. It started boiling all by its self. I think I must have got about 7ml from that run.
I've take pictures, so I'll up load them when I get on my PC next.

chloroform by the haloform reaction

NiCoLi_BrLiNTe - 21-7-2013 at 04:02

The following are pictures of the process I used to finally produce chloroform by the haloform reaction. I'm still not really sure what the problem was in the first place; although I think it may have been overkill on the cooling of reagents (i.e. using cryo-hol to try and keep everything at a constant -5⁰C or simply the incorrect ratios (although I'm uncertain how this would cause the reaction to stall). Whatever the cause, I think next time I'll skip over the reflux stage and just go strait to adding the reagents in a distillation setup with vigorous stirring of the reaction flask.
The dropping funnel was charged with 22 ml of acetone, the pictures are kind of self explanatory, but let me know if theres any part you're especially interested in, I have video of most of the distillation and will try to explain as best as I can. Finally I know adding the reagents strait into the distillation apparatus is probably risky, but I believe with a bit of planning and some precautions, it will be a worthwhile way to increase yield and save time.
Thank you all for your help and suggestions,

NiK

[Edited on 21-7-2013 by NiCoLi_BrLiNTe]

Trichloromethane from bleach and acetone.

Bot0nist - 24-7-2013 at 08:26

Trichloromethane
is prepared from household bleach and pure acetone nail polish remover.
The reaction that takes place is very exothermic, and this rise in heat
must be accounted for, especially when using larger amounts of reagents
and scaling up. I use a salt/ice bath and aggressive magnetic stirring.
Do this under a well ventilated hood or in a ventilated area, because
the acetone and chloroform are volatile and not especially good for you.
The rapid rise in heat during the reaction also makes it possible for
the volatiles to boil out of solution.

Reagents Used:

500ml 8.25% bleach (sodium hypochlorite in alkaline solution.)
12ml Acetone.
Lots of Ice and salt.

Here are some pictures of the process.

Extra strength bleach

Pure acetone

500ml of bleach chilled to ~ -12°C

12ml acetone, chilled.

Acetone added with a sep. funnel drop-wise. Followed bbarts advice on allowing a top layer of acetone to form. Mag stirring at full speed at this point.

Last of the acetone is added.

Cloudiness and exotherm started.

Solution rose in temp to ~ 3°C with aggressive cooling and slow addition.

Solution losses its yellow color and temperature rise stops.

[Edited on 24-7-2013 by Bot0nist]

Bot0nist - 24-7-2013 at 08:31

A blob of chloroform seen at the bottom of the flask after 1 hour from first acetone addition.

Sep funnel shots.

Chloroform separated from the funnel.

Yeild ~10ml crude chloroform.

I did not do a distillation of the obtained, crude trichloromethane, so it likely contains some dissolved acetone. boiling point measured at ~ 60.5 °C.

Thanks for watching. Any comments, corrections or flaming welcome.

[Edited on 24-7-2013 by Bot0nist]

Magpie - 24-7-2013 at 11:14

That looks good - a very simple procedure with easily attainable reagents. Lots of good pictures. 10mL chloroform on 12mL acetone seems like a very good yield. I would like to see what the yield is after purification, ie, distillation.

I made chloroform some time ago using the procedure in Brewster (1960) at 1/2 scale. This procedure calls for 100g of HTH calcium hypochlorite and 37mL of acetone. Equipment is a 1000 mL RBF with a reflux condenser. Workup was by steam distillation followed by simple distillation. I combined 3 each 1/2 batches and after the final distillation had just 25mL chloroform on 55.5mL of acetone.

[Edited on 25-7-2013 by Magpie]

Bot0nist - 27-8-2013 at 17:19

Sorry, I somehow missed your respone Magpie.

I thought the crude yield looked to good to be true. Unfourtunatly, my last 14/20 condensor broke months ago (hope Dr. B has some left, when I have the $...).

This has left me with only a 24/40 option. This would make the mechanical losses quite high I fear. I could add a few jugs of bleach to the ol' grocery list in a couple weeks and do several, similarly scaled and cooled reactions in sequence and collect the nonpolars together for a 24/40 run. Im worried that x5ing the scale would make for serious challenges with exotherm.

I vaguely remember reading something about up to 20% acetone impurity can be hard to remove. I need to ustfs real quick to find it.

I'll try and get a more precise yeild data posted soon.

Thank you for the interests. I actually find trichloromethane pretty useful as a solvent, when I have it handy at least.

[Edited on 28-8-2013 by Bot0nist]

Chloroform precipitate and invisible chloroform?

Fantasma4500 - 11-6-2014 at 03:25

i have longer time ago succesfully made a few mL of chloroform using store bought bleach and a small amount of acetone, it gave a large amount of white fluffy like precipitate, but below all of this there was a very heavy liquid which seemed to be chloroform

now again i have tried with some homemade bleach made from salt water electrolysis for roughly 12 hours, i let the thing run with 5V 40A power supply, on average this thing can put out 150g NaClO3 after 3 days runtime when i just let the solution cool to room temperature, but the bleach solution i ended up with im not sure about, but it smelt decently strong

approx 1.5L of bleach solution, homemade, cooled with icecubes had 20mL of acetone added to it and the lid was fitted onto it, it was given a shake and let stand
after a while a white cloudy layer formed at the top, and later on it settled down (3 hours?)
i have before taken this strange powder and dried out, to then test it for melting point.. it was above 200*C if it even has one
i have read about a precipitate with an existing melting point when this is done with Ca(ClO)2 and acetone but this is not the case when using NaClO

i saw a video where one guy used 400 mL bleach solution and 20 mL acetone, so i suppose 20 mL acetone is not too much when having 1.5L? ive thought about whether the acetone could dissolve the chloroform, where the acetone would be soluble in water and thus making the chloroform soluble aswell

i see small bubbles clinging to the side of the jar, and when i shake it the chloroform smell gets pretty intense

could it be my acetone had some crap added to prevent chloroform synthesis??

Zyklon-A - 11-6-2014 at 06:02

I haven't heard of adding things to acetone to make it unsuitable for simple chloroform synthesis. Have you used that acetone for anything else?
You might want to distill it just to be safe.
My guess to your problem is just because you stoichoimetry was to far off. You definitely should evaluate the concentration of your NaOCl (aq).
About the precipitate, have you tested it at all yet? Check it's solubility in water, whether it burns in air (indicating an organic compound), melting point, maybe even boiling point.

Fantasma4500 - 11-6-2014 at 14:57

yeah i have
used it for dissolving organics mostly, had alot of uses for it but its like i dont recall them, strange?
anyways the datasheet from the manufacturer says its pretty high quality, it leaves nothing behind when let evaporate off, its perfectly clear and burns very well

but yes the stoichiometry might be why, ill let my cell run during the night aswell to get a very high concentration, hopefully..
about the % of the acetone im fairly certain its at least 95%, but the NaClO's concentration i cant tell and i dont know how i could possibly accurately titrate it?? reacting it with more or less anything is like dumping several fuels oxidizers and catalysts together and burning it off.. if you have a suggestion for NaClO titration please share

actually i see its possible to do with thiosulfate, but i used the last i had long time ago in an desperate attempt on creating manganese sulfate

Zyklon-A - 12-6-2014 at 06:34

What about using hydrogen peroxide to test the conc.
the reaction is: H₂O₂ + NaOCl → NaCl + H₂O + O₂. So if you use in excess of peroxide, you can just measure the oxygen generated to get the conc. pretty accurate.

[Edited on 13-6-2014 by Zyklonb]

Fantasma4500 - 15-6-2014 at 03:21

1mL 1% i recall to give 3.33mL O2
ill need to first run a quick gas titration of my H2O2 first, estimated to be 12%
no problem, ill try it soon

Fantasma4500 - 21-6-2014 at 14:04

my 12% H2O2 seems to be... 28.5%?!
if 1 mL 1% H2O2 yields 3.33mL of gas
then 1 mL 12% H2O2 should at most generate 39.9 mL gas

but i got 95 mL

i used a 1 mL pipette to drop the H2O2 into a 250 mL flask, coupled up to a ground glass flask adapter to silicone hose, then to 100 mL graduated cylinder filled with water under water
it was as if i could just keep adding bleach, but after about ... 40 mL it stopped giving off gas

i decided to try and find out what % my H2O2 has simply using excess NaClO solution

anyhow, i seems to get the result that i need about 40 mL bleach solution for 1 mL H2O2

titration with MnO2 gives me 65 mL of gas per mL H2O2

[Edited on 21-6-2014 by Antiswat]

blogfast25 - 22-6-2014 at 05:11

Quote: Originally posted by Antiswat

my 12% H2O2 seems to be... 28.5%?!
if 1 mL 1% H2O2 yields 3.33mL of gas

[...]

titration with MnO2 gives me 65 mL of gas per mL H2O2

I don't know how you arrive at these numbers.

1 ml of 1 w% H2O2 is 0.01 g of pure H2O2.

Assume oxidation as:

H2O2 === > O2 + 2 H+ + 2 e-, so 1 mol of O2 per mol of H2O2, 24.5 L O2 per mol O2 at STP.

So 0.01 g H2O2 / 34 g/mol H2O2 x 24.5 L/mol = 0.720 L = 72 ml.

Not 3.3 ml. The 72 ml corresponds roughly with your MnO2 result.

Gazometer measurements can be tricky and imprecise: try iodometry instead; hydrogen peroxide oxidises iodide to iodine, then titrate iodine with thiosulphate and starch indicator.

[Edited on 22-6-2014 by blogfast25]

Metacelsus - 22-6-2014 at 08:37

Yesterday, I performed the haloform reaction with 600 mL 8.25% NaClO and 15 mL acetone. The reaction proceeded as expected, with an exotherm raising the temperature to 30 C despite external cooling. However, I have a cloudy emulsion of chloroform that refuses to separate (even after 24h in the funnel). How could I recover the chloroform?

Ideas I have:
1) Distillation - it is a heterogeneous mixture, so the steam distillation concept applies.
2) Adding a salt to break the emulsion - not sure if this would work.
3) Extracting with a high-boiling nonpolar solvent (xylene) and then distilling.

Metacelsus - 22-6-2014 at 16:49

I've decided to extract with 2x20ml xylene. The extraction went very slowly, but the layers did separate. I think there was some kind of surfactant in the bleach. I will soon distill the organic layer.

Fantasma4500 - 25-6-2014 at 11:20

but...
if 1 mL 1% H2O2 will yield 72 mL of oxygen by being reacted with MnO2
and i used 1 mL 12% H2O2, then it would yield 12 times as being as 72 mL -- being 864 mL of gas

also i got the numbers from another thread
i looked through it to see if anybody would point out that 3.3mL would be wrong, but nobody did so.. which is unusual for false calculations

but 72 mL oxygen per 1% 1mL H2O2 is just unreal

i have found some place that 12% H2O2 should expand 40 times as in oxygen per volume
meaning 1 mL should give 40 mL of oxygen --- 39.9 mL

http://puu.sh/9JPEN.png

a thing i have considered whether could have fucked up my gas titration was that NaClO + H2O2 seemingly generates mono-oxygen (which has some specific name i ofcourse forgot)
this could potentially account for why i got so much more oxygen than i supposed i would get

Metacelsus - 25-6-2014 at 16:10

The xylene distillation yielded 4.3 mL chloroform.

I tried the same reaction with a different brand of bleach. No emulsion formed, and the yield was 73% (50.5 mL chloroform (after distillation) from 63 mL acetone).

blogfast25 - 26-6-2014 at 04:56

Quote: Originally posted by Antiswat

but...
if 1 mL 1% H2O2 will yield 72 mL of oxygen by being reacted with MnO2
and i used 1 mL 12% H2O2, then it would yield 12 times as being as 72 mL -- being 864 mL of gas

also i got the numbers from another thread
i looked through it to see if anybody would point out that 3.3mL would be wrong, but nobody did so.. which is unusual for false calculations

but 72 mL oxygen per 1% 1mL H2O2 is just unreal

i have found some place that 12% H2O2 should expand 40 times as in oxygen per volume
meaning 1 mL should give 40 mL of oxygen --- 39.9 mL

http://puu.sh/9JPEN.png

a thing i have considered whether could have fucked up my gas titration was that NaClO + H2O2 seemingly generates mono-oxygen (which has some specific name i ofcourse forgot)
this could potentially account for why i got so much more oxygen than i supposed i would get

Stop getting these basic date from 'another thread' and do the basic calculations yourself.

Mono-oxygen? Atomic oxygen IMMEDIATELY combines to 'normal' diatomic oxygen.

"but 72 mL oxygen per 1% 1mL H2O2 is just unreal"

Why is it 'unreal'? And what does 'unreal' mean in your strange little world?

Nicodem - 26-6-2014 at 07:27

Quote: Originally posted by blogfast25

Assume oxidation as:

H2O2 === > O2 + 2 H+ + 2 e-, so 1 mol of O2 per mol of H2O2, 24.5 L O2 per mol O2 at STP.

So 0.01 g H2O2 / 34 g/mol H2O2 x 24.5 L/mol = 0.720 L = 72 ml.

The reaction equation is only valid for the H₂O₂ oxydation reactions (like with NaOCl), but not also for the disproportionations (like with MnO₂ or iodide catalysis) where the ratio is 2H₂O₂ -> O₂.

Besides, the result of your calculation is 12.25 mL (at least according to my calculator).

Fantasma4500 - 1-9-2014 at 08:15

true, the calculations i found was somehow incorrect, found 12% H2O2 to must yield 41.4 mL oxygen using MnO2 to catalyse it into free oxygen

recalling 40 mL homemade NaClO solution was used to fully react 1mL 12% H2O2 it would yield me somewhat about 0.65% NaClO solution which seems absurd
perhaps i should try electrolysing a NaCl solution well cooled down for longer time until i reach less than 10 mL NaClO solution per mL of H2O2 needed to react it fully

but still doesnt make up for white precipitate, in which seems not to occur when using store bought NaClO solution

Aquakidney - 1-9-2014 at 13:46

Hello,

Long-time lurker, first-time poster. I was doing this same prep a few months ago and found this procedure online for determining OCl^- concentrations. Does require thiosulfate, but getting some would be a lot less trouble than trying to collect and measure a gas.

Incidentally, my particular lot of Clorox Concentrated Bleach actually contained 8.62% NaOCl, not the advertised 8.25%, so that was nice.

The next project is 2,4-dinitrophenylhydrazine to see if there's any acetone left in the CHCl₃.

Attachment: hypochlorite lab.pdf (57kB)
This file has been downloaded 2015 times

Fantasma4500 - 2-9-2014 at 07:21

inverted measuring cylinder filled with water in water, its not that bad, but you will probably get wet hands somehow
supposing the too high concentration may be because they expect it to decompose, in which it ofcourse does over time

anyhow, happy first post and welcome to sciencemadness

franklyn - 9-9-2014 at 23:34

Attachment: Was Chloroform produced before 1831.pdf (89kB)
This file has been downloaded 1157 times

On somewhat larger scale

bbartlog - 13-9-2014 at 16:39

Did this preparation on a larger scale, since I want to do some experiments using chloroform as a solvent. Yield suffered somewhat.

Experimental

Four containers of 8.25% sodium hypochlorite bleach (Chlorox), each containing 3.78 L of liquid, were chilled for four hours in a freezer, at which point some ice began forming. Estimated average temperature was slightly below freezing.
The contents were poured in to a 23 L glass carboy.

Empties

342g of acetone at room temperature were measured in to a 500ml beaker and then poured in to the carboy. The acetone remained on top of the bleach, forming a separate layer that gradually turned yellow. The carboy was repeatedly tipped and rotated in order to agitate the reaction mixture and mix the layers. After a few minutes, the upper part of the mixture grew warm and very small droplets of chloroform began to fall through the solution and accumulate at the bottom of the carboy.

The yellow layer

After repeated swirling (which should be done with caution if you are using a large glass carboy, as breakage can lead to deep lacerations), the top layer became mixed with the remainder of the solution. The carboy was left to stand overnight so as to allow the reaction to finish, and any suspended droplets of chloroform to settle to the bottom. Subsequently the top aqueous solution was decanted to the extent possible (without loss of the bottom layer), the remaining 750 ml or so poured via a funnel in to another narrow-mouth glass container, some further top layer removed by decantation and the remainder poured in to a 500ml separatory funnel:

Separation

Following separation the chloroform was weighed, yielding 401g of crude product (57% yield based on acetone).
The chloroform was then stored for several days in a screw top container, which (due to inadequate closure and evaporation) resulted in the loss of 15g of chloroform on re-weighing. Subsequently, the chloroform was stored overnight with an equal weight of 95% sulfuric acid, after which the sulfuric acid and chloroform were poured in to a 1L round flask with flat bottom and set up for distillation (yes, I know that's not supposed to be a Hempel column in the middle there...):

Discolored acid

Misuse of Hempel column

The sulfuric acid darkened noticeably as a result of this (it had been a very light amber color). Though it should not react with chloroform to an appreciable extent, I assume that impurities in the chloroform formed condensation products that led to this discoloration.
Distillation yielded approximately 235ml of chloroform, a yield of about 50% (I haven't weighed it yet).

Nearly full

This is a slightly disappointing yield. Previous runs on a smaller scale have yielded up to 75% or so. I would assume that the difficulties in mixing the acetone with the bleach when using vessels of this size may result in more extended contact between chloroform and acetone, resulting in more losses to the chlorobutanol side reaction. Adding the acetone in portions or else using some kind of mechanical stirring might mitigate this.