| Pages:
1
2
3 |
mycotheologist
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
How to kick start the haloform reaction
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
Hazard to Others
Posts: 404
Registered: 30-12-2009
Location: In the lab w/ Dr. Evil
Member Is Offline
Mood: devious
|
|
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
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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]
|
|
|
Nicodem
|
Thread Moved
22-4-2012 at 04:14 |
mycotheologist
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
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.
The less you bet, the more you lose when you win.
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 3227
Registered: 13-7-2003
Location: Wisconsin
Member Is Offline
Mood: Rock n' Roll
|
|
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
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
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.
The less you bet, the more you lose when you win.
|
|
|
Magpie
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
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.
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
mycotheologist
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
| 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.
The less you bet, the more you lose when you win.
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
| 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.
The less you bet, the more you lose when you win.
|
|
|
mycotheologist
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
| 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.
|
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.
[Edited on 23-4-2012 by mycotheologist]
|
|
|
mycotheologist
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
| 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
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
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
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
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
Hazard to Others
Posts: 154
Registered: 16-3-2012
Member Is Offline
Mood: No Mood
|
|
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
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
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.
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
| 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.
The less you bet, the more you lose when you win.
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
| 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.
The less you bet, the more you lose when you win.
|
|
|
| Pages:
1
2
3 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
