Sciencemadness Discussion Board - Chloroform disposal

Chloroform disposal

McLovin382 - 28-12-2008 at 18:11

I have a small amount of undistilled chloroform that I'd like to get rid of. I don't want to put something like that down the drain or into nature or anything but itd also be a pain to dispose 'properly' of such a small amount.

Is there a way to destroy chloroform (and other organic waste a chemist is likely to have to deal with) yourself?

UnintentionalChaos - 28-12-2008 at 18:22

Let it sit for a long time with concentrated lye solution in a tightly capped polyethylene bottle. The end product is a mix of formate and chloride, which can both be safely flushed.

granitestaterecovery - 28-12-2008 at 19:11

chloroform can be contained in a handkerchief and disposed with the intended abducted.and if there is money in there pocket this pays for itself. or donate it to the science lab anonymously.

Honestly you can probably. ask some one in chemistry dept of school

McLovin382 - 28-12-2008 at 21:41

lol granite. Well I'm not in school atm but I guess I could ask someone there anyways. They might look at me funny :p

What if I put it in NaOH solution like UnintentionalChaos said, and add some hydrogenation catalyst like magnesium shavings or something...would that potentially catalyze the decomposition? (Don't suppose I have any true reducing agents that could supply a steady stream of H2 in there)

[Edited on 28-12-2008 by McLovin382]

woelen - 29-12-2008 at 02:07

How much chloroform do want to get rid of? If there only is a few 100's of ml, then take an old towel or a big heap of paper tissue, put this outside and pour the chloroform on the towel and let it evaporate. The towel can be thrown away with normal waste. Do this on a windy and dry day! Chloroform disperses quickly in the atmosphere and soon is hydrolysed to harmless chemicals (CO2 and chlorides). In this way I get rid of small amounts of organic volatile chemicals, usually with quantities in the 1 ..5 ml range.

UnintentionalChaos - 29-12-2008 at 02:10

It's not a hydrogenation that is occuring. It's a hydrolysis. The net effect is to replace all three Cl with OH, which dehydrates to formate. The only way I can think to help it along is a phase transfer catalyst if there is some chloroform that isn't dissolved.

saintmichael - 29-12-2008 at 08:55

You might want to be careful with UnintentionalChaos' method, since I've heard of Chloroform waste bottles detonating in the presence of a base especially if there's any acetone with it.

BromicAcid - 29-12-2008 at 10:29

The danger is if chloroform comes into contact with solid sodium or potassium hydroxide. That is when the exothermic even can become unmanageable. Also, vigorous stirring will help since it will be biphasic.

Aluminium/nickel alloys in the presence of base can chew up chlorinated organics.

I believe fentons reagent (H2O2 with catalytic Fe3+

can also take care of chlorinated organics.

Thanks for being environmentally concious, most people forget that chloroform is a priority pollutant and a D-listed chemical.

vulture - 30-12-2008 at 02:24

Quote:

If there only is a few 100's of ml, then take an old towel or a big heap of paper tissue, put this outside and pour the chloroform on the towel and let it evaporate.

Then you might as well pour it down the drain! Letting it evaporate releases it into the environment too.

McLovin382 - 30-12-2008 at 07:25

It's not an amount that would necessarily harm the environment...just a bit of the stuff made in a small bottle via a basic reaction between NaOCl and EtOH.

Thanks for responding guys

woelen - 30-12-2008 at 11:02

Quote:

Originally posted by vulture

Quote:

If there only is a few 100's of ml, then take an old towel or a big heap of paper tissue, put this outside and pour the chloroform on the towel and let it evaporate.

Then you might as well pour it down the drain! Letting it evaporate releases it into the environment too.

For me there is a difference between pouring it down the drain and letting it evaporate. If it is poured down the drain, then it is released in high concentration and it may harm aqueous life (it also damages your tubing if this is made of PVC). If it is released into the atmosphere then it becomes so dilute that no animal will be harmed.

I can freely say this, because the number of people who release chloroform in the atmosphere is VERY low. If this were common practice for every household, then it would be no good, but if this is done by a few hobbyists, who do this once in a lifetime or maybe a few times in their life, then I see no real harm. It has to be put in perspective.

DJF90 - 30-12-2008 at 12:35

Chloroform is involved in the depletion of the ozone layer due to photochemical reactions that occur in the atmosphere. 1-2 mls is "acceptable", "a few 100's of ml" is not, especially if there is a method of destroying the chloroform rather than releasing it into the environment.

Diablo - 26-6-2012 at 09:36

It could be further oxidized to chlorobutanol.

AndersHoveland - 29-6-2012 at 23:59

Chloroform can be refluxed (heating in a special flask to cycle the boiled vapors back) with aqueous sodium hydroxide solution for 2 hours to form sodium formate and sodium chloride.

AJKOER - 3-7-2012 at 13:12

To your CHCl3 add dilute H2O2 and Na2CO3 (or just add OxyClean and a little water), shake and leave in an open container outside in sunlight.

Source for my suggested recipe: http://www.postech.ac.kr/see/art/journal/es&t31.pdf

To quote from the left hand bottom of page 93: "the CHCl3 dechlorination rates (Figure 3a) increase rapidly at pH > 10 in O2 saturated suspensions". The author also states that photo-oxidation in the presence of dissolved oxygen is a major pathway for haloform degradation. The hydrolysis reaction per the author's equation (16) is:

CHCl3 + H2O --> CO (g) + 3 HCl (aq)

which, I would argue, is moved to the right in the presence of light and Na2CO3. In the presence of light and H2O2, my take on the reaction:

CHCl3 + H2O2 --> CO2 (g) + 3 HCl (aq)

based on the author's equation (11).

So, if proved effective, my recipe would be simple and relatively cheap (and most likely mother approved) with the final product salt water.

[Edited on 4-7-2012 by AJKOER]

kristofvagyok - 3-7-2012 at 14:25

It could be easily disposed by let it evaporating, like as woelen said. It is not that toxic and not as harmful for the ozone layer as you think.

Also another short info: night butterflies are collected by a lamp. These lamps are placed somewhere out, far from cities in complete darkness. The butterflies come to the light and they die because a circa 100cm3 chloroform ampoulle is placed in the center of the "lamp". When the sun raises someone goes there, collects the "loot" and places a new chloroform ampoulle. With this they could identify how is the butterfly population.

So several liters of chloroform is evaporating every night....

quantumcorespacealchemyst - 7-12-2014 at 04:52

link is dead, this is archive https://web.archive.org/web/20040327063908/http://www.postec...
attached is pdf

i read that sodium cinnamate does something about dissolving things between layers, and that it is photo activated, i have found little data about it

although i havent found an inexpensive source of cinnamaldehyde, i am wondering about it for the production of sodium cinnimate in haloform with naocl.
i am wondering if this is a good decompostion catalyst of chloroform when alone and/or with NaOH solution

also if dry NaOH and CHCl3 react vigorously, what about passing the vapor of chloroform over a bed of NaOH in a reaction vessel? for gas products for other reactions

Attachment: es&t31.pdf (321kB)
This file has been downloaded 557 times

Amos - 7-12-2014 at 06:53

The NaOH method is tried and true, safe, and easy to conduct, and it really doesn't take very long. It's so fast in fact that the chloroform produced in the haloform reaction is constantly being eliminated, and if you wait too long to remove it, you'll end up with a severely damaged yield.

There's no reason to try anything else, IMO. By the time you can no longer smell any chloroform after shaking the container nor see any phase boundary within the mixture, it's probably a safe bet that the NaOH has done the job.

[Edited on 12-7-2014 by No Tears Only Dreams Now]

unionised - 7-12-2014 at 07:09

I suspect that, even if no deliberate attempt was made to dispose of it, the stuff will have evaporated in the intervening 6 years.

Amos - 7-12-2014 at 07:41

If the thread was revisited 6 years after the OP, who's to say that it won't still be visited by people with the same question?

S.C. Wack - 7-12-2014 at 13:17

Quote: Originally posted by No Tears Only Dreams Now

The NaOH method is tried and true, safe

Not necessarily if CO is a side product. Not if CHCl3 and and NaOH are put in a sealed container.

Why would anyone use solid NaOH? Surely ethanolic would be more effective? Surely it would work in water? Still, it is alleged that CO is a product, especially at temperatures between room and chloroform's bp, reflux being most favorable for formate.

Quote: Originally posted by AJKOER

To your CHCl3 add dilute H2O2 and Na2CO3 (or just add OxyClean and a little water), shake and leave in an open container outside in sunlight.

Source for my suggested recipe: http://www.postech.ac.kr/see/art/journal/es&t31.pdf

Quote: Originally posted by quantumcorespacealchemyst

link is dead, this is archive https://web.archive.org/web/20040327063908/http://www.postec...
attached is pdf

I predict that this will work without the addition of any chemical to the chloroform whatsoever.

Chloroform has close enough to zero ODP and GWP.

[Edited on 7-12-2014 by S.C. Wack]

Amos - 7-12-2014 at 15:27

So what you gleaned from my statement, in which I clearly said that there would be no reason to seal the container, is that the method might not work because the container is sealed.

And I might've assumed that people trying to react away highly volatile carcinogens maybe would do so OUTSIDE, where CO wouldn't be too much of an issue. But that's just me.

And not that I'm in anyway saying you're wrong, S.C. Wack, but this is the first I've heard of carbon monoxide being generated from the elimination of chloroform using a hydroxide.

S.C. Wack - 7-12-2014 at 17:13

Quote: Originally posted by No Tears Only Dreams Now

So what you gleaned from my statement, in which I clearly said that there would be no reason to seal the container

It should be clear that your quote was used by me to comment on safety. As usual I did not feel like quoting from every post in this thread to respond to each snippet. See the other posts for relevance to my statements.

You are not expected to know that the usual equation giving formate need not apply (this is from over 100 years ago (Geuther, Thiele and Dent, Self)), or that I have a tendency to reply to/be a Yossarian.

AJKOER - 14-12-2014 at 09:05

OK, the reason that even small 100 ml amounts of CHCl3 is a problem is because, per my recollection of the atmospheric chemistry, owes to its acting in the role of a catalyst for ozone destruction.

As a result, it only stops destroying O3 when the catalyst effectively burns out. Again, I remember a figure of something like this occuring after 50,000 molecules of ozone are consumed for each molecule of CHCl3.

I may be able to dig up sources.

S.C. Wack - 14-12-2014 at 10:35

Quote: Originally posted by AJKOER

I may be able to dig up sources.

Or not.

How do you explain chloroform being missing from Montreal and ODP lists? Stupidity? Ignorance? Bribes from CHCl3 users and manufacturers? Being naturally occurring, not photoreactive, and with a short atmospheric lifetime and general inability to reach the ozone layer?

Oxidation by ozone would actually be a great thing, since we'd be talking about ground level ozone aka pollution. Maybe China should be stocking up for next time they hold some international prestige event so they don't have to shut everything off?

morganbw - 14-12-2014 at 10:46

Quote: Originally posted by S.C. Wack

Quote: Originally posted by AJKOER

I may be able to dig up sources.

I think he is going by memory and is probably confusing it with methyl chloroform. Good chance anyway.

S.C. Wack - 14-12-2014 at 12:49

Quote: Originally posted by morganbw

I think he is going by memory and is probably confusing it with methyl chloroform. Good chance anyway.

Yes, I'm just being a dick with pent-up frustration at all the threads and posts from many of our um...newer...and/or more prolific members...

The majority of the chloroform in the atmosphere is natural and oceanic in origin, which is a good enough excuse for me to...well actually I'd never dispose of chloroform...on purpose.
http://www.tellusb.net/index.php/tellusb/article/view/14614

[Edited on 14-12-2014 by S.C. Wack]

AJKOER - 15-12-2014 at 04:03

One reference (see http://scorecard.goodguide.com/chemical-profiles/def/odp.htm...):

"Ozone depleting substances (ODS), including chlorofluorocarbons (CFCs), halons, and several other chemicals, are responsible for thinning the stratospheric ozone layer. When these substances reach the stratosphere, UV radiation from the sun breaks them apart to release chlorine or bromine atoms which react with ozone, starting chemical cycles of ozone destruction that deplete the ozone layer. One chlorine atom can break apart more than 100,000 ozone molecules, while a bromine atom can destroy about 4,000,000 ozone molecules"

Then this reference on page 13.4 (https://www.google.com/url?sa=t&source=web&rct=j&... ):

"As such, the total lifetime must take into account all of the processes determining the removal of a gas from the atmosphere, including photochemical losses within the troposphere and strato sphere (typically due to photodissociation or reaction with OH), heterogeneous removal processes, and permanent removal following uptake by the land or ocean."

Followed by the authors comment on page 13.5:

"From this total atmospheric lifetime, together with the evaluated loss lifetimes of CH3CC13 due to the ocean (about 85 years, with an uncertainty range from 50 years to infinity; see Butler et al., 1991) and stratospheric processes (40 ± 10 years), a tropospheric lifetime for reaction with OH of 6.6 years can be inferred (±25% ). The lifetimes of other key gases destroyed by OH (i.e., CH4, HCFCs, and hydrofluorocarbons [HFCs]) can then be inferred relative to that of methyl chloroform (see, e.g., Prather and Spivakovsky, 1990) "

And conveniently, per Table 13.1 on page 13.6, CH3CCl3 and CHCl3 are listed consecutively, with the relative hazard of Chloroform specifically, only one tenth of methyl chloroform. So apparently CH3CCl3 is more readily prone (possibly owing, in part, to its relative molecular size increasing the likelihood of a collision, I would guess) under atmospheric conditions to attack by the hydroxyl radical primilary (from the action of intense uv radiation on water vapor in the upper atomsphere), as was suggested by S.C. Wack reference above, than CHCl3 followed by photodissociation to the Cl radical, a catalytic agent in ozone destruction.

[Edit] Here is some background comments per Wikipedia http://en.m.wikipedia.org/wiki/Ozone_depletion:

"The Cl and Br atoms can then destroy ozone molecules through a variety of catalytic cycles. In the simplest example of such a cycle,[4] a chlorine atom reacts with an ozone molecule, taking an oxygen atom with it (forming ClO) and leaving a normal oxygen molecule. The chlorine monoxide (i.e., the ClO) can react with a second molecule of ozone (i.e., O3) to yield another chlorine atom and two molecules of oxygen. The chemical shorthand for these gas-phase reactions is:
Cl· + O3 → ClO + O2: The chlorine atom changes an ozone molecule to ordinary oxygen
ClO + O3 → Cl· + 2 O2: The ClO from the previous reaction destroys a second ozone molecule and recreates the original chlorine atom, which can repeat the first reaction and continue to destroy ozone."

And finally, my comment of CHCl3 ceasing to destroying O3 occuring after some 50,000 molecules of ozone are consumed should be most likely be revised down to say 10,000 (assuming the above cited 100,000 being applicable to CH3CCl3 as a source of the Cl radical, and the noted 1/10 relativity for CHCl3 specifically) or fewer, which some may still find troubling.

[Edited on 15-12-2014 by AJKOER]