Sciencemadness Discussion Board

CHCl3 and CCl4 Prep Overview

Sauron - 3-4-2008 at 09:16

Given an ample supply of materials for a reliable smooth chlorine generator, next we turn to what to do with it.

Chloroform and carbon tetrachloride are often needed, but often hard to get - essentially impossible where I am.

Conventional wisdom is, make CHCl3 via the haloform reaction, and make CCl4 by chlorination of CS2.

However, the haloform reaction of hypochlorite and acetone is IMO a rather laborious affair. A large volume of liquid reaction mixture has to be manipulated to obtain a miniscule yield of chloroform.

Similarly, carbon disulfide is flammable (very), rather costly, and its chlorination produces a large amount of stinking sulfur chloride byproducts that are unpleasant to handle, and difficult to separate from the CCl4. So unless one has a use for SCl2/S2Cl2, they are a nuisance.

I'd like to look at the alternatives.

CHCl3 can be made from CCl4 so let's focus on CCl4.

The main alternatives to chlorination of CS2 are twofold:

-- chlorination of methyl thiocyanate

-- preparation of trichloroacetyl chloride.

Methyl thiocyanate is rather expensive to buy but easy to make. Essentially one needs sodium or potassium thiocyanate and a methyl halide, 95% ethanol being the solvent.

A reaction I regard as overlooked and underappreciated is the second. H.C."Boron" Brown noted in a JACS article that benzoyl chloride is a general reagent for preparation of volatile acid chlorides. But in the cases of dichloro and trichloroacetic acids, the usual very high yields dropped sharply. The acyl chlorides of these acids fall apart readily to CCl4 and a little CHCl3.

I suggest that these acids are inexpensive enough (especially compared to CS2 and MeSCN) to buy them and go straight to Brown's procedure.

If more economy is desired, and an embarassment of Cl2 is at hand, then acetic acid (GAA) is dirt cheap and readily chlorinated beyond the monochloroacetic acid stage.

The fire hazards associated with CS2 are absent.

Another entry point to these acids is through chloral, from chlorination of anhydrous ethanol.

It is non-intuitive that ethanol and chlorine can get you to CCl4, or that acetic acid and chlorine can as well, but the chemistry is valid and the economics compelling.

The methyl thiocyanate route is the most interesting because it can be tweaked in a variety of directions, but the acid chloride route is apparently the most attractive for lab use and from cost effectiveness considerations.

The byproduct, dichloroacetyl chloride or trichloroacetyl chloride, are useful. Or they can be hydrolyzed back to the acids and reacted again. By this tactic, the per cent of conversion to CCl4 can be very high overall.

CCl4 by this method is entirely free from S-containing contaminants which is not the case when starting from CS2 or MeSCN.

I'd like to know if anyone else wants to explore this concept as well.

Klute - 3-4-2008 at 09:43

Do you have a ref for the methyl thiocyanate chlorination? I suppose anhydrous conditions have to be used for CCl4, as I have read several articles that state that chlorination of alkyl thiocyanates in aqueous mixtures yields sulfonyl chlorides and CNCl...

CH3-S-CN --Cl2/H20--> CH3-SO2-CN --Cl2--> CH3-SO2Cl + ClCN

Could the carbon tetrachloride come from further chlorination of CNCl?
Or perhaps in absence of water the C-S bond is cleaved on the methyl side, leaving a chlorinated -SCN moiety, before or after replacing the hydrogens with chlorines...
I would be interested in reading any litterature you could have on hand on that subject.

Sauron - 3-4-2008 at 09:59

Here's the paper by James in J.Chem.Soc., on which my remarks are based.

One of the ways the reaction can be manipulated is to produce cyanuric chloride. Since that is the solid trimer of CNCl you are right. But anyway read on.

James isolates a variety of intermediate chlorination products but the ultimate product of continuede chlorination is mainly CCl4.

I would think that EU residents who can't readily obtain cyanuric chloride (TCT, CC) would be even more intrigued by the possibility of making this by this method, than making CCl4, which anyway can be done by other means.

Whereas the only other way to make TCT is from CNCl, no fun, or from chlorination of CA (cyanuric acid) with PCl5, also no joy.

3 CH3SCN + 11 Cl2 -> C3N3Cl3 + 2 CSCl4 + CSCl2 + 9 HCl

C3N3Cl3 is cyanuric acid.

CSCl4 is trichloromethyl sulfenyl chloride and this on further chlorination gives CCl4 and SCl2.

CSCl2 is thiophosgene and this on further chlorination also gives CCl4 and SCl2.

So, 3 mols methyl thiocyanate ultimately form a mol of TCT and 3 mols of CCl4 alog with 3 mols SCl2 and a pile of HCl.

I am suggesting that this is a very profitable series of reactions. TCT is very useful, so is CCl4. Thiophosgene is also useful, and is a pain to make. See the low yielding and rather cumbersome Sn/HCl reduction of CSCl4 as exemplified in the Thiophosgene monograph in Org.Syn.

SCl2 can be oxidized to SO2Cl2. S2Cl2 can be used in prep of acetic anhydride.

3 mols MeSCN is c.220 g.

From this we get about 70% of a mol of TCT (c.100 g); theoretically 3 mols CCl4 (c.450 g) and 3 mols SCl2 (306 g) and all that gained mass is due to the cheap and plentiful Cl2 from TCCA and HCl. 11 mols Cl2 are consumed in the chlorination of MeSCN (3 mols) and a further 8 mols in the chlorination of the thiophosgene and trichloromethyl sulfenyl chloride.

So this takes about 6 mols TCCA, that is roughly 1.5 Kg or about $6 US worth out of my 50 Kg drum.

Not bad.

[Edited on 4-4-2008 by Sauron]

Attachment: JamesJCS.pdf (193kB)
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MagicJigPipe - 3-4-2008 at 10:54

"So, 3 mols methyl thiocyanate ultimately form a mol of TCT and 3 mols of CCl4 alog with 3 mols SCl2 and a pile of HCl."

This is amazing! Never before have I seen such a variety of useful reagents produced from a single reaction. I have been looking for feasible routes to TCT and CCl4 for the longest time.

It's unfortunate, however, that methyl thiocyanate seems to be unavailable to me. This is just a preliminary observation, though.

Thanks for the paper, Sauron!

Klute - 3-4-2008 at 11:25

Very interesting! Especially it seems very easy to seperate the TCT out of the reaction mixture.. But it would require large amounts of chlorine, and complete chlorination to CCl4 and SCl2 even more so!
I guess treating the distillate with water after complete chlorination and seperation of TCT would suffice to obtain SCl2-free CCl4, after fractionnation.

"The Action of Chlorine on Thiocyanates"
Treat B. Johnson and Irwin B. Douglass
JACS, (61)9, 2548 - 2550 (1939)

available at the ref thread thanks to Solo, middle of page 21, details the preparation of mesyl chloride from methyl thiocyanate. I think there is mention of the article you provided.

microcosmicus - 3-4-2008 at 11:35

Quote:

It's unfortunate, however, that methyl thiocyanate seems to be unavailable to me.


Not a problem. As Sauron mentioned above, it can be made from an alkali metal
thiocyanate and a methyl halide in ethanol. As I mentioned the other day in another
thread, the alkali metal thiocyanate can be made from sulphur, alkali metal
carbonate, and a ferrocyanide (or other source of cyanide). Looking further
back in the forum archives, methyl chloride can be made from methanol and
hydrochloric acid.

http://www.sciencemadness.org/talk/viewthread.php?tid=1509&a...

Hence, with a bit more work, one can instead start with
some quite common chemicals:

CH3OH
CH3CH2OH (as solvent)
S8
Na2CO3
Na4Fe(CN)6
HCl

[Edited on 3-4-2008 by microcosmicus]

woelen - 3-4-2008 at 12:58

Sounds interesting, but how do you get your methyl halide. Thiocyanates are no problem for me, they are easy to obtain and very cheap (e.g. at photography raw chemical suppliers) and are not watched or suspicious in any way. Methyl chloride is another issue. I am not well educated in organics, how could this be made easily?

[Edited on 3-4-08 by woelen]

MagicJigPipe - 3-4-2008 at 13:47

I would much rather use methyl bromide (BP 3.5C) or methyl iodide (liquid at room temperature) than methyl chloride which is a semi-low BP gas. Unless it is easily soluble in EtOH and could be bubbled through it. Is it?

I should be able to obtain an alkali thiocyanate so that's not a problem. Bromomethane and iodomethane (I would probably go with MeBr because HBr is easier to make with more readily available materials than HI).

On wiki it says that one must use I2 and red phosphorus to make MeI from MeOH. I've noticed this method used on various iodo compounds. Is it just easier that way or is there a reason why HI can't be used?

I still can't believe I have to fear making something like HI though. I almost forgot about that crap. HBr + MeOH it is.

microcosmicus - 3-4-2008 at 13:48

Quote:

Methyl chloride is another issue. I am not well educated in organics, how could this be made easily?


What about the methyl chloride thread cited above? Doesn't that answer the question?

[Edited on 3-4-2008 by microcosmicus]

Sauron - 3-4-2008 at 13:52

The easiest way I know to prepare methyl chloride (gas) is to react TCT and methanol, liquifying the MeCL with a dry ice-acetone dewar condenser and storing it in a lecture bottle till needed.

That is of course rather circular if you are going to use the MeSCN to make TCT, but if your main target is the CCl4, and you can obtain TCT, it's the way to go.

Failing that, methyl iodide. Or methyl bromide.

Acros sells MeSCN for about $50/100 ml. So 3 mols is about$100. What price CCl4? It's banned here for "environmental" reasons, while CHCl3 is banned because of supposed drug-cookery. So buying MeSCN is not entirely out of the question economically.

I think making CCl4 from trichloroacetic acid is cheaper, that acid is cheap, also rather easy to make. Dichloroacetic acid also works, and can be made by chlorination of acetic acid or by starting with trichloroacetaldehyde (chloral), itself easily made by chlorinating anhydrous ethanol. The prep. of dichloroacetic acid from chloral is in Org.Syn.

I have not yet crunched the numbers but, I think these are the cheapest routes to CCl4, particularly if you are willing to start with GAA - another photographic chemical, I used to run a custom B&W darkroom -- and as I mentioned there are no smelly toxic sulfur compounds involved. Only corrosive chloroacetic acids and a little benzoyl chloride.

MJP, put not your faith in Wiki. There are other ways to make MeI than MeOH and red P and I2. Return to the light, read Org.Syn. instead. It has that procedure, but it also has one IIRC using KI and dimethyl sulfate. USE A HOOD! but the same is true for all the methyl halides. Brain tumors are nature's way of telling you to lie down and die.

[Edited on 4-4-2008 by Sauron]

LSD25 - 3-4-2008 at 18:02

Might be interesting reading:

http://www.pubmedcentral.nih.gov/picrender.fcgi?artid=117288...

The reaction of CNBr + methionine gives methyl thiocyanate, so the question then becomes, is this restricted to CNX?

Like for instance what would happen if TCCA was used to decarboxylate cysteine to the correspondent nitrile? You'd thereotically end up with an S and a CN attached to one carbon (same as you'd get from the dehydration of taurine - although with taurine you'd have a sulfonic acid group free at the end to bind to the CN of another degraded taurine unit). The question is, would this bond, and where would it cleave (C-S or C-N)? Someone should have a good reason why this wouldn't work...

len1 - 3-4-2008 at 18:17

Quote:
Only corrosive chloroacetic acids and a little benzoyl chloride.


?? My understanding is that its at least stoichiometric. Using an expensive reactive precursor to make an inexpensive un-reactive non-precursor - albeit hard to purchase at times, although an interesting reaction is uninteresting economically. CCl4 was discovered by chlorinating CHCl3, so that reaction although old-hat and boring is at least on economically. Of course the more chlorinated the carbon, the harder it is to chlorinate further - but I dont think the yields for this process have been published anywhere.

[Edited on 4-4-2008 by len1]

Sauron - 3-4-2008 at 19:47

Where I am, I cannot buy CHCl3. And I cannot buy CCl4. But I need both. Who says CCl4 is unreactive? It is plenty reactive when properly motivated. Same with CHCl3.

Sure, I can get DCM, and maybe chlorinate that with some UV kicker to obtain the others. I believe this is done industrially all the way from methane and in the vapor phase all the way.

It is however, brute force and devoid of elegance.

Look at the three reactions I am contemplating:

CS2 -> CSCl4 + SCl2 Separate these, reflux w/Fe, and you have CCl4. The SCl2 is by product.

MeSCN -> TCT, separate (filter) Chlorinate the liquid mixture -> CCl4. Or fractionate it and you have CSCl2, and CSCl4

You can set the thiophosgene aside and purify it, and you can turn the trichloromethyl sulfenyl chloride into more thiophosgene if you want, or you can chlorinate it to CCl4. An embarassment of choices.

Acetic acid - > ClCH2COOH ->Cl2CHCOOH - > Cl3CCOOH (all one operation) then use Brown's method (PhCOCl) and you get roughly equal amounts of CCl4 and Cl3CCOCl.

Or chlorinate EtOH to Cl3CCHO and convert that to Cl2CHCOOH and enter the previous series in midpoint. Hint: chlorinating GAA requires UV, chlorinating ethanol does not.)

All of these get the job done. All pay bonuses in terms of side products. All are interesting. (Some would say, too interesting, if you are not eqipped to contain the "perchloromethyl mercaptan" and the thiophosgene.)

Looks good to me. I wouldn't mind chlorinating DCM, photolytically. But it has all the sex appeal of watching wet paint dry on a damp day.

LSD25's suggestion of cyanogen bromide and methionine on the other hand is definitely too interesting, particularly on a macro scale.



[Edited on 4-4-2008 by Sauron]

len1 - 3-4-2008 at 20:18

Thiophosgene is about as toxic as phosgene - though has much lower vapour pressure, and with the lack of warning and latent phase, I dont think we should touch it at an amateur setting - unless looking for a painful way to commit suicide. The method is elegant I grant that.

PhCOCl method is relatively safe but is expensive unless you want to make the PhCOCl as its schedule II.

CS2 is method is cheap but you have the thiophosgene problem there again I believe.

garage chemist - 3-4-2008 at 20:59

Very interesting, I didn't know that there is a method to make TCT that doesn't involve handling pure HCN or uses cyanuric acid and chlorinating agents.

DMS is indeed somewhat available to us, thanks to PainKilla's synthesis, and thiocyanates can be made from ferrocyanide.

I wonder why you still get both thiophosgene and perchloromethyl mercaptan from the MeSCN chlorination. If one would exhaustively chlorinate MeSCN, shouldn't the thiophosgene all be converted to to CCl3SCl, seeing that the CCl3SCl probably arises from chlorination of the thiophosgene in the reaction mix?
And isn't CCl3SCl eventually completely converted to CCl4 upon chlorination in the presence of anhydrous FeCl3, which is the way industry chlorinates CS2 when CCl4 is the sole desired product?

It seems like if CCl4 is the desired product besides TCT, one should chlorinate pure MeSCN first, separate the TCT by crystallization and then add FeCl3 or iron filings and continue chlorination until you have a mix of CCl4 and SCl2.
This way thiophosgene and the mercaptan are avoided as products.

len1, thiophosgene has very strong warning properties, making it a whole lot less dangerous than phosgene.
It's just like ketene, this one is immediately irritating as well.

[Edited on 4-4-2008 by garage chemist]

Sauron - 3-4-2008 at 21:14

Ketene, being a carcinogen, is a lot scarier.

If you exhaustively chlorinate MeSCN, you isolate no CSCl2 and no CSCl4 either. Fe or FeCl2 is not necessary to convert the latter to CCl4.

I suppose that the exhaustive chlorination can be carried out without removing the TCT, but, it would not be a big problem to use a transfer line to move the liquid mixture to another flask leaving TCT behind, without ever exposing the liquid to the lab environment. This in fact has to be repeated till no more TCT crystals are deposited.

woelen - 3-4-2008 at 22:39

Quote:
Originally posted by microcosmicus
What about the methyl chloride thread cited above? Doesn't that answer the question?
Sorry about not being sufficiently clear. I read the thread, but the methods presented do not seem like easy methods. The method with CH3OH and HCl seems the easiest one, but there is the effect of back-reaction and I have bad experience with that kind of reactions (crappy yield, messy, hard to isolate). Probably my bad experience tells more about me and my organic chemistry skills, than about the reaction though :P

Would it be better to add CH3OH to concentrated sulphuric acid and adding some NaCl? Then there is dry HCl. Or does this result in formation of (extremely toxic) methyl sulfate?

len1 - 4-4-2008 at 00:01

Quote:
Thiophosgene is NOT produced in the chlorination of CS2.


Quote:
CSCl2 is prepared in a two-step process from carbon disulfide. In the first step, carbon disulfide is chlorinated to give trichloromethanesulfenyl chloride, CCl3SCl:

CS2 + 3 Cl2 → CCl3SCl + S2Cl2
Reduction of trichloromethanesulfenyl chloride produces thiophosgene:

CCl3SCl + M → CSCl2 + MCl2
Typically, tin is used for the reducing agent M.


There is likely to be suffiecient reduction if this process is carried out in metal or with impurites such as are always present.

Yes CSCl2 smells as do most C-S compounds I did not mean to extend that aspect to it, but the effects are delayed as per COCl2, in that sense its just as insiduous. The warning is at much lower concentration, but what good is that if one decent whiff is enough to put you away?

These chlorinations are equilibrium reactions and never complete - there are always products at different stages of chlorination present - so I do not believe that just because no Cl2 is being absorbed theres no more thiophosgene.

[Edited on 4-4-2008 by len1]

len1 - 4-4-2008 at 03:17

Sauron - HCN was a dismal failure in WWI - the French filled great quantities of gas shell with it and started its use at the Somme in 1916, however it produced few casualties and the French wasted 4000 tons - it has far too high volatility.

So can we use this to recommend 'one of the most virulent poisons known' to the amateur?

On the other side of the scale diphenylcyanarsine was the most toxic chemical of WWI - yet it was ineffective - the particle size wasnt right.

I take your point about Cl2 - but Cl2 is used by housewides to disinfect floors - it was used in WWI as a demonstration of principle - because it was on hand, what followed were much more potent gases and the fact that CSCl2 was one of them should make one think.

I know you are just theorising about these things, but there are people out there who take these ideas on board and actually do this as practical stuff. That is why I post these 'additions' to posts. Further on that CCl4 is carcinogenic

[Edited on 4-4-2008 by len1]

trilobite - 4-4-2008 at 03:58

Reaction of dichloro- and trichloroacetic acid with benzoyl chloride, giving chloroform and carbon tetrachloride you say? I bet it goes like this:

Cl3CCOOH + PhCOCl --> Cl3CCOCl + PhCOOH
Cl3CCOCl --> CCl4 + CO

Don't take chances with the carbon monoxide.

Jor - 4-4-2008 at 04:16

That's interesting Sauron. Is HCl the second? I work and breath it almost everyday. How so? Can you give me the list or something og chemicals with most fatalities?

trilobite - 4-4-2008 at 06:25

I didn't bring up the subject of carbon monoxide because of lack of trust in your equipment or your confidence Sauron. I think it is important to actually mention possible toxic byproducts of a reaction so that one can be prepared. I checked the article by Herbert C. Brown, J. Am. Chem. Soc. 60(6) 1325-1328 (1938), where my suspicion is verified.

[Edited on 4/4/2008 by trilobite]

LSD25 - 4-4-2008 at 13:23

Sauron,

The predominent reason cited that CW weapons were not used more extensively and/or operationally, in the PTO was due to volatility and the susceptibility of existing protective equipment to compromise through humidity/sweat transfer. The increased lethality of such agents (predominantly Mustard & Lewisite) also came as a shock to the volunteer's who participated in the trials (predominantly UK, US & Aust, conducted in Nth Queensland) as did the inneffectiveness of the protective equipment supplied for that purpose.

The overview of a book, 'Keen as Mustard' appears second in this link:

http://www.awm.gov.au/journal/j32/bookreviews.htm

Here is another look at the events:

http://www.opcw.org/synthesis/html/s7/p20.html

These show that playing with chemical weapons and the agents used therein, in tropical climates (which I'd imagine would include Thailand?) can have extremely different results and outcomes than would be expected in cooler climates.

PS Nitrates caused more deaths than any other chemical (although the actual cause of death probably could have been more directly attributable to Fe or Pb) during WWI. Germany did not use chemical weapons again in WWII as the prevailing winds were counterproductive (as they found in WWI). The only true chemical weapon used in quantity in WWII was WP (actually misnamed - called WP because the smoke is white, the actual phosphorus used in shells is RP, WP is far too difficult to store safely). Of side interest, the Geneva conventions prohibits the use of proximity (time) fuses on WP rounds, this despite the fact that the fuckers fit on the pointy end of the shell!!!:D This is because the shells are made to the standard pattern design and use the same case as the HE shells, on which the proximity fuse is designed to fit (interchangeable with the standard instant (HE-Q)/delay (HE-D) fuses normally found on shells straight out of the box).

As an addition (post-edit) is it possible to get this article:

http://www3.interscience.wiley.com/cgi-bin/summary/117873144...

It deals with, inter alia, thiocyanuric acid - as a derivative of cyanuric acid - I'd be interested to see how this occurs.

[Edited on 4-4-2008 by LSD25]

len1 - 4-4-2008 at 16:45

My thread is not about benzyl chloride preparation but benzaldehyde and benzoic acid - benzyl chloride got made due to the irresponsible way in which some preparative manuals are written - as Ive outlined there. In fact I warn people about this compound and provide a procedure by which it is never isolated. I even have a section on health effects there - and Im proud of it - I think every practical prep dealing with dangerous substances should have one.

Having said that, I dont know why you like to pick problems. It was never my intention that you or others shouldnt follow your proposals (if you actually decide to do it). I think they are very interesting reactions, and you are one of the few who contribute to this forum constructively - I wish you continue that. My first two posts were just pointing the dangers - adding to rather than subtracting from what you had to say. After reading that, everyone is warned, and welcome to continue - same as on my BzCHO post. I myself would never isolate CSCl2, HCN, COCl2, exploding carcinogens, but thats just my preference. I have sufficient respect for others, and I wish my comments be taken in that light.

[Edited on 5-4-2008 by len1]

len1 - 4-4-2008 at 18:49

If you dont want people responding to your pageful of unprovoked accusations at them then I suggest you mark your posts as such, I also suggest you dont use expletives when reasoning runs out. As a defence of that other guy, its you who introduced CWC to the thread, and wrote more about it than all other combined. Have a good day

[Edited on 5-4-2008 by len1]

Civility? Do you speak it?

chemoleo - 4-4-2008 at 20:12

Sauron's last post just above was deleted as it amounted to nothing but an utterly pointless flame war Blitzkrieg, totally unnecessary and out of proportion.

I'm not even going to apologise for taking such a step.

This is out of order and has no place in this forum.

Please let's get back ON TOPIC.

A shame such measures are necessary given the quality of the posts here, including those of the originating author.

I recommend taking the 'E' out of the 'EGo'!

Thank you.

[Edited on 5-4-2008 by chemoleo]

Sauron - 4-4-2008 at 20:52

I have deleted as many of my posts as possible in this thread because a forum moderated in so biased a manner does not deserve good chemistry.

I'm sure chemoleo will delete this as well.

I will communicate developments in this area (CCl4/CHCl3 preps) to such members as are capable of carrying them out safely.

The hysterical overreaction of some disgusts me.

[Edited on 5-4-2008 by Sauron]

chemoleo - 4-4-2008 at 21:21

Good chemistry hardly goes amiss here, Sauron, and as you know no-one will debate that.

Picking on minute issues such as the detailed toxicity and war gas capability of SCCl2 is not (welcome), particularly if it erupts in animosity EVERY time when someone happens to contradict YOU ever so slightly.

It's a pattern I observe every single time.

Bias is not the issue here, as I wholeheartedly appreciate your *scientific* efforts here, as well as your knowledge.

But if your knowledge and effort comes at the expense of online fighting/insulting/cursing/judging/confrontation every single time someone dares to disagree or contradict, I feel we'd rather be deprived of that knowledge and live in something resembling peace.

As I said in the U2U (which I'm sure you aren't going to read), there are ways of dealing with disagreements with any members that don't automatically result in arguments, name-calling etc. Please try, you'll see yourself how truly friendly this forum is once you get the hang of it ....

Sauron - 4-4-2008 at 21:30

Your absurd distortion of my statements and demeanor simply underscores your inability to be impartial.

THAT is bias, personal bias in regard to me, and unbecoming of a moderator, a position to which you should never have been appointed.

LSD25 - 4-4-2008 at 21:37

Sauron,

If you suggest that CWC type agents are of some use to you in a tropical climate - I'm fucked if I see any reason why I shouldn't point out the problems with that approach. I am on shaky ground here, 'cos of the fact that Sauron is involved and I'm just gunna get slapped down, so I'll shut the fuck up NOW:D

Sauron - 4-4-2008 at 21:44

What CWC type agents?

Thiophosgene is not on CWC and I said so. It is a commercially available reagent used for making thiocyanates.

If you want to characterize it as a CWC type agent for your own reasons, then I am here to say you are wrong, and question your agenda.

Tropical climate is pertinent to what? The inside of my fume hood is not tropical.

Thiophosgene preparation is described in Org.Syn., q.v. Their method is clumsy and low yield. Their apparatus is too large to fit in a 2 meter long, 1 meter high, 1 meter deep hood.

I have a nice review on the useful synthetic chemistry of thiophosgene from Synthesis, which I will post here in the possibly futile hope that this compound be seen to be useful for something other than flinging at one's enemies. It is pretty useless for that.

In short you are either misinformed or mendacious, which is it?

This thread was about alt routes to CHCl3 and CCl4.

Thiophosgene came up only as an intermediate in one of the three reactions, and need not have been isolated. By continued chlorination, both intermediates were converted quantitatively to CCl4. The only other product was TCT and that is a sought after material for many members. The sulfur chlorides produced are mostly a nuisance but can be converted to more useful reagents. This is the same dilemna faced by every industrial maker of CCl4: what to do with all the SCl2/S2Cl2?

If someone wishes to stop at the intermediate stage and fractionate off the thiophosgene, before continuing chlorination of the other intermediate (CSCl4)

[Edited on 5-4-2008 by Sauron]

LSD25 - 4-4-2008 at 22:15

What, who me? What? Never...

Next thing you'll be suggesting I'm sarcastic:P

Sauron - 4-4-2008 at 22:42

Ah. Mendacity then, it is.

Jokes belong in Whimsy.

However, as promised,

"Thiophosgene in Organic Synthesis"

from SYNTHESIS.

[Edited on 5-4-2008 by Sauron]

Attachment: s-1978-24896.pdf (1.3MB)
This file has been downloaded 4057 times


Sauron - 5-4-2008 at 11:40

To resume the topic of this thread:

Conventional wisdom is that chloroform is best prepared by the haloform reaction, and CCl4 by chlorination of CS2,

The haloform reaction requires very large scale apparatus if the yield is to be more than 50-100 ml. The reaction can be hazardous and is always laborious.

The chlorination of CS2 has its own drawbacks. CS2 is rather costly, and very very flammable. Chlorination produces a toxic intermediate, CSCl4, and a nasty byproduct, SCl2 admixed with S2Cl2. The difficulties in separating and purifying CCl4 from these are pretty well known and discussed in the chemical and technological literature.

So I went looking for alternatives.

One such is chlorination of methyl thiocyanate. Like CS2, MeSCN is rather expensive. But unlike CS2, it is amenable to synthesis. See downthread.

The first stage of MeSCN chlorination produces cyanuric chloride as acrystalline precipitate. This is separated easily by filtration. The liquid mixture remaining contains the following components

CSCl2 thiophosgene
CSCl4 trichloromethyl sulfenyl chloride (same as in CS2 chlorination)
CCl4
SCl2 (same as in CS2 chlorination)

Continued chlorination of this mixture gives CCl4 and SCl2. It is possible if desired to seperate out the thiophosgene but not mandatory.

So to summarize, MeSCN can be economically synthesized as below, and gives CC (TCT) on chlorination as well as thiophosgene and products identical to those from CS2 chlorination. Ultimately the products can be as few as three (TCT, CCl4 and SCl2).

This is more interesting than CS2 chlorination but still contains the problem of separating CCl4 and SCl2, so let's move on to the other alternative.

Chlorination of trichloroacetic acid to trichloroacetyl chloride was reported by H.C.Brown to result in a significant amount of decomposition of the acid chloride to CCl4 and some CHCl3.

To a lkesser extent this also occurs when preparing the acyl chloride of dichloroacetic acid by Brown's method, benzotl chloride.

Glacial acetic acid is an inexpensive starting material and its chlorination through chloroacetic acid is well known. It is also available commercially at rather low cost.

I reasoned, it is also possible to chlorinate ethanol to chloral (trichloroacetaldehyde, which unlike acetic acid is a process that does not require UV. Chloral can be converted to dichloroacetic acid, see Org.Syn. The trichloroacetyl chloride process can be entered at that point advantageously.

Things turn out to be even simpler than that.

A number of researchers studied the chlorination of chloral and chloral hydrate under UV and found that trichloroacetyl chloride is formed and at ordinary temperatures decomposes by a radical process to CCl4 and gaseous products (CO, HCl) quantitatively.

At -50 C it was possible to recover 55% trichloroacetyl chloride.

At -10 C only 7.5%

At 10 C to 90 C, no acid chloride was recovered and only CCl4 remained.

This represents the conversion in two stages of ethanol (anhydrous) to CCl4 quantitatively. In first stage, to chloral by chlorination first at 0 C and ultimately at 90-100 C. The procedure is described in a Chem.Rev. article as well as in a number of patents. Anhydrous ethanol is cheap and ubiquitous. Yields are high 70-85%).

The second stage requires UV photolysis. Unless you rely on sunlight, photolytic reactions require special equipment and impose safety hazards (eye protection), and this is the downside of this method. However, this is compensated for by the highest economy of any of the methods and by the absence of noxious sulfur-containing intermediates. hose who are unwilling, unprepared or unequipped to deal with CSCl2, CSCl4 and sulfur chlorides should employ this method. The CCl4 product will be devoid of ill smelling impurities.

[Edited on 6-4-2008 by Sauron]

Sauron - 5-4-2008 at 12:18

Apologies for double post.

The preparation of methyl thiocyanate can be conducted by two methods.

Per Merck Index, but without references, MeSCN is prepared from barium thiocyanate and dimethyl sulfate according to the equation

Me2SO4 + Ba(SCN)2 => BaSO4 + 2 MeSCN

We are left to presume that solvent choices are, per Ullmann's, ethanol or DMF.

As garage chemist pointed out, PainKilla has furnished a method of preparing DMS for those who cannot purchase it. BE WARNED! Dimethyl sulfate is a proven human carcinogen, specifically a brain carcinogen. I personally do not relish its use and would not prepare it.

Barium thiocyanate is prepared per a procedure in Inorganic Syntheses vol 3, p 24, from ammonium thiocyanate and barium hydroxide. This procedure looks a little tedious and the commercially available barium thiocyanate is horribly expensive, almost $900/Kg from Alfa.

Ammonium thiocyanate is also a little expensive, but can be itself prepared from less dear starting materials.

The other methos is more advantageous.

An alkyl halide such as MeBr or MeI is reacted per Vogel, in 95% ethanol with potassium or sodium thiocyanate.

Methyl iodide (iodomethane) like DMS is a hazardous methylating agent, but less dangerous than DMS. It is not too expensive commercially, and can also be prepared, per Org.Syn., from DMS and potassium iodide.

I have on hand dimethyl sulfate, potassium iodide, and potassium thiocyanate. Therefore I can make my own MeI and from that, MeSCN. To me this appears better than the longer and more laborious NH4SCn to Ba(SCN)2 to MeSCN method.

But others may have different priorities.

The productivity of chlorination of MeSCN is extraordinary and due to the large amount of chlorine absorbed.

Overall equation is

3 CH3SCN + 15 Cl2 -> C3N3Cl3 + 3 CCl4 + 3 SCl2 + 9 HCl

The yield of TCT is 70% according to James (J.Chem.Soc.)

Sauron - 5-4-2008 at 13:00

Once again mea culpa for double post. In my defense I am restoring in a more concide form the content of my posts in this thread which I gutted earlier today. Hopefully the logic of my remarks will be more apparent than previously.

The chlorination of dry ethanol does not require photolysis and is not improved by it. This is quite unlike the chlorination of glacial acetic acid.

Chlorine initially oxidizes ethanol to acetaldehyde and monochlorinates that. This phase of the reaction needs to be chilled hard () C). As chlorination progresses the temperature can be raised and at the end is in the 90-100 C range. Yields are typically 70-85%.

CH3CH2OH + 3 Cl2 -> Cl3C-CHO

Thus 46 g ethanol will typically produce more than 100 g chloral.

The further and photolytic chlorination of chloral is quantitative.

C2HOCl3 + Cl2 -> Cl3CCOCl + HCl

And the trichloroacetyl chloride falls apart via the radical

Cl3CCOCl -> Cl3CCO. + Cl. => CCl4 + CO

So the overall conversion is Cl3COCl -> CCl4 + CO + HCl

I should point out that Cl2 in presence of CO is likely to combine to phosgene. Indeed, one of three investigators of this reaction as cited in the Chem.Rev. article on chloral's chemistry, observed phosgene among the products. Appropriate precautions need to be taken against CO and COCl2 as well as Cl2 of course.

I suppose I should mention nin passing that trichloroacetaldehyde (chloral) and its hydrate and alcoholates are controlled substances in USA and probably other countries. Chloral hydrate as an intoxicant really ceased to be a problem more than a century ago despite its ease of preparation, and lingered only as "mickey finn" knockout drops. So if you reside in a jurisdiction of concern, and wish to proceed, I'd advise processing the chloral into CCl4 immediately rather than storing it. Chloral is not the most stable of compounds anyway.

Remember, trichloroacetic acid, upon conversion to its chloride, in absence of photolysis still fell apart in the same manner, doubtless by same radical mechanism. The expected yield of trichloroacetyl chloride was 85-90%, but only 56% was obtained. (Brown, JACS.) Therefore it is reasonable to assume that trichloroacetic acid, upon photolytic chlorination, or chlorination in presence of a radical initiator, would give CCl4 in same fashion. This is probably practical, but would never be as economical as starting with ethanol except in terms of time saved.



[Edited on 6-4-2008 by Sauron]

Nicodem - 5-4-2008 at 13:12

Thanks for reviewing this topic. Being one of those lucky bastards that can just take a CCl4 bottle of the shelve, I don't have any interest in preparing it, but these reactions are new to me and quite interesting.

Anyway, if you are into discovering new things I have an idea you could try. Theoretically there is chance it might even work, though I would not bet on it.
Trichloroisocyanuric acid (TCCA) can be used in the Hell-Volhard-Zelinsky reaction (the relevant paper is uploaded at least in two posts: here and there). This reaction requires the enolization of the carboxylic carbonyl group which is done by having the acid equilibrate with a small amount of acid chloride (or bromide) which enolize without troubles. The small amount of acid chloride required is usually prepared in situ with any method of choice (red phosphorous, POCl3, another acid chloride, etc.).
So to keep it short, perchlorination of acetic acid with TCCA in the presence of an acid chloride generating reagent (for example cyanuric chloride, TCT), and possibly some catalytic amount of AlCl3 to speed up the decomposition of Cl3CCOCl, might give off CCl4 and CO as products distilling from the reaction mixture.
The drawback, if it works at all, would be the possible runaway due to highly exothermic reactions and the other is in the in the necessity of an inert and nonvolatile solvent since one would need at least 1 mol TCCA : 1 mol AcOH : 1/3 mol TCT : ~0.05 mol AlCl3 (an obviously impossible slurry to deal with, and furthermore one that can just become a small volcano). As for an appropriate solvent, there are only few choices and most are chlorinated alkanes which is a bit unreasonable for a reaction that is supposed to be used for the preparation of CCl4. But if anyone motivated enough wants to try it solventless on a small scale and report what happened…

Sauron - 5-4-2008 at 13:52

That is indeed interesting, thanks, Nicodem.

As it happens I had occasion to study the HVZ reaction a while back in connection with the alpha bromination of caproic acid, on the way to preparing racemic Norleucine for resolution via an enzyme to L-Nle.

The classical method is just Br2 with a little kicker from red P or in my case PBr3 (which I can get).

Clearly the P or P halide generates as you said, a little acid halide and that reacts, this process repeats and so on.

A slightly more au current variation is to formally prepare the acid chloride, and then a-brominate that with NBS (in CCl4!) to the a-bromoacid chloride, then liberate the acid.

My question is, does the HVZ method work beyond the mono-a-halo stage? Because I have not heard of any instance of such. Perhaps my interest in too narrow, as I was only looking for preparative routes to a-amino acids.

But I have no doubt at all that TCCA (a chloramide after all) would be a suitable alternative to N-bromosucconimide (a chloroimide).

Beware of trying to utilize all three chlorines on CC (TCT) because the final chlorine has a very different pKa and is far less labile. This usually requires higher temperature.

Therefore, I would advise that more than 1/3 equivalent TCT be used.

Better would be to preform the acid chloride by whatever means. Start with acetyl chloride. Then see if TCCA in the right solvent (whatever that might be) will go beyond the chloroacetyl chloride state. If it does, then indeed there ought to be stage set for radical decomposition to CCl4 and CO.

I would hesitate to try to go for broke with a one-pot gambit, and son't really see why AlCl3 is required.

Another reagent besides TCCA and the N-halosuccinimides that would be a candidate is yhe dihalogenated 5,5-dimethylhydantoin family. Like TCCA, cheap and ubiquitous.

Can you point me at the TCCA/HVZ article? The FSE and I often fail to have a meeting of the minds.

Again, thanks for the idea. By coincidence - I have to prepare chloroacetyl chloride sometime in not too distant future. This is for cyclization with 1,2-ethanediol on the way to BEDT-TTF, the tetrathiafulvalene. Chloroacetyl chloride is commercially available but very hard to ship, has to go by ocean. Only one company has a license to import it, and they don't want to bother, Apparently it has to be refrigerated in transit. The MOD here is paranoid about it, I had a hard time figuring why, but it turns out it is because they are afraid of chloroacetophenone! Oh well.

not_important - 5-4-2008 at 14:35

Quote:
Originally posted by Sauron
...

My question is, does the HVZ method work beyond the mono-a-halo stage? Because I have not heard of any instance of such. Perhaps my interest in too narrow, as I was only looking for preparative routes to a-amino acids.

...


Actually it can be difficult to stop at the mono-substituted stage in many cases. Chlorine is more likely to go further than bromine, because of crowding, but then acetic acid will have the least crowding issues.


Ah, March Advanced Org Chem 2-5

Quote:
When there are two alpha-hydrogens, one or both may be replaced, although it is often hard to stop with just one.



[Edited on 5-4-2008 by not_important]

Nicodem - 5-4-2008 at 14:56

The two posts where the Synthetic communication paper using TCCA for the alpha-monochlorination of carboxylic acids are linked in my above post.

Di- and trichlorination of acetic acid by the Hell-Volhard-Zelinsky reaction is possible, at least when using Cl2. Using TCCA would require acidic catalysis or prolonged heating though.

The proposition of adding some ~5mol% AlCl3 is for catalyzing all the processes involved:
- enolization of (chloro)acetyl chlorides, thus catalyzing the di- and trichlorination of chloroacetyl chloride (every new alpha-chlorine group reduces the reactivity of the enolized form via inductive effect);
- the activation of TCCA (TCCA gets much, much more electrophilic when activated by acids);
- the activation of TCT (protonated TCT should be more reactive toward carboxylic acids and might give off the last chlorine more readily; essentially you can act in two ways: increasing the nucleophilicity of AcOH by deprotonation with Et3N which is unfortunately not applicable in this case, or increasing the electrophilicity of TCT which is applicable);
- Cl3COCl should be easily decomposed by AlCl3 (just guessing, should check the literature but today I'm bored enough to make a lot of guessing).

Anyway, forget AlCl3. I remembered of an acid that would work as a great catalyst and a solvent at the same time. Sulfuric acid. TCCA is well soluble in it and it makes it incredibly electrophilic (there is a paper in J. Braz. Chem. where they actually did DFT calculations about it).

Edit: Damn, forgot that CCl4 decomposes to COCl2 in conc. H2SO4.

[Edited on 6/4/2008 by Nicodem]

Sauron - 5-4-2008 at 15:03

March did'nt happen to cite references for such reactions, did he? I just can't recall ever seeing a halogenation of acetic acid other than photolytic and usually with a "halogen carrier" as well.

But per March, it ought to be hard to stop at chloroacetyl chloride, or bromoacetyl bromide. You'd expect the dihalo product to predominate or at least be a major contaminant.

There ought to be a review of the Hell-Volhard Zellinski reaction in OR or Chem Rev and Merck would be where I'll look first in the named-reactions section. If unproductive, next the ACS search engine.

Thanks, Nicodem, for that Synth.Comm. paper. They got an 88% yield of chloroacetic acid using excess TCCA at 130 C for 19 hours. The acetyl chloride was generated first with PCl3 to the extent of about 2%. Purity by GC, 99.4% so no indication of dichlorination.

So far it sounds like a good cheap replacement for NCS and no one ought to be amazed that a chloamide can do the job of a chlorimide. I want to hit the lit looking for di- and trichlorination examples under HVZ conditions.

I found four reviews cited by Merck. Two in Chem Rev. and two in JACS. I've requested them and when they appear we will see what we will see.

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

So far I have obtained and read two of the four HVZ reviews cited in Merck Index. Neither of these makes any mention of poly-a-halogenation of carboxylic acids or acid halides. I will await the other two reviews which may be more pertinent.

[Edited on 6-4-2008 by Sauron]

Nicodem - 6-4-2008 at 02:07

I checked the literature about catalysts for decomposing CCl3COCl into CCl4 and only found two references at all about such decomposition:

- one is a thermolysis:
Quote:
Decomposition of Trichloroacetyl Chloride.-Several passages of this material through the furnace at 600°C caused about one-third to react. Carbon tetrachloride and hexachloroethane were formed in the ratio of about ten to one. Carbon monoxide and some phosgene were also produced.

From: J. Am. Chem. Soc., 61 (1939), 435-436. DOI: 10.1021/ja01871a061


- the other, even older, is the mechanistically already logical decomposition by AlCl3 (Unfortunately I don’t have access to this journal and the reaction conditions are not mentioned in the abstract, neither is this reaction the main topic of the paper.):
Quote:
The Modifications of Metachloral and the Decomposition of Chloral by Aluminium Chloride. Perchlorobutanal: CCl3CCl2CCl2CHO.
Boeseken, J.; Schimmel, A.
Recueil des Travaux Chimiques des Pays-Bas et de la Belgique, 32 (1913) 128-133.


I also checked for the preparation of trichloroacetic acid by the Hell-Volhard-Zelinsky reaction and found this patent US2613220 where 3mol% acetic anhydride is used as a reagent/source for enolization. I could find no other electrophilic chlorination, but besides the HVZ reaction there is also the radical chlorination of AcOH which does not require additions of reagents promoting the enolization (since the mechanism is completely different), yet it requires UV light for the homolytic dissociation of Cl2 (one such example is patent US2674620).

I guess it would be much better to start with chloroacetyl chloride which is relatively easily prepared from trichloroethylene (probably the most common chlorinated solvent that can be found OTC all over the world). With the appropriate corrosion resistant autoclave it is a quite easy preparation (I know there is a thread about it with the patent numbers and all somewhere on the forum). Chloroacetyl chloride could then be hopefully directly perchlorinated with TCCA/AlCl3 to CCl4 and CO (with all due safety requirements due to phosgene and other nasties).

Sauron - 6-4-2008 at 02:32

Better, don't you think, to chlorinate ethanol (anhydrous) to chloral? This does not require, nor is it promoted by UV.

It's high yield.

Once in the trichloroacetaldehyde form, only the final chlorination step to trichloroacetyl chloride needs photolysis. Surely preferable to irradiating it all the way from GAA?

As it happens one of my suppliers has a second hand Ace power supply for a Hanovia immersion lamp for sale. I will have price in next 24-36 hours. The special joint on Ace's photoreactor flasks is 60/40. Other glass companies sell 60/50 joints that will certainly work so I can have a flask made up locally to accept an Ace immersion well (quartz) and lamp to match the power supply.

Ethanol to CCl4.

Who'd have figured?

Here's the Chem.Rev. paper on chloral chemistry, its preparation, its photolytic halogenation, and a lot more.

Yes, Rec.Trav.Chim. is unavailable online so far. Too bad. The reference for decomposition on trichloroacetyl chloride in Brown's paper (JACS) is same one you cite and I have never been able to lay hands on it. Brown made no mention of AlCl3, so your efforts are not without reward. We have more information than before.

[Edited on 6-4-2008 by Sauron]

[Edited on 6-4-2008 by Sauron]

[Edited on 7-4-2008 by Sauron]

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chemoleo - 6-4-2008 at 06:07

I thought it worth noting that according to the excellent review, Chloral can also be produced by electrolysis of ethanol and NaCl or CaCl2 at 110-115 deg C, and yields of chloral rise to 61% in the presence of cyanuric acid as catalyst.
Even simpler and cheaper than straight chlorination, the only caveat being the temperature - the electrolysis vessel would be pressurised (EtOH BP: 78 deg C).

Then of course chloral can be decomposed to chloroform and formate with strong alkalis.

(THis is off topic but quite interesting - chloral also reacts with phenol to parahydroxy benzaldehyde and other goodies, see p 14)

[Edited on 6-4-2008 by chemoleo]

not_important - 6-4-2008 at 07:30

Quote:
Originally posted by Sauron
March did'nt happen to cite references for such reactions, did he? I just can't recall ever seeing a halogenation of acetic acid other than photolytic and usually with a "halogen carrier" as well.
[Edited on 6-4-2008 by Sauron]


No, not a single ref given in that paragraph. Several surveys, which you seem to be tracking down. But much of the original HVZ stuff seems to fall into the category of Really Old Chemistry, where stones axes were used to carve flasks out of tree trunks, and the nearest thing to a reference is a collection of lists of surveys.

Sauron - 6-4-2008 at 13:43

No, chemoleo, not OT at all. Chloroform is in the thread title, and my intent, given a facile route to CCl4, was to turn part of that into chloroform anyway (with wet Fe filings, one Cl is pulled off.)

So a direct transformation of chloral to chloroform is certainly of interest With formate for a bonus. I like it.

It looks like there is paydirt in the Sonntag review of aliphatic acid chloride chemistry (Chem.Rev. vol.52). In this long and rather comprehensive paper, halogenation of acid chlorides is buried in the middle and I almost overlooked it. But, indeed, it seems that acetyl chloride is chlorinated to dichloroacetyl and trichloroacetyl chlorides in absence of any powerful light source, by the addition of a small amount of iodine. If this reaction is carried out at too high a temperature (200 C) then hexachloroethane is produced. Clearly, although Sonntag makes no mention of CCl4 being formed, the hexachloroethane is resulting from the union of two CCl3 radicals rather than that of one CCl3 radical and a Cl radical. We will have to have recourse to Sonntag's references.

Attached is p 359 from Sonntag, the reaction he is talking about it chlorination of acetyl chloride with PCl3 at 190 C. which gives a mixture of ClCH2COCl, Cl2CHCOCl, and Cl3CCOCl. However, a mere 10 degree increase in reaction temperature is reported to yield instead hexachloroethane.

Three references are given, all accessible fortunately, in the 19th century German primary literature. One from Ber., one from Ann., and one from J.Prakt.Chem. Sounds like a job for References to me. I will go ferret out the DOIs.

While this helps save March's face, it is unlikely to provide a practical alternative to the chlorination of ethanol to chloral and the decomposition of that material to CCl4. The major product would have to be the trichloroacetyl chloride, TCCA or some other cheap reagent would have to replace PCl3 which is now so politically incorrect. But I will follow through and see where this takes us.


[Edited on 7-4-2008 by Sauron]

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Sauron - 6-4-2008 at 18:01

I got the first of the references from Gallica (BNF)

Attached below

He is proposing to prepare trichloroacetyl chloride from acetyl chloride and PCl5 (3 mols) according to the equation

CH3COCl + 3 PCl5 -> Cl3CCOCl + 3 PCl3 + 3 HCl

Three mols PCl5 is more than 500 g, to treat one mol acetyl chloride? It must have been nice in those days to be so extravagent. This method may be noteworthy as a method of obtaining PCl3 from PCl5 and therefore from red P.

As for the main thrust of the paper, : m-nitro-p-trichloroacetoluid is presumably N-trichloroacetyl-3-nitro-p-toluidine.
]
Trichloroacetyl chloride was also prepared conventionally by trating trichloroacetic acid with PCl5. In this case the byproduct is of course POCl3. No discussion of details, yields, analysis is given, the prep of the Cl3CCOCl is merely mentioned in passing as a preliminary matter.

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Nicodem - 7-4-2008 at 02:30

Iodine is indeed an excellent Lewis acid for certain reactions, perhaps just because it so mild as to prevent certain side reactions that would otherwise occur with such strong acids like AlCl3 (for example, direct chlorination of AcCl to CCl4 :P).

But otherwise acyl chlorides enolize so easily that for monochlorination you don't even need any iodine or other acid catalysts. However, if you will want to use TCCA for alpha-perchlorination of AcCl you will probably have to add an acid catalyst to get it go to CCl3COCl at some acceptable conditions, but only experiments can tell.

The mention of PCl3 as a chlorinating reagent is probably a mistake. PCl5 is the reagent used in such HVZ reactions as a replacement for Cl2 (PCl3 is generally not an oxidant). We'll see what the original references say.

Sauron - 7-4-2008 at 03:06

Yes, well, at least the sole reference that I have so far used

-- preformed acetyl chloride
-- stoichiometric PCl5 (3 mols) for the alpha-perchlorination

So Sonntag's entry is a little off.

But I have yet to see the other two references.

To my mind, ethanol is the better feedstock and chloral the more versatile intermediate.

Treat it with strong base and get chloroform

Hit it with UV and Cl2 and get CCl4.

And there's still the Boeseker report of decomposition by AlCl3.

So acetyl chloride, more precious than ethanol, is a less attractive feedstock. It might be interesting to establish the conditions under which TCCA would chlorinate chloroacetyl chloride to dichloroacetyl chloride. And that to trichloroacetyl chloride. In presence of and absence of catalysts. But looking at boiling points, and the time and temperature at which excess TCCA only formed the monochloroacetyl chloride, I suspect we may be looking at conditions of autogenous pressure in an autoclave.

Also easy enough to work on the conditions for decomposition of trichloroacetyl chloride. We already are taught that 200 C results in hexachloroethane, and we can safely assume that this is a competing reaction involving the same Cl3C, radical as formation of CCl4. So this represents a ceiling. We don't want hexacloroethane. Anyway we are told in the Sonntag review that CCl4 on its own at a certain temperature spits off a tenth of a mol per hour Cl2 and forms hexachloroethane. So that compound is always available easily if you have ample CCl4.

Nicodem - 7-4-2008 at 03:21

Chloral can also be a source to CCl3COCl possibly by chlorination with TCCA directly (I posted about the relevant patent in the thread dealing with benzoyl chloride preparation. I think chloral was not among the examples but I can't see any reason why it would not work as well.).
Quote:
Originally posted by Sauron
Also easy enough to work on the conditions for decomposition of trichloroacetyl chloride.

I think the best preparative option is to decompose it with AlCl3 since this should result in less side products, given that it is about a simple and well documented carbocation fragmentation (not involving radical species and no radicals can diffuse to couple into undesired side products like C2Cl6 and similar).

[Edited on 7/4/2008 by Nicodem]

LSD25 - 7-4-2008 at 05:23

I remember reading somewhere that trifluoroiodomethane is formed by the reaction of trifluoroacetic acid with silver iodide. I know where halodecarboxylation is concerned that there have been some serious moves away from silver salts to affect the reaction, so would other salts halo-decarboxylate the trichloroacetic acid?

https://www.chem.ubc.ca/faculty/wassell/CHEM415MANUAL/Experi...

https://www.chem.ubc.ca/faculty/wassell/CHEM415MANUAL/Experi...

ie. heat the stuff in a vessel with iodine - is it possible to swap it?

Nicodem - 7-4-2008 at 05:58

It is called the Hunsdiecker reaction and in the case of preparing CF3I surely requires I2 and silver trifluoroacetate (not trifluoroacetic acid and silver iodide!). There was a bromobenzene thread where it was recently discussed. It requires the formation of RCOOX (where X is a halogen) type of intermediate which decarboxylates to RX. Using it for preparing CCl4 would not be particularly practical unless one could use sodium instead of silver salt (and still for small quantities only or proof of concept). Besides I don't even know how well, if at all, it works with X=Cl.
However, for CCl3X (where X=Br or I), this surely is a good route.

LSD25 - 7-4-2008 at 06:34

I recently cited a number of articles on this very reaction where the author's used N-Halo reagents and other salts (lithium IIRC) to improve this reaction. I think not_important was suggesting that the non-metal variant was not actually a Hunsdiecker-Borodin-the rest, reaction, but a separate reaction altogether (that is if I understood correctly - which ain't always so). I recall that the various author's used predominantly halosuccinimide reagents for this, but what are the chances that TCCA would work?

Nicodem - 7-4-2008 at 07:32

Well, yes in theory (read: on paper) the reaction should work even on plain acid due to the equilibrium:
RCOOH + X<sub>2</sub> <=> RCOOX + HX

However halogens dissociate very badly in carboxylic acids and the above equilibrium is way to the left. Using a sodium salt of the carboxylic acid would help moving the equilibrium more toward the right by lowering [HX] by neutralization (that is, by introduction of another equilibrium: RCOO<sup>-</sup> + HX <=> RCOOH + X<sup>-</sup>;). Yet the [RCOOX] would still be too low for a rapid conversion to RX without side reactions prevailing (for example, when the decarboxylation takes too much time most of the formed RX will succumb to the nucleophilic substitution with RCOO<sup>-</sup> forming the RCOOR ester). Using a silver salt is optimal since it quantitatively and rapidly forms RCOOX due to the AgX precipitate forcing the equilibrium quantitatively to the right (and still you often end up with low yields).
Surely one could try to use the free acid and exploit the above equilibrium even if far to the left by instead working on making the next step faster:
RCOOX <=> [RCOO* + X*] <=> [R* + CO<sub>2</sub> + X*] => RX + CO<sub>2</sub>

Yet that requires irradiation with UV light and I can only guess what other side reaction would such a treatment promote (one for example, is the above thread talked about radical alpha-halogenation of carboxylic acids – the alternative to the HVZ reaction yielding the same products).

I checked if the Hunsdiecker reaction can be used in preparing alkyl chlorides and indeed it can be. However, mind that using silver carboxylates only works with X<sub>2</sub>. Here you can not just substitute TCCA instead of Cl<sub>2</sub> since then you ruin the reaction by not having the AgCl precipitation to drive the reaction. TCCA can only work in methods developed for N-chlorosuccinimide (NCS) if you can find one such (and if there exist I bet they are substrate specific rather than general).

PS: Though this is still somewhat on thread topic it is only barely so and with the potential of straying away given it has little relevance for CCl4 or CHCl3 production, so I suggest you open a thread dedicated to the "Hunsdiecker-Borodin-the rest" reaction if you are really interested in it. Surely several members would appreciate if you compile whatever information you have.

Sauron - 7-4-2008 at 08:51

Nicodem, the chloral review from CR did not mention side reactions, but stated that in the photolytic chlorination of chloral at temperatures above 10 C, no Cl3COCl was isolated, only CCl4 and the gaseous products from the radical fragmentation (CO, HCl, and sometimes small amount of COCl2.) This is very clean.

The entire two step process starting with ethanol consumes only Cl2 (from TCCA and HCl) and in final stage only, UV. No other reagents. None.

So the sole advantage of stopping at first stage and utilizing the chloral react with AlCl3 and TCCA would be that a UV source would not be required. The downside is that this is experintal. I hardly have anything against experimentalism (else I would not be here) but, it is not the path of least resistance.

Side reactions and C2Cl6 formation were the problem with the perchlorination of acetyl chloride with 3 mols PCl5, an absurdly expensive approach I'm sure even in the 19th century and insupportable today. 600g PCl5 to treat a mol of CH3COCl? That's seventy something grams substrate. And even so, the C2H6 problem on reared its head at 200 C. By comparison the photolytic chlorination of chloral runs quantitatively at ambient temperatures.

I must admit I have not yet looked at those patents but as you say, TCCA was not among the examples andin general patents are not things I like to base much of anything on. This is an old prejudice inculcated into me at university, that patents are about as reliable as Albanian politicians or Iranian bazaaris.

I wish we could obtain the damned Boeseker paper from Rec.Trav.Chim. that I have been chasing for a year ever since seeing it in H.C.Brown's footnote from 75 years gone. It details the use of AlCl3 to decompose chloral and/or trichloroacetyl chloride.

That decomposition of the acid chloride occurs under Brown's conditions with no UV and no AlCl3 to the extent of about 35-40%. Whether similar thermolytic fragmentation will occur with trichloroacetic acid and TCCA, I do not know. We have found no examples of dichlorination or perchlorination with TCCA, only monochlorination. The Trichlorination reactions on the acid chloride were only with PCl5 and do not seem to have ever seen any replication in the 20th century.

My working hypothesis, till the rest of the refs are in, and we can start to get some experience with this, is that the ethanol-chloral-then UV to CCl4 route is safe and sound.

Meanwhile the GAA or chloral to trichloroacetic acid route, is also sound, but the path from that point to CCl4 via AlCl3 is ill lit and requires some slashing at vines and hacking of trails, i.e., wet lab time.

So one is a solid preparative route and the other a good start to a research project. My concern is shortest distance from here to CCl4 with the least resistance. You, as you mentioned, are blessed with circumstances where CCl4 is off the shelf. Lucky you! Unfortunately I do not share this happy state. So while I share youe enthusiasm for thus new adventure, I prefer the surer path for the present.

Sauron - 7-4-2008 at 10:09

I read that US patent. Here it is in case anyone else wants to have a look.

Not very interesting.

Nicodem, the trichloroethane (I guess 1,1,1-trichloroethane) ubiquitous solvent you were talking about, There was a discussion of an old patent last year in regard to making Ac2O or acetyl chloride from this. But the patent was very broad and vague, and examples they gave required an autoclave. Chloroacetyl chloride was not a product.

That's the only discussion I recall involving 1,1,1-trichloroethane.



[Edited on 8-4-2008 by Sauron]

Attachment: US2613220.pdf (384kB)
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trilobite - 7-4-2008 at 15:12

I don't think that a radical pathway is needed to explain the decomposition of the halogenated acyl halides. I assume that the reaction proceeds through formation of a di- or trichloromethyl carbocation stabilized by the electronegative chlorine atoms. Obviously both Brönsted and Lewis acids catalyze the reaction.

These are not the only acyl halides where such phenomena can be observed. For example when pivaloyl chloride (aka trimethylacetyl chloride) is used in a Friedel-Crafts acylation of benzene most of the acyl chloride is decomposed with liberation of carbon monoxide and the resulting tert-butyl carbocation alkylates the substrate. If I remember correctly the experimental result was that if the temperature is kept low enough some pivaloylated product can be isolated, but most of the product consists of tert-butylbenzene.

I dug up some abstracts and references relating to the subject. Some of them discuss the reverse reaction, formation of the acyl halides through carbonylation, while others discuss the Friedel-Crafts reactions without decarbonylation.


Catalytic Action. III. 1. Friedel-Crafts Reaction.
Boeseken, J.
Delft. Rec. trav. chim., 29, 85-112 (1910)

Abstract
The author believes that catalysts do not react through intermediate addition products but by changing the affinities for each other of the ats. in the reacting substances, an at. or group of ats. being either (1) entirely eliminated from the rest of the mol. and reacting as such or (2) activated without being separated from the rest of the mol. E. g., (1) when SOCl2 reacts with PhCl and AlCl3, no dichlorodiphenylsulphone, but only polychlorobenzene, is formed; (2) BzCl and C6H6, in the presence of AlCl3, give BzPh. This dissociating action of AlCl3 was investigated with chloral and substituted Ac chlorides. Trimethylacetyl Chloride is decomp. very energetically by AlCl3 even at 0 Deg into HCl, CO and isobutene, C4H8. At ordinary temp. H2SO4, which in many respects is analogous to AlCl3 as a catalyst, liberates HCl from Me3CCOCl with formation of Me3CO.SO4H. On warming, CO is set free and the isobutene dissolves in the H2SO4 to form isobutenedisulphonic acid, C4H6(SO3H)2. Chloral is decomposed chiefly into CO, HCl and CCl2 which polymerizes to C2Cl4. If less than 1 mol. AlCl3 is used and the duration of the reaction prolonged, other products, such as COCl2, C2Cl6, C5H2O3Cl6 (chloralide), are formed in small quantity. Dichloroacetyl Chloride with excess of AlCl3 gives CHCl3 and CO. With smaller quantities of AlCl3, much HCl and little CO is evolved at first, then the amount of CO increases and a compd. C5H10, which is probably perchloropentamethylene, needles, m. 32 Deg, b45 132 Deg, is obtained. Trichloroacetyl chloride is very slowly decomposed by AlCl3 into CO and CCl4.


Reactions of carbon monoxide and sulfur dioxide with polychlorinated methanes.
Frank, C. E.; Hallowell, A. T.; Theobald, C. W.; Vaala, G. T.
Journal of Industrial and Engineering Chemistry (Washington, D. C.), 41, 2061-2 (1949)

Abstract
Chloroacetyl chlorides are obtained from CO and polychlorinated methanes at pressures of 850-950 atm. and in the presence of an AlCl3 catalyst. Optimum conditions for the prepn. of Cl3CCOCl comprise the use of 0.1 mole AlCl3/mole CCl4 at 200 Deg and 950 atm. CO pressure for 2 hrs. Beyond a certain min., time is not crit. Reaction begins at 100 Deg and a CO pressure of 50 atm. Polyhalogenated ethanes also react with CO, but the yields of acyl halides are low due to side reactions. SO2 can be made to react with CCl4 or CHCl3 to give SOCl2. With an excess of SO2 and in the presence of AlCl3, yields of 70-80% were obtained. The AlCl3 retains some activity after use.


Patents related to the invention are US2378048 and GB581278.


Reaction between aromatic compounds and derivatives of tertiary acids. V. The stability of the carbonyl group in certain acid halides and anhydrides
Rothstein, Eugene; Saville, Rowland W.
Journal of the Chemical Society, 1949, 1961-8

Abstract
The factors influencing the acylation or alkylation of aromatic compds. by derivs. of acids in the Friedel-Crafts synthesis are discussed. It is suggested that the stability of the acyl cation [R3CC+O] dets. the relative rates of the 2 processes. Loss of CO results in the formation of an electrophilic carbonium ion [R3C+], which can be substituted in the nucleus, yield an olefin, or undergo rearrangement. The effect of substituents on the relative ease of elimination of CO is also examd. and it is shown that the decompn. may be used for the detn. of the course of the reaction. Thus, condensation of diarylacetyl chlorides with C6H6 is attended by the simultaneous liberation of CO; this was not detected by previous workers in this field and, in consequence, an explanation advanced by McKenzie, et al. (C.A. 27, 81), of the formation of Ph2CH2 and Ph3CH from p-MeC6H4CH-PhCOCl is untenable. These diarylacetic acid derivs. provide the 1st instance of an unstable secondary acyl ion, Ar2CHC+O. Me2C(COCl)2 (13.5 g.), 22 g. AlCl3, and 140 cc. C6H6 give 52% Me2CBz2 and a negligible vol. of CO; 14 g. Me2C (COCl)2, 23 g. AlCl3, and 6.5 g. C6H6 in 120 cc. CS2 give 8.1% CO, 23% iso-PrPhCO, 8.1% Me2CBz2, and 0.5 g. 2,2-dimethyl-1,3-indandione(?); similar results were obtained with Et2C(COCl)2. ClCH2COCl and C6H6 give 87% BzCH2Cl and 17% CO; Cl2CHCOCl gives 79% BzCHCl2 and 3.6% CO; Cl3CCOCl gives 56% BzCCl3; Cl3CCOBr gives the same result. Me2CBrCOCl and Me2CBrCOBr give 2-methyl-1-indanone and Me2CBrBz. Camphoric anhydride (I) (10 g.) in 160 cc. C6H6, treated with 15 g. AlCl3, gives 89% CO and 67% 2-phenyl-1,1,2-trimethyl-5-cyclopentanecarboxylic acid and a small quantity of a hydrocarbon m. 80-1 Deg; PhOMe gives 3.4% CO and 89% 5(or 2)-anisoyl-1,1,2-trimethyl-2(or 5)-cyclopentanecarboxylic acid. PhMe yields 3% CO and 82% of the 2-toluyl deriv. PhCMe3 (10.5 g.) in 10 cc. CS2, added to 7 g. I, 10.5 g. AlCl3, and 70 cc. CS2 at 0 Deg and warmed, give 74.2% CO, 2.7 g. PhCMe3, a smaller quantity of p-C6H4(CMe3)2 and, after methylation, Me 2-(p-tert-butylphenyl)-1,1,2-trimethyl-5-cyclopentanecarboxylate, b0.2 110-15 Deg. PhNHAc yields mainly Me isolauronolate. p-MeC6H4CHPhCOCl and C6H6 with AlCl3 in CS2 give 70% CO and 33% Ph3COH; PhOMe yields little CO and 88.6% methoxyphenyl a-p-tolylbenzyl ketone, m. 107-8 Deg. Ph2CHCOCl and C6H6 give 68% CO and BzCHPh2. Et2CHCOCl and C6H6 yield 7% CO and 81% BzCHEt2 (2,4-dinitrophenylhydrazone, m. 115-16 Deg). AcNHCH2-COCl and C6H6 give 8% CO and BzCH2NHAc.


Decarbonylation of chloroacetic acid chlorides in the presence of aluminum chloride
Pomerantseva, E. G.; Kulikova, A. E.; Zil'berman, E. N.
Zhurnal Organicheskoi Khimii, 5(1), 187 (1969)

Abstract
Heating a mixt. of ClCH2COCl and AlCl3 gave CO, HCCCl, and a polymer (I) with conjugated double bonds. Similarly, Cl3CCOCl with AlCl3, gave CO, CCl4, ClCCCl, and a polymer. Heating CH2Cl2 with AlCl3 also gave I. A mechanism is proposed.


Direct carbonylation of polychloroalkanes to acid chlorides using metallic salt ternary systems: An example of multistep catalysis.
Monflier, Eric; Mortreux, Andre; Petit, Francis; Lecolier, Serge.
Journal of the Chemical Society, Chemical Communications, 1992(5), 439-41

Abstract
A catalytic cycle for the direct carbonylation of CCl4 to CCl3COCl, catalyzed by metal salt mixts., e.g., AlCl3/MCln/CuCl, under unexpectedly mild conditions, is proposed on the basis of FTIR, 17O and 27Al NMR spectroscopic studies.


Patents related to succesful Friedel-Crafts trichloroacetylation of substrates are EP189266, US4724267, US4731484, FR2677645. DE2648134 is about synthesis of trichloroacetyl chloride from CO and CCl4.

Sauron - 7-4-2008 at 20:39

Thanks. That argues against the use of AlCl3.

Those reactions do not proceed cleanly. Numerous side products form.

In comparison, the photolytic chlorination of trichloroacetaldehyde in absence of AlCl3, certainly a radical reaction, proceeds exclusively to CCl4 (and CO and HCl) and there is no reverse reaction.

This reaction takes place at temperatures from beloe ambient (such as 10 C) to 90 C.

Only at -10 C and below does any intermediate Cl3CCOCl survive. At 10 C, 7.5% and at -50 C, 55%.

Sauron - 8-4-2008 at 08:29

The preparation of chloral by reaction of 1,1,1-trichloroethane with aqueous hypochlorous acid is subject of a US patent. This may have been what Nicodem had in mind rather than that catalytic reaction.

Meanwhile I reread two patents from my files on chlorination of ethanol. These teach that 95% ethanol can be used instead of anhydrous, and unlike the chloral review, make no mention of cooling to 0 C in the initial (oxidation) stage of the reaction.

Chlorination of ethanol is in any case a much faster process than chlorination of glacial acetic acid to same degree of chlorine content and requires no UV irradiation.

Ethanol perchlorinates to chloral in about 6 hours compared to 60 hours for GAA to trichloroacetic acid.

Normally 2 mols ethanol are chlorinated to one mol chloral "alcoholate" (trichloroacetaldehyde ethyl hemiacetal) and chloral is then liberated with conc H2SO4 by distillation.

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

Back on MeSCN

The hangup in the DMS route to MeSCN is barium thiocyanate, which is very costly; and it can be made from ammonium thiocyanate, also relatively dear. Ammonium thiocyanate is made from ammonium dithiocarbonate, and that is made by adding NH3 to CS2 in water. Inorg.Syn. has the prep, I will have to look more closely to see if this is economically competitive with the alternative, which is to use DMS to make MeI, then react MeI with KSCN in EtOH. CS2 is after all rather costly.

[Edited on 9-4-2008 by Sauron]

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Sauron - 9-4-2008 at 20:29

Preparation of Chloral (Trichloroacetaldehyde) per Rosin's Patent

An apparatus is assembled consisting of a 1 liter RB flask with four necks, a termometer or thermocouple adjusted to nearly the bottom of the flask, a mechanical stirrer and an efficient reflux condenser and a gas inlet tube with coarse fritted tip. The flask is supported in a heating mantle. The condenser is topped with a gas outlet tube leading to a large caustic trap for receiving HCl produced in the reaction and any Cl2 not absorbed.

The gas inlet tube is connected to a chlorine gas generator charged with TCCA and equipped with a dropping funnel or other means of admitting 5 M hydrochloric acid at a controlled rate. The 5 M HCl is prepared by cuatiously mixing equal volumes of concentrated hydrochloric acid and water.

The chlorination flask is charged with 200 cc (3.37 mol) 95.5% ethanol, the stirrer is started, and chlorine is admitted at a rate of approx. 3.5 g per minute for about 2.5 hours. The temperature of the reaction mixture will have risen substantially during this stage of the chlorination, and condenser water is recirculated through a chiller to reflux any vapors. The reaction mixture at this point will weigh about 260-280 g and have a density of c. 1.29-1.35 @ 25 C. The material is now principally dichloroacetaldehyde ethyl hemiacetal. The caustic trap will have increased in weight by about 245 g.

120 ml water is now added to the reaction mixture.

The chlorine rate is now reduced somewhat to 2.5 g/min, and the reaction mixture is maintained at reflux by use of the heating mantle until the boiling point reaches 95 C. This process takes about a further 8 hours. The chlorination is complete for practical purposes when the weight of the reaction mixtures has reached 500-550 g and it has a density of 1.50-1.57 @ 25 C. Heating and chlorine generation are not halted. The reaction mixture is allowed to cool to room temperature and is then mixed with an approximately equal volume of concentrated H2SO4, thus liberating chloral from its hydrate. Upon distillation 364 g chloral is obtained. (73% yield).

Sauron's notes:

3.5 g/min x 150 min = 525 g Cl2 generated in initial phase.
2.5 g/min x 480 min = 1200 g Cl2 generated in second phase.
Total chlorine generated 1725 g

So we can calculate the amount of TCCA required, the amount of 50% dil HCl needed and the parameters of the caustic scrubber. I estimate it will require c. 5 Kg TCCA to generate 1725 g Cl2. That is based on 4 mols TCCA to the Kg (actually that figure is a little low) and about 90 g available Cl2 per mol TCCA (which is about right.) A Kg TCCA requires about 2.4 L of 5 N HCl, so about 12 liters (made from 6 liters conc HCl and 6 liters water) will do the trick.

It seems like > 30% of the generated chlorine is going to waste most likely in the latter stage of the chlorination, and ends up in the scrubber as chlorate and chloride.

Overall stoichiometry under anhydrous conditions

2 CH3CH2OH + 4 Cl2 -> Cl3CCH(OH)(OCH2CH3) + 5 HCl

Overall stoichiometry when water is added

CH3CH2OH + 4 Cl2 + H2O -> Cl3CCH(OH)2 + 5 HCl

Both of these and the production of chloral diethyl acetal are all happening here. The point of Rosin's patent is to maximize the chloral hydrate vs the chloral alcoholate so as to make better use of the ethanol. He claimed about 60% more chloral from a given amount of ethanol by this method.

One final comment regarding chlorine absorption and generation:

My estimates are based on Rosin's stated rates for the first 2.5 hours (3.5 g/min) and the following 8 hours) 2.5 g/min). The assumption, very likely faulty is that these are constant. I think in practice it is better to shoot for about 3 g/min throughout, this may lengthen the first stage a little and reduce time of second stage a little, those points are determined by weight and density not by the clock. Also the real limiting factor is the capacity of the reflux condenser.

It is highly likely that the chlorine generation will slow down over time and require gentle heating to drive it. Or the chlorine generator may need recharding with TCCA and HCl. Trial and error will determine the appropriate scale to achieve target 3 g/min flow rate and how to keep that rate for 10-11-12 hours. The correct flow rate is determined by chlorine absorption in the reaction, so the off gases need monitoring. The appearance of yellow green Cl2 means, slow down the Cl2 flow. This would be easily done from a tank, but from a Cl2 generator, that's another story. Trial and error will determine the best conditions.

The chloral once isolated from conc H2SO4 is ready to be either converted to CHCl3 with base, which is a reaction dependent on both pH and temperature; or charged to a photochemical reactor for irradiation and chlorination to trichloroacetyl chloride, which falls apart to CCl4, CO and HCl. I propose to do this in a 5 liter Ace setup using a 450 W Hanovia high pressure Hg-quartz arc lamp #7825-35 in a quartz immersion well, the amount of charge being sufficient to cover the full length of the arc. Ace has both conical and spherical reactors in this size. I have located a second hand power supply to match, a 7830-60 and this will save me some money.

Given the equipment at hand, the largest scale I can run is about 1 Kg chloral product per day so I will probably only do the photochemical step once every 3-4 days (and in the hood.) I have to build a containment for it to avoid UV exposure. The CCl4 product should about equal the chloral charge according to the mass balance. So maybe a week's work to produce 6-7 L CCl4. And that drum of TCCA will need replacing.

[Edited on 10-4-2008 by Sauron]

Sauron - 10-4-2008 at 23:21

Again on MeSCN:

The preparation of ammonium dithiocarbonate is described starting on page 48 of Inorganic Syntheses vol 3.

NH3 is bubbled into a solution of CS2 in isopropyl acetate at 25-30 C.

Yields are 92-96%.

According to Ullmann's the unstable dithiocarbonate is decomposed at 95 C to ammonium thiocyanate, H2S and S.

The preparation of barium thiocyanate from ammonium thiocyanate and barium hydroxide is described in same volume of IS starting on page 24.

This salt is reacted with dimethyl sulfate to give methyl thiocyanate.

In my opinion, it is better to utilize the DMS to prepare MeI from KI according to Org.Syn., then react the MeI with KSCN in EtOH according to Vogel.

The virtue of the second route is that no CS2 is required and fewer steps are involved. Both routes use DMS. I will have to look at the economics of this vs buying MeI or buying MeSCN, because if buying MeI is not so costly, then avoiding use of DMS is a good thing. I have the stuff, but. But.

woelen - 12-4-2008 at 02:43

According to my calculations, you will need almost 4 kg of TCCA (theoretically). In practice, probably you are right with your estimate of 5 kg, because at the end the reaction will become very slow, due to all the cyanuric acid, which makes mixing of dripped acid with the remaining TCCA more and more difficult.

If I were you (and if you have the right tubing), I would make a setup with 2 chlorine generators, connected to a T-shaped glass tube. First, you use one chlorine generator, and if this becomes too slow, you use the other and stopper the end on the T-tube for the exhausted one. In this way you can recharge the first one, while the second one is running. In this way, you could use alternating chlorine production with small generators. I think that nice steady flow of chlorine gas for so many hours will be very difficult to achieve, hence my suggestion of using alternate generators with smaller quantities of TCCA.

My experience with TCCA is quite good, when it comes to purity of the chlorine, but a constant rate is difficult. You'll observe a gradually decreasing rate, as more and more cyanuric acid builds up. With respect of constant production rate, tablets (not powder or granules!!) of Ca(OCl)2 is better.

Sauron - 12-4-2008 at 04:03

Thanks, the alternating Cl2 generator suggestion is very good. I appreciate the input.