Sciencemadness Discussion Board - Unexpected problems with p-TsOH synthesis

Unexpected problems with p-TsOH synthesis

Lambda-Eyde - 16-10-2011 at 15:18

Hello,

I finally got the opportunity to get in the basement and do some chemistry again - I decided to synthesize p-toluenesulfonic acid (p-TsOH), which I thought would be fun and easy. I couldn't have been more wrong.

I mixed 40 mL 96 % sulfuric acid (PA, Merck) with 80 mL toluene (hardware store, undistilled) in a 250 mL round-bottom flask, attached my new, unused Dean-Stark apparatus (no pictures, sorry...) and a reflux condenser on top. About five minutes after mixing the two reagents (without heating), the color turned a slight yellow. I thought that was normal. After heating (without stirring), the color turned to a deep, beautiful red, and later to opaque black. Then I noticed a strong, disgusting sulfurous stench (not hydrogen sulfide, I can recognize that) coming out from the top of the reflux condenser, so I attached a tubing adapter and led a hose to the fan (another Keck clip wasted, seriously; those things can barely handle salty water :mad:

).

Eventually, the mixture became darker and about 2-3 mL of a clear liquid collected in the Dean-Stark trap, which I assumed to be water. It wasn't. It had that same foul smell, only more concentrated. It also had a slight tinge of toluene mixed in with the horrible smell. I stopped the horrible failure of a synthesis at that point. I put it in a test tube, and tomorrow I'm going to talk to a professor about analyzing it. I forgot to bring a sample of the "toluene" with me. I think that could have helped in solving this mystery.

Here are some pictures of the mess in the flask:

(Not even the longest column in the world could have saved this mess

)

Now, I understand that the "black tar" phenomenon isn't unusual in organic chemistry. And I also understand that the toluene is the most likely suspect in this case. But really, look at picture number two: It has a consistency like that of crude oil! If there were any impurities in the toluene (organosulfur compounds), could there be so much to cause the 120 mL of clear liquid to transform into that slow-flowing goop? And what kind of reaction would cause this, especially the smell? The off-gas reminded me of sulfur dioxide, which would mean that the sulfuric acid has oxidized something in the mix. The tarry crap smelled slightly sulfurous, but also reminded me of n-butyl bromide. I can't really describe the smell with words (I emphasize the sulfur smell too much, the tar had a more "organic" odour to it which I can't put into words).

When I get the chance I'll fractionate the toluene and try the reaction again, on a smaller scale.

Does anyone know anything that could help me solve this mystery?

bahamuth - 16-10-2011 at 17:49

In Vogel 3rd. they mention in the preparation of p-toluenesulfonic acid that one should use thiophene-free toluene.
May be the cause as Vogel also mentiones that thiophenes are present in commercial toluene and that they cannot be removed by distillation.
To lazy right now to write the procedure, but might if you ask though..

Apparently thiophenes/thiotolenes sulfonate very easily but do not know if they char easily, or if the polymerization product polythiophene is formed as to produce the tar in your pictures.
I have done this particular preparation twice and both times the sulfonation went very well with minimal charring. Used Fluka analytical grade toluene and BDH analytic grade sulfuric acid.

DJF90 - 17-10-2011 at 00:16

Its worth your time purifying your toluene before use. I suggest perhaps using the procedure mentioned in "Purification of Laboratory Chemicals", which is washing the toluene with conc. H2SO4, then fractionating the toluene, IIRC. Thiophenes present likely polymerise in the presence of acid and heat so that might be your problem.

Nicodem - 17-10-2011 at 07:16

Quote: Originally posted by Lambda-Eyde

I mixed 40 mL 96 % sulfuric acid (PA, Merck) with 80 mL toluene (hardware store, undistilled) in a 250 mL round-bottom flask, attached my new, unused Dean-Stark apparatus (no pictures, sorry...) and a reflux condenser on top.

...

Does anyone know anything that could help me solve this mystery?

It does not look much of a mystery if you put some thought in the procedure you used. Since you did not provide any reference, I can only guess that you used some totally improvised reaction conditions that never were verified earlier. Is that so? In such case it is all your own fault!

But seriously, what makes you believe that toluenesulfonic acids can be refluxed? Have you checked their b.p. before setting up the reaction? I don't think they even have a b.p. at 1 atm! Their structure looks like they decompose before they boil. If you wanted to use a Dean-Stark trap, then you should have used conditions where the reflux temperature depends on the solvent (toluene at 111 °C) and not a mixture of H2SO4 and TsOH (probably way above the decomposition temperature). If you check your reactant ratios, you can see that there is no toluene left after full conversion and that there should be no water carried over to the trap, because H2SO4 is used in excess.

I hope you understand what I'm saying. The bottom line is: Always consider what you are doing before you do it (that's what theory and literature in general serve for). If you use a method relying on a Dean-Stark trap, then make sure the reaction conditions and reactant ratios are such that there can be water formed and that the reaction temperature is kept constant (and bellow decomposition temperature). An excess of H2SO4 instead of an excess of toluene completely defies the purpose of the trap.

If you learned the lesson, then your experiment was a success.

ScienceSquirrel - 17-10-2011 at 07:30

Organic chemistry is very much about 'recipes'.
The right proportions, solvent and reaction conditions can mean the difference between success and failure, getting the desired product or a big difference in yields.
The spectre of black gloop is always with the organic chemist.
Published multistep total syntheses are the result of hundreds of reactions carried out by gifted lab rats ( PhD students ) working for years while fueled by microwave pizza, cheap beer and thoughts of coping off with that fit final year undergraduate who only wants a chemistry degree so they can shoot off to a merchant bank and earn tens of times what any working chemist will ever make.

bfesser - 17-10-2011 at 07:33

Before disposal, however, see if you can isolate anything interesting from the tar. You might find <a href="http://en.wikipedia.org/wiki/Mauveine">something interesting</a>. If you find nothing, at least you'll have practiced your bench skills. Always learn from your mistakes and try to make the best of things. As Nicodem explained, try to follow a procedure like Vogel's for this reaction. It's a simple well characterized and established preparation--no need to re-invent the wheel here. Good luck!

[edit]
Vogel is available in the ScienceMadness.org <a href="http://library.sciencemadness.org/library/index.html">library</a>. DOWNLOAD IT NOW! If there is no god, Vogel makes a damn fine substitute.

[Edited on 10/17/11 by bfesser]

Nicodem - 17-10-2011 at 12:07

Quote: Originally posted by bfesser

Vogel is available in the ScienceMadness.org <a href="http://library.sciencemadness.org/library/index.html">library</a>. DOWNLOAD IT NOW! If there is no god, Vogel makes a damn fine substitute.

The results one gets when UTFSE on "toluenesulfonic acid" can't compare to Vogel's, but they ain't that bad either:

Quote: Originally posted by Nicodem

I forgot to mention, but you don't need any concentrated H2SO4 for the sulfonation of toluene. With the use of the Dean-Stark trap you can start with a diluted acid. You just use an excess of toluene over H2SO4 instead of the opposite. Once the mixture becomes homogeneous, you can consider the reaction complete. This might be more suited for those who can get the diluted H2SO4 for electrolyte use, but have troubles acquiring the concentrated one. It worked fine for me and I started with the 35% acid.

[rant] These younger generations! They just don't have a clue what it is like to spend whole days at the library. If things go on by this pace, the next generations will be literally illiterate. [/rant]

Lambda-Eyde - 17-10-2011 at 16:07

Quote: Originally posted by bahamuth

In Vogel 3rd. they mention in the preparation of p-toluenesulfonic acid that one should use thiophene-free toluene.
May be the cause as Vogel also mentiones that thiophenes are present in commercial toluene and that they cannot be removed by distillation.
To lazy right now to write the procedure, but might if you ask though..

I have the 5th edition of Vogel's which I used as a reference when planning the synthesis. Now that I read it more carefully, the 5th edition recommends to use "pure" toluene, referring to an earlier chapter on solvent purification which also mentions that commercial toluene contains thiophene. I remember reading about this a long time ago, but I didn't think of it when I planned the synthesis.

Quote: Originally posted by DJF90

Its worth your time purifying your toluene before use. I suggest perhaps using the procedure mentioned in "Purification of Laboratory Chemicals", which is washing the toluene with conc. H2SO4, then fractionating the toluene, IIRC. Thiophenes present likely polymerise in the presence of acid and heat so that might be your problem.

I can't think of anything besides a polymerization that could cause this hellish goo.

Vogel, 5th ed. (on further reading...) recommends the same procedure. Shaking with 15 vol% of conc. H2SO4 until the acid comes out colorless, taking care not to let the temperature exceed 30oC, then fractionating.

Quote: Originally posted by Nicodem

Heh... I did check Vogel, but I wasn't impressed with the experimental using the Dean-Stark, giving a lousy yield of 22 % (!). I therefore adjusted the sulfuric acid/toluene ratio a bit (Vogel uses 100 mL toluene/20 mL H2SO4

, thinking it would result in a little more. Somehow, I ended up using almost equimolar amounts. I intended to have the toluene in excess, but I must have had some sort of brainfart when I quickly calculated the moles (perhaps I didn't take into account the high density of the acid), but I don't have my notes here. I don't think I took any detailed notes either - after all, I thought this procedure would be quite uneventful.

Quote: Originally posted by Nicodem

Hold on a bit now, you're making me look more retarded than I am!

I think you misunderstood. Or maybe you see a huge flaw in my experimental that I haven't seen (perhaps the screwed up proportions I explained over). Anyways, it was never my intention to reflux the p-TsOH, what I intended was for the water/toluene azeotrope (79,8 % toluene, bp. 84,1oC) to collect in the Dean-Stark. The reflux condenser was placed on top of the trap and was only there to make sure the vapors ended up in it. By returning the upper layer of toluene to the flask, the equillibrium would be forced to the right side. That's what you use a Dean-Stark trap for, isn't it?

Blah! Sorry Nicodem, I didn't thoroughly read the rest of your post before writing the above. I did a few hours ago, but when I finally got around to writing my answer it was someplace deep in the back of my head. I won't rewrite the above, because I hope that it shows that I actually understand the concepts involved here. I completely understand your post now after reviewing my calculations.

What got me to mess with the reactant volumes in the first place was the lousy yield cited by Vogel, but after thinking it through it's actually quite misleading since the remaining toluene can be recycled.

Quote: Originally posted by Nicodem

If you learned the lesson, then your experiment was a success.

Of course.

Quote: Originally posted by bfesser

Before disposal, however, see if you can isolate anything interesting from the tar. You might find <a href="http://en.wikipedia.org/wiki/Mauveine">something interesting</a>.

No way. The soup smelled horribly and fumed like hell. I neutralized it with sodium hydroxide and banished it to the sewers. The glassware I dumped in alkaline bleach, hopefully I'll be able to make it shine again.
I would like to make Mauveine one day (shouldn't be too hard, provided your toluene isn't spiked with thiophene

) and dye a bow-tie with it. Now that's a conversation starter!

Quote: Originally posted by bfesser

I have the 5th edition on my PC and cherish it as a christian would a bible. Though, I must admit I haven't followed all its commandments to the letter...

Quote: Originally posted by Nicodem

The results one gets when UTFSE on "toluenesulfonic acid" can't compare to Vogel's, but they ain't that bad either:

I did search and I did read, but I keep forgetting the important bits.

Quote: Originally posted by Nicodem

Oh, I'll probably be just as bitter as you in a few years... Just wait.

Checklist for next try:

Purify the toluene acc. to Vogel
Run the synthesis acc. to Vogel
Report results and (hopefully) get a pat on the back.

[Edited on 18-10-2011 by Lambda-Eyde]

kmno4 - 17-10-2011 at 22:26

Recently I have done the same thing (acid:toluene molar ratio 1:6) with the same results: darkening (->black) acid layer with SO2 as gas product. For another experiment I have taken fresh tolune from a bottle. No black colour, no SO2, yield of raw p-TsOH is better than 90% (as dried monohydrate). I still had some acid in toluene, but I did not bother to recover it, because it is less pure than the first, large batch.

bfesser - 18-10-2011 at 09:31

Low grade toluene and sulfuric acid are cheap and relatively easy to come by, so don't be so concerned with the low yields reported by Vogel. Just purify your reagents carefully before starting the reaction à la Vogel, and you should have no problems.

I suppose you misunderstood my reference to Perkin's discovery. I didn't mean you'd recover toluene, but perhaps find something else interesting. I did have an opportunity to synthesize mauveine a few years ago, and I dyed a silk bow tie with it. It's one of my most cherished possessions.

starman - 18-10-2011 at 14:00

IIRC the Vogel procedure is for pTSA sulfate with H2SO4 in excess.Gattermans (also in library) gives the monohydrate with large excess toluene and a dean stark.

turd - 18-10-2011 at 14:31

Quote: Originally posted by kmno4

For another experiment I have taken fresh tolune from a bottle. No black colour, no SO2, yield of raw p-TsOH is better than 90% (as dried monohydrate).

Really? How do you do that? Even when refluxing commercial TsOH.xH2O with toluene on the Dean-Stark I get a very dark red solution. Is there something in my (hardware store) toluene?

kmno4 - 18-10-2011 at 23:22

Hm... I do not know I did it. Some kind of magic ??
My first experiment with sulfonation was conducted with "all purpose" toluene - combined residues (mainly) from extractions.
Of course, these toluene "slops" were distilled. However, mentioned toluene was not pure enough - I was dissapointed seeing as mixture with H2SO4 becomes yellow, red and finally carbon black.
Dissapointed but not very surprised....

I had to recover toluene, combine it with the rest (from the bottle) and fractionally distill it all (Vigreux, 40 cm). Indeed, I got some residue not passing between 109,5-111 C.
This residue may explain why earlier experiments (preparation of benzylsodium from phenylsodium) were also failed :mad:

Back to TsOH.
I took fresh (but 10 years old), not used, lab grade toluene from original bottle, stored in cold dark place... blabla.
.... and got transparent and weakly brown coloured solution of TsOH in toluene. Calculated amount of water was added..... blabla
... hydrated, raw TsOH separated, washed several times with toluene....
Dried monohydrate was obtained as white powder with some brownish pieces. It has unpleasant but weak odour of "sulfur-something". Water solution is clear and odourless. For my purposes it is perfect.
Toluene from separation and washings was colourless.
That is all magic

[Edited on 19-10-2011 by kmno4]

DJF90 - 11-9-2013 at 10:05

I ran a preparation of toluenesulfonic acid a few months ago. At the end of the reaction (dean stark contained theoretical quantity of water) the reaction mixture was concentrated to a dark residue in vacuo. This residue formed a semi-solid crystalline mass, thought to be p-TsOH (anhydrous mp = 38*C) with a residual amount of toluene.

Now I've done this reaction a couple times before. Crystallisation of the reaction mixture by addition of water gives a dark grey solid (TsOH*H2O) which is a pain to purify by crystallisation (I've tried the method using HCl (g), and the one using conc. HCl for recrystallisation (dissolve solids in minimal water and crystallise the TsOH by adding 3 volumes of conc. HCl?). Seeing as both methods are shitty, I decided to be clever about it and look for a literature boiling point. 140 *C at 20 mmHg. Seems doable...

So I had some time today (and I had just finished another vac. distillation so the pump was warm and the cold trap was cold (-40*C)) and decided to give it a shot. Using Sigma's interactive nomograph, I worked out that at 0.5 mmHg, the bp of the anhydrous acid is about 75*C. No problems. My McLeod gauge read 0.8 mbar (which is about the right pressure), and the oil bath reached 188*C with a very mild reflux in the lower joint of the stillhead. Hotplate could give out no heat (too little wattage for the large oil bath I guess), so I got the heat gun out. After blasting it, the stillhead thermometer eventually started to shoot up, and I obtained a colourless viscous distillate passing over at 162-170*C.

I gave up in the end because it was taking too long, but it raises an interesting point. Using the Nomograph, this extrapolates to a normal boiling point of 380-389*C, vs. the 260.5*C that the literature value gives. The distillate is very thick like glycerol. It hasn't crystallised yet.

Paddywhacker - 13-9-2013 at 15:53

Try seeding it with the tiniest amount of the original gunk, but YMMV.

starman - 13-9-2013 at 19:19

@DJF Interestingly using your recorded BP and lit value at atmospheric it gives a pressure of about 50mbar.How confident are you in the McLeod?

DJF90 - 15-9-2013 at 12:44

Well, considering the nature of how the McLeod guage works, I'm confident I reported the correct pressure in the system. I ended up adding a little water to the distillate (some had crystallised in the condenser) just for the sake of it, and it crystallised to a beautiful pure-white mass with evolution of a little heat.

starman - 15-9-2013 at 22:15

Quote: Originally posted by DJF90

Well, considering the nature of how the McLeod guage works, I'm confident I reported the correct pressure in the system. I ended up adding a little water to the distillate (some had crystallised in the condenser) just for the sake of it, and it crystallised to a beautiful pure-white mass with evolution of a little heat.

Well that sounds encouraging anyway. Recrystallising the hydrate with whatever form of HCl I've always found painful with less than satisfying end product.

DJF90 - 15-9-2013 at 22:36

That was the motivation behind distilling it. It would have worked excellently if I had set everything up in a mantle, but I didn't expect the temp to be so high, or for my hotplate to wimp out on me. Its definately the route I'll be taking when I actually need some TsOH; make sure to tare the receiver to determine yield and to enable you to calculate the correct mass of water to add to form TsOH*H2O. It would be convenient even on several hundred gram scale.

watson.fawkes - 16-9-2013 at 06:09

Quote: Originally posted by DJF90

Well, considering the nature of how the McLeod guage works, I'm confident I reported the correct pressure in the system.

I don't know the topology of your system, but what you can assume is that the gauge is reporting the correct pressure at the location of the gauge. If there's a leak in the system, it's possible the pressure is higher elsewhere. If the gauge is located next to the pump, this is more likely that if the gauge is at the farthest end away from the pump.

DJF90 - 16-9-2013 at 08:43

The guage is located in a suitable position relative to other components in the vacuum system. All ground glass joints are greased *properly* with dow corning high vac grease. I believe it provides a correct reading of the pressure in the system.

eidolonicaurum - 2-1-2014 at 04:48

In terms of purifying the toluene, I have heard that thiophenes can be sulphonated at room temperature, which is in contrast to toluene etc. This might help with reducing the amount of tar you get. I haven't given any thought to how the thiophene(s) sulphonate would be removed though.

bfesser - 2-1-2014 at 07:15

eidolonicaurum, there's a procedure for that in <a href="http://library.sciencemadness.org/library/books/vogel_practical_ochem_3.pdf">Vogel</a> <img src="../scipics/_pdf.png" />.

Pyro - 8-1-2014 at 15:23

bfesser, are you sure there is one for toluene? I just had a look under purification of common lab. reagents, there is no procedure for toluene. However there is one for benzene, which is exactly the same.

here it is because ''I pity the fool" without a paper copy of Vogel (you know, that stuff that burns

). PDF is so frustrating.

HeYBrO - 21-1-2014 at 23:14

Regarding the purification of toluene, i found a publication here. when it says wash the toluene with water and NaOH how much is recommended? I assume because the toluene is practically insoluble you can use a lot of water each time?Thanks.

bfesser - 22-1-2014 at 05:48

Quote: Originally posted by Pyro

bfesser, are you sure there is one for toluene? I just had a look under purification of common lab. reagents, there is no procedure for toluene. However there is one for benzene, which is exactly the same.

Yes, I'm sure. It's essentially the same procedure.<a href="http://library.sciencemadness.org/library/books/vogel_practical_ochem_3.pdf#page=195">

Quote:

16. Toluene. Toluene, free from sulphur compounds, may be purchased. If required dry, it is treated with anhydrous calcium chloride, followed by sodium wire (see under Benzene, 15).
   Commercial toluene contains methyl thiophenes (thiotolenes), b.p. 112–113°, which cannot be removed by distillation. It may be purified by shaking or stirring several times with about 10 per cent, of its volume of concentrated sulphuric acid as in the case of benzene (15), but, owing to the greater ease of sulphonation of toluene itself, some temperature control is necessary. The simplest procedure consists in mechanically stirring one litre of technical toluene with 100 ml. of concentrated sulphuric acid for about 30 minutes; the temperature is not allowed to rise above 30° by occasional cooling. The lower acid layer is removed, and the process is repeated twice. The subsequent working up is carried out as described under Benzene.
   Pure toluene has b.p. 110·5°/760 mm. <img src="../scipics/_pdf.png" />

</a>

Baffled - 24-1-2014 at 06:58

I did the synthesis from Vogel more than a year ago for college, but instead of salting out the pTsOH we decided to neutralize any remaining sulfuric acid with calcium carbonate and calcium hydroxide (seperately) after the x min of reflux mentioned in Vogel. After letting the neutralized reaction mixture stand for a few days in between labs we heated the reaction mix for a bit, filtered it and concentrated the filtrate to a small volume after which it was placed in an oven at 80° C to crystallize. I don't remember the specific yeild of calcium-p-toluenesulfonate (I must add identification of the product was difficult since our nmr quenched, and mp's were too high, we did get IR-spectra but couldn't find any references) but I do know we beat Vogel by a long shot. Both calcium carbonate and hydroxide worked pretty well too for sulfonating p-xylene and bromobenzene. http://library.sciencemadness.org/library/books/Weygand_and_Hilgetag_Preparative_Organic_Chemistry.pdf
We got a 70% yeild by refluxing 30 grams of toluene and 40 grams of sulfuric acid for 90 minutes at 110° C, adding 150 ml of dH₂O after the mix had cooled and adding to that 60g of calcium carbonate or 46g of calcium hydroxide.

[Edited on 24-1-2014 by Baffled]

[Edited on 24-1-2014 by Baffled]

Pyro - 24-1-2014 at 07:03

sorry about that, not only did I not find it, I posted the exact page that has the procedure!
now I'm mad at myself :mad:

Paddywhacker - 25-1-2014 at 23:03

Quote: Originally posted by Baffled

I did the synthesis from Vogel more than a year ago for college, but instead of salting out the pTsOH we decided to neutralize any remaining sulfuric acid with calcium carbonate and calcium hydroxide (seperately) after the x min of reflux mentioned in Vogel. After letting the neutralized reaction mixture stand for a few days in between labs we heated the reaction mix for a bit, filtered it and concentrated the filtrate to a small volume after which it was placed in an oven at 80° C to crystallize. I don't remember the specific yeild of calcium-p-toluenesulfonate (I must add identification of the product was difficult since our nmr quenched, and mp's were too high, we did get IR-spectra but couldn't find any references) but I do know we beat Vogel by a long shot. Both calcium carbonate and hydroxide worked pretty well too for sulfonating p-xylene and bromobenzene. http://library.sciencemadness.org/library/books/Weygand_and_Hilgetag_Preparative_Organic_Chemistry.pdf
We got a 70% yeild by refluxing 30 grams of toluene and 40 grams of sulfuric acid for 90 minutes at 110° C, adding 150 ml of dH₂O after the mix had cooled and adding to that 60g of calcium carbonate or 46g of calcium hydroxide.

With those ratios you neutralized *all* of the acids, including the TsOH.

How did you know that the calcium p-toluene sulfonate would not also precipitate with the CaSO4?

Generally, people want the free TsOH acid, and preparing it from the calcium salt would not be a shortcut.

On the other hand, if you determined the amount of free H2SO4 in your reaction mixture and just added the stociometric amount of calcium carbonate you might be on to a winning protocol, but I cannot see an easy way of making that determination.

Baffled - 26-1-2014 at 00:14

I found the solubility of the calcium salts in an old book full of solubility tables translated from Russian. I don't have it at hand right now.

Could one prepare the free acid by titration with HCl (or even CO₂?) and extraction the acid with a nonpolar solvent?

[Edited on 26-1-2014 by Baffled]

Paddywhacker - 26-1-2014 at 14:42

Quote: Originally posted by Baffled

...
Could one prepare the free acid by titration with HCl (or even CO₂?) and extraction the acid with a nonpolar solvent?

Then you'd be back to the same problem from before you neutralised it. The problem being how to produce pure crystalline TsOH. The standard way is to crystallize from concentrated HCl, which is awkward enough that anybody would want an improved method.

Baffled - 27-1-2014 at 05:55

Jes, sure. However, removing the remaining acids from the crystallized product would be a lot easier when there is no free sulfuric acid (there might be some sulfate, but that's minimal) left after workup. Just dry in vacuüm.
On an industrial scale sulfuric acid could be easily regenerated from the calcium sulfate residue, but that might not be of our concern.

HeYBrO - 27-1-2014 at 17:51

could someone suggest a method of disposal for the left of sulphuric acid and the waste washes? I don't really want to flush thiophenes down the drain if you get what i mean. is it possible to neutralise them or something?

[Edited on 28-1-2014 by HeYBrO]

Metacelsus - 27-1-2014 at 18:31

You could neutralize the acid with sodium bicarbonate. As for the thiophenes, I'm not too sure; you could look into burning them in a controlled manner (be careful!).

HeYBrO - 14-3-2014 at 17:13

Just a question (its really noobish, my apologies) But if i were to scale down this procedure by 10 times, obviously the water coming off is going to be much less, so how much of my 10 ml dean stark should i fill with toluene. I know that the toluene must be in excess and the amount of water will be probably under a ml ( according to the vocal fifth edition procedure ) I presume about 9 ml toluene in the dean stark from the begging would work to drip over and begin water collection? Or do i have it wrong?

Screen Shot 2014-03-15 at 12.09.52 pm.png - 159kB

[Edited on 15-3-2014 by HeYBrO]