Sciencemadness Discussion Board

The short questions thread (2)

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Sedit - 29-10-2009 at 17:07

So if I understand correctly then assuming the temperature of the flame is 2(were working figuratively here) then the real temperature is to the 4th power of two or 16? That is an insane amount of loss:o. Given that the radiant energy could be reflected back on a target wouldn't we be able to focus the energy back on the target and increase the temperature exponentially(theoreticaly).

gsd - 29-10-2009 at 17:59

Sorry for the confusion. I should have said: "radiation losses are PROPORTIONAL to the 4th power of absolute temperature of the flame"
The proportionality constant being Stephan-Boltzman constant. (=5.67*10^-8) J/(sec*M^2*K^-4))
Quote: Originally posted by Sedit  
So if I understand correctly then assuming the temperature of the flame is 2(were working figuratively here) then the real temperature is to the 4th power of two or 16? That is an insane amount of loss:o.


NO. The real temperature of the flame is its adiabatic temperature. The S-B law gives you amount of energy readiated by the hot object at a given temperature.
Quote: Originally posted by Sedit  

Given that the radiant energy could be reflected back on a target wouldn't we be able to focus the energy back on the target and increase the temperature exponentially(theoreticaly).


Yes. But again the maximum temperature you can achieve this way is the Adiabatic Flame Temperature.

gsd

Sedit - 29-10-2009 at 18:13

I got it now I think, thanks for the effort Gsd.

UnintentionalChaos - 29-10-2009 at 18:39

Quote: Originally posted by Sedit  
Its not really a factor of needing to graduate or anything of that nature its just that I used something simular as This to melt lead a few days ago for the reduction of Pottasium nitrate and it got me to thinking of the range of temperatures that could be achived using different fuels in a simular device. It worked very well and burned for 45 minutes or so on a single charge which kind of impressed me to say the least.

I was under the impression that there is an upper limit to the temperature at which a specific fuel can burn at assuming complete combustion and is independent of flame size.


Yes, there is a limit, but like I said, heat transfer lowers the temperature of smaller flames while lack of oxygen influx limits that of larger flames. There is a way to calculate a theoretical flame temperature (adiabatic flame temperature), but you need heat capacities over a very wide range of temperatures for the gases produced and the resulting number is not without flaws.

GLC - Kovats

Paddywhacker - 29-10-2009 at 20:05

The Kovats retention index is a useful way of systematizing GLC stationary phases. There is a useful list of RI here:- http://www.flavornet.org/flavornet.html

But this site also lists something called Ethyl Ester Retention Indexes

What are these? I would have thought they would be, analogously to the Kovats, relative to fatty acid ethyl esters, so that, on any given stationary phase, ethyl acetate = 200, ethyl propionate = 300, ethyl n-butyrate = 400, etc, but, from the values given in this site, that is not the case.

Edit:-
On close examination I think the Ethyl Ester RI data are bogus. Most compounds are represented twice with two sets of radically different data. My best guess is that these RI data ware computed from the Kovats data and that the deviations from the correct values are rounding errors. The duplicate entries must be human error.

[Edited on 30-10-2009 by Paddywhacker]

bbartlog - 29-10-2009 at 20:13

It's my understanding that if I have sodium bisulfate in aqueous solution, I can then (assuming it's concentrated enough) add ethanol and disproportionate it such that I end up with sodium sulfate in the aqueous phase and sulfuric acid in the ethanol.

Will the same thing work with ammonium bisulfate?

1281371269 - 1-11-2009 at 09:56

Would melting Pb in Sulphuric acid purify it, or would it react with the acid? Is there a simple way to purify it?
It's a pretty small bit, about 50g max.

Da_Boss - 1-11-2009 at 09:57

Does anyone know how to synthesize calcium hypophosphite from red phosphorous or phosphoric acid?

Sedit - 1-11-2009 at 10:37

Mossy how would you plan on keeping the Sulfuric acid from vaporizing when you melted the lead in the first place? I think worried about if it would react is a non issue here because it won't work the way you want it to anyway.

If you want to "clean" the lead I think you would be better just melting it down, Stirring and skimming the slag that forms on the top leaving a purer lead at the bottom. If its an alloy of lead such as solder is then your going to have to go a different way to seperate the alloy if you can at all.

1281371269 - 1-11-2009 at 11:28

Lead has a lower melting point than sulphuric acid's boiling point, by about 10 degrees. As the piece is small enough to fit in a test tube and be covered over with acid, it could be kept molten for at least a few minutes before the acid all boiled away.

Sedit - 1-11-2009 at 13:36

10 degrees is not much of anything. Odds are before you have a chance to melt all of the lead there will be a ten degree difference and your H2SO4 will start to boil away.

What exactly are you attempting to do with the lead?

1281371269 - 1-11-2009 at 14:31

There's no usage in mind, not currently anyway. I just wanted to purify it and wondered if this was viable.

Sedit - 1-11-2009 at 16:28

Sorry if I wasn't clear on what I ment. What the usage as in what do you plan on clearing out of the lead. Dirt, oxidation, or other metals alloyed with the Pb?

1281371269 - 2-11-2009 at 10:28

Well anything I can really - this would deal with other metals, but actually that's probably a minimal amount of the contamination anyway.

Sedit - 2-11-2009 at 11:31

In that case I would really think your best bet is to just melt it and skim any slag that forms off the surface until no more slag forms and the lead starts taking on a blue/blackish oxide coating. Let it cool and scape all the edges of the block off and you would be left with more then likely pure enough lead for most reasons.

Just my two cents on it maybe someone would be able to give you a better way to clean it. I have done this many times, Just the past few days matter of fact on some down right dirty lead covered in dirt and organic material and god knows what. Just stirring the molten Lead brings up all the crap and I would just periodicly scoop this out until it stoped comming out leaving me with lead I used to reduce some nitrate.

DJF90 - 2-11-2009 at 12:46

Can lead not be purified electrolytically like copper? Shouldnt be too hard to set up if its possible, and you should get some pretty pure lead out of it.

UnintentionalChaos - 2-11-2009 at 13:05

Just throw the lead in a crucible, melt, and pour the clean metal out from under the slag layer into water.

It's not very expensive to just buy pure material either...and 50g is maybe 40 cents worth of 99.9% material.

JohnWW - 2-11-2009 at 13:44

If you had a sufficiently large amount of Pb, and zone-refining heating equipment, you may be able to do it by zone refining, in which a narrow zone of Pb is kept molten in a long steel (or other impervious material) cylinder by heating coils, and slowly moving the heating coils along its length, so that highly pure Pb crystallizes behind the molten zone, and the impurities are swept along in the molten zone. However, this would work only if the metallic impurities formed eutectic mixtures with Pb with melting-points below that of pure Pb.

[Edited on 3-11-09 by JohnWW]

UnintentionalChaos - 2-11-2009 at 16:20

Quote: Originally posted by JohnWW  
If you had a sufficiently large amount of Pb, and zone-refining heating equipment, you may be able to do it by zone refining, in which a narrow zone of Pb is kept molten in a long steel (or other impervious material) cylinder by heating coils, and slowly moving the heating coils along its length, so that highly pure Pb crystallizes behind the molten zone, and the impurities are swept along in the molten zone.

[Edited on 2-11-09 by JohnWW]


That has got to be the most useless recommendation I've ever heard. Seriously?

12AX7 - 2-11-2009 at 21:13

Well, it might be economical for extracting values. Gold and silver form eutectics with lead, which would melt and be carried along the melt zone. Lead is valuable to precious metals recovery because it's cheap and sticks to everything of value (the gold and silver).

Tim

Paddywhacker - 2-11-2009 at 22:03

Quote: Originally posted by UnintentionalChaos  
Quote: Originally posted by JohnWW  
If you had a sufficiently large amount of Pb, and zone-refining heating equipment, you may be able to do it by zone refining, in which a narrow zone of Pb is kept molten in a long steel (or other impervious material) cylinder by heating coils, and slowly moving the heating coils along its length, so that highly pure Pb crystallizes behind the molten zone, and the impurities are swept along in the molten zone.

[Edited on 2-11-09 by JohnWW]


That has got to be the most useless recommendation I've ever heard. Seriously?

You could purify it one atom at a time. Use an electron beam to sputter atoms off a lead surface and ionise them. Then accelerate them in an electric field and use a magnetic field to deflect them according to their isotopic mass. Collect the different isotopic masses in various target bins.

You could probably modify an old mass spectrometer to do the job.

Run the machine for, well, rather a long time, and you eventually have isotopically pure lead that you can use for lead-lights, solder, or whatever.

merrlin - 3-11-2009 at 00:42

Quote: Originally posted by DJF90  
Can lead not be purified electrolytically like copper? Shouldnt be too hard to set up if its possible, and you should get some pretty pure lead out of it.


A keyword search of "Fundamental Aspects of Electrometallurgy" found this reference:

Dendritic electrocrystalisation of lead from lead nitrate solution. Surf. Technol. 1985; 26:177-83

I imagine that it would be a high purity, low volume process.

mr.crow - 3-11-2009 at 07:36

Quote: Originally posted by Paddywhacker  
Quote: Originally posted by UnintentionalChaos  
Quote: Originally posted by JohnWW  
If you had a sufficiently large amount of Pb, and zone-refining heating equipment, you may be able to do it by zone refining, in which a narrow zone of Pb is kept molten in a long steel (or other impervious material) cylinder by heating coils, and slowly moving the heating coils along its length, so that highly pure Pb crystallizes behind the molten zone, and the impurities are swept along in the molten zone.

[Edited on 2-11-09 by JohnWW]


That has got to be the most useless recommendation I've ever heard. Seriously?

You could purify it one atom at a time. Use an electron beam to sputter atoms off a lead surface and ionise them. Then accelerate them in an electric field and use a magnetic field to deflect them according to their isotopic mass. Collect the different isotopic masses in various target bins.

You could probably modify an old mass spectrometer to do the job.

Run the machine for, well, rather a long time, and you eventually have isotopically pure lead that you can use for lead-lights, solder, or whatever.


For solder, lol

Sounds like the Calutrons used to enrich the uranium for Little Boy. Don't let any governments find out!

12AX7 - 3-11-2009 at 08:46

Isotopically pure lead is rather valuable. The isotope descended from polonium, I think, decays enough to be a concern in flip-chips, where solder is bonded directly to the die. A single alpha decay can rip through a lot of silicon.

Tim

1281371269 - 3-11-2009 at 12:56

I think I'll go with melting it down, though it's barely enough even to do that. Point of interest - I got it from the inside of one of those used and glued back together bullets that you can buy in military museums.

Thanks for all the help.

DJF90 - 3-11-2009 at 15:02

Lead is very easy to get hold of. As you're a UK member as I seem to recall, theres a seller on ebay offering 1kg of lead milling media, for approx £3, with approx £4 for postage.

chemoleo - 4-11-2009 at 15:21

What are the best reaction conditions to achieve full condensation of a primary amine and benzaldehyde?

This of course forms a Schiff's base - but somehow the yields suck (20%, if that) if this done in aq medium, at pH 8 and RT.

[Edited on 4-11-2009 by chemoleo]

entropy51 - 4-11-2009 at 16:00

Are you trying to isolate the Schiff base or do you want to reduce to the secondary amine?

chemoleo - 4-11-2009 at 18:14

For now complete conversion to the imine is what is desired. Reduction afterwards maybe useful (Na cyanoborohydride) - but obviously is of little point if the initial yield of imine formation is so bad! Does anyone have protocols for the quantitative conversion of aldehyde to imine?

entropy51 - 4-11-2009 at 18:41

If you want to reduce, there's no need to isolate the imine. Shaking the amine and benzaldehyde in ethanol with catalyst and hydrogen will give the secondary amine.

chemoleo - 4-11-2009 at 18:48

Quote:
Shaking the amine and benzaldehyde in ethanol with catalyst and hydrogen will give the secondary amine

Catalyst?
If it serves to push the reaction to the right (reducing the initial imine to amine) then it may be useful.
Otherwise - again - what are the optimal conditions to condense benzaldehyde with a primary imine?
Many thanks.

Sedit - 4-11-2009 at 21:53

If your imine is of simple nature such as that formed from ammonia then hydrolysis back to the aldahyde in this case will be a problem. The reaction itself forms H2O and the equlibrium lies far to the left leaving you with your starting aldahyde however more stable forms of imine can be done with substituted amines such as methyl or ethylamine where they are more H2O tolerant and the shiffbase is stable.

My suggestion on increasing yeilds is to form the imine in dry alcohol in the presence of a mol sieve to absorb the generated H2O and push the equalibrium to the right although as I understand it without H2O at all the reaction will not proceed. Look on Rhodiums archives for 10,000 files or so on reductive amination reactions and the mechanics of such because if theres anything that them archives are good for this is it. The entire data base pretty much centers around amination of ketones and what not to yeild the amines.

Im quite sure there are others here that may either prove me wrong or be able to explain the whole concept much more gracefully then myself but personaly I say add an alcoholic NH3 or MeNH2 solution to your aldahyde to form the imine and if the results still appear to low start considering mol sieves and drying agents to shift the equilibrium to the right more.

Nicodem - 5-11-2009 at 01:24

If you use sodium cyanoborohydride or sodium triacetoxyborohydride (STAB), then you do not need to isolate any imine. NaBH3CN does not reduce benzaldehydes while STAB reduces them only very slowly compared to its speed at reducing iminium ions. NaBH3CN reductive aminations of aldehydes and ketones are actually performed in acidic H2O/MeOH media. You can not use the usual hydrogenation with Pd-C catalysis because much of the so formed benzylamine gets debenzylated, though it is actually possible to selectively stop the reduction at benzylamine stage if you slightly poison the catalyst with traces of thiophene. But with primary amines the formation of N,N-dibenzyl products can also be a problem even with a poisoned catalyst (it depends on the amine you will be using and its excess). With NaBH3CN or STAB this should not be a problem even with a stoichiometric ratio of reactants.

Imines of benzaldehydes are otherwise easy to form. You can either let the mixture stir over the night in the presence of MgSO4 in nonpolar solvents such as toluene, 1,2-dichloroethane, ethers, etc., then filter and rotavap. This generally gives a fairly pure imine. The other method is to use a catalytic amount of tosylic acid (or an amine salt) and remove water from the mixture of amine and aldehyde with water immiscible solvents such as hexane, benzene, toluene, etc by using a Dean-Stark trap (for benzaldehyde you don't even need any catalyst; see Klute's post on his preparation of the Schiff base from PhCHO and PhCH2CH2NH2). For hard to make imines (from ketones) generally SnCl4 is used as acid catalyst and water scavenger (but this is overkill for aldehydes which easily form imines and would probably just result in tar formation).

chemoleo - 5-11-2009 at 16:46

Thanks - so water is the problem - just what I feared :( Will report back if it can be improved.

Another question :)

Can someone suggest simple ways for getting from R-Phe-Br (para) to R-Phe-CH2-Br? I could think of doing a Grignard with CH2O and subsequent bromination, but surely there are better ways? R could be anything from Phe, NO2, CH3 etc - realising though that the nature of the para substituent R may have effects on the reaction that someone may hopefully propose :)

S.C. Wack - 5-11-2009 at 16:56

N-Methylbenzimine. Under a well-ventilated hood, a 250-mL Erlenmeyer flask is equipped with a magnetic stirrer and charged with an aqueous solution (40% w/w, 100 mL) of methylamine. Freshly distilled benzaldehyde (26.5 g, 0.25 mol) is added to the stirred solution of methylamine at room temperature. A mildly exothermic reaction occurs resulting in a milky white emulsion. The Erlenmeyer flask is stoppered and the mixture is stirred overnight (15 hr). The milky emulsion is transferred to a separatory funnel, diethyl ether (200 mL) is added, and the organic phase is separated and dried over potassium carbonate. The solids are removed by filtration, washed with ether (50 mL), and the combined filtrate and washings are concentrated on a rotary evaporator. The residue is distilled through a 10-cm Vigreux column using a water aspirator to give 28.3 g (95%) of imine as a water white liquid (bp 99-100°C/25 mm).

http://www.orgsyn.org/orgsyn/RxnTypes/section.asp?section=10...

[Edited on 6-11-2009 by S.C. Wack]

Molecular sieves

Bronstein - 11-11-2009 at 13:47

I read in one place that you must stir a liquid with molecular sieves in it very gently so that the sieves will not get destroyed. Is there any truth in this? Have seen another experimental procedure where they state to stir such a liquid vigorously.

(The drying should be more effective if you can stir it fast, so they sieves are suspended in the liquid instead of just lying on the bottom?)

marksev1 - 16-11-2009 at 03:19

I have a very stupid question, i used the search but didn't find any useful info so i'm asking it here....What is the advantage of having 29/32 joints compared to say a size smaller...i don't mean in terms that the 29/32 is the most used...but are there any advantages in speed or effectivity of distillation, i would think that only maybe very slighty faster if i'd use bigger joints...

User - 16-11-2009 at 04:11

If iam right..
Distilling for example one litre, it would take longer for the smaller setup to spit it out at the same purity (considering this equal to speed).

gsd - 16-11-2009 at 06:03

Molecular sieves
@ Bronstein

IMHO the Molecular sieves should be gently handled as they are brittle in nature. To ensure intimate solid liquid contact with such solids, packed bed, trickle bed, fluid bed or spouted bed contacting techniques are used instead of vigorous stirring.

gsd

bbartlog - 17-11-2009 at 13:46

Quote:
are there any advantages in speed or effectivity of distillation, i would think that only maybe very slighty faster if i'd use bigger joints...


Considering only the joint size in isolation, it would be unlikely to have much effect; the gas velocity through the joint would have to be pretty high before backpressure and/or gas compression and expansion effects were measurable.
Of course, the size of the joint may affect the size of the column that you can put downstream, and that will definitely matter. Look at the size of the various glassware components available in different joint sizes and you'll see what I mean. I think in choosing joint size you should look first at the overall size of the glassware you want (considering price and all) and then settle on one common joint size that lets you buy things in that general range.

Sedit - 20-11-2009 at 08:21

I have heard of Sulphur being used in place of lithium in birch reductions and it got me woundering about the structure of ammonia/sulphur complex that forms. I know Lithium forms Li(NH3)x structures but how exactly would sulphur bond and do you think it would be possible to form this complex simular to how Lithium bronze complexes are formed in a non polar solvent with anhydrous Ammonia feed into it.

It seems interesting because I can not find any information on this in the slightest only a few veuge references to it being used in S/NH3 systems.

schwabb - 21-11-2009 at 05:52

Hi guys!

i recently got a load of old chemicals and some of them are perculiarly packaged. My 'aniline' is a liquid, so i assume it is in solution of something, any ideas what? I think it has the words 'tonc vapour' written beneath the word aniline. It is in a brown bottle and the liquid looks opaque/ does not let light through.

what im really wondering about is my supposed thionyl chloride. It is packaged in a little glass bottle inside one of those large plastic jar things, specifically for the purpose (both recepticles are BDH manufactured)... BUT!! inside the plastic container, surrounding the glass bottle is a small amount of liquid. I was wondering if there was some common way of packaging like this, like how thiosulphate is used to package bromine, or if you guys had any idea what might be going on.

The bottles are really old, with the outer label barely legible (there is a newer one stuck on top) and the inner one missing completely, i dont think its a spillage as there is no noticable odour, although it presumably could have degraded with general humidity. I havnt opened the bottle of SOCL2 yet, and i would rather not at the moment.

Any ideas would be much appreciated.

DJF90 - 21-11-2009 at 06:02

Aniline is supposed to be liquid. It is likely dark in colour because thats something aniline does, but if you distill it it will come over colourless once more. It will however darken again upon standing for long periods I beleive.

schwabb - 21-11-2009 at 06:12

hah, well that was a stupid mistake! im not sure why but i had just go it into my head that aniline was solid, i think i must have been confusing it with something else. Thanks for that.. Fancy hazarding a guess as to what is going on with my thionyl chloride?

Sedit - 22-11-2009 at 08:45

I have a 90% DMSO 10% H2O mixture I wish to dry but as I understand partial pressure is needed in order to prevent decomposition of the DMSO and not having a vacuum means short path distillation is out of the question. However I have considered a couple of ideas that I wanted to get your take on which would be the better method.

Given the large difference in BP from the H2O and DMSO I thought about just distilling the water from the mixture and allowing it to cool in a dessicator. This sounds like it would be very easy but I am unsure how easy the DMSO decomposes to and what it decomposes to.

The other way I was thinking was to azeotropically dry it using EtOH or IpOH. Adding an effective amount of the given alcohol to the solution to aid in the removal of the water. Given the low BP of the H2O/-OH azeotrope this will keep the temperature pretty low as I dry it and should yeild pretty dry DMSO.

Anyone with first hand experiance here I would sure like to hear some suggestions.

sonogashira - 22-11-2009 at 09:31

From Armarego - 'Purification of laboratory chemicals'

Dimethyl sulfoxide (DMSO) 167-68-51 M 78.1, m 18.0-18S0, b 75.6-75.8O/12mm,
190°/760mm, d 1.100, n 1.479. Colourless, odourless, very hygroscopic liquid, synthesised from
dimethyl sulfide. The main impurity is water, with a trace of dimethyl sulfone. The Karl-Fischer test is
applicable. It is dried with Linde types 4A or 13X molecular sieves, by prolonged contact and passage through a
column of the material, then distd under reduced pressure. Other drying agents include CaH2, CaO, BaO and
CaS04. It can also be fractionally crystd by partial freezing. More extensive purification is achieved by
standing overnight with freshly heated and cooled chromatographic grade alumina. It is then refluxed for 4h over
CaO, dried over CaH2, and then fractionally distd at low pressure. For efficiency of desiccants in drying
dimethyl sulfoxide see Burfield and Smithers [J Org Chem 43 3966 1978; Sat0 et al. J Chem SOC, Dalton
Trans 1949 1986].
Rapid purification: Stand over freshly activated alumina, BaO or CaS04 overnight. Filter and distil over
CaH2 under reduced pressure (- 12 mm Hg). Store over 4A molecular sieves.

Sedit - 22-11-2009 at 14:06

The issues with all of those is that they require a final distillation using a vacuum and where as I admit that would be the best way to go I just do not have the means to perform that.

However I mainly needed to know in order to create NaDMSO and Im starting to think that I may possibly be able to mix DMSO and NaOH and slowly distill the H2O off as it slowly forms and push the equilibrium to the right more simular to the alkoxide formation in I think it was Klutes thread.

This could possibly work better for me then trying to dry it first as I would accomplish two steps in one process.

sonogashira - 22-11-2009 at 14:29

It happens I was reading of these anions recently and dimethyl sulfone anion appears to be preferred if one is to oxidize/reduce after the alkylation reaction... but I don't know if that is of interest?

Sedit - 22-11-2009 at 15:00

It was going to undergo pyrolysis to yeild the terminal alkene. Its basicly to extent the carbon chain.

Point being the Dimsyl ions are useful for a few different reactions and are a simple Super base but the synthesis of it is not as trivial as it first appears and one normaly needs hydrides or alkali amides to effect the deprotanation. I have reference stating low concentrations of NaDMSO being make from NaOH and DMSO but I feel this is due to it forming equilibrium with the formed H2O. Remove the water and the reaction should be pushed to the right yeilding NaDMSO.

aonomus - 22-11-2009 at 15:09

With a sump pump, 1/2hp-1hp, will I achieve any useful vacuum with an aspirator?

Nicodem - 23-11-2009 at 08:19

Quote: Originally posted by Sedit  
I have reference stating low concentrations of NaDMSO being make from NaOH and DMSO but I feel this is due to it forming equilibrium with the formed H2O. Remove the water and the reaction should be pushed to the right yeilding NaDMSO.

In DMSO the dimsyl ion is approximately 5000 times more basic than hydroxide. Therefore if you have a 1M solution of NaOH it will contain less than 50 ppm water (and the equivalent of dimsyl ions which would be less than 0.003 mol/L). Obviously you can not remove this water by using any type of azeotropic distillation. Perhaps by heating NaOH/DMSO over CaH2 or something like that. Or just buy NaH.

Sedit - 23-11-2009 at 08:42

I have a reference though that states above 150C the Dimsyl ion becomes a primary componate and it would seem that at those temperatures the H2O should distill from the mixture correct?
Quote:
Acid/base reactions of DMSO and hydroxide ion, which give ris
to dimsyl anions 26, may be prominent at 150°C. The
basicity of hydroxide increases dramatically with temperature.

http://smartech.gatech.edu/bitstream/1853/2796/1/tps-140.pdf

It would appear that even though the Dimsyl ion is 5000x more basic at room temperature then perhaps thats not so true at higher temperatures. As long as the reaction progressed enough to allow any form of efficient water removal then the equilibrium should Proceed to the right.

I have found a way using electrolysis of a soluble potassium salt to safely generate the dimsyl ion so there is no real need for the hydride anymore.

Nicodem - 23-11-2009 at 11:08

Quote: Originally posted by Sedit  
I have a reference though that states above 150C the Dimsyl ion becomes a primary componate and it would seem that at those temperatures the H2O should distill from the mixture correct?


It says "prominent", not "primary". It only appears to me that you misunderstood the claims in that paper. Why not checking the original paper they cite? I do not have access to it, but you can always ask for it:
http://www.informaworld.com/smpp/content~db=all~content=a762...
Quote:
It would appear that even though the Dimsyl ion is 5000x more basic at room temperature then perhaps thats not so true at higher temperatures. As long as the reaction progressed enough to allow any form of efficient water removal then the equilibrium should Proceed to the right.

Even if at higher temperatures the hydroxide would be only 1000 times less basic than dimsyl, this would not make much of a difference. You would still not be able to remove the water with toluene/water azeotrope or any such similar thing. Read on the theory of azeotropic distillation to understand why not.


This remained unanswered:
Quote: Originally posted by Sedit  
I have heard of Sulphur being used in place of lithium in birch reductions and it got me woundering about the structure of ammonia/sulphur complex that forms. I know Lithium forms Li(NH3)x structures but how exactly would sulphur bond and do you think it would be possible to form this complex simular to how Lithium bronze complexes are formed in a non polar solvent with anhydrous Ammonia feed into it.

It seems interesting because I can not find any information on this in the slightest only a few veuge references to it being used in S/NH3 systems.

If you would read on the mechanism of the Birch reduction you would understand why sulfur is nonsense. Maybe you misread Sr for S. It would be easier to spot the origin of the misinformation if you would use references. Sulfur is not even a metal and therefore has no free electrons in its crystal stucture. No electrons no ammonia solvated electrons no reduction.

Sedit - 23-11-2009 at 11:18

My reference for the Sulfur in ammonia to produce solvated electrons was lost when my computer crashes or else I would not have even needed to ask the question here because the paper stated the structure that formed solvating the electrons. It was a paper documenting the use of solvated electrons in a battery that employed liquid ammonia. I will see what I can do about recovering that file from those lost on my hard drive for you.

I will check in on the original reference stated in the paper as well. I must have over looked that citation thank you. Azeotropic distillation though means the use of a co-solvent to aid in the water removal process. That would not be needed here since the difference in boiling points between DMSO and H2O is so great. You are still more then likely correct in that it would not be pushed all the way to the right but I am still intested in what sort of concentrations can be obtained and the mechanics behind it all.

Panache - 23-11-2009 at 21:02

i just had my first piranha accident, it was very disconcerting. That stuff readily burns nice round holes in your skin.
I was having an excellent day, so well that i thought 'i know i'll attend to the bucket of glassware i have accumulating for piranha washing'
I recommend never using this solution at the end of a day.
Anyway i was making up as normal, 500ml of conc sulphuric, 300ml of 30%h202, except it was 50%, well i had a geyser out of the two litre beaker, i was out of the way quick enough except for one drop on my arm, however my beautiful cedar bench top is fucked and i have a shit of a clean up.
Question, i am surprised the reaction was this vigorous, normally i get nothing really when i add the two, is the higher concentration h202 the culprit?
Whats best for neutralising this solution?

bbartlog - 24-11-2009 at 13:20

I have no idea what the proximate cause of your eruption would be; I could imagine both metal impurities in the H2SO4, or else the possibility that once the amount of H20 in the mix is small enough the H2SO4 just tears the peroxide apart.
But I also have to wonder what's on your glassware that you need a cleaning agent this aggressive. Do you actually have a bucket of stuff that wouldn't get clean with acetone, dilute HCl, 10% ammonia, or plain old detergent and water? Or do you just like to employ maximum force cleaning?

Panache - 24-11-2009 at 16:23

Quote: Originally posted by bbartlog  
I have no idea what the proximate cause of your eruption would be; I could imagine both metal impurities in the H2SO4, or else the possibility that once the amount of H20 in the mix is small enough the H2SO4 just tears the peroxide apart.
But I also have to wonder what's on your glassware that you need a cleaning agent this aggressive. Do you actually have a bucket of stuff that wouldn't get clean with acetone, dilute HCl, 10% ammonia, or plain old detergent and water? Or do you just like to employ maximum force cleaning?


Of course metal impurities, i had just used the sulphuric to rinse through the coil of a condenser brown stained from rusty water, it had been ineffective however obviously had solvated enough metal ions to cause the h2o2 to decompose.

I rarely use pirahna, maybe twice annually, i accumulate bits and pieces that i have otherwise been unable to get clean any other way. Most commonly complex pieces with difficult to get to bits within them. I had probably a dozen small bits and pieces yesterday, and seeing as i'm talking about it i will bore everyone with what happened after the spill. So i spent probably around and hour and a half cleaning everything in the proximity of the ruined bench and floor, then a further 30 minutes cleaning everything i used to clean the area.
Then i made up some fresh pirahna, immersed the several pieces of glass within it and went to move the 2litre beaker across to another bench more out of the way AND THEN DROPPED THE BEAKER moving it across benches.
So the cleaning began again, plus the entire reason for making the solution in the first place was in thousands of pieces across the lab floor.
I laughed, i can't remember ever dropping a beaker, let alone within circumstances such as these.
We should start a 'Science Madness' sitcom with accumulated funny lab stories.

Hamilton - 25-11-2009 at 19:53

hi,
i have got an easy one for you,

i want to make substantial quantities (well something like 50 grams) of white lead oxide form lead metal, what is the best, (safest) way to transform it?

thx

manimal - 25-11-2009 at 21:46

Quote: Originally posted by Hamilton  
hi,
i have got an easy one for you,

i want to make substantial quantities (well something like 50 grams) of white lead oxide form lead metal, what is the best, (safest) way to transform it?

thx


If you're looking to make lead(ii) acetate, a good way would be displacement of aqueous copper(ii) acetate by lead metal.

entropy51 - 28-11-2009 at 13:39

Quote:
Anyway i was making up as normal, 500ml of conc sulphuric, 300ml of 30%h202, except it was 50%, well i had a geyser out of the two litre beaker


Quote:
Then i made up some fresh pirahna, immersed the several pieces of glass within it and went to move the 2litre beaker across to another bench more out of the way AND THEN DROPPED THE BEAKER moving it across benches


You spilled pirhana twice in one day? You were making 800 mL? For cleaning glassware? What's the deadline on nominations for the 2009 Darwin Award?

sonogashira - 28-11-2009 at 14:23

Hehe - butter fingers!

But I have had a (less serious) accident also and would please like some help...

I was going to add MeBr to benzylamine in methanol solution (with potassium carbonate) to make the quaternary salt, but unfortunately I was working on two experiments at once (due to limited time) and added the benzylamine to acetophenone :(

I only experiment on a small (boiling tube) scale and I have no way to separate these by distillation, and I have only a small amount of acetophenone anyway, so I would like to make use of it and not dispose. I have looked (but failed) to see if this reaction can be done in acetophenone; nearest I could find was acetone solvent but many examples do not use a base which seems silly(?!)

So my question is, can I bubble MeBr through benzylamine in solution of actetophenone and expect quaternary salt formation... and what base can I use in acetophenone?
Thank you for any help.
It is a shame - I tried to save time but now have wasted it; but there is a lesson to be learned I suppose!

[Edited on 28-11-2009 by sonogashira]

[Edited on 28-11-2009 by sonogashira]

sonogashira - 29-11-2009 at 14:07

Well I tried on a small portion in the absence of any suggestions, and I think I have some product - but I will have to try and purify further.

The problem is that I did not use an additional base because I don't know what can work, so I have mono- and di-alkylated contaminants.

Hopefully someone can recommend a base to use because this test result was far from ideal :(

[Edited on 29-11-2009 by sonogashira]

Question on chloroacetyl chloride reactivity with hydroxyls

chemoleo - 5-12-2009 at 17:46

Chloroacetic acid ClCH2COOH react with primary and secondary hydroxyls in aqueous solution.
I was wondering what would happen with the use of Chloroacetyl chloride, ClCH2-COCl - would the acyl chloride attack the hydroxyls preferably? (obviously then under nonaqueous conditions)?

I.e.what if I wanted to convert ethylene glycol to ClCH2-COO-(CH2)2-COO-CH2-Cl? Would this work without formation of mixed products?

Nicodem - 6-12-2009 at 09:11

Quote: Originally posted by sonogashira  
I was going to add MeBr to benzylamine in methanol solution (with potassium carbonate) to make the quaternary salt, but unfortunately I was working on two experiments at once (due to limited time) and added the benzylamine to acetophenone :(

Recover your benzylamine by:
- add the acetophenone/benzylamine mixture to water
- add HCl untill acidic reaction
- wash with diethyl ether, dichloromethane or whatever suitable solvent you have (and recover acetophenone from the washes)
- basify with NaOH
- extract the precipitated benzyl amine oil with dichloromethane
- dry over Na2SO4
- evaporate.

Don't waste time and material with an acetophenone/benzylamine mixture as you will never be able to efficiently recover the trimethylbenzylammonium bromide.
Quote: Originally posted by chemoleo  
I.e.what if I wanted to convert ethylene glycol to ClCH2-COO-(CH2)2-COO-CH2-Cl? Would this work without formation of mixed products?

Yes, only the mono and diester of ethylene glycol can form by solvolysis of chloroacetyl chloride in ethylene glycol. Without pyridine the reaction might require a long time and/or heating, but otherwise no O-alkylation can occur, just O-acylations.

manimal - 8-12-2009 at 00:59

I haven't been able to find a good preparative procedure for o-toluenesulfonic acid.

sparkgap - 8-12-2009 at 04:19

Have you checked references for the preparation of saccharin? I know for a fact that o-toluenesulfonic acid is an intermediate in saccharin synthesis (and the reason the para isomer is widely available is that it's an unavoidable side product).

sparky (~_~)

[Edited on 8-12-2009 by sparkgap]

kclo4 - 25-12-2009 at 10:09

I've read a lot of peoples opinions on bomex, pyrex, kimax, etc but does anyone have experience with this new-ish pyrex vista?

User - 26-12-2009 at 06:42

Is this correct ?
5 mol/L H2SO4 = 81.194 %

Damn my math is so rusty that i am not sure.
I dont think its correct :S


*edit*
Damn i feel stupid.

It should be something like 26%
Still i cant confirm this.



[Edited on 26-12-2009 by User]

gsd - 26-12-2009 at 07:50

5 moles/L of H2SO4 = 5 X 98 = 490 gm/L of H2SO4 solution.

So the concentration is 49 % (W/V)

In order to know the concentration in weight / weight (w/w) basis, we need to calculate the specific gravity of the solution.

Here we assume that there is no volume change when H2O and H2SO4 are mixed. And the mixing is done isothermally at 30 Deg C. :)

Perry-Chapter 2-108 gives density of 100 % H2SO4 @30 Deg C as 1.8205

So volume of 490 gm of Acid is 490/1.8205 = 269.16 cc. Therefore volume of balance water in solution is 1000 - 269.16 = 730.84 cc. Also weight of that water is 730.84 gm.

So total weight of 1 lit of solution is 490 + 730.84 = 1220.84 gm.

So concentration of Acid = 100 X (490 / 1220.84) = 40.14 % (w/w)

Hope this is correct :)

gsd

per.y.ohlin - 26-12-2009 at 14:00

In case anyone wants a very precise calculation, I assumed there could be some volume change upon mixing.

I used the table at the bottom of this page as a source.

I used excel to find the data points for the relationship between P (W/W%) and M (molarity). I assumed the concentration would be linear between the two closest points:
(M,P)=(4.879,37.26) and (5.096,38.58).
P=slope*M+b
slope=(P2-P1)/(M2-M1)=(38.58-37.26)/(5.096-4.879)=6.083
b=P-slope*M=37.26-6.083*4.879=7.581043
We want to know the concentration for M=5.000, so:
P=6.083*5.000+7.581043=38.00

So its 38%, not 40%.



[Edited on 27-12-2009 by per.y.ohlin]

Attachment: sa.xls (115kB)
This file has been downloaded 528 times

sonogashira - 31-12-2009 at 09:53

Quote: Originally posted by Nicodem  
Quote: Originally posted by sonogashira  
I was going to add MeBr to benzylamine in methanol solution (with potassium carbonate) to make the quaternary salt, but unfortunately I was working on two experiments at once (due to limited time) and added the benzylamine to acetophenone :(

Recover your benzylamine by:
- add the acetophenone/benzylamine mixture to water
- add HCl untill acidic reaction
- wash with diethyl ether, dichloromethane or whatever suitable solvent you have (and recover acetophenone from the washes)
- basify with NaOH
- extract the precipitated benzyl amine oil with dichloromethane
- dry over Na2SO4
- evaporate.

Don't waste time and material with an acetophenone/benzylamine mixture as you will never be able to efficiently recover the trimethylbenzylammonium bromide.


Thank you! I have only just seen the reply! Can I ask you (or someone else!) about similar situation (theoretical this time - no more mistakes!):

If I had acetophenone contaminated with 10% benzaldehyde, could I remove the benzaldehyde by adding DCM and washing DCM-Acetophenone-Benzaldehyde with water? DCM is the only solvent I can get really :(

I can't find information if benzaldehyde forms an azeotrope but I suspect so (though it is a complete guess!). And I believe it is more soluble in water than acetophenone, but maybe DCM dissolves it well also? It is acetophenone I want to recover, just to be clear.

[Edited on 31-12-2009 by sonogashira]

not_important - 31-12-2009 at 10:01

Probably best to oxidise the benzaldehyde to benzoic acid, and then extract that from DCM-acetophenone with aqueous sodium bicarbonate or carbonate, followed by a plain water wash.

sonogashira - 31-12-2009 at 11:22

Thanks. It is the (expected) product ratio from the oxidation of styrene with H2O2 and palladium chloride, so I don't know if adding oxidizer will be very good in this case (probably should have said before!).

Maybe I could extract both into DCM, wash with water to remove catalyst and excess H2O2, then boil-off DCM... dilute with water and add an aqueous oxidant etc... but it seems rather long procedure.

Preferably I would like to separate the benzaldehyde from the DCM-Acetophenone-Benzaldehyde. I thought maybe bisulfite can work but I think acetophenone will react as well.

I gather you do not think it (benzaldehyde) can be removed just with water washing of the DCM solution? Or by heating above the bp of benzaldehyde?

[I also saw something about 'extractive distillation' which I couldn't really understand (it was a patent!), so maybe there is something I could add (to Acetophenone-Benzaldehyde; after removing DCM) to form a low-boiling azeotrope with benzaldehyde, and leave pure acetophenone behind?... If only!]
[Edited on 31-12-2009 by sonogashira]

[Edited on 31-12-2009 by sonogashira]

crazyboy - 31-12-2009 at 22:04

How much is 1 mole of 35% formaldehyde?

gsd - 31-12-2009 at 22:44


Formaldehyde - HCHO - M.W. = 30

35 % solution = 35 gm in 100 gm solution = 1.1667 moles in 100 gm solution

So (100/1.1667) = 85.714 gm of 35 % solution of formaldehyde will contain 1 mole.

gsd

not_important - 1-1-2010 at 00:10

Quote: Originally posted by sonogashira  
...
Preferably I would like to separate the benzaldehyde from the DCM-Acetophenone-Benzaldehyde. I thought maybe bisulfite can work but I think acetophenone will react as well.

I gather you do not think it (benzaldehyde) can be removed just with water washing of the DCM solution? Or by heating above the bp of benzaldehyde?...


Check the relative solubilities of the two products

acetophenone 5.5 g/L at 25°C, 12.2 g/L at 80°C
benzaldehyde 0.6 g/100 ml (20 °C)

so you're looking at a ratio on the order of 9:10, which means counter current extraction a best, and multiple stages most likely.

You may want want to get the catalyst out of solution first, or extract them away from the aqueous phase, before attempting any further workup. Matter of fact, doesn't this take a cosolvent or emulsifier to get good results?

Bisulfite will work, the aromatic ketones ArC=O... do not add bisulfite very well, even the methyl ketone acetophenone. You'll have to experiment a bit, basic idea is to wash the organic phase with nearly saturated NaHSO3 solution one or more times. Check each phase, the aqueous one after destroying the addition product and extracting the organics back into DCM; perhaps use TLC to evaluate the separation.

You'll end up with a little of the ketone coming along with the aldehyde, and aldehyde with the ketone. Use oxidation to clean up the ketone, a water-acetone solution of it and drip in KMnO4 solution persists for 10s of seconds, then add a ml or 2 of isopropanol to destroy the small excess of KMnO4. Filter off "MnO2", distill off acetone, add NaHCO3, extract ketone out with DCM.

Note that there will likely be unreacted styrene and some polymers of it. If you distill the acetophenone add a chaser to it before distilling, some non-reactive substance it mixes with and that boils much higher than the ketone. Kerosene that has had its low boiling components removed in a fractionation (just fractionate out the low boilers, don't distill the high boiling stuff you want) works well. The chaser allows most to all of the acetophenone to be fractionated out without running the still pot dry, and any styrene polymers will dissolve in the hot chaser - easier to remove from the flask.

sonogashira - 1-1-2010 at 03:03

Hey, thanks very much! You're right, there will be a small amount of acetone co-solvent (few ml), but it should be removed with the water wash. Sodium lauryl sulfate apparently converts 100% of styrene and with 100% acetophenone selectivity (zero benzaldehyde) - but it is claimed that it "forms an emulsion which is very difficult to separate."

[In fact, any good ideas to separate this and palladium chloride from the acetophenone product would be very useful also! This method may be even more useful if there is an easy way to separate them which the authors overlooked!]

Yes, acetophenone does tend to help polymerization - this is the reason I want it, in fact :D - so maybe I will need to distill. Probably I could use petroleum jelly (vaseline) as 'hot chaser' to save distilling kerosene...

Thanks for all the tips! :) I wanted to be sure I could purify the acetophenone before I waste money on buying the palladium chloride!

[Edit: and if the bisulfite won't remove all of the benzaldehyde maybe I will just go straight to oxidation to benzoic acid instead - I thought I would add this incase anyone can see a problem with the method I intend.]

I may even try distillation of calcium salts also and compare the two processes - but first to get a distillation set! :(



Second edit: according to US4433173 diethylene glycol is the solvent to use. It increases the volatility of the contaminants and one can then boil-off the acetophenone from the (non-azeotrope) mixture of the two - Diethylene glycol having relative volatility of 0.1 compared with acetophenone. I think this is the implication, but I will have to read it a few times more to make sure I understand correctly because of all the silly language they use for some ridiculous reason :mad:
[Edited on 2-1-2010 by sonogashira]

[Edited on 2-1-2010 by sonogashira]

Panache - 1-1-2010 at 20:19

Quote: Originally posted by sonogashira  
Sodium lauryl sulfate apparently converts 100% of styrene and with 100% acetophenone selectivity (zero benzaldehyde) - but it is claimed that it "forms an emulsion which is very difficult to separate."

[In fact, any good ideas to separate this and palladium chloride from the acetophenone product would be very useful also! This method may be even more useful if there is an easy way to separate them which the authors overlooked!]


[Edited on 2-1-2010 by sonogashira]


Time separates all emulsions, however if you are time poor my most reliable technique is the microwave, play with power levels and times, it can break an emulsion within seconds at times. If i have no joy with the microwave i use the freezer, freezes the water out, then just invert your flask onto a course filter and leave the organic solvent to drip out. This is all done in the freezer of course (not the microwaving though)

not_important - 1-1-2010 at 22:22

The bisulfite will remove the majority of the aldehyde, but a small amount will remain unless you push things with several extractions with aqueous bisulfite. However as the ketone is slightly water soluble a bit will be leached out into the aqueous layer with each extraction. As it is easier to handle traces of aldehyde in ketone as opposed to the reverse, it seems better to leave a bit of aldehyde behind and scrub it out as the acid.

sonogashira - 2-1-2010 at 02:45

Thank you Panache - unfortunately in this case the freezing point of acetophenone is 20 C (although when pure it is typically liquid at much lower temperature - unless there is impurities.) It will be worth to try and I do have lots of time, so I shall try (with a test-sample) to microwave, and also leave it standing.

not_important - I do not want to be a pain but there is one thing I don't understand: why use the bisulfite when I will need to use KMnO4 anyway? Could I not oxidize all of the benzaldehyde (10% benzaldehyde and 90% acetophenone) after removing the DCM? Maybe by adding water and acetone with KMnO4, as you suggested? (Maybe acid is needed too? - I have rarely used KMnO4 for organic oxidations... But I believe base and KMnO4 may turn acetophenone into benzoic acid?)

I want to prepare this (acetophenone) as solvent for some polymerization reactions, so I do not mind if this would cause extra expense with the purification step - I just want it as pure as possible. Is there some benefit in using bisulfite?

Sorry to be a bother, and thank you (both) for the kind help! ;)

[Edited on 2-1-2010 by sonogashira]

[Edited on 2-1-2010 by sonogashira]

not_important - 2-1-2010 at 09:47

Well, because you said that you wanted to recover both if you could :-)

If you're just cleaning up traces you can 'titrate' with the KMnO4, adding it drop by drop until the colour remains for some seconds; little chanch of chewing up much of the ketone with a dilute solution. The reaction rates of aldehydes and ketones to KMnO4 are generally much different, so if you don't just dump in gobs of the oxidiser you can readily tell when the aldehyde has been used up without attacking the ketone.

Or you can use one of the other redox reactions that aldehydes undergo readily while ketones do not, acidified dichromate - not too concentrated, Fehling's or Benedict's solution, or:
using Oxone http://www.organic-chemistry.org/abstracts/literature/565.sh...
using sodium perborate+acetic acid http://www.organic-chemistry.org/abstracts/literature/266.sh...

those last two might be useful if you just want to skip the separation and instead oxidise away all the aldehyde as the reagents are fairly cheap, low toxicity, and don't stain stuff.

sonogashira - 2-1-2010 at 10:10

Oh I see :) Maybe I implied that - but really I meant the exact opposite! I thought perhaps you knew of some obscure side reaction which I hadn't considered! Anyway, nevermind...

Thanks very much for the help. I will look into appropriate oxidizer to use... and hope for the stated yields!
Best wishes

[Edited on 2-1-2010 by sonogashira]

manimal - 13-1-2010 at 18:08

I am considering the potential effects of water on the chlorination of toluene in the presence of tin(IV) chloride. The catalyst is soluble in toluene, but would water interfere with the interaction of catalyst and substrate?

Nicodem - 15-1-2010 at 05:48

Quote: Originally posted by manimal  
I am considering the potential effects of water on the chlorination of toluene in the presence of tin(IV) chloride. The catalyst is soluble in toluene, but would water interfere with the interaction of catalyst and substrate?

Water will not interfere with the electrophilic chlorination itself, but it will decompose the catalyst by hydrolysing it (SnCl4 + H2O => SnO2 + HCl). So you do not want water in there when using hydrolysable Lewis acids as catalysts.

When using NCS, trichloroisocyanuric acid or other chloroimides instead of Cl2, you can use protic acids as catalysts, and in this case the presence of moisture does not interfere much or not at all (in fact the reaction can even be run with water in a biphasic system).

manimal - 17-1-2010 at 13:40

Quote: Originally posted by Nicodem  
Water will not interfere with the electrophilic chlorination itself, but it will decompose the catalyst by hydrolysing it (SnCl4 + H2O => SnO2 + HCl). So you do not want water in there when using hydrolysable Lewis acids as catalysts.

When using NCS, trichloroisocyanuric acid or other chloroimides instead of Cl2, you can use protic acids as catalysts, and in this case the presence of moisture does not interfere much or not at all (in fact the reaction can even be run with water in a biphasic system).


Thanks. What I had in mind to do, was to add the entire measure of chlorinator (TCCA) to toluene and catalyst, followed by a small (catalytic) quantity of HCl. This would introduce chlorine in a rate-limiting manner as C3Cl3N3O3 + 3 HCl + 3 C6H5CH3 --> C3N3H3O3 + 3 C6H4ClCH3 + 3 HCl. Thus, the initial Cl2 formed would have to react with the toluene first, releasing HCl which would be re-oxidized to Cl2 (in lieu of adding chlorine or oxidizer gradually over a period of time).

Nicodem - 19-1-2010 at 13:03

Quote: Originally posted by manimal  
Thanks. What I had in mind to do, was to add the entire measure of chlorinator (TCCA) to toluene and catalyst, followed by a small (catalytic) quantity of HCl. This would introduce chlorine in a rate-limiting manner as C3Cl3N3O3 + 3 HCl + 3 C6H5CH3 --> C3N3H3O3 + 3 C6H4ClCH3 + 3 HCl. Thus, the initial Cl2 formed would have to react with the toluene first, releasing HCl which would be re-oxidized to Cl2 (in lieu of adding chlorine or oxidizer gradually over a period of time).

What you propose does not sound particularly rational chemistry-wise. Why using TCCA as oxidant when you can chlorinate toluene directly and even considerably more efficiently by using it as the electrophile?

Just add a couple of mol% of tosylic acid monohydrate, AlCl3, FeCl3, SnCl4 or any other strong enough acid to a solution of TCCA in excess toluene and let stir overnight, filter off the cyanuric acid and fractionate the filtrate. Or just follow the literature example as done by Axt for chlorination of benzene in this thread. Don't forget that electrophilic halogenations of toluene are quite exothermic so never add too much acid catalyst and never do the first trial above about 10 mmol. TCCA in the presence of acid catalysts is more electrophilic than Cl2 which reacts with toluene quite sluggishly, so take this into account when planing anything on a larger scale.

Are you planing to separate the ortho- and para-chlorotoluenes and if so how?

manimal - 21-1-2010 at 17:46

Quote: Originally posted by Nicodem  
What you propose does not sound particularly rational chemistry-wise. Why using TCCA as oxidant when you can chlorinate toluene directly and even considerably more efficiently by using it as the electrophile?

Just add a couple of mol% of tosylic acid monohydrate, AlCl3, FeCl3, SnCl4 or any other strong enough acid to a solution of TCCA in excess toluene and let stir overnight, filter off the cyanuric acid and fractionate the filtrate. Or just follow the literature example as done by Axt for chlorination of benzene in this thread. Don't forget that electrophilic halogenations of toluene are quite exothermic so never add too much acid catalyst and never do the first trial above about 10 mmol. TCCA in the presence of acid catalysts is more electrophilic than Cl2 which reacts with toluene quite sluggishly, so take this into account when planing anything on a larger scale.

Are you planing to separate the ortho- and para-chlorotoluenes and if so how?


Actually, I was going by US3000975 (http://www.google.com/patents?id=V6tJAAAAEBAJ) on the chlorination of toluene. It specifies quite high selectivities of >95% and a preference for the o-isomer (75/25%) with chlorine gas in the presence of a tin chloride catalyst.

I considered the use of TCCA directly w/H2SO4, but was put off by the statement that it carries yields of only 66% (with preference for the p-isomer) according to TCCA - A Safe and Efficient Oxidant. I realize the patent might be garnishing its claims, and may indeed yield no higher than 66% either. I'm not particularly set on separating the isomers, but rather just to see how the chlorination goes.

[Edited on 22-1-2010 by manimal]

Nicodem - 22-1-2010 at 14:28

I'm quite sceptical about that patent. Chlorination of toluene with Cl2 almost invariably gives a near to statistical ratio of ortho vs. para chlorotoluene (which would be 67% : 33%). Sometimes the ratio is less, but I don't remember ever seing examples where it would be less then 50% ortho-chlorotulene unless electrophiles other than Cl2 were used or under some special conditions. A 75:25 ratio is out of the statistical distribution which means there must be some directional effect from the catalyst. If this would be the case then SnCl4 would be known for it and there should be some paper about it (and not just a claim in a patent). Obviously if you follow that patent you need to introduce dry Cl2 from outside reaction. Adding TCCA in the presence of SnCl4 would just cause direct chlorination from this reagent instead of Cl2 while adding HCl(aq) would quench SnCl4. In either case you would not get the desired ortho regioselectivity increase (provided that the patent claim is true).

The claim that under such biphasic conditions TCCA gives 66% para regioisomer is also a bit uncertain. If you check the method used to determine the para/ortho ratio in the original paper (JOC, 35, 719–722), they did so by measuring the refractive index. This property is highly sensitive for any contamination of other compounds (residues of toluene, polychlorotoluenes, etc.) and I would not trust that data with much accuracy. They did not use GC like they should. I would expect that under such conditions one gets a near to 1:1 mixture of ortho- vs. para-chlorotoluene.

manimal - 22-1-2010 at 19:08

Quote: Originally posted by Nicodem  
Chlorination of toluene with Cl2 almost invariably gives a near to statistical ratio of ortho vs. para chlorotoluene (which would be 67% : 33%).


That's interesting. I was perusing the old literature and it said that iodine is a ring-chlorination catalyst (presumably by forming iodine trichloride which acts as electrophile) which favors almost exclusively the p-isomer. However, many more recent patents have been filed for increasing p/o ratio with wacky catalysts like ferrocene, which barely achieve 50:50. Why would they go to such trouble if catalytic iodine can facilitate it much more easily? Perhaps the early lit. references were mistaken.

[Edited on 23-1-2010 by manimal]

Nicodem - 24-1-2010 at 11:02

Quote: Originally posted by manimal  
That's interesting. I was perusing the old literature and it said that iodine is a ring-chlorination catalyst (presumably by forming iodine trichloride which acts as electrophile) which favors almost exclusively the p-isomer. However, many more recent patents have been filed for increasing p/o ratio with wacky catalysts like ferrocene, which barely achieve 50:50. Why would they go to such trouble if catalytic iodine can facilitate it much more easily? Perhaps the early lit. references were mistaken.

That iodine catalyses the chlorination is no surprise as it is after all a Lewis acid, but I never heard about its influence on the regioselectivity of toluene chlorination. I would really like to read more about it. Could you please dig out the reference describing all this?

Sobrero - 25-1-2010 at 11:08

What is this piece of glassware used for? I can't think of any setup where this would come in handy.
(It makes a good bong though :D)

DSC09636b_herschaald.jpg - 26kB

12AX7 - 25-1-2010 at 11:38

The spheres might be for drying agents. Don't ask how you're supposed to put the stuff inside them..

Tim

hector2000 - 27-1-2010 at 06:07

what is name of this componet?
this componet is possible?if yes how we can make it?

mfcd00009677.gif - 847B

12AX7 - 27-1-2010 at 07:20

Acetone chlorohydrin.

Looks unstable, and certainly irritating. I don't recommend making it.

Tim

Sedit - 27-1-2010 at 08:14

@Sobrero

I dont know much about glassware but suck back traps perhaps?

Seems like something I would love to have right now to feed NH3 into Ether and use the Bulbs as suckback traps.

cadra - 27-1-2010 at 08:22

Hi Everyone, this is my first post. I didn't know whether to post in short questions or someplace else. I have what I hope are a couple of quick questions for someone with more experience than me.

I have a Schiff base that I would like to methylate. I know that I can do this via a number of methylating agents such as dimethyl sulfate, methyl tosylate, or MeBr. However, pure curiosity has lead me to wonder if I can also methylate a Schiff base via Eschweiler Clarke reaction. After methylation I would like to hydrolyze to hopefully provide a mono-N-methylated amine and the starting aldehyde.

My first question is: Does this seem possible? Has anyone heard of this (I can't seem to find any references in my searches).

My second question: If this is possible, I am fearful that if I utilize oxalic acid dihydrate and paraformaldehyde to effect the Eschweiler Clarke as discussed in this thread:

http://www.sciencemadness.org/talk/viewthread.php?tid=7263

that the water from the dihydrate will hydrolyze the Schiff base prior to methylation, causing dimethylation of the amine (which is not what I want).

Here's another thread that I found on monomethylation of amines where the method was presented briefly but I didn't see anyone trying this on a Schiff base such as could be made by condensation of benzaldehyde with amine:

http://www.sciencemadness.org/talk/viewthread.php?tid=8375#p...

Note that to perform the Eschweiler Clarke utilizing oxalic acid dihydate and paraformaldehyde in the above mentioned paper, tempuratures of 100-120C need to be used to decompose oxalic acid to formic acid.

Thank you very much in advance!
Cadra

User - 30-1-2010 at 10:32

?

Does anyone know a good method for capturing CO.
The reaction I am looking at produces certain amounts of it and yeah as we all know CO is a bitch.
So can anyone suggest an absorbent/reactant to destroy/react it?
Thnx

entropy51 - 30-1-2010 at 12:26

Quote:
Does anyone know a good method for capturing CO.
One of my lab manuals gives a recipe for a CO abosrber: Dissolve 200 gm cuprous chloride and 250 gm NH4Cl in 700 mL of water and add 1/3 its volume of conc. NH4OH.

The lab manual is Introduction to Organic Laboratory Techniques by Pavia, Lampman and Kriz.

Roscoe's Treatise of Chemistry confirms that CO is soluble in ammoniacal cuprous chloride.

When I use CO I keep a CO detector running on the bench to be sure the hood is capturing it all.


[Edited on 30-1-2010 by entropy51]

Xenoid - 30-1-2010 at 13:31

Quote: Originally posted by User  
Does anyone know a good method for capturing CO.


This has been discussed at length before, about 2 years ago, you'll need to search the forum.
Hopcalite was used at one time in respirators - 50%MnO2(I assume the active form) copper oxide (30%) cobaltic oxide (15%) and silver oxide (5%).

Or you could bubble the CO through blood :D

"Destroying carbon monoxide in a gas stream"

The thread is here : http://www.sciencemadness.org/talk/viewthread.php?tid=10394#...

[Edited on 30-1-2010 by Xenoid]

User - 3-2-2010 at 02:35

What is a reasonable price for a second hand heating mantle?

Electromantel 100ml max 450 deg C.
Price 250 euro's

Isopad 250 ml , not temp specified.
price 250 euro's


Its seems like a rip off to me..
I am having a hard time finding a good deal for these devices.
And I dont do ebay, often the shipping is more than the original price.

*edit*
I mailed the company that had these offers.
I said that i found the prices where just too high.
Really laughing my ass off now i got a reaction.
The new the proposed price is 200 euro's for a model that costed 360.
Its a model with build-in stirrer.
It's still too much money for my liking but maybe iam not rational.
Still funny though.


[Edited on 4-2-2010 by User]
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