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ammonium isocyanate
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I'm not sure about recrystalization from water or ethanol (in part it would depend upon the concentrations of the two chemicals), but I know that PEG
is insoluble in diethyl ether and straight chain hydrocarbons. Paracetamol is slightly soluble in diethyl ether, so if you have some lying around,
it's worth a try. Another method would be to dissolve both in ethanol, evaporate/boil it away, and wash it with cold water (PEG is much more soluble
in water than paracetamol).
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DJF90
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Perhaps you could check a paper or something... Perhaps a paper like the one included below 
Attachment: Paracetamol synthesis and extraction.pdf (144kB)
This file has been downloaded 4610 times
Attachment: Paracetamol solubility.pdf (98kB)
This file has been downloaded 34459 times
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ammonium isocyanate
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Can carbon dioxide bubbled through an hot aqueous solution of alkali long chain carboxylates (i.e. from biodiesel saponification) acidify the salts to
the corresponding carboxylic acids and alkali carbonates?
I ask because this would be useful for recycling the alkali hydroxides (which can be made more easily from alkali carbonates than alkali chlorides)
used in the splitting of triglycerides into free carboxylic acids.
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UnintentionalChaos
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Quote: Originally posted by ammonium isocyanate  | Can carbon dioxide bubbled through an hot aqueous solution of alkali long chain carboxylates (i.e. from biodiesel saponification) acidify the salts to
the corresponding carboxylic acids and alkali carbonates?
I ask because this would be useful for recycling the alkali hydroxides (which can be made more easily from alkali carbonates than alkali chlorides)
used in the splitting of triglycerides into free carboxylic acids. |
No. The pKa of the long chain fatty acids is lower than the pka1 of carbonic acid. If this were not the case, bars of soap would react with air to
free fatty acids and sodium bicarbonate over time.
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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ammonium isocyanate
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Ok thanks.
I couldn't find a source for the pKa of specific acids in peanut oil (the feedstock I'm using) other than the really big ones (oleic and linoleic).
Do you know where I could find a reference? (Although now that I think about it yeah its pretty obvious that CO2 couldn't degrade alkali carboxylate
salts.)
I suppose I could use nitric acid and decompose the resulting nitrate to an oxide and hydrolyze that to the hydroxide, but HNO3 is pretty precious
stuff so it would probably only be worthwhile with lithium salts and even then I'm not so sure.
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UnintentionalChaos
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The question is why? NaOH and KOH are incredibly cheap feedstocks.
I suspect that all the long chain fatty acids have largely the same pKa values since they all have a terminal carboxylic acid at the end of a lot of
methylene units. Whatever double bonds the acid may have are so far downchain as to have negligible effect on the acidity of that terminal group
Department of Redundancy Department - Now with paperwork!
'In organic synthesis, we call decomposition products "crap", however this is not a IUPAC approved nomenclature.' -Nicodem
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Nicodem
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What other fatty acids besides linoleic and oleic acids do you get from peanut oil anyway. There are also some palmitic acid and arachidic acid
triglycerides in peanut oil along with some minor ones. For example, arachidic acid has the same pKa like all fatty acids (4.78 in water), but the basicity of fatty acid salts depends extremely on many factors, because the anions aggregate to higher structures
depending on concentration, counterion, temperature, etc. So it makes no sense to rely on these pKa numbers. Just imagine that the apparent pKa of
fatty acids ordered in monolayers or bilayers can reach up to 8 or more! It suffices to think of fatty acids in their molecular form as acids of
similar strength as their shorter aliphatic carboxylic acids (pKa ~ 4.7), though they are rarely or ever in molecular form when in aqueous solutions.
I assume the easiest way to recycle alkali hydroxides from your waste (NaCl or KCl) is to again make them like they are often made from chlorides in
the first place, via electrolysis in a separated electrolytic cell (see Kirk-Othmer or other sources for references). You get the aqueous solution of
NaOH (or KOH) which takes a lot of energy to dry. You can also recycle HCl (and some energy as heat) as well if you lead the Cl2 and H2 in a burner.
Of course, unless you have an industrial production this would make no sense to get involved in such a infrastructural endeavour. NaOH and HCl are
just too cheap and you will never beat the economy of the industry with anything you build for your small scale work.
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not_important
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| Quote: Originally posted by ammonium isocyanate | | Can carbon dioxide bubbled through an hot aqueous solution of alkali long chain carboxylates (i.e. from biodiesel saponification) acidify the salts to
the corresponding carboxylic acids and alkali carbonates? |
Yes, under moderately high pressure. Ten to twenty bar of CO2 should convert the sodium salts mostly to the free acids and NaHCO3. You need to
maintain the pressure until the carboxylic acids and the NaHCO3 have been separated.
-----
However this won't help much, simple base transesterfication tends to be wasteful of raw materials; the use of a large excess of alcohol with the
attendant need for recovery is one example, the problem with FFA is another. Small scale batch mode operations make it even worse, "backyard
biodiesel" is rather dirty and inefficient.
Reactive distillation is likely much more efficient, as FFA cause it no problem and the aqueous glycerol stream is salt free.
[Edited on 20-7-2009 by not_important]
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ammonium isocyanate
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True, but LiOH is not. I don't plan on producing biodiesel for use as a fuel, but instead I plan to seperate out various fatty acids as an experiment
in itself, and for use as reagents in their own right. As such, I would be performing a saponification reaction instead of transesterification.
I don't have a vacuum distillation aparatus, and don't plan on buying one. Therefore, the easiest way for me to seperate out polyunsaturated fatty
acids from other fatty acids is to prepare their lithium salts and then dissolve them in acetone ,lithium salts of polyunsaturated fatty acids are
soluble in acetone, but others are not (I don't have the reference on hand, but I could dig it up if you want). Problem is, all the fatty acids must
be converted to lithium salts, and thus alot of waste lithium would be produced in this process that I would prefer to recycle or use for other
experiment (I suppose preparing lithium chlorates/perchlorates could be interesting).
My proposed method would be as follows:
-Split triglycerides into glycerin and sodium salts using NaOH
-Acidify the sodium salts using HCl
-Prepare lithium salts by adding Li2CO3
-Seperate out the polyunsaturated fats using acetone
-Acidify the lithium salts using HCl
This produces alot of waste NaCl and LiCl. If I could recycle LiOH, it would be reasonable cheap for me to use this, thus cutting out alot of steps
and drastically reducing the amount of waste produced. Obviously I could regenerate LiOH from LiCl, but I don't have a suitable cell with a
semi-permeable membrane and can't find one at a reasonable price (and also I really, really don't want to use the mercury amalgam method).
Additionally, what would I do with all the Cl2 generated?
The only other way I can think of to regenerate LiOH without much of an aparatus would be to react LiCl with Na2CO3, precipitating the slightly
soluble Li2CO3, then react that with Ca2NO3, precipitating CaCO3, and decompose the resulting LiNO3.
[Edited on 20-7-2009 by ammonium isocyanate]
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Paddywhacker
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You could use an anion-exchange resin. Regenerate in the OH- form with dilute NaOH. Then pass your Li salt solution through.
The trouble is that you will have to do it in small batches to avoid overloading the resin capacity, and you will end up with a dilute solution
requiring a lot of CO2-free evaporation.
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not_important
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Prolonged boiling of fats with Li2CO3 in water will saponify the fats, as escape of CO2 drives the hydrolysis reaction towards completion.
Alternatively accept the NaCl waste formed, how big of a scale are you doing this on? Use acetone or MEK or even toluene to extract the free fatty
acids after removing most of the water under reduced pressure. Evaporate the solvent, and do a preliminary separation by cooling the FFA to remove
saturated acids, then mono-unsaturated ones. Repeat on each fraction to improve separation.
A trick sometimes used was to take the solids rich in higher melting acids, drop them in a Buchner supported by a flask, put the thing in an oven of
just large box with a small heater, and slowly raise the temperature. The lower melting stuff 'sweats' out of the higher melting, although some of
that does come with the sweat, and drips into the flask.
Rather than converting all of the rough cut of unsaturated acids to the Li salt, dissolve the free acids in acetone and add LiOH in IPA slowly. The
more saturated acids should precipitate as their Li salts, when precipitate stops forming you've a solution that's mostly the polyunsaturated acids
with small amounts of their Li salts and some of the more saturated acids per the solubility of their Li salts - those most be low but non-zero.
The combination of preliminary removal of the more saturated acids by cooling, and then fractional precipitation of their Li salts, should noticeably
reduce the consumption of LiOH over simply converting all the FFA to Li salts.
At that point the use of CO2 under pressure as a way of recovering the lithium becomes attractive. Treat the acetone solution with 20 bar CO2, bleed
the solution off through a filter to leave behind the lithium as carbonate. Treatment of that with a suspension of Ca(OH)2 in water gives a solution
of mostly LiOH; evaporate in stainless steel 'flask' and extract the LiOH out with IPA to get the solution used to ppt the saturated fatty acids as Li
salts.
Main consumables are NaOH, HCl or H2SO4, Ca(OH)2, and CO2. Waste is NaCl or Na2SO4, CaCO3.
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ammonium isocyanate
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Thanks not_important for all the suggestions!
I only plan on using about a gallon of peanut oil, so it's not worth building an expensive aparatus. The method I plan on using is as follows:
1. Saponify the triglycerides with dirt-cheap NaOH.
2. Seperate out the glycerin and possibly purify it.
3. Acidify with HCl, H2SO4, or H3PO4 (haven't decided).
4. Remove H2O and extract the fatty acids with acetone.
5. Evaporate acetone.
6. Remove saturated fats by cooling in an ice bath.
7. Seperated out all but the most saturated fats in a dry ice/isopropanol bath.
8. Dissolve the remaining liquid fatty acids in acetone.
9. Add the LiOH in isopropanol until phenolphthalein indicator turns pink.
10. Remove the precipitate and react with iodomethane to produce the methyl ester.
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Agent MadHatter
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Anyone have trouble buying this chemical?
Has anyone ever had trouble buying phosphorus oxychloride before? Has someone had the DEA come knocking just because they bought it?
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JohnWW
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There are various possible alternatives to POCl3 for organic chlorinations e.g. to form acyl chlorides, namely: PCl5, PCl3, AsCl3, SiCl4, BCl3, SCl2,
S2Cl2, although these may require different reaction conditions. If you need to make an acyl chloride or similar just as an intermediate, the
corresponding bromides should also work.
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Nicodem
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Threads Merged
22-7-2009 at 23:23 |
manimal
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| Quote: Originally posted by ammonium isocyanate | | I'm not sure about recrystalization from water or ethanol (in part it would depend upon the concentrations of the two chemicals), but I know that PEG
is insoluble in diethyl ether and straight chain hydrocarbons. Paracetamol is slightly soluble in diethyl ether, so if you have some lying around,
it's worth a try. Another method would be to dissolve both in ethanol, evaporate/boil it away, and wash it with cold water (PEG is much more soluble
in water than paracetamol). |
I carried out an ethanol extraction with good results. However, I used a brass rod to stir the solution around, and the solution turned red. Obviously
a complex of some kind, caused by residual copper corrosion. Anyone familiar with that sort?
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bfesser
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I'm refluxing an aqueous mixture of sodium hydroxide and acetylsalicylic acid to prepare salicylic acid. Under these conditions, is there any
possibility of forming residual benzene (concerned with safety, but my bet is no):
~1.55g ASA
~1.37g NaOH
~25 mL solution
Am I driving it too hard, i.e. to sodium phenoxide?
[Edited on 7/23/09 by bfesser]
[Edited on 7/23/09 by bfesser]
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crazyboy
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I suggest acid hydrolysis it is far more effective and cleaner. From my notes:
10g ASA are added to a 1L Erlenmeyer flask. 700ml hot water are added and the ASA dissolves with stirring and gentle heating. A small amount of
concentrated hydrochloric acid is slowly added until a pH of 2 is reached. Several boiling stones are added to promote even boiling and the mixture is
heated to a light boil for one hour. The solution is poured into a 600ml beaker and SA precipitates as the solution cools.
The fluffy mass is filtered and washed with 200ml ice cold water. Yield: 5.15g.
As for your original question, no it is very unlikely that you will produce benzene.
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DJF90
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You're not going to get phenoxide from a salicylic acid derivative unless it decarboxylates. The concentration of your NaOH solution is less than 1M,
so there is little danger of that! I would probably use a 2M solution myself, i.e. 2g NaOH/ 25mls.
Acid hydrolysis is an equilibrium reaction so you would have to distil out a component to shift the equilirbium and make it tend towards a
quantitative yield. Personally in this case I would use base hydrolysis as neither component of the ester (Acetic acid and salicylic acid) are
particularly volatile at the temperatures used.
[Edited on 23-7-2009 by DJF90]
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bfesser
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crazyboy:
Thanks for the advice, but I have no hydrochloric acid, only sulfuric. And I don't want to waste any of that preparing hydrochloric. Could I
substitute the sulfuric in for the hydrochloric without appreciable side reactions?
DJF90:
That was the last of my NaOH, unfortunately. Thanks for the advice, though.
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crazyboy
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| Quote: Originally posted by bfesser | crazyboy:
Thanks for the advice, but I have no hydrochloric acid, only sulfuric. And I don't want to waste any of that preparing hydrochloric. Could I
substitute the sulfuric in for the hydrochloric without appreciable side reactions?
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Never tried it but I don't see why not.
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crazyboy
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A few random questions:
1. How interchangeable are potassium and sodium salts? Assuming you adjust the amount to the right molarity can NaOH be substituted for KOH and K2CO3
be used in place of Na2CO3?
2. Is xylene a suitable substitution for toluene as an organic solvent?
[Edited on 24-7-2009 by crazyboy]
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ammonium isocyanate
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Crazyboy,
1. It all depends on the intended purpose of the salts. If, say, all you are trying to do is neutralize an acid and don't care about what salt is
formed, then they are interchangable. However, sodium and potassium salts do exhibit different solubility characteristics, so if you are trying to
perform a double-displacement reaction, it may be effected (although most would probably still work). Overall, they are pretty much the same for most
purposes.
2. Usually. Xylene is usually a mix of isomers, which would cause problems as a reactant but probably not as a solvent. The main difference is that
the bp of xylene is about 30*C higher, which may be good or bad depending on the application.
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kclo4
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Also potassium hydroxide, and potassium carbonate are stronger bases then their corresponding sodium salts. That could be important in some
situations.
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User
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I was wondering, is there any reaction between sodiumbicarbonate x hydrate and etOH.
I am asking this because i would like to neutralize/clean up an alcoholic solution.
Then it could be distilled so any residue would be eliminated.
Anyone any ideas about this.
What a fine day for chemistry this is.
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manimal
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No, barcarbonate will not react with ethanol. Nor is it soluble to any appreciable extent.
| Quote: Originally posted by manimal | Since concentrated ammonia solution is so useful but rather scarce, I was thinking that a promising way to prepare it would be to heat intimately
mixed ammonium sulfate and calcium hydroxide in a metal can and pipe the fumes into water.
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Regarding this, I found this method to work, but it required strong heating, was slow, and gave variable yields of ammonia with unknown amounts of
water. I also tried pyrolysis of urea, which was a disaster. It expanded in volume and flowed into the recieving flask, and set up into a rock-hard
chunk that made it necessary to discard my 'flask' (actually a metal can).
The best results I got were from heating 10% ammonia under reflux and piping the fumes into water. I would recommend this for the purpose of preparing
concentrated ammonia.
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