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Jor
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[*] posted on 18-12-2007 at 10:55
acetylsalicylic acid


I was wondering , besides it's medical use, what are the main uses of this compound in chemistry. I'm not really 'skilled' at organic chemistry, but can it be used as a weak acetylating agent? And does it have any coordination chemistry in for example methanol with metals? Is it more reactive than benzoic acid, and therefore easier to form esters from?, etc,etc. I have no clue at all...
I need it for my final experiment at school.

[Edited on 18-12-2007 by Jor]
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[*] posted on 18-12-2007 at 16:41


This would not be a good acetylating agent. There are a lot of good textbooks on organic chemistry, Steitwiser (one of CRX' professors), Morrison & Boyd (good chemists), etc. Read and work problems.
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[*] posted on 18-12-2007 at 18:53


How would one convert it (orthohydroxybenzylic acid) into, say, phenylethylamine? First of all, I would think, protect the -OH group by acetylation with acetic anhydride; then reduce the -COOH to -OH with LiAlH4, then oxidize the alcohol to -CHO with dichromate, then react the aldehyde with KCN to obtain a cyanohydrin, then acid-hydrolyze that to an alpha-hydroxy acid which can be reduced with smething like LiAlH4 to an alcohol with two carbons in the side-chain, then react with with PCl3 or similar to form an alkyl chloride, then react that with NH3 to form the amine. The acetylated phenolic -OH would still have to be removed somehow. (Which means that benzoic acid rather than salicylic acid is easier to start with). An interesting exercise for a backyard chemist.

[Edited on 20-12-07 by JohnWW]
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[*] posted on 19-12-2007 at 00:28


Quote:
Originally posted by JohnWW
How would one convert it (orthohydrocxybenzylic acid) into, say, phenylethylamine? First of all, I would think, protect the -OH group by acetylation with acetic anhydride; then reduce the -COOH to -OH with LiAlH4 , then oxidize the alcohol to -CHO with dichromate, then react the aldehyde with KCN to give a cyanohydrin, then acid-hydrolyse that to an alpha-hydroxy acid which can be reduced with smething like LiAlH4 to an alcohol with two carbons in the side-chain, then react with with PCl3 or similat to form an alkyl chloride, then react that with NH3 to form the amine. The acetylated phenolic -OH would still have to be removed somehow. (Which means that benzylic acid rather tythan salicylic acid is easier to start with). An interesting exercise for a backyard chemist.

Yes interesting for a amateur chemist , but only for advanced hobbyist as it sounds like a rather dangerous operation.
This ofcourse is a nice synthesis , but is there anything more practical for a school experiment? i doubt the school has Lithium alumiumhydride and working with cyanides and mercury salts is forbidden here (altho they have it in stock, like a kilo of cyanide and like 4 mercury salts ,250g each). Also the school-lab is not well enough equipped for such organic experiments...
Ofcourse there are way better alternatives for acetylsalicylic acid, but i could make up the scenario 'what if for example no benzoic acid/[alkali]benzoate is available?' . And arent there any not too complicated reaction that favor aspirine as starting compound?
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[*] posted on 19-12-2007 at 00:31


I'm pretty sure it can be nitrated via mixed acids (sulfuric/nitric) to picric acid (2,4,6-Trinitrophenol).

Not much use for picric acid except as an explosive, though.




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[*] posted on 19-12-2007 at 01:06


Quote:
Originally posted by MagicJigPipe
I'm pretty sure it can be nitrated via mixed acids (sulfuric/nitric) to picric acid (2,4,6-Trinitrophenol).

Not much use for picric acid except as an explosive, though.

So what reaction would that be? And how can I show the presence of picric acid. If i remember right, picrid acid is only dangerous when dry, and I would, after the experiment, immediatly destroy it, by dissolving it and reducing it or if you guys know a better option to neutralize it.... After reduction with for example zinc I yield 2,4,6-aminophenol right? or picramic acid?

Im quite new to organics, but I'm learning...

[Edited on 19-12-2007 by Jor]
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[*] posted on 19-12-2007 at 01:58


Picric acid has a very strong yellow color. You certainly will see that color.



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[*] posted on 19-12-2007 at 04:18


Quote:
Originally posted by woelen
Picric acid has a very strong yellow color. You certainly will see that color.

yes, but how do I destroy the picric acid? using hydrogen is no option, as I won't have the time to setup a gas bubbler,etc and as far as I know you would also need a catalyst.
But using a reducing agent, wich is a common one so no boro or alumiumhydrides, will that work? Using zinc or iron , formic acid or any suitable reductor?

Could anyone give a setup?
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[*] posted on 19-12-2007 at 10:12


Quote:
Originally posted by Jor
I'm not really 'skilled' at organic chemistry, but can it be used as a weak acetylating agent?

Well, phenyl acetate is a mild acetylating reagent, but it can only acetylate good nucleophiles like amines and anilines (and only those that are not stericaly hindered). Going by this analogy you could say acetylsalicylic acid might also work given its only difference is in the ortho-COOH group, however that gruop is acidic (pKa 3.5) and the amines would deprotonate it rendering themselves non-nucleophilic, as well as causing other unfavorable electronic effects in the deprotonated acetylsalicylic acid. But using a good nucleophile yet a weak base like aniline, you might actually prepare acetanilide this way.
Try refluxing 1.1 equivalents of aniline with acetylsalicylic acid in acetonitrile and follow with TLC. Isolation of acetanilide and its recrystallization is also quite simple, so this would be a nice demonstration of the "unknown powers of aspirin" for a school project.




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[*] posted on 19-12-2007 at 11:49


Quote:
Originally posted by Nicodem
Quote:
Originally posted by Jor
I'm not really 'skilled' at organic chemistry, but can it be used as a weak acetylating agent?

Well, phenyl acetate is a mild acetylating reagent, but it can only acetylate good nucleophiles like amines and anilines (and only those that are not stericaly hindered). Going by this analogy you could say acetylsalicylic acid might also work given its only difference is in the ortho-COOH group, however that gruop is acidic (pKa 3.5) and the amines would deprotonate it rendering themselves non-nucleophilic, as well as causing other unfavorable electronic effects in the deprotonated acetylsalicylic acid. But using a good nucleophile yet a weak base like aniline, you might actually prepare acetanilide this way.
Try refluxing 1.1 equivalents of aniline with acetylsalicylic acid in acetonitrile and follow with TLC. Isolation of acetanilide and its recrystallization is also quite simple, so this would be a nice demonstration of the "unknown powers of aspirin" for a school project.


If the acetylsalicylic acid's acidity is a problem maybe i could prepare the potassium salt, by dissolving aspirine in a potassium solution in ethanol and then let it evaporate to dryness. I would take a slight excess of hydroxide , so all acid is neutralized. Any carbonate in the hydroxide will precitipate out of solution i think. After evaporation i would have a solid mass of potassium acetylsalicylicate and a very small amount of KOH. However I am not sure if the potassium salt dissolves in acetonitrile or other suitable solvents.

A problem, however, is acetonitrile i think. I don't think school has it. And it's also quite toxic and volatile. The school's fumehood sucks about 3000 liters per second, so I think that's not sufficient. i will check tommorrow if any acetonitrile is at hand, but if not, what other solvent would be usable? Methanol or ethanol should work right if the potassium salt is used to prevent any ester formation with the -COOH group. however does potassium salt dissolve? I don't know! What solvents could be used besides acetonitrile?
Also what are the solubilities of acetaniline in acetonitrile. Evaporating such a toxic solvent doesnt sound like fun. crystallisation would be better.
Damn , I don't like organics, go inorganic.
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[*] posted on 19-12-2007 at 13:40


You can not make the potassium salt by using KOH. That would cause partial hydrolysis. The best way would be by using K2CO3 in methanol. But anyway, it is quite pointless to make the salt as that would do little to solve the problem. Yes, it would enable the aniline to remain unprotonated and thus nucleophilic, but it also would cause total deprotonation of the acetylsalicylic acid, which should in theory dramatically reduce its reactivity at the acetyl carbonyl due to the negative charge destabilizing the transition states (and also reducing the electrophilicity due to the inductive effect). Not to mention the solubility problems in using the potassium salt. Trust me, the best thing is to leave it to the equilibrium and thermodynamics to settle these things out. Often the best thing to do is also the simplest. Aniline has pKa 4.2 and acetylsalicylic acid pKa 3.5. Though these values are not for acetonitrile, but for water, they should be indications enough to demonstrate you have more than plenty of unprotonated aniline and non-deprotonated acetylsalicylic acid, in addition to their ionic forms, in order for them to react completely.

Acetonitrile is one of the relatively innocuous solvents found in an organic lab. That’s one of the reasons why I advised it. I use it daily and don't use any protection for it. Don't get scared by its MSDS, rather compare it the other solvents you consider safer. The other reasons for using it are in that it is quite dielectric, aprotic and inert toward both reactants and product. Acetanilide is well soluble in acetonitrile. If the TLC shows the reaction went to competition (if anything forms at all!), then isolate by rotavaping the solvent, dissolving the residue in ethyl acetate, washing with sat. NaHCO3, 1M HCl, water, rotavaping again and recrystallizing the acetanilide (this can be done in water).
If you don't have acetonitrile (which is unimaginable for an organic or analytical lab) then you can try other solvents. It needs to have a bp>70°C, preferably polar and aprotic (and of course inert toward species involved). Protic solvents might do, but transesterification might complicate things (you can try isopropanol). You might also try a nonpolar, aprotic solvent like toluene which refluxes at higher temperatures and might actually do better.

Edit: I just remembered that one option would be to do it solventless. Just heat the aniline with acetylsalicylic acid in a small flask at an oil bath of about 100-120°C for half hour and analyse the melt with TLC.

[Edited on 19/12/2007 by Nicodem]




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[*] posted on 19-12-2007 at 15:27


Quote:
Originally posted by Nicodem
You can not make the potassium salt by using KOH. That would cause partial hydrolysis. The best way would be by using K2CO3 in methanol. But anyway, it is quite pointless to make the salt as that would do little to solve the problem. Yes, it would enable the aniline to remain unprotonated and thus nucleophilic, but it also would cause total deprotonation of the acetylsalicylic acid, which should in theory dramatically reduce its reactivity at the acetyl carbonyl due to the negative charge destabilizing the transition states (and also reducing the electrophilicity due to the inductive effect). Not to mention the solubility problems in using the potassium salt. Trust me, the best thing is to leave it to the equilibrium and thermodynamics to settle these things out. Often the best thing to do is also the simplest. Aniline has pKa 4.2 and acetylsalicylic acid pKa 3.5. Though these values are not for acetonitrile, but for water, they should be indications enough to demonstrate you have more than plenty of unprotonated aniline and non-deprotonated acetylsalicylic acid, in addition to their ionic forms, in order for them to react completely.

Acetonitrile is one of the relatively innocuous solvents found in an organic lab. That’s one of the reasons why I advised it. I use it daily and don't use any protection for it. Don't get scared by its MSDS, rather compare it the other solvents you consider safer. The other reasons for using it are in that it is quite dielectric, aprotic and inert toward both reactants and product. Acetanilide is well soluble in acetonitrile. If the TLC shows the reaction went to competition (if anything forms at all!), then isolate by rotavaping the solvent, dissolving the residue in ethyl acetate, washing with sat. NaHCO3, 1M HCl, water, rotavaping again and recrystallizing the acetanilide (this can be done in water).
If you don't have acetonitrile (which is unimaginable for an organic or analytical lab) then you can try other solvents. It needs to have a bp>70°C, preferably polar and aprotic (and of course inert toward species involved). Protic solvents might do, but transesterification might complicate things (you can try isopropanol). You might also try a nonpolar, aprotic solvent like toluene which refluxes at higher temperatures and might actually do better.

Edit: I just remembered that one option would be to do it solventless. Just heat the aniline with acetylsalicylic acid in a small flask at an oil bath of about 100-120°C for half hour and analyse the melt with TLC.

[Edited on 19/12/2007 by Nicodem]

Ok all sounds good and well do-able.
However I do not have any books or tables giving Rf values for acetanilide , acetylsalicylic acid, aniline in acetonitrile or toluene. I think I will use paper (cellulose) as the stationary phase. Silicia is also possibel tho.

and why would you want to wash with sodiumbicarbonate first and HCl after? i guess bicarbonate is to neutralize any remainign acid, but why the hydrochloric acid? Sorry, I don't have years of experience, but planning to go study Chemical Engineering, so my noobiness is not for long!

[Edited on 19-12-2007 by Jor]
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[*] posted on 20-12-2007 at 09:24


I assume you have no ability to do a literature search so I did it for you to check what kind of related stuff has been reported up to now. I found only one reference where the N-acetylating ability of acetylsalicylic acid was studied and even this was not preparative but for use in an analytical method. In case you don't have access to ACS journals, here is the relevant excerpt:
Quote:
Anal. Chem., 58 (1986) 821-824; DOI: 10.1021/ac00295a037

Determination of Aspirin:
Preparation of Calibration Graph. In a 50-mL round-bottomed
flask fitted with a reflux water condenser, 100 mg of accurately
weighed aspirin is combined with about 200 mg of 4-aminophenol
and 20 mL of ethanol and the mixture is swirled to effect dissolution
and refluxed gently for about 20 min. After cooling, the
contents are transferred to a 50-mL calibrated flask and diluted
to the mark with acetone. ….
(irrelevant description of spectroscopic analysis follows)

Apparently you can use protic solvents, so try refluxing in ethanol or isopropanol.

Another reference uses 6-carboxymethyl-acetylsalicylic acid to N-acetylate p-toluidine:
Quote:
J. Am. Chem. Soc., 77 (1955) 5092-5095; DOI: 10.1021/ja01624a042

Reaction of the Methyl 3-Acetoxyphthalates with p-Toluidine.
–A solution of 1.19 g. (0.005 mole) of the ester
and 0.85 g. of p-toluidine in 1 ml. of pyridine was allowed to
stand for one hour. Water (10 ml.) was then added and
the precipitated solid collected after chilling.
In the case of the low-melting half-ester IV, an exothermic
reaction ensued and there was obtained 0.62 g. (83%) of
aceto-p-toluide melting at 144-146', undepressed by admixture
of an authentic sample. In contrast, the high-melting
half-ester V did not produce an exothermic reaction and led
only to the recovery of 0.80 g. (94%) of unchanged p-toluidine
melting at 42-43°C.
(Compound IV is 6-carboxymethyl-acetylsalicylic acid, quite more electrophilic at acetyl carbonyl when compared to acetylsalicylic acid, yet similar in that it has an –COOH orto to the phenolic group. On the other hand, compound V which is methyl acetylsalicylate and thus similarly electrophilic, was inert at the same reaction condition. Yet consider that this was tested only at room temperature. It is quite possible that the heat of neutralization in the case of IV caused enough temperature rising to make the acetylation reaction proceed, while V had no such advantage.)

The acidic wash in the workup was proposed just to remove the unreacted aniline which might interfere with recrystallization. If the yield is as good as the above two references indicate, then you can actually simplify the workup by just quenching the reaction mixture with 1M NaOH and vacuum filtering off the crude product.

You do not need any Rf data when doing TLC. One just cuts a silica plated TLC wide enough for 4 samples. The first is your reaction mixture, the second and third is acetylsalicylic acid and aniline respectively and the third a sample of authentic acetanilide (I'm sure almost every lab has some on the shelves). You can also add a fifth sample of salicylic acid which is the other product. Use an petroleum ether / ethyl acetate mobile phase (3:1 should do, otherwise try other polarities). This would be in case you want to determine the proceeding without doubts, but in practice (we) lazy chemists just do a TLC immediately after mixing the reagents, before starting the reflux (thus obtaining a TLC of time zero which should show two spots) and monitoring changes with doing a TLC from time to time during the reaction (any new spots indicate that a reaction is proceeding and the disappearance of one of the starting materials indicates the end of the reaction). Expect a lot of tailing characteristic for carboxylic acids, but since salicylic acid wont travel far, this is not that problematic. Otherwise just find someone familiar with doing TLCs (you must have a teacher at the school, don't you?) or just do the reaction blindly. Or, if you want to impress your schoolmates, do a kinetic study, by following the reaction with HPLC at regular intervals (assuming you have a HPLC at your school).

[Edited on 20/12/2007 by Nicodem]




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[*] posted on 21-12-2007 at 15:50


Quote:
Originally posted by MagicJigPipe
Not much use for picric acid except as an explosive, though.

Actually, it's also used in the preparation of some preservative solutions for insect specimens, if you're the bug collecting type.




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[*] posted on 22-12-2007 at 09:02


Salicylic acid readily complexes Iron(III), forming an intensely purple complex. The color is strong enough that you could probably use it to measure relatively small iron concentrations. If you really like coordination chemistry, salicylic acid can be electroreduced to the aldehyde, from which you could make salen ligands, for example.

Re picric acid, it is sometimes used to produce crystalline salts with difficult compounds, to make them easier to characterize.
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[*] posted on 11-2-2008 at 06:00


Sorry for the bump.

im a bit late, but after all I'm still going to try Nicodem's experiment. I Will heat acetylsalicylic acid with aniline in an oil bath for 20 mins.

But I have a remaining question
-How do i seperate the acetanilide and acetylsalicylic acid?
If acetylsalicylic acid dissolves in aniline and acetanilide doesnt , then its not no problem. But I dont have the -disliterature to see if this is true. I cant test it. I have asked
school and they have aniline but no acetanilide anymore.

If they are both soluble I was thinking about adding a large excess of water to the reaction mixture (slight excess aniline over aspirin...). This will precipitate both aspirine and acetanilide out (if both are noth soluble, i can filter straight away.). I will then filter, followed by one of the following:
-Melting point measuring.
-Add water to the mixture of solids and heat; acetanilide is soluble in hot water. Filter and evaporate.
-Add moderately concentrated sodiumcarbonate solution. This will dissolve the acetylsalicylic acid, while the acetanide wont dissolve (1g per 190mL).
-I wont use TLC, because I just hate that method and I cant explain in my work how the process works.

What do you guys think is the best method? Are acetanilide and acetylsalicylic acid soluble in aniline or not?
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[*] posted on 11-2-2008 at 07:24


Take the reaction mix up in a solvent that is not miscible with water, ether would be traditional but you can pick others based on the solubilities of all reagents and products.

Extract that solution with Na2CO3 to remove unreacted ASA and salicylic acid byproduct. Wash with a bit of water or brine, then extract with moderately strong hydrochloric acid to remove unreacted aniline.

It's easier to do a good extraction from a liquid phase, rather than from a solid phase, thus the use of a solvent.

TLC is your friend, especially if you don't have a personal GC or HPLC plus money to burn. A 2D TLC can give you lots of information fairly quickly, a 1D is rather nice for tracking reaction progress when at least one reactant or product has a reasonably different R value in whatever development solvent you use.
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[*] posted on 11-2-2008 at 09:01


Sorry if I have misunderstood anything you said not_important, I am dutch, and my English is not awesome (In my opinion):

The setup would be as following:
-after 20 minutes of heating the mixture of aniline and acetylsalicylic acid (without solvent) , add enough diethyl ether to dissolve the entire reaction mixture.
Then add an equal volume of 10% sodium carbonate (anhydrous) solution to the solution in ether. Let stand for 5 minutes and pippete away the ether layer, and pour it into an erlenmeyer. Add a few mLs of distilled water to wash the ether. Pipette away the water layer after 1 minute. Then add a some (approx 1/4 the volume of ether) 20% hydrochloric acid. Let stand for some minutes. Pipette away ether layer, pour into a evaporating dish, and let evaporate.

Is this what you actually ment in your post not_important, or did I miss something?

One more question remains:
Wouldnt a small amount of acetanilide be protonated by the hydrochloric acid, on the nitrogen atom, and dissolve in the water layer?

I will finally prove the reaction product is acetanilide by means of melting points measuring.

PS: I know TLC comes at handy, but I dont really understand how it works. You know, I have quite some knowledge of chemistry now, but not due to school. Im just in high school (17 years old), and we certainly dont learn a lot from our books. Instead I learn by experimenting in my own lab, and reading through a lot of theory on the internet and reading the boards. This way I have learned a lot about inorganics, and also quite some organics, but no things like TLC, because they simply don't interest me too much. I will learn how that stuff works when I go to university! :)
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[*] posted on 11-2-2008 at 09:47


[Trying not to get too far off topic, but...]

Just as not_important said, TLC truly is your friend. It's extremely cheap in comparison to HPLC or GC, and it takes less time and is generally less of a headache. You don't have to worry about potentially mucking up a 300$ column by not removing certain components of your sample well enough. It can also be more intuitive and therefore fun, especially when working with colored products that you can actually visually observe being separated. Solvents are <em>much</em> cheaper.

TLC is very similar to paper chromatography, but instead of using paper as your stationary phase you will use a layer of silica or alumina on an 'inert' backing. The solvent is drawn up by capillary action and slow evaporation, and as it moves upward, it carries the compounds to be separated. The compounds are separated based on size, polarity, etc. You can either buy the plates pre-made or you can even make them yourself. And in my opinion, pulling capillary micropipettes for TLC spotting is a lot of fun.

I highly recommend picking up a book on basic organic lab techniques from a library or used bookstore. Even if you're not planning on going into organic, you're going to have to take it, so why not get started now?

Maybe you could start by reading the Wikipedia entry on TLC and then using search engines to find other introductory articles. Here's a few to get you started:

http://en.wikipedia.org/wiki/Thin_layer_chromatography
http://orgchem.colorado.edu/hndbksupport/TLC/TLC.html
http://www.wooster.edu/Chemistry/analytical/gc/default.html

Undoubtedly you'll want to find articles in Dutch as well...
Hope this helps. If not, delete it.

[edit]
http://nl.wikipedia.org/wiki/Dunnelaagchromatografie
Admittedly, this article doesn't even compare to the English entry. I'll post in the Chemistry Wiki Project to see if anyone can help translate this.

[edit]
Ok, I didn't get very far, but you might if you're so inclined. Here's the <a href='http://nl.wikipedia.org/wiki/Portaal:Scheikunde'>Dutch chemistry portal</a>, if you can find a Dutch Wiki Project for chemistry from this portal, you could consider requesting an update to the article there. You also might consider posting in the <a href='http://nl.wikipedia.org/w/index.php?title=Overleg:Dunnelaagchromatografie&action=edit'>discussion page</a> for the article itself on the off chance that someone might see it.

(I'm only about 90% positive those are the correct links.)

[Edited on 2/11/08 by bfesser]
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[*] posted on 11-2-2008 at 10:20


The workup I advised some ten posts above will do just fine even if you do the reaction solventless (just skip the rotavaping). Essentially it is what not_important also advised you to do. Avoid using diethyl ether since acetanilide is not particularly well soluble in it. Ethyl acetate is better suited for amides.

Alternatively, you can directly recrystallize the whole reaction mixture from water, skipping the extraction and washings entirely. However, this might not give you a pure product after a single recrystallization. It's more of an try and see method and not so full proof like the other one.

Acetanilide can not be protonated by aq. HCl. It is a too weak a base for that (pKa of protonated acetanilide is 0.61, which is barely more basic than H2O). But anyhow, it is not normal to do washings with 20% HCl. The usual solution for acidic washings with HCl is 1M or 3-5% or whatever is handy and ready on the shelves.

PS: You would generally use separatory funnels for operations involving extractions and washes - not pipetes and such stuff. Check some book about laboratory practice before trying this out. I'm sure it will help you.




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[*] posted on 11-2-2008 at 10:40


ok , thank you very much. I know what to do now :)
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[*] posted on 14-2-2008 at 03:43


I preformed the experiment today. Everything went just OK, I think at least.

I took 5 grams of acetylsalicylic acid and 4 mL of aniline.
The big problem was, the aniline was brown. I wonder how, because it has been in a the complete dark for 10 years , in a brown glass bottle. There were tow bottles, both aniline was brown, so I took the newest bottle.

The reaction mixture became a brown syrup-like substance, but then I started to heat at 110C-120C, and everthing became a normal liquid again. After heating for 20 minutes, the liquid was still brown/dark red. I added approx 40mL of ethylacetate, and let dissolve. Again, a darkred/brown liquid results. Then I add approx 40mL of satured sodiumbicarbonate solution, and shaked well for 1 minute, after no bubbles evolved anymore (alot of them evolved, so it might just be nog acetylsalicylic acid has reacted...). I seperated the two layers, and added 30mL of 1M HCl to the still remaining brown liquid (now approx 30-35mL). Shaked again for 1 minute and seperated. Now , In the fume hood at school, the remaining 30mL of brownred liquid is evaporizing.

Im really worried about the brown color. What is this? Has this influence on the experiment?

Joris
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[*] posted on 14-2-2008 at 06:04


Aniline and its derivatives are very easily oxidized by air (just like phenol and its derivatives). The brown matter is some polymeric crap, containing many aromatic rings, attached to each other through oxygen, nitrogen or carbon atoms. One cannot speak of a single species. I have phenol (90%), aniline, aminophenol and catechol, all of which are red or brown/red, due to slight oxidation.

Making colorless aniline is not difficult though. Take some activated charcoal (I use granules, 1mm size, available in aquarium stores, used for filtering the water) and put a spatula full of this in 5 ml of aniline in a well-stoppered bottle or test tube. Leave this overnight. Next day, the aniline is colorless. Decant the liquid from the solid particles.
This also works for aqueous solutions of aniline salts (e.g. sulfate or hydrochloride), these also can easily be made colorless.

The brown stuff most likely does not affect your experimental outcome very much. You, however, need a purification step to get rid of this stuff.




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[*] posted on 14-2-2008 at 06:49


So how should I seperate the acetanilide (if formed at all) from the 'polymeric crap' ?

The crystals that have formed are shiny, brown crystals. Shall I redissolve in ethylacetate, add some activated charcoal, stand overnight (closed system, so it wont evaporate), and filter off next day?

Joris
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[*] posted on 14-2-2008 at 06:54


I think that is a good method. Normally I would do the treatment with the carbon before performing the reaction, but given your situation, you could try it afterwards. I'm not sure though if acetanilide is not absorbed by activated carbon. Big molecules can be absorbed more easily.

[Edited on 14-2-08 by woelen]




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