bbartlog
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Purple chlorophenol?
After reading one of the threads on this board concerning the possibility of extracting phenol from aspirin (among other ideas), I decided to try it.
Among other more well-considered attempts involving heating with basic compounds, I also did two experiments with sodium hypochlorite. The first
involved involved adding a small amount (2g) of salicylic acid (already hydrolyzed) along with roughly 6g of 31% HCl to a small quantity (roughly 50g)
of 6% sodium hypochlorite bleach.
This was enough acid to acidify the bleach and cause the production of some chlorine fumes. I quickly dumped in some sodium bicarbonate and left the
area; on returning I saw that the solution was now a deep purple(!) and had a sweet, somewhat cloying smell.
I wasn't sure of the purity of the salicylic acid (in fact, I know that it had some CaCl2 and acetic acid mixed in with it) so I decided to repeat the
experiment in a more simple form.
This time I added 200mmol of aspirin to 400mmol of sodium hypochlorite bleach (this is 112 325mg tablets into 480 ml of 6% bleach). Having not thought
through the stoichiometry properly, I thought this would be enough NaOCl to keep things basic. Of course, each mole of aspirin ultimately can produce
three moles of acid (acetic, carbolic, carbonic) as it comes apart, so soon after the aspirin dissolved and hydrolyzed I once again had chlorine
coming out of my flask, only in somewhat larger quantities. I quickly set it outside on my back porch. On bringing it in, I once again had a deep
purple solution, which I neutralized with sodium bicarbonate.
Finally, I added 3mmol (2 tablets) of aspirin to 40ml of bleach. This appeared to hydrolyze, but did not react further, leading me to believe that
either acidic conditions or free chlorine are needed for this reaction to take place.
Anyway - what is this purple phenol compound? The principal reaction that comes to mind for these conditions is some sort of Hunsdiecker reaction of a
salicylate salt, which should give a chlorophenol. But chlorophenols, mono- or otherwise, aren't purple AFAIK. A second thought that comes to mind is
that some unhydrolyzed acetylsalicylic acid could undergo something resembling a Krapcho decarboxylation, which would produce chloroethane (explaining
the sweet smell) and CO2 along with phenol, which might oxidize to quinone. But I doubt there was any unhydrolyzed aspirin in the first trial (it had
been sitting in HCl all night). Do I just have to assume I've produced some hard-to-identify quinone?
[Edited on 10-11-2009 by bbartlog]
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not_important
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First a minor point. It would be better to say synthesizing or preparing phenol from aspirin, "extracting" has the implication that what is being
extract already exists in some mixture or complex material and just needs to be dissolved out.
I am rather doubtful that you'll get any decarboxylation of salicylic acid in aqueous solution. The equivalents with 3 or more OH will, I believe even
2 phenolic OH groups allows slow decarboxylation in a boiling solution.
Second, both carboxylic acids involved are strong acids than HOCl, I don't know and am too lazy to look up the relative Ka for the phenolic part of
salicylic acid vs HOCl. So at least twice as much HOCl will be form as aspirin in the mix, assuming there's little free NaOH in the bleach.
HOCl can decompose to give chlorine, or react with the salicylic acid, and the Cl2 is certain to so reach. The salicylic acid could be oxidised to
the p-quinone, which will decarboxylate readily as well as reacting with the Cl2 and OCl. Quinones can react with other stuff, including the
decarboxylated transitional state. The quinones could even fragment. In short, don't be surprised to get a horrible mess for a result, unless excess
bleach pushes it to a limited number of end products.
Lastly aspirin tablets are not pure acetylsalicylic acid, some of the compounding ingredients could react, too.
Aqueous acetic acid and sodium acetate are not really reactive towards Cl2 or HOCl, I'd not look there for sources of unexpected results. Krapcho
decarboxylation? LiCl in DMSO at 150 C on substrates with electron withdrawing groups on the beta carbon, does sound much like a phenol ester of
acetic acid in water near room temperature.
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bbartlog
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Quote: Originally posted by not_important  | | First a minor point. It would be better to say synthesizing or preparing phenol from aspirin, "extracting" has the implication that what is being
extract already exists in some mixture or complex material and just needs to be dissolved out. |
Good point. I'm using the word in a sloppy way but it has a more specific meaning in chemistry...
| Quote: | | I am rather doubtful that you'll get any decarboxylation of salicylic acid in aqueous solution. The equivalents with 3 or more OH will, I believe even
2 phenolic OH groups allows slow decarboxylation in a boiling solution. |
Well, this is a poor test of the idea. I suspect some decarboxylation occurred, as not all of the evolved gas appeared to be chlorine; it wouldn't
surprise me at all if one of the reactions was a poorly-yielding Hunsdiecker, but since I wasn't aiming for chlorophenol (and the whole result is a
mess in any case) this doesn't really help.
In general I expect you're right. One of the trials I did was to heat (to boiling) acetylsalicylic acid with an excess (5:1 mole ratio) of Ca(OH)2 and
water for an hour. Ca(OH)2 was chosen in part because calcium carbolate is not white, so there would have been an easy visual indication of any
progress towards decarboxylation - but the mixture stayed white throughout.
However, this just means that the calcium salicylate salt is pretty stable; it's likely that salicylic acid by itself would be a little easier to
decarboxylate. I will probably try heating some salicylic acid in a CaCl2 brine (which should stay liquid even past the boiling point of phenol) as
another experiment.
| Quote: | | The salicylic acid could be oxidised to the p-quinone, which will decarboxylate readily as well as reacting with the Cl2 and OCl. Quinones can react
with other stuff, including the decarboxylated transitional state. The quinones could even fragment. In short, don't be surprised to get a horrible
mess for a result |
Yeah. Looking at it now I see that the third trial I did has also reacted (just more slowly) to form a brownish solution. I also suspect that the
purple port-wine color I saw may actually be a result of brown and red reaction products mixed together, not some single compound that's purple. I
need to do some TLC if I'm going to analyze this stuff further rather than just dump it.
| Quote: | | Lastly aspirin tablets are not pure acetylsalicylic acid, some of the compounding ingredients could react, too. |
Yes, I've filtered out the binder material sometimes, though not in the last two trials I described. It actually doesn't compose much of the aspirin
by weight and seems to be fairly innocuous (calcium stearate maybe?) so I haven't worried much about it.
| Quote: | | Krapcho decarboxylation? LiCl in DMSO at 150 C on substrates with electron withdrawing groups on the beta carbon, does sound much like a phenol ester
of acetic acid in water near room temperature. |
Sure, sure. Not to mention that the first reaction that is likely to happen in any of these situations is the hydrolysis, after which there is no
ester.
Thanks for the comments!
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garage chemist
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Phenol from salicylic acid is more involved that just heating it.
Ullmanns Encyclopedia says:
When heated at or above its melting point salicylic acid decomposes into phenol and carbon dioxide. Under a carbon dioxide atmosphere at 230 °C the
main product is phenyl salicylate. At 250 °C, xanthone is formed in parallel with phenol. For the behavior of salicylic acid as a function of
temperature see [9].
[9] M. Wesolowski, Thermochim. Acta 31 (1979) 133 –146.
So, you need to heat exactly to 230°C and not higher.
Since salicylic acid already starts to sublime at 76°C, this has to be done in a vessel that does not allow the vapors (but the CO2) to escape, like
a flask stoppered with a narrow hose adapter and immersed up to the joint in the oil bath.
Then, your product is phenyl salicylate which you have to hydrolyze, e.g. by boiling with 3 equivalents of NaOH in aqueous suspension until
homogenous. An alcohol cosolvent may be necessary to speed up the reaction if it is found to proceed extremely slow.
Acidification will then release the phenol, which can be steamdistilled.
Now, the question is whether salicylic acid is also volatile with steam- if yes (and I would guess so, given its 76°C sublimation temperature), you
have to adjust the pH in such a way that salicylic acid exists as nonvolatile salicylate while phenol stays free. A suitable buffer can do this.
Sodium bicarbonate will only deprotonate the salicylic acid, not the phenol, so that steamdistillation can then be applied.
But boiling will cause any excess bicarbonate to decompose into carbonate and CO2, and carbonate is a strong enough base to deprotonate phenol.
So only add bicarbonate untill fizzing stops, not more, and steamdistill.
Direct extraction (e.g. with ether) may work too, but this will not remove colored and oxidised material.
Extract the steam distillate with ether, dry with e.g. CaCl2, evaporate the ether and, maybe, distill the phenol if you have enough crude product.
Small scale steamdistillation is as simple as distilling the aqueous mixture and adding water from time to time so that the water level in the boiling
flask doesn't sink too low.
Blowing steam into the flask is only convenient when working on a relatively large scale, and even then it's not absolutely necessary.
I'd really like to see someone succeed with a proper preparation of pure phenol from salicylic acid, using a well thought-out method.
Good luck, and please, tell us all about it if you succeed!
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Nicodem
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A home chemist friendly method of phenol synthesis starting from salicylic acid was recently described at the Hyperlab forum in a dedicated thread.
I'm not going to search for it, but those interested should not have troubles finding it.
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bbartlog
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Thanks for the information! Yes, I gathered from the some of the previous posts that pure pyrolysis was a little tricky (requiring reasonable
temperature control) and not likely to result in quantitative yields. If I do try it it will likely be using an oil bath and some sort of improvised
reflux that allows CO2 to escape while recondensing phenol and/or salicylate. But there are a couple of other things still on my list to try first.
The steam distillation of phenol sounds like a pain - I've done steam distillation and it just seems unreliable in terms of its ability to separate
different constituents (though your suggestions about proper acidification will be really helpful if I do try it). If I end up with mixed
Na(hydroxide+phenolate+salicylate) I'll probably dry it, acidify with HCl, filter out most of the salicylic acid (solubility is quite low in water),
then do some sort of solvent extraction. I haven't looked at the possibilities closely but I'd expect NaCl and phenol to be separable fairly easily
once the salicylic acid is (mostly) gone.
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bbartlog
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After poking around in Gmelin and a few other places it sounds like I probably have a mix of chloroquinones, chloro(hydro?)quinone and
dichloro(hydro?)quinone and maybe even chloranilic acid. Given the stoichiometry this implies that in acid solution the first reaction (whatever it
might be) leads to further reactions on the product molecule. The brownish color of the third trial (basic, but with NaOCl in excess) suggests that
that reaction stopped with the a less chlorinated chloroquinone, whereas the wine-purple (from brown+red) color of the earlier trials implies the
creation of even more chlorinated products. Makes sense that acid conditions would make the ring more vulnerable to halogenation, likewise I can sort
of imagine how the first chlorination would shift the electron distribution to make further chlorination easier . Once I put together a proper trap
for the chlorine I think I'll try the reaction a little differently and see if I can produce mostly chloranilic acid (or whatever my red moiety might
turn out to be).
This orgsyn article is interesting in context. The dichloroquinone is described as tan, but the author mentions a dark red mixture/residue which I
can only assume (given his setup) is some further-oxidized aromatic compound.
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not_important
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| Quote: | | it wouldn't surprise me at all if one of the reactions was a poorly-yielding Hunsdiecker |
It would surprise me, the Hunsdiecker-Borodin reaction is the reaction of metals salt of carboxylic acids with elemental halogens to give
halides. The metal salt is commonly silver, sometime mercury as HgO and the free carboxylic acid, rarely thallium(I)Pb(IV) and halides can be
used, certain organic bases will work in place of silver. The reaction is sensitive to water, getting the silver salts dry enough can be a problem.
It has a radical intermediate step, on some substrates this leads to other products than the normal substitution.
SFAIK the alkali and alkaline earth salts of carboxylic acids undergo thermal decarboxylation more readily than the free acids, the alkaline earth
salts (and salts of other metals in a +2 state or higher) tend to give the ketone (RCO2)2Ca => RCOR + CaCO3. Copper, especially as Cu(I)
salts, is often catalytic for the decarboxylation. But the reaction is just not going to happen in water.
Halogenation of quinones is often an addition-elimination reaction, not a substitution.
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bbartlog
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Sure, in this case the proposed halide is chlorophenol. R-COOH -> R-Cl where R is phenol. Someone here posted a reference that gave quite a list of
Hunsdiecker reactions, including some that used alkali metal salts. On the other hand, given what you're saying about water, it can't have been that.
Nonetheless something similar may have happened; I'm not seeing any way to explain the wine-dark color I have without some chlorination of the ring.
| Quote: | | the alkali and alkaline earth salts of carboxylic acids undergo thermal decarboxylation more readily than the free acids |
Good to know. Makes thermodynamic sense. I will have to try this tonight. Pyrolytic decomposition of salicylic acid in CaCl2 brine was tried
yesterday; I was able to vaporize the salicylic acid (so temperatures >200C were attained) and saw some nice crystals climbing up the side of my
flask, rather than the usual liquid ring. I still have to see how much phenol is in the mix (probably not much). I know there is some (from the smell)
but this tends to be deceiving; I also smelled phenol after extended boiling of a sodium salicylate and sodium hydroxide solution, but the yield for
that was about 2%. In any case, the salicylic acid floats atop the brine, though no doubt some dissolved as well; this defeats one of the purposes of
the setup, which was to disperse the acid to discourage the formation of xanthone as opposed to phenol.
| Quote: | | the alkaline earth salts (and salts of other metals in a +2 state or higher) tend to give the ketone (RCO2)2Ca => RCOR + CaCO3. Copper, especially
as Cu(I) salts, is often catalytic for the decarboxylation. But the reaction is just not going to happen in water. |
Yes, as for acetone from calcium acetate. In the case of salicylates self-condensation (RCOR - H2O) leads to xanthone. Since xanthone is not my aim I
will probably use sodium (hoping that the monosalicylate will result in less ketone production), and try to devise some way to disperse it in some
medium (for the same reason). Unfortunately CaCl2 brine won't work due to the metathesis reaction. Maybe glycerol, if sodium salicylate is soluble.
I'll try the addition of copper separately, though none of the salts I have are in the +1 oxidation state.
I'll also try dropwise addition of NaOCl to an excess of acetylsalicylic acid in solution. All previous trials have been done with at least a 2:1 mole
ratio of hypochlorite to salicylate, with rapid addition of the salicylate to the hypochlorite, which makes identifying possible intermediate products
hard. There is actually a fair bit of fizzing (clear, not yellow/green) very soon after addition of aspirin tablets to bleach (before the evolution of
chlorine and subsequent colored products), which makes me wonder whether the decarboxylation doesn't happen as one of the first reactions.
| Quote: | | A home chemist friendly method of phenol synthesis starting from salicylic acid was recently described at the Hyperlab forum in a dedicated thread.
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If I get tired of goofing around, or end up actually needing phenol more than practice with my chemistry, I'll be sure to look it up :-). But I've
never read the Hyperlab forum; is it worthwhile?
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benzylchloride1
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I have experimented with the pyrolysis of sodium salicyclate about a year ago. The sodium salicyclate was prepared by exactly neutralizing 20g of
salicylic acid with 9.2 g of sodium carbonate in an aqueous solution, the water was evaporated, leaving a solid, whic hwas then finely powdered and
placed in a distillation flask attached to a distillation apparatus. The flask was strongly heated with a propane torch, a slightly brownish liguid
distilled for a while. This would crystallize upon cooling and had a very strong odor of phenol. After a while, a red colored oil distilled. The crude
phenol was dissolved in toluene and extracted with sodium hydroxide solution. The sodium hydroxide extracts were acidified, sodium chloride was added
to salt out the phenol, and extracted with fresh toluene. After drying with magnesium sulfate and removing the solvent with a rotary evaporator, a
light brown colored oil remained which solidified when chilled and was a mixture of liquid and solid at room temperature. In all 7.8g of crude phenol
was obtained, the theoretical yield was 13.6 g and the percentage yield was 58%.
[Edited on 13-11-2009 by benzylchloride1]
Amateur NMR spectroscopist
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aonomus
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While this might seem out there, the iron(III) complex with salicylic acid is deep purple, and perhaps if you were using a bleach solution OTC (6%
NaOCl), it might have had slight iron contaminants, enough to give colour to the solution. Since its not ACS or higher grade NaOCl, I wouldn't be too
surprised if it had iron contaminants from an industrial process.
Perhaps you could try making a small amount of the iron(III) salicylic acid complex for comparison? There are a few procedures online used as a
qualitative test for salicylic acid.
Just a proposal of course, seems like a simpler theory compared with complex polymerization products giving a purple colour.
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bbartlog
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Interesting idea... thing is, I'm not sure that Fe(III) has any way to stay in solution in NaOCl; wouldn't it be thrown down as Fe(x)O(y)H(z) (pretty
much all of those are insoluble AFAIK, whether it's magnetite or hydroxide or plain old rust)? It should also be noted that the wine-dark color is not
some evanescent thing; a 1cm thickness of this solution is basically black, even a few millimeters thickness shows the color strongly. I have a
saturated solution of Fe(II)Cl, which is significantly less opaque. Last but not least, if there were Fe(III) salicylate in solution, I'd expect it to
be transformed into sodium salicylate when I neutralized the solution with sodium bicarbonate - but there was no noticeable color change when I did
that.
Also, here's another trial I did. It doesn't directly disprove your idea (there is too much excess sodium for iron salicylates to form), but neither
does it yet disprove mine:
100mmol (55 tablets) of aspirin was added to 100ml of water. To this was added 200mmol (8g) of NaOH. This was stirred and heated gently (somewhat
short of boiling) to speed dissolution and hydrolysis. Finally, 100mmol of NaOCl in 6.25% solution (119g) was slowly added with constant stirring.
No evolution of gas was seen. The solution, which had been a colorless and slightly translucent, turned yellowish[1]. The same sweet odor previously
noticed (on all my previous hypochlorite+salicylate trials) was present. No red, brown or purple color was seen.
There are two primary hypotheses:
Na-COO-(C6H4)-OH + NaOCl -> Na2CO3 + Cl-(C6H4)-OH (2-chlorophenol). (the chlorophenol hypothesis)
Na-COO-(C6H4)-OH + NaOCl -> Na-COO-(C6H3)-(OH)2 + NaCl (the gentisic acid hypothesis)
Most other products for which the stoichiometry would work out either a) are colorless, or b) would require the evolution of CO2 and/or chlorine gas.
For some reason, I can't find reference to any quinone-type compounds that also have a carboxylic acid group hanging off the ring; maybe it's an
unstable configuration.
Anyway, simply acidifying the solution should allow me to eliminate one of the above ideas. HCl + Na2CO3 would give off CO2, while HCl plus sodium
gentisate would throw down the gentisic acid (the salt is water soluble but the acid basically isn't).
I did two other trials involving Bordeaux powder (CaOH+CuSO4), following on some ideas in the patents that garagechemist posted here, which I will
describe once I've had a chance to work up the products.
I also want to point out that no polymerization is required to get a red and/or purple color. Gmelin describes purple, red, and brown
clorophenol/chloroquinone compounds with just one ring.
[1] The NaOCl itself is also faintly yellowish. The yellowish color produced following the reaction is perhaps ten times stronger.
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bbartlog
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| Quote: | | In all 7.8g of crude phenol was obtained, the theoretical yield was 13.6 g and the percentage yield was 58%. |
58% is pretty good for dry pyrolysis, IMO. Any idea what the red oil you saw distilling over was? The redness in most of the mixtures I have comes
from calcium phenolate. Which, however, turns out to be thermochromic - the red color disappears by the time you get up to boiling temperatures,
making it much less useful than I hoped as an indicator. It also seems to be less stable than sodium phenolate.
Anyway I'm leery of trying to pyrolyze powders, I've too often found that the bottom burns before the top melts, and to me it feels like giving up on
quantitative yields before you've even started.
[Edited on 13-11-2009 by bbartlog]
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aonomus
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In basic conditions, it would form the hydroxide and precipitate out, and any free chloride ions would end up forming chlorine gas (which throws FeCl3
out the window).
So, you're right bbartlog, even if there was free Fe(III) (I didn't think about whether it would stay in solution or precipitate out), with sodium in
excess, it would form sodium salicylate preferentially, so a slight amount of Fe(III) wouldn't give something so dark its pretty much black... well,
it was a theory 
Although I must ask, any pictures?
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entropy51
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| Quote: | | Anyway I'm leery of trying to pyrolyze powders, I've too often found that the bottom burns before the top melts, and to me it feels like giving up on
quantitative yields before you've even started. |
You expect quantitative yields? In organic chemistry? LOL. ROFLOL. LMAO.
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bbartlog
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I didn't say I *expected* them (and it's not like I've achieved them, unless you count trivial things like NaOH+HCOOH)... just saying, pyrolysis seems
like an extra handicap. Though maybe I just need to take more care with heating and/or get thicker glassware.
| Quote: | | Although I must ask, any pictures? |
Not yet, but I'll see what I can do.
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bbartlog
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Acidifying the yellow solution that came from equimolar aspirin and NaOH plus bleach gives a white precipitate. So my chlorophenol idea is dead. In
fact, in this case the bleach didn't even decarboxylate the salicylic acid. It must have oxidized it to some dihydroxybenzoic acid, whose sodium salt
was soluble.
As for the earlier case with the dark solution, I found this in a (Google) book called 'The preparation of organic compounds' from 1891:
' The true quinones, such as benzoquinone, readily add on two molecules of a phenol to form an addition compound [...]
The reaction takes place very readily, the highly coloured addition products being produced when the solutions of the two components
are mixed. Hence the quinhydrones are the first products formed when the quinones are prepared by oxidising the hydroquinones [...]
One part of benzoquinone and 2 parts of phenol are heated under a reflux condenser in ligroin solution for a few minutes. On cooling (the solution
having been concentrated if necessary), the phenoquinone separates out as red needles with a green reflex. ..'
So red and brown (and probably other colors) don't require chlorination; presumably the chlorine ended up with the sodium as it usually does. I think
that's all the investigation of aspirin+bleach I care to do. But the phenol trials are still ongoing.
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not_important
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| Quote: | | I can't find reference to any quinone-type compounds that also have a carboxylic acid group hanging off the ring; maybe it's an unstable
configuration |
Look at just a part of a quinone with a carboxylic acid group:
O=C-C-CO2H
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and then consider acetoacetic acid CH3COCH2CO2H, normally only seen as its esters as it readily decarboxylates. The beta-keto isn't very stable, nor
to a lesser degree is =C-CO2H
If you could oxidise salicylic acid to quinone in useful yields it could be handy for the small scale experimenter, as benzoquinone and hydroquinone
have both become more difficult or more expensive to obtain in some regions. Potassium chlorate with HCl will convert salicylic acid to chloranil,
which has a number of uses.
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aonomus
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Interestingly enough, I had my ASA in a vial in a dessicator, and gave it a stir to break up lumps and get the wet stuff exposed to dry air. I used a
brand new (and just washed with dH2O) stainless steel spatula, and the next day I checked it and there was trace purple/red/brown. Presumably even
trace residues were enough to cause formation of a iron complex, which I find odd because the acetyl group is still attached, meaning the ASA would be
a monodentate ligand and form a completely different complex.
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Nicodem
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| Quote: Originally posted by bbartlog |
Finally, I added 3mmol (2 tablets) of aspirin to 40ml of bleach. This appeared to hydrolyze, but did not react further, leading me to believe that
either acidic conditions or free chlorine are needed for this reaction to take place. |
What do you mean by "appeared to hydrolyze, but did not react further". Have you followed the reaction by any mean? What kind of analytical method did
you use?
| Quote: | | Anyway - what is this purple phenol compound? The principal reaction that comes to mind for these conditions is some sort of Hunsdiecker reaction of a
salicylate salt, which should give a chlorophenol. But chlorophenols, mono- or otherwise, aren't purple AFAIK. |
What has the coloration of the reaction mixture has to do with whether or not any chlorophenol formed or not? The color is completely irrelevant. Only
analysis can say what formed and what not. The chlorination of salicylic acid proceeds in this order: 5-chlorosalicylic acid (and some 3-chloro
regioisomer), 3,5-dichlorosalicylic acid, 2,4,6-trichlorophenol (via ipso electrophilic substitution) and finally chloranil via exhaustive
chlorination.
Neither Hunsdiecker reaction and much less a Krapcho decarboxylation (you do not even have an aliphatic ester here!) have anything to do here.
Instead of doing careless experiments (200mmmol! WTF?) without thought (ever heard about literature and theory?), and without goal (where is any
analysis?), and then going wild with speculation (fairytales are for kids only!), how about an scientific approach that would give a result?
Don't get the impression that I'm trying to humiliate you by criticizing you this way. If I would not appreciate your eagerness I would not even
bother trying to direct you toward science. I promise you that if you try a scientific approach it will be much more rewarding than anything you did
up to now. The only results worth being proud of are those that are backed up with hard evidence.
For the beginning try to prepare 5-chlorosalicylic acid by chlorination of aspirin or salicylic acid with bleach. First try searching for all the
literature preparations that you can find. Report here anything you find and then I will help you and direct you trough a preparative
synthesis.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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bbartlog
International Hazard
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Thanks for the feedback. Yes, the use of 200mmol is pretty stupid as a trial, and I definitely need to do a lot more reading, both in general and
before any specific experiment that I know little about. As for the use of color as a guide to what may or may not have happened, it seems like a
decent enough fallback given that I don't have an IR spectrometer (or any other kind, other than the narrow-spectrum visible-light spectrometer at the
back of my eyeballs); the 19th century chemists found it useful enough to record the color of their products pretty assiduously.
That the named reactions I mentioned couldn't possibly have taken place I already realized.
| Quote: | | First try searching for all the literature preparations that you can find. Report here anything you find and then I will help you
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Thanks for the offer! I am probably going to shelve the hypochlorite and aspirin experiments for now, though, just because I need to build out some
lab equipment first... my wife insists on a fume hood for some reason, and with that I'll probably make a bench with some other rigging.
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Nicodem
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| Quote: Originally posted by bbartlog | | As for the use of color as a guide to what may or may not have happened, it seems like a decent enough fallback given that I don't have an IR
spectrometer (or any other kind, other than the narrow-spectrum visible-light spectrometer at the back of my eyeballs); the 19th century chemists
found it useful enough to record the color of their products pretty assiduously. |
That is exactly why I said you should do a preparative synthesis. Obviously I understand you have no HPLC, GC or NMR at hand. TLC plates are usually
the simplest thing an amateur can use, but they are not very cheap and, just as HPLC and GC, they would also not give all the answers unless you
already had of all the chlorophenolic products of the reaction to be used as reference standards.
However, by doing a preparative reaction you can prove the identity of the major reaction product, because you isolate it. Then a simple
measurement of a melting point is all that it takes. That is the scientific approach available to any amateur.
| Quote: | | I am probably going to shelve the hypochlorite and aspirin experiments for now, though, just because I need to build out some lab equipment first...
my wife insists on a fume hood for some reason, and with that I'll probably make a bench with some other rigging. |
It is a good idea to build/buy some equipment first. Besides when having a tendency to do experiment on such a huge scale without having enough
knowledge and experience, it is most certainly also a good idea to have a fume hood. Chlorophenols are rather toxic and moreover they stink terribly
and their stench hangs on for months on clothes and hairs. But in the monochlorination of salicylic acid you should be able to avoid unhealthy, stinky
decarboxylated products. 5-Chlorosalicylic acid does not stink, is not as toxic, can be efficiently purified by recrystallization, and this experiment
is actually quite nice for a beginner. You would learn a lot.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
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benzylchloride1
Hazard to Others
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The red oil that distilled after the phenol was possibly phenyl salicyclate; the oil dissolved in toluene and when the toluene was extracted with
sodium hydroxide, it remained in the toluene layer. This was almost a year ago and I did not have my infrared spectrophotometer at that time, also I
do not have the oil to run characterization tests, so the identity of this compound is purely conjecture. I have prepared 4-chloro and 4-iodo phenol
from the acid hydrolysis of p-acetamidophenol and subsequent diazotization, and the respective Sandmeyer reactions. The products were both awful
smelling with a clinging odor.
[Edited on 16-11-2009 by benzylchloride1]
Amateur NMR spectroscopist
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