Sciencemadness Discussion Board - Isolating Methyl Salicylate

Isolating Methyl Salicylate

mericad193724 - 18-10-2006 at 14:24

I have been spending the last few days doing Methyl Salicylate synthesis. I used 20g Salicylic acid, 60ml Methanol and 6.25ml Conc. H2SO4. I refluxed, boiled off the methanol and added this to about 200ml of water as the Vogel procedure says...the problem is I don't see a layer of Methyl Salicylate, it should be around 15 ml. It smells really strongly of Wintergreen but no layers! Just a clear solution!

My question is can I add food coloring (or something) to the solution so that it will dissolve in the water and make it colored, but not dissolve in Methyl Salicylate so I can see the layers distinctly?

Help would be really appreciated

thanks

Mericad

guy - 18-10-2006 at 14:43

Try adding salt to separate it.

The_Davster - 18-10-2006 at 14:52

It floats on water, and as such with a small ammount it is hard to see the layers because the meniscuses(sp?) are so close together. Using a smaller diameter vessel may make it easier to see.

mericad193724 - 18-10-2006 at 16:31

I thought it sinks on water...it is more dense(1.184 g/ml) and Vogel says to remove the lower layer in a separatory funnel as the Methyl Salicylate. Am I confused or are you confused

The_Davster - 18-10-2006 at 16:43

I am not familiar with the vogel procedure...I only remember an upper layer when we made it in a similar fashion in a high school lab years ago. Perhaps the unreacted methanol is preferentially soluble in the wintergreen decreasing the density? Not sure.

(dang you...I have a midterm tomorrow....now I am gonna go try this

)

EDIT: Ok, yeah it does seem pure methyl salicylate sinks, wheras with methanol it floats. You boil your methanol off, therfore it should sink, I just crash mine into water, therfore it floats. I would try the salting out method then.

[Edited on 19-10-2006 by rogue chemist]

[Edited on 19-10-2006 by rogue chemist]

mericad193724 - 18-10-2006 at 17:28

Thanks for trying it out Rogue!

I just tried salting it out (NaCl) and that didn't work so I added food coloring as a last stand...and that didn't work either.

There is one thing I never explained, the first time I diluted by reflux mix with water I immediately had about 10ml of a dark black liquid at the bottom. I let this sit in the freezer for ONLY few minutes and the black stuff solidified so I though it was just insoluble crap. (Its black because I used Drain Opener H2SO4). It may not be a coincidence that I expected a yield of 10ml and this black liquid was 10ml. I should try to distill this "crap" and see if I get the methyl salicylate out of the dark mashed potato looking stuff.

It has a high BP in the 220s. I always hear people distilling it under vacuum, can you distill it without a vacuum or will it decompose?

Mericad

BeerChloride - 20-10-2006 at 06:55

Hi mericad. That black stuff was your product I believe.
I'm synthesizing methyl salicylate (MS), too. I'm still working on getting some pure conc. H2SO4 from the nasty brown stuff. Battery acid is pretty pure - I concentrated some to 88.0 %, and used that. It worked, but I don't know what my yield was for the methyl salicylate - need some ether for extraction. Don't try naphtha. It doesn't work.

Methyl Salicylate: solubility: 0.08 g/100 ml H20 @20 C. mp: -9 C

I think it might help to neutralize the excess H2SO4 (not exactly sure why). After I did that, when I added water (I didn't evap the excess MeOH) the MS appeared as oil droplets. It's quite insoluble. Here's the thing: MS evaporates! It's vapor pressure is over 1mm Hg at 54 C. So even though the bp is over 200 you can lose it by heating. The ether method is good because ether is much more volatile than MS, so you can remove it. I suppose you wouldn't lose too much by boiling away the MeOH, but then you still have to separate from the aqueous. It would seem that without ether (or something else that might work) it's best to just add water and pull off what settles on the bottom.

mericad193724 - 20-10-2006 at 19:30

BeerChloride,

you are right the black stuff was the methyl salicylate. I neutralized this with Na2CO3 and added salt...I had two layers form. I added water and extracted the black Methyl Salicylate 5ml (Methyl Salicylate on top). GOOD DAM BLACK H2SO4 ruins everything! I tried to extract clear methyl salicylate from this by distillation. I ended up with some milky stuff that LOST ITS WINTERGREEN SMELL! It smelled like burnt rubber.

Did it decompose??? I can't find decomp temperature anywhere...can someone please let me know what it is?

(Clear H2SO4 is a MUST for Methyl Salicylate synthesis)

Mericad

not_important - 20-10-2006 at 21:16

There's no problem with distilling methyl salicylate at ordinary pressure - bp 220 to 224 - provided there's nothing nasty mixed with.

Milky sounds like it was not dry, and had acid or base in it. Smell like burnt rubber is not triggering any synapses.

I've made methyl salicylate using drain clearer H2SO4, but I used less acid and flood refluxed through molecular sieves to force the reaction.

neutrino - 20-10-2006 at 23:39

Maybe something in the acid decomposed when heated to this high temperature. Try pre-boiling the acid for a while to destroy any pesky organics ahead of time.

YT2095 - 21-10-2006 at 01:23

an Idea occured to me and I don`t know how feasible it is, but rather than risk partial hydrolysis of the ester when the base passes through the floating layer of the ester to neutralise the acid underneath, if you were to put in a peice of reasonably reactive Metal such as Mg or Zn into the mixture, that should react with only the acid part and leave the product alone.
although I have no idea what action the H2 liberated would have on the ester?

does this sound like a workable (less messy) method of neutralising for seperation?

I have a batch of Ethyl salicylate on the go at the moment (I don`t have any MeOH), I`ll give this a test anyway, if it works I think this may be nice method that adding liquid Bases, and also as a solid the reaction will stop anyway even if you added too much, and the product is a salt too, killing 2 birds with one stone.

Just a Thought

unionised - 21-10-2006 at 03:02

Using an insoluble base like CaCO3 works too.

BeerChloride - 21-10-2006 at 07:25

Sorry you lost your product, mericad. I think I smelled something similar when I tested a distillation of that H2SO4 on microscale. It turned black. I'm going to buy some dry ice and try to freeze-purify some of it. It SHOULD work...

I just did another synthesis of MS. It worked really well. It was a bitch, though because with the quantities I used my product was about 0.6 ml.

After the reaction, the reaction flask was cooled with ice water. Then I added aqueous sodium carbonate (NOT saturated) to the reaction flask. A bubble of MS formed easily in the bottom. When I adjusted the ph right to 7, it turned a little cloudy which I realized later was probably unreacted salicylic acid. The MS will be on the bottom unless you add enough salt to your aqueous part to make its density greater than MS. However, it does do funny things with the surface tension, and bubbles of MS always tend to float on the surface. So I finally poured the mixture into a sep funnel with a small teflon valve. I then spent some time swirling and generally working with the mixture gently to cause as much MS as possible to sink into the blob at the bottom. The blob was cloudy from some unreacted salicylic acid. So I taped a tiny piece of paper towel to the outlet tube as a filter, and carefully let the blob out into a vial. This worked and I recovered clear MS with no water in it.

I wonder if I hadn't completely neutralized the excess H2SO4 whether the MS would have been clear the whole time? One other thing, actually I used a pipette to remove almost all the aqueous layer, then added about 10 ml of very saturated aqueous calcium chloride which is very dense. The MS was then less dense than the aqueous and floated, but half of it was still stuck on the bottom. So I diluted the whole thing with 20 ml of water. When the MS blob accumulated again, it was clear, but at that point I could tell that the bottom half of the MS blob had fine white crystals making it look milky, which is why I made the filter. I'm not exactly sure at this point what effect the calcium chloride really had. Perhaps it did pull some water out of the MS/SA blob.

not_important - 21-10-2006 at 08:26

It's not too good of an idea to toss reactive metals into an organic reaction mix. In your case it would not likely cause problems, but in other cases it or the H2 could reac with the product; try a more general solution.

Use mild bases - sodium bicarbonate for example, and keep the solution cool or cold. With esters, don't try to use aqueous ammonia though.

With small quantities of reactants/products, use solvents to prevent losing product through bits being trapped in the aqueous layer or sticking to glassware. Also use salt water for all but the last wash, to minimise the loss through solution.

Saturated CaCl2, or even saturate NaCl solutions, will generally pull some water out of the organic layer.

To clean up that dark H2SO4, try heating it to 80 C or so, and adding H2O2 (even 3%) drop by drop until the acid clears. Alternatively add a persulfate, a small bit at a time. When you're done with that you will want to heat the acid up more to remove added and formed water. If you used persulfates you may find that the matching sulfate crystallises out when the acid is cooled, just decant the acid from the crystals.

Dariusrussell - 3-11-2013 at 15:20

I just tried to separate my Methyl Salicylate dissolved in MeOH and I had some issues...
I poured the Methyl Salicylate/MeOH into a sep funnel and it was looking nice. I then poured in a saturated NaCl solution, upon addition a 'plume' of fine white precipitate formed. Along with that a very nice emulsion formed with it. After circa 1 hour passed there was about 5-7mL of a liquid at the bottom that had phase separated.
Has anyone experienced this white precipitate?
Im leaving it overnight to see if will settle out or do something. It seems to be nonpolar but I havent been able to collect any of it.
EDIT: Sorry I forgot to add my pictures...
Also could this precipitate be excess salicylic acid? It seems like a lot more than I used.

[Edited on 3-11-2013 by Dariusrussell]

[Edited on 3-11-2013 by Dariusrussell]

adamsium - 4-11-2013 at 05:37

I had a similar experience. I believe I ended up deciding - as you mentioned - that it was likely excess starting material. I think I removed it by washing with a dilute bicarbonate solution... I got a very nice clean and clear product in the end. I will have to check back over my notes if more details are needed, as it was a while ago now.

[Edit:] I will have to check my notes. I do remember that I got crystals in my product, so then did the aqueous base washes to remove it. I'm pretty sure that I decided it was almost certainly salicylic acid, but I can't recall if i had crystals at the point of the initial separation. As I said, though, I did end up with a very clear and clean product in the end.

[Edited on 4-11-2013 by adamsium]

Dariusrussell - 4-11-2013 at 06:12

I filtered and did a base wash, unfortunately I lost almost all my product at some point along the way. Later today I will make some mere Methyl S, and leave it to reflux overnight to ensure that it's completed and then do the workup again. I'm new to doing workups, should I do the base wash first, or NaCl wash?

adamsium - 5-11-2013 at 04:10

Okay, I've looked at my notes and here is what I did to isolate the product:

Added the reaction mixture to DCM and chilled water in a separatory funnel. According to my notes, a "flocculent white precipitate was observed" at this point. The precipitate was in the lower (organic) layer after shaking.
Additional DCM was added, and the precipitate dissolved.
Organic layer was run out into a flask, and another lot of DCM was added to the funnel for another extraction. The organic layers were combined.
The aqueous layer was "very clear", but "there was a slight cloudiness to the organic layers in the erlenmeyer (although they appeared clear whilst in the sep. funnel)" (according to the notes).
The organic layers were returned to the sep. funnel and shaken with aqueous sodium bicarbonate until gas evolution appeared to have ceased. The organic layer was collected in a flask again.
Anhydrous sodium sulfate was added to the organic layer in the flask until free flowing.
The mixture was gravity filtered.
The filtrate was placed in a pear flask, along with a boiling chip, and the DCM was distilled off.
After distilling off the solvent and allowing the product to cool, "small, needle-like crystals crashed out [in the product]. The liquid was almost colourless, with perhaps a very slight hint of yellow/brown."
For some reason, my notes stop there, but I remember that I basically just continued with more bicarbonate washes and extractions (after dissolving the product in DCM again to transfer it from the flask and perform the washes/extractions). After distilling off the DCM again, I ended up with a quite clear (just a slight yellowish tint, depending on the lighting) liquid as my product, with a very strong smell of wintergreen.

I've deliberately not mentioned any quantities, as I did this on a minuscule scale (in 2 mL of methanol, and even that was a >1000% excess for my 4 mmol of salicylic acid).

Here's a picture of my final product, after ampouling:

[Edited on 5-11-2013 by adamsium]

Dariusrussell - 6-11-2013 at 16:28

I finally finished my rerun and the yields could be better.
-I refluxed a mixture of 2 grams Salicylic Acid and MeOH for 6 hours to ensure close to completion.
At that point I observed a noticeable layer floating on top
-Added chloroform (closest similar solvent I had) and worked up
I got hit with the unfilterable crystals again, but I had an epiphany, I recrystallized them and then filtered.
I ended up with ~1mL of a yellow oil, it is pretty impure. I might try a 20g/~300mL run with some HEET to make a large enough batch to vacuum distill.
Does anyone have any tips on running this large(ish) scale?

ChemSwede - 26-12-2013 at 08:50

I'm borrowing this thread.

The other day I tried to synthesize methyl salicylate, but I consider it a failure.

I used:
5,6g sodium salicylate
20ml methanol
ca 2ml conc. H2SO4

I mixed the reagents in a 100ml RBF and added some boiling chips.
I then connected it to a 20cm Liebig condenser with cooling water and refluxed at around 70 C for about 1½h. During reflux the mixture never became clear, it remained milky white.
After reflux I let the mixture cool and then added about 20ml of 40% NaCl-solution. Then I added about 30ml of ethyl acetate.
Two layers formed, both milky, but most of the white powder was in the aqueous layer.
Here I decided that the experiment was a failure.
I added some NaHCO3-solution too see if any acid was present, but no bubbles of CO2 were formed.

What went wrong?

I assume that the white, milky powder was unreacted sodium salicylate or salicylic acid, and there was a lot of it.
Should I use salicylic acid instead of sodium salicylate? I thought the sulfuric acid would convert the salicylate to salicylic acid.

Should I use more methanol?

Longer reflux time, higher reflux temp?

Any advise on how to improve the synth is greatly appreciated.

I actually smelled some methyl salicylate mixed with ethyl acetate, but it was a very weak smell.

subsecret - 26-12-2013 at 10:48

When I attempted this reaction yesterday, I used 9 g of salicylic acid and 30-40 mL of methanol, with 3 mL of concentrated sulfuric acid as a catalyst. I figured it would be better to keep the methanol in excess, as I'd prefer to have excess liquid reagents rather than solid ones. The mixture was refluxed with calcium sulfate boiling chips (just 2 or 3) for approximately 40 minutes at approximately 67 C. I noticed the strong smell of wintergreen when wafting from the flask. The boiling chips had dissolved to small specks.

After these forty minutes had passed, the mixture was cooled slightly and and an excess of sodium hydroxide solution was added to neutralize the catalyst. Upon addition of this sodium hydroxide, a white precipitate formed. After about an hour, colorless, needle-like crystals began forming on the surface of the (now settled and compacted) precipitate. The total volume of this solution was almost 200 mL. The solution is in a beaker in my lab, covered with aluminium foil.

No layer separation was observed, and a real work-up was not performed due to the precipitate.

Causes for this?

Perhaps the reaction was not heated for long enough, and the yield was not very high.
As ChemSwede said, there is a lot of the precipitate, and I'm not certain that it's excess salicylic acid. I think there is more material here than I had started with. (I could be wrong, and I may eventually dry and weigh the material).

I will eventually do a work-up of the solution with the precipitate, once I figure out what's going on.

Nicodem - 26-12-2013 at 10:53

Quote: Originally posted by ChemSwede

I used:
5,6g sodium salicylate
20ml methanol
ca 2ml conc. H2SO4

Don't you think that you need to at least compensate the amount of sulfuric acid for the sodium salicylate. You used 37 mmol of sulfuric acid per 35 mmol of the salicylate, which makes it almost no sulfuric acid present, just NaHSO₄. You need a prety good amount of sulfuric acid for this specific esterification. See the discussions above and in the other methyl salicylate threads. You can use aspirin directly, if you have no salicylic acid or don't want to bother with its preparation from the sodium salt.

ChemSwede - 26-12-2013 at 12:01

[/rquote]
Don't you think that you need to at least compensate the amount of sulfuric acid for the sodium salicylate. You used 37 mmol of sulfuric acid per 35 mmol of the salicylate, which makes it almost no sulfuric acid present, just NaHSO₄. You need a prety good amount of sulfuric acid for this specific esterification. See the discussions above and in the other methyl salicylate threads. You can use aspirin directly, if you have no salicylic acid or don't want to bother with its preparation from the sodium salt.[/rquote]

Thank you for your answer.
I suspected that I made that mistake. If I understand correctly, what I should have done was calculate the amount of H2SO4 needed to convert 5,6g salicylate to salicylic acid, and then also add an excess as catalyst.

The lack of CO2 bubbles when I added NaHCO3 also proves that there was no acid left.

Before my next attempt I think I will convert some salicylate to salicylic acid, using HCl. I can't find a thread about it.
It's a bit OT, but what method would give the best yield?

Hexavalent - 26-12-2013 at 12:30

Quote: Originally posted by ChemSwede

Before my next attempt I think I will convert some salicylate to salicylic acid, using HCl. I can't find a thread about it.
It's a bit OT, but what method would give the best yield?

Dissolve your sodium salicylate in water, and drip 12M hydrochloric acid into the solution until no more crystallizes out. There exists an equilibrium between chloride, hydrogen, sodium and salicylate ions, which is driven to the right, forming solid salicylic acid, due to the formation of this weaker acid and its poor solubility in water. When no more salicylic acid precipitates out, vacuum filter the solids, wash with ice-cold water, and recrystallize if necessary.

[Edited on 26-12-2013 by Hexavalent]

ChemSwede - 26-12-2013 at 13:52

Quote: Originally posted by Hexavalent

Quote: Originally posted by ChemSwede

Before my next attempt I think I will convert some salicylate to salicylic acid, using HCl. I can't find a thread about it.
It's a bit OT, but what method would give the best yield?

Thanks!
I suppose it will be reasonably pure after just rinsing with cold water.
If I was to purify it further, what solvent would be most suitable for recrystallization? Salicylic acid seems to be soluble in many organic solvents.
Should I just dissolve the acid, filter and then evaporate the solvent?

I'd like to do a recrystallization with minimum losses.

DraconicAcid - 26-12-2013 at 14:12

Recrystallize it from water. It dissolves in hot water, precipitates from cold.

Paddywhacker - 26-12-2013 at 14:52

While brine is often useful for assisting separations, I think that in this case it is causing problems.

Firstly, the excess methanol is decreasing the solubility of salt, causing it to crash out of solution, and secondly, brine is denser than water, making the aqueous phase closer in density to the heavy methyl salicylate. The closer the difference in density, the more difficult the separation.

ChemSwede - 26-12-2013 at 16:10

Quote: Originally posted by Paddywhacker

While brine is often useful for assisting separations, I think that in this case it is causing problems.

Firstly, the excess methanol is decreasing the solubility of salt, causing it to crash out of solution, and secondly, brine is denser than water, making the aqueous phase closer in density to the heavy methyl salicylate. The closer the difference in density, the more difficult the separation.

That's good information! Maybe some of the milky stuff was NaCl that crashed out of solution.
But shouldn't methyl salicylate combine with ethyl acetate anyway, leaving methanol and NaCl in the aqueous layer?
I clearly saw two phases when I added ethyl acetate.

I'm actually a bit uncertain about how to extract my product with the best yield.
There are a few threads about Me-salicylate here on SM, which I have read.
Some suggest boiling off the methanol before adding water/brine. Others state that that would concentrate the H2SO4, leading to problems. I suppose the acid could hydrolyze the ester.
Some just add water and then ethyl acetat. The methanol goes in the water phase and the product into the ethyl acetate.
What's optimal?

Also I'm wondering about increasing yield during reflux. Since it's an equilibrium I suppose that removing water would create more product. For instance, I have a Dean Stark trap, but I'm not sure if it's any good here, since methanol is soluble in water.
Adding more H2SO4 might get rid of some water, but then the ester might hydrolyze.
How about adding some drying agent?

ChemSwede - 26-12-2013 at 16:12

Quote: Originally posted by DraconicAcid

Recrystallize it from water. It dissolves in hot water, precipitates from cold.

Sounds good if that works. I can't seem to find any data on solubility in hot water, but it would be much more convenient to use that method.

[Edited on 27-12-2013 by ChemSwede]

DraconicAcid - 26-12-2013 at 19:20

Quote: Originally posted by ChemSwede

Quote: Originally posted by DraconicAcid

Recrystallize it from water. It dissolves in hot water, precipitates from cold.

Sounds good if that works. I can't seem to find any data on solubility in hot water, but it would be much more convenient to use that method.

[Edited on 27-12-2013 by ChemSwede]

It works- my grade-11-equivalent students do it every year, to separate it from a mixture containing salt, sand, MnO₂ and glycerol. I can look up the solubilities when I'm back in my office.

ChemSwede - 27-12-2013 at 02:24

Quote: Originally posted by DraconicAcid

Quote: Originally posted by ChemSwede

Quote: Originally posted by DraconicAcid

Recrystallize it from water. It dissolves in hot water, precipitates from cold.

Sounds good if that works. I can't seem to find any data on solubility in hot water, but it would be much more convenient to use that method.

[Edited on 27-12-2013 by ChemSwede]

It works- my grade-11-equivalent students do it every year, to separate it from a mixture containing salt, sand, MnO₂ and glycerol. I can look up the solubilities when I'm back in my office.

I just found this page:
http://www.ce.gxnu.edu.cn/organic/net_course/content/Crystal...
It says that the solubility for salicylic acid in water at 100 C is 6,67g/10ml.
So I will prepare some pure salicylic acid and then go for the methyl salicylate again.

ChemSwede - 9-3-2014 at 12:29

Attempting to synthesize methyl salicylate again.
I've kind of failed once using 5,6g salicylic acid (prepared from sodium salicylate), 20ml methanol and 2ml conc. H2SO4.
I refluxed for 1½h at 70C.
Afterwards there was a strong smell of methyl salicylate, so I added ethyl acetate and water. Separated the organic layer and evaporated the ethyl acetate. I was left with a lot of white crystals with a strong scent of my ester.
i could not isolate any product since there was so much unreacted salicylic acid left.

I have now dissolved the unreacted salicylic acid from my last attempt in 40ml methanol. I added 3ml H2SO4 and a few boiling chips.
The mixture is refluxing at 70C, and I'm planning to do that for 2h this time.
Any ideas on how to make all the salicylic acid react? Temp, reflux time, adding more methanol during reflux?

How do I isolate the product most efficiently? Adding water and recover the ester at the bottom or boil away the methanol and add water f ex? There are a few suggestions on the forum.
Some say that boiling away the methanol concentrates the remaining acid, which is not good for the product.

Would be nice to isolate the product this time. Wonderful smell.

Mailinmypocket - 9-3-2014 at 13:02

Try using dichloromethane or chloroform, if you can, instead of ethyl acetate

ChemSwede - 9-3-2014 at 13:05

Update on the synthesis.
I left the reaction for ½h, and when I returned the boiling had stopped, but the temp was still 70C.
Did the methanol react or boil away?
Perhaps I should run the reflux at 60C since methanol boils at around 70C. My Liebig condenser is just 20cm so perhaps it cannot keep the methanol from boiling away.

I added another 25ml of methanol and the boiling started again. Still refluxing at 70C.

ChemSwede - 9-3-2014 at 13:19

Quote: Originally posted by Mailinmypocket

Try using dichloromethane or chloroform, if you can, instead of ethyl acetate

I have DCM, but I'm a bit reluctant to use it since it's hard to get, and since it's cancerogenous.
It should work with ethyl acetate I guess. I know it has a higher boiling point than DCM, and it can be hard to make it form two layers with water sometimes.

Mailinmypocket - 9-3-2014 at 13:55

You might want to consider it. Check out this link here for solubility of salicylic acid in ethyl acetate vs dichloromethane:

http://lxsrv7.oru.edu/~alang/onsc/solubility/allsolvents.php...

You will barely get any unreacted salicylic acid in your solvent phase. The low boiling point of DCM makes it easier to strip from the methyl salicylate afterwards. But if it is hard to get you can always redistill the DCM and recycle it...

[Edited on 9-3-2014 by Mailinmypocket]