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DutchChemistryBox
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Hello,
This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.
It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.
It should yield about 25g, after the reaction the methanol will be boiled of.
After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse
when I'm going to wash it with water?
How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?
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xfusion44
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@greenlight
@forgottenpasword
Thanks!  Unfortunately I'm very busy these days, so I'll try again, when I'll
have more time
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xfusion44
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Quote: Originally posted by DutchChemistryBox  | Hello,
This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.
It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.
It should yield about 25g, after the reaction the methanol will be boiled of.
After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse
when I'm going to wash it with water?
How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?
|
Before you boil off methanol, there should be about 13.8% by weight sulfuric acid, but I don't know this reaction, so I don't know if H2SO4 reacts
with something, or is there just to bind water to itself? However, if it doesn't react, after boiling, there should be 34.5% by weight H2SO4.
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DutchChemistryBox
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| Quote: Originally posted by xfusion44 | | Quote: Originally posted by DutchChemistryBox | Hello,
This weekend I'm going to make methysalicylate with the procedure from the book of Arthur Vogel.
It says:
28 g. of salicylic acid and 64 g. (81 ml.) of absolute methyl alcohol, and add 8 ml. of concentrated sulphuric acid.
It should yield about 25g, after the reaction the methanol will be boiled of.
After a rough calculation I see that about 25% of the mixture will consist of sulfuric acid. Isn't this wayyyy to much? I'm afraid it will hydrolyse
when I'm going to wash it with water?
How is such amount of sulfuric acid possible? Can somebody explain me why I'm wrong?
|
Before you boil off methanol, there should be about 13.8% by weight sulfuric acid, but I don't know this reaction, so I don't know if H2SO4 reacts
with something, or is there just to bind water to itself? However, if it doesn't react, after boiling, there should be 34.5% by weight H2SO4.
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It is just a fisher esterfication, the sulfuric acid acts as a catalyst. It will also force the equilibrium to the right due to the hygroscopic
caracter of the acid.
Till that part, it won't be rocket science.
I'm trying to understand why such a big amount of sulfuric acid won't hydrolyse my product.
[Edited on 11-12-2014 by DutchChemistryBox]
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Amos
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DutchChemistryBox, you need all that sulfuric acid in order to sequester water produced as a result of the esterification, which will produce a
molecule of water for each molecule of methyl salicylate esterified. The percentage of sulfuric acid will decrease as the amount of water increases,
though why percentages matter so much I'm not sure. It won't hydrolyze as long as it's in the methanol.
Now, in regards to your fear of hydrolysis occuring, it has been my experience that if you decide to boil off the methanol from your reaction mixture
IT WILL hydrolyse your product, and then more bad stuff happens. Trying this yielded a lot of foul-smelling black goo, probably due to the hydrolysis
of the ester followed by an unfavorable reaction of the salicylic acid with so much now-concentrated sulfuric acid. In order to prevent this, I would
recommend adding calcium carbonate to your reaction mixture when you believe it is sufficiently reacted. This will effectively remove the dangerous
sulfuric acid, and the calcium sulfate formed should absorb water present in the mixture to form the dihydrate, which should take care of the water as
well. Since it's insoluble, you can filter it off and squeeze any methanol/methyl salicylate mixture through the filter paper, and then simply boil
off the methanol. There'll be some mechanical losses on such a small scale, though.
[Edited on 12-11-2014 by No Tears Only Dreams Now]
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DutchChemistryBox
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| Quote: Originally posted by No Tears Only Dreams Now | DutchChemistryBox, you need all that sulfuric acid in order to sequester water produced as a result of the esterification, which will produce a
molecule of water for each molecule of methyl salicylate esterified. The percentage of sulfuric acid will decrease as the amount of water increases,
though why percentages matter so much I'm not sure. It won't hydrolyze as long as it's in the methanol.
Now, in regards to your fear of hydrolysis occuring, it has been my experience that if you decide to boil off the methanol from your reaction mixture
IT WILL hydrolyse your product, and then more bad stuff happens. Trying this yielded a lot of foul-smelling black goo, probably due to the hydrolysis
of the ester followed by an unfavorable reaction of the salicylic acid with so much now-concentrated sulfuric acid. In order to prevent this, I would
recommend adding calcium carbonate to your reaction mixture when you believe it is sufficiently reacted. This will effectively remove the dangerous
sulfuric acid, and the calcium sulfate formed should absorb water present in the mixture to form the dihydrate, which should take care of the water as
well. Since it's insoluble, you can filter it off and squeeze any methanol/methyl salicylate mixture through the filter paper, and then simply boil
off the methanol. There'll be some mechanical losses on such a small scale, though.
[Edited on 12-11-2014 by No Tears Only Dreams Now] |
Thank you for the input!
It still makes me wonder why Vogel does it that way.
In his procedure he washes his product with water after boiling of the methanol. Would it be possible to forget about whole the idea of boilling of
and just start directly with washing with water and bicarbonate?
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Amos
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If you left the methanol around it would probably greatly increase the solubility of the methyl salicylate in the solution. I don't see how you could
possibly boil off the methanol without charring and destroying the methyl salicylate, so the Vogel method really doesn't seem viable to me. But I
haven't tried it, so I could be wrong.
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gdflp
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Washing the acidic ester with saturated NaHCO3 to remove excess acid, then washing with saturated NaCl to remove water would most likely eliminate
most mechanical losses rather than adding calcium carbonate. The solution should then be able to boil without fear of decomposition.
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Amos
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| Quote: Originally posted by gdflp | | Washing the acidic ester with saturated NaHCO3 to remove excess acid, then washing with saturated NaCl to remove water would most likely eliminate
most mechanical losses rather than adding calcium carbonate. The solution should then be able to boil without fear of decomposition.
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You can't merely wash the ester until you've isolated the ester from all of the methanol and acid it's stuck in. If you boil off the methanol without
neutralizing the acid, your ester will decompose. If you wash the whole contents of the reaction vessel with sodium bicarbonate solution, you'll be
stuck with an aqueous/methanolic solution of methyl salicylate that is rapidly hydrolyzing thanks to all of the water. At that point I don't think you
could boil off the water and methanol without completely ruining the ester, as removing methanol while water is present will shift the equilibrium
even further to the left.
Given that sodium sulfate is insoluble in methanol, you could maybe add solid sodium bicarbonate to avoid introducing water, but the neutralization
still produces water. When you went to boil the methanol off, your now hydrated sodium sulfate would melt, which might cause hydrolysis of the ester
at such a high temp; not sure.
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gdflp
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I have done this before, an organic layer of methanol and methyl salicylate will form while the sulfuric acid migrates to the aqueous layer and
reacts. Also, water will not hydrolyze methyl salicylate on its own, it needs an acid catalyst for the hydrolysis to occur. After the ester has been
washed, it can simply be distilled to separate the methanol and methyl salicylate.
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Romain
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Hi, is CaCl2 soluble in dichloromethane? I couldn't find any data on that online.
I want to remove water and ethanol from dichloromethane, but I can't distill the mixture with the CaCl2 (don't have the apparatus), so I thought I'd
simply decant the CaCl2/ethanol/water and retrieve the dichloromethane. Now if the DCM is contaminated with CaCl2 that's not exactly ideal...
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DraconicAcid
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| Quote: Originally posted by Romain | Hi, is CaCl2 soluble in dichloromethane? I couldn't find any data on that online.
I want to remove water and ethanol from dichloromethane, but I can't distill the mixture with the CaCl2 (don't have the apparatus), so I thought I'd
simply decant the CaCl2/ethanol/water and retrieve the dichloromethane. Now if the DCM is contaminated with CaCl2 that's not exactly ideal...
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No, it is not soluble in dichloromethane.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Justin Blaise
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I've followed the Vogel procedure for methyl benzoate, which also calls for distilling the methanol off at the end of the reaction if I remember
correctly, and achieved the reported yields.
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DutchChemistryBox
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| Quote: Originally posted by Justin Blaise | | I've followed the Vogel procedure for methyl benzoate, which also calls for distilling the methanol off at the end of the reaction if I remember
correctly, and achieved the reported yields. |
So it seems to work, I've been searching for people who followed the Vogel procedure. I've heard good stories about it.
I'm quite confused now. Almost everbody seems to think that it won't be good practise due to the risk of hydrolysing the product. But it does seem to
work.
The question is why. Maybe there is not enough water present for hydrolysing (due to the hygroscopic caracter of Sulfuric acid)?
That it will only hydrolyse if there is more water or less acid present? That boiling of the methanol only is possible because of the high amount of
acid?
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xfusion44
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I have a few questions about chloroform...
I've made probably about 14ml of CHCl3, using 1l 67g/l NaClO and approx. 14ml of acetone.
First: balanced equation says, that there should be 1 mole of chloroform, when using 3moles NaClO and 1 mole acetone - in my case .9 and .3 moles of
each, so I should get about 35g of chloroform. I didn' measure it yet, but it looks more like 15ml (approx. 20g), than 35, so how is that? I've also
heard, that there should be about as much chloroform as the amount of acetone, that was used, is that true?
Also, how would I neutralise reaction mixture afterwards? Can I get anything useful from it? NaOH or sodium acetate maybe? Probably I can't just pour
it down the drain, so there must be a way to neutralise it.
And the last question is about stabilising it. Would 1% bw of ethanol be enough?
Thanks and sorry for my english
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xfusion44
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PS: also, what color should reaction mixture be? Last time I used excess acetone and after 3h mixtute was completely clear - like water, but now it's
not clear at all and a little yellowish in color - did I use too little acetone this time?
Thanks!
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DrMario
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I am vaguely familiar with the haloform reaction, and how it can be used to make iodoform. But what would happen if one added HCl or other strong acid
to acetone + iodine, instead of a strong base?
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confused
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im guessing that nothing would happen seeing as HCL and iodine will not react together, the iodine might dissolve a bit in the water from the HCL but
that should be it
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DrMario
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I started with a dark brown solution, and after adding HCl I have now a clear solution with a light coloured precipitate (can't tell the colour
exactly as it's in an amber bottle.
There was some potassium sulfate and traces of potassium iodide in the solution as well, so this isn't a well controlled experiment. I will have to
repeat it.
EDIT: after pouring in a pyrex beaker, I see that the liquid is light yellow. The precipitate is brown.
[Edited on 14-12-2014 by DrMario]
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greenlight
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@ Xfusion,
I have always made chloroform similar to the way you specified with 12.5% Sodium hypochlorite solution which I buy in the pool section at the hardware
store.
My ratios are 1000ml NaCl to 45.5ml Acetone but sometimes I scale it up and do 4 litres of NaCl at a time in a bucket if I need a larger amount.
I chill the Sodium hypochlorite to about 0-5 Degrees Celcius, then add the acetone, stir and keep it on ice for a couple hours afterwards before I
seperate it in a sep funnel.
The yields aren't that amazing.
My solution usually goes from the original yellow-green colour of the NaCl to yellow to a milky white with a blob of chloroform on the bottom of the
vessel.
I only make it when I need it now and rarely store it, but when I did i added a small amount of Ethanol to be safe.
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xfusion44
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| Quote: Originally posted by greenlight | @ Xfusion,
I have always made chloroform similar to the way you specified with 12.5% Sodium hypochlorite solution which I buy in the pool section at the hardware
store.
My ratios are 1000ml NaCl to 45.5ml Acetone but sometimes I scale it up and do 4 litres of NaCl at a time in a bucket if I need a larger amount.
I chill the Sodium hypochlorite to about 0-5 Degrees Celcius, then add the acetone, stir and keep it on ice for a couple hours afterwards before I
seperate it in a sep funnel.
The yields aren't that amazing.
My solution usually goes from the original yellow-green colour of the NaCl to yellow to a milky white with a blob of chloroform on the bottom of the
vessel.
I only make it when I need it now and rarely store it, but when I did i added a small amount of Ethanol to be safe. |
Thanks! Today I did it again and noticed, that I've made mistake yesterday in calculations. I should use about 17.4ml, not 14ml, so maybe that's why
it was still little yellowish in color.
How's about rxn mixture, if you know, maybe? Can I just pour it down the drain? Probably not?
Thanks 
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gdflp
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Once you've separated the chloroform, simply let the water portion of the reaction mixture sit for a few days. The haloform reaction produces 2
equivalents of sodium hydroxide which will slowly destroy the residual chloroform and convert it to a mixture of chloride and formate which can be
poured down the drain.
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xfusion44
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| Quote: Originally posted by gdflp | | Once you've separated the chloroform, simply let the water portion of the reaction mixture sit for a few days. The haloform reaction produces 2
equivalents of sodium hydroxide which will slowly destroy the residual chloroform and convert it to a mixture of chloride and formate which can be
poured down the drain. |
Thanks! I've just saw that in one of yt videos on how to make chloroform
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I Like Dots
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Hey guys, today I was in a thrift store, and a shiny butter knife caught my eye.
The knife is silver plated, on copper on an unknown base metal/alloy, The purpose of this post is identifying the base metal/s. Im fairly sure its
pewter, but just curious anyway. Im trying to design an experiment to identify the different elements.
Physical properties
•The base metal is pliable. The knife blade was able to bend somewhat easily, but snapped once bent too far.
•It has a density of a ~6.66 g/cm2
•Has a low melting point. Not sure exactly, but easily melts under a propane torch.
•Melted drops are reflective and shiny, reminiscent of fresh zinc.
Chemical properties
•Reactive towards silver nitrate. Black powder forms on surface.
•Reactive to copper sulfate, black color forms on surface, with copper underneath.
The metal in question
The knife, melted and broken in half
The plating layers visible on the knife
The Knife
The emblem on the knife
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Metacelsus
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I want to turn p-nitrotoluene into p-aminobenzoic acid. I plan to oxidize the methyl group (with permanganate) and reduce the nitro group (with
Sn/HCl).
I've read a reference which calls for acetylation with acetic anhydride before the oxidation (http://pubs.acs.org/doi/abs/10.1021/ed033p71). I have some acetic anhydride, but would rather not use it for this, as it was not cheap to get. Is
this acetylation strictly necessary?
Could I oxidize the methyl group first, and then reduce the nitro group?
(After more research, yes, I can. I guess I answered my own question.)
[Edited on 22-12-2014 by Cheddite Cheese]
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