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Author: Subject: Diethylether (sulfuric acid saving method)
Diachrynic
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[*] posted on 14-5-2024 at 12:05
Diethylether (sulfuric acid saving method)


The preparation of ether is very old. To give some perspective, in europe, diethylether is older than the potato.[5] It this therefore surprising that most documented preparations of ether on YouTube use much more sulfuric acid than actually needed. I was made aware of this process via a good friend and subsequently found a paper that detailed the limits of this procedure.

From the videos I have seen on YouTube, ether yields from 1-2.5x the volume of sulfuric acid have been reported.[1] There is a great post by len1 on this board which mentions most of the important insights from the paper which I will follow here, although without giving any literature and his yield is in the same ballpark as the YouTube ones mentioned before (although he does mention the sulfuric acid in the flask can probably be reused for about 3-4x the volume of ether).[2]

The limits of just how much ether a given amount of sulfuric acid can make was explored by Evans and Sutton in their remarkable 1913 paper.[3] According to them, a given volume of sulfuric acid can make up to 40x its volume of ether.

Synthesis:
Two trials were conducted. A 500 mL three-necked flask containing 40 mL 95% denaturated EtOH + 40 mL 95-98% H2SO4 (Note 1), a stir bar and two ceramic boiling stones (from a broken Büchner funnel) was used for the first trial and it was reused as is for the second trial. It was fitted with a 500 mL addition funnel (Note 2), an internal thermometer, and a simple distillation setup with vapor thermometer and a jacketed coil condenser with 2-7 °C cooling water, as well as a 1 L receiving flask in an ice bath. The synthesis was run in 8-12 h periods with about equally long interruptions for the nights. The oilbath temperature was maintained by a thermometer inside the oil that gave feedback to the hotplate. The temperature in the cooling water was maintained initially with ice, later with a recirculating water chiller.

ether_synthesis.png - 1.9MB
Fig. 1: Sketch of the apparatus.

First trial: The oilbath was maintained at 150-154 °C. When the internal temperature reached 140 °C, the addition of ethanol was started, at first very slowly. In total 734 g 95% denaturated EtOH was added during 26 hours, at a rate that kept the internal temperature at around 140-145 °C (Note 3). Still head temperature rose from 67-84 °C during the reaction (Note 4). The distillate weighed 734 g, was fractionally redistilled (Note 5) and the fraction boiling at 34-40 °C was collected.
Yield: 289 g of ether. (52% of the theory)

The flasks contents were black and contained a substantial amount of tar, but it didn't seem to affect anything, so it was used as is for the second preparation.

Ether_tar_residue.jpg - 1.5MB
Fig. 2: The residue in the distilling flask.

Second trial: The oilbath was maintained at 170-174 °C to allow for faster ethanol addition. When the internal temperature reached 140 °C, the addition of ethanol was started. 1500 g 95% denaturated EtOH was added during 17.5 h, at a rate that kept the internal temperature at around 140-145 °C. Still head temperature rose from 85-92 °C during the reaction. The distillate weighed 1515 g (Note 6), was fractionally redistilled and the fraction boiling at 34-40 °C was collected.
Yield: 545 g of ether. (48% of the theory)

The combined distillation residue of both reactions weighed 1301 g and was then further fractionally distilled (setup as before) to recover the unconsumed ethanol. There was obtained 1015 g of ethanol boiling in the range 78-82 °C and a residue of 282 g, which was dirty water containing some insoluble oil, it had a nauseating smell and was highly acidic. Several spoons of KOH were required to neutralize it. This residue was discarded.

Overall: 40 mL (ca. 72 g) 95-98% H2SO4 and 2234 g 95% denaturated EtOH were employed, of which 1219 g were consumed, yielding 834 g (ca. 1.1 L) of ether, or a 90% yield based on consumed ethanol (that is assuming the recovered ethanol was also 95%).

Notes:
1. The denaturant in the alcohol is likely 1% methyl ethyl ketone, 1% isopropanol and denatonium benzoate. The sulfuric acid was food grade.
2. Evans recommend an addition funnel with a long stem that reaches above the liquid surface, but does not deliver the ethanol onto the internal thermometer. In the experiment, an addition funnel was used that delivered the falling drops of ethanol just on the edge of the liquid in the flask opposite the thermometer.
3. It is often said to match the addition rate to the distillation rate, but in my experience, if the ethanol is added such that the internal temperature stays around 140-145 °C, the distillation rate more or less automatically adjusts to match the addition rate.
4. The temperature in the still head can be pretty much ignored, as it doesn't seem to correlate to the ether content at all.
5. The fractional distillation used a 60 cm vacuum insulated vigreux column, a relux divider still head, a jacketed coil condenser with 2-7 °C cooling water and a receiving flask in an ice bath. The distillation was done slowly over the course of 2-4 hours.
6. The amount of ether present was estimated as follows: About 3 g of CaCl2 and about 10 g of water were mixed (exothermic) and chilled in a fridge to 5-10 °C. A sample of the distillate was placed in a 10 mL measuring cylinder, some of the calcium chloride solution was added and the whole mixture carefully mixed. In this case 1.8 mL of distillate and 2 mL of the calcium chloride solution gave after shaking 0.8 mL of upper organic layer, indicating an ether content around 40-45 vol%.

Discussion:
The combination of stirring and ceramic boiling stones prevented bumping effectively, it was never an issue during the entire experiment.

It can also be seen that the dilution of water is not limiting the reaction, as it distills off alongside the ether and alcohol in more or less a stable equilibrium, if it is allowed to do so.

The effiency does not seem to substiantially drop over the course of the synthesis, going slower than in the second trial brings no real benefit, and the amount of ether produced was about 29x the volume of sulfuric acid. (It seems very likely that it can still keep making even more ether - the limit is not yet reached. It would be indicated by the flask running empty and being practially impossible to maintain at 140 °C, according to Evans and Sutton. Testing the distillate as discussed in Note 5 should also indicate completion.)

There is another thing worth mentioning. Where does the sulfuric acid actually go? According to the 1913 paper,[3] only 15-20% could be accounted as SO2, formed due to oxidation by the sulfuric acid. However, in their 1917 followup,[4] their new analysis shows the following: In the distillate and gaseous exhaust, about 2% of the sulfur in the acid leaves as sulfur dioxide and about 89% remains in the sulfate oxidation state, of which 47% was present as sulfuric acid, 8% as ethyl sulfuric acid, 34% as diethyl sulfate and about 5% as unspecified sulfonic acids and sulphonates. The remaining sulfur is found in the charred distillation residue in the flask.

Whether or not the sulfuric acid and diethyl sulfate actually slowly distill or are merely aerosolized and carried over as droplets from the bursting bubbles during the boiling in the distillation flask is unclear.

But given the toxic nature of diethyl sulfate, it is save to say that the crude distillate should be treated as dangerous not just for flammability reasons. These impurities will likely also be present in the "usual", less sulfuric acid efficient preparations!

Finally, I plan on purifying the ether further by drying over KOH and redistilling, but I anticipate the the majority of the loss will come from handling. I intentionally distilled very slowly, the column is very easy to overheat and flood, so the ether should already be more than adequate for extractions and such.

Literature:
[1] - (a) NileRed, YouTube 2014, "Making Diethyl Ether", https://www.youtube.com/watch?v=6Z2oE8-uthU, (b) myst23YT, YouTube 2010, "Make Diethyl ether", https://www.youtube.com/watch?v=ytdO3YzXNkQ, (c) Amateur Chemistry, YouTube 2023, "Turning Vodka into Diethyl Ether", https://www.youtube.com/watch?v=mot8RrJbRko, (d) Chemiolis, YouTube 2022, "Making Diethyl Ether", https://www.youtube.com/watch?v=cbCbq2OIyPA, (d) Thy Labs, YouTube 2021, "Making Diethyl Ether", https://www.youtube.com/watch?v=Qysm48HiKQo
[2] - len1, SciMad 2008, "Diethyl Ether - Illustrated Practical Guide", https://www.sciencemadness.org/whisper/viewthread.php?tid=9747
[3] - P. N. Evans, L. M. Sutton, "The efficiency of the preparation of ether from alcohol and sulfuric acid", J. Am. Chem. Soc. 1913, 35, 6, 794-800, https://doi.org/10.1021/ja02195a018
[4] - P. N. Evans, G. K. Foresman, "Sulphur By-Products of the Preparation of Ether", Proc. Indiana Acad. Sci. 1917, 27, 211-216, https://journals.indianapolis.iu.edu/index.php/ias/article/view/13303
[5] - Wikipedia, "History of the potato", https://en.wikipedia.org/wiki/History_of_the_potato




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[*] posted on 14-5-2024 at 21:33


Nice experiment and writeup, and a cute diagram (how?)



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[*] posted on 14-5-2024 at 21:48


A good question would be: is it possible to replace sulphuric acid by some other acid? At least partially. Like, say, even sodium bisulphate or sulphamic acid, or a mix thereof.
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[*] posted on 15-5-2024 at 00:10


Sulaiman, thank you. The drawing was done on paper first with pencil, then with black ink, and finally scanned.

Keras, it's an interesting idea. If I remember correctly, heating ethanol and sulfamic acid does produce ethyl sulfamate, CH3-CH2-OSO2-NH2, which would be analogous to the ethyl sulfuric acid that is the intermediate in the classical synthesis. I couldn't find anything whether ethyl sulfamate can form diethyl ether though. Maybe it is worth some testing.
In the 1913 paper cited, they also do a trial where they intentionally dilute the sulfuric acid first to about 30% and proceed as usual. When the mixture concentrates enough to maintain 140 °C, the reaction proceeds as with concentrated acid. Obviously it would be more practical to concentrate dilute acid by boiling most of the water off in a beaker or something before using it in this reaction, but it shows that >95% concentrated sulfuric acid is not a requirement at all.




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[*] posted on 15-5-2024 at 04:03


Quote: Originally posted by Diachrynic  
Keras, it's an interesting idea. If I remember correctly, heating ethanol and sulfamic acid does produce ethyl sulfamate, CH3-CH2-OSO2-NH2, which would be analogous to the ethyl sulfuric acid that is the intermediate in the classical synthesis.[…]


TBH, the idea here is to completely replace sulphuric acid by something else, although 15% sulphuric acid is still available in Europe. Since (at least theoretically) the acid is here only as a catalyst, it should be possible to substitute sulphuric acid for any (strong?) acid. Hydrohalid acids will evaporate, but crystalline acids like sulphamic or even citric acid should not. Bisulphate is quite acidic by itself, since apparently you can make hydrogen chloride by mixing bisulphate with salt.
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[*] posted on 15-5-2024 at 08:00


One candidate that has been mentioned as an option for "strong mineral acid" is phosphoric acid. How does it compare?
(In case of sulphuric acid, reduction to SO2 is one of the side reactions. Phosphoric acid is resistant to that one.)
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[*] posted on 15-5-2024 at 11:59


One substitute I read about somewhere (I can’t remember where) that you could also use is anhydrous zinc chloride.



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[*] posted on 15-5-2024 at 12:06


Phosphoric acid is weak (pKa 2.3) but polyphosphoric acid is stronger (H4P2O7 pKa1 ≈ 0.5, H3P3O9 pKa1 ≈ -12) but it tends to attack pretty much every available material except fused quartz (and possibly even that).

A combination of phosphoric acid with something like methanesulfonic acid might do it.




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[*] posted on 15-5-2024 at 13:26


The following paper has a nice set of experimental data beginning on page 4. Various salts and acids were tested. It seems that van Alphen was not in his best mood back then.

Attachment: J. van Alphen - The formation of ether from alcohol.pdf (428kB)
This file has been downloaded 131 times




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[*] posted on 15-5-2024 at 14:35


Quote:
When reading the original publications of Williamson, one is
immediately impressed by the fact that the only thing he tries to
prove and actually succeeds in proving, is that one molecule of
ether is formed from two molecules of alcohol


Thanks for this gem bnull, some very cool bits of information in here for people looking to tinker.




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[*] posted on 16-5-2024 at 03:18


Thanks for the article. It is quite encouraging. Since acidic salts seem to work, probably sodium bisulphate can be used. I’ll test that next week.

I've tried 75% phosphoric acid. It doesn’t work, at least if assembled the classical way. At 80 °C circa, the ethanol distills out and that’s it. It might be necessary to use a reflux with ice cold water to trap ether and ethanol vapours.
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[*] posted on 16-5-2024 at 09:49


Quote:
The best catalyst was ferric sulfate

A line from the paper I won't soon forget




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[*] posted on 17-5-2024 at 00:00


OK, so far I've tried, in a simple micro 10 mL erlenmeyer connected to a vertical tube (acts an air-cooled condenser), solid sodium bisulphate + ethanol; solid sodium bisulphate + 23% HCl + ethanol; dissolved sodium bisulphate in 37% sulphuric acid.

None of those gave off any ether smell. As far as I’m aware, the ethanol simply boils off at 80 °C circa and then, that’s it. The high boiling point of concentrated sulphuric acid keeps the ethanol bounded despite the temperature. In the paper cited above the experiments must've been conducted inside an autoclave or some high-pressure-bearing system, but definitely not what we usually do, i.e. a distillation apparatus where ethanol is fed in a continuous run.
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[*] posted on 17-5-2024 at 03:02


I'm fairly sure that adding a concentrated aqueous solution of sodium bisulphate to ethanol gives a precipitate of sodium sulphate and a solution of sulphuric acid. (anyone got the Merck index handy? That's where I think I read it but mine's not to hand)
I doubt the reaction goes to completion; but it might go far enough to produce something from which you can distill ether
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[*] posted on 17-5-2024 at 03:50


Quote: Originally posted by unionised  
I'm fairly sure that adding a concentrated aqueous solution of sodium bisulphate to ethanol gives a precipitate of sodium sulphate and a solution of sulphuric acid. (anyone got the Merck index handy? That's where I think I read it but mine's not to hand)
I doubt the reaction goes to completion; but it might go far enough to produce something from which you can distill ether


I’ll try again with a concentrated solution of sodium bisulphate and see what's going on. I’ll keep you in touch.
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[*] posted on 17-5-2024 at 05:02


Quote: Originally posted by Keras  
In the paper cited above the experiments must've been conducted inside an autoclave or some high-pressure-bearing system, but definitely not what we usually do, i.e. a distillation apparatus where ethanol is fed in a continuous run.


From the above reference:

J. van Alphen - The formation of ether from alcohol.

"If not stated otherwise the reaction mixture was heated for eight hours in a closed tube in a Carius oven at a temperature of 155-160°C".

Yeah, I guess it has to be a tightly-closed system, with the exception being sulphuric acid. At these temperatures and pressures, I guess it is a little less surprising that such an array of salts can give ether. But still a potentially good way of making small amounts of diethyl ether if you don't have access to sulphuric acid.
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[*] posted on 17-5-2024 at 05:11


Quote: Originally posted by unionised  
I'm fairly sure that adding a concentrated aqueous solution of sodium bisulphate to ethanol gives a precipitate of sodium sulphate and a solution of sulphuric acid. (anyone got the Merck index handy? That's where I think I read it but mine's not to hand)

"[D]ec by alcohol into sodium sulfate and free H2SO4." (Merck Index, 1972)

Quote: Originally posted by Keras  
In the paper cited above the experiments must've been conducted inside an autoclave or some high-pressure-bearing system, but definitely not what we usually do, i.e. a distillation apparatus where ethanol is fed in a continuous run.

It was a sealed tube in a furnace:
Quote:
If not stated otherwise the reaction mixture was heated for eight hours in a closed tube in a Carius oven at a temperature of 155°-160°.


I shared the paper because of the data set. Maybe a strating point for an alternative procedure.

Edit: I'll leave the typo as it is, as we would be building a new method on top of the previous ones, just like the geological strata are formed.

[Edited on 17-5-2024 by bnull]




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[*] posted on 17-5-2024 at 06:15


Could evaporative loss of ether be reduced by using a smaller receiving flask? It seems to me that the rate of evaporation is directly affected by how much ether is exposed to air. As well, an ether/air mixture is a fuel air explosive where the fuel's flash point is below the temperature of the receiving flask. A smaller vessel means a smaller explosion.

Love the "where does the sulfur go?" discussion. Byproducts are products too.

GREAT POST.

[Edited on 17-5-2024 by digga]
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[*] posted on 18-5-2024 at 00:57


Great experiment !!!
I had also similar experience when synthesizing 1,4 dioxane from ethylene glycol, I reduced the amount of H2SO4 about thrice, the reaction was smooth, yield good, side reactions reduced https://www.sciencemadness.org/whisper/viewthread.php?tid=65...
The vapor escaping from reaction flask carries out some unreacted ethanol. I wonder whether adding an intermediate reflux condenser (Liebig type) with cooling liquid about 50 C (which could condense part of the unreacted ethanol but not the diethylether) could improve efficiency according the ethanol used. Then only dietheylether enriched and unreacted ethanol depleted vapor could leave this intermediate condenser and enter the final spiral condenser - some unreacted ethanol is returned back into the reaction flask (from the reflux Liebig condenser with 50 C cooling liquid, its bottom joint connected to the reaction flask and its upper joint connected to the final spiral condenser).
In ether synthesis they add some sand into the reaction flask. Could someone explain me the role of the sand? In the post by len1 he used a tube to deliver the etanol into the sand to the bottom of the reaction flask instead dripping onto the reaction surface - so the role of the sand is to prevent ethanol boiling out unreacted as b.p. of ethanol is much lower (78 C) than reaction temperature (140 C) ? Len1 wrote that in his experiment only 12% of the ethanol passed unreacted from the reaction flask. Diachrynic used magnetic stirring that could have similar effect - quickly mixing dripped ethanol into the reaction mixture thus reducing its evaporation from the surface.
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[*] posted on 18-5-2024 at 11:24


Quote: Originally posted by unionised  
I'm fairly sure that adding a concentrated aqueous solution of sodium bisulphate to ethanol gives a precipitate of sodium sulphate and a solution of sulphuric acid. I doubt the reaction goes to completion; but it might go far enough to produce something from which you can distill ether


So I tried that today. Same setup, 10 mL erlenmeyer with straight tube on top acting as air-cooled condenser. Prepared a saturated solution of sodium bisulphate, 2 mL + 3 mL ethanol added. No dice. The ethanol just boils off with a lot of bumping (since I did experiment on the back of an envelope I used my hob for heating, no magnetic stirring). I stopped after a fair amount of liquid spurted from the top of the tube.

I then had another idea: I put 3 mL of saturated sodium bisulfate solution and let it boil until there was obviously less than 1 mL left – my reasoning being that the super-concentrated solution would boil at a temperature much higher than 100 °C, and probably in the 140 °C target range. Of course I had to eyeball all this, and once I estimated that the temperature was high enough I dropped ethanol from a pipette down the air condenser. But nothing really happened. Most of the drops just evaporated off instantly, and when I flooded the erlenmeyer with, say, 2 mL of ethanol, I still had no distinctive smell of ether, despite both liquid mixing (as evidenced by a precipitate of sodium bisulphate).

The solution was so concentrated that the small magnetic stir bar I had put in the erlenmeyer in hope it would avoid bumping (despite having no magnetic stirrer) was floating!

So once more, it’s a no for me. Might try with sulphamic acid next week.
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[*] posted on 18-5-2024 at 14:24


Everything that's old is new again:
http://www.sciencemadness.org/talk/viewthread.php?tid=79548

I still think you'd get a much better yield of sulfuric acid from KHSO4, because K2SO4 is less soluble in water (12% w/w) than Na2SO4 (25% w/w) and does not form hydrates. Ideally you would cool the solution to maximize precipitation, then filter, and only then try to distill ether.




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[*] posted on 18-5-2024 at 22:07


Quote: Originally posted by clearly_not_atara  
Everything that's old is new again:
http://www.sciencemadness.org/talk/viewthread.php?tid=79548

I still think you'd get a much better yield of sulfuric acid from KHSO4, because K2SO4 is less soluble in water (12% w/w) than Na2SO4 (25% w/w) and does not form hydrates. Ideally you would cool the solution to maximize precipitation, then filter, and only then try to distill ether.


Let's assume we should use potassium bisulphate. Can it be made by reacting sodium bisulphate with potassium chloride?

[EDIT] I will try the IPA method described by Tjerk using 99+% IPA. This might lead to reasonably fairly concentrated sulphuric acid.

[Edited on 19-5-2024 by Keras]
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[*] posted on 19-5-2024 at 01:48


Consider the major wanted and unwanted reactions and unreactions involved.
The wanted reactions:

  1. 2C2H5OH+HX=(C2H5)2O+H3O++X-
  2. (C2H5)2O distilling over (neat bp 35)
  3. H2O distilling over (neat bp 100)

The unwanted reactions and unreactions:

  1. C2H5OH distilling over unreacted (neat bp 78)
  2. HX distilling over (depends on HX identity)
  3. C2H5OH+HX=C2H5X+H2O. A side route if followed by C2H5OH+C2H5X=(C2H5)2O+HX. A side reaction if C2H5X distils over
  4. C2H5OH+HX=C2H4+H3O++X-, with C2H4 distilling over (neat bp -104)
  5. 2C2H4=C4H8, and followups to tars
  6. Reactions altering the X, such as reducing X

So how do you choose X and other reaction conditions to minimize all the unwanted reactions and unreactions?
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[*] posted on 19-5-2024 at 11:04


Quote: Originally posted by chornedsnorkack  
The unwanted reactions and unreactions:

  1. C2H5OH distilling over unreacted (neat bp 78)
  2. HX distilling over (depends on HX identity)
  3. C2H5OH+HX=C2H5X+H2O. A side route if followed by C2H5OH+C2H5X=(C2H5)2O+HX. A side reaction if C2H5X distils over
  4. C2H5OH+HX=C2H4+H3O++X-, with C2H4 distilling over (neat bp -104)
  5. 2C2H4=C4H8, and followups to tars
  6. Reactions altering the X, such as reducing X

A is unavoidable since the reaction proceeds at ~140 °C, and ethanol can always be recovered and reused in another batch. B only happens if HX forms an azeotrope with boiling point lesser than or equal to 140 °C, or if it is gaseous at that temperature. I'm not sure about C (probably because every time I see an X in a chemical equation I think of halogen; blame it on the books), but it could happen in case a volatile compound were formed with ethanol; boric acid, for example, forms volatile esters with some alcohols. D happens much above 140 °C, so temperature must be not be much higher than 140 °C. E may be a problem at temperatures above 140 °C. F is sure a problem. The important points really are B and F.

Quote: Originally posted by chornedsnorkack  

So how do you choose X and other reaction conditions to minimize all the unwanted reactions and unreactions?

The formation of ether by acid catalysis involves the protonation of one molecule of ethanol, a nucleophilic substitution (SN2), and a deprotonation. So HX must not be a good oxidizer under the conditions of the distillation (F), must be a fixed substance (B), and a good proton donor. The temperature must be well controlled because of side reactions and such.

* * *


I suppose that carborane acid (H(CHB11Cl11) or H(CHB11F11), or the whole class) could work, maybe at lower temperatures. It's more of a guess than researched material, of course, and I'm aware that it is not easily obtainable by amateurs or hobbyists. Carborane acid is the strongest known acid (C. A. Reed, 'Carborane acids. New "strong yet gentle" acids for organic and inorganic chemistry'), capable of protonating CO2 (S. Cummings, H. P. Hratchian and C. A. Reed, "The Strongest Acid: Protonation of Carbon Dioxide") and alkanes (M. Nava et al. "The Strongest Brønsted Acid: Protonation of Alkanes by H(CHB11F11) at Room Temperature"). It most probably protonates ethanol.

Edit: Corrected a typo (again), added one more reference.

[Edited on 19-5-2024 by bnull]




Quod scripsi, scripsi.

B. N. Ull

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[*] posted on 19-5-2024 at 11:26


Phosphoric acid seems to be the obvious alternative to prevent reduction of the acid from being a problem.
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