Sciencemadness Discussion Board - Diethylether (sulfuric acid saving method)

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.

Precipitates - 17-5-2024 at 05:02

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.

bnull - 17-5-2024 at 05:11

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

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]

digga - 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]

Fery - 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.

Keras - 18-5-2024 at 11:24

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.

clearly_not_atara - 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.

Keras - 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]

chornedsnorkack - 19-5-2024 at 01:48

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

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

The unwanted reactions and unreactions:

C₂H₅OH distilling over unreacted (neat bp 78)
HX distilling over (depends on HX identity)
C₂H₅OH+HX=C₂H₅X+H₂O. A side route if followed by C₂H₅OH+C₂H₅X=(C₂H₅)₂O+HX. A side reaction if C₂H₅X distils over
C₂H₅OH+HX=C₂H₄+H₃O⁺+X^-, with C₂H₄ distilling over (neat bp -104)
2C₂H₄=C₄H₈, and followups to tars
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?

bnull - 19-5-2024 at 11:04

The unwanted reactions and unreactions:

C₂H₅OH distilling over unreacted (neat bp 78)
HX distilling over (depends on HX identity)
C₂H₅OH+HX=C₂H₅X+H₂O. A side route if followed by C₂H₅OH+C₂H₅X=(C₂H₅)₂O+HX. A side reaction if C₂H₅X distils over
C₂H₅OH+HX=C₂H₄+H₃O⁺+X^-, with C₂H₄ distilling over (neat bp -104)
2C₂H₄=C₄H₈, and followups to tars
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.

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]

Jenks - 19-5-2024 at 11:26

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

Keras - 19-5-2024 at 12:12

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

I already tried 75% phosphoric acid, but ethanol just boils unchanged from it.

Jenks - 19-5-2024 at 18:51

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

I already tried 75% phosphoric acid, but ethanol just boils unchanged from it.

I'm surprised. This is useful to know. Maybe the dehydrating ability of sulfuric acid is driving the reaction. But Diachrynic wrote, "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." If the water is removed by distilling it as the azeotrope, the sulfuric acid is not needed for this. Maybe the ester intermediates are the key, and the mechanism is not as simple as protonation of ethanol.

[Edited on 20-5-2024 by Jenks]

[Edited on 20-5-2024 by Jenks]

Keras - 19-5-2024 at 23:31

I already tried 75% phosphoric acid, but ethanol just boils unchanged from it.

Yeah, there is obviously something specific about sulphuric acid.
I will try again in the following days and keep you posted.

bnull - 20-5-2024 at 05:04

If the water is removed by distilling it as the azeotrope, the sulfuric acid is not needed for this. Maybe the ester intermediates are the key, and the mechanism is not as simple as protonation of ethanol.

Protonation of ethanol is the crucial step (see, for example, https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Or...); otherwise there is no reaction. Sulfuric acid is needed because it is a strong acid and consequently a good proton donor, and is also non-volatile. Other common strong acids (HCl, HNO₃) are volatile and boil off well before the reaction temperature. Phosphoric acid is fixed but weak.

There are no ester intermediates. If that were indeed the case, phosphoric acid would work. Even boric acid would do.

[Edited on 20-5-2024 by bnull]

chornedsnorkack - 20-5-2024 at 09:19

Indeed. Compare the boiling point of 85% phosphoric and 85% sulphuric acid. They have the same molar mass - both are 50% by mole. The boiling point of 85% H₃PO₄ is quoted as 158 degrees - cannot quickly find a quote for 85% H₂SO₄ boiling point but from graphs looks around 215 degrees. Equimolar amount of sulphuric acid protonates water much more fully that phosphoric acid.
Strong monobasic acids like HClO₄, CF₃SO₃H, ClSO₃H and FSO₃H are more volatile themselves because they have less bonds for hydrogen bond network. (And ClSO₃H and FSO₃H hydrolyze in water anyway). The one acid whose strength is close (slightly lower) and volatility similarly small is H₂SeO₄, and it is more oxidizing.

RU_KLO - 22-5-2024 at 05:10

Quote: Originally posted by clearly_not_atara

Quote:

The best catalyst was ferric sulfate

A line from the paper I won't soon forget

I want to test the ether "pipe" procedure.

This is the main idea:

For this Im going to use a 1 inch threaded gas pipe (its used for Home gas instalation)
(check picture)

96º EtOH + Fe2SO4 will be poured inside the pipe, screwed (with teflon tape).
An then put in a oil bath at 150ºC (it will be put at ambient temperature and then heated)

after 8 hours, it will be removed from the bath, cooled (ice water), filtered and meassured.

Now the problems:
(I dont want to have a solvent granade in hot oil....)

1) the pipe:
I meassured 1mm thickness at the thread root. (from the rust I think they are made of iron, but regulations states: it should be form steel and test are 50 bars at room temp for 5 s. and for 1 inch pipe, thickness 2.9mm)
The thread is another problem. Checked without teflon tape (for volume with water) and it was like a wathering can.
From what I read Teflon will hold on 150C. Also maybe I will use high temp silicon seal (up to 240ºC)

Could the pressure be calculated? 150ºC + 70 ml EtOH (some converted to ether - lets hope - maybe some ethylene

Here Im in a field I dont know, so if mistakes are made, please correct.

form this calculator:
http://ddbonline.ddbst.com/AntoineCalculation/AntoineCalcula...

Temperature [°C] Pressure(1) [mmHg] Pressure(2) [mmHg]
150 T > Tmax 7355.67

from google:
7355 mmHg -> 9.8 Bar

so if this should pass a 50 bar test (although 5 s), I think its safe.

Will it hold?

If there is a leak because bubbles are seen from the thread.(because of pressure)
How will react ether/ethanol gas to 150C vegetable oil? (ignite? boil off? dissolve?)

Does someone could share broken glassware - high flamable solvent(if ether better) - oil bath experience, to expect the worst?
How it is mitigated?

thanks

pipe.jpg - 86kB

bnull - 22-5-2024 at 06:13

Quote: Originally posted by RU_KLO

Good Lord! I'm glad you asked before proceeding. Ferric sulfate (Fe₂(SO₄)₃) will corrode the steel pipe. The reaction, as far as I can see, is $$Fe_2(SO_4)_3+Fe^0\rightarrow3FeSO_4.$$ There's a big chance the threads will corrode from the inside to the end and you know what happens. Also, the sealed or closed tube is a glass tube with thick walls into which the substances are poured and the open end is sealed in the same manner as an ampoule. Then the tube is put inside a steel pipe or something similar to contain an eventual explosion.

Consult references on the sealed tube or Carius tube or the pipe bomb technique before actually trying. It seems there are plenty in the old books in the Library; selaed tube techniques are older than Quantum Physics.

As I wrote somewhere else, you can always buy new equipment but can't buy new fingers. I may be--and probably am--overestimating the risks. But safety is safety.

RU_KLO - 22-5-2024 at 06:22

Good Lord! I'm glad you asked before proceeding. Ferric sulfate (Fe₂(SO₄)₃) will corrode the steel pipe.

Do you think that 10 gr Fe2(SO4)3 in 70ml EtOH will make a hole (or compromise seal) in a 1mm steel pipe 8hs 150C?

bnull - 22-5-2024 at 06:27

Yep.

Keras - 22-5-2024 at 11:13

So I decided to give a go at the standard procedure using a less concentrated sulphuric acid. I took a bottle of 37% battery sulphuric acid, distilled it at the max temp. of my hot plate/stirrer (280 °C, but I suppose the liquid was not that hot), got about half the water out. So that should’ve left me with 74% acid. Let's say 65% is more reasonable.

I used that ~ 15 mL to try get ether from 10 mL of ethanol. Definitely I got ether, judging by the smell. But the distillation took forever and what I collected was ether in water, not water in ether. Nothing useable. Re-distillation of the product yielded nothing, no 36 °C perceptible stop, barely a small notch around 50 °C.

So, what seems to happen when you don’t use concentrated acid able to pull out water from the mix is that you have to face Le Châtelier's law. Since ether creation also makes one molecule of water, any water present drives the reaction backwards. Concentrated sulphuric acid both serves as a source of proton for protonating ethanol AND as a water absorbant. Once its dehydrating properties are lost (~85% ???) the reaction is much less vigorous.

Rainwater - 22-5-2024 at 17:02

When doing the pipe bomb method with other experiments i install a small brazed tube with lots of length, . I like the 1/8ID soft copper tubing. Connected to one of the caps, it allows you to have a pressure gauge and valve to collect products and monitor the reaction.
Forming a gas product inside a sealed reaction vessel makes a positive feedback loop,
increase in pressure, which increases temperature, which increases reaction speed, which increases pressure which goes boom. This process can also be illegal in some states without a license and the reactor may also require state certification. A 1in tube can pack a serious punch, you may wish to start on a smaller scale. The conversion from liquid water to steam will be in the ball park of 1ml_liquid = 1700ml_gas not sure how to calculate what your numbers will be, google the ideal gas laws if you want to figure out the pressure. It will be a lot.
Once you have accurate numbers, factor in a 80-120% safety margin and you can safely build a reactor.

Edit:
If all else fails, do your experiments in a hole in the ground. Ensure it is deep enough so the top of your appratus forms an angle of inclination that will limit the fallout to an acceptable area.

[Edited on 23-5-2024 by Rainwater]

clearly_not_atara - 23-5-2024 at 05:41

If you want to use a phosphoric-based catalyst for this rxn you should dehydrate the phosphoric acid to polyphosphoric acid first and then add ethanol, clearly 85% H3PO4 will not work but 100% H4P2O7 just might.

you can always buy new equipment but can't buy new fingers.

ding ding ding

Keras - 23-5-2024 at 11:12

I’ll try to make 100% phosphoric acid with, say, 25 mL of my 75% bottle and add phosphorus pentoxide.

If I’m not mistaken: 3 H₂O + P₂O₅ → 2 H₃PO₄.

So, d = 1.58. 25 mL are therefore ca. 40 g. So 10 g of water, M = 18, 0.67 mol / 3 = 0.222 mol of phosphorus pentoxide, M = 142 g/mol, that makes 31.5 g.

Still quite a lot into a measly 25 mL.

clearly_not_atara - 23-5-2024 at 14:09

I’ll try to make 100% phosphoric acid with, say, 25 mL of my 75% bottle and add phosphorus pentoxide.

If I’m not mistaken: 3 H₂O + P₂O₅ → 2 H₃PO₄.

So, d = 1.58. 25 mL are therefore ca. 40 g. So 10 g of water, M = 18, 0.67 mol / 3 = 0.222 mol of phosphorus pentoxide, M = 142 g/mol, that makes 31.5 g.

Still quite a lot into a measly 25 mL.

I don't think this will be nearly strong enough! The diphosphoric acid H4P2O7 is an interesting reference point because it forms a pure crystal with a relatively high mp while solutions of polyphosphoric acids which are more or less concentrated tend to form thick gels at rt. Pure H3PO4 essentially does not exist; it disproportionates into a mixture of water, phosphoric acid and polyphosphoric acids (though the majority of this solution is H3PO4). 85% is the highest concentration of H3PO4 which can be achieved that does not contain significant amounts of H4P2O7. See e.g.:

https://pubs.acs.org/doi/pdf/10.1021/i260053a014

"high-purity, crystalline, free-flowing pyrophosphoric acid" sounds much nicer than the usual sticky crap

[Edited on 23-5-2024 by clearly_not_atara]

chornedsnorkack - 23-5-2024 at 14:17

Quote: Originally posted by clearly_not_atara

I gather that pure H₃PO₄ does exist, because the disproportionation you mention is kinetically slow. This is why H₃PO₄, like S, has two melting points, too...

clearly_not_atara - 23-5-2024 at 14:28

I see you've put in time in the Central Bureaucracy

Mateo_swe - 24-5-2024 at 01:51

According to the included paper alcohol + anhydrous ferric sulphate is most effective with a yield of 75% of ether.
"Three cm3 alcohol and 13 cm3 ether were obtained from 4 g ferric sulphate and 20 cm3 alcohol (96 %). The yield was 75 %."
Its an old but intresting paper.

Only problem is the procedure,
"the reaction mixture was heated for eight hours in a closed tube in a Carius oven at a temperature of
155°C- 160°C. When cooled down, the tube was opened, the mixture fractionated by means of a Vigreux column."

Will this work in a non closed vessel like ordinary glassware with reflux condenser?

Attachment: J. van Alphen - The formation of ether from alcohol.pdf (428kB)
This file has been downloaded 67 times

Rainwater - 24-5-2024 at 10:08

Quote: Originally posted by Mateo_swe

Will this work in a non closed vessel like ordinary glassware with reflux condenser?

Likely but controlling the losses do to evaporation will be harder, the difference in pressure will have an effect of the reaction kinetics.

A simple pressure estimate for a closed vessel.
Couldnt find a one stop source for these values https://www.engineeringtoolbox.com/
Ethanol at 160c 10-15bar or 217psi
https://en.m.wikipedia.org/wiki/Diethyl_ether_(data_page)
Diethyl ether @ 156c 15200 mmHg or 295psi
So the regular black steel pipe will hold that pressure at that temperature(7500psi 1" sch 40). But as you noted before the caps are the weak point. A welded fitting would exceed the pipe rating.
A threaded fitting, standard at best will be 300psi at room temperature, then downgraded 5% per 10c so your looking at
150c-30c = 120
Every 10 degrees derate 5%, so thats 12 derateings
So 300_{start pressure} * (0.95_{derated factor}^(12_{# of deratings})) drum roll.... 162psi before applying a safety factor which would knock it down to abou 130psi.
Im going to go ahead and say this can not be safelydone with standard fittings from the hardware store. You can order high pressure/tempature fittings what will easily handle 1000+psi. They cost about 4x as much, but for a simple 1in reactor your looking at $40 + S&H

[Edited on 24-5-2024 by Rainwater]

RU_KLO - 27-5-2024 at 09:26

Quote: Originally posted by Rainwater

Quote: Originally posted by Mateo_swe

Will this work in a non closed vessel like ordinary glassware with reflux condenser?

What about a CO2 Fire extinguisher tube?
"CO2 gas is held in the fire extinguisher under pressure, normally around 55 bars at room temperature." (i.e: 797.708 Psi)

NFPA 10 :
"8.6.3.2 Dry chemical, dty powder, water, foam, and halogenated
agent discharge hose assemblies requiring a hydrostatic
pressure test shall be tested at 300 psi (2068 kPa) or at service
pressure, whichever is higher.
8.6.3.3 Low-pressure accessor·y hose used on wheeled extinguishers
shall be tested at 300 psi (2068 kPa) .
8.6.3.4 High-pressure accessory hose used on wheeled extinguishers
shall be tested at 3000 psi (20.68 MPa).
"

RU_KLO - 6-6-2024 at 08:47

recovering H2SO4 from Diethyl ether making.

It there a possibility to get/recover any usable H2SO4 from the "tar" of an classic H2SO4 - ethanol ether syntesis?

unionised - 6-6-2024 at 09:35

Can I just check?
Are you planning to cause a bang by allowing the ferric sulphate to attack the steel of the fire extinguisher?
There also seemed to be some talk of using pyrophosphoric acid, which not only (probably) won't work but attacks steel and glass and, for extra fun...

... makes the nerve as substitute tetraethyl pyrophosphate.

I suspect the yield would be bad.

But not bad enough.

https://en.wikipedia.org/wiki/Tetraethyl_pyrophosphate

BAV Chem - 7-6-2024 at 01:09

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

I already tried 75% phosphoric acid, but ethanol just boils unchanged from it.

Making ether using phosphoric acid is actually possible and i did that at one point. This was more than 3 years ago so I don't remember everything. Also I didn't write anything down back then so the following might not be the complete truth.

The acid I used back then was mostly polyphosphoric acids that I made by boiling it down as far as i could using my heating mantle (yes, I etched the flask a couple times). Making ether with that stuff doesn't work very well though. It starts out by first boiling off a bunch of ethanol which is probably because phosphoric acid doesn't have nearly as great an affinity for ethanol as sulfuric acid would. Somewhere at or above 140C (if memory serves) there was some ether produced but this didn't last long before all the ethanol was gone. On addition of more ethanol the mix boiled and splashed violently and most of the enthanol just distilled off. Pretty soon the production of ether was very slow and I think this was because most of the polyphosphoric acid was already hydrolyzed to a critical degree. Upon heating the mix way beyond 150C a whole load of ethylene was generated and the acid slowly lost its water again. After that point I think I tried adding more ethanol and doing the reaction at 140C as before and that did afford some more ether but the total yield was still terrible. I ran something like 500ml of ethanol through it and probably spent two days on it and got around 50ml of ether. I likely made more ethylene than ether and all in all it absolutely sucked in comparison to using sulfuric acid.

Keras - 7-6-2024 at 02:43

I have just tried to produce ether with the 100% phosphoric acid I made by dissolving the stoichiometric quantity of phosphorus pentoxide into 75% commercial phosphoric acid and I got exactly the same thing (picture). Most of what distills above 140+ (nothing does under) is ethanol, at ~75 °C. There is definitely ether, because the ethanol strongly smells of it. But when I’m redistilling the product (which I’m currently doing), everything passes at 75 °C+. Probably phosphoric acid doesn’t bind with ethanol, or something of the same ilk. It was promising though, because when I added an initial 5 mL of ethanol into the 5 mL of phosphoric acid, the solution warmed up. The phosphoric acid remained clear all along, there is no charring or tar like what we get with sulphuric acid (second picture).

I'd say that what is left after redistillation (I cut at 80 °C) is… smelling like pure ether, except it can’t be. I got strictly nothing at 36 °C, the solution was not even boiling.

Is there any test I can carry out to identify the ether contents, apart from smelling?

PS: If I remember correctly, ether is industrially made by passing ethanol vapours over solid phosphoric acid. Very different conditions.

[Edited on 7-6-2024 by Keras]

What have we got so far

bnull - 7-6-2024 at 05:50

I thought it would be good to condense what has been discussed, what works and what does not.

Sulfuric acid is the classical catalyst, although it degrades and side reactions occur. The temperature must be right, around 140 °C; too low and either ethanol just boils off or the yield of ether is very poor; too high and you'll be losing ethanol as ethylene. Not to mention The Tar.
The formation of ether using sulfuric acid starts with protonation of ethanol. It is the first and most crucial step. No esters are involved since they appear because of the side reactions. Sulfuric acid is both a good protonator and a good dehydrating agent and that's why it works. Other mechanisms are possible with other catalysts but, since we're dealing with sulfuric acid, protonation is still the first step.
Sulfuric acid is also fixed (non volatile), it doesn't distill along with the products.
Phosphoric acid and its cousins are fixed but not strong enough to be proposed as substitutes for sulfuric acid, and the yield is terrible. Other stong acids are either volatile (HCl, HNO₃) or good oxidisers (H₂SeO₄, HClO₄), or both. You must avoid the oxidisers at all costs because you're interested in ether, not aldehydes or carboxylic acids or "chemically-promoted relocation of lab equipment".
The reaction may be carried out in sealed tubes, which are thick-walled glass tubes, other materials being attacked by hot acids. Other substances work as catalysts, such as ferric sulfate (I guess I know why but I don't want to make an ass of myself again without a good reason). The tubes are sealed because of the high pressures needed. A metal container, such as a capped pipe section or a fire extinguisher, will be corroded; if you're lucky, the result is a sludge of iron oxides and sulfates in ethanol with perhaps a smidgen of ether inside a pitted container; if you're out of luck, mayhem befalls.
The side reactions range from obnoxious (The Tar) to wasteful (ethylene) and, in the case of pyrophosphoric acid, TEPP and a wink from Fritz Haber.

Am I missing something?

If we're to keep using sulfuric acid, so far the best catalyst in ether production, we'll need an additive. I list here some desirable characteristics of the additive. It is not an exhaustive list. If someone else has any other characteristic to suggest, please do.

It must not react with ethanol, unless the product formed is in the pathway of ether formation.
It must be stable in the presence of hot sulfuric acid.
It must not react with sulfuric acid, unless the product is in the pathway of ether formation.
It must be fixed so we don't need to separate it from the distillate.
It must prevent or at least reduce the oxidation of ethanol by sulfuric acid.
It must not form volatile esters with ethanol.
It may be either solid or liquid. If liquid, it should be completely miscible; the reaction flask looking like a lava lamp could be a problem.

On the other hand, if we had two substances A and B which react with ethanol and among themselves, it could be a solution. Let's suppose that A reacts with ethanol to form compound EtOA^@, and B reacts with ethanol to form compound EtB^$. If we then add EtOA^@ to EtB^$, the products are ether and A^@B^$, being ether volatile and A^@B^$ fixed. A bonus would be if we could regenerate A and B from the latter, perhaps by electrolysis or reactions involving less regulated reagents. A Williamson synthesis, for example, uses sodium as A and a halogen as B. But it is more cumbersome and expensive (and probably dangerous) than simply boiling ethanol in sulfuric acid. Again, if someone can suggest a pair (A, B) of substances, please do.

RU_KLO - 7-6-2024 at 05:59

Can I just check?
Are you planning to cause a bang by allowing the ferric sulphate to attack the steel of the fire extinguisher?

No, no. It was more like an "academic" question. After Bnulls post, understand it.

The problem is that I need some ether for inorganic analytical chemistry, and also concentrated sulfuric acid. Both regents are difficult to get in my country, they are restricted, special costly permits, you need an chemical engineer to sign the papers, etc. (Of course concentrated sulfuric acid is way more needed than ether.)

So any process for getting ether without sulfuric acid, is welcome.

On searching found that: Diethyl Ether and Sodium Chloride are formed when Sodium Ethoxide and Chloroethane react (Williamson Ether Synthesis)

Sodium Ethoxide: NaOH + ethanol
Chloroethane : HCl + ethanol

(Chloroethane: Toxic/flammable -

Chloroethane is produced by hydrochlorination of ethylene:
C2H4 + HCl → C2H5Cl
At various times in the past, chloroethane has also been produced from ethanol and hydrochloric acid, from ethane and chlorine, or from ethanol and phosphorus trichloride, but these routes are no longer economical.

Is this route viable for amateur (or is very dangerous - more dangerous than ethanol sulfuric acid?

found this in the forum:
http://www.sciencemadness.org/talk/viewthread.php?tid=79850&...

From PrepChem, with references:

Preparation of ethyl chloride:

To the round bottom flask fitted with properly cooled reflux condenser and gas inlet 100 g of ethyl alcohol and 50 g of anhydrous zinc chloride are placed. The top of reflux condenser is connected to two washing bottles. The first has water and the second concentrated sulfuric acid. Finally apparatus is connected to a cooled flask for condensing the reaction product. A dry stream of hydrogen chloride gas is passed through the boiling mixture. The vapors of ethyl chloride is washed with water, concentrated sulfuric and condensed in a flask which is cooled with freezing mixture of ice and salt. As ethyl chloride boils at 12° C it must be kept in sealed glass tubes. The yield is almost quantitative.

Preparation of organic compounds, E. de. Barry Barnett, 71, 1912

So its not an easy task.... (instead of the second H2SO4 washing flask, sodium ethoxide would be used...)

[Edited on 7-6-2024 by RU_KLO]

bnull - 7-6-2024 at 09:43

Quote: Originally posted by RU_KLO

So its not an easy task.... (instead of the second H2SO4 washing flask, sodium ethoxide would be used...)

The first flask is to absorb unreacted ethanol and hydrogen chloride. The second is to absorb water. Without the second flask, the sodium ethoxide will decompose.

I bought some 20 years ago (I'm quite old, you see) a small bottle of ether:ethanol 50:50. It was sold to remove surgical tapes (because, if you try to remove them, depending on the way you pull them and where they are stuck, you take out a nice patch of skin still glued to the tape; not much fun and hurts like hell meanwhile and afterwards. And yes, that was the time when I set my hand on fire accidentally). It's possible you can find the same kind of solution in a drugstore, a pharmacy, a chemist's (English is a mess), or a supplier of surgical materials, only remaining the question of separation of ether and ethanol.

Keras - 7-6-2024 at 11:03

As far as I know, Williamson’s ether synthesis has mediocre yields in the best cases.
(See anisole by ChemPlayer or NileRed)
Might be better with sodium ethoxide, which must be a stronger base than sodium phenoxide.

RU_KLO - 7-6-2024 at 13:48

As far as I know, Williamson’s ether synthesis has mediocre yields in the best cases.

To help mitigate this issue microwave-enhanced technology is now being utilized to speed up the reaction times for reactions such as the Williamson ether synthesis. This technology has transformed reaction times that required reflux of at least 1.5 hours to a quick 10-minute microwave run at 130 °C and this has increased the yield of ether synthesized from a range of 6-29% to 20-55% (data was compiled from several different lab sections incorporating the technology in their syntheses)

Baar, Marsha R.; Falcone, Danielle; Gordon, Christopher (2010). Microwave-Enhanced Organic Syntheses for the Undergraduate Laboratory: Diels−Alder Cycloaddition, Wittig Reaction, and Williamson Ether Synthesis. Journal of Chemical Education, 87(1), 84–86. doi:10.1021/ed800001x

as Keras said, 30% is not much for all the trouble...
Also putting flammable liquids into a microwave doesnt seem to safe...

Attachment: ed800001x.pdf (611kB)
This file has been downloaded 78 times

Keras - 8-6-2024 at 00:20

Quote: Originally posted by RU_KLO

This technology has transformed reaction times that required reflux of at least 1.5 hours to a quick 10-minute microwave run at 130 °C and this has increased the yield of ether synthesized from a range of 6-29% to 20-55% (data was compiled from several different lab sections incorporating the technology in their syntheses)

Thanks for the pointer. The yield remains moderate at best and, above all, I don’t know about any commercial (kitchen) microwave oven that allows to precisely control the temperature of the heated compound inside – and I don’t think that the common way to regulate the average power, i.e. duty cycling, is really adapted vs. a continuous power-regulated microwave source.

chornedsnorkack - 10-6-2024 at 08:39

How much heat does solvation of ethanol achieve by itself, and at which concentration?
Water gives about 160 Celsius around 85 % sulphuric acid. Ethanol has a bigger molar heat capacity than water (if lower heat capacity by mass).
Also sulphuric acid/water solution reaches boiling point of about 165 Celsius, which is 65 degrees above pure water, at 70% by mass... which is about 2,3 moles of water per mole of sulphuric acid. At what concentration can ethanol/sulphuric acid solution be heated to 140 Celsius (62 degrees above pure ethanol) without boiling?

[Edited on 10-6-2024 by chornedsnorkack]

RU_KLO - 12-6-2024 at 05:25

How much heat does solvation of ethanol achieve by itself, and at which concentration?
Water gives about 160 Celsius around 85 % sulphuric acid. Ethanol has a bigger molar heat capacity than water (if lower heat capacity by mass).
Also sulphuric acid/water solution reaches boiling point of about 165 Celsius, which is 65 degrees above pure water, at 70% by mass... which is about 2,3 moles of water per mole of sulphuric acid. At what concentration can ethanol/sulphuric acid solution be heated to 140 Celsius (62 degrees above pure ethanol) without boiling?

[Edited on 10-6-2024 by chornedsnorkack]

I think is more complicated than that.

Found this, but under reduced pressure:

"For the sulfuric acid + water + ethanol system, the
boiling points increase with decreasing mole fractions of
ethanol in the solution at fixed ratios of sulfuric acid/water,
and they also increase as the ratio of sulfuric acid/water is
increased."

But for the same paper, it seems that Sulfuric acid could be recovered at lower temperature with butyl acetate (Acid + Water + Butyl Acetate + Ethanol Quaternary System

how difficult is to get/make Butyl acetate? From wikipedia: It is a component of fingernail polish.[8]
(I found ethyl acetate un fingernail polish, not butyk

the paper talks about models, but can someone "translate" this for non engineer people? in the sense, how do you apply it and what are the benefits.
Does the sulfuric acid destilates with the butyl acetate (and how to separate it) or the butyl acetates remains in the boiling flask with water?

Determination and Modeling of Vapor–Liquid Equilibria for the Sulfuric Acid + Water + Butyl Acetate + Ethanol System
Geng Li, Zhibao Li*, and Edouard Asselin

https://pubs.acs.org/doi/10.1021/ie303197a

Attachment: ie303197a.pdf (1MB)
This file has been downloaded 75 times

jackchem2001 - 13-6-2024 at 07:19

After experimenting with this procedure a little bit I would like to put something into text that is probably quite obvious - you can't just feed all of the distillate back into boiling flask and redistill to try enrich it in ether. You just end up fractionating the mixture. The matching of the addition rate and distillation rate is done because a high concentration of sulfuric acid raises the boiling point of the mixture (a likely requirement for a different catalyst in a simple distillation setup).

Unrelated, but has anybody else observed white smoke upon addition of ethanol to the boiling flask? Also, even though this contradicts the pdf posted on page 1 which casts doubt on alkylation agents as intermediates in this reaction, has anybody tried adding sodium sulfate to potential catalyse formation of sulfuric esters? (https://www.sciencemadness.org/talk/viewthread.php?tid=78772)

[Edited on 13-6-2024 by jackchem2001]

jackchem2001 - 3-7-2024 at 01:04

Continuing from my last post suggesting the addition of sodium sulfate to catalyse the formation of sulfuric esters (allowing a lower temperature to be used to avoid tarring), in ethyl bromide procedures diethyl ether is observed as a major impurity. Evidently, ethanol does react with some alkylating agents at much lower temperatures to give diethyl ether.

As I hate the smell of ether so I likely won't do this procedure again, but if anybody happens to be doing a run then consider adding a small amount of sodium sulfate to the boiler and report the results

MrDoctor - 5-7-2024 at 00:06

I have attempted this reaction and noticed something strange. thermal runaways.

At first i thought some sort of hydrolysis was occuring which could explain things, since i was using 95% ethanol, but i see now that water simply cannot accumilate in this reaction, which increases the mystery.

What i observed was that at seemingly arbitrary temperatures between each runaway, my PID controller would suddenly find itself unable to heat the acid solution further. it would get stuck at say, 140C. no reflux or enough distillate was comming over to explain the heat loss, i also re-attempted without a fractionating column thinking it was trapping the water leading to exothermic splitting/reformation of ethyl sulfate <-> sulfuric acid + ethanol. energy is being applied but it is not being reflected in a change of temperature, nor, any notable distillate although the head temp would begin to increase, but only after the runaway began, not leading up to it.

Suddenly, the column head temperature would skyrocket from the boiling point of ether to the boiling point of a low-ethanol +water steam, and the distillate would reflect as much. the addition of more ethanol during this runaway would instantly and increase the temperature.

eventually my acid tarred up and became so thick it would foam so i had to stop, but this phenomena occured long before any viscosity changes or clarity occured.

During a runaway, over the course of a minute, regardless of if the PID was still applying power, the temperature would suddenly rise from wherever it began, in this case, 140C, and it would rise up to around 156C. this temperature change did not correspond to any thermal-momentum, the lag between heat being generated by the mantle, and then transferring to where the thermometer could read it. it also occured in a span that almost matched how quickly the temperature would rise with a manunally set 100% duty cycle. (also yes i did verify the PID wasnt simply defective and suddenly locking on 100%)

ethanol was pumped in consistently by peri pump, shutting off automatically when the temperature had risen above 60C as i was concerned that ethanol would simply boil on the surface of the acid without reacting.

Eventually the runaways would accompany foaming that caused the ejection of the black thick acid so i had to stop. there was a major viscosity change after i tried to vacuum concentrate the acid, but this phenomena was present prior to that, i had just assumed the acid had become too diluted with water or something, since runaways would end with noticeable amounts of water in the condenser.

Does anyone know what was happening?

A successful semi-automated ether preparation and an exercise in tar chemistry

BAV Chem - 13-7-2024 at 05:53

So apparently the process described by Diachrynic does not work for everyone so far. I've done the procedure myself with a few minor changes and for me it worked out pretty much perfectly on the first try.

Setup: The reaction was carried out in basically the same way as described on page 1. To ensure a successful experiment only the highest quality reagents were used. The sulfuric acid was tar grade and recovered from previous experiments and as such it contained the boiled down remains of a hydrazine sulfate preparation (azine method) and a failed nitration mix amongst other things (Note 1). It was the nastiest sample of H2SO4 that could be found in the back of the storage shelf. This reagent's purity was determined to be a seven on the browning index as defined by Kahverengi et al.^[1]
Similarly part of the ethanol was recovered from previous experiments or distilled from some exceptionally bad booze. Most of it however was store bought denatured alcohol (95%).

Procedure: A 3-neck 500ml RBF was charged with 50ml of sulfuric acid and 50ml of ethanol (Note 2). Heating was accomplished using a bath of melted candle wax with the hot plate set to nearly 200°C surface temp. The mix was constantly stirred and ethanol was added using an arduino controlled peristaltic pump system. This was set to run for 500ms at a time, which is equivalent to adding roughly 0.3 - 0.4ml of EtOH. Additions were automatically made whenever the reaction temperature was above 142°C with a minimum dead time of 15s between additions. A type K thermocouple was used to monitor the temperature inside the reaction flask. This technique worked out flawlessly and the temperature stayed within 0,5°C of the set temperature of 142°C for the majority of the experiment, that is assuming there was no drift in the thermocouple measurements. 4000ml or 3200g of 95% ethanol were added using the pump system before the reaction was stopped. Toward the end of the experiment there was quite a large amount of tar buildup in the flask and the mixture started to foam a little. It is also important to note that even at this point there was still a considerable amount of sulfuric acid left in the reaction flask and the experiment could've easily been continued.

Workup: The distillate obtained was fractionally distilled as described on page 1. Diethyl ether was collected up until 40°C and ethanol between 75°C and 80°C. A small amount of distillate came over while the temperature rose from the boiling point of ether to that of ethanol and this was simply included in the next batch.

Acid recovery: According to Evans and Foresman^[2] about 79% of the sulfuric acid that is lost is still contained in the distillate as either sulfuric acid or diethyl sulfate. The latter of these two should at least partially hydrolyze to give the former so it was attempted to recover some of the lost acid by boiling down the residue from the fractional distillation. However this only yielded a small amount of black sludge that presumably consisted of mostly charred organics and some sulfuric acid. The corresponding browning score would have been an eight to nine.
The leftover sulfuric acid in the RBF was treated in a similar manner. It was boiled down and heated until most of the impurities were digested and the material returned to its typical viscosity. This way about 33ml of H2SO4 were recovered which was of similar quality as the starting material.

Results: In total 4050ml or 3240g of ethanol were used which was converted to 2936g or roughly 3500ml of crude distillate. Fractional distillation of the latter yielded 1175g or 1650ml of diethyl ether, 1500ml or 1200g of ethanol and 430ml of a watery residue (Note 3). With the distilled ethanol being the same concentration as the starting material this equates to a conversion efficiency of 75% (Note 4). Assuming the other 25% have been converted to ethylene the amount of water that would be expected to form can be calculated to be 473.8ml. The 430ml of H2O obtained thus equate to 91% of the expected amount so these numbers seem to make sense. 50ml of tar grade sulfuric acid were employed of which 33ml could be recovered. This puts the amount of acid that was actually lost to just 17ml. Without recovering the acid the volume of ether produced would've been 33x that of the sulfuric acid employed. Accounting for the recovered sulfuric acid this factor goes up to about 97 (Note 5). Also it seems that the quality of the acid used doesn't make too much of a difference.

Notes:
1) It is obviously not advisable to boil down nitration mixtures unless it can be done safely.
2) Interestingly the addition of ethanol to the acid led to some white fumes appearing in the reaction flask, exactly as described by jackchem2001^[3], which is strange because no vapor phase reactions should occur here. I suspect it's just ethanol being vaporized from the exotherm of mixing H2SO4 and EtOH.
3) Diethyl sulfate which is apparently a byproduct of this reaction is said to have a peppermint like odor and this residue had a very weird smell that did in fact share some similarities with that of peppermint.
4) I was hoping for something more like 90% but oh well...
5) This figure may be subject to various measurement errors and might actually be somewhat lower.

References:
1) K. Kahverengi, M. de la Mierda, R. von Braun and G. Schlonk (2021) '50 Shades of Brown: An Index of Chemical Misery', J. Immat. Sci., 1, https://jabde.com/wp-content/uploads/2021/12/50-Shades-of-Br...
2) 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/v...
3) https://www.sciencemadness.org/talk/viewthread.php?tid=78772

BAV Chem - 20-7-2024 at 15:03

We might not even need to search for an alternative catalyst although admittedly that would be a very interesting topic to look at. I recently tried my hand at the process of making sulfuric acid using ethanol and sodium bisulfate and things worked out quite well (see this thread).

tl;dr I obtained about 50ml of reasonably pure sulfuric acid from 250g of pool grade NaHSO₄ which should be more than good enough to make liters of ether.

maybe better with a hot tube and H2SO4 on pumice?

Organikum - 27-7-2024 at 05:00

lugh from the HIVE once provided a German paper which I sadly do not have anymore but I have a still well working memory and the paper told this:

Ethanol vapors (~90% EtOH was used IIRC) were passed through a glass-tube heated to 135° - 150° C and filled with pumice pebbles soaked with 70% H2SO4, and the resulting mixture of water, ethanol and ether was condensed, the ether removed and the ethanol distilled from the water and returned to the feeding vessel.
They claimed up to 50% yield on ether per run whereby the higher the temperature the higher the yield but also more crap deposition on the pumice leading to shortened standing time of the catalyst. But at lower temperatures with 35% yield they claimed next to infinite (weeks) of activity.

This is the part I am sure about, more shaky is the following:
Tube was made of glass (I am sure of this), 3 to 5 cm diameter, pumice pebbles like 2-3 mm diameter and tube-length like 50 cm heated, catalyst filled.

To avoid massive temperature fluctuations it would be for sure best to immerse the tube into a liquid heating medium or maybe metal block and to use a rather thin walled tube. I wager thinner tube beats thicker tube one might consider a bundle of thinner tubes for a production setup.

But generally it works this is sure, I know it as I tried back then it in a quick and dirty test and there is absolutely no reason why it should not and with clean alcohol a high TOF seems plausible, a small setup running for days can produce copious amounts of ether I suppose. Put a flame-arrestor between boiling flask and hot tube and this should be quite save. As far as ether can be called save of course.
The integrated ethanol recovery - well I would skip this and I also would use absolute ethanol which I can buy for ethanol ovens cheap and easy (just condense out the MEK before use, add NaOH or KOH and let sit in a warm place for some weeks and use as is, I mean why distill when it gets distilled anyways?)

It for sure saves a lot of H2SO4 and the temperature in the hot tube is reasonable.
There is some back-pressure building up in the boiling flask that is understood, the pumice i must fill the tube completely to ensure contact, so clamp joints properly down and when using an addition funnel to refill, use a pressure equalized one.
Darn! I am tempted to build one myself, ether is still my favorite solvent, well as long the floor is not covered in eeri blue flames all of a sudden.....

Hope somebody finds this interesting and there is not already a thread about it, I did not search tbh.

regards
/ORG

bnull - 27-7-2024 at 15:51

lugh from the HIVE once provided a German paper

I found it. It's a note in French (which at least I can read without Google): Jean-Baptiste Senderens, "Déshydratation catalytique, en phase gazeuse, des alcools forméniques en presence de la ponce sulfurique et phosphorique",Compt. Rend. 192, 1335-7(1931) (Source gallica.bnf.fr / BnF). It is short (3 pages) and small, so it is also attached (if someone happens to be too lazy or not curious enough to check out the Comptes Rendus).

I don't know what are the rules of Sciencemadness regarding content from The Hive but, since it is not cookery, here's the (sort of) original post.

Attachment: Senderens - Synthesis of ether using pumice stone soaked in sulfuric acid in a hot tube (more or less) (363kB)
This file has been downloaded 32 times

[Edited on 28-7-2024 by bnull]

wg48temp9 - 28-7-2024 at 03:06

Just in case, nobody has seen the papers on Silica sulfuric acid as Solid Acid Catalyst. It's used for esterification and for producing ethers.

Sorry, i could not find the paper. I will try again later

Organikum - 28-7-2024 at 05:02

lugh from the HIVE once provided a German paper

Thanks for the paper, nice you can read it, most can not so I tell the essence of it. And lo and behold! It is - so true - looking like the holy grail of ether production!

Pumice granules soaked in H2SO4 of 62% to 64% strength catalyze the dehydration of ethanol to ether at 135° C to 140 °C whereby the the dehydration is complete (if the stream of ethanol is adjusted accordingly and no side-reactions occur at 135°C, at 140° C small amounts of ethylene are formed.

They ran the reaction 8 days, 8 hours a day t 135° C with no degradation or tarring of the catalyst.
The pumice was in "nacelles" what I interpret as a bag of wire and heating was done electrically.

Now thats pretty neat, isn't it?

Oh, they also mention that sodium bisulfite does NOT work, what hopefully ends this myth once and for all.

Whats your problem with The HIVE? SCM has no limitations as long the posts are trying to follow scientific rules in wording and this refers to all posts be it energetic (physically) energetic (mentally) or washing your dishes.

regards
/ORG

PS: And for "cookery" let's quote Alexander Shulgin:
"Orbitals are for mathematicians, Organic Chemistry is for people who like to cook!"

bnull - 28-7-2024 at 09:01

Thanks for the paper, nice you can read it, most can not so I tell the essence of it.

Being a Romance speaker has its advantages. Thanks for mentioning the paper anyway.

Pumice granules soaked in H2SO4 of 62% to 64% strength catalyze the dehydration of ethanol to ether at 135° C to 140 °C whereby the the dehydration is complete (if the stream of ethanol is adjusted accordingly and no side-reactions occur at 135°C, at 140° C small amounts of ethylene are formed.

They ran the reaction 8 days, 8 hours a day t 135° C with no degradation or tarring of the catalyst.
The pumice was in "nacelles" what I interpret as a bag of wire and heating was done electrically.

Now thats pretty neat, isn't it?

Oh, they also mention that sodium bisulfite does NOT work, what hopefully ends this myth once and for all.

Sodium bisulfate, but here Senderens was talking about passing alcohol vapor over heated sodium bisulfate, which doesn't mean that other processes involving it are doomed to failure. The main issue with bisulfate is that the catalytic dehydration of vapors ocurrs from 175 °C upwards, where ethylene, not ether, is the main product. In liquid phase, ethanol boils off unchanged in a common distillation setup at ambient pressure (Senderens again) but it is quite possible that it converts to ether in a sealed tube (a thought that occurred after reading van Alphen). The latter is not amateur-friendly, so while not exactly a myth until someone tries out the sealed tube method, the necessary conditions make it more interesting from a theoretical/academical point of view than a proper preparative method of large quantities of ether for the home chemist. It seems feasible with the right equipment but not at home with lab glassware or threaded pipes. And the "nacelles en porcelaine" are simply combustion boats made of porcelain, by the way.

Anyway. The process is quite neat and can be easily automated, with Arduino and stuff. An apparatus the size of a tabletop water cooler where you insert ethanol and receive ether. With a replaceable "sulfuric pumice" cartridge perhaps. Not bad. Not a holy grail but not bad either.

Whats your problem with The HIVE?

With the chemical aspects, none at all; I love Chemistry, and science in general. As for the recreational and commercial aspects and their legal developments, I don't effing care. I don't do or cook drugs or have problems with the Law itself or its supposed representatives. I don't really effing care. The problem is, there are delusional elements within the forum that become so paranoid at the briefest mention of drugs or drug-related content and sources (most of The Hive has been archived by Erowid, you see; no problem with that too) that they begin fearing for their lives and freedom as if the "drug police" were monitoring them 24/7 and ready to fall upon them in a sec just because they had read that content. In such state, they would start spewing their unreasonable fears (check out Detritus) in this thread while being annoying and unhelpful, and cheesing me and everyone else off in the process. Hence my excessive caution, like the monkey.

chornedsnorkack - 28-7-2024 at 23:18

Is it feasible to set up reactive distillation so that "no reaction" results in nothing distilling over - unreacted ethanol just flows back to the still?

Organikum - 29-7-2024 at 06:11

Is it feasible to set up reactive distillation so that "no reaction" results in nothing distilling over - unreacted ethanol just flows back to the still?

This makes not more sense on closer inspection then it made on first glance, maybe you try sorting your thoughts and understanding the process before posting?

Organikum - 29-7-2024 at 06:21

....

Well we are talking about doable processes so bisulfite is obviously out - if you want to chase dead cows, well do it, but do not send others to do do it.

The HIVE as mentioned here shows that one does not have to be a druggie to be a paranoid fuck, this is long known, same the the bad effects drugs like LSD and Crystal Meth have on people who do not consume them but only came in contact via TV shows and YouTube.

regards
/ORG

bnull - 29-7-2024 at 09:36

Is it feasible to set up reactive distillation so that "no reaction" results in nothing distilling over - unreacted ethanol just flows back to the still?

No. As explained in the second paper by Senderens (Catalytic dehydration of alliphatic alcohols by alkaline bisulfates, in French), while a proper proportion of ethanol:sulfuric acid can be used so the solution can reach the reaction temperature (~140 °C), the same can't be done with sodium bisulfate because the increase in boiling temperature is negligible given the amounts of bisulfate used. This was verified by @Keras several posts ago. Even a paste made of bisulfate and ethanol would just dry up and the alcohol boil off unchanged.

So we're still dealing with sulfuric acid (with or without additives) and other possible substitutes at about 1 atm.

chornedsnorkack - 29-7-2024 at 11:28

Is it feasible to set up reactive distillation so that "no reaction" results in nothing distilling over - unreacted ethanol just flows back to the still?

This makes not more sense on closer inspection then it made on first glance, maybe you try sorting your thoughts and understanding the process before posting?

You have offered two options:

Entirely open apparatus - vapour is free to get out and liquid flow out at constant pressure regardless of any reaction with volume change - with free return inflow of air
Rigidly closed apparatus - glass welded closed against pressure, no change of volume till the glass violently breaks

How about apparatus which allows only a small difference of pressure between inside and outside, which the glass and joints are designed to fit? Like a liquid lock?
Since the likely vapours to reach the gas lock (ethanol, ether, water) may dissolve in water and/or oil, a liquid lock proof of these might be a metal one (Hg or Ga based).
If there were no catalyst, or if the catalyst were completely insoluble in the liquid (like a solid superacid), bringing the reaction
2C₂H₅OH=(C₂H₅)₂O+H₂O to completion (but equilibrium would not be completion), the liquid would separate, and its boiling point would drop from 78,3 degrees to 34,2 degrees of ether/water azeotrope - meaning that at constant thermostat temperature, the ether would boil off and bubble through the liquid lock.
What we want is ethanol/sulphuric acid solution of such concentration that its boiling point is 140 degrees. What concentration is that? And if the reaction goes on, how much does its vapour pressure increase, at constant temperature?

bnull - 29-7-2024 at 15:57

You have offered two options

I guess it was I who offered those options, @chornedsnorkack. And yes, it was either a distillation setup open to the atmosphere, which is the original idea of the thread and the only safe (as hot sulfuric acid goes) option so far, or a sealed tube, which seems feasible with the right equipment (i.e., professional lab or industry) but not at home with lab glassware or threaded pipes, as I pointed a page ago. Since the quantities produced by the latter are inevitably too small (< 20 mL), it remains of theoretical interest.

So, we're stuck with (1) round-bottom flasks and condensers, (2) the sulfuric pumice boat method of Senderens, or (3) the solid acid catalyst mentioned by @wg48temp9. I have a vague impression of having read about (3) in the forum; memory fails me now.

What we want is ethanol/sulphuric acid solution of such concentration that its boiling point is 140 degrees. What concentration is that? And if the reaction goes on, how much does its vapour pressure increase, at constant temperature?

Don't forget the water. There are so many compositions EtOH:H₂SO₄:H₂O with the same minimum boiling point (it is a minimum because we need the solution to reach at least the reaction temperature). I remember seeing something in the forum, possibly by you, about those mixtures. Again, memory fails me.

BAV Chem - 30-7-2024 at 01:06

Just an alkaline bisulfate mixed with ethanol doesn't catalyze the dehydration of the alcohol to ether but I would argue this is only because the bisulfate itself does not raise the boiling point enough, and is not acidic enough to catalyze the reaction at an appropriate temperature. Also there's never gonna be enough free sulfuric acid in solution to do this because the eqilibrium of NaHSO₄ <---> Na₂SO₄ usually lies too far on the left.
However if you were to filter off the sodium sulfate before doing the reaction it would likely work. I have done a similar thing but isolated the sulfuric acid instead of using it straigt away recently (see somewhere above) and it certainly seems practical. By adding something like 250g of sodium bisulfate to 500ml of ethanol, stirring, filtering and distilling it should be possible to obtain an impure but usable sulfuric acid concentrate that would be perfectly suitable for making ether. Yes, it's just the method from the first post with extra steps but it would allow us to start from bisulfate and not sulfuric acid. You could simply distill off the ethanol until the temp reaches 140°C and then add more ethanol as usual.

Is it feasible to set up reactive distillation so that "no reaction" results in nothing distilling over - unreacted ethanol just flows back to the still?

Wouldn't that be the same as a fractional distillation where the ethanol recondenses but ether distills over? Some people actually do it that way and get terrible conversion efficiencies out of their sulfuric acid because of water build up. However this brings me to a theoretically possible but frankly impractial idea. If you were to set up a distillation of sorts such that the ethanol would be returned without returning the water it should be doable. Tbh I have no good idea of how to do this. All I can think of would be a series of condensers, one running at 80-90°C to take out all the water, another one at like 60°C to take out the ethanol and one to finally condense the ether. That would basically be like building a miniature chemical plant that you would feed with ethanol and get out ether and water at the other end.

bnull - 30-7-2024 at 08:54

There's also that.

However if you were to filter off the sodium sulfate before doing the reaction it would likely work. I have done a similar thing but isolated the sulfuric acid instead of using it straigt away recently (see somewhere above) and it certainly seems practical. (...) Yes, it's just the method from the first post with extra steps but it would allow us to start from bisulfate and not sulfuric acid. You could simply distill off the ethanol until the temp reaches 140°C and then add more ethanol as usual.

The "somewhere above" is this other thread, The simplest preparation of sulfuric acid?! Another instance of my memory failing.

Tbh I have no good idea of how to do this. All I can think of would be a series of condensers, one running at 80-90°C to take out all the water, another one at like 60°C to take out the ethanol and one to finally condense the ether. That would basically be like building a miniature chemical plant that you would feed with ethanol and get out ether and water at the other end.

Something like the drawing below? The collectors look familiar...

Pumukli - 30-7-2024 at 10:12

There"s a synth description in Org. Synth somewhere (phorone? diacetone-alcohol?) in which before they let wet acetone dripping back from the reflux condenser to the flask they let it percolate through dry, anhydrous BaO in a perforated "cup" or something like that.

Something similar, maybe filled with anhydrous Na2SO4 might work with wet ethanol.

EF2000 - 31-7-2024 at 06:12

Like that: https://www.orgsyn.org/demo.aspx?prep=CV1P0199

Synthesis of diacetone alcohol. They use Soxhlet extractor with thimbles filled with Ba(OH)2.

Organikum - 8-8-2024 at 03:05