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Tsjerk
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I had another go at it, with isopropanol this time.
This article didn't give me much hope in getting the method to work with methanol.
I again made a 50 ml solution with 0.1 mol NaHSO4 and added 50 ml IPA. This time there was no immediate precipitation, but after half an hour in the
fridge there were some very nice plate like crystals.
After filtering and rinsing with a bit of IPA I dried the sodium sulfate filter cake in a microwave and found 6 grams of Na2SO4, instead of the
expected 7.1, so there must be some NaHSO4 left in the acid
I distilled the filtrate until the temperature in the head reached 100 degrees. I don't know if there was any dehydration of the alcohol, but at least
I didn't notice anything. I determined the concentration of the IPA by density and it was around 85% with a 100% recovery of IPA. I titrated the water
and the acid (in the form of H2SO4 and NaHSO4) yield was around 100%.
Next I dissolved 0.125 mol NaHSO4 in 30 ml and added the 85% alcohol, but this just formed two layers.
There might be a sweet spot in between these two conditions, but I think it would be hard to find and you will always be left with some sodium in the
acid. The recovery of the alcohol is easy though. I will try with ethanol sometime soon.
The solubility data on Wikipedia (280 g/l) is wrong. Judging on the size of the water layer in the second experiment I think this data (670 g/l) is correct.
[Edited on 26-1-2021 by Tsjerk]
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Tsjerk
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Actually the IPA/NaHSO4 method does work. The water layer I observed turned out to be a saturated sodium sulfate solution, sodium sulfate has the
property to have its maximum solubility exactly in the amount of water needed to crystallize as the decahydrate.
When I checked the beaker this morning the water layer had completely crystallized, but the crystals could easily be broken up and I didn't see a
water layer anymore. To be sure I placed the beaker in the fridge for half an hour. I filtered the mixture, which is easy to do as the crystals are
nice and big, and dried the crystals in the microwave. This gave exactly the expected amount of Na2SO4 and titration of the acid gave the expected
amount of H2SO4.
So a 30 ml solution containing 0.125 mol NaHSO4 with 58 ml 85% IPA and about ten minutes of stirring gives an easy to filter suspension of crystals
and sulfuric acid low in sodium. The IPA is available as rubbing alcohol and can easily be recycled. Once you distill off the IPA the sulfuric acid is
already around 25%.
Edit: These amounts can probably be optimized, but even with these numbers you could easily run 1 mol batches in a one liter distillation setup giving
0.5 mol of sulfuric acid each time.
[Edited on 27-1-2021 by Tsjerk]
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Fantasma4500
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@clearly_not_atara very low yields, best yields is in bark or even better leaves- or, once again better thistles. there used to be some site having
data for potassium content in different types of wood, bark, leaves and thistles- its gone. sad.
its all approximates since a plant doesnt just magically manifest potassium ions into existence, they pull it from the ground
and if the ground is a complete zero in potassium- then so will the tree
if people are interested in getting potassium, the pool chemical "caroat" is a tripple salt mix of KHSO4, K2SO4 and KHSO5, name comes from coroats
acid H2SO5 made by conc H2SO4 and H2O2, browsing about i see its less and less available, only 500g available on fleabay in europe. as for
fertilizers, we dont have your gardening center, it varies a lot. i was once able to find some calcium nitrate, impure with some nitrogenous hydrogen
molecule too, and it was bound into some sort of complex- while americans would swear that you can buy it all in 5N grade because they can at their
local kmart or whatever- and one other user in this thread mentioned that you can *just* go and buy battery acid
this is barely possible anymore, last i looked around you have to buy 20L jugs and its just .. sub20% - this will get worse no doubt. were not seeing
politicians getting busy with writing laws that promote freedom- i dont even watch TV and i know this for a fact.
acid salts and the corresponding acid can act a bit weird, i found out some years back that you can in a very specific percentage range of H2SO4
dissolve calcium sulfate, and i managed to produce a fistsized bunch of crystal cake with long crystals, it was formed as i casually diluted the acid
over time as i was using some of the acid for cleaning, we have seen same effect with NaOH and hydroxides, chromium hydroxide for instance. once the
sulfuric acid gets well concentrated this should not be an issue, i know with iron sulfate the solubility actually decreases a lot, even with fairly
low conc H2SO4, i believe they also use HCl to precipitate NaCl out of solution for making crystalline salt for cooking
very neat project in this thread, especially if you go ahead and buy bulk of pH minus, guaranteed to raise zero suspicion if the cashier is a young
woman whose only concept of acid is a psychoactive drug she heard stories about.
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BAV Chem
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This method of producing sulfuric acid might be quite laborious but it seems to scale up remarkably well and I actually had great success with it.
I started with 250g of pool grade sodium bisulfate and with heating dissolved it in 200ml of water. 300ml of ethanol were heated to near boiling and
the bisulfate solution was also brought to roughly the same temperature. The intention behind this was to use as much bisulfate with as little water
as possible. The amounts of water and ethanol were chosen with the help of [1] according to which the solubility of Na2SO4 drops
to negligible levels in ethanol water mixtures with ethanol concentrations greater than 50% w/w. Under strong stirring I added the hot ethanol to the
bisulfate solution which caused in a bunch of fine white precipitate to appear. Upon standing at room temp some more sulfate crystallized. This was
then filtered off and the solution was chilled in the freezer to -18°C. By doing so another batch of solid precipitated in the form of small
platelets which gave the solution a gelatinous consistency. These were also filtered off and washed with a small amount of EtOH. The filtrate was then
fractionally distilled to recover the ethanol and finally boiled down to concentrate the acid.
The crystallized sodium sulfate appeared to be quite hygroscopic, especially the second crop of crystals, so I suspected it still had some bisulfate
or other acidic species in it. Because of this I mixed it with 150ml of hot water (not everything dissolved), added 300ml of hot ethanol to it and
proceeded as before. This time the two crops of sulfate obtained were not hygroscopic and didn't have much of an acidic reaction towards sodium
bicarbonate. After again recovering the ethanol and boiling down some more acid was obtained.
The first run yielded roughly 50ml of sulfuric acid whereas the second one only gave something like 15ml. Obviously doing a second run doesn't
improve the yield much and is kind of pointless. The two batches of crude product were combined and concentrated further. This was done by boiling the
acid in a beaker wrapped in rock wool insulation with a round bottom flask on top. This way the remaining water boiled out until the sulfuric acid
itself started to readily reflux in the beaker. In theory this should get it up to nearly azeotropic concentration. After this about 50ml of hopefully
very concentrated acid was left, weighing 92g. Assuming this is the azeotrope at 98% the yield comes out to be 101,5% (wait what?). A quick (and
likely inaccurate) density measurement gave a density of 1,86. 98% H2SO4 has a density of 1,84 so something is off. I suspect
there's still some sodium bisulfate dissolved in the acid. According to [2] at 20°C a liter of concentrated sulfuric acid can dissolve up to 87g of
Na2SO4 which is epuivalent to 147g of NaHSO4. This means that my 50ml of sulfuric could contain as much as 7,4g of
bisulfate. Really I have no idea what my yield is on this but it seems to be upward of 80 or even 90%.
Perhaps one could improve the efficiency of a single run by using even less water, maybe even just melted NaHSO4 * H2O and/or a
little bit more EtOH. I didn't want to do the latter because i wanted the filtrate to all fit in a 500ml boiling flask.
Literature:
1) Toro, Dobrosz-Gómez & García (2014) 'Sodium sulfate solubility in (water + ethanol) mixed solvents in the presence of hydrochloric acid.
Experimental measurements and modeling' Fluid Phase Equilibria, 384(), 106–113. doi: 10.1016/j.fluid.2014.10.025
2) J. J. Stöckley, R. Bartunek (1934) 'Process for the separation of sodium sulfate from sulfuric acid', US Patent US1812310A
[Edited on 20-7-2024 by BAV Chem]
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chloric1
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Quote: Originally posted by BAV Chem  | This method of producing sulfuric acid might be quite laborious but it seems to scale up remarkably well and I actually had great success with it.
I started with 250g of pool grade sodium bisulfate and with heating dissolved it in 200ml of water. 300ml of ethanol were heated to near boiling and
the bisulfate solution was also brought to roughly the same temperature. The intention behind this was to use as much bisulfate with as little water
as possible. The amounts of water and ethanol were chosen with the help of [1] according to which the solubility of Na2SO4 drops
to negligible levels in ethanol water mixtures with ethanol concentrations greater than 50% w/w. Under strong stirring I added the hot ethanol to the
bisulfate solution which caused in a bunch of fine white precipitate to appear. Upon standing at room temp some more sulfate crystallized. This was
then filtered off and the solution was chilled in the freezer to -18°C. By doing so another batch of solid precipitated in the form of small
platelets which gave the solution a gelatinous consistency. These were also filtered off and washed with a small amount of EtOH. The filtrate was then
fractionally distilled to recover the ethanol and finally boiled down to concentrate the acid.
The crystallized sodium sulfate appeared to be quite hygroscopic, especially the second crop of crystals, so I suspected it still had some bisulfate
or other acidic species in it. Because of this I mixed it with 150ml of hot water (not everything dissolved), added 300ml of hot ethanol to it and
proceeded as before. This time the two crops of sulfate obtained were not hygroscopic and didn't have much of an acidic reaction towards sodium
bicarbonate. After again recovering the ethanol and boiling down some more acid was obtained.
The first run yielded roughly 50ml of sulfuric acid whereas the second one only gave something like 15ml. Obviously doing a second run doesn't
improve the yield much and is kind of pointless. The two batches of crude product were combined and concentrated further. This was done by boiling the
acid in a beaker wrapped in rock wool insulation with a round bottom flask on top. This way the remaining water boiled out until the sulfuric acid
itself started to readily reflux in the beaker. In theory this should get it up to nearly azeotropic concentration. After this about 50ml of hopefully
very concentrated acid was left, weighing 92g. Assuming this is the azeotrope at 98% the yield comes out to be 101,5% (wait what?). A quick (and
likely inaccurate) density measurement gave a density of 1,86. 98% H2SO4 has a density of 1,84 so something is off. I suspect
there's still some sodium bisulfate dissolved in the acid. According to [2] at 20°C a liter of concentrated sulfuric acid can dissolve up to 87g of
Na2SO4 which is epuivalent to 147g of NaHSO4. This means that my 50ml of sulfuric could contain as much as 7,4g of
bisulfate. Really I have no idea what my yield is on this but it seems to be upward of 80 or even 90%.
Perhaps one could improve the efficiency of a single run by using even less water, maybe even just melted NaHSO4 * H2O and/or a
little bit more EtOH. I didn't want to do the latter because i wanted the filtrate to all fit in a 500ml boiling flask.
Literature:
1) Toro, Dobrosz-Gómez & García (2014) 'Sodium sulfate solubility in (water + ethanol) mixed solvents in the presence of hydrochloric acid.
Experimental measurements and modeling' Fluid Phase Equilibria, 384(), 106–113. doi: 10.1016/j.fluid.2014.10.025
2) J. J. Stöckley, R. Bartunek (1934) 'Process for the separation of sodium sulfate from sulfuric acid', US Patent US1812310A
[Edited on 20-7-2024 by BAV Chem] |
This is absolutely splendid! I do want to mention there is a salt noted in the literature that is basically trisodium bisulfate. It’s basically a
double salt of neutral sulfate and bisulfate. Seems methanol and ethanol in 80% solutions in water break up that double salt taking more free
sulfuric acid with it. Personally, I use the tainted 93% drain cleaner acid to make nitric, hydrobromic, and hydrochloric acids. Mostly nitric
though. And the nitric acid and some NO2 readily destroys the dyes and organicsnin the drain cleaner acid. So 400 grams of potassium bisulfate waste
product could be converted to a considerable amount of clean sulfuric acid. Having roughly 10% bisulfate contamination would not be an issue here as
it can go right back into making more nitric acid using less drain cleaner acid. I suspect a batch of sulfuric acid reused like this multiple times
would balance itself out in the sense that only so much bisulfate can dissolve in concentrated acid.
Fellow molecular manipulator
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chempyre235
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That's not bad. By doing the math:
250g (NaHSO4) --> ~50ml H2SO4
The cheapest sodium bisulfate I can find online sells for about $2/lb (in bulk) in my area in the US, so 1lb (453g) makes about 90ml of the sulfuric
acid. So, since 90ml of acid costs about $2, the acid runs almost $25/liter.
This is still more expensive than drain opener, which is sold in 32floz. bottles (1L ~ 33.8floz.) runs for about $12. However, this is viable in cases
where shipping or purchasing sulfuric acid is an issue.
"However beautiful the strategy, you should occasionally look at the results." -Winston Churchill
"I weep at the sight of flaming acetic anhydride." -@Madscientist
"...the elements shall melt with fervent heat..." -2 Peter 3:10
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MrDoctor
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dont forget bisulfate is available in 20/25kg bulk bags too if you know where to look. cleaning chem shops often also have really bulk pool supplies.
What im curious to know though is, exactly what reactions are taking place, and why can you boil down that mixture without it forming diethyl ether?
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clearly_not_atara
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It's really funny that my most impactful contribution here might be this method of making sulfuric acid that I just ran across one random afternoon.
I'm really impressed by all the work everyone has done. You took an idea about solubility differences and worked out a scalable method for making a
very important reagent.
I haven't figured out what to do about residual sodium. One possibility is to mix the sulfuric acid with dry alcohol and cool to very low temperatures
to hopefully precipitate most/all of it (either as Na2SO4 or NaHSO4). Then alcohol is boiled off again.
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MrDoctor
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polyphosphoric acid might work here to just dehydrate all the contaminated sulfuric acid into SO3, and then dissolve that into oleum if theres no
potential for oleum itself to exist stable while the polyphosphoric acid still exists at that hydrate-state, it can be made by dehydrating phosphoric
acid, but i think prefferebly starting from MAP, heated in a copper crucible with a torch, and based on what ive heard, it should just throw off SO3
fumes as you drip sulfuric acid onto it, with no liquid pool, so it should be completely efficient although the phosphoric acid might not be super
reuseable on account of the sodium presence if its final form is ultimately the bisulfate, depending on how that affects the copper crucible at high
temps.
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teodor
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Tsjerk's method of mixing IPA and NaHSO4 is quite interesting. I did myself experiments with mixing NaHSO4 and glacial AA but it doesn't work by some
reason I don't quite remember now quite well in details. Also I suspended my experiments with that after discovering NaHSO4 + NaPO3 method of SO3
generation. With relatively easy built source of SO3 turning it into H2SO4 probably could be a minor issue.
But that IPA + H2SO4 discovery is very nice and important for understanding the fascinating chemistry of H2SO4.
As for my metaphosphates studies I hope one day I will have time to carefully describe all the results, just have a lot of things on my plate now.
Surely I will have a lot to report but will do that after additional study - the amount of experimental details generated now is exceeding my ability
to refine the experiment conditions. During refinement more interesting details have arised.
As for NaHSO4 + NaPO3 I think it works much better than heating NaHSO4 alone, the only thing that the word "much" should be converted to some
numerical values I am not ready to provide. For example, I unable to say exactly what is the lowest temperature when the decomposition starts but
still would rate this as the probably "simplest" method, at least after some experimental improvements.
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Belowzero
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Personally I hate distilling sulfuric acid but I don't really see a reason why the H2SO4/Na2SO4 mixture could not be distilled. Leaving behind the
initial product too which can in turn be used again.
Thanks BAV chem for sharing this finding with us.
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clearly_not_atara
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https://patents.google.com/patent/US3364202A/en
| Quote: | 1000 k-g./ hour of ammonium bisulfate and 120 kg./hour of Water are fed to stirring vessel A via conduit 1, while a mixture of 2000 kg. of methanol
and 128 kg. of water is fed to the said vessel via conduit 3. 703 kg. of double salt,
consisting of 327 kg. of NH HSO 335 kg. of (NH SO 26 kg. of H 0 and 15 kg. of methanol, are separated off in filter device P. They are mixed with 415
kg. of water in stirring vessel B, and, after that, extracted with 2000 kg. of methanol and 128 kg. of water in stirring vessel C. 502 kg. of ammonium
sulfate, still containing 14 kg. of Water and 5 kg. of methanol, are separated off in filter device Q. This water and methyl alcohol are removed in
drying drum T, so that 483 kg. of ammonium sulfate are discharged via conduit 11.
Distillation of column X receives 2545 kg. of solution, consisting of 1985 kg. of methanol, 222 kg. of H 0 248 kg. of H 80 and 90 kg. of NH HSO via
conduit 13,
and, 2744 kg. of solution, consisting of I 2010 kg. of methanol 555 kg. of H 0,
110 kg. of H SO and 69 kg. of NH HSO via conduit 14.
The top product issuing from this column consists of 3995 kg. of methanol and 242 kg. of H 0. The bottom product, consisting of 1052 kg. of a solution
containing 40.5% by weight of H SO calculated to the sum of the free sulfuric acid and water presentand furthermore 15% by Weight of ammonium
bisulfate, is concentrated in distillation column Y.
540 kg. of sulfuric acid, contaminated with ammonium sulfate and being composed of 79.0% by weight of H 50 16.8% by Weight of (NH SO and 4.2% by
Weight of H 0 is discharged as bottom product. |
It appears that the reaction of ammonium bisulfate with
methanol requires only a relatively small amount of water.
But really I found this while chasing down a claim from a book that monoalkyl sulfate esters could be produced by the reaction of the alcohols with
| Quote: | | ammonium bisulfate in the presence of hydrazine or phosphoric acid [54] |
https://books.google.com/books?hl=en&lr=&id=B-27_lCJ...
The idea that methylsulfuric acid could be readily obtained from methanol and NH4HSO4 is pretty interesting.
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bnull
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| Quote: | But really I found this while chasing down a claim from a book that monoalkyl sulfate esters could be produced by the reaction of the alcohols with
| Quote: | | ammonium bisulfate in the presence of hydrazine or phosphoric acid [54] |
|
There is a mistake in the book. It is not phosphoric acid but hypophosphorous acid (the old German name is unterphosphoriger säure, which is
what appears in the patent). The first mentions of hydrazine and a phosphorus acid appear in the patent at column 4, lines 25-29 (from the patent page
at Espacenet):
| Quote: | | Als farbstabilisierende Maßnahme hat sich außer der Verdrängung des Luftsauerstoffes durch Inertgase der Zusatz von 0,01 bis 1,0% an Hydrazin oder
unterphosphoriger Säure bewährt. |
As far as I understand, hydrazine and hypophosphorous acid are used to avoid discoloration of the product, no phosphoric acid involved at all. So, you
don't need these if you don't mind having to purify the product. Nice, eh?
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