Sciencemadness Discussion Board - The simplest preparation of sulfuric acid?!

The simplest preparation of sulfuric acid?!

clearly_not_atara - 6-1-2018 at 21:22

"Adding ethanol to the solution [KHSO4 in water] precipitates potassium sulfate"

http://en.wikipedia.org/wiki/Potassium_bisulfate

If true, sulfuric acid must remain in solution. This is quite remarkable, as preparing sulfuric acid from bisulfates is, I thought, rather difficult. But the quote is unsourced, and it's not something I'd expect to find articles about. Can anyone verify that this works?

j_sum1 - 7-1-2018 at 00:42

I can give it a try using NaHSO4.
I'll get back to you.

Edit
Missed the H. Kinda important.

[Edited on 7-1-2018 by j_sum1]

NEMO-Chemistry - 7-1-2018 at 02:46

I doubt this is relevant, but at 170C ethanol and Sulphuric acid gives ethene, so it might be reasonable to assume an eye on temperature would be a good idea.

I have methanol and no ethanol, so i cant try this, Looking at wiki however it seems very specific about the salts being potassium. So if JS1 gets a blank result with the sodium salt, i wouldnt be too quick to dismiss validity of what you found.

If its correct, then alot of home chemists are going to be happy bunnies in the UK!

Also isnt potassium one of those salts that will salt out ethanol from water? and sodium carbonate dosnt work as well?

Sounds interesting though dosnt it.

On reading the entry for the sodium salt, there is no mention of ethanol and salting out, it also mentions the sodium salt being hygroscopic. This kind of reminds me of POKs potassium, where the conditions and reactants need to be spot on, but JS1 results should clear that up

This might be useful
http://pubs.acs.org/doi/pdf/10.1021/acs.oprd.7b00197

Take a look at the attached file, looks like with the correct amounts your spot on the money, again the salt mentioned is the potassium one, whats interesting however is the Alcohol dosnt have to be ethanol.

Seems i only have the Sodium salt as well. I will order the potassium one.

Attachment: bisulphate.pdf (181kB)
This file has been downloaded 929 times

[Edited on 7-1-2018 by NEMO-Chemistry]

j_sum1 - 7-1-2018 at 03:53

Well, this indeed is interesting.
I made a saturated solution of NaHSO4 and added some pure ethanol at test tube scale. Precipitate formed quickly.
I tried the same in a 50mL beaker with methylated spirits (95% ethanol + junk - mostly water) and got the same result: a thick, almost gelatenous precipitate filling most of the liquid space. It needs to be filtered and distilled. Then, if successful it will be a process to discover the best proportions to use. But indicators are promising. It wouls scale up to bucket scale better than many preparations, the ethanol should be recoverable and it looks to be a cheap exercise.

unionised - 7-1-2018 at 05:13

The first thing to do is check that the ppt isn't the starting material.

SWIM - 7-1-2018 at 08:23

the potassium salt gives less gelatinous results.

Ether isn't much of a worry if you can filter the solids out and then dilute with water to break any reaction complexes and hydrolyze the organosulfates.

Really it's just another way to get a bunch of dilute sulfuric acid which then would have to be boiled down.

Buy battery acid, boil it.

ave369 - 7-1-2018 at 12:06

Could anyone filter and dry the precipitate, dissolve it in water and use a pH test paper to check if it's sulfate or still bisulfate? I could try this tomorrow. I've got a lot of KHSO4 as a byproduct of Glauber's classic nitric acid synth, a lot of ethanol as well.

[Edited on 7-1-2018 by ave369]

NEMO-Chemistry - 7-1-2018 at 12:38

I can confirm it works with methanol and IPA. I used sodium salt. the paper i attached above gives amounts of alcohol to add as a rough guide.

Looks really promising.

NEMO-Chemistry - 7-1-2018 at 12:39

Quote: Originally posted by SWIM

the potassium salt gives less gelatinous results.

Ether isn't much of a worry if you can filter the solids out and then dilute with water to break any reaction complexes and hydrolyze the organosulfates.

Really it's just another way to get a bunch of dilute sulfuric acid which then would have to be boiled down.

Buy battery acid, boil it.

Battery acid by end Q2 is going to get hard to get hold of in UK

j_sum1 - 7-1-2018 at 13:54

Quote: Originally posted by ave369

Could anyone filter and dry the precipitate, dissolve it in water and use a pH test paper to check if it's sulfate or still bisulfate? I could try this tomorrow. I've got a lot of KHSO4 as a byproduct of Glauber's classic nitric acid synth, a lot of ethanol as well.

[Edited on 7-1-2018 by ave369]

Did this.
Solution was quite acidic. Not sure if it is NaHSO4 or some H2SO4 that was locked in the crystals.

Crystals are a lot more volumous than starting material. They appear a lot whiter with a crunchy texture like nice powder snow after they come off the Buchner. So it looks like there is some change.

I am yet to test the filtrate.

j_sum1 - 7-1-2018 at 21:40

I am here to report that this seems to work. I will do another run at a larger scale but indicators are that we now have another route to sulfuric acid that is pretty feasible and low tech/low expense.

I prepared some saturated NaHSO4 solution -- about 20mL. To this was added around 25mL of store-bought methylated spirits. Precipitation was fairly rapid. I left it for several hours for crystallisation to complete. Crystals occupied most of the volume of the beaker.

After crushing with a glass rod the crystals were vacuum filtered and washed with further methylated spirits.

The crystals are white and fluffy and easily soluble. They test acidic in solution. This means they containe either unreacted NaHSO4 or entrapped H2SO4 or, according to the paper that NEMO gave, some species like NaH3(SO4)2 or Na3H(SO4)2. They are of unknown hydration (obviously). I will dry them out and test further to find out what I can about them.

The filtrate was distilled using short-path condenser.

First fraction was taken off when puffs of white smoky vapour began to appear. First fraction is flammable, smells like ethanol with hints of something else. It tests neutral pH.

Second fraction was taken off when the distillate began to come over in large oily drops instead of a free-flowing glass-hugging liquid. It is also flammable but not as much. It tests acidic with pH paper and smells quite fruity. Evidently some esterification has occurred.

Fraction remaining in the flask is reasonably viscous. No noticable odour. No precipitate occurred although I expect one will form on further boiling down. It is discoloured quite yellow which can be attributed to impurities in the starting material. It is very acidic.
Putting a few drops on a spoon and subjecting to a flame gave interesting results. Volume dectreased as remaining water boiled away. White powder began to appear. This gave a characteristic Na Flame colour. An oily green liquid remained: presumably sulfuric acid contaminated with spoon plus whatever impurities were present at the start.

Conclusion
1. The product looks suspiciously like crude H2SO4. Nice.
2. Exact proportions and concentrations need to be determined to minimise waste and unwanted byproducts.
3. Some but probably not all of the alcohol is recoverable for further batches.
4. It seems that there is considerable esterification going on. And possibly some ether produced too. Vacuum distillation might prevent these.
5. Product will require distillation to be of use -- either that or much higher grade starting reagents.

Given that distillation of H2SO4 is not for the faint hearted and that SO3 can be produced from heating bisulfate it might be that little is to be gained by the alcohol salting process. OTOH, if impurities are not a problem, H2SO4 is otherwise unavailable and you need something stronger than bisulfate then this could be a rough and ready route. It might also be a sensible way to use up the bisulfate byproduct of nitric acid distillation as well as getting rid of the unwanted bottle of vodka that great uncle Vlad gave at Christmas.

In any case, this has shown a viable process and one that could undoubtedly be refined further.

[Edited on 8-1-2018 by j_sum1]

Assured Fish - 7-1-2018 at 21:58

j_sum1 Could you try doing a melting point determination of the fluffy crystal precipitate?
NaHSO4.H2O melts at 58*C and thus if it was meltable below 100*C, then there would still be considerable bisulfate present.
If you are unable to melt it below 200*C then i think we can safely assume the bisulfate has reacted entirely.

This is a rather marvelous find.

j_sum1 - 7-1-2018 at 23:03

MP exceeds 190°C. My thermometer tops out at 200 and I did not want to pop another one.
So no bisulfate remaining -- which is a good sign. That does not explain why it is so acidic though.
I will give it a bit more of a wash and see what happens.

I am just recrystallising some NaHSO4 at the moment so I can attempt with a purer reagent. I will use vacuum distillation this time around to keep the temp down.

j_sum1 - 7-1-2018 at 23:56

Precipitate washed several times with ethanol. Mixed with distilled water and still acidic. It melts higher than 190C with no signs of any decomposition before that. I am not sure what it might be.

I'll try hitting it with a butane flame next.

clearly_not_atara - 8-1-2018 at 01:50

Not SES? I guess there might be a phase like Na3H(SO4)2, which happens with eg potassium tetraoxalate KH3(C2O4)2

j_sum1 - 8-1-2018 at 02:46

K3H(SO4)2 and a few other similar-looking species were mentioned i. The paper excerpt that NEMO cited. But that just widens the list of possibilities and identifying it positively becomes more problematic.

Since the product is in the filtrate that is where I'll concentrate my efforts. +he main reason for analysing the filter residue was to confirm that the reaction had actually occurred. I think we can safely say that now.

NEMO-Chemistry - 8-1-2018 at 03:45

Sorry i posted the half arsed version of the paper! Here is the full one with the missing bits. Sorry JS1

[Edited on 8-1-2018 by NEMO-Chemistry]

Attachment: mientka2008.pdf (228kB)
This file has been downloaded 763 times

GrayGhost- - 8-1-2018 at 07:57

Hi, what concentration have you acid J-sum?

A aproximation is messure density.

[Edited on 8-1-2018 by GrayGhost-]

18thTimeLucky - 8-1-2018 at 08:30

I doubt measuring the density will give anything remotely accurate for the concentration due to the impurities, but maybe with a couple recrystallisations of the starting materials you could, we will see.

I am confused how any esterification occured, where did a carboxylic acid come from? Any ideas?

Also how sure are we it is sulfuric acid apart from an acidic pH? Would the simple sulfuric acid test of adding some of the filtrate to some sugar work?

GrayGhost- - 8-1-2018 at 09:02

Bubbling SO2 in hydrogen peroxide H2O2 obtain sulfuric acid, but is pure o contain sulfurous acid? I can buy max conc. peroxide 10volumes in farmacy.

SWIM - 8-1-2018 at 09:51

Quote: Originally posted by 18thTimeLucky?

I doubt measuring the density will give anything remotely accurate for the concentration due to the impurities, but maybe with a couple recrystallisations of the starting materials you could, we will see.

I am confused how any esterification occured, where did a carboxylic acid come from? Any ideas?

Also how sure are we it is sulfuric acid apart from an acidic pH? Would the simple sulfuric acid test of adding some of the filtrate to some sugar work?

Diethyl sulfate??? Should mostly break down to ether if this was an atmospheric pressure distillation, but beware if that smell is at all minty!

I see from the CRC that the sodium acid sulfate is 'slightly soluble' in alcohol, and breaks down in it as well, so when I get the chance tonight I'll try putting a few samples of NaHSO4 in methylated spirit and let them sit a few days to see what happens.

I've done slow reactions like this before and you don't need much solubility to get it to go to completion in a few days if the reaction in solution is fast.

I suppose it may work until the acid content reaches some critical level and either stops the reaction or just dissolves the remaining bisulfate

Haven't looked it up, but I assume it is soluble in sulfuric acid. If not, does it decompose to sulfate and acid in those conditions?

Jjay reported some precipitate from boiling down drain cleaner that he thought was sodium sulfate. If it is, then maybe just adding the acid sulfate to the acid and heating, or letting it sit for a few weeks, would do it.

If the acid sulfate had to be added as a saturated solution to dissolve it that'd just mean a little more boiling down.

Of course that wouldn't be a way to make H2SO4 from scratch, but it might be a convenient way to make more so you never run out.

Also, once you make a bit from some other, possibly less convenient, process you could use that acid for this process if it was better suited to your situation.

Edit: dried some bisulfate in the oven as it was a bit cakey, and then added 36 grams to 50ml of Clean Strip Green, which is supposed to be just ethanol and methanol.

Added another 36 grams to 50ml of sulfuric acid based drain cleaner that claims 98% on the label.

Neither is generating any noticeable warmth. The alcohol sample is free-flowing powder after a few hours, but the sample in the acid caked and fused into a solid porous mass which was broken up with a stir rod.

The crystals in the acid may have changed in character to some extent.

Both will be left for a few days with occasional shaking too see if anything develops.

[Edited on 9-1-2018 by SWIM]

j_sum1 - 9-1-2018 at 03:30

Second attempt.
The plan was to recrystallise some NaHSO4 so that I was working wth something a bit more pure. My recrystallisation did not play nice so I went ahead with the hardware grade.

28g of NaHSO4 was added to 100mL water. This is a bit below saturation at the temperatures I was working with. I might have been better with a saturated or supersaturated solution but I wanted to ensure that any precipitation was the result of the alcohol.
I elected to use methylated spirits again which AFAIK is 5% water plus bittrex plus a few other simple organics on top of the 95% ethanol. If this is ever going to be a practical procedure it will need to work on cheap OTC reagents.
Addition of methylated spirits slowly turned the solution cloudy but no substantial precipitate until a considerable portion had been added. I ended up using a gross excess of about 300mL.
Some fine precipitate made it through the vacuum filter. I let it settle and decanted the clear liquor.
I set up for vacuum distillation. Not much of a vacuum -- my pump sucks (or doesn't suck for those who want to take me literally.) The idea was to keep the temperature low to avoid unwanted side reactions. I pulled off three fractions (photographed)

Fraction 1 appears to be a clean mixture of alcohol and water. Neutral.
Fraction 2 smells odd. It is reasonaby acidic but I didn't titrate. There is no significant ester odour like there was in the first run.
Fraction 3 came over after puffs of white vapour began to appear in the distiling flask. No smell. Turned pH paper deep red. It might be reasonable quality dilute sulfuric.
The residue in the flask appeared a thick and oily discoloured sludge. Evidently there is some organic material still in it. Cooled to 50°C it still emitted acidic vapours from the mouth of the flask. After leaving for a couple of hours it resembled a gritty paste with crystalline material in it.

Conclusion.
That's the end of the road for me. The idea is intriguing and has some merit. But the work up is a dog and indications are that it will continue to bark like one. Yield, impurities and time investment are all on the wrong side of practical to replace any of the established methods and I still cannot be certain that it is sulfuric -- it might still be dissolved bisulfate.

NEMO-Chemistry - 9-1-2018 at 03:58

I have the potassium salt ordered, i will double what you have done 'just in case'. I will also try pure methanol and see what happens.

Great work JS1

I know the difference between potassium salt and Sodium should be non existent, but everything I have read nags me about it. The other nag i have is methylated spirits, apart from cleaning and a small alcohol burner, i never had much luck with it in UK.

Just maybe they use recycled solvent in it, no idea. I found some other information and again it mentions using Potassium, 8it wont make much sense if the results i get differ from yours, but the POK thing is nagging away at me.

I got a great template to work from now

, really cool experiment you did.

NEMO-Chemistry - 9-1-2018 at 07:36

Reading this a bit more....

I think we are going to find a big difference between potassium Bisulphate and sodium Bisulphate.
On the wiki page is the following
Potassium bisulfate is also formed by the union of sulfuric acid with potassium sulfate. It goes on to mention that much of it is from the production of Nitric acid, so again this would be Potassium Sulphate.

I think the sulphuric acid salting out, is very much more a potassium bisulphate favored than the salting out from sodium Bisulphate. Again wiki mentions potassium sulphate dropping out on salting.

Also mentioned is the following

Aqueous solutions of potassium bisulfate behave as two separate, uncombined compounds, K2SO4 and H2SO4. Adding ethanol to the solution precipitates out potassium sulfate

I cant find anything similar relating to sodium bisulphate, while i dont dispute its possible to some extent, i seriously think we are missing something here.

While the sodium salt seems to need some coaxing etc, all the papers and info from the potassium salt, all seem to suggest the sulphuric acid literally just falls out of solution, or rather the potassium sulphate does.

I have looked for something that says sodium bisulphate behaves like separate uncombined compounds in solution, i cant find any. This would be key to it working, if the sodium bisulphate dosnt behave as separate compounds in solution then obviously it wont be as straight forward.

Anyone think i am barking mad with this?

[Edited on 9-1-2018 by NEMO-Chemistry]

SWIM - 9-1-2018 at 10:34

I think it's worth checking.

But the more I think about the differences between the salts, the more confusing it gets.
Usually poatassium salts are a bit more soluble in alcohol than sodium salts, so if the Potassium precipitates better what's going on?

Maybe the potassium salt lattice structure doesn't accept acid or alcohol inclusions as much and forms denser crystals.

But I say always check everything within reason. I'm not the kinda guy who tries to save a few hours of research time by spending days struggling in the lab.

Status report: Oven dried sodium bisufate in alcohol still looks much the same and is free-flowing.
Same bisulfate in '98% strength' drain cleaner has dissolved appreciably, about a third of the material appears to be gone. Remainder is stuck together in a loose, friable mass.

If nothing else, this shows that H2SO4 is a hell of a lot better solvent than I thought. It may well be capable of holding more dissolved bisulfate than water does.

NEMO-Chemistry - 9-1-2018 at 17:44

Its mainly wiki that got me thinking, looking at the solubility and the fact the potassium salt acts like two separate substances in solution, also potassium bisulphate is often made as potassium sulphate with sulfuric acid added, the sodium salt is very different, no solubility data for alcohols etc.

But more than that, every paper where its mentioned has only mentioned the potassium salt, none mention the sodium. Then finally we get to potassium carbonate salts ethanol much much better than sodium carbonate. There should be as much difference as there is.....

I got a knock back from the chem company! no idea why but apparently they are not stocking it now??? Strange as they said it was dispatched.

So will get potassium nitrate and make some nitric acid

clearly_not_atara - 11-1-2018 at 17:09

Potassium salts are generally less hygroscopic than sodium salts, which may be an advantage in this case. It's possible that sodium sulfate catalyses the formation of ethylsulfuric acid and its salts. Potassium bisulfate is also almost twice as soluble in water as the sodium salt, which probably increases the yield significantly -- both sodium and potassium sulfate are very insoluble in alcohol.

Nice work j_sum1. Methylated spirits usually contain a few percent butanone these days IIRC. I don't think that would interfere with the rxn too much but it's there to prevent distilling them to get ethanol.

EDIT: Ammonium bisulfate may work even better -- the aqueous solubility is very high, about 200g / 100 mL at 300 K, and in some cases the addition of excess ammonium bisulfate to water precipitates (NH4)3H(SO4)2 (Beyer&Bothe 2006, attached), before any alcohol is even added! Ammonium sulfate is famously insoluble in organic solvents, as well.

[Edited on 12-1-2018 by clearly_not_atara]

[Edited on 12-1-2018 by clearly_not_atara]

Attachment: beyer2006.pdf (227kB)
This file has been downloaded 645 times

SWIM - 13-1-2018 at 13:19

The bisufate in methylated spirits has increased in density by 15% and become rather viscous.

A sample showed strong acidity and a considerable reserve of it.

This is consistent with the description of the bisufate as slightly soluble in alcohol, but decomposed by it, in the CRC manual. The total acidity of the solution exceeded anything that the slight solubility of the bisulfate could have contributed by many times.

The solution was returned to the bisulfate to see if any further reaction occurs, as monitored by testing the density.
I think the possibilities of bisulfate systems for obtaining sulfuric acid merit some thought and attention. At least a bit more than they've received so far.

@clearly_not_atara, Where did that reference for sodium bisulfate being 'very insoluble' come from?
Doesn't seem like there'd be any reaction in that situation to me, at least not on a reasonable time scale.
Just can't see how the reaction could proceed into the crystals much without some degree of solvation.
Is there some phenomena I may be unaware of here?

JJay - 13-1-2018 at 20:06

If you reflux sodium bisulfate and ethanol, it forms some sodium sulfate and some ethyl sulfate. This was claimed to be an extremely high-yielding reaction in a U.S. patent but didn't seem to be when I tried it with pool-grade sodium bisulfate and anhydrous ethanol, and I'm not exactly sure why. It definitely does work to some extent, though.

I don't see why you couldn't distill off the ethanol and then heat the ethyl sulfate to 130 or so and distill off ether, leaving behind sulfuric acid. You probably won't get all of the ethanol out that way, and as the temperature increases you'll distill off ethene and then ethyl sulfate (probably some diethyl sulfate also). There may be some sulfuric acid left when you reach its boiling point, though.

[Edited on 15-1-2018 by JJay]

NEMO-Chemistry - 14-1-2018 at 02:33

I cant get hold of potassium bisulphate at the moment, my normal sources dont have it or wont sell it!
I have sodium bisulphate, how can i convert it?

As a side note I have real trouble getting most potassium salts including carbonate!! I am slightly restricted in which companies i can use, but could do with confirmed sources in the UK for potassium Bisulphate. Ebay isnt an option for this one, i have to buy this through the company, on advice i shouldnt buy chemicals for the company via ebay.

AJKOER - 22-2-2020 at 18:27

Two comments:

First, on the alcohol/bisulfate mix, H2SO4/KHSO4 may be a case of a so-called acid salt, reported with sodium acetate and acetic acid, see this reference at https://pubs.rsc.org/en/content/articlelanding/1919/ct/ct919... . For a more recent science perspective, see https://www.ncbi.nlm.nih.gov/pubmed/17619064 .

Note, mentioned above is a possible salt, NaH3(SO4)2 which I re-write as NaHSO4(H2SO4).
-----------------------------------------------------------------

Second, an interesting nonredox reaction equilibrium, per a source, 'Redox and nonredox reactions of magnetite and hematite in rocks' at https://www.researchgate.net/publication/283343347_Redox_and... Namely:

Fe(III)2O3 (hematite) + Fe(II) + H2O = Fe(II)Fe(III)2O4 (magnetite) + 2 H+

If the ferrous salt is FeSO4, then some presence of H2SO4 implied. Likely, not a practical path to sulfuric acid, but one may be able to use it to create associated sulfates of low solubility starting with a target metal or its oxide.

[Edited on 23-2-2020 by AJKOER]

G-Coupled - 23-2-2020 at 11:52

Quote: Originally posted by NEMO-Chemistry

I cant get hold of potassium bisulphate at the moment, my normal sources dont have it or wont sell it!
I have sodium bisulphate, how can i convert it?

As a side note I have real trouble getting most potassium salts including carbonate!! I am slightly restricted in which companies i can use, but could do with confirmed sources in the UK for potassium Bisulphate. Ebay isnt an option for this one, i have to buy this through the company, on advice i shouldnt buy chemicals for the company via ebay.

Can't you make a private purchase from eBay?

clearly_not_atara - 23-2-2020 at 12:18

I guess this prep would work okay with sodium bisulfate. As noted above, I think ammonium bisulfate has more potential.

But I'm just confused when people say they can't get potassium. It's one of the three essential elements in fertilizer! "NPK" is the name of the game. I can go to the gardening store five minutes away and buy a 10kg bag of K2SO4. Nobody would even raise an eyebrow if I walked out with ten of them.

What do you actually see in the fertilizer aisle? Do they even sell fertilizers? Why are UK gardeners recommending high-potash fertilizer if it supposedly doesn't exist there?

https://ferndalegardencentre.co.uk/what-is-high-potash-plant...

In a post-apocalyptic nightmare, you can still get potassium salts by simply burning a bunch of wood and leaching the ashes. Annoying, but effective.

[Edited on 23-2-2020 by clearly_not_atara]

Electroplating waste acid

MadHatter - 23-2-2020 at 14:03

Do you have access to CuSO₄, carbon rods and a battery
charger ? This is how I produced dilute H₂SO₄ in the past
and concentrated it by boiling off the H₂O. CuSO₄ in the US is
available as a root killer. Last check, $29.99 + tax for 10 LBS
from Walmart. Carbon rods from dry cell batteries or my
favorites, arc welder gouging rods(lookup my (per)chlorates
methods), are used.

When my favorite nephew came to me and asked me to help
him with his senior year(high school) final project in science, I
was elated !

His teacher's only requirement was that it had
to involve electricity. I suggested electroplating.

Fellow mad scientists feel free to jump in at any point with tips,
possible side reactions or if it's clear that I have my head up my
ass. I welcome any criticism negative or positive.

The balanced equation(I think), with electricity, is as follows:

2CuSO₄ + 2H₂O --> 2Cu + 2H₂SO₄ + O₂

Anyway, dissolve CuSO₄ in H₂O and electrolyze with the carbon
rods of your choice. The solution, which starts as a deep blue,
will eventually turn clear with Cu being electroplated on the
cathode. This leaves the H₂SO₄ in the solution.

BTW, this is the process IIRC, for producing 92 - 93% H₂SO₄
drain cleaners such as Rooto.

I hope this helps somebody. HAVE FUN !

AJKOER - 2-3-2020 at 19:21

Quote: Originally posted by AJKOER

....
Second, an interesting nonredox reaction equilibrium, per a source, 'Redox and nonredox reactions of magnetite and hematite in rocks' at https://www.researchgate.net/publication/283343347_Redox_and... Namely:

Fe(III)2O3 (hematite) + Fe(II) + H2O = Fe(II)Fe(III)2O4 (magnetite) + 2 H+

If the ferrous salt is FeSO4, then some presence of H2SO4 implied. Likely, not a practical path to sulfuric acid, but one may be able to use it to create associated sulfates of low solubility starting with a target metal or its oxide.

[Edited on 23-2-2020 by AJKOER]

Per this paper https://pubs.acs.org/doi/pdf/10.1021/acs.jpcc.5b10949 , a reaction of interest forming FeSO4 from the action of O2 on FeS2:

FeS2 + 7/2 O2 + H2O = Fe(2+) + 2 SO4(2-) + 2 H+

So, for a school project, use water and oxygen to turn FeS2 and added Fe2O3 into H2SO4 and magnetite.

Tsjerk - 5-9-2020 at 06:31

This reaction works and is nearly quantitative. I added 100 ml methanol to 100 ml solution containing 49 gram KHSO4 (a little heating was used to get everything in solution). The precipitate was vacuum filtered and rinsed with a bit of methanol. The extra methanol didn't cause extra precipitation in the filtrate, so probably less methanol would suffice.

The total volume of the filtrate was 180 ml of which 50 ml was titrated with 0.096 mol of NaOH, corresponding with a sulfuric acid yield of 96%.

The K2SO4 was dried and weighed and found to be 32,3 gr. Exactly the weight of the K2SO4 plus the the 4% yield loss in KHSO4. The KHSO4 is probably the reason people find their precipitate to be acidic.

I didn't bother to concentrate the acid as I have plenty, but I guess one could recover the methanol by simple distillation. I don't think methyl hydrogen sulfate or dimethyl ether would form because of the large amount of water present.

njl - 20-1-2021 at 17:05

What concentration can be obtained by distilling off methanol? Is there an equally simple way to remove/destroy enough methanol to get +98% acid, or even acid free off organic contaminants?

clearly_not_atara - 21-1-2021 at 08:57

In principle, you should have no trouble distilling off all of the methanol. However, if dimethyl ether is produced, you will have to contend with the fact that this is a "heavy" (concentrates near the ground), flammable gas that is now floating around your laboratory. Shouldn't be a problem with proper ventilation (including floor vents/open door) but definitely not something to ignore.

Fyndium - 21-1-2021 at 09:40

In case of any very volatile gases can be formed, leading the exhaust tube directly into the ventilation duct is a good manner. I do this every time I deal with something smelly, toxic or flammable.

katyushaslab - 21-1-2021 at 10:20

This has interesting potential for "recovering" H2SO4 consumed in making nitric acid from nitrate salts of alkali metals, given the "remaining product" is XHSO4 (X being whatever alkali metal). Instead of this being a waste product, you could expend cheap* alcohol on regenerating it for future reactions?

Certainly one to consider experimenting with in the future, will be very interested to hear how others fare with this whole process.

* Methylated spirits being cheap enough, but given you are assumed to already have distillation apparatus, home-made is also a viable option...

njl - 21-1-2021 at 10:35

Isn't the methanol also recoverable? It's just changing the solubility of KSO4, right? Not actually being consumed except in side reactions.

beta4 - 24-1-2021 at 13:27

Quote: Originally posted by clearly_not_atara

In principle, you should have no trouble distilling off all of the methanol. However, if dimethyl ether is produced, you will have to contend with the fact that this is a "heavy" (concentrates near the ground), flammable gas that is now floating around your laboratory. Shouldn't be a problem with proper ventilation (including floor vents/open door) but definitely not something to ignore.

I'm not sure, but can't the sulfuric acid and methanol combination form dimethyl sulfate (https://en.wikipedia.org/wiki/Dimethyl_sulfate)? That alone would be a reason not to try this experiment.

Tsjerk - 24-1-2021 at 15:11

I don't think the methanol and sulfuric acid will react to any significant extent before all methanol is evaporated. After that it is just a normal sulfuric acid in water solution.

If anything would come to be it would first be methyl hydrogen sulfate, but for that you already need pretty anhydrous conditions.

woelen - 25-1-2021 at 01:51

Quote: Originally posted by Tsjerk

This reaction works and is nearly quantitative. I added 100 ml methanol to 100 ml solution containing 49 gram KHSO4 (a little heating was used to get everything in solution). The precipitate was vacuum filtered and rinsed with a bit of methanol. The extra methanol didn't cause extra precipitation in the filtrate, so probably less methanol would suffice.

The total volume of the filtrate was 180 ml of which 50 ml was titrated with 0.096 mol of NaOH, corresponding with a sulfuric acid yield of 96%.

The K2SO4 was dried and weighed and found to be 32,3 gr. Exactly the weight of the K2SO4 plus the the 4% yield loss in KHSO4. The KHSO4 is probably the reason people find their precipitate to be acidic.

I didn't bother to concentrate the acid as I have plenty, but I guess one could recover the methanol by simple distillation. I don't think methyl hydrogen sulfate or dimethyl ether would form because of the large amount of water present.

This really is an interesting option if you cannot get your hands on H2SO4. Recovering the methanol should indeed be easy and you don't need it pure. Even if you get it recovered, together with let's say 20% of water, then you can use it again for the next batch of making H2SO4.

You can make a lot of H2SO4 with this process. Dissolve KHSO4 in water, add methanol and then distill off methanol, plus maybe a little water to be sure to recover nearly all of it. Then continue boiling outside to drive off more water and the last remains of methanol, until the acid starts fuming. Then you have 75% or so acid.

The methanol then can be used again to make a new batch of H2SO4, as described above. I can imaging that with a certain batch of methanol you can make 5 or even more batches of acid.

It would be interesting to see whether this also works for NaHSO4. That chemical is even more accessible than KHSO4. J_sum1 tried with this, but the results look somewhat messy and not fully satisfactorily. Maybe if methanol is used instead of cheap OTC methylated spirit, which may contain all kinds of oily crap. NaHSO4 also can be obtained at decent purity in many countries as pH-minus for swimming pools. Where I live, the stuff can be purchased in 5 kg and 10 kg buckets, and it is a nice white more or less free flowing powder, which gives clear solutions and does not seem to be contaminated with organics or colored metal salts.

[Edited on 25-1-21 by woelen]

Tsjerk - 25-1-2021 at 05:19

I repeated the experiment with NaHSO4 and methanol and got a sulfuric acid yield of 95%. I haven't weighed the sodium sulfate yet, but I also don't know what hydrate precipitated out. The experiment was done on a 0.1 molar scale with 50 ml water / 50 ml methanol.

I do think the filtration would be very hard without a vacuum system though. I don't really remember how the potassium sulfate looked but I can't remember it looking like a gel as the sodium sulfate does. But then again a Chinese vacuum funnel and a hand pump or a water aspirator comes cheap.

I haven't tried, but I guess this should also work with ethanol. When you buy denatured 96% ethanol and reflux it for some time with NaOH and distill it, it should be just fine for this purpose.

Fyndium - 25-1-2021 at 05:59

Methylated spirit aka denatured ethanol tends to contain nil to a LOT of additives, depending on source. The cheapest one for automotive use contains almost 20% of detergents and other stuff. I have therefore always strip distilled my ethanol, resulting in azeotropic product according to alcometer.

Difference between methanol and ethanol is that former does not form azeotrope, being readily recovered in quite pure form, while ethanol is a lot trickier. If available, I would go with methanol.

Sludge of sodium sulfate should be manageable with good frit filter, like said, they come reasonably cheap in china. I have never seen potassium bisulfate anywhere, but sodium is available in every superstore.

woelen - 25-1-2021 at 07:29

@Tsjerk: That sounds really promising. This may be a very good route for making H2SO4 for a decent price. I have been looking for 15% H2SO4. Some seller carry this, but it is really expensive. A 1 liter bottle of 15% H2SO4 is only slightly cheaper than a 1 liter bottle of 96% H2SO4, but if you look at the weight of acid, it only contains appr. 80 ml of conc. H2SO4, added to nearly one liter of water (you should keep in mind that conc. H2SO4 has a densitiy of almost 1.84 gram/liter, while 15% H2SO4 is just over 1.1 gram per liter).

See https://wissen.science-and-fun.de/chemistry/chemistry/densit...

Tsjerk - 25-1-2021 at 08:22

I spoke too soon, I just weighed the Na2SO4 and even if it anhydrous it is way too little. So at least with these proportions (50/50 water/methanol) it doesn't work. I will try again tomorrow to see if a bit more alcohol works.

i will try 40/60 with 0.1 mol NaHSO4

[Edited on 25-1-2021 by Tsjerk]

woelen - 25-1-2021 at 13:03

Maybe NaHSO4 is too soluble in methanol. Longer chain alcohols might work better in that case, but the disadvantage of such longer chain alcohols also is more facile reaction and formation of brown or black goo on heating of the resulting dilute acid.

metalresearcher - 25-1-2021 at 13:04

I did a quick test.

I put a chunk of KHSO4 obtained by distilling HNO3 off KNO3 + H2SO4 drain cleaner in a test tube, dissolved it completely and added some water. It did not dissolve completely, so I poured the solution into another test tube. I checked the pH and it appeared to be 2.0, rather acid.
Then I added ethanol to the solution and it got cloudy.

Tsjerk - 26-1-2021 at 13:20

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]

Tsjerk - 27-1-2021 at 03:06

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]

Fantasma4500 - 21-2-2022 at 13:15

@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.

BAV Chem - 20-7-2024 at 14:52

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]