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Author: Subject: The simplest preparation of sulfuric acid?!
NEMO-Chemistry
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[*] posted on 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 :D
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clearly_not_atara
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[*] posted on 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)
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[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 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?




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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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?
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clearly_not_atara
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[*] posted on 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]




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 23-2-2020 at 14:03
Electroplating waste acid


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 !:D 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 !:D




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[*] posted on 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.


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[*] posted on 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.
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[*] posted on 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?
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[*] posted on 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.



[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 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.
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[*] posted on 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...
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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...




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[*] posted on 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]
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[*] posted on 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.



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[*] posted on 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.


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