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clearly_not_atara
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[*] posted on 6-1-2018 at 21:22
The simplest preparation of sulfuric acid?!


"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?




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




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[*] posted on 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 :D

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 27 times

[Edited on 7-1-2018 by NEMO-Chemistry]
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[*] posted on 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.




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[*] posted on 7-1-2018 at 05:13


The first thing to do is check that the ppt isn't the starting material.
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[*] posted on 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.







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[*] posted on 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]




Smells like ammonia....
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[*] posted on 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.
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[*] posted on 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
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[*] posted on 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.




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[*] posted on 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]




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




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[*] posted on 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.




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[*] posted on 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



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




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[*] posted on 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 28 times

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[*] posted on 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-]
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[*] posted on 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?




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




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[*] posted on 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)

IMG_20180109_210233.jpg - 1.4MB

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.




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[*] posted on 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 :D, really cool experiment you did.
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[*] posted on 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]
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[*] posted on 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.





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