I am on my eternal quest of reproducing H2SO4 and HNO4 from common fertilizers and might have found a route to H2SO4.
By extraction of NH4HSO4 with EtOH (NH4)2SO4 preciptates and leaves a liquid mix of EtOH and H2SO4. These react to form the ester Et-HSO4,
ethylhydrogensulfate.
this decompose upon heating into sulfuric acid, ether and water (classical labb route to ether)
However , does anyone know if the ester forms readily at roomtemperature?
My idea was to vaccum distill of the EtOH , leaving H2SO4 left , but if Et-HSO4 is present this needs to be removed in some way. Either by
heating/decomposition or distillation.
the ester is completely miscible with water and EtOH as far as I have found out.
Adding water will of course hydrolyse the ester into alcohol and acid by shifting the equlibrium , but this means diluting the acid , which i prefer
not to.
/rickard
I'm a little confused
Polverone - 30-10-2002 at 14:50
Most people consider vacuum distillation to be less easy/convenient than atmospheric pressure distillation. If you don't vacuum distill then sulfuric
acid should be the only thing left behind. Of course this method isn't very optimized without recycling of materials. You would use up a lot of
ethanol, I think, to prepare your H2SO4. Is this why you wanted to do vacuum distillation, so you would recover ethanol instead of water and ether?
This is just a guess, but I'd try removing as much of the liquid as possible under reduced pressure, and then strongly heating at atmospheric pressure
whatever is left behind. You're going to have to do this at some point anyway to get concentrated acid, right?
Although I've seen you post about it before, I didn't realize that H2SO4/HNO3 production was one of your personal quests (like my cyanide/phosphorus
quests). It does present some interesting challenges. Are you determined to stick to fertilizers, or will any inexpensive, unregulated starting
materials do? Sulfurous acid is easy to produce; bubble the fumes from burning sulfur through water. It's easy to oxidize it to sulfuric acid with
hydrogen peroxide, but hydrogen peroxide isn't necessarily inexpensive and unregulated.
It's impractical, industrially, to convert SO2/H2SO3 to SO3/H2SO4 without catalysts and sophisticated apparatus/conditions. However, a home
experimenter doesn't have deadlines to meet and profit margins to make. Is it possible that without a catalyst or with some simple catalyst H2SO3
could be oxidized to H2SO4 using only atmospheric oxygen, even if it took days or even a couple of weeks?
I know that I have seen large quantities of ferrous sulfate for sale for use on soil, and its dry distillation will yield sulfur trioxide. Is that a
method you've considered, or do you think it's too expensive to go that route?rikkitikkitavi - 30-10-2002 at 22:21
You are right on it Polverone.
The main reason for vacuum evaporation/distillation the liquid is to minimize losses ot EtOH. However I will do some further testing with various
amounts to calculate yields e t c to see if it is feasible.
One possible route is
SO2+H2O+H2O2 => H2SO4
major problem is that H2OS is expensive!
However , I have a few patents where the SO2 is oxidized by air in the aqueous phase whith V,Mn,Fe catalysts present. Up 40 % it seems to work.
I have also seen references where Cl2 is used
2 H2O+SO2 + Cl2 => 2HCl + H2SO4
Generating large amounts of Cl2 is cumbersome and dangerous. perhaps ClO3- could be used?
Another route to SO3
NH4HSO4 + K2SO4 => NH3 + 2KHSO4(250-400C)
2KHSO4 => K2S2O7 + H2O (400C)
K2S2O7 = K2SO4 + H2SO4 (450C)
all reactions by heating. I have seen a german patent about this.
Not only fertilizers are interesting, all cheap suggestions are welcome, and sometimes not so cheap! I aim high and hit low...
/rickardvulture - 31-10-2002 at 05:39
My chemistry encyclopedia states that when sulfur is burned in air, up to 40% of SO3 is produced. Maybe if you mix it stoichiometrically with KMnO4 to
produce
SO3, you'll end up with nearly 100% SO3?
I once tested a pyrotechnic mix of KMnO4/S and i noticed it burned quite fierce with rather high light output. It also didn't smell as much to SO2 as
blackpowder does.Marvin - 31-10-2002 at 09:23
I'm a tad confused here, if sulphur is available why not simply use the bell jar method to make sulphuric acid?
Either KNO3 or NaNO3 should be available everywhere. The amount required for a given amount of acid then just depends on how much time you are
willing to spend on the process. If nitrates are in extremely short supply, you can always make it yourself from animal dung/soil/wood ash.
The process will yeild up to about 70% sulphuric acid while consuming a very small amount of nitrate, which can be concentrated the normal way. Oleum
on the other hand requires contact process, which is not hard to do, a million times easier than oxidising ammonia for example, platinised rockwool
being slightly exotic, or distillation of iron sulphate, which requires high temperatures/clay rhetorts.rikkitikkitavi - 31-10-2002 at 09:49
when sulfur is burned in air only SO2 is formed. This is oxidized by V2O5 in the contact process, but this is very difficult to accomplish at home.
Thus we need some form of ozidizer to make the S(IV)=> S(VI) transition. The method described by Marvin is a version of the lead-chamber process,
which I also have considered. (this process is not longer in use since the contact process produces stronger acid for less price)
I planned on using a setup with 50 l glassjars (damejeanne) originally suited for whine making. You need a large chamber volume so that NO=> NO2 and
reoxidize further SO2.
The NO=> NO2 is a slow reaction at the concentrations we are talking about here( a few %) taking several minutes to complete.
The synthesis of HNO3 from NH4HSO4 and Ca(NO3)2 is quite forward, thus giving access to large amounts of HNO3.
Nitrates are cheap and easily accesible as fertilizers, Ca(NO3)2 is one. This can be converted to Cu(NO3)2 with CuSO4 , which is more easily
decomposed into NO2 and CuO. (or FeSO4 for that matter)
The major drawback is that even I have access to sulfur, this is not cheap. But I am still on the idea generation stage so any input is welcome.
My next consideration is, concentration of H2SO4? Up to about 80 % virtually only water is evaporated , but after that the H2SO4 concentration in the
vapour is rising sharply. This leads to losses of H2SO4 unless some type of closed evaporator system is used. I believe reading through patents will
reveal a lot here.
And the temperatures involved? Pyrex flasks e t c at 300+ C ? Is this feasible?
Especially a fragile item like distillation flasks with condensers e t c?
I cant imagine many other materials H2SO4-resistant @ 300 C...
/rickard
Sulfur is expensive?
Polverone - 31-10-2002 at 11:54
What's your sulfur source? I have relatives who use agricultural dusting sulfur on their vineyard. They buy it in 100 pound sacks and I don't think it
costs that much. I imagine that one sack could supply your experimenting needs for a long time.
Pyrex will resist boiling conc. H2SO4, but in larger quantities I would really worry about bumping. Having a flask of boiling H2SO4 shatter is near
the top of Things You Want to Avoid.rikkitikkitavi - 31-10-2002 at 13:28
100 pounds would yeild more than 300 pounds of H2SO4, 98 % and that is probably more than I will spend in a lifetime.
But now I have some clues of where to look...it could be classified as fertilizer
I know one thing that tops that :
having a flask of boiling H2SO4 shatter and forgetting to remove my jar of NaClO3 standing just under it...
you probably have to be very very careful during heating making sure that no thermal stress is present. I might consider trying to purchase a
quarts-flask, but glassware in quartz are very expensive.
Thanks for all the input.
/rickardrikkitikkitavi - 5-11-2002 at 09:08
I would just like to add that I have found several interesting patents covering of how to make sulfuric acid from SO2 gas+ O2(air) , either by
homogenous catalyzation in the liquid phase (HNO3 , 0,1-15 M) , electrolysis of a aqeous SO2 solution by graphite electrodes e t c..
I will probably make a compilation if there is interest for this, that way I can get a better overlook myself too.
I will focus on the patents utilising SO2, air, O2 electricity and easily accessible
catalysts or other easily accessible compounds.
I dont think I need all the trouble of vacuum distilling ethylhydrogensulfate or a messy solution of this. But keep the topic alive of course!
The main reason why these patents arent in use in a industrial process is that they all produce diluted sulfuric acid < 60 %, which has to be
concentrated to 98 % .
This concentration process is expensive and uses lots of energy. A metric tonne of 98 % sulfuric costs about 3-40 US $ in bulk! (calculate the
overprice for drain cleaner he he , theres got to be a LOT of middle men taking their cut)
The usual contact process actually produces 98% acid direct and a lot of energy , both as electricity and hot water.
A 300000 tpa plant outputs about 30 MW! of energy which can be sold off.
But energy is not of a concern for us of course. that is cheap and plenty.
/rickard
let's see the info!
Polverone - 5-11-2002 at 10:25
I am definitely interested in seeing what you've dug up. Even though I don't think I will have to use any of these methods any time soon, they still
interest me. Non-electrolytic methods are of greater interest to me because electrolysis is tedious and messy, though versatile. I'd certainly like to
hear more about the HNO3-catalyzed method. Have you yet had any luck in locating an inexpensive source of sulfur?rikkitikkitavi - 5-11-2002 at 12:22
the patent describing a HNO3-liquidphase oxidation is US patent :5788949
It seems like a very simple method to do on a lab scale.
I have found out that at www.depatisnet.de
(german patent office) you can search all major national databases , however only the first page (or just a patent number ) is viewable.
You can even search for words in the abstract text (H2SO4 gave over 4500 hits!)
At USPTO you can only search for patent# or class for patents older than 1982.
It seems like the amount of information is overvelming, so many patents... argh I m drowning in information. It will take some time to compile all
this
/rickard rikkitikkitavi - 7-11-2002 at 11:14
I have to correct myself about the patent search @ www.depatisnet.de
You can view all the pages in the documents, making it a invaluable tool for patents searches world wide!
/rickard
Sulfur
Richy - 4-8-2003 at 23:44
"Have you yet had any luck in locating an inexpensive source of sulfur?"
Sulfur should be avaliable at almost every nursery you visit. 2kg for $5.
Back to sulfurous acid... 1g of H2O2 would oxidise 2g of H2SO3 to 2g H2SO4, right?
The readily avaliable H2O2 is a 2% soln.... 200ml for $3 AUS.... you'd need to pay $3 to oxidise 4ml of H2SO3 to H2SO4. Rediculous.
I've heard that a 40% H2O2 soln may be avaliable from pool stores. $20 AUS for 1lt. 400g H2SO4 for $20 (considering water and sulfur cost
basically nothing) sounds okay to me.
How about oxidation by the O2 in out atmosphere? I suppose it would be practical to leave sulfurous acid, prepared from SO2 + H2O, then leave it in a
dish with a wide surface area for 10+ days.
Then one could distill off the H2SO3 (i was looking for the boiling point but couldnt find it, does anyone know... i suppose it would be lower, less
stable molecule) and water for nice conc. H2SO4.
Then one could re-expose the H2SO3 distilled off to the atomsphere once again for more. Nice to know 100 parts S yields 300 parts H2SO4!
peace
Richy
[Edited on 5-8-2003 by Richy]blip - 5-8-2003 at 14:59
I read about destroying chlorate with SO<sub>2</sub> and SO<sub>3</sub><sup>2-</sup>. It's likely dangerous,
but also likely produces sulfates/sulfuric acid along with just salt leftover if you don't use too much excess. Seeing as I can produce lots of
chlorate easily, that seems like a feasible solution for me. How well does electrolyzing H<sub>2</sub>SO<sub>3</sub> work?BromicAcid - 5-8-2003 at 17:42
When concentrating H2SO4 by boiling it down I use corning glass cookware. The sauce pans easily withstand burner to oven switching and they hold up
exceedingly well to concentrated acid, as another plus they are very thick and tough. As for the ultimate yet unwanted test for thermal shock I put
one of these pans though, I was making cast thermite using molten sulfur as a binder, nearly 400 g of the material went up in the pan shortly after
the incorporation, molten metal flew out and I flung the pan away. Sure, it was covered in drops of metal but it did not crack, the only damage was a
thin coating of glass that came out when I chipped the majority of the remaining mass from the bottom. Truely a mythic item of kitchen chemistry.
Na2S2O8
chemoleo - 6-8-2003 at 17:00
sodium peroxodisulphate (the stuff used for making electronic circuit boards etc to dissolve copper, so u can get it anywhere) produces vast amounts
of white SO3 when molten and heated. Maybe thats a nice route for the production of SO3...
by the way its also easy to produce Na2S2O8 via electrolysis of NaHSO4/SO4
[Edited on 8-8-2003 by chemoleo]
Apparatus
Richy - 8-8-2003 at 04:34
ive been talking to someone who claims that to form sulfurous acid via bubbling SO2 in to H2O would require sophisticated pressure equipment... is
this true?
i dont see why the two wouldnt react just by simply vaporising sulfur in a flask with a hose leading into water... the S forms SO2 with the air in the
flask and tube, and forms H2SO3 when it comes in contacts with the water...
wouldnt that be okay by itself????
1st attempt
Richy - 10-8-2003 at 03:18
Just attempted bubbling SO2 gas through water to form sulfurous acid. Heated the flask filled with sulfur to 200C and sulfur began to produce yellow
fumes in the flask. Bubbles appeared in the bottle of water, but the reaction was a complete failure. The bubbles were simply the air in the flask
being expanded through the hose. a pH test confimred that no sulfurous acid was formed at all, and now one of my flasks is rendered useless because
the molten sulfur hardened to form a moon-rock type substance which will never come out in a million years. I might try to re heat the rock to molten
and pour that shit out. ill see if i can locate sodium peroxodisulphate, would it be found at an electroplating supply store? if not where might it be
obtained? im also interested in the nitrates route, i have access to ammonia nitrate which could easily be converted to NaNO3, but how would H2SO4 be
created?
thanks for everyones input, ill have to keep trying
sulphur
mick - 15-10-2003 at 12:10
Leave the sulphur to soak in toluene or xylene, paint stuff, and it should come loose after a week.madscientist - 15-10-2003 at 19:04
Acetone works very well for loosening sulfur fused to glass.Microtek - 16-10-2003 at 02:04
Just heating the sulfur is not going to work; you need to burn it in a stream of oxygen ( or air ) and then pass the gaseous product through water.
SO2 dissolved in water smells extremely strongly of SO2, so I think the SO2 + H2O <--> H2SO3 equilibrium isn't very far to the right, which
means that heating this would expell SO2 from the soln and push the equilibrium even further to the left.Blind Angel - 19-10-2003 at 15:59
Maybe a bit off-topic, but is bubbling Cl<sub>2</sub> through H<sub>2</sub>SO<sub>3</sub> would yield
HSO<sub>3</sub>Cl and HCl ? or maybe the addition of HCl would give H<sub>2</sub> + HSO<sub>3</sub>Cl?
I'm searching a way to chlorosulfonic acidrikkitikkitavi - 20-10-2003 at 09:05
bubbling Cl2 through H2SO3 (or rather a solution of SO2 in H2O, since H2SO3 exists only in small amounts) gives
HCl(aq) and H2SO4 (aq)
/rickardGeomancer - 21-2-2004 at 19:31
Has anybody actually succeeded in using the contact process? I believe axehandle, among others, was thinking of it. I've been thinking of using
either a platinum or platinum/rhodium catalyst (I don't know where to get vanadium salts). The sticking point is temperature control. Materials
are also a problem. Anyone know a good source for stainless steel tubes?
Along similar lines, are there any good liquid phase oxdations of sulfur? I know that boiling HNO3 works, anything else?BromicAcid - 21-2-2004 at 19:58
Quote:
Materials are also a problem.
If you live within fifty miles or so of a large city there is a very good chance there is a scrap yard there. Although the temperature might be a
problem, lead has been used in H2SO4 production since way back in the day (chamber acid times). It holds up well to this, plus it is easily melted to
coat your apparatus. Currently lead in my closest scrap yard is $.21 a pound, of course, being a chemist when I found that out I bought forty pounds,
I've used about 100g so far in the last year. I have yet to find a scarp yard that has any quips about selling to an individual.
