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Author: Subject: Sulfuric Acid at Home
agorot
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[*] posted on 25-1-2010 at 19:05
Sulfuric Acid at Home


As many here are, I am a home chemist, and I therefore don't have access to a chemical like sulfuric acid with great ease. Sure you can buy it some places, but who knows for how much longer. I live in America and we just had another attempted terrorist bombing over Christmas break, and so I'm trying to devise a way that I can make sulfuric acid at home.

Not an easy task. And dangerous. Very dangerous. But with careful planning and extensive collaboration, I think we can come up with a "device" that can help manufacture the acid at home in very small amounts. Right now when many of us can still buy the acid, this doesn't seem like such a high priority project, but soon when we won't be able to purchase H2SO4 online anymore (and other sources like car batteries become just too expensive for the amount of acid at relatively weak concentrations in them), we will need to be able to produce it at home.

Lead metal is resistant to corrosion from sulfuric acid, and I'm assuming by extension that it is relatively resistant to SO3, or at least more so than many materials. Lead was used in the Lead Chamber Process to manufacture H2SO4.

So at home I was going to make SO3 by doing simple distillation. I have a very good set of distillation glassware, all 24/40 ground glass joints. I have a 1000 mL boiling flask, a meter long grahm condenser, and all the right joints, stoppers, and bent adapters to make it a closed system. I was going to heat NaHSO4 to

1) Dehydration at 315°C:
2 NaHSO4 → Na2S2O7 + H2O (and then get rid of all the water once dehydration is complete)
2) Cracking at 460°C:
Na2S2O7 → Na2SO4 + SO3

I would then condense the SO3 in another 1000mL round bottom flask, possibly by using dry ice in an acetone bath to get the temperature VERY cold and limit the amount of SO3 that could build up pressure on the receiving end of the system).

So once I have the SO3, I would need to make oleum. I do have about 100mL of conc. H2SO4, so I could add the SO3 to that and start very small. The problem is I need to do this in a very controlled way.

I was thinking of buying some lead metal sheets and copper pipe, and I have a small propane welding torch at home (available at most hardware stores). The melting point of lead is only about 330 degrees, so I was going to get some copper pipe and line the inside of it with lead by melting the lead onto the copper (copper melts at about 1000 degrees, so I wouldn't have a problem with it actually liquefying, but I would just need to be sure that the metals don't mix too much like when you are making an alloy so the copper could be corroded by the acid).

The shape of this copper/lead pipe would be such that it would have three compartments--one for conc H2SO4, one for SO3 (s), and one for H2O, and I could separate all three and operate a lead valve to open and close each compartment from the other two. I can't explain it well in this message, nor can I draw it, but just imagine that I have such a system devised shaped somewhat like a Y with the starting places of the chemicals as the conc sulfuric on bottom and the SO3 and H2O on the top two compartments.

So I would fill the compartments with their respective reactants with their separating valves closed. Then I would seal it off completely and put it in an ice bath. Then I would open the valve so that the SO3 could fume and start to fall down the Y-shaped "device" and dissolve in the conc H2SO4 to become oleum.

Once the oleum was made, I would flip the Y shaped device upside down and open the valve so that the oleum could slowly drip into the water and then collected and drained as conc H2SO4.

Of course, I would use correct stoichiometry to measure everything out correctly. The device would remain completely closed while in operation so no nasty fumes or anything like that to deal with.




What do you think? Do you think that the temperature will get so great even in the ice bath that the device would explode? I would have thick, lead-lined copper pipe...and I don't think the reaction could get so hot on the inside that it would begin to melt the lead if I had it in an ice bath especially.

Again, if you don't have expierience dealing with this sort of stuff, its scary. Look at this video of SO3 on a piece of chicken, and imagine if you replaced the chicken with your hand....
http://www.youtube.com/watch?v=WqFj8xuaH7M

[Edited on 26-1-2010 by agorot]

[Edited on 26-1-2010 by agorot]
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[*] posted on 25-1-2010 at 19:33


First of all, I believe you can simply pipe the gaseous SO3 into the H2SO4 (contact process); you don't need to condense it first.

If you do decide to make parts of your apparatus out of metal, your biggest issue will be joining the metal to the glass in a way that will resist SO3 while still allowing for thermal expansion. There is some good information in some of the old books in the scimadness library on joining glass and metal (I would start with Brauer). However, I think it would be simpler to just make the whole thing of glass.

Oh, another very important thing: you seem to be thinking of making this a closed system (no outlet for overpressure). Doing that would be insane. Even if you have a very cold receiver, you can't guarantee that other gases (air) don't overpressurize the system, or that small amounts of impurities wouldn't result in other gases building up, or for that matter that the heat transfer is rapid enough for condensation to keep pace. You would want a pipe leading out through a bottle (or bucket, depending on your scale) of some neutralizing solution like sodium bicarbonate.

Also you seem to be thinking of baking some dry solid in your glassware at 500C or so. How are you planning to clean the baked-on sodium sulfate from your glass? Do you think that heating your 1000ml flask to such temperatures is safe, or for that matter that you can get the solid inside the flask to 500C before the borosilicate glass begins to soften?

I believe some member here has posted their own manufacture of oleum; a search would turn it up. There is also a stickied thread on the manufacture of sulfuric acid via the lead chamber process, which has a lot of useful information. Finally, if you have NaHSO4, you can disproportionate it to H2SO4 and Na2SO4 via ethanol and water, and then carefully concentrate the sulfuric acid (boiling off the ethanol); which, so long as you only needed concentrated H2SO4 and not actual oleum, would be simpler than your method.

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[*] posted on 25-1-2010 at 19:39


Mmm. I just love chicken l'oleum. :D Search and study the work of garage chemist and len1 in making oleum.



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agorot
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[*] posted on 25-1-2010 at 19:45


I was planning on fist making the SO3 in the distillation apparatus and then adding the SO3 (after its all made and cooled) into the metal Y device (entirely lead-lined copper).

Now that you understand that the only place where pressure is going to be an issue is in the metal "Y" which would be quite durable, do you think I would still have the overpressure problem? I guess I could install another valve or something, but then I would lose some oleum, and that would not be good. I was planning on doing this at low temperatures and in small amounts, but perhaps you're right. I should try to fix that.

Sodium sulfate is soluble in water, and although it would be baked in well to the glassware, I was thinking that letting water sit in the flask and scrubbing, even heating it a little would probably dissolve it after several rinsings.

I think that my borosilicate round bottom flasks could handle the near 500 degree temperatures necessary for a time. I've heated the flasks before on a very hot butane flame before, and I'm sure its gotten quite hot. The flasks are heavy duty.
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agorot
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[*] posted on 25-1-2010 at 20:39


i just found len1's post. you don't need to respond to this anymore :)
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[*] posted on 2-2-2010 at 11:45


Making H2SO4 just doesn't get any easier than this:

http://www.youtube.com/watch?v=okvvD3-DF9U

No high temps, no SO3 in the atmosphere. Metabisulphite is easy enough to acquire since home wine makers and home brewers use it for cleaning their equipment and bottles.
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agorot
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[*] posted on 2-2-2010 at 14:19


Yeah, but that method is neither efficient nor does it make a concentrated solution. Plus, the hydrogen peroxide is not exactly cheap nor is it readily available in decent concentrations. I want to use relatively easily obtainable materials to make this acid as concentrated as possible at home, and I have the access to SO3 and high temps.

I could use another oxidizer, but its just not that efficient, and metabisulfite is something I can't buy at a regular store, but I can get bisulfate and make so3 that way

I do really wish there was a video though of someone working with oleum and diluting it or even (this would be somewhat dangerous) SO3 directly into water. I wouldn't try that second method before I saw someone else do it successfully.

[Edited on 2-2-2010 by agorot]
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[*] posted on 2-2-2010 at 15:31


Try searching youtube :p http://www.youtube.com/watch?v=QOKX6Dn-K_w

I hope you change your mind the moment he opens the bottle




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[*] posted on 2-2-2010 at 15:32


Quote:
I do really wish there was a video though of someone working with oleum and diluting it or even (this would be somewhat dangerous) SO3 directly into water.

AFAIK, when solid SO3 is added to water directly a violent reaction occurs which spatters hot acid in all directions and when gaseous SO3 is bubbled into water there is little absorption but a sulphuric acid mist forms and this mist lingers and is difficult to condense.
H2SO4 however, readily absorbs SO3 vapour to form oleum; this is then diluted to 98% by adding it to water.






[Edited on 3-2-2010 by hissingnoise]
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[*] posted on 2-2-2010 at 16:26


Quote: Originally posted by agorot  
Yeah, but that method is neither efficient nor does it make a concentrated solution. Plus, the hydrogen peroxide is not exactly cheap nor is it readily available in decent concentrations. I want to use relatively easily obtainable materials to make this acid as concentrated as possible at home, and I have the access to SO3 and high temps.

I could use another oxidizer, but its just not that efficient, and metabisulfite is something I can't buy at a regular store, but I can get bisulfate and make so3 that way

I do really wish there was a video though of someone working with oleum and diluting it or even (this would be somewhat dangerous) SO3 directly into water. I wouldn't try that second method before I saw someone else do it successfully.

[Edited on 2-2-2010 by agorot]


In the original post, you wrote: "But with careful planning and extensive collaboration, I think we can come up with a "device" that can help manufacture the acid at home in very small amounts".

How much do you really want to make, and how important is efficiency? If you don't want to bother with heating sulphates, or don't want to deal with SO3 (understandable considering how nasty that stuff is) then running SO2 into hydrogen peroxide or nitric acid fits the description: make small amounts at home easily. (The latter doing pretty much what the "lead chamber" process does. In either case, you can always boil it down to get the concentration you need.)

As for nitric acid, you can make that as well if you have a high voltage xfmr from a microwave oven or a neon sign xfmr.

If you're talking making liters at a time, then catalytic oxidization of SO2 really is your best choice.
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[*] posted on 2-2-2010 at 17:07


Efficiency isn't the most important thing, just as long as I'm not wasting a ton of expensive reactant. That's why the bisulfate method is adventageous because its $10 for 5 lbs at my local grocery store. I actually was thinking of making a lead-lined copper reaction vessel because I thought maybe it would be the safest bet to control heat and pressure. The whole apparatus would have to be kept cold of course.

Did you read my first post? I was thinking of dissolving SO3 in conc H2SO4 (I have like a quarter of a liter I got from distilling drain cleaners to separate the acid from the buffers in a 24/40 distillation system using a graham condenser--this was expensive and very inefficient because the buffers significantly increased the boiling point of the acid and after a few minutes no more acid would come over at ~450 degrees, and I didn't want to push my glassware further).

This would make oleum. I want to do small batches at a time, especially at the beginning. I was going to make like 20 percent oleum and then cool this significantly in a dry ice/acetone bath, then I would open up a valve in the container remotely so that oleum could drip into equally cold distilled water.

If I got the valves to function properly and this was done in a lead lined vessel welded together and done at very low temperatures, I think I have a chance. But I want to plan this out well first so I don't kill myself doing it.

So do you think this is feasible?
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[*] posted on 2-2-2010 at 21:49


Is there a way to make sulfuric acid with nitric acid ?

Because those plasma reactor chambers that some guys on this board have made look really amazing.
Electricity,water and air in.. Nitric acid out :)

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[*] posted on 3-2-2010 at 00:41


I wonder if electrolysis, at a suitable voltage with suitably resistant electrodes, of aqueous sulfite or metabisulfite solutions, which could be obtained by burning sulfur to SO2 gas and leading this into an alkaline solution, would result in oxidation to sulfate? Has anyone tried this?
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[*] posted on 3-2-2010 at 05:18


Sulfurous acid (SO2 in water - makes a clathrate as I learned when I did it;)) oxidizes in water during freeze/thaw cycles to a much stronger acid than it starts at. I got the idea from the abstract of a Japanese article and the end result, after it was warmed to RT and no clathrate remained, sent my pH paper off the scale. About the easiest route I know, especially here - we ain't in Kansas anymore - you know, where H2SO4 just ain't available, neither is HEET.
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[*] posted on 3-2-2010 at 05:29


bquirky, could you post a link to the thread you're talking about. I'm new to this forum and haven't read that thread yet but I'd love to

johnww, I could try electrolysis of bisulfate this weekend. definitely something i'm doing outside just in case. I haven't built myself a fume hood yet.
this is what I think i'd get

I have h+,na+, and SO4-2
anode: place where possilby H+ and Na+ would be reduced to H2(g) and then maybe hydroxide? depends on what is produced at the cathode
cathode: SO4-2 to SO3? if so3 was produced it would react violently with the water...and if there are any hydroxide ions floating around then the acid would be immediately neutralized. maybe these ion interactions could be stopped if I used a salt bridge or similar?
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[*] posted on 30-4-2010 at 20:23


Sulfuric acid can be produced by burning sulfur and absorbing the SO2 in water
to form sulfurous acid using the setup shown in the second part here _
http://www.youtube.com/watch?v=2gXByJkg0iY
The liquor produced is then further oxidized into sulfuric acid in a second step.

Note * Using H2O2 instead of plain water one can get an initial concentration
SO2 + H2O2 => H2SO4
the resulting percentage of H2SO4 depending on the H2O2 percentage used.

http://books.google.com/books/download/Modern_inorganic_chem...
Modern Inorganic Chemistry - Mellor 1912
Chapter XXIII pg.418
Compounds of Sulphur with Oxygen
Preparation: Sulphur dioxide is formed when sulphur burns in air . .
Between 6 - 8 per cent of the sulphur is simulataneously oxidized
to sulphur trioxide."
Absorbed into 12 % ( 40 volume ) H2O2 , 18 - 20 % of the solution will
become sulfuric acid to start with , the balance being sulfurous acid.

________________________________

In the second step the liquor is heated in a sealed tube ( pipe bomb ) at over
150 ºC. the solution deposits sulphur as the sulfurous acid becomes H2SO4
3 H2SO3 => 2H2SO4 + H2O + S
Precipitated sulfur can be harvested , washed and dried to be burned again.


A solution of sulfurous acid, heated in the absence of oxygen,
disproportionates into sulfuric acid and free sulfur.
This is the only mention of it that I can find _
http://www.terrapub.co.jp/journals/GJ/pdf/0101/01010045.PDF

.
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[*] posted on 1-5-2010 at 00:46


Quote: Originally posted by JohnWW  
I wonder if electrolysis, at a suitable voltage with suitably resistant electrodes, of aqueous sulfite or metabisulfite solutions, which could be obtained by burning sulfur to SO2 gas and leading this into an alkaline solution, would result in oxidation to sulfate? Has anyone tried this?


Sodium-metabisulfite is a food-processing raw-material, 20 $ for 25 kg ...
==> It gives SO2 with acids ...

About the electrolysis I wondered too ..
==> Might lead-electrodes work ?
==> Probably carbon-rods should work, since these even withstand the chlorate-electrolysis for a while ...

The charge-consumption would be 2 electrons for SO2 > SO3
==> Efficiency ??

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[*] posted on 1-5-2010 at 10:38


Quote: Originally posted by franklyn  
... Note * Using H2O2 instead of plain water one can get an initial concentration
SO2 + H2O2 => H2SO4
the resulting percentage of H2SO4 depending on the H2O2 percentage used.


You don't want to use a H2O2 that is too strong though, that could result in a steam explosion or worse. SO2 bubbled into 35% conc. H2O2 can get very hot (at one point I've measured 105 C on a Hg thermometer). Forming H2SO3 alone and attempting to oxidize or decompose seems like the bigger waste since SO2 only has such a limited solubility in water.
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[*] posted on 1-5-2010 at 14:33


@ Formatik
Your hands on experience trumps what I may speculate,
this is what I have been trying to convey _

1. Bubble SO2 into very cold H2O2 while maintaining cold.

2. Seal acidified liquor into pipe and heat for some hours at ~ 200 ºC
_ see reference for particulars.

40 - 50 % H2SO4 certainly seems achievable , and that can be fortified.

* I don't have reference or data to relate at hand , but I believe SO2
is better solvated by even low concentration H2SO4.

* 30 % H2O2 can be had at reasonable cost but still is as expensive as
H2SO4 would be. 40 volume is very much OTC.

If a home brewed method of oxidizing SO2 to obtain SO3 can be
optimized , the problem is solved.

.
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[*] posted on 1-5-2010 at 14:38


Did anyone bother to read this thread?
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[*] posted on 2-5-2010 at 01:12


In the video he mentions concentrating the sulfuric acid to 95% concentration. Is this verified? I think I read somewhere the maximum achievable concentration through evaporation under atmospheric pressure is around 80%.
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[*] posted on 9-5-2010 at 00:20


When heating to concentrate Ive read several places 80% is about max concentration before a considerable amount of H2SO4 is lost to the atmosphere and 93.3% is maximum concentration by this method @ atmospheric pressure,maybe why tech grade is this concentration. FwIW ,so concentrating by heat to 95%/ 98% as Ive heard elsewhere seems to defy the laws of chemistry.
Please correct me if Im wrong!

[Edited on 9-5-2010 by grndpndr]

[Edited on 9-5-2010 by grndpndr]
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[*] posted on 9-5-2010 at 03:09


http://www.sciencemadness.org/talk/viewthread.php?tid=2824

It's stickied for you on the forum Sulphur to Sulphuric by the old fashioned method. Once lead was the only material to build the equipment now a polythene drum is good enough.

Cycle the process using a demijohn (wine making!) or other moderate size container and get sulphuric out at the bottom.

[Edited on 9-5-2010 by Contrabasso]
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[*] posted on 9-5-2010 at 04:25


grndpndr, I've always assumed that ~98% H2SO4 is what you get by boiling dilute acid.
I haven't, so far, read anything that contradicts this.
At its b. p., H2SO4 dissociates to SO3 + H2O and this would account for the ~2% water.
I also assume that the acid will absorb *some* H2O from moist air as it cools.
Doing the boiling when the air is dry/cold should give a higher concentration. . .


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[*] posted on 9-5-2010 at 09:19


Quote: Originally posted by grndpndr  
so concentrating by heat to 95%/ 98% as Ive heard elsewhere seems to defy the laws of chemistry.
Please correct me if Im wrong!
In his Treatise on Chemistry (page 332 of volume I) Sir Henry Roscoe states that sulfuric can be concentrated to 98% by driving off the water by heating in Pt or glass vessels.
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