Sciencemadness Discussion Board

Sulfuric Acid at Home

agorot - 25-1-2010 at 19:05

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]

bbartlog - 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.

Magpie - 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.

agorot - 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.

agorot - 25-1-2010 at 20:39

i just found len1's post. you don't need to respond to this anymore :)

Foss_Jeane - 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.

agorot - 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]

mr.crow - 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

hissingnoise - 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]

Foss_Jeane - 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.

agorot - 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?

bquirky - 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 :)

JohnWW - 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?

un0me2 - 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.

agorot - 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?

franklyn - 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

.

chief - 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 ??

Formatik - 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.

franklyn - 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.

.

Lambda-Eyde - 1-5-2010 at 14:38

Did anyone bother to read this thread?

Cloner - 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%.

grndpndr - 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]

Contrabasso - 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]

hissingnoise - 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. . .


entropy51 - 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.

grndpndr - 8-6-2010 at 23:45

Not a 100yr old book+ but a modern chemistry treatise w/references not as "i was led to believe" OR ANOTHER METHOD beyond simple heating in a beaker as is so often discussed here to concentrate H2so4 to 98% unequivacly
Dont blow me off on this one guys several modern sources state unequivacly the azetrope of a h2so4/water is 93.3% @338C at ambient pressure.Others vjust as insistent azetrope
is 98.3 at the same temp.

Beins this is an honest chemistry forum above the average lets have honest answers.How do you beat the stated azetrope with heaT and a beAker as has been so often reccomended possible?:(
If azetrope means what i assume above this temp as much h2SO4 is lost as H20.Others are commenting on this apparent innacuracy.Please dont shove me off to a treatise ill never find but use qoutes (RECENT) and explanations as I have mine simple sweet and I think, irrefutable given the circumstances.If the inflated info's BS let it die,this is a science forum.Totse/utube needless to say arent acceptable evidence.

To p0revent useless bickering ive cut to the chase it seems from what ive found theres a big discrepancy between modern information on the net (reliable sources) listing 98.3 and 93.3? aS THE AZETROPE AT IDENTICAL TEMPS,AMBIENT PRESSURE WASNT MENTIONED EXCEPT IN THE 93.3 FIGURE.
Now whaT?I Hope in a civil manner WHATEVER the outcome.Ive really no dog here ID JUST LIKE TO SEE THE TRUTH OR IS THERE A PROBLEM WITH METHODS OF MEASUREMENT/pressure?aLTHOUGH A 5% DIFFERENCe cant be ignored.

In retrospect there does seem to be more ofa consensis towards 98.3 but no pressures mentioned.No offense anyone just a interest in the truth.:)

New info added/some incorrect retracted.

[Edited on 9-6-2010 by grndpndr]

BenZeen - 9-6-2010 at 01:48

electrolysis of copper sulfate soln

grndpndr - 9-6-2010 at 02:12

I dont pretend to understand this but the handbook of chemistry also Design of adibatic Flash units, Ekberg.
Of course pressure has a difference maybe 1.5 % or thereabouts between 1bar and .05bar.Seems the main difference is again something I dont grasp is the composition measurements.The lower 93.3 azetrope seems measured like this,( mole %H2SO4), the higher 98.3,
(%w/w H2SO4), If anyone grasps this I would appreciate an explanation.If its a wrong turn also appreciate it.I was under the impression this was quite straighforward with some misunderstanding?
This was taken from theDesign of Aliphatic Flash Units table 1 Ekberg. HELP!Not that the small difference will have alot of difference in my backyard chem its an interestin point?

edited ;revisions content as better came to light.

[Edited on 9-6-2010 by grndpndr]

[Edited on 9-6-2010 by grndpndr]

hissingnoise - 9-6-2010 at 02:37

grndpndr, the figure 93.3% for conc. H2SO4 is obviously a typo - 98.3% is the correct value. . .

grndpndr - 9-6-2010 at 07:31

ok?THERE SURE IS ALOT OF TYPOS OUT THERE TO INCLUDE MODERN TEXTS.SAY APPROX 1/3 CONTAIN THE AZETROPE OF h2SO4 AS 93.3 @AMBIENT TEMP.tHE VERY SAME TEXT/SAME TABLE LISTED THE 2 AZETROPES AT 93.3 THE OTHER SIDE 98.3.tHE ONLY DIFFERENCE A MEASUREMENT DIFFERENCE AS i TRYED TO EXPLAIN ABOVE.iTS CONFUSING AT BEST ONE WOULD THINK A STD WOULD BE THAT HARD TO RECOGNIZE.:(

[Edited on 9-6-2010 by grndpndr]

[Edited on 9-6-2010 by grndpndr]

rrkss - 9-6-2010 at 07:38

I've done concentration of sulfuric acid using boiling. When titrated using a mohr pipet and Fisher 1.00 N NaOH solution, 1 mL of H2SO4 requires 37.0 mL of NaOH solution to make the phenolphalein turn pink. To me this gives me proof that my solution is at the 98% concentration and not 93%.

[Edited on 6-9-10 by rrkss]

agorot - 9-6-2010 at 08:06

its 98% max. really.

bbartlog - 9-6-2010 at 08:21

As you suggested, the 93.3% figure is based on mol/mol rather than weight/weight. So in one hundred moles of the azeotrope there will be 93.3 moles of H2SO4 and 6.7 moles of water. Such a mix would contain

93.3 x 98 ~ 9144 g H2SO4
6.7 x 18 ~ 114g H2O

for a total of 9256g, of which less than two percent by weight is water. On the other hand, my calculation suggests that this would end up being somewhat more than 98.3% by weight, so either I've made a mistake or else I really don't understand the mol% system; but I'm quite sure that the different percentages are related to the different measures, not to some mistake. Presumably there are circumstances where the mol/mol measure is more convenient for calculation.

grndpndr - 9-6-2010 at 10:58

Thanks all and particularly barrtlog for the expl i was looking for couldnt express well,(could it be a diff in atmospheric pressure)Also Q.S for the quick dirty Idea of a 1.84 auto hydrometer.Thanks all:)Auto supply hyd to cheap or expensive sciplus $9hyd and my own 100ml grad cyl.

Im curious the difference might not be bar's? Between 1bar and .o5bar the difference in is 1.5% concentration give or take.
im at 4500ft so..well see when the hydrometer gets here.

\One last thing do I have to achieve boiling to concentrate my h2so4 to azetropic.My inex. worn hotplate wont manage even the 290C BP?:( concerned I may need to go to a propane burner/asbestos.:(


[Edited on 9-6-2010 by grndpndr]

MttLsp - 11-7-2010 at 19:59

Here's a method of making H2SO4 I just successfully completed. I'm not sure it's economical, but I'm new to chemistry so at least I'm learning something.

Basically I just add stoichiometric amounts of Copper and Sodium Bisulfate in a test tube and heat it and the following reaction takes place:
4NaHSO4 + heat ---> 2Na2S2O7 + 2H2O
2Na2S2O7 + heat ---> 2Na2SO4 + 2SO3
2SO3 + Cu ---> CuSO4 + SO2

Then I run the sulfur dioxide through hydrogen peroxide (3%) to make about a 2.9% H2SO4 solution and cook it down.

In the end I'm left with a nice chunk of copper (II) sulfate in the test tube I can grow crystals and concentrated sulfuric acid.

If anybody is interested here are the amounts I use as a guide:
NaHSO4 = 1 g
Cu = .132 g
H2O2 = 2.362 mL

PS. I strip copper electrical wire instead of using copper flashing since the wire has more surface area.

Lambda-Eyde - 11-7-2010 at 23:35

Err - why not bubble the sulfur trioxide directly into water to make sulfuric acid?

Contrabasso - 12-7-2010 at 01:33

Sulphur trioxide doesn't easily dissolve in water without forming an acid mist that gets everywhere you don't want it! So it's usual to form or buy sulphuric acid and add sulphur trioxide to it -which in practise works better. The product is a strong sulphuric acid or oleum, this is then extracted and the concentration adjusted to usable or saleable specifications.

Formatik - 13-7-2010 at 11:27

Quote: Originally posted by MttLsp  
Here's a method of making H2SO4 I just successfully completed. I'm not sure it's economical, but I'm new to chemistry so at least I'm learning something.

Basically I just add stoichiometric amounts of Copper and Sodium Bisulfate in a test tube and heat it and the following reaction takes place:
4NaHSO4 + heat ---> 2Na2S2O7 + 2H2O
2Na2S2O7 + heat ---> 2Na2SO4 + 2SO3
2SO3 + Cu ---> CuSO4 + SO2


The copper may have reacted with the bisulfate directly:

4 NaHSO4 + Cu = CuSO4 + 2 Na2SO4 + SO2 + 2 H2O

This is not an efficient use of the sulfate. Crystallization of some Na+ salt isn't hard, but you'll still have impure CuSO4.

If the pyrosulfate intermediate was needed then the reaction might need something like a quartz vessel to be useful, because when you are doing the reaction in larger amounts, decomposition of pyrosulfate happens only significantly with blow torch temperatures or an electric heating coil, etc. Don't know if the copper changes this.

Carbon can be used to reduce Na2SO4 to SO2 (alongside some Na2S), but this needs temperatures above 1000 C to give off SO2 at a significant rate. NaHSO4 might reduce similarily. Though I'm sure there are reductants that should work at lower temperatures. It should be tested carefully, e.g. at red glow aluminium decomposes Na2SO4 under detonation - Ch. Tissier, A. Tissier (C.r. 43 [1856] 1187; J. pr. Ch. 71 [1857] 76).

NaHSO4 should be converted to H2SO4 with aq. HCl. I've done it with Na2SO4 and HCl, but could only convert a bit less than half of its sulfate to sulfuric acid. I had to also distill the H2SO4.

[Edited on 14-7-2010 by Formatik]

MttLsp - 15-7-2010 at 18:32

Oh wow I didn't even think about how Na2SO4 is left over. And that HCl + NaHSO4 is a good idea, but the distillation seems dangerous unless you had a temperature controlled hotplate or mantle. I actually used table salt and hydrochloric acid to make the sodium hydrogen sulfate. But I have a question:

Is there a way to seperate the Na2SO4 and CuSO4 in the reaction i suggested. Perhaps by combining the SO3 from the NaHSO4 with Cupric Oxide in a seperate flask and then running the SO2 into the H2O2 as previously stated? Or would this still be inefficient?

Formatik - 16-7-2010 at 07:56

Quote: Originally posted by MttLsp  
Oh wow I didn't even think about how Na2SO4 is left over. And that HCl + NaHSO4 is a good idea, but the distillation seems dangerous unless you had a temperature controlled hotplate or mantle.


You would need a strong hot plate to distill even small amounts of H2SO4. I was actually heating less than 2mL under the bunsen burner, got impatient and starting heating with an additional second burner! Its high boiling point makes this problematic at regular atmospheric pressure.
Quote:
I actually used table salt and hydrochloric acid to make the sodium hydrogen sulfate.


Usually H2SO4 and table salt, or Chile saltpeter is used to make it. But it makes no sense to start out with H2SO4 only to end up later with H2SO4 with a lower yield.
Quote:
Is there a way to seperate the Na2SO4 and CuSO4 in the reaction i suggested.


Crystallization as mentioned. Again, it only leaves you with partly pure CuSO4. You could also find a solvent that one is soluble in, but the other is not (good luck).
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Perhaps by combining the SO3 from the NaHSO4


Again, try decomposing large amounts of NaHSO4 by itself, and you'll see its pyrosulfate doesn't decompose significantly until you start heating with temperatures that soften and ruin normal borosilicate flasks, beakers, etc (see thread on decomposing NaHSO4 in prepublications on why quartz is used).
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with Cupric Oxide in a seperate flask


That ought to leave you with CuSO4. SO3 + CuO = CuSO4.

AndreiChim - 15-6-2011 at 11:56

Hi. Could you give more details to the NaHSO4-ethanol/water method you described earlier? It sounds interesting. Thank you.

As for availability...

albqbrian - 15-6-2011 at 16:08

Luckily I see SA as a chemical that will be widely available. If for no other reason than its use in biodiesel production. Check out suppliers in that area for a variety of chems. Like:

www.dudadiesel.com