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Author: Subject: H2SO4 by the Lead Chamber Process - success
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[*] posted on 17-6-2009 at 16:04


Catalysts are reaction specific, i.e. what used as a catalyst in a catalytic converter (platinum) for the reaction of CO with O2 to form CO2 will not necessarily be the same as that used for the reaction with 2SO2+O2 - > 2SO3. I don't think (but I may be wrong) that the catalyst in the converter will work for this reaction.
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entropy51
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[*] posted on 17-6-2009 at 16:18


Indeed you may be wrong. :o Catalytic converters contain platinum, palladium, rhodium and other catalysts. At least some of them will catalyze SO2 to SO3. Probably not a practical method to manufacture H2SO4, but I wouldn't rule it out, except poisoning of the catalyst is probably an issue, just as it can be in the contact process for H2SO4 manufacture.

See http://pubs.acs.org/doi/abs/10.1021/i200033a031
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[*] posted on 17-6-2009 at 17:33


I have a book that has a lab scale contact process in it, using platinised something or another. Prohibitively expensive, but just shows that platinum will catalyse the oxidation of SO2.
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[*] posted on 26-8-2009 at 03:13


According to Industrial electrochemistry By Derek Pletcher, Frank Walsh. If SO2 is continuously bubbled into a cell of water with two PbO2 electrodes with a P.D. of 1.4V across each, conc sulpuric can be the resulting product;
H2SO3 + H2O --> H2SO4 + H2
This could be a usefull way to conc H2SO4
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entropy51
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[*] posted on 26-8-2009 at 05:37


Quote: Originally posted by Picric-A  
According to Industrial electrochemistry By Derek Pletcher, Frank Walsh. If SO2 is continuously bubbled into a cell of water with two PbO2 electrodes with a P.D. of 1.4V across each, conc sulpuric can be the resulting product;
H2SO3 + H2O --> H2SO4 + H2
This could be a usefull way to conc H2SO4

Instead of posting a brain fart, why don't you make a whole bunch of concentrated H2SO4 and then tell us about it.
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[*] posted on 26-8-2009 at 06:09


Quote: Originally posted by DJF90  
I have a book that has a lab scale contact process in it, using platinised something or another.

Platinised asbestos was used in those old processes but fairly pure SO2 was required to minimise catalyst-poisoning.
Chloroplatinic acid, reduced, supplied the finely divided Pt.
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[*] posted on 26-8-2009 at 07:10


Yes I know this. Its no hassle to generate SO2 from metabisulfite, which should be fairly pure - send it through an approprate washbottle or two to remove impurities and drying train to remove moisture and it should be pure enough for this application. I believe it was platinised kaowool that they used, although I'll have to double check this - they might even have the catalyst preparation in the experimental procedure.
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[*] posted on 26-8-2009 at 08:55


Platinum catalyst poisoning from the SO2 was problematic when iron pyrites was burned to produce the SO2. It is much less problematic when pure sulfur is burned to supply the SO2. Metabisufite would seem like a good SO2 source for the contact process.

Vanadium pentoxide catalyst is much less subject to poisoning.
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[*] posted on 27-8-2009 at 08:19


Quote: Originally posted by entropy51  
Vanadium pentoxide catalyst is much less subject to poisoning.
From talking to vendors, it seems that the main poison for vanadium oxide catalysts is arsenic, which can be a problem when converting off-gas from smelting sulfides.
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[*] posted on 11-10-2009 at 03:20
Sulphur burner


The contact and the lead chamber processes require lots of SO2 from burning sulphur. Thoughts on a burner container went through steel, lead and ceramic, then thoughts on a wick wandered through paper and steel and stainless steel (pyro sieve mesh!?!) then I wondered about using a relatively fine stainless mesh prepared with V2O5. Would the sulphur flame be hot enough to get the V2O5 up to the region of exothermic catalysis? Could a one shot process be controlled to be safe, stable and efficient at producing SO3 directly?
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[*] posted on 11-10-2009 at 06:30


Quote: Originally posted by Contrabasso  
Could a one shot process be controlled to be safe, stable and efficient at producing SO3 directly?
In industry, at the start of the campaign, the catalyst bed often receives supplementary heat to get the SO2 -> SO3 oxidation going. Since that oxidation is exothermic, once it gets going it's self-heating, to the point that it later requires external cooling, which they do by using it as a source of process steam.

For a small-scale synthesis, the surface-area to volume (square-cube) ratios are all different, and you're going to be in a much different thermodynamic regime. If you do get into the self-heating regime, you're probably making more SO3 than you can use and more hazard than you can handle. At the very least, when prototyping, consider using both external heat for the catalyst and thermocouple to monitor its temperature.
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[*] posted on 11-10-2009 at 16:19


Is there some part of H2SO4 by the Lead Chamber Process that I don't understand? Some people really should raise their standards for posting, in the right thread or anywhere else here. Or do you want to remake TOTSE? I see that the meth syntheses posts are back, so I guess you do.

Another simple JCE illustration of this. They had a couple for the SO3 process as well, back in the day.

Attachment: JCE1930p1668.pdf (1.9MB)
This file has been downloaded 1068 times




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[*] posted on 27-12-2009 at 17:22


How about this:

http://www.youtube.com/watch?v=5dUSF9Gl0xE

Start with dead cheap easy to get copper sulfate, and produce pure sulfuric acid. You just need a platinum coated (or a pure Pt) electrode, and it is very easy.
You could in theory convert 1kg of copper sulfate into about 230mL of concentrated pure sulfuric acid. Ofcourse during boiling down the acid, there may be some losses, as H2SO4 fumes.
Will take a lot of time ofcourse but it should be a promising path to pure acid.
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[*] posted on 27-12-2009 at 18:14


Using copper sulfate does have the advantage that you can do the electrochemistry without a membrane.
In any case, you don't need a platinum electrode; PbO2 will also work. In fact when I did this I just started with lead, which when used as an anode under these conditions (dilute sulfuric acid) acquires a PbO2 coating pretty quickly. It doesn't hold up all that well (tends to shed bits over a period of days) but it's adequate. I used Na2SO4 and MgSO4 in my two runs, though.
I also wouldn't regard it as a 'promising path to pure acid' unless your dead cheap and easy to get CuSO4 also happens to be reagent grade.
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[*] posted on 28-12-2009 at 13:45


Quote: Originally posted by Jor  
How about this:

http://www.youtube.com/watch?v=5dUSF9Gl0xE

Start with dead cheap easy to get copper sulfate, and produce pure sulfuric acid. You just need a platinum coated (or a pure Pt) electrode, and it is very easy.
You could in theory convert 1kg of copper sulfate into about 230mL of concentrated pure sulfuric acid. Ofcourse during boiling down the acid, there may be some losses, as H2SO4 fumes.
Will take a lot of time ofcourse but it should be a promising path to pure acid.



NurdRage is amusing but shouldn't be followed, he's a mediocore chemist at best and will lead you entirely down the wrong way of doing things. He's got little regard for practicality or cost and sometimes he's blatantly wrong. He's a clown.
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[*] posted on 4-1-2010 at 06:57


SC Wack:

I regeards to your post below, could you reload that file? The server is saying the file is "damaged and could not be opened"
------------------------------
Is there some part of H2SO4 by the Lead Chamber Process that I don't understand? Some people really should raise their standards for posting, in the right thread or anywhere else here. Or do you want to remake TOTSE? I see that the meth syntheses posts are back, so I guess you do.

Another simple JCE illustration of this. They had a couple for the SO3 process as well, back in the day.

Attachment: JCE1930p1668.pdf (1.9MB)
This file has been downloaded 88 times
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[*] posted on 4-1-2010 at 14:27


Quote: Originally posted by jgourlay  
SC Wack:

I regeards to your post below, could you reload that file? The server is saying the file is "damaged and could not be opened"


It still works for me.




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[*] posted on 13-6-2010 at 20:53


Quote: Originally posted by Formatik  
... Conc. H2SO4 will not attack iron, but dilute acid will, so if the sulfates are distilled in iron, at best they should be made anhydrous before proceeding to a higher heat. Quartz and Vycor can handle higher heat. ...


Reading over the thread on SO3 from NaHSO4 in prepublication, using most common metal tubes probably won't work, e.g. hot conc. H2SO4 does attack steel forming SO2, and for SO3, as mentioned from Gmelin hot SO3 is reduced by iron forming sulfide.
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[*] posted on 14-7-2010 at 08:02


Hi folks!

Sorry to ask irrelevant stuff, maybe I would have better luck in short question thread.

What could one get from pyrolysis of NOHSO4 (chamber crystals)?

If someone could kindly provide a reference, or a reference of a refrence, I would be very gratefull!

I made some, and wondering what uses it has other than generating N2O3.

BTW I've already used the FSE:)

[Edited on 14-7-2010 by Jimmymajesty]
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[*] posted on 14-7-2010 at 21:53


Heating NOHSO4 forms dinitrosyl sulfate ((NO)2S2O7): 2 NOHSO4 <- -> (NO)2S2O7 + H2O (A. Michaelis, O. Schuman, Ber. 7 [1874] 1077). NOHSO4 solubilized in conc. H2SO4 (or a soln. on NaNO2 in H2SO4) reacts differently with organics under certain conditions, nitrosation, nitration, diazotization or oxidation can occur. Doesn't look to spectacular IMO, or anything that HNO3, mixed acid, HNO2, etc. couldn't pull off.
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[*] posted on 21-7-2010 at 11:38


Formatik thanx for the info!

What about heating the nitrosyl sulphuric acid with sulphur? Can you foresee any spectacular?

Sorry for my brainfarts, but I am not at home at the moment so I cannot make experiements, It is easier to ask the skilled in the art:)
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[*] posted on 15-8-2010 at 00:52


It's been said plenty of places & I'll add it here (H2O + H2SO3 + Cl2 ==> H2SO4 + 2HCl), simply because a LOT of new chemists don't realise that there is, in fact, an easier route.

For those who wish to play around with the contact process, go for your life. For those who want pure Halogen Acids and Sulfuric Acid, use the sensible route.;)

[Edited on 15-8-2010 by un0me2]




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[*] posted on 16-8-2010 at 05:20


un0me2 : how do you get the two separated once they are mixed like that?
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[*] posted on 16-8-2010 at 10:44


HCl is more volatile, so all you have to do is heat it, and the H2SO4 stays behind. This ought to depend on how much H2SO4 there is in solution to begin with, since conc. H2SO4 already drives HCl out of solution sans external heating. There must be some kind of chart somewhere showing at what concentrations HCl and H2SO4 coexist in solution, etc.

In terms of yield, instead of using a H2SO3 solution, it might be better to simultaneously bubble Cl2 and SO2 into water, because if you just dissolve SO2 in H2O you can lose some sulfur, because SO2 is difficultly soluble in water (a bit better when cold). There may be some info on the reaction in Gmelin.

One would have to work in a fume hood, or outside in a safe area with chemical respirators (better is a gas mask, or something that protects eyes from fumes also) since these two are hefty and lethal gases. An excess of ammonia destroys either gas.

And yes, chlorine is the oxidizer much more readily preparable from ubiquitous material.;)
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[*] posted on 16-8-2010 at 13:39


SO2 is soluble as all fuck in water - the clathrate is insoluble till it melts (about 15-20'C IIRC) but collection of the clathrate should give a strong solution of SO2/H2O in about the right proportions. Actually, dissolving SO2 in water is endothermic, so be aware of that. I've seen STRONG solutions of SO2 in water, they stink like hell (yum SO2), but a strong solution is quite workable. Adding more SO2 as the oxidation proceeds should be feasible.

As Formatik said, HCl is a gas - no worries whatsoever there the equilibrium is one-sided, the gaseous reduction product leaves the reaction and the aqueous solution of the oxidized product stays behind. I'd be interested to see what is the maximum strength of H2SO4 that could be made by this process (which, while oxidizing the acid, also dehydrates it).

Dissolve the HCl gas given off in distilled water (and use distilled water for the H2SO3) and all of a sudden you have two pure acids without visiting chemical supply houses. Also works with I2 & Br2.




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