| Pages:
1
..
6
7
8
9
10
..
12 |
1281371269
Hazard to Others
Posts: 312
Registered: 15-5-2009
Member Is Offline
|
|
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.
|
|
|
entropy51
Gone, but not forgotten
Posts: 1612
Registered: 30-5-2009
Member Is Offline
Mood: Fissile
|
|
Indeed you may be wrong.  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
|
|
|
DJF90
International Hazard
Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
Picric-A
National Hazard
Posts: 796
Registered: 1-5-2008
Location: England
Member Is Offline
Mood: Fuming
|
|
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
|
|
|
entropy51
Gone, but not forgotten
Posts: 1612
Registered: 30-5-2009
Member Is Offline
Mood: Fissile
|
|
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.
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
| 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.
|
|
|
DJF90
International Hazard
Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
entropy51
Gone, but not forgotten
Posts: 1612
Registered: 30-5-2009
Member Is Offline
Mood: Fissile
|
|
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.
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
Contrabasso
Hazard to Others
Posts: 277
Registered: 2-4-2008
Member Is Offline
Mood: No Mood
|
|
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?
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
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 1283 times
|
|
|
Jor
National Hazard
Posts: 950
Registered: 21-11-2007
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
bbartlog
International Hazard
Posts: 1139
Registered: 27-8-2009
Location: Unmoored in time
Member Is Offline
Mood: No Mood
|
|
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.
|
|
|
Alexein
Harmless
Posts: 35
Registered: 14-7-2007
Member Is Offline
Mood: Metastable
|
|
| 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.
|
|
|
jgourlay
Hazard to Others
Posts: 249
Registered: 9-7-2008
Member Is Offline
Mood: No Mood
|
|
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
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
| 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.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
| 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.
|
|
|
Jimmymajesty
Hazard to Others
Posts: 153
Registered: 9-7-2009
Member Is Offline
Mood: No Mood
|
|
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]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
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.
|
|
|
Jimmymajesty
Hazard to Others
Posts: 153
Registered: 9-7-2009
Member Is Offline
Mood: No Mood
|
|
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
|
|
|
un0me2
aliced25 sock puppet
Posts: 205
Registered: 3-2-2010
Member Is Offline
Mood: No Mood
|
|
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]
quam temere in nosmet legem sancimus iniquam
|
|
|
jgourlay
Hazard to Others
Posts: 249
Registered: 9-7-2008
Member Is Offline
Mood: No Mood
|
|
un0me2 : how do you get the two separated once they are mixed like that?
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
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.
|
|
|
un0me2
aliced25 sock puppet
Posts: 205
Registered: 3-2-2010
Member Is Offline
Mood: No Mood
|
|
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.
quam temere in nosmet legem sancimus iniquam
|
|
|
| Pages:
1
..
6
7
8
9
10
..
12 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
