Sciencemadness Discussion Board - SO2 Production

SO2 Production

Waffles SS - 13-1-2011 at 03:39

This is possible to produce SO2 by this plan?

gsd - 13-1-2011 at 09:33

Yes why not?

BTW you will have to deal with sublimed Sulphur - lots and lots of it - in this scheme.

gsd

Waffles SS - 13-1-2011 at 11:40

I want to use this plan for making Hydroxylamine.HCl
In this plan i use air( NOT Pure oxygen) and i should control flow rate of air by air pump(high rate may lead to SO3 and low rate cause flame off)
Air pump will be aquarium pump(it is simple

)

Stainless steel 316L is resistance to SO2 but i think i should use better alloy
http://www.engineeringtoolbox.com/metal-corrosion-resistance...

Yes i have lot of sulfur in this plan(this is my mistake in drawing

)
There is better plan for produce SO2?

garage chemist - 13-1-2011 at 12:18

The burning sulfur will constantly tend to go out in this scheme unless the air stream is quite fast and the liquid sulfur is absorbed onto a porous refractory substance to provide a large surface area for burning.
I have observed that pieces of kaowool (ceramic fibre) impregnated with sulfur and set alight are almost impossible to extinguish by blowing out, and will burn (the blue flame is only barely visible in daylight, so watch out) for quite a long time. This observation could come in useful here.
Another problem is that the SO2 will always be mixed with an excess of air when it is generated by burning sulfur. This is a big problem when it is to be used for aqueous chemistry- in the industrial manufacture of Na and K metabisulfite, the oxygen oxidises almost half of the formed bisulfite to sulfate. The salts have to be separated by fractional crystallisation.
This process is only economically viable due to the cheapness of sulfur, and when a market for the sulfate byproduct exists.

For generating SO2 at home you should buy some K or Na metabisulfite and add HCl to that in a gas generator. The SO2 so obtained is pure and only has to be dried when it is required anhydrous.
The reaction of aqueous HCl with metabisulfite is very endothermic and the gas generating flask gets very cold (far below 0°C)- so cold actually that a large amount of the SO2 stays dissolved in the aqueous chloride solution. The flask should be immersed in a warm water bath and either shaken manually or magnetically stirred to liberate all of the SO2.

[Edited on 13-1-2011 by garage chemist]

Waffles SS - 13-1-2011 at 14:06

Thanks garage chemist
I will have problem If i use smaller amount of sulfur in this plan ?
The reaction of aqueous HCl with metabisulfite is very endothermic???!!!

Quote:

H2SO4 (aq) + Na2S2O5 (aq) → 2 SO2 (g) + Na2SO4 (s) + H2O (l)

This is an exothermic reaction.

http://en.wikipedia.org/wiki/Sulfur_dioxide

Wiki wrote sulfuric acid not hydrogen chloride

garage chemist - 13-1-2011 at 14:12

If you drip conc. sulfuric acid into a slurry of bisulfite in water it may be exothermic enough not to get the flask covered with frost like with HCl.
But I would not trust that information from wikipedia unless you have verified it yourself.
I use HCl rather than H2SO4 because it is cheaper and easier to get. It's not that much of a problem to provide a means of heating for the gas generator flask.

ScienceSquirrel - 13-1-2011 at 14:18

It does matter which acid you use. The reaction with sulphuric acid will be exothermic as dilution with the water formed in the reaction will generate a lot of heat.
If you use concentrated sulphuric acid, quite dry sulphur dioxide will pour off.

Eclectic - 13-1-2011 at 15:34

I was thinking of the same basic arrangement as at the top of the post as primary sulfur evaporator using a small jet of pure oxygen, then taking the output gasses into a burner in pure oxygen. Might work as a source for SO3 production using V2O5 catalyst. It effectively would be a flame of S, SO2 burning in an atmosphere of SO2, O2.

Waffles SS - 14-1-2011 at 00:16

Quote: Originally posted by garage chemist

Another problem is that the SO2 will always be mixed with an excess of air when it is generated by burning sulfur. This is a big problem when it is to be used for aqueous chemistry- in the industrial manufacture of Na and K metabisulfite, the oxygen oxidises almost half of the formed bisulfite to sulfate. The salts have to be separated by fractional crystallisation.
[Edited on 13-1-2011 by garage chemist]

I belive amount of oxygen (mixed with sulfur dioxide) will decrease if i control the flow rate of air by air pump,but this is difficult to remove all of it.
Also SO2 can be separate if i do fraction distil at -10c.