| Pages:
1
2
3
..
12 |
axehandle
Free Radical
Posts: 1065
Registered: 30-12-2003
Location: Sweden
Member Is Offline
Mood: horny
|
|
H2SO4 by the Lead Chamber Process - success
I was going to hold on to this a bit further - until I have more observations to report - but I decided otherwise since the more that try the process,
the more improvements might be invented - and many of you may have suggestions and ideas to optimize the process.
Also, I can't resist the temptation of bragging about finally finding a successful, simple way of making
H<SUB>2</SUB>SO<SUB>4</SUB> with easily obtained tools and chemicals...
The Lead Chamber Process, invented in 1746 by the Brit John Roebuck, was the first feasible industrial manufacturing process for
H<SUB>2</SUB>SO<SUB>4</SUB>, and wasn't replaced with the contact process (using a Pt catalyst to oxidize
SO<SUB>2</SUB> to SO<SUB>3</SUB> until the first half of
the 1900s. Lots of historical information exists online, I won't go into the historical aspect more than strictly necessary.
In the lead chamber process, 7 part sulfur (note 2) is burned together with 1 part sodium or potassium nitrate inside a lead-lined chamber (note 1)
with water covering the floor. SO<SUB>2</SUB> is generated in abundance together with lesser amounts of NO from the saltpetre. Over time
(several hours in my experience) the NO catalyses the oxidation of the gaseous SO<SUB>2</SUB> to the trioxide SO<SUB>3</SUB>;
which combines with the water on the floor, as well as with water vapour (note 2) inside the chamber, to form dilute
H<SUB>2</SUB>SO<SUB>4</SUB>. The reaction mechanism is described by the following, unbalanced reactions:
(a) 3S(s) + 2KNO<SUB>3</SUB>(s) --> K<SUB>2</SUB>S(s) + 2SO<SUB>2</SUB>(g) + 2NO(g) ;sulfur +
KNO<SUB>3</SUB> reaction
(b) S(s) + O<SUB>2</SUB>(g) --> SO<SUB>2</SUB>(g) ;combustion inside the chamber
(c) 3NO(g) + 3/2O<SUB>2</SUB>(g) --> 3NO<SUB>2</SUB>(g) ;spontaneous at NTP
(d) 3NO<SUB>2</SUB>(g) + 3SO<SUB>2</SUB>(g) --> 3SO<SUB>3</SUB>(g) + 3NO ;catalyzed oxidation
(e) SO<SUB>3</SUB>(g) + H<SUB>2</SUB>O(l) --> H<SUB>2</SUB>SO<SUB>4</SUB>(aq) ;absorption
In (a,b) 1 part of the sulfur combined with the nitrate during the combustion, and the other 6 parts combines with oxygen from the air inside the
chamber. (c, d) constitute the catalyzed step.
Now, the "1 part nitre to 7 parts sulfur" ratio is, according to the sources I've read, the one that was commonly used in plants at the
time. I'd guess it was chosen to save on the amount nitrate needed while still making the reaction go fast enough. The nice thing here is that
the amount of nitre used only influences the rate of the reaction, not its completion. Lesser amounts of nitre makes (c, d) proceed more slowly,
larger amounts make them faster. The nitrate (or rather the NO produced when combined with sulfur) acts as catalyst, not as sulfur oxidizer. Note that
the oxygen content of the chamber must be in excess to the sulfur burnt for (c, d) to have oxygen to draw from.
Armed with this knowledge I proceeded with trying out the process on a much smaller scale.
1) The lid to a 25-liter fermentation tank made of PP was fitted with a glass pipe going through it that when the lid was fitted ended about 50mm from
the bottom of the tank.
500ml of deionized water was poured into the otherwise empty tank.
A small glass bowl with KNO<SUB>3</SUB> in the bottom was placed in the bottom of the tank.
Solid pieces of sulfur were prepared by casting purified sublimed sulfur to a puck shape in a can, then crushing the puck.
To initiate a burn, a piece of solid sulfur was placed on the bed of KNO<SUB>3</SUB> and set on fire with a powerful butane lighter. Then
the lid was put on. Once molten, it burns nicely in a puddle on top of the KNO<SUB>3</SUB>. When it gets warm enough after about 30
seconds, the KNO<SUB>3</SUB> under the sulfur / mixed with the sulfur will start decomposing, and a black powder like burn will to
self-perpetuate, going on until all the sulfur is burnt. The whole reason for the pipe for those that can't guess it, is so that the
SO<SUB>2</SUB> + NO rising upward from the burning sludge in the bowl won't start to leak out until the chamber is filled; I wanted
to really fill the chamber.
Now, this deviates from the 1/7 ratio in the original process and seems like a horrible waste of KNO<SUB>3</SUB>, but I have 4 reasons:
1) The chamber size is very small and has limited oxygen. Thus a stochiometric amount of KNO<SUB>3</SUB> will make all the oxygen for (a,
b) come from the KNO<SUB>3</SUB>, not the chamber air -- leaving the oxygen of the latter intact for reactions (c, d).
2) It makes (c, d) go at a faster rate.
3) It's easier, and I'm very lazy.
4) I've got 70kg of the stuff sitting in my wardrobe, and can get more very easily.
After the burn, all that's needed is to wait a few hours. The cloudy gas from the burn will clear up as the SO<SUB>2</SUB> inside the
chamber is oxidized to SO<SUB>3</SUB> and absorbed into the water. If the chamber was hermetically sealed, it would collapse as the gas
volume inside decreases. I've tested the process with a plastic Coke bottle which was sealed, and it imploded slowly over a few hours. After the
waiting period, the lid is removed, and the NO/NO<SUB>2</SUB> allowed to vent. Then the process is repeated.
I had done over 10 burns with my first batch when the glass bowl used as burner cracked from the heat, spilling KNO<SUB>3</SUB> into the
dilute H<SUB>2</SUB>SO<SUB>4</SUB>. Before thinking, and quite pissed off, I flushed all the fluid down the toilet to start
over using a ceramic bowl instead. I should have saved it for analysis -- all I can say is that it fizzled a lot when coming into contact with
whatever salt deposits are always building up in the the toilet bowl. Normally I remove the deposits using HCl -- the HNO<SUB>3</SUB> from
the KNO<SUB>3</SUB> in the H<SUB>2</SUB>SO<SUB>4</SUB> seemed to work just as well... I have a new batch going
and have done 2 burns in it so far.
I'll report the progress.
----
note 1: Lead being the only cheap material resistant to H<SUB>2</SUB>SO<SUB>4</SUB> available at the time.
note 2: Yes, I know I'm mixing American and British spelling, and no, I won't stick with one of them. I prefer parts of both...
My PGP key, Fingerprint 5D96 E09E 365D 1867 2DF5 C2FE 4269 9C19 E079 CD35
\"Verbing nouns weirds the language!\"
|
|
|
frogfot
Hazard to Others
Posts: 212
Registered: 30-11-2002
Location: Sweden
Member Is Offline
Mood: happy
|
|
If I got it right, you didn't mix the two components to keep reaction speed down? Why not make solid S/KNO3 cakes in an iron can? Maby hang it
below the lid to avoid the heat (making sure the can doesn't melt/react..).
Keep up the testing! 
[Edited on 13-11-2004 by frogfot]
|
|
|
Tacho
National Hazard
Posts: 582
Registered: 5-12-2003
Member Is Offline
Mood: No Mood
|
|
H-O-L-Y C-O-W !!!
Congratulations axe! Bigtime congratulations!
A practical setup for making H2SO4 in the kitchen is something remarkable.
Dilute H2SO4 can be concentrated, and with concentrated H2SO4, one can make most (all?) mineral acids!!!
Woo Hoo!
|
|
|
Mr. Wizard
International Hazard
Posts: 1042
Registered: 30-3-2003
Member Is Offline
Mood: No Mood
|
|
Very nice explanation of how it works, thanks for the details. Do you have any pictures? You have brought the most cumbersome industrial process
down to laboratory size! What effect does the dissolving of SO2 in water have on the reaction? Does too much water pull the SO2 out of the container
volume, and pull in fresh air? Does it slow the catalytic oxidation of the SO2 down, or does the NO2 work on the SO2 solution? Would heating the
water keep the SO2 out of the water, without effecting the SO3 absorption? Thank you for sharing with us.
|
|
|
axehandle
Free Radical
Posts: 1065
Registered: 30-12-2003
Location: Sweden
Member Is Offline
Mood: horny
|
|
| Quote: |
If I got it right, you didn't mix the two components to keep reaction speed down? Why not make solid S/KNO3 cakes in an iron can? Maby hang it
below the lid to avoid the heat (making sure the can doesn't melt/react..).
|
Well, I actually did try that, but found that if the S and the KNO<SUB>3</SUB> are mixed when burned, the burning gets too quick, making
much of the sulfur overheat, boil and contaminate the entire reaction vessel (not that that is a disaster, it would be easy enough to filter out
afterwards, but still there's a danger of fuel-air explosion (I assume)).
| Quote: |
Congratulations axe! Bigtime congratulations!
|
Thank you.
| Quote: |
.... Do you have any pictures? You have brought the most cumbersome industrial process down to laboratory size! What effect does the dissolving of SO2
in water have on the reaction? Does too much water pull the SO2 out of the container volume, and pull in fresh air? Does it slow the catalytic
oxidation of the SO2 down, or does the NO2 work on the SO2 solution? Would heating the water keep the SO2 out of the water, without effecting the SO3
absorption? Thank you for sharing with us.
|
No pictures. Is it needed? It's a big plastic bucket with a lid, a hole in the lid with a glass pipe shoved through, a food bowl with
KNO<SUB>3</SUB>, a bit of water at the bottom... aaah, I suppose I could take a picture later...
About dissolving SO<SUB>2</SUB>: Not sure there. Have only tried it with very small amounts of water. Wouldn't the
SO<SUB>2</SUB> so dissolved carry over to the next burn, making the water already saturated provided the burns take place at intervals
close enough to make the autodecomposition of SO<SUB>2</SUB>(aq) insignificant? I too would love to know if the catalysis takes place in
aqueous soln. as well. Self, I suspect the formation of nitrosylsulfuric acid HNOSO<SUB>4</SUB>.
|
|
|
axehandle
Free Radical
Posts: 1065
Registered: 30-12-2003
Location: Sweden
Member Is Offline
Mood: horny
|
|
Another observation
When doing the burns my style, i.e. setting fire to a blob of molten sulfur on a bed of KNO<SUB>3</SUB> powder, some of the sulfur will
vapourize during the burn, once the fire gets hot enough to make the KNO<SUB>3</SUB> join into the reaction... It seems to condense on the
inside walls of the chamber and in the water/dilute acid at the bottom. I can't see this being a problem though -- sulfur shouldn't react in
any way with H<SUB>2</SUB>SO<SUB>4</SUB>, right? And it's easy to decant off the liquid, thus separating out the
sulfur...
|
|
|
neutrino
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
Are you sure that NOx is formed? It should react with the water and oxygen, forming nitric acid, which would then oxidize the sulfur into sulfuric
acid. It seems odd that there would be sulfur left over.
|
|
|
darkflame89
Hazard to Others
Posts: 255
Registered: 1-3-2004
Location: With probability 1, "somewhere" in this
Member Is Offline
Mood: No Mood
|
|
Congratulations, axehandle. Anyway I found this, similar to what axehandle is doing(in fact the same):
"Pop bottle Process"
Ignis ubique latet, naturam amplectitur omnem.
|
|
|
Esplosivo
Hazard to Others
Posts: 491
Registered: 7-2-2004
Location: Mediterranean
Member Is Offline
Mood: Quantized
|
|
Congratulations Axehandle!
Congratulations axe! I was also thinking about the production of H2SO4 but the setup of the production is different, although its still an idea which
will need improvements. I would just like to ask, could the container for burning the Sulfur/Nitrate mix be made of glass? If not what are other
possible substitutes? I don't like lead a lot, it has the tendency of melting at 'low' temperatures which I quite dislike.
I would also like to ask, is the KNO3 really necessary? I didn't know that a reaction could occur where a sulfide forms on reaction with a
nitrate, giving off NOx fumes. I'm not doubting what you said axe, but it seems strange to me. Can't excess air oxidise most, if not all, of
the sulfuric (IV) acid (sulfurous acid)? Thanks.
Theory guides, experiment decides.
|
|
|
axehandle
Free Radical
Posts: 1065
Registered: 30-12-2003
Location: Sweden
Member Is Offline
Mood: horny
|
|
| Quote: |
Are you sure that NOx is formed?
|
I'd recognize that horrible smell anywhere -- yes, I'm very sure.
| Quote: |
Congratulations axe! I was also thinking about the production of H2SO4 but the setup of the production is different, although its still an idea which
will need improvements. I would just like to ask, could the container for burning the Sulfur/Nitrate mix be made of glass? If not what are other
possible substitutes? I don't like lead a lot,
|
The first containers used after the process was invented were actually made of glass, but glass imposed limits on how big they could be constructed so
it wasn't until lead saw use that the process became industrially feasible. I can't see any reason why any material resistant to sulfuric
acid couldn't be used (using PP plastic myself...).
| Quote: |
it has the tendency of melting at 'low' temperatures which I quite dislike.
I would also like to ask, is the KNO3 really necessary? I didn't know that a reaction could occur where a sulfide forms on reaction with a
nitrate, giving off NOx fumes. I'm not doubting what you said axe, but it seems strange to me. Can't excess air oxidise most, if not all, of
the sulfuric (IV) acid (sulfurous acid)? Thanks.
|
I suppose one could get away with leaving out the nitre, having all the H<SUB>2</SUB>SO<SUB>4</SUB> come from
autodecomposition of H<SUB>2</SUB>SO<SUB>3</SUB>, but I think the process would be horribly inefficient -- days instead of
hours, and very dilute acid. BTW, the reaction
S + KNO<SUB>3</SUB> --> SO<SUB>2</SUB> + KNO
also takes place, AFAIK. I don't know the extent of the
3S + 2KNO<SUB>3</SUB> --> K<SUB>2</SUB>S + SO<SUB>2</SUB> +2NO
reaction, only that it <b>does</b> take place. The "cloudyness" clearing up inside the chamber over a couple of hours combined
with the fact that the fluid does <i>not</i> smell of SO<SUB>2</SUB> upon opening the container is proof.
Oh, and most nitrates would probably work. In the old days, NaNO<SUB>3</SUB> was used. Perhaps even
NH<SUB>4</SUB>NO<SUB>3</SUB> could substitute .
[Edited on 2004-11-14 by axehandle]
|
|
|
TheBear
Hazard to Self
Posts: 78
Registered: 17-10-2002
Location: Sweden
Member Is Offline
Mood: distilled
|
|
rotten eggs... you mean H2S. Probably just a typo.
But, doesn't H2SO4 react with elemental sulfur, producing SO2? I recall reading that in some related thread.
EDIT: Just realiazed that it won't matter.
Anyways, great work!
[Edited on 14-11-2004 by TheBear]
|
|
|
axehandle
Free Radical
Posts: 1065
Registered: 30-12-2003
Location: Sweden
Member Is Offline
Mood: horny
|
|
Was more of a "thinko" than a typo --- didn't mean rotten eggs at all. Edited.
My PGP key, Fingerprint 5D96 E09E 365D 1867 2DF5 C2FE 4269 9C19 E079 CD35
\"Verbing nouns weirds the language!\"
|
|
|
Esplosivo
Hazard to Others
Posts: 491
Registered: 7-2-2004
Location: Mediterranean
Member Is Offline
Mood: Quantized
|
|
| Quote: |
I don't know the extent of the
3S + 2KNO3 --> K2S + SO2 +2NO
reaction, only that it does take place. The "cloudyness" clearing up inside the chamber over a couple of hours combined with the fact that
the fluid does not smell of rotten eggs (SO2...) upon opening the container is proof.
|
I suppose that the NO dissolves in water producing a dilute HNO<sub>2</sub> solution which readily oxidizes the
H<sub>2</sub>SO<sub>3</sub>. But I don't know if such a reaction does occur. NO sometimes acts as a reducing agent. One
would need the reduction potentials of both half reactions to work out if such a rxn could occur. Does anyone have any idea if the reaction mentioned
occurs?
Theory guides, experiment decides.
|
|
|
Eclectic
National Hazard
Posts: 899
Registered: 14-11-2004
Member Is Offline
Mood: Obsessive
|
|
Hey Axehandle
Did you ever get your Vanadium catalyst H2So4 reactor working? I have an old book on H2SO4 manufacture that lists a few catalyst formulas.
Most of them are about 30% sulfur and some cesium salt as a promoter. The reaction apparently happens in a liquid phase of CsVSO4 phase on the
carrier.
I saw your project pictures while trying to connect to your FTP. (I need the Handbook of Preparative Inorganic Chemistry for the needy folks on
emule)
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
It seems to me that it may be possible to produce oleum by the chamber method
if the chamber material is compatable with the fuming sulfuric acid produced . It also seems possible that the process could be made to run
continuously by a forced flow of the the sulfur combustion gases into the chamber . A paddle stirrer could be used to maintain a cyclonic flow of the
gases in something like a large glass water carboy , which would centrifuge the condensing fog of acid droplets against the walls and aid the speed of
the reaction and separation . The sulfur could be preburned as a sulfur lamp in a separate chamber , where the sulfur is melted in a pot and burned on
a glass fabric wick , the combustion air supplied to the sulfur lamp by a small pump like an aquarium pump .
If a small chamber equipped with a spark plug was placed in the air line to the sulfur lamp , and a continuous arc maintained in the airflow , there
may be a sufficient catalytic amount of nitrogen oxides produced to maintain the process . This would have to tested and it may require several spark
plugs driven from a sign transformer to produce the needed amount of nitric oxides . By regulating the moisture content of the incoming air , the
process can be controlled ,
all theoretically of course 
By such a scheme , the only precursors required for oleum would be sulfur , air ,
and electricity .
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
This sounds very good. Great work, axehandle!
Rosco, you gave me the idea of using an electric arc to make the NO2 rather than using a nitrate. Now the required chemicals would really only be
water, air and sulfur.
My MOT produces enormous amounts of NO2 when an arc is drawn (2,4 kV @ 0,7A- alot of power, I know), when an arc is drawn for only 10 seconds, the
entire room very noticeably smells of NO2.
Even an NST can fill a glass jar with red NO2 when an arc is set up in it. This will be the key to the oxidation of SO2.
In the industrial lead chamber process, the NO2 is recycled- this won't be necessary in the homemade apparatus.
We only need a reliable sulfur burner where compressed air can be pumped in and SO2 + air can be taken out under slight pressure.
It might be a good idea to try to liquefy the SO2 and store it as a liquid under pressure.
The boiling point is -10°C, this is definately achievable with cooling mixtures or even the freezer. It could be stored in thick-walled PET
bottles, these things can stand a lot of pressure.
Then the SO2 can be mixed with the right amount of air (weigh the SO2 bottle to know how much has been used) and NO2 and left to react in a large
plastic drum (300 litres) with a small amount of water at the bottom.
I will definately try to liquefy SO2 someday, it is one of the few gases that can be liquefied very easily.
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
There probably is no need for the added step of quantifying the SO2 or liquifying it .
So long as an excess of air is being supplied , the reaction should proceed ,
rate limited by the wick area and flame size of the sulfur lamp which will operate
at maximum SO2 output so long as it is receiving sufficient air , plus the extra amount of air required for the further oxidation of the SO2 to SO3
which will
occur in the chamber .
One way of making a sulfur lamp would be
something like a four liter resin kettle kept one quarter full of molten sulfur with
a floating wick carrier . Once the sulfur is
molten and the wick is wetted with molten sulfur , the wick is ignited , and
the cover with fresh air inlet and exhaust
ports is put in place . The correct wick size
should result in sufficient output while keeping the sulfur molten by it's own heat . Solid sulfur could be added through the cover , to
replenish the sulfur
as it is burned away . The process could
be automated to run continuously .
Alternately to burning sulfur , hydrogen sulfide gas could be generated and fed to a burner , which would produce the exact amount of water required
to form ordinary sulfuric acid . But the removal
of some of the water vapor by running the gases through a condenser and trap would be necessary if oleum was the desired end product . IIRC simply
heating gently a mixture of paraffin and sulfur can be used as a simple method of producing
hydrogen sulfide . Depending upon what
hardware may be available either of these
methods of producing SO2 should work ,
but the direct burning of the elemental sulfur would be more economical .
An alternate method for hydrogen sulfide could be converting the sulfur to sodium or calcium polysulfide and slowly infusing it
with hydrochloric acid . If sodium nitrite is added to the polysulfide solution , the nitrous fumes produced concurrently may
be an alternative to other sources for the
catalytic amount of nitrogen oxides .
Heating sodium bisulfate and sodium nitrate together causes a reaction which produces nitric acid but because of the high temperature the nitric
largely decomposes to nitrogen oxides , and this could be another source for separate generation of the catalyst .
[Edited on 15-11-2004 by Rosco Bodine]
|
|
|
Theoretic
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
axehandle, KNO3 => KNO reduction probably does take place, however, first KNO3 => KNO2 then KNO2 => KNO, the NO- anion very quickly pairing
to form the hyponitrite N2O2--, this then disproportionates like this:
3N2O2-- => 2N2 + 2NO2- + 2O--. The oxide could then gobble up a considerable amount of your SO3 (and nitrogen oxides), and its formation could be
prevented by using enough sulfur (which means KNO doesn't hang around long enough to decompose). As a sidenote, reduction of a nitrite by a
sulfide (at high heat) can be a useful source of K2O and Na2O:
8NaNO2 + 3Na2S => 4N2 + 4Na2O + 3Na2SO4.
|
|
|
Rosco Bodine
Banned
Posts: 6370
Registered: 29-9-2004
Member Is Offline
Mood: analytical
|
|
factors to be considered
# 1 on the list is nitrosylsulfuric acid ,
which can become a huge impurity dissolved in the sulfuric acid .
I found a patent which gives some good insight into the process of practical conversion of SO2 to H2SO4 as it is
done on a commercial scale .
The patent is a good read for gaining insight into the reaction conditions and
what sort of sequence is required for
efficiently performing the manufacture of H2SO4 using Nitrogen oxides and SO2 .
Attachment: US3649188 Sulfuric Acid Manufacture.pdf (665kB)
This file has been downloaded 3279 times
|
|
|
Nerro
National Hazard
Posts: 596
Registered: 29-9-2004
Location: Netherlands
Member Is Offline
Mood: Whatever...
|
|
@Rosco
The nitrosylsulfuric acid will be present but since most members here will not use this H2SO4 for anything else than to slurp up water from reactions
I don't think it will be too much of a nuisance. If they need PA H2SO4 they should just buy it...
I just read a moment ago that in the old H2SO4 factories a rust-coloured slurry remained in the lead chambers. This slurry contained up to 15% Se! So
if any member intends to build and use a lead chamber and then he might want to consider using cheap impure S afterall. The Se might be usefull or at
least slightly profitable. If I can get my hands on the right tools and a huge chunk of lead in the near future I will definately build a lead
chamber. Even if the H2SO4 won't be too pure it will still be a nice cheap source considering the rediculous prices I pay now. (€15 for 1L)
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
|
|
|
Taaie-Neuskoek
Hazard to Others
Posts: 222
Registered: 14-5-2004
Location: Zermany
Member Is Offline
Mood: Botanical!
|
|
nitrosylsulfuric is a serious contaminant as it can diazotize benzene-ring like structures, which can be unwanted reactions.
Nice project btw, I planned to make an SO3 reactor myself, I'm not very interested in making H2SO4, but more in oleum...
I though first of using the microwave method Axehandle abandoned because he didn't had a microwave, but this shines another light on the matter.
How hot does the mixure burn, would it be possible to do this reaction in an RBF, with a long column and a condensor connected to it...??
[Edited on 27-5-2005 by Taaie-Neuskoek]
Never argue with idiots, they drag you down their level and beat you with experience.
|
|
|
Jome
Hazard to Others
Posts: 154
Registered: 10-6-2004
Location: Soutwest sweden
Member Is Offline
Mood: desiccated
|
|
How stable is nitrosylsulfuric acid? If acid is boiled down, maby it breaks down. Or if H2O2 is added its wrecked by formation of HNO3 or something?
|
|
|
Phel
Hazard to Self
Posts: 72
Registered: 20-11-2004
Location: Scandium
Member Is Offline
Mood: viscous creep
|
|
| Quote: |
If acid is boiled down, maby it breaks down.
|
According to Russian Journal of Applied Chemistry Issue Vol.74 Issue.1 Page no. 167-169
Nitrosylsulfuric acid is hydrolysed when temperature and concentration is increased.
Journal attached:
|
|
|
Phel
Hazard to Self
Posts: 72
Registered: 20-11-2004
Location: Scandium
Member Is Offline
Mood: viscous creep
|
|
Damnit, forgot attachment, sorry about the double post.
Attachment: NSA.pdf (31kB)
This file has been downloaded 3051 times
|
|
|
Jome
Hazard to Others
Posts: 154
Registered: 10-6-2004
Location: Soutwest sweden
Member Is Offline
Mood: desiccated
|
|
The compound is turned back into (which?) gas and sulfuric acid when hydrolysed?
I've got an idea about how these "lead chambers" could be constructed in a slightly different way. One could let the
KNO<sub>3</sub> and S burn in a separate burn-chamber connected to the main tank by a hose. That way there'd be no chance of
dropping ashes or melt into the water/sulfuric acid in the bottom. The fuel could be changed easier too.
I guess a metallic or ceramic container is needed (or bs glass?) for this "burner", my plastic (soda 50cl) bottle got burned to hell when I
tried this today.
|
|
|
| Pages:
1
2
3
..
12 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
