Sciencemadness Discussion Board

H2SO4 by the Lead Chamber Process - success

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axehandle - 12-11-2004 at 13:01

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...

frogfot - 13-11-2004 at 01:35

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]

H-O-L-Y C-O-W !!!

Tacho - 13-11-2004 at 03:12

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 - 13-11-2004 at 07:02

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 - 13-11-2004 at 09:55

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>.

Another observation

axehandle - 13-11-2004 at 19:00

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 - 13-11-2004 at 20:14

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 - 14-11-2004 at 00:31

Congratulations, axehandle. Anyway I found this, similar to what axehandle is doing(in fact the same):

"Pop bottle Process"

Congratulations Axehandle!

Esplosivo - 14-11-2004 at 02:44

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.

axehandle - 14-11-2004 at 06:01

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 - 14-11-2004 at 06:15

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 - 14-11-2004 at 06:27

Was more of a "thinko" than a typo --- didn't mean rotten eggs at all. Edited.

:)

Esplosivo - 14-11-2004 at 06:40

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?

Hey Axehandle

Eclectic - 14-11-2004 at 12:29

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 - 15-11-2004 at 07:49

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 :D

By such a scheme , the only precursors required for oleum would be sulfur , air ,
and electricity .

garage chemist - 15-11-2004 at 09:04

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 - 15-11-2004 at 09:24

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 - 27-4-2005 at 09:28

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.

factors to be considered

Rosco Bodine - 27-4-2005 at 11:00

# 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 2812 times


Nerro - 25-5-2005 at 04:08

@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)

Taaie-Neuskoek - 27-5-2005 at 12:14

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]

Jome - 27-5-2005 at 14:32

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 - 27-5-2005 at 15:12

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 - 27-5-2005 at 15:13

Damnit, forgot attachment, sorry about the double post.

Attachment: NSA.pdf (31kB)
This file has been downloaded 2613 times


Jome - 31-5-2005 at 11:19

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.

Pyridinium - 31-5-2005 at 19:51

Quote:
Originally posted by Jome
The compound is turned back into (which?) gas and sulfuric acid when hydrolysed?


HSO4NO (or if you prefer, ONOSO3H) forms NO, NO2, H2SO4, and probably SO2/SO3 when decomposed.

Merck states the decomposition temp. is 73.5 C but curiously, it also says the NOx will form above 50 C (so which is it?)

Water is said to accelerate the decomp., so even if H2SO4 stabilizes it, you could dil. the contaminated H2SO4 with an excess of water, then boil off the water. The NOx / HNO3 formed would volatilize, leaving the H2SO4.

Another idea, since ONOSO3H can crystallize... maybe cool the H2SO4 down greatly so it crystallizes out, then decant the liquid? I don't know what temp you'd need.

Sorry if this already was mentioned in the thread. It's getting late and I'm tired.

haydz - 22-9-2005 at 13:47

This sounds very interesting, has anyone had any good amount of H2SO4 come from it? Has anyone tried concentrating it?

neo_90 - 28-6-2006 at 13:40

I know that this thread has been dead for over 9 months.. but I have a questions..

the H2SO4 produced this way, can it become pure enough to make HNO3?
and/or will it be pure enough to work as a catalyst when producing nice smelling esters?

12AX7 - 28-6-2006 at 19:08

I don't see why not. But you may have trouble getting it dry enough to be valuable as a dehydration agent.

Tim

neo_90 - 29-6-2006 at 02:17

when you say dry enough, do you mean a higher %?
cant I just boil the water out?

and, is there a way to finde out what precentage the acid is?
to make HNO3 I've read that it has to be 96%..

enhzflep - 29-6-2006 at 02:35

Yeah, you can _just_ boil the water out. It's just that it takes a lot of energy. The process now used adds SO3 to water/acid mix and can make anhydrous acid without additional heating (and the associated product loss)

To check out the strength of your acid, reffer to this thread.
http://www.sciencemadness.org/talk/viewthread.php?tid=5817#p...

Nitric may be made with sulphuric of less than 96% concentration. The only thing is that the distillation will need to be run at a higher temp, will produce acid with some water in it (hence higher temp needed), in addition to these two points, you will have a lower yield as a result of the greater amount of nitric being decomposed at the higher distillation temp.

That said, I've successfully made nitric from boiled car battery acid + nitrate salts. Though it still wasn't as potent as that produced by the distillation of comercial 70% with comercial 98% sulphuric.

ps - don't wear too much cotton (t-shirts, jeans etc) :P

tupence_hapeny - 23-4-2007 at 19:12

Ummm,

I finally found full-text access to the journal article regarding the oxidation of sulfurous acid to sulfuric acid via freezing the sulfurous acid to -10C (x3 freeze-thaw cycles), which apparently converts sulfurous acid to sulfuric in 100% yield (NB best would be ~40% as this is the saturation point of sulfurous acid in water).

Unfortunately, I neither read nor write in Japanese (and the article is in Japanese), full text is available here:

http://www.jstage.jst.go.jp/article/nikkashi/2001/2/2001_125...

Now what I would like to know, is:

(1) The procedure - namely, is the dissolved oxygen just the oxygen that is already in the solution or is H2O2 (or other agent) used?

(2) Whether the 40% H2SO4/H20 solution can be added to with more sulfurous acid - which is then oxidized to more H2SO4 - and if so, whether another round could be done, converting 80% H2SO4 & 40% H2SO3 to 97% H2SO4 & 23% SO3 via the same route?

Ideally I would like someone to post a translation of this article, I know it is a big ask... However - it would be seriously important - veritably changing forever the ability of amateur chemists to access a range of reactions cheaply and easily. Even if it cannot be further concentrated by that method, 40% H2SO4 would be easily prepared - boil it down to concentrate it - then freeze the SO3 (~10C) and filter (I would suggest a porcelain frit funnel). Add 40% H2SO4 solution to the frozen SO3 and all of a sudden have 80-120% (1:1 or 1:2) H2SO4.

This could alter amateur chemistry for ever...

NB A similar procedure is also apparently possible for the oxidation of nitrous acid to nitric

[Edited on 24-4-2007 by tupence_hapeny]

Attachment: Honda, 'Acelleration of Oxidation of Sulfurous Acid by Freezing' (2001) 2 Chem Soc Japan 125.pdf (307kB)
This file has been downloaded 2422 times


Aqua_Fortis_100% - 20-6-2007 at 06:19

sorry by up again this thread, but a few things still disturb me..

Quote:
originally posted by axehandle :

(a) 3S(s) + 2KNO3(s) --> K2S(s) + 2SO2(g) + 2NO(g) ;sulfur + KNO3 reaction
(b) S(s) + O2(g) --> SO2(g) ;combustion inside the chamber
(c) 3NO(g) + 3/2O2(g) --> 3NO2(g) ;spontaneous at NTP
(d) 3NO2(g) + 3SO2(g) --> 3SO3(g) + 3NO ;catalyzed oxidation
(e) SO3(g) + H2O(l) --> H2SO4(aq) ;absorption




and ,from the way which axe have made, what happen to K2S? my worry is which some of it can be oxidised to more SO2 and K2O by the saltpeter and some of this can fall into H2SO4, impurifying it... any ideas?


EDIT: tupence_hapeny , great document.. has anyone tried this?

Bromic: i've seen another great document which you have posted at another thread from catalytic oxidation of SO2(aq) in presence of some MnSO4 ...

what about build a *somewhat* different and more expensive lead chamber?.. replacing the KNO3 by KMnO4 (i'm sure which at least some MnSO4 are produced) and fire it inside a chamber with temperature and pressure control (to insure full absorpition of the SO2 generated)...

unfortunatelly by this line of thought , the sulfuric acid whcih can be produced will be very impure , expensive and diluted... but is an idea :D

[Edited on 20-6-2007 by Aqua_Fortis_100%]

[Edited on 20-6-2007 by Aqua_Fortis_100%]

497 - 3-11-2007 at 18:29

i found a page with a ton of good info on H2SO4 production:

http://www.sulphuric-acid.com/TechManual/LeadChamber/Lead_Ch...

497 - 3-11-2007 at 21:05

i found this reaction interesting:

2 HNO3 + 3 SO2 + 2 H2O ---> 2 NO + 3 H2SO4

or

4 HNO2 + 2 SO2 ---> 2 H2SO4 + 4 NO

it works according to United States Patent 4155989. seems like it would work awful nicely for making so homemade H2SO4. and i also like because you don't need the big lead (or plastic) chamber for the gases to react. i'll have to look into this some more.

[Edited on 3-11-2007 by 497]

497 - 3-11-2007 at 22:47

so this is my idea for H2SO4 production. seems doable, no 500C temps, no big chambers, no KNO3, no sulfates. sorry about the crude picture... can anyone find something wrong with it?

[Edited on 3-11-2007 by 497]

[Edited on 3-11-2007 by 497]

Attachment: sulfuric acid.ppt (40kB)
This file has been downloaded 1955 times


12AX7 - 4-11-2007 at 01:00

Either you're adding HNO3, or you'll run out of oxidation. You'll need an air bubbler in the NOx column at least.

What's wrong with bubbling air through H2SO3? Doesn't work?

Tim

497 - 4-11-2007 at 11:50

as far as i've seen H2SO3 doesn't convert on its own. i think if it did we'd all be making our own H2SO4 :P

and why would i have to add more HNO3? the reaction works just fine with HNO2 also. in fact that may end up doing the majority of the work. thanks for the input.

i have a couple of problems with it though. i am skeptical its going to get much gas exchange with simple bubbling, so i was thinking i would use an aquarium airstone that produces very fine bubbles. but the problem is i'm not sure where i can get one that will hold up in concentrated sulfuric and nitric... or maybe theres another way to get good exchange...

also i'm not sure what to use as containers. something fairly tall would definitely be better but is has to have a sealed lid, so glass is not an option. my first thought was PVC but i'm not sure how well that would hold up in those acids. if it can handle them i might even get clear PVC. if all else fails i could use some SS 316 pipe, i can weld it too... but in very big diameters that shit is expensive!

well i checked ebay... 5 ft of 2 1/2 inch diameter SS 304 going for $60 plus $20 shipping... not too bad i suppose. 304 should work too shouldn't it?

[Edited on 4-11-2007 by 497]

Armistice19 - 4-11-2007 at 19:26

Correct me if I'm wrong, but wouldn't bubbling pure oxygen gas through H2SO3 oxidize the acid into the desired H2SO4? If so, Home Depot sells Brazing kits for 50$. This includes an oxygen, and mapp gas adapter, that mixes the gases into one general outlet via the seperate tubes. Obtaining pure oxygen is simply a matter of seperating the tubes.

Just a thought,

Armistice.

497 - 4-11-2007 at 19:51

no i don't know for sure that straight O2 won't oxidize H2SO3. but it doesnt quite make sence since this entire thread is based on producing H2SO4 and nobody ever mentioned bubbling O2 through it. well actually tupence_hapeny did talk about in the above post. the fact that honda is freezing and thawing it 3 times to oxidize it kinda gives me the idea it quite that easy.

12AX7 - 4-11-2007 at 20:51

Ok, I'll run you through your equations:

Burner: S + O2 = SO2(g).

The SO2 is pumped into a water tank, along with residual O2 and inert gasses N2, CO2, Ar, etc. diluting it.

The SO2 dissolves, and reacts with HNOx from the following step:
SO2(aq) + 2HNO3 = H2SO4 + 2NO2(aq/g) (or hydrogen as units of H2O)
H2SO3(aq) + NO2(aq) = H2SO4 + NO(g)
SO2(aq) + 2HNO2 = H2SO4 + 2NO(g)

The NO gas is passed off into water, where it presumably forms HNO and HNO2. No HNO3 is formed, because no NO2 is present. (If there is, as a result of circulation it will soon be reduced to NO.)

If NO also has enough oxidizing power to produce sulfuric acid, then you will continually lose N2 gas in the process.

If oxygen is admitted either to the nitrate solution or to the NO gas, you might be able to sustain something here.

Note that the nitrate bath is going to be very rich in water, and you need to add it to the sulfate bath in proportion to the amount of oxidation required at any given moment in time. Your "sulfuric acid" will very quickly become little more than mere acid rain. One solution would be to aerate the NO gas and bubble NO2 back into the solution, but this is silly, because to add air, you must remove the inert gasses! You would need pure oxygen to maintain such a catalytic cycle.

Armistice: are you referring to those 1 pound propane bottle sized oxygen cylinders? The ones that give you eight, count them eight minutes of burn time? And cost about a dollar per minute of use? Ouch.

Tim

497 - 4-11-2007 at 21:31

why couldn't you supplement in some air in the SO2 feed? wouldn't the O2 pass through the sulfur column and oxidise the NO on its way to the nitrate column? from what i've read NO is rather quickly oxidized. and i was thinking the water input would be slow, hopefully resulting in fairly concentrated acid.

revised version

497 - 7-11-2007 at 11:59

so i decided a batch process would be much more effective. coments?

sulfuric acid.bmp - 759kB

497 - 7-11-2007 at 12:00

and here's the info on it

Attachment: notes.txt (1kB)
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Armistice19 - 8-11-2007 at 18:26

I was actually expecting an answer like this one. I was attempting to add a bit of speculation to an old oxidation theory which I combined with your previous statement about bubbling air through sulfurous acid. I was taking a risk, but now I honestly believe it was worth the informative response. I also found that my method would actually prove quite practical, but only in the presence of a Vanadium Oxide catalyst, furthermore using that method would not be categorized as the lead chamber process at all, it’s the contact process. I despise the lead chamber process. I would much rather roast pyrite from a local “Treasures of the earth” store. Some of you are probably thinking “What? The $10 oxygen cylinders and now this?!?!” well I might as well end your confusion by stating the fact that never buy my equipment or supplies, I find that chemistry in general is too expensive for my taste, but that never stopped me from doing it.

bilcksneatff - 12-11-2007 at 16:12

Quote:
Originally posted by 497
so i decided a batch process would be much more effective. coments?


That looks like a very good idea, but I've heard that there's a better way to produce SO2. SO2 is used as a preservative in winemaking, and it is produced by putting tablets of potassium metabisulfite (or sodium metabisulfite) into the wine. You could find these in any winemaking kits. Haven't tried it myself, though.

[Edited on 12-11-2007 by bilcksneatff]

S.C. Wack - 12-11-2007 at 16:46

Not rocket science

Attachment: jce_7_1138_1930.pdf (116kB)
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chemkid - 12-11-2007 at 19:22

Another entirely unrelated question: Could ammonium nitrate be substituted for potassium or sodium nitrate?
Furthermore, a mixture of pottasium nitrate and sulfur would be the same as for gun powder correct? Essentially i would be heating gun powder until it burns?

Chemkid

The_Davster - 12-11-2007 at 19:30

Quote:
Originally posted by 497
coments?


Unless you make the tube going into the first absorbtion collumn longer and higher than the water level it will just flow into the sulfur burner.

How do you plan to keep the sulfur alight? Some form of wick would be beneficial.

bilcksneatff - 13-11-2007 at 04:19

Quote:
Originally posted by chemkid
Another entirely unrelated question: Could ammonium nitrate be substituted for potassium or sodium nitrate?
Furthermore, a mixture of pottasium nitrate and sulfur would be the same as for gun powder correct? Essentially i would be heating gun powder until it burns?

Chemkid


Ammonium nitrate could probably be used, but there is an easy way to convert it to sodium nitrate. You mix a 2:1 molar ratio of NH4NO3 with sodium carbonate (Na2CO3) and place it over low heat (actually, I put it next to my woodstove). One mole of NH4NO3 is 160 grams, and one mole of Na2CO3 is 105 grams.

2NH4NO3 + Na2CO3 --> 2NaNO3 + NH3 + H2O + CO2

The KNO3/S mixture is not really gunpowder. Gunpowder is a mixture of KNO3, sulfur, and charcoal all ground together.

[Edited on 13-11-2007 by bilcksneatff]

Aqua_Fortis_100% - 13-11-2007 at 05:44

Some months ago I was just experimentating with this interesting subject.. The first experiment I did was the described in a link posted above ('pop bottle process') . I've used a PETE 2.5 L pop bottle , some sublimed sulfur , KNO3 , 10mL or so of water, and a encurved teaspoon adapted to the lide of the pop bottle.

This is a extremely mess process, because you often became exposed to nauseous SO2. I simply have putted the KNO3 in the encurved spoon and heated with until melted,and the adding the S8 (I've found that is much better doing this than heating both simultaneously,because less sulfur is just burned in open air and hence less is wasted(since when mixed the sulfur ignites pretty fast)) which in turns gave a red-brown melt and ignited instantaneously (but smoothly ;) ) and quickly puting this in the pop bottle.. I did this 10 times(waiting ~30 mins between each burn,and letting some O2 come into the bottle), but sometimes the sulfur-KNO3 melt have fallen in the solution and causing more mess..:mad::mad:

after this , I've filtrated the suspension of sulfur in the acid misture .. then I used the procedure described in a link posted by tupency h. wherein a cooled (-10*C) solution of H2SO3 + O2 can yield H2SO4.. So I've put the filtrate in a container and this in a salt-ice bath for 20mins.. when I back to the place where the solution were standing, I've noted a white precipitate (KNO3 from the mess or any salt formed in situ like KHSO4, KHSO3, etc ¿¿¿)..

This gave me a great grief and foul thinkings :mad::mad:

so I dumped the contents on the cement floor , in backyard..and the stuff made some bubbles and the caracteristic 'whoooosh' from acids on the things.. So , at least this gave me some happy thinkings..:D (however this can also probebly by due by the H2SO3)...:(

bilcksneatff - 13-11-2007 at 13:25

Quote:
Originally posted by Aqua_Fortis_100%
Some months ago I was just experimentating with this interesting subject.. The first experiment I did was the described in a link posted above ('pop bottle process') . I've used a PETE 2.5 L pop bottle , some sublimed sulfur , KNO3 , 10mL or so of water, and a encurved teaspoon adapted to the lide of the pop bottle. ...:(


I suggest using a copper endcap instead of a spoon. Less chance of producing a mess of KNO3 in the resulting acid. See http://cavemanchemistry.com/cavebook/chsaltpeter3.html

chemkid - 13-11-2007 at 15:49

One last idea i omitted in my prior post. Would using N2O work as well?

Chemkid

bilcksneatff - 13-11-2007 at 16:25

Quote:
Originally posted by chemkid
One last idea i omitted in my prior post. Would using N2O work as well?

Chemkid


N2O would definitely not work. The nitrates produce NO2, not N2O.

497 - 13-11-2007 at 21:45

sorry i've been away so long...

@The_Davster, the water won't flow back into the sulfur burner because its under pressure from the pump. the wick is a good idea, i have yet to attempt to build the burner. i've lit a pile of it outside and it burned quite nicely just in a puddle.

@bilcksneatff, the metabisulfite is a good idea, especially if the sulfur burner turns out to not work... but i don't think i can easily get ahold of any, im not exactly in wine country. and i would guess that its quite a bit more expensive that plain old sulfur. worth a try anyway.

thanks for the input. i'm sure theres plenty more to improve on the overall design. unfortunately it might be a while until i have to time to build one... but i will.

[Edited on 13-11-2007 by 497]

The_Davster - 13-11-2007 at 21:51

Open the burner to add more sulfur? Power failure? A couple more feet of pipe could prevent a nasty/annoying accident.

I once cast some sulfur around a thick cloth wick for a sufur candle. Whether or not you would like to use cloth is up to you, but a tight coil of some metal mesh would have the same effect without the potential contamination of the produced acid.

497 - 13-11-2007 at 21:58

you are right, having it run back in would be a pain in the ass. as soon as i have a chance i will test some different designs for sulfur burners.

Aqua_Fortis_100% - 14-11-2007 at 03:14

Quote:
Originally posted by bilcksneatff :
I suggest using a copper endcap instead of a spoon. Less chance of producing a mess of KNO3 in the resulting acid. See http://cavemanchemistry.com/cavebook/chsaltpeter3.html


Yes that's the link I'm refering in the post. I've used the spoon solely because in the time I did the procedure I didn't have ANY copper/iron wire or endcap laying around, and lazy to go to the local (6 km far from my house :mad: ) hardware store... So I've did the necessary improvises..

[Edited on 14-11-2007 by Aqua_Fortis_100%]

bilcksneatff - 14-11-2007 at 04:05

Quote:
Originally posted by Aqua_Fortis_100%

Yes that's the link I'm refering in the post. I've used the spoon solely because in the time I did the procedure I didn't have ANY copper/iron wire or endcap laying around, and lazy to go to the local (6 km far from my house :mad: ) hardware store... So I've did the necessary improvises..

[Edited on 14-11-2007 by Aqua_Fortis_100%]


I understand, I was just thinking a copper endcap would be less likely to spill KNO3 or sulfur into your acid.

Quote:
Originally posted by 497

@bilcksneatff, the metabisulfite is a good idea, especially if the sulfur burner turns out to not work... but i don't think i can easily get ahold of any, im not exactly in wine country. and i would guess that its quite a bit more expensive that plain old sulfur. worth a try anyway.

[Edited on 13-11-2007 by 497]


I would check www.midwestsupplies.com/Products/ProdByID.aspx?ProdID=4938

They have a one pound bag of potassium metabisulfite for U.S. $14.50.

[Edited on 14-11-2007 by bilcksneatff]

[Edited on 14-11-2007 by bilcksneatff]

not_important - 14-11-2007 at 04:34

Note that in N America sulfur should cost $1 to $1.50 per pound for garden grade, which should be washed with water to remove (organic) wetting agents before use.

Burning a pound of sulfur will give you about 2 pounds of SO2, while adding acid to a pound of bisulfite will give you less than a pound of SO2.

The trick is in getting a small size sulfur burner to work well.

bilcksneatff - 14-11-2007 at 06:12

Quote:
Originally posted by not_important
Note that in N America sulfur should cost $1 to $1.50 per pound for garden grade, which should be washed with water to remove (organic) wetting agents before use.

Burning a pound of sulfur will give you about 2 pounds of SO2, while adding acid to a pound of bisulfite will give you less than a pound of SO2.

The trick is in getting a small size sulfur burner to work well.


One pound S --> ~1.5 pounds SO2
One pound K2S2O5 --> ~1 pound SO2

Didn't really think about it that way!

Quote:
Originally posted by 497
so i decided a batch process would be much more effective. coments?


I thought of another idea. Would 10-15% H2O2 be cheaper for you than HNO3? You can simply bubble SO2 through 10-15% H2O2 to get H2SO4. The only problem I see is the fact that the HNO3 will be regenerated, whereas the H2O2 will not.

[Edited on 14-11-2007 by bilcksneatff]

Upscaling...

Aqua_Fortis_100% - 14-11-2007 at 17:51

This Drawing is my new backyard chamber..(I dont have a cam at the momment, So I did a fast drawing)..This was did quite a some days, and didn't post before together with the earlier posts because of the hurry I'm in..

I've did some simple calculation on gases before the pratice and will present below. Please, if anyone find in my calculations some weird or wrong thing, feel free to tell and correct me..

(I)

Descripition:

PETE bottle is a fairly resistant material against many diluted and weak acids, I often store filtrated battery acid in it without any problem, and also is a fairly inexpensive material to get(usually free if you like drink or 'hunting' in the weird and funny neighbourhood trash).. That's the reason I've did chosen and used this material in my experiments.

My chamber consists of a (more or less irregular :D ) serie of pop-bottles straight on the floor (the generator are also in the same position..My drawing mess up in this also ; be aware to not think the pop-bottles are upside-down in a weird position )and with some wet cloth on it ,( to help cooling the inlet gases). Each of them being connected to the other by means of 8 cm piece of common aquarium tubing and with the help of hotmelt glue.

At one extreme of the chamber, the aquarium tube is connected in a side of a empty kerosene metal can(with some PTFE, to give good seal), which in turns holds a pierced cork stopper with generous amounts of PTFE around it ( to both protect the cork stopper and to improve seal) with an radio antenna Al tubing which is jointed together with a old refrigerator copper tubing (PTFE is used also here, since the temperature, IMO, isn't that great at the point). The copper tubing is connected in the Rojek lid of my 'reaction vessel' ( a tomato 'tin' can ) and some plaster of paris (I just used this because is the most common and the only that I know somewhat..I've great problems with sealing and soldering tubes to vessels (speciallity the first),because the hardware store here recently don't appear to be an friendly place to visit..high prices and very few interesting things). The edges of the tin can were wrapped with some Al foil, to help keep the lid in the correct place(the rojek lid seem to have tiny amounts of a plastic material in it..bad news) and the lid pressured in the place... The holder of the reaction vessel in the drawing is a iron type paraphernalia, but in my experiment was simply 2 common bricks and the alcohol burner under the tin can...

The reactants were impure (CaCO3 - gardening) sulfur ,tap water (10mL in each pop-bottle) and double saltpeter (NaNO3/KNO3,1:1 - without wax additives, which were removed-, which I just considered as KNO3 to dont increase the trouble..)
..

(II)

Consulting Axehandle rxns an overall one can be made, based on the natural sulfur allotrope...

S8 + 16/3 KNO3 --------> 8/3 K2S + 16/3 SO2 + 16/3 NO


S8 => 32.1*8 = 256,8g
16/3 KNO3 => (16/3)*101.1 = 539.2g
------------------------------------------
8/3 K2S => (8/3)*110.3 = 294.1g
16/3 SO2 => (16/3)*64.1 = 341.9g
16/3 NO => (16/3)*30 = 160g

========================

(III)

So, the sum of gaseous products is NO + SO2 , what means 16/3 + 16/3 mols => 32/3 mols of gas

arbitrarily (because neither of my thermometers are reliable to measure) considering the final temp (after passing the last pop-bottle and venting to outdoor) as 30°C, PV=nRT can be used to determine the final volume of all gases in the system:

1*V = (32/3)*0.082*303
V= ~265.02 L

======================

(IV)

Next step is calculating the total volume of all pop-bottles connected to each other...(desprezing minor volumes like the tubing and can volumes,water, etc)..

pop-bottle type: PB Amount: Volume(L):

1.5 L........................3.....................4.5

2.0 L........................23...................46

2.5 L........................1.....................2.5
Total PB: 27 Total Vol: 53

So, by dividing the total vol of the chamber by the arbitrary value found at (IV) we will find the theoretical necessary amount of reactants needed to fill all the chamber:

53L/265.02L = ~ 0.2

so,

16/3 SO2 * 0.2 = ~ 1.067 SO2
16.3 NO * 0.2 = ~ 1.067 NO

Now looking also (II):

If to produce 16/3 moles of SO2 we need 256.8g of S8, what is the amount of S8 needed, if we need only 1.067 mols of SO2 ? AFAIK: x = ~51.4g S8

Doing a similar calculation with the KNO3 we will find: ~ 107.87g KNO3..

So again arbitrarily, I've just cutted the amounts aproximately by half: 25g S8 and 55g KNO3 by shoot...

the results of my experiments I will post tomorrow, because its getting late here.. (the results were very weird at principle..but just funny) :D

[Edited on 14-11-2007 by Aqua_Fortis_100%]

h2so4[1].bmp - 1.8MB

497 - 15-11-2007 at 12:02

why not do the same thing but in water, like my design? it seems much simpler and easier, not to mention more compact. and if using HNO3 you don't need to screw around with KNO3. and if you do decide to do a gas phase reaction, there are much more effective containers (5 gallon HDPE buckets.) it will be a nightmare getting all those tubing connection to seal, belive me i've been doing alot of that for another project.

H2O2 might should work, but i like the HNO3 better because its not consumed.

on another note, i've had some more success producing HNO3. more details in the thread in technochemistry.

[Edited on 15-11-2007 by 497]

Aqua_Fortis_100% - 15-11-2007 at 17:57

@497, yes your method seems to have much less labour than my one..but you are using HNO3 and this isn't very much OTC in most countries, because of its very interesting and 'popular' properties/uses..;) .. Additionally most people here and in other places wants H2SO4 actually to make HNO3 , not the opposite(well..the lead chamber process can be used just for fun, like my project, or then for real production of H2SO4, which is very hard to worth, because which sulfuric is still a widely avaliable chemical, even if will not be OTC in concentrated form..also is difficult to have near 100% no losses in HNO3 and hence , at a given time you will need to replace some, more on that latter).

Ironically, the initial results of my 'lead chamber' were... NITRIC ACID ! While I've just carefully watching the heating of the rxn vessel, large brown/red clouds evolved during the start and filled the kerosene can , the first and 2nd pop-bottles..I´m very sure that this stuff is NOx, not only by their colour but because also of it foul and toxic smell.. after, when the rxn vessel started to evolve white fumes of SO2, the improvised sealing on the rxn can failed and started to leak out..what forced me to remove the can immediatelly from the heat source and allowing to cool down..inside the can a black boiling mass..

This is weird for me, because the sulfur tends to react in open air with KNO3 in a manner completelly diferent: burning imediatelly and evolving SO2, and I did tried this many times.. but I just thinked before that the KNO3 itself are able to give some oxygen to S8 and start a oxidant/combustible type reaction and burning without O2..,But inside the can the only thing happened was the boiling of the mass (without fire..even when the lid was removed).. So my next try would be adaptating a hose from the aquarium pump to the reaction vessel and allowing lo force some air along it..this will prevent this mess to happen , fire the mix , reduce the nitrate consumption and also give a more homogeneous distribution of the gasses along the bottles..But also maybe requires more volume to work properly, without leaking out too much gasses before they react (which means more bottles attached) or lower amounts of reactants per shoot..


Quote:
and if you do decide to do a gas phase reaction, there are much more effective containers (5 gallon HDPE buckets.) it will be a nightmare getting all those tubing connection to seal, belive me i've been doing alot of that for another project.


497, I didn't found make these plastic connection as mess to do(was even fun!), although the sealing between the metalic, heat exposed parts was a bitc* ..If anyone can advise me about how to give good seal these parts I will appreciate a lot..

And about the HDPE containers, I know most people here like of it, but I insist in several PETE connected together, because ,chiefly , a much more effective reaction area is obtained, whereas using a HDPE container you will have only a tiny area of the liquid at the bottom. And I found experimentally, also by accident, that the surface area of the water is even MUCH more increased when the bottle chamber is exposed to the sun : the water evaporated condenses on whole wall of the PETE containers , creating many tiny drops and huge surface area to reactions take place..



ah... For some reason I'm just happy somewhat with my very diluted < 1-2%(?) nitric acid..maybe because was created too weirdly..

497 - 16-11-2007 at 18:11

well why don't you just use that nitric you got in my process??? :P but really you have a point, nitric isn't the most OTC.... one reason i'm working on a nitric acid producing apparatus, see: https://sciencemadness.org/talk/viewthread.php?tid=1518&...

and as for the surface area issue, i agree the PETE bottles will have more. but i think that could be solved by agitating the water in an HDPE bucket. either way you're having fun, and thats what really matters. :D

another thing, i don't think you really need to heat the reaction KNO3/S8 reaction. from what i've read earlier in the thread all you need to do is light the mixture with a lighter and it burns just fine after that. 3S + 2KNO3 --> K2S + SO2 +2NO, etc.

anyone know a good way to measure the concentration of nitric acid? i have a pH/conductivity meter, but i don't trust it's accuracy above very dilute acid.

The_Davster - 16-11-2007 at 18:24

Quote:
Originally posted by 497


anyone know a good way to measure the concentration of nitric acid? i have a pH/conductivity meter, but i don't trust it's accuracy above very dilute acid.


Titration with NaOH solution(standardize against an acid of known concentration first) using the pH meter as an endpoint indicator.

How accurate is your scale? Density is convenient to measure the concentration of acid. http://www.efma.org/publications/NitricAcid/Section15.asp

12AX7 - 16-11-2007 at 18:51

Ah, make your own! My large balance is accurate to at least 0.5g, at weights up to 10 kg! A small scale one can be readily made with little more than coathangers. Your accuracy is only limited by measurements on the beam, friction and your test weights.

Tim

497 - 16-11-2007 at 18:58

heh, you just read my mind, i was just thinking about what materials i could find to make something like you describe.

also i just found online some sodium metabisulfite for $2 a pound plus shipping when you buy 23 pounds. thats pretty good, it seems a bit easier than burning sulfur.

23 pounds = 55 mols Na2S205 = 3.5kg SO2 = 5.4 kg H2SO4

pretty good for 50 bucks

and also the resulting Na2SO3 can be made to produce another mol of SO2 by Na2SO3 + 2HCl ---> 2NaCl + H20 + SO2. if you can HCl thats cheap enough.

[Edited on 16-11-2007 by 497]

[Edited on 16-11-2007 by 497]

bilcksneatff - 19-11-2007 at 08:53

497, where did you find that Na2S2O5?

This is a random question, but if you burned KNO3 or NaNO3 with charcoal, would you get K2CO3, CO2, and NO?

2KNO3 + 2C --> K2CO3 + CO2 + NO



[Edited on 19-11-2007 by bilcksneatff]

497 - 19-11-2007 at 12:02

http://www.chemistrystore.com/sodium_metabisulfite.htm

i've never heard of that reaction, i dont know...

bilcksneatff - 24-11-2007 at 10:52

Here's something very interesting from http://www.ucc.ie/academic/chem/dolchem/html/comp/h2so3.html...

"Reactions of Sulphurous acid
The solution when heated in a sealed tube at 150 deg.C. deposits sulphur.



3H2SO3 ===> 2H2SO4 + H2O + S


Sulphurous acid can be oxidised by the use of strong oxidising agents.

Oxidising of Sulphurous acid by Oxygen



2H2SO3 + O2 + 4H2O ===> 4H30(ion) + 2SO4(ion)


Sulphurous acid solution is slowly oxidised by atmospheric oxygen to sulphuric acid.


Oxidising of Sulphurous acid by Permanganate ions

When Sulphurous acid is added to permanganate ion which is coloured purple, SO2 will decolourise the MnO4(ion) when it is reduced to the colourless Mn(ion).


2MnO4(ion) + 5H2SO3 + 4H2O ===> 2Mn(ion) + 4H3O(ion) + 5SO4(ion) + 3H2O"

chemkid - 25-11-2007 at 10:39

Here is another sulfuric acid making idea that is somewhat close to this thread. I am not sure if this works but here it is....

Na2SO4 + 2CH3COOH --> 2CH3COONa + H2SO4

Would this work?

Chemkid

[Edited on 25-11-2007 by chemkid]

not_important - 25-11-2007 at 10:53

Compare the Ka (acid strength) values for acetic, H2DO4, and HSO4(-). Also think as to which has the lower boiling point and so could be distilled away from a reaction mix.

bilcksneatff - 25-11-2007 at 11:14

Quote:
Originally posted by chemkid
Here is another sulfuric acid making idea that is somewhat close to this thread. I am not sure if this works but here it is....

Na2SO4 + 2CH3COOH --> 2CH3COONa + H2SO4

Would this work?

Chemkid

[Edited on 25-11-2007 by chemkid]


I know that it works in the opposite direction.

H2SO4 + 2CH3COONa --> 2CH3COOH + Na2SO4

If it even works at all, it would probably require a lot of energy.

Pixicious - 16-2-2008 at 13:05

Why wouldn't this work?

SO2 + H2O ==> H2SO3
3H2SO3 ==> 2H2SO4 + H2O + S

Sulphur is burnt in a small container with a little air coming in from the sides, then sealed up when smoke starts coming out and the flame allowed to die..

SO2 is produced and water is at the bottom of the container. SO2 is soliable in water. (See first equation) - or does this need a special temperature. (somewhere between 0-100C?)

There we have H2SO3, this is then placed inside a sealed tube and heated to 150C.

We have a diluted H2SO4 with a little sulphur. Is sulphur solible in Sulphuric acid? If not could it be seperated?

Btw I'm new here but I've been reading through your site with some interest.

I guess what I am asking is.. How is sulphur dioxide disolved in water?


[Edited on 16-2-2008 by Pixicious]

microcosmicus - 16-2-2008 at 14:47

Dissolving SO2 in water is easy --- you just mix the gas with water and it dissolves.
Like most gases, it dissolves better at low temperature. More specifically, at room
pressure, the solubility in g/L at different temperatures is as follows:

0C 220
10C 150
20C 110
30C 80
40C 65
50C 50
60C 40
70C 35
80C 34

Your idea of burning sulphur in a container with water, then closing it sounds like
how I proceed. Once the sulphur is burnt and the container closed, I then shake it
around to get the SO2 well mixed with the H2O. As the table shows, you will
get more gas into solution if you use cold water.

The hard part is getting it to disproportionate. I have also heard about the sealed
tube at 150C, however the question is at what rate this reaction proceeds. I once
tried it, heating the tube for an hour but didn't notice any appreciable change.
However, the metal bottle I used was hardly the best vessel for the purpose,
it is just what happened to be at hand. One day, when I buy or make a glass
ampoule, I plan to seal some SO2 solution in it and try again. I suspect that
it will take a long time. Even if it not a practical way of making H2SO4, it would
be interesting to see this reaction happen.

As for solubility of sulphur, since the sulphuric acid is going to be quite dilute,
it should precipitate out just fine, so separation would be a matter of filtering.
Unreacted SO3 could be eliminated by heating, sulphuric acid having a
rather high boiling point.

[Edited on 16-2-2008 by microcosmicus]

not_important - 16-2-2008 at 16:53

Don't even need to heat it. Fill a gas tube with SO2 gas, put it in sunlight or other bright light. Ever so often look through the tube lengthwise or as near as you can get, it will become foggy.

3 SO2 <=> 2 SO3 + S

A problem with working with gases is the low density. Three moles of SO2 is going to occupy about 67 liters at STP. This would yield one mole of sulfur, roughly 16 cc, and two moles of SO3, between 50 and 90 cc depending if solid or liquid and which form, or about 110 ml of H2SO4. Now it's true that you start out with only about 55 cc of sulfur to make the SO2, it's just the the volume of gas is inconvenient.

I can envision a flask, strongly illuminated , with a fractionating column - lots of surface area, both held at about 100 C. As sulfur forms it drifts to the surfaces, collects, and molten runs down to the bottom of the flask. The SO2 and SO3 pass through the column, the SO3 is condensed in another flask at say 20 C. The system is sealed, unreacted SO2 just drifts about until it reacts. It could be pressurised, which should help due to the 3:2 volume ration.

Some SOx will dissolve in the sulfur, some SO2 will dissolve in the SO3, how much I don't know.

It's a very slow process, you would want a concentrated source of SO2, say liquefied from a sulfur burner and kept under pressure, and a way to feed it into the system whenever the pressure dropped low enough. It's not very practical, more of a intellectual joyride.

497 - 16-2-2008 at 23:27

Could you bypass the issues of large volumes by liquefying the SO2? It boils at -10C or 34 PSI at room temp. Hopefully it can actually decompose in liquid phase. It would need to be contained in glass to decompose it with light. Shouldn't be a problem in large borosilicate tubing or the like. That would be pretty awesome if you could just set a tube of liquid SO2 out in the sun (or 150 watt metal halide that I happen to have :P) for a few hours/days/weeks and come back a have a tube full of SO3. Maybe even use a little parabolic reflector to speed things up. Somehow I doubt it would be that easy.

Also I've found that 32%H2SO4 + 50%H2O2 reacting with copper metal produces quite a lot of nice clean SO2 (maybe some O2 too) and quite a bit of heat, that tends to start to boil it after a while.

Edit: I found pressure specs for borosilicate tubing, 38 mm dia 4 mm wall can handle 200 psi, so there should be no problems with pressure as long as the end caps can be affixed strongly. Having a means of getting the liquid in and out that is resistant to SO3 might be a challenge though. Or you could just not worry about it and break the tube open when you get the SO3 out.


[Edited on 16-2-2008 by 497]

Pixicious - 17-2-2008 at 18:43

I put together a fosters beer can and make an attempt.

I have something with a ph 2 and something which turns limus paper red.

What tests are there to determine whether what I have is H2SO3 or H2SO4?

I can see myself having either H2SO4 or H2SO3 because from reading it is said that SO3 in small amounts is given off and is the smoke you see and the ph is higher than I would expect for H2SO3.

I have a flask which can be made air-tight, do I fill this with hydrogen before sealing or do I keep it with normal air before heating?

I'n going to give that a go this evening.

497 - 17-2-2008 at 18:58

I'm pretty sure the pH of real H2SO4 is going to be quite a bit lower. Dip a paper towel in it, it will start to dissolve if its H2SO4 of any decent concentration, I'm pretty sure it won't dissolve in H2SO3. How much water did you add? The beer can is only going to hold like 0.01 moles of SO2, with very much water its going to be very diluted.

From what I've read its not that easy to convert H2SO3 > H2SO4 without some more extensive manipulation.

I have no idea what you're saying about the hydrogen...

Pixicious - 17-2-2008 at 19:43

About 20ml of water was put at the bottom of the can and the experiment has been run 4-5 times to increase the concentration. From 1-10g of sulphur have been burn with an unknown percentage escaping before the cap was put on.

I wish I had done it with a little more accuracy.

I will try the paper towel test right now.

Pixicious - 17-2-2008 at 20:09

The paper towel was placed inside for 10-15 seconds and hadn't dissolved. I'm worried I may have lost a little since the ph is now reading at about 3..

The only reason I spoke about the hydrogen is due to a reaction with the air when heating inside a sealed container.

I will nake a couple batches of liquid in this way and put to one side until a) they all have a ph of 2 and I'm satified I have enough before I can make an half decent attempt at making sulphuric acid in this way.

I'm wondering if I can buy a 2l coke bottle and burn more sulphur to increase the production rate. I will also be able 'to see' when the flame dies and all the SO2 and SO3 has been dissolved,


[Edited on 18-2-2008 by Pixicious]

497 - 17-2-2008 at 22:19

Hey, anyone got any I2 laying around? ;) I just realized that you could use part of the iodine-sulfur process to make good concentrated H2SO4.

I2 + SO2 + 2 H2O -> 2HI + H2SO4

Then recycle the iodine:

2HI -> H2 + I2

I also found out that liquid sulfur reacts with SO3 to make SO2... not good. Hopefully its not the case for solid sulfur.

[Edited on 17-2-2008 by 497]

not_important - 17-2-2008 at 23:18

Quote:
Originally posted by 497
Hey, anyone got any I2 laying around? ;) I just realized that you could use part of the iodine-sulfur process to make good concentrated H2SO4.

I2 + SO2 + 2 H2O -> 2HI + H2SO4

Then recycle the iodine:

2HI -> H2 + I2

I also found out that liquid sulfur reacts with SO3 to make SO2... not good. Hopefully its not the case for solid sulfur.

[Edited on 17-2-2008 by 497]



You're referring to the Bunsen reaction for the first part. It's one way of making aqueous hydrogen iodide in the lab.

The reaction is an equilibrium, too high of a concentration of the acids drives it towards I2 + SO2 + H2O. This includes distilling the HI-H2O azeotrope away from the mix.

In the IS cycle to produce H2, the separation is accomplished by using an excess of I2 and some starting HI to form two phases, HI-I2-H2O and H2O-H2SO4. The excess iodine, plus an excess of water, is used to help drive the reaction to HI + H2SO4.

In that case it typically is run at 110 to 130 C, with cooling as it is exothermic.

I'm attaching part of a material balance sheet for the Bunsen section of a ISH plant.

section 115 is the feed into the Bunsen reactor, run under a bit of pressure.
116 is the flow out of the reactor into a phase separator
117A is the off gases
118A is the H2SO4 phase
119A is the HI-I2 phase

Note that both acid streams are fairly dilute, the H2SO4 also contains a little SO2 which is no big deal. However the acid concentrations, plus the large of iodine used, might discourage you; this is especially true in the USA and Oz where the element is contraband.

Bunsen.png - 84kB

497 - 17-2-2008 at 23:56

Yes I just found out about the concentration issues on my own... how frustrating. I suppose that process is pretty much out of the question.

I found an interesting reference in the attached pdf on the "Westinghouse" process that involves electrochemically producing H2SO4 from SO2. I imagine the concentration would not be high though.

SO2 (g) + 2H2O(aq) &#1048774; H2SO4 (aq) + H2 (g)

Also does anyone know anything about the possibilities of SO2(l) -> SO3(l) + S(s) ? It seems like a good way to solve the problems of large volumes of SO2 gas...

Attachment: lect5.pdf (2.2MB)
This file has been downloaded 1296 times


497 - 18-2-2008 at 00:04

Here's some good info on the Westinghouse cycle. They talk about electrolyzing acid up to 80 or 90 w-%, and they're more worried about efficiency, it probably could go even higher since electricity costs wouldn't be much of an issue. I'm kind of surprised I've never heard that this could be done.

[Edited on 17-2-2008 by 497]

Pixicious - 18-2-2008 at 03:43

Then what I currently have (and gathering) in a jar is H2SO3 which when heated will produce H2O + SO2.

I do have some iodine. Could it be worth mixing the iodine and H2SO3 into a sealed container and heat upto 120C? If this is a result the two should be pretty pure, if the reaction is allowed to complete and correct quantites used.

I will have a method of producing pure (as close as after opening) H2SO4, in maybe a reasonable quanity.

[Edit] I knew I bought that iodine for a reason.

[Edited on 18-2-2008 by Pixicious]

not_important - 18-2-2008 at 05:27

As I already said, the reaction is reversible - mix H2SO4 and HI and you will get some I2, H2O, and SO2. An excess of I2 and H2O is used to force the reaction, the concentration of the acids produced is less than 20%. Use less water, more SO2 and I2 will remain unreacted.

Pixicious - 19-2-2008 at 08:06

I heard you.

What you are saying is I should release the gas at the end of the experiement and boil the excess water forming a concentrated H2SO4.

497 - 19-2-2008 at 13:26

That might work, the problem is you will end up with very little. I'd say at most 5 grams from a 2 liter bottle. If I were you I'd try like a plastic 55 gallon drum or a big trash can that you can seal. In theory 55 gallons of SO2 gas could yield around 910 grams H2SO4. I doubt you could get that in real life, for one thing you're not going to be able to fill the container completely with SO2 by burning it in the container although you could do better with straight O2 injected.

I think aqueous conversion would be much better, that same 55 gallons of SO2 gas could be dissolved in just 3 liters of cold water. Its been mention in this thread I think, and it sounds worth trying, freezing the aqueous SO2 to convert it. I think they say 3 freeze-thaw cycles and its completely converted. I'm not sure how concentrated the SO2 can be though. You'd still have to concentrate it a lot too, you'd only get around 20 or 30% concentration initially.

What I may do some time when I have the chance is fill a 5 gallon bucket with ice water, build a good sulfur burned and pump straight SO2 into that water. It should dissolve about 3 kilos of SO2, yielding in theory about 5 kilos of H2SO4. Thats a lot of boiling to concentrate that all... at least 20 kilowatts of energy. And considering I can buy 5 gallons of battery acid (8 kilos H2SO4) for $15, it seems like too much work... unless maybe you had an enormous sulfur burner that saturated a 55 gallon (makeing 70 kilos H2SO4!) drum of water with SO2... sigh... maybe some day if H2SO4 becomes regulated.

Pixicious - 20-2-2008 at 02:52

Quote:

...and it sounds worth trying, freezing the aqueous SO2 to convert it.

Are you saying freezing the aqu SO2 will convert it to SO3?

not_important - 20-2-2008 at 02:58

Quote:
Originally posted by Pixicious
I heard you.

What you are saying is I should release the gas at the end of the experiement and boil the excess water forming a concentrated H2SO4.



The liquid will be a mixture of water, H2SO4, and HI. Concentrating it will result in the reverse reaction, at least to some extent. So as you start to boil off water, you'll be driving SO2 and some I2 off.

This is why the sulfur-iodine cycle projects use a major excess of I2 and have some HI in the input stream, the two liquid phases formed allow the separation of the H2SO4 from most of the HI.

497 - 21-2-2008 at 21:34

Pixicious, yes you are right it should if done correctly convert SO2 to SO3 which will immediately be converted H2SO4, probably fairly dilute, but at least it'd be cheap... assuming you can make an effective sulfur burner.

Pixicious - 22-2-2008 at 02:11

SO2 + 2H2O -> H2SO3 + H2O -> 2H + SO3 + H2O -> H2SO4

I assume then the reaction must be something like that.

I'll give it a try this evening. I won't doubt you I just don't see it.

[Edit] So far so good. I made a new batch of H2SO3 this morning with excess water. It is in the freezer at the moment. I'm assuming an excess of water is needed for the reaction. It had a ph of 2-3. Limius was slightly red.

There is a solution which is taking longer to freeze than the water. won't try filtering it off until it has been frozen and allowed to thaw three times. The lid has been and will be sealed throughout. I will seperate the two out on the fourth freeze and see what they both are. It would be interesting to me to find out why.

497: My suphur burner is comprised of a 2litre coke bottle cut in two with the bottom of a fosters can supported half way by two metal pieces of a clothes peg. There is a side window and I'm using selotape as a sealent. Works rather well. It is lit and sealed and the sulphur burns for maybe five minutes (I use a mini blowtorch to light as much sulphur as I can) I can run in batches, after 3-5 a PH of 2 is averaged. The amount of water I use is about 100ml.



[Edited on 23-2-2008 by Pixicious]


[Edited on 23-2-2008 by Pixicious]

Pixicious - 4-3-2008 at 06:21

It didn't work I am afraid. Sorry 497 gave several attempts. (my previous post has no edit button)

Would it be possible -at all- to produce SO3 from Na2SO4?

LSD25 - 6-4-2008 at 22:17

Here is something interesting and funny at the same time:

http://blog.modernmechanix.com/2008/03/05/dangerous-acids-ma...

BTW The trick with the freeze/thaw aqueous variant is to keep it in the freezer for a whole lot longer than they say, until the whole of the solution is is absolutely solid (took hours in my home freezer) and then defrost the lot (clathrate included - around 15-20C). I am in the process of defrosting one of my attempts, the liquid in this, admittedly dilute, sample when half-defrosted is at PH.1-2 (PH 1-14 paper). That sort of suggests it ain't likely to be sulfurous acid, that and the precipitation of calcium sulfate as the conversion took place.

garage chemist - 6-4-2008 at 23:25

In aqueous solution, SO2 is oxidised by aerial oxygen.
You could try bubbling air into your sulfurous acid, very slowly, to avoid expelling too much of the SO2 and give the oxygen time to react.

LSD25 - 7-4-2008 at 02:24

I was actually referring to that Japanese article, where they use dissolved oxygen to oxidise H2SO3 in aqueous solution. They claim 100% conversion via three 1 hour freeze cycles, I did not find this to be accurate, in fact, overnight freezing - three to four times - was required to change the properties of the material when allowed to defrost at room temp (the clathrate increased then decreased to nothing on the fourth defrost cycle). I really should use a hydrometer to work out the spec.grav of the suspected sulfuric acid. It would also have helped to have checked the PH of the H2SO3 mixture on each thaw cycle to see if it decreased or increased during that time.

PS Can I use Calcium chloride instead of barium to ascertain the identity of the acid?

Armistice19 - 7-4-2008 at 10:55

We DON'T need people advocating criminal activity here. Kindly take your "business" elsewhere. This is a CHEMISTRY forum.


Edited by vulture


[Edited on 7-4-2008 by vulture]

497 - 8-4-2008 at 20:44

Quote:

Would it be possible -at all- to produce SO3 from Na2SO4


https://sciencemadness.org/talk/viewthread.php?tid=10217

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