Sciencemadness Discussion Board - H2SO4 by the Lead Chamber Process - success

H2SO4 by the Lead Chamber Process - success

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 H₂SO₄ 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 H₂SO₄, and wasn't replaced with the contact process (using a Pt catalyst to oxidize SO₂ to SO₃) 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. SO₂ 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 SO₂ to the trioxide SO₃; which combines with the water on the floor, as well as with water vapour (note 2) inside the chamber, to form dilute H₂SO₄. The reaction mechanism is described by the following, unbalanced reactions:

(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

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 KNO₃ 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 KNO₃ 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 KNO₃. When it gets warm enough after about 30 seconds, the KNO₃ 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 SO₂ + 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 KNO₃, but I have 4 reasons:
1) The chamber size is very small and has limited oxygen. Thus a stochiometric amount of KNO₃ will make all the oxygen for (a, b) come from the KNO₃, 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 SO₂ inside the chamber is oxidized to SO₃ 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/NO₂ 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 KNO₃ into the dilute H₂SO₄. 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 HNO₃ from the KNO₃ in the H₂SO₄ 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 H₂SO₄ 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 KNO₃ 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 KNO₃, a bit of water at the bottom... aaah, I suppose I could take a picture later...

About dissolving SO₂: Not sure there. Have only tried it with very small amounts of water. Wouldn't the SO₂ 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 SO₂(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 HNOSO₄.

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 KNO₃ powder, some of the sulfur will vapourize during the burn, once the fire gets hot enough to make the KNO₃ 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 H₂SO₄, 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 H₂SO₄ come from autodecomposition of H₂SO₃, but I think the process would be horribly inefficient -- days instead of hours, and very dilute acid. BTW, the reaction
S + KNO₃ --> SO₂ + KNO
also takes place, AFAIK. 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 SO₂ upon opening the container is proof.

Oh, and most nitrates would probably work. In the old days, NaNO₃ was used. Perhaps even NH₄NO₃ 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 HNO₂ solution which readily oxidizes the H₂SO₃. 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

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 975 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 800 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 KNO₃ 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#pid66270

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)

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/_artic...

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 574 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

[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_Chamber.h...

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)
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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

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

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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:

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:

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..

(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 http://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

) 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)

[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???

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&page=7.

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.

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#Sulphu...

"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

) 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?

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) 􀃆 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 16 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-made-safe...

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

Contrabasso - 10-4-2008 at 11:41

Would it be reasonable to try Axehandle's process with a lead burning cup in a 200 litre plastic drum and feed it with a low flow of oxygen - just enough to make bubbles come out of the exhaust into water.

I'm thinking of making a thing like a HUGE distilation kit out of drain pipe so that the main reaction happens in a downward sloping 3 inch tube say 3 metres long dipping into a flask of water.

The aim being to burn say a kilo of sulphur and make a kilo of acid at a time.

Yes I have a source of O2 in cylinders!

497 - 10-4-2008 at 21:07

Are you planning on using a nitrate in combination to produce SO3? Or freeze/thaw method? You won't need a huge volume like a barrel if you're simply collecting the SO2 for later conversion/use. That definitely would be cool to be able to produce that much at a time.

Another thing about burning sulfur. Its a pain in the ass with pure oxygen, I've done it. The heat produced vaporizes the sulfur and it condenses all over, plugging things up and being a general annoyance. It is nice to have pure SO2 gas produced, but you could save your money by burning it with air and then using dry ice (that should be cheap) to condense the SO2 in pure liquid form. Then you can use it however you want.

[Edited on 10-4-2008 by 497]

LSD25 - 10-4-2008 at 22:10

I posted a link on the previous page, it appears to have been overlooked, but it deals DIRECTLY with the topic of this thread and provides a FOOLPROOF method of making H2SO4 by the contact process using a supported iron oxide catalyst (it also details how to make it):

Quote:

To prepare sulphuric acid, you will need some sulphur, water, calcium chloride, and iron (ferric) oxide. The experiment is a simple one and requires only homemade apparatus consisting of a bottle, a flask, glass tubing, a few corks, a glass funnel, a gas burner, and rubber tubing. The parts should be arranged as shown in the illustrations. Flowers of sulphur placed in the shallow lid from a tin can is burned under the funnel at the extreme right. The sulphur dioxide formed together with some air is collected by the funnel and then passes through a drying bottle, containing the calcium chloride, to the horizontal tube of hot iron oxide. The presence of the hot iron oxide causes the sulphur dioxide to steal oxygen from the air and become sulphur trioxide. Because in this reaction, it induces a chemical change in another substance and is unchanged itself, the iron oxide is said to be a catalyst.

Finally, the sulphur trioxide formed is bubbled through water in the absorbing flask at the left. Being soluble, it combines with the water and a weak solution of sulphuric acid results.

Unaided, the original sulphur dioxide formed by the burning sulphur would not follow the desired course through the various tubes and bottles. To pull it through the system, suction must be applied to the mouth of the absorbing flask. This can be done by allowing water to siphon from a gallon jug and applying the suction formed in the jug to the absorbing flask by means of a length of rubber tubing as shown in the drawing.

To prepare the iron oxide catalyst for this experiment, soak some asbestos fiber or pumice stone in iron chloride or some other iron chemical solution until the mass is well saturated. Then add ammonium hydroxide (ordinary household ammonia will serve). This will precipitate iron hydroxide in the pores of the asbestos or pumice. The liquid then can be poured off, fresh water added and shaken and also poured off.

Next heat the impregnated pumice or asbestos in a crucible or tin-can lid over a gas burner. This final operation will convert the iron hydroxide into the desired iron oxide. The finished catalyst then is placed in the horizontal tube and heated gently with a gas burner as the sulphur dioxide is pulled through.

After burning about a teaspoonful of the sulphur, remove the absorber from the system and test the liquid with a piece of blue litmus paper. If an acid is present, the paper will turn pink. To prove that it is sulphuric acid, place a small quantity of the liquid in a test tube and add two drops of hydrochloric acid followed by several drops of barium chloride solution. If sulphuric acid is present, a white precipitate will be formed.

Although sulphuric acid made by this simple process will be weak, it should dissolve bits of magnesium and attack pieces of zinc to produce tiny bubbles of hydrogen gas. Of course, the concentration of the liquid can be increased by boiling but even then the home chemist will find that the acid will be too weak -to be of any great value for experimental purposes. ‘ It is interesting to note, however, that this same type of contact process is used commercially to manufacture sulphuric acid. Of course, a more expensive substance, usually a form of platinum, is used as the catalyst.

Now to find some glass tubing and get the borer out...

Does anyone else find it funny that the article covers the making of fairly good sulfuric acid - then says you need acid so strong that it cannot be made by simply boiling that acid down (ie. fuming)... Suggesting to adolescents that they need fuming H2SO4 for home use.... wouldn't I loved to have lived then (this is not gone because of drugs, but cos of the insurance industry I'd suggest).

[Edited on 10-4-2008 by LSD25]

[Edited on 10-4-2008 by LSD25]

xlg_acids_1.jpg - 46kB

not_important - 10-4-2008 at 22:23

What they said is

Quote:

Of course, the concentration of the liquid can be increased by boiling but even then the home chemist will find that the acid will be too weak -to be of any great value for experimental purposes.

It's a real dilute solution of sulfuric acid that is made, the amount of concentrated H2SO4 you'd get by boiling down would be too small to be worth the effort. BTW, unless some care is taken just boiling down dilute H2SO4 results in the lose of some acid long before the 98% stage is reached.

iron oxide is a poor catalyst for the SO2 + O2 => SO3 reaction, much SO2 remains unconverted. If this was not so, the platinum based contact process would not have been in use at that time, nor would V2O5 catalysts taken over, as vanadium is considerably more expensive than iron. If you have cheap to free sulfur, it might be OK, but I think the neighbors are going to suspect you of hosting satanic visits from all the SO2 released.

LSD25 - 10-4-2008 at 22:32

I got sulfur everywhere here, we got horses and everybody I know considers the shit to be the answer to just about every question that is ever asked regarding keeping 'em healthy and fixing 'em up.

Want to see something insane from the same author's? Try this - Atomic Energy experiments for the home chemist... Fuck, I was born 50 years to fucking late

[Edited on 10-4-2008 by LSD25]

med_chemcraft_atomic.jpg - 49kB

497 - 10-4-2008 at 22:44

I have seen this and it certainly is not foolproof.

Quote:

Finally, the sulphur trioxide formed is bubbled through water

This is evidence enough that the process is very inefficient, if the gas coming out of the catalyst tube had very much SO3 at all the reaction with water would be quite violent. I bet it has less that 2% SO3 in it. The only thing I can see this possibly being useful for is if you needed to prepare a small amount of SO3, it might work. This could never be cheap enough to be done on a very large scale to produce H2SO4.

First off, because of the every low efficiency you're going to need a lot of sulfur for a small amount of acid. In addition, you're going to need a ton of propane to keep that catalyst tube hot for very long. Then you're going to need even more propane (or electricity) to boil down the extremely dilute acid to a useful concentration. There's no way that could be cost effective.

I have recently been thinking about using ozone to oxidize SO2. I know it can be done when water is present, I'm not sure about in anhydrous conditions. If it would work in anhydrous environments then it might be a decent way to produce pure SO3. In any case an ozone generator could be built and run continuously cheaply and easily.

In the end I doubt it is possible for the amateur to beat the lead chamber process in cost effectiveness. I still like to think about improvements though.

Edit: not_important, you beat me to it...

[Edited on 10-4-2008 by 497]

LSD25 - 11-4-2008 at 00:19

Why not just run more SO3 into the liquid? By the look of it, the SO3 originally produced is run into H2SO3 (well, if all of the SO2 is not being oxidised, then H2SO3 would exist, no?), not water, but if more SO3 was run into the H2SO4 by continued running of the process, wouldn't that provide a stronger acid (perhaps cooling would be necessary)? The lead-chamber process & the contact process, don't do so primarily (or so it appears) because of the difficulties inherent in handling or containing the strong acid - here the acid is being made in glass (borosilicate at that), so this restriction would not apply.

For mine this is unlikely to be economical in the USA where you have H2SO4 on tap, but here, where it ain't so available (only real option is to order it in), a useful method is necessary and even metabisulfite is cheap in comparison. Given my location, it is fucking near impossible for me to access useful quantities of H2SO4, so the amount of effort, time and money is really not at issue. This looks to consume less energy than the multiple freeze cycles of the other method, so it may be a goer. ALthough, considering the increasing strength of the H2SO4 content does not seem to harm the outcome of the freeze-thaw approach, this might be conceivably run in tandem with it, thus converting the H2SO3 to H2SO4 as well.

But the point is, it is fucking difficult for me to access, thus it must be made. This approach is workable (for mine), unlike those using expensive metal or quartz materials, even more expensive oxidants and hyperexpensive catalysts. I also ain't a fan of using ultra-hard to access nitrates as the oxidant for this process. So my options are limited, as I'd suggest are the options of many others. If anyone has a good reason why the continued running of SO2/SO3 into dilute H2SO4 (and then H2SO3/H2SO4) would not work, let me hear 'em.

not_important - 11-4-2008 at 02:15

Doing some reading turns up references stating that as concentrations and acidity increase, SO2(aq)/H2SO3 switches from being a reducing agent to being an oxidiser, cpnverting Fe(II) to Fe(III), mercurous to mercuric, and even redoxing itself to give H2SO4 and sulfur.

This is suggestive that the oxidation of SO2 by O2 may slow down under the same conditions (but I've no direct documentation). If true this would limit the concentration of H2SO4 formed, at least from a practical standpoint.

It also suggests that rather than an absorption chamber it might be better to have a fractionating column with air and SO2 introduced at the bottom and water or dilute H2SO4 refluxing through the column. As H2SO4 was formed it would be carried down the column to collect in the still pot, unreacted SO2 being stripped out to mostly residing in the cool upper portion.

[Edited on 11-4-2008 by not_important]

LSD25 - 11-4-2008 at 04:07

I was actually considering whether it would be practicable - particularly if starting from alkali sulfites - to condense the off-gasses from the catalyst tube - collect the liquified SO3 (if it is kept dry) and then run the gaseous SO2 into an alkalli solution - thus giving back the starting material less that which was oxidised (well, except for the muriatic, but that is a whole 'nother story).

PS Doesn't SO2 liquify at about the same temp that SO3 solidifies?

not_important - 11-4-2008 at 07:20

Quote:

PS Doesn't SO2 liquify at about the same temp that SO3 solidifies?

SO2 BP : -10 C

SO3 MP: 16.8(gamma), 62.3(alpha), 32.5(beta)

both beta and alpha require traces of water to form.

497 - 11-4-2008 at 17:03

Quote:

Why not just run more SO3 into the liquid?

I think you are seriously overestimating the amount of SO2 that would be oxidized. You would need to run a setup like that for a *long* time to get any acid of useful quantity and in the process use *a lot* of fuel.

There is a reason the lead chamber process was used for many years. It worked. Unless you need oleum of course.

If I were you, this is what I'd do:

-Get a 55 gallon plastic barrel with a screw on lid, they're not too expensive here, you can probably get them
-Add a liter or two of water
-Hook a pump thats intake is at the bottom of the container, the outlet is a mister or spray nozzle near the top (to greatly speed absorption)
-Fill it with O2 (or air, it will be much less efficient)
-Light a can of 250g KNO3/S8 (1:7 ratio) and suspend it inside the container
-Let it burn out
-Start the spray pump and run it for a few hours (At this point the amount of gas in the container will drop, to keep it from imploding you could just allow air in, or preferably, refill it with a stoichiometric 2:1 SO2/O2 mix and keep running it until no more gas is absorbed.)
-Open up the lid, let it air out... poor environment..
-Boil down the acid, every 100g of sulfur burned should give at least 125ml concentrated acid.

The acid shouldn't end up too dilute, probably 10-30% unless you add way too much water. Acid of this concentration is not hard to boil down, I do it all the time. It takes at most a half a kilo of propane for a liter of 34% acid to be boiled to azeotropic concentration.

I feel your frustration, I am in a similar position with chemical availability, although not with sulfuric fortunately, many other chems are very hard to get here (Alaska) affordably because of shipping etc. Good luck.

[Edited on 11-4-2008 by 497]

Contrabasso - 12-4-2008 at 10:21

Given a 205 litre plastic drum, on it's side with a few litres (say 5) in the bottom, it should be possible to feed a lead fire basket with the KNO3/S mixture and blow a stream of air or oxygen in and have a steady supply of acid, especially if you can arrange to feed pressed pucks of fuel in occasionally through a small hatch.

The drum lid could be modded to have all the works affixed hearth, oxygen feed, fuel feed, water feed and acid draw off.

Even bigger use a large plastic header tank and a plastic lid. water in the bottom lead hearth there too, plastic lid on with tape to seal.

Formatik - 13-6-2008 at 16:39

Fe2O3 catalyst is actually not that bad to use, after platinum and vanadium it is one of the best oxides to use, being second best after chromium oxides. This is a table of effectiveness of catalysts from Gmelin:

for Fe2O3 at a temperature of 625º there is a 69.5% conversion rate at a flowing velocity of 150 cm3/min. Though platinum is clearly the best catalyst to use, asbestos containing 7% Pt has a working temperature of 425ºC and the conversion rate is 99.5% at a flowing velocity of 150 cm3/min. The lower temperature for the platinum is also most ideal for the equilibrium favoring the SO3, thus also the high conversion rate. Platinum catalyst can be made e.g. by absorbing hexachloroplatinic acid from aqua regia and platinum with water absorbed onto a porous substance like pumice, diatomaceous earth, or the like and then heating to glow.

A way to H2SO4 is from estimated amounts of H2O2 and SO2. H2O2 will oxidize SO2 even in the cold to form H2SO4. Sources of SO2 is by burning and roasting sulfides (pyrite (FeS2); sphalerite and ZnS; chalcopyrite (CuFeS2), galena (PbS), etc), sulfur, or decomposing sulfites, or thiosulfites with a dilute acid. As is mentioned in this Gmelin Handbuch, SO2 through light, heat and electricity forms H2SO4 and sulfur, in air its aqueous solutions are only slowly oxidized. In addition to H2O2, other oxidizing agents which oxidize SO2 are mentioned by the Gmelin: I2, Br2, Cl2, HClO, HNO3, metal salts like MnSO4, Hg(NO3)2, mercury salts, AuCl3, etc. But with concentrated H2SO4, H2O2 forms H2SO5 and H2S2O8, as described in Gmelin S[B], p. 777.

I’ve decided to try the above oxidation, but aborted the procedure because the reaction got too violent. I added 110 g of a mixture of sodium hyposulfite and metabisulfite to a 500 mL flask, the flask had a rubber stopper with a 50 mL separatory funnel and also a tube running out of it. The tube lead into 50 mL of 35% H2O2 in a 100 mL graduated cylinder. Then the separatory funnel filled with 16.8% pure HCl. The acid was then let drip in slowly and portionwise with occasional stirring.

At first the bubbling of SO2 proceeded smoothly for several minutes, and the reaction between H2O2 and SO2 is highly exothermic reaching around 105ºC at some points. Though after some volume reduction, at some point the SO2 generator did something unexpected, without any warning whatsoever e.g. effervescence, foaming, etc. as SO2 was bubbling into the H2O2, the stopper blew off violently from the flask and the tubing shot out of the graduated cylinder, the acid/peroxide mixture spattered all over even on my arms and over the gas mask. After washing off, I came back and tried an even slower addition, but even then the exact same thing happened. I thought maybe the acid was too strong and diluted it with around 2 times the volume with water. The same thing happened! So I halted the procedure.

The following is from the nice interesting experimentor chemistry book "Chemie selbst erlebt" by Erich Grosse. The Lead chamber process: 52. The contact process from pyrite: 53, 54. Acid from plaster: 55, and from kieserite mineral (MgSO4.H2O): 56, 57.

497 - 14-6-2008 at 14:35

While Fe2O3 might work well enough, why not use V2O5? Its not very hard to get nor very expensive.

I'm glad someone finally tried the H2O2 route, I keep reading that it would work but until now I've never seen an actual account of it being done. I'm not sure what is happening with your SO2 generator, but it is unfortunate that the experiment was never completed.

For me, simple production of sulfuric acid is of much less interest than production of high concentration (>97%) sulfuric acid or oleum. I can buy battery acid easily and cheaply, I have no use for dilute acid. This may not be the case for others in other countries, but I doubt there are many occasions that concentrated H2O2 is less valuable than dilute H2SO4.

So my dream is that someday (soon) I will be armed with a 3kw induction furnace and some V2O5 and be able to produce all the oleum I could ever need.

Formatik - 14-6-2008 at 20:43

Quote:

Originally posted by 497
While Fe2O3 might work well enough, why not use V2O5? Its not very hard to get nor very expensive.

Hadn't given vanadium too much thought. Do you know any good common sources?

Quote:

I'm glad someone finally tried the H2O2 route, I keep reading that it would work but until now I've never seen an actual account of it being done. I'm not sure what is happening with your SO2 generator, but it is unfortunate that the experiment was never completed.

I really have no explanation. I thought at first maybe some of the peroxide and acid mixture suctioned into acid sulfite solution through the tube, but later separatley adding a H2O2/H2SO4 mixture to an acid sulfite solution effervesces noticeably. The blow off occured everytime after only a portionwise addition of the hydrochloric acid from the funnel into the liquid sulfite mixture, even if just let sit. It's really strange because it wasn't reacting like this when copious amounts of SO2 gas were generated and bubbling in earlier.

Quote:

For me, simple production of sulfuric acid is of much less interest than production of high concentration (>97%) sulfuric acid or oleum. I can buy battery acid easily and cheaply, I have no use for dilute acid. This may not be the case for others in other countries, but I doubt there are many occasions that concentrated H2O2 is less valuable than dilute H2SO4.

Sulfuric acid is a universal chemical, so it may be simple to purchase. It has been in wide use for many years, and that will not change much, but its availability might. I'm interested in any dilution not too low, as H2SO4 is probably the easiest acid to concentrate.

Quote:

So my dream is that someday (soon) I will be armed with a 3kw induction furnace and some V2O5 and be able to produce all the oleum I could ever need.

If you're interested in oleum, then you can distill sulfates at strong heat to get SO3 directly. Namely, iron sulfates. These were used in the times of old to prepare oleum. One reason for use was low decomposition temperature, so any sulfate which has a low decomposition temperature and forms sulfur trioxide would be most suitable. For this reason, iron (II)- a.k.a. green vitriol and iron (III) sulfates. Iron (II) sulfate is less preferable to the higher oxidation compound as some SO2 is also produced: 2 FeSO4 = Fe2O3 + SO3 + SO2 compared to Fe2(SO4)3 = 3 SO3 + Fe2O3. The decomposition temperature of Fe2(SO4)3 is 480 deg.C. Also, from CuSO4 the last hydrate of 5 H2O is removed at 200º, and then anhydrous CuSO4 above 340ºC is said to decompose to CuO and SO3.

I have a method for a vanadium catalyst used for SO2 oxidation from an unknown reference. Pumice and V2O5 in mass ratio of 2:1 is made into a dough with water and then vacuum dried, then heated in a drying closet at 120º for 30 min, or to little pieces of pumice or asbestos fibers so much concentrated ammonium vanadate solution is added as much as can be absorbed, then it's dried and glowed weakly, or pieces of pumice or asbestos fibers are rolled and mixed around in V2O5 powder (this gives a lesser, but still good working catalyst). They say the favorable temperature is between 400 and 500ºC, but say it even starts working at 200ºC. However, from the temperature you need for Pt and V as catalysts, one can already more easily break down the sulfates.

[Edited on 14-6-2008 by Schockwave]

497 - 15-6-2008 at 04:43

Quote:

Hadn't given vanadium too much thought. Do you know any good common sources?

Good old United Nuclear sells it at a good price IIRC.

And yes while H2SO4 is easy to concentrate to 97% or so, if you need any more than that you're out of luck. Can you get dilute sulfuric easily? If you can then the ability to produce it is only useful if it is cheaper or if it becomes unavailable. I doubt it could be done much cheaper and while it is definately a useful capability if it does become unavailable, I seriously doubt it will in the near future. So dilute sulfuric (ie. lead chamber, H2O2 route, etc.) is of little use to me or most others as far as I can tell.

And as far as decomposition of sulfates, I was under the impression that it didn't work too well. I haven't looked into it too deeply though. The Fe2(SO4)3 route looks interesting. But my question is, if it is apparently so easy why is it not being done?

[Edited on 15-6-2008 by 497]

LSD25 - 15-6-2008 at 05:43

Why not just grab an old catalytic converter - the oxidising part thereof is a honeycomb of Pd/Pt on cerium and alumina. It should oxidise SO2 rapidly and quantitively (that is what it does to nitrogen oxides / carbon monoxide and what it was designed to do). They are comparatively high throughput and work at fairly low temperatures.

Len1 & Garage Chemist have done sterling work on the preparation of the same from sulfates, I think that is probably the way to go on a small-scale (especially for oleum).

Formatik - 15-6-2008 at 15:44

Quote:

Originally posted by 497 Good old United Nuclear sells it at a good price IIRC.

Alright thanks, it looks like they are temporarily sold out of it.

Quote:

And as far as decomposition of sulfates, I was under the impression that it didn't work too well. I haven't looked into it too deeply though. The Fe2(SO4)3 route looks interesting. But my question is, if it is apparently so easy why is it not being done?

For a hobbyist it is easier to put sulfates in some pipes or tubes and heat, than the catalytic set-up and preparation.

[Edited on 15-6-2008 by Schockwave]

497 - 16-6-2008 at 15:22

Now that I have researched the decomposition of sulfates more extensively it does seem to be viable. But, it does not appear to be as effective as you make it out to be. First off, the temperature required to decompose Fe2(SO4)3 at reasonable pace is more like 800*C rather than 480*C stated. At these higher temperatures at least 60-80% of the SO3 decomposes into SO2 and O2. From garage chemist's writeup on decomposing Fe2(SO4)3:

http://img295.imageshack.us/img295/5236/diagrammsg6.jpg

Also, the decomposition of sulfates seems to require a glass vessel that can withstand the high temperature and corrosive nature of SO3/H2O. From what I understand a steel pipe would not withstand that much abuse. I may be wrong.

Iron and sulfuric acid being cheap as they are, low yields shouldn't be a big problem. The containment on the other hand.. I don't happen to have a quarts flask, so if we can figure out an alternative, I might just have to give it a try.

@LSD25 - I would imagine a cat would work if you could manage to keep the whole thing hot enough. Also I wonder how well the honeycomb and housing would stand up to 500* SO3... It'd be worth a try especially if you wanted a larger quantity of SO3.

Formatik - 17-6-2008 at 04:41

SO3 will begin forming at below 500º from both iron (II) and iron (III) sulfates as mentioned in a dissertation document in this thread, but yes that could take a while and some patience for higher yields. Though aluminum sulfate is already said to decompose at 500º as much as over 80%, but its entire decomposition occurs above 800 deg. Mn, Co, and Cu sulfates are said by the same decompose about 650ºC. Ni and Zn sulfate begin 750ºC, and Mg sulfate begins at 850º. According to that information, so far aluminum sulfate seems the best choice for rapid high yield.

Conc. H2SO4 will not attack iron, but dilute acid will, so if the sulfates are distilled in iron, at best they should be made anhydrous before proceeding to a higher heat. Quartz and Vycor can handle higher heat. Borosilicate glass can handle lower temperature, around that for most of the aluminum sulfate.

[Edited on 17-6-2008 by Schockwave]

497 - 18-6-2008 at 18:49

Hmm I've seen in a patent that at least 900*C is required for "quick" decomposition of anhydrous aluminum sulfate. I'm not sure what they mean by quick, the main goal of the process in the patent is not to produce SO3. I don't know all the details, its a long patent and I don't have time to read the whole thing. Here it is:

http://www.google.com/patents?id=NZdKAAAAEBAJ&printsec=abstract&zoo...

Edit: Ok I think I've found a critical part of the process that has been omitted by the others who have attempted to decompose sulfates to get a substantial yield of SO3. US patent #2413492 stated that ferrous sulfate is completely decomposed at 560*C in a current of air. Later it goes on to say that a temperature of 700*C is optimum for speed. When garage chemist decomposed ferric sulfate he did not get rapid decomposition at 700*C. I think the key is the oxidation of FeSO4 to Fe2O(SO4)2 (basic sulfate).

The reaction stated in the patent goes as follows:

2FeSO4-H2O + O2 ---(167-455*C)--> Fe2O(SO4)2 + 2H2O

Fe2O(SO4)2 ---(492-560*C)--> Fe2O3 + 2SO3

Is this old news? Because it sheds a whole lot of light on things for me.

I think running a setup using a high temperature air current like in the patent might be a little challenging to build, but the yields would be so much higher and more importantly for me, the temperatures would be lower. The air current serves to slow that decomposition of SO3 and reduce its partial pressure in addition to oxidizing the FeSO4. I think it would be doable.

Here's the patent:

http://www.google.com/patents?id=hXtjAAAAEBAJ&printsec=abstract&zoo...

[Edited on 18-6-2008 by 497]

[Edited on 18-6-2008 by 497]

Formatik - 19-6-2008 at 21:50

It seems like we have come to some of the same information via different sources. Below are some scans from Gmelin. Mostly relevant information concerning Fe2(SO4)3: under air absence, SO3 tension is unnoticeable at 400º, at 500º it becomes measurably large. By little air ingression, heated in a tube closed at one end the decomposition temperature of Fe2(SO4)3 is 705º. In a stream of air, the decomposition of Fe2(SO4)3 begins at 550º; in a stream of 5% SO2 and 95% air, it is at 620º, in a N2 stream noticeable decomposition at 660º. So the high temperature garage chemist needed, likely was due to no air stream.

For FeSO4, it’s noted that the info from the literature concerning its course and temperature of the decomposition agree little, and that they are dependent on the conditions of the attempts. Though this should apply to most, if not all sulfates. In a glass tube with little air ingression, FeSO4 at 500 to 585º remains constant in mass for 10 to 20 minutes, at 590º begins decomposition, forming SO3, which decomposition at 625 to 635º in 2 hours is only 3%. By an unrestricted air ingression: FeSO4 barely oxidizes at 245º, but rapidly at 440º; dry and absolutely anhydrous salt decomposes over 300º very rapidly. In an open crucible between 300 and 535º increases in mass due to oxidation to ferric salt, though remains mass constant at 535º for several minutes, but over 535º it loses mass due to decomposition. In dry air stream: decomposition begins at 550º, at 580º this is only little stronger, but at 600º there is a sudden, rapid increase there. At temperatures up to 960º, no further decomposition is noted. According to Warlimont the decomposition begins at 470º. The roasting of FeSO4 in a dry air stream in 3 hours at e.g. 550º is 100% decomposition (see chart below on p. 399).

Concerning the Al2(SO4)3, there is also some variation here. In an air stream, decomposition begins at 590º, the other values not specific of a air conditions, are over that, up to 620º. The complete decomposition varies from 750º in a vacuum, to 770º in a one side closed pipe, to e.g. over 960º.

I’ve also looked at several other metal sulfates, decomposition temperatures of alkali (Na, K, Cs) and alkaline earths (Sr, Ba, etc) are of course ridiculously high, though the latter might be able to be lowered like with the CaSO4 by the addition of C. I think SnSO4 (Gmelin Sn 63) could serve well as an SO2 source at even 378º (but below this, decomposition is insignificant) it completely decomposes to form SnO2 and SO2. Though some other indications say 500 to 600º is needed for complete decomposition.

Gmelin (S [B] 356): completely dry SO3 will not attack Sn, Pb, Cu, Ag, Zn, Cd, Ni, Mg (not even powdered Mg, if it’s absolutely dry. SO3 will not further attack either Mg or Al after it has formed a layer on the metal). At regular temperatures, liquid SO3 or SO3 vapors will not attack Fe. A fine Fe wire when heated in a glass tube with liquid SO3 gets covered with a black layer, which when warmed with HCl solubilizes to a yellow color and gives H2S formation.

Fe2(SO4)3: I, II.
FeSO4: I, II.
Al2(SO4)3: I, II.

497 - 19-6-2008 at 23:33

I think iron is the way to go. If you didn't see in the patent, the whole purpose of decomposing the sulfate is to separate if from almost any other contamination, they all decompose higher, aluminum being the next lowest.

So my current process would be:

-Take FeSO4*xH2O and heat it to 500*C or so in a crucible with stirring until no more weight is gained
-Put the dry oxidized product into a steel tube and pump a CaCl2 then H2SO4 dried stream of preheated 500-600*C air through for an hour or two while bubbling the exiting air through more H2SO4 and then NaOH.

I like it. Yields *should* be above 70% I think. And 600*C or 700*C shouldn't be too hard with propane.

Also, while in an ideal world SO3 wouldn't attack said metals, what if there's just a tiny bit of water? That could cause some big problems... I suppose if the system was flushed with dry 600*C air for a while you wouldn't have to worry?

[Edited on 19-6-2008 by 497]

[Edited on 20-6-2008 by 497]

497 - 2-7-2008 at 18:36

I was just revisiting the old oleum/SO3 thread and found the discovery by garage chemist that HPO3 will dehydrate H2SO4 to SO3. Phosphoric acid is usually OTC and a copper crucible can be used, so HPO3 is easy. Also IIRC dehydrated boric acid could also dehydrate H2SO4. Why are these methods not used? It seems easy enough to me, and doesn't require too high temperatures.

Also, with the recent developments in the phosphorus thread, it looks like it would not be too hard to make a few hundred grams of white P and oxidize it to P2O5 and use that to dehydrate the H2SO4 at even lower temperatures. P is also useful for so many other things, armed with white P and SO3 what more could a person want?

Picric-A - 21-7-2008 at 09:58

are you sure phosphoric acid, HPO3, can dehydrate H2SO4? i thought only phosphorous pentoxide, P2O5, was capable of doing that?
If it can that should be an easy way to oleum =)

12AX7 - 21-7-2008 at 12:31

Read the metaphosphoric acid thread. Glassy fused HPO3 has been used to prepare SO3.

Tim

497 - 28-9-2008 at 13:10

Here's another patent on the decomposition of basic iron sulfate to SO3 and Fe2O3. It gives a little different reaction scheme than the other patent did. I tend to believe this patent as it is much newer.

4FeSO4 + H2SO4 + O2 --> Fe4O(SO4)5 + H2O

Then

Fe4O(SO4)5 --> 2Fe2O3 + 5SO3

The first step proceeds in a slurry of FeSO4 in semi-concentrated H2SO4 just below the boiling point of the mixture (which depends on water content of the H2SO4).
The second step occurs at 500-700*C, apparently without air.

So...

Mix 600g anhydrous FeSO4 (or equivalent hydrated) with 300g 33% H2SO4 battery acid (or equivalent concentrated) and heat with stirring to over 150*C in a stream of air, maybe with a heat gun. It should eventually harden into a solid cake of about 700g after all the water is driven off. Crush up the cake and load it into a makeshift retort and heat to ~600*C for a while. Yield: 300-350g SO3!

Sounds pretty good to me. Anything wrong with that process?

DJF90 - 28-9-2008 at 13:32

It looks too good to be true, but I guess the only way to find out is to try it

Picric-A - 6-10-2008 at 10:25

Are you sure the ferrous sulphate can be oxidised that easily?
I would of thought it would take hours to oxidise it simply with air...
If not it is an extremly easy way to lots of SO3

Since making that post i have found out that the oxidation of FeSO4 occurs rapidly at high Ph so Reacting FeSO4 with hot H2SO4 whilst bubbling air through it should oxidise it pretty quick!

[Edited on 6-10-2008 by Picric-A]

not_important - 6-10-2008 at 17:23

High pH means neutral to alkaline, not acidic. Increased temperature will speed the reaction, though.

Picric-A - 6-10-2008 at 23:09

Of course, sorry for that stupid mistake. Too tired last night

Rosco Bodine - 27-10-2008 at 00:11

For the precursor desired for the pyrolysis ...

I am wondering if perhaps a synthesis of ferric sulfate
from ferrous sulfate via H2O2 and then partial hydrolysis,
then dehydration of precipitated copiapite may also work.

4 FeSO4 + 2 H2O2 + 2 H2SO4 ---> 2 Fe2(SO4)3 + 4 H2O

Fe2(SO4)3 + 2 H2O ----> 2 Fe(OH)SO4 + H2SO4

Since half the H2SO4 required for the first reaction is regenerated via hydrolysis, the algebraic sum would possibly adjust the initial reaction minimum H2SO4 stoichiometric requirement to 1 H2SO4

2 Fe(OH)SO4 + Fe2(SO4)3 + 17 H2O ----> Fe4(OH)2(SO4)5-17 H2O ( copiapite precipitate )

Fe4(OH)2(SO4)5 - 17 H2O -----> Fe4O(SO4)5 + 18 H2O

H2SO4 from ozone

Formatik - 28-10-2008 at 15:34

I was reading the wiki entry on O3 and it mentioned the following reaction between elemental sulfur and ozone to form SA:

S + H2O + O3 = H2SO4

But no details. Does anyone know more about this like reaction time and conditions?

12AX7 - 28-10-2008 at 20:36

Well, that's trivial, but ozone is rather harder to generate in quantity than SO2, or SO3 for that matter. And one could argue the SO3 is safer, pound for pound, than that much ozone.

Tim

Rosco Bodine - 28-10-2008 at 22:09

Using H2O2 instead of air oxidation for the process described
by 497 above ....

I'm still wondering if the following summary reaction derived from my reactions above wouldn't be the easiest way to form a precursor for pyrolysis to SO3.

4 FeSO4 + 2 H2O2 + H2SO4 + 15 H2O ---> Fe4(OH)2(SO4)5-17 H2O ( copiapite precipitate )

Fe4(OH)2(SO4)5 - 17 H2O -----> Fe4O(SO4)5 + 18 H2O

Fe4O(SO4)5 ----> 2 Fe2O3 + 5 SO3

[Edited on 29-10-2008 by Rosco Bodine]

497 - 29-10-2008 at 18:17

That would be nice if that worked. I don't see why it wouldn't. According wikipedia it can be prepared with nitric acid as the oxidizer. They call it ferric subsulfate or basic ferric sulfate. It is apperently used as some sort of medical treatment for certain skin problems.. I'll do some more looking around.

Rosco Bodine - 29-10-2008 at 23:35

Yeah this sort of reminds me of the line of thinking which I had going when contemplating a pyrolysis precursor for
calcium cyanamide. If you can get right to the immediate precursor via some preliminary workup which eliminates
rotary kilns and other steps which are more convenient for
industry ....then you are a lot closer to a worthwhile lab scale method. The temperature and pH will probably affect the density of the precipitated copiapite ...and the only concern I have there is possible gelling .....but boiling and agitation would probably break that higher hydrate.
Sometimes those "superhydrates" are unstable transition
species which form nicely crystalline lower hydrates or
even anhydrous derivatives....and I agree it would be nice if this one behaves in that way. It would probably be the
easiest route to SO3 and oleum which has been proposed so far in any discussions here. So then SO3 would be
"pyrolitic distillate of copiapite anhydride"

[Edited on 30-10-2008 by Rosco Bodine]

497 - 29-10-2008 at 23:53

Well I found lots of stuff on using H2O2 to oxidize ferrous to ferric sulfate, but so far not much on making basic sulfate. Since O2 can do it under the right conditions, I see no reason H2O2 couldn't. It doesn't look like gels would be a problem, so far all everything I've found has said "crystalline" I haven't had a lot of time to work on this lately.. I'll try to post more soon.

US 2563623 is kind of interesting..
US 3529957 might be useful

Edit:
According to this, H2O2 will in fact work. They used it as a control for studying some bacterial oxidation crap, but it should still be useful data. Apparently the precipitate will have a greatly varying Fe:SO4 ratio depending on various conditions. At pH 2.5 (with 2%H2O2!) they got a ratio of 6:1 Fe:SO4, which is far from ideal.. But I think at a lower pH and more concentrated H2O2 (and maybe some other additives), the ratio will get much better. The decomposition should go similarly at different ratios right? It definately requires some testing though. If you wanted you might even be able to optimize the ratio to get the lowest decomposition point, I don't know.

Another idea that's probably less practical, but still interesting: an acid solution of FeSO4 is known to reduce N2O4 to NO and form Fe2(SO4)3. I wonder if the conditions were right it could be oxidized further to a basic sulfate (or subsulfate, oxysulfate, hydroxysulfate, or whatever else you want to call it)? It would be an interesting process because the NO could easily be recycled.

The other thing one could work on would be figuring out an easy effective high temperature air oxidation route that could be used instead. Still it's hard to see how that could be easier that mixing a couple solutions and having your precourser all ready to go..

[Edited on 30-10-2008 by 497]

Rosco Bodine - 30-10-2008 at 01:35

The reactions I listed are valid and previously reported reactions, so I was already confident about the reactions before proposing them . The significant unknowns I have there are with regards to the rate of reaction, the most favorable pH , temperature and the physical form of the end product. It's like... this should work, but how well it will work I'm not sure, but it would be worth an experiment.
I think the concentration of the reactants would probably
be dilute and the solutions would be hot and perhaps even brought to boiling, basically an open simmering cauldron
reaction where the H2O2 simply accellerates what air would do given more time.

[Edited on 30-10-2008 by Rosco Bodine]

497 - 30-10-2008 at 01:44

According to the paper I linked the reaction time is apparently pretty fast.. But as far as optimum temp, pH, etc, that is anybodies guess. I seriously doubt there is much if any available information on these specific conditions. So experimentation is the name of the game as far as I can tell. And I certainly do think it is worth experimenting with, as you said before, this could be the easiest route to SO3 yet. And I really would like to have some SO3!

here's a couple of patents

Rosco Bodine - 30-10-2008 at 01:53

US3078180 H2O2 oxidation of Ferrous Sulfate

Attachment: US3078180 H2O2 oxidation of Ferrous Sulfate.pdf (356kB)
This file has been downloaded 109 times

and the second patent

Rosco Bodine - 30-10-2008 at 01:57

US3574599 Copiapite Basic Ferric Sulfate

A couple more of interest are

US2905533 Basic Ferric sulfate

US2413492 Purification Crystallization of Ferrous Sulfate

Attachment: US3574599 Copiapite Basic Ferric Sulfate.pdf (256kB)
This file has been downloaded 121 times

Rosco Bodine - 30-10-2008 at 02:01

Looks like I should just post those other two patents also
so here they are

Attachment: US2905533 Basic Ferric sulfate.pdf (184kB)
This file has been downloaded 153 times

Rosco Bodine - 30-10-2008 at 02:02

and here's the last one

Attachment: US2413492 Purification Crystallization of Ferrous Sulfate.pdf (249kB)
This file has been downloaded 178 times

497 - 30-10-2008 at 02:08

I saw all of those, except for the one about enameling, when I was searching earlier.. Some useful information but I didn't see anything too applicable or specific, at least in terms of synthesis of the basic sulfate via H2O2.

Edit;
Are there really that many people watching this thread? Each attachment you posted was downloaded about 10 times within a minute. Strange..

Edit2:
I just remembered, I neglected to post any info on an interesting patent I saw. It talked about using NaClO3 with iron sulfate to produce mixed oxysulfate salts. I didn't read the whole thing and I don't know how the decomposition properties are, but they looked interesting. Well, not too interesting to me since I have no easy source of chlorate, but maybe for someone else.

here it is.. ugh I'm accumulating such a huge mass of PDFs.. And they're all so disorganized, I feel sad just thinking about trying to organize them all.

[Edited on 30-10-2008 by 497]

Attachment: PRODUCTION_OF_IRON_OXIDE_AND_IRON_FREE_O-1.pdf (177kB)
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Rosco Bodine - 30-10-2008 at 02:10

You kind of have to pick the pieces parts from each one
and then interpolate

Yeah I saw all that interest and figured I better give it up

It looks to me like there would need to be some additional
H2SO4 in that sodium chlorate oxidation.....and it also
seems that reaction is adaptable to the use of ordinary bleach instead of sodium chlorate. Of course you end up
with salt solution as a byproduct.

[Edited on 30-10-2008 by Rosco Bodine]

497 - 30-10-2008 at 02:33

And here's another patent. I'm not sure how useful it is, but it does talk a little about alternative oxidizers.. Maybe ammonium persulfate would be useful? Or various peroxyhydrates?

Anyway, it was nice corresponding semi-instantly with you Rosco, but alas it is 2:30 AM and I am very tired. Talk to you later.

Yes bleach might be a good idea, have to try that. As long as you could do it without getting too much sodium caught along with the Fe. I considered that a while ago but dismissed it for some reason, can't remember why...

[Edited on 30-10-2008 by 497]

[Edited on 30-10-2008 by 497]

Attachment: Process_of_preparing_a_preferred_ferric_.pdf (228kB)
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497 - 30-10-2008 at 13:46

Would this be the reaction?

4FeSO4 + 2NaOCl + 3H2SO4 = Fe4(OH)2(SO4)5 + 2NaHSO4 + 2HCl

I wonder if something would have to be added to neutralize the HCl?

Edit:
I just found some patents on ferrate synthesis.. interesting stuff. FeO4-- is supposed to be a stronger oxidizer than MnO4--. It is said to be produced by oxidizing Fe3+ salts with concentrated hypochlorite. It releases O2 in acid solutions, maybe you could use it to synth basic sulfate? It looks useful for many other things too. Maybe via this:

3FeSO4 + Na2FeO4 + 3H2SO4 --> Fe4(OH)2(SO4)5 + Na2SO4 + 2H2O

Then you wouldn't have to deal with chlorine and other crap in there. It should be a strong enough oxidizer at least..

[Edited on 30-10-2008 by 497]

Rosco Bodine - 30-10-2008 at 19:03

I'll have to check further but I would expect that the
basic ferric sulfate is going to require moderately basic to
neutral or only very slightly acidic conditions to form ....
maybe pH 10 to pH 6.5 for example.

Your first equation proposed above would be too acidic.

You are going to have intermediate hydrates and hydrolysis
reactions to consider ....so it is going to be an algebraic
stoichiometry. I'll have to work it out later for the bleach.

There was a method I think in one of the lead oxide related
threads where bleach was used to precipitate lead oxide from lead salt solutions and the method for iron may be similar.

[Edited on 30-10-2008 by Rosco Bodine]

497 - 30-10-2008 at 20:35

Quote:

would expect that the basic ferric sulfate is going to require moderately basic to neutral or only very slightly acidic conditions to form .... maybe pH 10 to pH 6.5 for example.

I don't know about that, according to that paper I linked to, they got a precipitate that was like 99:1 Fe:SO4 when precipitated at pH 9, a 20:1 ratio at 6, and a 6:1 ratio at 2.5 (with H2O2).So you'd have to have it below pH 1 probably, to get any decent ratio? Unless there's some other factors that could be changed? I wonder how different oxidizers effect the precipitate differently?

Here's the important part of the paper I was talking about..

[Edited on 30-10-2008 by 497]

Attachment: precipitated iron.doc (94kB)
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Rosco Bodine - 30-10-2008 at 23:29

I am just looking at this intuitively and making an educated guess. We are intending making ferric sulfate and then allowing it to hydrolyze 50% which results in an addition compound between that 50% which is hydrolyzed
and that other 50% which is not hydrolyzed. Too acidic conditions will be stabilizing against the desired hydrolysis and prevent precipitation of the desired addition compound. Read the copiapite related patent US3574599, column 3, line 66.

[Edited on 31-10-2008 by Rosco Bodine]

497 - 31-10-2008 at 00:00

You might be right, have a look at this. It gives specifics on the preparation of basic ferric sulfate of the formula 3Fe2O3*4SO3*9H2O on page 446. Pretty much they heated a neutral 0.125 M Ferric sulfate solution to 140*C for a couple hours, or longer at a lower temperature. The formula is not quite what we want, but it's close, maybe running the hydrolysis in a slightly acid solution (maybe more concentrated?) would do the trick.

Edit: I found another reference to hydrolysis, I quote:

"The ferric sulfate is hydrolyzed to basic ferric sulfates,
the ratio of iron, hydroxyl, and sulfate depending upon
the dilution and acidity during hydrolysis. The manner
of hydrolysis is represented by the reaction:
Fe2(SO4)3 + 2H20 --> 2Fe(OH)S04 + H2SO4
Actually, the hydrolysis may proceed until practically
complete with formation of ferric hydroxide, Fe(OH)3.
The buffer capacity of the streams has an additional
determining effect upon the extent of hydrolysis. In
the absence of acid, ferrous sulfate, also, is oxidized
to basic ferric sulfate:
4FeSO4 + 02 + 2H20 --> 4Fe(OH)SO4"

So I guess it comes down to fine tuning the pH, temperature, and concentration to give the best formula. Do we even know what the best formula is? What we really need is a graph of decomposition temperature versus Fe:SO4 ratio.. And while I'm wishing for things, a graph of Fe:SO4 ratio versus pH, temp, and conc. would be great

[Edited on 30-10-2008 by 497]

Attachment: AM51_443.pdf (663kB)
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497 - 31-10-2008 at 00:34

Another good reference. It's a kinetics study of precipitation of various iron/sulfate ratios. All the tests were done with very dilute solutions but I think it still has some valuable information..

Attachment: iron precip kinetics.pdf (824kB)
This file has been downloaded 168 times

Rosco Bodine - 31-10-2008 at 00:48

I'm not sure what would be the optimum conditions but it looks like it is similar to the reactions of Bismuth and that
the ferric salt will very easily hydrolyze even in strongly acidic conditions particularly at elevated temperatures,
going by that autoclave reaction of the neutral ferric sulfate. It looks like the copiapite may be a no go as an open beaker reaction if it requires geological pressures and temperatures for its formation in fairly extreme acid condition. That paper or the patent could either one or both be wrong however ......I'm just not sure on this one . The water absortion for the hydrate formation should be a good indication for the copiapite if and when it is achieved. There may be a specific dilution and pH
where that is the only product or it may not happen except
under extreme conditions.....I just don't know ....and I think I qualified that earlier as a possible impediment .
I regarded that copiapite as a probable gell and unstable intermediate. This is one of those contemplated reaction schemes which would definitely require experiments,
unless the unknown process variables are already published somewhere. Yeah a graph would be nice

497 - 31-10-2008 at 01:11

Maybe if someone could get this paper it could shed some light on things.

I definately wouldn't want to have have to use autoclave conditions, so if it doesn't work at or below 100*C, I don't think it would be worth the effort. Basically all the information I've found so far has been focused on industrial scale stuff, I don't think there is much info on less cost effective reactions that may be most suitable for us. I really like the idea of using H2O2, I think it is most promising. I would imagine that mixing a concentrated boiling solution of ferric sulfate with hot/boiling (maybe acidified to pH 1-3) 3% H2O2 would give a usable product. Just have to try it I guess.. And hope it doesn't explosively decompose the H2O2...

I'm still curious as to how much the Fe:SO4 ratio affects decomposition..

Rosco Bodine - 31-10-2008 at 07:27

I get too tired to think straight sometimes. And I'm sure it shows. What I was hoping was that the copiapite would be the first thing to drop out as a stable precipitate as a
solubility limited reaction. The autoclaved neutral solution of ferric sulfate is a rather extreme hydrolysis
giving a more basic product. So it could very well be
that a copiapite precipitate appears as an intermediate
and then in further reaction with the superheated water
in the autoclave, the product which we would want instead of simply being dehydrated is further hydrolyzed
with the loss of SO4 which we would prefer to keep.

They were on the right track but went too far with the hydrolysis reaction .

So, if for example we were to simply mix the correct
concentration and pH of precursor solutions.....it may well be that the conditions are favorable for the copiapite precipitate to be the principal product. It may be straightforward, and all of this controversy is an imagined potential problem that doesn't exist.

Magnetite can be made in an open beaker at 75C, so it would seem likely that copiapite should be doable.
And I could be wrong, but I just don't see an active metal
like iron giving up more sulfuric acid on the loss of that
water of hydration on drying. Iron would seem less inclined to further hydrolysis particularly after having dropped out of solution.

I don't recall that autoclave being transparent but was described as "glass lined" so their visual
observation only of an end product has not ruled out that
there was copiapite there as an unobserved intermediate which was not harvested, but was further destructively hydrolyzed through subsequent products before the autoclave was opened to see what was the end product.

They hardboiled and pressure cooked the ferric egg until they ended up with a more modified material than the copiapite which we want. The autoclave literally water leached the H2SO4 life right out of copiapite intermediate
and converted it to a more hydrolyzed and more basic
product. So it is milder conditions which we want in terms of temperature and water.

The copiapite related patent indicates that pH is controlling
and that tracks with what I am thinking. So this still looks possible as an open beaker ...or bucket reaction.

[Edited on 31-10-2008 by Rosco Bodine]

Fe4O(SO4)5 Monsel's Salt, Powder Styptic U.S.P.

Rosco Bodine - 31-10-2008 at 11:17

I was thinking more on this Fe4O(SO4)5 and it seemed familiar and it should to every man who shaves carelessly sometimes and nicks the skin , ..ouch .

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

Attached is the file for the pharmaceutical preparation

So indeed it can be made under ordinary conditions.
This process is likely also possible using different reagents
which may produce a similar reaction condition.

A solution of this Fe4O(SO4)5 is called Monsel's Solution
and the crystals obtained from cooling or evaporation are Monsel's Salt .

[Edited on 31-10-2008 by Rosco Bodine]

Attachment: Basic Ferric Sulfate preparation Principles_of_Pharmacy.pdf (197kB)
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a larger scale preparation

Rosco Bodine - 31-10-2008 at 12:06

Here is a patent method which shows a
larger scale preparation. See example 1 .

US50111693 ( attached file )

Searching for Monsel's Solution or Monsel's Salt
may bring up alternate methods.

Here's another excerpt from a medical chemistry reference
which tends to support my original idea that supposed the
copiapite intermediate, and this reaction may very well work with H2O2 in the same way as it works with HNO3

http://www.sciencemadness.org/scipics/Ferric%20Sulfate%20and%20Mons...

[Edited on 1-11-2008 by Rosco Bodine]

Attachment: US5011693 Preparation of Monsels Solution.pdf (103kB)
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497 - 31-10-2008 at 15:36

Good information. Now the question is, will H2O2 or some other oxidizer be a suitable substitute? I really don't want to have to deal with nitric acid if I can possibly avoid it, so I hope H2O2 will work. Although, it looks like the amount of HNO3 needed is relatively small and not concentrated, so could probably deal with that. And if it requires more concentrated H2O2 to work, that's fine too, I have about 3 gallons of 50%. Others might have a harder time getting high concentrated H2O2 though...

This looks very promising.

Another problem I need to figure out is how to get substantial amounts of FeSO4.. Of course you can make it by reacting H2SO4 with Fe or Fe2O3, but I wonder if there's a better way? I guess I'd like to avoid using my sulfuric acid if I can.. I'll have to look in the garden store to see if they have it.

Edit:
On the internet it sells for about $5 per 4 pound bag. Not bad.
Or 50 pounds for $23, that's about 80 mols. Even better. That's a lot of SO3!

[Edited on 31-10-2008 by 497]

Rosco Bodine - 31-10-2008 at 15:53

Yeah I think H2O2 will work like I was originally proposing, and the copiapite doesn't or may not precipitate but is a transitional or theoretical intermediate. In either case,
it seems likely to work using H2O2 whether the copiapite precipitates and must be dried or whether the end result
is a solution of Monsel's Salt , the same amount of water
will have to be evaporated away and ultimately the
anhydrous Monsel's Salt is the expected end product.

Basically you do the same process to get Monsel's Salt
as you would do to convert Ferrous Sulfate to Ferric Sulfate only you use one-half the amount of added H2SO4.

Copperas (ferrous sulfate) is a common garden fertilizer .

Iron filings are collected by the bucketful at garages which
turn brake drums and rotors. Battery eletrolyte is sold in five gallon poly bags in a heavy cardboard carton having a rubber dispensing hose. 27% H2O2 is sold by the gallon
as a spa and pool substitute for chlorine.

Rosco Bodine - 31-10-2008 at 16:07

Actually the oxidation of Ferrous Sulfate to Ferric Sulfate
should go okay using an aerator in the acidified solution
of Ferrous Sulfate , probably some heating required also,
and the reaction should proceed fine just more slowly
than using H2O2 as the oxygen source.

497 - 31-10-2008 at 19:52

So *theoretically* it takes 98g H2SO4 + 68g H2O2 + 607g FeSO4 to make 720g Fe4O(SO4)5 which in turn makes 400g SO3 + 320g Fe2O3. Not bad at all, I like how little H2O2 it requires.

That comes out to
240ml 33% H2SO4 battery acid ($0.25)
180ml 35% H2O2 ($5) or 2.3 liters 3% H2O2 ($2)
733g Copperas ($1)
Some propane (>$2)
Depending on your price of copperas it might be slightly cheaper to recycle the Fe2O3 with battery acid.

I don't know about you, but I sure wouldn't mind having some $25/kg SO3...

I wonder what concentrations would be best? In the patent they use about a liter per kilo copperas, so I suppose that would be the place to start? If 3% H2O2 was too dilute you might be able to bring it up to 10 or 20% by freezing out some of the water..

Alternatively you might be able to use 800ml battery acid + 400g calcium nitrate fertilizer, filter off CaSO4 and use in a similar manner to the patents. It might even be a little cheaper. But then you have the additional problems of dealing with the N2O4 and filtration of that damn CaSO4...

[Edited on 31-10-2008 by 497]

497 - 1-11-2008 at 23:02

This has some really great information on basic iron sulfates. Much more detail than I've seen anywhere else. Pages 1965 through 1981 are useful.

According to the above book there are only three distinct basic sulfate salts. They are Fe2O3:SO3 ratios of 1:2, 3:4, 2:5.

Woops, here it is

http://books.google.com/books?id=4cAGAAAAYAAJ&pg=PA1965&dq=ferric+s...

[Edited on 1-11-2008 by 497]

Rosco Bodine - 2-11-2008 at 01:25

I think the reactions I wrote before are probably correct.
I notice the article shows a previous report of 1 hydroxyl on the copiapite when there should be 2 hydroxyls for stoichiometric balance there. Maybe a typo. I notice
that they couldn't seem to be certain about the analysis
of what was water of crystallization apart from the
water bound as hydroxyl. Is it 16 + 1 or is it 17 + 1 ....
it won't matter either way when the dehydrated material
is the Monsel Salt

The journal article is describing crystalline hydrates and they indicate that the copiapite is unstable above 90C. So the dehydration to the Monsel Salt should be done above 90C.

[Edited on 2-11-2008 by Rosco Bodine]

497 - 2-11-2008 at 01:56

Supposedly concentrated H2SO4 and SO3 don't attack iron metal... Is this true in practice? Because it would really ruin my day if I built a retort out of steel and suddenly ended up with a hole in it and a hundred grams of SO3 and/or red hot sulfate on the floor.

But really, this seems too easy. Not long ago I would have never imagined being able to get my hinds on something like SO3/oleum. It just seems like there must be something we're neglecting that will screw up the whole process... At least that always seems to be how things end up for me

. But I might just have to give it a try.. I think I can construct a steel retort without too much trouble. I suppose it would be good to start small and run a small batch to make sure it will work first. Then I'll go industrial scale!

Just kidding... I don't even know what I could use that much SO3 for anyway.. Not to mention the danger in dealing with much of it.

Edit:
I realized recycling the Fe2O3 byproduct with battery acid would give you ferric sulfate rather than ferrous. I'm still not clear on whether you can get the copiapate product directly from Fe2(SO4)3 without and oxidizer.. I know the stoiciometry of the reaction doesn't need oxygen or sulfuric acid, but I wonder how it would proceed, because all the preparations talk about using ferrous sulfate... would just be a matter of hydrolyzing ferric sulfate and getting the temperature and concentration optimized to give the copiapate? It would be nice not to have to use H2O2 after the initial batch.

Edit2:
Today I looked around for ferrous sulfate. Lowes has a granular moss killer that says "10% iron" and is about 35% "ferrous sulfate monohydrate" along with a bunch of unknown "inert ingredients". 3 pounds for $8. Not a good deal in my opinion, especially since it would require purification. They also had 99% zinc sulfate hydrate at $3/lb..

Then I checked at the local feed store. They had it pure in 50 pound bags for $60... I don't really want 50 pounds of it, and I don't really want to pay $60 but I suppose there's not really another option..

[Edited on 2-11-2008 by 497]

Rosco Bodine - 2-11-2008 at 17:32

It is available in 4 or 5 lb bags for a few dollars.

I have a 4 lb bag of Hi-Yield brand Copperas which states
analysis of 11% sulfur as combined sulfur and
19% iron derived from ferrrous sulfate.

I have seen that brand and others at various garden centers and feed stores, and Ace hardware I think has it too in a different brand.

Would have to do the math to figure out what is the level of hydration there corresponding to that analysis.

[Edited on 2-11-2008 by Rosco Bodine]

497 - 2-11-2008 at 18:29

Sadly in Alaska many things that are available to everyone else are not available.. The local Ace might have had it but they just went out of business (Lowes and Home Depot moved in). And the feed store I went to is probably the only one within a couple hundred miles...

The only other place I need to check is the local greenhouses, but I'm afraid most if not all of them have closed for the winter..

497 - 12-11-2008 at 22:53

I just found a very useful looking page on microwave casting of metals. If it can melt stirling silver it sure as hell can make SO3!

http://home.c2i.net/metaphor/mvpage.html

Apperently the use of a Fe3O4 + C powder mixture that is applied as a sort of stucco works quite well. Melted 50g Ag in 15 minutes (at 850 watts). It is capable of heating up to the melting or iron, but not much higher because the absorber couldn't handle it.

If one were to make a bunch of Fe3O4 + C loaded ceramic beads (or whatever shape) and mix them in with some Fe4O(SO4)5 in a glass or ceramic flask (that could also have the surface coated with the same mix), I think the results could be quite nice.

Saerynide - 10-12-2008 at 01:31

Woah! That is SO cool.... though I'm sure my housemates would not be very fond of the idea of me casting bars in their microwave essential to making hot pockets

Lead chamber volume

Contrabasso - 4-3-2009 at 10:18

Looking at developing Axehandle's work, What is a reasonable volume of "lead" chamber to use in a home setup. I do actually want to make a few litres of conc H2SO4. I reckon about a Kilo of sulphur with 150 - 200 g of nitrate should yield about 2.5 kilos of conc acid.

First I thought of a 20ish litre plastic bottle, then I thought 100litre plastic dustbin.

Has anyone actually scaled Axehandle's work up to production size yet? What size chamber did you use?

Mossydie - 1-6-2009 at 13:56

I don't know about production size, but the perfect chamber would be, I think, a 25l glass fish tank or bigger if you wanted to make lots - it's def. going to be waterproof and a lid could be made by simply putting a piece of glass sheet on top. A hole could then me made in this and a tube put through to stop the chamber collapsing. A clay pot could be used for the burn.

Fish tanks come up often on freecycle, so as soon as I can get my chemical proof gloved hands on one I'll post results.
How many burns would one carry out to get conc. acid - or would the resulting acid require boiling down?

Mossydie - 9-6-2009 at 10:39

Just set off my chamber today, and hopefully it all goes well and I can take in a few hundred ml of conc acid before my chemistry exam and prove my teacher wrong

.
A side note:
http://www.gardendirect.co.uk/sulphur-powder-p-750 - a site courtesy of PhZero, 25kg of pure sulphur for £60
Then with 3-4kg of KNO3 - a few pounds
and two suitable vessels, one for the process and one for storage, one has 150kg of H2SO4 because the rest of the chemicals are free from nature (with a little purification of course).
Of course, the result would not be reagent grade and energy and time would also be used up but in raw materials, assuming very little was lost, that's 150kg H2SO4 for about £70. Is there a market for such an item as kg bottles of non reagent 98% H2SO4 (that wont get one raided in minutes)?

[Edited on 9-6-2009 by Mossydie]

entropy51 - 9-6-2009 at 14:24

Mossydie, I hope you're joking about selling kg bottles of 98% acid made using the chamber process.

Looking at this thread, if anybody has made any sulfuric acid at all, I can't tell.

It's not an easy process. You can make a few mL of dilute acid as an experiment. To make signifcant amounts requires an industrial scale.

And the chamber process doesn't make 98% acid. About 64% acid is as good as it gets, as you'd know if you read some of those chemistry books I keep nagging you about. In the old days, they boiled and distilled the chamber acid to make 98%.

You should try it as an experiment in a 4 Liter or so container, but don't waste too much time and materials trying to go commercial.

Mossydie - 9-6-2009 at 14:57

Of course. But dilute acid can be concentrated with ease.

I was taught about the various processes of sulphuric acid manufacture in my GCSE course actually but also that one could never make any at home. This link (courtesy of you!) about 'Dangerous ACIDS MADE SAFELY BY Home Chemist ' describes the use of ferric oxide as a catalyst in something similar to the contact process (well, it would be if the resulting SO3 were added to H2SO4 instead of H2O):
http://blog.modernmechanix.com/2008/03/05/dangerous-acids-made-safe...
Another method would be to bubble SO2 through H2O2, I might try that out if I can set up an apparatus for it.

The links / idea were for general interest - I don't have £60 to waste on 25kg of Sulphur! And I was sort of impressed by the idea that so much could be made so cheaply and I got carried away with the idea. I also know that industrially the acid costs less than water...

However, if one were to find a suitable vessel and was not bothered about tiny levels of impurities then they could use this as a good source of sulphuric acid (it's cheaper than electrolyte for sure)

jgourlay - 17-6-2009 at 10:45

High, kindergartner walking in amongst the Ph.D.'s here! Would feeding sulphur dioxide + oxygen through an automotive catalytic converter give you what you want?

Mossydie - 17-6-2009 at 16:04

Catalysts are reaction specific, i.e. what used as a catalyst in a catalytic converter (platinum) for the reaction of CO with O2 to form CO2 will not necessarily be the same as that used for the reaction with 2SO2+O2 - > 2SO3. I don't think (but I may be wrong) that the catalyst in the converter will work for this reaction.

entropy51 - 17-6-2009 at 16:18

Indeed you may be wrong.

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

See http://pubs.acs.org/doi/abs/10.1021/i200033a031

DJF90 - 17-6-2009 at 17:33

I have a book that has a lab scale contact process in it, using platinised something or another. Prohibitively expensive, but just shows that platinum will catalyse the oxidation of SO2.

Picric-A - 26-8-2009 at 03:13

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

entropy51 - 26-8-2009 at 05:37

Quote: Originally posted by Picric-A

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

Instead of posting a brain fart, why don't you make a whole bunch of concentrated H2SO4 and then tell us about it.

hissingnoise - 26-8-2009 at 06:09

Quote: Originally posted by DJF90

I have a book that has a lab scale contact process in it, using platinised something or another.

Platinised asbestos was used in those old processes but fairly pure SO2 was required to minimise catalyst-poisoning.
Chloroplatinic acid, reduced, supplied the finely divided Pt.

DJF90 - 26-8-2009 at 07:10

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

entropy51 - 26-8-2009 at 08:55

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

Vanadium pentoxide catalyst is much less subject to poisoning.

watson.fawkes - 27-8-2009 at 08:19

Quote: Originally posted by entropy51

Vanadium pentoxide catalyst is much less subject to poisoning.

From talking to vendors, it seems that the main poison for vanadium oxide catalysts is arsenic, which can be a problem when converting off-gas from smelting sulfides.

Sulphur burner

Contrabasso - 11-10-2009 at 03:20

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

watson.fawkes - 11-10-2009 at 06:30

Quote: Originally posted by Contrabasso

Could a one shot process be controlled to be safe, stable and efficient at producing SO3 directly?

In industry, at the start of the campaign, the catalyst bed often receives supplementary heat to get the SO2 -> SO3 oxidation going. Since that oxidation is exothermic, once it gets going it's self-heating, to the point that it later requires external cooling, which they do by using it as a source of process steam.

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

S.C. Wack - 11-10-2009 at 16:19

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

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

Attachment: JCE1930p1668.pdf (1.9MB)
This file has been downloaded 170 times

Jor - 27-12-2009 at 17:22

How about this:

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

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

bbartlog - 27-12-2009 at 18:14

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

Alexein - 28-12-2009 at 13:45

Quote: Originally posted by Jor

How about this:

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

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

NurdRage is amusing but shouldn't be followed, he's a mediocore chemist at best and will lead you entirely down the wrong way of doing things. He's got little regard for practicality or cost and sometimes he's blatantly wrong. He's a clown.

jgourlay - 4-1-2010 at 06:57

SC Wack:

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

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

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S.C. Wack - 4-1-2010 at 14:27

Quote: Originally posted by jgourlay

SC Wack:

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

It still works for me.

Formatik - 13-6-2010 at 20:53

Quote: Originally posted by Formatik

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

Reading over the thread on SO3 from NaHSO4 in prepublication, using most common metal tubes probably won't work, e.g. hot conc. H2SO4 does attack steel forming SO2, and for SO3, as mentioned from Gmelin hot SO3 is reduced by iron forming sulfide.

Jimmymajesty - 14-7-2010 at 08:02

Hi folks!

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

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

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

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

BTW I've already used the FSE

[Edited on 14-7-2010 by Jimmymajesty]

Formatik - 14-7-2010 at 21:53

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

Jimmymajesty - 21-7-2010 at 11:38

Formatik thanx for the info!

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

Sorry for my brainfarts, but I am not at home at the moment so I cannot make experiements, It is easier to ask the skilled in the art

un0me2 - 15-8-2010 at 00:52

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

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

[Edited on 15-8-2010 by un0me2]

jgourlay - 16-8-2010 at 05:20

un0me2 : how do you get the two separated once they are mixed like that?

Formatik - 16-8-2010 at 10:44

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

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

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

And yes, chlorine is the oxidizer much more readily preparable from ubiquitous material.

un0me2 - 16-8-2010 at 13:39

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

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

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

Lambda-Eyde - 16-8-2010 at 13:45

Quote: Originally posted by un0me2

Also works with I2 & Br2.

Are you sure that iodine would be a strong enough oxidant for this?

Formatik - 16-8-2010 at 19:47

Quote: Originally posted by un0me2

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

Around 20 deg., about 10g SO2 will solubilize in 100g H2O. Which isn't that much. If you were bubbling SO2 into just H2O to form H2SO3, most SO2 could be lost (depending on your ratios), or unless it was recovered (maybe by setting up a series, but it's too much work), so it might work better to just oxidize the SO2 directly using Cl2 and water. That's if you care about the sulfur loss. If not then it would be easier to gas the H2SO3 solution with Cl2 because you would be working on generating and maintaining gas flow for only one gas.

If you were to solubilize the SO2, you could do it like this: first solvate estimated SO2 in an large excess of cold water so you don't need to worry about solubility. Then take the solution and gas it with Cl2. The fumes that come over collect with water which would then contain HCl, some HClO and Cl2. Then boil the H2SO3 oxidized solution just to get out the water (probably not until white fumes form), some more HCl should come over then. Finally, boil the aq. HCl containing solution to purify it, around 20% concentration will be reached, when you go beyond this, HCl strength will actually decrease on boiling.

Alcohol and ether solubilize at least over two times more SO2 than H2O, and I've handled those solutions. No fun, no fun at all. Up there with liquid ammonia.

Quote: Originally posted by Lambda-Eyde

Are you sure that iodine would be a strong enough oxidant for this?

Br2 and I2 are one of the oxidants which are listed to oxidize SO2 in Gmelin's Handbuch (it was mentioned on page 5 of this thread).

[Edited on 17-8-2010 by Formatik]

497 - 16-8-2010 at 21:14

Do you think bubbling Cl2 + SO2 into water could get H2SO4 concentrated past azeotropic? I don't remember if there are any side reactions that could occur in a mixture of HCl, conc. H2SO4, Cl2, SO2 and H2O... I can't think of any off the top of my head, besides formation of sulfuryl chloride, but that should not occur without a catalyst right?

[Edited on 17-8-2010 by 497]

S.C. Wack - 17-8-2010 at 02:53

What does this have to do with the lead chamber process?

I note that chlorine was well known in the past, and they chose to use the dearer nitrous fumes.

The solubility of SO2 in water is of course discussed in Mellor.

EDIT: It may well have been used at the later stages, and was patented earlier.

Early patents didn't seem to catch on. References from Lunge: http://books.google.com/books?id=RAhCAAAAIAAJ
1854: http://books.google.com/books?id=SgALAQAAIAAJ&pg=PA503
1863: http://books.google.com/books?id=-T4oAQAAIAAJ&pg=PA39
1904, German:
http://v3.espacenet.com/publicationDetails/originalDocument?CC=DE&N...
http://v3.espacenet.com/publicationDetails/originalDocument?CC=DE&N...

But I think the reference you want is:
http://dx.doi.org/10.1002/ange.19230365503

[Edited on 18-8-2010 by S.C. Wack]

Formatik - 19-8-2010 at 21:37

References in Gmelin verify the reaction goes as thought: when SO2 and Cl2 are led into water, this exotherms a bit and accumulates the H2SO4 as the HCl concentration decreases. Neumann described the reaction is going rapidly and almost completely (95-100% theoretical amounts were converted), the sulfuric and hydrochloric acids result immediately as fine droplets/fog, these are difficult to absorb and also pass over, as gases and water initially interact.

The patent mentioned of Stolle, leads same parts SO2 and Cl2 into water, eventually raising the temperature to 250 deg., yielding 90% H2SO4 and conc., free from Cl2 and SO2, aqueous HCl. Neumann's process is much more descriptive.

Neumann also described despite having used a Cl2-excess, a significant amount of SO2 got solubilized in H2SO4, since SO2 solubility increases with H2SO4 concentration. Though experiments also showed conc. H2SO4 which had Cl2 or SO2 solubilized in it, after blowing in air for 15 minutes, were almost completely removed.

Quote: Originally posted by S.C. Wack

What does this have to do with the lead chamber process?

It seems this thread is the designated stickied sulfuric acid thread. I would retitle it as the sulfuric acid preparation thread, or remove the non-Chamber discussions and sticky those with said title instead. Good eye on that reference, I also found it through Gmelin.

Quote: Originally posted by 497

Do you think bubbling Cl2 + SO2 into water could get H2SO4 concentrated past azeotropic? I don't remember if there are any side reactions that could occur in a mixture of HCl, conc. H2SO4, Cl2, SO2 and H2O... I can't think of any off the top of my head, besides formation of sulfuryl chloride, but that should not occur without a catalyst right?

I doubt it's of concern. Neumann described that after the reaction heat slows down, that the gases come out ununited. This heat is especially large when water is first consumed in the reaction. Their later experiments used additional heat (60-92 deg), to make the reaction go much faster.

Concerning the concentration of H2SO4 obtained by combination of SO2 and Cl2 with H2O, Neumann says it is that of the Chamber acid or Glover acid (66-88%). That's the raw figure then, it can be concentrated further by regular means. For practical purposes, instead of H2O, conc. HCl was recommended. Then when a specific gravity of 1.6 is reached, the hydrochloric acid content has been nearly completely removed.

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[Edited on 20-8-2010 by Formatik]

un0me2 - 24-8-2010 at 00:29

Nice route to clean H2SO4 & clean HX acids but... Anyway, the only reason it was posted was to allow those who didn't realise the alternative existed. Personally I like the construction tips on the Lead Chamber-type processes, but yeah...