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Author: Subject: H2SO4 by the Lead Chamber Process - success
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[*] posted on 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!
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[*] posted on 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]
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[*] posted on 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...:o 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]

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[*] posted on 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.
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[*] posted on 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]

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[*] posted on 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]
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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?




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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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).




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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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=a...

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=a...

[Edited on 18-6-2008 by 497]


[Edited on 18-6-2008 by 497]
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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? :P
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[*] posted on 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 =)
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