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

H2SO4 by the Lead Chamber Process - success

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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=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]

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 898 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 880 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 1332 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 2003 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 1031 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...

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

Attachment: US5011693 Preparation of Monsels Solution.pdf (103kB)
This file has been downloaded 1739 times

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&...

[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?

1281371269 - 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?

1281371269 - 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.

1281371269 - 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-ma...
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?

1281371269 - 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 1204 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"
------------------------------
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 88 times

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.

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