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Author: Subject: H2SO4 by the Lead Chamber Process - success
Rosco Bodine
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[*] posted on 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]
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[*] posted on 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 :P

[Edited on 30-10-2008 by 497]

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[*] posted on 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..

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[*] posted on 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 :D
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[*] posted on 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..
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[*] posted on 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]
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smile.gif posted on 31-10-2008 at 11:17
Fe4O(SO4)5 Monsel's Salt, Powder Styptic U.S.P.


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]

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[*] posted on 31-10-2008 at 12:06
a larger scale preparation


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]

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[*] posted on 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]
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Rosco Bodine
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[*] posted on 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.
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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 :D

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]
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[*] posted on 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. :D

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! :P 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]
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[*] posted on 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]
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[*] posted on 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..
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[*] posted on 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.
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[*] posted on 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 :D



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[*] posted on 4-3-2009 at 10:18
Lead chamber volume


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?
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[*] posted on 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?
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[*] posted on 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 :D.
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]
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[*] posted on 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.
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[*] posted on 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)
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[*] posted on 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?
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