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Author: Subject: Sulphuric acid.... AGAIN
Titanium100
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[*] posted on 28-5-2015 at 03:28
Sulphuric acid.... AGAIN


Hi all,

I am hoping someone could help me with some questions that I have not been able to find the answers for, despite using TFSE.
I hope they are not too idiotic?

Oh, .... and Blogfast ...please be gentle :)

Anyway, suppose you generate sulphur dioxide and dissolve that in water, I believe that you get sulphurous acid, and that to get sulphuric you need to add? another oxygen. I have read that air oxidation will work, but I am guessing that it would be very slow with a container of sulphurous, and limited surface area. Assuming that hydrogen peroxide and nitric acid were both unavailable, would bubbling ozone through the solution be a viable way to do the oxidation? Would bubbling air through it work at all? Also whatever method was used to oxidise the sulphurous, what test could be used to determine if the reaction was complete?
Thanks for any assistance.
Tim
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byko3y
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[*] posted on 28-5-2015 at 04:09


I know no effective way of oxone generation. Sulfur dioxide is a reducing agent, so there's is a lot of reagents capable of oxidating it.
I've already posted somewhere on this forum an overview of different methods, including kinetics study of manganese catalyzed oxidation of SO2 with oxygen. Maybe somewhere in HCl + H2SO4 thread. If you find it - post link here pls.
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hissingnoise
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[*] posted on 28-5-2015 at 04:21


Quote:
Would bubbling air through it work at all?

Yes it would, albeit slowly ─ constant stirring of the sol. will work too . . .

Ozone, though difficult to produce in quantity, is a very strong oxidiser and will rapidly oxidise SO2 in the gas-phase and in solution!

See here?

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woelen
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[*] posted on 28-5-2015 at 05:20


Another option is to bubble Cl2 through your solution of SO2. The reagents, needed for making Cl2, can be obtained in nearly every country. In your solution you get H2SO4 and HCl:

2H2O + SO2 + Cl2 --> H2SO4 + 2HCl

By boiling the solution, you concentrate the H2SO4 and drive off the HCl.

Cl2 is absorbed very well by a solution of SO2 in water. As soon as absorption of Cl2 stops and it bubbles through, you are ready and have converted all SO2 to H2SO4.

As I wrote earlier, making SO2 can best be done with solid Na2S2O5 (sodium metabilsulfite) or solid K2S2O5, to which 10% HCl is added, followed by gentle heating.




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byko3y
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[*] posted on 28-5-2015 at 17:11


woelen, it's okay, but as you might know, to make Cl2 you actually need an acid.
So it's HOCl + HCL -> Cl2 ... + SO2 + H2O -> H2SO4 + HCl
And generation of SO2 by using another acid looks even more dummy, because there's easier ways of converting HCl to H2SO4.
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macckone
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[*] posted on 28-5-2015 at 19:08


I am interested in an easy way to convert HCl to sulfuric acid. It is news to me that they exist. Going from sulfuric to hydrochloric is easy but I have not seen a good way to go the reverse direction.

On the original topic iron oxide will catalyze the oxidation of sulfur dioxide and sulfurous acid. If it is an acceptable contaminant then iron may be the way to go. In solution iron will complex but mainly exist as the ion.
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byko3y
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[*] posted on 28-5-2015 at 21:52


PbSO4(s) + 4HCl(conc.) = H2PbCl4(aq) + H2SO4
172 g of CaSO4 + 278 g of PbCl2, add hot 60°C water, separate the precipitate of PbSO4 weighting 303 g. The solution is discarded. To the PbSO4 added 100-150 g of HCl solution and the mixture is heated to 60°C. Yield - 100 g of dilluted H2SO4. The precipitate of PbCl2 weighting 278 g can be reused.
Lead sulfate is toxic by inhalation, ingestion and skin contact. It is a cumulative poison, and repeated exposure may lead to anemia, kidney damage, eyesight damage or damage to the central nervous system (especially in children). It is also corrosive - contact with the eyes can lead to severe irritation or burns. Typical threshold limit value (above which the substance is harmful) is 0.15 mg/m3.
Some of these might also work:
Sb2(SO4)3 + 8HCl(conc.) = 2H[SbCl4] + 3H2SO4
Bi2(SO4)3 + 8HCl(conc, cool) = 2H[BiCl4] + 3H2SO4
SnSO4 + 3HCl (conc) = H[SnCl3] + H2SO4
Sn(SO4)2 + 6HCl (conc.) = H2[SnCl6] + 2H2SO4
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[*] posted on 29-5-2015 at 00:46


Intriguing procedure.
Have you had any experience with it? How good are the yields? How dilute is the product?
It sounds simple: almost too good to be true. I am surprised that I haven't come across it before. Or is there a catch?

Could the same be done with MgSO4? High purity epsom salt is essier to find than plaster of paris without impurities -- and a bit easier tyo work with too.
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[*] posted on 29-5-2015 at 01:51


@byko3y: Your methods of converting HCl to H2SO4 are crap. The reactions certainly do not proceed nicely and smoothly. At best you get partial conversion and I expect the procedures to be very messy.

What I propose in my earlier post is a method for making H2SO4, using ONLY easy to obtain chemicals. You only need 10% HCl, metabisulfite and some swimming pool chlorine variant or bleach. The procedure is not the most simple one, but it sure can be done by a home chemist. Passing SO2 through water, then leading Cl2 through that, then passing more SO2 in that, leading more Cl2 through that. In this way I expect you can obtain a concentration of well over 10% H2SO4. Then you need to boil down.

The methods with the salts are MUCH more difficult. Not easy to separate. I even have doubts on compounds like Sb2(SO4)3. If you cannot obtain sulphuric acid, then you certainly cannot obtain Sb2(SO4)3. This definitely is not something you easily make at home (just read up on Sb-chemistry). The same is true for SnSO4, Sn(SO4)2 and Bi2(SO4)3. All these methods look like purely theoretical, but they will never work in practice. The salts will hydrolyse in water and separation of the resulting compounds is very hard.

I also have doubts on the PbSO4 method. PbCl2 is more soluble than PbSO4.

Summarizing: There is no easy way of converting HCl to H2SO4. What I propose is doable, I do not claim it to be super easy, but if you really have access to only the most basic chemicals, then I think it is the best there is.

There is practical chemistry and armchair chemistry . . .




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[*] posted on 29-5-2015 at 03:00


Thanks woelen. I had some suspicions that it might not proceed as well as it looked on paper. Still, I would love byko3y to give some more details on his experience. I have done some more thinking it over and I begin to see some of the issues.

The thing that I noticed is the large number of insoluble species. CaSO4 is not highly soluble. Nor is PbSO4. And PbCl2 is not that soluble either. (This is from memory and so correct me if I am wrong.) So, either you have exceptionally large amounts of water -- more than was quoted, or you have a displacement reaction involving solids as both reactants and products:
CaSO4(s) + PbCl2(aq) --> PbSO4(s) + CaCl2(aq)
I can see difficulty in knowing when the displacement is complete. I would be nervous with lead in being absolutely certain in knowing where it is. I would want to know that it was all gone from the solution and had entered the solid precipitate. Just thinking about that, I don't know how I could be sure.
So... you start off with white sludge mixed with a lead solution. And you finish with a white sludge and a solution that you hope contains no lead. Then you filter something that is likely quite messy. Add HCl in excess. Drive off the excess HCl by boiling and continue to reduce the volume because you want to make sure all of your PbCl2 precipitates. Then filter again -- something that is quite acidic. Hopefully you don't kill your filter paper with the sulfuric acid. And then boil off some more to concentrate.
You have mechanical losses in your filtering. You have no real way of knowing where your lead is going and whether you have any CaSO4 left in you sludge. Lots of boiling and filtering. And then you have the issue of decontaminating lead compounds off all of your surfaces.

I guess all of this could be streamlined and if the losses could be minimised then it might be workable. But I don't think it is something that I would relish.

I think I will continue my electrolysis of copper sulfate and concentrating the product. Tedious, but it works.
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byko3y
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[*] posted on 29-5-2015 at 03:35


woelen, this is a hundred years old procedure from a russian chemistry book. Probably the sufuric acid will not be very pure, but it will be obtained.
I have no experience on this procedure, because I don't want to mess around with lead.
Lead is easy to obtain for me, it's not regulated like mercury.
In case you can use relatively complex aparatus and have a good chlorine source, like TCCA or cheap permanganate (because NaOCl is a bad, usually very dilluted source), you can make sulfuryl chloride and distill it (it's relatively resistant to hydrolysis by water). THen add water and you will get a pure sulfuric acid and HCl gas. SO2 forms addicts with camphor, acetic acid and maybe some other compounds, thus creating a buffer to avoid a need of precisely measure chlorine and SO2 flow.
I really would be glad if someone tried metal-catalyzed oxidation of SO2 with air in water, which is safe, but probably slow.

[Edited on 29-5-2015 by byko3y]
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[*] posted on 29-5-2015 at 04:02


Again, that is all really interesting byko3y. But again, I am not sure that I am inspired to attempt the routes involving sulfuryl chloride.
I also would be interested in knowing a bit more about SO2 oxidation to SO3 with probably an iron catalyst. I do need better equipment before I start playing around with SO2 gas. It is not great stuff to breathe.
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[*] posted on 29-5-2015 at 04:54


You need a real good equipment to play around with gaseous SO3 in fact. On leakage it creates a cloud of sulfuric acid particles destroying everything in its way. Despite the fact I only started some uncertain experiments of SO3 production, quickly I found out that this compound is not something amateur chemists wants to deal with (there's a thread about its production from MgSO4+NaHSO4 or persulfate http://www.sciencemadness.org/talk/viewthread.php?tid=5495&a... ). I'd better try a direct SCl2 oxidation with oxygen ( http://www.sciencemadness.org/talk/files.php?pid=133835&... ), instead of messing with SO3.
Making water solution of SO3 from SO2 looks more safer to me, but I'm too lazy to try it myself.
http://www.sciencedirect.com/science/article/pii/09601686919...
http://www.sciencedirect.com/science/article/pii/00092509818... (I've posted it somewhere on the forum)
http://pubs.acs.org/doi/abs/10.1021/j100218a027
More on topic: http://pubs.acs.org/action/doSearch?text1=Oxidation+of+aqueo...
http://pubs.acs.org/action/doSearch?text1=Oxidation+of+aqueo...

[Edited on 29-5-2015 by byko3y]
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hissingnoise
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[*] posted on 29-5-2015 at 11:22


Quote:
Making water solution of SO3 from SO2 looks more safer to me, but I'm too lazy to try it myself.

SO3 does not dissolve in water, as has been said here innumerable times!
Both compounds superheat on contact, producing a mist of 98% H2SO4 which is largely incondensible!

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[*] posted on 29-5-2015 at 11:33


But Sulfur Trioxide does dissolve in concentrated sulfuric acid making noxious Fuming Sulfuric Acid. This can be diluted back down to concentrated. You need some sulfuric to make some sulfuric.
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[*] posted on 30-5-2015 at 02:07


Quote:
You need some sulfuric to make some sulfuric.

Well, if you're sufficiently adventurous/foolhardy, you can just dump the solid SO3 you collected from your coldtrap into water and stand well back?



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[*] posted on 2-6-2015 at 09:49


All of this talk leads to an interesting possibility for making lead sulfate. For those with Birkeland Eyde reactors lead nitrate is easy to make. Also Epsom salt is easy to get. Mixing solutions of the two should give a relatively pure precipitate of lead sulfate which is otherwise difficult to make.

Mixing lead sulfate with hydrochloric acid doesn't seem like an efficient way to sulfuric acid given that the chloride is more soluble than the sulfate. But it is probably easier than cooking various sulfates to produce the trioxide which is hard but not impossible to combine with water safely.
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[*] posted on 2-6-2015 at 10:37


Quote: Originally posted by macckone  
For those with Birkeland Eyde reactors lead nitrate is easy to make.


All TWO of'em, macckone? :D

Not to mention that lead sulphate is such a bore!




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macckone
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[*] posted on 2-6-2015 at 16:51


I think there are more than 2. I have seen some pictures. And then there is the one I have. As for usefulness. It can be used to make ... err ... discharged lead acid batteries.
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[*] posted on 3-6-2015 at 16:42


Code:
1. Al + S8 => Al2S3 [ burning ] Al2S3 + H2O => H2S + Al(OH)3 CxHy + S8 => H2S + CxHySz 2. CuSO4 + H2S => CuS + H2SO4 CuS + 4H2SO4 (conc., hot) => CuSO4 + 4SO2↑ + 4H2O Some ideas on regeneration: CuS + 8HNO3 => CuSO4 + 8NO2 + 4H2O or http://link.springer.com/article/10.1007%2Fs10973-005-0601-1 2CuS + 2.5O2 => Cu2O + 2SO2 [ 280-475°C ] Cu2O + SO2 + O2 => 2CuO + SO3 [ 280-475°C ] Cu2O + O2 + SO2 => CuO*CuSO4 [ 475-680°C ] 2CuO + SO2 + 0.5O2 => CuO*CuSO4 [ 475-680°C ] Cu2O + 2SO2 + 1.5O2 => 2CuSO4 [ 475-680°C ] PS: CuSO4 starts decomposing at 700°C. Hydrogen peroxide facts: CuS + NaOH + H2O2 => Cu(OH)2 + NaSO4 + H2O [ ph > 9 ] H2S + H2O2 => S8 + H2O2 Copper(II) sulfide: Solubility in water 0.000033 g/100 mL (18 °C) Soluble in HNO3, NH4OH, KCN insoluble in HCl, H2SO4


[Edited on 4-6-2015 by byko3y]
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[*] posted on 5-6-2015 at 12:57


Titanium100,

What's your motivation to make H2SO4? If you're not interested in making it, then just buy it as drain cleaner. It's by far the cheapest way - about $6.2/kg when you buy a 1 liter bottle of drain cleaner.

Have you considered electrolyzing copper (II) sulfate? That pretty simple and doesn't involve messy intermediate chemicals. As much as I like electrolysis, this is slower (needs 53.6 amp * hours per mole of H2SO4 produced) and more expensive ($25 for 10 lbs of CuSO4 * 5H2O -> $14 /kg acid, not including shipping or electricity)

Of course, if you're stranded on an island like Cyrus Smith from The Mysterious Island and want to make the acid from scratch, go for it. He and his fellow castaways roast "shistose" pyrites to make iron sulfate, which are roasted again to yield SO3. Very ingenious, but I don't know how realistic it is?



[Edited on 5-6-2015 by UncleJoe1985]
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byko3y
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[*] posted on 5-6-2015 at 13:42


Roasting FeSO4 is a nice way of making sulfuric acid, but it requires some special equipment, like porcelain retort and furnace capable of reaching 900-1000 C.
I really want to see something new in OTC sulfuric acid production, because of everything community knows the only actually viable route is a contact process :/ This one would be nice in case I did not have to deal with SO3 and 450 C of catalyst bed temperature. Well, probably pyrex glass + acid resistant clay will help me to make the final few parts of apparatus (from catalyst to H2SO4).
And currently if I needed SO3 I would have used Na2S2O7+MgSO4 method, but I prefer to not need SO3 :)
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[*] posted on 6-6-2015 at 06:06


Actual decomposition temp of ferrous sulfate is 680C. And steel is fairly resistant to dry sulfur trioxide. Of course you can apply a coating of porcelain without too much difficulty. I have been looking for porcelain retorts with little success.


[Edited on 6-6-2015 by macckone]

[Edited on 6-6-2015 by macckone]
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[*] posted on 6-6-2015 at 16:24


Quote: Originally posted by macckoneOf course you can apply a coating of porcelain without too much difficulty.[/rquote  

How do you do this? The only ways I know of to coat something with a ceramic are very high temperature, butif you know a nice easy way please share.
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[*] posted on 6-6-2015 at 23:28


Coating requires a kiln although it has been done with oxygen acetylene.
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