| Pages:
1
2
3
..
5 |
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Small scale production of H2SO4 in the amateur lab
Sulfuric acid is hard to get your hands on in certain countries and there also is upcoming regulation in the EU to limit its availability for the
general public. For home chemistry purposes, this chemical, however, is so basic and of such great importance, that many things will become
impossible, or at least very hard, if this chemical is not available anymore. For this reason we start a sticky thread on small scale production of
H2SO4. There already is a thread on the lead chamber process, but that is just one specific method of making sulfuric acid, which may not be suitable for everyone.
In this thread any serious method can be discussed. Even better would be descriptions of practical attempts, even if they failed. Learning from
failures can lead to future attempts with more success.
In order to limit the number of sticky threads, the lead chamber thread will not be sticky anymore, but it can be found directly by means of the above
link.
|
|
|
woelen
|
Thread Topped
15-8-2020 at 13:22 |
nzlostpass
Hazard to Self
Posts: 66
Registered: 3-9-2019
Member Is Offline
|
|
One of my first experiments was sulphuric acid by electrolysis of copper sulphate solution....I used carbon rods from batteries which discoloured it
but seemed to filter out ok. Im unsure how scaleable it is and I also am unsure what concentration can be achieved before boiling it down to
concentrate.
|
|
|
Abromination
Hazard to Others
Posts: 432
Registered: 10-7-2018
Location: Alaska
Member Is Offline
Mood: 1,4 tar
|
|
Because of hazard shipping issues, it is practically impossible to find or purchase sulfuric acid in my state. In the past I have used a sulfur
dioxide generator with 15 percent hydrogen peroxide, although this process is a lot of work for very little acid.
One method that I do not see talked about very often is the reaction of oxalic acid and iron (iii) sulfate in aqueous solution. Both reagents are
cheap, if you can find oxalic acid, and should theoretically produce iron free acid due to the extreme insolubility of iron oxalate. The solution can
then be boiled down or distilled to produce concentrated acid.
This method should be much easier then the SO2 and should produce more acid at once due to the convenience of not needing a gas generator.
I personally have not tried this method yet, but I have oxalic acid coming in the mail during the next few weeks and will share some observations.
List of materials made by ScienceMadness.org users:
https://docs.google.com/spreadsheets/d/1nmJ8uq-h4IkXPxD5svnT...
--------------------------------
Elements Collected: H, Li, B, C, N, O, Mg, Al, Si, P, S, Fe, Ni, Cu, Zn, Ag, I, Au, Pb, Bi, Am
Last Acquired: B
Next: Na
-------------- |
|
|
j_sum1
Administrator
Posts: 6333
Registered: 4-10-2014
Location: At home
Member Is Offline
Mood: Most of the ducks are in a row
|
|
Quote: Originally posted by nzlostpass  | | One of my first experiments was sulphuric acid by electrolysis of copper sulphate solution....I used carbon rods from batteries which discoloured it
but seemed to filter out ok. Im unsure how scaleable it is and I also am unsure what concentration can be achieved before boiling it down to
concentrate. |
I did something similar. I did refine the process to something quite manageable. I will post it here later: including some photos.
|
|
|
Fulmen
International Hazard
Posts: 1725
Registered: 24-9-2005
Member Is Offline
Mood: Bored
|
|
I'm seriously considering buying a 25kg bag of sulfur. I can get it in 1/2kg quantities from a home gardening supplier, but at a obscene price. Even
including shipping the bag will be cheaper than 2kg from home garden supplier.
I just need to convince a farming friend that he won't get on any terrorist watch list by ordering it. Shouldn't be a problem, as he's producing
ecological food and sulfur is approved for this. It actually looks like a very versatile compound that can be used both for soil improvement and as a
fungicide.
The process will most likely be the contact process. Vanadium is the typical catalyst, but IIRC platinum could also be used. It's said to be sensitive
to arsenic, but fertilizer grade sulfur should be pretty low in that so it might be worth a look. Perhaps a broken automotive catalyst could work?
We're not banging rocks together here. We know how to put a man back together.
|
|
|
Bedlasky
International Hazard
Posts: 1241
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
NurdRage have few videos on how to make sulfuric acid:
Copper(II) chloride method
Electrobromine process
Electrolysis of copper(II) sulfate
SO2/H2O2 method
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
A suggested route to very dilute H2SO4 based on a Wikipedia commentary on aluminum sulfate (https://en.wikipedia.org/wiki/Aluminium_sulfate ), to quote:
"When dissolved in a large amount of neutral or slightly alkaline water, aluminum sulfate produces a gelatinous precipitate of aluminium hydroxide,
Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment
insoluble...Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a
dilute sulfuric acid solution."
This method requires at least 3 steps, first obtaining or preparing aluminum sulfate, hydrolyzing it followed by a concentration step.
As to the first step, I was thinking of microwave heating of home available annealed Al foil in aqueous NaHCO3, as a pre-activation treatment. Rinse
off the treated Al metal and place in a concentration solution of CuSO4. Copper should plate out leaving a solution of aluminum sulfate.
An interesting alternate path requiring some investigation would be first prepare Al(OH)3 and add the freshly prepared salt to albeit only slightly
acidic aqueous MgSO4. May not react, or react slowly, or even may selectively react depending on the Al(OH)3 prep path (where high surface area, poor
crystallinity of the Al(OH)3, small particile size could be promoting factors, see, forexample, relatedly https://www.researchgate.net/publication/257224933_Effect_of...), or possibly result in a Mg-Al salt (see https://www.sciencedirect.com/topics/chemistry/magnesium-hyd... ).
The last step is futher dilute the obtained aluminum sulfate and followed by boiling to concentrate.
[Edited on 17-8-2020 by AJKOER]
|
|
|
Bedlasky
International Hazard
Posts: 1241
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
| Quote: Originally posted by AJKOER | A suggested route to very dilute H2SO4 based on a Wikipedia commentary on aluminum sulfate (https://en.wikipedia.org/wiki/Aluminium_sulfate ), to quote:
"When dissolved in a large amount of neutral or slightly alkaline water, aluminum sulfate produces a gelatinous precipitate of aluminium hydroxide,
Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment
insoluble...Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a
dilute sulfuric acid solution."
This method requires at least 3 steps, first obtaining or preparing aluminum sulfate, hydrolyzing it followed by a concentration step.
As to the first step, I was thinking of microwave heating of home available annealed Al foil in aqueous NaHCO3, as a pre-activation treatment. Rinse
off the treated Al metal and place in a concentration solution of CuSO4. Copper should plate out leaving a solution of aluminum sulfate.
An interesting alternate path requiring some investigation would be first prepare Al(OH)3 and add the freshly prepared salt to albeit only slightly
acidic aqueous MgSO4. May not react, or react slowly, or even may selectively react depending on the Al(OH)3 prep path, or possibly result in a Mg-Al
salt.
The last step is futher dilute the obtained aluminum sulfate and followed by boiling to concentrate.
[Edited on 17-8-2020 by AJKOER] |
Hydrolysis of aluminium sulfate is more complicated than that.
[Al(H2O)6]3+ + H2O <--> [Al(H2O)5(OH)]2+ + H3O+
[Al(H2O)5(OH)]2+ + H2O <--> [Al(H2O)4(OH)2]+ + H3O+
[Al(H2O)4(OH)2]+ + H2O <--> [Al(H2O)3(OH)3] + H3O+
[Al(H2O)3(OH)3] complex can react with other molecules of [Al(H2O)3(OH)3] to form polymeric species.
All positively charged complexes are soluble in water.
The major complex is [Al(H2O)5(OH)]2+ followed by [Al(H2O)4(OH)2]+ and the minor complex is [Al(H2O)3(OH)3].
When I prepared saturated solution of potassium alum I never saw any formation of aluminium hydroxide (even after few months). Aluminium sulfate may
behave differently, but I doubt that huge amounts of aluminium hydroxide are formed.
So this will be very inefficient way how to prepare sulfuric acid, because it will be very very dilute and highly contaminate with aluminium.
--------------------------------------------------------------------------------
I had a few ideas yesterday and I'll try it when I will have some time:
Oxidation of sulfur by hot conc. HNO3
Oxidation of sulfur by Fenton's reagent (which requires only catalytic amount of iron)
[Edited on 17-8-2020 by Bedlasky]
|
|
|
clearly_not_atara
International Hazard
Posts: 2799
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
By far the simplest method I have heard of is to use the disproportionation of potassium bisulfate in ethanol:
http://www.sciencemadness.org/talk/viewthread.php?tid=79548
You can just boil down the filtrate to obtain reasonably concentrated sulfuric acid! It is still effective but less efficient and more "sticky" when
using the sodium salt. The possibility of using NH5SO4 has also occurred to me but not been tested (I'll get to it someday...).
This method is also good for making a "seed" quantity of sulfuric acid that can then be treated with SO3 to give oleum and diluted etc. Using both
methods together you can make lots of sulfuric acid from just bisulfate, ethanol and water -- but the intermediacy of SO3 is always a little scary.
[Edited on 17-8-2020 by clearly_not_atara]
|
|
|
RogueRose
International Hazard
Posts: 1595
Registered: 16-6-2014
Member Is Offline
|
|
| Quote: Originally posted by Fulmen | I'm seriously considering buying a 25kg bag of sulfur. I can get it in 1/2kg quantities from a home gardening supplier, but at a obscene price. Even
including shipping the bag will be cheaper than 2kg from home garden supplier.
I just need to convince a farming friend that he won't get on any terrorist watch list by ordering it. Shouldn't be a problem, as he's producing
ecological food and sulfur is approved for this. It actually looks like a very versatile compound that can be used both for soil improvement and as a
fungicide.
The process will most likely be the contact process. Vanadium is the typical catalyst, but IIRC platinum could also be used. It's said to be sensitive
to arsenic, but fertilizer grade sulfur should be pretty low in that so it might be worth a look. Perhaps a broken automotive catalyst could work?
|
Do you have any Tractor Supply stores, Rural King or similar stores near you - or even any feed &/or seed stores? All of them should have "sulfur
Flour" in 50lb/25kg bags for very good prices. I think maybe even pottery supply shops might have it as well, but I'm not sure, it may have been only
available at a few places I looked.
|
|
|
Fulmen
International Hazard
Posts: 1725
Registered: 24-9-2005
Member Is Offline
Mood: Bored
|
|
Nah. I ain't living in the US, and there simply isn't any shops selling anything useful anywhere close.
We're not banging rocks together here. We know how to put a man back together.
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
If you can access epsom salt and plaster then you can also electrolyze magnesium sulfate with a plaster separator.
Sodium sulfate can also be electrolized in a similar chamber.
The electrolytic methods are less than ideal and produce weak acid with impurities for a lot of work.
The copper sulfate method is the easiest.
You can also superheat gypsum to high temperatures (1080C) to release sulfur dioxide.
Burning sulfur is going to be easier though possible not cheaper.
Burning DMSO is also a possibility as DMSO is available for use on horses.
Sulfur dioxide can be catalyzed to sulfur trioxide on a vanadium oxide catalyst or via a chamber method.
Both have been done by amateur chemists.
The chamber method is best done on a small scale using 5 gallon water bottles.
|
|
|
Tsjerk
International Hazard
Posts: 3032
Registered: 20-4-2005
Location: Netherlands
Member Is Offline
Mood: Mood
|
|
I would still go for the bisulfate/pyrosulfate route as suggested by Atara.
Attachment: SO3_and_oleum.pdf (198kB)
This file has been downloaded 1103 times
|
|
|
Mateo_swe
National Hazard
Posts: 548
Registered: 24-8-2019
Location: Within EU
Member Is Offline
|
|
It sure would be nice if there was a pretty easy process for sulfuric acid preparation.
We might be able to buy it now but there is already politicians deciding that the common citizen should not have access to any chemical that is
dangerous or can be used to make anything dangerous.
The ban is already in motion so any DIY way of preparing chemicals is good to have knowledge about.
I have tried the oxalic acid + Ferrous Sulfate method and it works.
FeSO4 + H2C2O4 --Heat --> FeC2O4 (s) + H2SO4
There is a thread about this method around here somewhere if i remember correctly.
If there is an easy way of making cation and anion membranes (maybe there is), there is a good way of making sulfuric acid and sodium hydroxide from
sodium sulfate in a 3 compartment cell.
In the old days they made sulfuric acid by heating Ferrous Sulfate until it decomposed.
It will need a high temp oven to experiment with this method.
Since Ferrous Sulfate is cheap in 25kg bags (32 Euros on ebay + shipping), methods using this is nice.
I add some pdfs about preparation of sulfuric acid.
Attachment: chemicals_sulfuric_acid_preparation_from_oxalic_acid_thompson1849.pdf (209kB)
This file has been downloaded 986 times
Attachment: h2so4_vogel_eng_H2SO4_from_Oxalic_acid_and_Iron_sulphate.pdf (255kB)
This file has been downloaded 957 times
Attachment: US5928488 - Electrolytic Sodium sulfate salt splitter comprising a polymeric ion conductor.pdf (981kB)
This file has been downloaded 889 times
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
For preparation of copper sulfate
Start with ammonium hydroxide and magnesium sulfate
To form ammonium sulfate excess ammonia then add to copper oxide prepared by heating copper till it blackens and flakes off when it cools strong
heating to drive off the ammonia from the complex to form copper sulfate
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Clearly Not Atara:
If I am reading a referenced source on salting-out correctly, NH4+ > K > Na....
So, heating (but below 350 C, which is a cited decomposition temperature forming sulfur oxides) ammonium sulfate (from, for example, the action of NH3
on aqueous Epsom's Salt, MgSO4) to create ammonium bisulfate (see http://www.sciencemadness.org/talk/viewthread.php?tid=198 ) and adding a select alcohol, may work as well.
There may also be other paths starting with ammonium sulfate, and I am looking at one now.
[Edited on 19-8-2020 by AJKOER]
[Edited on 20-8-2020 by AJKOER]
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
I believe you can produce oleum by simply heating sodium bisulfate until it decomposes to sodium pyrosulfate and then heating the sodium pyrosulfate
until it decomposes. It looks easy, but the high temperatures required cannot be safely achieved in ordinary borosilicate glass, and of course
anything involving oleum is extremely hazardous. Len1 discussed the procedure in some detail in his book, Small-Scale Syntheses of Laboratory
Reagents.
I could possibly try it out next weekend, but I'm not sure what to do with the oleum afterwards, and I'd have to do it outside, preferably in a
location that won't terrorize the local homeless population... anyway, I ordered some sodium bisulfate.
|
|
|
Mateo_swe
National Hazard
Posts: 548
Registered: 24-8-2019
Location: Within EU
Member Is Offline
|
|
Do you think Sodium metabisulfite (Na2S2O5) could be used as a replacement for the Sodium bisulfate in the process?
It is found in brew shops for ok prices.
[Edited on 2020-8-22 by Mateo_swe]
|
|
|
Tsjerk
International Hazard
Posts: 3032
Registered: 20-4-2005
Location: Netherlands
Member Is Offline
Mood: Mood
|
|
JJay, have a look at the pdf I attached above. Here the decomposition is done in borosilicate and persulfate is cheap and available as etchant for
circuit boards.
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
| Quote: Originally posted by Tsjerk | | JJay, have a look at the pdf I attached above. Here the decomposition is done in borosilicate and persulfate is cheap and available as etchant for
circuit boards. |
I saw it. The major drawback I see is that it requires anhydrous sulfuric acid as one of the starting materials. Also, sodium persulfate is much more
expensive than sodium bisulfate. I considered giving it a try (and why not just pyrolyze the cheaper sodium bisulfate instead of sodium persulfate?),
but I think the strongest sulfuric acid I have is 93%, so I'd like to know what the effect on yields is if there is a little bit of water present.
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
| Quote: Originally posted by Mateo_swe | Do you think Sodium metabisulfite (Na2S2O5) could be used as a replacement for the Sodium bisulfate in the process?
It is found in brew shops for ok prices.
[Edited on 2020-8-22 by Mateo_swe] |
It could be oxidized to sodium bisulfate, but sodium bisulfate is usually cheaper and easier to find.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Bedlasky:
Thanks for your comment, as I agree, the hydrolysis of Aluminum sulfate with the 'acid' water to create a separable amount of very dilute H2SO4 is
unlikely.
However, perhaps working with the slightly more powerful CO2 (aq) composition, leading to an insoluble salt product, may be worth exploring as a path
to some very dilute H2SO4 itself, based on a Science Direct commentary. More precisely, the cited reaction, occurring in natural waters, of the
bicarbonate ion acting on alum resulting in a precipitate of Al(OH)3 per the cited reaction (from a Science Direct compilation available at: https://www.sciencedirect.com/topics/chemistry/aluminium-sul...):
Al2(SO4)3.xH20 + 6 HCO3- (aq) <--> 2 Al(OH)3 (s) ....
but in a new suggested embodiment, exploring the use of slightly acidic carbonic acid. Note: the reaction is reversible, but using a reaction chamber,
like a test tube, where the surface area contact of the created Al(OH)3 precipitate is reduced, may be beneficial and avoid stirring.
Did find a supporting reference (https://thefishsite.com/articles/managing-high-ph-in-freshwa... ) relating to care for treating pH problems in fresh water ponds, with naturally
contained dissolved CO2 from air, to quote:
"Overtreatment with alum can cause a dramatic decrease in pH, possibly to levels more dangerous than the original high pH problem...A safer, longer
lasting way to reduce high pH is to add carbon dioxide, which acts as an acid in water...The careful use of alum is probably the safest and most
dependable emergency treatment."
Also, a more dated reference (at https://www.jstor.org/stable/pdf/41224990.pdf) claims that optimal efficiency relating to coagulation (Al(OH)3 formation) occurs not in alkaline
pH, but at around pH 5.5 (which is also cited in this 1989 thesis at https://scholars.unh.edu/cgi/viewcontent.cgi?article=2599&am... ). Further, the addition of acid assists (like H2SO4 or other), and that CO2 is
not particularly good. On the latter point to quote another source (https://drinking-water.extension.org/drinking-water-treatmen...):
"Citric acid and alum can be used instead, although they are more expensive."
The chemistry associated with the use of agents other than CO2 in transition metal and oxygen rich natural waters, however, may be related to the
presence of active redox reactions and also, per the cited thesis (Pages 39-40), aluminum polymer formation per the presence of fulvic acid and the
electrostatic properties of colloidals resulting in particle aggregation. But more insightful per a 2019 work (https://www.researchgate.net/profile/Wenzheng_Yu/publication...), to quote:
"When monomeric Al (alum) was added together with kaolin at pH 7, some [Al(OH2)(6-n)(OH)n](3-n)+ absorbed on the surface of kaolin, causing the
surface of kaolin to be fully covered with alum precipitates/nanoparticles (Duan and Gregory, 2003). Then the flocculation process started."
So, hydrated Al(OH)3 can be induced (with select acids, kaolin based clay,...) to introduce deposits of varying charged surface complexes to assist in
particle aggregation.
Other redox reactions in the presence of Fe ions are also possible. Further, with charged colloidal particles, solvated electron creation is possible,
which may result, in the presence of O2 again, the creation of a superoxide radical anion:
O2 (d) + e- (aq) = .O2- (or .HO2 at pH < 4.88)
These radical species have been reported to interact with Al3+, as I have noted previously on SM with references discussed at https://www.sciencemadness.org/whisper/viewthread.php?tid=96... , and further promote redox chemistry.
On the role of acids and Al speciation, note the following source https://www.researchgate.net/figure/Monomeric-and-polymeric-... , to quote:
"Numerous studies have reported that there are various Al forms in soils in monomer, polymer or solid phase, and that their concentration depends on
the degree and duration of hydrolysis of the Al compounds (Delhaize and Ryan, 1995) (Figure 1). Rout et al . (2001) found a significant correlation
between low pH and high concentrations of phytotoxic Al species, which is related to the reduction of exchangeable bases in the soil solution (Mora et
al ., 2006). Soil acidification is associated with inappropriate agricultural practices (Rengel, 1996), heavy winter precipitation that causes the
loss of bases (Na + , K + , Ca 2+ , Mg 2+ ) due to leaching (Mora et al ., 2006), use of ammoniacal fertilizer (urea) and nutrient uptake by plants
(Mora et al . , 2006)."
Still, assuming dilute H2SO4 is after all present, can the weak acid be effectively concentrated here?
The suggested theoretical investigation reminds me of an Oxalate acid path, but with a much weaker starting acid.
[Edited on 24-8-2020 by AJKOER]
[Edited on 24-8-2020 by AJKOER]
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
Thermal decomposition of gypsum to sulfur dioxide for use conversion to sulfuric acid by chamber or contact process.
See attached file.
Attachment: Decomposition of Calcium Sulfate.pdf (3.3MB)
This file has been downloaded 1045 times
|
|
|
Alkoholvergiftung
Hazard to Others
Posts: 188
Registered: 12-7-2018
Member Is Offline
|
|
Müller-Kühne process for Sulfuric acid and cement. works with an rotation stove and an mix of gypsum clay and sand the temperatur is around 700C.
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
Alkoholvergiftung,
Müller-Kühne process requires a final temperature of 1150-1250C to release most of the SO2 and carbon is also utilized.
|
|
|
| Pages:
1
2
3
..
5 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
