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Author: Subject: Small scale production of H2SO4 in the amateur lab
macckone
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[*] posted on 13-9-2020 at 10:03


Quote: Originally posted by clearly_not_atara  
If I were to guess, this is made possible because the ionic reactions are like this:

PbCl2 (Ksp 1.7e-5) + CaSO4 (Ksp 4.9e-5) >> PbSO4 (Ksp 2.1e-8) + Ca2+ (aq) + 2 Cl- (aq)

PbSO4 + HCl + Cl- >> PbCl2 + HSO4- (aq)

That is, the sulfate ions are protonated in the acidic hydrochloric acid solution, but not in the neutral solution. The success of the calcium-to-lead metathesis thus depends on a neutral pH, while the lead sulfate-to-chloride metathesis happens at acidic pH, where the sulfate ion is protonated to bisulfate (pKa ~2).


This changes with ion concentrations and removal as well.
Both reactions are equilibrium, so if you can remove the calcium chloride, the reaction proceeds reasonably well.
If you remove the sulfuric acid the second reaction proceeds as well.
Hence my thought on the soxhlet extractor where you have continuous if slow removal of reactants from one side of the equation.
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AJKOER
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[*] posted on 24-9-2020 at 08:02


Came across an interesting electrochemical based reaction to quote:

"2 FeCl2 + 2 H2O + 1/2 O2 --> Fe2O3 + 4 HCl (1)

Hydrochloric acid is regenerated, and commodity oxides of iron are obtained as a result of the reaction [8, 9]."

Source: http://rudmet.net/media/articles/Article_CIS_vol.15_18_pp.28...

Turning to sulfate chemistry, with respect to ferric sulfate by Atomistry at http://iron.atomistry.com/ferric_sulphate.html to quote:

"Upon dilution, ferric sulphate solutions readily undergo hydrolysis, precipitates being obtained which, however, have no well-defined composition.
A study of the electric conductivities of aqueous solutions of the salt indicates that the hydrolysis proceeds in two stages, embodying (1) a rapid change unaccompanied by precipitation, and (2) a slower change, progressing at a measurable rate, and accompanied by the production of a so-called basic salt. Colloidal ferric hydroxide does not appear to be formed during hydrolysis, the salt thus differing from ferric chloride and nitrate...In dilute solution ferric sulphate is reduced by metallic iron to ferrous sulphate."

More recent research on the action of O2 on iron salt solutions, for example, this fully free available 2008 reference: "Air Oxidation of Ferrous Iron in Water" by Ahmet Alıcılar, et al, at http://www.jieas.com/fvolumes/vol081-5/3-5-11.pdf.

“Abstract: Air oxidation of ferrous iron in water was studied...Thereafter, the experiment was successively repeated by blowing air to the solution without and with inert packing. Lastly, the catalytic effect of ferric hydroxide was investigated. While the maximum yield of 86 % is catalytically achieved by blowing air at a neutral medium, the oxidation was almost completed in an alkaline solution even at stationary oxidation was almost completed in an alkaline solution even at stationary atmosphere. The reaction was first order with respect to Fe2+. “

And further:

“Oxidation of iron is achieved by addition of chemical oxidants. However, it can be easily and low cost carried out by contact with air (Wong, 1984). During oxidation of Fe2+ salt aqueous solutions, poor soluble compounds including Fe3+ oxides are formed (Domingo et al., 1994). The composition of precipitate formed depends on numerous parameters such as temperature, pH, concentration, feed rate and anion nature (Das & Anand, 1995; Tolchev et al., 2002)....The oxidation kinetics of Fe(II)(aq) species has been previously reviewed by many workers (Wehrli, 1990; Zhang et al., 1992). The stoichiometry for the overall oxidation of Fe2+ ions by O2 is given by Eq. (1) (Burke & Banwart, 2002).

O2(aq) + 4Fe2+ + 6H2O ↔ 4FeOOH(s) + 8H+ (1) “

So, similar to FeCl2, the slow air pump oxidation of a RT dilute aqueous FeSO4 and metal Iron mix to the point of possible dilute H2SO4 creation may be a subject of my future investigations. Substituting (NH4)2SO4(aq)/Fe for FeSO4(aq)/Fe, may also be interest.

[Edited on 24-9-2020 by AJKOER]
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[*] posted on 30-9-2020 at 00:33


I just had a crazy idea of how to make sulfuric acid. As we all know sulfuric acid can be made by adding oxalic acid to Ferrous Sulfate. Ferrous Sulfate being essentially free oxalic acid is what makes this an expensive process. But what if you made the oxalic acid in situ? If you put out a bowl of Ferrous Sulfate, sucrose and a catalytic amount of nitric acid in the dark one might expect the following reactions to take place:

C12H22O11 + 12 HNO3 = 6 C2O4H2 + 11 H2O + 12 NO
NO + O2 + H2O = HNO3 (catalyzed by iron in solution)
C2O4H2 + FeSO4 + 2 H2O = FeC2O4·2H2O + H2SO4

With the overall reaction being:
C12H22O11 + 9 O2 + 6 FeSO4*5H2O = 6 FeC2O4·2H2O + 6 H2SO4 + 23 H2O

The main problem may be overoxidation of the sugar and the slow reaction between sugar and nitric acid. The former is catalyzed by light/iron couple which should be easy to avoid unless iron by itself can do it too. Oxalate may not be rremoved fast enough due to solubilizing effects from the sugar. The latter problem could probably be solved by adding more nitric acid. Nitrogen oxides may be lost but conditions should be very favorable for reoxidation so it may be manegable.

Just like with the oxalic acid process a nice bonus is the Ferrous Oxalate dihydrate byproduct which may be useful to make iron or iron oxide powder though it may be diffucult to contain it 100g+ scale.

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[*] posted on 5-11-2020 at 03:38


Quote: Originally posted by Tsjerk  
Adding methanol to a potassium hydrogen sulfate solution works and is nearly quantitative.

http://www.sciencemadness.org/talk/viewthread.php?tid=79548&...

That is of course a great find. But will KHSO4 remain available OTC when H2SO4 won't?
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[*] posted on 5-11-2020 at 04:32


Probably, considering it has commercial use as a fertilizer, pool chemical, food additive etc. Those tend to be difficult to regulate.
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clearly_not_atara
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[*] posted on 20-1-2021 at 16:15


Hypothetically speaking:

Suppose you wanted to use what seems to be the consensus option for conc. H2SO4: you start with 15%-ish H2SO4 and boil it down to around (80%)?, then you feed SO3 through it (via Na2S2O7) until you have roughly 50% oleum, then re-mix with 15% H2SO4 and repeat, etc.

What % do you need to concentrate to in order for it to be safe to pass SO3? You can't pass SO3 into water -- it boils -- but do you need to boil down to 50%? 80%? 90%?




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 21-1-2021 at 11:10


We'll look after 5 years, when EU has noticed that any chemicals left from other methods of banning remain at large, and decides to ban them.

The bisulfate method sounds by far the best method to date. The reagents are bearably priced and available, and the process is high yielding and part of the reactants can be reused. For concentrating the acid, efficient methods should be developed, because in general it is PITA.
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[*] posted on 25-1-2021 at 04:03


I recently seen a video of Epsom salt In a terracotta pot Inside a plastic tub being electrolyzed produced h2so4.cant remember the channel but it was an Australian guy

[Edited on 25-1-2021 by draculic acid69]
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[*] posted on 25-1-2021 at 15:52


It looks like ferric oxide is a great catalyst for oxidation of sulfur dioxide to sulfur trioxide:

Quote:

Ferric oxide possesses the power of catalytically promoting the combination of sulphur dioxide and oxygen at red heat. The action is perceptible at temperatures just above 400° C., attaining a maximum at 625° C. when 70 per cent, of the sulphur dioxide is converted into trioxide. The origin of the ferric oxide is of considerable importance, that prepared from the hydroxide being particularly active. Admixture of copper oxide increases the efficiency, as does also the presence of arsenic at temperatures above 700° C.

source: http://iron.atomistry.com/ferric_oxide.html
I'm not sure about activity of Fe2O3 from rust (I got large amount of that compound from iron scrap by oxidation under water, then roasting it on glowing charcoal) I'll try to use it as catalyst soon.
There is another reaction that I'm focusing on:
Quote:

Sulphur dioxide may be passed into a solution of a ferric salt for a similar purpose, or it may be generated in the solution by addition of an alkali sulphite and a little dilute mineral acid. Thus, ferric sulphate is reduced in accordance with the equation

Fe2(SO4)3 + SO2 + 2H2O = 2FeSO4 + 2H2SO4.
Source: http://iron.atomistry.com/iron_salts.html

The last reaction may be used for recovering sulfur dioxide that's no oxidised on the catalyst.





[Edited on 25-1-2021 by Piroz]
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[*] posted on 31-1-2021 at 17:11


Call me both a risk taker and lazy, but I’d be offering to pickup used car batteries and distilling before going through that much work for that little yield.
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[*] posted on 25-9-2021 at 13:20
A quick and relatively cheap way to make sulfuric acid in EU


The Best way is to mix NaHSO4 with HCl in stochiometric amounts, maybe excess HCl if you want to be on the safer side. Then boil it, let it cool down as much as possible to crystalize the NaCl. Filter out the NaCl. Then Distill the acid. First to come out is the HCl then the Sulfuric acid. Both of the chemicals are cheap to buy in the EU atleast and are a great substitute against other ways of obtaining sulfuric acid. I calculated that with my efficiencies of making I can get almost as low in price to commercially available sulfuric acid in the US. So I think its pretty good apart from the fact that you will always have the paranoia that the round bottom or the flat bottom will crack when you try distilling the acid.
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[*] posted on 25-9-2021 at 13:57


Or you crash out Na2SO4 with methanol/ethanol from a solution of NaHSO4.

Edit: KHSO4 works better.

[Edited on 25-9-2021 by Tsjerk]
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[*] posted on 26-9-2021 at 09:27


But how do these methods help in getting H2SO4? In the EU we still can buy H2SO4 (15%).

Is boiling down 15% H2SO4 more difficult than making the acid from HCl and NaHSO4 or making it from ethanol and KHSO4? Boiling down is a pain in the ass, but I'm afraid that making pure _concentrated_ acid from HCl and NaHSO4 is even more a pain in the ass.

If you really want concentrated acid (e.g. 96%), from chemicals which can be obtained legally in the EU, and without boiling down so much that you reach 96%, then I see only one method:
- boil down the 15% acid until it reaches 85% or so. At that concentration, further boiling will produce intense fumes and you will lose a lot of acid if you want to get it at 95%. This is not practical.
- add chlorosulfonic acid (slowly and carefully, while stirring) to the 85+ % acid. This will increase the concentration to 95% with production of a lot of HCl (which can be absorbed in e.g. methanol to create an interesting side-product).
This method is not for the faint of heart though. Chlorosulfonic acid is EXTREMELY corrosive and the reaction with 85% acid is very exothermic. It must be done very slowly with good stirring. I, however, think that this reaction is less dangerous than distilling or boiling appr. 95% acid.

The above method is nice for making a small amount of concentrated H2SO4, e.g. 50 ml. I would not use it for makin g liters of conc. H2SO4. But making liters of conc. H2SO4 is not wise anyway. Just make what you need for experiments you do now and maybe a few experiments to come. For many more experiments boiled down acid to e.g. 40% is suitable and that is not an issue at all, giving only water vapor and no need to heat to insanely high temperatures like 300 C.

[Edited on 26-9-21 by woelen]




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[*] posted on 6-10-2021 at 14:08


100 ml of 85% H2SO4 contains 26.7g H2O. Let say, we try to get 100% H2SO4 (to keep the math a bit simpler). The reaction of this water with chlorosulfonic acid will require 173g of the acid. It will give 145g H2SO4, so the result will be 296g H2SO4 total. So, we will get 160 ml of 100% H2SO4. So, ~170 g HClSO4 -> 160 ml 100% H2SO4. And for 95%, roughly 85g HClSO4 -> 120 ml (and + 100ml 85% = a lot of 15 % H2SO4). This is not the most economical method.
On the other hand, if you can buy somewhere SiCl4 I would give a try to another method.



[Edited on 6-10-2021 by teodor]
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[*] posted on 10-10-2021 at 09:24


I tried to prepare ~100% H2SO4 by mixing 2.4 ml HClSO3 and 6.5 ml 96% H2SO4. Eventually, it doesn't result in HCl evolution except for several small bubbles. Probably the preparation requires heating. I used the mix to prepare 6M H2SO4 solution in glacial acetic acid. When I've added the acetic acid the solution became warm & highly saturated with HCl and was boiling probably because of HCl evolution. After some time the pressure stopped building. Now it is a fuming liquid with the mixed smell of AcOH and HCl.
So, I think the mixing of chlorosulfonic acid and sulfuric acid doesn't cause a fast reaction and easy HCl elimination if the concentration of H2SO4 is already high.
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[*] posted on 10-10-2021 at 09:46


That's a quite surprising result. If you look at the reaction between water and chlorosulfonic acid, then I would expect that mixing chlorosulfonic acid with something, which contains a little water quickly destroys that little water. Getting rid of dissolved HCl is not an issue, simply heat and you drive it off, but if the destruction of water is not fast and complete with chlorosulfonic acid, then that indeed is a problem.

I have done the prep of 100% H2SO4 from 20% oleum and 96% H2SO4 and that works really nice (although mixing of these still produces noticeable heat) and results in a slightly fuming acid. So, I expected that the reaction with chlorosulfonic acid also would be fast and complete, but apparently the HCl, bound to the SO3, spoils the reaction to quite some extent.




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[*] posted on 10-10-2021 at 23:13


I think the reason is that water is not quite "free" in 96-98% H2SO4. For example, according to https://doi.org/10.1021/ja01630a063 in H2SO4 with a concentration of 90%+ the water exists only in the fully ionized form (H3O+). And probably H3O+ and HClSO3 can coexist. (I will elaborate more on this in my thread about the reaction of chlorosulfonic acid with metals. I have mixtures of 96% H2SO4 and HClSO3 in dichloromethane which I keep in closed bottles for several days - there is no pressure).
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[*] posted on 11-10-2021 at 01:06


If this hypothesis is true, then making pure H2SO4 from HSO3Cl and 90% H2SO4 could be done, if also some SO3 is added:

In 90% H2SO4 nearly all water is present as H3O(+):

H2SO4 + H2O <--->>> HSO4(-) + H3O(+)

SO3 adds to sulfate (and bisulfate): SO3 + HSO4(-) ---> HS2O7(-).
This compound strongly attracts hydrogen ion: HS2O7(-) + H3O(+) --> H2S2O7 + H2O

Then the free H2O can react with HSO3Cl to form H2SO4 + HCl.

This may be interesting from a theoretical point of view, I do not see this as a practical method for preparing H2SO4. Getting SO3 (or strong oleum) is much harder than getting HSO3Cl. At least, oleum is MUCH more expensive, if you can get it at all. I once read the term 'poor man's oleum' for chlorosulfonic acid, as it can be used as a substitute in some special cases, but it definitely is not a snap-in replacement for true oleum. Some people over here succesfully made small quantities of oleum, but making and isolating it safely is HARD. I do have a small quantity of oleum, but I only use that for really special experiments (e.g. making chloryl compounds), because of its very high price and the good chance that I'll never get an opportunity to get my hands on it again.




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[*] posted on 12-10-2021 at 02:55


Last time I bought 30% oleum (you probably can guess where) at the same price per kg as chlorosulfonic acid. But the point is if you can buy oleum you probably can buy 98% H2SO4 also. So, this method is definitely not "scalable".

As per equilibrium, I can cite the book "Sulfur in organic and inorganic chemistry, Volume 2" (A.Senning):
"Sulfuric acid is slightly self-dissociated into sulfur trioxide and water
(1) H2SO4 <-> H2O + SO3
... Water is nearly completely ionized as a base
(2) H2O + H2SO4 <-> H3O(+) + HSO4(-)
Sulfur trioxide is completely converted to disulfuric acid, H2S2O7. This acid is ionized as a moderately weak acid
(3) H2S2O7 + H2SO4 <-> H3SO4(+) + HS2O7(-)
Thus, since the ions H3SO4(+) and HSO4(-) are in equilibrium as a consequence of the autoprotolysis reaction, it follows that the ions H3O+ and HS2O7- must also be in equilibrium
(4) 2H2SO4 <-> H3O(+) + HS2O7(-)
... The complete self-dissociation reaction in the sulfuric acid solvent system can be described then by the above equations". (Then the 4 equilibrium constants are given for 25C).
This matter is a bit hard for me, but probably I would understand that with a help of some visual or mathematical model. Because all 4 equations are in a single system, I am unable to model the situation you have described without the understanding of this 4 equations system first.

As a practical method for a basic laboratory (no distillation), I probably will do some experiments following Tsjerk suggestion of using KHSO4/NaHSO4 and alcohol.

Because getting and isolating SO3 by all known methods is hard in an average home lab I have a plan to continue my experiments about "oxidation of SO2 in pyrosulfuryl chloride" but of course it is only the field of thoughts and experiments yet.

Another field of possible experiments could be separating the mixture of H2SO4 and H3PO4 with a help of some organic solvent. If it is ever possible, it will allow concentrate H2SO4 with P2O5 which is cheaper and more available than chlorosulfonic acid or oleum. But I doubt based on the same considerations about water ionization.

As a practical approach, I see another way. We probably can use 85% H2SO4 + P2O5 in many reactions which require 95%+ H2SO4. Because we know how to use 36% HCl and can avoid using 100% HCl in most cases by some tricks. Now probably is the time to invent more tricks for using 85% H2SO4 instead of concentrated acid.
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[*] posted on 27-10-2021 at 13:11


Quote: Originally posted by woelen  
If this hypothesis is true, then making pure H2SO4 from HSO3Cl and 90% H2SO4 could be done, if also some SO3 is added:

In 90% H2SO4 nearly all water is present as H3O(+):

H2SO4 + H2O <--->>> HSO4(-) + H3O(+)

SO3 adds to sulfate (and bisulfate): SO3 + HSO4(-) ---> HS2O7(-).
This compound strongly attracts hydrogen ion: HS2O7(-) + H3O(+) --> H2S2O7 + H2O

Then the free H2O can react with HSO3Cl to form H2SO4 + HCl.

This may be interesting from a theoretical point of view, I do not see this as a practical method for preparing H2SO4. Getting SO3 (or strong oleum) is much harder than getting HSO3Cl. At least, oleum is MUCH more expensive, if you can get it at all. I once read the term 'poor man's oleum' for chlorosulfonic acid, as it can be used as a substitute in some special cases, but it definitely is not a snap-in replacement for true oleum. Some people over here succesfully made small quantities of oleum, but making and isolating it safely is HARD. I do have a small quantity of oleum, but I only use that for really special experiments (e.g. making chloryl compounds), because of its very high price and the good chance that I'll never get an opportunity to get my hands on it again.


I'm not sure how's that related to the original goal of this thread.
Unless bottles of chlorosulfuric acid are sold at grocery stores and I missed a memo, what's the point if the starting reagent is even harder to get than the final product.
I wouldn't even know where to source that apart from sigma or shady ebay sellers (which would probably sell you h2so4 regardless of the ban).
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[*] posted on 27-10-2021 at 15:34


What about thermal decomposition of NaHSO4?

NaHSO4 decomposes at 280°C to form H2O and at 480°C to form SO3, so the temperatures are achievable. If you bubble the gases in a small amount of water (there's already the stoichimetric amount present in the reaction), you should be able to get quite a high concentration of sulfuric acid.
You also don't need to seperate byproducts.
I don't know how efficient this process would be, but low effiency should be that of a problem, as technical NaHSO4 is available in kg quantities in the hardware store as pH- granules for swimming pools (at least in Germany, I guess it is similar in other (EU) countries).
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[*] posted on 27-10-2021 at 23:45


How about SO2Cl2 to get rid of those last pesky few % of water?
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[*] posted on 28-10-2021 at 00:18


SO2Cl2 reacts with water very slowly. Besides that, it does not dissolve in H2SO4, nor in water, so you only have a reaction at the contact leyer between the two solutions. SO2Cl2 also is not OTC, but making it seems doable for the more advanced amateur.

@zerodan: In my previous post I already wrote myself that the described method is interesting from an academic point of view, but should not be considered a practical method for making H2SO4 in larger quantities.

Where I live, however, chlorosulfonic acid, can be obtained legally, while getting H2SO4 at more than 15% is illegal. The world is a strange place indeed :P but reality for me is that chlorosulfonic acid can be ontained more easily than conc. H2SO4. A few weeks ago I actually purchased some HSO3Cl from a respected supplier (not some shady eBay seller) who sells to private individuals.




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[*] posted on 28-10-2021 at 02:10


Oh, and I thought being able to buy kilos of red phosphorus and a few 100g of KMnO4 but heavily restricted KNO3 in Germany was weird :D
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[*] posted on 28-10-2021 at 03:15


Also, I just ordered p-toluenesulphonic acid. It is still legal to buy it in any quantity. It is quite handy when you need sulphuric acid for catalysis. I wonder if one can use it to make ether out of ethanol the classic way. I will try that when I put my hands on it.
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