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garage chemist
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Sulfuric acid production from sodium sulfate/bisulfate and HCl
In the german forum, someone just wrote that one can obtain H2SO4 by adding Na2SO4 or NaHSO4 to conc. HCl and filtering off the precipitate of NaCl.
I have used this very method of NaCl precipitation about 3 years ago to make HClO4 from NaClO4 cell liquor. It works without any problems. Pour
saturated NaClO4 solution into a large amount of 37% HCl, stir, suction filter the NaCl precipitate, boil in a still to expell the HCl gas (capture by
dissolving it in water), distill off the excess azeotropic HCl and vacuum distill the residue to obtain HClO4 free from residual dissolved salts.
The solubility of NaCl in water strongly decreases with increasing HCl concentration. One can precipitate NaCl from brine by gassing it with HCl. If
another sodium salt is used, the liquid will then contain the corresponding acid, even if that acid is stronger than HCl, as is the case with HClO4.
If someone has a diagram of NaCl solubility in hydrochloric acid of various concentration, please share it with us here.
I don't know why I didn't make the mental connection back then that this method can be used to obtain all kinds of acids, weak and strong, from their
sodium salts in a preparative manner. Including sulfuric acid.
This information now comes three years too late, as it seems that amateur activity in such a basic field of reagent preparation as sulfuric acid
manufacture has very much decreased. But maybe someone will still find this info useful and write about his experience with this method.
A sensible procedure would be to prepare a saturated solution of NaHSO4, pour it into a tenfold excess of conc. HCl, filter off the NaCl and distill
the filtrate (with HCl gas capture) until nothing more comes over and the thermometer in the pot reads over 300°C. The residue would then be conc.
H2SO4.
The only drawback is that the H2SO4 so produced will contain residual NaHSO4. Vacuum distillation would be required to obtain it pure. Whether one can
do this economically is the important question.
Alternatively, the raw filtrate could be gassed with HCl until saturated to further decrease the NaCl solubility. Again, a diagram of NaCl solubility
vs. HCl concentration would be most helpful.
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hissingnoise
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Quote: Originally posted by garage chemist  | | In the german forum, someone just wrote that one can obtain H2SO4 by adding Na2SO4 or NaHSO4 to conc. HCl and filtering off the precipitate of NaCl.
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I can't see this working well if at all---the best you could hope for would be an equilibrium mixture containing sulphate, chloride, HCl and a small
quantity of H2SO4. . .
The difference with HClO4 is that NaCl is stable in its solutions; it won't last long in H2SO4---unfortunately.
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setback
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You can use metabisulfate and HCl to produce SO2 (by slowly adding the HCl drop by drop via an addition funnel). Then bubble the SO2 through an
oxidizing liquid (conc HNO3 or 30% H2O2) in an icebath.
Then you would boil the acid to concentrate it and to remove unreacted SO2. I would do all of this outside.
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Formatik
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| Quote: Originally posted by garage chemist | | In the german forum, someone just wrote that one can obtain H2SO4 by adding Na2SO4 or NaHSO4 to conc. HCl and filtering off the precipitate of NaCl.
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That was me. It was part of my NaHSO4 and NaCl separation thread. http://www.sciencemadness.org/talk/viewthread.php?tid=11490 Kind of hidden in that thread I guess. Verborgene Schätze, halt. I should have tried
a density reading on the distilled acid and that I regret. But I did use the liquid for a qualitative test by reacting with an alkali chlorate to
further prove it was sulfuric acid (reaction, odor is basically the same as with conc. H2SO4). I haven't tried it with the bisulfate. Though should
work also. I used solids in both attempts because I wanted as little NaCl to solubilize in the aq. HCl as possible. I don't know if this is
energetically meaningful method of preparation. Below is a NaCl in HCl solubility table from the Dictionary of Chemical Solubilities I used
to gather some thoughts on the experiment.
[Edited on 27-7-2009 by Formatik]
Attachment: NaCl vs. HCl.pdf (184kB)
This file has been downloaded 1292 times
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DJF90
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setback: I think you are refering to metabisulfIte - Na2S2O5. I dont think there is such a thing as metabisulfAte - I presume the equivalent would be
pyrosulfate - Na2S2O7.
Hissingnoise - I think if the preparation is done with a solution of Sodium bisulfate then the concentration of the formed H2SO4 will be too low to
readily liberate HCl from the precipitated NaCl. It will be the precipitation of the NaCl which will drive the reaction forward. Once all the NaCl has
been filtered out, the solution can be placed in a still as suggested and heated till the thermometer reads about 300*C.
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setback
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| Quote: Originally posted by DJF90 | setback: I think you are refering to metabisulfIte - Na2S2O5. I dont think there is such a thing as metabisulfAte - I presume the equivalent would be
pyrosulfate - Na2S2O7.
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No, you're right, it's metabisulfite. Brain fart, and I should know that one, I use it all the time for brewing  .
I think that process would be a sort of twist on the lead chamber process. Except instead of burning sulfur to get SO2 you are using metabisulfite.
Also, instead of using a catalyst in the gas phase, you are running it through an oxidizing liquid.
You have to concentrate down the acid by boiling, and I've never scaled it up, but and of course you can always buy the acid via biodiesel stores or
whatever. It's cool to know you can make it though.
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setback
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I would do it outside or in a hood, especially if you use HNO3 as the oxidizer.
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hissingnoise
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| Quote: Originally posted by DJF90 | | I think if the preparation is done with a solution of Sodium bisulfate then the concentration of the formed H2SO4 will be too low to readily liberate
HCl from the precipitated NaCl. |
Yes, but concentrating such dilute solutions is so intensive (and Formatik's product so contaminated) that the process is of little more than academic
interest. . .
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Formatik
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The brown material might have just been some humic molecules. The difference on the action on the tissue comes from a higher water content, ergo I
didn't evaporate off enough water.
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JohnWW
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Dictionary Of Chemical Solubilities (1921)
This, dated 1921, can be downloaded from:
http://www.archive.org/details/dictionaryofchem00comerich
or, to give the links:
http://www.archive.org/download/dictionaryofchem00comerich/d... 181 Mb
http://www.archive.org/stream/dictionaryofchem00comerich/dic... 75 Mb
It appears to have a lot of solubility data that the International Critical Tables, the Handbook Of Chemistry & Physics, and Perry's Chemical
Engineers Handbook (see References section) do not have. About 1,180 pages.
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mysteriusbhoice
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You can just electrolyse the Na2SO4 with a salt splitter cell to produce sulfuric acid
im currently producing sulfuric acid from Gypsum CaSO4 using sodium hydroxide as a catalyst
I first react the CaSO4 with NaOH to produce Ca(OH)2 and Na2SO4
then I place the Na2SO4 solution into a Membrane Cell to produce H2SO4 in the anode using Pb electrodes
and NaOH is regenerated in the cathode
CaSO4 + 2NaOh --> Na2SO4 + Ca(OH)2
Na2SO4 + 2H2O + e --> 2NaOH + H2SO4
source:
https://www.google.ch/patents/US5928488
[Edited on 6-10-2017 by mysteriusbhoice]
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chloric1
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| Quote: Originally posted by hissingnoise | | Quote: Originally posted by DJF90 | | I think if the preparation is done with a solution of Sodium bisulfate then the concentration of the formed H2SO4 will be too low to readily liberate
HCl from the precipitated NaCl. |
Yes, but concentrating such dilute solutions is so intensive (and Formatik's product so contaminated) that the process is of little more than academic
interest. . .
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I disagree! I am about to run a third trial, and after all HCl solution is distilled off and white fumes are produced, I get a sulfuric acid roughly
of 79~80% concentration. The level of contamination is corresponding to amount of sodium chloride collected before vacuum filtration. If too much
sodium bisulfate is unreacted, just recycle with more water then add additional sodium bisulfate to saturate and react with HCl again to build up a
“reserve” until a satisfactory yield is obtained. Even a 10% bisulfate contamination is still quite useful for making HCl HBr or nitric acid. I
will be posting a report of my findings shortly!
Fellow molecular manipulator
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teodor
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Hm, let's see what other people already did:
 
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chloric1
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Running second batch now
I have found that the sodium bisulfate needs to be dissolved in the absolute minimum of boiling water. A slight presence of sulfuric acid helps
sodium bisulfate lower its melting point so you can practically liquify almost to a molten salt mixture. Two moles of excess HCl are helpful. The
picture I share is my second batch where I more or less melted sodium bisulfate as well as dissolved in water. It’s funny if you squirt 5 ml into
the 400 ml of 10M (32%) hydrochloric acid nothing happens. Squirt a few more ml and it start snowing cubic sodium chloride crystals! The flask was
so hot so I basically dumped it into the acid after that.
Update-After vacuum filtration through fritted glass Büchner funnel twice, I washed solids with 99% isopropyl alcohol and put into lab oven a little
over 100 degrees C while I slept. This morning the dried salt weighed 100 grams. The crude bisulfate came from my dry HCl run I did over the weekend
to increase my HCl solution to 37% concentration. I started with 120 grams of NaCl so I basically ended up with ~83% return except I found a few
grams of sodium bisulfate crystals in my first batch of reclaimed sulfuric acid so I threw those in instead of wasting them. My acid filtrate is just
under 600 ml so I may want to run some HCl gas to see if more sodium will precipitate as chloride. Or I could distill off maybe 200 ml first since my
distillate will be almost HCl and water anyways and it may mean using less dry HCl to begin with.
[Edited on 9/17/2025 by chloric1]
Fellow molecular manipulator
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teodor
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Interesting, chloric1. Which bisulfate do you use. There could be 4 types of "bisulfate"
1. NaHSO4
2. NaHSO4 * H2O
3. Na3H(SO4)2
4. NaH3(SO4)2 * H2O
What is your source of the salt?
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Sulaiman
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very interesting to follow
I am fascinated by the molecular manipulations being performed,
I liked the preparation of sulphuric acid using the 'weaker' oxalic acid exploiting the oxalate insolubility,
but this is next level !
thanks for sharing.
CAUTION : Hobby Chemist, not Professional or even Amateur
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Texium
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Thread Moved
17-9-2025 at 07:00 |
chloric1
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| Quote: Originally posted by teodor | Interesting, chloric1. Which bisulfate do you use. There could be 4 types of "bisulfate"
1. NaHSO4
2. NaHSO4 * H2O
3. Na3H(SO4)2
4. NaH3(SO4)2 * H2O
What is your source of the salt? |
Pretty sure it was mostly sodium bisulfate with maybe 5% excess sulfuric acid. I’ve noticed such mixtures have much lower melting points,
especially when water is added. Might be close to number 4
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woelen
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Standard pH-minus, which is the common source of bisulfate, is NaHSO4.H2O.
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chloric1
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Correct. The pH minus in the USA is listed as 93% sodium bisulfate the balance being sodium sulfate. I’ve been told the sodium sulfate makes
storage and handling easier. But my first batch of reclaimed sulfuric acid came from lab grade sodium bisulfate that seems to be mostly anhydrous.
My second batch currently in the works is sodium bisulfate product I made from pickling salt and 93% drain cleaner sulfuric acid. I set up a HCl
generator to increase distilled 17% HCl solution to 37% HCl. Sometimes I like to just open the bottle an and watch it fume and take pride that I made
it! Water clear 12 M HCl!
When is was distilling nitric acid before I learned sodium and potassium bisulfates can form not only low melting hydrates but low melting adducts
with free sulfuric acid.
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teodor
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It's always good to have a lab grade material as a reference, you are right.
May be I have to share some additional information. It could help to determine which type of product you get depending on H2SO4 concentration. As you
see, it could be different types of salts and solvates and even mixtures.
Unfortunately I don't have similar pictures for potassium.
I made some calculation yesterday. Given 4 basic types of sodium bisulfate salts (just repeat what I already wrote before):
1) NaHSO4
2) NaHSO4 * H2O
3) Na3H(SO4)2 = NaHSO4 * Na2SO4
4) NaH3(SO4)2 * H2O = NaHSO4 * H2SO4 * H2O
The following amount of 0.5M BaCl2 solution will completely percipitate sulfate from 2.5 g of unknown type of bisulfate salt:
1) 20.8 ml
2) 18.1 ml
3) 38.2 ml
4) 42.3 ml
So, this is basic type of analysis I will try to perform on my bisulfate before attempting to repeat those experiments of yours, chloric1. Any
intermediate values can be for the mixtures.
P.S. I recall our old discussion about peculiarities of some salt and particulary Na2SO4 solubility in water (why different hydrates have different
solubility when recalculated to anhydrous compound). As you see it is the same here. The amount recalculated for anhydrous Na2SO4 which H2SO4 solution
can hold jumps depending on the crystal product which is in equilibrium with the solution.
[Edited on 18-9-2025 by teodor]
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chloric1
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Oh thanks! I will take heed of this information! I have some lab grade BaCl₂ I should make up a solution so I can keep track of phases and
mixtures. Now, by taking total weight vs. sulfate determination by titration, I’d be able to get a better idea of sodium content. In many cases
of reacting sulfuric acid with a sodium salt I do like excess acid in that a lower melting byproduct is easier to handle and clear from flask.
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teodor
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| Quote: Originally posted by chloric1 | | Oh thanks! I will take heed of this information! I have some lab grade BaCl₂ I should make up a solution so I can keep track of phases and
mixtures. Now, by taking total weight vs. sulfate determination by titration, I’d be able to get a better idea of sodium content. In many cases
of reacting sulfuric acid with a sodium salt I do like excess acid in that a lower melting byproduct is easier to handle and clear from flask.
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You are welcome. This is not a complete plan of analysis, only the first rough step. To get the sodium content I think we need to convert sulfate to
chloride (by barium chloride), precipitate barium, filter, evaporate and weight the residue.
It's interesting what you are writing about lowering the melting point by acid/water. Sure, I need to try this. But my first step would be analysis of
starting materials, otherwise I will unable to get reliable data.
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chloric1
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Gotcha! I’ll be able to be more thorough this weekend and I’ll go back to my lab grade sodium bisulfate to establish a control. The hydrochloric
acid is hardware store variety 31.45% or approximately 10 Molar
[Edited on 9/18/2025 by chloric1]
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teodor
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Oops. English wikipedia article about sodium bisulfate is extremely inaccurate.
The real melting point of NaHSO4 is 182C, not 315C.
The decomposition temperature to Na2S2O7 is also much lower. It starts decomposition almost at melting point but slowly, at 240-250C the decomposition
happens during 4 hrs.
315 is another point, where SO3 begins to evolve.
It looks like wikipedia on other languages gives more correct information, at least the melting point.
[Edited on 18-9-2025 by teodor]
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Radiums Lab
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Ill get the info fixed soon.
Water is dangerous if you don't know how to handle it, elemental fluorine (F₂) on the other hand is pretty tame if you know what you are doing.
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