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 985 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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