Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
Author: Subject: Sulfuric acid production from sodium sulfate/bisulfate and HCl
garage chemist
chemical wizard
*****




Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline

Mood: No Mood

[*] posted on 27-7-2009 at 02:58
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.






www.versuchschemie.de
Das aktivste deutsche Chemieforum!
View user's profile View All Posts By User
hissingnoise
International Hazard
*****




Posts: 3940
Registered: 26-12-2002
Member Is Offline

Mood: Pulverulescent!

[*] posted on 27-7-2009 at 06:42


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.


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.
View user's profile View All Posts By User
setback
Hazard to Self
**




Posts: 50
Registered: 17-5-2009
Member Is Offline

Mood: No Mood

[*] posted on 27-7-2009 at 07:53


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.
View user's profile View All Posts By User
Formatik
National Hazard
****




Posts: 927
Registered: 25-3-2008
Member Is Offline

Mood: equilibrium

[*] posted on 27-7-2009 at 09:58


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.


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 1036 times

View user's profile View All Posts By User
DJF90
International Hazard
*****




Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline

Mood: No Mood

[*] posted on 27-7-2009 at 14:55


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.
View user's profile View All Posts By User
setback
Hazard to Self
**




Posts: 50
Registered: 17-5-2009
Member Is Offline

Mood: No Mood

[*] posted on 27-7-2009 at 16:48


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.


No, you're right, it's metabisulfite. Brain fart, and I should know that one, I use it all the time for brewing :P.

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.
View user's profile View All Posts By User
setback
Hazard to Self
**




Posts: 50
Registered: 17-5-2009
Member Is Offline

Mood: No Mood

[*] posted on 27-7-2009 at 16:51


I would do it outside or in a hood, especially if you use HNO3 as the oxidizer.
View user's profile View All Posts By User
hissingnoise
International Hazard
*****




Posts: 3940
Registered: 26-12-2002
Member Is Offline

Mood: Pulverulescent!

[*] posted on 28-7-2009 at 02:47


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. . .
View user's profile View All Posts By User
Formatik
National Hazard
****




Posts: 927
Registered: 25-3-2008
Member Is Offline

Mood: equilibrium

[*] posted on 28-7-2009 at 13:25


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.
View user's profile View All Posts By User
JohnWW
International Hazard
*****




Posts: 2849
Registered: 27-7-2004
Location: New Zealand
Member Is Offline

Mood: No Mood

[*] posted on 28-7-2009 at 14:17
Dictionary Of Chemical Solubilities (1921)


Quote: Originally posted by Formatik  
(cut)The Dictionary of Chemical Solubilities
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.
View user's profile View All Posts By User
mysteriusbhoice
Hazard to Others
***




Posts: 477
Registered: 27-1-2016
Member Is Offline

Mood: Became chemistry catboy Vtuber Nyaa

[*] posted on 6-10-2017 at 07:08


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]
View user's profile Visit user's homepage View All Posts By User

  Go To Top