Formatik
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The separation of NaHSO4 from NaCl
Typically, H2SO4 is needed for the preparation of NaHSO4. Formed from H2SO4 with NaCl or NaNO3, etc. But there is a way to get the hydrogen sulfates
without H2SO4.
K2SO4 is soluble in aq HCl and from a warmly-prepared solution, upon cooling crystals of KHSO4 and KCl precipitate. Na2SO4 reacts analogously. Mg, Fe,
Zn sulfates crystallize out of their solutions from dilute HCl unchanged (Ann. 19 [1836] 5).
I've looked for methods of separating these two but haven't found anything. Anyone have any good ideas?
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497
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Wow, that's interesting. It makes sense I guess, but I'd never thought about it working that way. I like the concept though, it could make synthesis
of H2SO4 even more OTC.
Supposedly NaHSO4 will form H2SO4 and Na2SO4 when added to ethanol, but with the presence of NaCl might just give you HCl again?
Also you could try it with ammonium salts, if it works you could precipitate out most of the chloride by adding methanol to the concentrated solution.
Apperently NH4HSO4 is substantially soluble in methanol solutions, but I haven't found any actual data.
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kclo4
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Interesting, I had always suspected that would happen since it seemed like it would be an easily reversible reaction but had never read anything on it
to this day.
I like what 497 has suggested - fractional crystallization.
The problem with that is I think it would be hard to predict the solubility of one salt from another since they are probably going to form an
equilibrium of some sort. Water and an Alcohol as a solvent is probably not possible, and an alcohol will probably just produce HCl and Na2SO4 again.
If it works with HCl, I wonder what other sort of acids it works with, be interesting if someone could get this reaction to work with Carbonic acid -
sounds possible under pressure and with reduced temperature to me at least.
I bet you it is going to be hard to dry the salt mixture with out it fuming off Hydrogen chloride.
If its the Chloride you want, that would make it a lot easier to extract - but I get the feeling it is the bisulfate that is of interest.
The best I have is to perhaps grow large crystals of each that allow them to be separated manually? Or maybe separate them via density some how?
It actually seems pretty difficult to me for such an easy to get chemical.
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Formatik
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| Quote: | Originally posted by 497
Wow, that's interesting. It makes sense I guess, but I'd never thought about it working that way. I like the concept though, it could make synthesis
of H2SO4 even more OTC. |
That is true, also since many sulfates can be used to obtain the alkali sulfates by double-decomposition. E.g. filtering off the MgCO3 or Mg(OH)2 from
mixing respective solutions of Na2CO3 or NaOH with aq MgSO4. I've done those too. Then precipitating alkali sulfate with alcohol, or boiling down the
solution and chilling.
| Quote: | | Supposedly NaHSO4 will form H2SO4 and Na2SO4 when added to ethanol, but with the presence of NaCl might just give you HCl again?
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I think it might also give HCl.
I have also looked up the deal on alkali hydrogen sulfates in alcohol in Gmelin a while ago. Basically: anhydrous alcohol only solubilizes very small
amounts of H2SO4 out of pure KHSO4. Water-containing alcohol decomposes it into K2SO4 and H2SO4. Absolute alcohol removes a part of the acid from dry
NaHSO4. Methanol, n-butanol, amyl alcohol, do the same thing. Repeated treatment with regular alcohol eventually leaves only Na2SO4 behind. Removing
H2SO4 from ethanol would be problematic. There is supposed to be something about the extraction of H2SO4 from KHSO4 and alcohol in Ber. pharm. Ges. 31
[1921] 194 from Th. Sabalitschka.
| Quote: | | Also you could try it with ammonium salts, if it works you could precipitate out most of the chloride by adding methanol to the concentrated solution.
Apperently NH4HSO4 is substantially soluble in methanol solutions, but I haven't found any actual data. |
It would be good if we could find more information on this. The only information I've found on this is that NH4HSO4 is said to be very sl. sol. in
alcohol according this. Usually if an inorganic compound is barely soluble in ethanol, the ease of solubility of it in methanol lurks closely to ethanol.
| Quote: | Originally posted by kclo4
Interesting, I had always suspected that would happen since it seemed like it would be an easily reversible reaction but had never read anything on it
to this day.
I like what 497 has suggested - fractional crystallization.
The problem with that is I think it would be hard to predict the solubility of one salt from another since they are probably going to form an
equilibrium of some sort. Water and an Alcohol as a solvent is probably not possible, and an alcohol will probably just produce HCl and Na2SO4 again.
If it works with HCl, I wonder what other sort of acids it works with, be interesting if someone could get this reaction to work with Carbonic acid -
sounds possible under pressure and with reduced temperature to me at least.
I bet you it is going to be hard to dry the salt mixture with out it fuming off Hydrogen chloride.
If its the Chloride you want, that would make it a lot easier to extract - but I get the feeling it is the bisulfate that is of interest.
The best I have is to perhaps grow large crystals of each that allow them to be separated manually? Or maybe separate them via density some how?
It actually seems pretty difficult to me for such an easy to get chemical. |
It is correct fractional crystallization with water won't work since NaHSO4 is a bit more soluble than NaCl: 50 g per 100 g H2O at 0 to 18.75º,
compared to: 35.75 g per 100 g H2O at 0.35º and 35.84 g per 100g at 15.20º.
HCl is the only acid I've read this works with. HNO3 might work.
Drying the crystallized mixture is not a problem. Just filter through glass and then leave in a ventilated area, or however one wants to dry.
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497
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I was going to say something about NH4HSO4 being substantially more soluble than NH4Cl, making fractional crystallization easier... Then I realized
you could save your HCl and make NH4HSO4 just by thermally decomposing (NH4)2SO4...
At 100*C KCl is 56g/100g versus KHSO4 at 122g/100g. That could probably be fractionally crystallized...
[Edited on 21-11-2008 by 497]
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blogfast25
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| Quote: | Originally posted by kclo4
If it works with HCl, I wonder what other sort of acids it works with, be interesting if someone could get this reaction to work with Carbonic acid -
sounds possible under pressure and with reduced temperature to me at least. |
That's in essence the first step of the Solvay process, but buffered with ammonia...
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Formatik
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| Quote: | Originally posted by 497
I was going to say something about NH4HSO4 being substantially more soluble than NH4Cl, making fractional crystallization easier... Then I realized
you could save your HCl and make NH4HSO4 just by thermally decomposing (NH4)2SO4... |
I think you may be onto something with the ammonium salts. Following USP3282646 and this article, heating (NH4)2SO4 between 200 to 300º evolves NH3 and forms NH4HSO4: 3 (NH4)2SO4 -> 3 NH4HSO4 + 3 NH3. The older Gmelin 7. Aufl.,
III, Tl.2, 545 describes decomposition begins at 280º and that this reaction attacks glass. Then further heating of the acid sulfate between 350 to
530º forms: 3 NH4HSO4 -> NH3 + 3 SO2 + 6 H2O + N2. That's also a source for SO2 like the patent describes.
Perhaps one could exchange bisulfate from ammonium bisulfate through potassium orthophosphate (K3PO4) which is very soluble in water, where the
(NH4)3PO4.3H2O is difficultly soluble.
| Quote: | | At 100*C KCl is 56g/100g versus KHSO4 at 122g/100g. That could probably be fractionally crystallized... |
This data shows part of the KCl will crystallize out, but KHSO4 remains in solution still largely contaminated. It is the same problem as by the
sodium salt mixture.
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kclo4
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| Quote: | | Quote: | Originally posted by Formatik
| Quote: | | At 100*C KCl is 56g/100g versus KHSO4 at 122g/100g. That could probably be fractionally crystallized... |
This data shows part of the KCl will crystallize out, but KHSO4 remains in solution still largely contaminated. It is the same problem as by the
sodium salt mixture. |
Actually, at those temperatures, and with all the salt dissolved you'd be kicking off major amounts of Hydrochloric acid/hydrogen chloride.
I've mixed Sodium Bisulfate with Potassium, or Ammonium Nitrate and distilled a decent amount of semi-dilute Nitric acid from it.
Now that's a thought! What if you could some how extract Nitric acid by mixing Hydrochloric acid, and the Nitrate and Sulfate salts of Potassium,
Sodium or Ammonium? - Sorry, that is a bit off topic I suppose.
[Quote]I was going to say something about NH4HSO4 being substantially more soluble than NH4Cl, making fractional crystallization easier... Then I
realized you could save your HCl and make NH4HSO4 just by thermally decomposing (NH4)2SO4... |
I've attempted the thermal decomposition of Fertilizer grade Ammonium Sulfate and it seemed very hard to achieve in any useable amount. The Ammonium
sulfate didn't ever melt fully, and then the Ammonium Bisulfate that seemed to have formed started to sublimate. In the end, I had nothing but a flask
full of junk that cracked when I tried to clean it.
I was hoping this would allow me to produce a decent amount of Ammonium Bisulfate to produce Nitric acid with via distilling with Ammonium Nitrate in
solution, and then also to collect the Ammonia formed. Seemed like a decent way to get two useful chemicals with out using very many chemicals. The
Ammonium sulfate formed from the producing the nitric acid could then be dangerously reused.
Anyways, I don't know if you have ever had luck with decomposing Ammonium Sulfate into Bisulfate, but I find it to be very tedious. So you might
actually be onto something by using these salts for the production of bisulfate via the fractional crystallization method mentioned. Plus, Ammonium
Chloride is nice to have around, it has a decent amount of uses. 
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Formatik
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Gmelin also mentioned sublimation when heating the sulfate. From what they are saying, it looks like the sublimation probably happens after the large
amounts of NH3 is driven off.
[Edited on 22-11-2008 by Formatik]
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kclo4
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Yeah it does, the reaction goes like this:
(NH4)2SO4 <=> NH4HSO4 + NH3
NH4HSO4 <=>H2O + NH3 + SO3
But the problem is you get localized heating, and so the bottom of your reaction vessel is producing NH3, H2O, and SO3, while the top hasn't even
melted and the middle section is producing NH3 and NH4SO4. So, it doesn't surprise me that Gmelin mentions the sulfate sublimating, in a perfect
world it wouldn't, but it would be hard to get it to decompose into the bisulfate with out it decomposing thus producing Ammonium Sulfate again.
Unless of course your working on a small scale using only a few grams, then I think you could pull it off.
Perhaps one could find a substance that helps it melt at a lower temperature and then have constant agitation as it was heated to its decomposition
temperature preventing the localized heating problem.
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S.C. Wack
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Looking at the older literature (Mellor vol 2), it seems that the chemistry of bisulfate might not be simple and straightforward. Perhaps there is a
good modern review out there that could lead to reference chasing.
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Formatik
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To get even heating distribution, use a heating mantle. But don't use a flask for this, too precious since it will get attacked.
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kclo4
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I've Got It!.. I bet
React the Potassium Sulfate with the concentrated Hydrogen Chloride. Then add the proper amount of Calcium Sulfate and allow it to react with the
Potassium Chloride and form Potassium Sulfate and Calcium Chloride.
Now you should have a solution of Potassium Bisulfate, Potassium Suflate, and Calcium Chloride, right? (the only reason I can see why this wouldn't
happen is if this turned everything into a sulfate and left Hydrochloric acid out, but with the potassium ions, I don't think it will.)
Calcium Chloride is highly soluble, while the Bisulfate and sulfate of Potassium are significantly less soluble. This will allow you to precipitate
both of them leaving the Calcium Chloride in solution.
Once you have precipitated the Potassium Sulfates, and have rid it from any remaining chlorides, you will be able to dissolve them into a hot
concentrated solution with out any fear of them reacting with each other. This will alow you to fractionally crystallize them from one another.
You can then of course re-use the extracted Potassium Sulfate from above to produce more Bisulfate.
Magnesium Sulfate might work as well.
If that doesn't work, go to Walmart or somewhere and get pH down, Sani-flush prills and use the Sodium Bisulfate found in that. 
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497
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| Quote: |
Then add the proper amount of Calcium Sulfate and allow it to react with the Potassium Chloride and form Potassium Sulfate and Calcium Chloride.
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I don't think that will work... Think about it, the reverse reaction happens. Try adding some CaCl2 to K2SO4 and you'll see.
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kclo4
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| Quote: | Originally posted by 497
| Quote: |
Then add the proper amount of Calcium Sulfate and allow it to react with the Potassium Chloride and form Potassium Sulfate and Calcium Chloride.
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I don't think that will work... Think about it, the reverse reaction happens. Try adding some CaCl2 to K2SO4 and you'll see. |
Are you sure? I've never done it, but why would it do that? Seems like I must be missing something here  I know reactions like this can happen with lead salts and the like, but I though Calcium Sulfate might be soluble
enough to not just be kicked out of the reaction.
Care to explain?
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smuv
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| Quote: | | I though Calcium Sulfate might be soluble enough to not just be kicked out of the reaction. |
Calcium sulfate is not very soluble at all. This is just a simple double replacement reaction. I have used a similar process to make large
quantities of calcium sulfate from sodium sulfate and calcium chloride...then I found out that these days blackboard 'chalk' is nearly pure calcium
sulfate...
"Titanium tetrachloride…You sly temptress." --Walter Bishop
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kclo4
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Yeah I know its a double replacement reaction, but I was really just focusing on that potassium is more reactive then Calcium, well dang with this
method I guess Magnesium is the only hope - oh well.
Smuv, Plaster of Paris is mostly Calcium Sulfate and is probably a lot cheaper then using Chalk. Shouldn't be to hard to purify it at all since I
think it contains Calcium Hydroxide as the other Ingredient in relatively low amounts.
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blogfast25
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@ kclo4:
Is it remotely possible that you're confusing electrochemical displacement reactions a la:
3 Cu<sup>2+</sup> (aq) + 2 Al (s) ---> 3 Cu (s) + 2 Al<sup>3+</sup> (aq)
with displacement reactions based on solubility a la:
K2SO4 (aq) + CaCl2 (aq) ---> 2 KCl (aq) + CaSO4 (s)
The only way to get rid of the KCl (in Formatik's interesting experiment) by means of a displacement reaction is to add something that precipitates as
an insoluble chloride, but doesn't precipitate as bisulphate or sulphate.
****
What you have in mind with plaster is no more or less possible with plaster than it is with Epsom salt (MgSO4), trust me...
No-frills brands of plaster (DryWall, wall-filler etc) usually don't contain any slaked lime at all. A simple pH test can reveal that...
[Edited on 24-11-2008 by blogfast25]
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Picric-A
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Washing plaster of paris with around 5% H2SO4 then drying produces nearly pure CaSO4 as Smuv says.
NaHSO4 can be got cheaply from pool supply stores at around £10 for 5kg ,
I am going to look into extracting H2SO4 from ethanol. Shurely even simply distilled methylated spirit would work?
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smuv
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@kclo4, when something is there to displace the equilibrium reactivity of the metal doesn't really make a difference.
@Picric A
Are you talking about the disproportionation of bisulfate in ethanol to sulfate and sulfuric acid?
It would be extremely hard to obtain pure H2SO4 from this, because in an ethanolic solution, H2SO4 exists as ethyl hydrogen sulfate and diethyl
sulfate. Trying to distill off the ethanol will produce ether...
There is a patent floating around here that uses Bisulfate in ethanol to produce ethyl hydrogen sulfate; I have tried it a few times with mixed
success. If you still want to give it a try, I recommend grinding the tiny pellets of pH down into a fine powder before proceeding (to increase
surface area).
"Titanium tetrachloride…You sly temptress." --Walter Bishop
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Formatik
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There are also salts which will precipitate other soluble salts out of solution. This is said to happen with K2SO4 and (NH4)2SO4, when adding one of
them to a saturated solution of the other, it dissolves precipitating the other, described here. If there could be some sulfate or other salt added which bumps out most of the less soluble chloride from a solution without interfering
with the bisulfate, that would be great.
But the route through ammonium salts holds more promise so far.
Gmelin mentions heating NH4HSO4 with NaCl is a method to the bisulfate:
NH4HSO4 + NaCl = NH4NaSO4 + HCl;
NH4NaSO4 = NH3 + NaHSO4
But this mixture needs to be heated around 400º to get rid of the NH3. For KHSO4, at a temp. between 200 to 400º: K2SO4 + NH4HSO4 <- -> 2
KHSO4 + NH3. The equilibrium goes to right with increasing temperature.
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Picric-A
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@ smuv - Yes all i have found so far are methods of using it to make ether and ethene ect...
So far, i have come to the conclusion that the only easy method is decomposing it and passing the SO3/H2SO4 through Conc H2SO4 followed by diluting
(or not,).
On wikipedia it states:
'Adding ethanol to a solution of potassium bisulfate precipitates out sulfate.'
I wonder if this works with NaHSO4...
Will report back on any findings.
[Edited on 26-11-2008 by Picric-A]
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Formatik
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| Quote: | | Originally posted by Formatik ... Removing H2SO4 from ethanol would be problematic. There is supposed to be something about the extraction of
H2SO4 from KHSO4 and alcohol in Ber. pharm. Ges. 31 [1921] 194 from Th. Sabalitschka. |
I haven't since found this reference anywhere.
Forgot where, but somewhere in Beilstein it's also stated ethysulfuric acid hydrolyzes: C2H5.HSO4 + H2O = H2SO4 + C2H5OH
So an idea of having some water mixed with ethanolic ethylsulfuric acid, and then distilling I think should leave behind mainly H2SO4 (plus a little
Na or K salt).
It looks like there was a related thread here some time ago: http://sciencemadness.org/talk/viewthread.php?tid=273
[Edited on 17-3-2009 by Formatik]
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Formatik
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Starting with 8.5 g finely powdered hydrated Na2SO4 (likely Glauber's salt) which was slightly impure with Mg(OH)2. Adding the Na2SO4 to 50 mL 31% HCl
caused an endotherm. This was shaken very vigorously for some time. Then allowed to stand until the solids settled and the liquid siphoned off. The
mostly clear liquid then evaporated on low heat eventually yielding a dark brown liquid with black particulates and a strange odor (maybe contaminants
from previous filtering and pulverizing). Heating the dark brown liquid on a stronger heat produces dense white fumes (unlike HCl). Letting it cool to
room temperature, and at 19 deg. it remained liquid. Viscous in nature (like conc. H2SO4), allowing contact to a cool surface caused some
crystallization. Mass: 3.1g. Taking this now liquid gel and applying it to a paper tissue caused it to slowly eat holes in it just like conc. H2SO4,
but at a slower rate (something neither oversatd. aq. NaHSO4 or conc. HCl can do). More contact with cool surface caused a lot more solidification.
And adding a small amount of H2O to this caused an exotherm, spiking at 32 deg.
Though I have not attempted to optimize on the amounts, to me this seems plain useless for any preparation. But I still think its interesting that any
H2SO4 formation occurred at all, which I wouldn't have thought very likely. I was able to form some impure conc. H2SO4 in the same way before but
neglected the results and attributed it to impurities. I decided to attempt to replicate the results after reading this paragraph from A textbook
of Chemistry by W.A. Noyes:
[Edited on 2-7-2009 by Formatik]
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Formatik
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I've repeated the experiment because I wanted to see if it makes a difference if the Na2SO4 is anhydrous or not. I also wanted to characterize the
sulfuric acid a bit better, i.e. yield. So basically, 7.4g anhydrous powdered Na2SO4 was added to 30 mL 31% HCl and this then shaken violently for
some time. Mixing caused no endotherm. It was let stand for around 40 minutes for the solids to settle. Then treated the same as above. But it was
thereafter put into a 500mL distillation flask using a long dropper. This was then heated under flame for some time, where a white solid remained
behind. Yield: about 2.3g of a dark brown, oily liquid (H2SO4). This didn't solidify even when chilled in ice. Though I question its purity still
because some of the same brown material distilled over, and it didn't eat holes in tissue as fast as commercial conc. H2SO4.
[Edited on 4-7-2009 by Formatik]
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