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Author: Subject: Recovery of K2SO4 from K2SO4/H2SO4/HNO3
freedompyro
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[*] posted on 2-1-2012 at 04:29
Recovery of K2SO4 from K2SO4/H2SO4/HNO3


I was wondering if anyone knew of a way to recover some of the waste products of making HNO3 from a nitrate salt.

I happen to have use for these Sulfates of Potassium, Sodium, Strontium, etc... But I'm not sure how to purify them from the remaining H2SO4 and trace amount of HNO3. Possibly just adding water to the solution and allowing it to evaporate slowly would create large sulfate crystals of decent purity that could be filtered out?
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[*] posted on 2-1-2012 at 05:01


You can concentrate the solution somewhat but not to dryness as H2SO4 is not a volatile compound. Upon cooling (in an ice bath) crystals of K2SO4 will deposit and can be filtered. Further recrystallizations from water will improve the purity.

[Edited on 2-1-2012 by kavu]
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[*] posted on 2-1-2012 at 05:02


Distill off the acids, crystals will stay. And they are probably (mostly) hydrogensulfates, not sulfates.



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[*] posted on 2-1-2012 at 05:26


Heating converts bisulphate to sulphate - small point, but still . . .


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[*] posted on 2-1-2012 at 05:32


Quote: Originally posted by hissingnoise  
Heating converts bisulphate to sulphate - small point, but still . . .




Maybe at around 350°C, but I am sure that boiling point of the acids is far lower than this :)




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[*] posted on 2-1-2012 at 05:38


337C for sulfuric, but anyways, just concentrate, cool slowly, and decant. Dissolve crystals in water and recrystalize. No need to distill sulfuric acid to get patassium (bi)sulfate from the mix.



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[*] posted on 2-1-2012 at 05:39


Quote: Originally posted by hissingnoise  
Heating converts bisulphate to sulphate - small point, but still . . .

Not according to garage chemist:

Quote:

Sodium hydrogen sulfate gives off water when heated at 300- 500°C and turns into sodium pyrosulfate:

2 NaHSO4 -----> Na2S2O7 + H2O


Quote: Originally posted by Adas  

Maybe at around 350°C, but I am sure that boiling point of the acids is far lower than this :)


Well, don't be so sure. Or at least make sure you're sure before saying you're sure - IE. google it or check Wikipedia before doing stupid mistakes. Sulfuric acid boils (with some decomposition) at 337 degrees C.




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[*] posted on 2-1-2012 at 06:52


Quote:

Sodium hydrogen sulfate gives off water when heated at 300- 500°C and turns into sodium pyrosulfate:
2 NaHSO4 -----> Na2S2O7 + H2O

Not in the presence of H<sub>2</sub>SO<sub>4</sub>, it doesn't!


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[*] posted on 2-1-2012 at 07:11


Quote:
Sulfuric acid boils (with some decomposition) at 337 degrees C.

One more small point ─ sulphuric acid boils @ ~290°C
It's the azeotrope that has a boiling point of ~339°C


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[*] posted on 2-1-2012 at 08:04


Moved from organic chemistry to general chemistry.



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[*] posted on 2-1-2012 at 14:46


As K2SO4 has a melting point of 1069 °C... I think I'm safe to just heat the crap out of it in a pyrex flask with a propane burner and turn it off when the H2SO4 stops subliming.

I found this on wikipedia:
"Potassium hydrogen sulfate or bisulfate, KHSO4, is readily produced by mixing K2SO4 with an equivalent number of moles of sulfuric acid. It forms rhombic pyramids, which melt at 197 °C. It dissolves in three parts of water at 0 °C. The solution behaves much as if its two congeners, K2SO4 and H2SO4, were present side by side of each other uncombined; an excess of ethanol the precipitates normal sulfate (with little bisulfate) with excess acid remaining."

[Edited on 2-1-2012 by freedompyro]
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[*] posted on 2-1-2012 at 17:05


Quote:
As K2SO4 has a melting point of 1069 °C... I think I'm safe to just heat the crap out of it in a pyrex flask with a propane burner and turn it off when the H2SO4 stops subliming.

- Look FP, I don't know how far you intend heating your flask, but 'heating the crap' out of a borosilicate vessel sure sounds pretty ominous to me!
And I don't know where you're getting your numbers, but anhydrous sodium sulphate has a melting point of ~886°C --- and the fully hydrated salt melts at temperatures not much above 30°C!
Are you, perhaps, mixing up Fahrenheit and Celsius?

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[*] posted on 2-1-2012 at 17:21


And BTW, where on earth did you get the idea that H<sub>2</sub>SO<sub>4</sub> sublimes?
It simply evaporates, as do the other mineral acids!

[Edited on 3-1-2012 by hissingnoise]
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[*] posted on 2-1-2012 at 23:35


All that information on boiling points was taken off the potassium sulfate wikipedia page.

With the boiling point of H2SO4 the borosilicate vessel should only reach 350-375C max. I had read somewhere before that H2SO4 didn't boil and just started subliming giant clouds of toxic sulfuric acid mist as it nears it's boiling point. I guess that info is wrong, and the stuff actually SO3. Kinda surprised the stuff actually boils. Never heated it quite that high before.
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[*] posted on 3-1-2012 at 13:33


Quote:
And I don't know where you're getting your numbers, but anhydrous sodium sulphate has a melting point of ~886°C ---

Ooops!
Quote:
As K2SO4 has a melting point of 1069 °C.

My apologies FP; for some silly reason I'd assumed you were referring to the sodium salt melting-point.
So let's just agree to blame Lambda-Eyde for bringing sodium into the equation in the first place? :D
As to the fragility, or otherwise, of your flask --- you might get away with heating it to over 300°C, but any slight draught as it cools will almost certainly cause it to shatter!
Borosilicate glass can actually be heated to ~500°C for short periods, but this requires a longish annealing process!


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[*] posted on 3-1-2012 at 14:50


It's possible that I've posted this info as hissingnoise, at some time, but, in industrial nitration processes, when spent sulphuric acid needed only to be reconcentrated after denitration, the sulphuric acid was boiled in silicon iron pots over a strong heat source!
Because of the relatively poor heat conductance of the metal, the pots sometimes failed catastrophically, dumping boiling acid into the furnace!
However, I've often reconcentrated sulphuric acid in heavy-walled casserole dishes without incident!
Quickfit RBFs are somewhat less replaceable . . .
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[*] posted on 3-1-2012 at 21:31


BR8805833A states that KHSO4 could be ppted as K2SO4 by adding ethanol to KHSO4 solution (it doesnt talk about other alcohols, but might work as well) , since the equillibrium

2 KHSO4 <--> K2SO4 + H2SO4

is pulled to the right, because K2SO4 is much less soluble in ethanolic solutions than KHSO4.. Since alcohol dillution and K2SO4 ppting release good amounts of heat, alcohol is gradually added (and with stirring) and temperature must be about room temp (~25°C) since otherwise the K2SO4 will ppt as 'pseudo-doble salt' K2SO4*nH2SO4 (where n is 1/3, 1/2, 3/4, 3/5, 3, 6/7, 4 or 6) and in this case you will have new mess to play with in order to get pure K2SO4..

The patent deals with LOW TEMPERATURE recover of K2SO4 from chemical reactors ( KCl + H2SO4 ---> KHSO4 + HCl ), just to dont mess with high temp corrosive acids in reactors.. The alcohol can be distilled and recovered and according with paper, only 2% of new alcohol must be added to make up losses in closed system.

Of course if that is presumed to be adapted to spent nitration mixed acids, you have to get rid of residual HNO3/NOx (with will react with alcohol; nasty), and still if dont react, it can ppt along with K2SO4 when adding alcohol.. This is the boring part (or maybe not if you actually want to eliminate nitrate bringing it to high temp fuzzing flame of a evil heavy metal torch in glassware :D just be sure to sign first a proper insurance contract though.. )




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[*] posted on 9-1-2012 at 14:43


Today I was cleaning my homelab and found a somewhat old bottle containing some KHSO4 from spent nitrating mix and I know there was no (or might be little) nitrate since in the procedure Ive gently boiled until NOx disapeared..

Tested the procedure of ethanol.

Added some water to the bottle.. You can see the large KHSO4 crystals in the bottom. (I had to break it since it formed a large mass of solidified KHSO4).





KHSO4 and ethanol (home cleaning grade - 92,8° INMP):





Ive heated somewhat the bottle and left it to cool to room temperature (A beaker was used as water bath to speed up that):





Measured 100mL of alcohol; six ~10mL portions were added with stirring and between these additions, it was let cool (the temp doesnt change much.. After each adition, it was found with a thermometer in the bottle that the temp rose from 25 to max of only 27°C).. After the first addition, crystals started to ppt:







After the fifth addition there was no more increase in ppt volume, so a last sixth portion of alcohol was added and then filtered.. Like before, the remaining portion (~40mL) were used in ~10mL portions to wash the K2SO4 on filter:





The mass of coulple of grams of free flowing white crystals of K2SO4 were dried, stored in plastic container, labbeled and stored.







A bit of this K2SO4 was dissolved in water and some NaHCO3 added. Not visible reaction...

So this would be a interesting way to recover potassium values from your spent nitrating mix.. And of course you could recover ethanol by distillation, left behind H2SO4 (probably at some moment this could make some impurities like diethyl ether and such, but I dont know).. Or you could give it another uses to filtered liquid (neutralizing some nasty high pH spills, or just use it as toilet cleaner..)..


EDIT: Unfortunatelly I cant tell the actual yield since I didnt knew the KHSO4 contents of spent nitration mix (As Ive said, this was made long ago, and dont recorded the amount of material on label).. But the 'chem' feel was that given the amount of KHSO4 crystals, the yields were aceptable/good..

[Edited on 9-1-2012 by Aqua_Fortis_100%]




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[*] posted on 10-1-2012 at 02:57


Peach Syndrome? :D

P
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[*] posted on 10-1-2012 at 03:47


Quote: Originally posted by Pulverulescent  
Peach Syndrome? :D

P


I really hope it's contagious.




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[*] posted on 10-1-2012 at 06:33


Quote:
I really hope it's contagious.

What? Pictures too big, so numerous, and taking up space better used? :P:D;)

P
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[*] posted on 11-1-2012 at 02:32


Very nice work. Thanks a ton Aqua Fortis!
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[*] posted on 11-1-2012 at 10:54


Quote: Originally posted by Pulverulescent  


Peach Syndrome? :D

P


Im SMDB member since end of 2006, but have reading this great forum earlier..IIRC end of 2005..And I sincerelly hope to stay continuing as member, helping and being help by the SMDB community.
From this time until today, in many threads I have read, noticed that there is a great trend: when someone did an experiment occasionally other people ask for pics/videos/schemes for several reasons, most of all IMHO, to have better understanding of the experiment itself (the more exciting an experiment is, the bigger the trend to people ask for pics).

You is the first one which I see that dont like pictures.
Whats wrong? The size of my pics (500x600) is well under the SMDB recommended max size, gives good visibility and you dont have to scroll any horizontal bar to read this thread (unless your computer is disconfigured), what is a very tedious when happen in some threads. Also, I dont like to post tiny 'click to enlarge' pics because when pics are many, is also too tedious to see every pic individually and takes an additional time of your life..


Quote:

taking up space better used? :P:D;)

P


like this http://www.sciencemadness.org/talk/viewthread.php?tid=18419 ? A whole thread to speak about ... ?

Sorry. Hope you dont take this as an agressive revenge. I enjoy your posts Pulverulescent/hissingnoise (specially in energetics section), but you should have to understand/agree that Im just defending myself, like anyone would do.

Peace and cheers.


Back to thread:

Hello freedompyro, thank you too, this thread was an inspiration to me to search and try out the recovery of K2SO4. Fert grade K2SO4 is substantially more expensive than KCl (as K2SO4 is made from it) and is becoming more and more as agricultural potassium needs rises and sedimentary KCl deposits are depleted.. And besides this, as K2SO4 can be used to make KOH/K2CO3, is interesting to not waste any K[+].


Has anyone more ideas on homelab uses for the spent alcoholic acid mix other than distill/use as toilet cleaner? Both ethanol and H2SO4 are very useful substances to just throw off..And to distill one has to waste some energy... Direct uses are interesting to me.


[Edited on 11-1-2012 by Aqua_Fortis_100%]




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[*] posted on 11-1-2012 at 13:38


Quote:
A whole thread to speak about ... ?

Touché!
. . . Yes AF, it is the ever-pretentious, moi!
It was, though, in Misc., (a feeble defense?) and my smartass comment was intended as witty more than anything!
I enjoy all your posts too, BTW!
And you do take the trouble 'some' of us seem to find too time-consuming!
So all in all, I wouldn't have you change a thing!

P

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