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Author: Subject: Crystallisation of divalent metal sulfates in excess of sulfuric acid
Vylletra Heart
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[*] posted on 21-12-2017 at 09:44
Crystallisation of divalent metal sulfates in excess of sulfuric acid


Hello people, I am currently in the midst of making some nickel sulfate(NiSO4) to crystallise and I was wondering if the acidity of the solution might affect the resulting composition of the crystals.

You see, sulfate salts of monovalent ions such as NH4+ and the alkali metals have the ability to exist as either M2SO4 or MHSO4, where M is the monovalent ion. I would expect that both types of sulfates have very different crystal structures as well.

I was wondering, since monovalent cations have the ability to exhibit this property, would cations of other charges such as 2+ and 3+ form bisulfates as well? I've been searching up the web and there has been nothing about bisulfates of metals of charges larger that +1.

The reason for this question is that I am planning to make large single crystals of nickel sulfate, and wish to save on my usage of nickel oxide so I am using a large excess of sulfuric acid in my production of nickel sulfate. I fear that the excess acid can protonate sulfate ions to form bisulfate ions. This may end up forming Ni(HSO4)2, which I do not want. I would really like to make pure crystals of NiSO4 so I am hoping that this does not happen. Would there be any telltale signs if it does? Or am I just being anxious over nothing and what I've just mentioned will never happen?

Simply put, I don't want to have 2 different crystal structures growing during crystallisation of the solution at constant temperature(at 25C where I live, if it helps). Thanks for reading this really long post!
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aga
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[*] posted on 21-12-2017 at 14:04


Which previously published paper are you following ?

Any of these ?

https://www.google.com/patents/US1936829
https://www.google.com/patents/US7364717
https://www.youtube.com/watch?v=SEtOiW4Lo8Q

Or have you got some of the rare mineral Retgersite ?

(I just googled those)

... or are you just mixing it all up in a bucket with the eggs on top ?

https://www.youtube.com/watch?v=aczPDGC3f8U




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kmno4
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[*] posted on 21-12-2017 at 14:32


Quote: Originally posted by Vylletra Heart  
I fear that the excess acid can protonate sulfate ions to form bisulfate ions.

In your 25 C world it will not happen.
See the paper:
Code:
http://pubs.acs.org/doi/abs/10.1021/j100898a014

.... and have pity on forum idiots, just wanting to show you "I exist" ;)




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Vylletra Heart
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[*] posted on 22-12-2017 at 06:12


Quote: Originally posted by aga  
Which previously published paper are you following ?

Any of these ?

https://www.google.com/patents/US1936829
https://www.google.com/patents/US7364717
https://www.youtube.com/watch?v=SEtOiW4Lo8Q

Or have you got some of the rare mineral Retgersite ?

(I just googled those)

... or are you just mixing it all up in a bucket with the eggs on top ?

https://www.youtube.com/watch?v=aczPDGC3f8U


Judging by the way you're answering me I assume you're mocking me. Anyway, when I made the nickel sulfate solution it was just a simple process of adding nickel oxide to concentrated sulfuric acid, and I do not think that I need to follow a paper to understand that. I am a student so I do understand what happens when you add a basic oxide to acid. The question I was asking was referring specifically to the existence of nickel bisulfate crystals, Ni(HSO4)2.

If you did give me those papers as genuine references, then I'm sorry for not carefully reading them as I have a hectic schedule. By quickly browsing through them, they seem to be about the dissolution of nickel alloys or related compounds in acid, which I do not have any issues with.

Nevertheless, thanks for trying to help.
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Vylletra Heart
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[*] posted on 22-12-2017 at 06:40


Quote: Originally posted by kmno4  
Quote: Originally posted by Vylletra Heart  
I fear that the excess acid can protonate sulfate ions to form bisulfate ions.

In your 25 C world it will not happen.
See the paper:
Code:
http://pubs.acs.org/doi/abs/10.1021/j100898a014

.... and have pity on forum idiots, just wanting to show you "I exist" ;)


Thanks for the heads up! I am currently a bit busy with other stuff at the moment so I will get about to reading it in the future.

Would I be right in saying that even in a solution of pure sulfuric acid, there would still be an insignificant amount of bisulfate ions? I read up the pKa of sulfuric acid and it seems that both values seem to be very low(hence a high likelihood of deprotonation).

So if we applied this knowledge to an extreme situation, let's say, dissolving 1 mole of anhydrous NiSO4 into 1 mole of pure anhydrous sulfuric acid(this would increase the amount of bisulfate species), would a crystallisation still result in only NiSO4 crystals formed?

Anyway, what do you mean by forum idiots? I'm still new to this forum so I may not know of any etiquette that everyone has to follow...
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chornedsnorkack
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[*] posted on 23-12-2017 at 10:36


Taguchi´s "Chemistry in Aqueous and Non-Aqueous Solvents" discusses concentrated sulphuric acid a bit.
He states that salts which are "appreciably" soluble in sulphuric acid are "of course" recovered as bisulphates... except many are further solvated.
So:
Li2SO4 - quite soluble, solid phase 2LiHSO4.H2SO4
Na2SO4 - quite soluble, solid phase 4NaHSO4.7H2SO4
K2SO4 - quite soluble, solid phase KHSO4.H2SO4
MgSO4 - poorly soluble, but solid phase Mg(HSO4)2.2H2SO4
CaSO4 - quite soluble, solid phase Ca(HSO4)2.2H2SO4
BaSO4 - quite soluble, solid phase Ba(HSO4)2.2H2SO4

Solid phases unspecified:
Ag2SO4 - quite soluble in H2SO4, poorly soluble in water
ZnSO4 - poorly soluble in H2SO4, quite soluble in water
PbSO4 - poorly soluble in either H2SO4 or water
CuSO4 - poorly soluble in H2SO4, quite soluble in water
FeSO4 - poorly soluble in H2SO4, quite soluble in water
NiSO4 - poorly soluble in H2SO4, quite soluble in water
HgSO4 - poorly soluble in H2SO4, quite soluble in water
Hg2SO4 - poorly soluble in H2SO4, quite soluble in water
Al2(SO4)3 - poorly soluble in H2SO4, quite soluble in water
Tl2(SO4)3 - poorly soluble in H2SO4, quite soluble in water

So... yes, in solution, HSO4- would be the prevalent ion.
But if for any reason in concentrated H2SO4, NiSO4 is a poorly soluble stable solid and "Ni(HSO4)2" would be very soluble, then the insolubility of NiSO4 would drive the reaction to prevent the existence of solid Ni(HSO4)2.
Seems that this is the case.
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Vylletra Heart
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[*] posted on 27-12-2017 at 09:27


Quote: Originally posted by chornedsnorkack  
Taguchi´s "Chemistry in Aqueous and Non-Aqueous Solvents" discusses concentrated sulphuric acid a bit.
He states that salts which are "appreciably" soluble in sulphuric acid are "of course" recovered as bisulphates... except many are further solvated.
So:
Li2SO4 - quite soluble, solid phase 2LiHSO4.H2SO4
Na2SO4 - quite soluble, solid phase 4NaHSO4.7H2SO4
K2SO4 - quite soluble, solid phase KHSO4.H2SO4
MgSO4 - poorly soluble, but solid phase Mg(HSO4)2.2H2SO4
CaSO4 - quite soluble, solid phase Ca(HSO4)2.2H2SO4
BaSO4 - quite soluble, solid phase Ba(HSO4)2.2H2SO4

Solid phases unspecified:
Ag2SO4 - quite soluble in H2SO4, poorly soluble in water
ZnSO4 - poorly soluble in H2SO4, quite soluble in water
PbSO4 - poorly soluble in either H2SO4 or water
CuSO4 - poorly soluble in H2SO4, quite soluble in water
FeSO4 - poorly soluble in H2SO4, quite soluble in water
NiSO4 - poorly soluble in H2SO4, quite soluble in water
HgSO4 - poorly soluble in H2SO4, quite soluble in water
Hg2SO4 - poorly soluble in H2SO4, quite soluble in water
Al2(SO4)3 - poorly soluble in H2SO4, quite soluble in water
Tl2(SO4)3 - poorly soluble in H2SO4, quite soluble in water

So... yes, in solution, HSO4- would be the prevalent ion.
But if for any reason in concentrated H2SO4, NiSO4 is a poorly soluble stable solid and "Ni(HSO4)2" would be very soluble, then the insolubility of NiSO4 would drive the reaction to prevent the existence of solid Ni(HSO4)2.
Seems that this is the case.


Thank you so much for the insightful reply! May I ask where can I find this source?

I have a few questions about this general trends of solubilities though. It seems that most of the non-monovalent cations that you have listed form sulfates that are poorly soluble in sulfuric acid. However, the group II cations seem to disregard this property and are able to form solid bisulfates. The other cations that obey the property are all transition metals. Are the solid phase bisulfate crystals solely dependent on the presence of the 3d shells, or due to something else entirely?

Another thing that I have also noticed is that since silver is the only transition metal you listed that forms a monovalent ion(unlike the Hg22+ ion which is still divalent), it apparently fits my hypothesis that all monovalent ions can form solid bisulfates. However, you did not mention of a solid phase silver bisulfate, which leads me to question of its existence. I have checked wikipedia, and it mentions that upon dissolution in sulfuric acid, silver sulfate forms soluble silver bisulfate which is the driving force for dissolving. I would like to know if crystallising silver sulfate from a solution of sulfuric acid will result in a bisulfate or just the sulfate?

I'm not too sure if I'm correct saying that the cation size plays a role in the solubility in sulfuric acid, due to the observed increasing solubility in sulfuric acid from Mg2+ to Ba2+.(This is a weird finding considering the sulfate of barium is less soluble in water compared to magnesium)
I thought of the lattice energy formula which states that the energy in a crystal lattice is proportional to the product of the charge of one of each of the component ions, and inversely proportional to the sum of the radius of one of each of the different ions(correct me if I am wrong). This suggests that charge has a larger impact on the lattice energy than the radius.
Considering a salt like nickel sulfate, the component ions would both have charges of 2+. This would cause the lattice energy of the salt to be significantly higher than per se, potassium sulfate. When potassium sulfate forms the bisulfate, the resulting charge of the ions would both be 1+, lowering the lattice energy further. Thus, this made me hypothesise that the smaller the lattice energy, the higher the solubility of the sulfate salt in sulfuric acid.
At this point, am I correct in saying that the H+ ions attacking the solid sulfate crystals to form soluble bisulfate is the mechanism for dissolution of all sulfates in sulfuric acid? This would fit the lattice energy theory because the H+ ions would not have enough energy to break apart a crystal that is insoluble in sulfuric acid due to its high lattice energy(nickel sulfate is again a great example here).

My apologies if I am spouting nonsense at this point! I would really love to hear more feedback about my logic. It's hard to find great help when I'm miles ahead of my college's chemistry syllabus hahahahaha!
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chornedsnorkack
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[*] posted on 27-12-2017 at 12:49


Quote: Originally posted by Vylletra Heart  
Quote: Originally posted by chornedsnorkack  
Taguchi´s "Chemistry in Aqueous and Non-Aqueous Solvents" discusses concentrated sulphuric acid a bit.
He states that salts which are "appreciably" soluble in sulphuric acid are "of course" recovered as bisulphates... except many are further solvated.
...
Solid phases unspecified:
...


Thank you so much for the insightful reply! May I ask where can I find this source?


I found it at Google Books. Not all pages reachable. So I did not attempt tracking down his sources in turn. (Even if I had the references, chances are I could not access them).
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The Volatile Chemist
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[*] posted on 27-12-2017 at 18:34


Quote: Originally posted by chornedsnorkack  

BaSO4 - quite soluble, solid phase Ba(HSO4)2.2H2SO4

What?
I lean towards doubting the 'quite soluble' unless it's just a relative term to BaSO4's solubility in water. I can check, but I really doubt BaSO4 being 'quite soluble' in much of anything...




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chornedsnorkack
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[*] posted on 27-12-2017 at 21:41


Quote: Originally posted by The Volatile Chemist  
Quote: Originally posted by chornedsnorkack  

BaSO4 - quite soluble, solid phase Ba(HSO4)2.2H2SO4

What?
I lean towards doubting the 'quite soluble' unless it's just a relative term to BaSO4's solubility in water. I can check, but I really doubt BaSO4 being 'quite soluble' in much of anything...


Yes.
While the solubility in water is indeed tiny, Taguchi, in page 122, Table 4.10 specifies solubility in 100 % sulphuric acid, in units of mole %.
The value is given as 8,85 %.
It would be interesting to have Taguchi´s source tracked down and find at which concentration barium sulphate does start to dissolve well.
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kmno4
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[*] posted on 28-12-2017 at 01:50


Quote:

The value is given as 8,85 %.

This value is taken from the reference:
Code:
http://pubs.acs.org/doi/abs/10.1021/ja01438a002

, also available from the board.




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[*] posted on 28-12-2017 at 02:02


An other note on this subject

Attachment: xxxxxtrenner1930.pdf (451kB)
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