Chemical Forums
Chemistry Forums for Students => Inorganic Chemistry Forum => Topic started by: covek11 on September 01, 2009, 01:00:09 PM
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Yesterday I precipitated
Ca(NO3)2 + H2SO4 = CaSO4*2H2O + HNO3
After washing it with plenty of water and air drying I found yield of only 62%.
How is that possible when CaSO4 should precipitate quantitatively ? I made 10.7 kgs of it. What is more interesting is that I tested supernatant with a solution of NaF and no CaF precipitate appeared. But, when I neutralized filtrate with Na2CO3 I found white precipitate !
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Are you sure you have used stoichiometric amounts of reactants?
What was pH of the solution?
Do you know Ksp values of all involved salts?
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Yesterday I precipitated
Ca(NO3)2 + H2SO4 = CaSO4*2H2O + HNO3
After washing it with plenty of water and air drying I found yield of only 62%.
How is that possible when CaSO4 should precipitate quantitatively ? I made 10.7 kgs of it. What is more interesting is that I tested supernatant with a solution of NaF and no CaF precipitate appeared. But, when I neutralized filtrate with Na2CO3 I found white precipitate !
CaSO4 dissolves in strongly acidic solution:
CaSO4 + H+ <--> Ca++ + HSO4-
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I had right amounts of reactants for sure and I had 5% excess of H2SO4.
O, nooo. So, what would be a solution to stop dissociation of CaSO4 ? Add some ammonia after all H2SO4 is added? Is there a possibility that I could precipitate Ca(OH)2?
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To precipitate Ca(OH)2 you would need pH in the 12 range or something like that - so no much risk. I would try to neutralize solution against pH strips with any strong base, around pH 7 concentration of HSO4- is neglectable.
Excess SO42- seems the best idea to precipiate as much calcium as possible.
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Thank you Borek.
We still haven't explained why no precipitate appeared when I mixed a sample of supernatant with solution of NaF. CaF should precipitate under acidic conditions as well as far I know.
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Not necesarilly. HF is much weaker than HSO4-. If pH was low enough to increase solubility of sulfate, it was also high enough to increase solubility of fluoride.
I don't HAVE to be right, I am just trying to estimate equiilbrium concentrations looking at Kso/Ka values.
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Thank you Borek.
We still haven't explained why no precipitate appeared when I mixed a sample of supernatant with solution of NaF. CaF should precipitate under acidic conditions as well as far I know.
Just a question: do you have to necessarily use H2SO4 to precipitate Ca++ ? If I could choose I would use Na2SO4.
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Na2SO4 would be more reliable, but it leaves Na+ behind. I am not allowed to have Na in ppm quantities in CaSO4.
When I make CaF2, I simply react Ca(NO3)2 with HF and I have almost 100 % yield. Product of that reaction is HNO3, which is stronger acid than HF and on that basis all HF should be in undissociated form, but precipitation still takes place. ???
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Na2SO4 would be more reliable, but it leaves Na+ behind. I am not allowed to have Na in ppm quantities in CaSO4.
What about (NH4)2SO4 ?When I make CaF2, I simply react Ca(NO3)2 with HF and I have almost 100 % yield. Product of that reaction is HNO3, which is stronger acid than HF and on that basis all HF should be in undissociated form, but precipitation still takes place. ???
If you are in water solution you cannot write HNO3, only H3O+: this is the strongest acid present there.
To understand what happens you have to write the equilibriums involving weak acids/bases and insoluble/(almost insoluble) salts/complexes.
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I just wanted to let you know that I repeated CaSO4 run. This time I adjusted pH at 7 before addition of H2SO4 into Ca(NO3)2. This time I had yield of 75 %. It is weird that I lost 25% of CaSO4*2H2O. When I looked at Ksp value I calculated that I should lose only 200-300g which is 3% of theoretical value. Is it possible that I precipitated CaSO4*0.5H2O or anhydrous CaSO4 ? I know it sounds crazy, but I don't see any better explanation.
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Is your nitrate anhydrous?
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No, I used water solutions.
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How have you checked calcium nitrate solution concentration?
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I haven't checked it. I relied on Certificate of Analysis of 99.5% pure CaCO3 that I dissolved in Nitric Acid.
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I haven't checked it. I relied on Certificate of Analysis of 99.5% pure CaCO3 that I dissolved in Nitric Acid.
Sounds reasonable.
Could be your water content is not what you expect it to be, I am afraid I can't be of help any further.
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I haven't checked it. I relied on Certificate of Analysis of 99.5% pure CaCO3 that I dissolved in Nitric Acid.
Did you boil the solution after having dissolved CaCO3 with HNO3? Otherwise you can have HCO3- ions which could keep Ca++ ions in solution (if I remember well). Boiling the solution eliminates dissolved CO2 and so H2CO3 and HCO3- ions.
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No, I didn't boil it.
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Otherwise you can have HCO3- ions which could keep Ca++ ions in solution (if I remember well). Boiling the solution eliminates dissolved CO2 and so H2CO3 and HCO3- ions.
In the presence of excess sulfuric acid HCO3- won't survive long enough in the solution.
Besides, I think you may be mixing two things - HCO3- doesn't 'keep' Ca2+ in solution, but Ca(HCO3)2 is much better soluble in water. Thus water kept over solid CaCO3 will have higher concentration of Ca2+ when saturated with CO2.
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In the presence of excess sulfuric acid HCO3- won't survive long enough in the solution.
Yes, but he said to have adjusted the pH to 7.
Besides, I think you may be mixing two things - HCO3- doesn't 'keep' Ca2+ in solution, but Ca(HCO3)2 is much better soluble in water. Thus water kept over solid CaCO3 will have higher concentration of Ca2+ when saturated with CO2.
You're right.