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Author: Subject: Analytical Chemistry - Precipitate
Kogor
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[*] posted on 4-12-2010 at 12:00
Analytical Chemistry - Precipitate


Does PH influence the solubility of a precipitate? Examples if it does?

Thanks, sorry for my english :P

[Edited on 4-12-2010 by Kogor]
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MagicJigPipe
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[*] posted on 4-12-2010 at 12:57


Of course! An aliphatic amine would be a good example. Once the amine is protonated by an acid it will become soluble (or more so) so, in this case, the solubility is somewhat pH dependent.

Also, another example would be "high" molecular weight carboxylic acids like benzoic acid. It is only slightly soluble in water but if the solution is made basic it dissolves readily due to the formation of a benzoate salt. I'm not sure if this is what you mean, though.

Or perhaps you mean something like AgCl? I don't think pH affects the solubility of strong acid/strong base salts in any significant way but it should affect basic/acidic salts like bicarbonates, carbonates, ammonium salts, carboxylates etc...

This is as much as I know on this subject without looking it up so, I hope this helps.




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Kogor
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[*] posted on 4-12-2010 at 13:40


Quote: Originally posted by MagicJigPipe  


Or perhaps you mean something like AgCl? I don't think pH affects the solubility of strong acid/strong base salts in any significant way but it should affect basic/acidic salts like bicarbonates, carbonates, ammonium salts, carboxylates etc...

This is as much as I know on this subject without looking it up so, I hope this helps.


Yes, i mean inorganic compounds such as AgCl, Ag2CrO4, BaSO4, CaC2O4...

Thanks for the help :D
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DJF90
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[*] posted on 4-12-2010 at 14:51


A more important effect is the ionic strength of the solution. But I'll let you go read about that one on your own!
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Lambda-Eyde
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[*] posted on 5-12-2010 at 09:29


Addition of HCl will reduce the solubility of AgCl. Have you learned about the common-ion effect?
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Kogor
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[*] posted on 5-12-2010 at 10:29


Quote: Originally posted by Lambda-Eyde  
Addition of HCl will reduce the solubility of AgCl. Have you learned about the common-ion effect?


Yes, i know that effect.
I need to know only about the PH, for example: If i add HCl to a solution with a precipitate of BaSO4, will it change its solubility? Or nothing will happen?
I think it only depends on the Kps of the precipitate, the ions concentration and the temperature, but i'm not sure if having an acid or basic solution can affect solubility (at least for inorganic compounds).



[Edited on 5-12-2010 by Kogor]
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[*] posted on 5-12-2010 at 13:24


Quote: Originally posted by Kogor  
I need to know only about the PH, for example: If i add HCl to a solution with a precipitate of BaSO4, will it change its solubility? Or nothing will happen?
I think it only depends on the Kps of the precipitate, the ions concentration and the temperature, but i'm not sure if having an acid or basic solution can affect solubility (at least for inorganic compounds).

Yes, of course, it depends only on the Ksp, but have you bothered checking the equation for this constant? If you do, you will see it is about ion activity or, incorrectly simplified, about ion concentration. The concentration of some ions is highly dependent on pH. The classical pedagogic example of huge solubility change from a small change in pH, which is thought already in the elementary school due to its significance in geology, is that of CaCO<sub>3</sub> in water. The activity of hydrated calcium cations in water, [Ca<sup>2+</sup>(aq)], is more or less constant in a wide range of pH, but the same is not true for the activity of the carbonate anions, [CO<sub>3</sub><sup>2-</sup>]. The activity of these changes dramatically when the pH drops by a couple of units, for example, from 8 to 6. This is because the pKa1 of the carbonate makes it a relatively strong base (pKa1 = 10.35, pKa2 = 6.33). This means small acidifications (droping in pH) will protonate the carbonate anions to form bicarbonate anions, thus lowering the [CO<sub>3</sub><sup>2-</sup>], which on turn is part of the Ksp equation. Thus the solubility of CaCO<sub>3</sub> is highly dependent on pH - it is almost insoluble at pH > 8, but solubility increases at pH bellow 7.

So, how do you know when does the pH influence the solubility of a ionic compound? Simple, you check the pKa of the ions it dissociates into.
You can think of the change in pH as change in the concentration of H3O+ or OH- ions in the solution. If these interact with any ion derived from the dissociation of the solute, then the solubility will change. Similar as you do when evaluating the influence on the solubility of adding any third party additive (you check if it forms forms coordination compounds, if it precipitates a certain ion due to even more unfavourable Ksp, or if introduces a redox equilibrium, etc.).




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Kogor
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[*] posted on 5-12-2010 at 20:39


Quote: Originally posted by Nicodem  
Quote: Originally posted by Kogor  
I need to know only about the PH, for example: If i add HCl to a solution with a precipitate of BaSO4, will it change its solubility? Or nothing will happen?
I think it only depends on the Kps of the precipitate, the ions concentration and the temperature, but i'm not sure if having an acid or basic solution can affect solubility (at least for inorganic compounds).

Yes, of course, it depends only on the Ksp, but have you bothered checking the equation for this constant? If you do, you will see it is about ion activity or, incorrectly simplified, about ion concentration. The concentration of some ions is highly dependent on pH. The classical pedagogic example of huge solubility change from a small change in pH, which is thought already in the elementary school due to its significance in geology, is that of CaCO<sub>3</sub> in water. The activity of hydrated calcium cations in water, [Ca<sup>2+</sup>(aq)], is more or less constant in a wide range of pH, but the same is not true for the activity of the carbonate anions, [CO<sub>3</sub><sup>2-</sup>]. The activity of these changes dramatically when the pH drops by a couple of units, for example, from 8 to 6. This is because the pKa1 of the carbonate makes it a relatively strong base (pKa1 = 10.35, pKa2 = 6.33). This means small acidifications (droping in pH) will protonate the carbonate anions to form bicarbonate anions, thus lowering the [CO<sub>3</sub><sup>2-</sup>], which on turn is part of the Ksp equation. Thus the solubility of CaCO<sub>3</sub> is highly dependent on pH - it is almost insoluble at pH > 8, but solubility increases at pH bellow 7.

So, how do you know when does the pH influence the solubility of a ionic compound? Simple, you check the pKa of the ions it dissociates into.
You can think of the change in pH as change in the concentration of H3O+ or OH- ions in the solution. If these interact with any ion derived from the dissociation of the solute, then the solubility will change. Similar as you do when evaluating the influence on the solubility of adding any third party additive (you check if it forms forms coordination compounds, if it precipitates a certain ion due to even more unfavourable Ksp, or if introduces a redox equilibrium, etc.).


I understand it now.
Thanks for all the answers :D
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