Kogor
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Analytical Chemistry - Precipitate
Does PH influence the solubility of a precipitate? Examples if it does?
Thanks, sorry for my english 
[Edited on 4-12-2010 by Kogor]
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MagicJigPipe
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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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that the only way to detect it is to be free to inquire. And we know that as long as men are free to ask what they must, free to say what they think,
free to think what they will, freedom can never be lost, and science can never regress." -J. Robert Oppenheimer
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Kogor
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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 
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DJF90
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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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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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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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Nicodem
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| 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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| 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
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