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Dissolving or solubility of K2HPO4 in 0.01M CaCl2? CaHPO4...
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Dissolving or solubility of K2HPO4 in 0.01M CaCl2? CaHPO4...
It might be worth checking the pH during your experiments, just to be sure that it does not change too much during the reaction, as Ca adsorptin varies with pH (H+ ions compete with Ca2+ ions for the same binding sites). Good luck Peter Shellus
It might be worth checking the pH during your experiments, just to be sure that it does not change too much during the reaction, as Ca adsorptin varies with pH (H+ ions compete with Ca2+ ions for the same binding sites). Good luck Peter Shellus
Thank you, Dr. Schmidt, Today, I filtered the solution and check with ICP-OES, 40 PPM gave me 27.5 ppm, this means only about 69% of P solubilized. My concern is during adsorption experiment, 1) difficult to quantify how much P precipitated with Ca from solution and Ca from adsorbent (let say soil or sediment) 2) if we have enough amount of Ca from the adsorbent, the P precipitated with Ca and will give me wrong adsorption capacity for the sorbent. What do you recommend, if I changed CaCl2 to KCl? will it precipitate? Walter Schmidt
Thank you, Dr. Schmidt, Today, I filtered the solution and check with ICP-OES, 40 PPM gave me 27.5 ppm, this means only about 69% of P solubilized. My concern is during adsorption experiment, 1) difficult to quantify how much P precipitated with Ca from solution and Ca from adsorbent (let say soil or sediment) 2) if we have enough amount of Ca from the adsorbent, the P precipitated with Ca and will give me wrong adsorption capacity for the sorbent. What do you recommend, if I changed CaCl2 to KCl? will it precipitate? Walter Schmidt
Unless there is a compelling reason for using CaCl2 as a background electrolyte it is better to use NaCl or KCl, which will not cause precipitation of phosphates. Using any solution containing Ca (or otheralkaline earth metals) runs the risk of precipitation and hence loss of phosphatefrom the solution. I presume the system includes a buffer, as adsorption of phosphate by the substrate will tend to change the pH.
Unless there is a compelling reason for using CaCl2 as a background electrolyte it is better to use NaCl or KCl, which will not cause precipitation of phosphates. Using any solution containing Ca (or otheralkaline earth metals) runs the risk of precipitation and hence loss of phosphatefrom the solution. I presume the system includes a buffer, as adsorption of phosphate by the substrate will tend to change the pH.
Robert Peter Shellis Thank you Dr. Shellis, No compelling reason to use CaCl2, I was wondering many papers published their work on P adsorption on soil using 0.01M CaCl2 as background solution. I don't have any buffer, pH is not my concern for this specific work. I will go with KCl.
Robert Peter Shellis Thank you Dr. Shellis, No compelling reason to use CaCl2, I was wondering many papers published their work on P adsorption on soil using 0.01M CaCl2 as background solution. I don't have any buffer, pH is not my concern for this specific work. I will go with KCl.
You are making a saturated solution and yes calcium phosphate is the excess calcium phosphate. The solubility product from mixing calcium ions [from any source] and phosphate from any source is always a constant [at steady state/ equilibrium]. Thus 1) removing the precipitate [no matter how much forms] will result in the solution concentration of calcium and phosphate being a constant, and 2) if you change the temperature at which you do the experiment, the solution concentration will be different. If it were colder at night, a precipitate would /could form overnight and the concentration of ions in solution when warmed up may not be the same as one had thought. Warming up the sample to dissolve the precipitate within a specific concentration range would work, except this would no longer be the concentration one expected. Cannot get good [low variance] P sorption data unless one begins with a constant [reference] amount of P . . .
You are making a saturated solution and yes calcium phosphate is the excess calcium phosphate. The solubility product from mixing calcium ions [from any source] and phosphate from any source is always a constant [at steady state/ equilibrium]. Thus 1) removing the precipitate [no matter how much forms] will result in the solution concentration of calcium and phosphate being a constant, and 2) if you change the temperature at which you do the experiment, the solution concentration will be different. If it were colder at night, a precipitate would /could form overnight and the concentration of ions in solution when warmed up may not be the same as one had thought. Warming up the sample to dissolve the precipitate within a specific concentration range would work, except this would no longer be the concentration one expected. Cannot get good [low variance] P sorption data unless one begins with a constant [reference] amount of P . . .
It might be worth checking the pH during your experiments, just to be sure that it does not change too much during the reaction, as Ca adsorptin varies with pH (H+ ions compete with Ca2+ ions for the same binding sites).
Good luck
Peter Shellus
It might be worth checking the pH during your experiments, just to be sure that it does not change too much during the reaction, as Ca adsorptin varies with pH (H+ ions compete with Ca2+ ions for the same binding sites).
Good luck
Peter Shellus
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Thank you! Robert Peter Shellis
Thank you! Robert Peter Shellis
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Thank you, Dr. Schmidt,
Today, I filtered the solution and check with ICP-OES, 40 PPM gave me 27.5 ppm, this means only about 69% of P solubilized. My concern is during adsorption experiment, 1) difficult to quantify how much P precipitated with Ca from solution and Ca from adsorbent (let say soil or sediment) 2) if we have enough amount of Ca from the adsorbent, the P precipitated with Ca and will give me wrong adsorption capacity for the sorbent. What do you recommend, if I changed CaCl2 to KCl? will it precipitate? Walter Schmidt
Thank you, Dr. Schmidt,
Today, I filtered the solution and check with ICP-OES, 40 PPM gave me 27.5 ppm, this means only about 69% of P solubilized. My concern is during adsorption experiment, 1) difficult to quantify how much P precipitated with Ca from solution and Ca from adsorbent (let say soil or sediment) 2) if we have enough amount of Ca from the adsorbent, the P precipitated with Ca and will give me wrong adsorption capacity for the sorbent. What do you recommend, if I changed CaCl2 to KCl? will it precipitate? Walter Schmidt
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Unless there is a compelling reason for using CaCl2 as a background electrolyte it is better to use NaCl or KCl, which will not cause precipitation of phosphates. Using any solution containing Ca (or otheralkaline earth metals) runs the risk of precipitation and hence loss of phosphatefrom the solution. I presume the system includes a buffer, as adsorption of phosphate by the substrate will tend to change the pH.
Unless there is a compelling reason for using CaCl2 as a background electrolyte it is better to use NaCl or KCl, which will not cause precipitation of phosphates. Using any solution containing Ca (or otheralkaline earth metals) runs the risk of precipitation and hence loss of phosphatefrom the solution. I presume the system includes a buffer, as adsorption of phosphate by the substrate will tend to change the pH.
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Robert Peter Shellis Thank you Dr. Shellis,
No compelling reason to use CaCl2, I was wondering many papers published their work on P adsorption on soil using 0.01M CaCl2 as background solution. I don't have any buffer, pH is not my concern for this specific work. I will go with KCl.
Robert Peter Shellis Thank you Dr. Shellis,
No compelling reason to use CaCl2, I was wondering many papers published their work on P adsorption on soil using 0.01M CaCl2 as background solution. I don't have any buffer, pH is not my concern for this specific work. I will go with KCl.
More
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You are making a saturated solution and yes calcium phosphate is the excess calcium phosphate. The solubility product from mixing calcium ions [from any source] and phosphate from any source is always a constant [at steady state/ equilibrium]. Thus 1) removing the precipitate [no matter how much forms] will result in the solution concentration of calcium and phosphate being a constant, and 2) if you change the temperature at which you do the experiment, the solution concentration will be different. If it were colder at night, a precipitate would /could form overnight and the concentration of ions in solution when warmed up may not be the same as one had thought. Warming up the sample to dissolve the precipitate within a specific concentration range would work, except this would no longer be the concentration one expected. Cannot get good [low variance] P sorption data unless one begins with a constant [reference] amount of P . . .
You are making a saturated solution and yes calcium phosphate is the excess calcium phosphate. The solubility product from mixing calcium ions [from any source] and phosphate from any source is always a constant [at steady state/ equilibrium]. Thus 1) removing the precipitate [no matter how much forms] will result in the solution concentration of calcium and phosphate being a constant, and 2) if you change the temperature at which you do the experiment, the solution concentration will be different. If it were colder at night, a precipitate would /could form overnight and the concentration of ions in solution when warmed up may not be the same as one had thought. Warming up the sample to dissolve the precipitate within a specific concentration range would work, except this would no longer be the concentration one expected. Cannot get good [low variance] P sorption data unless one begins with a constant [reference] amount of P . . .
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