2. Solubility Equilibria
W e will assume that when atypical ionic solid dissolves in water it dissociates completely into separate hydrated cations and anions. In the beginning no cations and anions are present. Ultimately dynamic equilibrium is reached. The solution is said to be saturated. The equilibrium expression for the following reaction, according to the law of mass action.
Bi2S3(s) Bi3+(aq) + 3S2–(aq)
[Bi3+] 2and [S2-]3 are expressed in mol/L
Ksp = solubility product for equilibrium expression Bi2S3 is a pure solid and not included in the expression. The amount of excess solid does not effect the position of the solubility equilibrium.
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3. Solubility is an equilibrium position.
Solubility product (Ksp) is an equilibrium constant; and has only one value for a given solid at a given temperature.
Solubility is an equilibrium position.
In a pure water at a specific temperature a given salt has a particular solubility.
On the other hand, if a common ion is present in the solution the solubility varies according to the concentration od the common ion.
In all cases the product of the ion concentration must satisfy the Ksp expression. Table 16.1
The solubility product is a n equilibrium constant and has one value for a given solid at a given temperature. On the other hand solubility is a n equilibrium position.

4. Relative Solubilities,, pH and solubilities
Salt’s Ksp gives us the information about its solubility.
We mist be carefull in using Ksp values to predict the relative solubilities for a group of salts.
1. The salts being compared produce same number of ions .
Salt = [cation]+[anion]
Ksp = [cation]+[anion]
[cation] = x [anion] = x
Ksp = [cation]+[anion] x2
x =√Ksp = solubility
2. The salts being compared produce different number of ions. Ksp vales cannot be compared directly to determine the relative solubilities. Table 16.2 shows the order is opposite to the order of the Ksp values.

5. Common ion effect, pH and solubilities
When the water contains an ion in common with the dissolving salt.
pH and solubilities
Mg(OH)2(aq)↔Mg2+(aq) + 2OH-(aq)
Addition of OH- ion (an increase in pH) by the common ion effect will force the equilibrium to the left, decreasing the solubilities of Mg(OH)2.
On the other hand addition of H+ ion (a decrease in pH) increase the solubility , because OH- ions are removed from the solution by reacting with the added H+ ions. In response to the lower concentration of OH- , the equilibrium position moves to the right.
Suspension of solid Mg(OH)2 (Milk of Magnesia) dissolves as requires in the stomach to combat excess of acidity.

6. Explain. If yes, explain and verify. If no, provide a counter-example.
CONCEPT CHECK!
In comparing several salts at a given temperature, does a higher Ksp value always mean a higher solubility?
Explain. If yes, explain and verify. If no, provide a counter-example. No
No. In order to relate Ksp values to solubility directly, the salts must contain the same number of ions. For example, for a binary salt, Ksp = s2 (s = solubility); for a ternary salt, Ksp = 4s3.
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7. EXERCISE!
Calculate the solubility of silver chloride in water. Ksp = 1.6 × 10–10
1.3×10-5 M
Calculate the solubility of silver phosphate in water. Ksp = 1.8 × 10–18
1.6×10-5 M
a) 1.3×10-5 M
b) 1.6×10-5 M
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8. The solubilities are the same.
CONCEPT CHECK!
How does the solubility of silver chloride in water compare to that of silver chloride in an acidic solution (made by adding nitric acid to the solution)?
Explain.
The solubilities are the same.
The solubilities are the same. Since HCl is a strong acid, it is completely dissociated in water.
There are no common ions between AgCl and HNO3.
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9. The silver phosphate is more soluble in an acidic solution.
CONCEPT CHECK!
How does the solubility of silver phosphate in water compare to that of silver phosphate in an acidic solution (made by adding nitric acid to the solution)?
Explain.
The silver phosphate is more soluble in an acidic solution.
The silver phosphate is more soluble in an acidic solution. This is because the phosphate ion is a relatively good base and will react with the proton from the acid (essentially to completion). The phosphate ion does not react nearly as well with water. This is an example of the effect of LeChâtelier's principle on the position of the solubility equilibrium.
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10. The Ksp values are the same.
CONCEPT CHECK!
How does the Ksp of silver phosphate in water compare to that of silver phosphate in an acidic solution (made by adding nitric acid to the solution)?
Explain.
The Ksp values are the same.
The Ksp values are the same (assuming the temperature is constant).
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11. Calculate the solubility of AgCl in:
EXERCISE!
Calculate the solubility of AgCl in:
Ksp = 1.6 × 10–10
100.0 mL of 4.00 x 10-3 M calcium chloride.
2.0×10-8 M
100.0 mL of 4.00 x 10-3 M calcium nitrate.
1.3×10-5 M
a) 2.0×10-8 M Note: [Cl-] in CaCl2 is twice the [CaCl2] given mL is not used in the calculation.
b) 1.3×10-5 M
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12. Precipitation (Mixing Two Solutions of Ions)
How to predict whether a precipitate will form when two solutions are mixed.
We will use ion product Q just like Ksp for a given solids, except that initial concentrations are used instead of equilibrium concentrations.
Q > Ksp; precipitation occurs and will continue until the concentrations are reduced to the point that they satisfy Ksp.
Q < Ksp; no precipitation occurs.
Problem 16.49
Some times we want to do more than simply predict whether precipitation will occur. We may want to calculate the equilibrium concentration in the solution after the precipitation occurs. Problem 16.5
Selective Precipitation: Mixture of metal ions in aqueous solutions are often separated by selective precipitation.
It is done by using a reagent whose anion forms a precipitate with only one or a few of the metal ions in mixture.
Suppose we have a solution containing both Ba2+ and Ag+ ions. If NaCl is added to the solution, AgCl precipitates as a white solid, but since BaCl2 is soluble , Ba2+ ions remain in solution. Problem 16.57
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13. Selective Precipitation (Mixtures of Metal Ions)
Use a reagent whose anion forms a precipitate with only one or a few of the metal ions in the mixture.
Example:
Solution contains Ba2+ and Ag+ ions.
Adding NaCl will form a precipitate with Ag+ (AgCl), while still leaving Ba2+ in solution.
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14. Separation of Cu2+ and Hg2+ from Ni2+ and Mn2+ using H2S
At a low pH, [S2–] is relatively low and only the very insoluble HgS and CuS precipitate.
When OH– is added to lower [H+], the value of [S2–] increases, and MnS and NiS precipitate.
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15. Separation of Cu2+ and Hg2+ from Ni2+ and Mn2+ using H2S
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16. Separating the Common Cations by Selective Precipitation
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17. Complex Ion Equilibria
Complex ions Charged species consisting of a metal ion surrounded by ligands.
Ligand: Lewis base H2O, NH3, Cl-, CN-. The number of ions attached to a metal ion is called coordination number. Co(H2O)62+ . Coordination number = 6
Metal ions add ligand one at a time in steps characterized by equilibrium constant called formation (stability) constant.
Equilibrium constant for each step of the formation of a complex ion by the addition of an individual ligand to a metal ion or complex ion in aqueous solution can be calculated.
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18. Complex Ion Equilibria
Be2+(aq) + F–(aq) BeF+(aq) K1 = 7.9 × 104 BeF+(aq) + F–(aq) BeF2(aq) K2 = 5.8 × 103 BeF2(aq) + F–(aq) BeF3– (aq) K3 = 6.1 × 102 BeF3– (aq) + F–(aq) BeF42– (aq) K4 = 2.7 × 101
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19. Complex Ions and Solubility
Two strategies for dissolving a water–insoluble ionic solid.
If the anion of the solid is a good base, the solubility is greatly increased by acidifying the solution.
In cases where the anion is not sufficiently basic, the ionic solid often can be dissolved in a solution containing a ligand that forms stable complex ions with its cation.
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20. Ksp (AgCl) = 1.6 × 10–10 0.48 M CONCEPT CHECK!
Calculate the solubility of silver chloride in 10.0 M ammonia given the following information:
Ksp (AgCl) = 1.6 × 10–10
Ag+ + NH AgNH K = 2.1 × 103
AgNH3+ + NH Ag(NH3)2+ K = 8.2 × 103
0.48 M
Calculate the concentration of NH3 in the final equilibrium mixture.
9.0 M
a) 0.48 M
b) 9.0 M
This problem is discussed at length in the text.
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