In this problem, we are given a solution containing Cl- ions, and we need to calculate the minimum concentration of Ag+ ions required to initiate the precipitation of AgCl. The solubility product constant (Ksp) of AgCl is also given, which will be used to calculate the minimum concentration of Ag+ ions.



The precipitation of AgCl occurs when the product of the concentrations of Ag+ and Cl- ions exceeds the solubility product constant of AgCl. Mathematically, we can express this as:

[Ag+][Cl-] > Ksp

Since the concentration of Cl- ions is given as 0.020 M, we can rearrange the above equation to solve for the minimum concentration of Ag+ ions required:

[Ag+] > Ksp / [Cl-]

Substituting the values, we get:

[Ag+] > (1.6 x 10^-10) / (0.020) = 8 x 10^-9 M

Therefore, the minimum concentration of Ag+ ions required to initiate the precipitation of AgCl is 8 x 10^-9 M.



The solubility product constant (Ksp) of AgCl is a measure of the extent to which the compound dissociates into its constituent ions in water. When Ag+ and Cl- ions are present in a solution in concentrations that exceed the Ksp of AgCl, the excess ions combine to form solid AgCl, which precipitates out of the solution. In this problem, we used the Ksp of AgCl and the concentration of Cl- ions to calculate the minimum concentration of Ag+ ions required to initiate the precipitation of AgCl. This calculation is based on the principle of equilibrium, which states that a system will reach a state of equilibrium when the rates of the forward and reverse reactions are equal. In this case, the forward reaction involves the dissociation of AgCl into its constituent ions, while the reverse reaction involves the recombining of Ag+ and Cl- ions to form solid AgCl. At equilibrium, the product of the concentrations of the ions in solution is equal to the Ksp of AgCl, which determines the extent to which the compound will dissociate.

1.6×10^-10÷0.02