Final answer:
The correct [Ag+] when PbCl2 begins to precipitate from a 0.15 M Pb2+ and Ag+ solution upon the addition of Cl- ions cannot be determined without the specific Ksp values. Normally, once PbCl2 precipitates, we would then assess the possible precipitation of AgCl, but the question does not provide the necessary information to make this calculation.
Step-by-step explanation:
The student's question asks at what concentration of Ag+ will lead(II) chloride (PbCl2) begin to precipitate when chloride ions (Cl-) are added to a 0.15 M solution of both Pb2+ and Ag+. This is a classic problem of selective precipitation based on solubility products (Ksp).
The problem requires the use of solubility rules and Ksp values for silver chloride (AgCl) and lead chloride (PbCl2) to determine the saturation point.
To solve the problem, we must compare the ion product (Q) to the solubility product constants (Ksp) for PbCl2 and AgCl, knowing that when Q reaches Ksp, precipitation begins.
Let's consider the Ksp of PbCl2 and AgCl. Since PbCl2 has a smaller Ksp, it will precipitate first when a common ion like Cl- is added.
At the point when PbCl2 begins to precipitate, we can calculate the concentration of Ag+ remaining in solution. The key is to use the Ksp values to find out what concentration of Cl- will cause PbCl2 to precipitate and then check to see if this concentration of Cl- is enough to also start precipitating AgCl.
The given scenarios do not provide specific Ksp values or tell us the amount of Cl- being added, so typically we cannot calculate the exact concentration of Ag+ without this data.
However, since the question is multiple choice, we can determine that the correct answer is one of the provided options, even though we cannot solve it with the given information.