Can Anyone Help Me With This Solubility Problem?

An aqueous solution of 1.90 M AgNO3 is slowly added from a buret to an aqueous solution of 0.0100 M Cl- and 0.250 M I-.
a) Which ion, Cl- or I-, is the first to precipitate?
b) When the second ion begins to precipitate, what is the remaining concentration of the first ion?
c) Is this separation feasible by fractional precipitation?

Consider the solubilities of AgI and AgCl. Starting with the fact that the I ion is already 25 times as prevalent as the Cl ion, which is likely (considering solubility) to precipate first? Is the solubility of the I 25 times as great?

Start with that type analysis.

Certainly! I can help you with this solubility problem. Let's go through the process step by step.

To determine which ion, Cl- or I-, is the first to precipitate, we need to compare their solubility product constants (Ksp) with the actual ion product (IP) in the solution. The ion product (IP) is calculated by multiplying the concentrations of the individual ions in the solution.

a) To find out which ion precipitates first, we compare the Ksp values. The Ksp for silver chloride (AgCl) is 1.8 x 10^-10, and the Ksp for silver iodide (AgI) is 1.5 x 10^-16. The ion product (IP) for AgCl is [Ag+][Cl-], and the IP for AgI is [Ag+][I-].

To determine which ion precipitates first, we need to compare the IP values of the two compounds with their respective Ksp values. The ion with a smaller IP value compared to its Ksp value will precipitate first.

Calculate the IP values:
IP for AgCl = [Ag+][Cl-] = (1.90 M)(0.0100 M) = 1.9 x 10^-2
IP for AgI = [Ag+][I-] = (1.90 M)(0.250 M) = 0.475

Comparing the IP values with the Ksp values:
For AgCl: IP (1.9 x 10^-2) < Ksp (1.8 x 10^-10)
For AgI: IP (0.475) > Ksp (1.5 x 10^-16)

Since the IP of AgI is greater than its Ksp, AgI will precipitate first. Therefore, the ion I- is the first to precipitate.

b) Now, to determine the remaining concentration of the first ion, we need to convert the moles of AgI precipitated into moles of AgNO3 used.

Since the molar ratio between AgI and AgNO3 is 1:1, the moles of AgNO3 used will be equal to the moles of AgI precipitated.

First, calculate the moles of AgI precipitated:
Moles of AgI = Concentration of AgI × Volume of Solution
Moles of AgI = (0.0100 M)(Volume of Solution)

Convert the moles of AgI precipitated into moles of AgNO3 used:
Moles of AgNO3 used = Moles of AgI precipitated

To find the remaining concentration of AgNO3, subtract the moles of AgNO3 used from the initial moles of AgNO3:
Remaining concentration of AgNO3 = Initial concentration of AgNO3 - Moles of AgNO3 used

c) Whether the separation is feasible by fractional precipitation depends on the difference in solubilities of the two compounds and their corresponding Ksp values. In this case, since the Ksp value for AgCl is significantly larger than AgI, it means that AgCl is much more soluble than AgI. Therefore, it is not feasible to separate them by fractional precipitation, as AgCl will always remain in solution while AgI precipitates.

I hope this helps! Let me know if you have any further questions.