Question
A solution is 0.15$M$ in both $\mathrm{Pb}^{2+}$ and $\mathrm{Ag}^{+}$ . If $\mathrm{Cl}$ is added to this solution, what is $\left[\mathrm{Ag}^{+}\right]$ when $\mathrm{PbCl}_{2}$ begins to precipitate?
Step 1
The first one is $\mathrm{PbCl}_{2}$ which dissociates into $\mathrm{Pb}^{2+}$ and $2\mathrm{Cl}^{-}$ with a solubility product constant (Ksp) of $1.6 \times 10^{-5}$. The second one is $\mathrm{AgCl}$ which dissociates into $\mathrm{Ag}^{+}$ and $\mathrm{Cl}^{-}$ Show more…
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A solution is $0.15 M$ in both $\mathrm{Pb}^{2+}$ and $\mathrm{Ag}^{+}$. If $\mathrm{Cl}^{-}$ is added to this solution, what is $\left[\mathrm{Ag}^{+}\right]$ when $\mathrm{PbCl}_{2}$ begins to precipitate?
A solution is $0.020 M$ in $\mathrm{Pb}^{2+}$ and $0.020 \mathrm{M}$ in $\mathrm{Ag}^{+} .$ As $\mathrm{Cl}^{-}$ is introduced to the solution by the addition of solid $\mathrm{NaCl}$, determine (a) which substance will precipitate first, $\mathrm{AgCl}$ or $\mathrm{PbCl}_{2},$ and $(\mathrm{b})$ the fraction of the metal ion in the first precipitate that remains in solution at the moment the precipitation of the second compound begins.
A solution is 0.15 M in both Pb2+ and Ag+. If Cl– is added to this solution, what is [Ag+] when PbCl2 begins to precipitate?
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