KI(aq) ???? ??? ?????? ?? [Pb2+] =[Ag+] = 0.10 M (?) ???? ???? ???? ??????? ????, ?? Agl? Pbl2
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For AgI, Ksp = 8.3 x 10^-17 For PbI2, Ksp = 7.9 x 10^-9 The smaller the Ksp, the less soluble the compound is. Show more…
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$\mathrm{KI}(\mathrm{aq})$ is slowly added to a solution with $\left[\mathrm{Pb}^{2+}\right]=$ $\left[\mathrm{Ag}^{+}\right]=0.10 \mathrm{M} .$ For $\mathrm{PbI}_{2}, K_{\mathrm{sp}}=7.1 \times 10^{-9} ;$ for $\mathrm{AgI},$ $K_{\mathrm{sp}}=8.5 \times 10^{-17}$. (a) Which precipitate should form first, $\mathrm{PbI}_{2}$ or AgI? (b) What $\left[\mathrm{I}^{-}\right]$ is required for the second cation to begin to precipitate? (c) What concentration of the first cation to precipitate remains in solution at the point at which the second cation begins to precipitate? (d) $\operatorname{Can} \mathrm{Pb}^{2+}(\mathrm{aq})$ and $\mathrm{Ag}^{+}($ aq) be effectively separated by fractional precipitation of their iodides?
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?
$\mathrm{KI}(\mathrm{aq})$ is slowly added to a solution with $\left[\mathrm{Pb}^{2+}\right]=$ $\left[\mathrm{Ag}^{+}\right]=0.10 \mathrm{M} .$ For $\mathrm{PbI}_{2}, K_{\mathrm{sp}}=7.1 \times 10^{-9} ;$ for $\mathrm{AgI}$ $K_{\mathrm{sp}}=8.5 \times 10^{-17}$ (a) Which precipitate should form first, $\mathrm{PbI}_{2}$ or AgI? (b) What $\left[\mathrm{I}^{-}\right]$ is required for the second cation to begin to precipitate? (c) What concentration of the first cation to precipitate remains in solution at the point at which the second cation begins to precipitate? (d) $\operatorname{Can} \mathrm{Pb}^{2+}(\mathrm{aq})$ and $\mathrm{Ag}^{+}(\mathrm{aq})$ be effectively sepa- rated by fractional precipitation of their iodides?
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