For the reaction, ["Ag(CN)"_2]^(-)-> "Ag"^+ + "2CN"^- the equilibrium constant at 298 K is 4.0 × 10^-10. The silver ion concentration in a solution which was originally 0.12 M in KCN and 0.04 M in "AgNO"_3 is?
1 Answer
I got
The way I see this is that the
K_f = 1/(4.0 xx 10^(-10)) = 2.5 xx 10^9 .
So, this is first a limiting reactant problem, and then it becomes a common-ion effect problem with leftover excess reactant.
We assume that these solutions are those concentrations after they were mixed, so that the solution volume can be arbitrarily
2"KCN"(aq) + "AgNO"_3(aq) -> "K"["Ag"("CN")_2] (aq) + "KNO"_3(aq)
I assume you can realize that
So, our ICE table begins with
"Ag"("CN")_2^(-)(aq) rightleftharpoons "Ag"^(+)(aq) + 2"CN"^(-)(aq)
"I"" ""0.04 M"" "" "" "" ""0.00 M"" "" ""0.04 M"
"C"" "-x" "" "" "" "" "+x" "" "" "+2x
"E"" "(0.04 - x)"M"" "" "x" M"" "" "(0.04 + 2x)" M"
This gives an equilibrium expression of:
K_D -= K_f^(-1)
= 4.0 xx 10^(-10) = (["Ag"^(+)]["CN"^(-)]^2)/(["Ag"("CN")_2^(-)])
= (x(0.04 + 2x)^2)/(0.04 - x)
When I use the small x approximation here, I get:
4.0 xx 10^(-10) = (x(0.04)^2)/(0.04)
= 0.04x
=> color(blue)(x = ["Ag"^(+)]) = K_D/0.04 = color(blue)(1.0 xx 10^(-8) "M")