Answer
$K_{net} = 3.6 \times 10^{3}$
Since the constant value is greater than 1, the reaction is product favored.
Work Step by Step
1. Write the $K_{sp}$ chemical reaction:
$AgCl(s) \lt -- \gt Ag^+(aq) + Cl^-(aq)$ _____ $K_{sp} = 1.8 \times 10^{-10}$
2. Write the chemical reaction that produces $[Ag(S_2O_3)_2]^{3-}$:
$Ag^+(aq) + 2S_2O{_3}^{2-}(aq) \lt -- \gt [Ag(S_2O_3)_2]^{3-}(aq)$
$K_f = 2.0 \times 10^{13}$
3. Now, combine these equations:
$AgCl(s) + Ag^+(aq) + 2S_2O{_3}^{2-}(aq) \lt -- \gt Ag^+(aq) + Cl^-(aq)+ [Ag(S_2O_3)_2]^{3-}(aq)$
- When you do this, you will have to multiply the equilibrium constants:
$K_{net} = 1.8 \times 10^{-10} \times 2.0 \times 10^{13} = 3.6 \times 10^{3}$
Since the constant value is greater than 1, the reaction is product favored:
