The equilibrium constant for the reaction

$$2 \mathrm{Fe}^{3+}(a q)+\mathrm{Hg}_{2}^{2+}(a q) \rightleftharpoons 2 \mathrm{Fe}^{2+}(a q)+2 \mathrm{Hg}^{2+}(a q)$$

is $K_{c}=9.1 \times 10^{-6}$ at 298 $\mathrm{K}$

(a) What is $\Delta G^{\circ}$ at this temperature?

(b) If standard-state concentrations of the reactants and products are mixed, in which direction does the reaction proceed?

(c) Calculate $\Delta G$ when $\left[\mathrm{Fe}^{3+}\right]=0.20 M,\left[\mathrm{Hg}_{2}^{2+}\right]=0.010 M,$ $\left[\mathrm{Fe}^{2+}\right]=0.010 M,$ and $\left[\mathrm{Hg}^{2+}\right]=0.025 M$ . In which direction will the reaction proceed to achieve equilibrium?

## Discussion

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## Recommended Questions

Given that the $\Delta G_{\mathrm{f}}^{\circ}$ for $\mathrm{Pb}^{2+}(a q)$ and $\mathrm{Cl}^{-}(a q)$ is $-24.3 \mathrm{kJ} / \mathrm{mole}$ and $-131.2 \mathrm{kJ} / \mathrm{mole}$ respectively, determine the solubility product, $K_{\mathrm{sp}},$ for $\mathrm{PbCl}_{2}(s)$

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