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

At $900^{\circ} \mathrm{C}, K_{\mathrm{C}}=0.0108$ for the reaction

$$\mathrm{CaCO}_{3}(s) \rightleftharpoons \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)$$

A mixture of $\mathrm{CaCO}_{3}, \mathrm{CaO},$ and $\mathrm{CO}_{2}$ is placed in a 10.0 - $\mathrm{L}$ vessel at $900^{\circ} \mathrm{C}$ . For the following mixtures, will the

amount of $\mathrm{CaCO}_{3}$ increase, decrease, or remain the same as the system approaches equilibrium?

\begin{equation}

\begin{array}{l}{\text { (a) } 15.0 \mathrm{g} \mathrm{CaCO}_{3}, 15.0 \mathrm{g} \mathrm{CaO}, \text { and } 4.25 \mathrm{gCO}_{2}} \\ {\text { (b) } 2.50 \mathrm{g} \mathrm{CaCO}_{3}, 25.0 \mathrm{g} \mathrm{CaO}, \text { and } 5.66 \mathrm{g} \mathrm{CO}_{2}} \\ {\text { (a) } 30.5 \mathrm{g} \mathrm{CaCO}_{3}, 25.5 \mathrm{g} \mathrm{CaO}, \text { and } 6.48 \mathrm{g} \mathrm{CO}_{2}}\end{array}

\end{equation}

For the reaction below, $K_{\mathrm{p}}=1.16$ at $800 .^{\circ} \mathrm{C}$

$$\mathrm{CaCO}_{3}(s) \rightleftharpoons \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)$$

If a 20.0 -g sample of $\mathrm{CaCO}_{3}$ is put into a 10.0 -L container and heated to $800 .^{\circ} \mathrm{C},$ what percentage by mass of the $\mathrm{CaCO}_{3}$ will react to reach equilibrium?