Calculate the standard Gibbs free energy change of reaction in the following reaction using the standard molar values for Gibbs free energy change given here. Comment on whether the reaction is Spontaneous or not, under standard conditions.

${\mathrm{NH}}_4{\mathrm{NO}}_3\left(\mathrm{s}\right)\to {\mathrm{N}}_2\mathrm{O}\left(\mathrm{g}\right)+2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right)$

Values for

$$\begin{gathered} {\mathrm{G}}^{⊝} \mathrm{in} \mathrm{kJ} {\mathrm{mol}}^{-1}; \\ {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)=-273.2 \\ {\mathrm{N}}_2\mathrm{O}\left(\mathrm{g}\right)=+104.2 \\ {\mathrm{NH}}_4{\mathrm{NO}}_3=-184.0. \end{gathered}$$

Calculate the standard Gibbs free energy change of reaction in the following reaction using the standard molar values for Gibbs free energy change given here. Comment on whether the reaction is Spontaneous or not, under standard conditions.

$2\mathrm{Mg}\left(\mathrm{s}\right)+{\mathrm{O}}_2\left(\mathrm{g}\right)\to 2\mathrm{MgO}\left(\mathrm{s}\right)$

Values for

$$\begin{gathered} {\mathrm{G}}^{⊝} \mathrm{in} \mathrm{kJ} {\mathrm{mol}}^{-1}; \\ \mathrm{MgO}\left(\mathrm{s}\right)=-569.4 \\ \mathrm{Mg}\left(\mathrm{s}\right)=0 \\ {\mathrm{O}}_2\left(\mathrm{g}\right)=0. \end{gathered}$$

Calculate the standard Gibbs free energy change of reaction in the following reaction using the standard molar values for Gibbs free energy change given here. Comment on whether the reaction is Spontaneous or not, under standard conditions.

${\mathrm{C}}_3{\mathrm{H}}_8\left(\mathrm{g}\right)+5{\mathrm{O}}_2\left(\mathrm{g}\right)\to 3{\mathrm{CO}}_2\left(\mathrm{g}\right)+4{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$

Values for

$$\begin{gathered} {\mathrm{G}}^{⊝} \mathrm{in} \mathrm{kJ} {\mathrm{mol}}^{-1}; \\ {\mathrm{CO}}_2\left(\mathrm{g}\right)=-394.4 \\ {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)=-273.2 \\ {\mathrm{O}}_2\left(\mathrm{g}\right)=0 \\ {\mathrm{C}}_3{\mathrm{H}}_8\left(\mathrm{g}\right)=-23.4. \end{gathered}$$

Calculate the standard Gibbs free energy change of reaction in the following reaction using the standard molar values for Gibbs free energy change given here. Comment on whether the reaction is Spontaneous or not, under standard conditions.

$3 \mathrm{Mg}\left(\mathrm{s}\right)+{\mathrm{Fe}}_2{\mathrm{O}}_3\left(\mathrm{s}\right)\to 3\mathrm{MgO}\left(\mathrm{s}\right)+2\mathrm{Fe}\left(\mathrm{s}\right)$

Values for

$$\begin{gathered} {\mathrm{G}}^{⊝} \mathrm{in} \mathrm{kJ} {\mathrm{mol}}^{-1}; \\ \mathrm{MgO}\left(\mathrm{s}\right)=-569.4 \\ \mathrm{Fe}\left(\mathrm{s}\right)=0 \\ \mathrm{Mg}\left(\mathrm{s}\right)=0 \\ {\mathrm{Fe}}_2{\mathrm{O}}_3\left(\mathrm{s}\right)=-742.2. \end{gathered}$$

Calculate the Gibbs free energy for the following reaction and state. Giving a reason for the position of equilibrium for the reaction.

$$\begin{gathered} \mathrm{Sn}\left(\mathrm{s}\right)+{\mathrm{Pb}}^{2+}\left(\mathrm{aq}\right)\to {\mathrm{Sn}}^{2+}\left(\mathrm{aq}\right)+\mathrm{Pb}\left(\mathrm{s}\right) \\ {\mathrm{Sn}}^{2+}\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\rightleftharpoons \mathrm{Sn}\left(\mathrm{s}\right); {\mathrm{E}}^{⊝}=-0.14 \mathrm{V} \\ {\mathrm{Pb}}^{2+}\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\rightleftharpoons \mathrm{Pb}\left(\mathrm{s}\right); {\mathrm{E}}^{⊝}=-0.13 \mathrm{V} \end{gathered}$$

Calculate the Gibbs free energy for the following reaction and state. Giving a reason for the position of equilibrium for the reaction.

$$\begin{gathered} {\mathrm{Fe}}^{3+}\left(\mathrm{aq}\right)+\mathrm{Ag}\left(\mathrm{aq}\right)\to {\mathrm{Fe}}^{2+}\left(\mathrm{aq}\right)+{\mathrm{Ag}}^{+}\left(\mathrm{aq}\right) \\ {\mathrm{Fe}}^{3+}\left(\mathrm{aq}\right)+{\mathrm{e}}^{-}\rightleftharpoons {\mathrm{Fe}}^{2+}; {\mathrm{E}}^{⊝}=+0.77 \mathrm{V} \\ {\mathrm{Ag}}^{+}\left(\mathrm{aq}\right)+{\mathrm{e}}^{-}\rightleftharpoons \mathrm{Ag}\left(\mathrm{aq}\right); {\mathrm{E}}^{⊝}=+0.80 \mathrm{V} \end{gathered}$$

Calculate the Gibbs free energy for the following reaction and state. Giving a reason for the position of equilibrium for the reaction.

When ethanol undergoes combustion, carbon dioxide and water are formed.

${\mathrm{C}}_2{\mathrm{H}}_5\mathrm{OH}\left(\mathrm{l}\right)+3{\mathrm{O}}_2\left(\mathrm{g}\right)\to 2{\mathrm{CO}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$.

Calculate the standard entropy change for this reaction.

Values for

 $$\begin{gathered} {\mathrm{S}}^{⊝}\mathrm{in} \mathrm{J} {\mathrm{K}}^{-1}{\mathrm{mol}}^{-1}; \\ {\mathrm{C}}_2{\mathrm{H}}_5\mathrm{OH}\left(\mathrm{l}\right)=160.7, \\ {\mathrm{O}}_2\left(\mathrm{g}\right)=205.0, \\ {\mathrm{CO}}_2\left(\mathrm{g}\right)=213.6, \\ {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)=69.90. \end{gathered}$$

Calculate the Gibbs free change for the given reaction.

When ethanol undergoes combustion, carbon dioxide and water are formed.

${\mathrm{C}}_2{\mathrm{H}}_5\mathrm{OH}\left(\mathrm{l}\right)+3{\mathrm{O}}_2\left(\mathrm{g}\right)\to 2{\mathrm{CO}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$.

Values for

 $$\begin{gathered} ∆{\mathrm{G}}_{\mathrm{f}}^{⊝}\mathrm{in} \mathrm{kJ} {\mathrm{mol}}^{-1}: \\ {\mathrm{C}}_2{\mathrm{H}}_5\mathrm{OH}\left(\mathrm{l}\right)=-174.9, \\ {\mathrm{O}}_2\left(\mathrm{g}\right)=0, \\ {\mathrm{CO}}_2\left(\mathrm{g}\right)=-394.4, \\ {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)=-237.2. \end{gathered}$$