330433 Which from the following is the correct relationship between standard Gibbs energy change and standard cell potential?

330434 Following two half cells form a complete cell which has \(\Delta \mathrm{G}^{\circ}\) (in \(\mathrm{kJ}\) ) value
\({\text{2}}{{\text{H}}^{\text{ + }}}{\text{ + 1/2}}{{\text{O}}_{\text{2}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {{\text{H}}_{\text{2}}}{\text{O; E}}^\circ \,{\text{ = + 1}}{\text{.23 V }}\)
\({\text{F}}{{\text{e}}^{{\text{2 + }}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {\text{Fe(s);E}}^\circ {\text{ = - 0}}{\text{.44 V}}\)

330435 \(\Delta \mathrm{G}^{\circ}\) for the reaction, \(\mathrm{Cu}^{2+}+\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+\mathrm{Cu}\) is: [given : \(\mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\circ}=+0.34 \mathrm{~V}, \mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.44 \mathrm{~V}\) ]

330436 The potential of standard hydrogen electrode is zero. This implies that

330433 Which from the following is the correct relationship between standard Gibbs energy change and standard cell potential?

330434 Following two half cells form a complete cell which has \(\Delta \mathrm{G}^{\circ}\) (in \(\mathrm{kJ}\) ) value
\({\text{2}}{{\text{H}}^{\text{ + }}}{\text{ + 1/2}}{{\text{O}}_{\text{2}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {{\text{H}}_{\text{2}}}{\text{O; E}}^\circ \,{\text{ = + 1}}{\text{.23 V }}\)
\({\text{F}}{{\text{e}}^{{\text{2 + }}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {\text{Fe(s);E}}^\circ {\text{ = - 0}}{\text{.44 V}}\)

330435 \(\Delta \mathrm{G}^{\circ}\) for the reaction, \(\mathrm{Cu}^{2+}+\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+\mathrm{Cu}\) is: [given : \(\mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\circ}=+0.34 \mathrm{~V}, \mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.44 \mathrm{~V}\) ]

330436 The potential of standard hydrogen electrode is zero. This implies that

330433 Which from the following is the correct relationship between standard Gibbs energy change and standard cell potential?

330434 Following two half cells form a complete cell which has \(\Delta \mathrm{G}^{\circ}\) (in \(\mathrm{kJ}\) ) value
\({\text{2}}{{\text{H}}^{\text{ + }}}{\text{ + 1/2}}{{\text{O}}_{\text{2}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {{\text{H}}_{\text{2}}}{\text{O; E}}^\circ \,{\text{ = + 1}}{\text{.23 V }}\)
\({\text{F}}{{\text{e}}^{{\text{2 + }}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {\text{Fe(s);E}}^\circ {\text{ = - 0}}{\text{.44 V}}\)

330435 \(\Delta \mathrm{G}^{\circ}\) for the reaction, \(\mathrm{Cu}^{2+}+\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+\mathrm{Cu}\) is: [given : \(\mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\circ}=+0.34 \mathrm{~V}, \mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.44 \mathrm{~V}\) ]

330436 The potential of standard hydrogen electrode is zero. This implies that

330433 Which from the following is the correct relationship between standard Gibbs energy change and standard cell potential?

330434 Following two half cells form a complete cell which has \(\Delta \mathrm{G}^{\circ}\) (in \(\mathrm{kJ}\) ) value
\({\text{2}}{{\text{H}}^{\text{ + }}}{\text{ + 1/2}}{{\text{O}}_{\text{2}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {{\text{H}}_{\text{2}}}{\text{O; E}}^\circ \,{\text{ = + 1}}{\text{.23 V }}\)
\({\text{F}}{{\text{e}}^{{\text{2 + }}}}{\text{ + 2}}{{\text{e}}^{\text{ - }}} \to {\text{Fe(s);E}}^\circ {\text{ = - 0}}{\text{.44 V}}\)

330435 \(\Delta \mathrm{G}^{\circ}\) for the reaction, \(\mathrm{Cu}^{2+}+\mathrm{Fe} \rightarrow \mathrm{Fe}^{2+}+\mathrm{Cu}\) is: [given : \(\mathrm{E}_{\mathrm{Cu}^{2+} / \mathrm{Cu}}^{\circ}=+0.34 \mathrm{~V}, \mathrm{E}_{\mathrm{Fe}^{2+} / \mathrm{Fe}}^{\circ}=-0.44 \mathrm{~V}\) ]

330436 The potential of standard hydrogen electrode is zero. This implies that