Calculate the value of E°_(cell) and Δ G° that can be obtained from the following cell under the…
(a)[3]
Calculate the value of $E^\circ_{\text{cell}}$ and $\Delta G^\circ$ that can be obtained from the following cell under the standard conditions at $25^\circ\text{C}$:
Given $E^\circ_{(\text{Zn}^{2+}/\text{Zn})} = -0\cdot76\text{ V}$; $E^\circ_{(\text{Sn}^{2+}/\text{Sn})} = -0\cdot14\text{ V}$
$1\text{ Faraday} = 96500\text{ C mol}^{-1}$
(b)[2]
How much electricity in Faraday is required for the complete reduction of $\text{MnO}_4^-$ ions present in $500\text{ ml}$ of $0\cdot5\text{ M}$ solution to $\text{Mn}^{2+}$?
Formulas used: Standard emf of a cell: E0cell = E0cathode - E0anode with E0anode = -0.76 V, E0cathode = -0.14 V, so E0cell = 0.62 V
Gibbs energy from emf: dG = -n*F*E with E = 0.62 V, F = 96500 C mol^-1, n = 2, so dG = -119.66 kJ mol^-1
Zn is the anode and Sn the cathode.
$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode} = -0\cdot14 - (-0\cdot76) = 0\cdot62\ \mathrm{V}$
$\Delta G^\circ = -nFE^\circ_{cell} = -2 \times 96500\ \mathrm{C\ mol^{-1}} \times 0\cdot62\ \mathrm{V} = -119660\ \mathrm{J\ mol^{-1}}$
