Calculate the emf and ΔG of the cell reaction for the following cell at 25°C: $$\mathrm{Mg(s)\,|\,Mg^{2+}(0.001\ M)\,||\,Sn^{2+}(0.1\ M)\,|\,Sn(s)}$$ ($E^{\circ}_{\mathrm{Mg^{2+}/Mg}} = -2.34$ V; $E^{\circ}_{\mathrm{Sn^{2+}/Sn}} = -0.136$ V; 1 F = 96500 C mol$^{-1}$)

CBSE Class 12 · 2003 · Delhi · Redox Reactions and ElectrochemistryShort Answer

Calculate the emf and ΔG of the cell reaction for the following cell at 25°C:

Mg (s) ∣ M g^{2 +} (0.001 M) ∣∣ S n^{2 +} (0.1 M) ∣ Sn (s)

(

E_{M g^{2 +} /Mg}^{\circ} = - 2.34

E_{S n^{2 +} /Sn}^{\circ} = - 0.136

V; 1 F = 96500 C mol

^{- 1}

)

Answer: E_cell ≈ 2.26 V; ΔG ≈ −437 kJ mol⁻¹

E_{ce l l}^{\circ} = - 0.136 - (- 2.34) = 2.204 V, n = 2

Nernst equation:

E_{ce l l} = 2.204 - \frac{0.059}{2} lo g \frac{1 0 ^{- 3}}{0.1} = 2.204 - 0.0295 \times (- 2)

= 2.204 + 0.059

E_{ce l l} = 2.263 V

Δ G = - n F E_{ce l l} = - 2 \times 96500 \times 2.263

Δ G = - 4.37 \times 1 0^{5} J = - 436.8 kJ mo l^{- 1}

CBSE Class 12 Board · 2003 · Delhi · Subjective