Question

What is the ΔG⁰ for the following reaction?

$$\mathrm{Al}{\phantom{A}}_{(\mathrm{s})}+\mathrm{Fe}{\phantom{A}}_{(\mathrm{aq})}^{3+}⟶\mathrm{Al}{\phantom{A}}_{(\mathrm{aq})}^{3+}+\mathrm{Fe}{\phantom{A}}_{(\mathrm{s})}$$; $$\mathrm{E}{\phantom{A}}_{\mathrm{cell}}^0$$ = +2.43

पर्याय

• −703 kJ

• −469 kJ

• −235 kJ

• −173 kJ

Answer 1

−703 kJ

Explanation:

For the reaction, n = 3.

ΔG⁰ = −n F ${\displaystyle {\text{E}}_{\text{cell}}^0}$

= −3 × 96500 × 2.43

= −703485 J

= −703 kJ