Given the following spontaneous reaction occurring in an electrochemical cell un

Question

Given the following spontaneous reaction occurring in an electrochemical cell under standard condition (25 degree C, 1 atm, 1 mol/L...) 3 Cl_2 (g) + 2 Fe (s) rightarrow 6 Cl^- (aq) + 2 Fe^3+ (aq) Cl_2 (g) + 2e^- rightarrow 2 Cl^- (aq) E degree = +1.36 V Fe^3+ (aq) + 3e^- rightarrow Fe (s) E degree = -0.04 V Write the cell notation when the above reaction Eq. 1 serves as a battery rxn. Is Eq. 1 a spontaneous rxn? Is it a redox rxn? Which species are oxidant and reductant respectively? Use the standard half-cell potentials listed above to calculate the standard cell potential. What is the standard free energy change (Delta G degree) of the battery reaction? If the above battery reaction reaches its equilibrium (a dead battery), then what is its equilibrium constant K?

Explanation / Answer

1)The cell notation is

Fe/Fe+3 || Pt,Cl2/2Cl-

2) It is a spontaneous redox reaction , Cl2 is the oxidant/oxidising agent

Fe is the reducing agent/reductant

3) E cell = SRP of right hand electrode - SRP of left hand electrode

=+1.36 -(-0.04)

= 1.40V

4)We know that

Delta G = -nFE

= -6 x96500 C x 1.40V

= 810600 J

= 810.6 kJ

5) When the battery reaches equilibrium , Ecell = 0 and delta G = 0

then the relation is Delta G = -2.303x RT log Keq.

Since Delta G=0, E = 0 then Keq = 1.0

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