Problem: Using standard electrode potentials calculate ΔG° and use its value to estimate the equilibrium constant for each of the reactions at 25 °C.Br2(l) +2 Cl-(aq)→2 Br-(aq) + Cl2(g)

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Br2(l) + 2 Cl-(aq) → 2 Br-(aq) + Cl2(g)

Recall:

Lose               Gain
Electron         Electrons
Oxidation       Reduction

cathode → reduction → oxidation number decreases

anode → oxidation → oxidation number increases

For ions: charge of the ion = oxidation number
For neutral atoms/compound: oxidation number = 0

For Br: There is a decrease in oxidation number, therefore, it undergoes reduction and is the cathode.

Br2(l) + 2e- → 2 Br-(aq) E° = 1.09 V

For Cl: There is an increase in oxidation number, therefore, it undergoes oxidation and is the anode.

Cl2(g) + 2e- → 2 Cl-(aq)  E° = 1.36 V

E°cell=E°cathode-E°anodeEcell=1.09V-1.36 V  

  E°cell = -0.27 V

G°=-nFE°cell

ΔG° = Gibbs Free Energy, J
n = # of e- transferred
F = Faraday’s constant = 96485 J/(mol e-)
cell = standard cell potential, V

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Problem Details

Using standard electrode potentials calculate ΔG° and use its value to estimate the equilibrium constant for each of the reactions at 25 °C.

Br2(l) +2 Cl-(aq)→2 Br-(aq) + Cl2(g)

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