Consider the electrochemical cell with the following half-cell reactions: Fe3+ (aq) + e- → Fe2+ (aq) E° = +0.771 V and Cd2+ (aq) + 2e- → Cd (s) E° = -0.403 V What is the overall cell potential, Ecell, under standard conditions? Which species will be reduced at the cathode? Which electrode is the anode? Justify your choices using the given standard reduction potentials.

Question

Consider the electrochemical cell with the following half-cell reactions: Fe3+ (aq) + e- → Fe2+ (aq)    E° = +0.771 V          and Cd2+ (aq) + 2e- → Cd (s)    E° = -0.403 V What is the overall cell potential, Ecell, under standard conditions? Which species will be reduced at the cathode? Which electrode is the anode? Justify your choices using the given standard reduction pot

Answer 1

To determine the overall cell potential  Ecell  under standard conditions, we need to identify which half-cell reaction will act as the reduction  cathode  and which will act as the oxidation  anode . The species with the higher standard reduction potential  E  will be reduced, while the species with the lower standard reduction potential will be oxidized.In this case, Fe3+ has a higher standard reduction potential  +0.771 V  than Cd2+  -0.403 V . Therefore, Fe3+ will be reduced at the cathode, and Cd2+ will be oxidized at the anode. The half-cell reactions will be:Cathode  reduction : Fe3+  aq  + e-  Fe2+  aq     E = +0.771 VAnode  oxidation : Cd  s   Cd2+  aq  + 2e-    E = +0.403 V  reverse of the given reaction Now, we can calculate the overall cell potential  Ecell  by adding the standard reduction potentials of the two half-cell reactions:Ecell = E cathode  - E anode Ecell =  +0.771 V  -  +0.403 V Ecell = +0.368 VThe overall cell potential under standard conditions is +0.368 V. The species reduced at the cathode is Fe3+, and the electrode acting as the anode is the one where Cd is oxidized to Cd2+.