Using a table of standard electrode potentials (as in Appendix M of your text), calculate the standard cell potential, Eo, for: 2 Fe2+(aq) + Cl2(g) 2 Fe3+(aq) + 2 Cl(aq).
Look up Eo values for
Fe^3+ + e ==> Fe^2+ and change the sign.
Look up Cl2(g) ==> 2Cl^-(aq)
Add the two for Eocell.
dont cheat
To calculate the standard cell potential, Eo, for the given reaction, you need to use a table of standard electrode potentials.
First, let's find the standard electrode potentials for the half-reactions involved in the overall reaction:
1. The half-reaction for the reduction of Fe2+ to Fe3+ is: Fe2+(aq) → Fe3+(aq)
From the table of standard electrode potentials, you will find that the standard electrode potential for this reaction is +0.77V.
2. The half-reaction for the oxidation of Cl- to Cl2 is: 2 Cl-(aq) → Cl2(g) + 2 e-
From the table of standard electrode potentials, you will find that the standard electrode potential for this reaction is +1.36V.
Now, we need to reverse the reduction half-reaction and adjust for stoichiometry to cancel out electrons:
Fe3+(aq) + e- → Fe2+(aq) (multiply this reaction by 2 to cancel out the electrons)
By reversing the reduction half-reaction, the sign of the standard electrode potential changes, so the new standard electrode potential for this reversed reaction is -0.77V.
Now, we can add the two half-reactions together:
2 Fe2+(aq) + Cl2(g) → 2 Fe3+(aq) + 2 Cl-(aq)
To calculate the standard cell potential (Eo), we simply add the standard electrode potentials of the half-reactions:
Eo = Eo(reduction) + Eo(oxidation)
= -0.77V + 1.36V
= 0.59V
Therefore, the standard cell potential (Eo) for the given reaction is 0.59V.