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I'm working through TPR MCAT 2015 book and am stuck on the following question:
The voltage of a galvanic cell composed of one Zn and one Fe electrode, in 1 M aqueous solutions of their respective ions is measured to be 0.33 V, confirming the tabulated half-reaction values below, referenced against the standard hydrogen electrode
Fe2+ + 2e --> Fe E0=-0.44
Zn --> 2e + Zn2+ E0=+0.77
If enough H2SO4 is added to each chamber of the cell to make the solution 1 M in H+, how will the cell's deltaG change?
A.) -85 kJ/mol
B.) -148 kJ/mol
C.) +85 kJ/mol
D.) +148 kJ/mol
I've read over the explanation as well as the electrochemistry chapter for the second time, but I still don't understand. Any help with this problem would be appreciated!
The voltage of a galvanic cell composed of one Zn and one Fe electrode, in 1 M aqueous solutions of their respective ions is measured to be 0.33 V, confirming the tabulated half-reaction values below, referenced against the standard hydrogen electrode
Fe2+ + 2e --> Fe E0=-0.44
Zn --> 2e + Zn2+ E0=+0.77
If enough H2SO4 is added to each chamber of the cell to make the solution 1 M in H+, how will the cell's deltaG change?
A.) -85 kJ/mol
B.) -148 kJ/mol
C.) +85 kJ/mol
D.) +148 kJ/mol
I've read over the explanation as well as the electrochemistry chapter for the second time, but I still don't understand. Any help with this problem would be appreciated!
