# electrochemistry experiment

something strange. i made an electrochemistry experiment in which i have measured the potential of cell: Ag, AgCl | (1M)KCl ║ [Fe(CN)6] 3-, [Fe(CN)6] 4- | Pt

with different concentrations of Fe+2/Fe+3 . from the results i have found the faraday constant: 96181 Coulon /mol
very close to the published value (99.7% accuracy).

the intercept of the graph of E (v) against ln(Fe+2/Fe+3) is:

intercept:    0.198985714  V


(the intercept is also the standard potential for the cell)

yet, when i tried to find the standard potential of

Fe3+ +e- -> Fe2+


i got the result of +0.423 V which is almost half of the published value (+0.77 V)

E_cathode =E^° + E_anode

E_cathode=0.198985+0.224= +0.423 V


how can it be that the faraday constant is almost accurate and in the same time the result of the standard potential is far from the real one.? i'm sure i don't have any calculation mistakes, if it's necessary i can describe the method of work here.

thanks!

What you are measuring is not the potential $$Fe^{2+}/Fe^{3+}$$. It is the potentiel of the couple $$[Fe(CN)_6]^{4-}/[Fe(CN)_6]^{3-}$$ which is equal to +0.36 V in the table.