If acidic H2O2 is mixed with Fe^2+, which reaction occurs: the oxidation of Fe^2

Question

If acidic H2O2 is mixed with Fe^2+, which reaction occurs: the oxidation of Fe^2+ to Fe^3+ or the reduction of Fe^2+ to Fe? Give the balanced redox reaction for each possibility and explain your reason for selecting the reaction which will occur. Assume standard solution concentrations.

For a) I have:

red: H2O2 + 2H+ + 2e- ---> 2 H2O E=+1.78 V

oxi Fe2+ ----> Fe3+ + e- E=-0.77V

Ecell = 1.01V

I'm stuck on b) but for the explanation I think I understand that the one which is more positive is more thermodynamically favoured, but if I'm wrong on this I'd appreciate a pointer in the right direction.

Explanation / Answer

red: H2O2 + 2H+ + 2e- ---> 2 H2O E=+1.78 V

oxi Fe3+ + e-  ----> Fe2+ E=+0.77V

TRUE, the most positive is likely to reduce (i.e. follow the reaction shown) and the most negative, is likely to xidizde (reverse reaction)

therefor; the Fe reqction must be inverted so it shows the oxidation

red: H2O2 + 2H+ + 2e- ---> 2 H2O E=+1.78 V

Fe2+ ----> Fe3+ + e- E= -0.77V

In here

Ecell = Ered + Eox = 1.78 + -0.77 = 1.01 V

E°cell = 1.01 V

NOTE: E°cell must be always positive for reaction to be spontanous, therefore, if you get a negative value, this reaction is not favoured (thermodynamically)

NOTE that the reduction of Iron was not likely since the potential of H2O2 is higher

The explanation may be gone as simlpe as E°cell > 0

or

G = -nFE°cell

n = positive integer number of the numebr of electrons moving

F = 96500 C/mol farady constant (positive)

Then onlyE°cell may take either positive or negativ value

For any reaction to be spontanous

G < 0

so; in this case, E°cell > 0 in order to get a G< 0

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