Question

About General Chem Electrochemistry (Electron transer) chapter.

I have a hard time understading which to flip when determing the oxidation-reduction equations.

By looking at the half equations, How would you decide which one is the cathode or anode? I know that oxidation occurs at the anode which needs to be filped when writing an overall equation.

I know how to decide when there is an overall equation because you can decide which one is oxidized or reduced by calculating the oxidation numbers.

BUT, How would you decide when you only have the half- equations?

Because Sometimes, Oxidation occurs with the smaller E value ( the value you find from the standard reduction potential) and Sometimes It's NOT....

Please clarify this as soon as possible. Thank you !!!

Answer 1

Remember that each species will have a specific reduction potential. Remember that this is, as the name implies, a potential to reduce. We use it to compare it (numerical) with other species.

Note that the basis if 2H+ + 2e- -> H2(g) reduction. Therefore E° = 0 V

All other samples are based on this reference.

Find the Reduction Potential of each reaction (Tables)

Zn2+ + 2 e− <-> Zn(s) Ered = −0.762 V

F2(g) + 2 e− <-> 2 F−   Ered = +2.87 V

The most positive has more potential to reduce, it will be reduced

The most negative will be oxidized, since it will donate it selectrons

For total E°cell potential:

E°cell = Ered – Eox

Eox = -Ered of the one being oxidized

E°cell = 2.87 - (-0.7618) = 3.6318 V

E°cell = 3.6318 V

special tips:

if you want the reaction to be spontaneous, then you must have E°cell value POSITIVE

which implies that we need E°cell = Ecathode - Eanode > 0

this is achieved always flipping the MOST NEGATIVE