Determine what redox reaction, if any, occurs (at 25° C) when lead metal (Pb) is added to 1.0-M solution of NiCl2
Ecell = Ecathode- Eanode
Here is standard reduction potentials given for Pb and Cl
Pb2+(aq) + 2e- → Pb(s) E°(V) = -0.13
Cl2(g) 2e- → 2Cl-(aq) E°(V) = +1.36
The Attempt at a Solution
According to my book, the half reaction with the greater reduction potential is the one that is the cathode, so that must be +1.36. So using the equation above, the 1.36 - (-0.13) = 1.49, which is a positive number. According to my book, that reaction should occur. But the answer in the back of the book says "no reaction occurs". I don't understand.
Am I supposed to be using the Nickel in the reaction instead of the Chlorine?
The standard reduction potential they give for Nickel is:
Ni2+(aq) + 2e- → Ni(s) E°(V) = -0.25
Edit: The next question has me adding Pb to HCl, but in the answers, it shows that the Hydrogen is being used in the equation instead of the Cl. I really don't understand what's going on. How do I know which one to use?