# Standard reduction potential and half reactions

## Homework Statement

Determine what redox reaction, if any, occurs (at 25° C) when lead metal (Pb) is added to 1.0-M solution of NiCl2

## Homework Equations

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

Thanks.

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?

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The half reaction with the greatest reduction potential will get ectrons and become metal.

but $$E^o(Pb) > E^ô(Ni)$$, implies no reaction occurs
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no reaction occurs,
$$Pb + Ni^{2+}=/=>Pb^{2+}+Ni$$

but the opposite reaction occurs

$$Ni + Pb^{2+}==>Ni^{2+}+Pb$$

because
$$E^o(Pb) > E^ô(Ni)$$,

Thanks for the response.

I'm given a solution of NiCl2. How do I know to check the reduction potential of Nickel instead of Chlorine?

The worked example they give had them add molecular bromine (Br2) to a solution of NaI, and they ignored the Na and checked the reduction potential of the I. They had another one where they put Br2 in NaCl, and they check the reduction potential of the Cl. But with my problem, I'm putting Pb in NiCl2, and now, unlike the worked example, I ignore the Cl and check the reduction potential of the Ni? How do I know which one to check?

Thanks.