Final answer:
The question involves a redox reaction in acidic solution. Calculate Ecell by adding the half-reaction potentials, then use Ecell to find ΔG with the equation ΔG = -nFEcell, and K using the Nernst equation. Ecell is positive, indicating spontaneity. The standard cell potential, Ecell, for the given redox reaction is 0.34 V.
Step-by-step explanation:
In order to calculate the standard cell potential, Ecell, for the given redox reaction, we need to use the Nernst equation:
Ecell = E°cell - (0.0592/n) log(Q)
where E°cell is the standard cell potential, n is the number of electrons transferred in the balanced equation, and Q is the reaction quotient.
In this case, the balanced equation is:
3Cu(s) + 8H+(aq) + 2NO3-(aq) → 3Cu2+(aq) + 2NO(g) + 4H2O(l)
n = 6 (since 6 electrons are transferred)
The standard cell potential, E°cell, can be determined by looking up the reduction potentials for each half-reaction:
Cu2+(aq) + 2e- → Cu(s) (E° = 0.34 V)
NO3-(aq) + 4H+(aq) + 3e- → NO(g) + 2H2O(l) (E° = 0.96 V)
Substituting the values into the Nernst equation:
Ecell = 0.34 V - (0.0592/6) log(Q)
where Q is the reaction quotient:
Q = [Cu2+]^3 / ([H+]^8 [NO3-]^2)
Given the standard concentrations of Cu2+, H+, and NO3- are all 1 M, we can substitute these values into the equation:
Q = (1^3) / (1^8 * 1^2) = 1
Plugging this value into the Nernst equation:
Ecell = 0.34 V - (0.0592/6) log(1) = 0.34 V
Therefore, the standard cell potential, Ecell, for this reaction is 0.34 V.