To calculate the value of Ecell for the given voltaic cell, we need to use the Nernst equation. The Nernst equation relates the cell potential (Ecell) to the concentrations of the species involved in the cell reaction.
The voltaic cell is represented as:
Ag(s)|Ag^+,(satd Ag2CrO4)||Ag^+(0.110 M)|Ag(s)
To use the Nernst equation, we need to identify the half-cell reactions that occur at the anode and the cathode.
Anode: Ag(s) ā Ag^+(satd Ag2CrO4) + e^-
Cathode: Ag^+(0.110 M) + e^- ā Ag(s)
The standard reduction potential (EĀ°) for the half-cell reactions can be obtained from a standard reduction potential table.
The Nernst equation is:
Ecell = EĀ°cell - (0.0592/n)logQ
where Ecell is the cell potential, EĀ°cell is the standard cell potential, n is the number of moles of electrons transferred, and Q is the reaction quotient.
In this case, the number of moles of electrons transferred (n) is 1 for both half-cell reactions.
Now, let's calculate the reaction quotient (Q) using the concentrations of the species involved:
Q = [Ag^+(satd Ag2CrO4)] / [Ag^+(0.110 M)]
Given that Ksp = [Ag^+(satd Ag2CrO4)]^2, we can solve for [Ag^+(satd Ag2CrO4)]:
Ksp = (1.1 Ć 10^-12) = [Ag^+(satd Ag2CrO4)]^2
[Ag^+(satd Ag2CrO4)] = sqrt(1.1 Ć 10^-12)
Now we can substitute the values in the Nernst equation to calculate Ecell:
Ecell = EĀ°cell - (0.0592/n)log([Ag^+(satd Ag2CrO4)] / [Ag^+(0.110 M)])
= EĀ°cell - (0.0592/1)log(sqrt(1.1 Ć 10^-12) / (0.110))
Remember to also consider the sign conventions for the half-cell reactions when plugging in the values.
Using the reference table and the concentrations provided, you can calculate Ecell.