We are asked to find the **emf** of the given reaction. We will use the **Nernst Equation** to calculate the cell potential with the given conditions.

The Nernst Equation relates the concentrations of compounds and cell potential.

$\overline{){{\mathbf{E}}}_{{\mathbf{cell}}}{\mathbf{=}}{\mathbf{E}}{{\mathbf{\xb0}}}_{{\mathbf{cell}}}{\mathbf{-}}\mathbf{\left(}\frac{\mathbf{0}\mathbf{.}\mathbf{0592}\mathbf{}\mathbf{V}}{\mathbf{n}}\mathbf{\right)}{\mathbf{logQ}}}$

E_{cell} = cell potential under non-standard conditions

E°_{cell} = standard cell potential

n = number of e^{-} transferred

Q= reaction quotient = [products]/[reactants]

**Let’s first determine how many electrons were transferred in the reaction:**

A voltaic cell is constructed with two silver-silver chloride electrodes, each of which is based on the following half-reaction:

AgCl(s) + e^{-} → Ag(s) + Cl^{-} (aq).

The two cell compartments have [Cl^{-} ] = 1.51×10^{−2} M and [Cl^{-} ] = 2.60 M , respectively.

What is the cell emf for the concentrations given?

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