**The Nernst Equation at 25****°C:**

$\overline{){{\mathbf{E}}}_{{\mathbf{cell}}}{\mathbf{}}{\mathbf{=}}{\mathbf{}}{{\mathbf{E}}^{\mathbf{o}}}_{{\mathbf{cell}}}{\mathbf{}}{\mathbf{-}}{\mathbf{}}\mathbf{\left(}\frac{\mathbf{0}\mathbf{.}\mathbf{05916}\mathbf{}\mathbf{V}}{\mathbf{n}}\mathbf{\right)}{\mathbf{logQ}}}$

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

E°_{cell} = standard cell potential

n = number of e^{-} transferred

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

We have to determine the E°_{cell} first and as well as the **anode (oxidation)** and **cathode (reduction)** in the concentration cell and the **number of electrons transferred (n)**.

Calculate the cell potential for the following reaction that takes place in an electrochemical cell at 25 degrees C.

Sn(s) │ Sn2+ (aq, 0.022M) ║ Ag+ (aq, 2.7M) │Ag(s)

Write the net cell equation. Phases are optional. do not include the concentrations

Calculate the following values at 25.0 C using standard potentials as needed.

E°_{cell} =

E_{cell} =

ΔG°_{rxn} =

ΔG_{rxn} =

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