Calculate the cell potential for the following unbalanced reaction that takes place in an electrochemical cell at 25 °C when [ ] = 0.500 M and [ ] = 2.00 M Mg(s) + (aq) → (aq) + Fe(s)
E°(Mg²⁺/Mg) = -2.37 V and E°(Fe³⁺/Fe) = -0.036 V

Question

Calculate the cell potential for the following unbalanced reaction that takes place in an electrochemical cell at 25 °C when [

] = 0.500 M and [ Calculate the cell potential for the following unbalanced reaction that takes place in an electrochemical cell at 25 °C when [ ] = 0.500 M and [ ] = 2.00 M Mg(s) + (aq) → (aq) + Fe(s) E°(Mg2+/Mg) = -2.37 V and E°(Fe3+/Fe) = -0.036 V A) +2.09 V B) -3.18 V C) +2.35 V D) +0.36 V E) -1.51 V

] = 2.00 M Mg(s) + Calculate the cell potential for the following unbalanced reaction that takes place in an electrochemical cell at 25 °C when [ ] = 0.500 M and [ ] = 2.00 M Mg(s) + (aq) → (aq) + Fe(s) E°(Mg2+/Mg) = -2.37 V and E°(Fe3+/Fe) = -0.036 V A) +2.09 V B) -3.18 V C) +2.35 V D) +0.36 V E) -1.51 V

(aq) →

(aq) + Fe(s)
E°(Mg²⁺/Mg) = -2.37 V and E°(Fe³⁺/Fe) = -0.036 V

Quizplus · Accepted Answer

First, we need to balance the redox reaction as follows:

Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

Now, we can use the Nernst equation to calculate the cell potential at non-standard conditions:

Ecell = E°cell - (RT/nF) ln Q

where E°cell is the standard cell potential, R is the gas constant (8.314 J/mol K), T is the temperature in kelvin (25°C = 298 K), n is the number of electrons transferred in the balanced equation (2 in this case), F is the Faraday constant (96,485 C/mol), and Q is the reaction quotient, which is the product of the concentrations of the products raised to their stoichiometric coefficients divided by the product of the concentrations of the reactants raised to their stoichiometric coefficients.

Q = [Mg2+]/[Fe2+]

At equilibrium, the cell potential is zero, so we can set Ecell = 0 and solve for Q:

0 = -2.37 V + (-0.036 V) - (RT/2F) ln [Mg2+]/[Fe2+]

ln [Mg2+]/[Fe2+] = (2F/RT) (2.37 V - 0.036 V)

[Mg2+]/[Fe2+] = exp[(2F/RT) (2.37 V - 0.036 V)] = 4.53

Now we can substitute this value of Q into the Nernst equation:

Ecell = -2.37 V - (0.0257 V) ln (4.53) = +2.35 V

Therefore, the answer is (C) +2.35 V.

Calculate the Cell Potential for the Following Unbalanced Reaction That