Question

1. Consider a galvanic cell consisting of the following two redox couples:

           

            Ag⁺(0.010M) + e^- --> Ag(s)    E⁰ = +0.80 V

         Cr³⁺(0.010M) + e^- --> Cr(s)   E⁰ = -0.74 V

        

Write the equation for the half reaction occurring at the cathode

Write the equation for the half reaction occurring at the anode.

Write the equation for the cell

What is the standard cell potential E⁰_cell for the cell?

Realizing the nonstandard concentrations, what is the actual cell potential, E⁰_cell for the cell? What is the value of the Nernst equation

Answer 1

anode reaction: oxidation takes place

Cr (s) -------------------------> Cr+3 (aq) + 3 e-   ,   E⁰_Cr+3/Cr = - 0.74 V

cathode reaction : reduction takes palce

3 Ag+ (aq) + 3e- ----------------------------->3 Ag (s) , E⁰_Ag+/Ag = + 0.80 V

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net cell reaction:

Cr(s) +3 Ag+ (aq) -------------------------> Cr+3 (aq) + 3 Ag (s)

E⁰_cell= E⁰_cathode- E⁰_anode

E⁰_cell= E⁰_Ag+/Ag - E⁰_Cr+3/Cr

  E⁰_cell = 0.80 - (-0.74)

E⁰_cell = 1.54V

nernest equation

Ecell = E⁰_cell -2.303RT/nF* log [Zn+2]/[Fe+2]

Here R= universal gas constant 8.314 J/K mol

T = absolute temperature =25(0C)= 298k

F= faraday = 96500 Coloumb/mol

     n   = no of moles of electrons are transfered =2

2.303RT/F= 0.0591

Ecell = E⁰_cell -(0.0591/n)* log [Cr+3]/[Ag+]^3

Ecell = 1.54 - (0.059 x1/3) * {log 0.01/ (0.01)^3}

Ecell   = 1.46 V

cell potential =Ecell   = 1.46 V