Given the following spontaneous reaction occurring in an electrochemical cell un

Question

Given the following spontaneous reaction occurring in an electrochemical cell under standard condition (25C, 1 atm, 1 mol/L)

1. Write the cell notation when the above reaction serves as a battery rxn.

2. Use the standard half-cell potentials listed aboce to calculate the standard cell potential.

4. If the above battery reaction reaches its equilibrium, then what is the equilibrium constant K?

Given the following spontaneous reaction occurring in an electrochemical cell under standard condition (25C, 1 atm, 1 mol/L) 3Cl2(g)+ 2Fe(s) DeltaG) of the battery reaction? 4. If the above battery reaction reaches its equilibrium, then what is the equilibrium constant K? --> Fe(s) E= -0.04 V 1. Write the cell notation when the above reaction serves as a battery rxn. 2. Use the standard half-cell potentials listed aboce to calculate the standard cell potential. 3. What is the standard free energy change (---> 2Cl- (aq) E= +1.36 V Fe3+(aq)+ 3e- --> 6Cl -(aq)+ 2Fe3+(aq) Cl2(g) + 2e- ---> 2Fe(s)

Explanation / Answer

1) cell notation Fe(s)/Fe3+(aq) // Cl2(g)/Cl- (aq) ,Pt

2)    E0 = E0Cl2/Cl- - E0Fe+3/Fe

     = +1.36- (-0.04)

   = 1.40V

3) detla G0 = -nF Ecell0

here n= 6 moles of electrons transfered

detla G0 = -nF Ecell0

                 = -6 x 96500 x1.4

              = -810.6kJ/mol

4) ?G?= - 2.303RT logKc

-810.6 = -2 303 x 8.314 x10^-3 x 298 x logKc

Kc= 10^142.03

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