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

ID: 882431 • Letter: 1

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 E0cell for the cell?

Realizing the nonstandard concentrations, what is the actual cell potential, E0cell for the cell? What is the value of the Nernst equation

anode reaction: oxidation takes place

Cr (s) -------------------------> Cr+3 (aq) + 3 e- , E0Cr+3/Cr = - 0.74 V

cathode reaction : reduction takes palce

3 Ag+ (aq) + 3e- ----------------------------->3 Ag (s) , E0Ag+/Ag = + 0.80 V

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net cell reaction:

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

E0cell= E0cathode- E0anode

E0cell= E0Ag+/Ag - E0Cr+3/Cr

E0cell = 0.80 - (-0.74)

E0cell = 1.54V

nernest equation

Ecell = E0cell -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 = E0cell -(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

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