The document covers the fundamentals of electrochemistry, focusing on oxidation/reduction (redox) reactions and electrode potentials. It discusses standard reduction potentials for various oxidizing and reducing agents, illustrating their roles in voltaic cells and the computation of cell potential. Additionally, the Nernst equation is introduced to explain the effect of concentrations on electrode potentials.

1. Electrochemical Cells and Electrode
Potentials
Dr. S. Sreenivasa
Associate Professor and Chairman
DOS and R in Organic Chemistry
Tumkur University, Tumakuru

2. Electrochemistry
Oxidation/Reduction Reactions
• “Redox” reactions involve electron transfer from
one species to another
• Ox1 + Red2  Red1 + Ox2
• Ox1 + ne-  Red1 (Reduction ½ reaction)
• Red2  Ox2 + ne- (Oxidation ½ reaction)
• “Reducing agent” donates electrons (is oxidezed)
• “Oxidizing agent” accepts electrons (is reduced)

3. Electrochemistry
Oxidation/Reduction Reactions
• Typical oxidizing agents: Standard Potentials,V
– O2 + 4H+ + 4e-  2H2O +1.229
– Ce4+ + e-  Ce3+ +1.6 (acid)
– MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O +1.51
• Typical reducing agents:
– Zn2+ + 2e-  Zno -0.763
– Cr3+ + e-  Cr2+ -0.408
– Na+ + e-  Nao -2.714

4. Fig. 12.1. Voltaic cell.
The salt bridge allows charge transfer through the solution and prevents mixing.
The spontaneous cell reaction (Fe2+ + Ce4+ = Fe3+ + Ce4+) generates the cell potential.
The cell potential depends on the half-reaction potentials at each electrode.
The Nernst equation describes the concentration dependence.
A battery is a voltaic cell. It goes dead when the reaction is complete (Ecell = 0).
©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

5. Electrochemistry
Standard Reduction Potentials
• Half-Reaction Potentials:
• They are measured relative to each other
• Reference reduction half-reaction:
• standard hydrogen electrode (SHE)
• normal hydrogen electrode (NHE)
• 2H+(a=1.0) + 2e-  H2(g 1atm) 0.0000 volts

6. The more positive the Eo, the better oxidizing agent is the oxidized form (e.g., MnO4
-).
The more negative the Eo, the better reducing agent is the reduced form (e.g., Zn).
©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

7. Electrochemistry
Reduction Potentials
• General Conclusions:
• 1. The more positive the electrode potential, the
stronger an oxidizing agent the oxidized form is
and the weaker a reducing agent the reduced
form is
• 2. The more negative the reduction potential, the
weaker the oxidizing agent is the oxidized formis
and the stronger the reducing agent the reduced
form is.

8. Electrochemistry
Oxidation/Reduction Reactions
• Typical oxidizing agents: Standard Potentials,V
– O2 + 4H+ + 4e-  2H2O +1.229
– Ce4+ + e-  Ce3+ +1.6 (acid)
– MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O +1.51
• Typical reducing agents:
– Zn2+ + 2e-  Zno -0.763
– Cr3+ + e-  Cr2+ -0.408
– Na+ + e-  Nao -2.714

9. Electrochemistry
Oxidation/Reduction Reactions
• Net Redox Reactions: Standard Potentials,V
• MnO4
-  Mn2+
• MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O +1.51
• Sn4+ + 2e-  Sn2+ +0.154
• Balanced Net Ionic Reaction:
• 2MnO4
- + 16H+ + 5Sn2+  2Mn2+ + 5Sn4+ + 8H2O

10. Electrochemistry
Voltaic Cell
• The spontaneous (Voltaic) cell reaction is the one that
gives a positive cell voltage when subtracting one half-
reaction from the other.
• Eo
cell = Eo
right – Eo
left = Eo
cathode – Eo
anode =Eo
+ - Eo
-
• Which is the Anode? The Cathode?
• Convention:
• The anode is the electrode where oxidation occurs 
the more negative half-reaction potential
• The cathode is the electrode where reduction occurs
 the more positive half-reaction potential
• anode  solution  cathode

11. Electrochemistry
Oxidation/Reduction Reactions
• Net Redox Reactions: Standard Potentials,V
• MnO4
-  Mn2+
• MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O +1.51
• Sn4+ + 2e-  Sn2+ +0.154
• Balanced Net Ionic Reaction:
• 2MnO4
- + 16H+ + 5Sn2+  2Mn2+ + 5Sn4+ + 8H2O
• Eo
cell = Eo
cat – Eo
an = (+1.51 – (+0.154)) = +1.36 V

12. Electrochemistry
Nernst Equation
• Effects of Concentrations on Potentials:
• aOx + ne-  bRed
• E = Eo – (2.3026RT/nF) log([Red]b/[Ox]a
– Where E is the reduction at specific conc.,
– Eo is standard reduction potential, n is number of electrons
involved in the half reaction,
– R is the gas constant (8.3143 V coul deg-1mol-1),
– T is absolute temperature,
– and F is the Faraday constant (96487 coul eq-1).
• At 25oC(298.16K) the value of 2.3026RT/F is 0.05916
• Note: Concentrations should be activities

13. Electrochemistry
• Calculations:
• MnO4
- + 8H+ + 5e-  Mn2+ + 4H2O Eo = +1.51 V
• For [H+] = 1.0M, [MnO4
-] = 0.10M, [Mn2+] = 0.010M
• E = Eo – 0.05916/5 (log ([Mn2+]/[MnO4
-][H+]8)
• E = +1.51 – 0.1183(-1) = +1.63 V vs NHE
• Note: This is more positive than Eo
• Greater tendency to be reduced compared to standard
conditions.

14. Electrochemistry
• Calculations:
• Silver electrode/silver chloride deposit/0.010M NaCl
• AgCl + 1e-  Ago + Cl- E = ?
• Ag+ + 1e-  Ago Eo = +0.799 V
• AgCl  Ag+ + Cl- Ksp= 1.8 x 10-10
• AgCl + e-  Ago + Cl-
• E = Eo - (0.05916/1) Log (1/[Ag+])
• [Ag+] = Ksp/[Cl-] = 1.8 x 10-10/(0.010) = 1.8 x 10-8
• E = +0.799 – (0.05916)(7.74) = +0.341 V

15. Thank You