Section 10
Electrochemical Cells and Electrode Potentials
ElectrochemistryOxidation/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)
ElectrochemistryOxidation/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
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).
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)
ElectrochemistryStandard 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+(1.0) + 2e- H2(g 1atm) 0.0000 volts
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).
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)
ElectrochemistryReduction 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.
ElectrochemistryOxidation/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
ElectrochemistryOxidation/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
ElectrochemistryVoltaic Cell
• The spontaneous (Voltaic) cell reaction is the one that gives a positive cell voltage when subtracting one half-reaction from the other.
• Eocell = Eo
right – Eo
left = Eocathode – 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
ElectrochemistryOxidation/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
• Eocell = Eo
cat – Eoan = (+1.51 – (+0.154)) = +1.36 V
ElectrochemistryNernst 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
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.
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