Introduction to Electrochemistry

1

• Introduction to Electrochemistry
• Formal potentials are the potentials of
  half-cells relative to the SHE measured under
  conditions that the ratio of the analytical
  concentrations of reactants and products are
  exactly unity and the concentrations of other
  species in the system are carefully specified.
• Symbol E0
• Example for Fe(CN)63- e- Fe(CN)64-
  E0 0.36 V
• in 1 M HClO4 or 1 M H2SO4, E vs. SHE 0.72 V
• H reacts with the Fe(CN)63- and Fe(CN)64- to
  form protonated forms
• The Hn Fe(CN)nn-3 is a stronger acid than
  HnFe(CN)6n-4 and the ratio Fe(CN)64-/Fe(CN)6
  3- will be less than 1 even though the ratio of
  analytical concentrations is 1
• Use formal potentials when appropriate and be
  careful about other species in solution

2

• Titration Curves for Redox Systems
• The system
• Add the titrant to the analytical solution which
  is part of a galvanic cell and allow the
  chemistry to come to equilibrium
• Generally the titrant is an oxidizing agent
• Often one of the electrodes is made of an inert
  material such as Pt and the other electrode is
  a secondary standard electrode such as Ag/AgCl or
  Hg/Hg2Cl2 separated from the analytical solution
  by a salt bridge
• Measure the potential in volts or millivolts of
  the inert electrode vs. the reference electrode
  after equilibrium is established
• At equilibrium in the analytical solution if and
  Eref 0.000 V,
• Ecathode Eanode Esystem
• In calculation of the titration curve one may
  calculate Esystem from either the anodic
  reaction or cathodic reaction
• Use the half-equation that is most convenient
• Example consider the titration of Fe2 with Ce4
• Fe2 Ce4 Fe3 Ce2
• Prior to the equivalence point, Fe2 and Fe3
  are easy to calculate but Ce4 requires an
  equilibrium constant calculation
• After the equivalence point, Ce4 and Ce3
  are easy to calculate but Fe2 requires an
  equilibrium constant calculation
• At the equivalence point, the Nernst equation
  does not provide enough information to
  calculate the potential

3

• Titration Curves for Redox Systems
• The system
• At the equivalence point, but . . .
• Fe3 Ce3 and
• Fe4 Ce4 and are very small and require
  Keq to calculate
• However,
• Ce4 e- Ce3
• Fe3 e- Fe2

4

• Titration Curves for Redox Systems
• Another Example Cr2O72- 6Fe2 14H
  2Cr3 6Fe3 H2O Cr2O72- 14H 6e-
  2Cr3 7H20
• Fe3 1e- Fe2

5

• Titration Curves for Redox Systems
• Derive the titration curve for the titration of
  50.00 mL of 0.0500 M Fe2 with 0.01667 M
  Cr2O72-, both 1.0 M in H2SO4
• Equivalence point occurs at 25.00 mL added
  Cr2O72-
• Initial point
• No Cr2O72- has been added, therefore no Cr3 or
  Fe3 has been produced
• Cannot calculate the potential of the system

6

• Titration Curves for Redox Systems
• Derive the titration curve for the titration of
  50.00 mL of 0.0500 M Fe2with 0.01667 M
  Cr2O72-, both 1.0 M in H2SO4
• Add 5.00 mL Cr2O72- make use of the Fe2/Fe3
  couple
• At the equivalence point use earlier result

7

• Titration Curves for Redox Systems
• Derive the titration curve for the titration of
  50.00 mL of 0.0500 M Fe2with 0.01667 M
  Cr2O72-, both 1.0 M in H2SO4
• Beyond the equivalence point 25.10 mL added
  Cr2O72-
• Make use of the Cr3/Cr2O72- couple
• At 30.00 mL