1. Introduction to Electroanalytical Chemistry
Potentiometry, Voltammetry, Amperometry, Biosensors

2. Applications Study Redox Chemistry Electrochemical analysis
electron transfer reactions, oxidation, reduction, organics & inorganics, proteins
Adsorption of species at interfaces
Electrochemical analysis
Measure the Potential of reaction or process
E = const + k log C (potentiometry)
Measure the Rate of a redox reaction; Current (I) = k C (voltammetry)
• Electrochemical Synthesis
Organics, inorganics, materials, polymers

3. Electrochemical Cells
Galvanic Cells and Electrolytic Cells
• Galvanic Cells – power output; batteries
• Potentiometric cells (I=0) read Chapter 2
– measure potential for analyte to react
current = 0 (reaction is not allowed to occur)
Equil. Voltage is measured (Eeq)
Electrolytic cells, power applied, output meas.
The Nernst Equation
For a reversible process: Ox + ne- → Red
E = Eo – (2.303RT/nF) Log (ared/aox)
a (activity), related directly to concentration

4. Voltammetry is a dynamic method
Related to rate of reaction at an electrode
O + ne = R, Eo in Volts
I = kA[O] k = const. A = area
Faradaic current, caused by electron transfer
Also a non-faradaic current forms
part of background current

5. Electrical Double layer at Electrode
Heterogeneous system: electrode/solution interface
The Electrical Double Layer, e’s in electrode; ions in solution – important for voltammetry:
Compact inner layer: do to d1, E decreases linearly.
Diffuse layer: d1 to d2, E decreases exponentially.

6. Electrolysis: Faradaic and Non-Faradaic Currents
Two types of processes at electrode/solution interface that produce current
Direct transfer of electrons, oxidation or reduction
Faradaic Processes. Chemical reaction rate at electrode proportional to the Faradaic current.
Nonfaradaic current: due to change in double layer when E is changed; not useful for analysis
Mass Transport: continuously brings reactant from the bulk of solution to electrode surface to be oxidized or reduced (Faradaic)
Convection: stirring or flowing solution
Migration: electrostatic attraction of ion to electrode
Diffusion: due to concentration gradient.

7. Typical 3-electrode Voltammetry cell O O e- R R Reference electrode
Counter
electrode
Working electrode O
Reduction at electrode
Causes current flow in
External circuit O e-
Mass transport R R
End of Working electrode
Bulk solution

8. Analytical Electrolytic Cells
Use external potential (voltage) to drive reaction
Applied potential controls electron energy
As Eo gets more negative, need more energetic electrons in order to cause reduction. For a reversible reaction:
 Eapplied is more negative than Eo, reduction will occur
if Eapplied is more positive than Eo, oxidation will occur
O + ne- = R Eo,V electrode reaction

9. Current Flows in electrolytic cells
Due to Oxidation or reduction
Electrons transferred
Measured current (proportional to reaction rate, concentration)
Where does the reaction take place?
On electrode surface, soln. interface
NOT in bulk solution

10. Analytical Applications of Electrolytic Cells
Amperometry
Set Eapplied so that desired reaction occurs
Stir solution
Measure Current
Voltammetry
Quiet or stirred solution
Vary (“scan”) Eapplied
Indicates reaction rate
Reaction at electrode surface produces concentration gradient with bulk solution
Mass transport brings unreacted species to electrode surface

11. Cell for voltammetry, measures I vs. E
wire
potentiostat
insulator
electrode
material
reference N2
inlet
counter
working electrode
Electrochemical cell
Output, I vs. E, quiet solution
Input: E-t waveform
reduction
E, V
time
Figure1

12. Polarization - theoretical
Ideal Non-Polarized Electrode
Ideally Polarized Electrode
reduction
No oxidation or reduction
oxidation

13. Possible STEPS in electron transfer processes
Charge-transfer may be rate limiting
Rate limiting step may be mass transfer
Rate limiting step may be chemical reaction
Adsorption, desorption or crystallization polarization

14. Overvoltage or Overpotential η
η = E – Eeq; can be zero or finite
E < Eeq  η < 0
Amt. of potential in excess of Eeq needed to make
a non-reversible reaction happen, for example
reduction
Eeq

15. NERNST Equation: Fundamental Equation for reversible electron transfer at electrodes
O + ne- = R, Eo in Volts
E.g., Fe e- = Fe2+
If in a cell, I = 0, then E = Eeq
All equilibrium electrochemical reactions obey the
Nernst Equation
Reversibility means that O and R are at equilibrium at all times, not all
Electrochemical reactions are reversible
E = Eo - [RT/nF] ln (aR/aO) ; a = activity
aR = fRCR ao = foCo f = activity coefficient, depends on ionic strength
Then E = Eo - [RT/nF] ln (fR/fO) - [RT/nF] ln (CR/CO)
F = Faraday const., 96,500 coul/e, R = gas const.
T = absolute temperature

16. Ionic strength I = Σ zi2mi,
Z = charge on ion, m = concentration of ion
Debye Huckel theory says log fR = 0.5 zi2 I1/2
So fR/fOwill be constant at constant I.
And so, below are more usable forms of Nernst Eqn.
E = Eo - const. - [RT/nF] ln (CR/CO) Or
E = Eo’ - [RT/nF] ln (CR/CO); Eo’ = formal potential of O/R
At 25 oC using base 10 logs
E = Eo’ - [0.0592/n] log (CR/CO); equil. systems