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

2. Uses To 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); C = concentration

3. Electrochemical Cells
Galvanic Cells and Electrolytic Cells
• Galvanic Cells – power output; batteries (car)
• Potentiometric cells (I=0)
– 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), proportional 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 (ET), oxidation, reduction
Faradaic Processes. ET 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
Conductive Au, Pt, C
May have redox film coating 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?
At electrode surface-soln. interface (soln or film)
NOT in bulk solution
Can put a catalytic or other kind of film on electrode
Counter
electrode
Refer-
ence O
potentiostat e-
solution R
End of Working electrode

10. Analytical Applications of Electrolytic Cells
Amperometry
Set Eapplied so that desired reaction occurs
Stir or flow solution
Measure Current
Voltammetry
Quiet or stirred solution; or redox film on electrode
Vary (“scan”) Eapplied
Indicates reaction rate
Reaction at electrode surface
Mass transport brings reactive species to electrode surface
For film, mass transport within and outside of film

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. 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

13. 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

14. Linear Potential Sweep Voltammetry
Ox + ne– ⇄ Red
E(t) = Ei – vt
At Ei nonfaradaic current
v ranges from 10 mV s-1 to ~1000 V s-1 for conventional electrodes; up to 106 V s-1 for UMEs.
Supporting electrolyte present in excess to minimize migration.
Solution unstirred to min. convection.
O2 is removed by bubbling N2.
Stationary electrodes: HDME, graphite, Pt, etc.

15. Cyclic Voltammetry Ox + ne–  Red Ox – ne–  Red
For reversible electron transfer:
DE = |Ep,a – Ep,c| = 59 mV/n
|Ep – Ep/2| = 57 mV/n
ip,a/ip,c = 1
Reversal technique analogous to double-potential step methods.
Experiment time scale: 103 s to 10-5 s.