Introduction to
Electroanalytical Chemistry
Potentiometry, Voltammetry,
Amperometry, Biosensors
Applications
• Study Redox Chemistry
– 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
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
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
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.
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.
Typical 3-electrode
Voltammetry cell
Reference electrode
Counter
electrode
Working electrode
O
e-
O
Mass transport
R
End of Working electrode
R
Bulk solution
Reduction at electrode
Causes current flow in
External circuit
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
• 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
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
– Measure Current
• Indicates reaction rate
• Reaction at electrode surface produces concentration
gradient with bulk solution
• Mass transport brings unreacted species to electrode surface
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
E, V
time
reduction
Polarization - theoretical
Ideal Non-Polarized Electrode
Ideally Polarized Electrode
reduction
No oxidation or reduction
oxidation
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
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
NERNST Equation: Fundamental Equation
for reversible electron transfer at electrodes
O + ne- = R,
Eo in Volts
•E.g., Fe3+ + 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)
aR = fRCR
ao = foCo
;
a = activity
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
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