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 SynthesisOrganics, 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 = areaFaradaic 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

Counterelectrode

Reference electrode

Working electrode

End of Working electrode

O

R

O

R

e-

Bulk solution

Mass transport

Reduction at electrodeCauses current flow inExternal 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

E, V

time

Input: E-t waveform

potentiostat

Electrochemical cell

counter

working electrode

N2

inlet

reference

insulator electrodematerial

Cell for voltammetry, measures I vs. E

wire

Output, I vs. E, quiet solution

reduction

Polarization - theoretical

Ideally Polarized ElectrodeIdeal Non-Polarized Electrode

No oxidation or reduction

reduction

oxidation

Possible STEPS in electron transfer processes

Rate limiting step may be mass transfer

Rate limiting step may be chemical reaction

Adsorption, desorption or crystallization polarization

Charge-transfer may be rate limiting

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) ; 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

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)

OrE = 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