Introduction to Electroanalytical Chemistry

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

2. 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

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 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,Velectrode 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 • Measure Current • 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 electrode material insulator 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., 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

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