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Introduction to Electroanalytical Chemistry

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Presentation on theme: "Introduction to Electroanalytical Chemistry"— Presentation transcript:

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

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