Introduction to Electroanalytical Chemistry Potentiometry, Voltammetry, Amperometry, Biosensors

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

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

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 (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.

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

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? 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

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

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

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) 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

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.

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.