Analytical electrochemistry wang

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Analytical electrochemistry wang

  1. 1. E-book biếu tặnghọc sinh, sinh viên Việt NamHội Khuyến Học Thanh Niên Việt NamDownload tại www.ttxva.org
  2. 2. ANALYTICALELECTROCHEMISTRYSECOND EDITIONAnalytical Electrochemistry, Second Edition. Joseph WangCopyright # 2000 Wiley-VCHISBNs: 0-471-28272-3 (Hardback); 0-471-22823-0 (Electronic)
  3. 3. ANALYTICALELECTROCHEMISTRYSecond EditionJOSEPH WANGA JOHN WILEY & SONS, INC., PUBLICATIONNew York = Chichester = Weinheim = Brisbane = Singapore = Toronto
  4. 4. Copyright # 2001 by Wiley-VCH. All rights reserved.No part of this publication may be reproduced, stored in a retrieval system or transmitted in any formor by any means, electronic or mechanical, including uploading, downloading, printing, decompiling,recording or otherwise, except as permitted under Sections 107 or 108 of the 1976 United StatesCopyright Act, without the prior written permission of the Publisher. Requests to the Publisher forpermission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 605 ThirdAvenue, New York, NY 10158-0012, (212) 850-6011, fax (212) 850-6008,E-Mail: PERMREQ @ WILEY.COM.This publication is designed to provide accurate and authoritative information in regard to the subjectmatter covered. It is sold with the understanding that the publisher is not engaged in rendering professionalservices. If professional advice or other expert assistance is required, the services of a competentprofessional person should be sought.ISBN 0-471-22823-0.This title is also available in print as ISBN 0-471-28272-3.For more information about Wiley products, visit our web site at www.Wiley.com.
  5. 5. Dedicated to the memory of my parents, Elka and Moshe Wang
  6. 6. CONTENTSPREFACE xiABBREVIATIONS AND SYMBOLS xiii1 FUNDAMENTAL CONCEPTS 11-1 Why Electroanalysis? = 11-2 Faradaic Processes = 31-2.1 Mass Transport-Controlled Reactions = 41-2.1.1 Potential Step Experiment = 71-2.1.2 Potential-Sweep Experiments = 81-2.2 Reactions Controlled by the Rate of Electron Transfer = 111-2.2.1 Activated Complex Theory = 161-3 The Electrical Double Layer = 181-4 Electrocapillary Effect = 22Supplementary Reading = 25References = 26Questions = 262 STUDY OF ELECTRODE REACTIONS 282-1 Cyclic Voltammetry = 282-1.1 Data Interpretation = 30vii
  7. 7. 2-1.1.1 Reversible Systems = 312-1.1.2 Irreversible and Quasi-Reversible Systems = 322-1.2 Study of Reaction Mechanisms = 332-1.3 Study of Adsorption Processes = 362-1.4 Quantitative Applications = 392-2 Spectroelectrochemistry = 402-2.1 Experimental Arrangement = 402-2.2 Principles and Applications = 412-2.3 Other Spectroelectrochemical and Spectroscopic Techniques = 442-3 Scanning Probe Microscopy = 462-3.1 Scanning Tunneling Microscopy = 462-3.2 Atomic Force Microscopy = 472-3.3 Scanning Electrochemical Microscopy = 492-4 Electrochemical Quartz Crystal Microbalance = 52References = 54Examples = 56Questions = 593 CONTROLLED-POTENTIAL TECHNIQUES 603-1 Chronoamperometry = 603-2 Polarography = 623-3 Pulse Voltammetry = 673-3.1 Normal-Pulse Voltammetry = 673-3.2 Differential-Pulse Voltammetry = 683-3.3 Square-Wave Voltammetry = 723-3.4 Staircase Voltammetry = 743-4 AC Voltammetry = 743-5 Stripping Analysis = 753-5.1 Anodic Stripping Voltammetry = 763-5.2 Potentiometric Stripping Analysis = 793-5.3 Adsorptive Stripping Voltammetry and Potentiometry = 803-5.4 Cathodic Stripping Voltammetry = 823-5.5 Applications = 843-6 Flow Analysis = 843-6.1 Principles = 873-6.2 Cell Design = 883-6.3 Mass Transport and Current Response = 903-6.4 Detection Modes = 92viii CONTENTS
  8. 8. References = 94Examples = 96Questions = 984 PRACTICAL CONSIDERATIONS 1004-1 Electrochemical Cells = 1004-2 Solvents and Supporting Electrolytes = 1024-3 Oxygen Removal = 1034-4 Instrumentation = 1044-5 Working Electrodes = 1074-5.1 Mercury Electrodes = 1084-5.2 Solid Electrodes = 1104-5.2.1 Rotating Disk and Ring-Disk Electrodes = 1114-5.2.2 Carbon Electrodes = 1134-5.2.2.1 Glassy-Carbon Electrodes = 1144-5.2.2.2 Carbon-Paste Electrodes = 1154-5.2.2.3 Carbon-Fiber Electrodes = 1154-5.2.3 Metal Electrodes = 1174-5.3 Chemically Modi®ed Electrodes = 1184-5.3.1 Self-Assembled Monolayers = 1184-5.3.2 Sol-Gel Encapsulation of Reactive Species = 1204-5.3.3 Electrocatalytic Modi®ed Electrodes = 1214-5.3.4 Preconcentrating Electrodes = 1214-5.3.5 Permselective Coatings = 1234-5.3.6 Conducting Polymers = 1244-5.4 Microelectrodes = 1284-5.4.1 Diffusion at Microelectrodes = 1294-5.4.2 Con®gurations of Microelectrodes = 1304-5.4.3 Composite Electrodes = 131References = 135Examples = 138Questions = 1385 POTENTIOMETRY 1405-1 Principles of Potentiometric Measurements = 1405-2 Ion-Se1ective Electrodes = 1475-2.1 Glass Electrodes = 1475-2.1.1 pH Electrodes = 147CONTENTS ix
  9. 9. 5-2.1.2 Glass Electrodes for Other Cations = 1515-2.2 Liquid-Membrane Electrodes = 1525-2.2.1 Ion-Exchanger Electrodes = 1525-2.2.2 Neutral Carrier Electrodes = 1545-2.3 Solid-State Electrodes = 1565-2.4 Coated-Wire Electrodes = 1605-3 On-Line and In-Vivo Potentiometric Measurements = 161References = 165Examples = 167Questions = 1696 ELECTROCHEMICAL SENSORS 1716-1 Electrochemical Biosensors = 1716-1.1 Enzyme-Based Electrodes = 1726-1.1.1 Impractical and Theoretical Considerations = 1726-1.1.2 Enzyme Electrodes of Analytical Signi®cance = 1766-1.1.2.1 Glucose Sensors = 1766-1.1.2.2 Ethanol Electrodes = 1786-1.1.2.3 Urea Electrodes = 1816-1.1.2.4 Toxin (Enzyme Inhibition) Biosensors = 1816-1.1.3 Tissue and Bacteria Electrodes = 1826-1.2 Af®nity Biosensors = 1836-1.2.1 Immunosensors = 1836-1.2.2 DNA Hybridization Biosensors = 1856-1.2.3 Receptor-Based Sensors = 1866-2 Gas Sensors = 1886-2.1 Carbon Dioxide Sensors = 1896-2.2 Oxygen Electrodes = 1906-3 Solid-State Devices = 1916-3.1 Microfabrication of Solid-State Sensor Assemblies = 1936-3.2 Microfabrication Techniques = 1936-4 Sensor Arrays = 197References = 198Examples = 201Questions = 202INDEX 205x CONTENTS
  10. 10. PREFACEThe goal of this textbook is to cover the full scope of modern electroanalyticaltechniques and devices. The main emphasis is on electroanalysis, rather thanphysical electrochemistry. The objective is to provide a sound understanding ofthe fundamentals of electrode reactions and of the principles of electrochemicalmethods, and to demonstrate their potential for solving real-life analytical problems.Given the impressive progress in electroanalytical chemistry, and its growing impacton analytical chemistry, this work offers also an up-to-date, easy-to-read presentationof recent advances including new methodologies, sensors, detectors, and micro-systems. The book is suitable for a graduate-level course in electroanalyticalchemistry or as a supplement to a high-level undergraduate course in instrumentalanalysis. It should also be very useful to those considering the use of electroanalysisin their laboratories.The material is presented in six roughly equal chapters. The ®rst chapter isdevoted to fundamental aspects of electrode reactions and the structure of theinterfacial region. Chapter 2 discusses the study of electrode reactions and high-resolution surface characterization. Chapter 3 gives an overview of ®nite-currentcontrolled-potential techniques. Chapter 4 describes the electrochemical instrumen-tation and electrode materials (including new modi®ed and microelectrodes).Chapter 5 deals with the principles of potentiometric measurements and variousclasses of ion-selective electrodes, while Chapter 6 is devoted to the growing ®eld ofchemical sensors (including modern biosensors, gas sensors, solid-state devices, andsensor arrays). I have tried to provide numerous references to review literature at theend of each chapter. By discussing the very latest advances, it is hoped to bridge thecommon gap between recent research literature and standard textbooks.xi
  11. 11. This second edition of Analytical Electrochemistry is extensively revised andupdated, and re¯ects the rapid growth of electroanalytical chemistry during the1990s. It contains a number of new topics, including self-assembled monolayers,DNA biosensors, sol-gel surface modi®cation, detection for capillary electrophor-esis, single molecule detection, and micromachined analyzers (``Lab-on-a-Chip).Other topics such as the fundamentals of faradaic processes, principles of potentio-metric measurements, spectroelectrochemistry, modi®ed and microelectrodes, scan-ning electron microscopy, electrical communication between redox enzymes andelectrodes, and enzyme and immunoelectrodes, have been greatly expanded. Theentire text has been updated to cover the very latest (as of 1999) developments inelectroanalytical chemistry. Numerous new worked examples and end-of-chapterquestions have been added to this edition. The organization of the book has beenchanged somewhat, by moving the study of electrode reactions forward to Chapter 2.In the ®ve years since the ®rst edition I received numerous suggestions, many ofwhich have been incorporated in the second edition.Finally, I wish to thank my wife, Ruth, and my daughter, Sharon, for their loveand patience; the editorial and production staff of Wiley Inc. for their help andsupport; and the numerous electrochemists across the globe who led to the advancesreported in this textbook. Thank you all!Joseph WangLas Cruces, New Mexico, U.S.A.xii PREFACE
  12. 12. ABBREVIATIONS AND SYMBOLSa ActivityA AbsorbanceA Area of electrodeAb AntibodyAdSV Adsorptive stripping voltammetryAES Auger electron spectroscopyAFM Atomic force microscopyAg AntigenASV Anodic stripping voltammetryB Adsorption coef®cientC ConcentrationCdl Differential capacitanceCHg Concentration in amalgamCSV Cathodic stripping voltammetryCWE Coated-wire electrodeCME Chemically modi®ed electrodeCV Cyclic voltammetryCZE Capillary zone electrophoresisD Diffusion coef®cientDC Direct currentDNA Deoxyribonucleic acidxiii
  13. 13. DME Dropping mercury electrodeDPV Differential pulse voltammetryE Potential (V)DE Pulse amplitude; step heightEB Binding energy (in XPS)Eeq Equilibrium potentialEStandard electrode potentialE1=2 Half wave potentialEp Peak potentialEpzc Potential of zero chargeEC Electrode process involving an electrochemical reaction followed by achemical stepECL ElectrochemiluminescenceEQCM Electrochemical quartz crystal microbalanceESCA Electron spectroscopy for chemical analysisEXAFS X-ray adsorption ®ne structureF Faraday constantFET Field-effect transistorFIA Flow injection analysisf Activity coef®cient; frequencyDf Frequency change (in EQCM)DG Free energyDGz Free energy of activationHMDE Hanging mercury drop electrodei Electric currentic Charging currentil Limiting currentit Tunneling currentDi Current differenceIHP Inner Helmholtz planeIRS Internal re¯ectance spectroscopyISE Ion-selective electrodeISFET Ion-selective ®eld-effect transistorJ Fluxkpotij Potentiometric selectivity coef®cientkStandard rate constantKm Michaelis±Menten constant; mass transport coef®cientxiv ABBREVIATIONS AND SYMBOLS
  14. 14. l Film thicknessLCEC Liquid chromatography=electrochemistryLEED Low-energy electron diffractionm Mercury ¯ow rate (in polarography); electron mass (in STM)Dm Mass charge (in EQCM)M MediatorMFE Mercury ®lm electrodeN Collection ef®ciencyNADH Dihydronicotinamide adenine dinucleotiden Number of electrons transferredNP Normal pulseO The oxidized speciesOHP Outer Helmholtz planeOTE Optically transparent electrodePAD Pulsed amperometric detectionPSA Potentiometric stripping analysisq ChargeQCM Quartz crystal microbalancer distance; radiusR Resistance; gas constantRDE Rotating disk electrodeRe Reynolds numberRRDE Rotating ring disk electrodeRVC Reticulated vitreous carbonS Barrier width (in STM)SAM Self-assembled monolayersSc Schmidt numberSECM Scanning electrochemical microscopySERS Surface enhanced Raman scatteringSTM Scanning tunneling microscopySWV Square-wave voltammetryT Temperaturet Timetd Deposition timetm Transition time (in PSA)U Flow rate, stirring rateABBREVIATIONS AND SYMBOLS xv
  15. 15. v Potential scan rateVHg Volume of mercury electrodeVmax Maximum rateW1=2 Peak width (at half height)WE Working electrodeWJD Wall jet detectorXPS X-ray photoelectron spectroscopya Transfer coef®cientG Surface coveragee Dielectric constant; molar absorptivityg Surface tensiond Thickness of the diffusion layerdH Thickness of the hydrodynamic boundary layerZ Overvoltagem Ionic strengthn Kinematic viscosityo Angular velocityxvi ABBREVIATIONS AND SYMBOLS
  16. 16. CHAPTER 1FUNDAMENTAL CONCEPTS1-1 WHY ELECTROANALYSIS?Electroanalytical techniques are concerned with the interplay between electricity andchemistry, namely the measurements of electrical quantities, such as current,potential, or charge, and their relationship to chemical parameters. Such use ofelectrical measurements for analytical purposes has found a vast range of applica-tions, including environmental monitoring, industrial quality control, and biomedicalanalysis. Advances in the 1980s and 1990sÐincluding the development of ultra-microelectrodes, the design of tailored interfaces and molecular monolayers, thecoupling of biological components and electrochemical transducers, the synthesis ofionophores and receptors containing cavities of molecular size, the development ofultratrace voltammetric techniques or of high-resolution scanning probe microsco-pies, and the microfabrication of molecular devices or ef®cient ¯ow detectorsÐhaveled to a substantial increase in the popularity of electroanalysis, and to its expansioninto new phases and environments. Indeed, electrochemical probes are receiving amajor share of the attention in the development of chemical sensors.In contrast to many chemical measurements that involve homogeneous bulksolutions, electrochemical processes take place at the electrode±solution interface.The distinction between various electroanalytical techniques re¯ects the type ofelectrical signal used for the quantitation. The two principal types of electroanaly-tical measurements are potentiometric and potentiostatic. Both types require at leasttwo electrodes (conductors) and a contacting sample (electrolyte) solution, whichconstitute the electrochemical cell. The electrode surface is thus a junction betweenan ionic conductor and an electronic conductor. One of the two electrodes responds1Analytical Electrochemistry, Second Edition. Joseph WangCopyright # 2000 Wiley-VCHISBNs: 0-471-28272-3 (Hardback); 0-471-22823-0 (Electronic)
  17. 17. to the target analyte(s) and is thus termed the indicator (or working) electrode. Thesecond one, termed the reference electrode, is of constant potential (that is,independent of the properties of the solution). Electrochemical cells can be classi®edas electrolytic (when they consume electricity from an external source) or galvanic(if they are used to produce electrical energy).Potentiometry (discussed in Chapter 5), which is of great practical importance, isa static (zero current) technique in which the information about the samplecomposition is obtained from measurement of the potential established across amembrane. Different types of membrane materials, possessing different ion-recogni-tion processes, have been developed to impart high selectivity. The resultingpotentiometric probes have thus been widely used for several decades for directmonitoring of ionic species such as protons or calcium, ¯uoride, and potassium ionsin complex samples.Controlled-potential (potentiostatic) techniques deal with the study of charge-transfer processes at the electrode±solution interface, and are based on dynamic (nozero current) situations. Here, the electrode potential is being used to derive anelectron-transfer reaction and the resultant current is measured. The role of thepotential is analogous to that of the wavelength in optical measurements. Such acontrollable parameter can be viewed as ``electron pressure, which forces thechemical species to gain or lose an electron (reduction or oxidation, respectively).TABLE 1-1 Properties of Controlled-Potential TechniquesTechniqueaWorkingElectrodebDetectionLimit (M)Speed(time percycle) (min)ResponseShapeDC polarography DME 10À 53 WaveNP polarography DME 5  10À 73 WaveDP polarography DME 10À 83 PeakDP voltammetry Solid 5  10À 73 PeakSW polarography DME 10À 80.1 PeakAC polarography DME 5  10À 71 PeakChronoamperometry Stationary 10À 50.1 TransientCyclic voltammetry Stationary 10À 50.1±2 PeakStripping voltammetry HMDE, MFE 10À 103±6 PeakAdsorptive strippingvoltammetryHMDE 10À 102±5 PeakAdsorptive strippingvoltammetrySolid 10À 94±5 PeakAdsorptive-catalyticstripping voltammetryHMDE 10À 122±5 PeakaDC ˆ direct current; NP ˆ normal pulse; DP ˆ differential pulse; SW ˆ square wave; AC ˆ alternatingcurrent.bDME ˆ dropping mercury electrode; HMDE ˆ hanging mercury drop electrode; MFE ˆ mercury ®lmelectrode.2 FUNDAMENTAL CONCEPTS
  18. 18. Accordingly, the resulting current re¯ects the rate at which electrons move across theelectrode±solution interface. Potentiostatic techniques can thus measure any chemi-cal species that is electroactive, in other words, that can be made to reduce oroxidize. Knowledge of the reactivity of functional group in a given compound can beused to predict its electroactivity. Nonelectroactive compounds may also be detectedin connection with indirect or derivatization procedures.The advantages of controlled-potential techniques include high sensitivity,selectivity towards electroactive species, a wide linear range, portable and low-cost instrumentation, speciation capability, and a wide range of electrodes that allowassays of unusual environments. Several properties of these techniques are summar-ized in Table 1-1. Extremely low (nanomolar) detection limits can be achieved withvery small sample volumes (5±20 ml), thus allowing the determination of analyteamounts of 10À13to 10À15mol on a routine basis. Improved selectivity may beachieved via the coupling of controlled-potential schemes with chromatographic oroptical procedures.This chapter attempts to give an overview of electrode processes, together withdiscussion of electron transfer kinetics, mass transport, and the electrode±solutioninterface.1-2 FARADAIC PROCESSESThe objective of controlled-potential electroanalytical experiments is to obtain acurrent response that is related to the concentration of the target analyte. Thisobjective is accomplished by monitoring the transfer of electron(s) during the redoxprocess of the analyte:O ‡ neÀ„ R …1-1†where O and R are the oxidized and reduced forms, respectively, of the redox couple.Such a reaction will occur in a potential region that makes the electron transferthermodynamically or kinetically favorable. For systems controlled by the laws ofthermodynamics, the potential of the electrode can be used to establish theconcentration of the electroactive species at the surface [CO…0; t† and CR…0; t†]according to the Nernst equation:E ˆ E‡2:3RTnFlogCO…0; t†CR…0; t†…1-2†where Eis the standard potential for the redox reaction, R is the universal gasconstant (8.314 J KÀ1molÀ1), T is the Kelvin temperature, n is the number ofelectrons transferred in the reaction, and F is the Faraday constant (96,487coulombs). On the negative side of E, the oxidized form thus tends to be reduced,and the forward reaction (i.e., reduction) is more favorable. The current resultingfrom a change in oxidation state of the electroactive species is termed the faradaic1-2 FARADAIC PROCESSES 3
  19. 19. current because it obeys Faradays law (i.e. the reaction of 1mole of substanceinvolves a change of n  96,487 coulombs). The faradaic current is a direct measureof the rate of the redox reaction. The resulting current±potential plot, known as thevoltammogram, is a display of current signal (vertical axis) versus the excitationpotential (horizontal axis). The exact shape and magnitude of the voltammetricresponse is governed by the processes involved in the electrode reaction. Thetotal current is the summation of the faradaic currents for the sample and blanksolutions, as well as the nonfaradaic charging background current (discussed inSection 1-3).The pathway of the electrode reaction can be quite complicated, and takes placein a sequence that involves several steps. The rate of such reactions is determined bythe slowest step in the sequence. Simple reactions involve only mass transport of theelectroactive species to the electrode surface, the electron transfer across theinterface, and the transport of the product back to the bulk solution. More complexreactions include additional chemical and surface processes that precede or followthe actual electron transfer. The net rate of the reaction, and hence the measuredcurrent, may be limited either by mass transport of the reactant or by the rate ofelectron transfer. The more sluggish process will be the rate-determining step.Whether a given reaction is controlled by the mass transport or electron transfer isusually determined by the type of compound being measured and by variousexperimental conditions (electrode material, media, operating potential, mode ofmass transport, time scale, etc.). For a given system, the rate-determining step maythus depend on the potential range under investigation. When the overall reaction iscontrolled solely by the rate at which the electroactive species reach the surface (i.e.,a facile electron transfer), the current is said to be mass transport-limited. Suchreactions are called nernstian or reversible, because they obey thermodynamicrelationships. Several important techniques (discussed in Chapter 4) rely on suchmass transport-limited conditions.1-2.1 Mass Transport-Controlled ReactionsMass transport occurs by three different modes: DiffusionÐthe spontaneous movement under the in¯uence of concentrationgradient (i.e., from regions of high concentration to regions of lower concen-tration), aimed at minimizing concentration differences. ConvectionÐtransport to the electrode by a gross physical movement; such¯uid ¯ow occurs with stirring or ¯ow of the solution and with rotation orvibration of the electrode (i.e., forced convection) or due to density gradients(i.e., natural convection); MigrationÐmovement of charged particles along an electrical ®eld (i.e., thecharge is carried through the solution by ions according to their transferencenumber).These modes of mass transport are illustrated in Figure 1-1.4 FUNDAMENTAL CONCEPTS
  20. 20. The ¯ux (J) is a common measure of the rate of mass transport at a ®xed point. Itis de®ned as the number of molecules penetrating a unit area of an imaginary planein a unit of time, and has the units of mol cmÀ2sÀ1. The ¯ux to the electrode isdescribed mathematically by a differential equation, known as the Nernst±Planckequation, given here for one dimension:J…x; t† ˆ ÀD@C…x; t†@xÀzFDCRT@f…x; t†@x‡ C…x; t†V…x; t† …1-3†where D is the diffusion coef®cient (cmÀ2sÀ1), @C…x; t†=@x is the concentrationgradient (at distance x and time t), @f…x; t†=@x is the potential gradient, z and C arethe charge and concentration, respectively, of the electroactive species, and V…x; t† isthe hydrodynamic velocity (in the x direction). In aqueous media, D usually rangesbetween 10À5and 10À 6cmÀ2sÀ1. The current (i) is directly proportional to the ¯ux:i ˆ ÀnFAJ …1-4†FIGURE 1-1 The three modes of mass transport. (Reproduced with permission fromreference 1.)1-2 FARADAIC PROCESSES 5
  21. 21. As indicated by equation (1-3), the situation is quite complex when the threemodes of mass transport occur simultaneously. This complication makes it dif®cultto relate the current to the analyte concentration. The situation can be greatlysimpli®ed by suppressing the electromigration or convection, through the addition ofexcess inert salt or use of a quiescent solution, respectively. Under these conditions,the movement of the electroactive species is limited by diffusion. The reactionoccurring at the surface of the electrode generates a concentration gradient adjacentto the surface, which in turn gives rise to a diffusional ¯ux. Equations governingdiffusion processes are thus relevant to many electroanalytical procedures.According to Ficks ®rst law, the rate of diffusion (i.e., the ¯ux) is directlyproportional to the slope of the concentration gradient:J…x; t† ˆ ÀD@C…x; t†@x…1-5†Combination of equations (1-4) and (1-5) yields a general expression for the currentresponse:i ˆ nFAD@C…x; t†@x…1-6†Hence, the current (at any time) is proportional to the concentration gradient of theelectroactive species. As indicated by the above equations, the diffusional ¯ux istime dependent. Such dependence is described by Ficks second law (for lineardiffusion):@C…x; t†@tˆ D@2C…x; t†@x2…1-7†This equation re¯ects the rate of change with time of the concentration betweenparallel planes at points x and (x ‡ dx) (which is equal to the difference in ¯ux at thetwo planes). Ficks second law is valid for the conditions assumed, namely planesparallel to one another and perpendicular to the direction of diffusion, i.e., conditionsof linear diffusion. In contrast, for the case of diffusion toward a spherical electrode(where the lines of ¯ux are not parallel but are perpendicular to segments of thesphere), Ficks second law has the form@C@tˆ D@2C@r2‡2r@C@r …1-8†where r is the distance from the center of the electrode. Overall, Ficks laws describethe ¯ux and the concentration of the electroactive species as functions of positionand time. The solution of these partial differential equations usually requires theapplication of a (Laplace transformation) mathematical method. The Laplacetransformation is of great value for such applications, as it enables the conversion6 FUNDAMENTAL CONCEPTS
  22. 22. of the problem into a domain where a simpler mathematical manipulation ispossible. Details of use of the Laplace transformation are beyond the scope ofthis text, and can be found in reference 2. The establishment of proper initial andboundary conditions (which depend upon the speci®c experiment) is also essentialfor this treatment. The current±concentration±time relationships that result fromsuch treatment will be described below for several relevant experiments.1-2.1.1 Potential-Step Experiment Let us see, for example, what happens ina potential-step experiment involving the reduction of O to R, a potential valuecorresponding to complete reduction of O, a quiescent solution, and a planarelectrode imbedded in a planar insulator. (Only O is initially present in solution.)The current±time relationship during such an experiment can be understood from theresulting concentration±time pro®les. Since the surface concentration of O is zero atthe new potential, a concentration gradient is established near the surface. The regionwithin which the solution is depleted of O is known as the diffusion layer, and itsthickness is given by d. The concentration gradient is steep at ®rst, and the diffusionlayer is thin (see Figure 1-2, for t1). As time goes by, the diffusion layer expands (tod2 and d3 at t2 and t3), and hence the concentration gradient decreases.Initial and boundary conditions in such experiment include CO…x; 0† ˆ CO…b†(i.e., at t ˆ 0, the concentration is uniform throughout the system and equal to thebulk concentration; CO…b†), CO…0; t† ˆ 0 for t 0 (i.e., at later times the surfaceconcentration is zero); and CO…x; 0† 3 CO…b† as x 3 I (i.e., the concentrationincreases as the distance from the electrode increases). Solution of Ficks laws (forFIGURE 1-2 Concentration pro®les for different times t after the start of a potential-stepexperiment.1-2 FARADAIC PROCESSES 7
  23. 23. linear diffusion, i.e., a planar electrode) for these conditions results in a time-dependent concentration pro®le,CO…x; t† ˆ CO…b†f1 À erf‰X=…4DOt†1=2Šg …1-9†whose derivative with respect to x gives the concentration gradient at the surface,@C@xˆCO…b†…pDOt†1=2…1-10†when substituted into equation (1-6) leads to the well-known Cottrell equation:i…t† ˆnFADOCO…b†…pDOt†1=2…1-11†That is, the current decreases in proportion to the square root of time, with …pDOt†1=2corresponding to the diffusion layer thickness.Solving equation (1-8) (using Laplace transform techniques) yields the timeevolution of the current of a spherical electrode:i…t† ˆnFADOCO…b†…pDOt†1=2‡nFADOCOr…1-12†The current response of a spherical electrode following a potential step thus containstime-dependent and time-independent terms, re¯ecting the planar and sphericaldiffusional ®elds, respectively (Figure 1-3), becoming time independent at long timescales. As expected from equation (1-12), the change from one regime to another isstrongly dependent upon the radius of the electrode. The unique mass transportproperties of ultramicroelectrodes (discussed in Section 4-5.4) are attributed to theshrinkage of the electrode radius.1-2.1.2 Potential-Sweep Experiments Let us move to a voltammetricexperiment involving a linear potential scan, the reduction of O to R and a quiescentsolution. The slope of the concentration gradient is given by (CO…b; t† À CO…0; t††=dwhere CO…b; t) and CO…0; t† are the bulk and surface concentrations of O. Thechange in the slope, and hence the resulting current, is due to changes of bothCO…0; t† and d. First, as the potential is scanned negatively, and approaches thestandard potential (E) of the couple, the surface concentration rapidly decreasesin accordance to the Nernst equation (equation 1-2). For example, at a potentialequal to Ethe concentration ratio is unity (CO…0; t†=CR…0; t† ˆ 1). For a potential59 mV more negative than E, CR…0; t† is present at 10-fold excess(CO…0; t†=CR…0; t† ˆ 1=10; n ˆ 1). The decrease in CO…0; t† is coupled with anincrease in the diffusion layer thickness, which dominates the change in the slopeafter CO…0; t† approaches zero. The net result is a peak-shaped voltammogram. Such8 FUNDAMENTAL CONCEPTS
  24. 24. current±potential curves and the corresponding concentration±distance pro®les (forselected potentials along the scan) are shown in Figure 1-4. As will be discussed inSection 4-5.4, shrinking the electrode dimension to the micrometer domain results ina sigmoidal-shaped voltammetric response under quiescent conditions, characteristicof the different (radial) diffusional ®eld and higher ¯ux of electroactive species ofultramicroelectrodes.(a)(b)FIGURE 1-3 Planar (a) and spherical (b) diffusional ®elds at spherical electrodes.FIGURE 1-4 Concentration pro®les (left) for different potentials during a linear sweepvoltammetric experiment in unstirred solution. The resulting voltammogram is shown on theright, along with the points corresponding to each concentration gradient. (Reproduced withpermission from reference 1.)1-2 FARADAIC PROCESSES 9
  25. 25. Let us see now what happens in a similar linear scan voltammetric experiment,but utilizing a stirred solution. Under these conditions, the bulk concentration(CO…b; t†) is maintained at a distance d by the stirring. It is not in¯uenced by thesurface electron transfer reaction (as long as the ratio of electrode area to solutionvolume is small). The slope of the concentration±distance pro®le‰…CO…b; t† À CO…0; t††=dŠ is thus determined solely by the change in the surfaceconcentration (CO…0; t†). Hence, the decrease in CO…0; t† during the potential scan(around E) results in a sharp rise in the current. When a potential more negativethan Eby 118 mV is reached, CO…0; t† approaches zero, and a limiting current (il) isachieved:il ˆnFADOCO…b; t†d…1-13†The resulting voltammogram thus has a sigmoidal (wave) shape. If the stirringrate (U) is increased, the diffusion layer thickness becomes thinner, according tod ˆBUa…1-14†where B and a are constants for a given system. As a result, the concentrationgradient becomes steeper (see Figure 1-5, curve b), thereby increasing the limitingcurrent. Similar considerations apply to other forced convection systems, e.g., thoserelying on solution ¯ow or electrode rotation (see Sections 3-6 and 4-5, respec-tively). For all of these hydrodynamic systems, the sensitivity of the measurementcan be enhanced by increasing the convection rate.Initially it was assumed that no solution movement occurs within the diffusionlayer. Actually, a velocity gradient exists in a layer, termed the hydrodynamicboundary layer (or the Prandtl layer), where the ¯uid velocity increases from zero atthe interface to the constant bulk value (U). The thickness of the hydrodynamiclayer, dH , is related to that of the diffusion layer:d Dn 1=3dH …1-15†where n is the kinematic viscosity. In aqueous media (with n 9 10À2cm2sÀ1andD 9 10À5cm2sÀ1), dH is $ 10-fold larger than d, indicating negligible convectionwithin the diffusion layer (Figure 1-6). The above discussion applies to other forcedconvection systems, such as ¯ow detectors or rotating electrodes (see Sections 3-6and 4-5, respectively). d values of 10±50 mm and 100±150 mm are common forelectrode rotation and solution stirring, respectively. Additional means for enhancingthe mass transport and thinning the diffusion layer, including the use of powerultrasound, heated electrodes, or laser activation, are currently being studied (3,4a).These methods may simultaneously minimize surface fouling effects, as desired forretaining surface reactivity.10 FUNDAMENTAL CONCEPTS
  26. 26. 1-2.2 Reactions Controlled by the Rate of Electron TransferIn this section we consider experiments in which the current is controlled by the rateof electron transfer (i.e., reactions with suf®ciently fast mass transport). The current±potential relationship for such reactions is different from those discussed (above) formass transport-controlled reactions.FIGURE 1-5 Concentration pro®les for two rates of convection transport: low (curve a) andhigh (curve b).FIGURE 1-6 The hydrodynamic boundary (Prandtl) layer. Also shown (as dotted line), isthe diffusion layer.1-2 FARADAIC PROCESSES 11
  27. 27. Consider again the electron-transfer reaction: O ‡ neÀ„ R; the actual electrontransfer step involves transfer of the electron between the conduction band of theelectrode and a molecular orbital of O or R (e.g., for a reduction, from theconduction band into an unoccupied orbital in O). The rate of the forward(reduction) reaction, Vf , is ®rst order in O:Vf ˆ kf CO…0; t† …1-16†while that of the reversed (oxidation) reaction Vb, is ®rst order in R:Vb ˆ kbCR…0; t† …1-17†where kf and kb are the forward and backward heterogeneous rate constants,respectively. These constants depend upon the operating potential according to thefollowing exponential relationships:kf ˆ kexp‰ÀanF…E À E†=RTŠ …1-18†kb ˆ kexp‰…1 À a†nF…E À E†=RTŠ …1-19†where kis the standard heterogeneous rate constant, and a is the transfer coef®cient.The value of k(in cm sÀ1) re¯ects the reaction between the particular reactant andthe electrode material used. The value of a (between zero and unity) re¯ects thesymmetry of the free energy curve (with respect to the reactants and products). Forsymmetric curves, a will be close to 0.5; a is a measure of the fraction of energy thatis put into the system that is actually used to lower the activation energy (seediscussion in Section 1-2.2.1). Overall, equations (1-18) and (1-19) indicate that bychanging the applied potential we in¯uence kf and kb in an exponential fashion.Positive and negative potentials thus speed up the oxidation and reduction reactions,respectively. For an oxidation, the energy of the electrons in the donor orbital of Rmust be equal or higher than the energy of electrons in the electrode. For reduction,the energy of the electrons in the electrode must be higher than their energy in thereceptor orbital of R.Since the net reaction rate isVnet ˆ Vf À Vb ˆ kf CO…0; t† À kbCR…0; t† …1-20†and as the forward and backward currents are proportional to Vf and Vb, respectively,if ˆ nFAVf …1-21†ib ˆ nFAVb …1-22†12 FUNDAMENTAL CONCEPTS
  28. 28. the overall current is given by the difference between the currents due to the forwardand backward reactions:inet ˆ if À ib ˆ nFA‰kf CO…0; t† À kbCR…0; t†Š …1-23†By substituting the expressions for kf and kb (equations 1-17 and 1-18), one obtainsi ˆ nFAkfCO…0; t† exp‰ÀanF…E À E†=RTŠÀ CR…0; t† exp‰…1 À a†nF…E À E†=RTŠg …1-24†which describes the current±potential relationship for reactions controlled by the rateof electron transfer. Note that the net current depends on both the operating potentialand the surface concentration of each form of the redox couple. For example, Figure1-7 displays the current±potential dependence for the case where CO…0; t† ˆ CR…0; t†and a ˆ 0:50. Large negative potentials accelerate the movement of charge in thecathodic direction, and also decelerate the charge movement in the oppositedirection. As a result the anodic current component becomes negligible and thenet current merges with the cathodic component. The acceleration and decelerationof the cathodic and anodic currents are not necessarily as symmetric as depicted inFigure 1-7, and would differ for a values different than 0.5. Similarly, no cathodiccurrent contribution is observed at suf®ciently large positive potentials.When E ˆ Eeq, no net current is ¯owing. This situation, however, is dynamic,with continuous movement of charge carriers in both directions, and equal opposinganodic and cathodic current components. The absolute magnitude of these compo-FIGURE 1-7 Current±potential curve for the system O ‡ neÀ„ R, assuming that electron-transfer is rate limiting, CO ˆ CR, and a ˆ 0:5. The dotted lines show the cathodic (ic) andanodic (ia) components.1-2 FARADAIC PROCESSES 13
  29. 29. nents at Eeq is the exchange current (io) which is directly proportional to the standardrate constant:io ˆ ic ˆ ia ˆ nFAkC …1-25†where ic and ia are the cathodic and anodic components, respectively.The exchange current density for common redox couples (at room temperature)can range from 10À6mA cmÀ2to A cmÀ2. Equation (1-24) can be written in terms ofthe exchange current to give the Butler±Volmer equation:i ˆ i0 exp…ÀanFZ=RT† À exp‰…1 À a†nFZ=RTŠÈ Ʌ1-26†where Z ˆ E À Eeq is called the overvoltage (i.e., the extra potential beyond theequilibration potential leading to a net current i). The overvoltage is always de®nedwith respect to a speci®c reaction, for which the equilibrium potential is known.Equation (1-26) can be used for extracting information on i0 and a, which areimportant kinetic parameters. For suf®ciently large overvoltages (Z 118 mV=n),one of the exponential terms in equation (1-26) will be negligible compared with theother. For example, at large negative overpotentials, ic ) ia and equation (1-26)becomesi ˆ i0 exp…ÀanFZ=RT† …1-27†and hence we getln i ˆ ln i0 À anFZ=RT …1-28†This logarithmic current±potential dependence was derived by Tafel, and is knownas the Tafel equation. By plotting log i vs. Z one obtains the Tafel plots for thecathodic and anodic branches of the current±overvoltage curve (Figure 1-8). Suchplots are linear only at high values of overpotentials; severe deviations from linearityare observed as Z approaches zero. Extrapolation of the linear portions of these plotsto the zero overvoltage gives an intercept that corresponds to log i0; the slope can beused to obtain the value of the transfer coef®cient a. Another form of the Tafelequation is obtained by rearrangement of equation (1-28):Z ˆ a À b log i …1-29†with b, the Tafel slope, having the value of 2.303 RT=anF. For a ˆ 0:5 and n ˆ 1,this corresponds to 118 mV (at 25 C). Equation (1-29) indicates that the applicationof small potentials (beyond the equilibrium potential) can increase the current bymany orders of magnitude. In practice, however, the current could not rise to anin®nite value due to restrictions from the rate at which the reactant reaches thesurface. (Recall that the rate-determining step depends upon the potential region.)14 FUNDAMENTAL CONCEPTS
  30. 30. For small departures from E, the exponential term in equation (1-27) may belinearized and the current is approximately proportional to Z:i ˆ i0 nFZ=RT …1-30†Hence, the net current is directly proportional to the overvoltage in a narrowpotential range near E.Note also that when inet ˆ 0 (i.e., when E ˆ Eeq) one can obtain the followingfrom equation (1-24):CO…0; t† exp‰ÀanF…Eeq À E†=RTŠ ˆ CR…0; t† exp‰…1 À a†nF…Eeq À E†=RTŠ …1-31†Rearrangement of equation (1-31) yields the exponential form of the Nernstequation:CO…0; t†CR…0; t†ˆ exp‰nF…Eeq À E†=RTŠ …1-32†expected for equilibrium conditions.The equilibrium potential for a given reaction is related to the formal potential:Eeq ˆ E‡ …2:3RT=nF† log Q …1-33†where Q is the equilibrium ratio function (i.e., ratio of the equilibrium concentra-tions).FIGURE 1-8 Tafel plots for the cathodic and anodic branches of the current±potentialcurve.1-2 FARADAIC PROCESSES 15
  31. 31. 1-2.2.1 Activated Complex Theory The effect of the operating potentialupon the rate constants (equations 1-18 and 1-19) can be understood in terms of thefree energy barrier. Figure 1-9 shows a typical Morse potential energy curve for thereaction: O ‡ neÀ„ R, at an inert metallic electrode (with O and R being soluble).Because of the somewhat different structures of O and R, there is a barrier toelectron transfer (associated with changes in bond lengths and bond angles). In orderfor the transition from the oxidized form to occur, it is thus necessary to overcomethe free energy of activation, DGz. The frequency with which the electron crosses theenergy barrier as it moves from the electrode to O (i.e., the rate constant) is given byk ˆ AeÀDGz=RT…1-34†Any alteration in DGzwill thus affect the rate of the reaction. If DGzis increased,the reaction rate will decrease. At equilibrium, the cathodic and anodic activationenergies are equal (DGzc;0 ˆ DGza;0) and the probability of electron transfer is theFIGURE 1-9 Free energy curve for a redox process at a potential more positive than theequilibrium value.16 FUNDAMENTAL CONCEPTS
  32. 32. same in both directions. A, known as the frequency factor, is given as a simplefunction of the Boltzmann constant kHand the Planck constant, h:A ˆkHTh…1-35†Now let us discuss nonequilibirum situations. By varying the potential of theworking electrode, we can in¯uence the free energy of its resident electrons, thusmaking one reaction more favorable. For example, a potential shift E from theequilibrium value moves the O ‡ neÀcurve up or down by f ˆ ÀnFE. The dottedline in Figure 1-10 displays such a change for the case of a positive E. Under thiscondition the barrier for reduction, DGzc , is larger than DGzc;0. A careful study of thenew curve reveals that only a fraction (a) of the energy shift f is actually used toincrease the activation energy barrier, and hence to accelerate the rate of the reaction.Based on the symmetry of the two potential curves, this fraction (the transfercoef®cient) can range from zero to unity. Measured values of a in aqueous solutionshave ranged from 0.2 to 0.8. The term a is thus a measure of the symmetry of theactivation energy barrier. An a value of 0.5 indicates that the activated complex isexactly halfway between the reagents and products on the reaction coordinate (i.e.,FIGURE 1-10 Effect of a change in the applied potential on the free energies of activationfor reduction and oxidation.1-2 FARADAIC PROCESSES 17
  33. 33. an idealized curve). Values of a close to 0.5 are common for metallic electrodes witha simple electron transfer process. The barrier for reduction at E is thus given byDGzc ˆ DGzc;0 ‡ anFE …1-36†Similarly, examination of the ®gure reveals also that the new barrier for oxidation,DGza is lower than DGza;0:DGza ˆ DGza;0 À 1 À a… †nFE …1-37†By substituting the expressions for DGz(equations 1-36 and 1-37) in equation(1-34), we obtain for reductionkf ˆ A exp‰ÀDGzc;0=RTŠ exp‰ÀanFE=RTŠ …1-38†and for oxidationkb ˆ A exp‰ÀDGza;0=RTŠ exp‰…1 À a†nFE=RTŠ …1-39†The ®rst two factors in equations (1-38) and (1-39) are independent of the potential,and thus these equations can be rewritten askf ˆ kf exp‰ÀanFE=RTŠ …1-40†kb ˆ kb exp‰…1 À a†nFE=RTŠ …1-41†When the electrode is at equilibrium with the solution, and when the surfaceconcentrations of O and R are the same, E ˆ E, and kf and kb are equal:kf exp‰ÀanFE=RTŠ ˆ kb exp‰…1 À a†nFE=RTŠ ˆ k…1-42†and correspond to the standard rate constant k. By substituting for kf and kb (usingequation 1-42) in equations (1-40) and (1-41), one obtains equations (1-18) and(1-19) (which describe the effect of the operating potential upon the rate constants).1-3 THE ELECTRICAL DOUBLE LAYERThe electrical double layer is the array of charged particles and/or oriented dipolesthat exists at every material interface. In electrochemistry, such a layer re¯ects theionic zones formed in the solution to compensate for the excess of charge on theelectrode (qe). A positively charged electrode thus attracts a layer of negative ions(and vice versa). Since the interface must be neutral, qe ‡ qs ˆ 0 (where qs is thecharge of the ions in the nearby solution). Accordingly, such a counterlayer is made18 FUNDAMENTAL CONCEPTS
  34. 34. of ions of opposite sign to that of the electrode. As illustrated in Figure 1-11 theelectrical double layer has a complex structure of several distinct parts.The inner layer (closest to the electrode), known as the inner Helmholtz plane(IHP), contains solvent molecules and speci®cally adsorbed ions (which are not fullysolvated). It is de®ned by the locus of points for the speci®cally adsorbed ions. Thenext layer, the outer Helmholtz plane (OHP), re¯ects the imaginary plane passingthrough the center of solvated ions at their closest approach to the surface. Thesolvated ions are nonspeci®cally adsorbed and are attracted to the surface by long-range coulombic forces. Both Helmholtz layers represent the compact layer. Such acompact layer of charges is strongly held by the electrode and can survive even whenthe electrode is pulled out of the solution. The Helmholtz model does not take intoaccount the thermal motion of ions, which loosens them from the compact layer.The outer layer (beyond the compact layer), referred to as the diffuse layer (orGouy layer), is a three-dimensional region of scattered ions, which extends from theOHP into the bulk solution. Such an ionic distribution re¯ects the counterbalancebetween ordering forces of the electrical ®eld and the disorder caused by a randomthermal motion. Based on the equilibrium between these two opposing effects, theconcentration of ionic species at a given distance from the surface, C…x†, decaysexponentially with the ratio between the electrostatic energy (zFF) and the thermalenergy (RT), in accordance with the Boltzmann equation:C…x† ˆ C…0† exp…ÀzFF†=RT† …1-43†The total charge of the compact and diffuse layers equals (and is opposite in sign to)the net charge on the electrode side. The potential±distance pro®le across the double-FIGURE 1-11 Schematic representation of the electrical double layer. IHP ˆ innerHelmholtz plane; OHP ˆouter Helmoltz plane.1-3 THE ELECTRICAL DOUBLE LAYER 19
  35. 35. layer region involves two segments, with a linear increase up to the OHP and anexponential increase within the diffuse layer. These potential drops are displayed inFigure 1-12. Depending upon the ionic strength, the thickness of the double layermay extend to more than 10 nm.The electrical double layer resembles an ordinary (parallel-plate) capacitor. For anideal capacitor, the charge (q) is directly proportional to the potential difference:q ˆ CE …1-44†where C is the capacitance (in farads, F), the ratio of the charge stored to the appliedpotential. The potential±charge relationship for the electrical double layer isq ˆ CdlA…E À Epzc† …1-45†where Cdl is the capacitance per unit area and Epzc is the potential of zero charge(i.e., where the sign of the electrode charge reverses and no net charge exists in thedouble layer). Cdl values are usually in the range 10±40 mF cmÀ2.The capacitance of the double layer consists of combination of the capacitance ofthe compact layer in series with that of the diffuse layer. For two capacitors in series,the total capacitance is given by1Cˆ1CH‡1CG…1-46†FIGURE 1-12 Variation of the potential across the electrical double layer.20 FUNDAMENTAL CONCEPTS
  36. 36. where CH and CG represent that capacitance of the compact and diffuse layers,respectively. The smaller of these capacitances determines the observed behavior. Byanalogy with a parallel-plate (ideal) capacitor, CH is given byCH ˆÀe4pd…1-47†where d is the distance between the plates and e the dielectric constant (e ˆ 78 forwater at room temperature.) Accordingly, CH increases with decreasing separationbetween the electrode surface and the counterionic layer, as well as with increasingdielectric constant of the intervening medium. The value of CG is strongly affectedby the electrolyte concentration; the compact layer is largely independent of theconcentration. For example, at suf®ciently high electrolyte concentration, most of thecharge is con®ned near the Helmholz plane, and little is scattered diffusely into thesolution (i.e., the diffuse double layer becomes small). Under these conditions,1=CH ) 1=CG, 1=C 9 1=CH or C 9 CH. In contrast, for dilute solutions, CG isvery small (compared to CH ) and C 9 CG.Figure 1-13 displays the experimental dependence of the double-layer capacitanceupon the applied potential and electrolyte concentration. As expected for theparallel-plate model, the capacitance is nearly independent of the potential orconcentration over several hundreds of millivolts. Nevertheless, a sharp dip in thecapacitance is observed (around À0.5 V; i.e., the Epzc) with dilute solutions,re¯ecting the contribution of the diffuse layer. Comparison of the double layerwith the parallel-plate capacitor is thus most appropriate at high electrolyteconcentrations (i.e., when C 9 CH ).The charging of the double layer is responsible for the background (residual)current known as the charging current, which limits the detectability of controlled-potential techniques. Such a charging process is nonfaradaic because electrons arenot transferred across the electrode±solution interface. It occurs when a potential isapplied across the double layer, or when the electrode area or capacitances arechanging. Note that the current is the time derivative of the charge. Hence, whensuch processes occur, a residual current ¯ows based on the differential equationi ˆdqdtˆ Cdl AdEdt‡ Cdl…E À Epzc†dAdt‡ A…E À Epzc†dCdldt…1-48†where dE=dt and dA=dt are the potential scan rate and rate of area change,respectively. The second term is applicable to the dropping mercury electrode(discussed in Section 4-2). The term dCdl=dt is important when adsorption processeschange the double-layer capacitance.Alternately, for potential-step experiments (e.g., chronoamperometry, see Section3-1), the charging current is the same as that obtained when a potential step isapplied to a series RC circuit:ic ˆERSeÀt=RCdl…1-49†1-3 THE ELECTRICAL DOUBLE LAYER 21
  37. 37. that is, the current decreases exponentially with time. E is the magnitude of thepotential step, while RS is the (uncompensated) solution resistance.Equation (1-48) can be used for calculating the double-layer capacitance of solidelectrodes. By recording linear scan voltammograms at different scan rates (using thesupporting electrolyte solution), and plotting the charging current (at a givenpotential) versus the scan rate, one obtains a straight line with slope correspondingto Cdl A.Measurements of the double-layer capacitance provide valuable insights intoadsorption and desorption processes, as well as into the structure of ®lm-modi®edelectrodes (6).Further discussion of the electrical double layer can be found in several reviews(5,7±11).1-4 ELECTROCAPILLARY EFFECTElectrocapillarity is the study of the interfacial tension as a function of the electrodepotential. Such a study can provide useful insight into the structure and properties ofFIGURE 1-13 Double-layer capacitance of a mercury drop electrode in NaF solutions ofdifferent concentrations. (Reproduced with permission from reference 5.)22 FUNDAMENTAL CONCEPTS
  38. 38. the electrical double layer. The in¯uence of the electrode±solution potentialdifference upon the surface tension (g) is particularly pronounced at nonrigidelectrodes (such as the dropping mercury electrode discussed in Section 4-5). Aplot of the surface tension versus the potential (like those shown in Figure 1-14), iscalled an electrocapillary curve.The excess charge on the electrode can be obtained from the slope of theelectrocapillary curve (at any potential), by the Lippman equation:@g@E const: pressureˆ q …1-50†The more highly charged the interface becomes, the more the charges repel eachother, thereby decreasing the cohesive forces, lowering the surface tension, and¯attening the mercury drop. The second differential of the electrocapillary plot givesdirectly the differential capacitance of the double layer:@2g@E2 ˆ ÀCdl …1-51†Hence, the differential capacitance represents the slope of the plot of q vs. E.An important point of the electrocapillary curve is its maximum. Such maximumvalue of g, obtained when q ˆ 0, corresponds to the potential of zero charge (Epzc).The surface tension is a maximum because on the uncharged surface there is norepulsion between like charges. The charge on the electrode changes its sign after theFIGURE 1-14 Electrocapillary curve (surface tension g vs. potential).1-4 ELECTROCAPILLARY EFFECT 23
  39. 39. potential passes through the Epzc. Experimental electrocapillary curves have a nearlyparabolic shape around Epzc. Such a parabolic shape corresponds to a linear changeof the charge with the potential. The deviation from a parabolic shape depends onthe solution composition, particularly on the nature of the anions present in theelectrolyte. In particular, speci®c interaction of various anions (e.g., halides) with theFIGURE 1-15 Electrocapillary curves for different electrolytes showing the relativestrength of speci®c adsorption. (Reproduced with permission from reference 5.)FIGURE 1-16 Electrocapillary curves of background (j), ethynylestradiol (), b-estradiol(m) and morgestrel (r). (Reproduced with permission from reference 12.)24 FUNDAMENTAL CONCEPTS
  40. 40. mercury surface, occurring at positive potentials, causes deviations from parabolicbehavior (with shifts of Epzc to more cathodic potentials). As shown in Figure 1-15,the change in the surface tension, and the negative shift in the Epzc, increase in theorder: IÀ BrÀ CNSÀ NOÀ3 OHÀ. (These changes are expected from thestrength of the speci®c adsorption.) Such ions can be speci®cally adsorbed becausethey are not solvated. Inorganic cations, in contrast, are less speci®cally adsorbed(because they are usually hydrated). Similarly, blockage of the surface by a neutraladsorbate often causes depressions in the surface tension in the vicinity of Epzc(Figure 1-16). Note the reduced dependence of g on the potential around thispotential. At more positive or negative potentials, such adsorbates are displaced fromthe surface by oriented water molecules. The electrocapillary approach is notsuitable for measuring the Epzc of solid (i.e., rigid) electrodes.SUPPLEMENTARY READINGSeveral international journals bring together papers and reviews covering innova-tions and trends in the ®eld of electroanalytical chemistry:Bioelectrochemistry and BioenergeticsBiosensors and BioelectronicsElectroanalysisElectrochimica ActaJournal of Applied ElectrochemistryJournal of Electroanalytical and Interfacial ElectrochemistryJournal of the Electrochemical SocietyLangmuirSensors and ActuatorsUseful information can be found in many prominent journals that cater to allbranches of analytical chemistry including The Analyst, Analytica Chimica Acta,Analytical Chemistry, Talanta, Analytical Letters, and Fresenius Zeitschrift fuÈrAnalytical Chemie. Biennial reviews published in the June issue of AnalyticalChemistry offer comprehensive summaries of fundamental and practical researchwork.Many textbooks and reference works dealing with various aspects of electro-analytical chemistry have been published in recent decades. Some of these are givenbelow as suggestions for additional reading, in alphabetic order:W.J. Albery, Electrode Kinetics, Clarendon Press, Oxford, 1975.A.J. Bard and L. Faulkner, Electrochemical Methods, Wiley, New York, 1980.J.M. Bockris and A. Reddy, Modern Electrochemistry, Vol. 1,2, Plenum Press, New York, 1970.A.M. Bond, Modern Polarographic Methods in Analytical Chemistry, Dekker, New York,1980.SUPPLEMENTARY READING 25
  41. 41. C. Brett and A.M. Oliveira Brett, Electrochemistry: Principles, Methods and Applications,Oxford University Press, Oxford, 1993.D. Diamond, Chemical and Biological Sensors, Wiley, New York, 1998.E. Gileadi, Electrode Kinetics, VCH Publishers, New York, 1993.P. Kissinger and W. Heineman, Laboratory Techniques in Electroanalytical Chemistry, (2nded.), Dekker, New York, 1996.J. Janata, Principles of Chemical Sensors, Plenum Press, New York, 1989.J. Koryta and J. Dvorak, Principles of Electrochemistry, Wiley, Chichester, 1987.P. Rieger, Electrochemistry, Prentice-Hall, Englewood Cliffs, NJ, 1987.D. Sawyer and J. Roberts, Experimental Electrochemistry for Chemists, Wiley, New York,1974.M. Smyth and J. Vos, Analytical Voltammetry, Elsevier, Amsterdam, 1992.A.P. Turner, I. Karube and G. Wilson, Biosensors, Oxford University Press, Oxford, 1987.J. Wang, Electroanalytical Techniques in Clinical Chemistry and Laboratory Medicine, VCHPublishers, New York, 1988.REFERENCES1. J.R. Maloy, J. Chem. Educ., 60, 285 (1983).2. M.G. Smith, Laplace Transform Theory, Van Nostrand, London, 1966.3. R.G. Compton, J. Eklund, and F. Marken, Electroanalysis, 9, 509 (1997).4. P. Grundler, and A. Kirbs, Electroanalysis, 11, 223 (1999).4a. J. Alden, and R.G. Compton, Anal. Chem., 72, 198A (2000).5. D. Grahame, Chem. Rev. 41, 441 (1947).6. A. Swietlow, M. Skoog, and G. Johansson, Electroanalysis, 4, 921 (1992).7. D.C. Grahame, Annu. Rev. Phys. Chem., 6, 337 (1955).8. D. Mohilner, Electroanal. Chem., 1, 241 (1966).9. OM. Bockris, M.A. Devanathan, and K. Muller, Proc. R. Soc., 55, A274 (1963).10. R. Parsons, J. Electrochem. Soc., 127, 176C (1980).11. H.B. Mark, Analyst 115, 667 (1990).12. A.M. Bond, I. Heritage, and M. Briggs, Langmuir, 1, 110 (1985).Questions1. Show or draw the concentration pro®le/gradient near the electrode surfaceduring a linear scan voltammetric experiment in stirred a solution. (Use 5±6potentials on both sides of E.) Show also the resulting voltammogram,along with points for each concentration gradient (in a manner analogous toFigure 1-4).2. Describe and draw clearly the structure of the electrical double layer (with its26 FUNDAMENTAL CONCEPTS
  42. 42. several distinct parts).3. Use the activated complex theory for explaining clearly how the appliedpotential affects the rate constant of an electron-transfer reaction. Draw freeenergy curves and use proper equations for your explanation.4. Use equations to demonstrate how an increase of the stirring rate will effect themass transport-controlled limiting current.5. Derive the Nernst equation from the Butler±Volmer equation.6. Explain clearly why polyanionic DNA molecules adsorb onto electrode surfacesat potentials more positive than Epzc, and suggest a protocol for desorbing themback to the solution.7. Which experimental conditions assure that the movement of the electroactivespecies is limited by diffusion? How do these conditions suppress the migrationand convection effects?8. Explain clearly the reason for the peaked response of linear sweep voltammetricexperiments involving a planar macrodisk electrode and a quiescent solution.9. The net current ¯owing at the equilibrium potential is zero, yet this is a dynamicsituation with equal opposing cathodic and anodic current components (whoseabsolute value is i0). Suggest an experimental route for estimating the value ofi0.10. Explain clearly why only a fraction of the energy shift (associated with apotential shift) is used for increasing the activation energy barrier.Questions 27
  43. 43. CHAPTER 2STUDY OF ELECTRODE REACTIONS2-1 CYCLIC VOLTAMMETRYCyclic voltammetry is the most widely used technique for acquiring qualitativeinformation about electrochemical reactions. The power of cyclic voltammetryresults from its ability to rapidly provide considerable information on the thermo-dynamics of redox processes, on the kinetics of heterogeneous electron-transferreactions, and on coupled chemical reactions or adsorption processes. Cyclicvoltammetry is often the ®rst experiment performed in an electroanalytical study.In particular, it offers a rapid location of redox potentials of the electroactive species,and convenient evaluation of the effect of media upon the redox process.Cyclic voltammetry consists of scanning linearly the potential of a stationaryworking electrode (in an unstirred solution) using a triangular potential waveform(Figure 2-1). Depending on the information sought, single or multiple cycles can beused. During the potential sweep, the potentiostat measures the current resultingfrom the applied potential. The resulting plot of current versus potential is termed acyclic voltammogram. The cyclic voltammogram is a complicated, time-dependentfunction of a large number of physical and chemical parameters.Figure 2-2 illustrates the expected response of a reversible redox couple during asingle potential cycle. It is assumed that only the oxidized form O is present initially.Thus, a negative-going potential scan is chosen for the ®rst half-cycle, starting froma value where no reduction occurs. As the applied potential approaches thecharacteristic Efor the redox process, a cathodic current begins to increase, untila peak is reached. After traversing the potential region in which the reductionprocess takes place (at least 90=n mV beyond the peak), the direction of the potential28Analytical Electrochemistry, Second Edition. Joseph WangCopyright # 2000 Wiley-VCHISBNs: 0-471-28272-3 (Hardback); 0-471-22823-0 (Electronic)
  44. 44. sweep is reversed. During the reverse scan, R molecules (generated in the forwardhalf cycle, and accumulated near the surface) are reoxidized back to O and an anodicpeak results.The characteristic peaks in the cyclic voltammogram are caused by the formationof the diffusion layer near the electrode surface. These can best be understood bycarefully examining the concentration±distance pro®les during the potential sweep(see Section 1-2.1.2). For example, Figure 2-3 illustrates four concentrationgradients for the reactant and product at different times corresponding to (a) theinitial potential value, (b) and (d) the formal potential of the couple (during theforward and reversed scans, respectively), and (c) to the achievement of a zeroreactant surface concentration. Note that the continuous change in the surfaceconcentration is coupled with an expansion of the diffusion layer thickness (asFIGURE 2-1 Potential±time excitation signal in cyclic voltammetric experiment.FIGURE 2-2 Typical cyclic voltammogram for a reversible O ‡ neÀ„ R redox process.2-1 CYCLIC VOLTAMMETRY 29
  45. 45. expected in quiescent solutions). The resulting current peaks thus re¯ect thecontinuous change of the concentration gradient with the time. Hence, the increaseto the peak current corresponds to the achievement of diffusion control, while thecurrent drop (beyond the peak) exhibits a tÀ1=2dependence (independent of theapplied potential). For the above reasons, the reversal current has the same shape asthe forward one. As will be discussed in Chapter 4, the use of ultramicroelectrodesÐfor which the mass transport process is dominated by radial (rather than linear)diffusionÐresults in a sigmoidal-shaped cyclic voltammogram.2-1.1 Data InterpretationThe cyclic voltammogram is characterized by several important parameters. Four ofthese observables, the two peak currents and two peak potentials, provide the basisfor the diagnostics developed by Nicholson and Shain (1) for analyzing the cyclicvoltammetric response.FIGURE 2-3 Concentration distribution of the oxidized and reduced forms of the redoxcouple at different times during a cyclic voltammetric experiment corresponding to the initialpotential (a), to the formal potential of the couple during the forward and reversed scans (b, d),and to the achievement of a zero reactant surface concentration (c).30 STUDY OF ELECTRODE REACTIONS
  46. 46. 2-1.1.1 Reversible Systems The peak current for a reversible couple (at25C), is given by the Randles±Sevcik equation:ip ˆ …2:69  105† n3=2ACD1=2v1=2…2-1†where n is the number of electrons, A is the electrode area (in cm2), C is theconcentration (in mol cmÀ3), D is the diffusion coef®cient (in cm2sÀ1), and v is thescan rate (in V sÀ1). Accordingly, the current is directly proportional to concentra-tion and increases with the square root of the scan rate. The ratio of the reverse-to-forward peak currents, ip;r=ip;f , is unity for a simple reversible couple. As will bediscussed in the following sections, this peak ratio can be strongly affected bychemical reactions coupled to the redox process. The current peaks are commonlymeasured by extrapolating the preceding baseline current.The position of the peaks on the potential axis (Ep) is related to the formalpotential of the redox process. The formal potential for a reversible couple iscentered between Ep;a and Ep;c:EˆEp;a ‡ Ep;c2…2-2†The separation between the peak potentials (for a reversible couple) is given byDEp ˆ Ep;a À Ep;c ˆ0:059nV …2-3†Thus, the peak separation can be used to determine the number of electronstransferred, and as a criterion for a Nernstian behavior. Accordingly, a fast one-electron process exhibits a DEp of about 59 mV. Both the cathodic and anodic peakpotentials are independent of the scan rate. It is possible to relate the half-peakpotential (Ep=2, where the current is half of the peak current) to the polarographichalf-wave potential, E1=2:Ep=2 ˆ E1=2 Æ0:028nV …2-4†(The sign is positive for a reduction process.)For multielectron-transfer (reversible) processes, the cyclic voltammogramconsists of several distinct peaks if the Evalues for the individual steps aresuccessively higher and are well separated. An example of such a mechanism is thesix-step reduction of the fullerenes C60 and C70 to yield the hexaanion products C6À60and C6À70 . Such six successive reduction peaks are observed in Figure 2-4.The situation is very different when the redox reaction is slow or coupled with achemical reaction. Indeed, it is these ``nonideal processes that are usually ofgreatest chemical interest and for which the diagnostic power of cyclic voltammetryis most useful. Such information is usually obtained by comparing the experimental2-1 CYCLIC VOLTAMMETRY 31
  47. 47. voltammograms with those derived from theoretical (simulated) ones (1). Propercompensation of the ohmic drop (see Section 4-4) is crucial for such diagnosticapplications of cyclic voltammetry.2-1.1.2 Irreversible and Quasi-Reversible Systems For irreversibleprocesses (those with sluggish electron exchange), the individual peaks are reducedin size and widely separated (Figure 2-5, curve A). Totally irreversible systems arecharacterized by a shift of the peak potential with the scan rate:Ep ˆ EÀRTanaF0:78 À lnkD1=2‡ lnanaFvRT 1=24 5…2-5†where a is the transfer coef®cient and na is the number of electrons involved in thecharge-transfer step. Thus, Ep occurs at potentials higher than E, with the over-potential related to kand a. Independent of the value k, such peak displacementcan be compensated by an appropriate change of the scan rate. The peak potentialand the half-peak potential (at 25C) will differ by 48=an mV. Hence, the voltam-mogram becomes more drawn-out as an decreases.FIGURE 2-4 Cyclic voltammetry of C60 and C70 in an acetonitrile±toluene solution.(Reproduced with permission from reference 2.)32 STUDY OF ELECTRODE REACTIONS
  48. 48. The peak current, given byip ˆ …2:99  105†n…ana†1=2ACD1=2v1=2…2-6†is still proportional to the bulk concentration, but will be lower in height (dependingupon the value of a). Assuming a value of 0.5, the ratio of the reversible-to-irreversible current peaks is 1.27 (i.e., the peak current for the irreversible process isabout 80% of the peak for a reversible one).For quasi-reversible systems (with 10À1 k 10À5cm sÀ1) the current iscontrolled by both the charge transfer and mass transport. The shape of the cyclicvoltammogram is a function of k=paDp(where a ˆ nFv=RT). As k=paDpincreases, the process approaches the reversible case. For small values ofk=paDp(i.e., at very fast v) the system exhibits an irreversible behavior. Overall,the voltammograms of a quasi-reversible system are more drawn-out and exhibit alarger separation in peak potentials compared to those of a reversible system (Figure2-5, curve B).2-1.2 Study of Reaction MechanismsOne of the most important applications of cyclic voltammetry is for qualitativediagnosis of chemical reactions that precede or succeed the redox process (1). Suchreaction mechanisms are commonly classi®ed by using the letters E and C (for theredox and chemical steps, respectively) in the order of the steps in the reactionscheme. The occurrence of such chemical reactions, which directly affect theavailable surface concentration of the electroactive species, is common to redoxprocesses of many important organic and inorganic compounds. Changes in theshape of the cyclic voltammogram, resulting from the chemical competition for theelectrochemical reactant or product, can be extremely useful for elucidating thesereaction pathways and for providing reliable chemical information about reactiveintermediates.FIGURE 2-5 Cyclic voltammograms for irreversible (curve A) and quasi-reversible (curveB) redox processes.2-1 CYCLIC VOLTAMMETRY 33
  49. 49. For example, when the redox system is perturbed by a following chemicalreaction, that is, an EC mechanism,O ‡ neÀ„ R 3 Z …2-7†the cyclic voltammogram will exhibit a smaller reverse peak (because the product Ris chemically removed from the surface). The peak ratio ip;r=ip;f will thus be smallerthan unity; the exact value of the peak ratio can be used to estimated the rate constantof the chemical step. In the extreme case, the chemical reaction may be so fast thatall of R will be converted to Z, and no reverse peak will be observed. A classicexample of such an EC mechanism is the oxidation of the drug chlorpromazine toform a radical cation that reacts with water to give an electroinactive sulfoxide.Ligand exchange reactions (e.g., of iron porphyrin complexes) occurring afterelectron transfer represent another example of such a mechanism.Additional information on the rates of these (and other) coupled chemicalreactions can be achieved by changing the scan rate (i.e., adjusting the experimentaltime scale). In particular, the scan rate controls the time spent between the switchingpotential and the peak potential (during which the chemical reaction occurs). Hence,as illustrated in Figure 2-6, i is the ratio of the rate constant (of the chemical step) tothe scan rate, which controls the peak ratio. Most useful information is obtainedwhen the reaction time lies within the experimental time scale. For scan ratesbetween 0.02 and 200 V sÀ1(common with conventional electrodes), the accessible180 120 60 0 1800.40.20.0Ð0.2Currentfunction(E EV2)n (mV)k/a = 500 100.1, 0.010.10.01–FIGURE 2-6 Cyclic voltammograms for a reversible electron transfer followed by anirreversible step for various ratios of chemical rate constant to scan rate, k/a, wherea ˆ nFv=RT. (Reproduced with permission from reference 1.)34 STUDY OF ELECTRODE REACTIONS
  50. 50. time scale is around 0.1±1000 ms. Ultramicroelectrodes (discussed in Section 4-5.4)offer the use of much faster scan rates and hence the possibility of shifting the upperlimit of follow-up rate constants measurable by cyclic voltammetry (3). For example,highly reactive species generated by the electron transfer, and living for 25 ns can bedetected using a scan rate of 106V sÀ1. A wide variety of fast reactions (includingisomerization and dimerization) can thus be probed. The extraction of suchinformation commonly requires background subtraction to correct for the largecharging current contribution associated with ultrafast scan rates.A special case of the EC mechanism is the catalytic regeneration of O during thechemical step:O ‡ neÀ„ R …2-8†R ‡ A „ O …2-9†An example of such a catalytic EC process is the oxidation of dopamine in thepresence of ascorbic acid (4). The dopamine quinone formed in the redox step isreduced back to dopamine by the ascorbate ion. The peak ratio for such a catalyticreaction is always unity.Other reaction mechanisms can be elucidated in a similar fashion. For example,for a CE mechanism, where a slow chemical reaction precedes the electron transfer,the ratio of ip;r=ip;f is generally larger than unity, and approaches unity as the scanrate decreases. The reverse peak is usually not affected by the coupled reaction,while the forward peak is no longer proportional to the square root of the scan rate.ECE processes, with a chemical step being interposed between electron transfersteps,O1 ‡ neÀ„ R1 3 O2 ‡ neÀ3 R2 …2-10†are also easily explored by cyclic voltammetry, because the two redox couples can beobserved separately. The rate constant of the chemical step can thus be estimatedfrom the relative sizes of the two cyclic voltammetric peaks.Many anodic oxidations involve an ECE pathway. For example, the neurotrans-mitter epinephrine can be oxidized to its quinone, which proceeds via cyclization toleukoadrenochrome. The latter can rapidly undergo electron transfer to formadrenochrome (5). The electrochemical oxidation of aniline is another classicalexample of an ECE pathway (6). The cation radical thus formed rapidly undergoes adimerization reaction to yield an easily oxidized p-aminodiphenylamine product.Another example (of industrial relevance) is the reductive coupling of activatedole®ns to yield a radical anion, which reacts with the parent ole®n to give a reducibledimer (7). If the chemical step is very fast (in comparison to the electron-transferprocess), the system will behave as an EE mechanism (of two successive charge-transfer steps). Table 2-1 summarizes common electrochemical mechanisms invol-ving coupled chemical reactions. Powerful cyclic voltammetric computationalsimulators, exploring the behavior of virtually any user-speci®c mechanism, have2-1 CYCLIC VOLTAMMETRY 35
  51. 51. been developed (9). Such simulated voltammograms can be compared with and®tted to the experimental ones. The new software also provides movie-likepresentations of the corresponding continuous changes in the concentration pro®les.2-1.3 Study of Adsorption ProcessesCyclic voltammetry can also be used for evaluating the interfacial behavior ofelectroactive compounds. Both the reactant and the product can be involved in anTABLE 2-1 Electrochemical Mechanisms Involving Coupled Chemical Reactions1. Reversible electron transfer, no chemical complications:O ‡ neÀ„ R2. Reversible electron transfer followed by a reversible chemical reaction, ErCr mechanism:O ‡ neÀ„ RR 23ktkÀ1Z3. Reversible electron transfer followed by an irreversible chemical reaction, ErCi mechanism:O ‡ neÀ„ RR 23kZ4. Reversible chemical reaction preceding a reversible electron transfer, CrEr mechanism:Z 23k1kÀ1OO ‡ neÀ„ R5. Reversible chemical reaction preceding an irreversible electron transfer, CrEi mechanism:Z 23k1kÀ1OO ‡ neÀ„ R6. Reversible electron transfer followed by an irreversible regeneration of starting materials,catalytic mechanism:O ‡ neÀ„ RR ‡ Z À3kO7. Irreversible electron transfer followed by an irreversible regeneration of starting material:O ‡ neÀ„ RR ‡ Z À3kO8. Multiple electron transfer with intervening chemical reactionÐECE mechanismO ‡ n1eÀ„ R R „ YY ‡ n2eÀ„ ZAdapted with permission from reference 88.36 STUDY OF ELECTRODE REACTIONS
  52. 52. adsorption±desorption process. Such interfacial behavior can occur in studies ofnumerous organic compounds, as well as of metal complexes (if the ligand isspeci®cally adsorbed). For example, Figure 2-7 illustrates repetitive cyclic voltam-mograms, at the hanging mercury drop electrode, for ribo¯avin in a sodiumhydroxide solution. A gradual increase of the cathodic and anodic peak currentsis observed, indicating progressive adsorptive accumulation at the surface. Note alsothat the separation between the peak potentials is smaller than expected for solution-phase processes. Indeed, ideal Nernstian behavior of surface-con®ned nonreactingspecies is manifested by symmetrical cyclic voltammetric peaks …DEp ˆ 0†, and apeak half-width of 90.6=n mV (Figure 2-8). The peak current is directly proportionalto the surface coverage (G) and potential scan rate:ip ˆn2F2GAv4RT…2-11†Recall that Nernstian behavior of diffusing species yields a v1=2dependence. Inpractice, the ideal behavior is approached for relatively slow scan rates, andfor an adsorbed layer that shows no intermolecular interactions and fast electrontransfers.The peak area at saturation (i.e., the quantity of charge consumed during thereduction or adsorption of the adsorbed layer) can be used to calculate the surfacecoverage:Q ˆ nFAG …2-12†This can be used for calculating the area occupied by the adsorbed molecule andhence to predict its orientation on the surface. The surface coverage is commonlyFIGURE 2-7 Repetitive cyclic voltammograms for 1  10À6M ribo¯avin in a 1 mM sodiumhydroxide solution. (Reproduced with permission from reference 10.)2-1 CYCLIC VOLTAMMETRY 37
  53. 53. related to the bulk concentration via the adsorption isotherm. One of the mostfrequently used at present is the Langmuir isotherm:G ˆ GmBC1 ‡ BC …2-13†where Gm is the surface concentration corresponding to a monolayer coverage(mol cmÀ2), and B is the adsorption coef®cient. A linearized isotherm, G ˆ GmBC,is obtained for low adsorbate concentrations (i.e., when 1 ) BC). The Langmuirisotherm is applicable to a monolayer coverage and assumes that there are nointeractions between adsorbed species. Other isotherms (e.g., those of Frumkin orTemkin) take such interactions into account. Indeed, the Langmuir isotherm is aspecial case of the Frumkin isotherm when there are no interactions.When either the reactant (O) or product (R) is adsorbed (but not both), oneexpects to observe a postpeak or prepeak, respectively (at potentials more negativeor positive than the diffusion-controlled peak).Equations have been derived for less ideal situations, involving quasi-reversible andirreversible adsorbing electroactive molecules and different strengths of adsorptionFIGURE 2-8 Ideal cyclic voltammetric behavior for a surface layer on an electrode. Thesurface coverage, G, can be obtained from the area under the peak. (Reproduced withpermission from reference 11.)38 STUDY OF ELECTRODE REACTIONS
  54. 54. of the reactant and product (11±14). The rates of fast adsorption processes can becharacterized by high-speed cyclic voltammetry at ultramicroelectrodes (15).Two general models can describe the kinetics of adsorption. The ®rst involves fastadsorption with mass transport control, while the other involves kinetic control ofthe system. Under the latter (and Langmuirian) conditions, the surface coverage ofthe adsorbate at time t, Gt, is given by.Gt ˆ Ge 1 À exp…ÀkHCtt†Â Å2-14†where Ge is the surface coverage and kHis the adsorption rate constant.The behavior and performance of chemically modi®ed electrodes based onsurface-con®ned redox modi®ers and conducting polymers (Chapter 4) can alsobe investigated by cyclic voltammetry, in a manner similar to that for adsorbedspecies. For example, Figure 2-9 illustrates the use of cyclic voltammetry for in-situprobing of the growth of an electropolymerized ®lm. Changes in the cyclicvoltammetric response of a redox marker (e.g., ferrocyanide) are commonlyemployed for probing the blocking=barrier properties of insulating ®lms (such asself-assembled monolayers).2-1.4 Quantitative ApplicationsCyclic voltammetry can also be useful for quantitative purposes, based on measure-ments of the peak current (equation 2-1). Such quantitative applications require theestablishment of the proper baseline. For neighboring peaks (of a mixture), thebaseline for the second peak is obtained by extrapolating the current decay of theFIGURE 2-9 Repetitive cyclic voltammograms illustrating the continuous growth ofpolyaniline on a platinum surface.2-1 CYCLIC VOLTAMMETRY 39
  55. 55. ®rst one (in accordance with tÀ1=2). Background reactions, primarily those asso-ciated with the double-layer charging and redox-surface processes, limit the detec-tion limit to around the 1  10À5M level. Background-subtracted cyclicvoltammetry can be employed for measuring lower concentrations (16). In particular,fast-scan (1000 V sÀ1) background-subtracted cyclic voltammetry is seeingincreased use for the in-vivo monitoring of neurotransmitters (such as dopamineor serotonin) in the brain. Such coupling of digital background subtraction and fastvoltammetric measurements provides the subsecond temporal resolution necessaryfor detecting dynamic concentration changes in the micromolar range occurring inthe extracellular environment of the brain. The good temporal and chemicalresolution of such in-vivo cyclic voltammetric experiments offers improved under-standing of the chemistry of the brain. These repetitive scanning in-vivo experimentsgenerate large quantities of data that are best represented as three-dimensional(potential, current, time) color contour images (17). For example, the temporalrelease of dopamine following electrical stimulation is evidenced from the rapidincrease in color around its peak potential. The ultrafast scanning also eliminatesinterferences from adsorption processes and chemical reactions that are coupled tothe primary oxidation reaction of catecholamine neurotransmitters (18):HOHOR ROO+ 2H + 2e …2-15†For more detailed information on the theory of cyclic voltammetry, and theinterpretation of cyclic voltammograms, see references (1,7,19,20).2-2 SPECTROELECTROCHEMISTRYThe coupling of optical and electrochemical methodsÐspectroelectrochemistryÐhas been employed for over two decades to investigate a wide variety of inorganic,organic, and biological redox systems (21,22). Such a combination of electroche-mical perturbations with the molecular speci®city of optical monitoring successfullyaddresses the limited structural information available from the current response. Itcan be extremely useful for the elucidation of reaction mechanisms, and for thedelineation of kinetic and thermodynamic parameters. A variety of informativeoptical methods have thus been coupled with electrochemical techniques. While thefollowing sections will focus primarily on transmission=absorption UV-visiblespectroscopic procedures, powerful spectroelectrochemical data can be obtained inre¯ectance experiments (in which the light beam is re¯ected from the electrodesurface), using vibrational spectroscopic investigations, as well as from lumines-cence and scattering spectrochemical studies.2-2.1 Experimental ArrangementOptically transparent electrodes (OTEs), which enable light to be passed throughtheir surface and the adjacent solution, are the keys for performing transmission40 STUDY OF ELECTRODE REACTIONS
  56. 56. spectroelectrochemical experiments. One type of OTE consists of a metal (gold,silver, nickel) micromesh containing small (10±30 mm) holes, which couples goodoptical transmission (over 50%) with good electrical conductivity. Such a minigrid isusually sandwiched between two microscopic slides, which form a thin-layer cell(Figure 2-10). The resulting chamber, containing the electroactive species insolution, contacts a larger container that holds the reference and auxiliary electrodes.The OTE is placed in the spectrophotometer so that the optical beam is passeddirectly through the transparent electrode and the solution. The working volume ofthe cell is only 30±50 mL, and complete electrolysis of the solute requires only 30±60 s. Alternately, the OTE may consist of a thin (100±5000 AÊ ) ®lm of a metal (e.g.,gold or platinum) or a semiconductor (e.g., tin oxide), deposited on a transparentmaterial such as quartz or glass substrate. The ®lm thickness is often selected as acompromise between its electrical conductivity and optical transmission.Improvements in cell designs have been reported in recent years, including theuse of ®ber optics for the illumination and collection of light near electrode surfaces(23), the fabrication of long-path-length OTEs via drilling of a small hole through asolid conducting material for sensitive optical monitoring of weakly absorbingspecies (24,25), and the incorporation of open porous materials (particularlyreticulated vitreous carbon) within a thin-layer compartment (26).2-2.2 Principles and ApplicationsThe primary advantage of spectroelectrochemistry is the cross-correlation ofinformation from the simultaneous electrochemical and optical measurements. Ina typical experiment, one measures absorption changes resulting from speciesproduced (or consumed) in the redox process. The change in absorbance is relatedto concentration and optical path length. Careful evaluation of the temporalabsorbance response (A±t curve) during the electrochemical generation (or consump-tion) of an optically active species can yield extremely useful insights on reactionFIGURE 2-10 Thin-layer spectroelectrochemical cell. OTE ˆ optically transparent electrode.2-2 SPECTROELECTROCHEMISTRY 41
  57. 57. mechanisms and kinetics. Such experiments are particularly useful when the reactantand product have suf®ciently different spectra.Consider, for example, the general redox process:O ‡ neÀ„ R …2-16†When the potential of the OTE is stepped to a value such that reaction (2-16)proceeds at a diffusion-controlled rate, the time-dependent absorbance of R is givenbyA ˆ2COeRD1=2O t1=2p1=2…2-17†where eR is the molar absorptivity of R and DO and CO are the diffusion coef®cientand concentration of O, respectively. Hence, A increases linearly with the square rootof time (t1=2), re¯ecting the continuous generation of R at a rate determined by thediffusion of O to the surface. Equation (2-17) is valid when the generated species isstable. However, when R is a short-lived species (i.e., in an EC mechanism), theabsorbance response will be smaller than that expected from equation (2-17). Therate constant for its decomposition reaction can thus be calculated from the decreasein the absorbance. Many other reaction mechanisms can be studied in a similarfashion from the deviation of the A±t curve from the shape predicted by equation(2-17). Such a potential-step experiment is known as chronoabsorptometry.Thin-layer spectroelectrochemistry can be extremely useful for measuring theformal redox potential (E) and n values. This is accomplished by spectrallydetermining the concentration ratio of oxidized to reduced ([O]=[R]) species ateach applied potential (from the absorbance ratio at the appropriate wavelengths).Since bulk electrolysis is achieved within a few seconds (under thin-layer condi-tions), the whole solution rapidly reaches an equilibrium with each applied potential(in accordance to the Nernst equation). For example, Figure 2-11 shows spectra forthe complex [Tc(dmpe)2Br2] ‡in dimethylformamide using a series of potentials[dmpe is 1,2-bis(dimethylphosphine)ethane]. The logarithm of the resulting concen-tration ratio ([O]=[R]) can be plotted against the applied potential to yield a straightline, with an intercept corresponding to the formal potential. The slope of thisNernstian plot (0.059= n V) can be used to determine the value of n.Besides potential-step experiments, it is possible to employ linear potential scanperturbations of an OTE (27). This voltabsorptometric approach results in an opticalanalogue of a voltammetric experiment. A dA=dE vs. E plot (obtained by differ-entiating the absorbance of the reaction product with respect to the changingpotential) is morphologically identical to the voltammetric response for the redoxprocess (Figure 2-12). Depending upon the molar absorptivity of the monitoredspecies, the derivative optical response may afford a more sensitive tool than thevoltammetric one. This concept is also not prone to charging-current backgroundcontributions and holds considerable promise for mechanism diagnosis and kineticcharacterization of coupled chemical reactions.42 STUDY OF ELECTRODE REACTIONS
  58. 58. FIGURE 2-11 Spectra for a series of applied potentials (mV vs. Ag=AgC1) during thin-layer spectroelectrochemical experiment on 1:04  10À3M ‰Tc…III†…dmpe†2Br2Š‡. Medium isdimethylformamide containing 0.5 M TEAP ˆ tetraethylammonium perchlorate. (Reproducedwith permission from reference 27.)FIGURE 2-12 Plot of dA=dt vs. E for 1:55  10À3M methyl viologen at tin oxide opticallytransparent electrode, using scan rates of (A) 25, (B) 50, (C) 97.2 and (D) 265 mV sÀ1.(Reproduced with permission from reference 28.)2-2 SPECTROELECTROCHEMISTRY 43
  59. 59. Spectroelectrochemical experiments can be used to probe various adsorp-tion=desorption processes. In particular, changes in the absorbance accruing fromsuch processes can be probed utilizing the large ratio of surface area to solutionvolume of OTEs with long optical path length (29). Additional information on suchprocesses can be obtained from the Raman spectroelectrochemical experimentsdescribed later.In addition to transmission experiments, it is possible to use more sensitivere¯ectance protocols. In particular, in internal re¯ectance spectroscopy (IRS) thelight beam is introduced to the electrode at an angle, and the spectrum is recordedfrom the re¯ected beam at the solid±solution interface. Prisms are used to allow theradiation enter and leave. In addition to its higher sensitivity, IRS is less prone tosolution resistance effects.Infrared spectroelectrochemical methods, particularly those based on Fouriertransform infrared (FTIR) spectroscopy can provide structural information that UV-visible absorbance techniques do not. FTIR spectroelectrochemistry has thus beenfruitful in the characterization of reactions occurring on electrode surfaces. Thetechnique requires very thin cells to overcome solvent absorption problems.Besides its widespread use for investigating the mechanism of redox processes,spectroelectrochemistry can be useful for analytical purposes. In particular, thesimultaneous pro®ling of optical and electrochemical properties can enhance theoverall selectivity of different sensing (30) and detection (31) applications. Suchcoupling of two modes of selectivity is facilitated by the judicious choice of theoperating potential and wavelength.2-2.3 Other Spectroelectrochemical and Spectroscopic TechniquesIn addition to UV-visible absorption measurements, other spectroscopic techniquescan be used for monitoring the dynamics of electrochemical events or the fate ofelectrogenerated species. Particularly informative are the couplings of electrochem-istry with electron spin resonance, nuclear magnetic resonance, and mass spectro-scopy. A variety of specially designed cells have been constructed to facilitate suchstudies, and several reviews have been published (32±36). Electrochemilumines-cence (ECL) is another useful technique for studying the fate of electrogeneratedradicals that emit light. It involves the formation of light-emitting excited-statespecies as a result of fast and highly energetic electron-transfer reactions of reactantsformed electrochemically (37,38). Various organic and inorganic substances (e.g.,polycyclic hydrocarbons, nitro compounds, luminol, Ru…bpy†2‡3 ) can produce ECLupon electron transfer from electrodes, in connection with the formation of radicalions. The electrogenerated radicals behave as very strong redox agents, and theirreactions with each other or with other substances are suf®ciently energetic to beable to populate excited states. ECL experiments are usually carried out by recordingthe spectra of the emitted light using a deoxygenated nonaqueous medium (e.g.,highly puri®ed acetonitrile or dimethylformamide). Analytical applications of ECLrelying on the linear dependence of the ECL intensity and the reactant concentrationhave also been realized (39).44 STUDY OF ELECTRODE REACTIONS

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