INTERNATIONAL SOCIETY OF ELECTROCHEMISTRY (I.S.E.) w (formerly C.I.T.C.E.) SOClÉTÉ INTERNATIONALE D'ELECTROCHIMIE (S.I.E.)
-i (anciennement C.I.T.C.E.) f
G lMlS-r>v^f- 6?2
OJ IB f (ö -I 0 ö I o- l 23rd MEETING
$ m, Ds n< Z ; c ; »i' EXTENDED ABSTRACTS
STOCKHOLM, SWEDEN August 27th-September 2nd, 1972 We regret that some of the pages in the microfiche copy of this report may not be up to the proper legibility standards, even though the best possible copy was used for preparing the master fiche. INTERNATIONAL SOCIETY OF ELECTROCHEMISTRY (I.S.E.) SOCIETE INTERNATIONALE D"ELECTROCHIMIE (S.I.E.)
23rd MEETING including
SYMPOSIUM ON ELECTROCHEMICAL ENGINEERING (114th EVENT OF THE EUROPEAN FEDERATION OF CHEMICAL ENGINEERING)
and
SYMPOSIUM ON ACCELERATED CORROSION TESTING WITH ELECTROCHEMICAL METHODS (66th EVENT OF THE EUROPEAN FEDERATION OF CORROSION)
EXTENDED ABSTRACTS
STOCKHOLM, SWEDEN AUGUST 27th - SEPTEMBER 2nd, 1972 This volume has been produced by off-set printing from originals as received from the authors» Technical Editor: Mr. Jaak Berendson, Met.Eng., The Royal Institute of Technology, Stockholm, Sweden. Printing Office: AB Realtryck, Stockholm, Sweden. Edition: 600 copies. CONTENTS
SYMPOSIA
ELECTROCHEMICAL ENGINEERING
Optimization of Electrolytic Processes (Plenary Lecture) N. Ibl
Design and Materials of Construction for Electrolytic Cells (Plenary Lecture") Robert B. MacMullin
Some Fundamencal Aspects of Industrial Electro lytic Processes (Plenary Lecture) Charles W. Tobias
Dimensionally Stable Anodes in Chlorate Electrolysis R.T. Atanasoski, A.O. Filip, B.Z. Nikolic, M.M. JakSic and A.R. Despic
Cells for Electro-Organic Synthesis Fritz Beck
Basic Problems of the Electrochemical Production of Chlorine Giuseppe Bjanchi and Giuseppe Faita
Experimental Realisation of a Clamped Ionic Concentration Gradient Michel Delmotte. and Jacques Chanu
Pulsating Potential Electrolysis in Deposition of Copper Aleksandar Despic and Gordana Savic Maglic
Electroneutrality and Steady State during the Transfer of Electric Current through Graded Electrolyte Solutions Chaim Forgacs II
New Concepts in the Scaleup of Chlorine-Caustic Mercury Cells P. Gallone 25
Performance Studies of a Bipolar Packed Bed Cell for the Production of Propylene Oxide Francis Goodridge and Ola Osifade 28
Some Performance Characteristics of Three- Dimensional Electrodes Martin Fleischmann and Francis Goodridge 30
A Model for Membrane Fouling in Electrodialysis G. Grossman and A.A. Sonin 33
Mass Transfer Determined Reactions in Flui- dised Bed Cells Ewald Heitz and Siegfried Pionteck 36
Some Surface Studies of Electrochemically Machined High-Temperature Alloys James P. Hoare, Armand J. Chartrand and Mitchell A. LaBoda 38
The Influence of Foreign Atoms on the Proper ties of Electrolytic Zinc Powder I.N. Justinijanovie, J.N. Jovicevic and A.R. Despic 41
On the Anodic Dissolution of Metals at High Current Densities V.D. Kashcheey, B.N. Kabanov and A.D. Davydov 42
Sulphate Electrolysis Based on Anionic Proton Migration Joseph Kerti 43
Some Features of a Novel Bi-Polar Electrolytic Flow Cell Christopher J.H. King 46
Electrochemical Investigations of the Cylindrical Ro'tating Electrode Dj. Matic, B. Lovrecek and D. Skansi 49
The Glanor Diaphragm Cell: .A New Tool for the Chlorine Industry Vittorio de Nora 51 Theory of Electrochemical Machining S.K^Rangarajan, T.G. Ramesh and S. Ramabhadran 52 Study of Mass Transfer in Flowing Electrolyte with Quasi Porous Electrode Branko LovreSek, Darke Skans i and Djani Matic 53
Electrokinetic Salt Rejection by Porous Materials - Theory and Experiment G. Jacazio, R.F. Probstein, A.A. Sonin and D. Yung 54
Design Approach to the Chlorate Cell M.M. Jaksic, N.M. Jaksic, M.D. Spasojevic and B.Z. Nikolic " " 57
Concentration-Polarization at Ion-Exchange Membrane/Solution Interfaces C. Forgacs, J. Leibovitz, J. Sinkovic and K.S. Spiegler 58
Engineering Aspects of Gold Plating for Electronic Devices Dennis Turner 62
The Effect of Superimposed Alternating and Direct Current on the Internal Stresses in Nickel Deposits Jean Vereecken and René Winand 63
Ionic Mass Transfer by Free Convection with Simultaneous Heat Transfer A.A. Wragg and A.K. Nasiruddin 65
Calculation of Local Current Densities and Terminal Voltage for a Monopolar Sandwich Electro- lyzer; Application to Chlorate Cells I. Rousar, V. Cezner, J, Hostomsky, M.M. Jak'sic, M. Spasojevic and B.Z. Nikolic 68
The Decomposition of Sodium Amalgam in a Horizontal Decomposer J. Hostomsky, I. Rousar and V. Cezner 70
Sodium Amalgam Decomposition in a Vertical Tower S. Rajasekaran4 I. Rousar, J. Hostomsky and V. Cezner 72 ACCELERATED CORROSION TESTING WITH ELECTROCHEMICAL METHODS
Controlled Potential Methods of Corrosion Testing (Plenary Lecture) Norbert D. Greene 77
Use of Methods of Electrochemistry and Radio active Indicators for Determining Low Corrosion Rates (Plenary Lecture) Ya.M. Kolotyrkin 79
The Polarization Resistance Method of Corrosion Testing (Plenary Lecture) Milan Prazäk 81
The Radiographical Method of the Examination of Permeability of Paint Coating Modified with Corrosion Inhibitor L. Chromy and A. Zimnoch 84
Pitting Potential Measurements on Stainless Steels, as Compared "with Rest Potentials and the Electrochemical Behaviour of Passive Films J.M. Defranoux 85
Intergranular Corrosion of Austenitic Stainless Steels in the Sensitized or Solution-Annealed Condition André Desestret, Michel Froment and Pierre Guiraldenq 87
Electrochemical Methods for Corrosion Evaluation of Low-Alloy Steels, Copper Alloys and Stainless Steels Lucio Giuliani, Alberto Tamba and Roberto Bruno 90
Accelerated Stress-Corrosion Cracking of Ti-8%A1- l%Mo-l%V in the Presence of a Hydrogen-Recombina tion Poison J..A.S. Green and A.J. Sedriks 93
Accelerated Method for the Investigation of the Susceptibility of Stainless Steels to Intergranular Corrosion by Potentiodynamic Polarization Lajos Hackl, G. Muller, J. Horvåth and F. Märta 96
Basic Principles of Electrochemical Corrosion Testing Michel Keddam 99 v
55 59 • The Application of Fe and Fe * to Corrosion Testing Pance Kirkov and Zarko Radosavljevic 102 The Electrochemical Behaviour of Nimonic and Tita nium Alloys in Molten Salts David G. Lovering 103
An Electrochemical Test for Quality of Conversion Coatings on Al Alloys F. Mansfeld and E.P. Parry 106
On the So-Called Linear Polarization Method for Measurement of Corrosion Rates K.B. Oldham and F. Mansfeld 109
Evaluation of the Susceptibility of Austenitic Stainless Steels to Stress C.C. by Chemical or Electrochemical Methods Giuseppe Bianchi, Alessandro Cerquetti, Francesco Mazza and Sandro Torchio 112
Comparison of Corrosion Rates in Sulfuric Acid Obtained Using Resistance-Polarization, Tafel- Extrapolation, and Weight-Loss Techniques James R. Myers 115
Electrochemical Studies of Corrosion-Erosion during Hydraulic Transport of Solids in Pipes J. Postlethwaite and E.B. Tinker 116
Electrochemical Corrosion Testing of Metals in Alkali Sulphate Melts A. Rahmel 119
Electrochemical Investigations on Pitting Giinter Herbsleb and Wilhelm Schwenk 120 Transient Straining Electrodes for Studying Stress Corrosion Cracking R.W. Staehle, J.B. Lumsden and S. Abe 123
Experiments on the Acceleration of the Intergranular Stress Corrosion Cracking of Type AISI 304 Stainless Steel by Means of Potentiostatic Polarization Pee ter Tarkpea, Kjell Lundberg and Walter Hiibner 126
Corrosion Tests with the Aid of Electrochemical Measurements on Nickelbase Alloys under Pressure Water Conditions Norbert Wieling 129 SECTION MEETINGS
SECTION 1 -ELECTROCHEMICAL THERMODYNAMICS i • i—t i i i —• —— —• Thermodynamics of Surface Layers. (Invited Paper) Hermann Göhr 133
The Adsorption of Acetone at the Mercury - Solu tion Interface Zofia Borkowska and Barbara Behr 136
Entropies of Silver Halides in Water, Molten Calcium Nitrate Tetrahydrate and Anhydrous Nitrate Melts Brian Burrows and Soefjan Noersjamsi 139
Thermodynamics of Mixed Aqueous Solvents: Transfer Parameters Gerard Douhéret 141
The Transference Numbers in Aqueous Solutions of 1:1 Type Salts as a Function of Temperature and Concentration Nadezda Jakovljevic Halai 144
Polarographic Determination of Hydrogen Ion Activities in Strongly Acidic Media: A New Acidity Function Jiri Janata and Geertje Jansen 147
Stability of Laminar Flow of Liquid Mixture with Temperature Gradient Claude Klapisz and Jacques Chanu 149
Adsorption of Tetrapropyl- and Tetrabutyl- Ammonium Cations and Higher Alkylsulphonate Anions on Mercury J. Kuta and I. Smoler 152
A Method of Estimating the Reaction of Electrodes at the Surface of Sulphide Minerals in Conditions of Flotation Janusz Lekki 156
Determination of the Equilibrium Constant of the Side Reaction Occuring in the Concentration Cell Zbigniew Moser and Krzysztof Fitzner 159 VII
Quantum Mechanical Estimation of Heats of Adsorp tion of Atomic Oxygen on some Metal Surfaces Walter Sarholz, Detlef Baresel and Gunter Schuiz-Ekloff 162
Thermodynamics and Kinetics of Metal Ion Adsorption J.W. Schultze 163
EMF of the Cells (Pt)Ho/H Hal, Hg2Halo/Hg in Organic Solvents and their Mixtures with Water Kurt Schwabe 166
Coupled Self and Chemical Diffusion in Electro lyte Solutions. Diffusion Potential and Corre sponding Reciprocal Relations Pierre Turq and Marius Chemla 169
SECTION 2 - ELECTROCHEMICAL KINETICS
Kinetics of Thin Oxide Layer Formation on Noble Metals (Invited Paper) K.J. Vetter 175
Kinetics of the Electrochemical Formation of Parathiocyanogen Films A.J. Arvia and A.J. Calandra 178
Activities or Concentrations in Equations of Electrode Kinetics M.M. Jaksic, R.T. Atanasoski, A.R. Despic and V. Nakic 180
Electrode Reaction of Cu(II) - EDTA on Mercury. Effects of Double Layer and Ion Pair Formation A. Baric, M. Branica and J. KSta 181
Identification of Transient Phenomena during the Anodic Polarization of Iron in Dilute Sulphuric Acid G.J. Bignold and M. Fleischmann 184
The Electrode Potential of Metal Dendrite Growth at the Limiting Current Aurelian Cälu§aru 187 The Electrochemical Reduction of Sulfur in Organic Media Raymond Bonnaterre and Georges Cauquis 190
Electrochemical Redox Processes of Uranium in Aqueous Solutions of Acetylacetone B. Cosovic, Lj. Jeftic and M. Branica 192
Nucleation Mechanism and Electrochemical Transport Processes in the Oxide Formation during the Anodisation of Aluminium Paul Csokän 195
Laplace Plane Analysis of Electrochemical Systems: Application to Hg/Hg2 Karl Doblhofer and Arthur A. Pilla 198
Studies on Kinetics of Electrophoretic Deposition M. Elmas 200
Galvanostatic Study of Follow-Up Chemical Reactions Oldrich Fischer, Josef Bezdék and Vladimir Komärek 203
Influence of Adsorbed Electroinactive Substances on Charge-Transfer Processes Rolando Guidelli 206
Versuche zur Bestimmung des pH-Wertes Konzentrier- ter Saurer Lösungen uber die Kinetik der Wasser- stoffelektrode S. Ernst and C.H. Hamann 207
Electrodeposition of Palladium Powder N. Ibl, G. Gut and M. Weber 208
Electrolytic Reduction of Nitric Oxide to Hydroxylamine L.J,J, Janssen and J.G. Hoogland 209
Influence of pH on Electrical Behaviour of Barrier Layer on Aluminium Olga Korelic and Branko Lovrecek 211
The Influence of Halide Ions on the Kinetics of Iron Corrosion N.A. Darwish, F. Hilbert, W.J. Lorenz and H. Rosswag ~ 213
Electrochemical Investigation of Valve Metals M. Metxkos, B. Lovrecek and B. Jaric 216 Etude du Mecanisme de l'Oxydo-Reduction du Pt en Solution H2SO, , N Guy Bronoöl and Eliane Momot 217
Specific Role of Chloride in Anodic Reactions on Gold B. Lovrecek, K. Moslavac and R. Radeka 220
Adsorption and Inhibition in the Hydroqui- none/Quinone Redox-System on Mercury W.J. Plieth, I. Stellmacher and B. Quast 222
Electrochemical Stability and Electrocatalytic Behaviour of Sodium-Tungsten Bronzes Jean-Paul Randin and Ashok K. Vijh 224
The Influence of some Anions on the Anodic Dissolution of Copper John C. Reeve 227
Hydrogen Peroxide Behavior on Silver and Nickel Electrodes in Alkaline Solution N.A. Shumilova, N.D. Merkulova, G.V. Zhitaeva, E.I. Khrushcheva, G.P. Samoilov and V.S. Bagotzky 230
The Influence of Surface Oxides of Platinum on the Rate of Anodic Oxygen Evolution in Solutions of Sulphuric Acid Otomar Spalek and Jan Balej 231
Use of Rotating Disc Cathodes for Studying Hydrogen Permeation through Iron and Steel Zuzanna Szklarska-Smiaiowska, Michal Smialowski and Tadeusz Zakroczymski 234
Kinetics of Zinc Reduction in Water-Methanol Solutions of NaCIO, Joanna Taraszewska 237
Electrode Kinetic and Preparative Investiga tion of the Direct Anodic Oxidation of Olefins on Carbon Anodes in Non-Aqueous Solutions M. Katz and H. Wendt 240
Contribution to the Study of the Mechanism of the Anodic Dissolution of Cu2S Patrick Brennet, Jean Vereecken and Rene Winand 243 A.C. Polarographic Study of D.C. Polarographic Non-Additive Diffusion Current Phenomena and Underlying Fast Homogeneous Redox-Reactions H. Yamaoka 244
The Effect of Chloride Ions on the Electro chemical Behaviour of Nickel at Mercury Electrodes v Jean Chevalet and Vera Zutic 247 Redox Processes of Uranium(VI) Peroxo Com plexes in Alkaline Hydroxide Solutions Vera Sutic and Marko Branica 250
SECTION 3 - EXPERIMENTAL METHODS IN ELECTROCHEMISTRY Recent Progress of Impedance Measurements at Very Low Frequencies (Invited Paper) Israel Epelboin, Claude Gabrielli, Michel Keddam and Hisasi Takenouti 255
Polarographic Behaviour of Isocinchomeronic Acid Luigi Campanella, Pierluigi Cignini and Giorgio De Angelis 258
Experimental Studies of Complex Reaction Mechanisms of Uranium in Acidic Media Dunj a Cukman, M. Vukovic and V. Pravdic 261
Theory and Applications of Microelectrodes F. Aimeur, M. Daguenet, F. Kermiche and M. Meklati 264
Adsorption of Organics on Non-Noble Metals P.M. Dragic and N.R. Tomov 266
Evidence of the Passivity of Cobalt in Sulphuric Acid Medium Israel Epelboin, Claude Gabrielli and Philippe Morel 268
Applications of Programmed Sampling Pulse Techniques in Electrochemical Analysis and Research Jacques Tacussel and Jean-Jacques Fombon 271 XI
Polarographic Studies on Electrode Kinetics and Formation Constants of the Complexes Formed by Zn^+ with Ortho and Meta-Toluate Ions J.N. Gaur, D.S. Jain and Anand Kumar 274
Faradayic Impedance of Electrochemical Systems with Charge-Transfer Overvoltage A.V. Gorodisky, Y.K. Delimarsky, A.V. Panov, N.H. Tumanova and B.F. Ometsinsky 277
Logarithmic Analysis of Two Overlapping D.C. Polarographic Waves Using Computer I. Ru£ic and Lj. Jeftic 279
Anodic Behaviour of Nickel in Sulfuric Acid B. Dubois, A. Jouanneau and M.C. Petit 282
An Evaluation of the Activity Coefficients of HCl in Water-Dioxan and Water-THF Mixtures by Use of the Conductimetric Method Marie-Claude Justice and Jean-Claude Justice 285
An Extension towards Lower Dielectric Constant Solvents of the Conductimetric Method for Electrolytes Studies Christian Micheletti and Jean-Claude Justice 287
Advanced Techniques for E.S.R. Investigations of Electrochemically Generated Radicals Bertel Kastening, B. Gostisa-Mihelcic and J. Divis'ek 290
The Elimination of Resistance Polarization in Potentiostatic Investigations B. Lengyel, jun., J. Devay and J. Mészäros 291
Recording of Z-Shaped Passivation Current- Voltage Curves Claude Gabrielli, Michel Keddam and Jean-Claude Lestrade 294
Polarography of Lanthanides in N-N-Dimethyl- form amide Alain Levéque and Robert Rosset 297
Pulse Polarographic Instrumentation and Methodology Janet Osteryoung, R.A. Osteryoung and J.H. Christie 300 Electrometric Determination of Stability Constants of Zinc, Cadmium and Lead Complexes in Buffer Solutions of 3-Hydroxypropionic Acid and a-, 8- and ^-Hydroxybutyric Acids I. Piljac, S. Nushi, B. Bach-Dragutinovic, B. Grabaric and I. Filipovic 302
Studies of Complex Electrochemical-Chemical Reaction Mechanisms with the Rotating Ring- Disc Electrode of Platinum, Amalgamated Platinum and Nickel Nikola Bonacci and Velimir Pravdij: 304
A General Approximative Method for the Inter pretation of Kinetic D.C. Polarographic Waves I. Ruzic 307
The Exchange Current Density vs.Concentration Relation and its Use in a Rigorous Determina tion of Solution Purity Sigmund Schuldiner and Murray Rosen 310
Some Investigations of Experimental Methods for Rapid Acquisition of Faradaic Admittance Frequency Response Profiles Donald E. Smith, Samuel C. Creason and John W. Hayes 312
Electro-Optical Study of Interfacial Polarisation and Hydrodynamic Motion in an Electrodialysis Cell Elisabeth Steyger , 315
Polarographic Reduction of Ni/II/ in the Presence of Ethylenedithioldiacetic Acid Solution Desanka Suznjevic and Milenko Susie 318
The "VOCTAN", an Apparatus for Electrochemical Techniques Involving Voltage/Current/Time Relationships Jacques R. Tacussel 319
Cyclic Chronopotentiometry. Digital Simulation and Experimental Verification of Model Systems of Higher Order Chemical Reactions Coupled with Electron Transfer Marijan Vukovic and Velimir Pravdic 322 XIII
SECTION 4 - HIGH TEMPERATURE ELECTROCHEMISTRY
Electrochemical Processes with Rate-Determining Preceding Acid-Base Reactions in Molten Electro lytes (Invited Paper) Y.K. Delimarsky, V.I. Shapoval and V.F. Grishchenko 327
Electrolysis of Plutonium and its Alloys Pierre Aury 329
Factors Influencing the EMF Measurement with Solid Electrolyte Cells in Liquid Metals W.A. Fischer and D. Janke 332
Calorimetric Study of the Heat of Mixing of NaF and ZrF^ at 1030°C André Fontana and René Winand 333
Electrolytic Deoxidation of Liquid Metals W.A. Fischer and D. Janke 336
The Determination of the Potentials of the Points of Zero Charge on Liquid in Contact with a Solid Electrolyte Sergey Karpachov 337
On a Unified E.M.F. Series in Molten Salts and other Solvents J.A.A. Retelaar 338
Structure and Transport in Aqueous Melts David G. Lovering 340
Isotope Effects of Electromigration in Molten Alkali Nitrates Arnold Lunden and Isao Okada 343
Electrolytic Plasmas in Molten Salts and Aqueous Solutions Pierre Mergault, Jean-Claude Valognes, Denis Letemturier, Jocelyne Garbarz-Olivier and Christian Guilpin 345
The Coulometric Deoxidation (Refining) of Liquid Metals K.E. 'Öberg, L.M. Friedman, W.M. Boorstein and R.A. Rapp ' 348
Electrochemical Studies in Aluminium Chloride Melts H.L, Jones, L.G. Boxall and R.A. Osteryoung 351 A Contribution to the Study of the Pb/PbSO^,H2S04 Electrode up to 250°C F. Letowski and G. Pinard-Legry 353
The Cathode Reaction on Aluminium in NaF-AlF3-Al203 Melts Jomar Thonstad and Sverre Rolseth 35 6
Determination of Electrode Parameters for the Chlorine/Carbon Electrode in Chloride Melts T. Berge, K.A. Paulsen and R. Tunold 358
The Electrochemistry of Oxygen and its Solubi lity in Molten Alkali Nitrates E. Desimoni, F. Paniccia and P.G. Zambonin 360
SECTION 5 - CORROSION
A Comparison of the Dissolution-Precipitation and Solid State Mechanisms of Passivation R. D. Armsj: r on g and J. A. Harrison 365
The Anodic Dissolution of Iron. V. Some Obser vations Regarding the Influence of Cold Working and of Annealing of Iron on the Anodic Dissolution Behaviour of the Metal Greg%rs Bech-Nielsen 367
About the Electrolytic Oxygen-Corrosion of Mild Steel Konrad Bohnenkamp 370
An Electrometric Study of the Tarnish Film Formed on Cu-Zn Alloys in Stress Corrosion Environments Based on Ammoniacal CuSO, Solutions Charles Booker and Muhammad Salim 373
Pitting Corrosion in Stainless Steel: Mechanism Sven Brennert and Göran Eklund 374
Electrochemical Studies Concerning Caustic Stress Corrosion Cracking of Alloyed Steels jLennart Dahl and Tommy Dahlgren 377
El ectrochemical Reactions and Corrosion in Molten Alkaline Sulfate and Vanadate J. Richard and J. Dubois 380 XV
Investigation on the Second Anodic Current Maximum on the Polarization Curves of Commercial Stainless Steels in Sulphuric Acid Liljana Felloni, G. Paolo Cammarota, Silvana Sostero and G. Luigi Zucchini 381
Anodic Oxidation of Rhenium A. Giraudeau, P. Lemoine and M. Gross 384
Influence d'un Milieu Bacterien sur l'Equilibre Electrochimique de 1'Aluminium Jean Brisou, Marie-Josephe Croissant, Jane Grimaudeau Irene Guillaume and Gabriel Valensi 386
High Temperature Potential/pH Equilibrium Diagrams of Metal/Sulphur/ Water Ternary Systems with Applications to Corrosion in Aqueous Hydro gen Sulphide Environments J. Horvath, G. Bencze and F. Märta 389
Some Electrochemical Aspects of Paint Coatings Romuald Juchniewicz 390
Corrosion of Stepwise Dissolving Metals; Corrosion Mechanism of Indium V.V. Losev and A.P. Pchelnikov 392
Influence of Cold Plastic Deformation on Corrosion Resistance of Austenitic Stainless Steels in Acid Aggressive Media and on their Susceptibility to the Pitting Corrosion Dany Sinigaglia, Pietro Pedeferri, Luisa Peraldo- Bicelli, Bruno Mazza and Gianpaolo Galliani 395
Pitting Evolution on a Stainless Steel in an Oxidizing Chloride Medium H. Coriou, A. Monnier, G. Pinard-Legry and G. PIante 397
The Electrochemical Behaviour of Molybdenum in some Corrosive Environments Octavian Radovici and Paula Matei 400
Corrosion Investigations of some Electro- catalytic Materials G. Richter, K. Mund and E. Weidlich 403 Corrosion Studies with a Twin Electrode Thin Layer Technique K. Ruber and E. Schmidt 406
Analysis of the Influence of Hydrogen on Pitting Corrosion and Stress Corrosion of Austenitic Stainless Steel in Chloride Environment A.A. Seys, M.J. Brabers and A.A. Van Haute 408
Synergistic Action of N-Dodecylamine with N-Caprynic Acid on the Corrosion, of Steel in Sulphuric Acid Solution Grzegorz Wieczorek and Zuzanna Szklarska- SmiaJTowska 410
The Effect of Using Road Salt on the Corrosion Climate for Vehicles Ulf Ulfvarson and Kurt Johansson 413
The Dissolution Kinetics of Lithiated NiO in Aqueous Acid Solutions Chin-Ho Lee, Alan Riga and Ernest B. Yeager 414
SECTION 6 - BATTERIES
Comparative Study of Fuel Cell Electrodes and Technology (Accomplishments 1968-1972, Outlook) (Invited Paper) Karl V. Kordesch 419
Discussion of Various Types of Test Electrodes for the Evaluation of Catalysts for the Electro- reduction of Oxygen H. Behret, H. Binder, A. Köhling and G. Sandstede 421
Crude Gas / Air Fuel Cell with Bipolar, Non-Noble Metal Electrodes and Acid Electrolyte Lothar Baudendistel, Harald Bohm, Gerhard Louis and Franz A. Pohl 424
Chemical Preparation of p-Mn02 and Comparison of its Electrochemical Properties with those of 3- and Y-Mn02 Marc Beley and Jean Brenet 426 XVII
Elaboration of a New Type of Oxide Electrode and Examination of Electrochemical Behaviour of Mixed Oxides of Transition Elements J. Ruch, J.F. Koenig and J. Brenet 429
Lithium/Sulfur Secondary Cells E.J. Cairns, H. Shimotake, E.C. Gay and J.R. Selman 432
High Energy and High Rate Lithium Organic Electrolyte Batteries M. Eisenberg 435 Spatial Current Distribution in Non-Isotropic Porous Electrodes Karl-Joachim Euler 436
The Oxidation of Pb in H^O^ G. Archdale and J.A. Harrison 437
Inorganic Electrolyte Based, Room Temperature, Lithium/Chlorine, Sodium/Chlorine, Lithium/Sulfur, Sodium/Sulfur and Lithium/Cupric Fluoride Cells Adam Heller, Kenneth W. French and James J. Auborn 438
A New Rechargeable Mercury/Mercury Oxide-Electrode Margarete Jung 441
The Oxygen Electrode Made of Ag2C03/Zn/PTFE Mixture M. Cenek, 0. Kouril, M. Caläbek, L. Komärek, J. Sandera and J. Vanåcek 444
Reaction Layers and Electrochemical Properties in Tubular Electrodes due to Engineering Design Alfons Lindholm 447
Rechargeable Metal-Air Battery Systems Lars Carlsson, Göte Granath, Gunnar Lindström, Olle Lindström, Leif Sköld, Lars Welin and Ingvar Åkerblom 450
Electrolyte Distribution and Shift through Carbon Electrodes for the Electroreduction of Air Oxygen Jifl Mrha, Jiff Jindra and Miroslava Musilovä 453
On the Structure of Catalysts and Electrodes for Electrochemical Conversion of Gaseous Reactants Konrad Mund 454 XVIII
Electrodes for Iron-Air-Secondary Cells M.W. Nippe, H. Cnobloch, D. Gröppel, G. Siemsen and F. v. Sturm 457 Selfdischarge of Iron Electrodes with Hydrogen Evolution in Alkaline Solutions Lars Öjefors 460
Ion Selectivity and Diffusion Potentials in Corrosion Layers. - Pb'SO, Films on Pb in H-SO^ Paul Ruetschi 463
Continued Development of Hybrid Power Cells for Cardiac Pacemakers Alvin Salkind, Allen Hahn, John Cassel and Victor Satinsky 466
Influences of Foreign Metal Ions on the Performances of Nickel Positive Electrodes Kuranobu Sugita and Siro Ohkuma 467
Design of High-Temperature Solid-Electrolyte Fuel-Cell Batteries for Maximum Power Output per Unit Volume E.F. Sverdrup, C.J. Warde and R.L. Ehack 470 SYMPOSIUM on
ELECTROCHEMICAL ENGINEERING
TEE 114th EVENT OF THE EUROPEAN FEDERATION OF CHEMICAL ENGINEERING Norbert Ibl 3
OPTIMIZATION OF ELECTROLYTIC PROCESSES N. Ibl Technisch-chemisches Laboratorium, Swiss Federal Institute of Techno logy5 Universitätsstr. 6, 8006 Zurich, Switzerland. Economic constraints are a characteristic feature of electro chemical engineering. Optimization plays an essential role. An electrochemical system can be optimized from various view points. One can, for instance, optimize the distance between the electrodes or the electric circuit (length and cross section of the bus bars, optimum use of rectifiers). In many cases the optimization presents itself as a balance bet ween energy and investment cost. This applies to the aforementioned optimization of the bus bars. It is also the case in the very central and important problem of the optimization of the current density i in an electrolytic process. If one increases the current density the total electrode area needed and therefore the investment decrease but the energy consumption increases. The optimum current density is that at which the sum of the investment and energy cost is a minimum. Let M be the amount of product to be made in time t and Q the quantity of electricity needed according to Faraday's law. The energy cost K e is K = bEAit = bEQ (1) where b is the price of the unit of electric energy (#/kwh), E the applied voltage and A the total electrode area needed. The invest ment costs can be split into the fixed costs Kf (which are essentially independent of i, such as the cost of most of the control instruments) and variable costs K^ which depend on i. The latter include mainly the cost of the cells and of the electrodes. Optimization is carried out by deriving the total cost (K = K + K^ + K^) with respect to current density i and equating the derivative to zero
dK ^ Q = i + d(bEQ) di di di In order to complete the calculation K- and E must be expressed as a function of i. As a first approximation we can use linear relation ships for K. and E
Ki = aAt = -^ (3) E = E° + Ri (4) where E° is a constant, R the resistance of the cell (Ohm.m ) and a the specific investment cost (per unit of electrode area and .per unit time). The value of a takes into account both the interest on and the amortization of the invested capital. An average has to be taken because of the progressive amortization of the initial investment. In spite of Tafel's law the linear relationship for E is very well full- filled in industrial chlorine electrolysis, for example. If Q can be assumed to be independent of i (constant current yield) a combination of eqs. (2) (3) and (4) leads to simple relation- Norbert Ibl 4
ships for the optimum current density i and the minimum total va riable cost K l = (5) K = 2Q 6 oV AIS oP ^^ < > Under optimum operating conditions the energy and the variable in vestment cost are equal. The equations presented will be illustrated by numerical examples. In the case of a battery, where the product is electric energy rather than a substance, the calculation of the optimum conditions is quite similar but it may be more convenient to consider the voltage, instead of the current, as the independent variable. In principle, the operating cost can be lowered by decreasing the energy cost through the application of a variable i to make use of surplus night electricity. We have computed the optimum ratio of day to night current for the two cases a) that one optimizes the power plant together with the electrochemical one b) that the night and day electricity is sold at a fixed rate. The optimum current calculated by the above procedure is in a sense an ideal value, which one should attempt to realize or at least to approach although this may not always be possible. A very general restraint is the limitation imposed by the mass transfer. If the cal culated optimum current is above the limiting current for natural con vection conditions, it can be realized only by stirring. The cost of the stirring has then to be included into the calculation. The opti mization of the stirring involves two aspects a) calculation of the optimum flow rate for a given hydrodynamic model b) comparison of various modes of stirring and selection of the most economic one. Numerical values will be shown for two ways of stirring: 1) pumping of the solution through the cell 2) bubbling of gas through the electrolyte. In principle, the optimization should be as complete as possible. It is then a complex problem involving the simultaneous consideration of various aspects and their interaction. In addition to the influence of the stirring the optimization of the current density has to be linked with other optimization calculations such as those mentioned at the beginning (for instance, with the determination of the optimum number of cells which results from a consideration of the rectifiers cost). Further, once the optimization has been carried out for a given system a comparison with other possibilities should be made. The final aim of the electrochemical engineer is to obtain a given product in the most economic way. This involves a consideration of alternative cell and electrode design or materials, of other electrolytes or even of a completely different route. For instance, one may envisage the use of porous, fixed or fluidized bed electrodes etc. This modifies in general both the values of a and R. According to the simplified equa tion (6), if one succeeds in working under optimum conditions, the pro duct of a and R is a decisive criterion for the evaluation and the com parison of the variable cost. R. B. MacMullin 5
DESIGN AM) MATERIALS OF CONSTRUCTION FOR ELECTROLYTIC CELLS Robert B. MacMullin R.B. MacMullin. Associates, 45 Falls St., Niagara Falls,N.Y. 14303,USA.
Introduction. The art of cell design has long preceded the science. Electrochemical engineering is the discipline of applying science to the art. The special problems of designing electrochemical, as dis tinguished from chemical, reactors will be described. Electrolytic processes may be classified by industrial application, as shown in Table 1. Cells may be categorized in various ways, according to func tion, as shown in Table 2. Every practical cell represents a partic ular combination of categories, and several hundred thousand different combinations are possible. Industrially important types are few in deed, and there are many opportunities for new, innovative designs. The role of cross fertilization of ideas between various segments of electrochemistry will be highlighted. Techniques of Design and Scale-up. General procedures will be review ed briefly. Some specific recent advances in techniques include: Porous electrodes (flow-by), Extended surface electrodes (fluidized bed), Forced circulation of electrolyte (as in EHD cells), Coping with anode effects, Current reversal and pulsing, Coping with polyanions that reduce efficiency (Br2 cells), Enhancement of heat transfer on gas-evolving electrodes (ECF cells), Conformal plotting as a means of estimating current distribution (U cells), Shock removal of metal from cathodes (Mn cells), Computer simulation in cell design (Chlor-alkali cells), Use of directive salts for controlling selectivity of electrode products (EHD cells), Electrowinning and electrodeposition of metals in magnetic fields. Materials of Construction. Designing cells around properties of avail able materials vs. designing materials for new modes of operation, will be discussed. Some new, or at least different, materials now available for cell components include: Anodes, - new forms of carbon and graph ite, magnetite, lead dioxide, ruthenium dioxide, composites, and Pt plated on Ti, Ta substrates. Cathodes, - specific metals to control selectivity. Membranes, - new types of ion exchange. Diaphragms, - permeable. Solid Electrolytes, -f3 and 0 " alumina for Na+, doped Zr02 ?or (F, RbAg4.l5 for Ag+. Baths, - Hydrotropes, non-aqueous sol vents, directive salts, buffers, molten salts. Flexible seals and covers, - various corrosion resisting elastomers. Tanks, - new plas tics and FRP, modified concrete, Ti and Ta, cladding for metals and concrete. Insulators, - new forms, such as BN. Bus, - Al, Ag, vs. Cu. Some Notable Achievements in industrial electrolysis, which illustrate new approaches to design principles and materials of construction will be briefly described: For chlor-alkali, new types of mercury and dia phragm cells, featuring metal anodes and other improvements. For chlorate and perchlorate, mono and bipolar cells. For sodium hypo chlorite, in-situ sterilization of water. For adiponitrile, EHD cells, some without diaphragms. For fluorocarbons, a new type of ECF cell. For H2, high temperature electrolysis of IfeO vapor, using a solid electrolyte. For uranium metal, pure molten TJ by reduction of U02 in a Hall type cell. R. B. MacMullin
Table 1. Some Industrial Applications of Electrolysis
Category Examples 1. Inorganic chemicals, preparative Chlor-alkali; chlorate 2. Inorganic chemicals, purification NaOH 3. Metals, electrowinning Al, Mg, Na, Ma, Ni 4. Metals, electrorefining Cu, Pb, Zn 5. Metals, electroplating Ag, Cu, M, Cr 6. Metals, electroforming Steel, alloys 7. Metals, electromachining Steel, alloys 8. Metals, electropolishing Steel, alloys 9. Metals, anodizing Al, Ti 10. Organic chemicals, preparative Adiponitrile; p-aminophenol; Fluorocarbons, TEL 11. Electrodialysis Desalting/Conc'n. saline water 12. Batteries Primary, secondary 13. Fuel Cells H2/O2, Na/S 14. Films, electrophoretic deposition Rubber latex
Table 2. Cells Classified by Various Functions
By Mode Example Charge Electrowinning cells, power consumed- Discharge Prim. Bat., fuel cells, power generated Reversible Sec. Batteries, power stored
By Polarity Example Monopolar assemblies Hooker CI Cell Bipolar assemblies Erco Chlorate Cell Combination assemblies Dow CI Cell
By Type of Anode Cell Type Anode Consumable Al pot Carbon Soluble, solid Ni cell Nickel matte Soluble, lig.. Al ref. Molten Al Insoluble, solid Cl/Alk. cells Graphite, DSA Insoluble, porous Phillips ECF Carbon Insoluble, packed particles Nalco TEL Pb Insoluble, fluidized particles Newcastle, CN removal Metal, graphite
By Type of Cathode Cell Type Cathode Rigid Metal Metal win., refin. Start, metal/plated metal Rigid Metal Cl/Alk. (dia.) Steel mesh " porous solid Fuel, H2/O2 C, graph., metal Packed particles Cl/Alk., Hg. Graphite in decomposer Fluidized particles Newcastle Metal, graphite Liquid Metal Cl/Alk. (Hg) Amalgam Liquid Metal Chlor" metal (Na) Molten Pb/Na Liquid Metal Ford Na/S bat. Liquid Na R. B. MacMullin 7
Table 2. (Continued)
By Type of Electrolyte Examples Aqueous solutions Brine Non aqueous organics THF,DMS0,Et20 plus support, electrolyte Non aqueous inorganic KH2F3 Homogenous combinations Monsanto EHD, H2O, QAS, AN/ADN Heterogenous emulsions Phillips EHD,H20,bufrer salt, DS,AN/ADN Fused salts NaCl, CaCl2, MgCl2, LiCl/KCl Solid electrolytes,(liq./liq. )Ford Na/S, beta AI2O3 Solid electrolytes,(gas/gas) G.E. H2/O9, stab. Zr02
By Type of Divider Type of Cell Divider
Undivided cells Chlorate cells None Partly divided Curtain Walls MEL Mg Ceramic Baffles, screens Downs Na St. steel mesh Isolated by fluid metal Cl/Alk., Mercury Chlormetals Na Molten Pb Permeable diaphragms Rigid, porous Wet batteries, fuel cells Ceramic, metal Flexible, porous Cl/Alk.; Mn Asbestos; textiles Ion selective membranes Organic Electrodialysis,desalting Anionic, cationic Electro hydrodimerization Cationic Solid electrolyte Ford Na/S battery Beta alumina, Na+ Solid electrolyte G.E. H2 generator Stabilized Zr02, Cr
By Hierarchy Cell Type Products Cell products = electrode products Chlor-alkali Clp , NaOH, H2 Cell products formed by secondary Hypo cells Na 0C1 chem. react, away from By Change of State at Elec trode Examples At Anode At Cathode No change, prods, sol. in bath Oxidation Reduction Gas is evolved 0%, CI2 H2 New liquid phase forms Heavier than bath Al, Mg Lighter than bath S (Ford bat.) Na, Mg Solid phase forms Adherent deposit PbOg Mn Non adherent deposit MnC»2 MS(0H)2 1 rt 0} I E»-^• S-,— 1 •> cu C t— i— 0) 4- a> i ro »>i— o CD >> >j CD 00 E "O S- 1— s_ 00 CU 00 = >j S- CU O cu JZ o ra = a> ra O "O -E •i— i— r— DO O -E 1 CU 00 i— 4-> E i- r— E CU O E O ra o OJ -P OJ OJ J2 O P— s- $- 4* 01 CU ~0 •i— i. OH •r- 3 E 4-J 00 S_ E CU E •r- 03 i— »i— cj »i— ra 4— E 00 >> O CU JD >> P 3 i— CU to E a +-> O +J 2 4- ra +J cu ra o o • 1— 4-> -a ra 4-> ra 4J o oo r— c s- cu ra S- a +J 4- cu ~a oo =3 cu a 4-> E JZ 00 4- 00 »r- CJ E 3 O 4- 3 CU 00 •r~ O) CU O CU 4- S- 4-> •!- o a; o O -r- CU a. • CU 00 ra cu o cu o 00 E -P E 3 3 10 3 i- r— Oil— 4-> 4-> U S- J- = CU >i>— 4-> ra 4- > O CU to CU r— S- • Q. E J= 4-S cu cu ra •!- E T3 o 3 -a 3 o s- o a >)0+> CD >>r- CU 03 aj 4- 00 O O O 00 E cu cu cu •i— O 4-> " 4J O r— U cu E J3 E oo oo ra o ra p— t-o1»- ra o s- S- 3 CU S- S- CU CU T- -a r— U -i- CU S- O J2 •!- J^ o o oo JZ P t- o O- OJ 4- 4- p- S- oo o JZ E JZ s- a. CLJZ JZ i— E E ra •r- 4-> 4- i— "O ra i_ O r— CU ra o CU « O -r- O 4- S- i p- O 4-> UJ JD P ra O O r— P +J Q. CU 3 o •P rfl 0(0 C Q. O. U- O. U 4-J JZ X O "O 4-> p- CU TJ CJ TJ oo 03 •r— E ra *p~ *p™ •i— > oo s_ CD o ra aac 4-> 3 4-> T3 "O CU -a o i— CU >> OJ E oo _! 3 o > o cu 4- oo u •i— ra T- 4-> oo cu ra cu O) i— ra o i— cu JZ • E OJ CU aS-r- rfl LU n S- (flPh O aj w DO >>-C D-4-3 E E JD i— r— JZ CU J= oo 4- s_ ra E -P 00 ra p— o cu E >> ra cu CJ >> T3 E o.^: = S_ •!- r— u 00 CU C O i— ro i— JD 4-> E -a P 4-> •(- cu O) JZ O 4-> "O 4-J CD CU C7) O- CU +•> *—' E •r- TD CU E CU E T- -r- OJ CU O CU JO 00 !—+>•!- E E o cu -i- ra -o cu #t OL r— E •r- ra -Q i/) o 3 ra +J cu ra > D_ 4-> T3 O O O E o +J JZ ••- cu E E cu to 4- ra 4-> t. 3 cu cu • D_ +J oo 4-> 4- i— O "O cu >> o cj4- ra cu ojz-p- oo J^ • cu cu r- E (OP CU E JE 3 CU 'i- i— DO CU 3 Q. E ra >» •r~ CU JZ CJ •!- o 4-> •p-OS-JZJZ!4->ECU O s- In the study of convective transport processes at electrodes two techniques have proven to be particularly useful: limiting current measurements, and optical observations of the boundary layer (inter- ferometry). The two techniques are complementary: limiting currents give only indirect information about the nature of the boundary layer, while interferograms allow direct visualization of the concentration field below the limiting current. The latter technique however is not suitable for the study of three dimensional flows, or when the electrolyte contains more than two ionic species. Experimental techniques as well as those of interpretation have become progressively more sophisticated. In the past 25-30 years limiting rates and their distribution have been investigated for a variety of cell geometries and flow regimes. In most cases the results are presented in some generalized form; equations resulting from the solution of the convective diffusion equation--or as dimensionless correlations. Further progress in this area should be encouraged in view of the fact that rates (current densities) in future cell processes will have to be significantly increased to reduce capital and labor charges. Transport limitations are particularly evident in process applications where the reacting species are present at very low concentration levels (e.g., removal of trace impurities from plant effluents). While transport can be enhanced by increasing flow races or the rate of movement of the electrode surface, practical considerations (excessive pressure drops and/or power requirements) may require that other means be employed to improve the supply of reactants or the removal of products. This is especially evident if one considers the need to keep interelectrode distances small to minimize ohmic losses. Recent interest in fluidized bed electrodes, through-flow (porous) electrodes, slurry electrolytes, turbulence inducers, and wipers show awareness of the need to find alternates to the simple forms of convective flow between parallel plates. Very high current density processes have been commercially introduced for the anodic shaping of metals; interelectrode distances in these are only fractions of milli meters. The study of transport of reaction products (in solution, as well as solids and gases)in these small gaps provides valuable insight relevant to high current density processes in general. Problems in primary and secondary current distribution in two- dimensional geometries no longer present a serious challenge;these can be solved to desired levels of accuracy by computer implemented numerical techniques. Most recently tertiary distribution' problems (i.e.,the distribution as influenced by both current density dependent interfacial potential difference and mass transport) have been successfully treated for the rotating disk and for planar electrodes in a channel. Current distribution and mass transport in three dimensional electrodes (porous-, fluidized bed) have been analyzed for a number of simplified models. The dilemma posed by the complex changing geometry 10 Charles W. Tobias and variable local composition in battery plates presents an as yet unsurmounted barrier to prediction of battery performance or the "state of charge." This problem area is intertwined with broader questions related to the need to greatly improve battery design concepts for urgently needed practical applications. One of the outstanding problems we must fact in cell process design is the dilemma of nonuniform accessibility of the electrode surface. The traditional simple cell geometries and flow arrangements do not assure uniform accessibility, and hence the current density, or electrode potential, varies over the electrode surface. While in some processes this may be tolerable, in others, particularly in organic synthesis processes,the electrode potential needs to be maintained within rather narrow limits. As we move toward increasing current densities, the adverse effects which result from nonuniform accessibility (loss of specificity, on metal deposit, etc.) tend to be accentuated. The design of electrochemical cells as we know them has been profoundly influenced by the transport properties of electrolytes and by the nature of conducting- and dielectric-materials available for hardware purposes. Design and operational concepts were shaped primarily by the low concentration and low mobility of charge carriers in aqueous electrolytes. Changing cost structures in the area of initial investment, depreciation, labor, and power present a new challenge to the electrochemical engineer. Methods available today for the development, design, and control of new electrochemical processes, and for the improvement of existing ones, are comparable in precision and convenience to those employed in ordinary chemical processing. The chemical and physical properties as well as the state of aggregation of materials available for cell construction(electrodes, diaphragms ..containers)allow a much broader range of design conceptions than reflected in current industrial practice. New processes appear economically attractive in the area of electroforming, machining, organic synthesis, and separations (membrane dyalisis, etc.). Use of organic solvent media may permit the realization of processes requiring higher electrode potentials than those feasible in aqueous media. In the area of electrochemical energy conversion and storage the demand for much higher performance levels and greater reliability must be met by imaginative hardware designs and by the utilization of new, high energy density, reactions. It is evident from reliable estimates of the availability of raw material and energy resources that electrochemical processing techniques as well as energy storage and conversion should receive increasing emphasis. Electrochemical engineering should be an exciting and rewarding field in the years to come. R.T. Atanasoski 11 DIMENSIONALITY STABLE ANODES IN CHLORATE ELECTROLYSIS R.T.Atanasoski/ A.O.Filip, B.Z.Nikolid, M.M.Jaksid and A.R.Despic" Institute for Chemistry,Technology and Metallurgy,Beograd and Faculty of Technology and Metallurgy,Beograd,Yugoslavia The use of dimensionally stable anodes (DSA) in the elec trolytic chlorate production has been investigated. They were all titanium based and then activated by platinum, iridium or ruthenium oxide. Activation with platinum and iridium has been done by the usually employed thermal de composition treatment, while those with ruthenium oxide have been obtained as such from the firm Oronzio De-Nora (Italy). Simultaneous measurements have been made of the amount of noble metal at the surface and of electrode activity, as functions of time of electrolysis at constant current and temperature. Amount of noble metal have been determined by a radio tracer technique. Electrodes have either been prepared with radioactive metal, or subsequently activated in a pile. The change in activity of longliving isotopes Ir 192 and Ru 103 with the time of electrolysis has been followed. The amount of the Pt-Ir (40%) alloy has been varied between 0.04 mg/cm2 and 0.9 mg/cm2 of the electrode surface. The measurements have shown that the rate of loosing the noble metal under constant conditions of electrolysis (300 mA/cm2, 60 C) has been lower at electrodes with smaller amounts of the alloy. Percentagewise, however,this loss is smaller at electrodes with larger amounts of the alloy, and hence, the working life of the latter is longer. It was found that considerable amount of the alloy is lost within the first fev/ weeks of electrolysis (10-20%). Later on the rate of loss becomes considerably lower and within several month of investigation all electrodes have main tained sufficient amount of noble metal to have reasonable electrochemical activity. Polarisation characteristics of all the electrodes with Pt-Ir alloy have been reasonable and compare favourably with the ones activated with ruthenium oxide. 12 Fritz Beck CELLS FOR ELECTRO-ORGANIC SYNTHESIS Fritz Beck Main Laboratory of the Badische Anilin- und Soda-Fabrik AG Ludwigshafen/Rh., F.R. Germany. An important precondition for the sucessful operation of an organic electrosynthesis on an industrial scale is the construction of electrochemical cells which are best adapted to the peculiarities of the organic systems. Two main problems exist .which can be handled by methods of cell design: 1. The specific conductivity of the reaction mixture is usually 10-1000 times lower than in the case of inorga nic systems. In order to operate the cells with high cur rent densities ( 5-100 A/dm2) at low voltages, the di stance of the electrodes must be minimized to values of a few tenths of a millimeter (1). The electrodes can be li quid permeable(Fig.1) or nonpermeable. In the latter case, the reaction mixture moves in a capillary flow between the electrodes. One of the designs of particular technical interest is the capillary gap cell with a pile of circu lar plates acting as bipolar electrodes (Fig. 2). The per formance of this cell is shown with some examples ( ca- thodic hydrodimerization of acrylonitrile, anodic Kolbe electrolysis of monomethyladipate ). The space-time-yield is rather high and can exeedlO h"^-. The concentration of the electrolyte can be lowered substantially, which re sults in a great simplification of working up conditions. Fritz Beck 13 Scale-up is achived according to the bipolar principle. In addition,cells with ion exchange membranes can be con structed. In this case the minimum distance between the electrodes is effected by contacting the liquid-permeabLe electrodes with the diaphragm (Fig. 3). organic 2. Side-reactions can lead to the formation of polymericV layers on the electrodes. Normally this electropolymeri- zation is of the ionic type and cannot therefore be in hibited by neutral molecules. Å continuous electrolysis is then impossible. To avoid this effect, the parameters of the electrolysis must be modified by periodic inter- uptions of currents;by changing the polarity of the elec trodes or by addition of special solvents. In severe ca ses, the electrodes must be cleaned continuously or periodically. Condit (2) has described the continuous washing of mercury cathodes in a by-pass. However, even in the case of solid electrodes, it is possible to re cover the primary activity by continuous exchange pro cesses, if the cells are designed as a collective of particles ( slurry electrodes of Gerischer, fluidized bed cell or packed bed cellf-Fleischmann and Goodridge). The swelling of polymeric materials, especially dia phragms, as the result of the organic components in the reaction mixture, can cause some trouble. This problem can be overcome by using polyolefins or perfluorinated materials. (1) F. Beck, H. Guthke, Chem. -Ing.-Techn. k\_, 943 (1969) (2) P.C. Condit, Ind.Engng.Chem. H8, 1?5? (1956) 14 Giuseppe Bianchi BASIC PROBLEMS OF THE ELECTROCHEMI^\ PRODUCTION OF CHLORI NE Giuseppe Bianchi and Giuseppe Faite Laboratory of Electrochemistry and Metallurgy of the Uni versity, University of Milan, Milan,ITALY. In the Laboratory of Electrochemistry and Metallurgy of the University of Milan a research program has been develjo ped with the aim of getting basic thermodynamic and kine tic data concerning .the electrolytic processes in HC1 and NaCl solutions: the results are reported in the present pji per. Thermodynamic data. Especially built hydrogen and chlorine reversible electro des have allowed the standard potential for the couple Clo/ ~ \/Cl~/ \, the equilibrium constants for the reac- 2 (gas) (acl) t i o n s C12(gas)+C1(aq) ^ C13(aq) C12(aq)+C1(aq) ~ C13(aq) and the mean molal ionic coefficients of aqueous HC1 in a wide range of temperatures to be obtained. On the whole these data enable one to calculate the reversible cell vo2 tage for the electrolysis of HC1 solutions as a function of the temperature and HC1 concentration.. E.m.f. data from the galvanic cell: -Pt/Na(Hg)/NaCl(aq)/AgCl/Ag/Pt+ have led to the standard potential for the Na /Na(Hg)couple to the activity coefficients of the sodium metal in the amalgams, and to the mean molal ionic coefficients of the NaCl solutions. The knowledge of all these data makes it possible to calcu late the reversible voltage corresponding to the inlet and outlet conditions of a mercury cathode cell. By taking into consideration the experimental results for the cathodic discharge of the Ca ions on Hg it is possi ble to demonstrate that for diluted calcium amalgams (2ppm) the discharge potential for the Ca++ ions is quite near the one for the discharge of the Na ions at the cell outlet, where both concentrated sodium amalgams and diluted NaCl solutions are present. Hence, it can be concluded that any factor leading to an excessive depletion of the ij>rine will cause the simultane- ous discharge of Na and Ca Kinetic data. The anodic discharge of the chlorides has been studied on titanium electrodes covered by a thin layer of electrocat^a lyti c materi als. "~ Giuseppe Bianchi 15 As it is well known this type of anode has allowed in the last few years some important advances in the tecnology of the chlorine diaphragm and mercury cathode cells. The results which are presented in this paper have been ob tained with two different kinds of coatings, that is plati num group metals coatings (Pt,Ir and Pt-Ir alloys) and oxi de coatings (mixed Ir and Ru dioxides), i) Pt,Ir and Pt-Ir alloys coatings. As regards the passivation phenomena it has been observed that: - the electrode potential for the chlorine discharge on Pt anodes increases quickly as a limiting value for the cur rent density is exceeded (passivation); - the limiting value for the current density is a function of pH and NaCl concentration; - with respect to the reversible potential of the Cl?, ./ ClT^oN couple the passivation of the Pt electrodes '9a s / begins at 0.1 volts, and is completed at 0.5 volts, more anodic; - the rate of the Pt passivation is also a function of the hydrodynamic conditions at the electrode surface (mass transport to and from the electrode surface); this fact has been tentatively ascribed to ClD""ions which form in the bulk solution as a consequence of the molecular chlo rine hydrolysis; - the Ir coatings do not passivate; the Pt-Ir alloys coat ings show a quite similar behaviour provided that the Ir content is at least 4% b.w.: a possible explanation resi des in the different properties of the oxygen layer adso_r bed on the Pt and Ir surfaces. As regards the kinetics of the chlorine evolution it has been concluded that the most probable reaction mechanism is the following one: 2C1 ad.s -*- Cl20 (rat* e determininga stepr )y both for Ir and Pt-Ir alloys and for non-passive platinum, ii) Mixed Iridium and Ruthenium oxide electrodes. This type of el ectrocatalyti c material is particularly int^e resting in view of the fact that a number of metal oxides' are characterized by both high electric conductivities (the temperature coefficients being negative) and remarka ble chemical inertia. Typical examples are TiO, Ir0«» RuOp etc. The data which are reported in the present paper concern the mixed Ir and Ru dioxides coatings and they show that: - the electrodes are characterized by a highly reprodubible behaviour also after prolonged electrolysis times (e.g. 15,000 hours) at high current densities; 16 Giuseppe Bianchi the tested electrodes do not passivate also at the high- est current densities and their behaviour is quite simi- lar to that of the Ir and Pt-1r coatings; the low slopes of the electrode potential vs. log of cur rent straight lines (e.g. 30-40 mV/decade of current) show that an electrochemical-chemical mechanism as for Ir, Pt-Ir alloys and non-passive Pt should be followed the chemical desorption step being rate-determining. Michel DELMOTTE 17 EXPERIMENTAL REALISATION OF A CLAMPED IONIC CONCENTRATION GRADIENT Michel DELMOTTE and Jacques CHANU Laboratoire de Thermodynamique des Milieux Ioniques et Biologiques Université PARIS VII, 2, Place Jussieu PARIS 5eme - Prance These last five years, the Thermodynamics of Irreversible Processes have been used in the analysis of a number of biological phenomena and have made the theoretical justification of the evolution of living systems possible particularly when convenient parameters of enzymatic cinetics are used. The existence of instable non-equilibrium steady states of open systems (lto3), mainly made up of ATPases of nervous cell membranes, in particular ionic environment, has been proved and this conclusion has incited us to elaborate a diffusion cell with a concentration ionic gradient kept constant and which is very similar to the Spiegler's transport cell. The principles of the construction of this apparatus will be presented before this technical description. Lastly, the first results obtained with this prototype will be exanhai CONSTRUCTION PRINCIPLES The purpose of th>* experimental realisation is to establish and clamp an ionic concent., -a .. ion gradient between the two surfaces of an artici- cial membrane 01' cellulose acetate. This membrane separates two homo geneous compartments which contain bulk ionic solutionsand is the support of nervous cell ATPases. First, in order to study the properties of this proteins in non- equilibrium ionic situation, the membrane is laid out horizontally with the most concentrated solution below and the less concentrated above so as to make possible the realisation of a pressure equilibrium. The osmotic pressure can be balanced by an artificial mechanic pressu re; the volume variation of the inferior compartment solution starts a feedback mechanism which modifies the artificial pressure. Secondly, the two different concentrations of ionic solutions are kept constant by only inlet or outlet of matter, i.e. by dilution for the half cell with lower concentration solution-inlet of pure water -and by superconcentration for the half cell with higher solution- inlet 18 Michel DELMOTTE of"superconcentrated solution"-. In the superior compartment the inlet of pure water is balanced by equivalent outlet of inner solution by simple owerflowing. For the inferior compartment, supply and withdrawal of matter are equilibrated in such a way that the total volume of the solution is constant because the existence of the pressure feedback mechanism-. In this condition, the volume v. of "superconcentrated so lution" which was added in by time unit is always different from the volume v of solution which the cell gives*up during the same time u- nit. Indeed, it can be written, acting partial molar volume V and V 0 • s w v. = n V ,. v + n V . v i s s(i) w w(i , ) One part of the n moles of salt crosses the membrane till the other 1 part n s with the n w moles of water suits to the inferior compartment solution and must be withdrawn out of cell v o = n' sV s(o, )s + n wV w(o, )v The ratio _i_ may be calculated; it is only function of the parameters of "superconcentrated solution" and inner inferior solution n _ __ n s(i) w(i) v.i _ w v n' __ ° §_ XT + Y n s(o) w(o) w In the technical description, we see how this condition may be reali zed in the facts. The last principles of elaooration of this apparatus are the particular importance of the mechanic homogeneisation of the bulk solutions and the isothermal maintenance of the two compartments. Finally, we shall see that the diffusion fluxes can be measured directly by "superconcen trated solution" consumption. 19 Michel DELMOTTE TECHNICAL DESCRIPTION -The cell- Each half of cell is a 200cm3 cylinder surmounted by a hemi sphere, hollowed in a single "Altuglas" block which supports all acces sory of measure and regulation. The inner diameter is 70mm and the outer l80mm.The membrane is a 90 cr 110mm diameter disk which is 0,03mm thick. -The concentration feedback mechanism- In each compartment, a condivity cell, associated to a Kohlraush bridge, allows us to measure the re sistivity of the solution. The unbalance signal is amplified and starts the special electrical alimentation of dosimetric pump. With the supe rior compartment a peristaltic dosimetric micropump is associated. With the inferior compartment a two sucker-heads pump is associated, which pushes the "superconcentrated solution" and pulls the inner inferior solution; the two sucker-heads of the pump are adjusted differently. The two feedback mechanisms give propertional answers, which allow a particular regulation. -The homogeneisation- Two magnetic stirrers similar to a spinning-top, are set in special cavity; the speed is adjustable up to lOOOrpm. -The temperature regulation system- In each half of cell a thermistor of about 130.000ohms at 20°C is associated to a Wheatstone bridge. The unbalance signal is also amplified and starts circulation pumps. Water, regulated at 19,30°C + 0j01oCa is led through a coil set in the inner surface of the cell. The feedback mechanism gives proportional effects In each compartment, it allows a regulation o.f + 0,02°C at 20°C-. THE FIRST RESULTS The first experimentations have been realized with NaCl. The lower con centration is clamped at about 0,1 M/l and the higher concentration at different values between 0,15 and 0,35 M/l. The steady states are par ticularly evident and the system proves to be very steady. Control ti trations by Mohr's method justify the results of this experimentation. In the study area, the Pich's laws are satisfied. 1-R. BLUMENTHAL. J. P. CHANGEUX> R. LEFEVER, J.Membrane Biol.2 351.374 (1970) 2-M.DELMOTTE, Vision Res.l0_ 671.678 (1970) 3-J. JULIEN; These Doc. 3eme Cycle PARIS 1971 20 Aleksandar Despid PULSATING POTENTIAL ELECTROLYSIS IN DEPOSITION OF COPPER Aleksandar Despic and Gordana Savic Maglid The Faculty of Technology and Metallurgy,University of Baograd and Electrochemistry Department of the Institute for Chemistry,Technology and Metallurgy,Beograd Theory of the effect of pulsating potential on the morpho logy of metal deposits has been developed recently(1). It has shown that in the case of potentiostatic pulsation,with square wave pulses of equal duration as the off-periods,and in the case of reversible systems,two diffusion layers develop. One is that corresponding to the instantaneous flux determined by the current produced by the potential pulse and the second one is that due to the net average cur rent of metal deposition. The second one develops in a normal way but with the lowest concentration at the inner boundary of 1/2 of that in the bulk. Hence, no total concentration polarisation can be attained with respect to this net flux and its limiting value is one-half of the limiting current density. The other, "effective" diffusion layer does not extend into solution above a certain thick ness determined by the frequency of pulsation.By increasing frequency this layer decreases in thickness and at suf ficiently high frequencies it can be made to follow very closely the microprofile of the surface. It is the high concentration polarisation in this layer which is respon sible for the amplification of surface roughness and eventual dendrite formation. The amplification is connected with the development of the diffusion layer and its divorce from the surface micro-profile. When by increasing puls ation frequency, this is not allowed, smooth deposits should be obtainable even at current densities in the limiting current region and causing substantial, or total concentration polarisation. This result may be of considerable practical interest and research has been undertaken to check it experimentally as well as to establish other important characteristic, of deposition process under such conditions (current efficiency energy consumption,etc.). Copper deposition from an electro lyte similar to this used in copper refining has been selected as the test system. The cell consisted of two electrodes only, the anode being much larger in size and sufficiently close to the cathode that pulsating cell volt age could be taken to represent the pulsating cathode potential. Deposition has been carried out over longer periods of time, by DC and by pulsating potential of different frequences. Alelrsandar I>esr?i6 21 It has been shown that indeed the maximum current density of deposition obtainable with pulsating potentials of high negative values is approximately 1/2 of the limiting cur rent of DC deposition. The deposit has been observed under the microscope,its roughness measured by talysur£instrument and the current efficiency has been determined by comparing the recorded overall DC current with that calculated from the results of anodic coulometric stripping of the deposit. Encouraging results are to be reported. References (1) A.R.Despic*,K.I.Popov, J.Appl.Electrochem. 1 (1971) 275-278 Chaim Forgacs 22 ELECTEDNEUTRALITY AND STEADY STATE DURING THE TRANSFER OF ELECTRIC CURRENT THROUGH GRACED ELECTROLYTE SOLUTIONS Chaim Forgacs The Negev Institute for Arid Zone Research, Beer Sheva, Israel, and Sea Water Conversion Laboratory, University of California, Berkeley, Ca., U.S.A. The rate of practical electrochemical transport processes (electrode processes, electrodialysis, etc.), is mostly controlled by the transport through a thin layer of graded electrolyte solution (boundary layer or diffusion layer). The description of the phenomena occurring in this domain are often based upon two assumptions: 1. Electroneutrality is obeyed. : 2. After a relatively short transient time, steady state is established. As the use of these assumptions could not offer a proper quantitative account of some phenomena in which we are interested (namely, con- ceatration polarization at ion-exchange membrane - electrolyte solur tion interface during current transfer), it was decided to critically : evaluate their range of validity. ' At the first stage of our research a relatively simple system was : treated: a graded layer of a 1:1 electrolyte (KC1) bounded by well i mixed solutions of the same electrolyte in different concentrations. . It has been shown, that electroneutrality can not be strictly observed in this system during the transfer of constant current unless ; the concentration gradient is zero. No particular assumptions [ regarding the mechanism of the transport processes in the system are ' needed to prove this. If this statement is applied to a Nemst-Planck type transport model, ! it follows that under the same conditions steady state cannot be established either, unless, again, the concentration gradient is equal to zero. i As a result of these findings we set out to develop a transport theory in which these assumptions are not made in advance. .; A transport theory in general is a collection of fluxes, rates of creation and rates of dissipation of the various components embedded .in the mass balance of the system. Its mathematical form (for one dimensional transport) is a set of pirtial differential equations. 3V ~ i » • • • / JWn ' • ••?•••?••• , fiZ. » » • »I ... »Pj f • • • * 2^? • • **3i* • • / where pj are the system variables (concentrations etc.) x the co ordinate in the direction of the transport, and q^ the system parameters. In the cases of our interest the equations (1) are second order parabolic and x does not occur explicitly in the functions F^. The system of differential equations (1) is accom panied by a proper set of initial end boundary conditions in the domain 0 4 x ^ 1, 0 ^ t < *». Chaim Forgacs 23 If the analytical solution of (1) can be found in the form of a set of functions Pj(x, t), the system will reach steady state provided that a set of time independent functions P^(x) is obtained when t approaches infinity, i.e. (2) J?±(x) = limpiCxjt) If the analytical solution can not be found, the following approach may be helpful: 1. Due to the well known uniqueness theorem for virtually all the F|-s of our interest, a time independent system describing the functions Pj^ can be obtained by equating (1) with zero. 2. Equations F^ = 0 can be transformed to a system of first order, ordinary differential equations (3) y, = GL (...y. ,...,...qj.„.) The system (1) has a steady state solution if (3) has a real solu tion for the given set of boundary conditions. Now, in the case of the functions G, and the domains of our interest, such a solution does exist if the boundary conditions are in the form (4) yk (0) = R^ where R are real numbers. Thus, a steady state solution does exist if the given boundary con ditions are equivalent with a set of boundary conditions in the form (4), i.e. if all the system variables and their first derivatives are kept constant at one end of the domain. There still does not exist a generalized algorithm to check this condition. Examples for the solution of particular problems are available. If a steady state solution does exist, the experimental measurables still may fluctuate around their average. The reason for these fluctuations in the amplification of small random changes in the system parameters by the particular structure of the transport process. A useful tool to predict the extent of these fluctuations is obtained in the following way: Equations (3) can be linearized by assuming certain, experimentally feasible values for y.. Then they will take the form {5) *k ~Uik}**k 24 Cnaim Forgacs where the matrix {a.,} contains the above-mentioned experimental variables. Bearing in mind that the solution of (3) and the experimental measurables calculated from the solution are functions of the eigenvalues of this matrix, good semiquantitative measures for this amplification effect in the system are the condition values of the matrix (Ref. 1} i.e. the relative changes in the eigenvalues caused by changes in the magnitude of the matrix elements. The dependance of experimentally found fluctuations during current transport through membranes on system parameters (Ref. 2) seems to confirm this explanation. Acknowledgement: Part of this work has been accomplished in the Membrane Group of the Sea Water Conversion Laboratory, University of California, Berkeley, directed by Prof. K.S. Spiegler, to whom the author is indebted for helpful suggestions, and critical discussions. References: (1) Computer Library Program F2 CAL SPECS, Computer Center, University of California, Berkeley. (2) C. Forgacs. Nature 190 (1961), 339. Patrizio Gallone 25 NEW CONCEPTS IN THE SCALEUP OF CHLORINE-CAUSTIC MERCURY CELLS P. Gallone Faculty of Engineering, University of Genoa, Italy Mercury cathode cells of conventional design embody an "elongated" cathode, its length being a multiple of the width. For commercial cells installed heretofore said multiple is never less than 5 and for most models it is greater than 8. Such elongated configuration has been re tained also for models of highest current rating (more than 300 kA) and highest current density (more than 10 kA/m2) as allowed by activated titanium anodes, with out consideration of some fundamental concepts based on hydrodynamics. This traditionalism is not justified, as new requirements must be satisfied to obtain optimum mer cury flow also at the highest current capacities and cur rent densities that characterize the ultimate trend. If cathodic current efficiency is taken as 95$» mercury flow rate Q (1/min) is related to cell load I_ (kA; and so dium concentration £ (per cent) by the equation: Q = 1/10 C (I) p Between cathode area S (m ), width W_ (m), length L (m) and current density i_ (kS"/m2) the following relations hold; S = W L = I/i (II) Combination of the former expressions gives: Hh ioo-g— C-^-)2 (in) The thickness t_ of the amalgam layer depends on slope J (mm/m) and ratio Q/W according to the equation: t = K Jm(-$-)n (IV) ' where K,m,n may be taken as constants, as they approxim ately depend only on the roughness state of the underly ing surface (1). The attached graph represents this func tion for a typically smooth (new) steel bottom and a typically rough (aged) bottom. On the other hand, the ratio Q/W, divided by amalgam viscosity, corresponds to the "thickness Reynolds Number',' on which depends the degree of turbulence of the mercury flow (2). Accordingly, and irrespective of cell size and slope, it is possible to establish experimentally the maximum value of Q/W above which the turbulence would be come excessive. Then also the slope will be established 26 Patrizio Gallone through eq. (IV), so as to obtain the minimum thickness that, from experience, is required for complete bottom coverage, while keeping a minimum holdup of mercury with in the cell of given capacity 1 and current density i_: minimum mercury holdup = t. .S = t . T/. /,,v u ^ mm mm. I/i (V; A typical set of operating parameters, based on long- established experience, is the following: C = 0.25^; Q/W = 40 1 min"1 m"1; J = 15 mm/m By introducing them into eq. (I'll) one obtains the proper values for L and W. The following two examples show that the typical L/W ratio characterizing the oblong cells of relatively old design and small capacity is not applicable to new models sized for highest current rating and current density. Example I (old design) Example II (new design) I = 100 kA; i = 8 kA/m2 I = 500 kA; i = 15 kA/m2 S = I/i = 12,5 m2 S = I/i = 38,4 m2 L/W = 12.5 L/W = 1.54 s = L w = 12.5 vr S = L W = 1.54 W2 W = (12,5/12.5)2 == 1 m W = (38.Vl.54)2 = 5 m L = S/W = 12.5 m L = S/W = 7.68 m A higher L/W ratio than required by satisfactory hydro- dynamic conditions is undesirable also in that it aggrav ates the effects due to mercury butter accumulation over the cathode surface. Such outlook about a different configuration to be given to new cells of highest capacity seems to be some how anticipated in some recent patent applications, des pite the paucity of arguments underlying their priority claims (3,4). Ref fpf^nc* e s (1) G.I. Volkov and E.V. Mulin, Khimic. Promysl, No.5, , N 1959, p.38. (2) D.B. Spalding, "Convective Mass Transfer", McGraw-HJll 1963, pp. 174-175. (3) Ninon Soda Co., Jap. App. SHO 36-22309 (1961). (4) Kureha Kagaku, D.A.S. 1, 567, 975 (Jan. 14, 1971). Patrizio Gallone 2 MERCURY LAYER THICKNESS VERSUS FLOW RATE PER UNIT WIDTH FOR SEVERAL SLOPES Smooth bottom i i ——. Rough bottom , i i i i _i l I I i ! : *-u_ 15 20 25 30 35 40 45 50 Q/W ( l min"1 m1) 28 Francis Goodridge PERFORMANCE STUDIES OF A BIPOLAR PACKED BED CELL FOR THE PRODUCTION OF PROPYLENE OXIDE Francis Goodridge and Ola Osifade Department of Chemical Engineering, University of Newcastle upon Tyne, Newcastle upon Tyne, NE1 7RU, England. In the preceding abstract it was mentioned that steep voltage gradients in the electrolyte could lead to bipolar three-dimensional structures. In the present case the steep voltage gradient is achieved by the use of a relatively dilute electrolyte, the particulate phase consisting of a packed bed of electronically conducting beads^ . Provided the applied voltage is high enough each of these conducting beads will act as an anode at one end and a cathode at the other. Continuous electronic contact between beads is avoided by using a mixture of glass and conducting beads. The reaction investigated in the present work is the production of propylene oxide from propylene gas and aqueous sodium bromide. This system was chosen partly as a test reaction and partly because of the intrinsic commercial interest of the product. The basic reaction sequence can be described as follows. In the anodic area, bromine is the main product and equilibrates with water to form hypobromite and bromide ions: Br2 + H2° ^~ ^ HBr0 + Br~ + H+ The hypobromite reacts with the dissolved propylene to produce bromo- hydrin: CH - CH = CH2 + HOBr > CHgCHOH CHgBr In the cathodic area the predominant reaction is the evolution of hydrogen: 2H20 + 2e > H2 + 20H~ Due to the intimate mixing of anolyte and catholyte the bromohydrin reacts with the cathodically generated base to form propylene oxide, and bromide ion re-enters the process: CH_CHOH CH0Br + 0H~ > CH CH CHQ + Ho0 + Br" \ i Ignoring side reactions the net process is, therefore, the production of propylene oxide from propylene and water. Bromide rather than chloride ions are used since at concentrations < 10~ molar little hypochlorite would be produced Experimental The equipment consists of a rectangular cell filled with a mixture of Francis Goodridge 29 glass and graphite coated glass beads 500 am. in diameter. The sodium bromide electrolyte is saturated with propylene gas in a cylindrical countereurrent packed absorption tower, one meter high and 5 cm in diameter, filled with Raschig rings. The electrolyte then enters the cell and flows down the bipolar bed which has a cross- section of 16 cm and a height of 8 cm. Current feeders to the particulate bed are arranged parallel to the flow of electrolyte. The latter, containing propylene oxide, is then returned to a reservoir via a cooler. So far no attempt has been made to extract the product continuously. The electrolyte emerging from the cell is analysed for propylene oxide, hypobromite, bromide, bromate, propylene bromides and glycols. Result s Parameters which have to be optimised include (i) the ratio of conducting to non-conducting beads and their relative size, (ii) the potential gradient across the bed, (iii) flow rates of electrolyte, (iv) temperature of the reacting system. Performance is measured primarily in terms of two quantities. Firstly the energy yield (quantity of propylene oxide produced per kWh) and secondly the space time yield (the quantity of propylene oxide produced per hour per unit volume of cell). Results indicate that the energy yield (propylene oxide and hypo bromite) increases with increase in the ratio of conducting to non conducting beads. For a value of 1:4 an energy yields of about 0.04 kg (kWh)"" have been obtained. Assuming minimal ohmic losses calculated energy yields of 0.45 kg (kWh)" are obtained. Using 10 molar sodium bromide and again a ratio of I :4 conducting to non- conducting beads, space time yields of 0.13 kg s m~" are obtained. Further results will be presented at the meeting. References 1. See extended abstract by Goodridge, F. and Osifade, O.B„ 2. Fleischmann, M., Oldfield, J.W. and Tennakoon, C.L.K., Symposium on electrochemical engineering, Newcastle 1971, Trans.Instn.Chem. Engrs. (In press). 3. Selvig, A., Promotionsarbeit, Eidgenössische Hochschule, Zurich, Switzerland, 1962, p.41. 30 Francis Goodridge SOME PERFORMANCE CHARACTERISTICS OF THREE-DIMENSIONAL ELECTRODES * Martin Fleischmann and Francis Goodridge Department of Chemistry, University of Southampton, S09 5NH and Department of Chemical Engineering, University of Newcastle upon Tyne , Newcastle upon Tyne. NE1 7RU, England. In this paper we define three-dimensional electrodes as structures made up of discrete electronically conducting entities, immersed in electrolyte. Electrical contact is made by suitable electrodes, termed current feeders. Common forms of these structures consist of packed'1^ or fluidised beds^ '. An essential feature of these elec trodes lies in the non-linear potential distribution along the direction of current flow. The potential distribution in the electro lyte can be accompanied by a corresponding potential variation in the particulate phase due to the flow of electronic current. If the potential gradient in the solution is sufficiently large compared with that in a single, electronically conducting entity, the latter will exhibit bipolar behaviour. We can therefore classify these three- dimensional structures into monopolar electrodes and bipolar cells. Monopolar Electrodes* These electrodes are characterised by the variation of the electrode potential, 0 - 0 along the direction of current flow (where 0m and 0S are the potentials of the electronically and ionically conducting phixses respectively). In a packed bed electrode 0m usually remains constant along the direction of current flow, but the marked varia tion in 0S leads to a corresponding distribution of electrode poten tial in this direction. In contrast^the fluidised bed electrode exhibits a variation in 0 due to the larger resistivity of the particulate phase and this can compensate to some extent for the change in 0g) depending on the position of the current feeder. In consequence^fluidised bed electrodes exhibit a more uniform electrode potential than packed beds. The effect of cell geometry on the potential distribution will be illustrated by some examples. Let us denote by B either the boundary or axis of the particulate electrode most remote from the counter electrode C. Most commonly B will coincide with the location of the current feeder. We can now define o* as the fraction of the electrode reaction taking place in a region of the particulate electrode within a distance n measured from B along the direction of current flow. We can write: o = I /I o (1) n n where I is the amount of the current carried by the electrolyte n across a plane a distance n measured from B along the direction of current flow, and I 0 the current across the same plane projected a distance n° (the boundary of the particulate electrode nearest the counter electrode) again along the direction of current flow. For a cylindrical symmetry we replace n by r and write: Francis Goodridge 31 I 2Krdr An f(E) j— -. ~p -v-, (2) = . 'o ri 2Ttr dr AD f(E) ;: fl K where A is the superficial area of the particulate phase per unit volume of the electrode, f(E) the experimental dependency of current density on electrode potential, and r - r, the total extend of the electrode in the direction of current flow. The integrals refer to unit height of the electrode. For planar symmetry 'X 3 J 5 dX Ap f (£) ( > A MODEL FOR MEMBRANE FOULING IN ELECTRQDIALYSIS* G. Grossman and A. A. Sonin Department of Mechanical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts, U.S.A. A theoretical model has been developed to describe certain types of surface fouling of ion exchange membranes in electrodialysis systems. In the model the fouling is represented by a thin surface film of an ion exchange material of sign opposite to that of the membrane itself. The salt is driven through the sandwich-type combination of membrane and fouling film under the action of an applied electric field and diffu sion. This model has been incorporated into a theory for the perfor mance of an electrodialysis system with laminar flow in the channels, and experimental results have been obtained which tend to Conlirm the theoretical predictions. The presence of a fouling film with ion exchange properties on the surface of an ion exchange membrane gives the membrane a layered struc ture. Composite membranes consisting of contiguous cation and anion layers (like "bipolar" membranes) have been studied experimentally by Frilett'e (J. Phys. Chem. 60, 435, 1956) and Ishibashi and Hirano (J_. Electrochenu Soc. Japan 26, 1956), and have been proposed by others as models for certain biological membranes. Unlike simple, homogeneous membranes, such layered membranes are known to have current-voltage characteristics which are anisotropic with respect to current direction and which show a tendency of the current to reach a saturation value in one or both directions. This gives rise to rectification properties similar to those observed in biological membranes. It also suggests a mechanism whereby the membrane structure, rather than concentration polarization in the solution outside, controls or at least influences the value of the limiting current, as has in fact often been observed with ion exchange membranes intended for application in electrodialysis systems. The model developed in this paper shows that an extremely thin ion exchange film having fixed charge of opposite sign to that of the mem brane itself can cause a significant reduction of the limiting current below the value expected on the basis of concentration polarization alone in the exterior solution. Consider, for example, a cation exchange membrane of thickness Ac sandwiched between two anion exchange layers, the left one having a thickness Aa£ and the right one a thick ness Aar. Based on certain simplifying assumptions (that the only mobile species are the ions of a fully dissociated salt; that the positive and negative ions have equal charge numbers Z and diffusion coefficients D; that the fixed charge concentrations in the ion exchange layers are large compared with the exterior salt concentra tions; and that the membranes are very thick compared with the Debye length), analytic expressions are derived for the current-voltage * This work has been supported by the Office of Saline Water, U. S. Department of the Interior ** Now at Avco Corporation, Wilmington, Massachusetts, U.S.A. 34 G. Grossman characteristics of the membrane. For example, for the special case where the anion layers are very thin compared with the central cation exchange membranes, one obtains . A C /_ . i ar a (1 + 5x te c 7F D —T"e ZF4> c , . r . a r „. = + + (1) ivT ZFTT i*cT i" • . Ä 0c c c & ,. i al a> a % Here, $ is the potential drop across the membrane, measured from a point just to the left of the membrane to a point just on the right, where the salt concentrations are C^ and Cr, respectively; Cc and Dc and Ca and Da are the fixed charge concentrations and diffusion co efficients in the cation membrane and the anion films, respectively; and j is the current density (positive to the right). Equation (1) shows that the presence of the anion surface films on both sides causes current saturation in both directions: ZFDaC£ ^ "*" A—C— ^or * "*" + °° (2) a£ a ZFD C2 ^ "*" " A—C— for * "*" ~ °° ^ ar a The positive limiting current is controlled by the left hand anion exchange film, and the negative limiting current by the right hand anion exchange film. If either film is absent, the current will not saturate in the corresponding direction. The current saturation occurs in this case because, in the anion exchange film on the side where the current vector enters the membrane, the co-ion concentration goes to zero at the boundary between the anion film and the cation membrane. This happens at a finite current. The Donnan potential across the junction then becomes infinite, and the current-voltage characteristic shows saturation. For an anion exchange membrane with very thin cation surface films, the subscript "c" for cation in Eq. (1) should be replaced by "a" for anion, and the subscripts "r" and "£" for right and left hand films should be interchanged. The ion exchange membranes in electrodialysis type demineralization systems are susceptible to various types of surface fouling which are known to reduce the limiting current. It has been suggested that some of the mote serious types of fouling are caused by colloidal particles with fixed charges which, when deposited on the membrane, give the mem brane a thin surface film with ion exchange properties, much as in our theoretical model. We have derived a solution for the limiting current in an electrodialysis system with laminar flow in unobstructed chan nels, assuming that the membranes have thin ion-exchange fouling films of opposite charge to that of the membrane Itself. Equation (1) is used for the membrane characteristics, and the hydrodynamic part [which controls the values of C0 and C,. in Eq. (1)] is treated as in an G. Grossman 35 earlier theory (Sonin and Probstein, Desalination 5, 293, 1968) where the membranes were assumed to be ideal (unfouled). The analysis shows that fouling films cause a reduction of the limiting current, as has been observed in practice. When the amount of salt removed in one pass through the system is small, it turns out that the fractional reduction in limiting current can be expressed as a function of the single para meter d2v AfCf D av,)!/3 (- (4) Df dC0 £D where A C, and D, are the thickness, fixed charge concentration, and ion saltV diffusion coefficient of the fouling films, d is the thickness of the flow channels, CQ is the salt concentration at the inlet of the dialysate channels, D is the diffusion coefficient in the solution, Vav is the average flow velocity in the channels, and I is the length of the current carrying part of the desalting channels. The theoretical curve for the reduction in limiting current is shown in Fig. 1. Also shown in Fig. 1 is a series of experimental results for the re duction in limiting current as a function of £. In these experiments the membranes were deliberately fouled with iron oxide, which ap parently causes a surface fouling film similar to an ion exchange material (Grossman and Sonin, Desalination, in print, 1972). For the data points shown the flow conditions were constant, while the value of £ was varied by varying the salt concentration CQ at the channel in lets. In the figure the ordinate represents the ratio of the measured limiting current to the theoretical limiting current for ideal, clean membranes (Sonin and Probstein, cited above). Since the properties of the fouling film were not known, a value had to be assumed for (AfCf/Df) in Eq. (4) before a comparison could be made, and the com parison is therefore not an absolute one. Nevertheless, the trend of the theory appears to be supported by the experimental results. l.Oj 60 C THEORY 3 O O •^1 EXPERIMENT 0 -3 -2 10 10 10 Fig. 1. Fractional reduction of limiting current as a function of the fouling parameter E 36 E. Heitz MASS TRANSFER DETERMINED REACTIONS IN FLUIDISED BED CELLS Ewald Heitz and Siegfried Pionteck Dechema-Institut, Frankfurt, Bundesrepublik Deutschland Basic work on fluidised bed cells has been published by Flaischmann, Goodridge and coworkers. The work presented here deals with some aspects of mass and charge transfer in such cells. An electrolyte flow circuit has been con structed with a channel flow electrochemical cell. The cell can be used with a fluidised bed electrolyte (non conducting particles) or with fluidised bed electrodes (conducting particles), the feeder electrode being part of the channel wall. The flow circuit enables different flow velocities with corresponding bed expansions to be achieved. The particles (glass beads with and without silver coating) had diameters from 0,25 to 0,90 mm . The electrochemical systems used were Ferrous-/Ferricyanide and Silver/Silver ions. The mass transfer was determined by measuring limiting currents a) in a one-phase flow and b) in a two-phase flow as a function of flow velocity and bed expansion with non conductingspheres (fig. 1 and 2). The maxima of the curves are interpreted in terms of collision numbers and collision intensities at various bead diameters. In the vicinity of the fluidisation point (F.P.) no discon tinuity can be detected. This finding is important with respect to mass and heat transfer in fluidised beds. Fig. 3 shows the "current/potential curves of the silver deposition in the presence of conducting and non con ducting spheres as well as in the one phase channel flow. The ohmic potential drop has been eliminated by using an interrupter method. The near linear polarisation curves (conducting particles) are interpreted on the basis of a current distribution (published by Fleischmann and Good ridge) under conditions of controlling mass transfer. Furthermore the influence is discussed of decreasing the exchange current density by a factor of 1000 (by cyanide addition). Conclusions are presented concerning the increase of the mass transfer by non conducting particles and the scale up of cells by use of conducting particles. E. Heitz 37 d-0,45 mm 0,5- 0,4- 0.3- 0,2- 0,1- 1 r 2 3 4 Fig, 1,2: cathodic limiting currents as a, function of different parameters, Ferricyanid = 2 • 10~ M 0% /10% / 22% Ag* 2-10*n ^ ^—z- u(cross-sect J! cm KNOy. 0,5 n diameter : 1mm 40- mA cm' \ % / , 30- 1-3 I / I 20- / •*\ glass beads t*~ p2 f / 10- / -1 / no beads -200 -400 -600 -800 -1000 q [mV] Fig, 3* current-potential curves of silver deposition 38 James P. Hoare SOME SURFACE STUDIES OF ELECTROCHEMICALLY MACHINED HIGH-TEMPERATURE ALLOYS James P. Hoare, Armand J. Chartrand, and Mitchell A. LaBoda Electrochemistry Department, Research Laboratories, General Motors Corporation, Warren, Michigan 48090 USA Since the electrochemical machining (ECM) operation is considered to be a controlled, high-speed, corrosion process, the electrochemical properties of the metal-electrolyte system as a function of potential emerge as one of the principal factors in the design of a machine which could produce an acceptably finished product. From an extensive screening of possible electrolytes suitable for the machining of hardened steels, LaBoda and coworkers (1) reported that high metal removal rates with excellent control of geometry and dimensions and with high quality surface finish were obtained on fully hardened steel by ECM machining in solutions of NaCl03. Solutions of NaCl gave cuts with rounded corners and those of Na2Cr20y did not produce any sig nificant cutting of steel by ECM. To understand the corrosion behavior of steel in these various electro lytes, steady-state polarization data under constant potential con ditions were obtained on steel microelectrodes in 0^ saturated so lutions of NaClO^, NaCl, Na£Cr207 among others in a dual Teflon cell with a saturated calomel probe-type reference electrode'and a Pt-gauze counter electrode (2). In addition, constant current stripping data were obtained to provide information about the presence of anodic surface films as a function of potential. The results of these studies were compared with those obtained on steel machined in an ECM test rig using the same solutions of NaCK>3, NaCl and Na2Cr207. From these studies it was concluded that those electrolytes which gave good dimensional control were those (e.g., solutions of NaC103 or NaN03) in which a passive film was formed on the steel anode. In those solutions (e.g., NaCl) in which protective films are not formed on steel, rounding of the corners (wild cutting) results. It was further concluded that the ECM process takes place in the transpassive region, and the closer the dimensional control, the narrower is the range of potential over which the transition from the passive to the transpassive state takes place. At points on the anodic workpiece remote from the face of the cathodic tool (sides of a cut), the long electrolyte path has a high resistance, the potential drop across the metal-solution interface is low corresponding to the passive region of the polarization curve, arid the metal removal rate through the pro tective anodic film is very low in this low current density (led) region. On the other hand, at points opposite the cutting face of the tool (bottom of cut), the short solution path has low resistance, the potential drop at the interface is high corresponding to the trans passive region of the polarization curve, and metal is removed at a high rate in this high current density (hed) region. In the hed region, the anodic film on steel in NaC103 solutions is reduced to a thin uniform porous layer through which metal ions pass rapidly and underneath which electropolishing of the metal surface James P. Hoare 39 takes place by the mechanism proposed by Hoar and Mowat (3). If the anodic film is removed entirely from the metal surface in the trans- passive region as in the case of steel in solutions o- ^aClO^, electropolishing of the surface does not occur but in& Lead a roughen ing of the sur/ace due to pitting in the led region. For solutions of Na2Cr207 the anodic film is so strongly protective that, even in the hed region, the film is thick enough to prevent rapid dissolution of metal. Because the great advantage that ECM holds over conventional machining methods is the ability to remove metal rapidly from high-strength, high-temperature alloys, and because such alloys contain significant amounts of elements other than iron and carbon, it was considered desirable to ascertain how the presence of these foreign elements mod ifies the nature of the protective film; and in turn, how the quality of the electrochemically machined part is affected. Samples of seven selected high-temperature alloys (H-13s tool steel; Inco 713-C, Ni- based; Inco 901, Ni-Fe alloy; GMR 235, Ni-based; HS-31, Co-based; A286, Ni-stainless steel; 410, Cr-5tainless steel) were machined in a laboratory, ECM, plunge-cut, machine and trepanned pegs from the plunge-cut were used as anodes on which the steady-state polarization and film-stripping data were obtained. All work was carried out in NaC103 solutions (350 g/1). It was found that the metal removal rate of high Ni-content alloys (713-C) is only somewhat lower than the carbon-steels but the removal rate of high Cr-content alloys (stainless steels) was very low. According to the limiting current-type of polarization curve for such alloy anodes, it is likely that the presence of Cr in the anodic film greatly increases the electronic conductivity of the film. Conse quently, a significant amount of current is consumed in the evolution of 02 rather than in the dissolution of metal, thus accounting for the low cutting rate. For those alloys whose polarization curves exhibit a sharp transition from the passive to the transpassive region, good dimensional control of metal removal was obtained with ECM even though the metal removal rate may be low. The smoothness of the machined surface was determined from the Proficorder tracings and the structure of the surface was assessed from scanning electron micrographs. The two stainless steel samples had RMS peak-to-valley roughness values of 5 to 10x10"6cm, yet the A286 surface was mirror bright, whereas the 410 surface was dull. The electron micrographs showed that the 410 surface had a sponge-like appearance and the dullness resulted from light being trapped in the cavities. On the other hand, metal was removed evenly on the grains of A286 with slight attack of the grain boundaries producing the highly reflecting surface. In GMR 235 the grains were preferentially attacked, leaving a network of raised grain boundaries (surface roughness = 2 to 40xl0~5 cm). When there are large differences between the dissolution rates of the metallurgical phases of an alloy such as 713C, the surface obtained with ECM is very grainy and rough (2 to 25x10"5 cm). 40 James P. Hoare (1) M. A. LaBoda and M. L. McMillan, Electrochera. Tech.,, _5, 34 (1967). (2) J. P. Hoare, et al, J. Electrochera. Soc, 116, 199 (1969); 117, 142 (1970); Corrosion, 27.,' 211 (1971). (3) T. P. Hoar, et al., Nature, 165, 64 (1950); 169, 324 (1952); Electrodepositors Tech. Soc, _2_å> 7 (1950). 41 I.N. Justinijanovic THE INFLUENCE OF FOREIGN ATOMS ON THE PROPERTIES OF ELEC TROLYTIC ZINC POWDER I.N.Justinijanovid, J.N.Jovidevid and A.R.Despid Institute for Chemistry,Technology and Metallurgy and the Faculty of Technology and Metallurgy,University of Beograd, Yugoslavia A study has been carried out of the conditions of obtaining zinc metal powder by cathodic deposition from alkaline zincate electrolytes. Electrolysis has bean carried out in cells which enabled maintaining well defined conditions of temperature, zincate concentration and cathode overpoten- tial. Hydrogen evolution rate was followed simultaneously by measuring the volume of evolving gas as a function of time. Thus, current efficiency could be worked out. The powder structure has been investigated by electron micro scopy and the specific surface area determined by the BET method. Powder formation was found to coincide with the attainment of diffusion limiting current density as expected from the theory. Conditions of spontaneous falling off of the powder from the cathode have also been investigated. Introduction of lead and tin ions into the electrolyte in concentrations much lower than that of zincate ( 0,29 3% v.s. zinc) had profound effects on the quality of the de posit. Zinc powder, from the lead containing electrolyte, preserved the dendritic macro-structure characteristic of the powder from the pure electrolyte. However dendrites are seen not to be single crystals, as usual, but rather a complex structure composed of numerous crystalites of high dispersion. This is reflected upon the specific surface area which has been found to be considerably larger than in powders from pure zincate solutions. The effect of tin could hardly be noticed on the micros copic appearance of the powder particles. Well developed dendrites with branching of up to the third order made crystals very similar to those from pure zincate. However, very large increase in specific surface area has been found due, obviously to sub-micro structure of the crystal grains. Attempts to interpret the obtained results are being made. 42 V.D. Kashch.eev OH THE ANODIC DISSOLUTION OP METALS AT HIGH CURRENT DENSITIES V.D«Käshcheev,B,N»Kabanov,A.DoDavydov Institute of Electrochemistry Academy of Sciences of the U.SoS.H,Moscow The electrochemical behavior of a number of metals and alloys has been studied in a wide potential range.The anodic polarization curves have been obtained up to the current densities 50 a/cm"# The effect of the anion nature concentration,the pH value and the intensity of stirring of an electrolyte on the shape of the polariza tion curves and the quality of the sample surface after dissolution has been demonstrated» The relationship between the shape of the anodic polari zation curve and the precision of the electrochemical machining has been analyzed» Joseph Kerti 43 SULPHATE ELECTROLYSIS BASED ON ANIONIC PROTON MIGRATION Dr Joseph Kerti Department of Chemical Technology, Technical University of Budapest, Hungary. The fundamental problems and difficulties of sulphate electrolysis carried out using insoluble anodes are due to the fact that the acid formed in the anode process is difficult to keep off the cathode owing to the great mob ility of hydroxonium cations. A considerable part of cathode processes - including ma jority of metal depositions - is maintainable only in the case if acidification of the solution in the cathode com partment as result of fundamental decomposition reaction of electrolyte: 0 w w MeSO^ + HpO = Me -t- Ha^2S0^4 +^ 0,>-5^ 02, 1 can be prevented or counteracted. In such cases sulphuric acid liberated on the anode must be separated from the cathode by means of a diaphragm as illustrated in Fig. !_• Fig. 1: Ion migration during electrolysis. It is, however, a relatively low current yield that methods like this make possible on account of the parti cipation of mobile hydroxonium ions in the current trans port through the diaphragm unless working potential of hydrogen evolution is negative enough e.g. in consequence of hydrogen overvoltage. These ions, having arrived the surface of the cathode and discharged may result in parti al water decomposition and, consequently, a current yield decrease the degree of which depends on the actual trans ference number of hydroxonium ions in stationary state of electrolysis. In order to solve key issue in question it has been pro posed using a rotating anode capable for separating the evolved sulphuric acid which is adherent to the surface of the anode, from the metal sulphate solution. Other ideas have been examined on principle of continuous over- 44 Joseph Kerti flowing of electrolyte from the cathode side of diaphragm into the anodic region. Efforts have "been used by applic- ating double-layer diaphragms and other structural devi ces for separating anodic compartments from cathodic ones without any result worthy of mention. It must be underlin ed that the use of permselektive membranes could not come up to expectations in this respect so far. All these ideas are very difficult to be accomplished on industrial scale in spite of use or combination tricks like mentioned and they are relatively expensive, recquire complicated appa ratus so that the electrolytic decomposition of metal sul phates according to reaction 1_ has not proved successful in practice. In many processes used in heavy industry me tal sulphate is formed as a by-product, for example, in the viscose process sodium sulphate or in the case of iron pickling iron sulphate in enormous quantities all over the world. Consequently, an economical conversion of these by-products to more valuable materials is highly desir able considering the extreme pollutions due to sulphate wastes in question as well. The reasoning which has so far hindered the use of anoly- tes concentrated for the salt considering "impurities" of sulphuric acid formed in anode reaction is proved to be false as the application of ternary electrolytes and spe cial combinations of electrolysis and original technology /e.g. production of viscose silk, iron pickling, working up separated ores soluble in acids according to the known Tainton»s method etc./ make possible using certain solu tion-structural and transport effects which enable the Tainton? s method to be applied to a wider field including the operations based on combinations in question. In sulphuric acid solutions containing excess bisulphate- -forming salt the equilibrium is strongly shifted towards the formation of HSO4. i°ns« These, like hydroxonium ions tend to release protons taking part in protonizational reactions e.g. neutralizations. As for migration, however, the behaviour of bisulphate ions is diametrically opposed to that of hydroxonium ions- during electrolysis. Due to their negative charge.bisulphate ions carry the protons in the direction of the anode. This phenomenon - the me chanism of which seems to be very complicated and not yet cleared up in every respect - was termed by the author as anionic proton migration. It has been proved that the migration of hydroxonium ions towards the cathode can be suppressed in warm, concentrated and acid solutions of sulphates as a result of the combination of a number of effects to an extent as to enable carrying out the electro lysis of sulphate salts using insoluble anodes directly without permselective membranes. Joseph Kerti 4 5 The practical and industrial applicability of the method "based on anionic proton migration can "be demonstrated by the example of the continuous regeneration of sulphuric acid pickling solutions with simultaneous cathodic iron refinement and that of the hath of viscose rayon produc tion. The continuous iron pickling and sulphuric acid regene ration system is represented in Fig. 2. The acid solu tion containing "bisulphate forming salts takes part in a circulation as shown in Fig. 2 in a definite direction and velocity depending on the productivity of pickling equipment and the current intensity. In this combination the following reactions could interpret the process. The predominant reaction of pickling: Fe + H2S04 = FeSO^ + H"20 2 The over-all reaction of electrochemical iron sulphate decomposition: FeSO^ + H20 = Fe + H2S04 + 0,5 02 j> The over-all process of reactions 2 and 2. is essentially the decomposition of rust: FeO = Fe + 0,5 02 4 It must be emphasized that all these approximative reac tions are only for the sake of characterization of the technical applicability the method to be published. «r ^ * -r >> + "^r —H t=^i i , rr~i _i Fig. 2: Iron pickling design connected with continuous regeneration of sulphuric acid and iron refinement. 1_: anode compartment; 2: cathode compartment, ^s diaph ragms; 4: pickling vats, g.: electrolysers. 46 Christopher J. H. King SOME FEATURES OF A NOVEL BI-POIAR ELECTROLYTIC FLOW CELL^1' Christopher J. H. King Depart\nen~of Chemical Engineering, University of Newcastle upon Tyne, Newcastle upon Tyne, NE1 7RU, England. The bi-polar electrolytic flow cell described here can be used for the electrolysis of organic and/or inorganic depolarisers where no diaphragm is required, and where it is desirable or inconsequential that the anolyte and catholyte should intimately mix. The cell is designed to function with electrolytes containing high concentrations of conducting species^ ', and also to cope with large amounts of evolved or reactant gases. It is also envisaged that in some cases the system could be run at the boiling point of a volatile product when the cell could act not only as a reactor but also as a separation stage. One of the primary intents of this design was simplicity of construction. The cell (fig. 1) consists of a vertical stack of horizontal cylindri cal electrodes (^1 cm in diameter), each electrode, A, being separated from the other one below it by insulating '0' rings, S, which produce a gap of approximately 1 mm. Other rod diameters and gaps can of course be used, depending on circumstances. The cylindrical electrodes are located in vertical slots machined into the cell body, B. The electrical power supply is connected to the upper and lower electrode in the stack. Electrolyte is fed to the cell via the upper electrode, which is hollow and perforated along its lower surface. The electrolyte passes through the perforations and cas cades over the tier of electrodes below, and in doing so fills the gap between the rods but falls in a thin film over the sides of the rods. The electrolyte falling off the lower electrode collects in a reservoir and may be recirculated. The space at the side of the tier of rods allows gas flow without producing a significant pressure drop, or difficulties in gas disengagement. To illustrate the effectiveness of the rods as bi-polar electrodes, one of the rods in the stack was split in two along its length and insulation placed between the two parts. Then by connecting an ammeter between these two halves the current flowing across the rod could be measured. Results in such a test for a total of 12 graphite rods using an electrolyte of 3% sodium chloride are given in table I. TABLE I EFFECTIVENESS OF BI-POIAR CELL USING 3% SODIUM CHLORIDE SOLUTION Current Current Voltage Effective Effective through across split across cell res film res cell, mA. rod, mA. cell, V. istance , Q . istance, Q . 53 18 25.6 484 730 152 106 35.3 232 768 280 230 38.8 139 776 Christopher, J. II. King ur*^ e 7 / / / i I i i i I ) i i f > / I I I > I /., i.. /.;./ + •+ n » »••-«- 2=3 + + + + + -f / [ft i I 11 III/ / // r/// / / / / / / / /////////// y / /y FIG. 1 . 48 Christopher J. H. King The results in table I indicate that, with suitable choice of conditions, by far the greater proportion of the current through the cell is carried by the bi-polar rods. The extent to which the bi polar effect operates will depend on the polarisation characteristics of the reaction in question and the resistance of the electrolyte. Taking as a trial reaction the formation of sodium hypochlorite from sodium chloride; the overall cell reaction is given by equations (1) and (2). The cell, containing 13 graphite rods of 1 cm diameter and 2Nacl + 2H0H > 2NaOH + Cl2 + H2 (1) Cl2 + 2NaOH v ^ NaCl + NaCIO + H20 (2) of effective length 12 cm, was used to electrolyse a 3% solution of sodium chloride. With an electrolyte flow of approximately half- litre per minute and 46 volts applied across the cell, the current flow through each bi-polar unit (as indicated by the split rod technique) was 260 mA., and that through the film round a rod only 50 mA.. Such an experiment gave a solution leaving the cell contain ing 0.07% in sodium hypochlorite by weight, the corresponding power efficiency being determined as 200 g kWh™ in terms of sodium hypochlorite. This would appear to be viable when compared to industrial units used for this reaction. This electrolysis reaction can be extended by the introduction of propylene into the gas space of the cell to form propylene oxide by a similar reaction sequence as described elsewhere' '. Other results from this bi-polar cell will be given at the meeting. References 1. Part of British Patent Application No. 60763/71, 30th December, 1971 by Fleischmann, M., King, C.J.H., Oldfield, J.W., Plimley, R.E. and Tennakoon, C.L.K. 2. See extended abstract by Fleischmann, M. and Goodridge, F. 3. See extended abstract by Goodridge, F. and Osifade, O.B. Dj .Matic . 49 ELECTROCHEMICAL INVESTIGATIONS OF THE CYLINDRICAL ROTATING ELECTRODE D.j.Katie» B.Lovrecek, D.Skansi Institute of Electrochemistry and Electrochemical Technology, Institute of Chemical Engineering Faculty of Technology, University of Zagreb, Yugoslavia The mechanism of the mass transfer on the cylindrical rotating electrode was investigated. Investigations were performed within the range from 3oo to 25oo rpm. The solution of ferri- and ferrocyanide of equimolar con centration in 1 N NaOH as supporting electrolyte was used as simple electrochemical system for the performance of these investigations. Two measuring methods were used: the steady-state current-potential and the galvanostatic pulse method. By the steady-state method curves with well-expressed limiting currents were obtained. It is, however, charac teristic that more or less irregular oscillations of current occur on the plateau of limiting currents,except at the lowest rotation rate. At the highest rotation ra tes, their frequency increases and amplitude decreases.By the galvanostatic pulse method it was possible to obtain e-?-t curves with a well-expressed and typical transition time. Measurements were performed with current pulses from 3.0 to lo.o mA, and in all cases it was possible to achi eve the typical transition time, except at the so-called critical rotation rate, respectively above it for a given value of the current pulse. The analysis of results was performed by connecting hydrodynamic and electrochemical conditions on the cylin drical rotating electrode. It was found that in a wide range of rotation rates there is a linear relationship between the Sh-number and the 0.6 power of the Re-number, this being otherwise suggested as criterion for the de veloped turbulence in systems with the cylindrical rota ting electrode. However, also at lower rotation rates, where this dependence cannot be described with the expo nent 0.6 above mentioned oscillations were observed on the plateau of limiting currents. These oscillations suggest that the influence of eddies is already present, therefore this would correspond to the so-called transition region. In the analysis of hydrodynamic conditions next to the electrode a modifica tion of Ta-number was introduced. This modification seems to describe better the system, which was investigated,and 50 Dj.Matic obtained numerical values of 3?a-numbers fit in general into the picture obtained on the basis of experimental results. Vittorio de Hora 51 THE GLANOR DIAPHRAGM CELL; A HEW TOOL FOR THE CHLORIHE INDUSTRY Vittorio de Nora Oronzio de Hora Impianti Elettrochimici, Lilian, Italy. A new bipolar diaphragm cell has been developed utilizing dimen- sionally stable anodes (DSA). Glsnor cells assembled to form an electrolyzer have important features particularly with regard to power consumption arid diaphragm life.The voltage drop in the various cell components is reported together with product quality at different current densities. One type of Glanor cell,the V 11/44,which produces up to 27 tons of chlorine per day is described and its economics discussed. The Glanor cell has been developed by Oronzio de Hora Impianti Elettrochimi,Milan-Italy in corporation v/ith PPG Industries,Inc., Pittsburgh Pa,USA. S.K. Rangarajan 52 THEORY OF ELECTROCHEMICAL MACHINING S,K« Rangarajan, T.G. Ramesh and S« Ramabhadran Materials Science Division and Aerodynamics Division, National Aeronautical Laboratory, Bangalore-17, India. This paper considers some of the problems associated with preci sion machining by the electrochemical method i.e.,removal of the metal from the work-piece by passing a current between it and the tool in a suitable electrolyte.The machining operations are very much controlled by the geometries of the work-piece and tool and Y-he relative movement of the tool. Fundamental electrochemical principles involved in the electro chemical machining operations are diseussed.The importance of current density distributions in such operations is emphasised. Several mathematical models describing the geometries of the tool, work-piece and the cell are considered and theory of processes like smoothing, shaping, deburring, cutting and cavity sinking are reviewed.Parameters signifying the job-precision are proposed. The results obtained by us in the analogy experiments on tool design are compared with the analytical results reported. Darko Skansi 53 STUDY OF MASS TRANSFER IN FLOWING ELECTROLYTE VITE QUASI POROUS ELECTRODE Branko Lovrecek, Darko Skansi, Djani Matic Institute of Electrochemistry and Electrochemical Technology Institute of Chemical Engineering Faculty of Technology, University of Zagreb, Zagreb, Yugoslavia In the continuation of our research of the electrolysis with a flowing electrolyte and a quasi porous electrode investigations with a quasi porous segmental electrode were performed. In the vertical flow cell three segments of the porous electrode, made of a rolled platinum grid,were placed. The segments were one behind another in the direction of the electrolyte flow. The solution of 2.oxlo~5 M/1 KxFe/ON/g in 1 N KOI in water was used as electrolyte. The steady-state method of measuring the current-potenti al relations was applied. By a suitable electrical circuit it was possible to mea sure the total current of the electrolysis, as well as the currents of each segment separately. The results obtained for each single segment are in accordance with our former interpretations based on a mo del, which should be electrochemically and hydrodynamica- lly equivalent to the investigated system. The results obtained with a multi segment electrode give information on the change of the initial concentration from the upper to the lower edge of the porous electrode. The decrease of concentration on one segment as well as along the whole segmental electrode does not depend on the initial concentration, if the potential at any point of the electrode corresponds to the plateau of the limi ting current. However, at higher flow rates, i.e. also at higher current densities in the vicinity of the lea ding edge, lower parts of the electrode remain on the lo wer potential than the potentials of the limiting current plateau due to the high IE drop. This appearance, which can be numerically analyzed, may be the limiting factor in the optimalization of electrochemical cell with a flo wing solution and a quasi porous electrode. 54 A. A. Sonin ELECTROKINETIC SALT REJECTION BY POROUS MATERIALS - THEORY AND EXPERIMENT* G. Jacazio, R. F. Probstein, A. A. Sonin and D. Yung Department of Mechanical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts, U.S.A. It has long been known that clays and solids have the property of partially rejecting salt when a saline solution filters through them. The same phenomenon has been evidenced by other relatively porous materials including cellophane membranes, porous glass, and ion exchange materials. In many such macro-porous materials, where the pore size is large compared with molecular dimensions, the rejection mechanism has been assumed to be an electrokinetic one, resulting from charge built up on the interior surfaces of the porous material when in contact with the salt solution. The salt-rejection phenomenon in flows through macro-porous media can be analyzed in terms of a relatively simple physical model, where the fluid is assumed to flow through the porous material via a series of uniformly distributed, straight, cylindrical pores. The problem can first be solved for a single capillary and then transformed back to apply to an actual porous medium by expressing the effective pore radius and pore length in terms of physically measurable bulk quanti ties like the permeability, porosity, and thickness of the porous bed. The physical mechanism of salt rejection in a flow through a capil lary with a charged interior wall is qualitatively clear. The surface charge gives rise to a potential field which extends a distance approx imately equal to the Debye length into the liquid within the pore. Within this region, there is an excess of ions having a charge opposite to that of the wall. If a pressure gradient is applied to make the salt solution flow through the pore, then because of an excess of charge of one sign within the pore liquid there results a net transport of charge and the buildup of the streaming potential The effect of the potential is to set up an electric field parallel to the surface which will increase the transport of one ion and reduce that of the other ion until there is an equal steady state transport of positive and negative ions, the feed and effluent solutions being insulated from each other so that there is no current flow. The presence of the elec tric field thus tends to reduce the average velocity of the ions in comparison with that of the water molecules, so that there is a net re jection of salt proportional to the difference between the two velocities. Related electrokinetic phenomena in a simple capillary tube have been analyzed in a large number of works. In general the earlier treatments were approximate and applied to the case where the Debye length XJJ was small compared with the tube radius a. More recently the capillary problem has been formulated quite rigorously by Cross and Osterle (J. Chem. Phys* 49, 228, 1968) for the general case, assuming * This work has been supported by the Office of Saline Water, U.S. Department of the Interior. A. A. Sonin 55 only that the Debye sheath can be described by the Chapman-Gouy model. The formulation is the usual phenomenological one in which transport of the ions and water molecules results from the effects of gradients across the membrane in pressure, concentration and electrical poten tial. Although applied to some fundamental electrokinetic problems, no explicit solutions were worked out for the problem of interest here. We have applied the capillary model to derive explicit solutions for salt rejection in a flow through a macro-porous medium with charged in terior surfaces. This analysis is for the case of a dilute fully dis sociated salt solution, and it is assumed that the potential of the pore wall has a fixed value independent of salt concentration. The solution for the salt rejection R, obtained by solving the system of equations, can be shown to be expressible in the functional form u - u ZF<$> - X „ s _ cf w uL DN u where u is the average water speed in the pore and u the average salt speed. The three parameters upon which the solution depends are the Debye length to radius ratio, the dimensionless surface potential, and a dimensionless flow rate or Peclet number based on the pore length L and the salt diffusion coefficient in the pore liquid D. In general, the solution for R can only be given numerically. It is possible, how ever, to express the result analytically in closed form for the par ticular case ^/a -** °°, where the rejection process takes place entirely in the sheaths formed outside the pore at the entrance and exit. An interesting feature of this solution is that it can be shown to be applicable for ^D/a a. 2. In order to check out the theory which was developed, experiments were carried out on the salt ^ejecting properties of compacted clay, through which saline solutions were forced under high pressures. Direct measurements were made of the wall potential 4> (related to the zeta potential) and all of the bulk properties which determine the parameters appearing in the theoretical solution. Figure 1 shows a comparison of the large Debye length solution with the experimental data. The different effective pore radii a were obtained by changing the compaction of the clay, while the different wall potentials were obtained by changing the pH of the feed solution. The agreement be tween theory and experiment is seen to be excellent. It is most im portant to emphasize that the comparison is an absolute one, since the theory has no adjustable constants or parameters. The effect of vary ing the Etabye length to radius ratio is shown in Fig. 2 which is for the case of large flow rates, i.e., uL/D » 1. Again the agreement between theory and experiment is an absolute one with no adjustable parameters in the theory and again the agreement is striking. A. A. Sonin i r 60 — 20^ \T. *\ 60 50 Theory. * •VJ =16.3 A 40 — £T 30 20 10 100 Ap atm Fig. 1. Percent salt rejection versus pressure drop across membrane ($ = ZFé /RT). w w • Q = 1.77 to 1.88 O C = 1.96 to 2.10 d) = 1.22 (30.5 mV) W $TT = i.60(40mV) 0.1 XD/a Fig. 2. Salt rejection versus Debye length to radius ratio for Q = uL/D > 1. Miroslav Spasojevic 5 DESIGN APPROACH TO THE CHLORATE CELL M.M.Jatesic, N.M.Jafcsié, M.D.Spaso.jevic and B.Z.Niteolic electrochemistry Department, Institute of Chemistry, Technology & Metallurgy, Beograd, Karnc-iijeva 4/IV, Yugoslavia An approach to the chlorate cell design based on mass- -transfer consideration and including both the effect of hydrodynamic operational factors and polarization chara cteristics of the electrodes is presented. The procedure relates to both unipolar and bipolar chlorate cells and is also applied to both cells with graphite and metallic anodes. Resulting optimizations are also given. 58 K.S. Spiegler CONCENTRATION-POLARIZATION AT ION-EXCHANGE MEMBRANE/SOLUTION INTERFACES C. Forgacs , J. Leibovitz, J. Sinkovic and K. S. Spiegler Sea Water Conversion Laboratory, University of CalifornTa", Berkeley, California, USA. The aim of this investigation is the study of polarization pheno mena in electrodialysis by electrochemical methods. This is being done by measuring current-voltage curves across ion-exchange membranes in different locations along the tortuous path of salt solutions flowing in planes parallel to the membrane surfaces. The results prove that for the system studied ("Mark II"-type, Ionics, Inc., Watertown, Mass.) that portion of the current which is due to the motion of the salt ions through the membrane, I , reaches a limiting value with increasing ter minal voltage, but this is not reflected in the existence of a limit ing-current plateau in all locations, because side-effects mask these plateaus in all but the last sections of the solution flow path. The values of the salt-carrying current, Is, are in fairly good agreement with computations based on the Nernst-Planck equations of ion flow; this is true both in the plateau region and at lower currents. Calculations of the current-voltage curves for a differential membrane element (of size small enough so the change of the specific resistance of the solution along the flow path is negligible) show that a limiting-current plateau, similar to a polarographic plateau should be reached when the electric potential difference across the membrane is gradually increased (1). These calculations are based on the Nernst-Planck equations of ion flow. In line with conventional electrochemical theory the voltage drop across the membrane is taken as the sum of the ohmic voltage drops (determined uniquely by the con centration distribution of the salt), plus membrane potentials, gov erned by (a) the concentrations at the membrane-solution interfaces and (b) the transport numbers of the ions in the membrane. The occurrence of a limiting-current plateau is due to the change of the ion transport numbers at the membrane/solution interface (1). The quantitative description of the phenomenon is smilar to that for metal/solution interfaces for which such plateaus have been reported repeatedly in the chemical literature (2-4). With some simplifying assumptions—applicable in dilute solutions--the current-voltage curves could be described in closed form. The problem is more complicated when the variation of the concentrations of the solutions in the com partments adjacent to the membrane along their flow paths is taken into account. The potential difference, AE (volt) across an anion-exchange membrane between two probe electrodes (e.g. Luagins) at distance x (cm) from the solution inlet is found to be (5) the following function of the current density, i (amp cm"2): Present address: Negev Research Institute, Beersheva, Israel. K.S. Spiegler 59 h< A£,'lhA a - I OM) I -^- ~ å ) / { Co + /?„ | idx + I L OM)f-^-<5J / ( cc0 - & J\dx j +K (1) 0 + [(£M) + C] In cto - fie ! «'dx - (<5i/B) + A, J idx + («3i/0) + - -i o o + (BIX) In Qo + Pi J «'dx o O J Here X is the equivalent conductance of the solution, ohm"1 cm2 eq- i the h's are the compartment thicknesses (cm), subscripts d and c standing for diluate and concentrate respectively, 6 (cm) the Nernst diffusion layer thickness, assumed to be uniform, and determined from the current-voltage curve of the last section (where a limiting cur rent plateau was found to exist), C^Q and cco the inlet concentrations to the diluate and concentrate compartments respectively and FL, the membrane resistance per unit area, ohm cm2. C is the preloganthmic factor in the membrane potential equation, (RT/30 (2t_ - 1) volt. t_ is the transport number of the anions in the membrane. The 3's are defined as *5 B = ?D(t_ - t_) (3, D and t_ being the diffusion coefficient and anion transport number of the electrolyte in free solution respectively. Equation (1) was derived for dilute KC1 solutions (t_ = 0.5), under the assumption of uniform coulombic efficiency: (TVth E l~ " *- (4) where tr_S is the anion transport number in the separator adjacent to the membrane. (For our almost non-permselective separator, this transport number is similar to that in free solution.) 60 K.S. Spiegler Computation of AEt^ as a function of i showed that in spite of the complexities of the situation, a limiting current plateau should be reached provided that the substitution of a hypothetical uniform average Nernst diffusion layer thickness for the periodically changing thickness is realistic (6). The validity of eq. (1) was indeed con firmed by a limited number of experiments in our Laboratory (5). Additional measurements are reported in Figure 1 which shows current- voltage curves measured in three of the eight sections making up the "Ionics" tortuous path spacer (Ref. 7, p. 250). The apparatus used is schematically shown in Figure 2*. It is seen that while the mea sured current density exhibits a reasonably flat plateau only in the last section, correction for the imperfect salt transport converts all curves to plateau curves. The aim of the correction method (5) is to isolate that portion, icorr of the current density which con tributes to the desalting process, while leaving out currents due to side-effects [e.g. participation of H+ and 0H~ in the current (7, Chapter 6 by L. H. Shaffer and M. S. Mintz)] 1corr E il^fc (5) Here rj^ is the coulombic efficiency determined from chloride titra tion olr influents and effluents in the absence and presence of the current. The pH measurements confirmed that at least part of the current above the plateau is due to electromigration of H+ and 0H~. Thus limiting ion transfer can be determined from measured cur rent-voltage curves even at anion-exchange membranes of relatively large size (45-50 cm), and the future analysis (for size) of polyions by "membrane polarography" does not seem impossible. The spacers were high-density porous polyethylene sheets of 1 mm thickness (Porvair, Ltd., King's Lynn, Norfolk, U.K.) covered with cellulose (cut from dialyzer tubing) facing the intermediate com partments. References (1) K. S. Spiegler, Desalination £, 367 (1971). (2) C. W. Tobias and R. G. Hickman, Z. Phys. Chem. (Leipzig) 229, 145 (1965). (3) N: Ibl and U. Braun, Chimia 2±, 395 (1967). (4) J. Sinkovic and D. Leskovsek, Electrochimica Acta J6_, 2125 (1971), (5) C. Forgacs, N. Ishibashi, J. Leibovitz, J. Sinkovic and K. S. Spiegler, Desalination }0_ (1972), in press. (6) A. Solan, Y. Winograd and U. Katz, Desalination 9_, 89 (1971). (7) K. S. Spiegler, ed. "Principles of Desalination", Academic Press, New York (1966}. ! K.S. Spiegler 4^-cr- 2 (Corrected) ,-S 4 (Corrected) ^f ~\ / A ^-D^8 (Corrected) .^.XV- I & j_ 0.5 1.0 1.5 20 2.5 30 AE (volt) Fig. 1. Current-Voltage Curves of Different Sections of Antony Exchange Membrane 11^ 8ZTT53r: Probe electrodes Ag/AgCI. Solution: 1 OTOI M KCI; flow velocities ud = 11.3, uc = 9.5 cm sec" . 25°C. Electrode Intermediate A.E. A.E. Comport Comport - Membrane Membrane ment ment Comport - Compart - Streams Streams ment Stream ment Stream Anodic Cathodic Note P = Pump F = Flowmeter PG= Pressure Gauge TH= Thermometer pH « Continuous pH Measurements In actual cell, membranes, separators, and spacers are horizontal. Diophrogm Valves Anion-Exchange Membrane m Cation-Exchange "* Membrane Separator Spacer Fig. 2. Schematic Flow Diagram of the Pilot-Size Electrodialysis Cell (5). P-pump, F-flow meter, pG-pressure gauge, TH-thermometer, pli-continuou5 pH measurements. In actual ceil, membranes, separators and spacers are horizontal• Solutions in intermediate compartments 0.05 M KC1, Uj ' 10.7 cm sec-'; in electrode compartments 0.7 M KC1 + 0.3 M KI + 0.1 M I, uel 26 cm sec- 25°C. Membrane size 50 x 45 cm. Spacers: ptrous polyethylene + dialyzer tubing. 62 Dennis Turner ENGINEERING ASPECTS OP GOLD PLATING POR ELECTRONIC DEVICES Dennis Turner Bell Telephone Laboratories, LTurray Hill, New Jersey, USA, Gold plating is used extensively in the preparation of electrical conductors on solid-state electronic devices.High quality deposits are required which must meet specific requirements of resistivity, hardness, porosity, smoothness and metal distribution on substra tes.The phosphate buffered gold cyanide bath was found to be best in meeting the engineering requirements of the deposit.The bath is well defined electrochemically.Efficient mass transport of gold cyanide ions to the cathode is important for rapid and uniform metal deposition.Novel techniques will be described to achieve uni form metal deposition both on a macro and micro scale. Jean Vereecken 63 THE EFFECT OF SUPERIMPOSED ALTERNATING AND DIRECT CURRENT ON THE INTERNAL STRESSES IN NICKEL DEPOSITS Jean Vereecken et René Winand Université Libre de Bruxelles, Service Métallurgie- Electrochimie, 50 av. F.D. Roosevelt, B-1050 Bruxnlles - Belgique. Plating is widely anplied to increase the resistance of metallic surface, to increase reflection power, to protect the base metal from corrosion, for decoration. However the properties of the deposits obtained by electrolysis are considerably influenced bv the internal stresses. These internal stresses are sometimes of considerable magnitude amounting for example to thousands kg/cm2 in the case of nickel. There are numerous factors which can give rise to an internal stress (1) ; these factors are : - occlusion of hydrogen - occlusion of foreign substances - alteration of the distance between the deposit crystals during deposition. Several studies tried to explain the influence of organic molecules on the internal stresses of electrolytic deposits (2). The effect of these organic molecules can be in several cases very important, but no obvious relationship was found between the nature of the organic molecules and their action on the internal stresses. There are many cases in which the internal stresses can be reduced considerably by carrying out deposition with a fluctuating current ,* for example an alternating current can be superimposed on the direct current (1). Our aim is to study the influence of a superimposed alternating current (frequency 50 H2) on the internal stresses of nickel deposits obtained on a copper cathode from a bath containing nickel sulfate NiF046Il2^ 330g/l, nickel chloride NiC^öI^O 4 5g/l and boric acid H3Bn3 3 0g/l. We investigated several electrochemical parameters such as the direct current density, the ratio alternating cur rent density/direct current density, the temperature and the stirring of the electrolyte. Amongst the numerous methods which have been developed for measuring internal stresses (1) we employed the strain gauge'method and the method of cathode bending. We concluded that without alternating current the inter nal stresses are minimum with a direct current densitv of 2 A/dm2. If we superimposed an alternating current with a fre quency of 50 Hz the internal stresses are minimum for Jac/^dc = 3. An increase of the temperature has a favourable effect 64 Jean Vereecken on the stresses but a stirring of the solution seems to have no influence. We also correlated the variation of the internal stresses with the parameters of the hydrogen evolution and with the structure of the deposit (form of the grains - orientation of the texture) to try to bring a contribution to the explanation of the nature of strains in metal deposits. 1- AT VAGFAMYAN, YS PETROVA The mechanical properties of electrolytic deposits - consultants Bureau - New York 1962. 2- RJ KENDRICK Trans. Inst. Met. Fin. 40 (1963) 2637. A. A. Wragg 65 IONIC MASS TRANSFER BY FREE CONVECTION WITH SIMULTANEOUS HEAT TRANSFER. A. A. Wragg and A. K. Nasiruddin Department of Chemical Engineering, University of Exeter, Exeter, U.K. and Department of Chemical Engineering, UMIST, Manchester 1, U.K. This work reports measurements of liquid-sol id mass transfer rates by free convec tion at upward facing horizontal plain surfaces in an unconfined fluid with simul taneous free convection heat transfer. The electrochemical technique is used in which limiting currents for the cathodic deposition of copper from acidified copper sulphate solutions at copper disc electrodes are measured. The results are corre lated in terms of a combined Grashof number for both mass and heat transfer. Apparatus and Experiments The apparatus consisted of a large rectangular perspex tank 40cm x 40cm x 46cm tall. Perspex plates holding the electrodes were fitted flush with the base plane of the cell and electrodes of 2, 4, 5, 10, 20, 30 and 51mm diameter were used. The electrodes were heated electrically by means of resistance wire cemented to the underside of the discs and electrically insulated from them. Two very fine chromeallumel thermocouples were embedded 1mm below thesurface of each cath ode, one in the centre and one of the edge, in order to measure the electrode sur face temperature, ts. Experiments were performed with solution concentrations of 0*005M, 0*01M, 0*05M, and 0«15M copper sulphate in 1*5M sulphuric acid. After fitting the required electrode into the cell the bulk liquid temperature, tb, was reduced to 10°C at which value it was maintained throughout the experiment. A polarisation curve was then plotted for isothermal free convection mass transfer. The electrode was then heated to 35 C and a new polarisation curve for combined heat and mass transfer was determined. Further readings were then taken for elec trode temperatures of 65 C and 91 C. In the case of the larger electrodes the max imum temperature attainable was 70 C for the 20mm and 45 C for the 30 and 51mm discs. Two sets of readings were taken for each electrode in each solution at each of the standard surface temperatures. Results A typical set of polarisation curves fordifferent values of ts is illustrated in Figure 1 showing the mass transfer enhancement due to the increasing thermal convection. Figure 2 depicts the trend in the value of the limiting current with ts-tb for various experimental conditions. A correlation of all the experimental results in terms of dimensionless groups using combined Grashof numbers as used for heat and mass transfer at vertical surfaces by Den Bouter et.al. and Marchiano and Arvia ' was attempted. Thus the following groups were calculated following den Bouter. Sh : Sc GR Gr + G = TO5 = W • m = m [wj \ 3 3 gde APc gde åPf where Gr = —~Ö : Gr = 3 : Pr = CpT/k 66 A. A. Wragg 200 400 600 00 20 40 mV t cb°c A logarithmic plotof Sh against Sc Grm yields the following correlating equations 25 9 Sh = 1-1 (Sc Grm)' (1) forScGR < 10 fi • ^3 9 and Sh = 0-155 (Sc Gr ) 7 (2) for Sc GR > 10 m' m In equation (1) the power on the Sc Grm group suggests laminar boundary layer and plume flowas found forthe isothermal case ' . In equation (2) the mass trans fer rate is independent of electrode diameter and the complex convection mech anism is probably one of thermals, interacting thermaIs causing turbulence and rad ial inflow at the disc edges '. Comparison with other work Some other equations for pure heat transferor pure mass transfer in laminar flow at horizontal surfaces are tabulated below. Author Equations «,„n Geometry Rotem and Claasen"'" Nu= 0*765 (Gr. Pr)n «_. Semi-inf. plate (theoretical) Clifton and Chapman'^ Nu=0*63 (Gr.Pr) n ^^Finite plate (theoretical) Kadambi and Drake11 Nu=0-795 Gr°^°Pr^]^Circu\arp\afe (theoretical) Kadambi and Drake '' Nu = 0'818 Gr^'^Opr Circular plate (experimental) 1 Fishenden and Saunders ^ Nu = 0*25 (GrPr)0-2n O5R Finite plate (experimental) Wragg Sh =0-64 (ScGr) Circularplate (experimental) .13 0-22 Band rows k i Sh =0-69 (ScGr) Circular plate (experimental) For combined heat and mass transfer solutions a re available forvertical plates such as those of Somers^, Wilcox and Mathers, Madden and Piret^ as followed by den Bouter et.al. Marchiano and Arvia^'^ give an independent solution resulting inadifferent formof thecombined Grashof number. Atheoretical so lut ion for com bined convection in horizontal boundary layer flow has recently been presented by Pera and Gebharr but is restricted to a Pr of 0*7 and low Sc values. A. A. Wragg 67 Results with nucleate boiling The surface temperatures of the 2, 4, 5 600 and 10mm electrodes were raised above the boiling point of the electrolyte so t • 125°C / s that local nucleate boiling occurred at the surface. A few polarisation curves : such as those depicted in Figure 3 were / t = 109°CV / s obtained for values of t«- of 109°C and 400 - 125 C. In some cases the mass transfer d • l-0cm rate at limiting current was 100 times that L /S for isothermal convection. This enhance // Cb - 0-15cm (ma) ment of mass transfer is caused by intense turbulence at the solution-electrode in 200 terface caused by bubble nucleation, - 91 •C.^/ growth, movement and collapse. With the larger values of t no plateau on the IIS > s i-E curves could be identified. This is / Isothermal v a novel way of producing high mass trans- III! _l • i ferrates in electrochemical systems with 0 200 400 600 800 out forced flow. Fig. 3 mV References ~T7 J. A. de Leeuw den Bouter, et.al. Chem. Engng. Sci. 23, 1185, (1968). 2. S. L. Marchiano and A. J. Arvia, Electrochim Acta, 137~1657, (1968). 3. S. L. Marchiano and A. J. Arvia, Electrochim Acta, 17, 741, (1969). 4. A. A Wragg. Electrochim Acta, K3, 2159, (1968). 5. A. A Wragg and R. P. Loomba, Int. J. Heat Mass Transf. 13, 439, (1970). 6. E. M Sparrow, ef.al. J. Fluid Mech. 41, 793, (1970). 7. L. Pera and B. Gebhart, Int. J. Heat Mass Trans. 15, 269.(1972) . 8. Z. Rotem and L. Claasen, J. Fluid Mech. 38, 173711969), 9. Z. Rotem and L. Ciaasen, Canadian J. Chem. Engng. 47, 461, (1969). 10. J. Clifton and A. Chapman, Int. J. Heat Mass Trans. 12, 1573, (1969). 11. V. Kadambi and R. M. Drake. Tech. Report Mech. Engng. HT-1, Princeton Univ. (1960). 12. M. Fishenden and O. A. Saunders, Engineering, 130, 193, (1930). 13. J. Brandowski, et.al. Chem. Stosow. Ser. B. 6, 455, (1969). 14. E. V. Somers, J. App. Mechanics. 23, 295, (T956). 15. W. G. Mathers, et.al. lndust. EngngT Chem._49, 961, (1957). 16. W. R. Wilcox, Chem. Engng. Sci. 13, 113, (1961). 68 CALCULATION OF LOCAL CURRENT DENSITIES AKD TERMINAL VOLTAGE FOR A MONOPOLAR SANDWICH ELECTROLYZER; APPLICATION TO CHLORATE CELLS. I,Rouäar,V«>Cezner, J*Hostomsky,M.M. Jakåic ,M.Spaso jevic and B.Z.Nikolic. Department of Inorganic Technology,Technical University, Prague,Czechoslovakia and Institute for Chemistry Techno logy ,and Metallurgy,Beograd,Yugoslavia* The present paper is a continuation of a series devoted to eiectrolyzers for chlorate pr o due tion [1 -4-] .Chemical engineering calculations are performed for a new type of a monopolar electrolyzer with power leads located on its sides[5] (sandwich-type electrolyzer). In a monopolar sandwich-type electrolyzer,the electrolyte is introduced at the bottom;a mixture of bubbles and ele ctrolyte leaves the cell at the top and enters the bubble separation section located above the cell.The current leads are located at the sides of the electrodes (Figs.l and 2)„The whole assembly can be devided into n identical cells.The total voltage tL, for one cell can be calculated using the assumptionfthat the lines of current flow are oeroendicular to the electrode sur£ace*Then the voltage distribution in the cell is given by local electrode T ootentials£v , &.,ohmic voltage drops in the electrolyte l M, and ohmic voltHge drops in the elect-nodes UA,U„* Using the procedure developed earlier[l-4] , thfe total voltage UT can be expressed in the form UT=(aA*bAln i)+(aK+bKln i) + 1^(1+0.75X3) + 2 S )+( S ^i{(2w /3) [(?A/SA>+(?K/sJ+2w P?AV A ^KV K?' (i is the aver, curr* density)»The terms at the right- -hand side of Eq.l. denote succesively the average anodic and cathodic potentials,the average voltage loss In th» gas-electrolyte mixture and the voltage loss in the electrode bodies.For the calculation of criterion K^, it is necessary to determine the flow rate of electrolyte V-y in a closed looo consisting of the electrolyser,bubble seoarator,reactor and cooler»The dependence between V™ and the total pressure drop in the loop can be obtained by an indeoendent measurement. References: loRousar I.,Cezner V.,RegnerA.:Coll•Czech.Cbem.Comnu 31,4193 (1966) 2oRoussr I.,Cezner V.:Ibid,32,1137 (1967) 3.RouSar I.,Cezner V.:Ibid,33,008 (1968) 4.Pouéar^1.:J.Electrochem Soc.,116,676 (1969) %Jaksic M.M* :Private communication- °snao jxtoj jo iCxqniassB aqq. jo MSTA doq. am, *z*%T& aezA'ioa^oaie ad^-qoiMpues 9V[% jo nao auo 'I'^T^ 69 70 THE DECOMPOSITION OF SODIUM AMALGAM IN A 'WIZONTAL DECOMPOSER J»Hostomsky,I.Pouåar,V"«Cezner Department of Inorganic Technology,Technieel University, Prague,Czechoslovakia In an industrial amalgam decomposer,the decomposition reaction Na(amalgam)+H2o=NaOH + 1/2H2 is carried out in a short-circuited galvanic cell Na(amalgam) / NaOH / H2(graphite). An element of the horizontal decomposer (i.e.with the horizontal amalgam flow)is shown in Fig»1.Along any closed sed line of electrical current,the reversible tension of the cell is compensated by the sum of electrode overvol- tages and of an ohmic loss in the electrolyte;the losses in the electrode materials and in the amalgam-graphite contact are negligible*The decomposition rate,which is equivalent to the electrical current I flowing between the anode and the cathode,can be evaluated by solving the Laplace equation for the potential?? in the electrolyte ( is related to a reversible hydrogen electrode in the same solution)•Boundary conditions are given in Fig.l; v is the sDecific resistance of the hydrogen-caustic mixture;U3is the reversible tension of the amalgam-hy drogen ceil calculated for the sodium concentration on the amalgam surface (i.e.U, includes a concentration overpotential due to the slow transport of sodium in amalgam).The decomDosition rate in the element is then The Laplace equation (1) and its boundary conditions ere numerically solved in a dimensionless form» I he decomposition rates were measured in a laboratory horizontal decomposer as a function of temperature,amal gam and caustic concentrations and amalgam flow rate;an example of results is given in Fig*2«Independently,the Tafel parameters for the hydrogen evolution on graphite were determined.Making use of the experimental decompo sition rates,adjustable parameters in the theoretical relationships were evaluated;average deviation between experimental and theoretical values constituted 13«7%« 71 3y U ^ = GfbTlog(lj|3) amalgam caustic graphite Fig»l»Element of a horizontal decomposer, TTT 1 !—I I I 1 I I T 1—TT (A/cm) 1 theory OS 0.2 - J I I I 1 11 J I I I 0.01 0D3 0.1 03 N(wt.VoNa) Fig<>2»Decomposition rate vs. bulk amalgam concentration; 80°C:20 wte« NaOH; amalgam flow rate 34cnrV(cm2.s) 72 SODIUM AMALGAM DECOMPOSITION IN A VERTICAL TOWER S.Rajasekaran,I.RÖuéar,J.Hostomsky,V.Cezner Department of inorganic Technology,TechnicalUniversity, Prague,Czechoslovakia» In a laboratory tower packed with graphite spheres,the reaction rate of sodium amalgam decomposition I\ia( amalgam) + HpO = NaOH + l/2Hp was measured as a function of amalgam concentration ana of amalgam flow rate for 80 C and 37»4wt»% NaOH.Theoreti- cal equations were derived describing the decomposition rate in the tower.The equations are based on the voltage balance between the reversible tension U of the amalgam- -hydrogen cell(which is the driving force of the decompo sition reaction)^electrode overvoltages "[, , "I,..,and ohmic losses in the bulk e lee tro lyte AIL.,, in thfe surface layer of electrolyte-gas mixture ^uV,T jfind in the contact resistance amalgam - graphite A IUT "UFa %+/V WE + ÄUEL + AUSL A schematic view of one current line is given in JFig.l. An approximate solution of the theoretical equations yielded analytical relationships for the decomposition rate,Tafel constants of hydrogen overvoltage on graphite, specific sui'face area of graohite packing ,and thickness of diffusion layer for sodium transport in amalgam» In comparison with earlier works of Hine (Electrochem. Technology 2,79(1964)),the present theory enabled the correlation~"of decomposition rates with flow rates of amalgam even in the region above the "optimum" flow rate» The adjustable parameters in the tneoretLcai equations were evaluated from the experimental decomposition rates and from the independently determined Tafel constants for the hydrogen overvoltage on graphite.An example of experimental and theoretical values is shown in .Fig.2. In comparison with other published data,it was found that the decomposition rates at lower amalgam concentrati on are determined by the sodium transport from the amal gam bulk to the amalgam surface.The absolute values of hydrogen overvoltage were found to decrease with increa sing concentration of sodium hydroxide. 73 GRAPHITE Fig.l, Schematic yiew of the amalgam droplet in contact with graphite and caustic. J£ P" T—rj T—TT T r O UJ 1 to a o 5: O EXPERIMENT • THEORY 0.4 l-L I I i J_L 10 10 1 N(wt.%Na) Fig.2•Decomposition rate vs0 bulk amalgam concentration; 17 . A wU% NaOH;80Ciamalgam flow rat 0*012 68 mol Hg /(cm .s) SYMPOSIUM on ACCELERATED CORROSION TESTING WITH ELECTROCHEMICAL METHODS THE 66th EVENT OF THE EUROPEAN FEDERATION OF CORROSION Norbert D. Greene 77 CONTROLLED POTENTIAL METHODS OF CORROSION TESTING Norbert D. Greene Department of Metallurgy, Institute of Materials Science, University of Connecticut, Storrs, Connecticut, U.S.A. Potentiostatic and other controlled potential techniques have been widely applied to corrosion studies during the past decade. The purpose of this presentation is to review these methods and to compare their advantages and limitations with more conventional methods of corrosion testing. Controlled potential methods are uniquely suited for investigating active-passive metals and alloys. Passivity is characterized by a dissolution rate which increases and then decreases as potential becomes more noble. Following this, dissolution rate generally remains low and constant (passive region) with further increases in electrode potential. Dissolution rate again increases with increasing potential at very noble potentials (transpajsive region) . The active-passive transition can be precisely described by the primary potential, E , and the critical anodic current density, Ic, which can be obtained from polarization data. Controlled potential methods are accomplished by means of a potentiostat, an electronic device. During the past decade, potentiostats have evolved from vacuum tubes to transitorized designs. The most recent innovations are composed of solid state, integrated circuit, operational amplifier modules. Potential step and pot tial sweep (i.e., potentiodynamic) methods are commonly employed during potentiostatic anodic polarization measurements. Unless carefully applied, both of these polarization techniques dan cause considerable distortion. Of primary importance is the voltage scan rate. This should be adjusted to match the intended application. Slow scans of 1-2 volts/hour are usually used for alloy evaluation studies while rapid scans (10-100 volts/ second) are employed for anodic protection current surge measurements. Today, anodic polarization methods are the basis for most corrosion- resistant alloy development programs. By comparing the case of passivation (i.e., passivity index) of different alloy compositions, their relative corrosion resistances can be estimated. Anodic protection feasibility is also determined by potentiostatic polarization measurements. Although pitting corrosion has been evaluated by various potentiostatic measurements, it appears that many of the observed correlations are fortuitous and are not consistent with electrode kinetics theory. Indeed, a major limitation of anodic polarization measurements is their general unsuitability for studies in halide-containing and/or near-neutral solutions (e.g., seawater). Controlled potential corrosion testing has been applied to a variety of corrosion problems. These tests are performed by potentxo- staticaily maintaining a specimen at constant potential during a conventional corrosion test. Applications of this procedure include investigations of intergranular corrosion, cathodic protection, Norbert D. Greene 78 selective corrosion and galvanic attack. As with polarization measurements» controlled potential corrosion tests are not generally- applicable to near—neutral, halide bearing solutions. Ya.M.Kolotyrkin 79" USE OP METHODS OP ELECTROCHEMISTRY AND RADIOACTIVE INDICATORS POR DETERMINING LOW CORROSION RATES Ya,M.Kolo tyrk in Karpov Institute of Physical Chemistry, Moscow, USSR Electrochemical methods do not always allow an estimate to be made of the metal dissolution rate, nor do they provide sufficiently full information on the corrosion process. Puller information can be obtained by combining electrochemical measurements with an analysis of the so lution for dissolution products. Among all the methods of analysis, the most sensitive is the method of radio active indicators. The paper examines the methodical aspects of the use of radioactive indicators in corrosion studies. Tests are usually made of a sample marked by tf -isotopes of the elements comprising it. Its dissolution rate is determi ned by the rate of transition into the solution of the radio-active components. The analysis can be conducted in regard to the "tf -isotopes of another element, if its transition into the solution takes place at the same comparative rate as the defining element. The simplest and most convenient method of introducing y-isotopes into the sample is by its neutron radiation in a nuclear reactor. Methods have also been developed for corrosion studies using pure /3 -radiators and the neutron-activa tion analysis of the dissolution products. Methods of scintillating Y« spectrometry are the most feasible for the analysis for' ^ -isotopes. Apart from enhancing the sensitivity and reliability of the analysis, this ensures the possibility of determining simultaneous ly the partial dissolution rates of the components of the alloy without preliminary separating the elements in the samples of the solution. The condition for the application of the method of radio active indicators is the absence of effects on the cor rosion process from radiation. As a rule, changes in the structure and properties of the sample caused by preli minary radiation and the effects of placing a radio-acti ve metal in an electrolyte are small and have no substan tial influence on the dissolution process. Methods of corrosion studies with the use of marked samp les are highly universal. Their very great sensitivity makes it possible to cut down the testing time and even in the case of extremely stable materials to study the kinetics of the initial stages of corrosion, the 80 Ya.M•Kolo tyrk in transition of the individual components of the alloy into the solution, and their accumulation on the surface of the sample. The data on the change in the corrosion rate in time are important not only for understanding the me- chani3m of the corrosion process, but in some cases for avoiding mistakes in determining the stationary dissolu tion rate of the metal. Studies of the early stages of dissolution are facilitated by using methods of continuous automatic registration of radio-isotopes in the electro lyte. The possibilities opened up by combining methods of elec trochemistry with radio-active indicators are illustrated by a number of examples taken from practical corrosion studies. The paper, for instance, gives data on the in fluence of the composition of the solution and the condi tions of polarisation on the corrosion resistance a pla tinum anode. The role of surface oxygen compounds of the metal in the dissolution process is discussed. The corro sion behaviour of platinum is compared with that of rhu- thenium. The results are given of the determination of the partial dissolution rates of the components of a number of high ly-alloyed metals testifying to the possibility of their non-uniform transition into the electrolyte. The selecti vity of the dissolution of hard solutions depends on the potential and, to a great extent, is determined by the corrosion properties of the components- The paper discus ses the data on the influence of therm0.! treatment of the kinetics of the dissolution of individual components of the alloys, which permits conclusions to be drawn on the mechanism of localised types of corrosion, and, specific cally, intercrystallite corrosion. It is shown on the example of electrically refined indium that comparative dissolution rates of metallic microadmixtures are deter mined not only by the degree of their electronegativity in regard to the aain metal, but also by the kinetics of the corresponding electrochemical reactions, as well as the diffusion rate of the microadmixtures in the sample to the metal-electrolyte boundary. Results are also given of measurements of non-stationary dissolution rates of passive iron in acid mediums in the transition period after a sharp change in the electrode potential or acidity of the solution. An analysis of the obtained results leads to the conclusion that the disso lution of iron in the passive range takes place accor ding to the electrochemical mechanism, not through the stage of the chemical dissolution of the surface phase filnu Milan Prazék 81 THE POLARIZATION RESISTANCE METHOD OF CORROSION TESTING Milan Prazåk G.V.Akimov State Research Institute for the Protection of Materials, Prague, Czechoslovakia. The plenary lecture refers to problems concerning the me thod of metals instantaneous corrosion velocity estimation by measuring their polarization resistance (resp.impedance or admittance). There are more reasons to expect that the method could be used as a common laboratory and plant test method. The base of the method was formulated by Stern and Geary(1) but some partial electrochemical problems arose which are limiting the possibilities and accuracy of the method. The discussion deals with following problems: 1. The usefulness of measuring the Tafel slopes /3a and /3c for evaluating the corrosion velocity (as icor or v^ froD1 the Rp value according the Stern's formula, m comparison with the evaluation by means of the analytically estimated relation of the empirical type v = B/Rp^ (16). The results (see Tab.l) do not seem to show the preference of the first way, which is moreover more complicated. 2. The effect of the exchange currents of the partial ano dic or cathodic reaction taking place simmultaneously with the corrosion reaction, as discussed by Prafcåk (16) and Mansfeld and Oldham (17). The possibility of eliminating errors caused by this effect in specific cases seems to have little perspective. 3. The causes and effect of spontaneous slow and fast fluc tuations of the corrosion potential Ecor occuring during the corrosion process. Elimination of the effect of slow fluctuations either by the use of alternating measuring current, or by a mathematical correction for the time de pendence of Ecor according Riggs (18) during permanent d.c. polarization. 4. The elimination of the ohmic resistance in poorly con ducting electrolytes either as proposed by Jones (19) for the three-electrodes connection, or by substracting the oh mic resistance component estimated at higher frequencies in the two-electrodes connection, or by a proper geometrical arrangement of the (two-) electrodes system. 5• The effect of the electrode capacity of the corroding metal, especially with respect to the application of the a.c. measurements of Rp. The high pseudocapacity of electro des coated by solid corrosion product films (as Fe in neut ral solutions)• In addition new perspectives of the method are discussed, especially for the region of low corrosion rates (icor < l^A/cnr), based either on evaluation of the char- Milan Pra£ék 82 ging curve E-t according Jones and Greene (2C), or on the impedance or capacity measurements according Schwenk and Biihler (8) . The possibility of application of the last method is probably based on the constancy of the RpC pro duct (the time constant) of corroding electrodes. References (tothe Abstract and Tab.l) 1. M.Stern,M.Geary : J.Electrochem.Soc.104, 56, 390, 561, 645 (1957) 2. S^Uneri: Communications de la Faculté des Sciences de 1'Université d'Ankara, T. 16 B, 37 (1969) 3. P.Neufeld; Corrosion Sci.,4, 245 (1964) 4. G.A.Marsh; Proceedings of the Second International Con gress of Metallic Corrosion, New York 1963, p.936. Ed. NACE, Houston 1966 5. L.I.AntropoVjM.A.GerasimenkoJ Zashchita metallov 2,115 (1966) 6. A.Cohen,R.V.Jelinek: Corrosion 22, 39 (1966) , 7. M.Pra2ék,B.Eremiéä,E.Kårnikova :Techn.rep. SVUOM 15/1970 8. W.Schwenk,H.-E.Buhler: Corrosion Sci.3, 261 (1963) 9. E.Fot,E.Heitz: Werkstoffe u.Korr.18, 529 (1967) 10. T.J.Butler,P.R.Carter: Electrochem.Technology 2> 157 (1965) 11. S.Evans,E.L.Koehler: J.Electrochem.Soc.108, 509 (1961) 12. J.Beran,J.Sulc: Techn.rep.äkoda ZJE 96/1970 , 13. M.Pra2åk,B.Eremiåä,E.Kårnikové: Techn.rep.SVUOM 65/1968 14. M.E.Indig,C.Groot; Corrosion 2j>, 171 (1970) 15. J.Beran,M.Krausové, J.Sulc: General Proceedings of the Conference on the Use of Zirconium Alloys in Nuclear Reactors,p.345 (in Czech). Mariånské Lézné 1966 16. M.Pragék: Werkstoffe u.Korr«19,, 845 (1968) 17. F.Mansfeld,K.B.01dham: Corrosion Sci.ll, 787 (1971) 18. O.L.Riggs Jr.: Corrosion 26,, 243 (197ÖT 19. D.A.Jones J Corrosion Sci.8, 19 (1968) 20. D.A.Jones,N.D.Greene : Corrosion 2J2, 198 (1966) Milan Pra2ék 83 Table 1. Relations for calculating the corrosion velocity from the polarization resistance Rp H • Relation Rp-cor.velocity v Comments CO p -P -P CD Rp in Jl.cm S cd icor = B/RP • •^ U x B = 26 mV /3s=120mV, /3c=120mV CD B = 52 mV /3a=l20mV, /3C= °° «H B = 17.4 mV /3a= 60mV, /3c=120mV H3 2 B = 26 mV compilation 3 B = 90 mV organic acids 4 B = 75 mV neutral solutions 5 B = 16-45 mV (20 recomend.) inhibited acids 6 B = 36 mV 54% LiBr, with inhibitor B = 20 mV without inhibitor 7 log v = 2.73 - 1.07 log Rp acid solns.;a.c. 0.25 Hz acid and neutr.sol.:d.c. v = 310/Rp approx. L at 10-* Si . cm2) 2 v '» g/m .h - mm/yr ? B = 30 mV approx. ( at 10? Si . cm^) r Ci> 8 B = 18 mV dil. H2S0A, d.c. (D 4 B = 4.9mV; 2.6mV; 1.3mV a.c.,f = 30;50;100 Hz (0 H 9 B = 22-41 mV 3% NaCl f 10 log v = 3.74 - 1.01 log Rp dil. H2SO4; v = mpy 1 3 B = 12 mV approx. 11 B = 29 mV pH = 2.5 - 3.5 12 v = (700+140)/Rp 40°C; v = mm/yr B = (64±13) mV c ale. 13 log v = 2.602 -1.05 log Rp pH=0.3-12; a.c; 0.25 Hz. v=g/m2.h; no HF present 2 f v = 800/Rp v=mm/yr; (at lO^il.cm ) 4 < B = 75 mV calc.; -•»- log v = 2.640 -1.02 log Rp pH=0.3-12; d.c.; no HF v = g/m2.h log v = 5.70 - 2.05 log Rp pHO;a.c.0.25Hz; HF pres. 1 i log v = 5.60 - 1.84 log Rp d.c.; v=g/m2.h; -»»- 14 B = 75.9 mV Zircalloy-2; LioH; 288°C * 15 v = 400/Rp ZrNb2.5; HNO3+H2SO4+HF tSJ v = mm/yr B = 35 mV |calc. 84 Ludwik Chromy THE RAD10GRAPHICAL METHOD OF THE EXAMINATION OF PERMEABILITY OF PAINT COATING MODIFIED WITH CORROSION INHIBITOR L. Chromy and A. Zimnoch Instytut Badawczo-Projektowy Przemysiu Farb i Lakierow, Gliwice, Poland and Uniwersytet Slaski, Katowice, Poland. The paint coatings have been examined with and without the addition of benzyl quinoline iodide used as corrosion inhibitor of steel. The coatings have been prepared from two polywinyl resins; copolymer of vinyl chloride and isobutyl ether of polyvinyl acetate and copolymer of vinyl chloride and vinylidene chloride. In order to determine the degree of film permeability, the isotopic method has been applied for measurements of the radioactive &<*Cu deposits in the pores of coatings tested. Compassing the results obtained it has been found that the anticorro- sive properties of protective coatings are better when modified with inhibitor. Additionally, the anticorrosive properties of the above mentioned coatings have been determined by means of electrochemical measure ments and compared with isotopic method The following parameters have been determined changes of electrode potential and current intensity calculated as a function of current density against time. The measure ments have been carried out with the addition of inhibitor and also without inhibitor in the presence of corrosion media at pH: 2, 4, 8, 10. Some relationships between the protective properties determined by means of accelerated electrochemical tests and the permeability of coatings examined by radiographical method have been estimated. The measurements of inhibition have been carried out by electroche mical method for five inhibitors:benzyl quinoline iodide, phenyl tiou- rea, triphenyl arsenic oxide, di-phenyl di-tiophosphine acid and natrium diethyl di-tiocarbamate. However, the isotopic method has been used for testing the best of them, namely, benzyl quinoline iodide. 85 J.M. Defranoux PITTING POTENTIAL MEASUREMENTS ON STAINLESS STEELS, AS COMPARED WITH REST POTENTIALS .AND THE ELECTROCHEMICAL BEHAVIOUR OP PASSIVE FILMS . J.M. Defranoux Sté Ugine Aciers (Groupe Péchiney Ugine Kuhlmann), Centre de Recherches Métallurgiques, 73-Ugine, Prance . Pitting potentials are now widely accepted as a valuable method of assessment of the resistance of stainless steels to pitting corrosion in halide media . Some recent findings show however, that it is sometimes necessary to take also into account the rest potential in the very conditions of the test : under certain circumstances, this rest potential can in fact rise notably and overtake the pitting potential. In such cases, pits initiate and grow spontaneously, without need for a concomitant lowering of the pitting poten tial . This applies particularly when the only oxidising species in solution is dissolved oxygen (sea water, de-icing products, etc). In that case, tests have shown that the kinetics of the oxygen cathodic reduction do not vary, at least in a quite wide range of steel compositions, surface treatments, etc; the rest potential is then determined by the anodic reaction, i.e. by the ease with which the anodic dissolution current can filter through the passive layer ("ionic transparency") . It is then possible to give further details on the mechanism of the phenomena described by Bianchi (l), as passive layers are heated in air at 15O-3O0°C, the resistance of stain less steel to pitting being thus notably damaged , It has been shown furthermore that such an impairment of the resistance to pitting can also result from a passivation treatment, e.g. in nitric acid, when the metal-solution poten tial during the treatment is too high, i.e. the nitric acid solution has too high a concentration . On the other hand, it has been attempted to find'a correlation between such properties and the electrochemical behaviour of redex systems, e.g. ferri-ferrocyanide, the stain less steel acting as an inert electrode, following the method used firstly by Meyer (2) for passive zirconium. The redox kinetics.give informations on the transfer of electrons between the metal and the solution through the passive film ("electro nic transparency") and the electrochemical double layer. It is thus hoped to get some insight into the semi-conducting pro perties of the film, and into the localization of potential differences, either in the passive film or in the double layer. When potential differences are entirely localized in the dou ble layer, the electrochemical kinetics do not correspond to phenomena occuring in the passive film and one cannot infer of them conclusions about the "bulk" properties of that film. The 86 J.M. Defranoux localization of the potential difference in the electrochemical double layer may be ascribed to the existence of electronic surface states on the passive film . (1) G.Bianchi et al., Corros. Sci., .1(0(1970) 19 (2) R.E. Meyer, J. Electrochem. Soc, W£ (i960) 847 Michel Froment 87 INTERGRANULAR CORROSION OF AUSTENITIC STAINLESS STEJv ^ TN THE SENSITIZED OR SOLUTION-ANNEALED CONDITION. X XX • • XXX André Desestret , Michel Froment , Pierre Guiraldenq Centre de Recherche de Creusot-Loire, 42 Unieux, France xx • • • Physique des Liquides et Electrochimie, Groupe de Recherche du CNRS, associé å 1'Université Paris VI, 11 quai St-Bernard, 75 Paris 5e XXX Laboratoir• e de Métallurgie• , Ecole Centrale de Lyon, 69 Ecully Intergranular corrosion tests for austenitic stainless steels are carried out by exposing test-pieces in acidic solutions (Huey, Monypenny-Strauss, Streicher test). These tests are rather time- consuming because they very often last several days and their results can be difficult to interpret for two main reasons : the complexity of the redox equilibria, which can be altered by progressive variations with time of the solution composition ; and the resulting variations of the metal potential due to the solution variations. On the other hand, depending upon the alloy composition, the intergranular attack can occur at different potentials. Thus, the comparison may be erro neous when a solution maintaining a fixed redox potential is used (1). Tests based on pote.ntiokinet.ic polarization curves are of interest as they allow covering a wide range of potentials. Furthermore, they can be carried out in a short period of time. During the lecture, we shall describe such a test applied to austenitic stainless steels con taining 18 % chromium, 12 % nickel, and from 0.009 % to 0.06 % carbon. Samples are either annealed (20 minutes at 1150°C, then water-quenched) or sensitized at temperatures between 500°C and 900°C. Polarization curves are obtained with a scanning rate of 9 volts per hour ; the electrolyte is a warm 5N sulphuric acid solution (70°C), agitated by bubbling nitrogen. Curves obtained when the potential is increased (upwards scan) do not show any evidence of sensitization. However, during the subsequent downwards scan, a current peak presumably related to the sensitization appears in the active domain. This last peak is found only when the samples have been maintained at a sufficiently high potential in the transpassive domain (+1300 mV/H»). Figure 1 shows polarization curves of steels with different carbon and titanium contents, the steels having been sensitized at 800°C for 8 hours. The area of the re activation peak increases considerably with increasing carbon contents. On the other hand, a titanium stabilised steel (Ti : 0.48 %) contain ing the same amount of carbon (C : 0.06 %) shows a smaller peak ; the polarization curve of this last steel exhibits also two other peaks related to the dissolution of titanium carbide and sulphur-rich phases, probably titanium sulphide (2). A systematic series of experiments performed with samples different compositions and different heat treatments shows the peak area to be closely related to the sensitization treatment. During the lecture, examples will be presented showing this test to be more sensitive than the conventional tests (i.e. Huey test). The test we have just described is based upon the preferential anodic dissolution of chromium carbides. This dissolution occurs when the sensitized samples are exposed to elevated potentials in the transpas- 88 Michel Froment sive domain (2). The initial carbide dissolution causes an increase in the area of chromium-depleted zones which undergo a strong prefe rential etching in the active domain during the subsequent downwards potential scan. The preferential dissolution of chromium carbides at high anodic po tentials may be used in the Huey test in order to detect the sensiti zation of a steel in a relatively short time. Due to the dissolution of chromium carbides, the metal-electrolyte interface is enriched in Cr ions which are oxidized to C^Oy in the concentrated acid solu tion. The reduction of dichromate ions must then replace the reduction of nitric acid as the cathodic reaction in such a way that the poten tial of the sensitized steels is shifted to higher values. Therefore, we followed the evolution of the potential measured against a reference electrode (S.C.E.) of steel samples immersed in a boiling 65 w/o nitric acid solution. An important decrease of the potential during the first ten hours of the test was generally observed for the annealed steels. As is shown in figure 2 for the steel with 0.049 % carbon, the potential decreases from +870 mV/SCE to +650 mV/SCE after 12 hours (curve 1). In most cases we observed that the potential reached a steady value between +650 mV and +700 mV. However, vith sensitized steels, there is a much less important decrease of poten tial. For instance, curve 2 of figure 2 shows that after a 20 minute treatment at 700°C the potential remains in the vicinity of +800mV/SCE. Depending upon the degree of sensitization the potential steadies out 'values between +725 and +825 mV/SCE. The initial rapid fall of poten tial is attributed to the rapid formation of N0£ from the reduction of HNO3. While the steel remains at a high potential at the beginning of the test, a transpassive attack occurs during which chromium car bides dissolve preferentially. Moreover, the same initial transpassive conditions can explain the preferential etching of £tain boundaries even in fully annealed steels. This last type of corrosion is probably related to the segregation of certain elements along again boundaries (3). We shall give some results pertaining to annealed steels containing 16 % chromium and 14 % nickel: we shall describe the role of elements such as carbon and silicon. Polarization curves recorded in sulphuric acid solutions exhibit a strong increase of the dissolution current in the transpassive domain as the carbon or silicon content is increased. Experiments with radio tracers show that the volume self-diffusion coefficient of Fe 59 increases when both of these elements are present. This means that silicon and carbon increase the concentration of point defects in the matrix and»consequently, the anodic dissolution rate. As shown by replica and transmission electron microscopy, the preferential anodic dissolution of grain boundary regions suggests a segregation of carbon and silicon in these regions. Besides, the activation energy of inter- granular self-diffusion of Fe 59 increases strongly when the steels contain large quantities of carbon or silicon. (1) A. DESESTRET et coll., Mém. Scient. du CETIM, sous presse. (2) VI. CIHAL, A. DESESTRET, M. FROMENT, G.H. WAGNER, Comm. a la 21e reunion du CITCE, Prague, octobre 1970, Resumes étendus p.287. Michel Froment (3) A. DESESTRET, M. FROMENT, P. GUIRALDENQ, Comm. au 4e Centres International de la Corrosion, Amsterdam, 1969. E(mV/H2) i » 0,017% C 0,023 % C 0,0<1 % C 0,06% C 0,059%C etO,48'/.Ti -J( A/0,5 cm2) Figure l EmV/EC5 . ann v^ 7nn J£ f 600 t heures 10 20 30 ELECTROCHEMICAL I-IETHODS FOR CORROSION EVALUATION OP LOW-ALLOY STEELS, COPPER ALLOYS AND STAINLESS STEELS Lucio Giuliani, Alberto Tamha e Roberto Bruno Centro Sperimentale Metallurgico - 00129 Roma (Italy) Electrochemical methods allow to obtain either absolute values of corrosion rates or comparative behaviours of similar materials in reference to peculiar corrosion properties. When actual corrosion rates are studied, the reproduction of the sy stem metal-environment is of paramount importancej on the other hand, when metallurgical compositions or comparative performances have to be evaluated, the most differentiating electrochemical conditions must be chosen. Some electrochemical procedures have been esta blished in our laboratory in order to screen new metallurgical com positions of self protective low-alloy steels for atmospheric exp£ sure, to evaluate copper alloys for desalination plants and to asses the quality of stainless steel surfaces, Low-Alloy Steels The most relevant property of low-alloy steels is -their capability to self-protection through compact and adherent rust layers. The nuclea-tion and growth of such rust layers are strongly dependent on the initial corrosion rates and thus on alloy composition at the interface metal-rust. The progressive potential ennoblement and the coverage passivability have been evaluated in several electrolytic solutions by means of potentiodynamic scans. The most relevant results have been obtained in Ea^SO* 0,1 M either on as-received surfaces and on artificially or naturally corroded surfaces. An evaluation criterium is represented in Pig, 1, where the current ratios Ia/lp are derived from potentiodynamic scans on polished and artificially corroded specimens (Fig, 2), The significance of such methods has been checked on some steel gra des with known atmospheric performances. Copper Alloys The spreading of desalination plants during the last years has prompted researches on copper alloys for tubing. The corrosion of copper in saline environments such as neutral brines depends on the supply of oxygen to the metal surface* This fact limits the validity of long range potentiodynamic scans and focuses the atten tion on the processes which occur in the neighborour of the corro sion potential. Micio Giuliani 91 The presence of electrochemical reactions other than Cu —• Cu+^ + 2é requires some modification of the linear polarization method (pola rization resistance)» For comparative purposes, the determination of anodic slopes, as a function of solution velocity and contamination, has proved meaning- ful. Pig. 3 shows the anodic slopes of Cu70Ni30Fe and Cu70Ni30Cr (UI- 732-X) at water velocity up to 4 m/s. Stainless Steels The corrosion 'behaviour of a stainless steel depends in a certain extent on its surface finish : this is particularly found when an insufficient pickling leaves a dechromized layer on the surface. The differences in surface finish are relevant to the stability of the passivation film. A comparative method for evaluation of stainless steels surfaces con stists in computing the ratios betveen passivity "breakdown times and passivation times obtained from potential decay curves for specimens anodically polarized in a suitable solution. An application of this method to AISI 410 stainless steel in 0.01 N H2SO4 is shown in Pig. 4» The compositional variations of the surface alloy responsible for the different depassivation kinetics have been detected by means of electron microprobe analyser. d.crtailng Incr«»»ing pstlniblllty patimbllity" Rg. 1 - Current ratios of anodlcally corroded (I ) and polished speciaens (T ) vs. potential showing low alloying effect on steel ueatherability. -1050 92 lucio Giuliani 500 ,"' Anodicatly o As-received rf S corroded & Pickled jf rS * Met. Polished _/0-^^ 0} e D Electropolished _^S^*S^^ i- c 250 >to W aID a> O n 1 1 i i Current density 50 100 Passivation Time (sec.) Fig. 4 - Passive fila breakdown tiaes vs passivation tiaes for A!SI steel Type 410 in N saturat^ ed 0.01N H S0t at 30 °C. 2 4 2 a/s 4«/s/ \«^ c Anodically 510 corroded o ex. ( b) ^ 102 50 Current density -S I 10 Fig, 2 - Tlpical anoitc behaviour in Ka S0t 0.1 H of a weathering (a) and a non weathering (b) stesl. -300 -200 .100 Potential oY (Ag/AgCJ) Fig. 3 - Vat«r valocity eff«ct on the anodic slope of 70 Cu 30 KiFe ( ) and 70 Cu 30 HiCr(lR« 732 - X) ( ) alloys . J.A.S. Green 93 ACCEUJRATED STRESS-CORROSION CRACKING OF Ti-8$41-l#Mo-l$V IN THE PRESENCE OF A HYDROGEN-RECOMBINATION POISON J.A.S. Green and A. J. Sedriks Research Institute for Advanced Studies, Martin Marietta Corp., Balti more, Md. 21227. Since 1965, when the transgranular stress-corrosion cracking (SCC) of high-strength titanium alloys exposed to aqueous chloride solutions was first reported, it has been established that susceptibility to transgranular cracking is associated with increased aluminum and oxy gen contents in the alloy, and heat treatments which favor the forma tion of the intermetallic, Ti-,Al, and that susceptible alloys exhibit coplanar dislocation arrays after deformation. The environmental fac tors associated with the transgranular cracking proce?.'.-, on the other hand, are not well understood. In fact, considerable controversy still exists as to the critical species in the failnre process. In aqueous chloride solutions Beck considers that the critical species are chloride ions, and that cracking occurs at a rate governed by the transport of these ions to the crack tip. Alternatively, Scully and co-workers have suggested that cracking results from some form of hydrogen embrittlement. Hence, a major issue to be resolved In the transgranular SCC of high-strength titanium alloys exposed to aqueous chloride environments is whether the primary embrittling species is the chloride ion or some form of hydrogen. To date, attempts to answer this issue have been inconclusive. Accor dingly, the objective of this work is to distinguish between the hydro gen and chloride models by testing a susceptible alloy in environments which independently favor attack by one or other of these species. In the present studies, the possibility that some form of hydrogen is res ponsible for cracking has been examined using electrolytes containing an hydrogen-recombination poison, i.e. arsenic. It is well known that the absorption of hydrogen by metals is increased by the presence of catalyst poisons in the electrolyte. Accordingly, if the cransgranular SCC of high strength titanium alloys is associated with the presence of hydrogen at the crack tip, the introduction of a hydrogen recombina tion poison in the solution should accelerate cracking. Tests were carried out on a commercial Ti-8 w/o Al-1 w/o Mo-1 w/o V alloy in the mill annealed condition. The composition and mechanical properties of this alloy were as follows: Composition: Al Mo V Fe N O C H Ti "(weight ^T~ 8.0 1-05 0-98 0.l6 0-001 0.008 0-03 COO1!- Balance Mechanical Properties: U.T.S- = 139,^00 psi, 0-2$ Y.S. = 12l+,000 psi, Tensile Elong. on 2" = 9-3$. Two types of stress-corrosion tests were used in the present study, subsequently referred to as the Instron test and the cantilever-beam test. In the Instron test, which measures crack velocity, tensile specimens 1 in. wide and O.06 in. in thickness containing a machined notch 1/8 in. deep with a 60 deg. included angle and 0*001 root radius were used. Specimens were tested in an Instron machine at a cross- head velocity of 0.002 in. per min in the presence of the environment. The cross-head motion was continued during crack propagation. The 94 J.A.S. Green velocity of cracking was measured by noting the elapsed time between the formation of a visible crack at the root of the notch, and the moment it passed a mark, previously scribed on the specimen at a distance of 2mm from the root of the notch. Luring testing, the poten tial of the specimen was potentiostatically controlled. The cantilever beam test measures the minimum fracture toughness in the presence of the environment, i.e. Kj/gcc* Specimens of dimensions 9 in. X 1 in. were cut in the long transverse direction from l/k in. thick mill-annealed Ti-8-1-1 plate. The specimens, then, were edge- notched with a 1/8 in. deep notch which contained at 6o° included angle and had a root radius of 0.001 in. To obtain a base-line value for the fracture toughness of the alloy, designated Kj/X, several specimens were loaded to failure in air. Other specimens, then, were loaded to various fractions of this base-line value in the presence of the environment, and the time to failure was recorded. Many of these tests were carried out under potentiostatically-controlled conditions. Standard formulae were used to convert the load and crack depth data to the stress intensity parameter, K. For convenience, the data ob tained in the presence of the environment, KJ/J_, and the minimum frac ture toughness obtained in the presence of the environment, Kj/scC; are presented as a fraction of the fracture toughness value in air, i.e. as the ratio %/i/^i/X' Using the Instron test, the relationship obtained between crack velo city and applied potential for the Ti-8-1-1 alloy in an aqueous solu tion containing 3-5% NaCl and 10 ppm As, added as sodium arsenite, NaAsOgjis shown in Fig. 1. Despite the fact that the velocity of cracking was lowered by the presence of As, indicating a less aggres sive environment, the load at which crack propagation commenced was ~ 30% lower in the arsenic-containing solution than in the plain salt solution. Furthermore, examination of the fracture surface by scan ning electron microscopy revealed larger stress-corrosion regions on specimens failed in the arsenic-containing solutions. Tests using the cantilever-beam technique, in which the potential was kept constant at -500 mV vs S.CE., showed that the addition of 10 ppm As lowered the value of K-j^Mnn "by ~ 20%, Fig. 2. This demonstrates that the addition of arsenic under controlled potential conditions in creases susceptibility to cracking. However, it was also found that under_open circuit conditions, the addition of arsenic did not lower •^I/SCC* "*"n •^ac'^} ^e stress level at which cracks propagate in arsenic- containing solution under open circuit conditions is ~ 10% higher than in the equivalent plain salt solution, Fig. 2. The observations that Kj^/cnn i-s lowered by the presence of arsenic at -500 mV vs S.E.C. and raised under open circuit conditions may be interpreted in terms of the equilibrium potential-pH diagram for the system arsenic-water. This diagram indicates that, under the condi tions employed in the controlled potential tests, arsenite would be re duced to metallic arsenic, and that this reduction is particularly favored under the acidic conditions which are known to exist in the environment adjacent to the tips of stress-corrosion cracks. Thus, under these conditions metallic arsenic would always be available at J.A.S. Green 95 the titanium surface to interfere with the hydrogen recombination re action. For the tests conducted at open circuit, however, meta"1 '.\c arsenic would not necessarily remain as the stable species on t. metal surface. In fact, under the highly reducing potential cor. .1- tions caused by the rupt ire of the oxide film at the titanium surface a hydride of arsenic, such as arsine, ASH3 would be the stable com pound. The formation of arsine, which consumes hydrogen, would serve to minimize the quantity of hydrogen available at the crack tip. Thl: may account for the observed increase in K^/SCC Fi§* 2J under open circuit conditions. ' These observations lend support to the view that the transgranular stress-corrosion cracking process in high-strength alpha-titanium alloys in aqueous solutions is a manifestation of some form of hydro gen-induced embrittlement. It should be noted, however, that since ar.\senic is thought to increase both the amount of absorbed hydrogen at the surface and the quantity of hydrogen entering the metal, the present observations do not distinguish between a mechanism requiring the entry of hydrogen into the lattice and one requiring its adsorp tion at the crack tip. It would be premature, however, to dismiss the role of the chloride ion on the basis of these observations. It is quite possible that the chloride ion may play an important subsi diary role either in aiding the ingress of hydrogen into the metal or in facilitating its accumulatior at the crack tip. -1—r 1.0 20- Ti-8% Al-1% Mo-l%V, Ti-8%AI.1%Mo-l%V 1 3.5%NoCI in HjO -500mV vj S.C.E. 15 0.8 3.5% NaCI and lOppm A», O y open circuit potential Z 2 KIi /KIx 1 1 _— 1 > 1 -J L .0.6 -0.6 -CM -0.2 0.0 10 100 APPLIED POTENTIAL, V v» S.C.E. TIME TO FAILURE, min Fig. 1 Fig. 2 96 Lajos Hackl ACCELERATED METHOD FOR THE INVESTIGATION OF THE SUSCEPTIBILITY OF STAINLESS STEELS TO INTERGRANULAR CORROSION BY POTENTIODYNAMIC POLARIZATION Lajos Hacklt G.M\iller|+ J.Horvåtht F.Marta+ +Institute of General and Physical Chemistry of the .University of Szeged, SZEGED, Hungary Materials Testing Laboratory, BUNA Works, Schkopau,GDR. Previous theories explaining the reasons of inter- granular corrosion of stainless steels are in agreement with the fact, that the susceptibility of the steels to this special corrosion are in.close connection with the new phases forming at the grain boundary of the solid solution and especially with the segregatum taking place at a certain heat-treatment. Furthermore these theories agree in the question, that the occuring corrosion process is an electrochemical one. As a consequence of the different chemical composition /decreased chrome content of the segregatum/ dissolution takes place along these zones. This process may be promoted by the stress appearing along the precipitating particles. As it comes from the electrochemical character of the corrosion process, it can be studied by a potentiostatic polarization. Considering the surface proportion and circumstances of zones, the polarization curve of the sensitized steel is composed of the curves with different chrome contents. Thus the presence of any zone, deficient in chrome, has an effect on the polarization curves. On the basis of the investigations of EDELEANU, PRAZÅK, T0MASHOV, TSHERNOVA as well as of others, this effect may be considered as the greatest in the passivation range. However, the observation of these changes is very difficult or even impossible, if we measure the potentiostatic or the potentiodynamic polarization curves starting from the more negative potentials through the active state of the whole surface. TOMASHOV and TSHERNOVA had proved experimentally, that starting from a potential range which ensures the passivity of the whole metal surface and proceeding toward the more negative potentials, the reactivation of the working electrode susceptible to intergranular corrosion takes place before the activation of the whole metal surface. Accordingly, the polarization curves obtained in the presence and in the absence of zones deficient in chrome are different from each other. According to this, it seemed possible, that the susceptibility of stainless steels to intergranular LajOP Hackl 97 corrosion can be definitely detect by electrochemical methods. Potentiodynamic polarization measurements were carried out at different austenitic stainless steels stabilized with titanium. In agreement with the experimental results of TOMASHOV and TSHERNOVA we established, that starting from the potential range, corresponding to the passive state of the whole surface, the electrodes were polarized towards the active region, reactivation took place at more positive potentials for steels which were susceptible than for those which were insusceptible to intergranular corrosion. The value of the reactivation potential was defined as the voltage at which the increase of current started. We examined the influence of surface pretreatment, composition of electrolytes, starting potential and potential sweep rate on the reac tivation potential of different austenitic stainless steels susceptible and insusceptible to intergranular attack. Best results have been obtained with 5>-10 % boiling sulfuric acid solution, Ey = 700 mV starting potential and about 1 sec/mV sweep rate. Parallel experiments were carried out with the STRAUSS-metliod too. The measuring cell and the results of a serial experiments are given in the figures; +• and - symbols refer to the results of samples found to be susceptible and insusceptible with parallel STRAUSS tests. The heat-treatment was the following: 1. Without heat-treatment. 2. Heat-treatment for 30 minutes at 1050°C, quenched. 3. Heat-treated, 4 hours, 900°C. 4. Heat-treated, 1 hour, 650°C. 98 La jos HacML 3. Heat-treatment for ~j>0 minutes at 1030°C, quenched, followed by 4- hours heat-treatment at 900°C. 6. Heat-treatment for 30 minutes at 1030°C, quenched, followed by 1 hour heat-treatment at 650°C. 7. 4- hours heat-treatment at 900°C, then 1 hour heat-treatment at 630°C. 8. Heat-treated for 30 minutes at 1030°C, quenched, followed by heat-treatment for 4 hours at 900°C, later on 1 hour heat-treatment at 650°C. It was established, that the reactivation potential is "EH = 230-430 mV, if the examined steels proved to be susceptible by the STRAUSS-method. The reactivation potential of steel samples insusceptible to inter granular corrosion is more negative than EH = 100 mV, Experiments were also carried out with stainless steel electrodes made of welded samples taken from different heat-treated zones» The values of reactivation potentials of steels susceptible and insusceptible to intergranular attack measured on electrodes taken from different zones of the welded joints, are presented in the following figure. +#9 mm *j%J) \j %& C1 I • ° INSUSCEPT. The conclusion can be drawn, that the susceptibility or welded joints to intergranular corrosion can be examined by potentiodynamic methods. Michel Keddam 99 BASIC PRINCIPLES OF ELECTROCHEMICAL CORROSION TESTING Michel Keddam Physique des Liquides et Electrochimie, Groupe de Recherche du CNRS, associé ä 1'Université Paris VI, 11 quai St-Bernard, Paris 5e, Fiance In order to accelerate the corrosion testing, electrochemical methods have been developed, which determine indirectly the corrosion rate. However, such methods have to be compared to the data from direct measurements, i.e. weight losses. It is also possible to evaluate the fiability of a test by means of analysis of the theoretical model on which it is based . The electrochemical nature of the metallic corrosion in an electro lytic medium led to several tests using electrochemical measurements. The aim of this paper is : - to discuss the validity of this electrochemical tests, - to present a new accelerated electrochemical corrosion test allowing to overpass some basic difficulties pointed out during the first part. 1. Validity of electrochemical corrosion tests Very evidently, electrochemical methods directly derived from homol ogous usual methods in electrochemical kinetics for Redox equilibria. The corrosion current is assumed to correspond to the exchange current and the corrosion potential to the equilibrium potential. In the fol lowing discussion, it is shown that some phenomenons generally dis regarded in the case of Redox processes are often prevailling in the corrosion processes. The corrosion rate is often determined by extrapolating the anodic and cathodic branches of the overall current-voltage curve to their intersection (corrosion current ICOrr» corrosion potential Vcorr). This extrapolation is done in a (V, log I) plotting since the reac tions are assumed to obey the so-called Tafel exponential law (one step reaction, onto a constant area, without any mass transfer con trol) . New, it is well known that reactions are likely to occur through several elementary steps ; they may be inhibited by adsorption of ions or molecules (and/or), they may be controlled by mass transfer Il| . The rate of each elementary process depends not only on V but also on the instant value of coverages 0. and concentrations c. : ö i i I = f(V, 0., c..) (1) Besides it is very difficult to obtain a neat straight V(log I) rela tionship. It is commonly assumed that measurements performed quite near corrosion potential may overcome these difficulties. The polarization resistance R = (v^Orr which is the slope of the P dI Vcorr steady state current-voltage curve at the corrosion potential itself is frequentlA" related to Icorr through a reciprocal proportionality. This relationship can be justified if tne anodic (Ia) and cathodic (Ic) components of the overall current obey the Tafel law. On the other hand, when this condition is not satisfied, the derivative of (1) gives : iUO Michel Iveddaui de dc AT - ^T - L (— ) K (2) R Wv we.,c. M0 dV bc. dV p corr i i i i i In the case of constant values of 0. and c, if the charge transfer process always obeys the Tafel law with respective exponents b and b , the first term of eq(2) is : c b I + b I = I (b+b) a a c c corr a c Consequently, this term will be used in the general case to calculate the corrosion rate rather than R . However, it cannot be obtained from p steady state measurements. Recent improvements in measurements and interpretation of electrochem ical impedance in very low frequencies \2\ allow us to separate different terms in eq(2). Each of them has a different dynamic behav iour with regard to the sinusoidal variation of potential around VCorr- Indeed, the variations of 0. and c. have finite rate determined by the rate constants and the diffusion coefficients. The charge transfer process instantaneously follows the potential variation. At sufficiently high frequencies, the impedance of the interface will be the so-called charge transfer resistance, Rt. At lower frequencies, the relaxation of surface coverage and of concentration appears as a phase shift between I and V. Finally, at zero frequency, the station ary slope Rp will be found. Figure 1 shows, as an example, the impedance of a pure iron specimen (0,2 cm^) at the corrosion potential (electrolyte IM, H2SO4 + 5mml-1 propargylic alcohol). The experimental complexe impedance has been plotted (real part R, imaginary part G) at each frequency value indicated in Hz. In high frequencies (between 400 and 1 Hz), the capacitive impedance is to be assigned to the electrochemical double layer. The value of Rt (450ft) can be determined at a frequency about 0,7 Hz. At lower frequencies values the impedance exhibits an in ductive behaviour that have been interpreted as the relaxation of surface coverages by FeOH (anodic dissolution of iron) |3| and inhib itor (cathodic process) |4| . The limit at zero frequency gives the usual polarization resistance (in this case 25ft). I JG(A) 200 100 .100 Michel Keddam 101 2. Determination of the corrosion rate by the transfer resistance Several authors already used periodic current in order to make easier the Rp determination |5| or to calculate the adsorption of inhibitor by the measurement of the double layer capacitance. The results shown in figure 1 explains the dependence of impedance on frequency repor ted in the litterature. We use the impedance measurement of the interface in a quite different way : to determine Rt. The corrosion rates calculated from R and Rt may be very different. These two methods have been compared with weight losses data in the study of the corrosion of pure iron in IM, H2SO4 with respect to the concentration of propargylic alcohol. The weight losses are mean values from 24 hours corrosion tests. The results are shown on figure 2 in term of inhibiting efficiency (H) vs_ inhibitor concen tration. jHT/.l 100 jnyft^*^—gg^. = -X© tr'~ ^sr Cone (m M ) 0,1 0,2 0,5 10 20 -100 © Figure 2 Curve l is deduced from weight losses, curve 2 from Rt and curve 3 from R . The Rfc measurement appears to be a satisfactory test for corrosion. At high concentrations, a strong stimulation of corrosion (H < 0) is foreseen from Rp though such a stimulation was not ob served by weight losses measurements. The interpretation of the electrochemical impedance at the corrosion potential on the basis of a mixed potential model, led to a theoretical demonstration that the transfer resistance is the entity the most closely correlated to the corrosion rate, and this has been verified by experiments |4|. Litterature 1 K.J. Vetter. Electrochemical Kinetics : Academic Press. 2 I» Epelboin, M. Keddam, H. Takenouti : This meeting. 3 I. Epelboin, M. Keddam. J. Electrochemical Soc. _ML§_> 1052, (1970) 4 I. Epelboin, M. Keddam, H. Takenouti. J. Applied Electrochem. To be published. 5 M. Prazak, M. Barton. Corr. Science, 7_, 159, (1967). T. Murakawa, S. Nagaura, N. Hackerman. Corr. Science, 7,79,(1967) (1 > "- ... THE APPLICATION OF Fe53 AND Fe3 TO CORROSION TESTING Pance Kirkov and Zarko Radosavljevic Center for the Isotopes Application in Science and Industry, Physical and Electrochemistry, Faculty of Technology and Metallurgy, Skopje, Yugoslavia. The rate and mechanism of the corrosion at the heavy alloying steels in the solutions of the strong electrolytes were investigated by Fe-^-> and Fe . The steel sample with and without passive oxide layer was been treated in the neutron reactors for Fe^9 (£n metal) and Fe-'-3 (in oxide layer) reached. The measurement on the concentra + + tion of Fe and Fe^ in Fe ^ and Fe 3 ±n the corrosion product and solutions have been used for determination the charges and the ions transfers. According the charges and the ions transfers dates the rate and mechanism of the corrosion have been postulated. By the comparative measurements of the corrosion at the native condition in the solutions of the strong electrolytes (NaCl, H„SO,, HCl etc) y and the charge transfer rate by Fe^5 ancj pe-> ? isotopes technic could be used as a very accurent method for the corrosion testing. 103 UdVlu. O. J-.ov<-i. J-ag THE ELECTROCHEMICAL BEHAVIOUR OF NIMONIC AND TITANIUM ALLOYS IN MOLTEN SALTS. David G.Lovering. Chemistry Department,City University»London EC1V 4PB,England. Nimonic and titanium alloys find increasing application particularly where strength and inertness at elevated temperatures are important. Condensed impurities on these materials under working conditions are likely to be composed mainly of liquid ionic salts. Accelerated corrosion testing of nickel and titanium metals,NI90, NIM115, and other materials was investigated in molten KN0~ at 350 C, by applying both anodic and cathodic potentials to specimens: the open- circuit potential-time variations were also noted. For nickel, nickel alloys and titanium, the initial open-circuit potentials (o.c.p.) rapidly moved a few hundred millivolts in the anodic direction during the first 10-20 minutes of immersion in the melt, then they drifted back in the cathodic direction, yielding stable potentials in the case of nickel and its alloys within 24 hours, but not reaching equilibrium for titanium. These results conflict with those of Piontelli , who apparently did not observe specimens for a sufficiently long period. It was confirmed that the stable oxide film formed on samples was a higher oxide of nickel i.e. NiO, , as expected on the basis of Zambonin's recent work . A competition between oxide growth and subsequent reaction with the melt is also consistent with oscillatory phenomena in the E-i curves observed in this and previous reports . The importance of the reaction between the melt and the (oxide-covered) specimens was demonstrated by polarising them at increasing cathodic potentials i.e. towards the region of anion decomposition . All alloys of nickel and titanium, as well as the metals themselves, dissolved in this region to form highly coloured (nitrosyl?) complexes in solution, similar to those reported by Bartlett and Johnson for noble metals. The apparent current efficiency of metal dissolution was < 10%,confirm ing a coupled chemical process (oxide reaction) accompanying anion electro-reduction. The absence of any insoluble(hence inhibiting)oxide of potassium ' in pure, molten KN0„, thus allows all metals and alloys to corrode without hinderance at relatively low cathodic potentials. Anodic polarisation of titanium specimens, led to film thickening and, eventually, gas evolution (02) at applied voltages of +30 ;a small electronic current through the film was apparent. Whilst nickel metal dissolved at small anodic applied potentials, without apparent over- potential or gas evolution,nimonic alloys invariably exhibited concomittant gas evolution (02) and dissolution as the anodic limiting reactions. Galvanostatic pulses applied to these materials yielded highly complex transitions - often as many as six- on the recorded transient signals. These transitions were current density dependent in a manner characteristic of kinetically-controlled electrode reactions. It was not the purpose of this investigation to study these reactions, per se. Similar corrosion testing was carried out for NI90, NIM115, titanium metal, IMI230 and other alloys in "wet" LiCl-KCl(41.5 mole %) melts at 400 C. All nimonic samples dissolved anodically at low cell voltages 104 Davia G. Lovermg '"° 5~lv> yielding diffusion-limited current plateaux at ^ 0.5A.cm . At higher anodic potentials chlorine evolution accompanied metal corrosion. Apparent current efficiencies for dissolution on the limit ing current plateaux always exceeded 100% by a significant amount. The erosion of intermetallic particles, observed visually, and their subsequent chemical reaction with the melt would account for these dis crepancies, however. Cathodic galvanostatic transients applied to nimonic specimens in this melt, commencing from the o^g.p., produced very long(20-40s at low current densities, < 0.5mA.cm )transition plateaux at ^ -3V v. 0.1m.Ag in KNOy glass membrane/...reference electrode. The samples appeared to be very bright and smooth in this region, consistent with oxide stripping to expose the bare metal.Removal of specimens from the melt whilst bright and with the current switched on, led to immediate oxidation and the production of a coherent,black film. oT The behaviour/titanium and its alloys in this chloride melt was disturbing - they dissolve spontaneously on open-circuit, with accompany- gas evolution. Condensates in realistic industrial environments are most likely to be chloride-bearing! The rate of corrosion could be slowed, but not stopped by potentiostating at appropriate cathodic (to the o.c.p.) potentials. Specificity of attack occurred at the three- phase boundary regardless of the presence of oxygen and/or water in and above the melt. This effect was reduced by the presence of small quantities of fluoride ions, which also led to increased anodic corrosion, but surprisingly, a small cathodic "passivation" plateau at ca. -1.1V v.ref., as well. Coulometric measurements and the purple colouration of the resulting solutions suggested that titanium(and its alloys) dissolved anodically to tervalent ions in this melt.However, the presence of two peaks in the anodic, potentiostatic transients is consistent with a two-stage electrochemical process . The presence of small amounts of ZnCl» (strong Lewis acid) in the melt led to very much enhanced corrosion rates of the titanium based materials, which also suffered attack by ZnCl„ in the vapour phase. The corrosion of these nimonic and titanium alloys was also briefly investigated in other molten media; the essential results are summarized as follows:- Equimolar NaOH-KOH at 215°C: Nimonic alloys not corroded under anodic or cathodic potential excursions. NH.N0o-NaN0o(20 mole %) at 150°C: ...... ,._ ,. . 4 3 3_ ' Nimonic alloys readily dissolve when anodically polarised, corrosion rates being enhanced in the presence of dissolved water, bromide or chloride ions. No passivation regions. Cathodic attack presumably limited by formation of melt- insoluble oxide of sodium . Titanium samples unaffected. CaCl„. öIUO at 35-62 C: Nimonic samples readily dissolved anodically. Titanium attacked slightly at three-phase boundary at cell voltages of + 20. Eutectics NaN03-KN03 at 250°C and LiNO^-KNO» at 320°C: Anodic gas 105 David G. Lovering evolution preceded corrosion; no cathodic attack for all nimonic alloys, NaNQ2-KN02(30 mole %) at 320°C: Gas evolution limited anodic and cathodic potential excursions for nimonics. AlClg. 6H2O + Al (NO 3) .3.9Hz, 0 at 180°C: Highly corrosive medium , spontaneously dissolving nimonic,titanium and stainless steel alloys as well as platinum and gold. Volatile melt. The results of this preliminary investigation are intended to be of immediate practical relevance, whilst detailed theoretical information will follow a closer examination of each individual system. Acknowledgements: I wish to thank my colleagues for their advice and Mr.P.A.Jackson for technical assistance in this project. 1 R.Piontelli Ann. N.Y.Acad .Sci. 79_, 1025 (1960) 2 P.G.Zambonin e.g. J.Electroanal .Chem.2_4_, App.25(1970) et seq. 3 G.J.Hills and K.E.Johnson "Advances in Polarography",p.97A»Pergamon London (1961). A H.E.Bartlett and K.E.Johnson Corr.Sci. 6_, 87 (1966). P.G.Zambonin J.Electroanal .Chem. 24_, 365 (1970) "Encyclopaedia of Electrochemistry" p.1130 Rheinhold,N.Y.(1964). C.A.Angell, lecture, Southampton University,18th May(1970). F. ndiib ieiu 106 AN ELECTROCHEMICAL TEST FOR QUALITY OF CONVERSION COATINGS ON Al ALLOYS F. Mansfeld and E. P. Parry North American Rockwell Science Center, Thousand Oaks, California, USA Accelerated corrosion of a metal immersed in a corrosive environment and electrically coupled to a more noble metal is one of the most common forms of corrosion. In many instances, a non-conducting material can be inserted between the dissimilar metals or other means can be used to prevent or minimize galvanic corrosion if a conducting path is not required between the metals. However, in certain applications in air craft and aerospace vehicles dissimilar metals (e.g. Al and Ti alloys) have to be connected in a way that the electrical resistance between them in minimal. In this situation, which is particularly dangerous from a corrosion standpoints chemical conversion coatings (1) have been used to protect the Al alloys. When applied properly, these coatings have been shown to have the necessary high conductivity, in addition to providing good corrosion protection. The testing of these conversion coatings is generally done using a 7 day salt spray test without electrical contact to a more noble metal. The corrosion tendency is evaluated visually after this time, and the conductivity of the coating is measured before and after the salt spray test. Because of the time required to perform this test, the qualita tive evaluation of corrosion susceptibility, and the fact that corro sion susceptibility in the salt spray test is not measured with a dis similar metal contact, it seemed desirable to examine alternative test procedures. One alternative appeared to be the direct measurement of the galvanic current flowing between the coupled dissimilar metals. Simultaneously, with galvanic current data, weight loss data were obtained to allow comparison between results from different test methods Experimental: Materials tested were the Al alloys 7075-0, 606I-T6 and 2024-T3, stain less steel 304 and Ti-6Al-W. The Al alloys were coated with Alodine 600, a chemical conversion coating with low electrical resistance. Dissolution rates of uncoated (bare) and coated Al alloys electrically coupled to SS304 or T t-6A1-4V were obtained from weight loss data and measurements of the galvanic current. The galvanic current between two dissimilar alloys was measured using the zero impedance ammeter described by Lauer and Mansfeld (2) for a 2k hour period by using a strip chart recorder. From these data the average current density was obtained and converted to a dissolution rate using the relation 1.23uA/ cm2 = 1 mdd (mg/dm2day). The electrolyte used for all tests was 3-5% NaCl at room temperature, stirred by air. For further experimental deta i 1s see Ref. (3). Results: Table I shows a comparison of dissolution rates calculated from gal vanic current data (r„) and from weight loss data (rw]) for bare and coated Al alloys. Pi scuss ion: The results of Table I show that the chemical conversion coating Alod ine 600 greatly reduces dissolution rates of Al alloys electrically coupled to stainless SS304 or Ti-6Al-4V and. immersed in air stirred Mpn^tp i n iU/ 3.5% NaCl. The damaging effect of the more noble metal is, in general, more pronounced for SS304 than for Ti-6A1-4V. Table I. Dissolution Rates [mdd] from Galvanic Current (rq) and Weight Loss Data (rw]) Bare Al Coated Al Couple »*wl fwl 7075/Ti 7.6 50.5 3.7 10.5 7075/ss 38.0 69.5 8.4 7.0 6061/Ti 6.3 16.5 0.-9t n .a . 6061/ss 28.0 32.5 22.Of 2k. 2t 202^/Ti 3-8 18.5 6. Of 3. 31 2024/SS 13.4 30.0 3.1 2.5 n a. = data not available; t = average of duplicate tests. A comparison of dissolution rates of coupled Al alloys obtained from galvanic current data (r^) and from weight loss data (rw]) shows that especially for bare Al alloys, where dissolution rates are high, rg values tend to be appreciably smaller than rwl values. These differ ences between galvanic current data and true dissolution rates have to be expected as pointed out recently (k). Most practical situations in galvanic corrosion can be classified in one of three cases treated theoretically by the authors (3,^) using the concepts of mixed poten tial theory. In Case I, it is assumed that the corrosion reaction is under diffusion control, while in Case II, it is assumed that Tafel behavior is obser ved for the oxidation reaction on the anode and the reduction reaction on the cathode. In Case ill, the anode is polarized only slightly due to coupling, Tafel behavior is therefore not observed for the oxidation reaction of the anode. Only for Case II is the dissolution rate cal culated directly from the galvanic current equal to the true dissolution rate of the anode. In Cases i and Mia correction has to be applied to the value obtained from galvanic current data as explained in Ref. (3). A decision about which Case applies to a given couple can be made by inspection of the potentiostatic polarization curves for the un coupled metals in the test electrolyte. In Table II it is indicated which case applies to the couples studied (3). Also listed in Table II are dissolution rates rj obtained from corrected galvanic current data for bare and coated alloys. For Case IN only a lower and an upper limit can be given for the value of r^. Good agreement is observed, in general, in Table II between weight loss and corrected galvanic current data. While galvanic current data still tend to be somewhat higher than weight loss data for Case I, very good correlation is observed for Case II. For Case Ml, weight loss data rw| fall between calculated upper and lower limit of rj. While the individual polarization curves of the alloys studied have to be known in order to apply the necessary corrections to galvanic cur rent data, continuous measurements of the galvanic current offer con siderable more information than weight loss data. The time dependence of galvanic current and potential allows predictions concerning long 108 F. Mansfeld time corrosion protection p rovided by conversion coatings, high gal- vanic currents which are in creasing with time being indicative of coat- ings with poor quality. As compared to the salt spray test, the tests reported here allow, in a s horter time, a ranking of coatings applied to various alloys or of coa tings prepared in different batches while the salt spray test only al lows "pass or fail" decisions. In addition, the galvanic current test p rovides information concerning the effect of coupling to a more noble al loy, information usually not available in the salt spray test. Table II. Dissolution Rate s [mdd] of Coupled Al Alloys from Weight Loss Data (rw|) S from Corrected Galvanic Current Data (rj) Couple Bare Coated rd wl rwl 7075 Case I 50.5 Case I 1 I 10.5 7.6-74 3.7-11.3 7075 Case 69.5 Case I 7.0 SS304 105 16.0 6061 Case I 16.5 Case I I I n.a. T~RT-4 6.3-21 0.9-5.5 6061 Case 32.5 Case I I 24. 2f SS304 43 22t 2024 Case I 18.5 Case I I I 3-3+ TTT-4 3.8-33 8 6.0-10.6 2024 Case I or I I 30.0 Case I I 2.5 SS304 13-4 to 43.4 3.1 n.a. = data not available; f = average of duplicate tests. Since conversion coatings have to pass the salt spray test in certa 1 n specifications, it is recommended to use the galvanic current test as additional quality control when coupling to dissimilar metals i s ex- pected. References: U) W. E. Pocock, Meta1- Progress 84_, 100 (1963). (2) G. Lauer and F. Mansfeld, Corrosion 26_, 504 (l 970) . (3) F. Mansfeld and E. P. Parry, submitted to Corrosion Scien ce. (4) F. Mansfeld, Corrosion 27, 436 (1971). F. Mans feld 109 ON THE SO-CALLED LINEAR POLARIZATION METHOD FOR MEASUREMENT OF CORROSION RATES K. B. Oldham, Trent University, Peterborough, Ontario, Canada, and F. Mansfeld, North American Rockwell Science Center, Thousand Oaks, Cali forni a, USA. The basis for the inappropriately named (l) linear polarization mel.c;', also called polarization resistance or Stern-Geary method, was given ';y Wagner and Traud (2) in 1938, who showed that the slope of a polariza tion curve at the corrosion potential can be used for an estimate of corrosion rates. Stern and coworkers (3,^0 showed later that "for most corroding systems, the corrosion rate can be estimated to within a factor of 2 by a simple measurement of the current required to polarize a few millivolts11 {k) . In the following years the mi - ~.'-"'cep;: i on has arisen that plots of current vs potential in the vicinity of the corr osion potential have to result in straight lines, the slope of which ("polarization resistance") is proportional to tne ;orrosion current. Sometimes, even portions of the polarization curve which are linear a few tens of millivolts away from the corrosion potential have been used to measure corrosion rates. In the following we are going to show that not only is there no theo retical reason for polarization curves to be linear at the corrosion potential $corr> the non-linearity possibly being severe, but that the polarization curve must display curvature in the vicinity of $corr for the Stern-Geary treatment to be valid. We will also introduce a new graphical method to calculate corrosion rates from polarization curves. Under the usual assumption of mixed potential theory the current i vs potential (j> relationship, known as polarization curve, has the mathema tical form: a plot of i vs 4> as well as of the slope di/d(f> of this curve using as parameters ba == 30 mV, bc = 120 mV - values which are typical for iron in acid media - and icorr = 1-00 uA. The polarization curve (solid line) is no_t_ linear at <£COrr» but when cj) is some 30 mV negative to (j)corr. This is brought out more clearly for the gradient of the curve (broken line), which is a variable quantity except in the neighborhood of ^"^corr = -30 mV, where it is virtually constant over quite a poten tial range. This explains why many workers have reported linear polar ization curves in the neighborhood of Geary method, it uses the polarization curve in the vicinity of positive and negative to cj>cor-r- (b) Determine the currents i/\ and \ Q and compute their geometric mean /i/^i Q. (c) Draw tangents to points A and C. From the crossing points of these tangents with i = 0 axis, determine the absolute pot ential differences | A^ | and jAQ| (see Fig. 2). (d) Compute the term: I*AI IAC! (e) Read the term U corresponding to W from a table given _? n Ref. = (7)- The corrosion current is given by iCorr ^ ^'A'C• The corrosion current can be obtained readily and with great accuracy by our graphical method, which should be especially helpful when corr osion rates have to be determined as a function of time, since it does not require knowledge of parameters other than the measured polariza tion curve. References: TO K. B. Oldham and F. Mansfeld, Corrosion 27, hlh (1971). (2) C. Wagner and W. Traud, Z. Elektrochem. Pf, 391 C1938)- (3) M. Stern and A. L. Geary, J. Electrochem. Soc. \0h, 56 (1957). (4) M. Stern and E. D. Weisert, Proc. ASTM 59., 1280~TT959). (5) F. Mansfeld and K. B. Oldham, Corr. Sci. U_, 787 (1971). (6) K. B. Oldham and F. Mansfeld, Corrosion (to be published). (7) K. B. Oldham and F. Mansfeld, to be submitted to Corr. Sci. F. Mansfeld 111 Figure Captions: Fig. 1: Polarization Curve (solid line) and slope of this curve (broken line) for ba = 30 mV, bc = 120 mV, and icorr = 1.00 yA. Fig. 2: Illustration of proposed graphical method for calculating corrosion rates. (Parameters as in Fig. 1.) T iM) • -•k("iV) Figure 1 1 / / +J It f / / +2 — A +1 >7 ^ // /I 1 0 /1 X" .>i , I 1 1 \\ \ C y* A ii/A_| -1 - - c J --J - IT A A -2 1 1 -iO -10 0 + 10 +20 *-»corr(mV) Figure 2 112 Francesco Mazza EVALUATION OF THE SUSCEPTIBILITY OF AUSTENITIC STAINLESS STEELS TO STRESS C.C. BY CHEMICAL OR ELECTROCHEMICAL METHODS Giuseppe Bianchi, Alessandro Cerquetti , Francesco Mazza and Sandro Torchio Laboratory of Electrochemistry and Metallurgy - University of Milan - Italy. Susceptibility of stainless steels to stress corrosion cracking is almost entirely evaluated by exposure to satu- rated magnesium chloride solutions at boiling temperature. This method can give rise to scattered results due to experimental difficul ties arising from handling hygroscopic chemicals as MgCl? and from operating at temperatures strongly influenced by the water content of the solution. In addition, the interpre tation of the mechanisms of s.c.c. results quite difficult being the satu rated magnesium chloride solutions at boiling tem perature far from being chemically defined. Sulphuric or perchloric or hydrochloric acid solu tions containing sodium chloride in very specific «r FULL SYMBOLS INDICATE CRACKING COMXHONS I 3 concentrations at 25°C ter 10,- 2 >-1 MtNaCfl can be succesfully used in to order to evaluate the sus ceptibility of austenitic Fig, 1 - Corrosion rates and sta i less steels to s,c.c. s. c, c. conditions for vari- These solutions (1-4) can ous materials in HCIO +NaCl produce within 24 r 48 solutions, hours, transgranular crack ing or, in some cases even intergranular cracking in solu- tion quenched specimens, in dependence of the solution composition. Transgranular cracking occurs at concentra tions corresponding to a well defined corrosion rate as displayed in Fig.l which shows the corrosion rates of U- bend specimens of various materials in 3M HCIO- containing various amounts of NaCl. In the case of HClO.+NaCl and HpSO. + NaCl solutions, intergranular cracking (I.G . A .) can occur on type 304 and 321 stainless steels and Fig, 1 in dicates the corrosion rates during I.G.A. development. Francesco Mazza 113 In all cases, potential measurements indicated that the specimens during the general corrosion or localized corro sion processes were in the active potential region indica ting that the corrosion process is under cathodic control of hydrogen evolution. External cathodic or anodic currents can markedly influence the initiation and development processes of stress c.c. pr_o duced in the above described solutions. In part icular , ano dic current (i.e. 0.4 mA/cm2 in 5NH2SO4 + 0.5 N NaCl ) sup plied by a galvanostatic or potent i os t at i c system can stroji gly stimulate the rate of initiation or development of the cracking process, being created more severe conditions for cracking occurrence. On the contrary very small cathodic currents can partially or completely protect the metal sur face increasing the time for cracking initiation or arre sting cracking in progress or even avoiding initiation.Even the cracking morphology can be modified by the application of external cathodic current so that cracking which occurs by intergranular path for exposure in 10 N H^SO^ + 0.1 M NaCl solution can be transformed in transgranular cracking when a cathodic current of 0,1 ;nA/cm2 flows through the specimen. A very useful tool in these studies has been found in the use of a polarization cell structured as the cell described by Greene (5) in which, however, the conventional platini zed platinum auxiliary electrodes have now been substitu ted with capillary-imbibition hydrogen electrodes, previou sly described (6), In this cell, the anodic reaction at the auxiliary electrodes, occurring during cathodic polariza tion runs in the acidic chloride media, does not lead to platinum dissolution or to chlorine evolution, but just to hydrogen ionization according to H^-*2H+ + 2e avoiding dangerous contamination of the testing solution especially during long term polarization experiments. Testing s.cc, susceptibility in acid-chloride solutions, in addition to testing in MgCl2 solutions can be useful under the necessity of enlarging the field of investigation within the spectrum of occurrence of s.c.c. phenomena. It can be very effective when room temperature testing is con venient or when an environment less aggressive than boiling MgClp solution is requested. In this case the aggressivity of the solution can be regulated by the solution concentra tion or'by the application of cathodic or anodic external currents. The possibility offered by these solutions to evji luate the susceptibility to I.G.A. of non~sensitized speci mens can give useful information on the structural unifor mity at the grain boundaries region independently upon any chromium depletion phenomena. 114 Francesco Mazza References 1) S.J. ACELLO and N.D.GREENE, Corrosion J8, 286t (1962). 2) F.MAZZA and N.D.GREENE, Proc. 2nd European Symposium on Corrosion Inhibitors, Ferrara 1965. p. 401 3) M.A.NIELSEN, Corrosion 27, 173 (1971). 4) G.BIANCHI,F.MAZZA and STTORCHIO, Nato Conference on the Theory of S.C.C. Ericeira, Portugal April 1971 5) N,D. GREENE, Experimental Electrode Kinetics - Rensselaer Polytechnic Institute TROY N.Y. 1965 G) G.BIANCHI,A.BAR0SI,G.FAITA,T,MUSSINI, J. Electrochem. Soc. 112, 921 (1965). James R. Myers 115 COMPARISON OF (JORRQbiUi\ RAICS iW SULFONIC ^CIC Or/TAIXEE USING Rl^iS':' ANCE -POLARIZATION, TAFEL- EXTRAPOLATION, AND WE IG IT-LOSS TECHNIQUES James R. layers Civil Engineering School, Air Force Institute of Teclinology, Wright- Patterson Air Force Base, Ohio 45433, U.S.A. Corrosion rates of four pure metals (Ni, Cr, Ti, and Al) and six ter nary alloys (Ni-15Cr-2Al, Ni-15Cr-4Al, Ni-20Cr-2Al, Ni-20Cr-4Al, Ni-15Cr-3Ti, and Ni-20Cr-3Ti) were determined in H2-saturated, IN H2SO4 at 22*1°C using resistance-polarization, Tafel-extrapolation, and weight-loss techniques. Corrosion rates predicted using the two electrochemical techniques were time dependent; in general 96 to 120 hours exposure to the electrolyte was required to establish a "steady- state" corrosion rate for each of the materials investigated. "Steady-state" corrosion rates predicted using the Tafel-extrapolation technique were in excellent agreement with corrosion rates obtained by weight-loss testing (Table 1). "Steady-state" corrosion rates pre dicted using the resistance-polarization technique were considered to be in good agreement with those obtained by weight-loss testing (Table 1). Table 1 Electrochemically-Predicted and Weight-Loss Determined Corrosion Rates in mdd (mg/dm2/day) for Ni, Cr, Ti, and Ni-Cr-Al/Ti Alloys in ^-Saturated, IN H2SO4 at 22±1°C Tafel- Resistance- Weight- Extrapolation Polarization Loss Material Corrosion Rate Corrosion Rate Corrosion Rate Ni 1.8 5.2 1.6 Cr 586 1130 820 Ti 24.2 18.5 26.1 Al 14.5 6.1 16.8 Ni-15Cr-2Al 2.6 6.8 1.2 Ni-15Cr-4Al 1.1 5.9 0.9 Ni-20Cr-2Al 3.1 1.8 1.6 Ni-20Cr-4Al 2.7 3.0 0.7 Ni-15Cr-3Ti 3.1 2.3 0.7 Ni-20Cr-3Ti 3.5 3.2 1.3 Corrosion rates for the materials investigated were considered to be in good agreement with those reported by earlier investigators for similar concentrations of deaerated, sulfuric acid. For example, McKay and Worthington (Corrosion Resistance of Metals and Alloys, pp. 138, New York, Reinhold Publisning Corp., 1936) reported the corrosion rate of aluminum in H2-saturated, 6% H2SO4 to be 18mdd, and Bond and Uhlig (J. Electrochem. Soc., 107, 488-493, 1960) reported that binary Ni-Cr alloys, having compositions similar to those used in the present study, corrode at a very low rate (about lmdd) when exposed for 10 days to H2-saturated, 5% (1.1N) H2SO4. ' ' .T . ^cs tl cthT-~ni t c ELECTROCHEMICAL STUDIES OF CORROSION-EROSION DURING HYDRAULIC TRANSPORT OF SOLIDS IN PIPES J. Postlethwaite and E. B. Tinker Faculty of Engineering, University of Saskatchewan, Regina Campus, Regina, Saskatchewan, Canada The hydraulic transportation of large quantities of solid raw materials, such as iron concentrate, coal and potash, over long distances in pipelines is receiving much interest at the moment as an alternative to rail transport. Such pipelines suffer metal loss both by chemical corrosion and mechanical abrasion (erosion). The dominant component varying with the slurry; for example it has been reported that with coal slurries chemical corrosion Is the manor factor whereas with phosphate slurries erosion is the major problem . Also of course the mechanical abrasion of corrosion products will affect the rate of corrosion. Electrochemical techniques offer the means of distinquishing between the two contributions to the total metal loss; by using potentio- static techniques to determine the instanteous corrosion rate and cathodic protection to measure the erosion in the absence of corrosion, over a period of time. This paper reports the application of potentiostatic techniques to the determination of the chemical component of the corrosion-erosion of 2" diameter carbon steel pipe in a flowloop, Fig. 1, containing: 30-50, and 60-100 mesh sand; 60-150 iron ore; 100-325 iron concentrate; and 20-60 potash suspended in aerated water at 70°F. The nominal analysis of the water was: pH 8.4j calcium hardness 350 ppm; total hardness 640 ppm, SOf 501 ppm; Cl 16 ppm, and of the carbon steel pipe, ASTM A53, 2" (SCH80) , wt % C 0.1, Mn 0.5, Cu 0.2, Ni 0.12, P 0.01, S 0.03. A single pipe sample was used to study the variation of the corrosion with slurry velocity and concentration. The pipe was masked with a polyester resin to leave a longitudinal strip, 6 x 1.25 cm, exposed to the slurry. Separate experiments were done with the strip at the top and the bottom of the pipe and with abraded and as received (with mill scale) specimens. Corrosion rates were determined for three slurry concentrations, 0.5, 0.75 and 0.85 of the maximum volume of solids the line could carry, and three velocities in the range, V to 2VC where Vc is the critical velocity at concentration c. Two controlled potential techniques were used to determine the rate of corrosion. These were both based on the fact that when oxygen reduction is the predominant cathodic reaction the corrosion rate can be determined by obtaining the rate of iron dissolution in dearated solution, at the previously determined corrosion potential in aerated solution. Of course with such a method of knowledge of Tafel constants is not required, as with linear polarization techniques which are necessary when hydrogen evolution is the predominant cathodic reaction. This is of considerable advantage when dealing with "impure" engineering systems. Method 1. The corrosion potential was determined for the iron pipe J. Postlethwaite 117 corroding in the slurry, with the slurry being continually aerated by passing a portion of the slurry through the aeration column. The potentiostat was set at this potential, EQQRR, Fig. 2, and the pipe brought under control at this potential, where the current was zero. The solution was then deaerated with commerical grade nitrogen. As the deaeration proceeded to completion the current which was recorded during this period rose to a maximum value, corresponding to the corrosion current, i(X)RR« With the system shown in Fig. 1, which had a working volume of 45 liters, deaeration was completed in 20-30 minutes. The residual dissolved oxygen concentration was around 0.03 ppm. When the system was reaerated the current fell back to close to zero, depending on the change in the corrosion potential during the time of the experiment. In fact the current following aeration was always < 10% and usually < 5% of the corrosion current; and an accuracy of 10% for such measurements is considered satisfactory. Method 11. Potentiokinetic anodic, E-i, curves were recorded for the specimen in deaerated slurry, starting at the deaeration rest potential and concluding above the aeration corrosion potential. Following sweeps at the three slurry velocities used, the slurry was reaerated and the aeration corrosion potentials determined for each slurry velocity. The corrosion current icoRR» was tnen read of the appropriate E-i curve. With a sweep rate of dE/dt = 50 mV/m the corrosion rates were much greater than those obtained by method 1, whereas with a sweep rate of 10 mV/m the results were usually in good agreement.jThe corrosion rates for the water alone were around 0.5 to 1 mm/year. The presence of sand, iron ore and iron concentrate increased the corrosion rate by an order of magnitude with maximum values around 30 mm/year; whereas with the potash reduced corrosion rates were obtained most probably because of the reduced oxygen solubility and scaling. For example at 0.75 C^x, and 1.5 Vc, the corrosion rates, in mm/year, at the bottom of the pipe (abraded) were: 60-100 sand, 4.6; 30-50 sand, 33; iron ore 19; iron concentrate 25; and potash 0.24. With the sands there was a large effect of velocity, especially with the lower concentration used, 0.5 Cmax. For example with 30-50 sand the corrosion at the bottom of the pipe (abraded) was 3.8 mm/year at Vc and 28 at 2VC; whereas the corrosion rate at 0.75 Cxnax and Vc was 32, and the same value at 0.85 Cmax and 2VC. There was a smaller effect of velocity with iron ore and the iron concentrate. The fact that such high corrosion rates can be obtained in slurry pipelines is shown by the recent data of Link and Tuason1, who reported a chemical corrosion rate of 44 mm/year and an erosion rate of 13 mm/year for a 4" iron pipe containing copper concentrate. Tomashov has reported a sharp decrease in the overpotential for iron dissolution with mechanical renewal. 1. J. M. Link and C. 0. Tuason, American Mining Congress, Las Vegas, 1971. 2. N. D. Tomashov and L. P. Vershinina, Electrochim Acta, 15, 501, 1970. J. Postlethwaite 3 m\ z % 1. Slurry Flowloop for Electrochemical Corrosion-Erosion Studies. 1, 2" diameter lucite tubing, 2, 2" x 2" rubber lined centrifugal pump; 3, aeration-deaeration column; 4, heat exchanger; 5, carbon steel pipe 6" x 2" diameter (SCH 80); 6, calomel electrode; 7, stainless steel counter electrode; 8, potentiokinetic equipment. -coRft -- 'CoRR É-o. t-cofm, <-OL <-CöRfc Lflu (a) Corrosion diagram (b) Method 1 (c) Method 11 Fig. 2. Determination of Corrosion Rate A. Rahrael 119 ELECTROCHEMICAL CORROSION TESTING OF METALS IN ALKALI SULPHATE MELTS Dr. A, Rahmel Dechema-Institut, Frankfurt am Main, FederaJ. Republic of Germany. By using in principle the same potcxitiostatic technique as normally used for corrosion testing in aqueous solutions one can get also weight loss-potential relationships or anodic partial current density- potential-curves for metals in alkali sulphate melts. Measurements have been done in the eutectic melt of Li2S04 , K2SO^ and Na2SO^ at 625 C on pure metals and technical alloys. Particularly Chromium containing Iron, Nickel and Cobalt based alloys as well as pure Chromium show in a wide potential range a passivity-like behaviour with a definite breakthrough potential depending on the Chromium content of the alloy. Such clear relationships can only be obtained by potentio- static measurements point by point and not by potentiodynamic measurements, because there is no simple relationship between weight loss and the total current for complex technical alloys. Measurements of the corrosion potential (without external current) show that in a neutral melt the corrosion potential of Chromium containing technical alloys is in the passive range and that the corrosion rate is low. In the presence of S03 in the atmosphere (acidic melt) the corrosion potential shifts into positive direction and into the transpassive range. On these conditions the corrosion rate is much higher than in a neutral melt. These findings can explain the corrosion of technical alloys under deposits reach in alkali sulphate in the presence of S03 in the atmosphere. Such conditions are - for example - realised in some coal fired boiler plants. This potentiostatic method probably can also be used as a short term test for testing the hot corrosion resistance of gas turbine alloys. i T-: "I ' T _, T „, Q , •>.. . ,.,1. «. J. lilc o. ill OWAII. Oiliv ELECTROCHEMICAL INVESTIGATIONS ON PITTING Giinter Herbsleb and Wilhelm Schwenk Mannesmann-Forschungsinstitut, Duisburg, Germany There is quite a number of papers which describe that pitt ing corrosion of stainless steel will only occur if the potential is more noble than a critical value, the letter being called "pitting potential" Uc. The meaning of this critical potential is still interesting in the research sector. Apart from the influence of steel or solution com position, there are two problems: 1) Test methods for determining Uc, and questions of precision and practical applications. 2) Time and hysteresis effects concerning the critical values for both pit formation and pit passivation. There are the following test methods: potentiokinetic (i), and potentiostatic methods without (ii) and with (iii) stepwise potential change, galvanostatic method (iv), and a method with a potentiostatic circuit connected with an ohmic resistance in line (v). Owing to (v) it should be possible to imitate the cathodic current density/potential behaviour of chemical corrosion without electrochemical control. If there is a hysteresis system, the test methods mentioned above should be applied for measuring the follow ing critical values: a) potential for pit formation (Uf): i, ii, iii; b) potential for pit passivation (Up): iii, iv, v. We investigated three very different systems to study the applicability of accelerated electrochemical methods for testing pitting susceptibility of stainless steels. I. Type 1.4301 (X 5 CrNi 18 9), 1 m NaCl with/without 1.5 m Na2S04 (inhibitor) at 25 °C. After water quenching from 1300 °C (diffusion heat treated) the pitting susceptibility of both cutting and rolling areas are equal each other with Uf = 0.25 V (hydrogen scalei By addition of sulphate inhibitor Uf is shifted to 0.55 V, as determined by methods (i, ii). Contrary to this result Up still remains at 0.25 V, as determined by method (iv). As to the results with method (iii) there are values between 0.25 and 0.55 V depending on the intensity of corrosion attack before potential change. The question arises whether the increase of Uf is or is not a notional effect only due to an increase in the induction period respectively to a decrease in pit formation rate. One may come to a decision by means of a long-term test at poten- V.'ilhelm Sch^enk 121 tials close to Uf. We did not find any pitting in a ö months' test by method (ii) at 0.4 V. As we have never heard of any induction period being longer than several days probably there is really an increase in Uf, therefore test methods (i, ii) are applicable. II. 35 % chromium steel, 1 m NaCl at 80 °C. Method (i) with a scan rate of 40 and 200 mV/h evaluated Uf at about 0.2 V. At more negative potentials we observed repassivating pitting. Thus we have also to consider a critical potential for the formation of repassivating pits Ur at about 0.1 V. By testing with methods (ii, iii) we got the same results. But applying a method (iv) with 20 uA/cm^ we found irregular potential oscillations with a Up value at about - 0.1 V. This does not correspond to the test results achieved by method (iii) after precorrosion at 0.25 V. Investigations by method (iv) and a method (v) with an open circuit potential U(i=o) = 0.3 V always re sulted in irregular oscillations of both potential and current. In addition, both the current density in the case of method (iv) and the ohmic resistance in the case of method (v) had a remarkable influence on corrosion poten tials which can be more negative than the Ur and Uf values. Concerning the Up values, we found that they were not con stant and depend on the circuit data. Thus only the methods (i, ii, iii) but not the methods (iv, v) were found to be applicable for testing pitting susceptibility of this system. This result does not correspond to the behaviour of other systems, e. g. type 1.4301. III. Austenitic stainless steels with different molydenum content (0 - 5 %) > 0.1 - 3 m NaCl at 25 °C Investigations by method (i) with scan rates from 20 mV/h up to 10 V/h resulted again in a shift of the Uc values to more noble potentials with increasing Mo content, decreas ing concentrations of NaCl, and increasing scan rate. There was, in particular, a difference up to 0.3 V, depending on the scan rate only. However, the reproducibility of the results as a whole was bad. In the case of all steel types and concentrations of NaCl with a scan rate of 0.1 V/h the Uc values were found in the range between 0.20 - 0.53 V. Taking into account the other scan rates applied, Uc was measured to range between 0.15 and 0.68 V. These small differences are due to the fact that in most cases by application of method (i) not the Uf, but the Ur values are achieved. There is a small difference between Uf and Ur only in the case of a steel with a low Mo content. The difference increases with increasing Mo content. Precise 122 Wilhelm Schwenk values of both Ur and Uf can only be measured by method (i) applying very small scan rates of about 20 mV/h. The Uf values of all steel types in 1 m NaCl were found between 0.25 V and 1.2 V; in the case of other concentrations of NaCl the values ranged between 0.17 V and 1.3 V. The be haviour of diffusion heat treated samples (1300 °C) was found to be similar. Compared with normally heat treated samples (1050 - 1100 °C) Uf is shifted to more noble potentials after diffusion heat treatment. Furthermore a precise differentiation between Ur and Uf is possible by applying a more rapid scan rate of about 100 mV/h. Con trary to earlier investigations with type 1.4301 we also observed repassivated pits in the rolled surface of the specimens, especially in the case of Mo containing steels. In view of these results we suppose that the beneficial effect of Mo is caused more by an increase of pit repassi- vation than by an inhibition of pit formation. The Ur values were found to be scattered at more negative poten tials in a wide range. The Up values obtained by method (iv) were nearly independent of steel type, heat treatment and concentration of NaCl. All values were found to range between 0.08 V and O.38 V. They have a similar meaning as the Ur values. There is only one exception, namely steel type 1.4449 with the highest Mo content after diffu sion heat treatment. Methods (i, iv) resulted in values of Ur, Uf and Up more noble than IV. We compared the test results of methods (i, iv) with those of methods (ii, iii). In the case of some systems under potentiostatic circuit conditions we found only repassivating pitting corrosion at potentials between Ur and Uf. All pits were passivated at potentials more negative than Ur. The results of method (iv) correspond to the Ur values. There are only small differences depending on steel type, heat treatment, and concentration of NaCl. From the results of systems II and III we can conclude that Uf is of practical interest to characterize pitting susceptibility. Test methods (i, ii, iii) are applicable. In order to differentiate between Ur and Up it is absolute ly necessary to apply very small scan rates in the case of method (i), e. g. about 20 mV/h for normally heat treated samples and about 100 mV/h for diffusion heat treated samples. By method (i) with a more rapid scan rate as well as by method (iv) one obtains Ur or Up values which from the practical point of view are not very interesting. K . rt . D l_ cl ii 11 J. *.: TRANSIENT STRAINING ELECTRODES FOR STUDYING STRESS CORROSION CRACKING R. W. Staehle, J- B. Lumsden, S. Abe Department, of Metallurgical Engineering, The Ohio State University, Columbus, Ohio, U.S.A. Models devised to explain the stress corrosion cracking behavior of many metal-environment systems assume that the central mechanistic issue involves the breaking and reformation of protective films. In those circumstances where the reformation rate of the film is suffi ciently slow one might expect circumstances which favor propagation of stress corrosion cracks. It is desirable to obtain quantitative measurements of the reforma tion rate of protective films. Two techniques which can be used to measure the amount of transient reaction are transient straining and the potential jump technique combined with ellipsometry. The. transient strain technique involves applying a transient strain, £ , to a wire specimen. The transient strains are usually in the range of 3.0% and are applied using a loaded spring. The potential i rs held constant and the resulting current transient is measured; the transient current and the strain transient are measured simultaneous ly using an oscillograph so that their instantaneous values may be compared. A peculiarity of this measurement is that the current- time response may be divided into two regimes. The first regime is associated with covering the non-protected surface made instanta neously available by the straining; the charge density of reaction in this portion, usually from zero to 20 ms, is designated as Qit . The remaining reaction until reaching the original background current is designated as Qix and relates primarily to the growth of the film. The effects of total strain on values of Q,;, and Qix are shown in Figure 1. Here the new area is taken as one-half the original area times the strain and Q" is compared as a function of £.^./,. The fact that the Ql, and Qlx are substantially constant and independent of €>xyt suggest that details of the surface structure involving slip morpho logy may not be too dominating. Figure 2 shows effects of strain rates. Again the values of Qit and Qt" seem little affected. Figures 1 and 2 suggest that the transient strain technique can be broadly applied to measure film growth phenomena and factors which affect film reformation kinetics. The potential transient technique monitors the current transient and ellipsometric transient simultaneously on an oscillograph. This per mits simultaneously monitoring the total reaction and the amount of film formed therefrom. The difference between these quantities is then the dissolution rate. Results from such calculations are shown in Figure 3. The dissolution rate becomes negative, i.e. precipita tion occurs for a short duration. This is then followed by a steady slate dissolution. 124 R.W. Staehle Type 304 E=+240mVH O. IM Na2S04(N2)pH7 é a7'25% /sec 10 1 i , 1 E 8 - •if u E 5 6 - •o """"""O" Qr^o °~ - o_0-o—o—-"""^ a _ 5 2 0 - 1 1 1 1 1 - 5 10 0 4 6 8 10 12 Time (msec) Total Strain (%) Figure 1. Type 304 stainless steel strained to various total strains. Current-time and extension-time at left. Ql and Qlx vs € at right. E= +240mVH O.IMNa2S04(Ns)pH7 £T=3.45 % 10 1 1 1 _ 8 E Q u i* E 6 - • © - O •a 4 -- 0i, o QL o a. 2 - a _cz O 0 - 1 1 1 0 200 400 600 800 Time (msec Strain Rate(%/sec) Figure 2, Type 304 stainless steel strained at various strain rates. Current-time and extension time at left. O; and o; vs £ at right. R.W. Staehle 125 Dissolution Rate vs. Time for Fe2 03 Film on Fe — Density - 5 g/cm o 0) in o< CO o CT. c o o Vi --leOmV. 0 0.20 0.40 0.60 Time (sec) o Dissolution Rate vs. Time for Fe203 Film on Fe - Density = 5g /cm 3 CD 0< 1240 mVH 'imfWBammAni^MN »A...- ..g _L 12 16 20 24 28 32 36 40 44 48 Time (sec) Figure 3. Dissolution rate of iron in a pH 8.6 sodium borate-boric acid solution determined at various potentials. 126 Peter Tarkpea EXPERIMENTS ON THE ACCELERATION OF THE INTERGRANULAR STRK .=•: -"ORuOSION CRACKING OF TYPE AISI 304 STAINLESS STEEL BY MEANS OF POTr,.w.'. ÖSTAT1C POLARIZATION. Peeter Tarkpea, Kjell Lundberg and Walter Hiibner. Section for Corrosion and Reactor Chemistry. AB Atomenergi, Studsvik, Sweden. It has been shown repeatedly during the last few years that intergranu- lar stress corrosion cracking (ISCC) may occur in austenitic Fe-Cr-Ni- alloys in various environments at elevated temperatures. Even in high purity water such cracking has been found. One general problem, when trying to .study the ISCC of austenitic alloys in high temperature water, is the very long time necessary for cracking to occur. We have started a series of experiments with the object of finding a method for potentiostatically accelerated testing for ISCC under con ditions, which as far as possible resemble practical high temperature applications, such as in water cooled nuclear reactors. Tests have been carried out with uniaxially stressed sheet specimens of commercial austenitic stainless steel AISI 304, with the following composition, %, C 0.043, Si 0.47, Mn 1.41, P 0.026, S 0.008, Cr 17.6, Ni 9.1, Mo 0.39, N 0.03. Both solution annealed, 1050 °C/5 min plus water quenching, and sensitized, 650 C/20 h, materials have been studied. The nominal stress was two times OQ,2 at 300 °C and the load was applied by direct loading of the sheet specimens with a lever system inside the autoclave. The experiments were carried out in high purity water containing 100 ppm oxygen at 300 °C in a refreshed 10 1 autoclave i.e., under flowing conditions, 1-2 1/h. The autoclave with its auxiliary equip ment used in these electrochemical tests is shown in figure 1. The water was analyzed frequently both before and after its passage through the autoclave. The chloride content of the water was less than 0.1 ppm. Sodium sulphate, 20 ppm, was added to the water in order to increase its conductivity. The tests were performed with potentiostatically controlled specimens, the specimens being coupled as working electrodes. The potentials were usually measured against an external mercury/mercury sulphate reference electrode. In some tests an internal mercury/mercury sulphate reference electrode was used. This electrode was kept under autoclave pressure but was mounted on the autoclave lid so that it was kept cool (30 °C). The liquid junction bridge, as well as the reference electro de, was car-efully cooled. A platinum net was used as a counter elec trode. This electrode was either placed in the autoclave or in a small separate autoclave, next to and quite near the outlet of the test autoclave. This was done in order to separate the anode and cathode compartments. In this way the effects of possible reverse electrode reactions could be excluded. Catodic and anodic polarization curves were recorded at 300 °C. They showed an active-passive transition. Figure 2. It was noticed that Peter Tarkpea 127 cracking was somewhat more pronounced in these transition regions. Thus cracks to a depth of 50-100 urn were noticed within less than 72 h with cathodic polarization to -670 mV and anodic polarization to 0, 100, 580, and 760 mV. One test specimen failed within 1 h with anodic polar ization to 580 mV. The material was polarized from the corrosion po tential. Cracks were only found in this material after sensitizing treatment, solution annealed material showed no cracking within testing times up to 150 h. In these experiments the counter electrode was placed in the test autoclave. In experiments, in which working- and counter electrode compartments where separated, no cracks were found within 120 h with anodic polar ization to 0, 300, and 600 mV or within 120 and 180 h at 800 mV. Only one pit was noticed in one specimen with 800 mV anodic polarization. The reverse electrode reaction was excluded in this case, which as far as we can see, seems to have decreased the susceptibility to ISCC. This phenomenon is under further study. In order to meet objections regarding the addition of sodium sulphate to the test solution some ISCC tests with unpolarized specimens in sulphate and sulphate-free enviroments have recently been started. Acknowledgement. This investigation was supported from the Swedish Board for Technical Development. FIG. 1. EQUIPMENT FOR ELECTROCHEMICAL TESTS A STANDS FOR AUTOCLAVE B OXYGENATION VESSEL C PUMP D SAMPLING POINTS FOR ANALYSIS E PREHEATER F COUNTER ELECTRODE G TEFLON CAPILLARY FROM REFERENCE ELECTRODE H TEST ELECTRODE J LOADING EQUIPMENT WITH LEVER SYSTEM K INSULATION OF ELECTRODE LEADS THROUGH AUTOCLAVE LID, WHERE Kl IS TEFLON INSULATED LIQUID JUNCTION WITH CELLULOSE PLUG AND K2 IS TEFLO INSULATION OF METAL WIRE LEAD L REFERENCE ELECTRODE CONNECTED TO LIQUID JUNCTION K THROUGH INTER MEDIATE VESSEL M POTENTIOSTAT N RECORDER 0 COOLER P CONSTANT PRESSURE REGULATOR i2B Peter Tarkpea -4*0 -532 {_ -1)000-- 103 I *l/ca2 5 6 7 8 910 3 4 567B9 10 •1 5 h r p 1 M Norbert Wieling 129 CORROSION TESTS WITH THE AID OF ELECTROCHEMICAL MEASUREMENTS ON NICKELBASE ALLOYS UNDER PRESSURE WATER CONDITIONS Norbert Wieling Siemens Aktiengesellschaft, Erlangen, Germany, It is at least problematical to predict the corrosion behaviour of metals in water of about 300°C with laboratory investigations at much lower temperatures and more aggressive media. Only when the test conditions agree with the service conditions of the materials in pressure water, more severe testing conditions at shorter times for the determination of the limits of the materials in long time service will become meaningfull. Electrochemical measurements in pressure water autoclaves make it possible to perform short time experiments under severe conditions. These experiments yield results, which are transferable and extrapolatable to the working conditions of the materials. For such experiments it is necessary to have a reference electrode, which gives a constant and known potential in the given temperature range. Electrochemical investigations on austenitic chromium-nickel-stain less steels and nickel-base alloys in pressure water autoclaves are described. For potential and polarisation measurements a newly de veloped high temperature reference electrode for pressure water was used. Design and performance of the reference electrode and the experimental set-up are described. The behaviour of the above men tioned materials as a function of oxygen and chloride contents and applied stresses were investigated and discussed. SECTION 1 ELECTROCHEMICAL THERMODYNAMICS Hermann Göhr lij THERMODYNAMICS OF SURFACE LAYERS Hermann Göhr Institut fiir Physikalische Chemie I, Erlangen, BRD The formation of layers on metal surfaces by reaction with a gas or - as treated here - by an anodic reaction within an electrolyte solution necessarily passes an initial stage where bonding of the non-metallic particle occurs on the atoms of the 'uncovered' surface of the metal. Ignoring polarization effects, the metal atoms remain at their sur face locations during the initial phase of this binding. Corresponding to the strength of the surface atom-particle bonding and of the particle-particle interaction different types of processes ma^ subsequently occur. In accordance with this, structurally different surface layers may result. Some thermodynamic aspects of such subrnonolayers will be discussed here. The type of surface layer that is formed depends primarily on whether or not surface atoms leave their places during the growth. They do so if the metal-metal bond is weakened enough by the metal-particle bond. In this case three- dimensional nucleation appeares. If, however, the metal atoms remain at their surface places, the coverage may take place in two" different v. . ys depending on the particle-par ticle interaction. In "Lha case of attraction insular regions of total coverage are favoured. Accordingly two-dimensional nuclei are growing. In the case of negligible interaction or repulsion, however, the particles will be distributed almost at random. Therefore a uniform degree of coverage of equivalent surface places results. The latter process may be called 'electro-chemisorptionf. The above mentioned three types of surface layers differ considerably with regard to their thermodynamic properties. In particular the entropy term causes characteristic dif ferences in the free energy of formation of the surface layer which depend on the average number of particles consumed per surface area. Since layer formation is an electrode reaction, we are able to estimate the free energy of formation by measuring the corresponding reversible potential. Thus, the different types of surface layers can be distinguished by studying the dependence of this reversible potential on the apparent degree of coverage, 0, which results from the charge in vested in the layer. A behaviour as follows is to be ex pected. The reversible potential decreases with growing three-dimensional nuclei. It will remain nearly constant during the growth of the two-dimensional nuclei. It in creases with increasing degree of coverage in the case of 134 T J ^ -v—, — -p ~< C. *-> T-» •*- electro-chemisorption. In the case of the three-dimensional nuclei the free energy approximately depends on the volume vM of a nucleus according to Thomsons formula - -1 / "\ AF = AF + Y V K v AJ:N ArMX l ? T VMX VN where V^ is the molar volume of layer substance MX,_K is a factor regarding for the shape of the nucleus,and Y is an average interfacial tension which depends on the interfacial tensions for spreading the nucleus on the metal surface in solution, and for increasing the nucleus-solution interface. Hence the difference of the reversible potentials with respect to nuclei and to bulk layer is e a 2 1/3 N " ^X Y VMX K.'(zT) vN~ where z is the number of faradays per formation of one mole MX. Thus, for a constant number of nuclei eN - ^X cc 0-1/3 On the contrary, the free energy of formation in the case of two-dimensional nuclei is nearly independent of the area covered^ since the surface is proportional to the amount of adsorbed substance, Consequently, € remains cimost constant until a monolayeris completed. Further growth is only possible at higher potentials. In the case of electro-chemisorption the electrode reaction M + X2" —> M-X + z e~ produces a uniform degree of coverage, ©, which depends on the activity a(X2"") of anions in solution according to an adsorption isotherm which, under certain conditions, may be described as 0/(1 - 0) = a(Xz~) K(e) exp(£U( 0)/RT) If the adsorption energy ZJU is independent of 0 a Langmuir isotherm results. Since the M-X-bond is certainly polar, the adsorption energy will become less negative with in creasing coverage, particularly when it approaches satu ration. As an additional important effect of polarity the electrode reaction is connected merely with partial trans fer of charge through the double layer. Therefore, only a fractional part of the change of potential contributes to the change of free energy. Thus z 6zF(£( 0) - e°) = RT (In 1 g Q - In a(X ")) - AU( 0 ) + AU° Assuming a low coverage and a constant concentration of Xz~ at room temperature we obtain the expression e(0~/n\) = 59mVnlog^ O~ +, cons.t A study of electro-chemisorption of 'oxygen' on silver in alkaline solution has provided results in good agreement with these conceptions. If we assume z = 1, from the experimentally estimated product 6z = 0.75 the reasonable value 6 = 0.75 is obtained. For 0 = 0.2 the reversible potential exceeds that of Ag~0 formation. But, as expected, no oxide is formed. Only at about half saturation does three-dimensional nucleation take place, and the reversible potential decreases. A probable example of two-dimensional Insular surface layer formation is found on Zirconium in perchloric acid solu tion. Here the adsorption of perchlorate ions on 7,r0_ surfaces is observed to occur at a certain constant potential. This reversible potential over a wide range is independent of the degree of coverage. 136 Zofia Borkowska THE ADSORPTION OE ACETONE AT THE :.ir;ricuii:r SOLUTION INTERFACE Zofia Borkowska and Barbara Behr Institute of Physical Chemistry, Polish Academy of Scien ces, Warszawa, Poland» The system water - acetone — Hg electrode has been inves tigated • The double layer capacity has been measured as the function of electrode potential. The electrode chargo and surfaoe tension curves have been evaluated by double integration. The zero charge potentials and the maximal surface tension values have been determined with the streaming mercury electrode and with the capillary elec trometer respectively. The problems of.suitable reference eleotrode and of the effect of the electrolyte used have been discussed. Acetone is a poor solvent for inorganio salts and there are only few electrolytes of sufficient solubility, e.g. NaClO. and Na- or NH.SCN which were used in this work. Up to 60 vol.$ of acetone HC1 has been used as the supporting electrolyte» The curves of the double layer capacity vs. electrode charge are shown in Pig» 1# It was found that in solutions of low water content the discharge of H* from HC1 solutions occurs at less negati ve potentials than in aqueous solutions. The same was observed for the discharge of NH* from NH. SCN. In HC1 solutions of constant activity equal to 0*1M the maximum of acetone adsorption occurs at ca« —5,w«oul/cnr /at nearly constant potential of oa« -730 mV vs. :1Bner Ag/AgCl elec trode/* The surface tension data have been analyzed according to the theory described by Randies and Behr4"^. At high aceto ne concentrations NaClO^ was used as supporting electroly te. For the uncharged Hg - solution interface the parame ters found to get the best fit to the experimental data /Fig. 2/ are: K = 0«8, £" = 5.1 where K denotes the ratio of the excess free energy of mixing of the solvent compo nents in the surface phase to that in the bulk phase of the same composition, and X - the thickness of the surface layer which is here about monomolecular. The relative sur face excess curves of acetone at the interfaces: solution - air and solution - Hg electrode at zero charge and at -5ucoul/cm2 are compared in Fig. 3. At low electrode charge the addition of Na - or NH. SCN to solutions of low acetone concentrations decreases and the addition of NaF increases the adsorption of acetone2\The vapour pressure of 20 vol* % of acetone was diminished by 3.5 mm Hg in the presence of 0.1 M NH.SCN and increased by 9 mm Hg in the presence of 0«1 M NäF ^ . Obviously the electrochemical effects of NaF can be interpreted as due to the increased acetone activity in the bulk phase, and Zofia Borkowska 137 the_effect of NH/( SCN as due to the surface activity of SCT: ions. At this p.cetone oo-nortTrc-v :• ».o-n *l_so the presen ce of HCl influences the adsorption of acetone. At higher acetone concentration /40-60 vol. %/ the electrolytes HCl, LiCl, NaClO. do not produce significant changes of either zero charge potential or surface tension. The double layer structure at the electrode-solution inter face in the case of various electrolytes has "been discus— sed» 1. J.E.B. Randiesj Ba Bear, J» Electroanalo Chem. in press 2. see also B, Behr, J. Dojlido, J0 Stroka, Coll. Czecho- slov. Chenu Communs. 2£, 1317 /1971/. 3. data of T. Treszczanowicz, 138 Zofla Borkowska Hg!h'oO-acetone k20 1 "ra^rå W >v ° \ /t=QS9 *iD0 o\ XP K=08o 330 X. nna -S/tC/art* 380 _ 1 ' • * jr am 0.1 éXfaJ,CO F*9-2 Jt/kHW* Bri an Burrows 1 NITRATE TETRAHYDRATE AND ANHYDROUS NITRATE MELTS. Brian Burrows* and Soefjan Noersjamsi** School of Chemistry, Macquarie University, North Ryde, N.S.W. Australia (^present address ; Battelle, Geneva Research Centre, Geneva, Switzerland : ** permanent address } Chemistry Department, Bandung Institute of Technology, Bandung, Indonesia). The properties of hydrate melts and concentrated aqueous solutions have been attracting increasing attention in recent years due, in part, to the fact that these solutions form an interesting link between the more conventional aqueous solutions (< 2-3 molal) on the one hand and ionic melts on the other. Experi mental data so far obtained, especially for systems of the type Ca(NO3)2.4H20 and MgCl2.6H20, suggests that these hydrate melts (1-5) resemble more closely ionic melts than aqueous electrolytes. In the work reported here the aim was to compare partial ionic entropies for the dissolution of silver halides in molten calcium nitrate tetrahydrate (obtained via solubility product measurements) with similar entropy data in water and anhydrous nitrate melts. In order, however, to make inter-solvent comparisons of thermodynamic data the choice of standard state is critical. This was first recognized by Gurney (6) and later by other workers (7, 8) in the context of ion- association. The choice of standard state is important since, in the case of solubility studies of silver halides for example, A S° values calculated from molal scale KSp data will contain an arbitrary constant, 2Rln(1000/M), which will vary with the molecular weight (M) of the solvent. By comparing, on the other hand, standard unitary entropies i.e. entropies based on a hypothetical mole fraction of unity standard state for the solute, one is comparing states of minimum configu- rational free energy. In other words the influence of the composition of the solution is removed and the entropies are primarily associated with the purely loca changes in the entropy of the solvent in the immediate neighbourhood of each added ion. 140 Brian Burrows 11 e poluhi 1. if V orodliof P of Aa^ 1 «OKI- PPO MCT i in mo i.r.en Ca(NO3)2•4H20 were determined potentiometrically over the temperature range 4 0° to 80°C. From the data standard partial ionic unitary entropies of solution were calculated and will be compared with similar entropies of AgX calculated from literature data in H2O (9), molten (Li,K)N03 eutectic (10) and molten (l-\!a,K)NC>3 eutectic (11). The same corresponding tempe rature (Tmp + 25) was chosen for the inter-solvent comparison. It will be shown that the relative partial ionic unitary entropies are markedly similar in the hydrate and anhydrous nitrate systems showing no significant trend from Cl~ to I- in contrast to the behaviour in aqueous solution. This suggests that the hydrate melt is composed of Ca(H20)g2+ complex cations and NO ~ simple anions in support of the conclusions of previous investigators (±-5) based on other experi mental approaches. References 1. C.A. ANGEL, J, Electrochem. Soc, 112, 1224 (1965) 2. J. BRAUNSTEIN et al, J. Chem. Eng. Data, 12, 415 (1967) 3. J. BRAUNSTEIN et al.,J. Phys Chem, , 10, 2734 (1966) C.A. ANGEL and D.M. GRUEN, J Amer, Chem, Soc. 88, 4. 5192 (1966) C.T. MOYNIHAN and A. FRATIELLO, J. Amer. Chem. Soc 5. 8£, 5546 (1967) R.W. GURNEY, Ionic Processes in Solution, McGraw- 6. Hill, New York (1953) pp. 90-97. 178-181 7. A.W. ADAMSON, J. Amer. Chem. Soc, 7(5, 1578 (1954) 8. H.A. BENT, J. Phys. Chem., 60, 12 3 (1956) 9. B.B. OWEN and S.R. BRINKLEY, J. Amer. Chem. Soc _60, 2233 (1938) 10. H.T. TIEN and G.W. HARRING TON, Inorg. Chem., 2_, 369 (1963) 11. S.N. FLENGAS and E R. RIDEAL, Proc. Roy. Soc. (London) A233, 4 43 (1956. Gérard Douhéret 141 THERMODYNAMICS OF MIXED AQUEOUS SOLVENTS : TRANSFER PARA METERS o Gérard Douhéret Laboratory of General Chemistry, University of Clermont-Fd France. The transfer coefficients, or medium activity coefficients or distribution coefficients, are usually turned to account in the purpose of setting a connection between activity scales when reference states are allowed to change, on a general point of view. On an other side, the operational procedures leading to activity measurements, lean to introduce quantities, called transfer parameters, which bind the "actual" value of the activity of an ionic species i, reported to a reference state, x, and the "apparent" value of the same, computed from experimental data» referring now to r. These two states characterize the transfer, t, which must not be looked on the limita tive aspect of a solvent change, as often done. Thus, the transfer coefficient appears as a theoretical limit of the transfer parameter. The following relations have been written, the transfer coefficient being defined as : t t yi gi = 6(ir} which means it includes, inside the 6 operator, the va riation of the standard chemical potential of i (reported to k, equals RT In 10). Such a definition may fit to va riations of k occurring during the transfer (for instance, nonisothermal processus). The transfer parameter becomes: (fcd ). = V - z.dr.t6(-^-^) + z.f'^g ,„.} g x i K a \sx) applied to a galvanic cell, involving a glass-membrane electrode, M, and a reference electrode, R, with a salt bridge, both giving place to diffusion and interface tensions, $.. and $ 'MR , Whereas <3>M is generally neglected, previous work , taking into account the differential equation of interface ten sions, has shown that $_. is able to be split into its JK. main components: a transfer tension, ($+-)R/ issued of the variation "of standard chemical potentials of the whole species j in solution, on both sides of interface, and a diffusion tension ($J)R r due to the heteroionic character 142 Gérard Douhéret ef the system * s ^o-r -• . f '• { c. .._,). iL reflects the possible variation of the salt bridge ions activity. So, we have : ( the latter terms being often small towards the two former. In the same way as t g. , (t d ). is related to free ent- i g i halpy variations of i occuring at the transfer, due to the definition of chemical potential. Moreover, implica tions corresponding to other ionic species j will affect it. Transfer parameters, originally regarded as terms applied to the calibration of peculiar apparatus, may be considered as transfer thermodynamic entities and genera lized to any ionic species. Both transfer coefficients and parameters exhibit a "compensation" effect, which allows to postulate for the late, enthalpic and entropic t t t 2 components, viz. { d, ) . and ( d ) . , likewise of g. So, ( d.). and ( d ). would be respectively representative of solute-solvent interactions, and of the "rigidity" of the final system as regard to the initial one, in relation to the nature of charged species. As a whole, ( d ). may nevertheless describe interionic forces and ion-solvent interactions. The variations of ( d ). in terms of any function of ionic g i J concentration, lead to an "extrapolated" value, namely ( d°)! , corresponding to infinite dilution (in fact, we g i must consider ions originating of the solvent autoproto- lysis). ( d ). and, at a smaller extent, ( d°)! depend upon the model chosen for the description of interionic effects, the most known remaining that of DEBYE-HUCKEL, where several adjustable parameters are used, e.g. ionic closest approache distance, a . Measurements proved quasi-linear laws (but non linear, as 3 t postulated by ARNAUD ), when ( d ). (i=H, in this case) is y -*- allowed to vary in terms of the above variables; depar tures from linearity became important for organic solvent rich systems, particularly for highly acidic solutions. The extrapolated parameter will enable us to follow the evolution of ion-solvent interactions on the whole range of composition of a binary system, in connection to its Gérard Douhéret 143 sLrucLural features . Consequently.- rr laws or vari :'cion of ( d°)! in terms of mole fraction x, of organic cosol- vent, have been analyzed for several mixed aqueous systems water-protic solvent and water-dipolar aprotic solvent; evidence has been mostly demonstrated of the stabilisation effect in water-rich regions^, corresponding to enhance ment of tridimensional water-lattice by insertion of organic molecules, giving place in some cases to the for mation of clathrates. In such media, the hydrogen ion undergoes a preferential solvation towards water decreas ing in the order : water-DMSO>-DMF>-C2H5OH>-CH3OH>-dioxan- ne^-acetone, THF>-ethyleneglycol>-ACN. A subsequent addi tion of solvent leads to a progressive destructuration through a microheterogeneity region-* (water-poor solvents) the proton being less solvated than in aqueous medium. 1G.DOUHERET and G.DURANTHON, J.Chim.phys., j68_,649 (1971) 2J.P.MOREL, J.Chim.phys., 62,895 (1970) 3R.ARNAUD, Bull.Soc.chim., 2782 (1971) F.FRANKS, Quart.Rev., 20,1 (1966) 5Yu.I.NABERUKHIN and A.ROGOV, Russ.chem.Rev.,40,297 (1971) 144 HadeSda Jakovl^evié Halai THE TRANSFERENCE NUMBERS JL)N AQu&vOS SCuuriGNS Or is i TYPE SALTS AS A FUNCTION OF TEMPERATURE AND CONCENTRATION Nadezda Jakovljevid Halai Faculty of Technology and Metallurgy,Beograd,Yugoslavia This is a mathematical analysis of the temperature and con centration dependence of transference numbers of cation in aqueous solution 1:1 type salts. A number of useful equations have been proposed for repre senting transference number data from the lowest concen tration (0.01 M) up to 0.2 M. The large number of temperature effects data upon the transference number can be found tabulated in literature but without any functional interrelationship. The commom effect of concentration and temperature, as we found, could be presented by an equation of the following type: nc,t - A'fc + B^ + nc,t=0 (1) where A= and B= <-$£>c Hn7c>t The coefficients of the equation (1) were evaluated from a statistical analysis of the experimentally known values in wide range of data (0-50 C, and molarities 0-0.2 M) . Some of the values of the temperature coefficients were estimated from the temperature dependence of limiting ionic conductance (eq.2). These are in good agreement with those obtained from a statistical Analysis. ± / + dn *> + - ,dlnAo dlnAo» /ON dt = n ' n (~~dt- " —dt"5 (2) The coefficient A,B or both for some salts are presented in -Table 1, where calc.denotes calculated values after eq.(1) and approx.denotes estimated values after eq.(2). Coefficient A does not depend upon concentration and B does not depend upon temperature within the experimental error. Calculated values of the cation transference numbers after eq.(1) are in good agreement with the values obtained ex perimentally (± 0.0003 units for NaCl-H20; ± 0.0002 units for KCl-H2Oj ± 0.0001 unit for KBr-H20 etc.) in the same temperature and concentration range. Table 1. The Coefficient A and B and n n of the Eq.(1) for Aqueous Solutions of some 1:1 Type Salts ' ~ 4 -1 A.10 unit deg concentration n ,_ unit salt B.lo\, z\ c,t= A0 calc. approx. unit ylit mol range Nacl 3.73-0.11 3,6 -2.42-0.09 0-0.1 0.3877-0.0002 KC1 -1.79-0.11 -1.9 -0.20-0.05 0-0.1 0.4954-0.0001 LiCl 5.77 -4.10 0-1.0 0.3221 NH.C1 1.93 -0.08-0.04 0-0.2 0.4861 4 KBr -0.79 -0.7 -0.66 <0.02 0.4868 0.44 >>0.02-0.2 0.4833* KJ -0.54 -0.19 <.0.05 0.4901 0.22 ^0.05-0.2 0.4887* KN03 1.0 0.97 0-0.2 AgNO 2.7 1.37-0.16 0-0.i 0.4548** NaAc 0.5 1.86-0.04 0-1.0 0.5491** KAc -2.5 3.99-0.19 0-0.1 0.6482** _ * extrapolated values r* estimated values 146 NadeSda Jakovljevié Halai REFERENCES 1. Allgood,R.W.,Gordon,A.R., J.Chem.Phys., 1J), 124,1942 2. Allgood,R.W.,Le Roy,D.J.,Gordon,A.R.,J.Chem.Phys., 8, 418,1940 3. Longsworth,L.G.,J.Am.Chem.Soc.,5^, 2741,1932 4. Keenan,A.G.,Gordon,A.R.,J.Chem.Phys.,11,172,1943 5. Longsworth,L.G.,I.Am.Chem.Soc.,57,1185,1935 6. Mc Innes,D.A.,Cowperthwite,J.A.,Chem.Reviews,11,171, 1932 7. Le Roy,D.J.,Gordon A.R.,I.Chem.Phys.,6, 398,1938 Jiri Janata 147 POLAROGRAPHIC DETERMINATION OF HYDROGEN ION ACTIVITIES IN STRONGLY ACIDIC MEDIA : A NEW ACIDITY FUNCTION Jiri Janata and Geertje Jansen Imperial Chemical Industries Limited, Petrochemical & Polymer Laboratory, Runcorn, Cheshire, England. A new electrochemical method for determining the activity of hydrogen ions in strongly acidic media has been developed. The potential of a glass electrode is measured against the half-wave potential of ferrocene at a dropping mercury electrode. In this procedure the glass electrode is connected to the reference electrode input (voltage follower) of a three-electrode polarograph. The oxidation (or combined) wave of ferrocene is then recorded for media of various hydrogen ion activities. The "apparent" shift of the ferrocene half-wave potential is due to the change of the potential of the glass electrode. The plot of the potential difference E against pH at 25°C has a theoretical slope of 59.1mV/pH. As there is no liquid junction potential involved in our method the measurement of E can be extended into concentrated acidic and mixed solvent solutions and the measured hydrogen ion activity can be referred to the aqueous standard state. With this procedure we have established a new acidity function, H , in several aqueous and mixed solutions (Table 1). 148 Jiri Janata Table 1 -H „ at 25°C GF M H2S04 HC104 H2S04 H2S04 H2S04 H2S04 20% 60% 40% 75% water water ethanol dioxan dioxan CH3CN 0.25 -0.64 -O.55 -1.33 -3.33 -2.15 -2.78 0.50 -O. 28 -0.12 -0.90 -2.98 -1.80 -2.25 l.O +0.31 +0.40 -0.46 -2.30 -1.19 -1.17 1.5 0.90 0.91 -0.29 -1.63 -0.66 -0.12 2.0 1.30 1.32 -0.10 -1.05 -0.28 +0.98 2.5 1.70 1.76 +0.22 -0.55 +0.27 2.07 3.0 2.11 2.24 0.56 +0.05 0.75 3.11 3.5 2.50 2.62 0.91 0.9O 1.38 4.21 4.0 2.84 3.02 1.31 1.76 1.92 5.28 4.5 3.25 3.46 1.75 2.58 2.43 6.38 5.0 3.62 4.00 2.12 3.26 2.96 5.5 4.09 4.48 2.60 3.85 3.52 6.0 4.45 4.95 3.10 4.10 6.5 4.85 5.46 3.69 4.61 7.0 5.26 6.04 4.28 O • J-. O 7.5 5.68 6.70 4.92 5.70 8.0 6.10 5.60 6.23 8.5 6.46 6.30 6.78 9.0 6.88 6.98 9.5 7.31 7.66 10.0 7.81 8.40 10.5 8.23 11.0 8.75 Claude KLAPISZ 149 STABILITY OF LAMINAR FLOW OF LIQUID MIXTURE WITH TEMPERATURE GRADIENT Claude KLAPISZ et Jacques CHANU Laboratoire de Thermodynamique des Milieux loniques et Bio- logiques - Université PARIS VII - 2, Place Jussieu PARIS 5eme - France When a vertical temperature gradient is imposed across an initially uniform liquid mixture confined between two rigid horizontal boundaries with a laminar flow, concentration differences generally appear. This phenomenon is know as thermal diffusion and called the SORET effect with laminar flow [1] . At the stationary state, the variation of concentration Ac is in the forme : Ac = D' e.z (1) —o O" In equation (1), c denotes the initial concentration, 3 the temperature gradient (3=-dT/dz), z the vertical coordi nate, D the isothermal coefficient of the solution, Df the thermal diffusion coefficient to the solution; D'/D is called the SORET coefficient. When 3 < 0 (ascendant gradient : Heated from above) and D'/D >0, it is well known that the relative variation of concentration increases with the temperature gradient. But, in the general case, equation (1) is available only when the absolute value |3|does not exceed a critical value. When |3|exceeds this critical value, the liquid layer becomes unstable. As shown by Lord RAYLEIGH, what decides the stability is the non dimensional parameter R _ gagd — (2) a - - ——~~KV ~ 150 Claude PCLAPISZ where g denotes the acceleration due to gravity, d- the depth of the layer, a the coefficient of volume expansion, v the kinematic viscosity, K the thermal diffusivity. The parameter Ra as defined here is called the Rayleigh number Rayleigh further showed that for a unic fluid (DT/D = 0), instability must set in when R exceeds a certain critical a value Rac. The detailed study of a unic fluid stability can be found in the very complete monograph by CHANDRASHEKAR [2] . All this considerations can be summed up with the above curves taken out of the recent book by GLANSDORFF and PRIGOGINE [3] . R CRITICAL STABILITY CURVES WITHOUT A FLOW Stability unstable R > 0 : Heated from a below (3>0) O DT/D R < 0 : Heated from area above (g<0) unstable So what we must find, are the critical Rayleigh numbers for different flow speeds (or different REYNOLDS numbers), Consequent ly we use the PRIGOGINE-GLANSDORFF variational technique : they established a general evolution criterion in macrosc opic physics, written as an integral, called "local pot ential", which can only decrease with time and thus takes 'its minimum value at the stationary state. It is then po ssible to answer a large class of problems in macroscopi c physics by solving a variational equation [4,5]. We written a general local potential for the study of the Claude KLAPISZ 151 stability of a liquid mixture heated l- m above or from below, with DVD positive or negative _nd with (or without) a låi.inar flow. With the use of very simple trial functions for the disturbance amplitudes, we can find the critical Rayleigh number in function of the concentration, the Reynolds number and the Soret coefficient. Our researches are based on electrolytic solutions, so we studied more specially the case of hydrochloric acid (Dr/D > 0), and we found that the Rayleigh number is always positive; it decreases when the concentration increases (with a Reynolds number kept constant) and increases when the Reynolds number increases too (with a concentration kept constant). REFERENCES [l] - G. THOMAS, J. de Chimie Physique, 1956, P.407 [2] - S. CHANDRASEKHAR, Hydro dynamic and Hydromagnetic stability, Clarendon Press, Oxford 1961 [3] - P. GLANSDORFF et J. PRIGOGINE, Structure Stabilité et Fluctuations, Masson, PARIS, 1971 [4] - J.K. PLÄTTEN These de Doctorat, Université de Bru- xel les, 1971 [5] - J.C. LEGROS, These de Doctorat, Université de Bru- xelles, 1971 152 Jaroslav Kuta ADSORPTION OF TETRAPROPYL- AND TETRABUTYL-AMMONIUM CATIONS AND HIGHER ALKYLSULPHONATE ANIONS ON MERCURY J. Kuta and I. Smoler J. Heyrovsky Institute of Polarography, Czechoslovak Academy of Sciences, Prague. From the effect of highly adsorbable surface active substances /SAS/ on electrode reactions at the DME some adsorption parameters have been obtained . For a better understanding of the inhibition or acce leration influence of these SAS on electrode processes complementary measurements using direct methods are very important. Hence capillary electrometer, drop-time technique and the measurements of differential capacities' /C-E curves/ were used. Simultaneously the applicability of each of these methods in the case of highly adsorbable SAS was followed. The problem of faulty measurements with capillary electrometer espe cially at potentials more positive than p.z.c. was reported for pure dilute electrolytes /up to 0.15 M/ several times and an interesting qualitative approach of understanding was given by Trasatti . A still more peculiar behavior is encountered with highly adsorbable SAS using capillary electrometer as e.g. extremly large time effects to reach adsorption equilibrium, deterioration of the properties of the capil lary with time, hysteresis, sticking^»^ of the meniscus etc. The com plexity of this problem will be discussed. The preferential adsorption of the surfactant on the layer of this surfactant firmly attached to the wall of the capillary might explain some of the faulty results. For such surfactants the drop-time technique offers mostly better results although its precision is smaller than that of the electro- capillarometer if it works properly. Another problem arises here for small concentrations of surfactants connected with the attainement of the equilibrium adsorption values during the drop-time. The differen tial capacity measurements give the most reproducible results. With tetrapropylammonium /TPA/ and tetrabutylammonium /TBA/ cations in 0.1N H2SO4 and 0.1N HCIOA and decylsulphonate anion /DS/ reasonably reproducible ^P -E curves could be obtained even using electrocapilla- rometer whereas with dodecylsulphonate /DDS/ the drop-time technique is to be preferred. Tetjrapropyl- and tetrabutyl-ammoniuix^c^tj^ons The p.z.c. is shifted towards more positive potentials with the con centration of TPA or TBA respectively and the plot p.z.c. vs log c^ fulfils the Esin-Markov equation. The maximum adsorption appears to be at^v-0.8 V /SCE/. For both surfactants the adsorbability is larger in perchloric that in sulfuric acid /Fig. 1/. This effect is apparent ly due to the increase of the induced positive charge in mercury which is higher in the case of sulfate than perchlorate anion. The J^ -value of TBA at p.z.c. determined from Gibbs equation at p.z.c. was about 5 % higher than the value reported by Damaskir. et al.°j' /2.8 x 10~^^mol.cm-2/. Also £^ -values obtained from the inhibition effect on i-t curves were practically the same. For a Jaroslav Kuta ^53 closely packed film of TBA, jj~ calculated from the geometrical mo del0 /r=4.94 Å/ agreed well with the experimental value. However, when inserting the value r=3.81 Å resulting from scaled particle theory of salt effect" the disagreement is quite large. This pro blem is still larger in the case of TPA. JO^ found experimentally /2.2 x 10~"-'-^mol .cm~^/ is nearest to the calculated value for r=4.5 Å °/2.6 x 10~l^mol.cm-^/, however, deviates considerably for other values of r given in the literature". Hence the conclusion if also in the case of TPA a close packed film is formed, is hampered by the uncertainty in revalues. Decyl- and dodecylsuj-phonate Higher alkylsulfates and to some extent sulphonates were studied using capacity measurements »•'•. With DS the *$ - E curves using capillary electrometer were still reasonably reproducible /Fig. 2/. In the plot 't* -log C^ a monoto nous decrease was observed. C-E curves also at the concentration 10""% did not show any anomalies /vide infra with DDS/. The experimental value of J^ in 0.1 N ^SOA was near to tne value 3 x 10~10mol.cm"^ at p.z.c. From the studies at water-air interface for most alkylsulfates /C«-Cng/ which are very similar in behavior to sulphonates, the area S per molecule for concentrations lower 2 than the critical concentration of micelle formation CK was S™49-52 Å independent of the chain length which corresponds to the calculated XJ = 3.3 x l0-10moi.cm~2/for S = 50 Å 2/. This value coincides with that on mercury interface. Because the cross-section-^ of the alkyl chain is 19.3 Å 2, the film of DS is not close packed if it is assumed that the sulphonate group is directed towards the solution. An inter action with water^ is apparently to be taken into account. As already mentioned$ the^ -E curves for DDS are reproducible using drop-time technique only. The $ -log C^ plot is up to 10"^M quite similar to that of DS, i.e. after small changes in 4* a sudden but almost constant change in the slope occurs. At the concentration near to 10"-% the curves start to be flattened out. Similar behavior in the case of Triton X-100 was explained by micelle formation^. On C-E curves of 10""-% DDS the differential capacity decreases already starting from positive potentials. At -0.25 V /SCE/ a sudden peak occurs, the height of which does not exceed the capacity value of the supporting electrolyte. This peak is falling to a minimum value of capacity in the potential range -0.4 to -0.9 V followed by a normal desorption peak at -1.15 V /SCE/. A similar positive peak with dode- cylsulfate11 at somewhat more negative potentials was attributed also to the micelle formation. The J[J* -value approximately determined from a sudden decrease in$* -log C^ plot where micelle formation does not take place in an appreciable way is again near to 3 x 10~10mol .cm-2 in accordance with the results at water-air interface below CK. 154 Jaroslav Kuta Literature 1. J. Kuta, Rev.Polarography /Japan/ 11, 62 /1963/. 2. R. de Levie, A.A. Husovsky, J.Electroanal.Chem. 20, 181 /1969/. 3. S. Trasatti, J.Electroanal.Chem. J31, 17 /1971/. 4. S. Sathyanarayana, K.G. Baikerikar, J.Ele.ctroanal .Chem. 21,449/1969/. 5. R.G. Barradas, F.K. Kinraierle, J.Electroanal.Chem.. JL^, 128 /1966/. 6. B.B. Damaskin, S. Vavricka, N.B. Grigor'ev, Zh.fiz.khiin. 36, 2530 /1962/. 7. B.B. Damaskin, R.I. Kaganovich, V.M. Gerovich , S.L. Dyatkina, Elektrokhim. _5, 507 /1969/. 8. R.A. Robinson, R.H. Stokes, Electrolyte Solutions, p. 120, But- terworthS;, London 1955. 9. W.L. Masterton, D. Bolocofsky, Tei Pei Lee, J.Phys.Chem. T5_y 2809 /1971/. 10. K. Eda, J.Chem.Soc./Japan/, 8£, 349, 461, 465, 708 /1959/; 8^, 689 /1960/. 11. B.B. Damaskin, N.V. Nikolaeva-Fedorovich, R.V. Ivanova, Zh.fiz. khim. 34-, 894 /1960/. 12. H. Kölbel, P. Kurzendörfer, Fort.Chem.Forsch. _12, 279 /1969/. 13. N.K. Adam, Proc.Roy.Soc. /London/ A1Q1, 452, 516 /1922/. 14. J.O'M. Bockris, D.A.J. Swinkels, J.Electrochem.Sdt. 111, 736 /1964/. Jaroslav Kuta 155 Fig, 1 4*— log C dependence 1/ TPA,O.IN HCIO E=~0.45V; 1/ TPA,O.IN H2SO4,E=-0.45V; 2/ TPA,O.IN HCIO, ,E—0.8V; 2/ TPA, O.IN H2S04,E=-0.8V 3/ TBA,0.1N HC104,e=-0.45V 3/ TBA,0.1N H2SO4,E=-0.45V 4/ TBA,0.1N HCIO,,E=-0.8V; 4/ TBA,0,1N H2SO,,E=-0.8V -2 log cA /SCE/. dynes Fig. 2 7t Capillary electrometer 420 - 1/ O.IN H„S0 • cone, of DDS 2/ lxlO~5M; 3/ 3xlO~5M; 4 4 400 4/ lxlO~ M; 5/ 3xlO~ M; 6/ lxlO~3M; 7/ 3xlO~3M; 8/ lxlO~2M; 9/ 3xlO~2M. 380 -7-0 WSCE; 156 J anus z LEK KI A METHOD OF ESTIMATING THE REACTION OF ELECTRODES AT THE SURFACE OF SULPHIDE MINERALS IN CONDITIONS OF FLOTATION janusz Lekki. Institute of Mineral Processing and Coal Preparation Technical University of Silesia, Gliwice, Poland The papers presented by Sato /l/ have made it possible to determine the reactions of electrodes occurring in the course of the first stage of oxidizing sulphide minerals. In the case of Sato's method the potentials of the mineral electrode E and of the platinum electrode Eh are measured simultaneously in a solution which contains a reagent, passing the oxide thorough this solution at a given pH.The equality of E and Eh justifies the attri bution of the value E to some given thermodynamic equilibrium and makes it possible to forsee the electrode reaction which is most likely to take place. The knowledge of electrode reactions is extremely important for the flotation of sulphides. Such a flotation system consists of the molecules of some sulphide in water and remains in continuous contact with air. In order to predict the electrode reactions which are possible at such conditions the following method has been applied. Measuremets of Eh and of the potential of the electrode E are taken at such conditions as exist during the microflotation tests , i.e. in 50 ml redistilled water containing 2g of fine-grained mineral 0,2 - 0,075 mm. The suspension is being stirred with a free access of air and the magnitudes of the potentials E and Eh are read off after 6 minutes on a millivoltmeter. Pyrite and galena were subjected to such investigations from deposits found in the neighbourhood of By torn, as well as synthetic chalcocite. The results of these measurements are presented in Figs. 1, 2 and 3. As may be gathered from Figs. 1 and 2 in the case of chalcocite and pyrite an equilibrium is achieved between the solu tion and the mineral : the results thus obtained confirm the reactions quoted by Sato, whosf equilibria have been denoted in the Figs, by the straight lines 1 and 2. These results allow us to postulate the creation of metal hydroxides at the surface of the investigated minerals. As may be seen from Fig. 3, in all the experiments with galena no equilibrium has been obtained. The values of the potentials Eh prove that the solution can oxidize the galena to lead hydroxide, but an equilibrium is not reached due to the formation of layers leading to electrode passivation, being themselves products of the oxidization of sulphur. The equilibrium curves for these reactions have been presented in Fig.3 in the form of the straight lines 2,3 and 4- after 157 Ja nusx LE K K I ['open and folun /2/. The relative coi ne ulene o el p.o ox pr r i mrr i a I data with the results obtained by Toperi and Tolun, who needed 2,3 hours for the attainment of rest potential, proves thai by investiga ting E and Eh in the presence of the mineral grains an equilibrium is readied in a rather shorter lime. V V .++ V , /J 1. Cu2S=oT+CuS+2e 1. FeS =Fe + S +2e 2 z (SCEi (SHE! (SCEXSHE + Eh =0.530 0.0295 lg(CiT) Eh =0.757+0.02 95 lg(Fl ) 0.6 03 0.6 + 2. Cu,S+2H20 + CuS+Cu(OH)+2rf* 2. FeS/3H20=Fe(OH)5+Sa+3H +3e 0.5 2e Eh =0.057 +0.0591 pH Eh=0.$02-0.0591pH 0.2 0.5 0Å 107 10 0.1 0.4 03 8 0.2 h 0.1 6 7 <& NaOH- "> Fig.l. The influence of pH on Fig,2. The influence of pil on the potential of a chal- the potential of a pyr i to cocite elect.E and a Pt elec ir. E and a Pt electr. Eh electr. Eh In the investigated systems no reagents have been introduced, which might change the redox potential. The evident potential changes prove that the oxide contained in the suspension is used up and that an equilibrium is reached between the sulphide minerals Cu^S and FeS2 on one hand and their hydroxides on the other. System may reach such equilibria determined by redox reactions by means of several intermediate reactions including chemical, ones. Thus it seems to be correct if the aforesaid results are being interpreted while taking into consideration the scmiconducti vc pro perties of sulphide minerals. Plaksin and Shafecv /3/ have shown that in the case of semiconductive sulphide minerals a cathode reaction takes place : |o2 + 2 e + H20 2 01 158 Janusz LEKKI Assuming that the formation of hydroxides is an anode reaction, both these reactions may be interpreted schematically, as it is the case in Fig. L. o Q S O 4CV2 X M \ S CA T M S e M S o MJ -OH Q f O <--S~> 2T M- -OH < 6 7 NaOH Fig.3.The influence of pH one the Fig.4. Diagram of the reaction potential of a galena electr. of oxidizing sulphide E and a Pt electr. Eh minerals in alkaline solutions . This diagram makes it clear that the process is a result of the catho- dic reaction, which has been postulated by /3/. The formation of hydroxides depends on the removal of sulphur from the place where the reaction is taking place, which - as we know - may be achieved in the case of chalcocite by diffusing CuS; in the case of pyrite it is molecular sulphur /l/. The total effect of these reactions is the attainment of a state of thermodynamic equilibrium, which may be proved by means of the suggested method of measurement. This is a general method and may be applied also for the purpose of inwesti- gating other sulphide minerals. In the conditions of flotation this method may be used in order to evaluate the reaction with xanthate. If the suggested interpretation of the measured magnitudes is accep ted, the competitive effect of xanthate ions on the hydroxide ions be comes obvious. References 1. Sato M.Economic Geology, 55, I960, 1202-30 2. Toperi D. and Tolun R . Trans . Instn.Min.Metall /Sect.C : Mine ral Process. Extr. Metall/ 78, 1969, C 191-7 3. Plaksin I.N. and ShafeevR.Sh. Trans. Instn. Min.Metall 72, July 1963, 715-22. Zbignlew Moser 159 DETERMINATION OF TEE EQUILIBRIUM CONSTANT OF THE SIDE REACTION OCCURING- IN THE CONCENTRATION CELL Zbigniew Moser and Erzysztof Fitzner Institute for Metal Research, Polish Academy cf Sciences, Krakow9 Reymonta 259 Poland Research on the evaluation of the thermodynamic properties of liquid metal solutions by the method of the e.m.fo of concentration cells results, when the alloy components take close position in electromotive series, in some error as side reactions are likely to occur* Wagner and Wer ner^ ) among others were concerned with the problem of side reactions in concentration cells» The influence of side reactions occuring between the com ponents of liquid electrodes and those of electrolyte result in decreased e.nuf. values and in the increased value of activity coefficients especially when the inves tigation deals with dilute solutions and higher temperatu res. This phenomenon was, for instance shown, when inves tigating dilute liquid zinc solutions in cadmiumv2) where the following reaction is possible Zn + CdCl2 *=*• Cd + ZnCl2 [l] in the concentration cell: C-)Zn\ZnCl2 in the eutectic mixture LiCl-KCl Zn-Cd(+) [2] The equilibrium constant of the reaction [l~\ while assu- ming X[cdl+ XLzn] = 1 and X^^y*- X(ZaC1 j = 1, is gi- 1 X ' " (CdCl0)\ KtCd LV J HZnCl0) 2' ' ^ ^—2' r X 3 CCdCl2) ' T[zn] H0d012) where: the brackets t Icorrespond to the metal phase and ( ) to phase salt respectively. On the basis of formula [3 J for the calculation of the equilibrium constant the following parameters ought to be verify: ^* AJI| or K of the reaction [1~] while assuming the solutions as ideal II» In l ig -i and lnjf rh^\ obtained from paper III* In (f (znCl ) an(^ "*"n ö (CdCl ) ^2^ ^s *^ne thermodyna- mic properties of liquid ZnClp and CdClp solutions, obtained experimentally in this paper. 160 Zbigniew Moser Research was carried out on liquid ZnClo and CdCl^ solu tions with small additions of both chlorides in the eutec- tic mixture LiCl-KCl. The method of cell formation of chlorides with the chlorine electrode^3' was used as the experimental method. While interpreting the experimental data with some approximation it was possible to take into consideration Oznci instead of VznCl when Z^nCl -* 0, and analogically ifC(iC2 instead of )(CdCi when XpdC2 —TT O respectively» Some further approximations are possible assuming that the problem of side reactions deals expecially with dilu te zinc solutions in cadmium, namely a «-, = 1 ; YQ^ = ^I and also If Zn — J^ and In £ Zn -* In j(° when Xz —^ O© These values were given in paper(27. Taking into account these simplifications the formula [3] takes the following form: X 1 X [Cd] Y " (CdCl9)\ ÖCdCl9 . n , --. 1 X X ÖZn " [Cd]A (CdCl2) J äznCl2 In formula [4l the equilibrium constant K correspond to side reaction [1~] while assuming the ideal ZnClg- CdCl2 and Zn-Cd solutions, while the constant X* correspond to real solutions respectively. . On the basis of the values K , ö znci » ^CdCl > dZn the K' data at temperatures of 773, 823 and 873°E were calculated and then the curve of equilibrium compositions (figure 1) of the side reaction [1J was obtained* References (1) C.Wagner and A.Werner, J.Electrochem.Soc. 110, 1963, 326 (2} Z.Moser, Bull.Acad.Pol.Sci.et lett. 17, 1969, 27 (3) C.Krohn and Z.Moser, Det KGL Eforske Videnskabers Sels- kabs Serifter 9, 19b7, 1 (4) R.Lorenz and G.Schulz, Z.anorg. und Allg.Chem. 170, 1928, 324' (5) K.Jellinek and H.Siewers, Z.Electrochem. 40, 1934? 871 Zbigniew Moser 161 CdCLj 0.9 Zn + CdCl2— Cd+ZnCl2 [1] III! 0.8 1 _ correspond to ideal solutions at temperature 873#K when the equilibrium constant of reaction [l] 0.7 [CdJ 0.6 2J3-correspond to real solutions (formulaUl)than %..,.,<= r—77> calculated respectively at temperatures 773°K (fine 2) '•c*-' CdCl2 ond 873»K (line 3). 0.5 O - data of Lorenr and Schulz » ' at 873°K . (5) O — data of Jellinsk and Siewers' at 873°K. 0A ZnCU Zn 0.1 Pig.1. Equilibrium compositions of the side reaction 1 compared with ideal solutions and literature &a*ca *9 *"^'« 162 V/alter Sarholz QUANTUM MECHANICAL ESTIMATION OF HEATS OF ADSORPTION OF ATOMIC OXYGEN ON SOME METAL SURFACES. Walter Sarholz, Detlef Baresel and Gunter Schulz-Ekloff Robert Bosch GmbH, Technisches Zentrum Forschung, Gerlingen bei Stuttgart, BRD Recently it was shown that the velocities of many electrocata- lytic reactions especially in fuel cells seem to be essential ly influenced by the heats of adsorption of the reacting and the inhibiting species. An optimal catalysis should arise, if the heats of adsorption of the reactants were of the same mag nitude and those of the inhibitors were low^'. Therefore we want to calculate the heats of adsorption of some fuels, of water and of the reaction products on several metals and refractory compounds quantum-mechanically. Usual reactants e.g. simple compounds (H21 CH5OH, CH2O) and water can adsorb via the oxygen, the carbon or the hydrogen atom. Therefore in a first step we calculated the heats of adsorption of oxygen on the metals of the above mentioned refractory com pounds. For this purpose only the bond between one surface atom and one oxygen atom was considered. A modified formula of Mulliken^ gave the covalent bond energies. The different atomic orbitals entering the bond were weighted by L.C.A.0.-coeffi cients, gained by a simple Huckel-calculation. This yielded also the charge distribution which permitted the estimation of the ionic part of bond energy. The results were compared with the experimental values referred by G.C.Bond5) and di scussed. References 1) G.Schulz-Ekloff, D.Baresel and J.Heidemeyer, Coll. Czechosl. Chem. Commun. 36, 928 (1971) 2) R.S.Mulliken, J. Chem. Phys. ^+6, *+97 (19^9) 3) G.C.Bond, Catalysis by Metals, p. k^Z Academie Press, London and New York, 1962 J.W.Schultze 163 THERMODYNAMICS AMD KINETICS OP *"ETAL ION ADSORPTION J.W.Schultze Institut fiir Physikalische Cheri e der ^reien tTni vers: tnt , Berlin-Dahlem, Thielallee 6^ The properties of the adsorption layers of metal ions Kez + on gold and olatinum electrodes in solutions of ^e7,+ (Cu?+ 1 , Cd2+, Pb^+, and Ri3+) with an excess of background electrolyte at potentials S- higher than the corresponding metal ion potential £]\/!e/p *. p = -i 'm (1) PM~Z+ ^6. -x O I lVTe /'?.<< z + 164 J.W.Schultze According to equation (1) the electrosorption valency V describes also the charge flow caused by adsorption Cb qm/ e) rT/jez + )£p . The surface concentration Fw^z* can therefore' ne calculated from coulometric data if ^ is evaluated from the potential dependency. The J'-values are ranging from 1.? to 2 in the case of Cu^+ /Pt, whereas they are 2 resp. 3 in the case of Pb^+/Au resp. Bi3+/Au. The maximum coverage 9max of adsorbed ions (related to 1 adsorbed ion per surface atom of Au or Pt) at the electrode + potential <£Me/Mez is greatly dependent on the radius of the metal atom. 9max is approximately O.S for the big lead and bismuth atoms, whereas it comes up to 1 for cadmium on gold, and it clearly exceeds 1 in the case of the little copper atoms on platinum or gold. The shape of the adsorption isotherm mostly corresponds to a logarithmic or Temkin isotherm. For the systems Cu2+/Pt, Pb2+/Au, and Cd^+/Au the adsorption isotherm is given by e + Me- = (RT/B) in aMez+ - %? & (2) where B is a constant dependent on the interaction of the adsorbed ions or on the surface heterogeneity. At potentials <£ -^ &wie/r/rez+ a metal deposition only takes place after an overvoltage of some 10 mV has been applied and after metal nuclei have formed. In spite of its high packing density, the adsorption layer does not have the properties of the bulk metal Me. The reason may be that the distances between the atoms in the epitactic adsorp tion layers seem to be mainly determined by the electrode metal and not by the adsorbate. The kinetics of the adsorption/desorption processes have been studied in the Cu^/Pt and Pb^ + /Au systems. In the case of copper on platinum the adsorption and desorption are strongly inhibited. The overvoltage 7£ of cathodic adsorption or anodic desorption of copper ions was measured in dependence on the current density i and the coverage 0. These values are plotted in Fig.3 for solutions of 0.1 M CuSOj, in 0.5 1* H2SO1,. At constant coverage © Tafel lines are obtained for the adsorption as well as the desorption. The results can be summarized in the equations v k Fe)/RT iad -- - r ad = " T ad °CU2 + -^[-(B^e • Ge2 J (?) and ^es = rvdes = ^kdes-exp[( ^c139 + i 1 1 r i 1 1 1——r Ads or pi ion Oeiorption o - »^ -A -1 O- ,* /\:r 1 , * A Fip;.3 The chemical ( c< , ft ) and electrochemical (o(e, fte) trans' fer coefficients0 o< and ft are combined bv the eauat'ions c< + ft = 1 and (cx + B )-2 = j( (5) c c e e The electrosorption valency /f obtained from the sum of transfer coefficients are in agreement vrith ^-values determined from Fip;.2 !. The exchange current densities of the adsorption process are similar to those of the corresponding rre£al electrode z+ + ?+ Me/Me . The i0-values determined for 0u2 /Pt or Cu /Cu are about 1 m.A/cm , whereas they are more than 1 A/cm? for Pb2+/Au or Pb2+/Pb systems. Therefore, the rate-deter mining step of the adsorption reaction seems to be very# similar to the rate determining step of the metal deposi tion. References: 1) J.W.Schultze, Ber.Bunsenp;es.physik.Chem. 7£, 7^ (ln70) 2) K.J.Vetter, J.W.Schultze, in preparation Kurt Schwabe 166 MF OF THE CELLS (Pb)H2/H Hal, Hg2Hal2/Hg IN ORGANIC SOLVENTS AND THEIR MIXTURES WITH WATER Kurt Schwabe, ForschungsInstitut Meinsberg, DDR Comparing the primary medium, effects of the reaction # f 1/2 H2 + AgCl(s)^Ag(sj + H + Cl with 1/2 H2 + 1/2 Hg2Gl2(s)^=^Hg(s) + H* + CI1 in different solvents and their mixtures with water gives small dif ferences, most values with Hg2^l2 are smaller than those with AgCl (Table 1). Discussion of these differences can not exclude an incomplete equilibrium of one of both re actions in the organic solvents. Measurements of the 11F of chains with transference: (Pt)H2/H Hal(H20)//H Hal(org„ solv. )/E2(Pt) (1) Ag/AgHal(s), HHal(H20)//H Hal(org.solv.)AgHal(s)/Ag (2) Hg/Hg2Hal2(s),HHal(H20)//HHal(org.solv.)Hg2Hal2(s)/Hg (3) and combinations of such cells prove, that the value of the so called phase boundary potential on the frontier between water and organic solvent is small with respect to the fflF of the cell. The large values of the FMF of the cells i, 2 and 3 at standard conditions found by ex trapolation therefore are caused by large differences between the standard galvani tension of the hydrogenelec- trode, the Hg2Hal2/Hg- and AgHal/.ag electrode in water and their standard galvani tension in organic solvents. In table 2 several values of the standard galvani tension in waterfree organic solvents refered to the standard hy drogen electrode in water are shown. That means, that the long known decrease of pH in organic solvents (esp. alco hols) compared with the pH-value in water at the same acid concentration corresponds to a real increase of H activity. This is confirmed by spectrophotometric measure ments of the Hammett function. The galvani tension of the above named electrodes with respect to their value in water depends from the water content of the organic solvent in high degree. In the most investigated solvent mixtures with water only at smaller water contents than 10 weight procent the diffe rence of the galvani tension with respect to that of the value in water becomes large (Figure 1). The effect of the increasing content of organic solvent on the galvani tension depends on the tendency of the solvent to associ ate with the ions of the electrode process. If a solid salt participates in the electrode reaction the change of the galvani tension is strongly influenced by the change of the solubility in the organic solvent mixture (Table 3)« In the case' of soluted ions (esp. H ) the strong change of the galvani tension seems to begin, when the content hurt 6chwa.be of the 'organic solvent is large enough to solvate the ions instead of their hydration. This is confirmed "by measurements of the dielectric constant in such electro lyte solutions with different water content. The results are discussed v/ith respect to their theo retical information and their practical applicability. Table 1. hrimary medium effect Ig fj^/^) (23"CJ> Organic solvent weight % Hg2Cl2 AgCl CH7OH 60 0.340 0.3433 -> 80 0.610 0.6190 90 0.920 0.9208 iso propanol 20 0.1361 0.1373 acetone 20 0.1233 0.1226 4-0 0.3087 O.3036 glycol 19.23 0.0947 0.100 30 0.2687 0.276 dioxane 10 0.1314- 0.1636 43 0.4862 0.4976 70 1.2932 v 1.3392 ,r 82 2.277 ( (Pfc)HP/HCl CH3OH + 134 ^ C2H5OH + 178 n-CzjHQOH + 181 ms _ 19 MAK + 222 Ag/AgCl NaCl CH5OH + 133.1 Ag/AgCl HC1 C2H5OH + 111.9 n-CMlQOH + 69.4 MAK — 18.6 Ag/AgBr NaBr CH^OH + 8.2 Kurt Schwabe 168 i ji^xbi. e 3* Influence of solubilicy of A-Cl on -alvo.nl tension o [mV1 weight % -(2.3 RT/F) AlsiCL[m"V] -At AgCl CH^OH CgH^OH CH^OH C2H50H 0 0 0 0 0 8.0 16,6 21.9 3 6*3 33*4 34.2 42.1 21 31 52.0 32.1 54.5 34 30 83-3 99.3 (24.8) 39.5 92.3 222 71 111 100 170 t 2S0 tso too so 20 40 60 SO a £ ot th» €•// Ag/AgCi,HCi/HtfPt> m doptndaoce of wtight-% miihanol Pierre Turq 169 COUPLED SELF AND CHEMICAL DIFFUSION IN ELECTROLYTE SOLUTIONS. DIFFUSION POTENTIAL AND CORRESPONDING RECIPROCAL RELATIONS. Pierre TURQ and Marius C HEM LA Laboratoire d'Electrochimie - Université de Paris VI - 9, Quai St. Bernard - Paris V - FRANCE. 1/ Introduction : Classical studies of diffusion in electrolyte solutions were made in two ways : 1) Chemical diffusion concerning the transport phenomena in an electrolyte concentration gradient. In this case both cations and anions diffuse at the same speed and the diffusion coefficient can be related to the sepai ate ionic mobilities by the Nernst-Hartley equa tion : 2 u u kT D = —— for a symmetrical electrolyte, u + u Ze an electrical field appears simultaneously to the chemical diffusion. 2) Self-diffusion. In this case, the motion of the ions is observed with radioactive tracers in the absence of chemical concentration gradient. The self diffusion coefficient is related to the ionic mobi- lityiby the Nernst-Einstein expression : _ kT u " ZeT We have generalised these processes to the case of a tracer ion dif fusing in a medium where a chemical concentration gradient and a tracer concentration gradient simultaneously exist. Considering the notion of self diffusion as a random-walk, it is noted that, in dilute solution, the isotopic diffusion coefficient is not greatly affected by the presence of other ions with or without gradient. Also it is well known that ionic interactions, such as the electr •oho- retic effect or the relaxation effect do not have much influence. On the value of the diffusion coefficient. Then D is negligibly affected by ionic strength. In each volume element of the solution there, is simultaneously a tracer flow (for self-diffusion) and an electrolyte flow (for chemical diffusion). However, the tracer ion can move in a counter-current of electrolytes and in several of these cases we have shown that the ion is accelerated by the counter-current. 11/ Expression of tracer flow. Electrolyte flow is not affected by tracer-flow, the tracer species concentration being negligible in comparison with the electrolyte concentration. We have for the tracer flow : J, = L, v" I L10 V ,-, 1 1 U IS u S where y is the chemical potential gradient of the tracer ion 1 and 170 Pierre lurq ^ 0 the chemical potential gradient of the supporting electrolyte. uS The phenomenological coefficient L describes the coupling between self-diffusion and chemical diffusion or, in other words, the chemi cal diffusion contribution to the tracer flow. L follows the Onsager's reciprocity relations and has been calculated directly from the properties of the ions at infinite dilution. Ill/ Experimental method and results. The open-end capillary method (Anderson and Saddington) has been used in this study. The capillary tube filled with the tracer species was immersed in another vessel containning the solution without isotope and at a different electrolyte concentration frorruthat of the solution in the capillary tube. Results for the tracers Na and CI in aqueous solutions of NaCl exhibited good agreement with the theory, showing particularly the influence of the diffusion potential electrical field. Electrolyte flow -• Electrical diffusion field ~* 1 Flow of Na - 22 increases J Na 22 The results are particularly 2 Flow of *T significan for CI where the 22 Na «- J Na same electrolyte flow accele decreases rates the tracer in the opposite 3 Flow of 36d - direction and decelerates the tracer in the electrolyte flow j 36ci decreases direction. 36 4 Flow of ci - j 36cx increases IV/ Conclusion . The first conclusion of this work is that small differences between chemical concentrations in the capillary and in the vessel will not affect greatly the tracer diffusion coefficient An other and more significant conclusion is that the diffusion potential (and the corresponding electrical field) is the principal factor governing the observed differences between coupled diffusion and pure self-diffusion. Pierre Turq 171 This study was an example of Onsagcr's reciprocal relations wich exhibit the mutual influence of the particles-flows and was one of the rare cases where the phenomenological coefficients could be calculated from the transport data of separate ions at infinite dilution. SECTION 2 ELECTROCHEMICAL KINETICS KINETICS OF TKIN OXIDE LAYER FOF^IATIOM 0:-J NOrLE "'E7ALS K.J.Vetter Institute of Physical Chemistry, pree University Perlin, Hernany When an anodic current is applied , most metals form pore- less oxide films in aaueous electrolyte solutions if the potential is positive enough. The thickness of these films varies from that of a monolayer by up to 100 8 or some 1000 8. The films on noble metals, such as Pt, Pd, Ir, Au, also Cr and others, reach thicknesses of one or fev? mono layers, whereas the typical "passive metals", such as Fe, Mij produce films of 20 to ^0 A thickness. Anodic metal films on Al, Ti. Ta, Mb, and similar metals reach thick nesses of 1000 Ä. These thicker oxide films start forming with the formation of monolayer films. Consequently, the discussion of the kinetics of thin oxide film formation is relevant also with regard to th-.. formation of the thicker films. From anodic and cathodic charging curves (galvanostatic potential-time relation) it follows that the formation and the reduction of the oxide film Is strongly impeded in the majority of cases. Generally, no reversible oxide potential can be measured, but the charging curves for a constant current density are shifting with increasing pH value to more negative potentials, while thepShapes remain unchanged. The shift is nearly ^0 mv/pH . For example, this effect could be stated at Pt, Pd, Ir, Au over nearly the full pK-scale. The distinct irreversibility of the charging curves indi cates that large overvoltages are involved. mhe potential or the overvoltage depends on the current density, on the pH value and on the coverage 9. The charge of the whole charging curve usually corresponds to one or two mono layers. In the case of Pt and Au, mafel relations £. = = E + b»loglil could be found anodically (b+) and catho- dically (h_). The Tafel factors b+ and b_ are a function ([ of the coverage according to the relation b = bo(l+a0) ^* . Because of (ös/öpH) ^ ^ 59 mV, an apparent electrochemical reaction order ZQ^- of the hydroxyl ions ZQH" = Cölogli// / "b log[0H~J)g:= 5° [mVj /b [mV] is obtained. Since the b values are ranging from 120 to 30 mV, the reaction orders vary from 0.5 to 2.0. "oreover, not only the rate constant but also these apparent reaction orders are depending conti- nously on the coverage. A simple explanation why such reaction orders are varying between 0.5- and 2.0 at different 9 cannot be given. The model of the film formation must be more complicate because it must also explain the shift of the charging curve in 176 K.J.Vetter the v/ide pH range between acid and alkaline electrolyte solutions, and further, an essential ageing effect of the oxide film. The model of the formation of thin oxide films given by Vetter and Schultze^»^ for Pt and Au consists of an oxide layer with a superposed electrosorpt.ion layer of oxygen or hydroxyl ions in the inner Helmholtz layer of the double layer. The oxide film is a film with metal and oxygen ions alternating at its surface, and it is covered with the electrosorption layer composed only of oxygen or hydroxyl ions. To explain the pH shift of the charging curves or of the potential for constant current density and the coverage respectively, an electrosorption equilibrium of oxygen or hydroxyl ions ^O-^rOH,^ + H+,aq ^0^ + 2H+«aq has to be assumed. This means that the potential difference faci-tpe between the inner (ad) and the outer (e) Helm holtz layer is independent of the current density and has a dependence of 5° mV/pH in both cases, namely an 0^~ or 0H~ electrosorption. The exchange current densities must be higher than the applied anodic and cathodic current densities. In this model the rate determining process is the place exchange reaction between electrosorbed oxygen ions and metal ions of the first surface layer of an oxide film not completely formed. The rate of this process depends on the potential difference Cp^- 1) F.P.Bowden, Proc .Roy .Soc. Ser .A 125,446(1929) and many subsequent papers of other authors, such as Butler, ? Frumkin, Ershler, Hickling, Breiter and coworkers,see 2) K.J.Vetter, D.Berndt, Z .Elektrochem. £2_,378 (1Q5?) 3) J.W.Schultze, K.J.Vetter, Ber.Bunsenges.physik.Chem. 75,^70(1971) 4) K.J.Vetter, J.W.Schultze, J.Electroanal.Chem. 34 ,131 (1972); 34,141(1972) AtJ-Arvia 178 KINETICS OF THE ELECTROCHEMICAL FORMATION OF PARATHIOCYANOGEN FILMS A.J.Arvia and A.J.Calandra Instituto de Investigaciones Fisicoquimicas Teoricas y Aplicadas, Division Electroquimica, Facultad de Ciencias Exactas, Universi- dad Nacional de La Plata,.La Plata, Argentina. The anodic oxidation of SCN ion either dissolved in acetonitrilet or dimethylsulphoxide or as molten KSCN, yields (SCN) on plati num electrodes. The reaction was studied from -3QC up to the bp of the solution in ACN; from 20QC up to the bp in DMSO and from 190 to 250 QC in the melt, by means of voltammetric and relaxation techniques. i) The anodic oxidation in molten KSCN is characterised by anodic E/I curves which present a current maximum. They can be reproduced with the theoretical kinetic parameters derived from the follow ing reaction pathways: I: Pt + SCN" = Pt(SCN) + e (la) Pt(SCN) + Pi*= Pt*(SCN) + Pt (rds) (lb) Pt**(SCN) + Pt(SCN) = Pt*(SCN) „ + Pt (Ic) _x _x+l and _ II: PtfPt**] + SCN = Pt[pt*] (SCN) + e (Ha) Ptfpt»] (SCN) = film formation (rds) (lib) where Pt corresponds to a nucleation centre on the reaction sur face. ii) The anodic oxidation in ACN solutions at T^OQC yields (SCN)» as reaction product which undergoes a slow polimerisation reac tion. Consequently the E/I curve for the reduction of (SCN)- was determined. Under these conditions the kinetics of the pseudo-ha logen couple (SCN) /SCN was established. The voltammetric charac* teristics fit the following reaction scheme: Pt + SCN~ -. Pt(SCN) + e (Ilia) Pt(SCN) + SCN = Pt + (SCN)2+ e(rds) (Illb) At T> 20QC the polimerisation of (SCN)2 on the electrode surface yields (SCN) which passivates the electrode. The rate of the film formation is then mainly controlled by the ohmic .resistance of the circuit. The behaviour of the I/time curves is roughly explained by Mueller model of passive film growth, iii) The oxidation of SCN~ in DMSO solutions follows essentially the pattern alrwady mentioned in ii). The quantity of electricity required to passivate the elec trode surface- is between 5 and 11 mC/cm for i) and between 516 and 2000 mC/cm for ii) and iii). The formation of the film invol ves an ohmic and a non-ohmic contribution; the latter becomes mo re significant as temperature increases. The kinetic contribution comprises the initial bidimensional growth, the steric effect of the solvent impeding bidimensional growth, the tridimensional growth, the diffusion and depletion of the ionic species inside A.J. Arvia 179 the pores which constitute the grov;ing film and, finally, the film sealing provoking the partial squeeze out of the solvent or its occlusion inside the film. The influence of the potential sweep rate on the increase of tridimensional growth can then be explained. 180 Radosiav Atanasoski ACTIVITIES OR CONCENTRATIONS BT EQUATIONS 01 ELECTRODE KINETICS li.I.i. Jaksié, RoT,, Atanasoski» A.R«, Despié and V. Naki6 Electrochemistry Department, Institute of Chemistry,Technology & Metallurgy, Beograd, Karnegijeva 4/IV, Yugoslavia. Amongst other possibilities,as it is well known,one could obtain the symmetry factor^ for the hydrogen evolution reaction from the following characteristic derivative ^ln i ( 1 ^r\ += - P H To distinguish whether activities or concentrations govern rela tions of the electrode kinetics,measurements of the cathode hydrogen evolution were undertaken at the constant hydrogen ion content,but varying a concentration of carrying neutral salt. Ante Baric 181 ELECTRODE REACTION OF Cu(ll) - EDTA ON MERCURY. EFFECTS OFDOUBLE LAYER AND ION PAIR FORMATION. A. Baric, M Branica and J. Kuta Center for Marine Research, Institute "R. Boskovic" Zagreb, Croatia, Yugoslavia and J. Heyrovsky Institute of Polarography Prague, Czechoslovakia The rate constants of two-electron charge transfer of the copper (II) - ethylenediaminetetraacetic acid complex on mercury in buf fered solution at pH 8 are markedly influenced by the composition and the concentration of the supporting electrolytes. The d. c. polarography and single-sweep voltammetry on a hanging mercury drop electrode were used in this study. The reduction of the copper(H)-EDTA complex takes place at the positive side cf the electrocapillary curve on mercury and its ra te is highly affected by the anions of the supporting electrolyte. In the presence of fluoride the rate constant is rather high and the charge transfer reaction is almost reversible. In chloride medium this reaction proceeds already with a small overvoltage ("quasi-reversible" reaction), and the apparent rate constant (ka) was determined by the usual procedure for quasi-reversible polarographic wave , as well as the corresponding reversible half-wave potential (Ef/2= -0.420 V vs. S. C. E. in 0.1 NaCl). In iodide solutions the rate constant of the charge transfer reaction decreases considerably, which is manifested by the large shift of the half-wave potential towards more negative values (E\r/2 = -0-565 V vs. S. C. E. in 0. 1 N Nal). The electrode reaction is then unidirectional (totally irreversible polarographic wave). The experimental values of charge transfer coefficientoc were near 0. 6. The apparent rate constants were evaluated by using the method of Koutecky . The obtained values of k0 " s are summarized in Table 1. Some a results were confirmed by using a single sweep voltammetry. The large changes in rate constants with the concentration and the anion of the supporting electrolyte can be explained by the double-layer effect and the ion-pair formation. By using capillary electrometer it was shown that in iodide solutions th,ere was no specific adsorption for the given concen tration of EDTA and of the components of the borate buffer. In chloride solution a small adsorption effect of EDTA was detected due to a smaller adsorption activity of chloride anions in com parison with that of the iodide ions. Therefore the corresponding values of the potential in the electrode double-layer for the 182 Ante Baric Table 1 Correction of the Appart--.. -..ate Constant for the Double Layer Effect = - ? - 1 6 3 3 3 3 Concn. E k°xl0 k°,xl0 k°,x.!0: / x 10 kk ° /xlOj k° /X10 1/2 a t$2>_ l *9\ ?2 N mV cm/sec cm/sec cm/seL/S c cm /sec cm/sec S. C. E. i 0, 05 -420: 11 2600 1600 33 23 0. -415; 0. 69 640 1500 10 13 0, 25 -40 5; 0.41 140 1100 1, 2 12 i 0. 50 -4oo; 0. 31 11 690 0. 39 6.4 1, 00, -395' 0. 32 38 850 0 17 9. 5 0. 10! -415 1600 26 23 11 11 1.00 -395 1500 520 23 96 12 1 Nal 2 NaCl 3 extrapolated reversible value 5 6 solutions of iodide and chloride can be used in the correspon ding correction of apparent rate constants. The true rate constant (k°) was calculated according to Frumkin's equation k° - k° exD (z -<*n) F^ t ~ \ GXP RT (1) For the charge of the electroactive species (z) the values -2 and -1 were successively inserted in equation (1) and for the poten tial of the closest approach^ the values of outer^? and inner G^, part of the double-layer were chosen. From Table f. follows that the constancy of k^. is obtained for z = - 1 and for poten tial s $ i . The lowering of the charge of the electroactive species may be attributed to the ion-pair formation. Ion-pair formation constant k = 89 in iodide solution was calculated by using the method proposed by Gierst . From this study it can be concluded that the large differences in the apparent rate constants k° for the reduction of Cu(ll)-EDTA complex can be explained by the double layer effect and the ion-pair formation between Cu-EDTA complex and the cation of the supporting electrolyte. Ante Baric 183 Literature 1. II. Matsuda and Y. Ayabe, Z. Elektrochcm. _63_ (1959) 1164. 2. J. Koryta, Electrochim. Acta 6> (1962) 67. 3. I. Ruzic, A. Baric and M. Branica, J. Electroanal. Chem. 2_9 (1971) 411. 4. J. Koutecky, Coll. Czech. Chem. Commun. 1_8_ (1953) 597. 5. D. C Grahame, Z. Elektrochem. 62 (1958) 264. Technical Report of the Office of Naval Research No 5, Amherst College 1957. 6. D. C. Grahame and R. Parsons, J. Am. Chem. Soc. J33_ (1961) 1291. 7. L. Gierst, E. Nicolas and L. Tyfcgat-Vandenbergen, Croat. Chem. Acta 42 (1970) 117. 184 G.J. Bignold IDENTIFICATION OF TRANSIENT PHENOMENA DURING THE ANODIC POLARIZATION OF IRON IN DILUTE SULPHURIC ACID G.J. Bignold and M. Fleischmann Central Electricity Research Laboratories, Leatherhead, Surrey, England and University of Southampton, Southampton S09 5NH, England. The anodic dissolution of iron in dilute sulphuric acid has been the subject of considerable work and discussion^"". When an iron electrode which is dissolving anodically in a dilute acid, is subjected to a sudden change of applied potential the current-time transient has the following characteristics: firstly a high charging current is observed which is due to the double layer capacity and electrode reaction pseudo-capacity of the electrode. This current decays within the first millisecond to a value ij which we shall refer to as the 'initial current'. There follows a period during which the current density rises to approach asymptotically a 'steady-state' value, ig. This behaviour has been observed by numerous authors1»2»3, The steady-state current has been shown to obey Tafel kinetics but there has been poor agreement between the various sets of reported data (see Florianovich et al. for a review^) which has led to some controversy over mechanistic interpretations. The work of Eichkorn and Lorenz^ has shown that variations in heat treatment account for some of the irreproducibility. The initial current and steady state currents follow different kinetic laws; the variation of ij with potential is apparently much more reproducible than ig. Several authors1»2,3 using different samples of iron of varying purity have found that the initial current at 25 C obeys the relationships: 1 (3E/31og iI)pH = 60 ± 5 mV decade" ... (1) and Olog ix/3pH)E = 1 ... (2) Heat treatment apparently has no effect upon these relationships. In this paper the transient behaviour is examined and possible causes for it are assessed. Measurements have been made of the variation of the initial current with potential. pH (in the range 1 to 3) and temperature (in the range 25 to 75 C). The results can be summarized, to within the limits of experimental reproducibility, by the equation: log iT = - ^ + 8.747 + PH - 2<3^RT ... (3) 2- (for solutions 0.05 Mw.r.t. SO, ions) The time taken for the rise in current to be 63% completed was also measured. The results show that the time constant is insensitive to pH and potential variations, but decreases significantly with increasing temperature. Two possible causes of the rising current/time transient have been advanced previously: Eichkorn et al. postulated that it was due to the relaxation of electrochemical pre-equilibria, whilst Bockris and G.J. Bignold 185 Kita-5 cited pH changes following the relaxation of these processes as being responsible. The former suggestion can be eliminated by consideration of the timescale over which relaxation would take place. It transpires that this should be very short (< 10~4 seconds) and dependent upon both pH and potential. The suggestion of Bockris and Kita-' leads, as they showed for the case of galvanostatic transients, to relaxations over approximately the correct time-scale. However the quantity of hydrogen released during the relaxation (assuming it to conform to equation (i) below) can be shown to be insufficient to account for the magnitude of the rise in current. Fe + H„0 -> (FeOH) + H+ - /" ... (i) 2 «- ADS Two further possible causes of the transient phenomena are put forward and examined in this paper. These are that the pitting geometry of the surface changes with the applied potential, and that the accumulation of ferrous ions around the electrode leads to a modifi cation in the pH at the surface, and thus causes variation in the current with time. The first of these suggestions has been developed into a model which shows that the time constant for relaxation of the pit shape must be expected to be potential dependent, and to vary within the range 30 to 6000 seconds for annealed iron (for cold worked iron which has a much greater pitting density this time constant range is reduced to 2 to 400 seconds). Thus the relaxations within the first 0.5 second which have been observed cannot be due to this phenomenon. Slow changes in the 'pseudo steady-state' value of the current can, in fact, be observed in the time change predicted for changes in pit shape. The second new sug^stion, that pH changes due to accumulation of dissolution produrts are responsible for the transients, poses severe mathematical probie »s if a rigorous treatment of the dissolution of iron in sulphuric acid/sodium sulphate mixtures is to be attempted. A simplified treatment for iron in an acid H+ A" has been adopted, after making the following assumptions: (a) Ions nove entirely under the influence of diffusional fields (electric fields being negligibly small). (b) The total ionic concentration of solution species is constant throughout the diffusion layer, and is the same as that in the bulk electrolyte (i.e. double layer effects are ignored). (c) Electroneutrality exists throughout the solution. Using these assumptions a model has been developed which predicts that the early part of the transient should follow the expression: 2 t ! (3) This expression has been tested for transients in 0.05 M H2SO4 over the temperature range 25-65 C. The form of the response observed :j.s correct over the first 50 ms and the slopes of plots of i versus t2 are within 30% of the values predicted by equation (3). This agree ment may be regarded as very good in view of the considerable G.J. Bignold 186 assumptions involved. The theory also predicts that the dissolution rate should be sensitive to perturbation of the diffusion conditions at the electrode. Using the same assumptions as above the response of a rotating disc of iron can be predicted to be given by: 2 ]- i ~ i + Ki /w2 ... (4) where K is a calculable constant involving the kinematic viscosity of the solution, the diffusion coefficient of the ferrous species, and the constants associated with the rotating disc system. Experiments in 0.05 M H2SO/ at 25 C gave results in good agreement with equation (4) and thus verify the validity of the theory. An important implication of the theory is that mechanistic interpre tations which have been placed on the 'steady state' data are in error because they do not take account of pH variation at the electrode. Measurements of the initial current, i-j-, are not subject to this criticism, and it is therefore upon these measurements that mechanistic interpretations should be based. The following sequence is suggested to account for the experimental observations. + Fe, . . + Ho0 + (FeOH). . . + H + e~ ... (ii) kink 2 <- kink (Fe0H)kink * (Fe0H)ADS '•• (iii) + + + ++ (FeOH)ATAO + H + FeOH + H + e" •> Fe + H_0 + e~ ... (iv) A distinction is drawn between iron atoms in kink sites (i.e. atoms which are available for reaction) and adsorbed ions on the remainder of the surface. The rate determining step, (iii), proposed which is the movement of an intermediate from its kink site out onto the surface of the iron, represents the largest change in the coordination of the iron atom during the sequence. The catalysts of this step by the presence of an anion accords with the observations of Kolotyrkin et al. on the effect of anions upon the reaction. The mechanism proposed predicts the Tafel slope and pH dependence of the dissolution reaction, as measured at the beginning of transients, and is consistent with modern views of dissolution and growth of crystals. The work reported was carried out at the Central Electricity Research Laboratories, and is published by permission of the C.E.G.B. References 1. K.E. Heusler, Zeits. fur Electrochemie. 62^ 582 (1958) 2. G.M. Florianovich, L.A. Sokolova and Ya. M. Kolotyrkin, Soviet Electrochem, _3> 917 (1967) 3. F. Hilbert, Y. Miyoshi, G. Eichkorn and W.J. Lorenz, J. Electro Chem. Soc. _U8, 1919 (1971) 4. G. Eichkorn. W.J. Lorenz, L. Albert and H. Fischer, Electrochimica Acta, JL3, 183 (1968) 5. J.O'M. Bockris and H. Kita, J. Electrochem. Soc. 108, 676 (1961) 6. G.M. Florianovich, L.A. Sokolova and Ya. M. Kolotyrkin, Electrochimica Acta. 12, 879 (1967) Aurelian Cälu§aru 187 THE ELECTRODE POTENTIAL OF METAL DENDRITE GROWTH AT THE LIMITING CURRENT Aurelian Cälusaru Institute of Atomic Physics, Bucharest, Romania. In the classical electrochemistry no distinction is made between the elementary act of the electrochemical reac tions before and after than the limiting current is at tained. However the study of the phenomena appearing at the limiting current showed that drastic changes may oc cur, depending on the electrode potential, although the current density do not vary and remains at the limiting value. The example in this field is represented by the metal dendrite formation at a characteristic potential on both solid (l) (Table 1) and mercury (2) (Table 2) elec trodes. Table 1. Potentials of dendrite formation on solid electrodes Metal Potential*) ,V Medium Au -0.95 HAuCl4(0.09-0.2 N) + HC1 0.1 N Cu -0.45 CuSO (0.02-0.2 N)+H2S04 0.1 N Cd -0.05-0.1 CdS04(0.05-0.2 N)+H2S04 0.1 N "V" ) The refference electrode is the considered metal in the above solution. Table 2. Potentials of dendrite formation on D.M.E. Metal Potential (SCE), V Medium Ag -0.3 AgNO +KN0_,2 M+gelatine 0.05 # Bi -0.31-0.71 Bi(N05)5+ HC1,4 "M+gelatine Cu -0.8 CuCl2+ KN0_,2 M+gelatine 0.05 # Pb -0.8 Pb(N03)2+KNQ5,2M+gelatine Cd -0.9 CdCl2+KN05,2 M+gelatine 0.05 Tl -1.3 T1N05+ KN03, 2 M For the explanation of these potential depending proces ses a quantum mechanical interpretation was given consi dering the tunneling phenomena at the metal-electrolyte (1) and metal-thin film-electrolyte interfaces (3). On D.M.E. the hydroxide precipitation at the limiting cur rent can be observed microscopically, giving the experi mental evidence on the thin film formation at the elec trode surface. The addition of the surfactants, especially the gelatine, has an important effect on the structure and thickness of the interfacial film. Owing to the high 188 Aurelian Cälusaru potential values in the limiting current region, the elec trons can easly tunnel across the barrier introduced by the thin films of the order of 50 Å, between the cathode and the metal ions at the film surface. In this manner both the nucleation and the insulating of the dendrites from the mercury surface may be explained. The dendrite growth seems to occur at a metal-electrolyte interface by electron tunneling from the metal cathode to the free electron levels on the ions in solution. Considering the condition of the constant transition probability of the electrons as a consequence of the limiting current, it is possible to find a relationship between the applied pote- tial and the neutralisation distance (3): exp (u ) const = A r D( s,u ) In 1+ex., p (,u* )v du (1) u m where ui = (Ei-EF~eV-r)/kT, A is a constant, D(s,ux) the transparence coefficient, ^i the electron energy, ^F the Fermi level energy, e_ the electron charge, V the applied potential, r the difference between the lowest free elec tron level in solution and Fermi level,J k_ Boltzmann con- stant, and T_ the absolute temperature. The integration limiT .j : tj- ~^m11 i-i s_ th4-T e energy of the lowest free level in Dropping mercury electrode ,r- 0.02V solution, when this level ti i n is considered as constant S-5Ä II S-/0Ä after a certain distance, § 15 the corresponding value >" u is m = (EF_r)/kT# if the energy of the lowest free electron level on ion varies as a function of the distance from the electrode it is necessary to consider the equation Fig. 1. Electrode potential (V) of the envelope curve of versus neutralisation distance the ground states of the (sv) for D.M.E. considering free levels on ions, the rectangular barrier. noted by f(s). With this function the limit um be 1- TQ = 4.53 eVfmercury work u E function); 2- TQ = 3-53 eV come m=( F -r-f(s))/kT. (l eV lower). Dotted lines cor Considering a definite respond to.the experimental shape of the barrier values (2). Values calculated D(s,ux) (for instance rectangular or parabolic) with I.B.M. computer. it is possible to cal culate the equation (1) in order to find the variation of the neutralisation distance as a function of applied potential, V. The relationship between the potential V Aurelian Caluparu 1S9 and the neutralisation distance sv, calculated with computer is given in Fig. 1 lor a rectangular bar- I.B.M. 1 rier and in Fig, T'm~ 2 T'RTP- nol i '" O If .v . t .-. "'"~ ."' " f the constant free electron level in solution, a''ter a oe- has been here considered. The relation- termined. distance D ship between V and sv. if. 2i strongly dependent on Z'.r- Drvfipiiia mercury electrode, r - 000 V barrier shape. The calcula 20 ted data for the case of variable free electron level energy in solution are not yet available, or.t it j p, no lOz doubt that in t • i p irodol an increase of sv • tn \/ will as =zrS: be obtained always. Thin "elastic barrier" during no the electron tunneling at the interface metal-electro Fig. 2. The same as Fig. 1, lyte means that the elemen considering the parabolic- tary act of the electroche- barrier. Values calculated raical reactions is different with I.B.M. computer. at the limiting current, compared with the same act below the limiting one. On this basis the mechanism of the metal den .rite formation at a characteritic potential in the limiting current region can be considered as a crystallisation process subsequent to the ion discharge. REFERENCES 1. A. Calu§aru, Thesis, Politechnic Institute, Bucharest 1957; ibid. Electrochim. Acta, 1_2, 1507 (1967). 2. A. Caluc;ar-u and J. Kuta, Nature, £11, 1080 (1966); ibid. J. electroanal. Chem., _20 , 383 (1969). 3. A. Cälusaru and A. Moldovan, Symposium on Modern Aspects of Double Layer Theory and. Elementary Act of the Electrochemical Reactions, Institute of Electro chemistry of U.S.S.R. Science Academy, wloscow, June 1971; ibid., Electrochim. Acta, in the press. 190 Georges CAUQUIS THE ELECTROCHEMICAL REDUCTION OF SULFUR IN ORGANIC MEDIA Raymond BONNATERRE and Georges CAUQUIS Labjratoire d'Electrochimie Organique et Analytique du Centre d'Etudes Nucléaires de Grenoble, Cedex 85, 38 - Grenoble-Gare, France. The present x^ork is devoted to the study of the different reduction states of sulfur in dimethylsulfoxide and in dimethylformamide by a combination of electrochemical and spectrometric methods. Solutions of polysulfides in organic media are generally strongly co lored but the exact assignment of the three absorption bands in the visible spectrum (band A : 618 nm, band B : 500 nm, band C : 430.nm) has given rise to much controversy. On the one hand, GIGGENBACH ' attributes the A band to the S2*~ anion radical and the C band to the S/ dimer, on the other hand, SAWYER (^) attributes the A band to the Sg' anion radical and the B band to the Sj£ dimer. They both consider the formation of a dimeric species of anion radical but the dimer is not characterised by the same absorption band in the two works. In this work, it is shown that the three absorption bands must be attributed to species of different reduction states and, more preci 2 2 - sely, to S6 ~ (A band), Sg ~ (B band) and S^ (C band). A solution of sulfur in DMSO or DMF gives rise; on an inert electrode (platinum or vitreous carbon) to two cathodic waves of nearly equal intensities. These two successive reductions steps are bielectronic steps and not one electron steps as previously assessed. Electrolysis at controlled potential on the limiting current of the first wave, together with spectrophotometry measurements, show that B band goes through a maximum of intensity as the number z of elec trons involved in the electrolysis is equal to two for one molecule Sg. The optical density of A band goes through a maximum when z = 8/3. This corresponds to a global process 2 3 Sg + 8 e ^=^r 4 S6 ~ which is interpreted as a bielectronic electrochemical reduction Sg + 2 e -5=*r Sg2" followed, in solution, by the dismutation The C band appears (for z > 8/3) when the reduction potential is chosen on the 'limiting current of the second wave and its intensity gradually rises until z = 4. These results can ;also show that the B band intensity decreases in the interval z = 2 to z = 8/3 whereas A band increases, as z increases, producing an isobestic point. As z values rise from 8/3 to 4, A band decreases while C band increases giving a second isobestic point. The relation between" the quantity of electricity consumed and optical density clearly shows that the three absorption bands cannot be attributed to species of the same reduction state and excludes any monomer-dimer equilibrium between the species which give these absorption bands. Georges CATJQUIS 191 ESR measurements of electrolysed solutions in DMF as well as in DMSO indicate the existence of two paramagnetic species. A first signal is symmetrical and its intensity goes through a maximum for z = 8/3. It is originated from S3*~ anion radical which is produced through S^~ partial homolytic dissociation. A second spectrum is observed, par tially superimposed on the first one, when z = 2, and arises from a S«*~ - ^4*"" mixture, the latter radical anion being formed through Sg2~ partial homolytic dissociation. A comparison of the A band intensity and the ESR signal intensity on both sides of their maximum observed for z = 8/3 is in accordance with these attributions. Bibliography : 1 - W. GIGGENBÅCH, J. Inorg. Nuclear Chem., 1968, 3£, 3189. 2 - M.V. MERRITT and D.T. SAWYER, Inorg. Chem., 1970, _9, 211. See also H. LUX and coll., Chem. Ber., 1961, 94, 1161 and 1968, 101, 2485. 192 Bozena Cosovic ELECTROCHEMICAL REDOX PROCESSES OF URANIUM IN AQUEOUS SOLUTIONS OF ACET YLACETONE B. Cosovic, Lj. Jeftic, M. Branica Center for Marine Research, "Rudjer Boskovic" Institute, Zagreb, Croatia, Yugoslavia and Z. Galus Institute of Fundamental Problems of Chemistry, University of Warsaw, Warsaw, Poland Electrochemical redox processes of uranium (VI) and uranium (V) in aqueous solutions of acetylacetone have been studied by means of polarography and cyclic voltammetry. Depending on the concentration of acetylacetonate ion (Acac ) ura- nyl ion exists in aqueous solutions as free ion (UO? ), mono- acetylacetonato (UO?Acac ), and bisacetylacetonato (UO^Acac?) complex ' . xA>ll the three species are at a mobile equilibrium and can be reduced on mercury in several steps, the first reduction step being the reduction of uranium (VI) to uranium (V)3'4. The presence of acetylacetone accelerates the rate of dispropor tionation of uranium (V), even under the conditions where uranium (VI) is not in the form of the acetylacetonato complex. To determine the rate of disproportionation reaction uranium (V) solution was prepared by electrolysis of uranium (VI) at constant potential (-0.5 V) and then a given amount of Acac was added to the known volume of the solution. Tne. decay of the uranium (V) oxidation current was followed by classical polarography and experiments with constant concentration (0, 2 M) of Acac showed that the reaction is of second order in respect to uranium (V) ion. With the increase of Acac concentration the rate of the disproportionation increases considerably this rate being influenced by Acac in a rather complicated way» Calculated rate constants (mole" min"1) for 0.04 M Acac and 0.30 M Acac are 36 and 550, respectively. Kinetic parameters (kg and oC n) for one electron reduction of uranium (VI) were determined from cyclic voltammetric curves and found to depend: - on the free ligand concentration resulting from distribution of uranyl species and - on. the concentration of acetylacetone resulting from its adsorption on mercury. All determined values for kg are higher than 10 cm sec The plot of E, /-, vs. the free ligand concentration gives an "S" - shaped curve. For the free ligand concentration lower than Bozena Cosovic ly-. log [Acac ] = -7. 5 the half-wave potential has the constant value ^\/2)l ~ "0.187 V (SCE) corresponding to the reduction of uo to UO2. For the free ligand concentration higher than log [Acac] = -4.5 the constant value (E, /2)TT = -0.455 V (SCE) indi cates that the complex UO^Acac^ undergoes one electron reduction to UC>2Acac2 without a change in the number 01 ligands. For the free ligand concentration between the mentioned values the first reduction step is the reduction of species UOiAcac + and U02Acac2 to UCU and UCXAcac , respectively, i. e. reduction is associated with the loss of one ligand. In the absence of the. complexing agent, the reduction of uranium. (V) occurs at about -0. 9 V (SCE). With gradual addition of acety- lacetone a new wave appears at -0. 5 V, while at the same time the reduction wave at -0. 9 V decreases. This new -wave is ir reversible and kinetically controlled. The relative height of this wave is strongly influenced by the uranium concentration. By increasing the concentration of uranium from 10" M to 10~ M the height of this wave increases from 45% to 85% in comparison with the height of the first reduction wave. The half-wave poten tial of the second wave shifts toward more positive potentials with increase of uranium concentration. The temperature coef ficient of the limiting current was calculated (0 - 40°C) as 3,06 kcal for the first wave and 2.42 kcal for the second wave. To investigate the nature of the second wave more closely the Kalousek commutator technique was used. The results obtained, which are independent of the frequency, show that the product of the second reduction wave is not oxidizable even at the potential of + 0„ 1 V. It can be concluded therefore, that uranium (IV) is the product of the second reduction wave. According to the litera ture uranium (IV) forms strong complexes with acetylacetone and this explains its electrochemical inertness . With gradual increase of the free ligand concentration the second wave shifts toward more negative potentials decreasing simultane ously in height until it disappears completely at log [Acac J } - -2. 5, i. e. under the conditions when the UO?Acac? reduces to UCUAcac-. It is evident from some of the above mentioned results that besides the adsorption of acetylacetone some reactants and/or products are also adsorbed. Adsorption phenomena were studied with Lippmann' s capillary electrometer and no adsorption of UCU and QC^Acac was observed. On the other hand strong adsorption of UO Acac. and uranium IV (product of the second reduction wave) was found. A pronounced adsorption was noticed in the whole investigated 194 Bozena Cosovic potential range (from 0.0 to -1.2 V) under certain conditions and it has been concluded therefrom that uranium (V) acetylacetonato complex is also adsorbed. Taking into account all the presented results the nature of the second wave can be explained by an ECE mechanism. Uranium (V) monoacetylacetone complex, formed either directly by reduction or by chemical equilibration between uranium (V) species, undergoes chemical change before the reduction to uranium (IV). This chemical reaction is strongly influenced by uranium concentration and by the presence of surface active substances. The most probable explanation of this reaction is dimerisation of uranium (V) monoacetylacetonate complex which is strongly adsorbed and therefore stabilized in the adsorbed layer of acetylacetone. REFERENCES 1. R. M. Izatt, B.P. Block, and W. C. Fernelius, J. Phys. Chem. 5_9 (1955) 80, 235. 2. J. Rydberg, Arkiv Kemi 8 (1955) 113. 3. Lj. Jeftic and M. Branica, Croat. Chem. Acta ^3_5_ (1963) 203. 4. M. Branica and J. Kuta, Coll. Czechoslov. Chem. Communs, jU (1966) 2833. 5. J. Rydberg, Acta Chem. Scand. 4 (1950) 1503; idem, Svensk. Kem. Tidskr. 6j (1955) 499. 3?au! Gso&an 195 NUCLEATION MECHANISM AND ELECTROCHEMICAL TRANSPORT PRO CESSES IN THE OXIDE FORMATION DURING THE ANODISATION OP ALUMINIUM Pro Paul Csokan General Design Institute for Engineering Industry, Budapest o Hungary We already stated that, the basic opinion in the classi cal KSLLER-HÖNTER--ROBINS0N theory is not in every respect correct© The apparently homogeneous expansion of oxide formation processes and the apparently even structure of the oxide layer is never in accordance with theoretical principles which can be generalised with regard to any variation of anodic oxidation but at most a valid pheno menon oocuring-under certain, closely limited experimen tal conditions* Our subsequent investigations demonstrated a basically different concept to the KELLER-HUNTER-ROBINSON model. Our microscopic motion-picture on documentary film taken during the process of anodic oxidation prove ondoubtedly our concept of the oxide formation mechanism on aluminium which may be divided into three stages: 1/ initial-period /formation of a chemisorbed oxigen-film/ 2/ nucleation-period /formation and growth of primary oxide nuclei/ 3/ formation-period /the oxide nuclei form a continous layer, formation of secondary oxidation zones and be gin to grow in depth, an upper porous layer is begin ning to form/. Therefore, no plain homogeneous basic layer is formed but primary oxide nuclei appear initially on predominant, e- nergetically preferre location. The number, size and dis tribution of primary nuclei and the phenomena of con - sequential processes depends strictly on experimental pa rameters. Under normal anodizing conditions /15-20% HpSO^. 12-18 ¥/ closely packed oxide nuclei form over the entire metal surface in fractions of a secundum, so that the surface is very rapidly covered by a continuous oxide film» If anodic oxidation is carried out by 20-25 V in dilute, e.g. 5% H2SO4, the process of nucleation is per ceptibly much slower and the primary, generally sphero- symmetrical oxide nuclei, which appear in considerably lesser amount, grow bigger. If the sulphuric acid concen tration is decreased down to 1% or even 0,1%, and a cell- -voltage of 40-60 V is applied, the amount of primary o- xide nuclei is rapidly decreasing, and their size reaches many times of the nuclei formed in concentrated solution. The primary nuclei grow up to a certain limit in three dimension. This process then ceases and a definite aniso tropic oxide formation in lateral direction begins 196 Paul Csokån in the boundary zone of primary nuclei and form terrace- -type secondary oxidation zones /Fig* 1-3<»A The spreading of secondary zones lasts until oxides touch eHoo other or grov; together and an oxide film forms which now dovers the entire surface o tfrom this moment onwards begins the stationary formation-periode* According to our experiments - together with other authors - now we can state that in the physically and chemically heterogeneous system of the anodic layers du ring the electrolysis the transport of the material par ticle is proceeding mostly by diffusion mechanism. The relation between the specific electrical conductivity and the diffusion coefficient of the ion-movement and the con nection with the defect-place can be expressed by the cor relation factor deduced from the NERNST-EINSTEIN equation» Experiments show that ion diffusion /e.g. diffusion of Al* and Al+++ ions, hidroxonium ions, etc/ developed by migration of defect-places /vacances/ has more importance like the interstition diffusion* Out of it the protomer ion motion is playing also important role in the system built up in the following variation; 0 Ål-0-AloO «*-*- Al-O-Al^ 0 <:—?> Al-OoooAl;^ <**-*> -*-> -0~A1~0-A1=0 In our view, light- and electron-microscopic picture re cords now available justify the virtually indisputable statement that pore formation is no way the decisive pre-requisite of oxide layer formation, but should be con sidered merely as a consequence of the selective chemical dissolution of the oxide in the electrolyte - a view also put forward some years ago by MURPHY and MICHELSON* The detail of the formation mechanism, the morphological characteristic /fibrous structure, basis-parallel diffe rent at ion of the oxide elementsf correlation with the theory from MURPHY and MICHELSON, and from other workers, such as GINSBERG, WEPERS, KADEN and BOGO3ÅVLÄSEI, etc •/, the surprising epitaxy-effect exerted by the quality of the base metal, the subsequent and sometimes strikingly uneven pore formation process are demonstrated by our documentary film under microscope ELYPOVIST, that be should like to show at the 23rd MEETING OP THE INTER NATIONAL SOCIETY OF ELECTROCHEMISTRY. /If the projection possible, presentation of the document- film /16 mm/, synchronised into English, projection time 20 minutes/. Paul Csokån 197 1 "-v?'^/v' $%*fy4' §1#SW |#lflSi :/;";";,«:i '!$###% 'MMlii ,•:• feJÉflé, Fig. 2 ./,-.:-« ,l *-V-s 1*4-»"' 'j *V »SV*** Fig. 3 :V * * J» 198 Karl Doblhofer LAPLACE PLANE ANALYSIS OF ELECTROCHEMICAL SYSTEMS: APPLICATION TO ++ Hg/Hg2 Karl Doblhofer and Arthur A. Pilla Electrochemistry Laboratory, ESB Inc, Yardley Penna. 19067 An analysis of the impedance of electrochemical systems must take into account the interactions of the various partial currents which may be involved in both faradaic and non-faradaic reactions. A description of the system in its most general form will include coupling terms among faradaic as well as non-faradaic partial currents in addition to the classical coupling between charging and faradaic currents. To illustrate the above, consider the simple adsorption of two elec- troactive species. For this system the essential quality is the variation of charge at the interface. At least three partial cur rents are necessary to describe this variation, each one given by an expression of the type: i] = i-j (r)s C-j, C2, T-j, IV,) (1) This particular quantity represents the rate of adsorption of species 1 having a bulk concentration C-j and surface excess V\. Examination of (1) shows that this partial current is explicitly coupled with any other partial current which is a function of any one of the same variables. This coupling has severe effects on the possibility of describing the system such that all kinetic quantities of interest are accessible experimentally. If the model is constructed such that the characteristic quality of impedance is obtained, it then becomes ambiguous to assiqn unique frequency behavior to the system: This means that the use of equiv alent electric circuits, even as shorthand notation to describe over all behavior, becomes useless in many cases since it is possible that an infinite number of minimum parameter (canonical) networks exist. The above generalities are treated in some detail in this study with special emphasis upon the useful limiting cases for impedance studies. For example, it is shown that the general lower frequency behavior for an electrochemical system is Z(s) = A + K//T (2) where s is the Laplace transform variable, K a function of the dif fusion coefficients and A is related to the charge transfer process. This of course is identical to Randles-Ershler type behavior. Other frequency ranges may be horribly complex, with the exception of the high frequency limit which always reduces to simple RC series behavior It is further shown that lower frequency behavior according to (2) in the freguency domain (with similar arguments holding in the time domain) need not necessarily mean that linear conditions (for true impedance) have been achieved. The manner in which this may be tested and circumvented is discussed in detail. Karl Doblhofer ias Application of the above approach has been mode to the mercury elec trode. Studies v/ere carried out in both the ideal polarized and faradaic potential regions usinq the transient impedance technique. Both real and imaginary axis impedance values were obtained at fre quencies up to 10$ rad/sec using an ultra fast potentiostat coupled with signal averaging technigues. Results in the double layer poten tial 'region showed no frequency dispersion. Results in the faradaic potential region showed that Hg/Hg2++ behaves following the straight forward Randles-Ershler scheme. No evidence for specific adsorption of Hg2++ was obtained nor was any unusual capacitance in the presence of the faradaic process. M. Elmas 200 STUDIES ON KINETICS OF ELECTROPHORETIC DEPOSITION M. Elmas Middle East Technical University, Ankara (present address: University of Karlsruhe, Karlsruhe). The electrodeposition of charged colloidal particles consists of their electrophoretic movement, in an ionic or polar liquid and their deposition on a surface. The difference between the electrodeposition processes taking place in aqueous and non-aqueous media is one of the degree but not of fundamental mechanism. To be able to use a coll oidal suspension as an electrodeposition cell, first it is required that the suspension should be stable: Yet at the same time one should be able to coagulate it on a required cite; two conflicting situa tions. Thus, one is concerned in stability of colloidal systems (1). DLVO (1,3) theory quantitatively predicts the way a particular system behaves. A number of other factors (to date not well under stood) affect the stability; i. Protection (stearic effects of ad sorbed polymers), ii. Sensitisation (ion-polymer interactions), iii. Adsorption (low polymer adsorption). The colloidal behaviour of aqueous systems is reasonably well established. However the stability of non-aqueous dispersions may be brought about by the joint action of first electrostatic repulsion, then the van der Waals forces and finally protection by added polymers. The surface charge causing interparticle repulsion mainly originates from dissociation of surface groups and surfactant adsorption (2). In these situations the relative acidity of solvent and that of particles are strongly coupled and determine the sign of surface charge. In practical electrodeposition cells the dispersion media used are organic alco hols of various carbon contents. It is known that the alcohols dis sociate to form RO~ and ROH* probably by the following mechanism (4): 2R0H =?--N R0~ + RO.i.C, with equilibrium predominantly lying to left and this is more so the larger the branching of alkyle chains get: Thus there is a mixture of ions and polar molecules in the solution, and when the fine particles to be electrodeposited (usually oxides) are introduced into it (specially into shorter alkylchains) the equilibrium shifts to the right.This is brought about by the adsorption of RO groups resulting in an effective surface charge (however a trace of H„0 present may be preferentially adsorbed and give rise to absorption of alcohol in liUO film formed) . In addition each oxide is known to exhibit relative acidic and basic properties. For in stance in £UO (an amphoteric solvent) SiO~ is acidic, A1«0„ is amphoteric and Mg~0 is basic. Relative to H~0, alcohols have somewhat similar properties but slightly more basic (ketones and amines also behave similarly); whereas nitro-paraffins are more acidic. Extending this similarity to alcohols SiOo will adsorp R0~, Mg~0; ROH+ and A^2°3' R0 » resulting resp. in negative, positive and A1„0„ will have a small positive charge (which will decrease with increase in alcohol chain length). From the above discussions it is clear that there are various mechanisms by which particle surface charges in non-aqueous media may be generated. In the literature one often finds different M. Elmas 201 proposals for the mechanism of electrodeposition in aqueous and non-aqueous suspensions in predicting time-current relationship as well as introducing critical time concept (5,6.7,8). Belcx^ we short ly describe what happens near the eiecLrode: (a) Liei_L rodeposi Lion in aqueous media: Basically the electrode to be coated is a "colloi dal particle" oppositely charged w.r.t. oncoming particles. Further more its charge as well as its ionic atmosphere may be changed at will by applying different potentials. This is not possible with the colloid particles; their surface charges car; only be changed by tempering with the solution and surface active agents. In the vicinity of electrode there exists electrostatic attraction and the electrostatic forces like van der Waals forces are attr icli ,-e . Therefore there is a favourable condition for the formation of the first layer. Eventually due to overlap of these oppositely charged ionic atmospheres a quasi-liquid molecular arrangement (causing an entropy decrease which creates an energy barrier preventing the final co-agulation) is set up. If kT-energy is sufficiently high the particles fall into attractive-potential-well and first layer forms. Meanwhile the electrolyte solution between electrode and particles will be dilated and squeezed out. However, in case of in sufficient kT-energy the electrostatic attraction could be made larger to cause final coagulation. Obviously this mechanism suggests that below a critical electrode potential first layer will not form. After the formation of first layer, the nature of electrode surface changes and it becomes a colloid having the same charge sign as the oncoming particles. The distribution of colloid particles (near the first deposited lciyer) is very similar to those of microions in the double layer of hydrophobic colloids. Further coagulation proceeds by the suppression of the double layer brought about by the excess electrolyte concentration due to the dilation between par ticles and to hitherto not well understood electrode reactions re sulting in ion transfer from the vicinity of electrode through de posited porous layer towards outer surface of the film where the collapse and the neutralisation of oncoming colloidal particles is taking place. Now since film's ohmic resistance is a function.of its porosity as well as of the interstitial mobilities of charge carry ing ions, the question whether ohmic drop or interstitial ionic diffusion is the rate determining process is immaterial. They should, in the end, lead to similar time-current och time-deposition rates (9,10). Thus, after instantaneous formation of first layer there is no critical time. However there is a critical electrode potential below which no deposition forms, (b) Electrodeposition in non-aqueous medåji. Again first layer deposit formation follows the same mechanism as in the case of aqueous electrodeposition. In such systems, however, because of the origin of surface charges and dielectric constants of medium, electrode's and particles' ionic atmospheres extend to large distances and the repulsion forces operate already at large separa tions; thus an energy barrier preventing further approach to the first layer is formed. Because the attraction forces causing coagu lation operate at short distances, for further deposition this !04 Oldfich Fischer Fig.l. Concentration of [Co en3^T £ 0.41mM, b_ 0.82mM, £ 1.64mM, free(enl.106 : 1 3.7, 2 4.9, 3_ 6,0, _4 8.2, \ 12.0, £ 12.2 M. Temp* 25°C. When the reaction (A) was studied as a chemical step preceding the oxi dation of Co(II) in the same solutions under galvanostatic conditions a value of K-^kT = *(9.4 +- 0.2).10>*'1 mol ^S4 0.5 1.0 1.5 was obtained (4). l/fif> , s-1/ 2 Also the anodic oxidation of p-phenylenediamine (PFD) dis solved in dry dimethylformamide (with 0,1M KC10, as the supporting electrolyte) is followed by an irreversible chemical reaction. The cation radical formed in the first electrochemical step undergoes chemical reactions for which following scheme can be proposed: NH2-C6H4-NH2 NH2-C6H4-NH + W (Bl) 2 NH2-C6H4-NH NH2-C6H4-NH-NH-C6H4-NH2 (B2) The two-step mechanism of the oxidation can be proved by means of cyclic voltammetry (Fig.2.) One electron con sumption for-the first step at the potential +0.55 V(SCE) was verified by coulometric analysis. The reaction (B2) was supposed to be the slower and therefore the rate de termining step. But the accurate evaluation of the rate Oldrich Fischer Fig.2. y 4 .x 1 5.10*"% PPD, 0.1M KC104, 0,31 V s" temp» 20°C. constants by galvanostatic electro lysis on the platinum electrode o with the area of 0.312 cm proved a first order irreversible chemical reaction obeying the relation(2/3): 2 (l+u)erf^k1T' = erf ^Itj+r') L J Results are summarised in The diffusion coef ficients of PPD were calculated from the transition time measu 1#2 8 0e4 w4 rements without °* y (sCE)"° current reversal. The temperaturre dependence of k-, has made possible to calculate the activation energy of this chemical step as E s 3.2 kcal mol » Temperature (°C) 15 25 35 46 k^.103 (a"1) 226 + 9 279 +14 3054; 7 434 +18 DPPDa°6(ca^ 8"1> 4*84 5.61 7.20 8.48 Literature: 1)D»Konrad,A.A.VlÖek:Collection Czech.Chem.Coram»,28»808(79 INFLUENCE OF ADSORBED ELECTROINACTIVE SUBSTANCES ON CHARGE-TRANSFER PROCESSES Rolando Guidelli Istituto di Chimica Analitica, Universita di Firenze, Firenze, Italia. The influence of specifically adsorbed electroinactive substances, either ionic or neutral, upon the rate of a charge-transfer process is considered to be due to two different effects, namely a blocking effect and an electrostatic effect. The blocking effect is accounted for by assuming that the charge- transfer process takes place both at the electrode surface covered with the surfactant and at the surface covered with the solvent according to a model similar to that proposed by Hush (J.Chem.Phys., _28,962,1956) . The ratio of the rate constant j_kf of the electrode process at the surfactant-covered surface to the corresponding rate constant kf at the solvent-covered surface is shown to be approxima tely given by the equation: k A G° . i f ads = exp RT o f where A G, o is the standard adsorption free energy of the surfac tant. The validity of the above equation, which holds for neutral surfactants, is demonstrated for several electroactive substances in the presence of several organic additives. The electrostatic effect is accounted for by introducing into the Frumkin equation the mean electrostatic potential cbi at the average position occupied by the discharging particle in the transition state. The potential O, is calculated according to a double layer model in which the potential drop across the compact layer is not regarded as a linear function of the distance from the electrode surface. The resulting equation for the rate constant kf as a function of the charge density q- at the inner Helmholtz plane, due to ionic specific adsorption, is applied to the discharge of H+ and Zn^+ on mercury in the presence of halide ions and to the discharge of In^+ in the presence of dodecylammonium ions. CR. Hamann 207 VERSUCHE ZUR BESTIMMUNG DES PH-WERTES KONZENTRIERTER SA'. RER LÖSUNGEN ÜBER DIE KINETIK DER WASSERSTOFFELEKTRODE S. Ernst und C.H. Hamann Institut für Physikalische Chemie der Universität Bonn, Bonn, Germany. Wie auf dem Treffen der Batterie-Sektion im April des Jahres 1970 in Strassburg von G. Clinckspoor, Bonn, bereits diskutiert wurde, ist die Kenntnis des Elektrolyt-pH-Wertes in der Batterieforschung von grossem Interesse. Angaben über diese Werte findet man jedoch selten, da in Batteriebau und -forschung häufig in starken Säure oder Laugenkonzentrationen gearbeitet wird, d.h. in Bereichen, in denen übliche Messverfahren (z.B. Glaselektroden) keine exakten Werte für die H+- resp. 0H~ -Ionenaktivitäten mehr liefern. Wie in der in Strassburg geführten Diskussion von unserer Seite bereits angemerkt wurde, liefert die pH-Abhängigkeit kinetischer Grössen eine Möglichkeit, auch in stark aktiven Elektrolytlösungen pH-Werte zu bestimmen. Untersucht man etwa die pH-Abhängigkeit der Austauschstromdichte der Wasserstoffelektrode, so erhält man aus der Gleichung der Durchtrittsstromspannungskurve die Beziehung In j0 = K + a' In a^ wenn man die Reaktionsgeschwindigkeitskonstanten sowie die Aktivi tät des adsorbierten Wasserstoffs als von der Protonenaktivität un abhängig ansetzt*' und hohe Ionenstärken voraussetzt. Es folgt x K § i0 - ' pH = - a' Durch Bestimmen der Austauschstromdichte in Elektrolytlösungen bekannten pH-Wertes können die konstanten K' und a' bestimmt werden, so dass die pH-Werte auch stark aktiver Lösungen der Messung unter Extrapolation der a*~- und K'-Werte zugänglich werden. Dieser Vortrag berichtet über bei derartigen Experimenten erhaltene Ergebnisse. x) Nur bei kathodischem Stromfluss gültig (Bedeckungsgrad Q = 1 über weite Strombereiche) ad 208 Norbert Ibl ELECTRODEPOSITION OF PALLADIUM POWDER N. Ibl, G. G-ut and M. Weber Technisch-chemisches Laboratorium, Swiss Federal Institute of Technology, Universitätsstrasse 6,Zurich, Switzerland Palladium was deposited from chloride solutions under poten- tiostatic and galvanostatic conditions. In the first case the current, in the second one the electrode potential va ries with time. The observed phenomena appear to be due to the very large solubility of hydrogen, in palladium on one hand, to an activation of the electrode surface on the other hand. In spite of this complication smooth steady state current voltage curves with a very well defined pla teau were obtained under specific conditionsu The roughness of the electrodeposited metal was measured with a mechanical profilometer, which follows the irregula rities of the surface with a tiny ball. The transition from an even to a very rough and powdery deposit takes place when one reaches the limiting current, in agreement with a general rule regarding the influence of mass trans port on the structure of electrodeposited metals. The catalytic activity of the electrodeposit was determined by measuring the rate of hydrogenation of cyclohexene under pressure and comparing it with that of a non-electrolytic powder. The surface area 9 was measured by the BET method and the grain size d by sedimentation. When increasing the current density the catalytic activity and 9 go through a marked maximum whereas d steadily decreases."The spectrum of the grain sizes obtained at a given current density is much broader in the region of the limiting current than at higher current densities. X-ray-diagrams revealed that the codeposited hydrogen first fcrms the a-H-Pd phase and at higher current densities the (3-phase. The maximum of the catalytic activity and of 9 corresponds to the transition from the a- to the p-phase. The interpretation of the maxi mum is discussed. Besides compact cathodes the metal was also electrodeposited in finely divided form on a suspension of active coal. This allows to obtain a catalytic material with a large active area as compared to the amount of the expensive metal used. The catalytic activity measured in the region of the maxi mum was approximately equal to that of the non-electrolytic catalytic powder used as test material. L.J.J . Janssen 209 ELECTROLYTIC REDUCTION OF NITRIC OXIDE TO HYDROXYLAMINE L.J.J. Janssen and J.G. Hoogland Eindhoven University of Technology, Eindhoven, The Netherlands. In acidic solution NO can be reduced to NH2OH by molecular #2 in the presence of a catalyst or by atomic hydrogen1). The electrochemical reduction of NO has been investigated only at a mercury dropping electrode. According to Masek^) the dissolved gas yields two polarographic reduction waves in a solution of pH 4.7, the sum of these waves corresponds with the reduction of NO to NH2OH. Ehman and Sawyer-^) found that the reduction of NO yields only one wave in a solution of pH 5 to 7; NO is reduced to N20. Our polarographic research shows that in strong acidic media (1-4 M HC1) the reduction of MO yields two distinctly separated waves. These waves, however, are not present in stronger acidic media (8.4 M HC1) whereas in weaker acidic media (0.4 M HC1) the first wave is clearly present and the second wave is overlapped by the wave of the hydrogen evolution. We found that in strong acidic media NO is reduced practically only to NH2OH. The reduction of NO cannot occur with high current densities at a smooth electrode because of the low solubility of NO and the large thickness of the diffusion layer. High geometric current densities are only possible with porous electrodes through which NO containing electrolyte flows4). A diaphragm cell was used. As cathode served a with electrolyte flow-through electrode, consisting of a number of metal gauzes lying on each other. The smallest mesh width of the available gauze was 4 5y. It is found that NO can be reduced to NH2OH on gauze of Pt, on amalgamated gauze of phosphor bronze or with Hg, Cd or Sn covered gauze of phosphor bronze. As cathode material only noble metals can be used and metals with a high hydrogen polarisation owing to the strong corrosive behaviour of a NO containing electrolyt. The current efficiency is dependent on pH. For an electrode consisting of 10 amalgamated phosphor bronze gauzes with a mesh width of 45u and a wire dimension of 35y, the current efficiency in sulphuric acid media at 200 mA/cm2 and 16°C is maximal in 2 M H2SO4. The current efficiency in 2 M H2SO4 is 83% and in 8 M H2S04 8%. For the formation of NH3 by reduction of NO# the current efficiency in 2 M H2SO4 is 0% and in 8 M H2S04 17%. 210 L.J.J. Janssen The current efficiency of the formation of NH2OH at 200 mA/cra2 and 20°C is 34% in 2 M H3PO4 with pH = 0.9 and 2% in a solution of 1.5 M K t^PO, and 1.5 M K2HPO4 solution of pH = 6.8. From 15 to 40°C the current efficiency is practically independent of the temperature. At a constant potential the current density of the NO reduction and the current efficiency strongly increase with decreasing mesh width. The number of gauzes per electrode also affects the maximum current density of the NO reduction; this maximum is 175 mA/cm^ for gauze with a mesh width of 45y and is reached at about 10 gauzes per electrode. The number of gauzes also strongly determines the relation between E and J for both solutions, with and without NO. The cathode potential is more positive in the presence of NO in the solution. The surface properties of amalgamated phosphor bronze gauze with a mesh width of 45y in 2 M H2S04 at 16°c» the experimental current efficiency, determined by titration of the NH2OH formed, is 33 mA/cm^ at total geometric current densities from 60 to 170 mA/cm2 . Taking into consideration a roughnessfactor of 2.75, the calculated thickness of the Nernst diffusion layer is 4.5u if the diffusion coefficient of NO is 10"^ cm2/s. The maximum geometric current density of the NO reduction in strong acidic electrolyte is mainly determined by the diffusion of NO, present in the pores of the electrode, to the inner electrode surface and by the penetration depth of the electric current into the pores of the electrode. It may be infered that by using gauze of a smaller mesh width than 45y the diffusion transport increases, the penetration depth of the current decreases however. The most favourable mesh width and wire dimension of the gauze of the electrode, for obtaining the highest current density of the NO reduction and the highest current efficiency, still have to be investigated. Literature 1. Gmelius Handbuch der Anorganischen Chemie, Verlag Chemie G.m.b.H, Weinheim, 8. Auflage, p. 718 2. J. Masek, Z. Anal. Chem. 224 (1967) 99 3. D.L. Ehman and D.T. Sawyer, J. Electroanal. Chem. 16 (1958) 541 4. L.J.J. Janssen and J.G. Hoogland, The Netherlands pat. appl. 6900493 Olga Korelic INFLUENCE OP pH ON ELECTRICAL BEHAVIOUR OP BARRIER LAYER ON ALUMINIUM Olga Korelié and Branko Lovrecek Institute of Electrochemistry and Electrochemical Technology, Faculty of Technology, University of Zagreb, Zagreb, Yugoslavia A number of Al/barrier layer/HJBO., solution systems were investigated within the pH range from 4.65 to 9.oo, The impedance was measured on the systems /a/ where the barrier layer was in contact with the electrolyte in which it was formed, and /b/ where the barrier layer was formed in one and the impedance measured in another electrolyte. Serial components of impedance Cs and Rs were determined in the frequency range from 5o c/s to 5o kc/s. The measured values show that Cs and Rs depend on fre quency, on pH of the solution in which the layer is for med, and on pH of the solution in which the impedance was measured.Such results point to the fact that the investi gated systems cannot be treated as simple non-r-ideal condenser, but that they are much more complex. The values obtained by measuring were analysed in details in order to get criteria for making conclusions about the mechanisms responsible for the given behaviour of the v phase layer as dielectricum. In that, values for tg *-£' were plotted. The plots of the above mentioned relations show that there is an essential difference in the behaviour of different systems in dependence on pH. The shape of cur ves for relations Rs-l/f and l/Cs-log f for the systems in the alcaline range points explicitly to theories which treat the phase layer as a cross-sectionally non-homo- geneus dielectricum. Accordingly, it was possible to calculate the change of specific resistivity as a function of the distance inside the oxide layer. Systems in the acid range could not be included into this theory. On the contrary, relations tgc» -logw, £*;-log \JJ , and $'* - 6J characterize the appearance of relaxation pro cesses for some systems in the acid range. Supposing the existance of relaxation of water dipols connected with the anodic layer /adsorbed or structurized/ heat of adsorpti on, according to Ebert and Langhamer, resp.free enthalpy of activation,according to Eyring, was calculated. xHigher Scool of Printing, Zagreb 212 Olga Korelié Between the acid and alcaline range there is a transition in which there are systems with mixed characteristics. With the increase of pH the predominance of relaxation phenomena vanishes gradually, and the properties charac teristic for a condenser with a cross-sectionally non- homogeneous dielectricum are more and more expressed. Systems having pH lower than 5»4- could not "be included in either of the above mentioned models. Anyway, with res pect to very high formation currents their oxide layers evidently do not represent typical barrier layers. The systems, where the phase layer is formed in the acid electrolyte and the impedance measured in an approxima tely neutral or alcaline range, possess mixed characte ristics, i.e. the ones of the system in which the layer is formed as well as those typical for the solutions in which the impedance is measured. W.J. Lorenz 213 THE INFLUENCE OF HALIDE IONS ON THE KINETICS OF IRON CORROSION N.A.Darwish, F.Hilbert, W. J. Lorenz and H. Roflwag Institute of Inorganic Technology and Analytical Chemistry, Technical University Graz, Austria and Institute of Physical Chemistry and Electrochemistry, University Karlsruhe, Germany. Introduction It is well known that the kinetics of iron corrosion are changed in acid solutions by the presence of halide ions. The two different dissolution mechanisms prevailing under normal conditions (1) are changed to ha lide inhibited mechanism. The kinetic data were up till now explained by the reaction sequence (2-4): + Fe+HzO _=, Fe(OH")ads+H [l] + Fe(OH-)ads+ X"ds H=i FeOH + x" + Z e" [z] From these equations the kinetic data (at 25 C) b = 60 mV , n, __ = + 1 , n , = - 1 + +, pH +, CI are calculated, if a Langmuir isotherm and ©v- —> 1 is assumed. Ex perimental results for solutions of pH 1 to 0, 5 and chloride concentra tions between 10 m and 2 m were in good agreement with this (2-4), In the present work we were interested in the dissolution behaviour of iron in concentrated NaCl + HC1 solutions. From our results the inva lidity of the above dissolution mechanism at concentrations of hydro chloric acid higher than 1 m follows. Experimental All experiments were done in oxygen-free solutions of A. R. grade che micals in conductance water. As electrode materials Ommet-iron (purity > 99, 98 /o Fe) and Armco-iron were used. The two materials gave identical results. The investigations were carried out in the sy stems I: Fe/HCl + NaCl (total-molarity 5, molarity of HC1 varied from 10 to 5 m), II: Fe/0, 1 m HCIO, + NaCIO, + NaCl 0 (total molarity b, molarity of NaCl varied from 2, 5 . 10 to 4,9 m) and 214 W.J. Lorcnz III: Fe/Zm HC1C) + NaClO, + NaCl ,44 -2 (total molarity 5, molarity of NaCl varied from 10 to 3 m). All experiments were done under potentiostatic control, with calomel reference electrodes. The potential values v/ere afterwards converted to the normal hydrogen scale. Results and discussion bt(rr*I o o The influence of increasing 0 0 0 HCl-concentrations (system o I) on the Tafel slopes of the 90 dissolution reaction is 80 shown in fig. 1. It can be o / o o seen that at pH 0 the slope 70 changes from about 60 mV 0 o 60 o o o to about 100 mV. This in —t fig. 1 crease we believe to be due -1 C^IMOL-L ! to a change in mechanism. This believe i-s confirmed by data shown in fig. 2. At the same pH the electroche mical reaction order with respect to c J_ changes s / from a negative value (as / calculated from the mecha nism eq. [l] , £2])to Ml V 03 05 1 2 3 i.5 fig. 2 n = 1 0 1 +,H30+ l' * ' ' That the observed changes are due to the combined in fluence of high concentra tions of chloride ions and hydrogen ions is proved by fig. 3, which shows the ex perimental results in 4, 9 m NaCl + 0, 1 m HCIO . Here the kinetic data of tne me TOO M0 500 »CO W0 200 B0 fig. 3 chanism eq. [l] , [2~\ are found. W.J. Lorenz 215 The experimental data found under the described conditions are entire ly new and very difficult to explain. Especially a positive reaction or der of hydrogen ions seems to be not in agreement with any of the pro posed dissolution mechanisms in literature. Very speculatively one could perhaps try to explain the results by a sequence like Fe + Cl" ===a FeCl , + e~ j~3 1 ads u J FeCl , +H ^^ FeClH n 4| ads ads u J FeClH* ===== FeCl+ + H+ + e~ |'5 ] ads aq L J from which, assuming a Temkin isotherm, the kinetic data b — 12 0mV; n .j. — + 1 ; n , si. -f 0, 5 are calculated. The measured reaction orders are related to the analytical concentrations and not to the acti vities of the ions. The activity coefficient of hydrogen ions is increa sing in concentrated solutions (5), which makes the value of n , questionable. ' In conclusion we would like to point out that we are far from having answers to most of the questions which arose during this work. Fur ther investigations are under progress in our laboratories, which we hope will give some of this answers. Literature 1) F. Hilbert, Y. Miyoshi, G. Eichkorn and W. J. Lorenz, J.Electrochem. Soc. 118, 1919, 1927(1971). 2) W. J. Lorenz, Corrosion Sci. 5_, 121 (1965) and references there. 3) L. L. Cavallaro, L. Felloni, G. Trabanelli and F. Pulidori, Electrochim. Acta 9_, 485 (1964). 4) K. Schwabe and C.Voigt, Electrochim. Acta L4, 869 (1969). 5) K. Schwabe, Z.phys.Chem. NF. _41 , 368(1964). 216 B„ lovrecek ELECTROCHEMICAL I1TV3STIGATI0LT OP VALVE IIETALS II. MetikoS, B. LovreSek, and B. Jari6 Institute of Electrochemistry and Electrochemical Technology, Faculty of Technology, University of Zagreb,Zagreb,Yugoslavia, In a wider program of our investigations on valve metals a special /V attention has recently been given to antimony.'7 e found earlier that in given circumstances the Schottky-Mott approximation can be applied to the anodic layer on antimony.However,a number- of phenomena,and particularly their quantitative evaluation,haven't as yet been given an adequate treatment what regards the anodic layer on antimony as semiconductor. In experimental work alternate current methods were applied for determining the impedance of the system/the bridge and the ellipse method/and stationary method current-potential. /1/ 22nd CITCE Meeting - Dubrovnik Investigation of Antimony as a Valve I.Ietal B. Lovrecek, B. Jarié and LI. MetikoS. ^îliane Î.Ioraot 217 ETUDE DU MECANISME DE L'OXYDO-REDUCTION DU Pt EN SOLUTION H SO, ,N Guy BRONOEL et Eliane MOMOT Laboratoire d'Electrolyse du C.N.R.S.- 92.BELLEVUE - FRANCE. Malgré de nombreux travaux sur ce sujet, le mécanisme de l'oxydation arvo,dique du platine ne peut, actuellement, être considéré comme éluci dé ; en particulier, Lorsqu'on veut rendre compte de la forme des cour», bes I =f(t), obtenues en réponse à une impulsion potentiostatique, on a pu montrer qu'un mécanisme simple ne pouvait fifere invoqué.çn L'analyse de mesures de la capacité différentielle de double cou che en fonction du potentiel et l'analyse des quantités d'électricité correspondant à l'obtention d'un recouvrement stationnaire,en fonction du potentiel, permettent de mesurer le taux de rugosité de l'électro de, la capacité spécifique du platine et le taux de recouvrement en oxyde. Des déterminations expérimentales ont été effectuées sur un échan tillon de platine polycristallin ayant subi, après polissage mécanique, 20 cycles de traitements électrochimiques du type Gilmanr' /„-. Si nous admettons suivant l'hypothèse classique de Frumkin que la capacité différentielle mesurée est reliée au recouvrement 8 suivant l'expression : c = c£ o (i-e) + c° oe (i) m M oxr C5. étant la capacité spécifique du métal (6=0) , C° , la capacité spé cifique du métal complètement recouvert (8=1) et p le taux de rugosité. On considère que les effets dus aux variations de recouvrement en es pèces ioniques adsorbées, sont négligeables devant les variations dues au recouvrement 9 en forme oxydée. D'après la courbe C = f(E-) on peut déterminer p(Tl et pC° . Puisque 9 égal q/q où q est la quantité d'électricité correspondant à un re couvrement maximal, L'équation (1) peut- s'écrire : Connaissant pCJU pC° , C et q qui est déterminé en intégrant, compte tenu des effets capacities, la courbe I=f(t) obtenue en réponse à une impulsion potentiostatique* on peut obtenir qQ et p. Les résultats obtenus nous permettent de conclure que le domaine biphasé Pt/Pt(0H) est compris entre 450 et 800 mV/ENH et que pour des potentiels plus anodiques que 800 mV il y a apparition d'une nouvelle forme oxydée. Nous avons pu déterminer la courbe ^ s f(E) et cons tater qu'au potentiel d'équilibre du système correspond un recouvrement de l'ordre de 0,215. Le taux de rugosité trouvé dans notre cas est de 2,25. Le taux important de rugosité qui a été trouvé étant vraisemblable ment dû aux traitements électrochimiques qui ont été effectués initia lement sur l'électrode, il nous a paru intéressant d'étudier, suivant cette même méthode l'évolution du taux de rugosité en fonction du 218 Eliaiie lïonot nombre de cycles de Gilman effectués sur l'électrode. L'analyse des résultats obtenus montre qu'après environ 20 trai tements, un état structural stationnaire est obtenu. L'action des traitements du genre Gilman, se traduit non seulement par un développement de la surface (x 1,79) , mais modifie l'énergie superficielle du platine non oxydé, donc entraîne des modifications tant dans les propriétés de la double couche (La capacité spécifique du platine varie de 20,5 à 16,5|jF par cm ) que des énergies mises en jeu dans les réactions d'oxydation superficielle (variation du poten tiel d'équilibre ). Nous avons complété ces informations en analysant la relation existant entre la quantité d'électricité q , à l'état stable et le po tentiel E de l'électrode. En effet, il est évident que la forme analy tique de la courbe q = f(E) dépend du mécanisme impliqué dans 1'oxydo- réduction de la surface. Nous avons analysé cette courbe q = f(E) en supposant différents mécanismes : soit un transfert simple, soit une adsorption préalable de dipôles d'eau, soit encore des effets d'interaction dans la couche adsorbée. Il s'avère qu'à l'état stationnaire un formalisme de la forme : i £Fn r + (l-cr)ZFr, t RT MLO ; C1—9 ) exp £=— exp - 2y6 = K' H ,9 exp - permet de rendre compte des courbes obtenues (entre 450 et 700 mV). Ceci revient à impliquer un phénomène d'interaction , et 1*ident ification des résultats conduisent à la détermination d'un coefficient d'interaction de l'ordre de 23. En ce qui concerne le comportement à l'état transitoire, c'est-à- dire l'étude des courbes I = f(t) obtenues en réponse à une impulsion potentiostatique, l'étude des domaines de recouvrement en oxyde dont nous venons de donner les résultats montre que î1on ne saurait obser ver un phénomène élémentaire qu'à la condition de se placer dans un domaine strictement biphasé qui, pour l'équilibre Pt-PtOH, est com pris entre 450 et 700 mV. Mais on constate que, même dans ces condi tions les formes obtenues sont d'une analyse difficile, puisque les processus impliqués se caractérisent par des phénomènes d'adsorption et d'interaction. BIBLIOGRAPHIE 1. E. M0MOT, M*. BONNEMAY, G. BRONOEL, C.R. Acad. Se. Paris, 270,1970,2108 271,1970, 334. 2. S. GILMAN, J. Phys. Chem'. ,£6, (1962),2657. 3. A.N.FRUMKIN, J. Phys. Chem. 35, (1926) 792. Eliane I'.omot 219 J ' 0 PtOH i / 0,4 / Pt + PtOII 0,3 ! + PtO i 0,2 i i l 0,1 • 1 /Sif\ + PtOH ! 400 600 800 mV/EIIK Pig» 1 — Variation du recouvrement en fonction du potentiel 2,5 cm* 2^ 1,5 h 15 10 10 20 nombre de trait s Pig. 2 - a) Variation du taux de rugosité /° en fonction du nombre de traitements de Gilnian b) Variation de la capacité spécifique Cg en fonction des traitements» 220 K.Iioslavac SPECIFIC ROLE OF CHLORIDE IN ANODIC REACTIONS ON GOLD B.Lovrecek, K.Hoslavac, R.Radekax Institute of Electrochemistry and Electrochemical Technologyj Faculty of Technology, University of Zagreb, Zagreb, Yugoslavia In continuation of our investigations of anodic reacti ons on gold the specific role of chloride was investiga ted.Investigations were performed by means of the galva- nostatic pulse method with additional "potential sweep" measurements» The following systems were investigated: 1. Au / o.99 N H2S0/+ , o.ol N KOI 2. Au / 0.99 N H2S04 , o.oo5 N KC1, o.oo5 N HC1 S0 3. Au / 0.98 N H2S0^ , o.ol N E2 4 > o«o°5 N KOI, o.oo5 N HC1 4. Au / o.97 N H2S0^ , o.o2 N K2S0^ , o.ol N HCl 5. Au / o.o97 N H2S0^ , o.o2 N K2S0^ , o.ol N KOI The common characteristic of electrolytes of all the systems investigated is H2SO4 as supporting electrolyte, as well as the constant ionic strength. For the sake of investigation of an eventual influence of cation / K+/ in some electrolytes a part of sulfuric acid was repla ced by potassium sulfate, but in a quantity which only slightly changed the pH value of the whole electrolyte. By means of the pulse galvanostatic, method characteris tic e-^-t curves were obtained, which, in investigated systems, dependent on the applied current density make two groups. One at low current densities where the cor responding e^-t curve shows only one step at the potenti al of about IV vs. standard hydrogen electrode. The other group consists of galvanostatic ei-t curves, obtai ned by applying higher current densities. It is characte ristic for them that the step at about 1 V disappears, but two steps appear at more positive potentials. The missing of the first step at high current densities does not mean necessarily that the corresponding reaction does not take place at all, because if a very small quantity of current 'is necessary for it, this step is not recorded under the given measuring conditions. 7R.R./ High Technical School of JNA, Zagreb KoMoslavac 221 QJhe obtained results were analyzed and they indicated that the rate determining processes in anodic reactions expressed by the appearance of different steps on the e.j_-t curve were not the same. A detailed analysis sug gested either a covering with a non-conductive layer or a diffusion as controlling factor in single anodic reac tions. VJ.J.Plieth 222 ADSORPTION AND INHIBITION IN THE HYDReQUlNQNE/QUINONE REDOX-SYSTEM ON MERCURY W.J.Plieth, I.Stellmacher, B.Quast Institut fiir Physikalische Chemie der Freien Universität, Berlin-Dahlen The adsoption behaviour and. electrochemical behaviour of the hydroquinone (QH2)/Quinone (Q) redox system on mercury were investigated under simultaneous conditions. Solutions of 0.5 M MaF and 0.025 M NaH2P04/Na2HPOi| with vailing ratios of NaH2P0i| and Na2HP0j| were used as the basic electrolyte to which different amounts of hydroquinone and quinone were added. The concentration of the redox components varied from 1«10~^ M to 1-lQ"2 M. The pH ranged from 5.5 to 7.2. The surface coverage^- can be calculated from the concentration dependences of the surface tension # . For the reversible system considered here, the derivation of Lippmann's electrocapillary equation leads to the follo wing three differential equations dy _ pt o_ H ^ÖHI " ^ 2(H20) " 2F dx - r a_ ~STJ^ " Q(H20) " 2F o& _ q_ dMH+ ~ P where )i is the chemical potential, q the free charge on anc3 the electrode, and ^QU9(HPO) - *"b.(H2°) the relative surface excess of QH2' and Q. One half of a possible relative surface excess of a semiquinone QH is contained in each of the P- values. When deriving the above equa tions adsorption? of fluoride', phosphate .and protons was neglected. Using these equations, the sum H^HPCHPO) + + IQ(£[ Q^and the single P -values can be calculated. The surfaci tension was measured, using the drop-weight method. An approximately constant surface coverage I5H2CH2O) + + T"Q(H?0) = l»l6»10" Mol/cm2 was found in these inve stigations. When one redox concentration was altered keeping the other redox concentration constant, only the ratio between hydroquinone and quinone in the adsorption layer changed. Thereby the increase in P of the oxidized component with more anodic equilibrium potential or the increase in P of the reduced component vjith more cathodic equilibrium potential was compensated by a simultaneous decrease of the P value of the other redox component. W.J.Plieth 223 In the electrochemical investigations the concentration dependence of the charge transfer resistance was measured. The experimentally determinable polarization resistance Rp is the sum of the transfer resistance RQ and a diffu sion resistance R& Rp = RD + Rd Working with a hanging mercury drop in a rotating electro lyte, Rd could be determined separately by measuring the diffusion currents i^. It is Rd = RT/2F • (l/idjQHp + l/idsQ). Then Rp can be calculated. Theoretically it can besshown that Rp is inversely proportional to the overall rate of the electrode process, regardless of the special redox mechanism. Considering the adsorption behaviour, the results obtained for the concentration dependence of Rp could be interpre ted by a mechanism QH2 «, (QH2)ad + \ e- + (QH2)ad *y (QH)ad + H + X2 e- (QH)ad **<*>ad + H + ^3 e" (Q)ad ^ Q + ELECTROCHEMICAL STABILITY AND ELECTROCATALYTIC BEHAVIOUR OF SODIUM- TUNGSTEN BRONZES Jean-Paul Randin and Ashok K. Vijh Hydro-Quebec Institute of Research, Vavennes, P.Q., Canada. Sodium-tungsten bronzes containing traces of platinum have been re ported to be good catalysts for the reduction of oxygen (1-7). In the present study, bronzes of the general formula NaxW03 (o (anodic) y PtO + NaxW03 -»- y Pt + NaxWO(3+ . [2] and (cathodic) y Pt + NaxWO,3+x -*• NaxW03 + y PtO [3] 226 Jean-Paul Randin In the reactions [2] and [3], Pt sites act merely as porters of oxy gen to and from the bronze sites by means of a path of low activation energy. The production of PtO consumed in the step [2] can take place not on ly by step [3] but also by means of pseudo-faradaic anodic steps such as: + Pt + H20 -»• Pt-OH+ H -+ e [4] followed by, Pt - OH -v PtO + H+ + e [5] In order to obtain 0 ~ for step [1] from PtO, one may envisage the cathodic steps (occurring perhaps within the bronze where reduction of PtO to PtOH would not appear to be feasible): PtO + e -> PtO" [6] followed by': 2_ PtO" + e -> Pt + 0 [7] It may be added that reaction [1] can also proceed by the direct dis charge of water on the bronze substrate as in the following step: + NaxW03 + yH20^Na W0,3+ . +2y H + 2y e [8] Since in the platinum-containing samples there is some*evidence that the presence of platinum catalyzes reaction [1], it seems preferable to suggest, therefore, a sequence of steps such as in Eqs [2] - [7] in which the production of 0 • occurs by participation of platinum sites. However, in the samples devoid of platinum, oxidation- reduction of the bronze would take place by a step such as in Eq. [8]. If the solid state process [1] is diffusion-controlled and if there is a simultaneous occurrence of the pseudo-faradaic processes suggested here, reaction [4] and [5], mixed kinetics would be expected as in deed is observed. References 1. D.B. Sepa, A. Damjanovic and J.O'M. Bockris, Electrochim. Acta 12, 746 (1967). 2. A. Damjanovic, D. Sepa and J.O'M. Bockris, J. Res. Inst. Catal. Hokhaido Univ. 1£, 1 (1968). 3. B. Broyde, J. Catalysis, 10_, 13 (1968). 4. J.M. Fishman, J.F. Henry and S. Tessore,Electrochim.Acta jL4_, 1314 (1969). 5. J.O'M. Bockris, A. Damjanovic and J. McHardy, Proc. 3rd Int. Symp. Fuel Cells, Brussels, Belgium, 16-20 June 1969, Press. Acad. Européennes, Brussels, 1969, p. 15. 6. R.A. Fredlein and J. McHardy, ^24th Power Sources Symp. P.S.C. Pub. Comittee, Red Bank, N.J. 1970, p. 175. 7. R.D. Armstrong, A.F. Douglas and D.E. Williams, Energy Conversion, 11, 7 (1971). John C. Reeve 227 THE INFLUENCE OF SOME ANIONS ON THE ANODIC DISSOLUTION OF COPPER. John C. Reeve Chemistry Department A, The Technical University of Denmark, Lyngby, Denmark. Anions forming notably stable complexes with Cu(I) in aqueous solution (e.g.chloride ions) have for a long time been known to have a marked influence on the anodic disso lution of copper by promotion of the loss of Cu(I). A pre viously undetected influence of the sulphate ion on the anodic Tafel slope was recently mentioned by the author[1J and is described now more fully; results demonstrating the stoichiometric kinetic involvement of sulphate in one of two parallel processes are also given. The electrodes (RDE) were made from pure, quenched, poly- crystalline, wet ground, nitric acid etched copper in rod form. Most experiments were carried out at 2 5°C using elec trolytes of 1M sulphuric acid, 1M perchloric acid (for both of which it is hoped the influence of the diffuse double layer will be small) or mixtures of these acids (analytical grade in thrice distilled water) and the cur rent response to a succession of single 100irV/sec. posi tive-going and return potential scans examined. An elec trode 'rest-period' (between the above scans) at c. 30 pA/cm2 anodic avoided the usual large drift in electrode properties, and highly reproducible, hysteresis free Tafel response was achieved over two current decades in sulphuric acid and over three current decades or more in perchloric acid; during a potential scan some 100 atomic layers are removed. The Tafel slope (b) in 1M perchloric acid is typi cally (61±3T) mV and in 1M sulphuric acid (47±3f) mV; ty pical currents at +305 mV (nhe) are 0.7 mA/cm2 and 7 mA/cm , respectively. The rotation rate (w) of the RDE was normally 1000 rpm, and the Tafel responses were independent of w. As a result of larger currents passing for longer periods the above Tafel slopes are reduced and the above currents increased. Examination of these Tafel responses over the range 0° to 7 5°C indicated no change in Tafel slope within ± c. 1 mV - the precision of measurement. This possible variation is much less than theoretically predicted for a simple model and thus complements the similar, much earlier cathodic result of Mattsson and Lindström [2]. The significance of these results is not considered further here; the influence of temperature in these and other electrolytes will be described elsewhere. Addition of 1M H2S04 to IM HCIO4 and plotting of reciprocal current (i~^) at a given potential versus b""-*- (Figure 1) gives indication of the involvement of largely independent f standard deviation for some 40 electrodes. 228 John C. Reeve parallel reactions [1] and a limiting Tafel law of 40-50 mV depending on the particular electrode. Because of the indicated involvement of parallel reactions log(i-ip) was plotted against log Ms ; ip is the current observed in IK perchloric acid, i the actual current ob served at the same potential and Ms the molarity of sul phuric acid in the mixed IK acids. The reaction order in dicated is 0.5 when the extrapolated Tafel slope (plots of i^vs.b--'-) has the typical value of c. 47 IPV (Figure 2), but shows a tendency to change towards unity when the extrapolated Tafel slope is closer to 40 mV. These results are all readily explained by an earlier sug gested model 11J of two parallel reactions. One reaction involves surface diffusion of Cu+ from steps with subse- quent electrodic oxidation of Cu to Cu on the inter-step surface, the other (giving the c. 60 mV Tafel response) involves either a potential independent rate-determining step subsequent to production of Cu or the less likely 'L3] direct two-electron transfer [4J production of Cu^+ (seve ral lines of evidence exclude loss of Cu+ as the explana tion). The current for the first mentioned path will ide ally exhibit a 40 to 47 mV Tafel response depending on the relative rate of surface diffusion and will be 'proportio nal to km where m changes from 1 to 0.5 between these li miting situations; k? is the rate constant for the, reaction Cu+ -> Cu2 + +e. Hence if k2 should be replaced by k^H^SOLj.], a not unreasonable involvement of HoSOq., the results are readily explained in a relatively simple manner which is in keeping with the observation by a number of workers that large areas of the electrode are often relatively inactive. The results of transient experiments by Brown and Thirsk [5] are apparently in rather better agreement with this model in which [Cu+j at the electrode should increase ac cording to a c. 80 mV-law rather than a c. 60 mV-law. These results could be of considerable importance, in ex plaining some phenomena related to the cathodic processes at copper electrodes. Cathodic experiments, which under some conditions indicate a slow potential independent step immediately after diffu sion of Cu from the bulk of the solution in perchloric acid and 'simple' behaviour in sulphuric acid [1] are to be described elsewhere; however, the diffusion coefficient of cupric ions in 1M perchloric acid was found to be some what greater than that in 1M sulphuric acid and the catho dic Tafel slope for the charge-transfer step was c. 14 0 mV (over c. 2 decades). References: [ljJ.C. Reeve, Coll .Czech .Chem .Comm.3_6_( 1971) 75 7 . [2]E.Mattsson and R.Lindström, in Proceedings Gth Meeting John C. Reeve 229 C.I.T.C.E. (Poitiers,1954), Butterworths, London(1955); Section 5.2.6, page 263. L3]B.E.Conway and J.0'M.Bockris, Proc.Royal Soc. (London) A 248(1958)394. [1]D.C.Grahame, Ann .Rev .Phys .Chem . ,6_( 1955)337. 15]0.R.Brown and H.R.Thirsk, Electrochim.Acta,10(1965) 383 log! i-ip ) /' i:/jA/0.2 cm2 at 400 mV ) ( i:pA/0.2 cm2at 400 mV ) 3.5 3.0 2.5 FIGURE 1. FIGURE 2. FIGURE 1: Influence of addition of 1M H2S04 to IM HC1O4. 10d/i versus 10^/b. The ringed point was obtained by addi tion of 1M Ba(C10t|)2 "to ^e fi-na-l solution. FIGURE 2: log i and log(i-ip) versus log(sulphuric acid molarity) illustrating a reaction order of c. 0.5 for the sulphuric acid catalysed parallel reaction. 230 N.A. Shumilova HYDROGEN PEROXIDE BEHAVIOR ON SILVER AND NICKEL ELECTRODES IN ALKA LINE SOLUTION N.A. Shumilova, N.D. Merkulova, G.V. Zhitaeva, E.I. Khrushcheva, G.P. Samoilov and V.S. Bagotzky Institute of Electrochemistry, Academy of Sciences, Moscow, U.S.S.R. The complexity and multistep nature of the reaction of electrochemi cal oxygen reduction on solid electrodes makes it necessary to use a variety of methods for its investigation. On the basis of the data obtained by various adsorption and pulse methods as well as by the rotating ring-disc electrode method, we can assume the oxygen reduc tion reaction on silver and nickel as well as on many platinum metals to follow two parallel courses, one of these being via an intermediate hydrogen peroxide formation. The behavior of the hydrogen peroxide formed is no less complex. It can undergo electrochemical reduction and oxidation as well as catalytic decomposition. With the help of the rotating ring-disc electrode method and a complex mathematical treatment of the experimental results proposed for the simplest case by Bockris and developed at our Institute we have been able to interpret the main hydrogen peroxide reactions and to esti mate the rate constants of these reactions, depending on the poten tial, the state of the electrode surface and the electrolyte compo sition. On a silver electrode at the steady state potential hydrogen per oxide undergoes electrochemical reduction, oxidation and catalytic decomposition (k», k~, k,) . The fact that these constants are equal indicates that catalytic decomposition of hydrogen peroxide at the steady state potentials occurs at conmensurable rates both by elect rochemical and purely chemical mechanisms. The constant k, depends little on the potential and is equal to about 1.10 cm/sec. When the potential shifts both in the anodic and cathodic directions the cata lytic decomposition reaction occurs mainly by a nonelectrochemical mechanism. On a nickel electrode at more cathodic potentials than' the steady state potential hydrogen peroxide practically does not decompose. The polarization curves of hydrogen peroxide reduction are of kinetic nature. The reduction rate depends strongLy on the state of the nickel electrode surface: on oxidized nickel it is by 1.0-1.5 orders of magnitude less than on reduced nickel. The oxidation state of the electrode practically does not affect the oxidation rate of hydrogen- peroxide on nickel. Some suggestions are made about a different effect of the state of nickel surface on the oxidation and reduction rates of hydrogen peroxide. Otomar Spalek 231 THE INFLUENCE OF SURFACE OXIDES OF PLATINUM ON THE RATE OF ANODIC OXYGEN EVOLUTION IN SOLUTIONS OF SULPHURIC ACID Qjcmar § pa lek and Jan Balej Institute of Inorganic Chemistry,Czechoslovak Academy of SciencesjPrague, GSSRe The rate of oxygen evolution depends strongly on the pre- treatment of electrode. The electrode properties are in fluenced mainly by the amount and the structure of surface oxides. In this work, we investigated the influence of two types of surface oxides of platinum in solutions of sulphuric acid on oxygen evolution in the potential region 1,6 5 - 2'„15 V(REE)o According to the potential and the time of anodic prepolarization, the surface of platinum electrode contains various amounts of chemisorbed oxygen" (oxide I) 2 and multilayer oxide (oxide II)0 The rate of oxygen evolution was studie"1 by measuring of polarization curves, after which the amount of cxic.es was determined by vcltamnietry0 Polarization curves of oxygen evolution have two branches (see Figol and 2). At lower potentials (section A),they are linear and have the slope b^OjlBS V. In higher poten tial region (section B), the polarization curves are no more linear and their slopes are the lower, the higher is the potential,, In Figoljpolarization curves of the electrode prepolarized at potential 1,9V, when only oxide I on the surface of the electrode is formed, are shown. In whole measured re gion, the overvoltage of oxygen evolution Increases with time of prepolarization. In. Fig.2, polarization curves of the electrode prepolari zed at 2,15V are shown» The slcpe of part B of polarization curves decreases with increasing time of prepolarizaticn0 .After selective reduction of oxide I, the slope of the curve in region B does not changes and only after re duction of oxide II initial polarization curve is re stored. The described decline of the slope of polarization curves was not observed in the experiments carried out with the electrode prepolarized at potentials higher than 2:,5V. This conforms the fact, that in this potential region no oxide II is formed. We considered possible causes of lower slope of polari zation curves in region B in comparison with region A. We supposed,that in region A the^rate determining step of oxygen evolution is following^ : + E2Q ^ H -*- OH (ads) + e (l) For the lower slcpe of polarization curves in region B we found the only explanation, according to which mole- ,32 C tornar äpale-k cular oxygen is formed simultaneously in another reaction way. According to this assumption, the total current den sities in section B were divided into the current densi ties of oxygen evolution by reaction path 1 (its rate de- termini r g step is reaction (1)) and that by reaction path 2a By this way,partial polarization curves shown in Figs o 3 and 4 were obtained,, The linearity cf all curves sup ports the presumption above mentioned. The concept of two paralel reaction path is also supported by the fact, that the presence of oxide II accelerates only the re action path 2* Because the reaction (1) is the rate con trolling step of reaction path 1, it cannot be included in reaction path 2n Oxidation of water by path 2 proceeds most likely by one of following reactions: + H20 - O(s.ds) + 2H + 2e (2) o HSC) + H.-.0 OH(ads)+ H" + HS07 (3) From the slope b = 59-68 LIV, it can be concluded, that each of these two reactions can be the rate determinjiig step of reaction path 2, References i lo T^Blegler,R.Woods:JoFlectroanal0Chem„2g,73-78 (1969), RoThacker^JoP.HoareiJ.Electroanal.Chem.^, 1, (1971) o 2. S„P.James:vJ.Electrochem.Soc. 116., 1681 (1969), JoBaioj,0,Spalek: Collection (in print), Extended Ab stracts of the CITGE Meeting in Prag 1970, p. 114. 3o Jo0'M.Bockris, A.K0M.S0Huq : Proc.JRoy.Soc. A^237J 277 (19 56")*. Otomar äpalek 233 J 1 3 logi fAjtrfJ "' ty'lA/cm ] ' Fig.l 5ig. 2 Figc] -Potentiostatic polarization curves of oxygen lut i o n on platinun. in IN- H..SC „ at 15 °C. Prepevoo - zation of the electrode st^ljSV: a- 1 min,, b lar -i - 5- min., c-60 mir_0 Curves measured from higher- to lower potentials. Time cLela.y at each potentia R. seco Fig.2 -Potentiostatic polarization curves of oxygen c e vo luticn on platinum in IN H?SQL at 15i°C. Prepo lar i- zation cf the electrode at 2915V: a- 5mii ,b-20mir_ c-1 h, d- 4hs Z2 1 i i i i 2.2 ' ~ 1 1 i E £ - - - •i?- M 2b^ •^d 2.0 2.0\- *^2c 2*^/ - _ - - t.8- 1.8 - - k/y - fe/5 - //la /la 1 16 1 i \ i IS 1 I I •J -3 logifiW -s logifAjHf1 Eig.3, Fig. A 3Fig3,4 -Partial polarization curves of oxygen evolution by -two paralel mechanisms. Prepolarization of the electrode at the potential 1,9V (Fig.3) and 2,15V (Fig.4) 234 Zuzanna Ssklarska-Smialowska USE OF ROTATING DISC CATHODES POR STUDYING HYDROGEN PERMEATION THROUGH IRON AND STEEL Zuzanna Szklarska-émialowska, Michal émialowski and Tadeusz Zakroczymski Institute of Physical Chemistry of the Polish Academy of Sciences, Warsaw, Poland The ability of hydrogen evolved electrolyticaliy to pene trate iron and steel membranes is controlled by its che mical potential which, in turn, *is an unknown function of various kinetic factors: cathodic current density, cata lytic activity of the metal surface, velocities of dif ferent electrode reactions occurring in the given elect rolyte etc. It is well known that pH plays an important role in the mechanism of cathodic reactions and that certain ele ments, particularly P, S, As, Sb, Se, Te, promote hydro gen entry into the cathode. This effect occurs already at negligible concentrations of the given element and, therefore, its nature is considered catalytic, but the exact manner by which promoters enhance the hydrogen entry into metals is in the dark. Previous investigations /1-3/ have shown that the forma tion of hydrides of the promoters is decisive for their effectiveness and that there occurs a relationship be tween the bond strength of the given hydride and its ability tc promote the entry of hydrogen into metals during cathodic polarization. The aim of the present work was to contribute to the knowledge of the mechanism of above phenomena. Technique of rotating disc cathodes and an electrolyte of low acidity were used to study the promoting effect both below and above the limiting current for discharge of hydronium ions. Main series of experiments were perform ed at 25°C on Fe 5% Ni alloy. This material instead of iron was chosen owing to its lower corrosion current density. The cathode had the form of a membrane, 0.1 mm thick, 14.5 mm. in dia., soldered to a steel cylinder, whose ex ternal surface was covered by a layer of polymethacryl- ate. The assembly rotated around its vertical axis at a speed of up to 35 rev.p.sec. The bottom of the membra ne, brought into contact with the electrolyte, was pola rized cath.odical.ly at a constant current. The volume of permeating hydrogen was determined by measuring the in crease in pressure within the steel cylinder. Prelimina ry measurements have shown that in spite of the formation of hydrogen bubbles on the cathode, its behaviour corres ponds to common rules established for rotating disc Zuzanna Szklarska-émialowska 235 2 2.5 3 J,5 4 4,5 o&um BOOT a casntet tsairt [«A/«?T 9QQ&BB BOOT OT CVBSBBt DSEJm [BA/CSI^1/3] Fig.1.Effect of current den Fig.2.Effect of current den sity on hydrogen permeation sity on hydrogen permeation through Pe5^Ni membranes in through Fe5>SN"i membranes in 0.1 N sodium sulphate solu 0.1 N" sodium sulphate solu tion with sulphuric acid, tion with sulphuric acid, pH=2.6,and promoters added: pH=2.6, and arsenic trio- a-arsenic trioxide,b-antimo - xide, at rotation speeds: ny trioxide,c-selenium dio a-4 rev.p.sec.b-9 rev.p. xide ,d-thiourea. sec.,c-16 rev.p.sec., Speeds9 rev.p#sec. d-25 rev.p.sec. electrodes. Some examples of experimental i esuits are given in Figs. 1-3» 3?ig.1 shows the effect of different promoters added in small amounts to the electrolyte on the permeation rate of hydrogen through 0.1 mm thick Fe5%&i alloy membranes. £n this case the rotation speed of the disc was constant &nd one measured permeation at different constant cathod- %c current densities. Under these conditions,initially, the permeation rate increased with time,attained a maxi mum, and decreased to an approximately constant final va lue. This was due to changes within the membrane. and on its surface. In Fig.1,the maximum permeation values are lotted.lt may be seen that below the limiting current tensity for the discharge of hydronium ions, in this case iL=12 mA p.sq.cm.,all promoters being used provoke much higher permeation values than above that limit.The drop in the permeation rate near i^ is particularly steep in the presence of selenium,curve c,and less pronounced in that of thiourea,curve d. At higher current densities than ij^the" hydrogen permeation rate in the presence of both Se and thiourea is very low. On the contrary, in the presence of As and Sb in the electrolyte, a quite rapid hydrogen permeation can still be observed at current densities much higher than i-r. 236 Zuzanna Szklarska-Smialowska ?00, The effect of changing the rotation speed of the disc in the presence of arsenic is shown in Pig.2. Pig.3 shows the effect of time on hydrogen permeation through a stationary catho de polarized in absence of promoters at current densi ties laying below, curve a, and above, curve b, the li miting ij, value» In this 15 ao n 50 XUS [•Hint.*] case, an electrolyte with Pig.3.Hydrogen permeation pH=1,75 was used,for which through a stationary Fe5%Ni il,=4- mA p.sq.cm. Under the membrane vs. time of catho- se conditions, particularly dic polarization at 0.25 mA above ij,, the permeation p.sq^cm.,curve a, and 9 inA rate was very low. p.sq.cm.,curve b, in 0.1 N No distinct correlation was sodium sulphate solution observed between the hydro with sulphuric acid,pH=1.75« gen permeation rate and the No promoter added. ability of the given promo ter to affect the course of cathodic polarization curves. Considerable permeation can likewise be observed at high and low overvoltages. The viewpoint of Bockris et al. /4/ that the rate of electrochemical charging with hydrogen may be associated with a certain overpo- tential could not be confirmed. Our results confirm earlier conclusions /1-3/ that the hydrides of the promoters play an essential part in the mechanism of hydrogen penetration into the cathode. The ability of the given element to promote the hydrogen entry is connected with the electrochemical stability of the respective hydride. Thus, AsHx and SbH>> are stable at both low and high pH values, while H2S and E^Se exist only within the acid range of electrode reactions. The drop in the hydrogen permeation rate at ij, is due to the change in the mechanism of the hydrogen evolution react ion, viz., passage frqm the discharge of hydronium ions below iip to the electrolytic decomposition of water mole cules above that limit. References. VK.Éiuialowski,Z.Szklarska-åmialowska, Bull.Acad.Pol.Sci. CI.Ill, 2,73 A95V; Roczniki Chem., 2J, 85 A955A 2/M.Émialowsk£, Hydrogen in Steel, Pergamon Press, Oxford, 1962, p*1lO. 3/J.P.Newman, L.L.Shreir, Corr.Sci• _% 631 71969/. 4/J.O'M. Bockris, P.K.Subramanian, Electrochim. Acta, 16, 2169 A97V. Joanna Taraszewska 237 KINETICS OF ZINC REDUCTION IN WATER-METHANOL SOLUTIONS OF NaClO^ Joanna Taraszewska Institute of Physical Chemistry, Polish Academy of Sciences, Warszawa Literature concerning the study of kinetics of metal reduction in mixed solvents is rather scarce /1,2,3,4/. The present work was undertaken to study the influence of the week adsorption of the organic substance as well as of the concentration of the electrolyte on the reac tion rate of Zn2*/Zn (Hg) • Methanol was used as an orga nic solvent, weekly adsorbed at the Hg electrode, as shown by Parsons /5/. As a salt exhibiting only insigni ficant specific adsorption NaClO. was chosen. Standard rate constants of. Zn2VZn(Hg) reaction in 0.1 to 3M NaClO, were determined in the whole composition range of H^O-CHUOH mixtures. For all the investigated solutions the capacity - potential curves were measured,, Impedance measurements were carried out using the Schering — type bridge according to Randies /6/ and the rate constants were evaluated with the method described in the paper of Bear et al /7/« The capacity values were twice inte grated to get the eleotrocapillary curves. Zero charge potentials were determined using the streaming electrode. The second integration constants - surface tension values - were determined with the capillary electrometer. As an example the charge - potential curves for solutions containing 0.1 M NaClO and various mole fractions of CH3OH are shown in Fig. 1. From the charge ~ potential curves, as well as from the surface tension data it follows, that the maximum adsorption of CHLOH occurs at ca. -650 mV in 0.1 M NaClO^ and shifts towards more negative potentials with increasing NaClO, concentration. The relative surface excesses of CELOH were calculated. In diluted NaClO, the surface exoes"s at the potential of Zn2+/Zn(Hg) reaction !•©. -lOOOmV is about the same as at the electrocapillary maximum. The standard rate constants k in pure CH^OH are smaller than in pure water at all NaClO. concentrations. The Ig k vs lg NaClO, concentration curves at several ClkOH mole fractions are shown in Fig. 2, It is seen that up to 18.4 mole % CH-OH this dependence is similar to that in pure water solutions. As is well known the effect of NaClO, concen tration in water is due to the ohange in the potential drop over the diffuse double layer. Therefore, it may be concluded that the double layer structure in the inve stigated HpO - CH3OH mixtures of low CEjOH content is similar to that in aq.ueous solutions» 238 J o anna T a r n. s z o v; s 2ca *- *i, <"'-'.« " - a Ä i • My. - • Jf I. - * - fW -l»r Fig. 1. The charge — potential curves in 0.1M NaClO, containing various mole fractions of CH^OH 1-C0!$ tap, OS^m Fig. 2. Curves of the Ig k vs lg NaClO. concentra tion at several mole fractions of CE-0H7 Joanna Taraszewsl^a " Literature 1. W« Jaenicke, P.H. Schweitzer, Z. physik. Chem. N.F0 52, 104 /1967/ 2. T. Biegler, E.R. Gonzalez, R. Parsons, Coll. 3j6» z*-^ /1971/ 3. I.O'M. Bockris, R. Parsons, Trans. Faraday Soo. 45, 916 /1949/ 4. B. Behr, J. Dojlldo, J. Stroka, Extend. Ahstr. C.I.T.C.E., 50 /1970/ 5. R. Parsons, M.A.V» Devanathan, Trans. Faraday Soc« 49, 673 /1953/ 6. TTE.B. Randies, private oommunication 7. B. Behr, J. Malyszko, Roczniki Chem. 41., 1589 /1967/ 240 Rartmut Wendt ELECTRODE KINETIC AND PREPARATIVE INVESTIGATION OF THE DIRECT ANODIC OXIDATION OF OLEFINS ON CARBON ANODES IN NON-AQUEOUS SOLUTIONS by M.Katz and H.Wendt from the Institut fttr chem.Technologie d. TH Darmstadt H.Schäfer [1] showed that by direct anodic oxidation of olefins these compounds can be coupled anodically to dimers. We decided to elucidate the reaction mechanism of the ano dic olefin oxidation in non aqueous solutions by combined electrode kinetic and preparative work. The electrode kinetic measurements were confined to anodic polarogranis and Tafe^plots of the olefins being measured at rotating carbon rod anodes in methanol. The polarograms show that all olefins (see Table l) are oxidized irreversibly in a rate determining 1-electron- transfer-step (but overall 2e-transfer for all olefins with the exception of butadiene) with a chargetransfer coeffi cient ofoc = 0.7- Table 1 Polarogr. data for the anod, olefin oxid. olefin Ej/2 " 7~J£ rT " shift of Tafel (vs.aq.SCE) $*•"*; tot.nr.of plot T«lfrff^* electrons Tos~vs.C10, I) methyl- + 1350 0.77 2 + l6o styrene £) styrene + 1^2o 0.70 2 + 80 1 1 3) ethyl- !1 t vinyl- + l46o 0.75 1.8 + 60 ether ^) buta diene + 1680 0.70 1.2 + 50 » As is show•• ••n- in figure 1'•' •fo r— styrene, the anodi. ' c wavI.. eI in pre• I I sence of sodiumperchlorate, as supporting electrolyte, is composed of an adsorption prewave and a main wave. In pre sence of sodium-tosylate there is only one wave with a i.iore positive half wave potential. Figure 2 shows the anion-influence on the position of the Tafel plots for styrene oxidation. Tosylate ions compared to perchlorate ions shift the current voltage curves re markably to':.iore anodic potentials. This shift is different for the various olefins, giving: + loo mV, + 80 mV, + 60 mV and + 50 mV for the olefins 1 to 4 (table 1) respectively. 241 Rartmut Wendt u[mV.wSCE]- So u [mv.,, SCE]- f-y ; c»od. T*flL fU+3 foV T rCHC ftfe©^ /Vd - /"»a* /a ^c The anion effect on the current voltage curve can be inter preted as follows: l) The olefin-species is oxidized in the adsorbed state (position within the IHP). This fact explains the relative- ly high charge transfer coefficient and the appearance of an anodic adsorption prewave. 2)Tosylate ions due to their aromatic character are strongly adsorbed on the anode and thus displace the only weakly ad sorbed olefins and so decrease their stationary surface con centration. However, the different olefins are displaced by tosylate ions to a different extent. Perchlorate ions how ever as spherical aniens penetrate much less into the inner Helmholtz plane. From the shift of the Tafellines a ratio c r 13°-<4erohlorate /™ > ^'v >tosylat -e - of loo,, lo, 5 and 4,5 is calculated for the olefins 1 to 4. l?he charge transfer coefficient of the olefin oxidation is not influence by the supporting electrolyte anion which has no influencevthe reaction mechanism. The supporting electro lyte anion however., has an effect on the rate of the charge transfer reaction (eq.2).For the anodic reaction we can wri te the following consecutiv.Livee steps-steps:. (l) (^CsC v )solv fc» y' ^ 'ad (2) (>c=c'v)ad-e —* v, ,ad Whereas toslyte anions onl0°y -chang Ol,e the velocity of the charge transfer, they qualitatively influence the further reaction of the radical-cation intermediate. From the composition of the products of the anodic olefin oxidation we know that two reactions paths exist for the ad sorbed radical cation: A) solvolysis B) electrophilic attack on unreacted adsorbed olefin. Solvolysis (A) with subsequent further oxidation and solvo- 242 Hartmut "tfendt lysis (ecec) leads to a monomer (1.2 substituted) oxidation product whereas electrophilic attack leads to the formation of dirners (or oligomers and polymers) equ. (4). According to this, the molar ratio of formed monomer products to the sum of thep* rldimen mpry ((annnrtd nnlvrnprpolymerl) r»T»niproducti n n t: sR will bh*e^ an funfunctioc n of the stationary olefin surface concentrationfe^. $) . Indeed, according to the ratio l00 an approximately r ( ;c=c-)clo4- /nyi-c; >tos« < m / loo fold increase in the ratio »'rnomonn/ •*• »»4 **„ is found when styrene is oxidized in presence of sodium tosylate instead of sodium perchlorate. Interesting enough with sodium-tosylate additionally unsaturated dimers are formed. Furthermore four other products are obtained. These compounds are supposed to be dimeric products formed by typical radical reactions of the primary oxidation product. Tosylate-anions not only decrease the surface concentration of the unreacted olefin but decrease even more the statio nary surface concentration of the radical cations, which are forced psxtially to desorb before being solvolized or added to an olefin. The displaced intermediates must then react in the solution forming products which cannot undergo further oxidation.and must react as radicals. The formation of unsaturated hydrocarbons in presence of tosylate anions is attributed to the proton-abstraction from the carbonium-ion-intermediates by the weak base tosylate{ ice •<|**.c€') ) . *m f) M »•» k. of s» <<4.-rnc~«) c- + HTos c-c- To* —- •-> = 6) C René Winand 243 COVT^TPUTTON TO THE STUDY OF THE MECHANISM OF THE ANODIC DISSOLUTION OF Cu2S. Patrick Brennet, Jean Vereecken and René Winand Université Libre de Bruxelles, Service Métallurnie- Electrochimie, 50, av. F.D. Roosevelt, B-1050 Bruxelles - Belgique. The mechanism of the anodic dissolution of Cu2S in sulfuric solutions was studied by electrochemical and mineralogical methods. The dissolution occurs in two steps separated by a short transition period. The first step is the formation at the surface of the anode of a layer of digenite (Cu^ QS) with liberation of Cu++ ions at low anodic potential and with a hicrh current efficiency according to the reaction +4 5 Cu2S » 5 Cul 8S + Cu - + 2 e During the transition period, the external surface of the digenite changes in a very thin laver of covelline blue-remaining Cu^ jS* Then, and simultaneously, this laver of Cu^ ]_S changes to true covelline CuS, the anodic potential rises sharply and the thickness of the layer of digenite does not increase anymore. During the last step, three reactions occur simulta neously with the same speeds, at high anodic notential : ++ at the interface Cu2S/CuifgS : 5Cu2>S—»4Cui 3S+Cu + 2 e s ++ at the interface Cui,s /CuS : 5Cux gS—**CuS + Cu + 2 e at the external surface of the anode : CuS—*Cu++ + S + 2e 244 H.Yamaoka. A.C.POLAROGRAPHIC STUDY OF D.C.POLAROGRAPHIC NON-ADDITIVE DIFFUSION CURRENT PHENOMENA AND UNDERLYING FAST HOMOGENEOUS REDOX-REACTIONS: H.Yamaoka. Chemistry Department A, Technical University of Denmark, Building 207, DK-2800, Denmark. When two (or more) electroactive species are present, the total d.c. polarographic limiting diffusion current is not always equal to the sum of the limiting diffusion currents of the individual species. It was once postulated that the non-additivity of diffusion currents can occur as a result of a very fast homogeneous redox-reaction in the diffusion layer between one electroactive species and the product resulting from the reduction of another, when the species have significantly different diffusion coefficients (ref.l).i: For the majority of polarographically reducible species and their reduction products, however, the magni tude of diffusion coefficients is of the same order under a comparable working condition. Hence the non-additive phenomena are usually more or less overshadowed by the experimental errors inherent in the conventional d.c. po larographic technique. On the other hand, scarcely any independent kinetic data were available for fast homo geneous redox-reactions in early 1950's. The subject did not attract an active interest among research workers over nearly 2 decades since then. Recently the present author has recognized the a.c. polaro graphic technique as a possible alternative method for the study of the d.c. polarographic non-additive diffusion current phenomena and the underlying fast homogeneous redox-reactions (ref.2). The figures a & b illustrate the essential point: the addition of monofluoro-pentamin- cobalt(III) does not seem to affect the d.c. wave of europium(III) (fig.a), whereas its addition obviously increases not only its own a.c. wave height but also that of europium(III) (fig.b). This latter increase in the a.c. wave height has been found to be due to a fast homogeneous :: The d.c. polarographic non-additive diffusion currents, sometimes called "latent --" or "consealed --", have different origins. Various examples are illustrated in standard textbooks of polarography. It is unfortunate that the usage "non-additive diffusion current" has by now been almost settled in the literature in spite of its ambiguity. The present work deals exclusively with such as defined in ref.1. K.Yc-mcioka 245 IV 16 / £ 3 < =1 2- C Q) i_ -••••T •> t— .JJU1* 3 O O 0, _1 L _1 I I 1_ J L < 0 -.2 -.4 -.6 -.8 -1.0 -12 D.C. Potential E vs. SCE (Volts) Fig. D.c.(a) & a.c.(b) polarograms of europium(III) ion in the presence of monofluoro-pentammin-cobalt(III) ion. 0 0.5 M perchloric acid + 0.5Msodium perchlorate 1 0 + 1.0 0 mM europium(III) ion II 1 + 0.95 mM monofluoro-pentammin-cobalt(III) ion III I + 1.8 5 mM do. IV I + 2.70 mM do. a.c. voltage amplitude 10 mV r.m.s. ; a.c, frequency 50 Hz; 25 C ; nitrogen-satd. H.Yarnaoka 246 redox-reaction in the diffusion layer between the in coming monofluoro-pentammin-cobalt(III) ions and the out going, electrogenerated europiumClDions . Analogous examples have been found for a series of fast second-order homogeneous redox-reactions of known origin (see detail in ref.2). Slow homogeneous redox-reactions have been found to give rise to no detectable increase in the a..c. wave height. On basis of the limited number of pilot systems tested, a value of at least 1,000 liter/mole per sec for the second-order rate constant seems required to detect an appreciable increase in the a.c. wave height under an ordinary working condition with d.m.e. This threshold value probably depends on the diffusion layer thickness (drop life time on a d.m.e., surface velocity on a streaming mercury electrode or rotation speed on a ro tating disk electrode) . It should again be emphasized that the d.c. polarographic non-additive diffusion current phenomena as treated here deal with two (or more) electroactive species in the po tential range of interest. This is in contrast with the d.c. polarographic catalytic waves which deal with one electroactive species in the presence of another which however is electroinactive if it were alone. Thus, the a.c. polarographic technique can be said to give infor mations about a more general type of fast second-order homogeneous redox-reactions than the conventional d.c. technique does. REFERENCES 1. S.L. Miller and E.F.Orlemann, J.Am.Chem.Soc., 75 (1953) 2001. 2. H.Yamaoka, J.Electroanal.Chem., 36 (1972) 457# Vera Zutic 247 THE EFFECT OF CHLORIDE IONS ON THE ELECTROCHEMICAL BEHAVIOUR OF NICKEL AT MERCURY ELECTRODES Jean Chevalet and Vera Zutic Laboratory of Electrochemistry, Faculty of Science, Paris, France, and Center for Marine Research, "Rudjer Boskovic" Institute, Zagreb, Croatia, Yugoslavia. The behaviour of Ni(ll) and Ni(O) has been studied in aqueous solutions of LiCl and LiN03 at chloride concentrations up to 14N, by using combined techniques (d, c. polarography, double potential step chronocoulometry^ ' and chronopotentiometry). Cathodic behaviour The kinetics of the reduction process is controlled by the chloride and free water contents. Therefore the series of experiments were performed at constant water activity (LiCl + LiNO_ = 10 M) or constant chloride concen trations. At high LiCl concentrations (10 M) a quasi-reversible wave is obtained with E. /_ = -0.47 V/SCE. This corresponds to the bielectronic. reduction^ the second electron transfer being the slow step \ with the apparent transfer coefficient of about 0,2 (Fig. 1.). With decreasing temperature the limiting current assumes partxy kinetic character. The effect of tetraalkyl ammonium, salts indicates that the reducing species is positively charged. With decreasing chloride content at ionis strength 10 the first wave becomes more irreversible and the total limiting current assumes kinetic character. This is indicated by the decreasing ratios of charge transfered at 0,04 and 4 sec, and the values of diffusion coefficients determined chronopotentiometrically^ ' (D = 2.3 and 3. 6 x 10"6 cm sec"1 in 10 M LiCl and 10 M LiNO^ respectively). These facts are best explained by assuming that the chemical reaction 2+ + [Ni(H20)6] + Cl"===[Ni(H20)6_nCl] + n ^O b preceeds the electroreduction proper. The first wave, when transfered to log\ - E coordinates and corrected for the parallel Ni(ll) aquocomplex reduction^ ' assumes limiting O slope and exhibits a chloride concentration order close to 1. Introduction of the stoichiometric equilibrium constant [NiCl J/ [Ni2+] [ci"J = 0.094, reported by Florence^5) leads to the k, and k, values of 200 1 mol sec"1 and 2100 sec" .respectively. 248 Vera Zutic Anodic behaviour At Ni(ll) concentrations below lO'-^M the double potential step experiments indicate that the reoxidation process is diffusion controlled. The recuperated charge values, up to 20% higher than theoretically predicted for the soluble redox couple^ ', were also observed for the amalgams of higher solubility (Cd, Zn, Tl, ). No specific adsorption of Ni(0) or Ni(ll) could be detected by double potential step chronocoulometry (6). When the Ni(ll) concentration exceeds 10" M, the reoxidation process is no longer diffusion controlled and depends on the potential selected for eiectrogenerating Ni(0), irrespective the composition of the solution. These results are interpreted by taking into account super saturation and powder formation. Reoxidation of cathodically generated Ni(0) gives Tafel slopes (Fig, 2.) which are characterized by an apparent °^ of about o. 4, very slight water effect, and a reaction order with respect to the chloride bulk concentration close to 1. This is explained by assuming that the rate determining step consists in the chloride assisted electron abstraction: Ni(Hg) + Cl" - e—-NiCl followed by fast oxidation to Ni(II). The exact values of the activities of water and chloride and of the amount of the chloride specific adsorption have been found to be essential in order to assess correctly r.he kinetic parame ters involved. References: 1) F. Kimrnerle, J. Chevalet, J. Electroanal. Chem., ZJ^ (1969) 237J 25 (1970) 275. 2) R. M. Hurd, J. Electrochem. Soc., 109 (1962) 327. 3) L. Gierst, Z. Electrochem. , _5j9 (1955) 784. 4) J. Dandoy, L,„Gierst, J. Electroanal. Chem. , 2_ (1961) 116. 5) T.M. Florence, Aust. J. Chem. , 1_9 (1966) 1343. 6) F. C, Anson, Anal. Chem. , _3_å (1966) 54. Vera Zutic 249 fOi'iuCl 3Q°C ^ -/ <0 ri> -3 4-JUL E/SCE -0.2 Kg•/. 'Co LU'Cth • % 5 +/ ± 3.5 v as 0 Q.25 \ \ ©S v, 0 + ^ \VVsN oov- >§ -/ ^\ -2- X 0\ J -4 -0,2. -0,4 F/gZ. B/SCB 250 Vera Zutic REDOX PROCESSES OF URANIUM(VI) PEROXO COMPLEXES IN ALKALINE HYDROXIDE SOLUTIONS Vera Zutic and Marko Branica Center for Marine Research, "Rudjer Boskovic" Institute, Zagreb, Yugoslavia Reduction processes of uranium(VI) peroxo complexes in aqueous alkaline hydroxide solutions have been studied using d. c. polarography, cyclic voltammetry and double potential step chronocoulometry at the mercury electrodes. The investigation has been limited to solutions of high pH (1-2 M LiOH), where, according to the available data'- ' stable monomeric uranyl peroxo species are the predominant forms. Pepending on the rLO./tJ ratio uranium(Vl) hydroxo (4). mono peroxo and triperoxo complexes have been characterized as electroactive species (Fig, 1.). Reduction of the uranium(VI) monoperoxo comples gives two waves -a quasi-reversible wave at -1,05 V/SCE and an irreversible one around -1.5 V, (Fig. 1. Curves 2-4). On the other hand, uranium (VI) triperoxo complex is reduced through a single seven-electron irreversible process E, /_ = -1.45 V. (Fig, 1. curve 5). In both cases the same uranium(V) hydroxocomplea has been identified as the stable reduction product. It can be concluded that the first step of the overall process is the one electron reduction of the uranyl group, the potential of which is determined by the stabili ty of uranyl-ligand bond. The following step shows the usual characteristics presented by the noncomplexed peroxo group. The main paths of these various electrode processes and the coupled chemical reactions can be sketched as follows: VI 1 __ V — ~ U II TI 2 ! n (o2 )i/(o2 - ne oC E1 /2/S( ^j X-J V 1) 1 0.49 -0.88 » V">3 2) 1 0.35* -1.05 3) 2*0.2 * -1. 5 4) 7 0.35* -1.45 apparent values Vera Zutic 251 The behaviour is complicated by the second order bulk redox reactions between U and any species containing the peroxo group. At the potentials around -0.1 - -0.3 V the uranyl peroxo com plexes produce a cathodic peak, which is caused by the direct reduction of the peroxo ligand of specifically adsorbed complex. In the presence of an excess of hydrogen peroxide this maximum is enhanced by catalytic regeneration of the peroxo species (Fig. 2. curve 2). References: 1) I.I. Chernyaev, Kompleksnie soedinenya urana, Nauka, Moscow 1964. 2) A.M. Gurevich, Radiokhimiya, l_ (1961) 321. 3) N. W. Alcock, J.Chem.Soc, (A.) 1968, 1588. 4) V. Zutic, M. Branica, J. Electroanal. Chem. , 2<8 (1970) 187. 1 mM U (VI) 1M Li OH HjOj/UtVl) u o lis, 2» 0*1 ' 3)1 t)2 —«"•*•** 5) V .jtirijn'S"i *" 0-, f\_s _y/r KI I. I I I I I I. I < -02 -(U -OS -0« -ID -12 -1* -W -llVn. SCE Fig. 1. Typical polarograms of the solutions 1 mM uranium(Vl)- - 2M LiOH - H2Q2 10°C, t=3 sec, var. 252 Vera Zutic 1 M UOM XPC 1) In* UIVIIOrWHjO, 21 ImH UWll-SmHHjO, J) ImM HjO, pH» Mam I tit, l , 1 _> ***%. -HI -03 -«i -«5 - 4- Fig. 2. Polarograms of the solutions of 1 mM [U02(0 )"] (1), 1 mM[UO(0,) 1+3 mM H2Oz (2), and 33 mM 3 H2Oz (3); 10°C,, t' A' 3 sec. SECTION 3 EXPERIMENTAL METHODS IN ELECTROCHEMISTRY Isra'el Epelboin 255 RECENT PROGRESS OF IMPEDANCE MEASUREMENTS AT VERY LOW FREQUENCIES Israel Epelboin, Claude Gabrielli, Michel Keddam, Hisasi Takenouti Physique des Liquides et Electrochimie, Groupe de Recherche du CNRS, associé ä 1'Université Paris VI, 11 quai St-Bernard, Paris 5e, France Since a long time theoretical investigations have predicted that it vas possible to obtain information on processes of mass transfer and adsorption at the electrode-electrolyte interface, by relating these processes to a so-called faradaic impedance [_1,2^ . However, impedance measurements have only provided few data on the mechanisms of electrochemical phenomena. Except for the Warburg impedance which was revealed long ago, the main use of alternating currents was restricted to the determining of the electrochemical double-layer capacitance. It is well known that data from solid electrodes are difficult to inter pret, particularly those concerning the dispersion of this capacitance [3J . Recently, impedances of electrochemical systems have been quantitatively related to mass transfer processes or reaction mecha nisms described by adequate models [4_j . This was mainly due to the fact that the measurement range was possibly broadened towards very low frequencies [4,5J. Such measurements are now easier thanks to certain novel experimental devices . In this lecture, we shall specify the origins of faradaic impedances and present the progress achieved in respect to techniques as well as certain characteristic results. Faradaic impedances. The current I through the interface is known to depend on potential V on account of activated processes (charge transfer) which are represented by Tafel 's exponential law. Current I also depends on the instant values of the concentrations (c) of species (adsorbed or in solution). The surface concentration is usually expressed in terms of a coverage coefficient 6. Consequently, I = f(V, Ci, 6i). (O The steady-state values of c^ and 0j_ are functions of the potential so that the I(V) law is more complex that Tafel's law. The dynamic behaviour of c^ and 8i values is controlled by equations accounting for mass conservation. These equations are obtained from the kinetic equations of mass transfer, chemical and electrochemical reactions as well as adsorption. Consequently, these laws are of the form : dc. d0. -r~ = D.Vc. +2k. c (2), j~ =EK 0 (3) dt ii ii at i i The variation rate of these parameters is hence finite and is directly related to the chemical (k^) and electrochemical (K^) rate constants as well as to the diffusion coefficients Dj_. Under linear conditions, i.e. at very small potential variations near a polarization point,we can introduce the concepts of time-constant and impedance. If the variation frequency of the potential is sufficiently high, C£ and 0i become frozen, and only charge transfer processes will then instantly vary with the potential. In that case, impedance is equal to a pure resistance, called charge transfer resistance. On the other hand, when the period of the sinusoidal signal is of the same magnitude than the time-constants corresponding to c£ and 0£, certain relaxation 256 Israel Epelboin phenomena appear, thus assigning a reactive behaviour to the impedance. Lastly, at sufficiently low frequencies, c^ and G j_ will simultaneously vary with the potential ; impedance is therefore equal to the slope of the I = f(V) steady-state curve. The response of the system to frequency variations is related to the existence of the various terms implied in the differential of eq. (1) hence : dl (4) dV vdV c ^dc e dv + lÖ9 c dV Impedance is function of the pulsation oj of the measurement signal. The function Z(jw) is calculated by linearizing equations (l)-(A). The faradaic impedance thus obtained accounts for all the processes involved at the interface. Under usual conditions (concentration, temperature) the above- mentioned time-constants are often greater than a second. Since the usual devices have been confined to acoustic frequencies, they cannot be used in relaxation studies. This explains why experimental data on diffusion and adsorption processes are so rare. Experimental. Impedance measurements at sub-acoustic frequencies are of high interest within other research fields, hence the progress achieved in constructing very elaborated electronic measurement devices[6] . When measurements are performed with sinusoidal signals, these devices are only limited to Transfer Function Analysers. We shall merely mention here the devices of this type we built in our laboratory for purely electrochemical purposes [7] . We shall describe in detail (see figure) a set-up involving a digital transfer function analyser TFA (Solartron JM 1600 and HF extension JX1639), which has been commercially available of late years. T. F. A. With such analysers the complex CORRELATOR AI AMPLIFIER quotient generator output signal/ correlator input signal, can be determined on the basis of a __ AV .POTENTIOSTAT GENERATOR correlation principle between 10 Hz and 160 kHz. However, in order to measure the impedance at the inter / face with this device, we had to h construct a specially designed ~*~—' potentiostat[8] . This potentiostat superimposes, without phase shift, the measuring alternating potential AV from the TFA generator on the direct potential between working electrode E and reference electrode R. The sinusoidal current ÅI is amplified without phase shift, then connected to the correlator. The impedance is directly deduced from the gain of the amplifier and the transfer function displayed by TFA. Applications. We studied the kinetics of mass transfer, in the case of fast reactions such as polarographic reactions (Io + 2e -> 31, K4 Fe(CN)6 -> K3 Fe(CN)fi). The uniformity of the diffusion layer thickness was obtained by the.use of a rotating disc electrode. On the basis of already published techniques L5J » we showed that the Nernst Israel Epelboin 257 layer model can satisfactorily account for the impedance of the diffusion layer. The above-depicted set-up allows more accurate measurements. The influence of convection on diffusion hence reveaLed was shown to have a relative value less than some hundredths^]. The above-mentioned model has subsequently been applied to the study of diffusion during processes partially limited by heterogeneous, reactions, such as Cu++ ion diffusion in copper depositionL'-OJ . Our results deal with the mechanism of multi-step reactions with adsorbed intermediates as well as the influence of adsorption during passivation and inhibition processes. Analysis of the faradaic impedance in terms of frequency, potential and pH showed that the dissolution of iron occurs in two stemps coupled by the adsorption of as a (FeOH)ac[s [4J . However, the role of (FeOH)ac[s catalyst was not possibly established, in opposition to what is still affirmed in the literature. The presence of (NiOH)a(js as an intermediate during the electrochemical and chemical deposition of nickel, has also been evidenced. Impedance measurements at very low frequencies allows to determine the mechanisms of certain passivation and inhibition processes. Coverage relaxations to which are referred these processes have been analysed. Considering corrosion as a mixed process, we theoretically and experimentally showed that corrosion rate is more accurately correlated to transfer resistance than to polarization resistance jj lj. Despite these results, an effort must henceforth be made within the field of interpreting by improving the models used, especially in the case of adsorption. As a matter of fact, certain theoretically predicted impedances could not be observed (hydrogen evolution), and, in other cases, the experimental faradaic impedance could not be accounted for by the usually accepted mechanism (silver electrode). ..iterature "~11 W. Warburg: Ann. Physik, 6_, 125,(1901). "X H. Gerischer and W. Mehl: Z. Elektrochem. 5_9, 1-49,(1955). 'X R- de Levie: Electroanal. Chem. 25_, 257,(1970). V I. Epelboin and M. Keddam: J. Electrochem. Soc. 118, 1052,(1970). 5 M.L. Boyer, I. Epelboin and M. Keddam: Electrochim. Acta JJ_, 221 , (1966). [6] T. Andre-Talamon, C. Deslouis, C. Gabrielli and J.C. Lestrade: Submitted to Electroanal. Chem. 7. T. Andre-Talamon and C. Gabrielli: Report ANVAR, n°47 95. '8 I. Epelboin, C. Gabrielli, J.C. Lestrade: Rev. Gen. Elec. 7^, 669, (1970). [9] C. Deslouis, I. Epelboin, M. Keddam and J.C. Lestrade: J.Electroan. Chem. ^8, 57,(1970) . [lO] I. Epelboin, F. Lenoir and R. Wiart'. J. Crystal Growth 13/14, 417, (1972). [ll] I. Epelboin, M. Keddam and H. TakenoutiL J. Appl. Electrochem. (1972) to be published. 258 Luigi Campanella POLAROGRAPHIC BEHAVIOUR OF ISOCIINCIIOHERGNIC ACID Luigi Campanella,Pierluigi Cignini and Giorgio De Angelis Institute of Analytical Chemistry ,University,Rome,Italy The polarographic behaviour of heterocyclic compounds containing an N atom in their molecule is quite complicated because in addition to normal waves these compounds can yield also catalytic ones,general ly ascribed to the hydrogen discharge.The shape and the height of these waves are largely varying according to the experimental con ditions »particularly those referring to pH,ionic strength,concentra tion,scanning rate of potential,presence and nature of the buffer system*A research is being, performed in our laboratory on the che mical and electrochemical behaviour of pyridinecarboxylic acids,aim* ing at getting useful information about'the equilibria in solution among the species of the acids and about the electrochemical activi ty of some of these species in comparison with others» We refer here particularly to the polarographic and oscillopolaro — graphic behaviour of 2,5-pyridinedicarboxylic acid (isocinchomero- ni c smelting point 238-237 »C,dissociation constants 4.95 10 and 8.77 10 ).The data reported in literature are scarce and only li mited to the analytical aspects of the problem»Particularly Volke and his coworkers 9 determined isocinchomeronic acid hy the po larographic method in a mixture containing several pyri dinemono- and di-carboxylic acids». The study was performed by us atiJL* 1 M (NaCIO ), T«25.0 +0.5 *C, C«10~ M in presence of 0*02 If citric buffer.The field of the very high acidities was also investigated,» this case fixing the ionic strength at 211 (lICIO . + NaCIO J. 4 4 Resulta The polarographic behaviour of isocinchomeronic acid is characteri zed by the presence of three waves,differently evidenced depending on pH.Only for pH{3 all' the three waves can be o 1 »served in the sa me polarographic curve.The first wave (three electron,reversible process) is present in the pH range l-3;its E y0 is a linear func tion of pH with slope 20 mV/pH and 60 mV/pH respectively for pll^l.2 and pH^l«2'|its height is decreasing in the pH interval 1—2 and: falls down to zero for pH « 3.The second wave (one electron) can be recorded in the pH range 1-7,where its height is firstly decreasing, then increasing,constant and decreasing again and its E_ y is linear ly depending on pll (A EL /_/^pll * 60 mV/pH for l{pH ^4 and - -120 mV/pH for 4^pH ^'7) »The existence field of the third wave Luigi Campanella 259 is pH l-6;the limiting current is decreasing as pH iacreases;E behaves as for the second wave,the breaking point in the E . wfe pH curves being in this case at pH«3.In alkaline solutions no of the three waves appears .The influence of the concentration on the limiting currents obeys generally to the Ilkovieh equation,the on ly irregularities being observed for pH^4 and concentration «J *£ 1*4 10 ..The Kheifets slope of the three waves (influence of the mercury column height on the limiting currents) results to be 0.75, O.JO and 0.80 respectively for the first,second and third wave.The temperature coefficient of the limiting current is 2*8 ,33 and 48 % for the three waves,that one of E , 0.3 ^ for the first wave and zero for the tiro o there. In very acid solutions the limiting current values are increasing with the acidity degree,the EL / of the first wave is firstly shifted to more negative values,then in the opposite sign,that one of the second wave is firstly constant then shifted to more negative values,the E . of the third wave is firstly constant,then becomes more positive; Oscillopolarographic; analysis yielded results well agreeing with the polarographic ones.So the oscillopolarographie peaks correspon ding to the three polarographic waves can be recorded;their peak currents are depending on; pH with a behaviour similar to that one of the polarographic limiting currents with the only exception of the third peak for solution at pH \ 6.The peak potential values are linear function of pll,with breaking points corresponding to varia tions of the slope's the corresponding EL . values »In solution at pH £ 4 the influence of the concentration on the peak currents obejs the Randl es-Sevcik equation (ensuring the proportionality between the concentration of the electro chemically active species and the peak current) and the peak potentials become more negative with in creasing concentration*At pH\4 the just said proportionality is still observed only for the second and the third peak,the potentials of which result not to depend on the con c en t ration* In very acid so lution the peak currents of the first and third peaks are proportio nal to the hydrogen ion concentration ,that one of the second peak decreases as the acidity degree of the solution is increased.The peak potentials of the first and second peak are shifted to more ne gative values as hydrogen ion is increased,that one of the third is unaffected.The scanning rate of the potential influences the peak current and potential,the latter according to log E -a «• b log V iE"m peak potential, V »scanning rate),where a and b are depending on pH and which peak is considered* 260 Luigi Campanella By the analysis of the described results it can be deduced that among the species under which isocinchomeronic acid can exist in aqueous solution the bivalent anionic completely deprotonated spe cies is the only one for which any polarographic activity can be ex eluded. As for pyridinemonocarboxylic acids also for isocincho meronic acid the polarographic behaviour is characterized by the high electrochemical activity of its'protonated cationic species which is reduced at the dropping mercury electrode either directly or through a kinetic process of equilibrium with the other species that are transformed into the cationic species at the electrode so lution interface,The rate of this transformation plays an extreme ly important role t indeed,if it is sufficiently high the pro tonat- ed cationic species can be emphasized in a wider range of pH than expected based on values of the equilibria constants among the spe cies of acids »On the other hand the waves and peaks assume,at least partial IT ,a kinetic natur-e.On the basis of the number of electrons involved in the electrode processes,of the^E /_/*^pH slopes,of the dependence of the height of waves and peaks-on the experimen tal conditions,particularly those referring to pH and temperature and of the value of the power of scanning rate in the Randles-Sev- cik equation ( i • k v , i -peak current) and of the results of different other experiments aiming at investigating how waves and peaks and u.v. spectra of the examined solutions are modified by a prolonged electrolysis,the following processes can be postulatedt H c H 1st) H*PC +H +3e —*l/2 2 + !? < 2) O^pH^l.l + HgPC + 3 H + 3 e-*l/2 H* + H*PC (Hg) l.l^pH,$2,l 2nd) H+PC + 1 e ~-+> l/2 H + H PC O^pH^l.l H?2 + H+«* Ht?C -^ •*/* Ho + VC 1«1*|H £4*6 HPC + 2 H ^iTPC JL»/ 1/2 H* • fi^PC 4.6{pH^7.0 3rd) H*PC + 1 e -*l/2 h* * H PC O^pH^l.l H~PC(zwitt) + H ^TH*PC «S^ 1/2 H '• H PC 1.1 ^pH <3,0 KPC"(zwitt) + 2^ H^HJ»C -2-> 1/2 H2 +' H PC 3.0$ pH^S.l where H PC,fiLPC,HPC are the eoniugated acids ot the base PC (bi valent anion pyridinedicarboxylato) jjswitt means zwitterion form and (H ) indicates a species reduced in the ring (di hydro derivative). Sm A*>out the nature of the processes through different experiments it can be concluded that all these are not only diffusion con tro 1- ledybeing evidenced an adsorption component for the first wave and a kinetic nature for the second and third. References t l)J.Volke and V.VolkovayColl.Czech.Chem.Comm. 20,1332 7l955); ^Tj.Volke and V.Volkova,Coll.Czech.Chem.Corara.,»0,908 (1955) 3)L«Campanella and G.Oe Angelis,Rev.Roumaine Chim.,16,545 (l97l) Dun ja Cukman 261 EXPERIMENTAL STUDIES OF COMPLEX REACTION MECHANISMS OF URANIUM IN ACIDIC MEDIA Dunja Cukman M. Vukovic, and V. Pravdic Laboratory of Electrochemistry, "Rudjer Boskovic" Institute, Zagreb, Yugoslavia Cyclic chronopotentiometry (CCP) has been proven as a sui table method in studying complex electrochemical-chemical reaction mechanisms. It is useful as a diagnostic technique for distinguishing between various possible reaction paths. Combined with numerical prediction based on digital simulation of various mechanisms, CCP is capable of accurate determination of rate coefficients of chemical reactions coupled to electron transfer. Studies of the uranium(Vl) - (V) - (IV) system in acidic media have been quite frequent. It has been taken as confirmed that the reduction of uranium(VI) to uranium(V) is followed by the disproportiönation reaction to uranium(lV) and (VI). Most studies have assumed the simple scheme and there was only some discrepancy as to the actual order of the reaction of disproporti- onation with respect to hydrogen ions (in acidic media) or some complexing ion, like carbonate (in alkaline solutions). CCP which tends to reveal deviations from simple mechanisms in subsequent reversals, has shed also some doubt on the reaction scheme of the uranium system. Fig. 1. shows a CCP-gram for one of the systems studied: a 3mM U(VI) solution in 2M NaCIO at pH=0. 87. The third relative transition time should be theoretically larger than the second, but experimentally the opposite was found. Quantitative data are shown in Table I. Deviations are beyond any experimental error. Using digital simulation a search was made for a plausible mechanism which would follow the experi mental pattern of a_ a_. Table II is a listing of a number of a few mechanisms with ap propriate model rate coefficients. None of the entries satisfies the experimental findings. The search for the alternative to the simple disproportionation has been continued proposing an E-C-E scheme. Fig. 2. has been constructed on the basis of model rate coefficients which yield the same second relative transition time a?. Full lines represent the even cycles for disproportionation, dashed thos*e for an ECE mechanism. Taking arbitrarily that experimentally a difference of 5% is required för discrimination, one can see that, except fbr the extremely slow and the extreme ly fast coupled reactions, there is quite a difference between these models. Quantitative interpretation escapes at present the capabilities of the CCP technique. It remains certain, that in highly acidic, high ionic strength media ion-pair formation yields 262 Dunj a Cukman reaction paths which, besides disproportionation of the uranium(V) species, involve at least one E-C-E mechanism. There is the possibility that some uranium(V) species is specifically adsorbed at the mercury electrode. 1 mH U(VI] E VS. SC.C. * 3 • • """B" M n«.i NUMBER OF CYCLES Fi1 * Fig. 1. CCP-gram for 3 mM U(VI) in 2M NaC104,. pH=0. 87, (HC104). Cathodic direction is to the right of figure. Fig. 2. Relative transition times for even (anodic) cycles for the disproportionation reaction (full line) and for the E-C-E mecha nism (dashed line). Dunja Cukman 263 Table I. Relative transition times for the reduction-oxidation reaction of 2 mM U(VI) in 2M NaC104>t different pH' s (HCLO4). Theoretical values are for the simple s&lfcproporti önation model. pH 1.0 2.0 3.0 a2 exp. 0.610 0.623 0. 666 a_ theor. 0.611 0.629 0. 667 a exp. 0.586 0.593 0.648 a_ theor. 0.648 0.654 0.667 Table II. Some possible model mechanisms. Relative transition times a_ and a, shown for the indicated model rate coefficients. Model Relative transition Model rate coefficients times O + e *R 0.258 0.415 k •= 7 x 10"3 2R-£-%0 + Z 0. 329 0.479 k = 2.7 x 10 k 0.271 0.421 k= 7xl0-3 kj = 3. 5x10-3 2R=±0 + Z 0.324 0.448 k=7xl0"4 k =8xl0-3 kl x 2R— O. + Z 0.226 0.275 k=7xl0"3 k^S.SxlO"2 kl O + R- —*E 3 2 0.279 0.328 k=7xl0" k =lxl0" kx> k References S.W. Feldberg, Anal. Chem., 16 (1964) 505 H. Imai, Bull. Chem. Soc. Jap. , 30 (1957) 873 H.N. Blount and H. B. Herman, J. Electrochem. Soc., 117 (1970) 504 H. B. Herman and H.N. Blount, J. Elect roanal. Chem., 215 (1970) 165 ^ Michel Daguenet 264 THEORY MB APPLICATIONS OP MICROELECTRODES P. ASmeur, H. Daguenet . P. Kerraiche, 11, Meklati. Laboratoire de Cinétique Physique et d'Electrochimie, Centre Universitaire de Perpignan, 66, Prance et Laboratoire d'Electrochiioie, Departement de 1 Chimie, Faculté des Sciences d Alger; 2 rue Didouche Mour^d, Alger, Algerief 1 - THE FJCROELiiCTRODES ON INERT SURFACES. Let us consider any flow on any inert surface S. Let's delimite on any point of S a circular micoelectrode whose diameter is Tery small in front of the characteristical length of the surface, but tall in front of the thickness of the boundary layer of diffusion. Calculation and experimentation show that the limiting flux of diffusion on the surface is proportional to the shear velocity elevated to the 2/3 power (l to 5). 2 - TEE DIiaCR0EL3dTR0DSS ON INERT SURFACES. Let us consider a first micro- electrodé EL and let's put besides it, in its wake, a second microelectrode JB . Let us suppose there is on E., the electrochemical reaction A —> I While on E_ the electrochemical reaction I -—> B. That Diinicroelectrode is similar to the so-called Ring-Disc Electrode. An approached calculation shows (6) that the proportion of the intermediary I which reaches E is given by a formula similary to formula obtained by Albery as regards the Ring^-Disc ALectrode (7). 3 - THE QUASI-UHIFORMLY ACCESSIBLE KICRQEL2CTR0DES. The normal velocity to the surface can be translated by: V& - 4 - THE PSiWDD-mttPOaHiY ACCESSIBLE iJLSCTRODflS. Let us consider a flo» on a electrochemical active surface fi. Let's deliirdte at a given point of E, by a thin.isolating sheath, a small part of surface so as to this part can constituted a electrocaly independent (of E) microelectrode (e). If the .isolating sheath is thin enough, we can consider in first approximation, /that the laicroelectrode is uniformly accesrible. }!?nce, itc pjrr.-.ts to ::c;_auré local vrJue of the diffusional flux density on a surface (l*). Michel Daguenet 265 5 - THE SPH5SXCAL MICROZL^GTRODES, The limiting flux on a portion of a sphere limited by a angle of Ilu* with regard to the direction of incidence (as concerns a superior angle, the boundary layer rises from the surface) is proportionnal to Up where U is the fluid velocity supposed constant outside the boundary layer. If the radius R is small enough, U represents the local mean value of the fluid velocity. After previous calibrating, the measure of the limiting flux gives the value of TJ (l3), 6 - Hence, these methods of measures permit to get local values of shear velocity, of flow velocity, of diffusional flux. For instance, in using these methods, we could measure the local shear velocity on a rough rotating disc and explain the diffusional flux variation on the surface (14 to 16). 7 - We can use these methods in conjunction with the electrochemical visualising method of parietal streamlines of a fluid (l7) and the whole can give informations about Fluids Eechanics, Chemical lsngTrieering, jslectrochemistry. (I) L.P. Reiss, J.J. Hanratty, A.I.Ch.K,, n* 2, 9, pp 154-60 (1963) (<=) K. Daguenet, F, Aouanouk, G. Cognet, CRAS t ^71 p 328 (1970) (3) M. Daguenet, F, Aouanouk, J. Chim. Phys, 67, n* 11-12 p 1956 (1970) (4) F, Aouanouk, M, .Daguenet, J, Chim. Phys. 67, n« 11-1^ p 1959 (1970) (5) F. Aouanouk, M. Daguenet, ^lectrochem. Acta, sous presses (6) F. Kermiche-Aouanouk, K. Daguenet,' a paraitre. (7) W.J. Alberyi H.L. Hitchman, Ring-Disc Electrodes, Oxford Sci. Res. papers (l97l) (8) V.G. Levitch, Physicochemical Hydrodynamics, Prentice-Hall (1962)- (9) Schlichting, Boundary Layer Theory,ftc Graw Hill (1967) (10) M. Lebouché*, These Doct. Nancy (1968) (II) F. Almeur, Rapport DEA, Alger, inedit (l97l) (12) M. Daguenet,-F. Almeur, J..Chim, Phys. sous presses (13) Bousgarbies, These Doct. Poitiers (l97l) (14) M. Daguenet, H. Meklati, G. Cognet, CRAS t 27<* p 1355 (l97l) . (15) K. Mekiati, H. Daguenet, CRAS, t ^< p 20^7 (l97l) (16) M, i^leklati, M, Daguenet, G. Cognet, a paraitre (17) 'Vim Daguenet, J.L. Peube, CRAS, t cl* p 351 (l97l) 266 ?.:;. Dr.?.2±6 ADSORPTION OF ORGANICS ON NON-NOBLE METALS P.M.Dragid and N.R.Tomov Institute for Chemistry,Technology and Metallurgy,Beograd and Faculty of Technology and Metallurgy,University of Beograd, Yugoslavia A newly developed technique for measuring adsorption of organic substances on non-noble metal electrodes will be described. The main feature of the technique is the use of a cylin drical cell (3) with a movable piston (2) inside the cylinder, as shown in fig.l. Vacuum deposited about one micron thick metal layer on the glass (1) is serving as the test electrode. When the piston is in the middle po sition and the valve in the center of the piston open,the electrode can be equilibrated with the bulk electrolyte containing up to 10"* M of organics at any desired poten tial, by connecting test electrode (1), counter electrode (4) and reference electrode (8) to a potentiostat. After equilibration the piston is positioned at 10 micron distance from the metal surface, what reduces volume of the adjecent electrolyte to ca.l0~*cm3 per cm2 of the electrode surface. By this operation the quantity of non- adsorbed organic in the solution becomes smaller or even negligible (for 10~4M solution 10"4 cm x 10"7mol cm~* « = io"'° molcm"1) compared to the adsorbed quantity of organic at the electrode surface (about 10~9molcm~*)• Then, the central valve is closed,! cm3 of organic free electro lyte sucked in through (7), thin metal layer dissolved by anodic stripping and the total amount of organics in samples taken through (7) (previously adsorbed at electrode surface + non-adsorbed from the remaining solution) deter mined by gas-chromatography. In a blank test without the electrode the remaining non-adosrbed amount can be sepa rately determined. By repeating the same procedure at different potentials and bulk concentrations the adsorption isotherms at dif ferent potentials can be obtained. On the basis of experimental data precision and limitations of the technique will be discussed. DJI. Dra5ié 267 -9 „l Fig.l Cylindrical cell for adsorption measurements l.Test electrode 5.Solution inlet 2.Teflon piston 6.Gaskets 3.Glass cell wall 7.Inlet and outlet tubes 4.Counter electrode 8.Reference electrode capilary 268 Israel Epelboi.n EVIDENCE OF THE PASSIVITY OF COBALT IN SULPHURIC ACID MEDIUM Israel Epelboin, Claude Gabrielli, Philippe Morel Physique des Liquides et Electrochimie, Groupe de Recherche du CNRS, associé ä 1'Université Paris VI, 11 quai St-Bernard, Paris 5e, France Although the mechanisms of the dissolution of cobalt and iron are known to be similar (1), the passivity of cobalt is still matter for discussion since it' "cannot be detected by simple potentiostatic current-voltage curves (2). In this paper, we shall show that the passivity of cobalt in sulphuric acid medium can however be detected by means of a regulating device with includes a negative resistance (3). The experimental apparatus (see Figure) includes a potentiostat P and an adjustable negative resistance (-R), synthesized from an active device which is called in electronics a negative impedance converter (4). Resistance -R is connected in series with the working electrode W. This device makes it possible to control the voltage between the reference electrode (Ref) and ground by imposing a load line of equation V = E + RI (L, in the figure). Consequently, although several values of the current at a given potential (Z-shaped curve (4) ) can be theoretically determined from the current-voltage characteristic f(I,V) = 0, it is possible to control the electrode polarization within the range where the potentiostat alone is ineffective. In order to obtain such a current-voltage curve, we must simultaneously adjust the voltage E and resistance -R. The current-voltage curve is then directly recorded by an XY tracer. The abscissae represent the poten tial difference between the working electrode (input Xj) and the reference electrode (X2). Values on the Y axis are equal to Rjl (I is the current through the cell and Rj a known resistance). As an example of the numerous possible uses of this device, we shall describe here a study of the passivity of cobalt. Since in this case the current density is high, the influence of mass transfer cannot be neglected. In order to obtain well defined hydrodynamic conditions, we used a rotating disc electrode. This electrode consists of the cross-section of a polycrystalline cobalt rod (Johnson-Matthey, 5 mm of diameter) whose lateral surface is isolated by an acrylic resin. The electrolyte, a molar sulphuric acid aqueous solution, was de-oxygenated by argon-bubbling and maintened at a temperature of 25 ± 0.1°C (H2SO4 : Merck RG ; water : purified by ion exchange resins p > 2Mft.cm) . . 1) We first used the device as a potentiostat, by taking -R equal to zero. The potentiokinetic current-voltage curve given in the figure was recorded at a sweep rate of 10 V.mn"1. The a-b-c-d-e part was obtained at increasing potential and the e-f-d-g-a part at decreasing potential. The maximum b and minimum c can be explained in terms of the formation of a sulphate salt on the analogy of the case of iron(5). In fact, microscopic examination of the electrode in situ, by polarized light, revealed the presence of crystals. At sufficiently slow sweep rates, these extrema disappear, but even in this case a stationary and Israel Epelboin 269 reversible curve can be obtained only within the a-g-d and e-f parts. Within the d-f range, which corresponds to the transition between active and passive states (about 150 mV wide, see shadowed part in figure) neither a stationary nor a reversible recording can be obtained. This study shows that the potentiostatic regulation cannot allow the detection of certain kinds of passivation. 2) On the contrary, we have been able to record the whole a-g-d-f-e current-voltage curve by associating a negative resistance -R (R > 0) with the potentiostat P. Such a stationary and reversible curve evidences the passivity of cobalt. However, the passivity range is as narrow as 100 mV (Figure). Within the a-g part, it is possible to obtain directly current- voltage curves that are corrected for ohmic drop, provided that -R is constant and equal to -Re, and input X| is connected to ground (R = resistance of the electrolyte). However, we avoided such a correction because current and potential have a non uniform distribu tion on the electrode surface within the d-f part. In this case, the definition of ohmic drop becomes a problem in itself. As a matter of fact, within the transition range'between passive and active states (d-f) we observed a disc-shaped black layer which was concentric to the electrode. Under this layer the electrode metal dissolves, but the remaining,electrode area keeps its original metal lic brightness and does not dissolve. The larger the current the larger the surface dissolved. If this highly brittle layer is distroyed by accident, the equilibrium of the system is,strongly perturbed, and reaches back the stationary state very slowly. Examination of the electrode surface state and studies of poten tial decay showed that the two above-mentioned areas are simultaneous ly present and that they,are wellrseparated ; the bright area corre sponds to passive cobalt and the black one to active cobalt. We studied the potential decay after interrupting the cell-current (estab lished with the same device). If the current is interrupted within the oxygen evolution or passivity field, two potential arrests can be observed : the first one lasts for some tens of milliseconds and is usually referred to passivity ; but the second is shorter and might be attributed to active state. Within the transition range (d-f), the larger the current the shorter the first potential arrest and the narrower the "passive" area. When the current value becomes close, to the value at point d, the first potential arrest disappears and only the second one remains. The whole electrode surface is then active. In the case of iron (4), the homologous point of f (see Figure) corresponds to the end of the potential arrest (Flade potential). The case of cobalt is different since the end of the potential arrest is much less anodic than the potential at point f (the difference being about 200 mV). 270 Israel Epélboin References (1) I. Epelboin, M. Micinic, Ph. Morel. Mem. Sc. Rev. Metall. (1971), 68, 10, p.727. (2) M.T. Tikkanen, T. Tuominen. Acta Pol. Scand. Chem. (1968), n°70. (3) C. Gabrielli, M. Keddam, J.C. Lestrade. Comm. this meeting. (4) C. Gabrielli, M. Keddam, J.C. Lestrade. C.R. Acad. Sc. Paris (1970), 271C, p.1428. (5) M. Froment, M. Keddam, Ph. Morel. C.R. Acad. Sc. Paris (1961), 253, p.2529. Figure Current-voltage curve of cobalt in a molar sulphuric acid aqueous solution and experimental apparatus. - thick line : stationary curve - thin line : potentiokinetic curve *? Jean-Jacques FOMBON 271 APPLICATIONS OF PROGRAMMED SAMPLING PULSE TECHNIQUES IN ELECTROCHEMICAL ANALYSIS AND RESERACH Jacques Tacussel and Jean-Jacques Fombon Centre de recherches CERAC, 72, rue d'Alsace, F-69 Villeurbanne, France. The availability of apparatus (!) designed in order to allow practical ly any voltammetrio technique and, more generally, any technique involving a voltage/current/time relationship to be applied, and including the most recent electronic advances, has permitted us to investigate new solutions to various electrochemical problems : 1 -ANALYSIS Since initial studies by Barker (2), pulse polarography methods re ceive continuously growing acceptance. It has been possible to improve the results in this field by using a potentiostatic circuit fitted with very low noise and short risetime solid state amplifiers. The provision for a large range of static or dynamic adjustments, has resulted in the best operating conditions for every particular analytical pro blem ; these adjustments were made according to the electrodes, solvents and supporting electrolyte employed, as well as to the kinetic characteristics of the studied phenomena. In particular, the appropriate choice of the samp ling window position and of the pulse polarity has. allowed complex analytical problems to be solved, taking into consideration differences of kinetic beha viour at the electrodes of the species to be separated. Direct analysis of electroactive species have been made, at the drop ping mercury electrode (DME), for concentrations as low as 10 "8 M. We believe that the only limit to the increase of sensitivity now lies in the relative imperfection of the electrodes and in the difficulty of obtaining very pure solvents, chemicals or materials. An original application of fast pulse analysis seems to consist in the detection and the continuous monitoring of ion traces passing into solution during a corrosion process. As the choice of the working electrode is not limited to the conventional DME, we have used, in this case, the ring-disk method, with detection and integration of the ring pulse currents, which allows the quantitative determination, in the nanogram range, of species generated on the disk electrode. 2 - RESEARCH The same apparatus have been used in the particular case where an external signal is needed, in place of the internally provided imposed signal. 272 Jean-Jacques FOMBON During the study of reactions -reversible or not -at the dropping mercury electrode, it has thus been possible to superimpose a single sine- wave to the sweep signal, this sineweve, of adjustable frequency and amplitu de, being triggered at time t0 with an exactly known delay from the beginning of the drop life. We have studied the influence of the current sampling window position along the sinewave (that window being sufficiently narrow, so that measurements carried out are practically punctual), for different selected values of the electrode imposed potential (fig, 1). The recording of the cur rent sinewave thus obtained allowed us to proceed to extremely accurate phase measurements and to "visualize" distorsions resulting from generation of harmonics. Finally, some preliminary measurements have shown the interest of the potentiostatic control of the working electrode potential by means of a complex signal (fig. 2).. We think that experimental studies of kinetic reactions at the electrodes will become easier using such aJtechnique, with provision for oh- mic drop compensation (3), as it makes possible to very accurately know the imposed parameters EE (t)3 and to measure the resulting current El (t)D with a comparable precision. It must be mentioned that using a suitable com plex signal (fig. 3) allows to automatically perform a controlled pre-treat ment of the electrode, by applying to it, immediately before the main impo sed signal, two calibrated potentiostic pulses of opposite polarity. This type of voltage signal sequence can be used both for single-sweep and repetitive operation. REFERENCES (1) J. R. TACUSSEL : The "VOCTAN", an apparatus for electrochemical techniques involving voltage/current/time relationships. "Peeper to be presented during the present Meeting, (2) 6.C. BARKER: Square wave and pulse polarography. In : P. ZUMAN and I.M. KOLTHOFF : Progress in Pqlarography; vol II. - New York, Intersciencé Publishers, 1962, pp. 411 - 427. (3) A.A. PILLA, R.B. ROE and C.C.; HERRMANN : High speed non-fara- dalc resistance compensation in potentiostatic techniques. In: J• elec- trochem. Soc, vol. 116, n° 8, august 1969, pp. 1105-1112. Jean-Jacques FOMBON E A FIG. E * FIG. Pre-treatment Measurement FIG. period period The current sampling window is open at time t. and clo at time t ; both are adjustable within entire drop li 274 JAGDISH N. GAUR POLAROGRAPHIC STUDIES ON ELECTRODE KINETICS AND FORMATION CONSTANTS OF THE COMPLEXES FORMED BY Zn2+ WITH ORTHO AND META-TOLUATE IONS J. N. Qaur. D.S. Jain & Anand Kumar Department of Chemistry, University of Rajasthan, Jaipur-4 (India). With recent developments of several methods for the determination of kinetic parameters (Ks & ) attention has been turned towards the polarographic study of the complexes of metals which are not reversibly reduced. Zinc falls under this category and its acetate and tartarate complexes have been investigated by Matsuda & coworkers1-%• In our laboratories, the zinc complexes with formate,succinate, adipate, glutarate, maleate and fumerate ions have been studied by Gaur & coworkers^-5. The present paper deals with the determination of kinetic parameters and the overall formation constants of the complexes formed by zinc with toluate ions. With both the ligands, the reduction waves of Zn2+ were well defined and diffusion controlled* The log plots were linear with a slope of the order of 36 mV, indicating quasi-reversible nature of the reductions* The Ej. was shifted towards more negative value with increasing ligand concentration, indicating complex formation with zinc ion. The reversible half-wave potentials, E?, were calculated from the observed EL by Gellings ^method6 which involves the following equation: The plots of EF VS- log C_ gave smooth curves with both ligähds indicating the formation of two or more complexes which are in equilibrium. The formation _ constants were calculated by De Ford and Hume's method'. Jagdish II. Gaur 275 o-toluate formed three complexes: +1 Zn[C6H4(CH3)C00j , Zn[c6H4(CH3)C0q32 and Zn[b6H4(CH3)000^3 having = 1#25 a d = Pi = 1.25, fi2 ^ /^3 13.25. The m-toluate ions formed two complexes species: +1 Zn[c6R4(CH3) C00] and Zn^CgE^CCH3)COQ^ with p Z is a measure of the degree of irreversibility and has a value and »A ^ Ä w The values of Ks were found to be of the order of 10"3 cm/sec showing that the reductions are quasi- reversible. 276 Jn-äisk II. O cur REFERENCES 1. H. Matsuda, Z. Electrochem., 65, 482 (1961); 62, 977 (1958). 2. H. Matsuda and Ayabe; ibid., 63, 1164 (1959). 3. J.N. Gaur and D.S. Jain, Review Polarog. Soc. Japan, 14, 206 (1967). 4. V.K. Sharma, Ph.D. Thesis, Raj as t han University, 1967 . 5. M.M. Palrecha, Ph.D. Thesis, Rajasthan University, 1968. 6. P.J. Gellings, Z. Electrochem., Ber. Bunsengesell- schaft Phy. Chem., 66, 477,481,789 (1962); 67, 167 (1963). 7. D.D. De Ford and D.N. Hume, J. Amer. Chem. Soc, 73, 5321 (1951). A. V# Gorodisky 277 FARADAYIC IMPEDANCE OF ELECTROCHEMICAL SYSTEMS WITH CHARGB-TRAUSFER OVERVOLTAGE A.V. Gorodisky, Y.K. Delimarsky, A.V. Panov, N.H. Tumanova, B.F. Ornatsinsky Institute of General and Inorganic Chemistry, Academy of Sciences of the Ukraine, Kiev, U.S.S.R. Electrochemical systems are nonlinear complex resistances with an active and a capacitive components. In the case of reversible systems the electrode capacity is connected with the concentration change on the electrode surface and the phase shift between the gradient (current) and its logarithm (potential). For irreversible systems without concentration polarisation, the electrode capa citance is attributed to the double layer capacity. In these systems the active resistance is taken to be faradayic one, the capasitance is considered to be non- faradayic resistance. Experiment shows that the active resistance cannot be identified only with faradayic processes and the capaci tance is due not only to the double layer. The impedance components change simbatically with electrode potential. At sufficiently low frequencies both components (active and capacitive) are derivatives of the current-yol"éage curve. Both components yield electrocapillary curve when frequency is high. Thus, in the former case we deal with the faradayic processes, in the latter case both impe dance components are associated with the double layer charging. \4vV From this point of view the elimination of the capacitive currents in a. c. polarography is not to be considered reasonable. The components of the faradayic impedance of the irrever sible system do not change their phase angle when the electrode potential varies, i.e. their behaviour is ana logous to that of the Warbu^ |ii^dänce> This 'analogy can be explained by the fact that the resistance of the electrochemical inaction is determined by Boltzmann energy distribution among the Ir^cti]^-particles^ When the equilibrium distribution; has;;béé^ disturbed the re distribution of energy does4nbt< take placedat once• The mechanism of energy transfer obeys the diffusion - heat conduction equations, i.e.; the Fick-Furier ones• Thus, activation polarisation without the change oft otal concentration can be described by the variation in con centration of active particles and their diffusion. Following this concept we can use ordinary mathematical methods of nonstationary diffusion calculation. The re sults of such calculations agree with experimental data. 278 A. V. Gorodisl-ry But the principal proof of the transport impediments of active particles is that the intensity of electrochemical reactions is influenced by hydrodynamic factors. REFERENCES 1. A.V. Gorodisky, Y.K. Delimarsky, A.V. Panov, Collec tion "Physical Chemistry and Electrochemistry of Molten Salts and Slags'*, "Chimiya" Publishing House, Leningrad, 1966, p. 3. 2. Y.K. Delimarsky, A.V. Gorodisky, V.S. Kikhno, J. Theor. Exp, Chem., 4, 554 (1968). 3. Y.K. Delimarsky, A.V. Gorodisky,' A.V. Panov, J. Theor. Exp. Chem., £, 454 (1971). 4. A.V.Gorodisky,Collection "Physical Chemistry and Electrochemistry of Molten Salts and Slags", "Naukova Dumka" Publishing House, Kiev,vol-2,p.42. Lj. J ef tid 279 LOGAT&THMIC ANALYSIS OF TWO OVERLAPPING D. C. POLAROGRAPHIC WAVES USING COMPUTER I. Ruzic and Li. Jeftic Center for Marine Research, "Rudjer Boskovic" Institute Zagreb, Croatia, Yugoslavia The equation of a pblarogram consisting of two overlapping d. c. polarographic waves can be given in the following way. »^ i = idx/(l + x) where x = (x| + x x )/(l + x x /x|) , i = i + i d2 ,(11 ) (2) x1= exp[- nxF (E - E ^) /RTJ , x2=exp[-n2F (E-E^^RTJ x*= (mx1 + x2)/(m +1) x*2 = (m+l)x1x2/(mx2+x1) and m= VSa From the experimental data values for log x = log i/(i -i) vs. potential can be calculated and plotted. The corresponding curve (Fig.) has two linear parts with end inflection point between them. 1 \ xt/xf x • ^^^ I Jr l log X / 1-M^r-x»p — d ^** -fL"/ •••' /r 'f % •0CHLJL-. —I_ - i i • •• .'•:• i i - X X2 o X2 E c<1> c<2) •••••£••• t1/2 ^1/2 280 Lj. J efti d Equations for these linear parts are given by expressions for x* and x* and corresponding values for the entire investigated potential range can be obtained by extrapolating the linear parts of the original curve. Values for x. and x? can be calculated from x, x* and x* (for the given potential) by using the. following expression: 2 X1.2 = (A^A - B^)/2 where A = x* + (x* - x)/(x - x* ) and B = 4x* (x* - x)/(x - x*2) Expressions for x. and x are actually equations for straight lines which represent the logarithmic analysis for each individual polarographic wave. The ratio of the limiting currents of individual waves can be evaluated from: m = (x*x - x2)/(Xl - x*x) wherefrom the limiting currents of individual waves can be calculated from the total limitng current. On the basis of the given theoretical approach the computer pro gram was made in FORTRAN language. The resulting output data consist of a complete logarithmic analysis of the composed and separated waves, the, corresponding limiting currents, half-wave potentials, slopes, transfer coef ficients and rate constants of electron transfer processes. The output results were presented either in the form of the Table or by using the printer as plotter, A plotting subprogram was made for the latter case to give an output plot of 120 x 60 character size. The information that can be obtained for a simple electrode pro cess depends on the rate of the charge transfer for the process involved . We have considered three possibilities for each wave, namely the reversible, the quasireversible and the irreversible charge transfer. This has given nine possible combinations for two waves. For all the combinations involving a reversible and an irreversible wave the half-wave potential for the reversible one and oCn. and k for the irreversible can be determined. s If there is a quasireversible wave in the combination it is relevant whether it precedes or follows the other. The reversible half-wave potential can be obtained for the preceding quasirever sible wave and c^n< and k for the following one. The same applies for two qua si reversible waves. If the quasireversible Lj. J efti 6 281 wave is sufficiently'larger than the other it is possible to evalu ate its reversible half -wave potential as well as ©^.n^ and k , The case of two-step electrode reaction has been treated in the following way: . for E > O the above mentioned method can be used , for E <. O the experimental logarithmic diagram must be interpolated by an antilogarithmic sum of three straight lines. From the slopes of these lines transfer coefficients for both steps can be calculated and the standard rate constants (if corresponding standard potentials are known) can be evaluated. REFERENCES: 1. I. Ruzid, and M. Branica, J. Electroanal. Chem. 2£ (1969) 243. 2. I. Ruzic, ibid.29 (1971) 440. 3. I. Ruzic", ibid.in press. 4. I. Ruzid, ibid. 25 (1970) 144. 282 Andre Jouanneau ANODIC BEHAVIOUR OF NICKEL IN SULFURIC ACID B.Dubois. A.Jouanneau and M.C.Petit Laboratoire de Mécanique-Physique Université de Bordeaux I FRANCE. ( 351, Cours de la Liberation 33 - TALENCE ) The anodic behaviour of nickel in sulfuric acid has been investigated, the polarization curves are determined by potentiostatic techniques. The total reaction is a multi- step process and it is possible that;several reactions are simultaneous : nickel dissolution,1passivation and oxygen discharge. To separate the different reactions, the current efficiency has been; obtained by the nickel weight losses, and by the volume of gas evolved. The figure I shows four curves : (1) the anodic potentiostatic polarization curve Itot ,= f(V),<2) the rate of nickel dissolution, (3) the rate of oxygen discharge and (4) the calculated polariza = tion curve Itotal f(V). To explain the shape of the anodic polarization curve, we have been supposed that, within the potential range inves tigated, three reactions take place simultaneous : metal dissolution, passivation and oxygen discharge. The mecha nisms of these reactions are similar : one or two electron transfer occurs to an intermediate ion OH" adsorbed on the electrode surface : 8i is the fraction of the total elec trode surface covered by the first intermediate adsorbed Ni(OH)ads, 62 by Ni.<°H)2ads and 63 by Ni(OH)3ads. The figu re II shows the particular disposition of 6i, 62 and 03 on the electrode surface. The first step is the formation of the intermediate Ni(OH) . bJy one electron transfer ads k Ni + OH" -^-> Ni(OH) , + e ads the rate of the electrochemical reaction is biv va = k-JOH"! (l-6i)e = k^l-9^ and k± =f(V,OH") After the first electron transfer, two possibilities depend upon the potential : nickel dissolution Ni<0H)ads Ni OH* + H+ > Ni++ + HgO *>2v v2 = k2 Ni(0H) + 2 0H Ni(0H) + 2 e ads k " —> 3 ads Ni(0H) Ni(0H) + + 2 3 ads —> ads 2 H + O + 2 e 7 b,V v5=k5(e1-e2^e3)|0H"re ° = K5 The values of 6x 62 and 63 are calculated with a IBM computer and presented on the figure III. Kl 2 + *1 *2 Kc + K, 3 5 KT+K7 " (K3+K4) K K 3 5 K K. V*6 " 3 6 6i and 63 = (6,-e2) K5 + K6 TC+KT " (K3 + *V 5 6 Within the potential range investigated, the total anodic current is the sum of the three partial currents Xtotal = Idissol. passiv. oxygen The values of the total anodic current are calculated and showed on the figure I. Itotal = F| 1^(1-61> + (K2+K5) (6,-62-63) + K6 631 284 André Jouanneau rrA Fig I 30- Fig IE Fig HE ©1 —7*—" / ö2 I 83 I / ®1 e /e3 / Ni / ><••'. , ; o. _ ,m „^£.j££££^ bOO 1000 1500 2000 Jean-Claude Justice 285 AN EVALUATION OF THE ACTIVITY COEFFICIENTS OF HC1 IN WATER-DIOXAN AND WATER-THF MIXTURES BY USE OF THE CONDUCTIMETRIC METHOD. Marie-Claude Justice and Jean-Claude Justice Equipe associée au CNRS, Laboratoire d'Electrochimie, Universite de Paris VI, 11 Quai St. Bernard, Båt. F, Paris V, France. The Debye treatment has led to the wellknown derivations of the activity coefficients law by Debye and Hilckel (D-H) and of the con ductance equation by Pitts and Fuoss and Onsager (F-O). Bjerrum has refined the D-H equation to take care of short range electros tatic interactions which had been somewhat neglected. The same generalisation has been proposed more recently for the F-O conduc tance equation. This leads to a consistent set of equations which makes use of the Bjerrum concept of association. l/2 1/2 /Z 3 A =&*[*. - Sc T + Ecpog cT+ J(R)cT+ Jj/2(R) r R 2 1-T 4~N f f e ^ 2. •: ,_. 2 .2 = KA =( exp/ — r dr . . . . (2) Hi cy'+ [1000 J VrDkT ,1/2 Ln y' = - _q —^ ... (3) t l + kRTVZ 2 where R is usually fixed at the Bjerrum criticajlvalue q = 2 DkT This system predicts the behaviour of electrolytes in dilute solutions ( kR"5" < 0,4) in solvents of low dielectric constants where pairwise ionic association is significant. Equations (2) and (3) may be used to interpret thermodynamic excess functions such as the activity coefficients. The entire set of eq (1), (2) and (3) is necessary to analyze conductimetric data . (A recent conductance treatment by Falkenhagen et al. has shown that the Bjerrum association concept resulted directly from a statistical mechanical approach). The close interconnection of ther modynamic and inductimetric variables in the above system allows to think that one experimental method might help to.derive infor mations concerning the other.- In order to test this assertion it was interesting to obtain conductance measurements on systems which had been studied thermodynamically before. The conductance of HC1 in water-THF and water-Dioxan has been measured in the-present work. The data are analyzed in terms of the adjustable parameters of the Fuoss-Hsia conductance equation. 286 Jean-Claude Justice In return the value of the association constants obtained are intro duced in equations (2) and (3) to derive, stoechiometric activity coefficients yt =yyt which are compared to existing litterature results by Harned and Morrison in water-Dioxan mixtures and by Roy and Sen in water - THF mixtures. For this purpose the original e.m.f. values are reanalyzed. The original treatments were based on a polynomial extropolation in order to reach the standard 's which were used again to cal culate the activity coefficients. This extrapolation method being invalid when association becomes important, the final results are thus more and more in error. A new analysis is presented which is consistent with the procedure used with conductance data. A final comparison and discussion is given. It is concluded that conductance which may be used on any electrolyte- solvent system with high precision is a very valuable method to derive reliable thermodynamic data.. Jean-Claude Justice ., 287 AN EXTENSION TOWARDS LOWER DIELECTRIC CONSTANT SOLVENTS OF THE CONDUC TIME TRIC METHOD FOR ELECTRO LYTES STUDIES. Christian Micheletti and Jean-Claude Justice. Equipe associée au CNRS, Laboratoire d'Electrochimie, Université de Paris VI, Båt. F, 11 Quai St. Bernard, 75 Paris 5e - FRANCE. Among the different features which characterise the theoretical equa tion for the prediction of the equivalent conductance of symmetrical electrolytes in dilute/solutions, its series expansion nature in salt concentration is the main factor which tends to limit its range of validity with respect to two experimental variables.. One/bf these variables is obviously the salt volume concentration itself. The other is the dielectric constant D of the solution (for a given salt concentration). This becomes obvious when one uses in the expan sion the more rational dimensionless variable T introduced by Fuoss as a "reduced" ionic concentration. T = q/ H'"1 = Hqy 1//2 ... (1). is the ratio of the Bjerrum critical distance q to effective Debye distance x'~ = H" T" ' where *~ is the stoechiometric Debye distance and ft* the fraction of free ions according to Bjerrum's concept of association : it is shown that the classical form of the conductance equation l/2 1/2 3/2 3/ 2 A=r(A„ - Sc J +EcTlogcr+ j(R)c]T+J3/2(R)c r ' ..) . .(2) can be expressed in terms of T according to 2 2 3 A =r(Ae - S'T+E'T log T+ J'T + J'3/2 T ...) ... (3) where the coefficients S', E1, J' and J' /_ now vary linearly with D through their hydrodynamic terms. (It is in addition assumed that the distance parameter R in eq. 2 was equal to the Bier rum distance q). Similarly the Debye-Bjerrum activity coefficient law for free ions can be expressed as Lny^ = - T /(1+ T ) . ... (4) The recent developments concerning the importance of the c term, the derivation of the J /_ coefficient and the meaning of the distance parameter R allow the expectation that the A equation may be used for solvents of dielectric constants lower than 1£. This was the value formerly recognized as a lower limit for D beyond which one could not expect to obtain valuable informations by comparing conduc tance data with equation 2. 288 Jean-Claude Justice The conductances of tetrabutylammonium and triisoamylbutylammo- nium tetraphenylborides (salts A and B respectively) have been mea sured in different binary solvent mixtures : methylisobutylketohe (MK) - THF (solvent Y) ; MK - Dioxan (solvent X)and THF-Dioxan (solvent Z) mixtures at 25°C. The variable T never exceeded 0.4 in any run.The data have been reduced by use of equation (2) in a recent derivation which takes care of the Chen effect, into three adjusted parameter (by a classical least square procedure) \, J , and = 1_ c ^A|/KA ( ^) ' ^ y'i where y» is the mean activity coefficient of free ions calculated by the Debye-Bjerrum equation (4).]R in J was fixed to the Bjerrum value ; R / was derived from the adjusted J , coefficient. ' The results are summarized in the following table. -• • ' • 1 1 5 r—t 5 ' — 1 2 ' o Salt Sol D q (A) Ao K a(A) *3/2(A) A io ah vent % MK 13, 05 2\A1 69, 56 18,9 668 2 6,9 X 10,75 26, 06 61,4 29 1500 6 7,5 X 8, 80 31,84 54,9 32 6000 4 7,0 A X 6,66 42, 07 47, 41 55 61600 2 7,0 z 6,47 43, 31 71,5 * 103600 4 6,8 X 6, 13 45,71 46, 09 * 178000 3 6,8 X 5,43 51, 60 44,1 * 840000 5 6,8 -WfoK 13, 05 21,47 68,16 19,3 564 2 7,6 Y 10, 32 27,15 72,42 25,9 1440 2 8,4 X 9,07 30,89 57,13 24,8 4400 2 7,2 Y 8,88 31,55 74, 0 38, 3 4050 6 7,7 Y 7,37 38, 02 51,75 29,2 20000 2 7,2 Z 7,25 38, 65 77,4 * 23500 6 7,2 B X 7,24 38, 70 49,3 45 20500 6 7,4 X 6,97 40,20 49,3 90 33500 3 7,2 Z 6,88 40,72 75,4 * 38900 4 7,1 X 6,59 42,52 47,72 * 57400 2 7,2 z 6,56 42, 71 70,8 * 62000 4 7,1 * i X 6,41 43, 71 45,14 67900 1 7,3 \ X * 6,26 44,76 46,45 * 98700 0, 5 7,2 X 6,25 44,83 46, 74 91200 2 7,3 L .- (a) % is the standard deviation in percent. A . The last column gives the value of the closest distance of approach of the associated ions obtained by solving the Bjerrum equation for a. The remarkable constancy of the contact distance, given the wide Jean-Claude Justice 289 range of variation for the observed association constants is stron gly in favor of the Bjerrum Concept of association. Whenever determinable, R , follows closely the values of q, except for the lower dielectric constants where the J , coefficients become positive (cf. * in column 6). This is probably due to the no longer negligeable ionic triplet :f or mation which should contribute posi tively (in c ) in the conductance equation. Nevertheless, adjusting empirically the J , makes this coefficient take care of this last effect and allows a good determination of the ionic pairwise association constant K. . A 290 Bertel Hastening ADVANCED TECHNIQUES FOE E.S.R. INVESTIGATIONS OP ELECTRO- CHEHICALLY GENERATED RADICALS Bertel Hastening , B. Gosti£a-Hihelci6 and J. DiviSek Forschungsabteilung Angewandte Elektrochemie,Kernforschungsan- lage Jiilich,Julich, Germany. Two techniques have been developed which allow for an improve ment of e.s.r. investigations in electrochemistry. The first of these methods aims at an- improvement in the internal ("intra muros") generation of radicals which has,in general, suffered from non-uniform current distribution across the gene rating electrode ' due. to the unfavourable sample cell geometry. This drawback is overcome by dividing the electrode into several independent strips each of which has its individual current supply,thus largely eliminating the iR drop in the electrolyte across the length of the sample cell.This allows for the appli cation of enhanced current densities at a comparatively large electrode area and,consequently,an increased sensitivity. The second technique applies the external generation of radicals in a separate electrolysis cell and is an improved version of the 2) flow technique ' formerly developed for the study of fast reac tions of radicals in a medium differing from that in which they were generated.The radical solution from the electrolysis cell and a. second reagent solution are periodically mixed and forced, by an all quartz-glass twopiston pressure pump,through a capil lary Inserted in the resonant cavity.The rapid turbulent flux velocity ensures uniform mixing as well as rapid transit of the mixed solution through the sensitive part of the e.s.r. cavity within a very short time after mixing. References: 1) J.K. Dohrmann and K.J. Vetter, J.Eleetroanal.Chem» 20 23 (1969) 2) B.Kastening, Z.anal.Chem. 224 196 (196?), Ber. Bunsenges.physikal.Chem. 72 20 (l968)jB.Kastening /- ' ' and S.Vavrieka, ibid. 72 27 (1968). B# Lengyelfjun# 291 THE ELIMINATION OP RESISTANCE POLARIZATION IN POTENTIO- STATIC INVESTIGATIONS B. Lengyel. .iun.. J. Devay, J, Mészåros Research Group for Electrochemistry of the Hungarian Academy of Sciences, Hungary A serious source of error arises in the study of the kinetics of electrode processes as a consequence of the ohmic potential appearing on the resistance of the solu tion layer between the reference electrode and the work ing electrode as well as on the resistance of the surface layer formed on the latter. The ohmic potential is espe cially large in the case of solutions of high resistance and when large currents are flowing in the cell as well as in the case of the formation of surface layer on the electrodes. An equipment was designed for the automatic compensation of the ohmic potential in the case of potentiostatic investigations. The working principle of the instrument is the following: an A.C. current proportional to the polarizing D.C. current is superimposed on the working electrode controlled by a potentiostat. The impedance of the capacity of the double layer is negligible as compar ed to the ohmic resistance when the frequency of the A.C. current is properly chosen. The A.C. voltage appearing between the reference electrode and the working electrode is proportional in this case to the ohmic potential de veloped on the resistance of the solution and that of the surface layer of the electrode. The A.C. voltage is amplified to such a degree as to obtain after rectification a D.C. voltage equal to the D.C. potential appearing on the ohmic resistance which has, to be compensated for. The control potential -of the potentiostat is increased by the above-mentioned D.C0 292 B.Lengyel, jun. voltage and thus it is ensured that the potential of the potentiostat independently from .the magnitude of the ohr> ic resistance and the variation of the latter during the measurement. "1 , , • r\. ^-^ * i i r 7&H> 6 Pig. 1. Our measuring equipment is shown in P±g# 1# The D.C# curr ent flowing through three-electrodes cell (l) is fed in the unit consisting of modulator (5) and A.C. voltage amplifier (6) , which generates an A.C. current propor tional to the D#C# current, and supplies the cell with Å.C. current through resistance (7) and condenser (8). Capacitor (8) serves to separate the direct current and alternating current circuits, and resistance (7) ensures the functioning of amplifier (6) as a current generator, i.e., the alternating current flowing through the cell is independent of the resistance of the cell* Reference electrode 6-4) is connected to the input of differential amplifier (2) through filters ÖL1,1^) , as well as to A.C. voltage amplifier (9) and demodulator CLQU The output voltage of the demodulator, which is equal to the voltage across the ohmic resistance, is added to the polarization voltage adjusted on potentiometer (ljj) » and fed to the second input of differential amplifier (2)• 293 B.Lengyel, jun. Filter circuit 0.1,-12*) prevents the A#C# voltage between the reference and the working electrodes to get to the input of the differential amplifier. Inductance '(4) pre vents the shunting of the cell by the output impedance of the potentiostat from the viewpoint of the A.C. current. The range of the A.C. measuring frequency can be determ ined from the frequency characteristics of the impedance of the measuring cell* In order to control experimentally the above measuring technique model systems were chosen where either a high resistance surface layer was formed on the electrode or where the solution between the working electrode and the reference electrode exhibited a high resistance. The study of the anodic polarization of silver in 0,1 n HgSCV and 0,1 n CH^COOH respectively were found very sui table for our purpose as in sulphate containing solutions the surface layer formed on the electrode surface had a high resistance which was changing in time while in acet ic acid solution no surface layer was formed and the mag nitude of the resistance polarization was defined by the ohmic resistance of the electrolyte only* The comparison of the polarization curves showed that the character of the latter was changed under the effect of the contnuous automatic compensation of the ohmic poten tial» Thus the above automatic compensation technique rendered possible the elimination of resistance polari zation when the latter was either constant or variable in time, A further advantage of the method consists in the fact that it is possible to place the reference electrode to a larger distance from the working electrode in order to eliminate the screening effect of the former electrode. 294 Jean-Claude Lestrade RECORDING OF Z-SHAPED PASSIVATION CURRENT-VOLTAGE CURVES Claude Gabrielli, Michel Keddam, Jean-Claude Lestrade Physique des Liquides et Electrochimie, Groupe de Recherche du CNRS, associé ä l'Universite Paris VI, 11 quai St-Bernard, Paris 5e,France If a potentiostatic regulation is used to record a current voltage curve, a single value of current must correspond to a given value of voltage. This requirement is not met with the Z-shaped carve shown in Fig.l, which obtains in the case of an iron electrode in 2N sulphuric acid (1). A load line such as v'-d', which charac terizes a potentiostatic regulation, intercepts the curve in three points and the observed current is not the same when the voltage is swept towards more anodic or more cathodic values. Then we need load lines such as v-d, which obtain with a voltage source in series with a negative resistance. IRON DISK ELECTRODE (05mm) 2N SULPHURIC ACID ROTATION SPEED: 750 R.P.M. 200 < E z UJ 100 a O -0.5 0 + 0.5 POTENTIAL (V/S.C E.) Figure l Jean-Claude Lestrade 295 The device known as a "negative impedance converter" (NIC) answers this purpose (2). It is shown schematically in Fig.2a, and is easily made with an operational amplifier (Fig.2b). Its main feature consists in the relations between input A and output B given in Fig.2a. The positive dimensionless constant K is calculated from Rj, R2 (Fig.2b) and the characteristics of the operational amplifier. Within a good approximation, it is equal to R2/R1• If a resistance R, in series with a voltage source E, is connected across the input, the output acts as a source E with a negative internal resistance -R/K. On the other hand, if a resistance R' is connected across the output, the input behaves as a negative resistance -KR1. NIC !£. '«JT li" »ft E^E, VKI, (a) (b) Figure 2 (a) Figure 3 A NIC may be used in the two ways shown in Fig.3. In the first instance (Fig.3a), the electrochemical cell is fed through the counter-electrode from a stabilized voltage source with negative internal resistance. The curve of Fig.l has been recorded with this- setting upV To be precise, the load line v^d corresponds to a linear relation between current and the potential drop from working to counter-electrode. Nevertheless, the experimental curve refers to the actual voltage between working and reference electrode which is measured independently and is stable with a proper choice of E and R. 296 Jean-Claude Lestrade In the other setting up (Fig.3b) a NIC is used to place a negative resistance between ground and the working electrode, and a potentios- tat controls the voltage between ground" and the reference electrode. The relation between t;he current and this voltage leads to a curve where the Z-shape feature disappears with a proper choice of R' for each current value. This setting up is particularly convenient when a known ohmic drop has to be corrected for ; so far, both devices have been used with equal success when applied to passivation studies, the former being generally prefered because of its simplicity. The main result obtained when studying the passivation of iron in acid sulphuric medium, consists in proving the existence of a continuous and reversible transition between active and passive state (Fig.l). This transition is hindered when a potentiostat is used as a control device (1). A detailed investigation (3), (4) shows that dissolution occurs on a ring-shaped zone of the rotating disk- electrode. On the other hand, a similar Z-shaped current-voltage curve has been observed in the case of cobalt, and is given in ano ther communication at this meeting (5). (1) C. Gabrielli, M. Keddam, J.C. Lestrade, C.R. Acad. Sci. 271C, 1428, (1970). (2) Vo Hoang Hien, G. Mesnard, Onde Electrique, 4£, 484, (1969). (3) C. Gabrielli, M. Keddam, J.C. Lestrade, H. Takenouti, C.R. Acad. Sci. 274C, 123, (1972). (4) I. Epelboin, C. Gabrielli, M. Keddam, J.C. Lestrade, H. Takenouti, submitted to J. Electrochem. Soc. (5) I. Epelboin, C. Gabrielli, Ph. Morel. This meeting. Alain Léveque 297 POLAROGRAPHY OF LANTHANIDES IN N- N- DIME THYLFORM AMIDE. Alain Léveque and Robert Rosset Laboratoire de récherches de Chimie analytique générale, associé au C.N. R. Si Ecole Supérieure de Physique et de Chimie de Paris, 10, Rue Vauquelin, Paris 5éme, France - In aqueous solutions only three lanthanides are reduced at the DME: europium, ytterbium and samarium whose oxydation state II + is stable. It is possible to extand the electroactivity range using non aqueous solvents and we choose N-N-dimethylformamide, basic electrolyte beeing tetramethylammonium perchlorate 0, 1 M. Half- wave potentials of lanthanides ions are given on figure 1.(reference electrode is Ag/AgCl whose potential taken to the half-wave poten tial of ferrocene in DMF is - 0, 456 V) SrWtt-1.86 V. Ybm -1.40 V. -Figure 1- Eum-0.67 V. Values are quite similar to those reported by G. GRITZNER (1, 2). All lanthanides give a single wave exeept for Eu-' , Sm^+and Yb^+ and there is a linear relation beetween limiting diffusion current and concentration. In the case of mixtures of lanthanides ions phenomenon appears to be much more, complicated : waves are not additive when the mix ture contains an yttric earth (gadolinium, terbium, dysprosium, hol- mium, erbium, thulium and lutecium). We observe also deforma tion of waves and classical polarography cannot be used for such mixtures. We show eerie earths ions (lanthanum, cerium, neody - num, praseodynum) are reduced in a single stage with amalgam for mation e.g.: La 3+ HgJ + 3 e La (Hg)| 298 Alain Lévéque For yttric earths it is not the l.anthanide ion which is reduced but solvation water of this ion, for instance : + Er (KzO)* + pe » Er (OH)3l + (n-p)H20 + -f"H^+ (p-3)OH~ In mixture with a ceric earth, hydroxyl ions produced at the D.M.E diffuse towards the solution and precipitate ceric ions (and yttric ones) : limiting diffusion current of ceric earth decreases and wa ve is distorted -(figure 2). Potentialvolts vslAg/AgCl| < •M . Ö -Figure 2 : Polarogram of a mixture lanthanum-lutecium Different evidences are given of this mechanism b\r studying lantha num - erbium and erbium - calcium mixtures (Caz is more diffi cult to reduce than Er^+). Reduction of solvation water of erbium give OH" ions that precipitates Ca (OHJ- insoluble in DMF and , again waves are not additive. The study of amount of water influen ce in solutions on half-wave potential and in the case of a mixture agree the proposed mechanism. Anodic stripping by pulse polarography demonstrates that ceric earths give amalgam but yttric earths do not give any redissolution peak. As precipitation reactions in the diffusion layer become less impor tant as concentration decreases, pulse polar ography applied to Alain Lévéque 299 Potential Figure 3 : Pulse-polarogram of a mixture lanthanum-erbium more dilute solutions, give good results for mixtures of ceric and yttric ions. Lanthanum - erbium mixture has been studied by this method (figure 3). BIBLIOGRAPHY (1) G.GRITZNER , V. GUTMANN et G. SCHOBER , Mh. Chem. 1965 , ^6 , 1056. (2) G.GRITZNER , V. GUTMANN and R. SCHMID , Electrochimica Acta , 1968, 13 , 919. LO o. x: o> rt o~ —|T3 o» -». o O" -? O O» O» O UI Ut rf O rfO O-CD —« "O -»• 3 -H 0 -a -a CL-n O O c -a O 1 fl) -J 3-C 3"0"0 3 C CD CD C 3 3 Hlf+Ol 3--I) 3-fli-i. 3-J. 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(3) At faster scan rates, peak current will be increased by more than expected theoretically by increasing the pulse amplitude, e_.g., greater than lOx increase by increasing pulse amplitude from 5 to 100 mV. (4) For a given pulse amplitude, the peak current will be di rectly proportional to the drop area if a constant scan rate times drop time product is used. Instrument design strategies which avoid or minimize the problems discussed above are suggested. 302 Ivan Piljac ELECTROMETRIC DETERMINATION OF STABILITY CONSTANTS OF ZINC, CADMIUM AND LEAD COMPLEXES IN BUFFER SOLUTIONS OF /3-HYDROXYPROPIONIC ACID AND*-,/3- AND r-HYDROXYBUTYRIC ACIDS I. Pil.iac, S. Nushi, B. Bach-Dragutinovic, B. Grabaric and I. Filipovic Laboratory of Inorganic Chemistry, Faculty of Technology and Institute of Inorganic and Analytical Chemistry, University of Zagreb, Zagreb, Croatia, Yugoslavia The stability constants of the complex species of zinc, cadmium and lead in aqueous buffer solutions of 4-hydroxypropionic acid and *-, /3-, and $*-hydroxybuty- ric acids have been determined by thrée-electrode pola- rography and by potentiometric titration. The ligand solution was delivered discontinuously, directly into the electrolytic"cell. The following mix tures were used as starting solutions: solution S^ (4-10~4 M Me, 0.01 M HL, 1.999 M NaClO/,) and solution 4 S2 (A-'IO- M Me, 0.01 M HL, 1.999 M NaL).To an initial volume of 10 ml Si, .specified volume increments of S2 were added, recording a polarogram after each addition; an initial polarogram was first recorded with pure S]_. The same procedure was followed in reversing the order of addition, i.e. adding S-, volume increments to 10 ml Sp. Ligand concentrations varied from 0 to 1.999 M in these experiments. Each system was prepared and recor ded in triplicate. Working electrode potentials and cor responding currents were measured at the end of a mer cury drop's lifetime. In measuring potentials the repro ducibility was 0.2 mV. Individual polarogram halfwave potential values of independently prepared identical solutions agreed within 1 mV; log i/(i^-i) vs. Eä#e^mV) plots were linear, with an average slope -29«5±0.5, suggesting reversibility of all polarograms. Potentiometric titration was carried out in a three- compartment electrolytic cell; junctions between the sol utions were formed by means of special porous glass. Measurements were made with a quinhydrone electrode. Potentials were measured by a,compensation method, using a precision potentiometer with reproducibility 0.2 mV. Refined halfwave potentials and diffusion currents were obtained by simultaneous treatment of graphically determined rough data for Ej/p and i^, utilizing a least-square program with an IBM 1130 computer. The same program was then applied for calculation of values of the function F0([L]), according to the method by DeFord and Hume and these values were subsequently weighted,following a suggestion by Momoki, Sato and Ivan Piljac 303 Ogawa. By least-square fitting of these weighted poly nomials, stability constants and corresponding standard errors were obtained,- The results from potentiometric measurements were used to calculate values for H and [L] according to the method by Pronaeus and from these values the stability constants were determined graphically. The latter were taken as approximations and refined by means of the weighted recurrent method by Gauss, using Tobias» Gauss Z program with a computer. Polarographically determined stability constants were the following: /*! ^2 ^3 Pb /5-hydroxypropionate 135*13 1263*121 3643*144 *-hydroxybutyrate 126*19 596*192 3712*280 4-hydroxybutyrate 149*19 996*206 5045*275 i^-hydroxybutyrate 189*23 1412*265 4380*301 4. ^2 ^3 >4 Gd /3-hydroxypropionate 14+2 59*19 180*22 c( -hydr oxy but y rate 17*3 136*30 176*7^ 279*46 >4-hydroxybutyrate 13*2 160*23 226*26 Jf-hydroxybut yr at e 28*2 158*21 267*26 Potentiometrically determined values for the stabi lity constants of the zinc <<-hydroxybutyrato complexes were A±= 48*6, A?= 1228*175, /%= 5860±615 and /3^= 19160*1437- 304 Velimir Pravdié STUDIES OF COMPLEX ELECTROCHEMICAL-CHEMICAL REACTION MECHAMISMS WITH THE ROTATING RING-DISC ELECTRODE OF PLATINUM, AMALGAMATED PLATINUM, AND NICKEL Nikola Bonacci and Velimir Pravdic Electrochemistry Laboratory, "Rudjer Boskovié" Institute, Zagreb, Yugoslavia A simple, novel construction of the rotating disk (RD) and rota ting ring-disk (RRD) electrode assembly is shown in two figures. Figure 1. shows the whole unit with the cell, and Fig. 2. the construction of the detachable electrode head. The electrode itself is a concentric, precisely machined skeleton, to which the electrode material (Pt, Ni, or any other metal) is silver-soldered or spot-welded. The whole unit is ^covered by extrusion with a paste of Araldite polyester resin which, upon hardening, is machined to final size and balanced. The result is a sturdy, precise unit with closely controlled concentric spacing and excellent balancing. The chemical resistivity of the body is that of the thermosetting polyester thereby suitable for most work, where traces of organics are not critical. Fig. 3 is a photograph of the electrode. The driving assembly consists of a D. C. motor connected to the electrode shaft through a friction disk transfer. The revolution rate is controlled by a tachometer-generator which is driven by gears on the electrode shaft. A potentiostat, fed from the tacho, controlls the D. C. motor. The unit is capable of working at a constant revolution rate, but the constant accelera tion/deceleration mode is preferred. The tachometer output is applied to one axis of an X-Y recorder, the electrical response of the electrode (current or voltage) to the other. Precise data, obeying the Levich equation, have been obtained up to 14, 000 rev/min. The electrode and the technique have been used in studying the U(VI) - U(V.) system in carbonate media. It has been postulated before that not the most stable form of tri-carbonato complexes undergo reduction, oxidation, and disproportiönation, but that there is a preceding dissociation step. Investigations into the reduction of U(VI) in NaHC03+NaC104 medium at pH=8. 55, show no kinetic control. Data shown in Fig. 4 and in Fig. 5 indicate diffusion limited currents up to revolution rates of 14,000 min"*. Data of the type shown in Figs. 4 and 5 have been used in calcu lations of the rate coefficients of electrochemical reactions. For the reduction of U(VI) at the amalgamated mercury electrode at pH=11.3, where disproportionation is of negligible rate, and the Velimir Pravdic 305 reduction reaction can be treated as a totally irreversible step, -7 the rate coefficient (corrected for ^ 2 potential) i kQ t= 9 x 10 cm sec" with a corrected overall transfer coeffi cient' CL = 0. 56. Further studies, involving the amalgamated ring-disk electrode, and measurements of the disproporti önation rate and kinetics, are in progress- Fig. 1. Fig, 1. General view of the electrolytic cell with the RD electrode* Fig. 2. Sketch of the RRD electrode unit. 30 Fig. 2. 306 Velimir Pravdid EM SC£. U (VI) 1iW3 M NQHCO3 1 NoClO^nHjOl H pH-«55 K) 20 Fig. 3. Fig. 4. Fig. 3. Photograph of the RRD electrode. Inside, the brass col U (VI) IxKT M NoHCO, 1 M lector rings for contacts are NaClOtxH,0 1 M pH-U5 shown. .-0.90 * 20 Fig. 4. The disk current vs. square root of angular velocity for reduction of 1 mM U(VI) in E NaHC03/Na2C03 solution at a u total concentration of Na+= 2 gion, pH=8. 55. 10H Fig. 5. The reciprocal current vs. reciprocal square root of angular velacity plot of data from Fig. 4. 5 i. •> iv wj-r.102(»tc*> Fig. 5. I. Ruzic 307 A GENERAL APPROXIMATIVE METHOD FOR THE INTERPRETATION OF KINETIC D. C. POLAROGRAPHIC WAVES I. Ruzic Center for Marine Research, Rudjer Boskovic" Institute Zagreb, Croatia, Yugoslavia The general scheme of an electrode process in which the charge transfer reaction coupled -with simple chemical reactions can be represented in the following way: A + Z sfe* B + Z' (I) ne -M$K E C ° + W 3* D + W' k/k' = K and J/J ' = <£ For pseudo-monomolecular chemical reactions the following con dition must be fulfilled: z or z* ^^a+b+c+d <£<£. w or w* , where a, b, c, d, z, z* , w, w' are concentrations of the corresponding spe cies. In this case the scheme (I) simplifies to: D (II) The equations of the limiting kinetic currents for this mechanism can be evaluated by using Brdicka-Wiesner approach. A more general expression for the kinetic current can be obtained by introducing a new term k z b and by adding the term (b*-b )'9f according to the treatment given by Kern : ^ \ i = nFq. l(T3y*(kza0-k' z'b°) +&(b*-b°) = = nFq-lo'^^fwc0 -j' w' d°) +je(c°-c*) (1) where q is the electrode surface,A an dit are the thicknesses of the reaction layers, X,)t' are the cathodic and anodic Ilkovic constants respectively and superscript refers to the concentra tion in the vicinity of the electrode. During the electrode reaction concentrations a and b decrease, while c and d increase at the electrode. Therefore: 0 id-i =)t(a ^Sa/Db+b°) and i-Id =jc! (c^^D^T^) (2) Combining the eqns. (1) and (2) with the Nernst equation: P = exp nF(E-E )/RT = b°/c° 308 I. Ruzic the following expression can be obtained: i-(Id+J0'/Dd/Dc w /w')/(l+^YDd/Dé w /V) y, (id+J^-/Da/Db. z' /ZK)/(l+^Da/Db z' /zK)-l *p HjfD /D (0'+l)w/w' 1+^-Jb /D z'/zK (3) = CJ 1+^'Yb /D w /w' 1-jJb /D ($# +l)z'/zK d c * d c From eqn. (3) the expressions- for the anodic (i ) and cathodic (i ) limiting currents follow as: i = (I +Jf6'VD /D w /w')/(l+(ZSyD /D w/w') J = c^Ac'HD./d d D c d*) (4) log(concn. inact. spec) Fig. A schematic presentation of the half-wave potential shift. The range of validity of Kern's approach (A), Koryta's ap proach (B) and of the method proposed here (C). I. Ruzic 309 and = 2 *c ^d^^T^b 'AK)/(l+^^7D"b z'/ZK) (4) j = b*i_/(b*+ D /D, a*) d a b From the eqn. (3) the ratio /($£/ P)presenting the shift of the half-wave potential can be written as: *¥(£/ P)= exp(nFAEl/2/RT) =(Wp.ft/£)/(tf / P) (5) wheref and Cj describe the influence of a preceding and a fol lowing ichemical reaction respectively. From the eqn. (5) the explicite relations for the calculation of k, k',^, andj^ rate constants were evaluated, while from the limiting currents only the ratios fk/k* and y/fj could be obtained. , The method presented here is more general than those of Koryta and Kern J (Fig. 1.). The limitations of the proposed method are: a) the solubility of the electroinactive forms participating in the reaction mechanism and b) pseudo-monomolecular character of the chemical reactions. REFERENCES: 1. Brdicka R. , and Wiesner K, . Coll. Czech. Chem, Comm, 1£ (1947) 138, 2. Brdicka R. , and Wiesner K. ,Coll. Czech. Chem, Comm, 1J2 (1947) 212, 3. Wiesner K. , Chem listy 41^ (1947) 6. 4. Kern D. M. H., J. Am. Chem. Soc. 15 (1953) 2473. 5. Koryta J., Coll. Czech. Chem. Comm. 24 (1959) 2903. 310 Sigmund Schuldiner THE EXCHANGE CURRENT DENSITY VS. CONCENTRATION RELATION AND ITS USE IN A RIGOROUS DETERMINATION OF SOLUTION PURITY Sigmund Schuldiner and Murray Rosen Electrochemistry Branch, Naval Research Laboratory, Washington, D.C., U.S.A. By lowering the concentration of a reactant or product (or both) of an equilibrium exchange, the effect of impurities on i0 can be de tected when the concentration of these impurities are significant in relation to the equilibrium potential of the reaction being studied. Thus the concentration of reactant and/or product that will be af fected by impurities can be determined. Interfacial changes in concentrations of reactants and/or products by polarization can give low enough concentrations so that effects of impurities on the overall kinetics should then be expected. In addition, if the im purities (either solution or electrode) influence catalytic processes, then the minimum concentrations at which equilibria are maintained gives a measure of the point at which impurities will significantly affect such catalytic processes. If a second equilibrium exchange below a given concentration of either O or R is suspected, further reduction of the other component should indicate whether such a change is due to impurities or not. For equilibrium surface processes that involve intermediates ad sorbed on a surface that control the rate of exchange, the exchange current density vs. concentration relation must, as shown by Par sons (1) for H atoms, be corrected using the proper adsorption isotherms. For such processes the fraction of available sites covered with the reactive intermediate may be estimated. In ad dition, the mechanism of the intermediate exchange may be indi cated. Experimental confirmations of this technique of determining solution purity have been published from this Laboratory (2,3) in the course of work related to the equilibrium hydrogen electrode. That work showed that for Pt/H2 and Rh/H2 the exchange current density was directly proportional to the hydrogen partial pressure. Assuming that the equilibrium exchange on the hydrogen electrode (Pt and Rh) involves H atoms as an intermediate and that the com bination/dissociation of H atoms and H2 molecules is the mechanism of H/H2 exchange Sigmund Sch.uldin.er .. 311 + 2H + 2e + 2Pt = 2Pt-H = 2Pt + HL,2 * Parsons equations (1) show that for the Langmuir adsorption iso therm (that holds for the equilibrium surface coverage, 6Q, being either « 1 or « 1) that a direct proportionality between i0 and PH 2 holds only when d^ <^ 1. When 0Q » 1, iQ would be directly pro- 5 portional to the (PH ) °* . As can be furtner shown by Parsons (1) equations, the assumption of a Temkin isotherm, applicable be 5 tween 0O = 0.2 to 0.8, would give an i0 dependence on (Pg ) °* for the combination/dissociation case. * For both the discharge and ion + atom control cases, the i0 for the Temkin cases would be inde pendent of Pjj . For the Langmuir cases of discharge and ion + atom control72the Pjj dependence would not be to the first power. The derivations given in this paper confirm that the mechanism for + the equilibrium exchange of H and H2 on both Pt and Rh include an exchange of adsorbed H atoms and H2. The evidence that only a small part of the total amount of hydrogen associated with the sur face is active adsorbed H atoms was the reason for the postulate previously proposed (4, 5, 6) that, in effect, the bulk of H associated with the reversible hydrogen electrode on a Pt metal is alloyed in the dermalayer. Only a relatively small number of active H atoms are adsorbed on the surface of the Pt-H alloy. References (1) R. Parsons, Trans. Faraday Soc. 54(1958) 1053 (2) S. Schuldiner, B.J. Piersma, and T.B. Warner, J. Electro- chem. Soc. ^13(1966) 573 (3) T.B. Warner and S. Schuldiner, ibid U2 (1965) 853 (4) S. Schuldiner, ibid 107 (1960) 452 (5) C.H. Presbrey, Jr. and S. Schuldiner, ibid 108 (1961) 985 (6) M. Rosen and S. Schuldiner, ibid 117 (1970) 35 312 Donald E. Smith SOME INVESTIGATIONS OF EXPERIMENTAL METHODS FOR RAPID ACQUISITION OF FARADAIC ADMITTANCE FREQUENCY RESPONSE PROFILES. Donald E. Smith, Samuel C. Creason, and John W. Hayes Department of Chemistry, Northwestern University, Evanston, Illinois, USA, and Department of Chemistry, The University of Sydney, Sydney, New South Wales, Australia» Very recently, several new experimental concepts have been ad vanced which enable rapid, automatic generation of the faradaic admittance frequency response profile(l-5). These techniques reveal the faradaic admittance vs. frequency response by the following means: (a) Fourier transformation of the transient response generated by an input pulse or step function ("Transient Admittance Analysis"); (b) Fourier analysis of the fundamental harmonic cell response signals resulting from simultaneous application of a set of discrete sinusoidal waveforms whose frequencies are related by non-integer coefficients ("Admittance Analysis in the Non-coherent Wave Frequency Multiplex Mode'")*' (c) Fourier analysis of the fundamental harmonic cell response signals resulting from simultaneous application of a set of dis crete sinusoidal waveforms whose frequencies are related by integer coefficients ("Admittance Analysis in the Coherent Wave Frequency Multiplex Mode"); (d) Fourier analysis of the cell response resulting from application of a broadband, nondeter min- istic signal such as white noise ("Noise Response Admittance Analysis"). In addition to Fourier analysis, each method relies heavily on digital data acquisition and analysis, preferable with the aid of small on-line laboratory digital computers (minicomputers). These tools allow one to utilize for calculation of the faradaic admittance the very general frequency domain relationship, if(«)E »(a))- Af(a)) = E(co)EVo)) {1) where Af(u) is the faradaic admittance, If(co) is the faradaic current, E(co) is the applied potential, and E*(u>) is the complex conjugate of the applied potential spectrum. Eqn. 1 is applicable to any generalized waveform and is the basis for data analysis with all above-mentioned forms of "Fourier Transform Faradaic Admittance Measurements." The basic measurement concept in question is quite old, but its application to nonsinusoidal wave- Donald E. Smith . 313 forms has been impractical until the recent advent of minicompu ters, together with discovery of the Fast Fourier Transform algorithm(6). Although all represent an improvement over conventional single-frequency sinusoidal measurements for acquisition of the faradaic admittance frequency response profile, each of the fore going techniques possesses a variety of advantages and disadvan tages relative to the others. For example, the techniques differ with regard to: (a) influence of far adaic nonlinearity; (b) applicability of the Fast Fourier Transform algorithm; (c) susceptibility to certain forms of measurement error associated with digital signal processing, such as the so-called "sampling window error"(7); (d) relative availability of signal sources; (e) applicability of applied signal spectrum. These and other relevant factors are sufficiently, numerous that a priori selection of the particular technique which offers the best combination of speed, accuracy and bandpass in acquiring faradaic admittance-frequency data is not clear-cut. Consequently, we have undertaken a detailed exper imental comparison of the measurement alternatives to obtain some empirical evidence to supplement guidelines provided by information theory regarding their relative merits. Each technique mentioned above has been implemented in our laboratory with an experimental set-up comprised of signal source, potentiostat, auxiliary preamplifiers and an on-line Raytheon 704 minicomputer with associated analog-digital inter facing. The computer effects digital data acquisition of applied potential and cell current time domain waveforms and converts them to their equivalent frequency domain representations by Fourier Transformation. Equation 2 is then invoked to obtain the faradaic admittance. With this system each technique was shown to yield faradaic admittance data with accuracy and precision at least equivalent to that obtained with single frequency measure ments. Techniques (a)-(d) were carefully compared under essen tially identical conditions (same redox system, same potentiostats, same date, etc.) in an effort to assess their relative merits. Our implementation approaches and experiences with computerized versions of Techniques (a)-(d) will be surveyed. Our observations and conclusions regarding relative fidelity, speed, and ease of implementation of each technique will be discussed. BIBLIOGRAPHY (1) A.A. Pilla, J. Electrochem. Soc, 117, 467 (1970). (2) K. Doblhofer and A.A. Pilla, J. Electrochem. Soc., submitted. 314 Donald E. Smith (3) S. C. Creason and D. E. Smith, J. Electroanal. Chem., submitted. (4) B. J. Huebert and D. E. Smith, Anal. Chem., submitted. (5) R. L. Birke, Anal. Chem., 43, 1253 (1971). (6) V. W. Cooley and J. W. Tukey, Math. Comp., 1£, 297 (1965). (7) Hewlett-Packard, Inc. "Fourier Analyzer Training Manual," Application Note 140-0, Hewlett-Packard, Inc., Santa Clara, California, 1970. Elisabeth Steyger 315 ELECTRO-OPTICAL STUDY OF INTERFACIAL POLARISATION AND HYDRODYNAMIC MO TION IN AN ELECTRODIALYSIS CELL. Elisabeth Steyger Laboratoire de Physique Expérimentale, Domaine Universitaire (38) Saint-Martin-d'Heres - FRANCE - A Michelson interferometer fitted with a laser light source is used to visualize, during electrodialysis of sodium chloride aqueous solu tions, the concentration changes due to electric migration and selec tive passage through ion-exchange membranes. The straight parallel in terference fringes observed with an homogeneous solution become curved along themembrane on application of an electric field : the thickness of the distorted fringe area (polarisation layer) is examined as a function of different parameters such as salt concentration, fluid velocity and appliod.voltage. A first electrodialysis call involved two membranes : an anionic mem brane andj a cationic .one, but the differences between their behaviour had led to study them separately. A new cell consisting of four com partments' has been used"extensively : the compartments contiguous to the investigated membrane .(anionic or cationic) are separated from the electrode;;:compartments by acetate, cellulose membranes in order to avoid disturbances due to' electrode reactions ; reference electrodes are in troduced on each side of the membrane t a stabilized, -power.--supply can assign a chosen intensity or voltage ; pumps make the fluid flow at well defined velocities which are measured by flow-meters. A cinematographic observation shows that the interference fringes un dergo turbulent and irregular motion when the fluid is in flow ; this can be attributed to the superposition of two phenomena : the change in concentration and the turbulence of the solution. It is therefore necessary to have a better knowledge of the hydrodynamic flow regime ; this lest one is studied at different velocities by three methods : (i) visualization by coloration : This consists of sending alternati vely water and potassium permanganate (because of its colour) through the cell and of following the motion of couloured trickles of water. This method is only qualitative and the transitory aspect of the ob served phenomena is a great disadvantage : indeed, the propagation of coloured waves can only be observed at the beginning of the experiment and we"allowed to suppose that a regime, which is turbulent at the turning on of the tap, can become more regular latere (ii) visualization by strioscopy : This optical method is based on the dependence of refractive index variations when an object is placed in a parallel beam. A small opaque screem, such as a razor-blade, stops much of the beams but not the light-waves diffracted by the optical inhomogeneities*of the object. In order to observe turbulences when water is in flow these inhomogeneities are made by small temperatures fluctuations artificially produced. A quantitative study, based on the distance covered by a trickle from one frame to the other, can be added to the filmed observation of the displacements of the trickles i the path of a trickle is made up point by point. The lengths between the "*are 316 Elisabeth Steyger points arB proportionnal to the local fluid velocities and allow us to determine the flow regime : indeed» the non-variation of velocities is typical of a laminar flow regime. In a turbulent flow regime the velocities are always changing. The scattering of local velocities around the'mean speBd calculated from the flow-meter indications is therefore significant of the flow regime as is shown by the fallowing graphs, obtained by platting for each interval between points the re- curence rate of that interval : recurrence rate recurrence rate = 4,5 cm/sec = 1,33 cm/sec 0.2 £>.<• 0.6 0.8 1,0 i#2 intervalo,2 o> o ,6 o,e i.o 1.2 i,i» XT lnterX/al in cm in cm (iii) measurement of pressure drop between two points-of the circuit located on either side of the cell : this allows us to detect the ap- pearence of turbulence in this part of the circuit : the variation of pressure drop with velocity obeys a law which is a function of v2 for a turbulent flow regime and a function of y for a laminar one (at first approximation). The change in flow regime is detected by a break in the logarithmic plot of pressure drop with velocity. These three methods are in agreement and lead to the conclusion that for a 16 cm path length the flow regime becomes turbulent, (in a 9 mm gap between parallel planes] when a 2-3. cm/sec velocity is reached. No parabolic flow distribution has been observed ; at lower speeds:strio- scopy has always shown a transitory rather than laminar flow regime. Addition of pH indicators in the cell shows a colour distribution cor responding^ local variation of pH at the membrane interfaces and pro ves experimentally that water electrolysis takes place when concentra tions of Na* and C£~ tend to zero. This makes ambiguous the interpre tation of fringes distortions as a change in the configuration of in terference fringes represents a variation of the optical path, that is to say of the refractive index of the solution. This index is linked to the concentration of the solution as well as with its nature. We have shown, with a series of NaCÄ, solutions at the same ionic strength Elisabeth Steyger 317 but at different pH, that the refractive index varies linearly with the pH and it is then impossible to assign the distortions of fringes uniquely on the concentration changes of Na+ and CI" ions. This leads us to mention that a complete study of the phenomena that occur in the vicinity of the membrane during electrodialysis requires the following simultaneous observations : the current-voltage relationship the local variation of pH the optical measurements of distorted fringes. Many authors describe the ion-tfransfer process through the Nernst-Fick relation : ..-•>••;. This formula shows that if i, D and the transfer numbers are constant. the concentration gradient in the boundary layer is linear. However, this simplification precludes the possibility of water splitting and it also neglects the convection phenomena. The non-stationnary hydro- dynamic flow regime and the participation of H* and OH" in the current transfer show that the Nernst-Fick relation cannot be applied in-the case under investigation. The distorted area really observed does not result from a pure diffusion layer but from a complex diffusion-con vection process. Moreover, it should be possible to relate experimen tally the thickness of the polarisation layer to the investigated pa rameters, but such a relationship has to be compared only to theore tical expressions involving the hydrodynamic flow regime such as the formulas of Sonin-Probstein or Solan-Winograd. 318 Desanka Suznjevic POLAROGRAPHIC REDUCTION OF Ni/II/ IN THE PRESENCE OF ETHYLENEDITHIOLDIACETIC ACID SOLUTION Desanka Suznjevic and Milenko Susie Department of Physical Chemistry, Faculty of Science, University of Belgrade and' Institute for Chemistry, Tech nology and Metallurgy, Belgrade, Yugoslavia. The polarographic reduction of Ni/II/ in aqueous solution containing ethylenedithioldiacetic acid /HpZ/exhibit a pre-wave before the free metal hexaquo complex wave. The E-. yp of this wave pis approximately 0.$0 V more posi tive inan the Ni/HpO/g background wave. The height of the pre-wave is independent of the mercury column height un der condition where pre-wave is- small fraction of the Ni /II/ aquo complex wave. The pre-wave limiting current at this condition is less than 25$ of the total Ni/II/ ion limiting current. The temperature coefficient of the pre- wave limiting current is rather large /5-10$ per degree/. All this data suggest that the pre-wave probably appears from a slow chemical formation reaction on the electrode, and reduction of the Ni/II/-HpZ complex. It was observed that the average value of kinetic current i/ gave straight ^ line with slope __ The pre-wave height is proportional to the total Ni/II/ ion concentration. _ _ The nature of anions investigated /C10^7 N2^~/ had only slight effect on i, • The dependence of logi2 from 1/T in the temperature range from 20-50 C gave a -straight li ne without inflections. The presence of HpZ does not affects the interfacial ten sion /drop time/ of mercury at investigated potentials. From the above results catalytic nature of the ;ore-wave mechanism, involving cyclic regeneration of ligan^during the Ni/II/-mono-HpZ complex reduction on the electrode, was assumed. The rate constant of a chemical reaction pre ceding .the electron transfer step was found from Kouteck^- theory'and from expression evaluated from Hanus investi gations concerning the pre-wave catalytic nature? Those value are: 0 -. n x x From Kouteck^- theory: k,,t=6.5xl0~2 T1 mo_.^i—1_-l s 1 /t/A=. 2orrO5 C// According to Hanus: kt..f== 3.9xK3.9x10r ^ »" " " » The value of k~was also evaluated from voltametry at con trolled current data?^3i|Cerence between this value and polarographic'aly obtained was found when the concentra tion of HpZ was in excess. •Kouteckt J*i Colin Czech.chem.Commun.78, 597 /1953/ 2.Tur»yan Ya.I.,Malyavinskaya,EIektrokhimiya,2,1185/1966/ 319 Jacques TACUSSEL THE "VOCTAN ", AN APPARATUS FOR ELECTROCHEMICAL TECHNIQUES INVOLVING VOLTAGE/CURRENT/TIME RELATIONSHIPS. Jacques R. Tacussel, S. O. L. E. A., 72 - 78 rue d'äl,*ace - F 69 - Villeur- banne - FRANCE. Starting from the AC polarograph (1") first presented at the 19 th mee ting of CITCE (Detroit, 1968) we had initially decided to develop a pulse polarograph for research (as well as analytical) applications. We believed that an important feature to be included for this purpose was the provision to measure the current through the cell under widely adjustable conditions, without the limitations encountered m conventional instruments with respect to the position of the applied pulse during drop life, and of the current sam pling window within that pulse. In fact, it appeared necessary to give the user complete freedom for the choice of a number of parameters : imposed drop time ; position, duration, amplitude and polarity of the pulse ; position and duration of the current sampling window. At this stage of the development, the instrument available was a very ver satile pulse polarograph, with several additional refinements (mainly inten ded for operation with a constant amplitude pulse superimposed on a DC ramp) such as a differentiation circuit for recording the curves Ai/AE, and a "chart speed controller ", a means of obtaining in a reduced time polaro- grams with sharper and higher peaks. For the latter, a special circuit was included, to increase the speed of scanning during the horizontal parts of the diagram, and to decrease it, in proportion to the amplitude of the peaks, when tracing parts of the curve where these are located. Another important feature was the possibility of displaying the current/voltage relationships either with a chart recorder or with a (memory) oscilloscope ; in the first case, the current measurements could be made in a "punctual " manner (sampling time can be no longer than 100 MS) or by integration during a longer time and on up to nine drops (the integration method giving a better signal-to-noise ratio and an increased sensitivity). Oscilloscopic display was interesting when dealing with relatively fast scanning rates (e. g. from 50 mV to 100 V. s"1) as in cyclic or programmed voltammetry. It then became evident that, using the electronic circuits and concepts employed in the instrument built-up as a prototype, it was possible to large ly exceed the field of conventional or even advanced pulse polarography, by introducing a technique not used until now, to our knowledge, in electroche mistry, which we suggest to name programmed sampling . This technique is suitable to investigate any repetitive phenomemon and is specially useful for those having a duration of say, some tens of microseconds to some se conds (for, instance : cyclic voltammetry curves, double-layer charging cur ves, oscillopolarographic curves, etc ). It is derived from the current mea surement technique used in pulse polarography ("punctual " current sampling) but differs from it in that the instant of sampling is progressively shifted 320 Jacques TACUSSEL along the time axis of the curve to be analysed, instead of being fixed at a definite position during the drop life, after the application of the superimpo sed pulse. This offers some analogy with the technique used in "sampling oscilloscopes " for visualizing phenomena up to frequencies well beyond the limits of amplifier bandwidth. In our case, the problem was not to improve the frequency response, as satisfactory results can be obtained with commercially available oscillos copes fitted with appropriate plug;-ins, but to greatly improve signal to noise ratio and increase precision over that normally obtainable with oscilloscopic readout. The system which we have developed combines the advantages of large bandwith oscilloscopic measurements with the precision and ease of usage of chart recording. The time analysis of the curve is effected by com paring a voltage, reflecting the real time scale'of the studied phenomenon, with a slowly varying scanning signal (in most cases, a triangle wave) sup plied by an auxiliary generator, the current sampling gate being operated each time these two voltages become equal ; by a proper choice of the scan ning rate, any phenomenon even lasting no more than a few tens of microse conds, can be expanded on several tens of centimeters of paper chart. The further step has been to use the capabilities of that apparatus for solving other problems where voltage/current/time relationships are invol ved, outside the field of polarography and voltammetry. These include chro- nopotentiometry, chronoamperometry, coulostatic techniques, etc. We be lieve that the instrument described (VOCTAN = voltage/cur rent/ time Ana lyzer ) affords to the electrochemist a valuable means of applying numerous methods ; some practical examples are given in another paper (2 ). The most important techniques which have been experimented using the "VOC TAN " are listed in the following table (see next page). REFERENCES CD J.R. TACUSSEL : A versatile apparatus for AC and DC polaro graphy, and drop-time measurements. - 19th Meeting of C.I.T.C.E. DETROIT, 1968, Extended Abstracts pp. 189-192. C2) J.R. TACUSSEL and J.J. FOMBQN : Applications of programmed sampling pulse techniques in Electrochemical analysis and reserach. Paper to be presented during the present Meeting. Jacques TACUSSEL 321 IMPOSED SIGNAL RESPONSE OBTAINED TECHNIQUE INVOLVED E=f(t) i = f(t) 1) Chronoamperometry. .31 L 2) Impedance measure 4 / ments by pulse 1)0.1 to 10s techniques. 2) S to 100 us 3) Chronocoulometry. 3) 0.1 to 10 s V E=f(t) 1) Pulse polarography. 2) Pulse tensammetry (both with superim AE 0.1 to 99s 5 to 50 ms posed pulses). <= 54 l-s— Er:f i = f(t) Erf (l) Pulse polarography with increasing 0.1to99s 5 to 50ms K >t amplitude pulses. Derivation of the curves can be effected electronically. i=f(E) E=f(0 Cyclic voltammetry. When using DME, the imposed signal is ap plied with an adjustable delay from the fall of the drop. |c 0.11 o 99 s ^i 5ms^Jo10s i=f(t) E=f(t) 0.1 to 99 s 0.1 to 10 s ; ± 5*1 1*=- ill Wi C hr onopotentiom etry. 322 Marijan Vukovic CYCLIC CHRONOPOTENTIOMETRY. DIGITAL SIMULATION AND EXPERIMENTAL VERIFICATION OF MODEL SYSTEMS OF HIGHER ORDER CHEMICAL REACTIONS COUPLED WITH ELECTRON TRANSFER Marijan Vukovic and Velimir Pravdic Electrochemistry Laboratory, "Rudjer Boskovic" Institute, Zagreb, Yugoslavia (1 2 3) Cyclic chronopotentiometry (CCP) has been extensivelyv ' ' ' used to study reduction-oxidation reactions involving coupled chemical reactions in solutions or at electrodes. CCP seems one of few methods capable of rendering information on the kinetics of higher order coupled chemical reactions. Digital simulation' ' of electrochemical processes serves as the theoretical formalism which is being fitted to working curves in coordinates of relative transition times CCn:t,=an) vs. number of cycles (transitions). Two cases are treated here: the coupled, following third order chemical transformation, and the preceding second order complex formation producing the electroactive species. The reduction of Ni^+ on the Hg pool.electrode in 0.1M perchlo- rate solution, pH=2. 0 fits the model of irreversible, following, third order reaction at lower, concn. of Ni + (0.45 and 0.90 mM). At 1.80 mM deviation has been observed, but deemed due to blocking of the electrode by the deposition of a Ni-Hg intermetal- lic compound. Figs. 1 and 2 show the quality of fit. Deposition of Ni on the surface is ruled out, because the experimentally found value of a£ is below that for a diffusion controlled process, and not above. The- experimental values deviate also from the model of two paralel second order reactions obtained by digital simulation. The apparent third order rate coefficients for the probable intermetallic compound formation following reduction of Niz+ in the mentioned» solution are: 0.45 mM Ni : k=1.4 x 10 ; 0.90 mM: k=l.l x 10 ; 1.80 mM: 3. 8 x 105 M"2sec_1. The reduction of Ti(IV) in thiocyanate media is preceded by a second order complexation reaction. Solution conditions were chosen such to allow the second order kinetics to prevail: the concentration of SCN was kept ten times the concentration of Ti(IV). Fitting of experimental curves by digital simulation is tedious: one has to chose discrete values for kf and kb respecting the order of magnitude (at least) for the equilibrium constant K from the literature. Although the single pulse chronopotentiometry plot of iQX vs. i0 ' gives a straight line indicating 2 as a number of the species involved in the reaction, and although the value for K=50 M seems reliable, CCP analysis is still short of giving reasonable values for kf and kfc. Marijan Vukovic 3.23 —i—r~ 0.1 M NoClOt 1X01 M HQOt x 45XW* M Ni2* 09 ? O aOxtt"* M Ni * IB x 10'3 M Y |2+ 0.1 M NaCIOt a7 001 M HCICi' IR» 4iax 0.5 UJ 0.6 3E X z 04 o UJ X zto P 02 < 1.1 Z te g *— UJ a9 z < > ar UJ 0.7 IL 0.6 ve > •- os < 0.4 _j UJ a3 a. 0.2 X 1 4 6 B 10 12 14 4 6 8 10 12 NUMBER OF CYCLES NUMBER OF CYCLES Fig. 1. Fig. 2. Fig. 1. Relative transition times vs. the number of cycles for the reduction and oxidation of Ni "*" in perchlorate medium. Full lines: theoretical calculations for a third order following reaction. Dashed lines: diffusion controlled process with no kinetic compli cations. Ratio of forward to reverse current density, i /i^^.=2. Fig, 2. Same for a different concn. of Ni and current density ratio, i /i =4, ' R' OX * References 1. H. B. Herman and A.J. Bard, Anal. Chem. , .36 (1964) 510 2. H. B. Herman and A.J. Bard, J. Electrochem. Soc. ,115 (1968) 1028 3. M. Vukovic and V. Pravdic, Croat. Chem. Acta, 42 (1970) 21 4. S, Feldberg, Electroanalytical Chemistry, Vol. 3, A. J. Bard ed, Marcel Dekker, Inc., New York, 1969, p. 199 SECTION 4 HIGH TEMPERATURE ELECTROCHEMISTRY Y.K. Delimarsky 327 ELECTROCHEMICAL PROCESSES WITH RATE-DETERMINING PRECEDING ACID-BASE REACTIONS IN MOLTEN ELECTROLYTES Y.K. Delimarsky, V.I. Shapoval, V.F. Grishchenko Institute of General and Inorganic Chemistry, Kiev, USSR For equilibria of the type (1) 2 a and others the importance of the acid-base properties of molten electrolytes for electrochemical kinetics has been shown. The rate- determining role of Reaction (1) in the following cathodic reaction MOn + ne—~ fl +£0*" <2) X has been proved. For the case of stationary polarisation such processes are rather difficult to study because they involve a self-inhibition resulting from accumulation of oxygen ions. In this case the limiting kinetic current is given by (3) l„ where $jt/|fO1 K = diffusion coefficient of electrochemically active species; tjå.*. = diffusion rate constant; K = dissociation rate constant ofthe "base"; K = stability constant of the "base"; C = activity (concentration). The polarogram equation also in dicates an uncommon relation between current and potential: L* y> = const- BI-^n (A) ctf^F L*~C* where G&-*= charge-transfer coefficient; OQ = number of electrons involved in the potential-determining stage. The inhibition can be reduced by addition of an "acid", e.g. P0~ to the melt. Polaro- graphic behaviour of W0~ in fused KCl-NaCl eutectic was experimental ly investigated. Cases have been found where the application of 328 Y.K. Delimarsky Equations (3) and (4) is possible. Attempts have been made to appro ximately calculate the,kinetic parameters and life time of reactant species when the rate of the process is -4 *>- T* -5 10 cm/sec ( £w0f ^5 1 sec; L^0j —» 10. sec). The application of nonstationary methods (oscillographic polarography, chronopotentiometry) permits to qualitatively identify the retardation of the acid-base reaction. The kinetic parameters can be quantitative ly estimated by chronopotentiometry. The investigation of electrolytic 2- 2- reduction of CO^ in KCl-LiCl, WO, in KCl-NaCl with excessive P0 showed that experimental relations cannot be described by the known Delahay-Berzins equation. For the case of very slow Reaction (1) (kinetic process) the diffusion of electrochemically active MO "2 species is suggested to be a preferred one. Relationships describing experimental data are discussed. Using them, the effect of tempera ture and P0Ö additives on the stability constants and kinetic para meters of preceding acid-base reactions has been studied ( %~ ~ lO^-lO"2 sec; T* ,- ~ 10_:L-1 sec at 400-510°C) . co2, CO3 Pierre AURY 329 ELECTROLYSIS OF PLUTONIUM AND ITS ALLOYS Pierre AURY Commissariat å l'Energie Atomique -BP 14- Is-sur-Tille - FRANCE - We practise electrolysis of Plutonium and its alloys in a double pur pose : 1°) - to prepare pure metal. 2°) - to recycle alloys and impure metal became unsuitables- for cer tains uses byithe ratio of impurities they contain. After having;applied the process in solid phase,(said with soluble anode) now operation is led in liquid phase.at 750° C. (Process said, "with melted;anode"). Electrolyte is a fused salt bath Kcl-Nacl mix ture with addition of Caf2 and PUF4. The apparatus is drawn of this used by LEARY & MULLINS from Scientific Laboratory LOS ALAMOS"(Univer sity of CALIFORNIA). The whole is shown in Fig. I in which may be no ticed : .1°) - The "inactive part", that's to say the part which is nöt conta minated by Plutonium and which is constituted by the furnace under the glove-box. 2°) - The "radio-active part", that's to say the electrolystical cell placed in the glove-box, in an argon atmosphere in consequence of the toxic radioactive and pyrophoric properties of Plutonium. Fig. 2 give the disposition of this cell. From the outside to the inside are met successively : 1°) - A stainless steel tube doubled to form a tight enclosure with a contact-equiped manometer. So is constituted a double security to pre vent either the failure of tho temperature regulator (high-alarm) or the failure of the furnace itself (low-alarm). 2°) - A tantalum container which in case of breaking of ceramic-cru cible avoid the contact between melted Plutonium and stainless steel of the tube indeed stainless steel and Plutonium form a low melting point alloy (414°C). On the other hand at 750°C. the solubility of Tantalum in Plutonium is about 650 ppm. 3°) - The proper electrolytic crucible is made of magnesia with 3 % of Yttrium oxyde, addition which permits a highly vitrified material free of any porosity. 4°) - The mobil part of the fitting including (Fig, 3) : - the current leading rods in tungsten (5), protected by magnesia . beads (6). - a stirring rod equally in magnesia (7) able to turn up to 1O0O t/mn. Geometry of the crucible and the mass of metals imploy an electrolytic current of about 20 A under a 4 volts voltage applied between elec trodes. According the third FARADAY low, this intensity leads to a 60 g/h deposit. An electronic device permits to switch out rapidly the applied voltage and to measure the back e.m.f. 200 us later. Frequency of this measure depends ei the evolution step of electrolyse. Operation is automatical ly stopped when back e.m.f. reaches the prechoiced value marked on the piloting "time counter. This value is function of the desired degree of purity of Plutonium. Initial metal may include from 2000 to 600O ppm in impurities (up to 2,5 % in case of an alloy) regards to firty tes ted elements ; after electrolyse this proportion turns about 200 ppm 330 Pierre AURY for a so called pure metal, 800 ppm for än intermediary quality. Next board gives for the principal elements the best degree of purification obtained : Element Al C Cr Cu Fe Ga N Ni Pb Ta U W Feed(ppm) 400 470 220 70 5500 14000 550 10000 150 350 13600 300 Product 10 20 . 2 15 <2 <10 10 5 <30 55 <10 (ppm) <1 General yield must be considered as the result of two partial yields : 1°) - The yield of "dissolution", that's to say the ratio : mass of Plutonium dissolved in thebath reported to the mass of Plutonium in cluded in the initial anode. For a given anode, we keep the control of this yield in the measure we can choice the' moment for stoping electro lysis. Accord the researched aim we sacrify yield to purity or inver sely. In the first case yield is about 90 % in the second it exceeds 95 %. We can overpass the theorical yield of 91 % given for a 1 % gallium alloy. At this stade, the begining solidification of the anode scrap, wich causes the breikage of the stirring rod. We must proceed to its exchange. So, we draw the solid part of the anode;, so doing, the most part of the impurities. We continue operation after having ligthy increase the temperature of the furnace. Eventually wo proceed to a partial reloading of the ano dic hole with impure Plutonium or alloy. 2°) - The decanting yield, wich is the ratio between the mass of Plu tonium recuperated at the cathode and mass of dissolved Plutonium. It may reach 95 % when all opportunities are resembled, that's to say : - to maintain a continuous argon circulation in the electrolysis cell, this one being primarly dried on molecular sieve which prevents oxyda- tion of Plutonium droplets dispersed in the bath and so make their coa- lesung easier. - to increase the temperature of the bath at the er.d of the operating so, to make it more flowing and to stir it frequently in the post hea ting stage. At last, in this time, we proceed to an addition of fluo- rite in electrolyte. Let us remark, that in direct consequence of the equipment conception cathode consists in melted Plutonium itself, so to say the metal is maintained to cathodic potential with the effect to embetter the decan ting output. So, the total output may vary from 80 to 90 %. The catho- dic-tore gathered after cooling, is conditionned by vacuum fusion. In terest of electrolysis can't be contested, when pure Plutonium is aimed at. For the preparation of intermediary quality metal, this way yet pre sents, compared to classic chemical way, real advantages and chiefly : purity of so obtained metal, quickness of the process, at last lower cost (about 70 % of the other). Automation of the apparatus reduces in a large scale, the personnel expenses. Pierre AURY 331 c~> && 4 Kg. S. 9 4- r&a 332 W.A. Fischer FACTORS INFLUENCING THE EMF MEASUREMENT WITH SOLID ELECTROLYTE CELLS IN LIQUID METALS W.A. Fischer and D. Janke Max-Planck-Institute fiir Eisenforschung, Diisseldorf, Germany High-temperature solid electrolyte cells for direct determinations of the oxygen potential in liquid metals have become a subject of global interest. The experience of numerous laboratory experiments is that reliable EMF values can only be obtained if certain fundamental sources of error are detected and as far as possible eliminated. The research work performed at the Max-Planck-Institute in Dusseldorf can be characterised by the development of two different types of oxygen concentration cells. 1) For the continuous registration of oxygen activity in liquid metals Zr02~CaO tubes closed at one end have been used with an inner Pt- or Ir-gas electrode. An outer lining of porous Zr02-CaO diminishes the thermo-shock sensibility of -the tubes and prevents metal oxides from diffusion into the electrolyte tubes at higher oxygen levels of the liquid metal. Cracks and micro-cracks in the electrolyte material can be avoided by heating the Tubes simultaneously with the metal in the furnace or by electrical preheating. Sufficient electrical contact of the inner.electrode will be achieved by sintering the Pt or Ir wire in a layer of Zr02 powder or by using Pt net electrodes. At high temperatures the cell voltage can be affected by a selective oxygen permeability of the electrolyte. A layer rich in nitrogen at the air ZrO~ interface and a layer rich in oxygen at the Zr02 liquid metal interface are resulting from the oxygen permeation through the solid electrolyte. These gradi ents of gas diffusion will be removed by maintaining a sufficient air flow in the tube and an intensive stirring in the melt, re spectively. 2) For single oxygen determinations in liquid metals quartz tube cells containing a Zr02~CaO plug as the electrolyte and a metal - metal oxide mixture as a reference electrode are applied. These cells may be introduced into the metallic, melt without be ing preheated. When used at high temperatures, however, chemical reactions between the zirconia, the metal oxide and the quartz tube, and also between the quartz tube and the oxygen containing metal give way to a rapid destruction of the oxygen sensor. The application of those cells is therefore restricted to short measur ing periods. André Fontana 333 CALORIMETRIC STUDY OP THE HEAT OF MIXING OF NaF and ZrF4 AT 1030°C André Fontana and René Winand Université Libre de Bruxelles, Service Métallurgie- Electrochimie, 50 av. F.D. Roosevelt, B-1050Bruxelles - Belgique. Previous work done in our laboratory shows the existence of high overvoltage in fused salt electrolysis. It was assumed that this phenomenon is caused by the presence of very stable complex ions in the melt. This paper concerns the first part of a study of the stability of complex ions in fused salts : the determination of the heat of mixing of molten fluorides by a calorimetric method. We built an isoperibol drop calorimeter able to work up to 1100°C under an inert atmosphere. The calorimetric cell contains a graphite crucible (where the studied reaction is evolved) inside of an isothermal nickel block. The heat devellopped in the crucible is measured by the difference of temperature between the crucible and the isothermal nickel block. This difference of temperature is measured by means of twelve chrome1-alumel thermocou ples with a differential voltmeter Hewlett Packard type 3420 A. The temperature of the nickel block is regulated by a Honeywell P.I.D. regulator (+ 0,5°C). The sensitivi ty of the calorimetric cell is 0,04°C by centimeter of the recorder, the stability being better than 0,004°C during an experience. At 1000°C, the constant of the calorimeter is 17,15 cal/°C, which causes a deviation of 10 mm of the recorder for an absorption or emission of heat of 0,688 cal in the graphite crucible. The calorimeter has been calibrated between 500°C and 110O°C by additions of silver (99,99%). By additions of small quantities of NaF (substance 1) and ZrF4 (substance 2) at 25°C in a fused mixture of NaF and ZrF4 of a given composition (substance 3, between 0 and 25 mole % ZrF4) at the temperature T, we measure absor ptions of the heat respectively of 1 3 2 3 2i - H* - Hie •+ pm. " and 22 - H$ - H$s + SL " The absorptions of heat comprises the heat content of the addition between the temperature T and 25°C (HT - H25) plus the partial molar enthalpy of mixing _of the addition in the fused salt at the temoerature T (AH^~3 and AH extrapolation of the curve of the heat content versus temperature up to 908°C, by estimation of the heat of fusion of ZrF4 at 908°C and of the heat capacity of li quid ZrF4 from 908°C to 1030°C, vre find a heat content of liquid ZrF4 at 1030°C of 36.300 + 6000 cal/mole. By the Bakhuis-Rooseboom method, v;e calculate the total molar enthalpies of the fused salts (AH^) (Fig 1). Then, ve determine the heat content of the fused mixtures at 1030°C so that it is possible to calculate the molar en thalpy of mixing of the mixtures at room temperatures. We find molar enthalpies between -1,5 and -8,5 kcal/mole at 25°C. Then, based on litterature data.concerning the vapor pressures of NaF and of ZrF4 above these fused mixtures, ve calculate the free energies AGiJ , -the entropies AP^ and the excess entropies of mixing (AS|)XS at 1030°C by the follv."ing relations. The partial molar free energies of NaF and ZrF4 in the fused salts are given respectively bv 3 3 AG^" = RT In a2 and AG^~ = RT In a2 Pi P? vzith a-i = S and a o = —— 1 TTl 2 7C2 The partial pressures pi and P2 of NaF and ZrF4 above the fused mixtures at 1028°C were measured by K.A. SENSE et aid). The molar free energies å0>J of the fused mixtures are calculated by the Bakhuis-Rooseloom Method. The molar entropy of mixing &S% are calculated by the relation : AG§ = AH£ - TAS§ (Fig 2) The excess entropies of the mixtures are given by the relation (AS^)XS =As£ - (6S3)_ideal These results shov; that the formation of mixtures of NaF and ZrF4 is highly exothermic. We measure heats of mixing as high as 7 kcal/mole for 25 mole % ZrF4 at 1030°C This could, possibly explain overvoltages as high as 100 mV during the electrolysis of fused NaF-ZrF4. They are pro bably due to the free energy of mixing of the fused salts. We are nov.r completing the results by e.m.f. measure ments of galvanic cells and studying the structure and the stability of the complex ions by absorption spectrophoto metry at high temperature. (1) K.A. SENSE, C.A. ALEXANDER, R.E. BOWMAN, R.E. FILBERT, J. Phys. Chem. 61 (1957) 337. André Fontana 335 10-: kcal/mole 5- mole /Q ZrF4 •Tr- —i— IO 15 20 25 fig.1. Enthalpies of fused NaF-ZrF4 mixtures at 1030°C. \ 2" cal/°K.mole ( ÄST)icjéal mole /o ZrF4 O —i— •T— -r 0 5 10 15 20 25 fig. 2. Entropies of fused NaF-ZrF4 mixtures at 1030 C. 336 D. Janke ELECTROLYTIC DEOXIDATION OF LIQUID METALS W.A. Fischer and D. Janke Max-Planck-Institute fur Eisenforschung, Diisseldorf, Germany The open-circuit EMF of an oxygen concentration cell is determined by the gradient of the oxygen potentials at the two interfaces of the solid electrolyte RT E = -v=r (In P_ (reference electrode) - In P (liquid metal)). The cell reaction represents the transport of ionized oxygen from the interface of high oxygen potential to the interface of low oxygen potential. When an external direct voltage is applied this reaction will be initiated. The direction of oxygen transport can be determined arbitrarily by the choice of polarity. A charge of 3.35 Ah corresponds to a transfer of l.g of oxygen. If the electro lytic cell described above is supposed to exhibit accurate per formance the externally applied voltage should not exceed the de composition voltage of the solid electrolyte. By means of a ZrO^-CaO tubular cell corresponding to an oxygen con centration cell Pt, air || Zr02 - CaO || liquid metal the electrolytic deoxidation of liquid iron, cobalt, nickel and copper at 1600 C and of liquid silver at 1400°C was examined. Oxygen minimum values of 10 to 30 ppm have been attained by this technique. Deviations from the theoretical oxygen transfer rates were observed because a slow reoxidation of the liquid metal had to be taken into account resulting from oxygen pick-up from the gas phase. Accurate performance of the electrolytic cells was proved over a period of several hours. Higher rates of oxygen transfer can be realized by using a cell with an enlarged surface or when several cells are applied simul taneously. The main advantage of electrolytic oxygen removal is based on the fact that the deoxidized metal is free of oxide inclusions. Sergey Karpachov 337 THE DETERMINATION OF THE POTENTIALS 03? THE POINTS OF &ERO CHARGE ON LIQUID IN CONTACT WITH A SOLID EEKJTROLYTE Sergey Karpachov Institute of Electrochemistry of the Urals Branch of the Academy of Sciences,Sverdlovsk,USSR. The determination of the potentials of the points of zero charge for- different metals in different electro lytes is very important for electrochemical kinetics and for the problem of electrochemical potential (Prof A.Frumkin).Therefore an attempt has been made to determine the potentials of the points of zero charge on liquid silver and tin in a solid zirconia-lime electrolyte.ZircoMa oxyde was stabilized by addition of the oxydes of ittrium and calcium.The little drop of liquid metal was placed on the plate of a solid electro lyte. On this plate there were two. platinum electrodes.. One was for a polarization of a metal drop and another- for measurements of its potential.All experiments were xarried out at the temperature 950°C, in the atmosphere of helium with traces of oxygen.At first we investigated the dependence of the cosine of a contact angle on the potential of metallic drop.At the potential of the paint of zero charge the cosine of a contact angle has its maxima value. It seems to us,it is the direct method to obUain the potential of the point of zero charge in the case- of solid electrolytes.Then was measured the active resis tance between the drop of a liquid metal and the plati- numelectrode which was situated on the back side of the plate of a solid electrolyte.This active resistance has maxima value at the potential at the point of zero charge because at this potential the area of the contact of a metallic drop with a solid electrolyte has its minima Value.The results of these both methods are in satisfacto ry agreement. At last we investigated the dependence of the capacity of a double layer on the electrode potential.The corres ponding curve has rather sharp minima. The potential of this minima is in agreement with the potential of zero charge.Such situation takes place in dilute aqueous so lution and it was explained. (Prof A.Frumkin and. collaborators) The reason of this coincidence in the case of solid electrolytes is not clear yet. 338 J.A.A. Ketelaar ON A UNIFIED E.M.F. SERIES IN MOLTEN SALTS AND OTHER SOLVENTS. J.A.At Ketelaar, Laboratory for Electrochemistry, University of Amsterdam, Amsterdam, Netherlands. In water at 25 C a unique series of standard electrode potentials can be set up with as reference electrode the standard normal according to Nernst convention. f* both are in their standard state, that is in the state of infinite dilution, the values for cations are in dependent of the nature of the anion and for anions independent of the cation respectively. Diffusion potentials should be eli minated e.g., by using a suitable secondary reference electrode. At other temperatures however, the reference electrode by defi nition is always the standard normal hydrogen electrode (at 1 atmosphere hydrogen partial pressure).' The same holds for the definition of standard electrode potentials in other solvents, including mixtures and solutions. Attempts have been made by several authors, Milazzo, de Ligny, Schwabe et al., to obtain unified series by creating a relation between the potentials of different reference electrodes. The measurements involve either non-isothermal cells (thermal cells) or isothermal cells with a contact between two different solvents.In both.^cases this means that a transfer of ions is present from one temperature to another tenperature in the same solvent or from one solvent to another solvent at the same tem perature. Theoretical .considerations and extrapolations e.g., to infinitely large ions have to be made.to obtain results. With thermal cells the difficulty arises that princi pally^dE^dT^ 4 dE/dT), _ and thus additional conventions or- supposi tions have to be invoked. In the case of the E.M.F. series in molten salts the same diffi culties arise to an even larger extent. Moreover for the E.M.F. of metals in their pure salts as many series arise as their are anions. In the case of either pure salts or of dilute solutions of the metal in another salt or salt mixture, the obvious choice of the primary reference electrode is the anion formation-electrode e.g., CI /Cl~, N02 + £02/N0~ S0$ + 02/So£~, CO + ^/CO^" etc. An effort has been made .by Boxall and Johnson (f) to refer all E.M.F. series in molten salts to the normal hydrogen electrode at 298 K. This involves',-a system of precise measurements they have made or used on both thermal cells and on cells with differ ent "solvents" e.g., different anions with transport. In principle both the -problems of E.M.F. series of molten salts and those with different solvents at ambient temperatures are formulated by the question how to refer the potential of the half-cell M/MZY at T to that of the half-cell M /MZX at T with X,Y beeing either different anions or solvents. An example is Ag/Ag+ in T1N0 at 86°C versus Ag/Ag+ in NaN03 at 300°C. The J.A.A. Ketelaar. 339 discussion is theoretically simpler for the cell: H/MZX/MZX/M/MZX/M Y/M - M There is no difficulty in comparing the complete cells: Z Z Z Z M/M Y /Y at T2 and M/M X /X at 11 as these give a potential of the form: ECMX,^) = -GCMX,^) /ZF and G at T-j_ can be easily reduced to T2 and vice versa or., to G{298°K)«, , , . nalicellsj: mentioned Thm e comparison betweenprirstrnowever, can be made along different pathways e.g., with a thermal cell between-T2 and T^. with MX or MY and the isothermal, two-"solvent" cell can be measured at- either T2 or Tj. . Now it can be shown that the free energies of the irreversible processes involved are different, for different pathways. This implies that the unified E.MCF. series obtained are not unam biguous and thus they have no simple physical meaning. Measure ments of thermal cells with anion electrodes, e.g., carbonate electrodes are discussed. From the observed potentials essential quantities characterizing the components are derived. However, no unified E.M.F. series, nor half-cell potentials can be found, without additional con ventions, just as in the earlier use of the generalized Nernst convention for each series separately. 1. L.G. Boxall and K.E. Johnson, J. Electroanal. Chem., 30(1971)25. David G.Lovering STRUCTURE AND TRANSPORT IN AQUEOUS MELTS • David G. Lovering . Chemistry Department,City University»London EC1V 4PB,England. In aqueous melts, such as Ca(NO )2.4H 0, there is only sufficient water to satisfy the primary coordination sphere of the cation. It has been suggested that tetrahedral* tetra-aquo calcium(II) species are transported as independent entities. Spectroscopic evidence suggests that substantial anion-cation contacting occurs, whilst quasi-lattice models have led to the conclusion that a composition and temperature dependent competition takes place between anions, and water molecules for coordination sites adjacent to the cations. Indirect tests of these proposals can be made using d.c.polarography»provided a suitable tracer-ion is available. Angell has demonstrated that Ca(N0 ) .4H 0 and Cd(N03) .4H 0 melts mix ideally in all proportions;hence Cd(Il) ions at low dilutions can be used to monitor the structure of molten Ca(N03)2.4H20. ^I, Figure 1 shows the polarograms obtained for Cd(II) ions in molten Ca(N03)2.4H20 over the temperature interval 46-94°C. The large limiting current changes are expected due .to viscosity changes.However, the small shift in Ei 'v 3mV is eyenless:j than expected from the RT, _ factor in the Heyrovsky-Ilkovic equationVciearly no extensive labilization of aquo groups occurs as predicted - at least, not within this temperature range. Figure 2 shows the polarograms observed when a Ca(N03)2.4H20 + 56 mole % KN03melt+ containing Cd(II)ions, is progressively diluted with water. ( K ions are relatively unhydrated" and serve mainly to lower the K20: N03 ratio). Large half-wave poten tial shifts are accompanied by large limiting current increases. Sufficient viscosity and density data are not available yet to enable diffusion coefficient changes to be calculated.Nevertheless»Figure 3 indicates that more stable complexes are being formed, and that these are in a stepwise equilibrium, since the half-wave shifts are a smooth, curvilinear function of water concentration, and are in the cathodic direction. This supports the contention that mixed aquo- nitrate complexes are present initially and that these are progressively trans formed to aquo species at elevated water concentrations. Analysis of the stability constants is problematic in this case, since the half- wave potential in the hypothetical anhydrous melt is unknown. The formation of bromo-cadmium complexes in molten Ca(N03)2.4H20 has already been reported. »' Those results were used to calculate the individual diffusion coefficients of the complexes,assuming additivity of their ionic fluxes. The final species in the series is reasonably assumed to be CdBr . The ratio of the diffusion coefficient of this entity to that intitially present(i.e.in the absence of Br ) was found to be 04/^ = 0.68. The mean ionic volume of the CdBr4 complex, based on Pausing radii, is 490^ . If the original s*pecies present is a mixed one e.g.Cd(N03).(H20)2, then the ratio of ionic volumes of this complex and CdB>T^2.-1 S also 0.68. Although the calculation is naive, it seems unlikely that the original species could be Cd2+ , Cd(H 0 ) *+, GdCNOg),, , or a dimer, since the ratios of ionic volumes and effusion coefficients would not correspond, so closely. David G.Lovering 341 Figure I. Ca(NO ) .4H 0 + 1.34xlO~3m.CdIi: at i)46.5°C,ii)54.0°C, iii)63.80C,iv)73.80C,v)93.50C. 3 11 Figure 2. Ca(NO ) .4H20 + 56mole%KN0 +(initially)0.85xlO~ m.Cd + added water:- i)Oiii)0.11,iii)0.39,iv)0.67,v)1.22,yi)1.78,vii)2.89,viii)5.11i all xlO m. and 50°C. Figure 3. Half-wave potential shifts as a function of water conc entration, for the data of Figure 2. Figure 4. Diffusion Coefficients as a function of metal-ion formal charge in Ca(N0 )2.4H20 at 50°C. *• 342 David G.Lovering The mode and rate of mass transfer of metal ions in aqueous melts will depend on the nature and size of their solution species. Figure 4 shows that, in Ca(N0_)2.4H20 at 50°C, the diffusion coefficients of metal ions are principally dependent^upon their formal ionic charge. Although this is an empirical relationship, some reasons for this behaviour are apparent. Celeda has shown that, in highly concentrated aqueous solutions, the charge density of the ionic atmosphere is a function solely of normality i.e. the number of charge equivalents per unit volume. Moreover, the thickness of the ionic atmosphere, which can influence the rate.of ion transport, is primarily a function of the formal ionic charge. The small variations amongst ions of like charge- in Figure 4 probably reflects the degree of covalency(or charge-quenching) in the interactions of the metal ions with water molecules and nitrate ions.For Tl ions in Ca(N03)2.4H20 melts, the Stokes-Einstein equation gives a fair approach to the observed diffusion coefficent when the simple,Pauling ionic radius is assumed. For multiply charged ions, agreement with this equation is usually within 20% provided radii are based on mixed aquo nitrate complexes. For In ions, it becomes necessary to consider the kinetic entity as having at least one "permanent" secondary coordination sphere as well as a primary sphere, however. Q The concepts of quasi-lattice site interchange ' , cooperative rearrangements1°»H and normality dependence^»H all appear to be common features in modern theories of ionic transport processes. i C.A.Angell J.Phys.Chem. 7£, 3988 (1966) D.E.Irish and G.E.Walrafen J.Chem.Phys.4£, 378(1967);Inorg.Chem.£, 3J.Braunstein J.Phys.Chem._n, 3402 (1967) - 425(1970) "C.T.Moynihan, C.R.Smalley,C.A.Angell and E.J..Sar* ibid 21,2287(1969) 5C.T.Moynihan and A.Fratiello J.A.C.S. 89_, 5546 (1967) 6D.G.Lovering thesis City University (1969) 7D.G.Lovering and D.J.Alner Chem.Comm. 570 (1970) 8F.A.Cotton and G.Wilkinson "Advanced Inorganic Chemistry"Interscience, London 1972. 9J.Celeda Scientific Papers of the Institute of Chemical Technology, Prague, B7, 3,(1966) 10G.Adam and J.H.Gibbs J.Chem.Phys.43_, 139 (1965); 46_, 4673 (1967) if C.A.Angell J.Phys.Chem. 70, 3988 (1966) Arnold Lunden 343 ISOTOPE EFFECTS OF ELECTROMIGRATION IN MOLTEN ALKALI NITRATES Arnold Lunden and Isao Okada Department of Physics, Chalmers University of Technology, S-402 20 Gothenburg 5» Sweden and Department of Chemistry, Faculty of Science, University of Tokyo, Bunkyo-ku, ,Tokyo, Japan In an electromigration'experiment the difference "between the mobilities of two ionic species (e.g. two cations) is determined, while the sum of the mobilities (of e.g. all cations relative to a common anion) is obtained in conductivity measurements. Thus changes in the mobilities of the participating ionic species might influence electromigration and conductivity in different ways. Electromigration isotope effects are currently expressed by means of the mass effect (|i)> which is the relative difference in mobility divided by the relative difference in mass: (i, = (Ab/b)/(Am/m). Electromigration in an ionic melt involves interaction between a number of ions, both like-charged and oppositely charged ones. By comparing bhe actual mass effects with predictions by various models, it is possible to estimate the degree of interaction and the number of interacting ions. It has previously been found for a halide system (LiCl-PbCl2) and a sulphate system (LigSO,- K^SO/) that the isotope effect tends to increase when a cation becomes sufficiently diluted. For the nitrate system KNO*- RbNO^ there is however not much of a concentration effect for either cation, and this appears to be the case also for LiNO?- KNO*. (Work is in progress with the latter system; final results should be available,, at the time of the conference.) A possible interpretation of this is that there is a much stronger anion-cation interaction in the nitrate melts than in systems with spherical anions (halides, sulfates). While it is well-known that the electrical conductivity of a molten salt increases when the temperature is raised, the temperature dependence of the electromigration isotope effect is more complex, and cases are known where the isotope effect increases (+ in the table), decreases (-), has a maximum (M), or is nearly independent of temperature (O). The results obtained so far by different groups for pure melts (and solid Li2S0^) are summarized in the table. Some work has also been done on the For cati ons (JJ, ) For anions (y, ) Ion* CI Br NO, SO, CI . Br 3 4 Li + 0 0 Na M K + + Rb + ( + ) M 0 Tl — Sn (-) Pb 0 344 Arnold Lunden temperature dependence in mixtures. In a KNOz- RbNOj mixture (90 % KNOz) "both isotope effects show a similar "behaviour to the pure melts, i.e. JJ, , has a maximum while JJ,„ increases over the whole temperature range. In a ITaNO^-KlTO^ mixture with 45 i° NaNO-z a maximum.was obtained for |A-M- > while ^ "was not studied. (In this latter work Wuhl et al. used a.technique that is different from ours,1 and "external'.'-, mass effects were measured, while we measure "internal" mass effects, i.e. in the first case a part of the apparatus serves as the reference frame of the mobilities while in our work the common anion serves as the reference frame.) Work is in progress on the temperature depend ence of |j,T . and p,^ in eutectic LiNCL-KNO-, and LiCl-KCl. Klemm has pointed out that there is a correlation between the temperature dependences of the mass effect and the Nernst- Einstein factor. Thus, in most cases the" temperature coeffi cients of (j, and of D/TV\_ have the same sign. References A. Klemm and K. Heinzinger, Advan. Chem. Ser. 82, 248 (1969) A.. Lunden and A. Bkhed, Z. Ha. turforsch. 24a, 892 (1969) N. Saito, K. Hirano, K. Okuyama and I. Okada, Z. Naturforsch., in press S. Wuhl, F. Lantelme and M. Chemla, J. Chim. Phys. 6£, 488(1968) Pierre Mergault 345 ELECTROLYTIC PLASMAS IN MOLTEN SALTS AND AQUEOUS SOLUTIONS Pierre Mergault, Jean-Claude Valognes, Denis Leteinturier, Jocelyne Garbarz-0livier and Christian Guilpin. Laboratoire de Physique desLiquides Ioniques, Université de Paris VI, Paris, France. It is well Known that small glowing discharges appear between anode- and bath during anode effect (AE) in molten salts. During anode or cathode effect in aqueous solutions (following Kellog (1 D , we shall speak of "aqueous anode effect" (AAED and "aqueous cathode effect" (ACE3), we can observe the same display between the concerned electrode and solution ( 2) . The I-U curves (I : current intensity, U : tension between electrodes] have in bulk the same shape represented in figure 1 •, there are three principal regions AB corresponds to conventional electrolysis, BC is the Kellog region (3) which is one of instabilities while I and U seem much more stable in the region CD. Thus, it is con venient to term these three phenomenas as "electrode effects" (EE). I and U values at points A, B, C, depend on EE, media, nature, size and shape of electrodes, experimen tal conditions such as tempe rature, pressure, current densities (cd) on electrodes, ratio of these cd, and so on.. Discharges are only ^D possible in an isolating media: we think this media is gazeous 0 in all the cases of EE. There V fore, there must be a gazeous Fig. 1 sheath surrounding the whole electrode concerned by EE and'U must be great enough in order that discharges can appear between the two walls of this sheath. Between the moment the gazeous sheath is completly achieved and the moment the first discharge appears, the two electrodes are electrically isolated and I vanishes : that can be seen on an oscilloscope and 1=0 during a time less than 1 JJS. This sheath is unstable and is. alternately destroyed and built up with random frequency from 1 kHz to far beyond. We study plasmas produced by EE by spectroscopically analysing the emitted light. Then, it is first necessary to adapt EE to this observation : bright enough light sour ce and possibility to transmit this light to the analysing instrument. Mo It en salts. We studied AE in cryolite and in LiCl-KCl mixtures of varying compositions from pure LiCl to pure KC1. The obtained spectra contain rays of neutral and 346 Pierre Mergault ionized atoms and molecular band spectra appear in the case of AE in cryolite. Rays are enlarged, shifted and sometimes forbidden rays (by selection rules] appear. In terpretation of these spectra .may only be done in the fra me of plasma spectroscopy theories. Griem's theory (4] permits us to show that it is necessary to introduce the Stark effect with field E. .. • 1 . - Cryolite. A graphite crucible was used as cathode. Anode is a 4 mm graphite rod immersed 5mm in molten cryo lite. With anodic cd from 30 A/cm2 to 56 A/cm2, 11 = 100 V and exposure times are smaller than 2mn with the spectro graph we used. A. particular ÄE was also obtained ; we kept AE with anodic cd 38 A/cm.2 (I-12A] during some minutes and suddenly reduced I to 3A (anodic cd = 9.4 A/cm2] ; then, the brightness of the anode was almost the same as under preceding experimental conditions during 10 mn. The larger anodic cd,the greater is the number of rays. Obtained spectra contain rays of Na,v Ca I, Ca II, Al, forbidden, rays and unidentified molecular bands. If anodic cd > 47 A/cm2, Fe rays appear. Some properties of anodic plasmas can be 3 determined : T=5000°K, electron den sity Ne~B 1021/m , ionic microfield F0 = 13 kV/cm, ionization ratio (a = N-j./Na/ N^ : ion density, Na : neutral atom density for same species) a(Na]=0.83, a(Al)=0,12, a(Ca]=0.36. Maximum pressure cor responding to these electron, ion and atom densities is only 0.06 atm instead of 1 atm : thus, there exist other particles in the plasma ; this is really the case (molecu lar bands]. 2. - LiCl-KCl mixtures. With a transparent- quartz cruci ble, 2mm graphite anode, 5mm iron cathode, U=70V, AE is stable and brightness is sufficient to observe the anode through the wall of the crucible. The anode image is di rectly formed on the slit of a spectrometer. Anodic plasma properties for LiCl-KCl eutectic are given elsewhere (5). For other compositions, properties are almost the same except the ratio of concentrations of Li and K in the plas mas are near these in the bath. Moreover, it seems that Griem's theory may no longer be used when the concentration of LiCl is too large. Aqueous solutions. We studied spectrometrically ACE (6] for U varying between 70 and 150 V while AAE was studied up to U=500 V ; all these values correspond to the region CD of figure 1, beyond C. Mean current intensity is small (10 to 100 mA) but on an oscilloscope we see I randomly fluctua ting at high frequency with an amplitude of about 10 A. In AAE, this amplitude increases with increasing U (beyond 380 V, I is alternately stable and fluctuating]. In all cases, the greater U and the brighter is the light source. 1.- ACE. Necessary voltage to have a sufficient bright ness is small compared to that of AAE. Even with refractory Pierre Mergault 347 metals, such as tantalum, melting of the cathode occurs for voltages smaller than 200 V if the cathode is a wire of 1 mm diameter and if ratio of cathodic and anodic surfaces is about 50. In all cases, metallic spectra of the cathode is obtained. Pt cathode in N KBr or K2C03 solutions permits the observation of rays of Pt r, Pt II, K I, K II, 0 1, OH and H (the 4861 8 ray is always enlarged and if U is large enough, the 4340 R ray is obtained). With a Pt C90 %) Rh C10 %) cathode, U = 110 V, in 2N H2SO4 solution, we obtain 135 Rh rays and 76 Pt rays between 3250 and 5000 K. With Au cathode and same conditions, 27 Au rays are obtai ned. Ag,cathode in LiCl solution gives 7 Ag rays, 3 enlar ged Li. rays . Cu cathode in LiCl, Na2C03, Na2S,04, BaCl2 solutions permits us to see Cu I and Cu II rays, Li, Nd or Ba rays according to the employed solution. 2. - AAE. The discharge aspect and brightness are c y p1 cal of the solution and concentration.' The solution is the cathode for the discharge and a cation spectra is obtained. Rays are always more or less enlarged and absorp tion occurs with sodium. Between 1N and 5N for halogenides, carbonates, sulfates of alcaline and alcaline-earth metals and between 10N and 15N for acids, with large enough volta ges, emitted light wasspectroscopically analysed. Bright ness increases very much for U s 3 8 0 V, whatever the concen tration. For U > 380 V , all rays of the cation and rays of 0 -II appear . (1 ) HH.KELLOG, J.Electrochem.Soc., 97, 1950, 133. (2) A.F0NTANA and R.WINAND, Electrochem.Acta, 16, 1971, 257 and 279. C3) A.HICKLING, in Modern Aspects of Electrochemistry n° 6 BOCKRIS and CONWAY, ed. Butterworths. C4) H.R. GRIEM, Plasma Spectroscopy, 1964, tic Grew Hill. (5) Sup . to t he J.Phys. 32, 1971, C5b-43 P. MERGAULT and J.C. VAL0GNES (6) P. MERGAULT, J.C.VALQGNES and C.GUILPIN, C.R. Acad.Sci. Serie C, 273, 1971,26 (7) P. MERGAULT, J.C.VAL0GNES,J.OLIVIER-GARBARZ and C.GUILPIN, C.R.Acad.Sci., Serie C, 274, 1972. 348 K. E. Oberg THE COULOMETRIC DEOXIDATION (REFINING) OF LIQUID METALS by K. E. Oberg, L. M. Friedman, W. M. Boorstein and R. A. Rapp, Department of Metallurgical Engineering, The Ohio State University, Columbus, Ohio, USA. Oxide particles, which degrade the mechanical strength and response of metal parts, are commonly formed upon the solidification of air-melted liquid metals. Likewise, in the deoxidation of liquid metals by the addition of re active elements such as aluminum, silicon and manganese, a fraction of the deoxidation products are unable to escape from the molten metal prior to freezing and are therefore trapped in the resulting ingot or casting. In this paper, a new electrochemical deoxidation method is described which should find particular industrial application in the con tinuous casting of alloys._, The ceramic oxide, stabilized-zirconia, is essentially an exclusive conductor of oxygen ions (and not electrons) at temperatures from about 700°C to its melting point (^2500°C). Thus, this ceramic can serve as an electro lyte in high temperature galvanic cells whereby the open- circuit voltage will correspond to the ratio of oxygen activities of the two electrodes according to the relation: RT , P62 (reference electrode) n O.C. ~ 4F PQ (in liquid metal) Likewise, when a voltage is applied between two electrodes which are excellent sources and sinks for oxygen, then oxygen can be transported according to Faraday's Law A~ . i (amps) L0 (moles/sec) = Chin /mnlac/can^ = 2F(amp-sec/mole- £-—- ) dt from the oxygen source electrode to the sink. In these studies, oxygen-contaminated melts of copper and lead were held in a stabilized zirconia electrolyte-crucible with an external porous platinum-in-air reference electrode to form the following electrochemical cell: K. S. Oberg 349 metal (jt) ZrO?(CaO) Pt-air (P' =0.21 atm) + solid °2 £ (dissolved) electrolyte reference at p electrode " ö2 The removal of oxygen from the melt via coulometric titration of oxygen ions through the electrolyte to air was effected by the application of a relatively large voltage, EapDi#, to this cell with the liquid metal at negative potential. In such a potentiostatic experiment, s the deoxidation current, ii0n» ^ related to the oxygen activity at the melt/electrolyte interface by the expression RT , „. RT E_appl, . - i ion *-*"*O *• io n = ^4 pI n PA<> ?- —4 p In P" *-*•** ion Lv Oo AV 0 where is the ionic resistance of the electrolyte. For stagnant (unstirred) melts, the potentiostatic experiment leads to a rapid depletion of dissolved oxygen (PgJ at the melt/electrolyte interface and a correspond ing decrease in iion» under these conditions the deoxida tion is controlled by the diffusion of oxygen in the melt is therefore ineffectual. However, if the melt is stirred vigorously by induction heating during deoxidation the oxygen concentration is maintained uniform throughout the melt except across a thin boundary layer at the metal/ electrolyte interface. Further, the vigorous induction stirring minimizes the boundary layer thickness and there by maximizes oxygen transport across this layer so that the rate controlling step during the deoxidation transient can be either the Faradaic oxygen transport in the electro lyte (at relatively high oxygen contents), or else diffusional transport of oxygen through the melt boundary layer (at relatively low oxygen contents). Most importantly, the continuous supply of oxygen to the electrolyte/melt in terface afforded by inductive stirring permits the use of pumping voltages far in excess of the decomposition voltage of zirconia because the applied voltage is man ifested as iion *^i ion ^roP (at hd-gner concentrations) instead of a reduction in P(J2 which would degrade the 350 K. E. Oberg electrolyte. By a separate electrochemical monitoring of the oxygen concentration during deoxidation, the course of rate control was found to obey these theoretical expectations, and information about oxygen transport through the boundary layer was obtained. The deoxidation of induct ion stirred melts of copper at 1150°C with Eappi# = 3 volts exhibited a Faradaic rate for oxygen concentrations in excess of 40 wt. ppm and a logarithmic rate at lower concentrations. In the deoxidation of liquid lead at lower temperatures with Eapp^# = 4 volts, an interfacial polarization step was observed which reduced the deoxida tion rate below theoretical Faradaic expectation. -b O» uoa>ii CT-». £ cr—i-ti o-o Vt t -1*3 rflo rf-a 3>0 3 m O 3 O -•• to -hiD m o m o a _c m O 3 o-<< -*• —» c C 3 Oi Dl 0<< —• -oc O fl) rf CD 3-3-0» C |3 3" CD — • -I» CD c+ 1 Oi X m rf O 3 • o» —• o CD • t— O O 3 O -$ rta< 3 rf N c :>-* -J rf 3 -J O» < CD CD -»• -»• D-lfD CD O* O rf Oi tO rf -»• . C -ST3 3 -S CD —i-a - m ar —• n> —• rf _J.fi) *< -*• CD CD rf 3- Vt Oi CT CD Vt O 3 CD • 7T —»rf -J O" Oi -*• O. in CL __i • -$ o o o» r o rf -*• O —•• rfto 3* o o < Q. I Vt 3" 01 O rf rf CD CD 3-CD fl» co -J -«•-»• 3 O P> (+-• in o -n -••rf "5 -S • —I 3" O m o 3- a> rf O» Oi IQ O CD C Oi C —•«/) 3 O • • -»-O —• -•• rf Oi CD o *< • CD -S O» rf TO ID -+ *< o -s rf rf n CD -$ Separate cyclic voltammetric studies in this solvent have shown that on oxidation, TMB yields a single redox couple at a potential identi cal to that for the couple above. However, the important feature of this voltammogram is the strong indication on cathodic sweep of a wave at 1200 mV, which at higher sweep rates (up to 100V/sec.) be comes a well defined peak with E .(cathpdic) = 1030-1040 mV.7 This peak which is^90-100 mV; cathodicpof the; main oxidation peak,' is, we feel, the reduction of the DMA radical, cation. These results at 175° are to be compared with previous studies at 25° on the oxidation of DMA in acetonitrile where the radical cation was not detected as such in electrochemical experiments.3 Thus-, we again conclude that the aromatic amine radical; cation is stabilized in the A1C13 containing melt. Almost identical cyclic voltammetric results to those for DMA were found when the oxidation of DPAmine was studied, the only apparent electroactive products being the'DPAmine radical cation and N,N'-diphenylbenzidine. These results again point to the stabilizing effect of the A1C13 containing melts. Several inorganic solutes - iron, silver, copper - have been investi gated. The behavior 1s'cons1stenCMth.;'.niprmal electrochemical behav ior in the 1:1 melt. Potential measurements in the 1:1 and 2:] alum inum chloride-sodium chloride melts suggest that the: Ag(I)/Ag elec trode couple may be a more suitable reference than the more commonly used A1(III)/A1 electrode; the silver potential is not as subject to change with variation in pCl as is the aluminum electrode. A study of melt equilibria has been carried out and the potential- composition curve has been fit with derived equilibrium parameters from the NaCl saturated melt to the 2:1 melt. This data may aid in understanding the behavior of the organic solutes in.this solvent system. Acknowledgement This work was supported by The Air Force Office of Scientific Research under Grant No. AFOSR 71-1995. References 1. R. S. Nicholson, Anal. Chem., 38., 1406(1966). 2. R. F. Nelson and S.W.F. Feldberg, J.Phys.Chem., 73, 2623(1969). 3. R. Hand and R. F. Nelson, J.Electrochem.Soc., 117_, 1353(1970). Ger ar d PINARD-LEGRY 353 A CONTRIBUTION TO THE STUDY OF THE Pb/PbSO , H SO ELECTRODE UP TO 250°C. 4 2 4 2fc stafc F. Letowski and G. Pinard-Legry Instytut chemii Nieorganicznej i Metalurgii Pierwiastkow Rzadkich Politechniki Wroclawskiej, Wroclaw, ul. Wybrzeze Wyspianskiego 27 Pologne. Commissariat å 1'Energie Atomique, Service d'Etude de la Corro sion et d'Electrochimie-BP. n°6-92 Fontenay-aux-Roses, France. The knowledge acquired on corrosion mechanisms in aqueous medium comes mostly from development of electrochemical methods. Yet the increasing use of high temperatures, particularly in the nuclear indus try, raises fresh problems as regards the technique of these measu rements. It is above all essential to define a reference electrode able to operate at the highest possible temperature. The Ag/ÄgCl, HC1 sys tem, whose potential up to about 300°C (l, 2) is well known, could be perfectly suitable. Nevertheless, a few traces of chloride ions are enough to seriously upset the corrosion phenomena of some materials, particularly stainless steel. In order to eliminate this contamination hazard, Lepeintre, Mahieu and Monjou (3) proposedtheuse of Pb/PbSO.i H^SO 4, electrode up tö.250°G. This electrode has been investigated again here in detail and then used for determining ,the current poten- tiaL.curves of. Fe-Cr-Ni' alloys in LiOHJat 250°C. 1. Technical process. The following cells were made : Pb/PbSO , H SO (0,05 m)/K SO (sat. 25°C)KCl(2m)/HCl(0, 1 m), AgCl/Ag andT^b/PbSO , H SO4(0,05m) K SO (sat. 25°C)H. SO (xm),PbSO /Pb. The various valUes^of x are, : 0.%5; 0. 01 0; 0. 1 a./d oS mole for 1000 g of H O. The e. m. f;measure ments were made in argon in an autoclave. The P. T. F. E. electrode are composed of two concentric compartments separated by a very fine passage filled with asbestos (fig. 1). The Pb/PbSO electrode is made of a rod of very pure lead (99. 999%) plunged into a given sulphu ric acid solution; the external compartment contains potassium sul phate satured at 25°C. The amounts of electrolyte are so calculated that at 250°C they are less than the volume of the two compartments. Finally, all the work of preparing the solutions and fitting the electro des is carried out in argon. The mounting of the Ag/AgCl electrode is identical to the preceding one. The silver and silver chloride coated (thermal process)platinum wire is plunged into a solution of HCl 0. lm (4) . The electrodes, which are partially immersed in a cell containing a solution of saturated K SO .are connected to an electronic millivolt- meter by means of insulate2 d 4an d leaktight passages in the lid of the autoclave. Three electrodes, two of which are identical, are actually used; a switch in the measuring circuit then enables the e. m. f. of the 354 Gérard PINARD-LEGRY two cellsto be recorded in turn. The measurement accuracy is 0..1mv. For each temperature a chek was made to ensure that the thermal and electro-chemical balances had been reached. After each set of measurements, the electrolytes in.the various electrodes were collec ted so as to check that their volume and pH had not varied during the test. 2. Measurements results. 2. 1. e. m. f. The e. m. f.* "values of the Pb/PbSOA, H SO fO.05m)//HCl (0. lm), AgCl/Ag cell are shown on figure 2 in. terms of the tempera ture. Determining the equation of this curve leads to the following re lations being proposedrfrom 1 0 to 120 °C:E=262+1.910 T-0.002937 T2 from 100 to 250°C:E=1855. 5rl 0.89 T+0.03074 T2-0.000029 T3. The e. m. f. variations of the Pb-Ag cell with .respect to various sulphuric acid solutions are also shown on figure 2.- 2. 2. Repeatability and stability. The e. m. f. variations of the cells studied never exceeded 5 mv after a period of heating at 250 °C. 2. 3. Polaris ability and stability. The tests were made by making a current circulate between the lead electrode and a platinum counterelectrode. From all the results of this study it is possible to conclude that the Pb/PbSO lead electrode may be used as reference electrode up to temperature of at least 250°C. The sulphuric acid concentration may be in the 0.005 and 0. lm bracket. 3. Electrochemical study of Fe-Cr-Ni alloys in LiOH at 250°C. By using the preceding electrode as reference the currents potential diagrams were determined for Fe-Cr-Ni (17% Cr,Ni variable)in LiOH (0.1m) at 250°C. The experimental system included a P. T. F. E. cell immersed in an autoclave enabling the medium to be renewe d (0.51.K"1) The wire sample was placed in the centre of a circular platinum coun ter-electrode. The results achieved highlight the beneficial effect of adding nickel on the behaviour of these alloys in the anode field. Thus the peak which corresponds to a change of nature in the oxidizing reac tion at a maximum current which drops markedly with the increase in the nickel content of the alloys used. References. (1) M. H. Lietzke and J. V. Vaughen J. Am. chem. Soc 77 (1955) 876 (2) R. S. Greeley, W. T. Smith Jr. , R. W. Stoughton and M. H. Lietzke J. Phys. Chem._64 (I960) 625 and 1445 (3) M. Le Peintre, C. Mahieu and J. Monjou C.R.Acad. Sci. 261 (1965) 3389 (4) E.J.G. Ives and G.J. Janz "Reference electrodes"Acad. Press New York (1961) Gérard PINARD-LEGRY 355 Fig. 1 : Pb/PbSO electrode arrangement a- Silver contact b- P. T. F. E. insulating sheath c- Pressure balancing ducts d- Lead rod e- Internal compartment (H SO ) f- External compartment (K SO ) g- Very pure asbestos fibre 500- Fig. 2 e. m. f. variations of the Pb/PbSO^, **2*^4 (xm )//HCl (0. lm),AgCl/Ag cells with temperature 356 Jomar Thonstad THE CATHODE REACTION ON ALUMINIUM IN NaF-AlF -Al 03 MELTS Jomar Thonstad and Sverre Rolseth . SINTEF, The Engineering Research Foundation, The University of Trondheim, Norwegian Institute of Technology, Trondheim, Norway. Most previous works on the cathode reaction on aluminium in cryolite- alumina melts concludes that there is little charge transfer over- voltage. It has "been claimed, however, that diffusion overvoltage occurs (l), due to the "fact that the sodium ion is the carrier of current and the aluminium ion is bound in complexes. Aluminium is more noble than sodium except in very NaF-rich melts, the difference being approximately 0.1 V at the cryolite (Na_AlF,-) composition at around 1000°C (2). Steady state measurements. The catho'de reaction proceeded smoothly below 1 A/cm2, whereas gas bubbles of sodium were formed at high cds, giving rise to current oscillations. - The gas bubbles could be observed when using cathodes made of iron, while the socalled metal fog obstructed the view of aluminium electrodes. When a constant current was applied to an aluminium cathode, the potential recorded against an aluminium reference electrode attained a stable value after about one minute. At cds below 0.15 A/cm2 the i-V curves were linear, and above 0.2 A/cm2 straight Tafel plots were obtained. At the cryolite composition the Tafel slope was 0.20 V with a standard deviation of 0.03 V. Occasionally considerably higher and lower values were found. The slope was .independent of the alumina content of the melt, while it decreased with increasing NaF/AlF^ molar ratio (C.R.), becoming 0.08 V at C.R.=6. An unex pected maximum of 0.28 V was found at C.R.=5» The overvoltage de creased slightly by application of mechanical stirring. In cryolite the overvoltage was -approximately 0.1 V at 0.5 A/cm2*. It is possible to interpret the steady state data in. terms of a multiple step slow charge transfer reaction for Al ions. Such a mechanism is not very probable, however, firstly because of the out come of non-steady state measurements and secondly because of the scatter in the experimental Tafel slopes. Non-steady state measurements. Double pulse measurements showed that the charge transfer process is rapid. The cd of the second pulse was varied from 0.1 to 1.9 A/cm2. The charge transfer resis tance was on an average 0.0030 ohm cm2, and the corresponding exchange currents calculated for one-, two- or three-electron steps were 36 A/cm2, 18 A/cm2 and 12 A/cm2 respectively.. Because the charge trans fer resistance was an order of magnitude lower than the ohmic resis tance in the measuring circuit, the precision of the data was low. This was even more true for results obtained with the ac impedance method. The faradaic impedance was evaluated as a series connection of a resistive and a capacitive term. In most results the resistive term attained a steady value at frequencies below 1000 Hz, the magnitude of which could vary from the same value as found by the double pulse method and up to 0.015 ohm cm2. The capacitive term was approxi mately zero at frequencies up to around 100 Hz and it then rose to Jomar Thonstad 357 a maximum at around 3000 Hz. The occurrance of this maximum can simply he due to experimental error in reading the phase angle. If the maximum is" real, it indicates the presence of a small reaction overvoltage together with the charge transfer. Altogether, these data yield ..exchange currents which are two orders of magnitude higher than the steady"state values. Potential decay .curves obtained by open circuit after steady state polarization.exhibited a.rather slow decay, reaching e.g. 20 mV after 20-50 seconds. .Plots of E versus log t yielded slopes of O.OU to 0*10 V/decade, which is considerably less than the Tafel.slopes. The origin, of the relatively .high overvoltage found ..in the. steady state measurements is not clear. It seems to be connected with an accumulation of reaction products in the proximity of the cathode surface. 1. R. Piontelli and G. Montanelli, Alluminio 22, 699 (1953) 2. M. Feinleib and B. Porter, J. Elecnsrochem. Soc. 103.; 231 (1956). 358 R. Tunold DETERMINATION OF ELECTRODE PARAMETERS FOR THE CHLORINE/CARBON ELECTRODE IN CHLORIDE MELTS T. Berge, K. A. Paulsen and R. Tunold Dept. of Industrial Electrochemistry, Technical University of Norway, Trondheim, Norway This work is part of a general study of halogene/carbon electrodes in halide melts. Steady-state cd/potential measurements have shown that a recombination of adsorbed chlorine atoms on the surface is likely to be the rate controlling step in the chlorine evolution reaction (l) (2) (3) (U). This was confirmed by double pulse measurements which showed that the discharge reaction is a fast one (3) (5)« - In this study, transient methods have been used to measure different- electrode parameters in chloride melts of various composition, from pure NaCl to binary mix tures of NaCl and AgCl, and NaCl and MgCl2. Galvanostatic single pulses were used to measure the double layer capa citance as a function of potential. The double layer capacitance, adsorption capacitance, charge transfer resistance and reaction resistance were calculated from AC-impedance measurements by a computer. This made it possible to suggest an equivalent circuit for the electrode. Cathodic stripping curves were performed to calculate the amount of adsorbed intermediates. The complex diagrams from the AC-impedance measurements confirm that a heterogeneous reaction probably is rate controlling. ' This reaction is coupled with diffusion at low cds. The reaction resistances measured with this method are in good agreement with the calculated linear re sistances of the steady state cd/potential curves. The uncertainty in determining the discharge resistance was great, but a qualitative agreement with data from the double pulse method (3) were obtained. There was also found a fair agreement between the double layer capaci tance measured with the two methods mentioned above. The results for the pure sodium chloride are in agreement with data from the litera ture (6). The capacitance/potential curves have three regions, one region with a nearly constant and.low capacitance, and two regions with high and increasing capacitance at increasing negative and positive potential. In the vicinity of Erev for the chlorine electrode a "hump" was obtained. The "hump" changed in magnitude and position on the poten tial axis depending- on the composition of the melt. This probably was due to adsorbed anionic species from the melt or to neutral inter mediates in the electrode reaction. H, Tunold 359 Cathodic stripping curves show that the number of chlorine atoms per 100 carbon atoms at the apparent surface increased from about 0.3 at zero cd to about 5 at a cd of 500 mA/cm2. This corresponds to a variation in surface coverage of chlorine atoms from 1-2% to about 20%. The calculations were based on a surface coverage equal 1 with a mono layer of chlorine atoms adsorbed in a hexagonal array oh the surface. A great uncertainty was involved in those measurements because of gas "bubbles adhering to the electrode at higher cds. Scattered measurements with the double pulse method in MgCl2-contain- ing melts showed that the exchange cd of the discharge reaction is ten times greater in these melts than in melts composed of NaCl and AgCl. This could be due to anion bridging by contact adsorbed nega tive complex ions. References (1) H. J. Vandenbroele and A. J. Arvia, An. Asoc. quim. argent., _55. 21 (1967) (2) . W. E. Triaca, C. Solomons and J. 0'M. Bockris, Electrochim. Acta, 13 19^9 (1968) (3) R. Tunold, H. M. B0, K. A. Paulsen and J. 0. Yttredal, Electrochim. Acta, 16 2101 (1971) {h) R. Tunold and K. A. Paulsen, Extended Abstracts, 22nd Meeting ISE, Dubrovnik Jugoslavia 1971, page 15 H (5) P. Drossbach, H. Hoff, Electrochim. Acta, lU 89 (1969) (6) E. A. Ukshe, N. G. Bukun, J. Electroanal. Chem., 32 283 (1971) 360 Pier Giorgio Zambonin THE ELECTROCHEMISTRY OP OXYGEN AND ITS SOLUBILITY IN MOLTEN ALKALI NITRATES. E. Desimoni, F. Paniccia, P.G. Zambonin Istituti di Chimica Analitica e Generale - Universita* Via Amendola, 173 - BARI (ITALY). The knowledge of the- solubility of oxygen and the availa bility of a sensitive analytical tool for its detection are of primary importance for several kinds of investiga tions in molten salts:first of all for a complete under standing of the oxide chemistry C1) and the behaviour(2) of the oxygen electrodes. Furthermore in the case of oxv. genated salts, such as nitrates and sulphates, the possi, bility to reveal lov/ oxygen concentrations can be a ra pid and usefull way to detect and/or follow the decompo sition of the melts themselves. Nevertheless, except some work relevant to the solubili ty of oxygen in molten carbonates, no experimental infor, mations are available, up to day, in this field. In the course of the present investigation the solubili ty of oxygen in the (Na,K)N03 eutectic melt has been de termined in the temperature range 520 - 600°K by using a high-sensivity (_+ Ö.02 torr) differential micrometric tecnique. In brief a solubility experiment was performed as follows. After a complete evacuation of the solvent- containig apparatus, oxygen was introduced in a free volume over the melt and the differential manometer rapi, dly calibrated versus the initial pressure (A* 1 atmosphe re). At this point a vigorous magnetic stirring of the melt was initiated and the pressure variations read until equilibrium was reached. Prom the pressure varia tions (after a suitable calibration) the number of dis solved moles was obtained and by knowing the final abso, lute pressure, the Henry's constant calculated (see Table l). Oxygen in a nitrate melt can be voltammetrically reduced (1a,1f)according to the electrode processes reported in Figure 1. The current steps relevant to the curves A+B (voltammetrically reversible) or C (voltammetrically irreversible) are diffusion controlled and can be used for analytical detection of oxygen in dry or wet melts Pier Giorgio Zambonin 361 respectively. A potentiometric study (2b) performed in the concentra tion range 1Cf ^ICT1 m of oxides has shown that both superoxide or hydroxide electrodes are potentiometrically reversible. The knowledge of oxygen and water (3) solubi lity permitted to calculate the relevant standard poten- tials: 02 + e = 0~ E° /0- = -0.645+0.005 V 02 + 2H20 + 4e =40H~ Eg fi Q/0H-=-0.495+0.005 V both referred to a reference electrode Ag/Ag (0.07m). TABLE 1: Thermodinamic parameters for the solubility of oxygen in the (Na,K)N03 eutectic melt. Temp. K 108 AH AS0 (°K) (mole ml x bar *) (Kcal/mole) (e.u. at 533°$ 508 0.84 533 1.03 4.2 -7.4 573 1.36 603 1.65 BIBLIOGRAPHY 1 )a-P.G. Zambonin and J. Jordan-J.Am.Qhem.Soc. t89i6365 (1970) fb-J.Jordan,W.B.Mc Carthy and P. G. Zambonin in "Characterization and Analysis in Molten Salts"G.Maman- tov, Ed .M.Dekker,New York,N.Y., 1969;c-P.G.Zambonin, J. Electroanal.Chem. 24.365. (1970) ;d-P.G. Zambonin and A. Cavaggioni.J.Am.Chem.Soc..9jt2854t(1971);e-P.G.Zambonin, F.Paniccia,A.Bufo, J.Phys.Chem.,in press;f-P.G.Zambonin, Anal.Chem. 41,1571,(1971). 2)a-P.G.Zambonin, J.Electroanal.Chern. g±, App.25,(1970); b-P.G.Zambonin, J.Electroanal.Chem. H, 243i (1971). 3) P. G. Zambonin, L. Cardetta, G. Signorile, J. Electroanal • Chem. 28,237,(1970). 362 Pier Giorgio Zambonin Figure 1: Current-potential curves recorded at a Plati num Rotating Disk electrode(0 =1.5mm; 600r.p.m.) at *v500°K: pure (Na,K)N03 under one atmosphere of dry Argon j melt under a pressure of one atmosplie re of dry oxygen,* melt under a pressure of one atmosphe re of oxygen containing traces of water —-—•«—. z O O X o o.o L^L oo o.o -LO -2.0 I— POTENTIAL v£. f^a/^9 (0.07 ml SECTION 5 CORROSION A COMPARISON OF THE DISSOLUTION-PRECIPITATION AND SOLID STATE MECHANISMS OP PASSIVATION R. D. Armstrong and J. A. Harrison Electrochemistry Research Laboratories, University of Newcastle upon Tyne, NEWCASTLE UPON TYNE, NET 7RU, England. The electrochemical passivation of metals has been gen erally ascribed to the presence of a thin film of some sort on the surface of the electrode . The mechanism whereby this film is formed has been the subject of much discussion. Muller seems to have been amongst the first to suggest that an initial film may be produced by a dissolution precipita tion process, the metal dissolving until a critical concen tration cÄ is reached in the vicinity of the anode when a precipitate is formed which blocks the metal surface. The other means of producing this film involves a solid state reaction. Anions (Xn~), which may derive from the solvent, e.g. 0^~ from HgO, react with the metal directly without metal cations entering the solution. This is analo gous to the gas phase oxidation of metals. In principle these two mechanisms can be distinguished by determining whether or not metal cations have entered the solution prior to the formation of a film. However, when metal dissolution occurs in parallel with a solid state reaction, i.e. metal dissolution occurs in the active region and during the onset of passivity, the distinction is not easily made and it is the purpose of this communication to consider the conclus ions which can be drawn from the various experimental pro cedures which are commonly used to explore the active- passive transition. It is first necessary to consider the various models in more detail . (A) Dissolution-precipitation In this model it is important to recognize that the metal can dissolve as metal cations or as complexed species which means that homogeneous reactions may play a part. The metal dissolution may be totally irreversible, as is the case with most of the transition metals, or it may be reversible to some extent (Zn, Cd, Hg,' Ag) . In what follows it will be assumed that the mass transport of metal cations within the metal, and of Xn~ to the surface of the metal, is unimpor tant . The concentration c* will generally be in excess of that thermodynamically in equilibrium with the precipitated material because of the need for supersaturation before pre cipitation can occur. The precipitate is assumed to .form in some way close to the electrode surface causing the rate of the metal dissolution reaction to be lowered. It should be noted that the nuclei of the precipitated material will 366 Ronald Dixon Armstrong always be three-dimensional so that this mechanism is not expected to lead to a monomolecular passivating film. It must also be recognized that a precipitate will form "near" the electrode surface and not on it, since otherwise pre cipitation is indistinguishable from solid state reaction. Near, the electrode should probably be taken to mean a dis tance which is much greater than the thickness of the elec trical double layer. If the rate constants for the precipitation are to be determined attempts must be made tö obtain the nucleation rate and crystal growth rate at the supersaturation which is implied by cx. These values and c36 must be independent of the potential or the current (whichever is the independ ent variable). (B) Solid state reaction in parallel with metal dissolution In this situation when a certain anodic potential (,E ) is exceeded a film starts to form on the metal surface through the direct attack of the anion on the metal. E will generally be more positive than the thermodynamically rever sible potential for the formation of the phase, because of the existence of a nucleation or crystallization overvoltage. At short times this film is non-uniform consisting of dis crete nuclei which may be two-dimensional or three-dimen sional in nature. At longer times these coalesce and form a continuous film varying in thickness from one monolayer to thousands of monolayers. Metal dissolution occurs at E <^EX at a .rate greater than that at B \ E*.. A and B are most likely to be distinguished by the use of a rotating disc electrode . In order to derive the neces sary diagnostic criteria, the diffusion equations for dis solution-precipitation at a disc electrode have been set up and solved in an approximate manner analytically, and by an iterative procedure on a computer. The influence of rota tion speed on the onset of precipitation has then been determined. This behaviour contrasts markedly with the lack of rotation speed dependence anticipated for the solid state reaction. Gregers Bech-Nielsen 367 THE ANODIC DISSOLUTION OF IRON. V. SOME OBSERVATIONS RE GARDING THE INFLUENCE OF COLD WORKING AND OF ANNEALING OF IRON ON THE ANODIC DISSOLUTION BEHAVIOUR OF THE METAL. Gregers Bech-Nielsen Chemistry Department A, The Technical University of Denmark Lyngby, Denmark. In previous publications[1,2]a particular behaviour of iron in the anodic potential range was described. Two current maxima, both preceded by a Tafel-region are observed. The slope of the first Tafel-line is always lower than that of the second one,while together the various slopes cover the entire range from c. 30 to c. 240 mV. This phenomenon has been observed with electrolyte soluti ons based on a number of inorganic as well as some carbox- ylate salts. Phenomena of a similar type have also been observed with Co[3] and NiLHJ. An explanation consistent with the observations assumes a shift between coupled,par allel reactions caused by a coverage by or an equivalent distribution of reacting intermediates, so that, when the reactions following the two Tafel-lines be designated IT and 12, the observed current between the Tafel-regions' is i = Il(l-0]_) + l20i (I)» whereas the current after the? 12-line is i=l2(l-02) (2), ©i and ©2 being the two coverages. From extrapolated values of I± and I2 together with measu red currents the values of ©^ and ©2 can be calculated. Reasonable fit with Langmuir adsorption isotherms is ob tained, and from the shift of the isotherms 'due to changes in pH, potential and anion concentrations;the following compositions have been suggested: For,-;the intermediate responsible of the I^-reaction: Fe^CX)^/'where X is the active molecule or anion in the electrolyte, and for the intermediate causing the ^-reaction F.e*M0H~)q (X), which at the onset of passivation is transformed to Fe-^I(OH")5. The compositions of the above postulated species were cal culated by a simple equilibrium treatment based on the Nernst equation, but although kinetic data are employed, the kinetic characteristics of the reactions are elimi nated, when equilibrium data are calculated from(l) or (2). It is, however, clear that if an expression like(l)can be used to establish the dependence of ©^ on pH, potential and anion concentration, then 1^ and I2 must also be functions of the same parameters.The potential dependence of current is given by the corresponding Tafel-slope, and the pH- dependence of the anodic dissolution of iron has been re cognized for a long time, whereas quantitative studies of the involvement of anions (other than halides and notori ous inhibitors) have been rather few until lately [5]. Since it was shown in previous work [1,2] that perchlorate ions hardly interfere in the presence of more readily 368 Gregers Bech-Nielsen adsorbed anions, all solutions were made up to 1 M by addi tion of sodium perchlorate. Another problem of importance in connection with an exami nation of the kinetics of the iron dissolution is the phys ical state of the metal surface. Lorenz et al. [6j have shown that iron electrodes in sulphuric acid will tend to follow the mechanism suggested by Bockris et al. [7j when completely annealed, but will come closer to the behaviour typical of the Heusler-mechanism [8] when strongly cold- worked. In order to trace a similar tendency in the pres ent case most experiments were made both with iron an nealed in an argon atmosphere at 900°C for 6 h and with cold-worked iron. Acetate solutions were used for all experiments reported here. First the influence of acetate was examined at pH 4 and at the potentials -400 mV (nhe) -for the I^-reaction and -200 mV for the ^-reaction. The reaction orders, de fined by ( 3log I/81og[E~]')pH £ » were found to be: for I^: -0.65 to -0.70, for I2: -0.2Ö to-0.25. The bi~values(Tafel- slopes for Ij_) were 30-32 mV independent of the treatment of the iron, while the b2~values were from 120 to 135 mV. The concentration of acetate was varied between antilog (-3.12) and antilog (-1.74) M. Variation of pH from 4.4 to 3.3 at constant acetate concen tration (=0.005 M) revealed a steady increase of b^ from c.3 0 to 4 3 mV for both kinds of pre-treatment of the metal. A direct plot of logli versus pH at a selected potential giyes results influenced by the choice of potential, when there is a variation of b-values. Use of the relation: Olog 1/3 PH)£E.J)E= -CSlogI/3.E)EE.])pH.(3É/a pK).E.j;I(3) does not include substantial errors from extrapolation of Tafel-lines, when a constant current level is selected within the portion of Tafel-region common for all sweeps. The plots of potential versus pH at constant current were strikingly linear over wide ranges of ph both for constant anion concentration and for a simultaneous variation of the latter (simultaneous change of pH and anion concentration should and does yield the difference between the individual orders). The above-mentioned linearity is an indication of a con stant product of reaction order and b-value, consistent with the idea of a simultaneous kinetic involvement of po tential and pH. In the present case of pH-variation the re action order for pH was calculated at each pH value by means of (3), and the range of b-values resulted in a range of values from 1.7 to 1.2. As in the experiments of Lorenz et al. [6j the lower b-values were combined with reaction orders approaching 2 and the higher b-values with reaction orders close to 1. Gregers Bech-Nielsen 369 For the ^-reaction the reaction order for pH was found to be 0.20 to 0.23, i.e. close to the numerical values found for the reaction order for acetate. Another group of experiments conducted in the pH-range 1.3 to 2.3 and consequently at very low acetate concentrations only included measurements of 1^ due to current limitati ons of the potentiostat. Here the differences between an nealed and cold-worked metal were more pronounced, but at the same time there was evidence of a very low reaction order for acetate; the influence seemed negligible below a concentration of c. 10-l+ M. In this case a correlation of a similar type as found by Lorenz et al. [6] .for Tafel- slope, reaction order with respect to pH and phy^ic-al state of the metal was noticed. The observation made by lofa et al. [9J that adsorption of inhibitors is less complete on the more perfect crystal faces may be confirmed by the fact that the annealed metal showed by far the least response to changes in the lower , acetate concentrations. Finally, it should be noted that i) A negative reaction order for both 1^ and I2 with respect to acetate ion can probably only be explained by the assumption that the starting material for the reactions should always contain acetate ions irrespective of changes in the bulk electro lyte (i.e. complete adsorption, independent of potential). This conclusion is compatible with earlier results for ad sorption equilibria [1,2]. ii) As a result of heteroge neity (involving surface diffusion, for example [10]) relatively simple reactions can involve fractional reac tion orders. References: 1 H.Nord & G.Bech-Nielsen, Electrochim.Acta !16, 849(1970). 2 G.Bech-Nielsen & J.C.Reeve, 6th Scandinavian Corrosion Congress, Gothenburg 1971, ch.8. 3 G.Bech-Nielsen, Unpublished results. 4 G.T.Burstein, Chemistry in New Zealand _3§., 159(1971). 5 G.M.Florianovich, L.A.Sokolova & Ya.M.Kolotyrkin, Electrochim.Acta 12., 879(1967). 6 F.Hilbert, Y.Miyoshi, G.Eichkorn & W.J.Lorenz, J.Elec- trochem.Soc. 118, 1919(1971) (& ref.s to earlier work). 7 J.O'M.Bockris, D.Drazic & A.R.Despic, Electrochim.Acta 4., 325 (1961). 8 K.E.Heusler, Z.Elektrochem. 629 582(1958). 9 Z.A.lofa, V.V.Batrakov & Yu.A.Nikiforova, Corr.Sci. £, 573(1968). 10 J.C.Reeve, Thesis, University of London 1971. p.444 & 445. Konrad Bohnenkamp 370 ABOUT THE ELECTROLYTIC OXYGEN-CORROSION OF MILD STEEL Konrad Bohnenkamp Max-Planck-Institut fiir Eisenforschung, Diisseldorf, BRD. In the corrosion of mild steel in nearly neutral electro lytes generally the reduction of oxygen is the cathodic partial reaction. A strict proportionality of corrosion to the bulk concentration of oxygen was observed at small stirring rates (1)(2). The rate determining step is the transport of oxygen to the electrode by diffusion and convection. On rotating disc electrodes the limiting current densities of oxygen reduction are locally con stant and exhibit a good reproducibility. Experiments on mild steel showed (3) that the corrosion rates are independent of time and correspond to the limiting cathodic currents only in solutions with a pH lower than 3.5. In neutral solutions rust lay'ejrs are formed, the corrosion rates are lower and depend on time. That should be studied in detail. Current density-potential curves showed cathodic limiting currents on electrodes of clean steel in 0.25 to 0.5 normal solutions of NaCl, Na2S04, and Na3CöH507 in equi librium with air or mixtures of oxygen and nitrogen. The same values were measured With electrodes of platinum and magnetite. The dependencies on rotation frequency (n = 1 to 25/sec) and concentration of oxygen (c(02)~ P(02J) were found in accordance to i ^ V^'PC^) • Up to 60OC the limiting currents under the condition p(02) ~ (1-p(H20)) atm were nearly independent of temperature in agreement with D(02)2/3y--1/6 c(02) = f(T) (v means kinetic vis- , cosity). The potentials for reaching the limiting current densities were found different for the different electrode materials. They also depended on the pre-treatment of the surfaces and in most cases they did not remain constant with time. Scattering values were measured on specimens with rust layers. At 20°C the limiting currents were not reached at potentials less cathodic than that required for significant decomposition of water (E^ » -800 mV). The corrosion rates at 20°C in acidified solution of Na7S04 (pH » 3) were nearly equivalent to the limiting diffusion currents of oxygen. Linear dependencies on partial pressure of oxygen and on square root of frequency were obeyed. Corrosion potentials were about EH = -350 mV. Corrosion rates in neutral solutions of NaCl and Na2S04 at 20°C showed a decrease with time and became nearly con stant after 10 to 20 hours at potentials around EH = -400 mV. At n = 25/sec and the oxygen pressure of air they were lower by a factor up to 4 than the values Konrad Bohnenkamp 371 corresponding to limiting cathodic current densities on clean surfaces. No significant differences were observed between specimens of mild steel, pure iron and a low alloyed steel. The dependence on the oxygen concentration disappeared at higher partial pressures. Even in Na- citrate solution where no rust was observed, corrosion rates at higher frequency were found too low at least by a factor 2. At 60°C the corrosion rates in the mentioned neutral solutions approximately reached the values cor responding to the limiting currents of oxygen reduction. A specimen pre-corroded at 20°C, then polarized to EH = -500 mV, showed stationary cathodic currents at 60oc higher by a factor 1.7 than at 20°C. In distilled water the specimens did not corrode at n = 25/sec. Without stirring corrosion rates were found the same as in the mentioned salt solutions. In sea water the corrosion rate decreased over longer times unlike in solution of NaCl. The same.effect was found more sharply in tap water of Diisseldorf. The corrosion at high frequency (n = 25/sec) was sometimes lower than at the low frequency (n = 1/sec). Corrosion also occured in tap water equilibrated with solid CaCO^ and a gas mixture with 7.5 Torr CO2. Even after ca,thodic pre-treatments (EJJ = -700 mV) in which a layer of CaC03 was formed depressing also the cathodic current, pitting corrosion was observed. The corrosion potentials were around EH = -300 mV. Formation of very effective tarnishing layers was observ ed in 3»10"4 normal solution of Na2S04. Without additions corrosion rates were found nearly the same as in the more concentrated solutions mentioned above. In equilibrium with CaC03 and 7.5 Torr CO2 no further corrosion was analysed at n = 25/sec after the first day. Then the corrosion potentials were more noble than E^ = +100 mV. With n = 1/sec corrosion rates and corrosion potentials were found similiar. to the experiments in sulfate solut ion without carbonates. Changing the frequencies of rotation after some days no changes of corrosion rates and potentials were observed. The transport of oxygen through pores of the rust layers does not seem to be sufficient for the corrosion rates in solutions of NaCl and Na2S04« The observed dependencies on rotation frequencies together with the calculated thicknesses of rust layers suggest that the electro chemical reduction of oxygen must also occur on the rust surface (4). That must be assumed at least at 60°C. It could be supposed, that the cathodic reaction rate on the rust layer and the electronic conductivity of the rust layer cause the increasing rate of corrosion with tempe rature. In the citrate solution the main reason for it is 372 Konrad Bohnenkamp probably the increasing anodic current lowering the corrosion potential remarkably (E^ (20OC) = -380 mV, EH (60°C) = -550 mV). The anodic reaction (Fe Fe++) can proceed only in the presence of pores in the rust layer. The transport of iron ions through the pores occurs more easily than that of oxygen because of their electrical charge and the possibility of higher concentration gradients (Fe + +) . In solutions containing carbonates tarnishing layers are formed, which raise the corrosion potential. This shows that the anodic process is hindered more extensively than the cathodic process, 1) Evans, U.R., T.P. Hoar: Proc.Roy.Soc. London, Ser. A, 137 (1932) 343 2) Uhlig, H.H., D. Triadis, M. Stern: J.Electrochem.Soc. 102 (1955) 59 3) Zembura, Z., W. Ziölkowska: Bulletin de l'academie polonaise des sciences, serie des sciences chimiques. Vol. XIII (1965) No. 3 4) Herzog, E.: Metaux 500, 42 (1967) 133 Charles Booker 373 AN ELECTROMETRIC STUDY OF THE TARNISH FILM FORMED ON Cu-Zn ALLOYS IN STRESS CORROSION ENVIRONMENTS BASED ON AMMONIACAL CuS04 SOLUTIONS Charles Booker and Muhammad Salim Department of Metallurgy and Materials, Sir John Cass School of Science and Technology, City of London Polytechnic, London, U.K. Thin films formed on a range of a-brass substrates, containing 10, 20 and 30% Zn, in a 'Mattsson's'solution were quantitatively reduced in a deoxygenated 0.1 N phosphate buffer solution pH 7.02 at a current density of 20 uA cm~2. The films formed in the range of pH 7.2-7.3 indicate the presence of CuO in the predominant CU2O that constitutes the tarnish. The film formation follows a linear rate law iip to an estimated thick ness of 200 nm for 10% Zn, 300 nm for 20% Zn and 400 ran for 30% Zn in the alloys studied. The tarnishing growth rates for the initial linear part of the curves at ambient temperature (298 K ± 1°) are 7.4 x 10"^, 130 x 10~9 and 170 x 10"9 g cm"2 s"1 respectively. The rate of film formation increases with temperature. Arrhenius plots derived from experiments at 273, 283, 293, 303 and 313 K(±0.5°) give apparent activation energies of 51.13 KJ mol~l for 10% Zn, 53.30 KJ mol"1 for 20% Zn and 33.60 KJ mol~l for 30% Zn alloy. These high rates of oxide growth are not in accord with a mechanism based on a solid state diffusion model. ' The magnitude of the activation energies determined in this study, suggest that the overall process of film formation is under mixed control, activation control predominating where the oxide growth rates are linear, while diffusion control intervenes when ion transport through the porosity of the tarnish layer becomes the rate controlling process. In other experiments, the 30% Zn alloy was polarized potentio- statically to potentials positive with respect to the normal limiting free corrosion potential (Eh = 0.25 V). At 0.26 V the rate of oxide formation is increased, whereas the overall oxide growth rate is reduced at 0.28 V. The rate of oxide formation is greater on annealed than on cold- worked substrates, while an increasing degree of cold work does not appear to further influence the rate of oxide growth. The advantages of the electrometric technique for estimation of the tarnish film as compared with other methods are outlined and the results obtained are discussed in relation to current advances in our understanding of the mechanism for the stress corrosion cracking of a-brass. 374 Sven Brennert PITTING CORROSION IN STAINLESS STEEL: MECHANISM Sven Brennert and Göran Eklund Swedish Institute for Metal Research, Stockholm, Sweden Potentiostatic experiments in 0,1-m NaCl-solutions (fig. 1) show, that the potential-current diagram can "be devided into three different zones The passive zone. The current peak zone, in which the frequency of the peaks as well as the amplitude seems to increase at higher potential values. The current in the peaks increase slowly and often drop sudden ly (1)« Sometimes a real pit starts hut fails to propagate. The pitting zone in which the current is irregular hut tend to increase. Even here current peaks with sudden drops are superimposed. It is known, that the metal is active at the bottom of an established pit, and that the anode, solution in the pit is rich in chloride and hydrogen ions as well as metal ions, e.-g. (2). Tests performed with ar tificial pits, from which the anode solution (more or less diluted with 0,1-m Nad) was isolated, showed that the pX was lower, the higher the corrosion resistance of the steel. Carbon steel pH 3»2 Stainless steel 14 Cr 2.3 11 18 Cr, 9 Ni J 2.0 " 17 Cr, 12 Ni, 2.7 Mo 1.7 Experiments were performed with stainless steel in synthetic anode solutions made by dissolving the same steel in HC1. The pH of the solu^- tion was adjusted with 0,1-m NaCl-solution. When the active steel was polarized in steps starting from the potential of the active steel (fig. 2), we could observe a typical "Plade effect" at low pH-values, compare e.g. (5). The Plade effect disapears at higher pH where the steel is passive over the whole potential range (fig. 3). Discussion The steel is active in the pit because of the acid corro sion products. The condition for a pit to propagate is that the acid corrosion products form at least as fast as they are removed by diffu sion and convection. - Paradoxically it seems to be, that pits can not be formed unless an acid solution is present but an acid solution can not be formed unless the steel is corroding under the formation of pits. Consequently a pitting should not be able to start on a perfect steel surface even if the potential considerably exceedes what is generally called the pitting potential. A "spark" is required. A sul fide inclusion is often supposed to initiate the process (4). At poten tials higher than 300 mV SCE a sulfide inclusion may form a very corro sive solution. This solution may be more or less acid but its content of sulphide ions is probably the main reason for its corrosive proper ties. This solution will sometimes activate the adjecent steel surface and a very small pit will start followed by a local potential drop. The pit can now either terminate or propagate. When the pit terminates the potential drop means, that the formation of sulphide and hydrogen ions will be considerably reduced for electrochemical reasons, and the diffusion process will dilute the solution why the current density decreases. The hydrolysis of the metal ions formed by dissolution of the metal can not maintain the low pH value. The potential in the pit increases and the steel ..will he passivated,. often followed by a steep Sven Brennert 375 drop in current» When the pit propagates, H-ions are formed to such an extent that the dissolution reaction can proceed without the help of sulphide ions. If a piece of steel is corroding under practical conditions in 0,1 Hi! KaCl, the potential in the pit is supposed to he about -300 mV (fig?* 3). At the. same time the potential of the passivated main part of the surface could he e.g. +100 mV. The ohmic resistance close to the pit is supposed to account for this large potential difference. This could be explained by the fact, that the resistance between the active sur face of the pit and a large surface at a great distance is approxi mativ _P_ . _1_ p = resistivity of the solution 2TT r r = diameter of the active surface The consequence is shown in fig» 4» The active surface (the whole pit or a part of it) will adjust to such a size, that the requirements are met» (ij Schwenk, Corrosion 1964, p. 133 (2) Hospadaruk and Petrocelli, J. Elechtrochem. Soc. 1966, p. 878 (3} Lizlovs and Bond, J» Electrochenu Soc. 19^9* p. 1130 ! (4) Szklarska-Smialowska, Corrosion 1971» P» 223 pitting p a s s 1 v| current peak zone zone zone failing pitting pitting start 25 pA Fig 1 Potentiodynamic diagram in Q,1-m NaCl Steel 18/12/2,7 Fig 2 Potentiostatic diagram in synthetic anodesolution pH 1 Steel 18/12/2,7 r\ r\ VK V K CO CM IV. CM CM CM rv. CM IT) m LT> in in m in CM in CM CO CO CO CM CO CM CO CM en CM CM I I I 1 1 1 I i I 376 Sven Brennert pH 0,5 ~1,5-m Cl Fig 3 Patentiostatic diagram in synthetic anodesolution with different pH. Steel 18/12/2,7 P= pitting starts. Schematically 0,1-m Cl 1000 1 T -200 200 300 mV E SCE Fig 4 Pitting Cell Equipotential lines 0,1-m NaCl + 0^ Stainless Steel 18/12/2,7 Cathode Anode E B, + 00 mV SCE 2" 1 ESCE -3D0 mV pH__ = 6 pH_ <1 CI = 0,1-m CI >1-m Lennart Dahl 377 ELECTROCHEMICAL STUDIES CONCERNING CAUSTIC STRESS CORROSION CRACKING OF ALLOYED STEELS Lennart Dahl and Tommy Dahlgren. Section for Corrosion and Reactor Chemistry, AB Atomenergi, Studsvik, Sweden. The stress corrosion cracking (SCC) of alloyed steels in alkaline en vironments has been given much less consideration than the similar attack on austenitic stainless steels, caused by chlorides. In order to investigate how and why steels crack in caustic solutions at high temperatures and pressures we have made some experiments on SCC in 20 % NaOH at 225 C. Most of the specimens in these experiments have been electrochemically polarized. This technique was chosen because it leads to short time to failure at suitable applied potentials and because our intention has been to study the electrochemical mechanism of the attack. The equipment (fig. 1) used consists of a 10 litre high-pressure Inconel-lined autoclave, designed for max. 350 °C and 170 atm. The autoclave is connected to an outer circulation system, which permits experiments at/flow-rates of 1-2 litres/h.. The electrochemical measure ments have been made using an Amel 551/SU-potentiostat, a platinum net being used as counterelectrode. An inner reference electrode was used, made with HgO(yellow)/Hg, which shows good stability and repro ducibility in alkaline solutions up to 225 °C or higher. Electric conductors in the autoclave were enclosed in teflon. A lever system of Inconel 600 has been constructed, to apply the desired mechanical load on the specimens, correction being made for bouyancy in the sodium hydroxide solutions. Potentiodynamic polarization curves at 225 °C (cf. fig. 2) have been recorded for an unstressed specimen in order to choose suitable potentials for SCC-testing. The SCC-tests have been made at the same temperature with the specimens stressed uniaxial ly at loads corresponding to the yield.limit at the test temperature of the 3 Cr-0.5 Mo-steel, at anodic, catodic and free corrosion potentials. Before testing the hardness (Ry IQQ) of each specimen has been measured, range 270 - 325 F.v IQQ. The weight of the specimen has been determined before and after testing and the calculated weight losses have been compared with those obtained from anodic current measurements by means of Faraday's law. In many cases the general corrosion rate calculated from weight loss is considerably higher than that calculated from the anodic current, which may indicate that local electrochemical processes have taken place. The hydrogen content at different places on the spe cimens has been 'analyzed after each run. The stressed region of the specimen generally contained more hydrogen than an unstressed region. The sodium hydroxide solution has been analyzed after each run with respect to 0H~, Cl" and total Fe. The mean composition (with max.error) was 6.1 +_ 0.2 N NaOH, 1.2 +_ 0.4 ppm Cl" and 3.3 +_ 2.1 ppm Fe. Fast stress corrosion fractures (<10 h) are obtained at cathodic (-900 mV) and at free corrosion potential. The same is true for speci mens under slightly anodic conditions (-720 mV), if the hardness of 378 Lennart Dahl the specimens is higher than ^320 Hv ^QQ. At more anodic potential (-500 mV) pits or shallow cracks form and at + 0 mV the main attack is general corrosion. Times to rupture versus potentials are plotted in fig. 3. By microsond analysis of cathodically polarized specimens it was found, that selective oxidation of iron occurs on the metal surface to a depth of about 15 urn. The experiments suggest, that the attack of 20% NaCH on stressed specimens of 3Cr-0.5 Mo-steel in the potential range -900 to -720 - which falls within the stability range of HFeO^ at the pH and tem perature considered - must primarily be anodic in nature, i.e. some general corrosion must occur. This is also shown by the experimental fact that only specimens which are subjected to a corrosion rate exceeding 5-10 g/m^h fail. The anodic attack will probably be oug- mented by the observed selective oxidation of iron, which can result in an enobling of the surface and an increase of the potential dif ference between the crack root and. the surface of the specimen. On the other hand the increasing hydrogen content of the specimen seem to shorten times to fracture, e.g. for specimens with a hardness of 300 - 310 H 1f)n a maximum of hydrogen content corresponds to a minimum of time to failure. It is then reasonable to suppose that caustic SCC of this steel is due to an interaction of anodic and cathodic processes, in which cathodically evolved hydrogen is of importance. Similar studies are being carried out with 13 Cr-, 13 Cr-1 Mo and 18 Cr-8 Ni-steels. Acknowledgement. This investigation was supported from the Swedish Board for Technical Development. Lennart Dahl 379 REFERENCE ELECTRODE COUNTER ELECTRODE WORKING ELECTRODE CONSTANT PRESSURE REGULATOR FIGURE 1 « CATHODE ANOOE i 5 . HAMMMlTMMVtCMM I HANOMUiMMMVICMM 0 MMONtallMMVNMIM 20mV/Mltt • «J _l . « _ '••••'* .1000 t -tOO , fc«w FK3UHE} ELECTHOOE POTENTIAL ImV) V.S.HtfH|3 r MUM j iLicrnoof HJTINIIAI MVI VA Nome 380 Jacques Dubois ELECTROCHEMICAL REACTIONS AND CORROSION IN MOLTEN ALKALINE SULFATE AND .VANADATE J» Richard and J« Dubois E»D.P. Direction des Etudes et Recherches» 78-Chat0U| Prance» High temperature corrosion mechanisms by molten alkaline compound$ in residual oil fired-boilers may be interpreted on electrochemi cal basis* In this purpose»results from studies related to acid-base reaction and oxidation and reduction properties of molten sodium metavana- date and NaVO, - Na2S0. eutectic mixture (14 i mole Na2S04) at .650 C are reported» The behaviour of various elements and heat-exchanger (superheater) materials are analysed in connection with the acidity conditions in these solvents» Liliana Felloni 381 INVESTIGATION ON THE SECOND ANODIC CURRENT MAXIMUM ON THE POLARIZATION CURVES OP COMMERCIAL STAINLESS STEELS IN SULPHURIC ACID Liliana Felloni, G.Paolo Cammarota,, Silvana Sostero and G.Luigi Zucchini. Istituto Chimico, Centro Studi Corrosione "Aldo Daccb", Universitå di Ferrara and Istituto di Metallurgia, Facoltå di Chimica Industriale, Universitå di Bologna, Italia. For stainless steels, the anodic curve and particularly the second current maximum,.observed in some circumstances,are still subject to question. As summarized recently (1), the occurrenas of a second maximum during anodic polarization of stainless steels in sulphuric acid solutions was explained by: i) surface nickel.enrichment during the pre-polarization exposure (2);, ii) oxidation of absorbed hydrogen, resulting from cathodic pre-treatment (3), iii) preferential attack along phase and grain boundaries of specimens containing f errite and martensite in the microstrue ture (.1), iv) sup posed precipitation of Cu or Mo in the martensitic stainless steels (4). The present study was undertaken: i) to find which of the above mentioned different processes could be responsible for the occurrence of the second maximum in the potentio- . kinetic polarization curves ( scan rate 2.4 V/h) of some commercial stainless steels examined, ii) to see if there could be one single explanation of experimental facts. The behaviour of AISI 430 (ferritic structure), 302, 304, 304L and 316 (austenitic structure containing some ferrite) D20 (Durimet, cast type, dentritic structure) and Ni (an nealed) was examined in 5 N and 1 N sulphuric acid saturat ed with nitrogen at 25°±1°C. The alloys studied were not heat-treated and they were used in the as-received condition. The chemical compositions, the percentage of ferrite per sq.cm measured on every type of 300 series, dissolution ra tes and electrochemical parameters are given in tables. The quantities of Fe, Cr and Ni in the corroding solutions were determined by atomic absorption spectrometric techni ques for samples kept at zero current potential and pola rized at a given potential. The fitness of the different methods used for the determi nation of the corrosion rate of the examined systems was analyzed. Experimental conditions where the second anodic current was found are described. For all types of alloys, the dissolution rates were measured at -600 mV vs SCE, and zero current potential. In some cases measurements were also made at -500 mV, -350 mV, -25O mV and + 1 V. The following results were obtained: 1 ) the amounts of Fe, Cr and Ni determined in the acidic 382 Liliana Felloni solutions were in the same ratio as'in the alloys. 2)The corrosion rates in 5 N sulphuric acid were related to the alloy nickel content and decreased in the order: 430 » 304 > 302 > 304L /• 316 ? Ni > D2o In the case of types 304 and 302, the inversion of the order might have been caused by a higher content of Mn in the former than one in the latter, while in Ni and D20 i"b "was obviously due to the presence of chrome. 3)In the systems examined, a cathodic pre-polarization did not actually influence the corrosion rates of the alloys. 4)The corrosion rates and the attack type depended on the imposed potential value. The shape of the anodic polarization curves was affected by: 1)the presence of the plastically deformed layer produced by abrasion treatment, 2)cathodic pre-polarization particularly for Ni and types D^o and 316 as regard the range of active-passive transi tion, 3)exposure time before the anodic polarization both in active-passive transition and the second maximum. Por type 430, the second anodic current maximum was always noted but its extent was dependent on the processes which had previously taken place on the sample surface. Por the 300 series, in the range of 2nd current maximum only one step appeared if the anodic curve was recorded starting from cathodic polarization, imposed immediately on the immersion of the specimen. The samples were polarized for 48 hrs at So» electrode po tential corresponding to the 2nd anodic current maximum, with three different starting states: i) on immersion, ii) at -600 mV after 1h, iii) a controlled time of free corrosion. As to the appearance of the. 2nd maximum, the time and the quantities of electricity required for reach- • ing the passive state as well as the microscopical observa tions of the etching morphology provided some helpful information and the following conclusions may be drawn from the data obtained: 1)A surface nickel enrichment is not confirmed experimental ly and, even if occurs, its contribution should be very small. 2)The desorption and oxidation of hydrogen are not directly responsible but these processes can play a role in the formation of an unstable passive film and their influence is closely related to the surface characteristics i.e. microroughness and chemical microheterogeneity. 3)The presence of different phases affects the absorption and diffusibility of hydrogen and the chemical composition and/or morphological structure of the passive film. 4)The data obtained suggest that the occurrence of the 2nd anodic current maximum is caused by different factors (simultaneous reactions, surface disarray, different phases) which all tend to produce an unstable, non-honio. geneus passive film. Liliana Felloni 383 5) -The appearance of the 2nd maximum is a sign of weakness that can or cannot actually exist, since it is controll ed "by the relative rates of simultaneous processes. REFERENCES 1) M.B.Rockel- Corrosion, 2J_, 95 (1971) 2) W.R.Prance and N.D.Greene- Corrosion, 24, 405 (1968) 3) C.D.Kim and B.B.Wilde- Corr.science, 10, 735 (1970) 4) T.Moisio and H.Mannerkoski- Corr.Science, £, 12.9(1969). 384 A. Giraudeau ANODIC OXIDATION OF RHENIUM A. GIRAUDEAU, P. LEMOINE, M. GROSS Laboratoire d'Electrochimie et Chimie Physique du Corps Solide, Université Louis Pasteur, 4, rue Blaise Pascal 67 - STRASBOURG - ESPLANADE (France) We have studied, by potentiodynamic and galvanostatic methods, the oxidation of rhenium in aqueous solutions in order to determine the re action mechanism which was unknown up to now. 1. General shape of polarization curves - When the electrolyte is a binary aqueous solution of an acid (H2SO4 or HC1), of a salt (Na2S04 or NaCl), or of a base (NaOH), the shape of oxidation curves is the same over the range of studied pH (Cologio CH+ = - 1,3 to Colog10 CH+ = + 14).. - In acid medium, the adjunction of the salt (NaCl 1M) has no effect and does not change the shape of the polarization curves. - In basic medium (Cologjo CJJ+ = + 11 to CologjQ Cg+ = +13, in other words CQH- = 10~3 M to CQH- = 10~1 M) and in the presence of a salt (NaCl 1M) a limiting current of diffusion appears, belonging to the 0H~ ions. The diffusion limiting current is directly proportional to the OH concentration. In any case, the position of the curves on the potential scale chan ges with the pH of the electrolyte..We have studied the evolution of the dissolution potential as a function of the pH of the electrolyte, for all the curves. This potential has been measured on the polariza tion curves for a current density of 1 uA/cm . The value of this po tential changes linearly with pH, in pH range of pH = 0 to pH = + 7 and pH = 10 to pH = 12, according the following relations : 0 Finally, in the range of negative values of Cologjo CH+ (0 to - 1,3) the slope of the curve U = f(pH) rises considerably, this shows an important slackening in the anodic corrosion mechanism. 2. Remarks During the electrolysis no coloured ion of rhenium appears. The surface of the electrode looses his bright metallic aspect and changes to a dark color. 3. Nature of the oxidation products The oxidimetric analysis of the electrolyte after the quantitative ano dic oxidation of a rhenium electrode shows the existence of a rhenium^ ion in the solution, with an oxidation degree of + 7. We have identified the perrhenate ion by spectrophotometric absorption in ultra-violet. On the other hand, the coulometric measurements show that the passage of metallic rhenium in perrhenate ion is quantitatively done with a 100 % yield, with the final reaction : + (I) Re + 4H20 •* ReO" + 7e + 8H (In acid medium) (II) Re + 80H" + ReOT + 7e + 4Ho0 (In basic medium). A. Giraudeau 385 This result is confirmed by the analysis of the electrolyte, after the oxidation of rhenium, by means of the atomic absorption spectrophoto metry. - Finally the rhenium electrodes, after polarization, have been stu died by electronic diffraction and we have only been able to obtain the diagram characteristic of. metallic rhenium. 4. Tafel Lines The semi-logarithmic representation of the experimental curves allows us to obtain the Tafel lines. In acid medium, for a range of current densities between 100 uA/cm^ to 15 mA/cm2 , these lines have all the same slope in an acidity range defined by - 1,3 < CologjQ CH+ < 0,3. Out of this interval, the slope decreases when the pH of the solution increases, until the pH = 7. - In basic medium, the semi-logarithmic representation is linear over a large range of current density and rhenium undergoes an anodic po lishing. The apparent values of an, obtained from the Tafel lines corresponding to different pH are : an = 1,2 - 1,3 < Colog]0 CR+ < 0,3 0,6 < an < 1,2 0,3 < pH < 7 an = 1,1 11 < pH < 12 The metallic rhenium is very strongly polarizable, as the weak values of transfer coefficients show it. 5. Time - potential equilibrium curves We have attempted to measure the potential U of the electrode as a function of time, at constant current density. In these experiments, the electrode potential increases rapidly and regularly, until the equilibrium is reached, corresponding to the value which can be read on the polarization curve, for the imposed value of the current in the corresponding medium. When the current density increases, then the electrode reaches the equilibrium more rapidly. 6. Discussion The time-potential equilibrium curves show a dissolution without for mation of stable film, or without formation of any film. These results confirm the studies carried out by electron diffraction on the surface of the electrode, which has not shown the appearance of an oxide film. Newertheless, the study of curves and Tafel lines makes us to think that the mechanism of the oxidation of the metallic rhenium is not as simple as show the final reactions (I) and (II) and we try to bring an explanation of the anodic oxidation. Bibliography 1) S. TRIBALAT et Dj. M0FIDI, C.R. Acad.Sc. Paris 1964, 258, 3477 2) V.O. LAVRENKO, Zhur. Fiz. Khim. 1961, 35, 198. 386 Irène Guillaume INFLUENCE D'UN MILIEU BACTERIEN SUR L'EQUILIBRE ELECTROCHIMIQUE DE L' ALUMINIUM Jean Brisou, Marie-Josèphe Croissant, Jane Grimaudeau, Irène Guillaume, Gabriel Valensi. Laboratoire de Thermodynamique Chimique et Electrochimie, Université de Poitiers, 40,Avenue du Recteur Pineau, 86 Poitiers, France et Laboratoire de Bactériologie,Université de Poitiers, rue Sainte Opportune, 86 Poitiers, France. Au cours de recherches antérieures, nous avons examiné les condi tions de corrosion et d'immunité du cuivre (l)(2) du nickel et du zinc (3) dans un milieu bactérien doux et marin. Les germes sélectionnés appartiennent aux différentes classes bactériennes et présentent des propriétés très variées. Nous avons également envisagé l'influence du milieu nutritif (2). Pour ce faire, nous avons utilisé successivement le stéphenson, la peptone, un mélange de succinate et de peptone, l'eu- gon additionné de glucose. Enfin nous, avons établi les courbes de croissance des bactéries ainsi que leur consommation d'oxygène en pré sence de sels métalliques (4). Nous avons ainsi pu établir que les germes réducteurs, convenablement nourris et se développant correcte ment sont inhibiteurs de corrosion. Cette étude a été étendue à l'aluminium dans lés mêmes conditions et pour les mêmes bactéries. Le pH initial de nos milieux étant uni formément fixé à 7, nous nous trouvons normalement dans les zones d'immunité ou de passivation du diagramme dé Pourbaix (5) valable en l'absence d'ions Cl" et de produits complexants. En fait, à 20°C, nous observons, soit une inhibition totale de corrosion, -soit un dépôt adhérent protecteur sur le métal. Les résultats sont plus variés à 37°C et nous nous limitons à ceux- ci dans le tableau ci-joint. La corrosion de l'aluminium se distingue nettement des phénomènes observés pour le cuivre, le nickel et le zinc. Pour que des ions Al"* apparaissent en solution, il faut que le pH s'abaisse à 4>7 ce qui est possible dans l'eugon. Généralement, la concentration en ions A13+ est nulle, et la corrosion apparaît du fait de la non homogénéité de la couche protectrice. L'attaque a lieu en certains points privilégiés où se forment des excroissances composées d'alumine et de cadavres microbiens. Ces ex croissances grossissent tandis que la plaque se creuse et que les bul les gazeuses y prennent naissance. Un tel phénomène est observé d'une part dans les trois milieux sté riles, d'autre part dans les bains où le germe est mort. Par contre, lorsque la bactérie subsiste, le métal est protégé. Par ailleurs, les courbes de croissance, la consommation d'oxygène ont montré que les bactéries sont très sensibles à la présence d'ions Al3+. Les taux minimaux d'ions susceptibles, d'une part de stopper Irène Guillaume 387 l'activité enzymatique et d'autre part de provoquer la mort des germes ont été établis. Les germes convenablement nourris et se développant normalement se raient là encore inhibiteurs de corrosion, mais ils sont beaucoup olus vulnérables aux ions Al3+. Une faible quantité de ceux-ci suffit pour faire disparaître la bactérie entraînant alors une corrosion dispersée mais profonde. (1) Electrochimica Acta 1970,15,1445-1454 (2) Compte rendu technique du CEBELCOR Janvier 1971 (3) Symposium on biological aspects of Electrochemistry Rome 1971- 18-195-204 (4) 1ère. Réunion mixte SFE-SIE Poitiers avril 1971 (5) M.Pourbaix Atlas d'équilibres électrochimiques Gauthier-Villars 1963.p.171 Légende du tableau : P = peptone S = succinate E = eugon V = bactérie vivante M = bactérie morte 3+ Pas. d'Al dans la solution Al3+dans la solution Bactérie Plaque intacte DépSt sur plaque Excroissances,bulles Excrois sances classe n° doux mer doux mer doux mer doux mer pH pH pt pH pH pH pH pH sterile P 7,4 P 7 S 7,4 S 7,2 E 5,7 E 6 A23789 PM 7,3 PM 6,9 SV 7,3 SV 7,3 EV 5,5 EM 4,9 A23860 PV 7,2 SV 6,9 SV 8,5 EV 6,5 EM 4,6 Blb3552 PV 7,3 PV 7,1 SV 7,1 SV 6,8 • EV 6,0 * EM 4,6 B26621 PV 6,9 PM 7,3 SV 7,5 SV 7,2 EV 6,8 EV 6,0 Cla3934 PV 7,0 PV 7,2 SV 7,4 SV 7,0 EM S.3 EM'^1 C2 5105 PV 7,1 PV 6,9 SV 7,3 SV 7,3 EV 6,2 EM 4,5 C1a3941 SV 7,2 SV 6,9 PM 7,4 PM 7,1 EV 6,5 EV 5,2 C 4011 SV 7,6 SV 7,1 PV 713 PV 7.0 EM 6,8 EM 5,3 D26023 PV 7,1 PV 6,9 EM 5,6 SV 8,1 EM 4,7 SV 7,3 Jozsef Horväth 389 HIGH TEMPERATURE POTENTIAL/PH EQUILIBRIUM DIAGRAMS OF METAL/SULPHUR/ WATER TERNARY SYSTEMS WITH APPLICATIONS TO CORROSION IN AQUEOUS HYDROGEN SULPHIDE ENVIRONMENTS J. Horväth, G. Bencze and F. Märta Institute of General and Physical Chemistry, University of Szeged, S zeged, Hungary Potential/pH equilibrium diagrams of metal/sulphur/water ternary systems developed for 25°C have been discussed in previous publica tions. In this paper potential/pH diagrams for the above systems are presented at 25, 50, 75, 100 and 200°C. The thermodynamic cal culations involved in deriving these diagrams make use of the Corre spondence Principle of CRISS and COBBLE. The method of the calcula tion performed by the authors was.similar to that used by BIERNAT and ROBINS in the case of high temperature potential/pH diagrams for the binary sulphur/water system. Potential/pH diagrams are presented indicating the stability domains of. the sulphides of Zn, Fe, Ni, Pb, Sn, Cu and Ag at different temperatures and ^S concentrations. On the basis of the changes in the stability domains of metal sulphides, especially in those of iron sulphides^it seems possible to explain certain corrosion phenomena observed in sour oil wells and oil re finery plants operating at high temperatures and/or pressures. To investigate the effect of sulphide layers on corrosion rate potentio- static polarization measurements were carried out in aqueous H2S solutions under different experimental conditions. Differences in the kinetics of corrosion reactions are interpreted on the basis of potential/pH diagrams. 390 Romuald Juchniewicz SOME EEECTROCHBMICAL ASPECTS OP PAINT COATINGS RomuaId Juchniewic z Gdansk Polytechnic, Gdansk, Poland» Many cases of accelerated destruction of paint coatings on some steel structures, especially ship hulls, have been observed in practice* Little or no attention has been paid to the fact that a given paint system can cause a corrosion macro-cell to form on a steel surface» It was reported/1/ that typical ship hull protection systems gave rise to potential^sdifferencesjfbetween%*hei underwater and above-water^partsy ^ with the underwåter^;par^as;;ja^ con dition of continuous mpisten^j|gix>f & a hull with sea water, the most intensive corrosion occurring in the water-line area. -•'••• :'h^P&*-' ' y> •. - • . Evans polarization curves testing of the cell consisting of two steel plates, one covered with chromate paint and the other with aluminium dust pigmented paint, immersed in 3% NaCl have yielded a potential difference of 150 milivolts and a corrosion current density of 0*22 micro- amps per sq*cm* Approaching the actual conditions further tests were made with a macro-cell consisting of one steel plate, one half of which was covered with passivating chromate paint and :thet-6tnei|^ pig mented paint* The greatest attack on the two paints border-line occurred when Al-pigmented coating was immersed while chromate coating was continuously sprayed with a jet of electrolyte* In one test series a Zn galvanic anode was inserted, into the part of plate covered with underwater Al-pigmented paint* Anode consumption rate was found to be increased fivefold as compared to the case of protecting the plate covered with Al-pigmented paint only* It is in a good accordance with analytical considerations of the tri- electrode system: steel with chromate coating/steel with Al-pigmented coating/zinc in sodium chloride solution* The above fact has an appreciable practical meaning since such an increase in galvanic protection current may be a cause of secondary electrode reactions on a steel surface leading to detrimental changes in paint coatings, e*g* saponification* External factors such as electrolyte aeration, tempera ture changes, and UV irradiation of coatings, have been found to have no essential influence on the operation of a corrosion macro-cell* Accelerated destruction has been also observed in cases of multilayer coatings composed for example of zinc dust pigmented priming and chromate pigmented top or inter mediate coatings* Romuald Juchniewicz 391 The above discussed facts need further investigation, nevertheless so far obtained results of laboratory tests and practical observations allow to put forward a con clusion that painting technology should avoid systems enabling corrosion macro-Cells to form on steel* References* 1# R. Juchniewicz, Corros. Sci.f 12, 91 /1972/. 392 V.Losev CORROSION OF STEPWISE DISSOLVING METALS; CORROSION MECHANISM OP INDIUM V.V.Losev and A.P.Pchelnikov Karpov Institute of Physical Chemistry, Moscow, USSR, Corrosion of a metal dissolving stepwise with formation of low-valency intermediates M+ and W1* ions as the final corrosion product proceeds by a complex electrochemi cal-chemical mechanism (EC-mechanism) comprising folio- , wing steps: e , (1) + (n-l)e, (2) + n+ M +(n-l)Ox -HP- M + (n-lOO", (3) 0x+e -*•> Ox (4) The anodic reactions (1) and (2) are compensated by the cathodic reduction of oxydant 0 (4), but the intermedia- te M can be further oxidized also by a direct chemical reaction with O (3) and can diffuse into the bulk of + the solution with the rate v^ = (D/£~) [M JB . At the corrosion potential the sum of anodic reactions rates is equal to the cathodic reaction rate: Fv + (n-DFVj^ s. Fv^ , (5) where v, v^ and v. are expressed in grammatom/cmsec. Under steady state conditions the formation rate of M+ ions which coincides with the corrosion rate i is equal to the sum of their further oxidation and diffusion rates v > v7 + v2 + Vo • (6) From (5) and (6) follows the general expression for the corrosion rate nFv s Fv4+(n-l)F(V2+v3). (?) There are two limiting cases: A. Vi^v; + v-a, i.e. M+ ions are oxidized only electro- chemically. From (6) we have v=v1 and (7) becomes nv a v.. (8) Hence the dissolution of one grammatom of the metal is accompanied by the reduction of n grammatoms of the oxy dant. Therefore the corrosion rate i (expressed in elect* rical units) is equal to the rate of the coupled cathodic 393 V. Lo se v process found e.g. by extrapolating the cathodic polariza tion curve to the corrosion potential (i®x*tr): i = nPv = Fv„ = i®xtr (9) c 4 c w/ and corrosion proceeds by the common electrochemical me chanism. B. vi«vp + v3» i*e* M ions are completely oxidized *— by the chemical reaction (3). From (6) we have v=v2+v- and (7) becomes v = v4 . (10), Hence in this case the dissolution of one grammatom of the metal is accompanied by the reduction of only one gramm atom of the oxydänt. Calculating the corrosion rate in the same manner is in the case A - from the expression i =nPv (i.e. assuming that the final corrosion product (Mn+) is formed only electrochemicallyxywe obtain from (10) i = nPv = nFv, » ni®xtr , (11) c 4 c I i.e. the corrosion rate is n-fold of 1* . Therefore the deviation of i^ext r from the directly measured true corro- c . sion rate ic is a criterion of the said complex EC-mecha nism of corrosion /l/. The extent of deviation of the ra tio iVi®xxr from 1 is a quantitative measure of contri- c c but ion of the EC-mechanism to the overall cprrpsion rate. Indium has been chosen for the experimental investiga tion since its anodic dissolution is known to proceed stepwise with formation of intermediates In+ which are oxidized by hydrogen ions with the formation of stable In^+ ions /2-4/. The corrosion rate (i_) of a radioacti- c ve indium electrode has been determined from the slope of continuosly recorded solution radioactivity-time curves in solutions xMHClO. + (3-x)M NaClO. (x = 0.01-1) and cathodic polarization curves have been obtained. As x can be seen from fig.l ic is much higher than i£ and xtr the ratio i c/i® c approaches 3 with an increasdii n the acid concentration. Hence in acid solutions the inter mediates In+ are practically completely oxidized by hyd rogen ions (3)» i.e. the corrosion proceeds by the x'This method of calculating i is somewhat arbitrary but it is convenient when the rates of (2) and (3) are comparable. 394 V.Losev EC-mechanism. When.0.2M NaCl is added i /iextr =1 (inde- c c pendent of pH), hence indium corrodes by the common elect rochemical mechanism. This effect is caused by the acce leration of the anodic oxidation of In+ by Gl~ ions /2,5/. 1. V.V.Losev, A.P.Pchelnikov, Elektrokhimia, 4, 264(1968). 2. A.P.Pchelnikov, V.V.Losev, Zashch.Metall.,~1, 482(1965) Proc. Third International Congress on Metallic Corro sion, vol.1, 111, Moscow (1969)» 3. R.E.Visco, J. Electrochem. Soc, 112, 935 (1965). 4. V.V.Losev, Electrochim. Acta, 15,"T(595 (1970). 5. A.P.Pchelnikov, V.V.Losev, Elektrokhimia, £, 284(1969). Y.v ' -Q6 -0,5 -QA -05 -6 -5 -A fyt (a/cm*) Fig.l, Cathodic polarization curves for different x: 1 - lxlO"2; 2 - 6xlO~2; 3 - 0,27; 4 - 1.0; I»-4» - corresponding i_ values; la - par- c tial polarization curve of hydrogen evolution for x = 0.01. 395 Bruno Mazza INFLUENCE OF GOLD PLASTIC DEFORMATION ON CORROSION RESI. STANCE OP AUSTENITIC STAINLESS STEELS IN ACID AGGRESSIVE MEDIA AND ON THEIR SUSCEPTIBILITY TO THE PITTING- CORRO SION Dany Sinigaglia, Pietro Pedeferri, Luisa Peraldo Bicelli, Bruno Iflazza and Gianpaolo Galliani Istituto di Chimica - Fisica, Elettrochimica e Me^ftallur- gia del Politecnico di Milano, Italy Centro di studio del C.N.R. sui processi elettrodici. The influence of cold plastic deformation on corrosion resistance of some austenitic stainless steels in acid aggressive media has "been studied with reference to the following points: - deformation type, i.e. by cold-drawing 6r by tensile stress; - relative orientation of the exposed surface with the deformation direction, i.e. parallel or perpofLicular; - type of austenitic slalnless steel, i.e. AISI 304, 304 L and 316; - environmental conditions, i.e. temperature, acidity and chloride ions content. The evaluation of corrosion rate in the active condition has been affected by traditional weight-loss and by Stern- Geary type methods. Anodic and cathodic polarization curves and characteristic parameters of passivity have been determined. The influence of deformation degree is always notable and is particularly strong in the case of AISI 304 L when the exposed surface is perpendicular to the defor mation direction. The susceptibility to the pitting corrosion has been te sted by chemical;etching in a suitable solution and valued by assuming as indicating parameter the pit den sity. In this case the influence of the deformation de gree has been studied in relation to the temperature of surface oxidation of the stainless steel (i.e. 25°C, 150°^ 300°C) as well as in relation to the stainless steel type and the relative orientation of the exposed surface with the deformation direction. Moreover the susceptibility to the pitting nucleation in the different conditions has been correlated to the elec- 396 Bruno Mazza tronic properties of the oxide films, i.e. their n-or p- type conductivity, by means of anodic and cathodic pola rization curves in the redox system Pe (CN)^~/ Pe (CN)|T The deformation degree has a very strong influence on the susceptibility to the pitting nucl eat ion in the case of AISI 304 L oxidized at 300°C with the exposed surface parallel to the deformation direction. Reference (1) G. Bianchi, A. Cerquetti, P. Mazza, S. Torchio, Corrosion Science 10, 19 (1970). Gérard Pinard-Legry 397 PITTING EVOLUTION ON A STAINLESS STEEL IN AN OXIDIZING CHLORIDE MEDIUM H. Coriou A. Monnier, G. Pinard-Legry and G. Plante Commissariat \ I'Energie Atomique, Service d'Etude de la Corrosion et d'Electrochimie, B. P. n° 6, 92 -Fontenay-aux -Roses -France 1. Pitting corrosion in an oxidizing chloride medium» Austenitic stainless steels may in some media give rise to a particu lar type of pitting corrosion. This is characterised by the formation of large cavities which are therefore particularly suitable for studying the phenomena occuring inside the pit. The experiments concerned very low. carbon steels of the Cr 18%, Ni 10%, and Cr 17%,Ni 13 % Mo 2. 5% to 3% types. Experiments were made at room temperature in an oxidizing solution (CI : 1 ion g. 1 ; Fe'+ : 10"1 ion g. 1" ; pH = 1. 7) on samples cut from a one millimeter thick sheet of stainless steel, annealed at 1100°C and air cooled. On the specimen whose surface was slightly cold-worked by abrasion with 120 mesh paper, the attack developped as triangular "cavities generally separated from the solution by a fine porous metal film. The corrosion products taken from these pits are composed of a green solution which is more dense than the attacking medium and by a par tially magnetic metal-like powder containing ferrite and carbides of the M7C3 kind (a = 14. 11 Å ; § = 0. 324). Further, the Cl~ content and pH are of the same value inside and out side the cavities. 2. Electrochemical study of pit growth. By means of an original device the exchanges of current existing bet ween the passive and active areas of the steel were evidenced. The samples were electrically connected through very low resistances (1 ohm), so that they could be considered as areas of the same sheet. On each sample, an anodic reaction occurs (dissolution localized within pits) as well as a cathodic one (reduction of Fe5 ). If dissolu tion prevails on one of them, that sample receives a positive current from the other specimen. On the contrary, a negative current would indicate that the cathodic reaction prevails. (The algebraic sum is constantly equal to zero). 2* 1. Fluctuations" artificially induced. The tests now described were conducted after 40 hours of immersion. When the film separating the cavity from the solution is destroyed, a sudden decrase of positive current on that sample is observed (fig. la, point A). Consequently, this destruction induces a selective standstill of the dissolution process, any slowing down of the local cathodic reaction being negligible. If we assume the active surface to be about 2 mm^ we should estimate the current density through that surface 398 Gérard Pinard-Legry 7 to be about 250mA. cm" . Such a value compares with those found by Rosenfeld and Danilov (l) . 1 „. ~+4- A .'"', *- * 44 < ,' fc .<• 5*2" UJ s 3cr A Figure 1 O 0 ' r "=rrrr 0-. -2. B »"*" u1 , • ,* _,<•**•""•—• - " i -4. . -* 40h TIME iQmn 40h TIME 10 "in 40 h TIME 10 mn a ' ' b' ' c Let us now consider a sample receiving a positive current and on which dissolution consequently prevails. If it is momentarily discon nected (between C and D,fig. lb), the resulting decrease of total cur rent shows that some initially active sites have been repassivated. Therefore, a sufficiently high current density is required to keep a pit in the active state, i By disconnecting a slightly attacked sample (between F and E, fig. lc), a reverse effect is induced : two hours later, the current reached + 10 mA. Before disconnection, the active areas of the other specimen were therefore providing sacrificial protection of this sample. 2. 2. Spontaneous fluctuations. The current fluctuations artificially induced in the previous tests help to determine a connection between the variations of the current passing through the sample and the activity of pits growing there. After this preliminary study we were able to elucidate the electrical exchanges occuring between the various areas of the same steel sample. Figure 2 shows the characteristic sections of recorded currents through each coupon. As the total current remains constantly equal to zero, every change of current on a sample induces variations in an opposite direction on the others. It is then possible to know which of them is causing variations. During the first hours of immersion, the samples are covered with many points of attack. A great number of fluctuations (fig. 2a) shows that repas sivation occurs even before the pits become visible with the naked eye. Some pits gradually stop growing. The fluctuations are not so frequent and are more clearly marked (fig. 2b, points G and H). Finally, only the bottom of the cavities previously described remain active* The variations are then few but important (fig. 2c). Gerard J?in.ard-Legry 399 ,, •4. • z -F Figure 2 +a- I Q =t= »^ :;-^ —V . : j_ -2- K_ •—^ -4 17 h TIME 10 mn TIME 10mn b c 3. A few chemical and electrochemical factors at the pit . The growth of the pits on a stainless steel plate immersed in the Fe CI3 - Na CI mixture used here, does not appear to be related to the dissolved oxygen content. Also, crevice corrosion phenomenon may start in completely deaerated Fe C^-Na CI solutions ; therefore the dissolved oxygen is not necessary for repas sivating the pits in a solution containing ferric ions. These tests were completed by making an Evans type cell with the same solution in a compartment and solutions of same pH and Cl~ concentration, but different oxygen and Fe3 concentrations in the other. It was noticed that if the oxygen and the ferric ion are both absent, there was no pitting corrosion but generalized attack on the related specimen. In another experiment, a sample was hung in a beaker containing in the bottom a green solution obtained by anodic dissolution of stainless steel (at high current density) and above ayel- low solution with Fe ions, both solutions having the same pH (1. 5) and the same chloride concentration (1. 3 ion. g. 1" ). Pits appeared only in the part of the sample immersed in the green solution. 4. Conclusion. The method described here permits the recording of exchange cur rents, whose fluctuations are related to pit changes. It particularly shows that passivation occurs at the very beginning of the attack. On the other hand, the absence of repas sivation inside the pits can be due to the particular composition of the solution of corrosion products obtained at high current density: so a change in local pH or in chloride concentration is not necessary. Reference : (1)I.L. Rosenfeld and I. S. Danilov, Corr. Sci. , 1_ (1967), 129-142. 400 Octavian Radovici THE ELECTROCHEMICAL BEHAVIOUR OE MOLYBDENUM IN SOME CORROSIVE ENVIRONMENTS Octavian Radovici and Paula Matei Institute of Physical Chemistry »Bucharest »ROMANIA and University of Bucharest, ROMANIA. Molybdenum and nickel molybdenum alloys (Hastelloys) has particularly good corrosion resistance in presence of mineral acids* Also all these materials are free from corrosion in fresh and salt waters, neutral and alkaline salts. Stainless steals|^lioyéft-w^tn}mÖ a better corrosion resistance towards diluted HC1 and halo- genide solution than other stainless steels. This unusual degree of corrosion resistance of molybde num and its alloys is due to the electrochemical beha viour of this metal in different electrolytic environ ments • The first indication about the electrochemical properties of Mo is given by Fourbaix diagram which shows the active character of this metal in all pH ranges. The cathodic and anodic behaviour especially in acidic solutions were studied by different authors (1-5)• In the present work«wt^atödied the anoaic and catholic polarisation curVes' Of specpure molybdenum in different acide, neutral and alkaline solutions under aerated and deaerated conditions. All the experiments were carried out poteritipstatically. , The polarisation data are divided into two separate groups. Group first contain all data concerning the be haviour of the electrodes in aerated solutions and the second group contain all data taken on the electrons in deaerated solutions. The results are exemplified in Fig.l which shows the anodic-cathodic behaviour of Mo in deaerated neutral and alcalin solutions. Similar shape of the anodic and ca thodic polarisation curves was obtained for acidic solu tions* Anodic polarisation curves have shown no passivation tendency of the metal in all environments* A Tafel slope RT (jngj) of 45~5o mV of anodic polarisation curve in acide and alkaline solutions is in agreement with those of other authors (5)» According to these authors the anodic dissolution processe is given by the following electrode reaction t Mo —* Mo(IV) + *» e~ —¥ Mo(VIJ+ 2 e" the rate determining step for this process being the Octavian Radovicl 401 reaction : Mo(IVj * Mo(VI) As shown in Fig.l there is a gradual decrease in the po tential with increasing cathodic current density, which continue UP to a current density of approximately 0.5 - 0.8 h-A cm~z in deaerated solution and 15-18 ^-A cm"2 in aerated solutions. At this point a limiting current density is observed for all solutions (acide,neutral and alkaline). Beyond this limiting current density, hydrogen gas was evolved with the corresponding Tafel slope of the h.e.r. This cathödic^b(9^viouri|of molybdenum electrode is the most interesting in describing the corrosion mechanism of the metal, because for spontaneous metallic corrosion to occur there must^ exist an electron sink^ Fbrlmblybde- num our experiments have shown that the cathodic process available to act as such a sink is the reduction of oxy gen even in strong acidic solutions accordingly to the following reactions * + 02 + 4 H + 4 e —» 2 H20 acide solutions 0~ • 2 BL0+ 4 e —• 4 0H~ neutral and alkaline * solutions. Because a limiting current density was observed and be cause this current does not depend on pH of solutions but only on oxygen concentration (aerated and deaerated solution) we suggest that the rate determining step for the spontaneous corrosion of Mo in all kind of solutions is the diffusion of oxygen toward the electrode which can be exprened by the very known diffusion equation A zFCD XL = "7" This mechanism can explain why molybdenum and its alloys contrary to iron and others constants metals does not corrode easy in very strong hydrocloxic acide solutions. BIBLIOGRAPHY l.G.Masing, G.Roth Werkstoff u.Korr, 3, 176 (1952). 2.N.D.Greene»S.B.Greene Slectrochem Techno1 1,276 (1963)* 3.T.Heumann, B.Hauck Z.Metallk. 56, 75 (1965). 4.N.Pentland,J.0,II Bockris J.Electrochem Soc.lo4,82 (1957) 5.L.L.Wikström, K.Kobe J.Electrochem. Soc. 116, 525 (1969) 402 Octavian Radovici &(#*) Fig.l. Poteabiogtatic polarisatioa curves of aolybdenua In de- aerated solutions i a) NaOH 1 M o; b) KC1 1U; c) Buffer solution pH = 7 < G. Richter 403 CORROSION INVESTIGATIONS OP SOME ELECTROCATALYTIC MATERIALS G. Richter, K. Mund and E. Weidlich Forschungslaboratorien der Siemens AG, Erlangen, West Germany Regarding corrosion.electrocatalysts are subjected to ex treme conditions.. On, one side the electrolytes employed are strongly corrosive and on the other the catalysts are characterized by a ;large vac$iy;e, surface area. Corrosion investigations are, .therefore,;•;"of/^fundamental importance in developing new catalystsand in;^establishing their working conditions like^temperajbureg^d^potential range. For acid fuel cells - barring the noble metals for eco nomic reasons - no catalyst is known; so far which is highr- ly active and sufficiently stable at the same time. Hence the activity and the stability of intemetallic and con ducting metal-metalloid compounds have been examined. A series of carbides, nitrides, sulphides and phosphides are stable in the potential range of anodic function. Some compounds like the carbides and nitrides of Ti, Zr and V with high formation energies are stable over a wide range of potentials*»2'. The stability of others particu larly of sulphides and phosphides is limited to a narrow range. The decomposition can take place either by reduc tion to metalloid hydrides or by oxidation. In the case of C02C a pronounced corrosion minimum is no ticed around 0.2 V(vaRHE). The carbon is thermodynamical- ly stable around this potential. Above this value the carbon is oxidized and reduction takes place at lower values. Till now a notable activity is observed only with com pounds like C02C, Ni3C, Ni^W and Ni^W which are all less corrosion resistant. . The corrosion current of WC^J determined by the potentio- dynamic sweep method in argon saturated 1M H2SO4 is 10~4 a/cm2 at room temperature and a potential of 0.98 V (vsJIHE). At 150 °C in concentrated phosphoric acid the current of 10~4 a/cm2 is already attained at a potential of 0.53 V. The oxidation leads to a sparingly soluble surface layer. The transport limitation in the growing surface layer of the electrocatalyst and the material release from the electrode can both cause a deactivation. If a dissolved species is in equilibrium with the catalyst surface a reversible corrosion mechanism follows and the corrosion is limited by the diffusion into the bulk of electrolyte irrespective of the large active surface area. The diffusion through the gas stop layer is the limiting step for gas diffusion electrodes. This basic mechanism 404 Gr. Richter is more thoroughly investigated for the corrosion of oxy gen electrodes with silver catalysts in alkaline electro lytes. The temperature and potential dependence of the equilibrium-concentration has been determined and the silver loss analyzed. At 60 °G and a potential of 1.080 V, near the equilibrium rest potential, the loss of a raney silver electrode ^ith an asbestos layer in 6M KOH amounts to 5.5» 10"' g cm~2'h~'. On the contrary the loss of silver under normal condi tions of operation is not measurable. The extrapolated value for a potential of 0.9 V is less than 10~8 g-cm"*2» h-1. The silver loss of the electrode is smaller than the cal culated value because the dissolved complex decomposes either spontaneously or by.means of light, foreign mate rials and reducing agents^-). The material loss alone is not responsible for the de activation of an electrode. It must also be considered that the recrystallization within the liquid phase de creases the active surface. The decomposition of the silver complex surprisingly leads to the re-formation of an active black catalyst powder and thus counteracts the recrystallisation. As catalysts for the oxygen electrodes in acid electro lytes organic complexes like Fe-phthalocyanine which were regarded stable have been investigated^/. It turned out however, that these compounds do not withstand the strong oxidative conditions atKthe oxygen electrode and gradual ly lose their activity-by hydrolysis and oxidation^»?). With monomeric phthalöcyanirié on a carbon substrate the formation of needle like crystals is observed which have been identified as phthalimrde. The most active complexes are particularly vulnerable. Carbon catalysts are more stable. As in the case of or ganic complexes the corrosion process is irreversible and increases with the active surface. Under potentiostatic . conditions the corrosion current decreases continuously^). Within 25 hours the current at 60 °C and at 950 mV drops to one third of its initial value. After this period it lies in the order of 10"'° a/cm^—of—BET—su-rface area for different kinds of carbon black and active carbon. On plain electrodes in oxygen free phosphoric acid at 950 mV a current of 10~5 a/cm2 is measured under potentiodynamic conditions. Even at a potential of 700 mV the anodic cur rent is too high to be tolerable. Therefore the applica tion of carbon catalysts for oxygen electrodes seems to be limited to temperatures below 100 °C. Since the activity of the carbon is determined by certain basic surface complexes which are converted into acid G-. Richter 405 groups under oxidizing conditions even a relatively small anodic current can lead to a considerable deactivation. The current of,oxygen reduction of a carbon electrode at a constant potential decreases in a manner analogous to the corrosion current. Initially a steep fall is observed which flattens after 100 to 200 hours of operation. At present the current obtained for an oxygen electrode with an active carbon catalyst load of 100 mg/cm2 at 60 OC and a potential of 0.7 V is 100 ma/cm2 which de creases to 40 ma/cm2 after 1000 hours of operation. REFERENCES 1) K. Mund, G. Richter and P. v. Sturm Collection Czechoslov. Chem. Commun., !>£, 439 (1971) 2) R.D.^Cowling and H.E. Hintermann J. Electroch'em. Soc, 117, 1447 (1970) 3) H. Bohm and P.A. Pohl Wiss. Ber. AEG-Telefunken, 41, 46 (1968) 4) T.P. Dirkse, D. Van der Hart and J. Vriesenga J. Inorg. Nucl. Chem., 27, 1779 (1965) 5) H. Jahnke Ber. Bunsengesellsch. Phys. Chem. 72, 1053 (1968) 6) A.H. Cook J. Chem. Soc, (1958) 1761-1780 7) A.A. Berlin and A.J. Scherle Inorg. Macromol. Rev., 1_, 235-270 (1971) 8) H. Binder, A. Köhling, K. Richter and G. Sandstede Electrochim. Acta, £, 255 (1964) 406 E. Schmidt CORROSION STUDIES WITH A TWIN ELECTRODE THIN LAYER TECHNIQUE K. Huber and E. Schmidt Institut fur Anorganische, Änalytische und Physikalische Chemie der Universität, Bern, Switzerland. A twin electrode technique is described by which solu ble electrolysis products of anodic oxydation processes at solid electrodes may be investigated. The anodized electrode .(Me) is separated by' a thin electrolyte layer from second ("scavenger") electrode (S) which is held at constant potential by an independent polarizing circuit. -3 The distance between both electrodes is about 5*10 cm. Any soluble electroactive species generated at Me (Me , Op, Hp etc.) will diffuse through the electrolyte layer and may be oxydized or reduced selectively at S, if a suitable potential difference between both electrodes is maintained. Under steady state conditions, the concentra tion profile of the diffusing species may be calculated from the current measured at S, if both the diffusion coefficient and the thin layer dimensions are known. The thickness of the oxide and/or chemisorbed oxygen layer deposited at Me .is obtained from the difference of the currents passing through both electrodes. This method has been used for examination of the cor rosion behaviour of Cd and Zn in 0.01 -'1M alcali hydrox ide solutions. At Cd electrodes, only about 1 % of the total anodic current in the passivation range close to the reversible Cd/CdO potential (E^-0.9 V/Hg,HgO) is consumed by the formation of soluble Cd species, the maximal Cdz+ con centration at the electrode corresponding to slight su- persaturation with respect to Cd(0H)p. A fractional oxy gen monolayer is observed in nonsaturated solutions above E. Schmidt 407 the nucleation potential of the passivating oxide layer. Saturation is maintained during the cathodic stripping of the oxide coverage which occurs in two well separated steps at 0.9V and - 1.4-V/Hg, HgO, respectively. In 0.01 and 0.1M hydroxide solutions containing an excess of sup porting electrolyte (NaCIO,), local breakdown of the pas sive layer is observed at E~0.5 V/Hg,HgO, large crystals of (3-Cd(0H)2 growing at insulated pitting sites. An un known soluble electrolysis product is generated simulta neously, which is reduced at scavenger electrode poten tials below -1.5 V/Hg,HgO. In the Zn system, large steady state currents are es tablished between the anodized Zn electrode and the sca venger electrode (amalgamated Ag), held at potentials below the Zn/ZnO equilibrium value (E<1.32 V/Hg,HgO) . Depending upon the Zn potential, two distinct limiting current plateaus with different Znz concentrations are found. In 0.01 M NaOH, the Zn concentration near the Zn electrode is 1.8 + 0.3-10"5 mMol cm-5 at 1.3 < B ANALYSIS OP THE INFLUENCE OF HYDROGEN ON PITTING CORROSION AND STRESS CORROSION OF AUSTENITIC STAINLESS STEEL IN CHLORIDE ENVIRONMENT A.A. Seys*, M.J. Brabers** and A.A. Van Haute * •Institute of Industrial Chemistry, Corrosion Division, University of Leuven, Heverlee, Belgium ••Department of Metallurgy, University of Leuven, Heverlee, Belgium. In a study on the mechanism of pitting corrosion a gas evolution in the pit has "been observed, "both in open circuit experiments and,, in anodic polarisation experiments as well in acidic.as "basic solutions. This gas has "been gaschromatographically identified as hydrogen. By means of vacuum extraction experiments a diffusion of hydrogen in the metal has "been shown, a diffusion which seems to progress logarithmi cally as a function of time.. The development of hydrogen is explained "by the local acidification of the electrolyte solution and "by the polarisation resistance in the pit, facts which have "been experimen tally proved. By use of pH indicators, it was shown that the pH in a potentiostatically growing pit of stainless steel in 0,1 N KHCO3 + 0,1 N KC1 (pH 8,4) is 2..3 whereas for nickel it was 6,5..7. Theore tically it is possible to calculate the pH a.t the "bottom of a pit for pure metals assuming that the solution inside a pit is saturated with the hydroxide of a given metal. These theoretical pH values agree with the experimental ones published by BrownO). The potential inside a pit has been measured in an artificial pit and also in a potentiostatic pit by use of a microprobe reference elec trode. A relation between the potential transients at the bottom of the pit, the current transients a,nd the moment of hydrogen emana.tion from the pit, has been shown. The emanation of a bubble from the pit coincides with a decrease of the potentialdrop in the pit and an in crease of the measured current. Measurements of the" potential inside a pit during the determination of the protection potential (2) revealed that at the protection potential the potential inside the pit equals the potential at the surface. Experiments have been carried out to determine the influence of the dissolved hydrogen. From the results it may be. concluded: - by replica techniques that the hydrogen embrittles the metal around the pit. This undoubtedly decreases the mechanical strength of the remaining metal. - by X-ray diffraction that hydrogen can cause phase transformation in the austenitic stainless steel. In the case of pitting corrosion such transformations have not yet been .confirmed. - by means of potentiokinetic polarisation curves that the anodic reactivity of stainless, steel charged with hydrogen is not increased;. This leads one to conclude that the development of hydrogen in the pit is not .of primary importance for the mechanism. One also finds a development of hydrogen at stress corrosion of auste nitic stainless steel in Cl~ environment. It has here been proved by Brabers (3) that hydrogen is indeed essential. The clear evidence therefore is that: - the top of a growing crack does not contain chloride ions A. Seys 409 - the fractography of the top of the crack points to a cleavage mode - the hydrogen in front of the crack causes vera*- high internal stresses under-the effect of the external lead. rrom the study of the mechanisms, the influence of external factors and of the electrochemical "behaviour of pitting corrosion and stress corrosion one comes to the conclusion that "both phenomena are analo- guous. Pitting corrosion is represented as a, more general form of the stress corrosion. There is one "big difference between pitting corro sion and stress corrosion: for the latter the hydrogen in the metal must diffuse to local areas in sufficient quantities to cause high triaxial stresses resulting in mechanical failure "while this is not so for pitting corrosion. This implies that some factors may affect the process in various ways. A concept common for stress corrosion and pitting corrosion is intro duced. This model may also he applied to the case of crevice corro sion. This general model has advantage from 'the theoretical (thermo dynamics, kinetics) and practical (common means of protection) point of view. ^ ; Brown B.F., Electrochemical Meeting Detroit, 5-12 oct.- 1969. (2) M. Pourbaix et al., Corrosion Science, Vol. 3, p. 239» 1963. (3) M.J. Brahers, Mechanism of stress corrosion cracking of austenitic stainless steel in chloride solutions, Nato Science Committee ERICEIRA Portugal, 1971. 410 Zuzanna Szklarska-Smialowska SYNERGISTIC ACTION OF N-DODECYLAMINE WITH N-CAPRYNIC ACID ON THE CORROSION OF STEEL IN SULPHURIC ACID SOLUTION Grzegorz Wieczorek and Zuzanna Szklarska-émialowska Institute of Physical Chemistry of the Polish Academy of Sciences, Warsaw, Poland There exists a limited number of works trying to explain, the nature of the synergism. The synergistic effect is expressed by various-authors in a different manner depend ent upon the criterion taken into account for the "addi tive action" of the inhibitor mixture.. In this work the synergistic effect /S/ of a mixture of two inhibitors has been determined by substraction of the effectiveness calculated from the equation given below, from that experimentally obtained i S = 0o__ - Qa % = Ö1 + «2 " 9192 ' ^ where Q* and 92 are surface fractions covered by inhibit ors 1 and 2. The protective properties of organic inhibitors,if they do not directly take part in the electrode reaction,are closely connected with their adsorption. Thus, synergism of the protective action occurs in such systems in which increase of the adsorption of one or another component of the mixture takes place on the surface of the corroding metal. This work was devoted to the study of the synergistic action of mixtures of primary aliphatic amines and acids. These compounds have a simple chemical structure and the interaction of their molecules between them on the metal surface is well known. It was found in previous work /1/ that for all primary aliphatic compounds /with 6-12 car bon atoms/ there exists a correlation between the effect iveness of the given inhibitor and the amount of its ad sorption on the metal surface. It was also found that the adsorption of primary aliphatic acids,alcohols and amines occurs accordingly to Frumkin's adsorption isotherm. The principal aim of this work was to investigate, the protective properties of n-caprynic acid with n-dodecyl- amine as corrosion inhibitors present in different con centrations and ratios in 1N H2SO4 solution,determination of the synergistic effect and finding out the dependence between this effect and the concentration of both compo nents in the mixture. The corrosion measurements were performed using mild steel of the same composition as in the former work. The procedure of preparing samples and solutions was also identical as previously described /1/. The value of coverage /Q/ of the steel surface by adsorb ed inhibitor molecules was determined applying the known dependence 9 = 1-.^ = ^ % where Uiand U are. corrosion Zuzanna Szklarska-iSmialowska 411 rates in solutions with and without inhibitor respective- 5' 3 1S inhi"bi'bor effectiveness. Fig.1 gives an example of dependence between q value versus concentration of dö de calamine.Five series of such measurements were perform ed. In each series the- concentration of caprynic acid was \fo e kept constant,and the concent 80 0-8 ration of dodecalamine was changed. In Fig.1 the adsorpt 60 06 ion isotherm for pure dodecal amine /the curve without 40 0-4 points/ is also plotted. The results can also be represent 20 02 ed in another form,namely as 0 versus cs /summaric concentra 10" 10" ifr» tion, 10"« cCmole/ll or c cs= 'Ac ÖJar""ea ^constantVJUöuauL' .• cAm sn^- Pig. 1 cAc are the concentrations e of amine and acid respecti 0-8 vely. Fig.2 gives an example of such a function for fixed con 06- centration of acid or amine equal 1x1CT4 mole/aj 0-4 0-2f The dependence between the mount of the synergistic/or Clan" 107T £* tagonistic/ effect and the con centration ratio of dodecylami- Fig. 2 ne to caprynic acid is given in Figo. As can be seen, the hig her the caprynic acid concent- ration, Ihe smaller the concentration ratio -of amine to acid at which occurs a maximum of the synergistic /or antagonistic/ effect. It results from this that the max imum of the synergistic effect is a function of the cap rynic acid concentration. The results suggest the existence of a certain regulari ty in the system: dodecylamine-caprynic acid-dilute H2S0^ - steel. Especially interesting feature is the discovery in the studied system of "concentration symmetry". This "symmetry" consists in the equality of the surface coverage when either is the amine concentra tion c.__ m the solution cÄ._ = X and the acid concentra tion c£j = Y, or is c V*,m an d c. = X. In both above cases the sum of Amchemical 'A^potentialc s for the mixture of components can be taken as constant. In Fig.4 the relationship is plotted between the degree of surface coverage 0g and log (c. c ) . Thermodynamic data obtain ed from the experimental adsorption isotherm /Frumkin 412 Zuzanna Szklarska-émialowska •U-S-W'4 CrnoU/l] Ad 2.10*40** H Q» 3.7-5«40*s • 5 450»W • 0-8 0-6 • QA 0-2 •w i& Io=*— .«/_$ Fig. 3 Fig, C4m»Cte tmole/lj adsorption isot herm: s_ K C e-xp U <0 = Am 'Ac / 1=5" mix are as follows: AG° /standard free adsorption energies for mixture/ = 13.§Kcal/mole ,ff"/for the mixture/ = -1.36 AG° /mix/ is approximately equal to the sum of standard free adsorption energies of n-dodecalamine / AG = 7«4 Kcal/mole/and n-caprynic acid / AG~ =6.5 Kcalfmole/. On the other hand, tff" coefficient for the mixture is arith metical mean of values for dodecylamine /f = -1.7/ and caprynic acid / f = -1.1/. On the basis of above results one can conclude that: 1/ in the corrosive medium under consideration, complexes are formed, composed of one molecule of n-caprynic acid and one molecule of n-dode- cylamine,2/. in the studied range of inhibitor concentra tion, those complexes are exclusively adsorbed on the me tal surface, since they have a higher adsorption energy than single components, 3/ the adsorption energies of complexes change, linearly with the surface coverage, 4/ the complexes lie flat on the metal surface and their adsorption energy for 0 «+• 0 is equal to the work neces sary for transferring their CH2 groups from the bulk of the solution to the surface, plus some small work ne cessary for transferring there the functional groups, 5/ the synergistic effect is caused by complexing pheno mena occurring in the corrosive solution and not on the metal surface. References 1/ Z.Szklarska-émialowska and G.Wieczorek, Corr. Sci., H, 843 /1271/ Ulf Ulfvarson 413 THE EFFECT OF USING ROAD SALT ON THE CORROSION CLIMATE FOR VEHICLES Ulf Ulfvarson and Kurt Johansson The National Institute of Materials Testing,Stockholm,Sweden. On nine places in Sweden "box shaped test specimens of car body- steel plate were mounted under cars and on fixed stations in the vicinity of the area operated "by the same cars.These pairs of mo bile respectively fixed stations were distributed all over Sweden from far- up in the north and down to the southern parts of the country .The specimens were exchanged for new specimens each month and each three month on both the mobile and fixed stations »The total exposure time was about one year and therefore all kinds of climatic conditions were covered.After exposure the test specimens were washed with de-ionized water and the wash water analyzed for sodium, calcium, chlorine and sulphur .The corrosion of the test specimens was estimated gravimetrically after derusting.The cli matic and chemical conditions of the atmosphere was registered in all test areas during the time of investigation and the amount of salting (sodium chloride for deicing during the winter season and calcium chloride during the spring, summer and autumn for dust binding) recorded. Corrosion of the test specimens during different parts of the year was compared with information about conditions mentioned.Since the investigation was recently ended it is not yet possible to give details of the result.It should be mentioned,however*that the same technique was applied in Gothenburg,the 2nd city of Sweden,and it was found that the salting during the winter season did not affect the corrosion of the test specimen.This surprising result was pro bably due to the high content of corrosion stimulants already pre- sent in the atmosphere,e.g. sulphur dioxide.As far as can be seen up to now on the other hand,the use of salts on the roads outside the big cities or other severely polluted areas will affect the corrosive conditions in which cars are driven. 414 Ernest B..Meager THE DISSOLUTION KINETICS OF LITHIATED MO IN AQUEOUS ACID SOLUTIONS Chin-Ho Lee, Alan Riga and Ernest B. Yeager Department of Chemistry, Case Western Reserve University, Cleveland, Ohio, USA, Luhrizol Corporation, Cleveland, Ohio, and Department of Chemistry, Case Western Resemre University, Cleveland, Ohio, USA. The present study has involved the examination of l) the dependence of dissolution on the electrode potential and the pH of the solution, and 2) the relationship between dissolution and the simultaneous residual current. Theoretical treatments of the dissolution kinetics of ionic compounds under the influence of electrode potentials have been attempted by Engell (l) and Vermilyea (2). The rate of dissolution with the back reaction negligible is given by Vermilyea as follows: r - £ (!) a z Fn a z Fn exp - ~w~+ + +exp - -=s~. where n = ferred are uncertain hut if z+ is +2 and z__ is -1, then for the H2SO4, a = 0.31 and a = 0.6?. The maximum could not he observed in HC1 because CI2 evolution but the Tafel slope of ^0.16 V/decade correr1 sponds to oc = 0.23. The change in the Tafel slope in the region 1.2 to 1.3V is probably a consequence of a large change in the Ni3+/Ni2+ ratio at the surface, + probably according to the reaction 2Ni0 + H2O -*- M203+2H + 2e. Only in the potential region below 0.8V is the dissolution rate dependent on the Li concentration in the oxide. This region has been found to correspond to exhaustive depletion in earlier studies (3) of the space charge properties of such electrodes. At more anodic potentials, the change in potential,drop; across the space charge region is small. .(2) Dependence of dissolution on the pH of the solution. The effect of pH on the dissolution was~ studied at 95°C in HC1 + KC1 with the total ionic strength held at unity. The Log r decreased line arly with pH, and (81og r/8pH)_ • -O.56, -0.53, and -0.32 at E = 0.9, 1.1, and 1.3V respectively. Theoretical prediction of (91og r/3pH)_j, can be made by introducing the relationship between the double layer potential with cation transfer rate controlling and negligible back reaction r = r+ - n+ k+ exp ^ (3) Combination of (2) and (3) yields Olog r/3pH)E * - a+ z+ (U) The predicted slope according to Eq. (U) using the values a+ = 0.23, z • 2 evaluated from Fig. 1 is -0.U6 as compared with an experimental value of * -0.5. Thus the agreement between experimental results and theory is satisfactory. (3) Other experimental evidence: a. The complexing agent (EDTA) was found to increase the dissolution rate linearly with concentration at 0.765 to I.165V. EDTA complexes with Ki2+. Thus the enhancement of dissolution by EDTA provides further evidence that the cation transfer is rate controlling at potentials cathodic to the maximum in Fig. 1. b. Although the residual currents during dissolution were not_very reproducible, statistically a small cathodic current (10 8 to 10 7 A/cm2) and a larger anodic current (10~6 to lO"1* A/cm2) were evident at potentials cathodic and anodic to M3.8V, respectively. The catho dic current was about 1% and the anodic current about 10$ of the dis solution rate expressed as charge of either sign per sec. The reaction corresponding to the cathodic current is postulated to be the reduction of lattice Ni3+ to lattice Ni2+, while at more anodic potentials the evolution of O2 (in the case of H2SO4) or CI2 (in the case of HCl) predominates even before the ordinary reversible potentials for forma tion of these gases from solution phase species have not been reached. 416 Ernest B. Yeager This anamolies results from the involvement of lattice oxygen in these reactions and the coupling of the NiO dissolution process. C. The dissolution was independent of the presence of the reducing agent T1C1 which was oxidized during dissolution at potentials of 1.265 to 1.U65-'. The proposal mechanism is as follows: (NiO) (H 0-) |lll^(Ni(0H) ) lattice 2 soIii 2'lattico e + + . + 2H 0 . (Ni(OH)2J lattice + 2H sol n !i2Wsoln 2o soln with either one of the following steps as r.d.s. (Ni2+) *(Ni2+) at E <1.68v lattice soln (OH"), .. . + H _ ->H 0 . at E~>1.68V lattice soln 2o soln Acknowledgement; The authors acknowledge the support of the research "by the U.S. Office of Naval Research. References: 1. H.J. Engell, Z. Physik. Chem. N.F. J, 158 (1956). 2. D.A. Vermilyea, J. Electrochem. Soc. 113, IO67 (1966). 3. D. Yohe, E.B. Yeager, R. Greef, and A. Riga, Electrochim. Acta 13, 1351 (1968). U. E.B. Sandell, "Colorimetric Determination of Trace Metal", 3rd Ed., Interscience, New York (1959)• Fig. 1: Dissolution of NiO(Li) in helium saturated 1 N_ HC1 and 1 N_ H2S01» solution at 95°C. Symbols corresponding to different group electrodes: o • Group 1, (C^)^ = 0.88 a/o; <£ Group 2, (C^)^ • 0.33 a/o; D «4Group 3, (CLi)Mg = 0.26 a/o; • 4 Group U, (CLi)MS - 0.12 a/o. Dashed line: Calculated from Eq. (l) with the following para meters: <*+z = 0.62, a_z_ = -0.62, r = 15,000 ug/cm2-hr, 1.0 2.0 <|>f = 1.68 V vs SHE. E (V) vs SHE SECTION 6 BATTERIES COMPARATIVE STUDY OF FUEL CELL ELECTRODES AND' TECHNOLOGY (ACCOMPLISHMENTS 1968-1972, OUTLOOK) Karl V. Kordesch Union Carbide Corporation, Consumer Products Division, P. O. Box 6116, Cleveland, Ohio 44101, USA. INTRODUCTION Work on fuel cells culminated during the years 1966-1968, but most of the innovations made during that time were not published till later. In particular, patents applied for during this peak development period only recently became public knowledge. It is, therefore, interesting to review these accomplishments. The largest area of "new" technology is concerned with problems of making cheaper electrodes, electrocatalysts, and with improving the life of cells and the reliability of systems. Since substantial prog ress has been made in all of these directions, one wonders now what led to the abrupt drop in fuel cell activities around 1969-1970, especially in the U. S. A. To put the conclusions ahead of the discussion, it is not enough to create a new technology; there must also be a commercial market for it to be successful. The military and space efforts produced special fuel cell power sources which did not have time to become practical enough for general use. The severe cutback in large scale govern ment support doomed commercial projects relying on "spin off" technology. With a few exceptions, long range programs were ter minated in the U. S. A. In Europe and Japan a little more optimism exists, but the efforts of companies to justify their work in the fuel cell field from a marketing viewpoint is evident. DISCUSSION 1. Cost Reduction Replacement of noble metals: Less expensive anode catalysts for acidic systems. Reduction of noble metal content: in alkaline systems to the extent that cost and availability become less significant. Use of low cost electrode structures: The often complicated proc esses for manufacturing multi-layered (precision) structures have heen superseded by cheaper ways: screen supports, sedimented layers, rolled or sprayed electrodes, and principles of paper-making techniques are used. Systems have been simplified in design and operation: The shift from alkaline cells with liquid electrolyte to acidic cells with matrix electrolyte is noticeable. Thus, C02 scrubbing units are eliminated, and complicated accessories for proper gas and humidity control are not require'd. The types of fuel cells have been reduced to a few. "Exotic fuels" and systems difficult to operate have been abandoned. 2. Improvements in Life Expectancy and Reliability The factors just mentioned for cost reduction fortunately had a 420 Karl V. Kordesch beneficial effect on longer life and higher reliability. The reason is the better reproducibility of larger-scale (cheaper) production methods and simplification. Non-noble metal catalysts have proven to be more stable in operation. 3. Fuels and Oxidants The fuels have been selected for practical applications. Liquid H2 and Oz are combinations suitable only for space applications. Hydro carbons, (direct and-indirect), hydrazine, alcohols, and hydrogen are the remaining candidates, with some exiting developments in H2- storage. Oxidants to be used: air or hydrogen peroxide. 4. The Search for Applications of Fuel Cell Technology Power supplies for military, underwater, and space efforts will certainly be built, and fuel cells rank high in exceptional service. The commercial market can only look to long range programs for power distribution and perhaps for electric vehicle propulsion as a distant possibility. For that reason, it is understandable that efforts are being made to use fuel cell technology in related fields, such as metal-air cells, regenerative cells, high energy density primary and secondary cells, nonaqueous batteries, organic depolarized cells, and electrochemical devices for analytical and medical applications. Hybrid systems which use some fuel cell electrodes have interesting possibilities. 5. Continuing Research and Development Fuel cell work has stimulated electrochemical research and its practical application. The theoretical concepts worked out for "con tinuous feed batteries" are applicable to other galvanic systems; also the test instruments and methods which evolved frorn fuel cell research are directly usable in areas of material evaluation and cor rosion testing. Practically all major companies in the field of batteries are continuing such work "related" to fuel cells. 6. Companies and Agencies Engaged in Major Fuel Cell Projects An up-to-date list of organizations which are involved in larger size fuel cell programs will be presented, together with the character istics of the system and its intended application. 7. Outlook The U. S. Government plans to spend more on research and developr ment than ever before in fiscal year 1973 ($17. 8 billion). Civilian programs that will get a particularly strong push include pollution free electric power and transportation. H. Behret 421 DISCUSSION OF VARIOUS TYPES OF TEST ELECTRODES FOR THE EVALUATION OF CATALYSTS FOR THE ELECTRO RE DUCT I ON OF OXYGEN H. Behret, H. Binder, A- Köhling and G. Sandstede Battelle-Institut e.V., 6 Frankfurt (Main), Germany For various solid substances it is a problem to evaluate their catalytic activity in a heterogeneous gas-liquid reaction. Differences in the properties of the substances, such as grain size, density, wettability, or porosity, complicate 'the 'comparison.' Comparison is also difficult .if catalysts with low conductivity have to be used for electrochemical reactions in gas electrodes. The necessity of comparing the electrochemical activities of differing substances' gave rise to the development of several types of test electrodes. Measurements with suspension electrodes were carried out by Shlygin, Gérischer, and others together with their coworkers. Bonnemay, v. Sturm, Bohm, Despic and their coworkers investigated different kinds of unbonded elec trodes. Fluidised bed electrodes - closely related-to suspension electrodes - are the object of research of Fleischmann, Janssen, Heitz, etc. As to bonded electrodes, it is impossible to mention the names of all the people engaged in this large field of research. Our own investigations were concerned with various types of test electrodes for the evaluation of catalysts for oxygen reduction. Catalyst research for the electro- reduction of oxygen is mainly necessary for air cathodes in fuel cells and metal/air batteries. The suspension electrode was useful for detecting and discriminating the catalytic activity of a series of metal chelates, e. g. phthalocyanines, tetraarylporphyrins and dibenzotetraazaannulenes, Being non-conducting materials these chelates had to be investigated by measurements on a conductive support, such as active carbon or carbon black. Our measuring apparatus for suspensions consists of a glass cell in which the material suspended in aqueous solution is whirled about by a stirrer. The investigations showed that only materials of nearly the same specific weight and almost the same grain size should be compared. Very small grains are not tossed against the gold electrode because of their low impulse 422 H. Behret and therefore are not active for electrochemical re duction. Porosity of the suspended material does not significantly influence the results because the reaction seems to take place preferably on the outer surface of the catalyst grain. The binder of the bonded electrodes which we compared consisted of polyethylene (PE) or polytetrafluoroethylene (PTFE) or gold. The porous PE-electrodes had been sintered together witli a catalyst of good conductivity, e. g. transition metal chelates with carbon support or transition metal sulphides without support. A comparison between bonded PE-electrodes and suspension electrodes reveals a relatively high activity -of the active carbon itself in the bonded electrode, which can be explained by the porosity of this material. All other results are in agreement with the results obtained on the suspension electrode, provided that the restrictive operating con ditions of suspensions are taken into consideration. Care must be taken in comparing the absolute current density of bonded electrodes with currents measured in suspensions. This is demonstrated in Figs. 1 and 2 for the active chelates Co(II)dibenzotetraazaannulene(CoTAA) on carbon black "Lurgi" (Fig. 1) and Co (II) te.tra- (p-methoxyphenyl) porphyrin (CoTMPP) on highly porous active carbon "Norit BRX" (Fig. 2). The values of the ordinate are normalised to the curve obtained for the pure support as electrode. mA current: current mA susp. el. density: 4 bonded el. cm bonded e 20 with CoTAA susp. el.— with CoTAA 10 10 without susp. e/.~| CoTAA [ bond.el.J 500 mV 700 potential Fig. 1: Carbon black electrodes 'He Behret 423 mA current: susp. el. current mA susp. el. with Co TMPP' density: bonded el. cm' JO bonded el. — 50 CoTMPP without CoTMPP 500 mV 700 potential Fig. 2: Active carbon electrodes Bonded PTFE-electrodes show the same characteristics as bonded PE-electrodes• Bonded gold electrodes are especi ally suitable for non-conducting materials. Porous gold electrodes for test purposes are easy to prepare by using pressure; results are. obtained quickly, but grain size and rheological behaviour under pressure often lower the mechanical strength of the electrode. The substance to be tested in an unbonded electrode (powder electrode) was poured on a piece of graphite felt. Substances of low conductivity were additionally ground with graphite powder. In both cases, the piece of gra phite felt with the powder was placed between porous metal plates acting as holders and electric current leads. In alkaline solution unbonded powder electrodes are suitable almost without reservation for the evaluation of catalysts for oxygen reduction. The same catalytic reduction in acid solution, however, is restricted by a limiting current. An explanation is given and methods of overcoming this limiting current are discussed. 424 Harald Bohm CRUDE GAS/AIR FUEL CELL WITH BIPOLAR, NON-NOBLE METAL ELECTRODES AND ACID ELECTROLYTE Lothar Baudendistel, Harald Bohm, Gerhard Louis and Franz A. Pohl AEG-TELEFUNKEN, Forschungsinstitut, Frankfurt /Main, Western Germany. Fuel cells with acid electrolytes are favorable supply ing air and a cheap fuel such as reformed gas or cracked methanol because of the insensitivity of the acid system against C02 . In the view of the economy and the poisoning by CO pla tinum should be avoided as the catalysts. For these reasons WC and activated charcoal may be used as cheap and stable electrode materials. The-hard material WC prepared for normal technical use exhibits very poor catalytic activity. But the carburization of W with CO at 700-800°C leads to suitable electrode materials wijth high activity for the oxidation of hydrogen. lOOmA/cm2 at a polarization of only ,100 mV can be reached at the WC gas-diffusion electrodes. The highest activity of WC is not attained at the stoichiometric composition but at a carbon content lower than 6.12%. This is caused by the presence of oxygen, maybe in form of surface oxides. Another important feature of the WC is its stability against the catalytic poison CO. Carbon monoxide is even oxidized on the WC-electrodes at potentials more than 300mV against the reference electrode, a H2-electrode in the same solution. This behavior of the anode cata lyst WC is a real advantage compared with platinum. As a precious metal free catalyst for the cathode only activated charcoal is known to be stable in acids over a long period.The too low activity of the purchasable charcoal must be increased by an activation process - the treatment of the carbon with ammonia at high tempera tures. Electrodes prepared from so activated charcoal show current densities up to 90 mA/cm2 at 600 mV. The performance of the WC/C - cell with small test electrodes 2 reaches about 50 mW/cm at 60°C in 2n H2S0^ . For technical application, however, the battery has to be constructed with thin, large-area electrodes; this leads to a compact assembly with a low weight / power ratio. Especially favorable is the use of bipolar elec trodes. This electrode consists of a PTFE-bonded carbon layer, a conductive, but gasimpermeable graphite foil separating the hydrogen and the. oxygen chamber and a PTFE-bonded WC-layer. Several bipolar electrodes are arranged side by side to a battery. The space between two bipolar electrodes formed by an plastic frame serves as the electrolyte chamber. The electrolyte is supplied Harald Bohm 425 through long and thin channels in these frames. The assembly of the electrodes in such a manner permits an electrical connection in series. 4-0 cells form a module with a weight of about 4- kg and a volume of 3 dnr. The gases reacting in the cell are hydrogen and oxygen. Because of the stable catalysts and the robustness of the acid system air and impure hydrogen may be used. In our case the crude hydrogen is supplied by catalytic cracking of methanol. CH5OH is converted on a purchas able catalyst for the methanol synthesis into a mixture of hydrogen and carbon monoxide. This mixture is di rectly supplied to the cell without any purification. Because of the endothermic character of the methanol cracking process the reactor has to be heated. This is performed^ by the combustion of the waste gas of the fuel cell mainly consisting of CO. By the complete burning of CO to C02 , carbon dioxide and water formed by the electrochemical reaction in the cell are the only pro ducts of the whole aggregate. In^ spite of the use of the methanol cracking gas and the air without any purification the cells show a good life time. Cells with large-area electrode run more than 5000 hours without any diminution of the catalytic acti vity of the electrodes. A battery with small electrodes had reached 35 000 hours in continuous operation with a minimum of maintenance. 426 Jean Brenet CHEMICAL PREPARATION OF p-MnU^ AND COMPARISON OF ITS ELECTROCHEMICAL PROPERTIES WITH THOSE OF 3-AND Y-MnO Marc BELEY - Jean BRENET Laboratoire d'Electrochimie et de Chimie Physique du Corps Solide Université Louis Pasteur - Strasbourg, France We have studied the chemical preparation of manganese dioxide pos sessing high electrochemical reactivity. Several of the methods em ployed have been described by one of us (French Patent n°1306706 - 1961 and 1525333 - 1967) ; others are unpublished (sealed letters deposited at the Academie des Sciences de Paris n° 12847.- 1951 and n°13408 - 1955). We were particularly interested in obtaining a sub-oxide by oxida tion of Mn(0H)2 in aqueous medium at moderate temperatures. This pro cedure leads to Mn30^ with hausmannite structure which can be easily converted to manganite, MnOOH at higher temperature. These conditions are in accord with previous results obtained by P. BRUNNER. Acidic disproportionation can be subsequently performed using either M^O^ or MnOOH. p-Mn02 is obtained in both eases at ordinary tempera ture in the first instance and at 60-80°C in the latter. However, if the operation is to be carried out at. ordinary temperature, it is neces sary to dry and maintain the Mri304 at temperatures close to 25°C. A consideration of the reaction kinetics at 25°C shows that the optimal conditions consist in acidification of a solution whose con centration is of the order of 2,3 moles/litre, but the reaction rate can be increased by increasing the excers of acid with respect to the "theoretical" amount of Mh^O^ corresponding to an overall, non - stoi + 2+ chiometric reaction represented by Mn30^ + 4H + 2S0^ -»• 2Mn + 2S0£ + Mn02 + 2H20. We then compared the cathodic reduction mechanisms of 3-Mn02 with those of Y-and pMn02 using the classical potentiodynamic procedure. The cathode element is prepared from the dioxide to be examined ; the elec trode, consisting of a mixture of Mn02 and very pure graphite, is pres sed as a pellet of 1-2 mm thickness at pressures of 200 kg/cm2. The proportions of graphite and Mn02 are chosen such that each sample will have a maximum coulomb/gram Mn02 ratio in the electrode. Trials have been carried out using pure J3-Mn02 obtained by dissociation of MnNO-j , Y-Mn02 produced industrially by anodic oxidation, and p-Mn02 prepared in our laboratory by the method already described. The curves corresponding to the expression I = f(U) are recorded by the potentiodynamic method using a 10 mV/min scanning speed for the potential. However, a plot for each curve could only be obtained once the equilibrium potential was reached, that is, about one hour after immersion of the electrode in the electrolyte consisting of a satura ted solution of ammonium chloride. Jean Brenet 427 The curve I = f (U) for p-Mn02 in the cathodic region is characteri zed by a peak at point A corresponding to a potential of 0.290V(ENH) and a overall current of about 13 mA. The area of the curve between the initial potential at zero current and 0.290V is proportional to the number of coulombs required during the cathodic reaction or reactions Up to the point A corresponding to the peak of the curve. This number is of the order of 360 cb/g. The reduction curve, represented on a se mi-logarithmic coordinate scale, cannot be expressed by a TAFEL rela tionship and has a point of inflection corresponding to a potential of 560 mV(ENH). It would thus appear that two successive reduction reac tions occur. An analogous curve is found for Y-Mn02 , but in this case, the peak lies at a lower potential than that measured for p-Mn02 and appears at D.110V (ENH). 550 cb/g are required and the point of inflection of the semi-logarithmic curve is at 0.515V (ENH). In this system also, therefore, at least two cathodic reactions take place and the results imply that p- and Y-Mn02 display analogous beha viour. a e The equilibrium potentials U, (p-Mn02 and U^Y-MaO^) * respectively 0,860V (ENH) and 0,815V (ENH). It is well known that the equilibrium potential U^ (3-Mn02) of £-Mn02 is lower than that of the other forms of this oxide, the value being 0.740V (ENH). Using an over-voltage of 0.200V (ENH) for each variety of oxide, we observe that the current is 0.2 mA for $-Mn02i 1 mA for p-Mn02 and 3 mA for Y-Mn02. For a given over-voltage, the current intensity is thus 5-fold for p-Mn02 and 15-fold for y-MnC^* Peak A of the curve for 3-Mn02 lies at - 0.245V (ENH) and corresponds to 435 cb/g g-Mn02. In order to compare the electrochemical reactivities of the three modifications, we have chosen a limiting potential of + 0.1V (ENH) , which corresponds to the order of magnitude of potentials for a catho de in a cell of the Leclanche type. In this case, the number of cou lombs/gram Mn02 is 150 for 3~Mn02> 670 for p-Mn02 and 610 for Y-Mn02. Thus, for comparable values of the p-and Y-forms, the "reactivity" of the 3-form will be four times lower. One of us (J.B.) was able to show already in 1951 that the reactivity is very probably dependent on the presence of OH-groupe in the Y-and p-modifications. The semi-logarithmic reduction curve for 3-Mn02 is also non-linear, but it does not show a point of inflection. This indicates that only one reaction governs the process of cathodic exchange, in contrast to the analogous systerns of Y-ahd p-Mh02. Judging from the concavity of the semi-logarithmic curves for p-Y-and 3~Mn02, it would appear that the governing reaction is the same for all three modifications in the region beyond the points of inflection characteristic of the Y-and p- forms. On the other hand, the reactions are presumably different in the case of Y-and p-Mn02 before the point of inflection and this dif ference can be attributed to the presence of OH groups. This study was continued with the aid of X-ray analysis in an at tempt to characterize the changes in crystalline structure in the va rious forms of Mn02. 428 Jean Brenet For Y-and P-M11O2, an effect of lattice expansion is observed which probably leads to the formation of groutite. We have observed and des cribed this effect previously (CITCE, Madrid, 1956). In the case of g-Mn02, first groutite and then manganite is formed. These results in dicate that different mechanisms are involved and our studies, toge ther with those of other authors, suggest that in Y-and p-MnC^ the acidic OH groups promote diffusion- in the solid, thus corresponding to .the formation of a solid solution of MnOOH in MnC^. Along with the re duction of Mn^+ ions in the solid phase, this process seems to be res- ponsable for the expansion of the crystal lattice. Disproportionation of the MnOOH phase formed in the initial step of the reaction occurs in the ft-MnOo system and in the second reduction reaction of y-and p-Mn02» This primary step is common to the reactions of all tree modifications. Further, we have compared tlie types of p-MnO^ obtained under various conditions of disproportionation. It is observed that for an increa sing excess of sulphuric acid, the value of the coulomb/gram MnO£ ratio increases up to the fixed limiting fixed potential. It appears proba ble that the concentration, of the solution and the amount of acid in excess also have an influence on the electrochemical reactivity of the dioxyde obtained. Examination of the first part of the reduction cur ves reveals that the current due to the diffusion affect in the solid is predominant. This confirms the importance both of the role played by diffusion kinetics in the solid and of the effect of acidic OH groups of the highly reactive p-and Y-dioxides on this diffusion! It is hoped that these first results will provide a better appreach to the optimal conditions in the chemical synthesis of highly reactive manganese dioxides and define more clearty the reactional behaviour of p-and Y-Mn02 in comparison with the relatively inert $-Mn02« This stu dy has been of value in the determination of the important part played by OH groups in the reactive dioxide forms. If still remains, however, to establish eventual relationships between the BET surface and the electrochemically active surface and ascertain the role played by the semi-conductivity of these dioxides resulting from the presence of OH groups or of an impurity which confers acceptor properties as we des cribed previously at the I.S.E. Meeting at Dubrovnick. If this goal can be attained, it is hoped that the problem of reactivity will serve to widen the scope of application of electrochemically active compounds employed as electrodes or electro-catalysts. Jean Brenet 429 ELABORATION OF A NEW TYPE OF OXIDE ELECTRODE AND EXAMINATION OF ELECTROCHEMICAL BEHAVIOUR OF MIXED OXIDES OF, TRANSITION ELEMENTS J. RUCH, J.F. KOENIG, J. BRENET Laboratoire d'Electrochimie et de Chimie Physique du Corps Solide Université Louis Pasteur - STRASBOURG (France) Considerable attention has been devoted to the study of mixed oxi des during the last few years, but only little is known of the elect rochemical properties of these materials. In the present work, we have examined the electrochemical behaviour of a number of mixed oxides in aqueous media at ambient temperature. The electrodes (lcm2 area) have been prepared by thermal decomposi tion of nitrates directly on a inert conductor. The process is descri bed in the French Patent, provisionally number 71.43642 at 1.12.71. The quality of the electrodes was tested by comparing them with fi-Mn02, which is generally considered as a relatively inert modifica tion. Using this material as a standard, we varied the different pa rameters likely to influence the reactivity of the deposited sample. Fig. 1 represents the re sults obtained for various electrodes prepared under the same conditions but subjected to different di scharge intensities. The electrolyte was normal su lphuric acid. Study of various mixed oxides The various products were studied by intensity potential measure ments and by reduction reactions at constant applied current. Normal sulphuric acid was. used as electrolyte in all cases and the equipment was of classical type. For potentiodynamic curves,the electrode was vibrated mechanically with 50 c.p.s. ; the electrolyte, maintained at 25°C and stirred mechanically, was out-gassed with oxygen-free argon. We examined first the oxides in t:he Ni-Mn-0 system, since the elec trochemical data are readily available for the individual oxides for med by these metals. Three well-defined oxides are known in this sys tem : Ni6Mn08 , NiMnO* and NiM^C^. We shall consider here only com pounds which crystallize in the spinel system, in this case NiM^O^ The shape of the intensity-potential curve of this compound (Fig.2) is characteristic in that the first peak A is independent of the voltage- scanning conditions, which is not true for dissolved species. Repeated scanning in the region between the equilibrium potential and a position beyond this first peak shows reversal behaviour patterns according to the thermal treatment applied : a) Well-crystallized products prepared directly at high temperature 430 Jean Brenet and in homogeneous phase show a progressive decrease in peak intensity as the scanning operation is repeated. b) For the products subjected to additional heat treatments (e.g., 4 h or 8 h at 800° C in air), but crystallographically equivalent to the preceding materials, the second or third scanning trial indicates a greater peak intensity than during' the first trial. In subsequent trials, however, the same progressive decrease in intensity occurs. c) Samples previously decomposed at 400° C and then heated at 800° C display an even more pronounced effect of the preceding type. This variation in the behaviour of the first peak.is accompanied by an intensity increase of the second peak (B). The reduction curves of these products at constant intensity are shown in Fig. 3. The measurements correspond to a variation of voltage as a function of the number of coulombs required per gram of product. These curves are easily derived from the previous one. As long as the applied intensity allows the reaction to progress according to the me chanisms giving rise to peak A, the reaction continues at a constant high voltage value until the product is consumed. If, following this point, the intensity is too high (higher than that corresponding to peak A, when allowance is made for the variation in the latter as a function of reaction time) a second level appears at an appreciably lower potential and the discharge occurs partially, if not almost en tirely, according to the reaction which gives rise to peak B. In a li miting case, it is possible to visualize discharge intensities which are sufficiently high to eliminate the level corresponding to mechan ism A entirely. The behaviour of the systems comprising the other coin- pounds is summarized as follows: CuMn20, . This compound is much more difficult to prepare in pure form. Its phase range with respect both to composition and temperature is very narrow, but it has the advantage of being the-best electrical conductor in the manganite series, and can thus be subjected to the highest current densities. However, its behaviour differs from that of the foregoing spinel, as shown from the sharpe of the discharge VS. intensity curves (Fig. 4). In contrast to the previous case, the vir tual disappearance of the first level is not observed at high intensi ties and the aspect of the curve in this region is completely diffe rent for the two systems. GoMn^O, . This compound is a quadratic spinel as opposed to the fo regoing cubic ones, its behaviour is characteristic when its variation in reactivity is studied as a function of the temperature employed for the additional heat treatment following initial heating at a lower temperature. The two cubic spinels exhibit maximum reactivity when this additional treatment is carried out at 800-850°C, that is within the range of thermal stability of the respective spinel phases. Both the reactivity and the crystalline character of Colto^O^ increase stea dily as a function of increasing temperature and additional thermal treatment. Considering the various results as a whole, it appears that there exists in these systems a correlation between higher reactivity and a greater degree of crystalline character, and this would seem to fee contrary to generally prevailling notions. Jean Brenet 431 Influence of cations We have attempted in the present work to establish the effect of metals such as Ni, Cu, Co, Al, Cd, Zn and Ag associated with manganese in different manganites. Unfortunately, the spinels formed by some of these metals have nearly negligible electrical conductivity; however, their equilibrium potentials all lie in the same region. further, we have also considered the cobaltites, which have the same structure but have cobalt in place of manganese. The equilibrium potentiel changes only slightly; for equivalent current intensities (often with better electrical conductivity of the products), the discharge level is situated at appreciably lower potentials. By carefully selecting the "foreign" metal in the system consider ed, it is possible to improve effectively the reactive properties of the compounds in certain cases (e.g., by adding Cu). We have also carried out successive substitution of the constituent elements using both manganites and cobaltites. Fig. 5 indicates the results obtained in this way when Co is progressively substituted "by Mn in spinel com-r pounds of the type (CUQ^S Mn^ C02 75-x ) °4* This work was performed with tne aid of the "Direction des Recher- ches et Moyens d'Essais du Ministére des Armees". 1 calhodiquc mA fig 2 IOÖ 200 300 , 400 50O 600 700 800 f.g3 Mg4 ng 5 432 Elton J. Cairns LITHIUM/SULFUR SECONDARY CELLS E. J. Cairns, H. Shimotake, E. C. Gay, and J. R. Selman Argonne National Laboratory, Argonne, Illinois, USA Introduction. In the United States there is serious concern over the supply of electrical energy during peak-load periods, and over the emission of air pollutants by motor vehicles. Both of these problems might be ameliorated by the use of high-performance rechargeable bat teries. A battery wi"th a specific energy of 220 W-hr/kg, together with a specific power of 220 W/kg, a cycle life of at least 1000 cycles and a lifetime of 3 years or more would be required before a signif icant application to off-peak energy storage or electric automobile propulsion could be made. Furthermore, the cost of the battery would have to be low - in the range of $10-20/kWh of energy storage capa bility. The lithium/sulfur cell1"5 shows promise for development into a. battery having these characteristics. Experimental. The lithium/sulfur cell was comprised of an electrode of molten lithium held by a porous stainless-steel or nickel current collector, an electrolyte of anhydrous molten alkali halides (e.g.j LiF-LiCl-KCl eutectic, m.p. ^347°C), and an electrode of molten sulfur held by porous carbon. The cells were usually operated at temperatures in the range 380-410°C. Two sizes of cells were investigated. The small cells had electrode areas near 2 cm2; the large cells had elec trode areas of 11 to 40 cm2. The construction of the cells was similar to that previously reported;4 the details of the construction of the larger cells will be reported at the meeting. The interelectrode dis tances were 0.5-3 cm. The preparation and operation of the cells was carried out in a helium atmosphere having <1 ppm H2O and <5 ppm each 7 N2 and 02. The voltage-current density, voltage-capacity density (A-hr/cm2), ca pacity density-cycle number, and capacity density-lifetime relation ships for the cells were studied as a function of the following vari ables: (i) discharge current density, (2) charge current density, (3) charge cut-off voltage, (4) operating temperature, (5) current-col lector geometry for the sulfur electrode, (6) identity and amount of various materials added to the sulfur to increase its conductivity and decrease its activity in the electrolyte. The electric performance ex periments were carried out at constant current with the aid of a dc power supply. Long-term discharge-charge cycling was carried out auto matically with the aid of meter-relays and timer-relays. Automatic data acquisition and computer data reduction and plotting were also used. Results and Discussion. The voltage-current density curve of the lith ium/sulfur cell for short-time operation (<10 seconds after closing the circuit) is a straight line from an open-circuit voltage of about 2.4 V (fully charged) to the short-circuit current density (typically, 8 A/cm2 at zero volts). The slope of this line can be calculated from the resistivity of the electrolyte and the interelectrode distance, when low-resistance current collectors are used. The maximum short- term power density is about 4 W/cm2. The voltage-capacity density curves were determined at a constant cur rent .density to a cut-off voltage of 1.0 V on. discharge, and to vari- Elton J. Cairns 433 ous voltages in the range of 2.5-3.2- V on recharge. The voltage- cathode composition curves are characterized by two main plateaus. The reactions taking place at the. sulfur electrode are complicated by the formation of immiscible liquid phases Li (nearly pure sulfur) and L^ (a solution of I^S in sulfur, corresponding to 37-40 at. % Li).8 The region of immiscibility results in an emf plateau near 2.2 V, and the region of Li2S precipitation results in a plateau near 2.1 V.8' At significant current densities, these plateaus appear at lower voltages and generally appear as one main plateau in the voltage-capacity densi ty curve. Because sulfur has a low electronic conductivity and a high viscosity, the overall reaction at the sulfur electrode is mass-transport limited. Therefore, the capacity density decreases with increasing current - density, increases with increasing charge, cut-off voltage, increases with increasing temperature, and increases with increasing intimacy of sulfur-electrolyte-current collector contact. Details of these effects will be presented and discussed at the meeting, including cur rent collector designs yielding the highest performance. Selected results from some of the cells are presented in the table below. Cell Cathode Area Depth A-hr/cmz A/cmz Cycle Life Additives No. TyPe (cm2) (cm) >1 V Life (hr) (at. %) 55 Enclosed 2.5 2.1 0.17 0.20 800 1100 — Laminated§ 56 Enclosed 1 1.3 0.52 0.27 600 750 ™ Reservoir** 57 Enclosed 2 1.9 0.70 2.0 >900 >4500 6Se, 2T1 Disk L-6 Enclosed 4 0.9 0.36 0.21 329 1075 Mixed L-7 Enclosed 4 1.0 0.85 0.29 350 2017 6Se, 2T1 Mixed S-l Shielded 30 0.3 0.40 0.30 250 800 6Se, 2T1 Disk S-10* Enclosed 13 0.6 0.74 0.15 16 100 6Se, 2T1 Disk , S-l3* Enclosed 20 0.6 0.27 0.31 130 500 14Se, 6T1 Disk "Sealed cell. §These results were presented in Reference 4. These results show that lithium/sulfur cells with molten alkali halide electrolytes are capable of delivering capacity densities in excess of 0.4 A-hr/cm2 for'hundreds of cycles and thousands .of hours. The life times have been extended beyond those previously reported by using: (1) "enclosed" cathode structures (to be described at the meeting), (2) electrolytes with small anions and cations,8 and (3) additives to the sulfur. Acknowledgments.. The authors are grateful to R. C. Vogel and D. S. Webster for support and encouragement, and to R. K. Steunenberg and 434 Elton J. Cairns M. L. Kyle for helpful discussions. Valuable assistance was provided by F. J. Martino and G. N. Vargo. This work was performed under the auspices of the Office of Air Programs of the Environmental Protection Agency and the U. S. Atomic Energy Commission. References. 1. H. Shimotake and E. J. Cairns, Presented at the Electrochem. Soc. Meeting, New York, May, 1969, Abstr. No. 206; See also Extended Abstracts of the Battery Division, 5_, 520 (1969). 2. E. J. Cairns and H. Shimotake, Science, 164, 1347 (1969). 3. H, Shimotake. M. L. Kyle, V. A. Maroni, and E. J. Cairns, in Proc. Fisrt Internat. Elec. Vehicle Symp., Elec. Vehicle Council, New York (1969) p.392. 4. M. L. Kyle, H. Shimotake, R. K. Steunenberg, F. J. Martino, R. Rubischko, and E. J. Cairns, in 1971 Intersociety Energy Conver sion Conf. Proceedings, Soc. Automotive Engrs,, New York, 1971, p.80. 5. E. J. Cairns, R. K. Steunenberg, and H. -Shimotake, in Kirk-Othmer Encyc!. of Chem. Tech., Supplement Vol., 2nd Ed., John Wiley & Sons, New York, 1971, p.120. 6. L. A. Heredy, N. P. Yao, and R. C. Saunders, in Proc. First Internat. Electric Vehicle Symp., Electric Research Council, New York, 1969. 7. C. E. Johnson, M. S. Foster, and M. L. Kyle, Nuclear Appl., 3_, 563 (1967). 8. E. J. Cairns, J. P. Ackerman, P. D. Hunt, and B. S. Tani, Pre sented at the Electrochem. Soc. Meeting, Cleveland, Oct. 1971, Abstr. No. 48; See also Extended Abstracts, p.118. M Eisenberg c 435 HIGH ENERGY AND HIGH RATE LITHIUM ORGANIC ELECTROLYTE BATTERIES M. Eisenberg Electrochimica Corporation, Menlo Park, California, USA. The ever increasing needs for greater portability of com munications and instrumentation equipment, both in civilian and military areas of application, has directed attention in recent years towards galvanic electrochemical couples employing active materials. Such couples, using aprotic solvent systems, offer the potential for achiev ing energy densities and performance characteristics unprecedented in the field of conventional batteries. The theoretical basis for high energy density organic electrolyte batteries is discussed with attention.to a merit, parameter for various anode candidate materials, including lithium. The development of high energy density lithium batteries operating in the vicinity of room temp erature' and capable of high rates of discharge are des cribed. The system employs lithium anodes and cupric chloride cathodes with an organic aprotic solvent electro lyte and has a theoretical energy density of 110 7 WH/k (based on the weight of active materials only) as compared to only 4 57 WH/k for the silver-zinc system, the highest . energy density battery known among aqueous systems. The cupric chloride-lithium organic electrolyte battery described in this paper has been developed in the direct ion of high discharge rate capabilities over the range of 10 to 6 0 minute discharges. In this range even small cap acity cells (.up to 10 ampere-hours) can yield energy densi ties from 3 5% to over 10 0% greater than equivalent silver- zinc cells at the same rate of discharge. Due to the broader range of liquidity of organic electrolytes com pared to aqueous systems, the new battery is also capable of operating efficiently at lower and higher temperatures than is possible with aqueous alkaline cells. Performance results will be shown over the entire range of tempera tures from -55°C to +74°C . Galvanostatic polarization studies have been conducted for both the cupric chloride cathode and the lithium anode. Using an exchange current density of the order of 0.3 x 10~3 amps/cm2, the anodic dissolution of lithium can be judged as fairly reversible. It is the cathode kinetics that is the limiting rate factor in the overall galvanic cell process. 436 Karl-Joachim Euler SPATIAL CURRENT DISTRIBUTION IN NON-ISOTROPIC POROUS ELECTRODES Prof. Dr.-Ing. Karl-Joachim Euler Fachbereich Physik der Gesamthochschule D 3500 Kassel, Bundesrepublik Déutschland In general, the spatial distribution ofcurrent or charge conversion has been treated so far for the case of uniform and isotropic electrodes. These simplifications are valid as long as only flat electrodes are investigated and as long as only general informations about the elec trode behaviour are needed. Recently, the influence of non-uniform ion and electron resistivities as well as non-uniform overvoltage has been investigated. Introducing these, two-phase electrodes show;-rather "strange" distribution of current conversion (1). The effects in triphase electrodes are less important (2). In all practical electrodes the current flow is multidirec tional. Thus, not only perpendicular but also tangential components of the electrical current density have to be considered. The mathematical treatment of non-isotropic resistivities in porous electrodes is rather complicated. However, se veral particular cases can be investigated, which show the direction as well as the magnitude of the effects to be expected. (1) K.J.Euler, Naturwiss. j>£ (1969) H. 6, S. 326/27 (2) K.J.Euler, 158th Natl. Meeting Am.Ch.Soc.Preprints 13 no 3 pp 176-202 (1969) J. A. Harrison 437 THE OXIDATION OP Pb IN H2S04 G. Archdale and J. A. Harrison The Electrochemical Laboratories, Chemistry Department, University, Newcastle-on-Tyne, England. The oxidation of Pb in H0SO4 has been studied by potentiostatic sweep, pulse, rotating disc, and rotating ring-disc. Thepresults show that Pb dissolves strongly as Pb and PbSO. ions in parallel with a solid state formation of PbSÖ*. However the solid state reaction is slow, and has a nucleation overpotential. This allows a potential range above the Pb/PbSQ4 potential in which Pb and PbSO. are observed to go into solution before the solid layer thickens. This suggests the possibility of a solution-precipitation mechanism in addition to the solid state reaction for the formation of solid Pb SO4. The fact that Pb , PbSO^ ions are precursors of solid PbSO^ makes this more feasible. The kinetics of the solid state formation of Pb304 have been analysed. Consequences of the solution- precipitation and solid state mechanisms will be discussed. 438 Adam Hel ler INORGANIC ELECTROLYTE BASED, ROOM TEMPERATURE, LITHIUM/CHLORINE, SODIUM/CHLORINE, LITHIUM/SULFUR, SODIUM/SULFUR AND LITHIUM/CUPRIC FLUORIDE CELLS Adam Heller, Kenneth W. French and James J. Auborn GTE Laboratories Incorporated, Bayside, New York and Waltham, Mass. USA. Introduction The chemistry and the physics of inorganic, aprotic ionic solutions has been a subject of study in our laboratory since we introduced in organic liquid lasers in 1966. The solvents used in these laser solutions were oxyhalides (such as SeOCl2 and POCI3) in which salts of Nd^+ were dissolved. Following our observation of the phenomenon of cathodic electroluminescence in 1970, our effort concentrated on the analysis of the process leading to light emission in inorganic ionic solutions.^»^ The study established three facts of significance in electrochemistry. First, that the application of an adequate external potential to metal electrodes immersed in an electrolyte, causes the injection of electrons from the Fermi level of the cathode into the electrolyte near the electrode; second, that the trapping of these electrons by solvated cations near the electrode is the primary step in electrodeposition of metals; and third, that some of the injected electrons acquire sufficient energy, when accelerated into the interface to excite solvated cations,.and thus to produce cathodic electroluminescence. In the specific case of solutions of trivalent rare earth cations (Re3+) in phosphorus oxychloride, the competing cathodic processes are excitation by inelastic collisions with injected electrons followed by luminescence and electrodeposition of rare earth metal. The anodic process leads to the formation of (dissolved) chlorine. Thus, application of an external potential' to electrodes immersed in the electroluminescent solution leads to rare earth metal/ chlorine cells with open circuit potentials of 3»^ -3-5 V. The forma tion of the high potential rare earth/chlorine cells, along with the observation that lithium metal is stable to extended boiling in phos phorus oxychloride, led us to develop a series.of high energy density cells with lithium and sodium anodes, some of which are rechargeable. Electrochemical Properties of Phosphorus Oxychloride Based Electrolytes Phosphorus oxychloride melts at 2°C and boils at 105°C. It is the first ambient temperature solvent which we know to be stable both to strongly reducing metals, such as lithium, and to strong oxidizers such as chlorine, bromine or sulfur. The conductivity of the pure solvent is less than 7x10"' ohm"^cm"' at 20°C. Conductive electro lytes are obtained by dissolving lithium (or sodium) salts with large anions in the solvent, or by in situ formation of such salts from sus pended lithium chloride (or fluoride) and a Lewis acid, such as AICI3, BCI3 or BF3. LiAlCl^, LiBClit, LiBFj^ and LiPFfc have, at 25°C, solu bilities of 0.9, 1.2, 1.5 and O.k moles/liter respectively, while the respective conductivities of the saturated solutions are of 7.8x 10"3, 8.7xlO~3, 1.1x10-3 and 3.7 xlO"3 ohm-lcm"!. The absolute viscosity of the solvent at 22°C is 0.97 cent?poise. Adam Hel ler 439 Rechargeable High Energy Density Cells with Chlorine/Platinum and Chlorine/Graphite Cathodes The desirable features of high energy density lithium/chlorine cells arejwelI_ recognized.5 Prior to the present work, high temperature lithium/chlorine cells based on fused salts (such as fused lithium chloride) were used. The high;temperature cells not only have high energy densities (> 1000 whr/kg) but also exceptionally high power densities. The Li/Cl2 fused salt cells run, however, into materials problems, which do not allow their operation for periods exceeding several hours. Our cells do not encounter such problems. Many standard structural materials, including Teflon, a range of metals, glasses and ceramics are stable to the electrolyte and to the cell constituents. The phosphorus oxychloride based Li/C12 ce^ls have an open circuit potential of 4.08 V, close to the potential calculated from the free energy of formation of solid lithium chloride from lithium metal and chlorine. Typical power densities are of 25 mw/cm2 with either platinum or porous graphite cathodes. The electrode reactions of the eel 1 are: charge Li. + + e <-r;—. Li (metal) (cathode) discharge v v charge AlClT t » A1C1, + 1/2 Cl. + e discharge AlCl, + P0C1 -• P0C1* + AlClT (anode) POCl* + Cl" -• POCl, the net anode reaction being charge Cl" ^icharni U2 Cl9 + e" discharge 2 Similar rechargeable cells, operating at ambient temperatures can be made also with ,sodium anodes, and with bromine/platinum and with bromine/graphite cathodes. The open circuit potentials of the various cells are, at 25°C, 4.08 V L1/POCI3: LiAlCl4/Cl2(Pt) 4.08 V Li/P0Cl3:LiAlClz/Cl2(graphite) V Li/POC^rLiAlBrj/B^Pt) 3.73 3.73 V L1./POCI3: LiAlBr/,/Br9(graphite) 3.84 V Na/POCy.LiAlCli/Cl^Pt) The potentials are close to the ones calculated from the free energies of formation of the product molecules. High Energy Density Cells'with Lithium Anodes and Solid Cathodes Cells having metal halide, metal sulfide and metal oxide cathodes and lithium anodes are known to yield high energy densities in organic electrolyte systems (such as propylene carbonate solutions of lithium perchlorate)£ Many of these cathodes are, however, either attacked 440 Adam Heller by, or slightly soluble in, phosphorus oxychloride based electrolytes. For this reason they do not lead to useful batteries. However, sulfur and nickel sulfide do not have such disadvantages, and stable, room temperature, lithium/sulfur, sodium/sulfur and lithium/nickel sulfide cells can be made. These cells have open circuit potentials much higher than expected for reactions in which lithium sulfide is the product. While the calculated potential for the reaction forming lithium sulfide from lithium and sulfur, is 2.3 V at 25°C, we measure an open circuit potential of 3.17 V in our lithium/sulfur cell, and 3.1^ V in our Li/NiS cell. To explain these potentials we postulate the cell reactions: Cls S% CI 2 Li + 1/2 S + 2P0C1- -+ 2 LiCl + P ?( (3-17 V) •* Cl'" " Cl CI S CI 2 Li + NiS + 2P0C1, -> 2 LiCl + ~Np' %P' (3.14 V) i Cl'rl'"U "II rCIl Ll 0 0 bl References 1. A. Heller, Physics Today, 20 (11), 34 (1967). 2. A. Heller, K.W. French and P.O. Haugsjäa, Chem. Phys. Letters JHO, 127 (1971). 3. P.O. Haugsjaa, A. Heller and K.W. French, Chem. Phys. Letters J£, 130 (1971). 4. A. Heller, K.W. French and P.O. Haugsjaa, J. Chem. Phys. 56, 2368 (1972). 5. For review see E.J. Cairns and H. Shimotake, Science 164, 1347 (1969). ' 6. For a review see R. Yasinski, Adv. Electrochem. and Electrochem. Eng. Ö, 233 097D. This work was supported in part by the Office of Naval Research, Washington, D.C. Mar gar e t e Jun g 441 A NEW RECHARGEABLE MEROURY/MERCURY OXIDE-ELECTRODE Margarete Jung VARTA Forschungs- und Entwicklungszentrum D-6233 Kelkheim/Ts., Federal Republic of Germany For some time mercury oxide has been employed already as an active mass in primary batteries (1,2,3;. Useful tech nical forms have been constructed by Rubens (4,5).Recharge ability of electrodes containing this oxide mixed with graphite is normally very difficult since on discharge the liquid mercury formed will coalesce to small droplets and on charging these, droplets will be covered with only very thin oxide layer.A number of procedures have been proposed in the past, to develop rechargeable Hg/HgO-electrodes (6, 11). However, due to their internal structure only a few cycles could be obtained. In cases where comparatively high quantities of precious metals were employed such electrodes proved to be uneconomical and have not yet attained technical importance. Hence, the task arose to search for a rechargeable Hg/HgO-electrode in which on the one hand such coalescence of mercury is completely suppressed and which on the other hand can be manufactured at a low cost by a simple process (12). The essential manufacturing steps are as follows: Mercury or mercury compounds are reacted with appropriate metals or metal alloys of the Raney type. Hereby the amalgams are formed. Especially Raney nickel has proved to be an excellent constituent for this amalgamation. Amounts up to 70 weight # of mercury can be fixed in the active mass. The remai ning is Raney nickel (18 - 20 weight fo) and oxides. The content of mercury can be varied in a relatively wide range depending on the conditions of activation and the type of mercury compound used. Generally, the amount of mercury increases with decreasing activation temperature. This can be explained by the fact that hydrogen, built into the Raney nickel structure, decisively determines the amount of mercury precipitated. Moreover, the mercury content increases.if salts with anions are used which can no longer be reduced by the nascent hydrogen. In order to increase the electronic conductivity, graphite or carbon black powder and alternatively nickel flakes are admixed. The Hg/HgO-electrodes are tested in combination with Cd- electrodes in sealed button cells, the construction of which is shown in fig. 1. Charge and discharge charac teristics of the cell are presented in fig. 2. The dimen sions of the cell are those of the wellknown 20 DK(20 mAh) Ni/Cd-cell. Charge utilisation is usually more than 97 per cent. On excessive overcharging small quantities of higher valency nickel oxides are formed, however, they decompose very rapidly on open standing and the electrode 442 Margarete Jung attains the reversible-Hg/HgO-potential. Essentially no marked drop in capacity was experienced on continuous cycling over 700 cycles. The active mass is also safe with respect to reversal of polarity, i.e. this sealed type of cell does not show any.change of dimensions due to gas formation on overloading and reversal. X-ray- analyses of the active masses showed that the peaks can be attributed to the presence of... the compound: NiHg^. On charging, this well crystallized amalgam is transformed to a distinctly x-ray amorphous intermediate phase before the well defined x-ray pattern of mercury: oxide, appears. A remarkable fact is that no peaks due .to; metallic nickel were observed at any charge or discharge state1. According to Davisson and G-ermer (13)» who studied the crystal structure of Raney nickel each fourth nickel atom is a surface atom. It is assumed that thes.e surface atoms are responsible for the amalgam formation. The rest of the nickel atoms are covered and are not detectable by x-ray. In addition Stereoscan-micrographs of the active masses in different states of charges and discharges are presen ted (fig. 3-5). Figure 3 shows freshly prepared and fig.4 deep discharged active mass, whilst in fig. 5 fully charged materials are shown. The regions of application of this new type of electrode may briefly be summarized: 1. In primary cells subjected to small rates of discharge where with commercial cells often short circuiting occurs due to the formation of mercury droplets. 2. In secundary cells' in combination with metal anodes such as cadmium, iron or zinc. 3. In standard references electrodes. Due to the good electronic conductivity of the active mass, these electrodes can be subjected to small current drains and still show the reversible Hg/HgO-potential. 4. In air-electrodes, since the mercury is being present in the active mass as an amalgam it can very easily react with oxygen, yielding mercury oxide. ?5 [V] charge- 1.0 discharge -| US Sctatn wrtpfwd potitiw tltctrod* [mAh] JO fig. 1 fig. 2 Margarete Jung 443 References: 1. Morrison US Pat. 975.885 (1910) 2. Ludvigsen. ' DR Pat. 290.748 (1915) 3. Heil US Pat. 1.195.677 (1916) 4. Ruben US Pat. 2.422.045 (1947) 5. Ruben US Pat. 2.422.046 (1947) 6. Edison US Pat. 704.303 (1902) 7. • Urry US Pat. 3.053.701 (1962) 8. Mallory Pr.Pat. 1.521.318 (1968) 9. Salcedo DAS 1.160.909 10. Ruben DBP 1.224.380 11. Ralston US Pat. 3.310.436 12. VARTA, M. Jung ,DAS 1.916.959 (1969) 13. Lavisson and Germer (Phys. Rev. 30(1927),705) 4*4 OldSicli Kou£il THE OXYGEN ELECTRODE MADE OF Ag2C03/Zn/PTFE MIXTURE. Cenek M.f Kouril O.. Calåbek M., Komårek L., Handera J. and Vanéöek J. The Faculty of Electrotechnical Engineering of the Brno Technical University, Czechoslovakia. The gradual drop of working voltage (of 0.2+0.3 V) is the greatest imperfection of low-temperature fuel cells. The cell voltage drop is affected,above all,by the cathode whether it contains silver or platinum catalyst; [l] • The attention was directed to the examination of alloys of transition elements (2,, 3} , to their choice from the viewpoint of application suitability as catalysts for electrochemical reduction of oxygeri, and to observation of the service life of the oxygen electrodes which con tain this catalyst. As catalysts of oxygen reaction? alloys of silver and magnesium [4] , silver with cadmium, tin, indium and antimony \5\ have been used. Silver of "Raney type" prepared of alloy of silver and rzinc [6] has been used, too, where zinc is usually not completely re moved from the alloy. The aforementioned data, as well as our own experience that metallic zinc affects favourably the parameters of the oxygen electrode of Ag/PTFE type[7] have induced us to examine the possibility to make oxygen electrodes of AgpCOo/Zn/PTFE mixture. Above all we have chooseo the ratio or components AgpCO^:Zn:PTFE =20:40:40 weight percent. Disk electrodes of Bacon tgpe were made of this mixture by pressing at 1,000 kp/cm pressure and they were sintered at the temperature of 391*393 C for 10 minutes. Electrodes made by this process indicated the potential of - 1,300 mV vs Hg/HgO in 7N KOH which is, approximately, the potential value of a zinc electrode in the same conditions. This potential is controlled, above all, by the reaction of Zn + 4 OH" = ZnOi" + 2HpO + 2e. When leaving the electrode in flooded state for a certain time (24 hours max.) and when applying oxygen under such pressure which causes its bubbling through the electrode, the result is, that the potential of this electrode subs tantially changes with time, as can be seen in Fig.l. Oxygen passing through the electrode pores is providing conditions for the origin of a local cell with oxygen depolarization. A certain value of mixed potential is corresponding to the degree of zinc corrosion and under gradual zinc corrosion the electrode potential gradually increases and reaches nearly the potential of a silver electrode in 7N KOH. Only the electrodes processed in this way are able to carry out the function of the oxygen electrode. The electrochemical parameters of the oxygen electrode depend 01d*ich Kou*il 445 substantially upon the degree of zinc corrosion, as seen in Fig.2. The electrodes were evaluated, above all, on the basis of comparing polarization curves and current densities i (mA/cm ) obtained at the polarization of -0.3V vs Hg/HgO; # = 25°C, 7N KOH, Po2 =/L0 + 30 cm of EjO column. Current densities of these electrodes were within 100 - 200 mA/cm . Furthermore, the structure of the electrode surface was examined with the aid of a scanning electron micro scope, as well as the degree of amalgamation and value of zinc corrosion. DTA and DTG were carried out with the initial material of the electrode and with the original mixture the electrodes were made of. The electrodes were subjected to service life tests in half-cell connection. The results of these tests which are given in Fig. 3 show a good stability of these electrodes. REFERENCES [l] COHN E.M. Report NASA N68-25640 (1968) [2] SAYED A., GREIN. C.T., RYBACK W.H., BIENSTOGK D. Report NASA, Contract No.W-12.300 (1968) [3] GINER J., PARRY J., SWETTE L., CATTABRIGA R., Report NASA, Contract No.NASW-1233(1968) U.S.Patent No. 3,458,360 (1969) [5] U.S.Patent No. 3,460,994 (1969) PRZYLUCKI J., BIELINSKI J., MADRY K., Third International Symposium on Fuel Cells, Brussels 1969, p.316; Pres. Acad. Eur., Brussels 1969 CENEK M., KOU&IL 0., §ANDERA J., VAMCEK J.; M Extended Abstracts, 22nd ISE Meeting, Dubrovnik, Jugoslavia 1971• 5 S 9 13 17 21 TIME OF 09 ACTION [HOURSj 4. I. FIG.1 446 Oldfcioh Koufcil Fig.l Time influence of oxygen action under pressure upon the potential change of Ag/Zn/PTFE electrode system polarized cathodically to the value of -1.5V vs-Hg/HgO for 20 minutes under flooded conditions in 7N KOH at 25 C. (Oxygen working pressure = 22 cm of HpO column). Fig.2 Dependence of current density of the oxygen electrode on the weight content of zinc in the electrode made of AgpCO-^/Zn/PTFE mixture = 20/40/ /40 weight percent. Zinc content was measured after 50 hours of electrode loading at 25 C; i (mA/cm2) measured at -0«3V vs Hg/HgO in 7N KOH. Fig.3 Service life of Ag/Zn/PTFE oxygen electrodes at 25°C in half-cell connection, in 7N KOH. (Ag?C07/ /Zn/PTFE = 20/40/40 weight percent; 62 mg Ag/ciir). Before applying oxygen the electrodes were flooded for 10 *:13 hours in 7N KOH. Oxygen working pressure = 18426 cm of HpO column; i (mA/cm ) measured at -0.3V vs Hg/HgO. "E .1 FIG. 2 t 20 30 *0 WEIGH PER CENT. 2n s 100 J) o o n y u o 4i h& A A FIG. 3 n TT-OD 90 - er° u u Q- t + =rt—i—^-+ "ifiäi rl—h eo + + —H H —Y 1000 2000 3000 4000 £000 TIME [HOURS] Alfons Lindholm 447 REACTION LAYERS AND ELECTROCHEMICAL PROPERTIES IN TUBULAR" ELECTRODES DUE TO ENGINEERING DESIGN' Alfons Lindholm Aktiebolaget Tudor, 440 41 Nol, Sweden Introduction The main purpose of the present work was to characterize the lead acid cell with special respect to the tubular electrodes in the positive plate. Because of the superior properties of the tubular electrodes, a great number of batteries in the world intended for tractionary and stationary use contain positive plates supplied with a certain number of parallel tubular elec trodes . Their outer envelope is usually of cylindrical shape and nowadays made of different acid resistant materials. The tube construction plays an important role for the capacity and the service life of the battery. With that in mind the construction must provide: high permeabili ty for electrolyte and gas, low electric resistance, good resis tance against oxidation attack at temperatures up to 60-70°C, high tensile strength, low deformability, sufficient bursting strength against internal pressure, and resistance to external abrasion. Being the container of the active material the wall structure of the tubes is considered to be of prime importance. The fibres which constitute the tubes may be braided, woven or felted. The macrostructure defined by the mesh size for braided and woven tubes and the density pro square cm of the fibres for felted tubes as well as the impregnating agent determine the access of electrolyte to the pore system of the active material (the ionic conductivity). The mechanical strength and deformability during operation as well as the apparent density of the active material determine to a high degree the electronic conductivity of the bulk material. In addition anomalous distributions of the discharge product, the cell geometry and the polarisation resistance are important factors which govern the discharge efficiency of the porous electrode. Extensive works (l-14) regarding kinetic studies and developed discharge and potential profiles inside porous electrodes have been published in recent years. Contributions (15-18) about morphological changes in the active material have shown new aspects of the discharge mechanism. The distribution of lead sulfate in flat plates using autoradiographic technique (19) and X-microprobe (20) have manifested the formation and variation of sulfate barriers throughout the active material. Tubes with good performance qualities have been used by several battery manufacturers since many years (21). The demand for new and still more qualified tubes has, however, increased in recent time. As a consequence this investigation describes the 448 Alfons Lindholm characteristic properties of already well established tube con structions as well as of tubes which are new products and the result of extensive experimental work. The influence of well modified physical tube construction, proper electrode geometry and balanced amount of active material is shown to be of highest importance for the proper and regular formation of electrochemi cal reaction layers throughout the electrode during its opera tion. The variations in internal resistance and polarisation properties under different conditions are also discussed. Experimental For statistical analysis of cell capacities we have worked with. up to 100 cells in each experimental series. The variables were: the material and the construction of the tubes, the inner dia meter of the tubes, the apparent density and the composition of the electroformed material, the diameter of the lead conductor and the available volume of the electrolyte. The general dis charge routine ended with a stepwise discharging process to get the current distribution profiles inside the electrodes. In a special made turning machine these electrodes were disproportio- nated into thin concentric layers, which were subjected to photometric analysis in respect to the reduction species. Elec trodes taken at random from the different series were then in vestigated for internal resistance and activation polarisation using square wave technique and for surface area using measure ments of the double layer capacitance. Results In this part of the work one well established double wall tube (Pg), i.e. an inner sleeving of braided glass fibres surrounded by a perforated PVC wrapping, is compared with a single wall tube (PgS) of braided glass fibres reinforced with axial warps. The fall of pressure at 10 lit air/min through the wall was for Pg about 90 and for PgS only 9 mm w.p. The electric resistance after 24 hours in sulphuric acid was for Pg 0.35 and for PgS 0.15 ohm x cm2. The internal resistances of some operating electrodes are shown in fig 1. It is evident that a high apparent density is favourable for the bulk resistance of the electrodes. The very small difference in resistance betv/een electrode PgS at 0°0 and Pg at +40°C is remarkable. The values of activation over- potential (fig 2) for electrode PgS at 0°C and -18°0 are of great interest according tö the charge acceptance at low temperatures. The observable higher ovérpotehtial of the Pg electrode is the consequence of nonuniform current distribution caused by. the shielding effect of the outer wrapping. Due to a pure activation polarisation the electrochemical reaction for the fully charged electrode results in a evolution of oxygen. On the discharged electrode the lead sulfate is transformed into lead dioxide. These divergencies in electrochemical properties may therefore be a good indication of the specific suitability of the diffe rent types of tubes for a variety of well defined practical Alf ons. Lindholm 449 applications in the lead acid cell. The reversible reactions, which occur on discharge, must consequently yield higher capaci ty outputs using tubes with low resistance and overpotential. The formation of reaction layers is presented in,fig 3, and 4. The tubes ($i=8#4 mm) containing different amounts of active material have in general quite distinctive distribution curves as well at 25 as at 50 percent of the discharge time. When scrutinizing the current distribution at final electrode .potential (at 100 a/o) it will become clear that the distribution of the reaction layers is highly dependent on the masstransport, the conductivities in the different phases and the thickness of the total active material. 21 references. Olle Lindström RECHARGEABLE METAL-AIR BATTERY SYSTEMS Lars Carlsson (a), Göte Granath (c), Gunnar Lindström (a), Olle Lindström (a,b,d), Leif Sköld (a), Lars Welln (a) and Ingvar Åkerblom (a);AGA Innovation Centre (a), Tudor AB (b), Swedish National Development Company (c) and Royal Institute of Technology (d), Stockholm, Sweden. The Interest In energy conversion during the sixties was In Sweden focussed on fuel cell systems for submarine propulsion. During th^e last years this effort was super seded by work on metal-air battery systems for powering Industrial vehicles used In mining, materials hand IIng,etct This paper will report results obtained In. a project Joint ly carried out by the above parties,with the heavy part of the development work carried out by the Power Sources Group of the AGA Innovation Centre, with support from the Swedish Board of Technical Development, The project was Initiated In 1968, got momentum In 1969 and prototype production started In 1971, At present a second generation of proto type systems Is In production and under test. Programs are now being made for production and marketing activities to follow In the middle of the seventies. The project has from Its beginning been oriented more towards hard-ware than to research, The goal was to manufacture a metal-air battery system on the basis of existing technology to achieve higher power and energy densities than present tract Ionary batte ries, Equally Important goals were that the new mass pro duced system should not cost more than the tractlonary lead acid battery and that It should have a similar life. The results obtained so far Indicate that these "goals will be reached In due time. Hard-ware projects to be carried out In a short time and with limited resources means among other things that a number of activities from electrode materials development to prototype production have to be carried out simultaneous ly. The major critical problems have to be Identified and solved whereas problems considered less Important have to be left for the next porject phase, A number of Important decisions had to be made early In the project. One such Important decision was the choice between the three candidate anode materials Iron, cadmium and zinc. Iron was chosen here on the basis of an Inter pretation of the available literature against the project specification. Another Important decision was type of cathode. The development of oxygen during charge Is detrl- 01 le Lindström 451 mental to many thin hydrophobic air electrodes developed for fuel cells. A third electrode may be used to over come this problem. However, we did not find the third electrode concept practical and consistent with the pro ject goals. We decided therefore to develop the classical double layer sintered nickel electrode for use In metal air cells with the fine layer serving as a kind of built- in third electrode. There were also some critical de cisions about the electrolyte systems though an alkaline electrolyte of course was required with-the chosen anode material. Circulating electrolyte was chosen here In spite of Jncreased compI ex Ity and larger leakage currents because of the process control advantages of this concept, Circulating electrolyte gives simplified control of water balance, carbon dioxide removal and cell check up, A choice also had to be made between monopolar and bipolar cell design. The monopolar design was chosen mainly be cause of its simplicity. The features Indicated above lead naturally to a filter press design of the type which Is quite popular In fuel cell technology. Electrode functions are of course most Important also In these battery systems though there are also a number of other functions of vital Importance to the performance of the system. The goal for the capacity density of the anode Is here preferably expressed as Ah/cm3 - the heavy Industrial applications we have In mind are In fact more volume critical than weight crltlea I, Anode.performance Is conveniently characterized by means of the pseudo- resistance parameters Introduced by Olle Lindström, ty pical figures are I Ah/cm-^ and I -O- » cm2, We have been able to reach these levels mainly by optimization of the structure of these sintered metal structures, A special production method had to be developed for these anodes, which at present are made In thicknesses from 0,5 mm to 4 mm. Work on air-cathodes, following similar lines as In anode development, has resulted In thin dual function cathodes with a high porosity and good neschanlcal strength, These electrodes operate at 0,2 bar differenti al pressure and use 1-10 mg/cm2 of a proprietary catalyst Current density at nominal power Is 100 mA/cm2. The pile design Is quite simple. The cathodes are welded Into polymer frames. Two frames are Joined to a double cathode air element, These air elements are stacked to gether with the anodes to a pile, The cell pitch Is at present 6,4 mm for a double cathode cell. The effective electrode area In a cell Is 400 cm , 452 Olle Lindström A number of design, production and systems problems had to be solved In this project; In fact the major project efforts have gone Into these activities» Among the func tions of the auxiliary systems can be mentioned the con trol of water balance, carbon dioxide, temperature and air circulation. One Interesting feature Is that the rate of air supply Is not governed by the oxygen demand, of the battery but by thermal requirements. The operating tempe rature of the system tras to be fairly high, between 40 C and 50 °C when the air temperature Is around 20 °C, to get efficient cooling. Cooling takes place by evaporation of water from the electrolyte to the air In the air ele ments. This requires tha.t evaporated water be recovered and returned to the system. One benefit of this feature Is that heat can be made available for Instance for heating the driver's compartment etc. The auxiliary sys tem Is weighing between 5 to 10 % of the weight of the battery p I Ies. Performance of batteries Is best represented by logarith mic power density versus energy density charts. Compari son In this way between data for commercially available batteries and our metal-air battery systems Indicates that the metal-air battery system may be 3-5 times better than the tractlonary lead acid battery on a weight basis over most of the power range. Several hundred cycles have been demonstrated with the first generation of prototype piles. Laboratory data Indicate that the 500 cycle goal may be achieved with the second prototype generation. Projections based on present technology furthermore In dicate that metal-air battery systems may become quite competitive on a cost per Installed kWh basis In large scale production. There Is still a long way to go, how ever, and the outcome of this stride Is depending not only on the sucessful development of the metal-air battery system but also on results achieved In the simul taneous development of conventional and unconventional competitive battery systems. Today It seems, however, reasonable to assume that there will be metal-air battery systems on the market and that there will be a market for these systems within not too many years. Jifi Mrha 453 ELECTROLYTE DISTRIBUTION AND SHIFT THROUGH CARBON ELECTRODES FOR THE ~ELECTROREDUCTION OF AIR OXYGEN Jifri Mrha.JiM Jindra and Miroslava Musilova* J.Heyrovsk^ Institute of Physical Chemistry and Electro- chemistry,Prague,Czechoslovakia. Some electrochemical properties of porous active carbon -teflon "air breathing"electrodes in alcaline solutions were studied. This electrodes consist from two layers with quite dif ferent hydrophobic properties.Active layer was prepared from a micture of finely ground active carbon (electro- catalyst) and polytetrafluorethylene (binding and impre gnating material).The gas layer was prepared from a pure polytetrafluorethylene• A special attention was paid to the electrolyte distri bution in the active layer as a function of cathodic load.Electrolyte distribution was measured by intrusion of mercury at low temperatures.The experimental data now a hand show no influence of cathodic load. Paralel to these measurements the electrolyte shift throu gh the active layer as a function of cathodic load has been studied.In contrary to electrolyte distribution measurements a strong influence of cathodic load on the electrolyte shift through the active layer was observed. Experimental data obtained were interpretated as a con sequence of KOH concentration gradient across the active layer(osmotic phenomena,surface tension changes). 454 Konrad Mund ON THE STRUCTURE OP CATALYSTS AND ELECTRODES POR ELECTRO CHEMICAL CONVERSION OP GASEOUS REACTANTS Konrad Mund Porschungslaboratorien der Siemens AG-, Erlangen, Deutsch- land Catalysts for electrochemical conversion of gaseous re- actants are finding technical interest in fuel cells. In general the performance obtained sdfar is poor if the reaction rate at the catalyst/electrolyte interface is considered in relation to real surface area. Por techni cal purposes, therefore, it is essential to prepare elec trodes with optimized porous structure resulting in a high utilization grade of the large surface area. Por the transport of gaseous reactants macro pores must be imparted to the electrodes so as to bring the gas as close to the periphery of catalyst particles as possible. Among the several methods of electrode preparation the one of supported electrodes is preferred in our laborato ry. Based on this principle, catalysts with different structures can be prepared which find application for various purposes. For instance, Raney Ni and Raney Ag electrodes for the reaction of H2 and O2 respectively in alkaline fuel cells and WC and C electrodes for acid elec trolytes. A simple model is proposed for the working process in such electrodes taking the influence of catalyst struc ture and thickness of the working layer into considera tion. The model is based on the assumption that the gas at first rapidly dissolves in the electrolyte and then diffuses to the walls of the micro pores resulting in ultimate sur face reaction. It can be shown that the ohmic drop inside the micro pores is small compared to the total polariza tion loss if technical current densities are drawn into account. The process inside the pore of the catalyst is, therefore, treated mathematically taking both charge transfer and diffusion into consideration and linearizing the problem. The polarization is a complex function of current, but a linear behaviour can be shown to hold good for polarization values ^ no* exceeding ^ = 6 RT/E. Thus the current I of a single pore is proportional to r3/2, where r is the radius of the pore, it is directly propor tional to the polarization ry and to the square roots of exchange current density i0, diffusion coefficient D and Konrad Mund 455 solubility c0 of gas in the electrolyte: Taking A for the surface area and R for the particle ra dius of the catalyst one can find the current ic produced by a certain amount of catalyst having a porosity p, a specific weight *fr and a torduosity factor x: u - ^Uv^-^-v- r^-%1 - V v ^'Y- (2) For certain catalyst particles \ describes the 'quality' of the material itself, including electrocatalytical ac tivity of the surface and the porous structure as well. Neglecting the electronic resistance between the parti cles it can be shown that for the thickness d of the work ing layer, the polarization resistance t\ ek/n2| 15*j-=20fj,g=57a-cm,\=168si'-g rl ~> ,-'.„-> W 40 sfi£50/J,9=29,3a-cm,\=55,5äg 2 6n KOH,60°C,4N-cm2 ~~Öl 02 03 mm CU Fig.1: Reciprocal polarization resistance of Raney Ni electrodes employing different thickness of working, layer and particle diameter Waldemar Nippe 457 ELECTRODES POR IRON-AIR-SECONDARY CELLS M.W. Nippe, H. Cnobloch, D. Gröppel, G. Siemsen, P. v. Sturm Forschungslaboratorium der Siemens AG-, Erlangen, Deutsch- land Theoretically the iron-air system looks promising as a secondary battery in automotive applications, especially concerning high energy density. With respect to high power density the electrodes of such a battery should be capable of delivering a current density of 50 mA/cm^ under sta tionary conditions and of 150 mA/cm^ for short intervals corresponding the acceleration time of a vehicle. Our cell will consist of three electrodes, one iron electrode and two air breathing electrodes. The latter serving also for the oxygen evolution process no further auxiliary elec trodes for charging are installed. The kinetics of the iron electrode is rather complicated. This might be recalled with aid of fig. 1, which presents a potential sweep diagram..Starting at the left lower cor ner where hydrogen evolves in positive direction the cur rent decreases and becomes positive. In peak I part of the occluded hydrogen is oxidized. In peak II iron is oxidized to the two valent state. It is known that this process in volves the solution phase with an anion HPe02 according to Pe + 3 OH" —• HFe02 + H20 + 2 e~ (1) The bihypoferrite-ion hydrolyses and iron hydroxide is precipitated. HFe02 + H20—•Pe(0H)2 + 0H~ (2) If the potential is shifted to more positive values, then in peak_III Fe304 is formed, initially over an intermedi ate Fe02-ion according to + HFe02 —•PeO^ + H + e" (3) and HPe02 + 2 Pe02+H20 —• Pe504 + 3 OH" (4) in addition to that at even more positive potentials in peak IV in a topochemical reaction 3 Pe(0H)2 + 2 OH" —• Fe^ + 4 H20 + 2 e" (5) Whereas this reaction scheme is largely proved by röntge- nographic investigations, this seems not to be the case in the reverse scan. We assume that in peak V Pe^O^is reduced over the solution and in peak VI topochemically to Pe(0H)2 which in turn is reduced to elementary iron (peak VIIJ. Prom these statements a simplified reaction- scheme can be derived (fig. 2). As iron II hydroxide is a voluminous product the pores necessary for ionic con- 458 Waldemar Eippe duction and diffusion of hydroxylions are soon stopped, thus blocking further oxidation of iron. Por this reason conventionally constructed iron elec trodes can be discharged by only small current densities. At high current densities the layers adjacent to the elec trolyte are oxidized first, and the pores are clogged. We achieve a more uniform oxidation by imparting the elec trode a special structure which is shown in fig. 3. In this structure the coarse pores perpendicular to the elec trode/electrolyte interface are responsible for transport processes. The coarse pores are formed by iron fibres, the fine pores by iron fibres with the active mass im bedded in them. The active mass consists of iron, acety lene black and a plastic resin as a binder. The black serves for good electric conduction even when the iron is already oxidized to a large extent.-The resin binds the components of the active mass. Its composition is 0.56 g/ cm2 iron, 0.05 g/cm2 black, and 0.02 g/cm2 resin. The electrode is encaged in expanded metal to avoid segrega tion of the structure when submitted to swelling during discharge. The electrodes are tested in half ceil arrangements. They are charged and discharged at a current density of 40 mA/ cm2. The discharge- is interrupted at -500 mV with respect to the Hg/HgO-electrode in the same solution. Then the discharge covers the two and three valent oxidation steps. With regard to the two valent step the active mass is oxi dized to a degree of 45 $. This corresponds to a 6 hours discharge rate. The charge density of a complete elec trode amounts to 0.3 Ah/cm2. With this type of electrode we achieved under the testing conditions over 200 cycles. In fig. 4 potential charge curves are given with the cur rent as parameter. The electrode can be discharged by high current densities. The air electrode is given a three layer structure. The layer adjacent to the electrolyte consists of porous nickel sheet, to which the .active carbon layer is attach ed,.; followed by a PTFE layer preventing the electrolyte to seep into the gas space. During the charging periods the oxygen replaces the electrolyte in the pores of the nickel layer, thus shielding the carbon layer from elec trochemical attack. This type of electrode has also been tested in half cells with 30 mA/cm2 alternating the current every 24 hours. So far, this electrode has run 200 cycles corresponding 4800 hours in each operating mode. The potential of the air electrode amounts to -200 mV at -30 mA/cm2, whereas -400 mV are measured at -150 mA/cm2. At a charging rate of +30 mA/cm2 we find a potential of +600 mV. 0 2 _ 2A6/3 |2^i^§ii2S_5i5SÉ HFeO; - FeO; liquid phase i Fe Fe(OH) Pe solid state H 2 "*" 3°4 0CD P H 6oo mv ~4oo Fig. 2. Oxidation mechanism of iron S3 Fig. 1. Potential sweep diagram of iron in 6 n KOH metallic fibers with active mass metallic fibers kkurrentcollector -SCsfenV Q2'Ah/un2 DA 0 —-~ expanded metal QQ50A Fig. 4. Potential-charge behaviour Fig. 3. Gross section of the iron elec of the iron electrode trode *0 460 Lars fijefors SELFDfSCHARGE OF IRON ELECTRODES WITH HYDROGEN EVOLUTION IN ALKALINE SOLUTIONS Lars Qjefors Dep. of Chemical Technology, Royal Institute of Technology, S-100 44 Stockholm 70, Sweden. The reaction with the electrolyte to produce hydrogen, particulary at elevated temperatures (40-60 C), is one problem associated with the use of iron electrodes in alkaline solutions. This reaction has two ; disadvantages, i.e. it is a selfdischarge process and the production of hydrogen is undesirable for safety reasons. The present work on the , corrosion of iron in alkaline solutions was undertaken to establish • basic information about the process and the influence of KOH-concen- : tration and additives to the electrode and the electrolyte. j i The experimental work has been carried out with porous iron electrodes t ; sintered in hydrogen atmosphere. ; A slow potentiodynamic method has been adopted (1) for investigation ,! of the discharge process and it has the advantage of showing plainly I the reaction processes which take place at the electrode. The curves : thus obtained are actually the curves of the electrode discharge, as j the rate is very slow (5,3 mV/min.). Fig. 1 shows that the discharge j process can be divided into three waves compared to two levels in | constant current discharge. Fig. 2 shows discharge curves of one ! electrode freshly charged (50 mA/g for 15 h) and one electrode gharged in the same way and then left on open circuit for 75 hours at 50 C. The amount of hydrogen (34,5 ml/gram at 5 0 C = 62,21 mAh/gram) corresponds relatively well with the difference in area (69,65 mAh/gram). This ^haws that the Ho-evoluting selfdischarge process is significantly greater than oxidation oy 0„ which is diffusion limited. It is well in accordance with results byTIersch (2), which showvthat the cathode process at 25 C to 90% is hydrogen evolution. Current I Fe—HFeCv-Fe(0H\ Electrode,freshly charned (50 mA/g for 15 hours) at 50": (mA/g) /Fe—FeO;—HFeOl—Fe(OH) B \Fe—FeCf Current Hydrogen evolved:34,5 ml/n (mA/g) =62.21mAh/o Difference in discharge capacity: 69,65 mAh/g Potential Potential vss Hg-Hg(O slectr. vs Hg-HgO *- (mV ) (mV) Potentiodynamic discharge of a porous iron Potentiodynamic discharge of one freshly electrode (thickness 1,1 mm) ir 5-M KOH charged and one selfdischarged electrode Scanning rate 5,3 mV/minute. (thickness 2,4 mm). Scanning rate 5,3 mV/min. Figure 1. Figure 2. Lars öjefors 461 Table 1 shows a number of reactions with hydrogen evolution which can occur spontaneously. Investigations by X-Ray diffraction show that Fe(OH)2 is the main selfdischarge product. At temperatures above 5 0°C small amounts of Fe^O* can be detected. It seems that reaction V pre dominates and reaction I or II, altough thermodynamically favoured only occurs at higher temperatures. This corresponds very well to Fig. 2, which shows that the Fe(OH)2-wave has almost disappeared after a 75-hour selfdischarge process at 5 0 C. Table 1 E.m.f. (volt) for hydrogen evoluting reactions (25°C) (3) Fe(OH) - 2H 0 = Fe + H 3 2 2 3 °4 2 + 0, 20 (I) 3 Fe + 4H20 = 3Fe3 p4 + 4H, +0. 085 (ID 2 Fe(OH)2 - H20 = Fe2 Os + H2 +0, 05 8 (III) + +0. 2 Fe + 3H20 = Fe, °3 3H, 051 (IV) Fe + H20 = Fe(OH)2 + H, +0,048 (V) This extensive selfdischarge on open circuit can be retarded either by influencing the anode process - which is undesirable - or by influen cing the hydrogen evolution. This latter effect can be obtained with additives to the electrode or electrolyte, which increase the hydrogen overvoltage. Fig. 3 shows decrease in capacity, calculated from H2-evolution, for charged iron electrodes in KOH-solutions of different molarity, during" a 200-hour period. The rate of H2-evolution is constant during the first 25 hours and decreases then slowly. The initial corrosion rate can also be calculated from the linear polarization technique which gives a better understanding of the effect from additives. Fig. 4Qshows a pola rization curve for a porous iron electrode (5-M KQH, 5 0 C) in the re gion of hydrogen evolution and active dissolution. Included for com parison is a theoretical curve (4) calculated for the values: 7 2 6 2 8 2 i°h = 2 • 10~ A/cm ; i°Fe = 5 • 10" A/cm ; id = 4 • 10" A/cm The agreement is very good below current densities of 4 • 10"3 A/g where limitations due to diffusion in the pores become neglible. Capaci ty Total capacity at constant current discharge(60 mA/a): Experimental curve . Potential ys Hg-Hg0 (mV) 930 s / S Corrosion pot.962 nV 950 / ^^. 970 . 7-H K0H ^^*^v:- 5-M KOH 3-M KOH 990 1-M KOH 1,40 lO'Vg ; \ 1010 id4 io- V Curren 150 HoursT Decrease' in capacity,calculated from hydrogen- ' (A;B) evolution in 5-H KOH for porous iron elect Polarization curves for an iron electrode rodes (thickness 2,4 mm). \l\K me) in 5-M KOH at 50*C. Figure 3. Figure 4. 462 Lars ÖJefors A plot of log rcorr v. s 1/T gives a linear relationship according to the formula: A e *corr = ' *P ART ) which is shown in Fig. 5 Additions such as K^S and HgCl2 lowers the preexponential A while the energy of activation does not cnange notaUy. The influence of selfdischarge inhibitors on the normal discharge pro cess has been studied by the slow potentiodynamic method (Fig. 6). . Currant * («A/g) no additives electrode•dopod in HgCU .Corrosion currant 0,175 g KtS/liter Mo" (A/o) i - - - - i 5-M KOH,no additives ! 0,80 g Na-mataarsenat/ liter 5-M KOH,the electrode 0,80 g EDTA/liter doped in HgClt 5-M KOH • 0,0175 g K£/l Potential vs Hg-Hg0 T»o — *i~ ~T*e »» —* («»V) Arrhenius plot of corrosion rates for Potentiodynamic discharge of a oorous iron porous iron electrodes (thickness 2,4 mm) slectrode (1,1 mm) at 50 C in 5-M KOH with calculated from hydrogen evolution,50'C.. different additives.Scanning rate 5,3 mV/min. Figure 5. Figure 6» The influences of a large number of additives have been studied with the mentioned methods as well as scanning electron microscopy and corrosion potential curves. References: (1) Labat, Jarrousseau, Laurent: Power Sources Conference 1970 (2) Hersch: Trans. Faraday Soc. 5.1, 1442 (1955) (3) Latimer: The oxidation states of the elements and their potentials in aqueous solutions, New York 1952 (4) Kaesche: Die Korrosion der Metalle Springer Verlag, Berlin (1966) Paul Ruetschi 463 ION SELECTIVITY AND DIFFUSION POTENTIALS IN CORROSION LAYERS. - PbSO. FILMS ON Pb IN H2S04 Paul Ruetschi m Leclanché S.A., Yverdon, Switzerland Anodic corrosion films on Pb in aqueous H?S0. exhibit, when formed within,a certain potential range, a characteristic multiphase structure. Such films then consist in their ex terior part (against the solution) of a layer of PbSO. and in their interior part of tetragonal PbO. Individual PbSO. crystals in the outer layer have dimensi ons in the order of 0,1 to 1/um, or more. When the layer has reached a thickness of 1/um or so, SOT2 and HS0~ ions are no longer able to penetrate across, since the pores are then practically closed. Further corrosion of Pb then occurs by dissociation of H?0 underneath the PbSO/ layer, precipi tation of the anodically generated Pb 2 in form of PbO, and migration of the H ions away from the electrode interface. The potential-determining couple, Pb/PbO, is thus exposed to a basic environment (e.g. pH 9,34, corresponding to equal concentrations of Pb 2 and HPbOp) Under these conditions one should expect the existence of a considerable diffusion potential across the PbSO. layer, this latter displaying the characteristics of a precipit ation membrane with pronounced ion selectivity. It appears that diffusion potentials in corrosion films have received little attention hitherto. Generally, it may be shown that membrane potentials due to concentration differences between opposite sides may, at vanishing matrix charge and vanishing current, be expressed >y < 2 (j/d = - (RT/F) fZ. (t±/Z±) d In. a. where t. = transport number Z. = valency a. = activity of ions of type i This equation follows from the condition of equilibrium for those ions i, to which the membrane is permeable. If the completely formed PbSO/ membrane becomes essentially non- permeable to SO 2, HSO. and Pb+ ions, so that the dif fusion potentialt|/^ arises entirely from the H , resp. OH ion concentration difference, one derives with (&„+) • (aoH~^ = const. 464 Paul Ruetschi l|)d = (RT/F) ln(aH+)1/(aH+)2 At 25° G the diffusion potential amounts thus to 59.16 mV per unit of pH difference. In the light of the foregoing, it appeared of interest to simulate corrosion film behaviour on Pb with artificially prepared precipitation membranes. The two-compartment cell used in the present study con sisted of 2 precisely machined Teflon blocks, clamped to gether by 8 threaded bolts. The electrolyte cavity in each block had a volume of 50 CUK; the horizontal connecting bore had a diameter of 15' mm. The cellophane membrane, DU PONT 193 PUDO, (thickness 12/urn) was sealed tightly be tween the two compartments by means of a Neoprene O-Ring. The precipitation membranes were formed by filling com partment I with 0,1 M HpS0., compartment II with 0,1 M Ba(C10J2 for BaSO, membraries, and with 0,1 M Pb(C10.) for PbSO, membranes. A minimum time of 24 h was allowed to complete the precipitation in the membrane. Then, the forming solutions were carefully removed by suction through evacuation pipes and the solutions to be studied were intro duced into the cell. For measurements on a BaSO. membrane, compartment I was filled with 0,1 M H?S0., equipped with a Hg/Hg^SO. reference, while compartment Ii was filled suc cessively with.more and more diluted solutions of HpS0A, and finally with increasingly concentrated solutions or Ba(0H)2, and was equipped with a H /H? reference electrode. The reference electrodes were previously equilibrated during 24 h with the corresponding solutions. The measured potential difference between the two elec trodes remained invariant at -0,735 V + 5 mV, irrespective of the pH in compartment II. Theoretically, the H+/Hp potential should have decreased by 59.16 mV per pH unit. + The pH-dependence of the H /H? electrode was therefore ex actly compensated by the pH dependence of the diffusion potential in the membrane,, the latter following the above equation 'over a range of 12 pH units. Similar results were obtained with PbSC, membranes and utilising PbO-saturated Ba(OK)p solutions. Paul Ruetschi 465 In order to carry the analogy to a corroding Pb electrode as far.as possible, it was tried to measure directly the potential difference of the following system: Ba(OH)p sat. PbSO._ H SO , Pb/PbO + PbO sat. BaSO^ Ofl ft -Hg SO./Hg pH 13.2 membrane pHl ' ^ 4 Indeed, the measured voltage of the above system, -0.49 V, was very close to the value of the characteristic depas- sivation plateau at -0.50 V observed when an anodized Pb electrode is left on open circuit in 0,1 M HLS0.. Theoreti cally, one expects for the above system, by considering the diffusion potential (13,2-1,0) x 59.16 mV, and by using the value of 0.676 V for Hg/Hg^SO. in 0.1 M H^SO. vs. STP Hp at pH 0 a value of -0,48 V. ^ 4 d 4 It is concluded that the assignment of voltage plateaus, observed during interrupted anodization or cathodic stripping, and the assignment of current peaks in potentio- dynamic studies, requires the consideration of possible diffusion potentials in the corrosion films. 466 Alvin Salkind CONTINUED DEVELOPMENT OF HYBRID POWER CELLS FOR CARDIAC PACEMAKERS Alvin Salkind, Allen Hahn, and John Cassel and Victor Satinsky ESB Incorporated, C. F. Norberg Technology Center, Yardley, Penna., USA, University of Missouri, Columbia, Missouri, USA, and The Hahnemann Medical College of Philadelphia, Penna., USA. Extensive investigation on materials for hybrid power cells as a long term energy source for cardiac pacemakers has been continuing using testing procedures previously described. Platinum black on tantalum screen has been examined as cathodes with aluminum and aluminum-amalgam (galvalum), and magnesium anodes. Over-voltage curves have been determined for all materials as well as long term in vitro stability. Cells have also been implanted in dogs intra muscularly (latissimus dorsi) with a four transistor multivibrator, the period of which varies as a function of input voltage. We have found that a platinum black-aluminum cell will produce a sufficient amount of energy to power a cardiac pacemaker for ex tended periods of time and maintain a relatively stable voltage (approximately 0.85 volts at 30 ]ia with a 25 cm^ cathode). Power cells implanted for as long as one year have produced little or no tissue or generalized reaction in dogs. The cathode material has proven to be remarkably stable in its physical characteristics. Accumulation of fluid usually occurs during the first week after implantation but is resorbed after that time. Kuranobu Sugita 467 INFLUENCES OF FOREIGN METAL IONS ON THE PERFORMANCES OF NICKEL POSITIVE ELECTRODES Kuranobu Sugita and Siro Ohkuma Railway Technical Research Institute, Tokyo, Japan 1 Adsorption of C-o 2+ orFe 2+ on nickel electrodes Test electrodes were prepared by the usual method in which active nickelhydroxides are deposited into the micropores of nickel plaques made by loose sintering. The electrode measured about 60mmx 80mmx 0.7mm. Four test cells were made each with one positive plate,two 'sintered type pure cadmium plates having a larger capacity, using glass rod spacers and '400cc of 7NK0H. The .capacity of the cells was ascertain ed under three cycles of charge and discharge. Then, the positives were washed in a discharged state thoroughly with flushing water and dried at 105°C.;Two cc of a 50$ KOH solution containing lOg/l of Co(N05)?-6H20(sol. A) or a 50$ KOH solution containing lOg/l of FeSÖ4*7H20(sol.B) was dropped through a pipet on each of the two electrodes. The 2cc corresponds to the volume for saturating the electrode. Two.other electrodes were separately dipped into sol.A or sol. B for 5 minutes at 90°C. After dried, 4 test cells were reassembled as described before. The variations of the cell capacities were measured as shown in Fig.l. The ,. „ capacities of the electrodes impregnated" with 2cc of Co - .or Fe2+-splution (curve 2 and curve 3 in Fig.l) changed a little but the increase of capacity in curve 1 and the decrease of capacity in curve 4 were more than 10$. Thus, the electrode .1 or 4 in Fig.l was considered to have a larger amount of Co or Fe than the electrode 2 or 3. The reporters, studied earlierthe adsorption-of Na on a sinter ed nickel electrode using 22^a ana observed that Na+ had been -adsorbed-by the electrode in a quantity which is more than 1.2 - 1»Impregnation of electrode with sol.A, 5min, 90°C. Ä < 2,2cc of the same sol. is dropped on the i»l .0 -p electrode, •H O 3,2cc of sol.B is cd PM • dropped on the elect cö o rode . 0 o 4,Impregnation of 13 5 7 electrode with, the Number of cycles same sol. 5min.,90 C Fig. 1 Change of positive .eletrode 2+ capacity with adsorption of Co or Fe 2+ Kuranobu Sugita necessary to saturate it.' From these results, it is concluded that nickel oxide electrodes are apt to adsorb foreign metal ions. The capacity of the electrode N0.1 in Fig.l increased gradually with repeated cycles after the Go-impregnation, while on the other hand the capacity of 1\T0.4 diminished immediately after the Fe-impregnation. The results seem to suggest that the reaction mechanism of Co is different from that of Fe. It is known well that the capacity of a nickel electrode is increased by addition of cobalt. Some workers reported about the formation of a solid solution between cobalt and nickel. The curve 1 in Fig.l seems to be a proof of this. It is well known too that the capacity of a nickel electrode is decreased by iron contamination. A sintered nickel electrode containing Z$ of f j'nely dispersed iron in its active materials lowers the charging overpotential of the electrode by lOmv below that * of an electrode without any iron. The. lowering of the charging overpotential is ..appa rently one of the reasons why the electrode capacity is decreased through iron contamination, although any other cause as hinted by Tichenor might be hidden. 2 Effect's of Co, Fe and Li on the durability of the sintered nickel electrode In durability tests, full care was taken to use sintered plaques having a good mechanical strength so as not to cause degradation or blistering of the positive.electrode. The mechanical strength was measured by a Martens scratch hardness tester with an aim at about half the- hardness of pure lead plate. The dimensions of test positive electrodes were much the same as in the previous experiment but their thickness was varied between 0.7mm and 2.Qmm. The cell assembly was also the same as in the previous experiment and was so designed that the cell capacity might be limited 2.5 1' 1 ,Cd + C, none of 1 Li+ ^2.0 1',Cd + C, LiOH 2.5 •PoI cd g/1 PM 4L ^ ocd 2 ,Cd + Fe, none of Li+ 1.5 2',Cd + Fe, LiOH 2.5 i 500 1000 1500 2000 g/1 Number of cycles Fig.2 Durability of sintered nickel electrodes coupled with various pocket negative plates 469 Kuranobu Sugita 1.25 1, Co 10$, LiOH 2.5 g/1 2, Ni(0H)2, none of 31.00 Li+ 2'.,Ni(0H)2, LiOH ->P5 2.5 g/1 •H 3, Fe 1%, LiOH 2.5 8,0.75 g/1 oCo Fig.3 Effects of 0.50 Co, Pe and Li+ on the durability of 1 200 400 600 800 sintered nickel 1000 electrodes , Number of cycles by the positive-throughout the tests. 2.1 Durability of sintered nickel electrodes coupled with various pocket negative plates. Pure nickel net was used for the positive plaque. Pocket negative plates were classified in two groups. The active material of the first group was cadmium plus carbon powder and that of the second, cadmium plus iron powder. A durability test was carried out in the absence or in the presence of LiOH, 1 g(2.5g/l). The cycl ing conditions were ; charging, 0.9A, 2h ; discharging, 1.2 A, lh ; temperature of air bath, 30° + 2°C. The results are shown in Fig. 2. The durability of group 2 in Fig.2 is much inferior to that of group 1. It is probably due to iron migrating from the negative active materials. LiOH is effec tive in all the groups. 2.2 Effect of foreign metal ions on the durability of sintered nickel electrodes. The following two experiments were executed using thinner plates than the.ones in experiment 2.1 and carried out under the following cycling condition : charging 0.9A, lh ; discharging, 0.6A lh ; temperature of air bath, 30 °± 2°C. The opposite electrodes were all sintered pure cadmium plates. (l) Nickel plated iron net for the positive plaque. The durability .of the electrodes with and without LiOH was stable during 2000 cycles showing a slight increase of capacity due to the addition of LiOH. (2) Pure nickel net for the positive plaque. Durability test on three kinds of positive electrodes was carried out with the addition of LiOH. A nickelhydroxide electrode without LiOH was compared, too. The results are shown in Fig. 3- Surprisingly, the durability of the electrode NO.2'in Fig.3 was the shortest of all. DTA--work, arid active oxygen measurement on NO.2'will testify to the fact that the active material of N0.2'became inactive after 800 cycles. 1, Sugita and Ohkuma ; Denki Kagaku,. 38 (12) ,900(1970) 470 C.J. Warde DESIGN OF HIGK-TEMPERATURE SOLID-ELECTROLYTE FUEL-CELL BATTERIES FOR MAXIMUM POWER OUTPUT PER UNIT VOLUME E.F. Sverdrup, C.J. Warde and R.L. Eback Westinghouse Research Laboratories, Pittsburgh, Pennsylvania, U.S.A. In a design of an integrated fuel-cell power system, conceived by Westinghouse under contract to the Office of Coal Research, it is important to maximize the power output per unit volume of fuel-cell batteries, because of heat-release and material-conservation con siderations. A battery-design study has been completed, in which the optimum film lengths on a. tubular substrate have been calculated for a range of values of electrolyte and interconnection resistivity- thickness products and electrode resistivity-thickness quotients. The analysis shows that power densities of greater than 5 kilowatts per cubic foot of battery volume would be possible, assuming that the batteries operate at 80% electrical efficiency and are constructed on tubes, one-half inch in diameter, which are spaced in the manner of shell and tube heat exchangers. ) t t t ( dg)i * gi + z.fC^gj} + z.tf. 6{—^} + z.f'^^l -0 3 CD 3 >-TO «< 3 O O» rt 3 rt(A SViWCitlArtS O O O. 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