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Pit initiation on passivated metal surfaces: crystallographic orientation and environmental effects Guo, Ruijin 1993

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PIT INITIATION ON PASSIVATED METAL SURFACES - Crystallographic Orientation and Environmental Effects by RUIJIN GUO B.Sc., Dalian University of Technology, China, 1982 M.Sc., Dalian University of Technology, China, 1984  A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in THE FACULTY OF GRADUATE STUDIES (Department of Metals and Materials Engineering)  We accept this thesis as conforming to the required standard  THE UNIVERSITY OF BRITISH COLUMBIA September 1993 © Ruijin Guo, 1993  In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the head of my department or by his or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.  (Signature)  Department of  meMc a—t fitated04  The University of British Columbia Vancouver, Canada  Date  DE-6 (2/88)  ^-e/Y`t/ler^  t  cr  7„  f.ApAy  11  ABSTRACT  The susceptibility to pitting corrosion of tin and zinc was found to be dependent on the crystallographic orientation of the surface in chloride solutions. For tin single crystals, the (111) face showed the lowest critical pitting potential among five differently oriented surfaces. The critical pitting potential of zinc single crystals decreased with the surface orientation in the following order: (1010) > (1120) > (0001). The critical pitting potentials for zinc and aluminum were found to be lower in chloride solutions than in bromide solutions.  The effects of the solution pH, the presence of buffers, and halide ion species on the pitting of polycrystalline nickel were investigated in 1.0 M halide solutions. It was found that the critical pitting potential (E cp ) was independent of the pH of unbuffered solutions in the pH region of 4.5 to 10.5. However, E cp was greatly increased at pH 12.5, and pitting corrosion was totally retarded at pH 14 in 1.0 M NaC1 solution. The addition of Na2 CO 3/NaHCO 3 or Na3PO 4/Na2HPO 4 buffer to the pH 10.5 solution raised the critical pitting potential of nickel. Furthermore, Ecp was lower in the chloride solution than in the bromide solution.  Using published thermodynamic data for halide complexes, potential-pH (E-pH) and halide concentration - pH (X-pH) diagrams were constructed for H 2 0-halide-metal (Sn, Zn, Al and Ni) systems. With increasing halide concentration, the formation of halide complexes is thermodynamically favoured and the zone of passivity is diminished. These new diagrams are unique and useful in the understanding of pit initiation.  A pitting theory is proposed, which emphasizes the change in local solution chemistry and the formation of halide complexes during the pit initiation process. Pits will be nucleated  111 when conditions are met for the formation of stable halide complexes in the local region. This theory successfully accounts for the effects of the solution pH, buffers, halide ions and the crystallographic orientations of single crystals on pitting behavior. It may also explain other aspects such as the effects of temperature, alloying elements, and fluid flow on the pitting corrosion of metals.  iv  Table of Contents ABSTRACT ^  ii  LIST OF TABLES ^  vii  LIST OF FIGURES ^  viii  LIST OF SYMBOLS AND ABBREVIATIONS ^  xiii  ACKNOWLEDGEMENTS ^  xv  1 INTRODUCTION ^  1  2 LITERATURE REVIEW ^ 2 2.1 Passivity ^ 2 2.1.1 Thermodynamics and Kinetics of Passivation ^ 2 2.1.2 Properties of the Passive Film on Metals ^ 6 2.2 Pitting Corrosion - Localized Film Breakdown ^ 10 2.2.1 Pitting Corrosion Phenomena and Criteria ^ 11 2.2.2 Effect of Halide Environments ^ 13 2.2.3 Effect of pH and Buffers ^ 16 2.2.4 Effect of Temperature ^ 17 2.2.5 Effect of Solution Flow ^ 18 2.2.6 Pitting of Metal Single Crystals ^ 18 2.2.7 Pit Morphology ^ 20 2.2.8 Alloying Element Effects ^ 21 2.2.9 Inclusions, Grain Boundaries and Dislocations ^ 22 2.2.10 Solution Chemistry inside Pits ^ 24 2.2.11 Application of E-pH Diagrams to Pitting Corrosion ^ 25 2.2.12 Theories of Pit Initiation ^ 28 2.3 Remaining Problems ^ 32 2.3.1 The Pitting Dependence on Crystallographic Orientations ^ 32 2.3.2 Solution pH and Local Solution pH Control ^ 33 2.3.3 Halide Complexes and Diagrams Dealing with Halide Complexes ^ 33 2.3.4 Pit Initiation Mechanisms ^ 34 3 OBJECTIVE ^  35  V 4 EXPERIMENTAL ^ 4.1 Materials ^ 4.2 Single Crystal Specimen Preparation ^ 4.3 Polycrystalline Nickel Specimen Preparation ^ 4.4 Test Temperature ^ 4.5 Test Solutions ^ 4.6 Electrochemical Test Cell and System ^ 4.7 Pitting Scan Test Technique ^ 4.8 Test Procedure ^ 4.9 Pit Morphology Examination ^  37 37 38 42 42 42 45 48 48 49  5 RESULTS ^ 50 5.1 Pitting Corrosion Behavior of Sn, Zn and Al Single Crystals ^ 50 5.1.1 Pitting Potential Dependence on Crystallographic Orientations ^ 51 5.1.2 Pit Morphology ^ 59 5.1.3 Effect of pH Buffer and Halides on Pitting of Zn ^ 75 5.1.4 Effect of Halides on Pitting of Al Single Crystals ^ 77 5.1.5 Overall Summary of Corrosion Behavior of Single Crystals ^ 85 5.2 Pitting Corrosion Behavior of Polycrystalline Nickel ^ 86 5.2.1 Polarization Behavior in Nitrate and Sulfate Solutions ^ 86 5.2.2 pH Effect on Pitting of Nickel in 0.1 M NaC1 ^ 89 5.2.3 Effect of Nitrate Inhibitors in Chloride Solutions ^ 90 5.2.4 Effect of Buffers in Chloride Solutions ^ 98 5.2.5 Comparison of Inhibitor with Buffers ^ 104 5.2.6 Halide Ion Effect ^ 107 5.2.7 Summary ^ 113 6 E-pH and X-pH DIAGRAMS ^ 6.1 Construction of E-pH and X-pH Diagrams ^ 6.1.1 Chemical Equilibrium ^ 6.1.2 Electrochemical Equilibrium ^ 6.1.3 E-pH Diagrams ^  114 114 114 115 116 6.1.4 X-pH Diagrams ^ 118 6.1.5 Thermodynamic Data and Assumption of Activity of Species ^ 120  vi 6.2 Diagrams for H2 0-Metal and H 2 0-Halide-Metal Oxide Systems ^ 121 6.2.1 Sn ^ 122 6.2.2 Zn ^ 126 6.2.3 Al ^ 132 6.2.4 Ni ^ 138 7 DISCUSSION ^ 144 7.1 Pit Initiation Stages and Governing Factors ^ 144 7.2 Electrode Kinetics ^ 149 7.3 Dynamic Nature of the Passive Film ^ 150 7.4 Local Solution Chemistry during the Film Breakdown/Reformation Process ^ 151 7.5 Halide Complex Formation and Pit Initiation Theory ^ 167 7.6 Evidence from Experiments on Environmental Effects ^ 172 7.7 An Explanation of the Pitting Dependence on Crystallographic Orientations ^ 179 7.8 Critical Pitting Potential ^ 182 7.9 Extension of the Proposed Theory to Other Aspects of Pitting ^ 184 8 CONCLUSIONS ^  187  9 REFERENCES ^  189  APPENDIX I ^  197  APPENDIX II ^  199  APPENDIX III ^  202  APPENDIX IV ^  204  vii  List of Tables Table 1. Aggressive ions in pitting corrosion of metals ^ 10 Table 2. Minimum halide concentration for pitting on various metals ^ 15 Table 3. Values of slope B in Equation (2.1) for various metals ^ 16 Table 4. Structure, melting points and growing orientations of single crystals ^ 37 Table 5. Electrochemical polishing parameters for Zn, Sn and Al ^ 41 Table 6. Test solutions for pitting corrosion of single crystals ^ 43 Table 7. Test solutions for pitting corrosion of nickel ^ 44 Table 8. Experimental summary for single crystals ^  50  Table 9. Critical pitting potentials obtained from five oriented faces of tin ^ 51 Table 10. Measured pit wall angles from the pits formed on the tin (001) surface ^  59  Table 11. Summary of corrosion behavior of single crystals ^ 85 Table 12. Critical pitting potentials and AE, p , for polycrystalline nickel in chloride solutions at pH 10.5 ^  105  Table 13. Standard chemical potentials of substances at 25 °C ^ 121 Table 14. Standard chemical potentials for Sn systems at 25 °C ^ 123 Table 15. Standard chemical potentials of substances for Zn systems at 25 °C ^ 127 Table 16. Standard chemical potentials of substances for Al systems at 25 °C ^ 133 Table 17. Standard chemical potentials of substances for Ni systems at 25 °C ^ 138 Table 18. Minimum halide concentration for the complex formation ^ 174  viii  List of Figures Figure 1 A Pourbaix E-pH diagram for Al - H 2 O system with corrosion, passivation and immunity zones ^  3  Figure 2 A typical polarization curve for the passive metal in a non-aggressive solution ^ 5 Figure 3 A typical cyclic polarization curve for the pitting corrosion of the passive metal in an aggressive solution (solid line) ^  12  Figure 4 An experimental E-pH diagram for pitting corrosion of iron in chloride solution [1111 ^  27  Figure 5 A X-ray back reflection Laue pattern from the tin (001) surface ^ 39 Figure 6 A schematic diagram of the specimen used in pitting tests ^ 40 Figure 7 A schematic diagram of the cell used in pitting tests ^ 46 Figure 8 A schematic diagram of the corrosion measurement system used in pitting tests ^  47  Figure 9 Potentiodynamic polarization results obtained on tin (011) and (111) faces in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 6.0 ^ 54 Figure 10 Variation of the critical pitting potential with crystallographic orientation of tin in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 6.0 ^ 55 Figure 11 Potentiodynamic polarization curves obtained on three differently oriented surfaces of zinc 0.1 M NaC1 at pH 9.2 ^  56  Figure 12 Variation of the critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 at pH 9.2 ^  57  Figure 13 Variation of critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 9.2 ^ 58 Figure 14 Pits formed on the tin (001) face (a, b, and c), and crystallographic facets (pit walls) in the pit (d) ^ Figure 15 A light interference photograph taken on the pitted (001) face of tin Figure 16 A schematic view of the pit on the tin (001) face and the determination of the pit wall angle ^  60 61 62  ix Figure 17 Relationship between the apparent crystallographic angle and the steps formed on pit walls ^  66  Figure 18 Pit morphology of the tin (111) face (a, b); crystallographic facets in the pit (c) ^  67  Figure 19 Pit morphology of the tin (011) face (a, b); crystallographic facets in the pit (c) ^  68  Figure 20 Pit morphology of the tin (110) face (a); crystallographic facets in the pit (b) ^  69  Figure 21 Pit morphology of the tin (100) face ^  70  Figure 22 Pit morphology of the zinc (0001) face ^  71  Figure 23 Pit morphology of the zinc (1010) face ^  72  Figure 24 Pit morphology of the zinc (1120) face ^  73  Figure 25 Pits formed on the cleaved basal plane (0001) of zinc ^ 74 Figure 26 Polarization curves of the zinc (0001) face in 0.1 M unbuffered and Na2 CO 3/NaHCO 3 buffered halide solutions at pH 9.2 ^ 76 Figure 27 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 3.4 ^  79  Figure 28 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 6.0 ^  80  Figure 29 Exposed (111) faces on the Al (100) surface corroded in 0.1 M HF at pH 3.4 ^ 81 Figure 30 A crystallographic pit formed on the Al (100) in 0.1 M NaC1 at pH 6.0 ^ 82 Figure 31 A thick salt film formed on the Al (100) in 0.1 M NaF at pH 6.0 ^ 83 Figure 32 Salt crystals formed on the surface of Al in 0.1 M NaF at pH 6.0 ^ 84 Figure 33 Potentiodynamic polarization curves of Ni in 1.0 M NaNO 3 , 1.0 M Na2 SO 4 and 1.0 M NaC1 at pH = 10.5 ^ 87 Figure 34 Potentiodynamic polarization curves of Ni in 1.0 M Na2 SO4 and 1.0  M NaC1 at pH = 2.5 ^  88  Figure 35 A potentiodynamic polarization result for Ni in 1.0 M NaCI at pH = 2.5 ^ 91  Figure 36 Potentiodynamic polarization results for Ni in 1.0 M NaC1 at pH = 4.5, 6.5, 8.5 and 10.5 ^  92  Figure 37 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 12.5 ^ 93 Figure 38 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 14.0 ^ 94 Figure 39 Effect of bulk solution pH on the critical pitting potential of Ni in unbuffered 1.0 M NaC1 solution ^  95  Figure 40 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaNO3 at pH = 10.5 ^  96  Figure 41 Effect of nitrate concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5 ^  97  Figure 42 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaHCO 3/Na2 CO3 buffer at pH = 10.5 ^ 99 Figure 43 Effect of carbonate buffer concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5 ^  100  Figure 44 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M Na2 HPO4/Na3 PO4 buffer at pH = 10.5 ^ 101 Figure 45 Effect of phosphate buffer concentration on the critical pitting potential of Ni in 1.0 M NaCl at pH = 10.5. ^  102  Figure 46 Potentiodynamic polarization results for Ni in unbuffered 1.0 M NaBr and Na2 HPO 4/Na3 PO4 buffered 1.0 M NaBr at pH = 10.5 ^ 103 Figure 47 Comparison of the increment in the critical pitting potential (AEc p) with the presence of NaNO 3 , NaHCO 3/Na2 CO 3 and Na2 HPO4/Na3 PO4 in 1.0 M NaC1 at pH 10.5 ^ 106 Figure 48 Polarization results for Ni in 1.0 M NaC1 and 1.0 M NaBr at pH = 10.5 ^  108  Figure 49 Polarization results for Ni in Na 2 HPO4/Na3 PO4 buffered 1.0 M NaC1 and 1.0 M NaBr at pH 10.5 ^ 109 Figure 50 Polarization curves for Ni in 1.0 M NaF and 1.0 M NaC1 at pH 10.5 ^ 110 Figure 51 Polarization curves for Ni in 1.0 M NaF at pH 6.0 and 1.0 M NaC1 at pH 4.5 ^  111  Figure 52 A polarization result for Ni in 1.0 M HF at pH 3.1 ^ 112  xi Figure 53 E-pH diagram for H 2 O - Sn system; activities of all solute species at 10 2 , 10 -4 , 10 ^ 124 Figure 54 E-pH diagram for H 2 O - C1 - Sn system; activities of chloride at 10 ° and 10 1 ; activity of other solute species at 10 -6 ^ 125 Figure 55 E-pH diagram for H 2O - Zn system; activities of all solute species at 1072 , 10 -4 , 10 ^ 128 Figure 56 E-pH diagram for H 2 O - CO 3 2- - Zn system; Activity of H2 CO3/HCO 3 -/CO 3 2- at 10-1 ; activities of other solute species at 10 -2 , 10-4 , 10 .6 ^ 129 Figure 57 E-pH diagram for H 2 O - C1 - Zn system; activities of chloride at 10" .3 and 10 1 ; activity of other solute species at 10 -6 ^  130  Figure 58 X-pH diagram for H 2 O - Cl - - Zn(OH) 2 system; activity of all solute species at 10', 10 -4 , 10 -6 ^ 131 Figure 59 E-pH diagram for H 2 O - Al system; activities of all solute species at 10', 10 -4 and 10 -6 ^ 134 Figure 60 E-pH diagram for H 2 O - Cr - Al system; activities of chloride at 10 ° and 10 1 ; activity of other solute species at 10 -6 ^ 135 Figure 61 X-pH diagram for H 2 O -F - Al 2 0 3 .3H2 0 system; activity of all solute species at 10', 10 -4 , 10 -6 ^  136  Figure 62 X-pH diasram for H 2 O -Cr - Al2 0 3 .3H 2 0 system; activity of all solute species at 10', 10 -4 , 10-6 ^  137  Figure 63 E-pH diagram for H 2 O - Ni system; activities of all solute species at 10 - , 10 -4 and 10 -6 ^ 140 Figure 64 E-pH diagram for H 2 O - Cr - Ni system; activities of chloride at 10 ° and 10 1 ; activity of other solute species at 10 -6 ^ 141 Figure 65 X-pH diagram for H 2 O -F - NiO system; activities of all solute species at 10', 10 -4 , 10-6 ^  142  Figure 66 X-pH diagram for H 2 O -C1 - - NiO system; activities of all solute species at 10 - `, 10-4 , 10 .6 ^  143  Figure 67 Several stages occurring during the pit initiation process ^ 145 Figure 68 Structure of the interface between metal substrate and bulk solution ^ 147 Figure 69 An active site with a radius of r during oxide film breakdown. ^ 153 Figure 70 Change in local pH with variation of r and i ^ 157  xii Figure 71 Change in local halide concentration with variation of r and i at a given bulk halide concentration of 0.01 M ^  159  Figure 72 Variation of local halide concentration with bulk halide concentration at a given anodic current density of 1 A/cm 2 ^  160  Figure 73 Local solution chemistry profile due to a diffusion process across the diffusion layer ^  162  Figure 74 Effect of i on the local pH for Ni, Zn, Sn and Al in a continuous film breakdown process at 8 = 10' cm ^  165  Figure 75 Variation of local halide concentration with bulk halide concentration in a continuous film breakdown process at 8 = 10 -2 cm ^ 166 Figure 76 A generalized X-pH diagram ^  168  Figure 77 A schematic diagram showing processes in pit initiation ^ 170 Figure 78 Change of the local solution pH with the bulk solution pH calculated for Ni using an one-dimensional diffusion model ^ 177  LIST OF SYMBOLS AND ABBREVIATIONS a^  electron transfer coefficient  3, Y^reaction orders  5^  diffusion layer thickness electrode overpotential wavelength of thallium light standard chemical potential  p^  density  D  diffusion coefficient  E  electrode potential  E°^standard electrode potential Epp^critical pitting potential Er^repassivation potential EDX^  energy dispersive X-ray analysis  F^  Faraday constant  AG°^  standard free energy change  h^  thickness of the oxide film  i^  current density  io^  exchange current density  i1^limiting current density  ip^passive current density  J^  diffusion flux  K  equilibrium constant  Km (m = 1,2, ...)^ stability constant  xiv  M^ M  molarity (mole/liter)  ^subscript for metals  ox^  subscript for oxides and hydrated oxides  Q^  electric charge  r^  radius of an active site  R^  reaction rate  SCE^  saturated calomel electrode  SHE^  standard hydrogen electrode  SEM^  scanning electron microscopy  State:  V  aq^ g^ 1^  ^  w^  aqueous gaseous liquid solid volume molecular weight  [xLin^  minimum halide concentration for the formation of dominant soluble complexes  z^  electric charge of species  XV  ACKNOWLEDGEMENTS  I would like to express my sincere thanks to my supervisors, Dr. Desmond Tromans and Dr. Fred Weinberg for their advice and guidance throughout this research. My thanks are also extended to other faculty and staff members, and fellow graduate students in the corrosion group for their help. Special thanks go to my wife and son for their support and understanding during this work. The author is also grateful for the financial support provided by the University Graduate Fellowships, and by the Research and Teaching Assistance of this Department.  INTRODUCTION^  1  1 INTRODUCTION  The passivity of metals is important in commercial applications. Many corrosion resistant alloys, including stainless steels, Ni-base alloys, aluminum alloys and titanium alloys, have been developed to provide materials which have good corrosion resistance based on the formation of protective (passive) films. However, in some cases, these materials can exhibit severe localized corrosion which can sharply restrict their use under certain service conditions. Several types of localized corrosion behavior include: (1) pitting corrosion, (2) crevice corrosion (3) intergranular corrosion and (4) stress corrosion cracking and corrosion fatigue. Localized corrosion contributes to a large percentage of corrosion problems in industrial situations. According to one Japanese report, related to the chemical process industry [1] , the corrosion they observed can be divided into the following categories: (1) general corrosion 15%, (2) stress-induced corrosion cracking 39%, (3) pitting corrosion 8%, (4) other localized corrosion 19%, (5) high temperature corrosion 6% and (6) others 13%. LaQue  [2]  , on the basis of his 44  years of experience, considered that localized corrosion was responsible for about 90% of metals which failed by corrosion. Extensive studies have been carried out on passivity, breakdown of passive films and localized corrosion phenomena in efforts to find better alloy/environment combinations to reduce the frequency of localized corrosion. Many investigations in the past 50 years  E3H5I  have  been undertaken to examine the fundamentals of passivity and its localized breakdown using chemical, electrochemical and physical methods. The reported results are very extensive, too large to be referred to comprehensively in this literature survey. The present review is confined to pit initiation on passivated metal surfaces, caused by localized electrochemical breakdown of passive films.  LITERATURE REVIEW  ^  2  2 LITERATURE REVIEW  2.1 Passivity Passivity of a metal surface was initially found on iron in concentrated nitric acid in 1790, according to Uhlig 161 . The passive oxide film was first isolated from the surface of iron, which was exposed in air or was passivated in chromate solution, by Evans in 1927 171  .  In the past 50 years the passive film has been studied extensively in-situ and ex-situ by electrochemical and physical techniques such as electrochemical reduction impedance techniques [101-112 1 optical elliopsometry ,  11314141  ,  [81 [91 ,  AC  and the surface analysis  techniques of Auger Electron Spectroscopy (AES), X-ray Photoelectron Spectroscopy (XPS) and Secondary Ion Mass Spectroscopy (SIMS)  [1514191 .  The thermodynamics and kinetics of  passivation have become more clearly understood, and many passive film properties have become known.  2.1.1 Thermodynamics and Kinetics of Passivation The thermodynamic approach to passivity is mainly the result of contributions by Pourbaix 1201 who proposed that a necessary condition for passivation is the formation of ,  a stable oxide, or hydroxide film. The most useful guide to passivation is via Pourbaix's potential-pH diagrams (E-pH diagrams) for metal-water systems calculated from aqueous thermodynamic equilibrium data. An E-pH diagram shows the stable regions for individual species, so that corrosion (stable soluble species), immunity (stable metal) and passivation (stable oxides) zones can be identified (Figure 1). However, E-pH diagrams  Al - Water system  1  0.5  0  -0.5  -1  -1.5  -2  -2.5  -3  0  ^ ^ ^ ^ ^ ^ 12 14 2 4 6^8 10  pH  Figure 1 A Pourbaix E-pH diagram for Al - H 2 O system with corrosion, passivation and immunity zones  LITERATURE REVIEW^  4  alone are only a guide to possible conditions for passivation. Other factors may be equally important, such as the presence of imperfections and defects in the film, and the adherence of the film to the metal. Passivation is better defined from a consideration of the corrosion kinetics. A metal is passive if it substantially resists corrosion in a given environment despite a marked thermodynamic tendency to react  16j .  The passivation phenomenon can be  described kinetically by an anodic polarization curve (Figure 2) obtained on a metal in an aqueous solution without aggressive species (such as halides ions). Three different potential regions are identified as active, passive and transpassive regions on the polarization curve. In the active region metal undergoes general corrosion until the passivation potential (E r ) is reached. After this point, the anodic current density markedly drops to an extremely low value (passive current density, ip , in the order of gA/cm2 ) in the passive region. It is in the passive region where many metals and alloys are designed to work to achieve their best corrosion resistance in aqueous environments. The passivation upper limit is determined by the onset of transpassive dissolution, involving the loss of the protective nature of the film by further oxidizing it into soluble species with a higher oxidation state  [21], [22].  Sometimes, if the transpassive potential is  higher than the potential for the decomposition of water, the transpassive behavior is marked by large anodic current increases corresponding to the evolution of oxygen.  LITERATURE REVIEW  5  Transpassive -------------------------  iP Current Density (i)  Figure 2 A typical polarization curve for the passive metal in a non-aggressive solution  LITERATURE REVIEW^  6  2.1.2 Properties of the Passive Film on Metals It is well known that a thin passive layer is responsible for the passivity of metals, but there has been a debate as to whether the passive layer is an adsorption layer or a solid-phase oxide layer (or hydrated oxide layer). The adsorption layer model  [6]  assumes that adsorbed oxygen atoms form a stable  two-dimensional structure of mixed 0 2- anions with metallic ions on the surface of a metal. At the optimum ratio of metal to oxygen ions, Uhlig  [23]  suggested that the  adsorption film is more stable than the 3-dimensional metal oxide film. Although such a mono-layer adsorption may lead to the subsequent growth of the oxide film, Uhlig proposed that passivity was very much dependent on the adsorption layer. However, the available experimental evidence suggests that the passive film is more likely to be a solid-phase oxide layer formed on the metal surface  [21], [24].  Strong  evidence has come from Evans' experiment which isolated an oxide film from iron m. Modern surface analysis techniques, such as AES, XPS and SIMS, have identified and measured the composition of the oxide film, its oxidation state and the thickness. There is now general agreement that the passive layer on the metal is composed of oxide layers with a thickness which varies from 1 to 10 nm  [25]-[33] .  (a). Composition and Thickness of the Passive Oxide Film: Kruger [25] and Cohen [9], [26], [27] have reported that the oxide film on the iron ,  surface has two oxide sub-layers, with an inner Fe 3 0 4 layer and a y-Fe 2 O 3 outer layer. The XPS results obtained by Strehblow [28] also confirmed this two-layer structure of the passive film on iron. Chromium enrichment has been found in passive films on stainless  LITERATURE REVIEW^  7  steels [15]-(18], an d Cr-rich mixed oxides are the main components of these films. The thickness of the passive film on iron and its alloys, such as stainless steels, varies from lnm to lOnm [281 The hydrated Ni (II) oxide was found on nickel passivated in both acid .  and alkaline solutions in the form of NiO•H 2 0 or NiOOH. The thickness of the oxide layer was found to be 2 - 8 nm  [291E301  . At more positive potentials, other types of oxides  with higher oxidation states, such as Ni 2 0 3 and Ni0 2 , were assumed to be present. The different modifications of Al(III) oxide and hydroxide are thermodynamically stable in neutral solutions [3114331 . Passive film oxides of Al 2 0 3 , Al 2 0 3 •H2 0 and Al 2 0 3 .3H2 0 have been reported on the surface of aluminum by many authors, the degree of hydration being dependent on the prevailing conditions. Under gaseous oxidation conditions, Al 2 0 3 is formed with a largely amorphous structure and a thickness up to 15 nm [311 Altenpohl and Post [331 reported that a film of Al 2 0 3 •H2 0 was formed above 75 .  °C, and Al 2 0 3 .3H2 0 below 75 °C, in double distilled water. The films consisted of two layers. Oxides were thick and porous in the outer layer of film, whereas the inner layer, 3nm thick, was non-porous and acted as a dielectric. The inner layer was considered to be the most protective. The oxide film on aluminum under anodizing conditions could grow up to 100 gm, however, only the inner layer, 3-5nm thick, was responsible for the passivation behavior. The outer layer, composed mostly of primary dissolutionprecipitation products, was non-protective [34]  .  Oxides films formed on zinc and tin have been less studied. In the near neutral pH range, ZnO and Zn(OH) 2 were reported and found to be relatively insoluble and protective [381 and Zn(OH) 2 is the most stable form according to Pourbaix ,  solutions, thick porous oxide or hydroxide films of zinc were formed  [361  ,  [201  .  In alkaline  which could  grow to a visible thickness. Little information appears to have been reported on the  LITERATURE REVIEW^  8  passive film of Sn. The film formed on a tin surface is complex because of the existence of two oxidation states, Sn(II) and Sn(IV). The Sn oxides, SnO and Sn0 2 , are thermodynamically more stable than their hydrated forms, Sn(OH) 2 and Sn(OH) 4 . (b) Crystalline and Non-crystalline Structure of Passive Films . Revesz and Kruger [371 suggested that non-crystalline passive films are more protective, less susceptible to breakdown by aggressive ions, exhibit greater repassivation rates, and are more ductile than crystalline films. Hoar  [381  also claimed  that the passivity of a non-crystalline film is superior to that of the crystalline film. The passive film formed on iron has been found to be crystalline, but the passive films on Fe-Cr alloys are reported to be amorphous [391 . It is believed that Cr in the Fe-Cr alloy promotes the formation of a non-crystalline film film formed on nickel is crystalline  [40], [41].  [391 .  It has been claimed that the oxide  The passive films formed on aluminum have  been reported to be non-crystalline [31], [421 Hydration of the oxide film (either as bound water or as Off bonds) facilitated the formation of the non-crystalline structure in Al and stainless steels [19],  ^Lawless [43] has reported the formation of crystalline oxide films  on Cu, Ni and Zn under low temperature gaseous oxidation conditions, but the structure of the oxide films might be changed in aqueous solutions at room temperature. (c) Conductivity of the Passive Film: The passive films formed on iron, chromium and stainless steels are considered by some authors to be n-type semiconductors  [21], [26],  having a higher electronic  conductivity. The higher electronic conductivity facilitates the kinetics of oxygen evolution at higher anodic potentials. However, the passive films formed on aluminum, zirconium and titanium have very low electronic conductivities  [211  , which hinders the  LITERATURE REVIEW^  9  oxygen evolution kinetics. Consequently, a high anodic potential results in the growth of oxide films on these metals. The passive oxide films on metals such as Zn and Sn have an electronic conductivity between the iron and aluminum groups  '  [443 1451  .  (d) Imperfections in the passive film: Imperfections exist in passive films. Grain boundaries, solute segregation and depletion, secondary phase particles and non-metallic inclusions in the metal substrate produce imperfections in the grown passive films 14614511 These imperfections are .  associated directly with the positions of the defects existing in the substrate, and are the main nucleation sites for localized corrosion of commercial alloys. Wood  [521  suggested  that defects in the passive films (film flaws) were produced when the oxide film formed over mechanical surface defects (such as scratches), and voids formed by vacancy coalescence in the substrate. Flaws may also occur in the growing film at substrate grain boundaries. According to Chao et al. [531 as-grown point defects are present in crystalline oxide films, and they suggest that these cation and anion vacancies in the passive oxide film play an important role in the growth and breakdown of the film.  LITERATURE REVIEW^  10  2.2 Pitting Corrosion - Localized Film Breakdown Pitting corrosion is a result of the localized breakdown of the passive film. Pitting corrosion occurs on the passivated metal surface in an environment containing aggressive ions such as Cl - , Br- , I- , C104 and SCN - . Of these, chloride ions are considered to be the most widely encountered and most aggressive anions. Table 1 lists the most common anion - metal systems where the localized breakdown of passivity occurs. Pitting corrosion is reviewed by Smialowska  [45]  in her comprehensive reference book titled "Pitting Corrosion  of Metals ".  Table 1. Aggressive ions in pitting corrosion of metals  Iron  C1, Br- , 1- ,C104  Nickel  Cl-, Br",r  Stainless Steels  Cl-, Bf, SCN  Aluminum  Cl-, Br , r, C104, SCN-  Titanium  Cl-, Br , r  Zinc  CL, Br , r, C104, Br0 3 -  Tin  Cl"  Cadmium  Cr, Br-  Zirconium  Cl-, Br , T  -  -  -  -  -  [451  LITERATURE REVIEW^  11  2.2.1 Pitting Corrosion Phenomena and Criteria It has been well recognized that a certain anodic electrode potential has to be reached in order for pits to initiate on a passive metal surface  '  [45] [54]  . A typical  single-cycle anodic polarization curve (solid line) in the aggressive solution is given schematically in Figure 3. There are two characteristic potentials in this figure: (1) the critical pitting potential (E cp ) and (2) the repassivation potential (E r). These two characteristic potentials divide the polarization curve into three regions. Above E cp , pits will be initiated and will grow; Between E cp and E 1 , existing pits will continue to grow, but no new pits will be nucleated; Below the repassivation potential E„ pits will be repassivated. As compared with the polarization curve in the non-aggressive solution (dashed line in Figure 3), the integrity of the passive film is destroyed by the localized breakdown in the aggressive solution, which gives rise to a sudden increase in anodic current and an anodic current hysteresis loop during the backward potential scan. The critical pitting potential, E cp , is considered to be a major parameter in the study of pit initiation. The critical pitting potential has been widely used in the following ways: 1. To define the conditions for the onset of pitting corrosion. The occurrence of pitting corrosion in a metal/environment system can be determined from the critical pitting potential. 2. To evaluate the pitting susceptibility of metals. The effects of environmental and metallurgical factors, such as temperature, pH, solution concentration, and alloy composition, can be determined from the measurement of the critical pitting potentials.  LITERATURE REVIEW^  12  ..  I  ..... ....  ..-  .0-  --- ....^Transpassive 00  .... ..-- .....  ..-  0-  Pit Initiation  Current Loop  E q.), Er Repassivation  Active  iP  Current Density (i)  Figure 3 A typical cyclic polarization curve for the pitting corrosion of the passive metal in an aggressive solution (solid line)  LITERATURE REVIEW^  13  3. To develop pit initiation mechanisms. An understanding of the factors governing the pit initiation process and critical pitting potential may lead to more effective development of pitting resistant alloy-environment systems. Besides the potential criterion, there are other useful criteria, such as the critical temperature for pitting [551,  ^and the critical (minimum) chloride concentration  1571  .  These criteria are less widely used in pitting studies.  2.2.2 Effect of Halide Environments Most metals are subject to pitting corrosion in halide solutions. There is a minimum concentration of halide for pit initiation of a given metal, below which pitting does not occur. These minimum halide concentrations are listed in Table 2. The aggressiveness of the halide species varies with different metals [451 For .  aluminum, iron, stainless steel and nickel, chloride ions are most aggressive, followed by bromide and iodide ions: Aggressiveness: Cl > Br > I -  -  However, for titanium and tantalum, bromide and iodide are more aggressive than chloride: Aggressiveness: Br > I > Cl-  -  LITERATURE REVIEW^  14  The critical pitting potential (E, p ) is dependent on the concentration of halides in the solution. An increase in the halide concentration will give rise to a decrease in the critical pitting potential. A relationship between E cp and halide concentration [X] has been established  [451E541  :  Ecp = A - Blog[X]^  (2.1)  Where A and B are experimental constants and [C] is the halide concentration (M). Values of slope B are listed in Table 3. They vary for different metals and electrolyte compositions.  LITERATURE REVIEW ^  15  Table 2. Minimum halide concentration for pitting on various metals * Metal  Halide Ion  Minimum Concentration, M  Iron  Cl- (a)  0.0003  Iron  Cr (b)  0.0005  Iron  CF (c)  0.003  Fe-5.6Cr  Cr (a)  0.017  Fe-11.6Cr  Cl- (a)  0.069  Fe-20Cr  Cr (a)  0.1  Fe-24.5Cr  Cr (a)  1.0  Fe-29.4Cr  Cr (a)  1.0  Fe-18.6Cr-9.9Ni  Cl" (a)  0.1  Nickel  CF (a)  0.001  Br  0.002  Titanium  -  (d)  * Source: Z. S. Smialowska, Pitting Corrosion of Metals, NACE Publication, 1986 (a)H 2 SO 4 + NaC1 solution (b)Phthalate buffer + NaC1, pH = 5 (c)Borate buffer + NaC1, pH = 8.4 (d)KBr solution  LITERATURE REVIEW^  16  Table 3. Values of slope B in Equation (2.1) for various metals * Metals  B, Volts  Iron  0.06 - 0.2  Iron-base alloys  0.04 - 0.068  Nickel and its alloys  0.071 - 0.078  Al-base alloys  0.05 - 0.13  Cadmium  0.03 - 0.18  Titanium  0.11  Zirconium  0.06 - 0.065  * Source: Z. S Smialowska, Pitting Corrosion of Metals, NACE Publication, 1986  2.2.3 Effect of pH and Buffers There are many studies dealing with the effect of bulk solution pH on pitting corrosion [6°]-4653 In general, the critical pitting potential is not affected, or hardly .  affected, in the pH range from acidic to slightly alkaline values. The critical pitting potential of iron was found to be independent of pH in chloride solution with pH 8-12.7, and in C104 solution of pH 0.6 to 7.7 [60] . Alvare and Gavele  [61]  found E cp was  unchanged on iron in chloride solutions of pH 7 - 10, but a higher critical pitting potential was obtained above pH 10. Similar results were reported for nickel, stainless steels and aluminum  [54], [62]-65] .  Buffers have been used by many authors in the investigation of pitting corrosion. The assumption that the buffer does not interfere with the pitting process is not valid. It  LITERATURE REVIEW^  17  was found that buffers affected the pitting corrosion behavior, and the presence of the buffer in the solution increased the critical pitting potential [66]  [661, [67]  Heusler and Fischer  reported that the critical pitting potential of iron increased with the concentration of  buffers in 0.01 M NaC1 solution. The increase in the critical pitting potential also depends on the types of buffers. The critical pitting potential was raised about 150 mV in borate buffer, but was raised only 30 mV in phthalate buffer, when the concentration of the buffers changed from 0.1 to 1.0 M in 0.05 M NaCl. Drogowska et al. [67] found that bicarbonate and phosphate buffers raised the critical pitting potentials of tin by a few hundred mV's in chloride solutions.  2.2.4 Effect of Temperature According to the Arrhenius rate equation, reactions proceed more rapidly at higher temperature. Consequently, it would be expected that pit initiation and pit growth occur more readily with increasing temperature. It was found that the critical pitting potential decreased with increase in temperature for stainless steels, Ni-base alloys, aluminum, and titanium [54], [68]-171] Toussek 1721 studied the pitting corrosion of 18Cr-10Ni stainless steel in 0.5 M NaC1, and found that the critical pitting potential (Ecp ) was a linear function of reciprocal absolute temperature (1/T), with a slope of 500 VK° in the temperature range of 6 - 40 C°. Above 40 C° the effect of temperature on E cp was less pronounced. Smialowska [68] reportedthat a linear function (E ci, = a - bT) existed between E cp and temperature (T) for ferritic and austenitic stainless steels with a slope of about 3 mV/°C, but the linear relation was no longer valid for Mo-containing stainless steel above 70 C°, where the critical pitting potential showed little change.  LITERATURE REVIEW^  18  2.2.5 Effect of Solution Flow The effect of solution flow on pitting corrosion has been investigated using a rotating electrode in an electrolytic cell. Powers and Wilfore [73] reported that the critical pitting potential of Ti-6A1-4V alloy was shifted from 1.8 V to 4.3 V (SCE) when the rotational speed was changed from 0 to 5000 rpm. Riskin and Turkovakaia  1741  found that  Ecr, for 18Cr - 8Ni stainless steel was shifted to a more positive value when the electrode was rotated at 3000 rpm in the chloride solution. Increases in E ci, were also reported for aluminum, where the critical pitting potential was raised about 50 mV when the rotational speed was changed from 0 to 2000 rpm  1753  . A 50 mV increase in the critical  pitting potential of duplex stainless steel was reported at 1590 rpm in chloride solutions [761  .  However, no effect of solution flow on the critical pitting potential was observed in  other investigations [77]  '  1781  .  2.2.6 Pitting of Metal Single Crystals Kruger [791 studied the effect of crystallographic orientation on the film breakdown tendency for single crystals of iron. His study showed that the resistance to pitting increases as the surface approaches the (100) orientation, with the lowest pitting resistance on the closest packed planes { 110}. It was also found that the pit density varied with crystallographic orientation, with the highest pit density on the {110} planes. This was in agreement with his finding that the {110} planes were most susceptible to pitting attack. A more detailed study of crystallographic orientation effects has been conducted by Yasuda et al. [W on aluminum single crystals in chloride-containing solutions, using  LITERATURE REVIEW^  19  potentiodynamic and galvanostatic methods. The results from potentiodynamic measurements showed that the critical pitting potential increased with the crystallographic orientation in the following order: { 111 } < {110} < { 1001. Based on galvanostatic testing results, the pit density and pit area were found to decrease with the crystallographic orientation in the pattern: { 111 } > { 1101 > { 100 }. This indicated that the close packed planes {111} had the lowest pitting resistance when evaluated either by the critical pitting potential or by pit density and pit area. However, the critical pitting potentials were reported to be independent of crystallographic orientations in Al-Cu alloy single crystals r 801 . The dependence of passivity on the crystal planes was reported for nickel single crystals by Latianision et al. [811 . The passive current densities on nickel single crystal surfaces decreased in 0.5 M H 2 SO 4 solution in the following order: { 100} > {110} > 11111, indicating that the passive film formed on the { 111 } faces was most protective. Garz et al.E 821 reported that the pitting resistance of nickel crystals is dependent on the crystallographic orientation of the tested surface. It has been found that the critical pitting potential in 0.5 M NiC12 solution increased in the order: {100} < {110} < {111} . Lei et al. [831 studied the breakdown of the passive film on single crystals and polycrystals of nickel in chloride solutions. Their study showed that the critical pitting potential of { 100 } faces was about 50mV higher than that of the polycrystalline nickel. The induction time for the pitting of Ni { 100} surfaces was reported to be slightly longer than that for a polycrystalline surface at the same potential. Their results suggest that Ni { 100} surfaces are more resistant to pitting corrosion. However, there are no data available on {110} and {111} oriented surfaces from their investigation to compare with { 100} orientations.  LITERATURE REVIEW^  20  It is worth noting that the pitting resistance of the {111}, { 110} and { 100} faces of Ni is in the reverse order from that reported for Al, although both metals have the F.C.C. crystal structure.  2.2.7 Pit Morphology Corrosion pits may be either crystallographic or non-crystallographic. A crystallographic pit exhibits pit walls, whose surface traces are consistent with the intersection of specific crystal planes with the specimen surface. Consequently, the pit morphology is related to the crystal structure. The walls of crystallographic - type pits are composed of crystalline planes, but the wall morphology may become quite complex if composed of multiple steps, where each step face has a different variation of the crystalline facet. Non-crystallographic pits are usually hemispherically shaped, an indication of an isotropic dissolution process within pits, with no evidence of crystallographic facets. Hemispherical pits were observed on iron in 0.5 M SO42- + 0.1 M Cl - solution by Vetter and Strehbolw [601 . Herbsleb and Engell [84] found a layer of sulfate on the bottom of hemispherical pits in iron, which they considered controls the isotropic dissolution in a diffusion-controlled process. They suggested that the sulfate anions enhance the formation of a viscous layer inside a pit, which controls the dissolution rate by a mass transfer process. Pickering and Frankenthal  [851  claimed that in a H 2 SO4 + NaCl solution  pits were polished into hemispherical pits in a later growth stage. Hemispherical pits were also observed on stainless steels and aluminum 186],[87].  LITERATURE REVIEW^  21  Crystallographic pits have been found on iron in NaC1 solutions HC1O 4 + NaC1 solution [851 . Brauns and Schwenk  [88]  [601  , and in a  observed crystallographic pits on  stainless steel in a chloride solution at a potential slightly above the critical pitting potential (E cp ). Latanision [811 , Garz et al. [82] and MacDougall [301 reported that the appearance of crystallographic pits on nickel varies with the crystallographic orientation of the metal surface. Pits were triangular on {111}, rectangular on {110} and square on {100}. The facets of the pits were shown to be {111} faces. Yasuda et al. [801 showed that pits formed on Al {100}, {110} and { 111 } faces were all crystallographic pits and the pit facets were {100}. Crystallographic pits, bounded by (1010) planes, were also found in zinc [891  .  The transition from crystallographic pits to hemispherical pits depends on the alloy composition, the solution chemistry and the applied electrode potential. For example, crystallographic pits were found on the iron surface in a phthalate buffer solution (pH 5) containing chloride, and hemispherical pits were formed on iron in 0.5 M sulfate + 0.01 M  chloride solution [601 Scully [901 found that the morphology of pits on Al depends on .  the kind of inhibitor added to the chloride solution. Crystallographic pits were found at the potential close to the critical pitting potential in stainless steels  [883  , but at higher  potential the pits tended to be non-crystallographic. The addition of alloying elements tended to change crystallographic pits into non-crystallographic pits on Al and Ni [801 ' [91] .  2.2.8 Alloying Element Effects Alloying additions to an alloy can change the pitting corrosion behavior of the material. Adding Mo and N to austenitic stainless steels dramatically increases the resistance to pitting of the steels  [92], [93]  . The addition of Cu to Al also raises the critical  LITERATURE REVIEW^  22  pitting potential [80] . Alloying elements can also affect the pitting behavior of metal single crystals. In contrast to pitting behavior of pure Al single crystal, Yasuda et al. [80] reported that the critical pitting potential was not affected by the crystallographic orientations of single phase Al-Cu single crystals, indicating that the addition of Cu results in the pitting independence of crystallographically oriented surfaces. Pits, which were initially crystallographic in pure metals, were reported to change to non-crystallographic pits when Cu (>0.5%) was added to Al 1801 , and Mo (>1%) added to Ni [9 ' ] .  2.2.9 Inclusions, Grain Boundaries and Dislocations The detrimental effect of sulfide inclusions on pitting resistance has been reported for carbon steels and stainless steels. In a study of pit nucleation conducted by Smialowski et al. [94] , it was found that pits nucleated preferentially at mixed manganese and iron sulfide inclusions which were present either in the form of separated particles or as shells surrounding Al, Cr and Mn oxides. Sulfide inclusions were identified as MnS, FeS, (Mn,Fe)S and CaS in carbon steels, and (Cr,Mn)S, (A1,Cr,Mn)S and TiS in stainless steels [50], [51], [94],  ^Pits were also seen to nucleate at other kinds of inclusions such as  oxides, silicates and carbides  1961 .  SEM observation of inclusions before and after pit  initiation demonstrated that pits could be nucleated directly at sulfide or oxide inclusions  Grain boundaries have been shown to be a preferred sites for pitting corrosion [46], [47], [80],  [97] .  In sensitized austenitic stainless steels chromium carbide precipitates  preferentially along grain boundaries and diminishes the content of Cr in the adjacent matrix, which makes it more susceptible to pitting corrosion [46] . Segregation of P at  LITERATURE REVIEW^  23  grain boundaries is also considered to cause preferential localized attack in stainless steels [471 A study on aged Al-Cu bi-crystals found that the precipitation of Al 2 Cu along .  the grain boundaries leads to pit nucleation and growth in the Cu-depleted region at grain boundaries [981 However, not all grain boundaries are sensitive to pitting corrosion. It .  was reported that there was no preferential pitting corrosion occurring at grain boundaries in solution-treated austenitic stainless steels in which the Cr-depletion zones along grain boundaries are eliminated during the solution treatment pitting at grain boundaries was detected on aluminum bi-crystals  [991 .  [981 .  No preferential  In contrast to aged  Al-Cu bi-crystals, it was found that grain boundaries were not the preferential locations for pitting to occur in the solution-treated Al-Cu bi-crystals, where the Al 2 Cu precipitates and Cu depletion zones were eliminated at the grain boundaries [981 . Therefore, the susceptibility of grain boundaries to pitting is due primarily to local chemical inhomogeneities such as precipitates, solute segregation and depletion at grain boundaries. It appears unlikely that alloys undergo a preferential pitting attack at grain boundaries when these chemical inhomogeneities are eliminated by proper heat treatments. The effect of dislocations on pit initiation in commercial alloys is less significant than inclusions and grain boundaries. Even in single crystals the effect of dislocations on pit initiation is not convincing. Wyon et al. E10°1 found that the number of pits formed on pure aluminum is much less than the dislocation density, whose observation showed that pit formation depended on the breakdown sites in the passive film rather than the emergent dislocations at the metal/passive film interface. Edeleanu et al. [Ion also found no correlation between pit initiation sites and dislocations in their study on aluminum. Kruger [791 using transmission electron microscopy (TEM), showed that the breakdown ,  LITERATURE REVIEW^  24  sites on iron single crystal surfaces are not related to dislocations. In a study of Fe-Cr alloy single crystals in MgCl 2 solution, Ahlers and Riecke  [1021  found that the pit density  in plastically deformed crystal specimens with a high dislocation density was the same as the unstrained crystal with a much lower dislocation density. A recent study on Al bi-crystals, which contained low angle tilt grain boundaries [981 that may be considered as a two-dimensional array of edge dislocations with their Burger's vectors parallel to the bi-crystal surface, showed no preferential boundary attack. Thus, the experimental evidence mostly indicates that dislocations do not directly act as preferential sites for pitting corrosion. However, Haruyama et al. [1031 studied the anodic behavior of iron whiskers substantially free of dislocations, and twisted whiskers, in a chloride containing solution. They claimed that the breakdown of the passive film could occur anywhere in the film, but pitting corrosion could occur only when the breakdown was located at an emergent dislocation. Hence, there is no general agreement on the effect of dislocations on pit initiation.  2.2.10 Solution Chemistry inside Pits Suzuki et al. [1041 showed that for three stainless steels, 304L, 316L, and 18Cr-16Ni-5Mo, the pH inside the pits dropped to 0.6-0.8, 0.06 -0.17 and -0.03 - 0.08 respectively in a neutral chloride solution. Peterson et al. [1051 found that the pH decreased to 1.2 -2.0 within pits for a stainless steel in seawater with a pH of 8.0. Butler et al.  [1061  measured local pH changes during pit growth in a bulk chloride solution with a  pH of 8.0. They found that the pH decreased to 2 inside a pit formed on an iron surface, and decreased to 1.5 in the pit on a Fe-7.5 % Cr surface. A lower pH was also observed  LITERATURE REVIEW^  25  within pits on aluminum (dropped to 3 from a bulk solution pH of 7), titanium (decreased to 1.7 from the bulk solution pH of 6.5) and copper (decreased to 5 from the bulk solution pH of 8)  [107], [109].  Halide ions can accumulate within growing pits, and their concentration can reach as high as 10 M. Suzuki et al. ""observed a considerable increase in the chloride concentration (to 6 M) in artificial pits. Mankowski and Smialowska  [110]  measured the  concentration of chlorides within naturally grown pits on a stainless steel in 0.5 M NaC1 + 0.1 N H2 SO 4 , and found that the chloride concentration depended on the stage of pit development. The maximum chloride concentration rose to more than 10 M at a pit diameter of about 0.5 mm, and then decreased to a steady value of 2.5 M as the pits grew larger. In general, the solution inside the pit can be significantly different from the bulk solution. The solution chemistry inside the pit exhibits a much lower pH and higher halide concentration than the bulk solution, which conditions facilitate the further growth of pits  2.2.11 Application of E-pH Diagrams to Pitting Corrosion Theoretical E-pH diagrams predict possible passivity conditions in metal - water systems. However, their application to aqueous solutions containing aggressive ions is limited when the passive film is subject to localized breakdown. Some attempts have been made to use E-pH diagrams to study localized corrosion in halide solutions. Figure 4 shows an experimental E-pH diagram for iron in the presence of chloride ions constructed by Pourbaix r ill] . The pitting corrosion zone, stable passive (perfect  LITERATURE REVIEW^  26  passivity) zone, and the conditions in active pits are indicated in Figure 4. These zones were determined experimentally and then plotted on a modified E-pH diagram for iron. Zuo et al. [112] also reported several experimental E-pH diagrams dealing with pitting corrosion of 18Cr-10Ni-Ti stainless steel at different temperatures. So far, only a few experimental E-pH diagrams have been constructed for application to localized corrosion, and none of them has considered the possible formation of complex species, particularly halide complexes. Therefore, it is difficult to use them to predict and explain pitting corrosion phenomena if there is the possibility of forming complexes with the aggressive anions.  LITERATURE REVIEW  27  10 12 19  0 2 y 6 8  . .^. .^. • . .^.^.^ .^.^.  _41  qr^  e^5^  , .^.  IIP\ •^  .  .^.^.^:--Pitting ....,.^  -2  ,  •  .^ . .^.  Imperfect pcsattlieity  _0_  aprotection  ct pAssiilley  _  immunity  2  ^4-4^6^  8  ^  IC) 12 19 pH  Figure 4 An experimental E-pH diagram for pitting corrosion of iron in chloride solution  11113  LITERATURE REVIEW^  28  2.2.12 Theories of Pit Initiation According to Kruger [211 , a successful theory for pit initiation should explain the following: 1. A critical pitting potential (E c d must be exceeded for pit initiation. 2. Aggressive ions are needed for the breakdown of the passive film on the metal surface. 3. Breakdown occurs at highly localized sites. Many theories have been proposed to account for pit initiation on the passivated metal surface. They are summarized as follows: 1. Adsorbed Ion Displacement Theory: This mechanism was suggested by Kolotyrkin [113] and Uhlig [541 . The passive film is considered to be an adsorbed film of oxygen. Passivity breakdown occurs when chloride ions are more readily adsorbed than passivating ions (0 2- , OH-). The chloride ions displace the passivating ions, resulting in pit initiation. Therefore, the critical pitting potential is the potential at which the adsorption of aggressive ions occurs, and pit initiation sites are the weakest points in the passive film where chloride ions are preferentially adsorbed. 2. Ion Penetration or Migration Theory: This theory was first proposed by Evans [7]  , and assumes that pores are present in the passive film. The C1 anion has a  sufficiently small diameter to penetrate through the pores in the film. Breakdown of the film occurs when the Cl - ions reach the metal surface.  LITERATURE REVIEW^  29  Hoar et al. [114] proposed that pit initiation is caused by the penetration of anions through the film under the influence of an electrostatic field across the film. Small ions more readily penetrate the oxide lattice, so that Cl - is more aggressive than Br- and P. Rozenfeld and Marshakov [115] suggested that the migration of a chloride ion is accompanied by the exchange of a passivating ion (e.g. OH-), which occurs at sites where the metal-oxygen bond is weakest. 3. Chemical-Mechanical Breakdown Theory: In 1967, Hoar E ll6iproposed a mechanical breakdown model for pit initiation. The adsorbing anions were postulated to replace adsorbed water on the film and reduce the interfacial tension or interfacial energy of the oxide/solution interface by the mutually repulsive force between the charged ions. Eventually, the interfacial tension is reduced so low that peptization by interfacial charge occurs. The cracks or splits thus produced in the film result in the breakdown of passivity. Another chemical-mechanical breakdown model was developed by Sato  [117]  .  He suggested that a high potential field could lead to mechanical rupture of the passive film by electrostriction pressure exceeding the compressive fracture strength of the film. The critical pitting potential in this model is the potential above which the film pressure exceeds the critical compressive strength of the film. The role of chloride ions is to retard the repassivation process after the breakdown of the passive film. They play a less important role in the rupture of the passive film. 4. Localized Acidification Theory: In 1937, Hoar  [118]  suggested that pits develop  because hydrolysis of corrosion products in the pits causes acidification. Galvele [119] developed a model, based on the assumption that metal cations hydrolyze inside micropits which already exist on the surface, and that the movement of cations out of pits  LITERATURE REVIEW^  30  is controlled by diffusion processes. Pit initiation is presumed to occur when a critical pH value is reached by the local acidification within micropits. In his model, Galvele emphasized the influence of acidification more than the specific role of chloride ions. Another model that considers the change in the local chloride concentration inside pits is due to Hisamatsu 11201 . He assumed that there was a critical chloride concentration in the pit electrolyte, C * , above which, the pit repassivation process is hindered and nucleated pits grow. 5. Depassivation - Repassivation Theory: This theory emphasizes the concept that passivity breakdown is a dynamic process. Videm  [1211  presumed the existence of  dynamic breakdown-repair events in the passive film. In the absence of aggressive anions, or below the critical pitting potential, film breakdown is followed by rapid healing, whereas in the presence of aggressive ions, and at the potential above E cp , the repassivation process is blocked. The role of chloride ions is to retard the repassivation process. Pitting is thus expected to occur when the rate of passivity breakdown is greater than that of repassivation. Wood's theory [521'  ^assumes the presence of pre-existing flaws in the passive  film. His theory may be classified with the ion penetration model, presuming that aggressive ions penetrate through the flaws. However, it is more likely that a dynamic crack-heal process occurs at the base of flaws. The flaws will be repassivated if the critical pitting potential is not met for pit initiation. 6. Chloride Nucleus Formation Theory: A critical chloride nucleus model has been developed based on the localization of chloride ions on the surface of the passive film.  LITERATURE REVIEW^  31  Such localization was found experimentally on the passive surface of iron and stainless steels by Janik-Czachor et al. [123], [124] anda was supported by the perturbation theory suggested by Okada " 251 . A chloride cluster model was first proposed by Hoar  [1261 ,  who assumed that several  chloride ions are adsorbed together on the passive film to form a localized chloride transitional complex with the cations in the passive film. Heusler and Fischer "271 suggested that the formation of a two-dimensional chloride nucleus causes the thinning and eventual breakdown of the passive film. The chloride nucleus theory  [581  " 281 suggests  that the formation of chloride nuclei is a dynamic process occurring randomly on the passive surface. When a pitting criterion (such as critical pitting potential) is met, stable critical chloride nuclei are formed and penetrate the oxide film layer, resulting in the breakdown of the passive film and pit initiation. 7. Point Defect Theory: This theory has been developed by Chao et al. [53] and Lin et al. [1291 . Point defects in the crystalline oxide film tend to accumulate to form a void at  the metal/oxide film interface. When the void grows to a critical size, the passive film collapses locally , resulting in the localized breakdown of the passive film. Based on this model, they have successfully derived a relationship between the critical pitting potential and halide concentration, and between electrode potential and the induction time for pitting.  LITERATURE REVIEW^  32  2.3 Remaining Problems Several remaining problems are outlined below, such as the pitting dependence of single crystals on crystallographic orientations, halide complex formation, and the change in local solution chemistry, which are the subjects of the research in the present thesis.  2.3.1 The Pitting Dependence on Crystallographic Orientations A pitting dependence on the crystallographic orientations has been observed on Fe, Al and Ni single crystals (see Section 2.2.6). However, this kind of research has been less systematically studied on single crystals of other metals. Therefore, it is still not clear whether orientation-dependent pitting is a common phenomenon on passivated metals. There is still no answer to the question why single crystals have different pitting resistances on the differently oriented surfaces. Pitting studies on iron and aluminum single crystals [79],[79] have suggested that the most closely packed surfaces ({ 110 } for iron and { 111 } for aluminum) exhibit the lowest pitting resistance, and these surfaces also have the highest dissolution rates [1301 ,[131]  ^the other hand, the pitting of nickel single  crystals indicated that the most closely packed { 111 } faces have the highest critical pitting potential [82],[133] , and the dissolution rate on the nickel {111} planes is the lowest [132] .  This may give some clue to the pitting dependence on crystallographic orientation,  but further work on the effect of crystallographic orientations is needed.  LITERATURE REVIEW^  33  2.3.2 Solution pH and Local Solution pH Control The critical pitting potential has been reported to be independent of bulk solution pH in an acidic to weakly alkaline pH region. However, there is a pronounced increase in the critical pitting potential when the solution becomes strongly alkaline (see Section 2.2.3). The local pH within growing pits has been found to be significantly different from the bulk solution pH, thus the local pH rather than the bulk solution pH is important in the pit initiation process (see Section 2.2.10). There is a lack of experimental evidence on local pH changes during pit initiation, possibly due to difficulties faced in monitoring the local surface pH. The significance of a local pH change during pit initiation is still unknown. Pitting corrosion behavior is anticipated to be different if a local pH is controlled by buffering the solution. Previous work has shown that the addition of buffers increased the critical pitting potential (see Section 2.2.3). However, the effect of the buffers was explained on the basis that the buffers acted as competitive adsorption inhibitors without attention paid to the local pH control. A recent study by Tromans and Sun [134] has suggested that the presence of buffers controls the local pH and prevents the breakdown of the film on Cu, rather than the buffer being competitively adsorbed on the surface.  2.3.3 Halide Complexes and Diagrams Dealing with Halide Complexes The interaction of halides with the oxide film had been long ignored until the formation of halide complexes was proposed by Hoar [135] for iron, and by Foley [136] for aluminum. Halide complex formation emphasizes the role of aggressive ions in the  LITERATURE REVIEW^  34  localized breakdown of passivity, and it can be presented in a diagram calculated thermodynamically in the same manner as Pourbaix's E-pH diagrams. Pourbaix's E-pH diagrams on simple H 2 0-metal systems give useful information on the possible conditions for passivity of metals, but their application is limited when they are applied to the localized breakdown of passivity in halide solutions. There are few theoretical E-pH studies which include halide complexes, though the thermodynamic data are available for the formation of halide - metal complexes. It is possible to introduce halide complexes into conventional E-pH type diagrams, or introduce them into diagrams that use halide concentration (X) and pH as the variations on the axes (termed X - pH diagrams). Such diagrams of water - halide - metal systems, based on thermodynamic considerations, could prove to be useful to an understanding of the possible role of complex halide ions in the pit initiation process.  2.3.4 Pit Initiation Mechanisms There is no generally accepted theory to account for all aspects of pit initiation. The adsorption theory fails to explain the breakdown of the solid phase oxide films. The local acidification model and chemical - mechanical breakdown model are faced with the difficulty of explaining the aggressive role of halides during film breakdown. Surface analysis techniques have not detected the penetration of chloride ions into the film. The point defect theory cannot be applied to amorphous films. The chloride nucleus theory fails to give information on the specific chloride complexes which lead to pit initiation. No generally accepted theory has been developed to account for crystallographic orientation - dependent pitting phenomena.  OBJECTIVE^  35  3 OBJECTIVE The purpose of the present investigation is to extend the fundamental knowledge of crystallographic orientation - dependent pitting phenomena and to enhance the mechanistic understanding of the pit initiation process in halide solutions. These goals are pursued by conducting experiments on pure metal single crystals of Sn, Zn and Al, and polycrystalline Ni in halide solutions, and by constructing thermodynamic diagrams to represent equilibria in several different water-halide-metal systems. The general sequence of procedures is outlined below. 1. Pitting Corrosion of Single Crystals The dependence of the critical pitting potential on the crystallographic orientation of Sn and Zn single crystals is investigated in aqueous chloride solutions. The pit morphology is examined using scanning electron microscopy (SEM). 2. Effects of Solution pH and Halides Buffers are used to help control the local surface pH and assess the significance of a local pH change during the pit initiation process on polycrystalline Ni. Pitting corrosion is investigated in F, Cr and Br solutions to correlate the aggressiveness of the halide species according to their tendency to form metal halide complexes on Zn and Al single crystals, and polycrystalline Ni. 3. E-pH Diagrams for Water-Halide-Metal Systems New E-pH diagrams are constructed for Sn, Zn, Al and Ni in halide environments, based on available thermodynamic data for halide complexes. X-pH diagrams are, for the first time, also introduced for the corresponding metal oxide - halide - water systems.  OBJECTIVE^ 4. Pit Initiation Mechanism An approach to the mechanism of pit initiation is made by taking into consideration the dynamic nature of the oxide film formation, local solution chemistry and halide - metal complexes. A theory is proposed to explain the effect of crystallographic orientations, solution pH, buffers and halide ions on pitting behavior.  36  EXPERIMENTAL  ^  37  4 EXPERIMENTAL  4.1 Materials Single crystals of zinc, tin and aluminum, with a purity of 99.995%, were used in the study of the dependence of pitting corrosion resistance and pit morphologies on crystallographic orientation. The single crystals were grown by Professor Fred Weinberg using the Bridgman technique with pre-oriented seeds in a horizontal furnace. The size of the grown crystals was about 8 x 6 x 150 mm. The crystal structures of zinc, tin and aluminum, their melting points and crystal growth directions are listed in Table 4.  Table 4. Structure, melting points and growth directions of single crystals  Metal  Structure  Melting point, °C  Growth Direction  Zn  HCP  420  [1010] and [1120]  Sn  BCT  232  [100] and [110]  Al  FCC  660  <001>  Electroplating grade polycrystalline nickel was also used in the study of environmental effects on pit initiation. The polycrystalline nickel, FCC structure, had a purity of 99.95%. Cold rolled sheets of nickel, about 3 mm thick, were used for the preparation of nickel specimens.  EXPERIMENTAL^  38  4.2 Single Crystal Specimen Preparation (a) Crystallographic Orientation Determination: The crystallographic orientations of Sn, Zn and Al single crystals were determined by the X-ray back-reflection Laue technique  [137j  .  A Philips Model PW 1830 X-ray generator  was used at a working voltage of 35 KV and a current of 20 mA. The distance between the single crystal specimen and X-ray film was 30 mm, and the exposure time was 30 min. Figure 5 shows a back reflection Laue pattern of a tin single crystal. The (0001) surface of zinc was also obtained by cleaving the zinc single crystal along the basal plane after being cooled in liquid nitrogen. Due to the difficulty in distinguishing between the (1010) and (1120) Laue patterns, the deformation twinning method was used to determine these orientations. The Zn single crystal was twinned along the well-defined twinning (1012) plane by applying a tensile stress parallel to the basal plane r 1381 . (b) Sampling and Mounting Once the orientations were determined, the crystals were cut into small pieces (about 10 mm long) along the desired crystallographic angles to obtain specifically oriented surfaces using spark machining or a jewelry saw (for the zinc single crystal, spark machining causes cleavage along the basal plane). Five differently oriented surfaces, (100), (110), (111), (011) and (001) were prepared for tin, while (0001), (1010) and (1120) oriented surfaces were prepared for zinc single crystals. The deviation of the oriented surface from the desired crystallographic orientation was ± 5°. A conducting wire (copper wire) was soldered to the sampled crystal, then mounted in cold-curing epoxy resin (Lecoset). A mounted specimen is shown schematically in Figure 6.  EXPERIMENTAL^  [110]  39  [100]  Figure 5 A X-ray back reflection Laue pattern from the tin (001) surface  Figure 6 A schematic diagram of the specimen used in pitting tests  EXPERIMENTAL  ^  41  (c) Mechanical and Electrochemical Polishing The specimen was mechanically polished up to grit 600 paper, and then to 1.0 gm using a water base diamond suspension (BUEHLER METADI). The surface was then polished electrochemically in order to remove the deformed layer resulting from mechanical polishing. Table 5 lists the electrolytes and conditions for the electrochemical polishing of zinc, tin, and aluminum single crystals.  Table 5. Electrochemical polishing parameters for Zn, Sn and Al Crystal  Zn * Al **  Electrolyte Cathode Current, A/cm2 Temperature, °C 0.15 HC1O4 200 mL Pt 30 (CH3 C0) 2 0 800 mL KOH 250 g Pt 0.22 20 H2 O 750 mL HC1O 4 62 mL Ethanol 700 mL Pt 3.85 20 2-butoxyethanol 100 mL H2 O^137 mL  Time, s 360 900. 20  * P. V. Shigelev, Electrolytic and Chemical Polishing of Metals, Freund Publishing House, 1974 ** Handbook of Metals, ASM, Vol. 9, P. 353, 1985  The electrochemically polished specimen was cleaned with distilled water and acetone and dried in cold air. In order to prevent preferential corrosion attack at the specimen/epoxy interface, the interface was coated with a cellulose acetate lacquer (red lacquer), leaving an exposed surface area varying from 0.2 -0.6 cm 2 for different samples.  EXPERIMENTAL^  42  4.3 Polycrystalline Nickel Specimen Preparation Nickel specimens (15x15 mm) were cut from 3 mm thick Ni sheet, then mounted and mechanically polished in the same manner as the single crystal specimen shown in Figure 6, without further electrochemical polishing. For nickel, the specimen/epoxy interfaces were protected by a GLYPTOL coating (an alkaloid enamel paint) instead of the red lacquer used for the single crystals. The exposed surface area of the test specimen was about 1 cm 2 .  4.4 Test Temperature All experiments were conducted at room temperature in the temperature range 20-23°C. Temperature changes within this range had no apparent effect on the experimental results.  4.5 Test Solutions Halide - containing solutions were used in the pitting corrosion tests. All the solutions were prepared using reagent grade chemicals and distilled water. The pH of the solution was adjusted to the desired value by addition of NaOH, HCl acid or buffers, and measured to ± 0.1 pH units by using a combination glass pH electrode and a Corning Model 125 pH meter. The halide - containing solutions which were used in tests for single crystals are listed in Table 6.  EXPERIMENTAL  ^  43  Table 6. Test solutions for pitting corrosion of single crystals.  Crystal  Solution No.  Test Solution  pH  Sn  A-1  0.1 M NaC1 + 0.5 M NaNO3  6.0  A-2  0.1 M NaC1  A-3  0.1 M NaC1 + 0.5 M NaNO3  A-4  0.1 M NaC1 + 0.1 M NaHCO 3/Na2 CO 3  A-5  0.1 M NaBr  A-6  0.1 M NaC1  A-7  0.1 M NaBr  A-8  0.1 MNaF  Zn  Al  9.2  3.4  6.0  In the study of the environmental effects of halides, pH, and buffers on the pitting corrosion of nickel, several types of solutions were prepared for experiments. Type I solutions, consisting of NaNO 3 or Na 2 SO 4 , were used as a basis for comparison with halide solutions. Their pH values were adjusted by the addition of NaOH (for B-1 and 2 solutions), or H 2 SO4 (for B-3 solution). Type II solutions were halide solutions with a pH range of 2.5 14, which were used to study the effect of bulk solution pH and the effect of different halide ions. Type III were chloride plus nitrate solutions. The pH values of Type II and III solutions were adjusted by the addition of NaOH or HC1. Type IV were halide solutions buffered by Na 2 CO 3 /NaHCO 3 or Na 3 PO4/Na2 HPO 4 , which were used to investigate the significance of pH control in pitting corrosion. All the solutions used for tests on nickel are listed in Table 7.  EXPERIMENTAL  ^  44  Table 7. Test solutions for pitting corrosion of nickel  Type  Solution No.  Composition  pH  Type I  B-1 B-2 B-3  10.5 10.5 2.5  Type II  B-4 B-5 B-6 B-7 B-8 B-9 B-10 B-11 B-12 B-13 B-14  1.0 M NaNO 3 1.0 M Na2 SO4 1.0 M Na2 S°4 1.0 M NaC1 1.0 M NaC1 1.0 M NaC1 1.0 M NaC1 1.0 M NaC1 1.0 M NaC1 1.0 M NaC1 1.0 M NaBr 1.0 M NaF 1.0 M NaF 1.0 MHF  2.5 4.5 6.5 8.5 10.5 12.5 14.0 10.5 6.0 10.5 3.1  Type III  B-15 B-16 B-17  1.0 M NaC1 + 0.1 M NaNO3 1.0 M NaC1 + 0.01 M NaNO3 1.0 M NaC1 + 0.001 M NaNO3  10.5 10.5 10.5  B-18  1.0 M NaC1 + 9.5x10 -2 M Na2 CO 3 + 5.0x10 -3 M NaHCO 3  10.5  B-19  1.0 M NaC1 + 9.5x10-3 M Na2 CO 3 + 5.0x10 -4 M NaHCO 3  10.5  B-20  1.0 M NaC1 + 9.5x10-4 M Na2 CO 3 + 5.0x10 -5 M NaHCO 3  10.5  B-21  1.0 M NaCl + 8.5x10 -2 M Na2 HPO 4 + 1.5x10 -2 M Na 3 PO4  10.5  B-22  1.0 M NaC1 + 8.5x10-3 M Na2 HPO4 + 1.5x10-3 M Na3 PO4  10.5  B-23  1.0 M NaC1 + 8.5x10 -4 M Na2 HPO4 + 1.5x10 -4 M Na3 PO4  10.5  B-24  1.0 M NaBr + 8.5x10 -2 M Na2 HPO 4 + 1.5x10 -2 M Na3PO4  10.5  Type IV  EXPERIMENTAL^  45  4.6 Electrochemical Test Cell and System The electrolyte test cell was a glass cell (EG&G C, K47 cell), Figure 7, containing five inlets for a working electrode (specimen), two platinum counter electrodes, a saturated calomel reference electrode (SCE) and a nitrogen gas purger. For fluoride solutions, a TEFLON (polytetrafluoroethylene) cell was used with the same configuration as the glass cell. A KC1 saturated agar-agar gel salt bridge was connected to the glass calomel reference electrode to prevent contact with the fluoride - containing solution. A microprocessor controlled potentiodynamic polarization system (EG&G Model 350A) was used for pitting corrosion testing. The accuracy was ±2mV in potential and ±0.5% of the maximum current in each current range. Data were transferred to a computer through a serial port (RS232C) and saved on floppy discs. Figure 8 shows schematically the overall system used for the electrochemical measurements.  WE - Working Electrode (Specimen) CE - Counter Electrode (Pt) SCE - Saturated Calomel Reference Electrode  Figure 7 A schematic diagram of the cell used in pitting tests  Nitrogen Gas Cylinder  Figure 8 A schematic diagram of the corrosion measurement system used in pitting tests  EXPERIMENTAL^  48  4.7 Pitting Scan Test Technique A cyclic potentiodynamic polarization technique, called the pitting scan test, was used in the present study. The test started at the corrosion potential (open - circuit potential), then the potential was increased in a positive (anodic) direction at a scan rate of 0.5 mV/s until a predetermined current threshold (5 mA/cm 2) was reached. Then the potential scan was reversed to the negative (cathodic) direction. The test was completed when the potential scanned back to 50 mV lower than the initial corrosion potential. During the backward scan, there is an anodic current hysteresis loop, which characterizes the occurrence of localized corrosion on the passivated surface. The schematic appearance of the resulting polarization curve is similar to that shown in Figure 3 (solid line). The critical pitting potential (E cp) is thus determined from the polarization curve. The measured value of the critical pitting potential is reported with reference to the Saturated Calomel Electrode (SCE). The conversion of potential values between the SCE and the Standard Hydrogen Electrode (SHE) is: E (SCE) = E (SHE) - 0.241 (Volts)  4.8 Test Procedure Dissolved oxygen was removed from the test solution by deaerating with N 2 for one hour before the specimen was placed into the cell. Subsequently, the specimen was put in the cell and immersed in the solution for another hour to let the corrosion potential (open-circuit potential) of the specimen stabilize, and then the pitting scan test was started at the corrosion potential. The nitrogen purge operated continuously during the test. After  EXPERIMENTAL^  49  each test, the specimen was washed with distilled water and ethanol, then dried and stored in a desiccator until required for pit morphology examination. Most of the tests were repeated two or three times to confirm the reproducibility and validity of the results.  4.9 Pit Morphology Examination Pit morphology was examined by scanning electron microscopy, using an Hitachi S-2300 SEM. Secondary electron imaging was employed, using an excitation energy of 20 KeV. In addition, a Zeiss Interference Microscope (Model 3001) was used to examine surface topography. This technique uses a monochromatic light source (thallium light source) to produce interference bands across the specimen surface. The interference patterns from a pitted surface enable the orientation of the facets in the pits to be determined in relation to the top surface of the specimen.  RESULTS  ^  50  5 RESULTS 5.1 Pitting Corrosion Behavior of Sn, Zn and Al Single Crystals Table 8 summarizes the pitting tests conducted on tin, zinc and aluminum single crystals. The differently oriented surfaces of the tin and zinc single crystals were tested in chloride solutions to investigate the dependence of the critical pitting potential on the crystallographic orientations. The effect of the Na 2 CO 3 /NaHCO 3 buffer was studied only on the zinc (0001) surface. Several halide solutions were used in the investigation of the pitting corrosion of the zinc (0001) and aluminum { 1001 surfaces in order to compare the aggressiveness of the different halide ions.  Table 8. Experimental summary for single crystals Effect  Test Solution  Crystallographic Orientation  A-1  Zn  (0001) (1610) (1120)  Crystallographic Orientation  A-2 A-3  Zn Zn  (0001) (0001)  Buffer  Halide  Al  { 1001  Halide  A-4 A-5 A-6 A-7 A-8  Single Crystal Sn  Surface Orientation (001) (011) (111) (110) (100)  RESULTS^  51  5.1.1 Pitting Potential Dependence on Crystallographic Orientations Figure 9 shows the potentiodynamic polarization curves obtained on the Sn (011) and (111) oriented faces in 0.1 M NaC1 + 0.5 M NaNO 3 solution at pH 6.0. The results clearly demonstrate the difference in their critical pitting potentials. The critical pitting potential of the Sn (011) face is about 25 mV higher than that of the (111) face. Five differently oriented surfaces of Sn single crystals have been tested in 0.1 M NaC1 + 0.5 M NaNO 3 solution at pH 6.0. The results of these pitting tests, with respect to effects on  the critical pitting potentials, are listed in Table 9 and plotted in Figure 10.  Table 9. Critical pitting potentials obtained from five oriented faces of tin  Tin Face  Critical Pitting Potential, V (SCE)  (001)  -0.285  -0.290  -0.288  (110)  -0.285  -0.290  -0.288  (011)  -0.285  -0.290  -0.289  (100)  -0.282  -0.285  -0.290  (111)  -0.305  -0.308  -0.310  The critical pitting potential for each Sn surface was measured three times, the scatter for each surface orientation being ± 5 mV. It is clear from Table 9 and Figure 10 that the critical pitting potentials for the (001), (110), (011) and (100) surfaces of Sn are very close and lie between -0.282 V (SCE) to -0.290 V (SCE). However, the (111)  RESULTS^  52  surface of Sn has a critical pitting potential that is about 22-25 mV lower than the other surface orientations. Therefore, the (111) face of Sn has the lowest pitting corrosion resistance among the five tested surface orientations. The critical pitting potentials of the Sn single crystals may be conveniently summarized: Eel, for Sn;  (001) (110) (100) (011)  Pitting tests were conducted on the (0001) (1010) and (1120) faces of zinc single crystals in chloride solutions. The potentiodynamic polarization curves measured in 0.1 M NaC1 at pH 9.2 are shown in Figure 11. In contrast to the polarization curves for tin  single crystals (Figure 9), there is no sharp potential point (no well-defined E cp ), where a sudden increase in anodic current occurs, on the polarization curve of zinc single crystals. Consequently, an extrapolation technique is used to determine the critical pitting potential on each orientated surface, as shown in Figure 11. The dependence of E cp on the crystallographic orientation of the zinc surface is shown in Figure 12. In these experiments only two repeated measurements were made for each surface and the reproducibility was ± 5 mV. The critical pitting potential is observed to progressively decrease as the crystallographic orientations are changed from (1010) —> (1120) —> (0001). The critical pitting potentials were also measured on the surfaces of zinc single crystals in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 9.2 (Figure 13). The results show that the presence of NaNO 3 raised the E cp to significantly higher values than those observed in 0.1 M NaCl at pH 9.2. However, the order of the dependence of E cp on  RESULTS^  53  crystallographic orientations in 0.1 M NaC1 + 0.5 M NaNO 3 exhibits the same trend as in 0.1 M NaCl. Therefore, in both 0.1 M NaC1 and 0.1 M NaC1 + 0.5 M NaNO 3 solutions at pH 9.2, the critical pitting potentials of Zn single crystals may be summarized: Ecp for zinc; (10 10) > (11 20) > (0001) -  -  -0.2  -0.3  -0.4  -0.6  -0.7  -0.8 0  2  4  6  8  Log(i), nA/cm 2  Figure 9 Potentiodynamic polarization results obtained on tin (011) and (111) faces in 0.1 M NaCl + 0.5 M NaNO 3 at pH 6.0  -0.26  -0.27  -0.28  IA 0 Cr)^-0.29 ..._ ...,'  a) 5  •  I  i  •  •  A  •  A  -0.3  •  I  0...^-0.31 C:Y) C IP 4-,  .CL^  -0.32  -0.33  -0.34  1  (001)  ^  (110)  ^  (011)  ^  (100)  Orientation  Figure 10 Variation of the critical pitting potential with crystallographic orientation of tin in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 6.0  -0.98 -1  1  5  3  LOG(i), nA/cm  7  2  Figure 11 Potentiodynamic polarization curves obtained on three differently oriented surfaces of zinc 0.1 M NaC1 at pH 9.2  -1.02 -1.03  GI  -1.04  O  -1.05  cr) >  Tz  -1.06  • •  -1.07  11'^-1.08  o^-1.09  • •  0CY)^-1.1 C  IP + a^-1.11 am "  -1.12 -1.13 -1.14  I^ (10 10)  ^  (1120)^ (0001) 1  Orientation Figure 12 Variation of the critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 at pH 9.2  -0.72  -0.73  -0.74  4E'  • •  -0.75  0  ta_  D)  -0.76  17 4-/  a-0.77  •  -0.78 (1010)  (1120)  (0001)  Orientation  Figure 13 Variation of critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 9.2  RESULTS^  59  5.1.2 Pit Morphology The shape of the pits on Sn crystals varied with the change of the crystallographic orientation of the surface. Figure 14 shows the crystallographic pits formed on the tin (001) face in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 6.0. These square pits show a 4-fold symmetry and pit edges are reasonably parallel to [100] or [010] directions. There are multiple steps formed on the pit walls in larger pits. In order to measure the slope of the pit walls and to determine their orientations, the interference microscope was used to obtain interference bands from the pits. Figure 15 shows an interference micrograph which was taken from the tin (001) face. By measuring the number of interference bands on the pit wall, the depth of the pit was calculated according to the equation: t = nX/2, where t is the depth of the pit, n is the number of bands on the pit wall, and X is the wavelength of the thallium light source (X = 0.45 pm). The pit wall angle (a) was thus determined by the equation: tana = t/l, where 1 is the half width of the pit (Figure 16). Table 10 gives some pit wall angles measured from the smaller pits where no obvious steps formed on the pit walls. Table 10. Measured pit wall angles from the pits formed on the tin (001) surface  n  t, gm  1, gm  a, °  10  2.70  5.1  27.8  15  4.05  8.7  25.0  18  4.86  9.7  26.6  RESULTS  (a)  (c)  ^  ^  60  (b)  (d)  [010]  Figure 14 Pits formed on the tin (001) face (a, b, and c), and crystallographic facets (pit walls) in the pit (d)  RESULTS  ^  Figure 15 A light interference photograph taken on the pitted (001) face of tin  61  [010]  6  .  0 ..._.  0  ,  14  top view  Interference Bands  1 )  A, side view  Figure 16 A schematic view of the pit on the tin (001) face and the determination of the pit wall angle  RESULTS^  63  The theoretical angles between the Sn (001) plane and either the (011), (101), (101) or (011) planes (they belong to the same family) are the same, and equal to 28.6°, as calculated form the lattice constants (a = b = 5.8197 A, c = 3.17488 A). The measured pit wall angles in Table 10 are close to 28.6°. The exposed (011), (101), (101) and (Oil) faces give a four-fold symmetrical pit on the tin (001) surface, and they intercept the Sn (001) surface with the traces parallel to [100] and [010] directions. The walls bounding the pits on the tin (001) surface are thus identified as (011), (101) (101) and (011) faces, which are referred to as {011} tetragonal bipyramidal planes in the following Sections. When multiple steps were formed on the pit walls, the apparent pit wall angle becomes smaller, as shown schematically in Figure 17. This was confirmed by the pit wall angle measurements performed on the larger pits with steps on the pit walls, where the angle varied from 15° to 20°. Unfortunately, no clear interference patterns could be obtained from the pits formed on other tin surfaces. Some calculations were made based on the interception traces of the pit walls with the tin free surface (the pitting tested specimen surface) in order to speculate on the possible orientations of the pit walls. On the Sn (111) surface, pits contained six walls (Figure 18). The measured angle between two surface traces at the upper part of the pit in Figure 18.b is 77-80°, close to the crystallographic angle, 77.8°, between the [101] and [011] directions, which are the interception traces of (100) and (010) planes with the tin (111) surface. Therefore, these two pit walls are speculated to be (010) and (100) planes, referred to as { 1001 tetragonal prismatic planes in the following Sections. At the lower part of the pit in Figure 18.b, the angle between the two surface traces was measured to be 133-135°. If the two corresponding pits walls are (011) and (101) planes, the interception trace of (011) with the tin (111) surface will be  RESULTS^  64  parallel to the [211] direction, and the interception of (101) with the (111) surface will be parallel to the [121]. The crystallographic angle between these two directions is 134.7°, which is in the range of measured angles. Therefore, the remaining four pit walls on the Sn (111) surface are predicted to be (011), (101), (011) and (101) planes ({011} tetragonal bipyramidal planes), as shown in Figure 18.c. Pits formed on the tin (011) surface have a two-fold symmetry with four pit walls and a pit base as shown in Figure 19. Assuming that the pit base is (011) and three of pit walls (excluding the right side pit wall) are (101), (101) and (011) planes as shown in Figure 19.c, the interception traces of the (101) and (101) on the tin (011) surface would be the [111] and [111] directions respectively. The measured angle between the traces on SEM photographs is 81-84°, very close to the theoretical angle 82.6° between the two [111] and [111] directions on the tin (011) surface. Thus, they are identified to be (011), (101), (101) and (011) planes (belonging to the {011} family). The remaining pit wall (right side pit wall in (Figure 19) is assumed to be a (010) plane, because its interception trace with the (011) surface is parallel to the [100] direction. A pit formed on the tin (110) surface (Figure 20) has two well-defined pit walls. The angle between these pit walls and the (110) surface is about 45°, measured by tilting the specimen while observing the surface in the SEM, and the surface interception traces are parallel to the [001] direction. Therefore, these two pit walls are assumed to be (100) and (010) planes ({100} family), as shown in Figure 20.b. The pits formed on the Sn (100) surface are rectangularly shaped, with crystallographic features (Figure 21). However, these pits are very shallow, and it was difficult to identify the pit walls.  RESULTS^  65  Figures 22 - 24 show the pit morphology on the (0001), (1010) and (1120) faces of zinc single crystals respectively in 0.1 M NaC1 at pH = 9.2. The shapes of the pits are different on the different faces. Pits close to a 6-fold symmetry were found on the zinc (0001) surface (Figure 22). The shapes of pits formed on the zinc (1010) and (1120) surfaces are irregular (Figure 23 and 24). Well-defined hexagonally shaped pits were found on the zinc surface cleaved along the (0001) plane (basal plane) after pitting in 0.1 M NaCl solution at pH = 9.2 (Figure 25).  The hexagonally shaped pits formed on the cleaved zinc (0001) surface are consistent with the hexagonal crystal structure of Zn (Figure 25). The pit edges are parallel to the interception traces of (1010), (1010), (0110), (6110), (1100), (1100) with the zinc (0001) surface respectively, and the pit walls are perpendicular to the zinc (0001) surface. Therefore, pit walls are identified as the (1010), (1010), (0110), (0110), (1100), (1100) planes (referred to as {1010} hexagonal prismatic planes) on the zinc (0001) surface. The pits formed on the zinc (1010) and (1120) surfaces exhibited crystallographic features to some extent (Figure 23 and 24). However, it is difficult to speculate on the possible orientation of the pit walls because no well-defined surface traces of the pit walls were observed.  less than 28.6  Figure 17 Relationship between the apparent crystallographic angle and the steps formed on pit walls  67  RESULTS^  (b)  (a) Mau 4,0  N  134.7  Figure 18 Pit morphology of the tin (111) face (a, b); crystallographic facets in the pit (c)  68  RESULTS^  a  -  0  Figure 19 Pit morphology of the tin (011) face (a, b); crystallographic facets in the pit (c)  RESULTS^  (a)  (b)  [00 1]  Figure 20 Pit morphology of the tin (110) face (a); crystallographic facets in the pit (b)  69  RESULTS^  70  77110  .,..11111.1101  , ...411nraT *MOW^  4,401010121  9119%.^::-.0.141.4^-77nis 1,--r•^- 7 r■_111114.446 1 orPIPIP ...up ...• -wo ^.... .,  -  ,  ,_,  ,............t  -10  1773111  "opa. "79110111r  4 gum  Figure 21 Pit morphology of the tin (100) face  RESULTS^  Figure 22 Pit morphology of the zinc (0001) face  71  RESULTS  ^  Figure 23 Pit morphology of the zinc (1010) face  72  RESULTS^  Figure 24 Pit morphology of the zinc (1120) face  73  RESULTS^  Figure 25 Pits formed on the cleaved basal plane (0001) of zinc  74  RESULTS^  75  5.1.3 Effect of pH Buffer and Halides on Pitting of Zn The effect of the NaHCO 3/Na2 CO 3 buffer on the pitting of Zn single crystals was investigated on the (0001) surface in the 0.1 M NaC1 + 0.1 M NaHCO 3/Na2 CO 3 solution at pH 9.2. The potentiodynamic polarization curves are presented in Figure 26, where they are compared with the unbuffered 0.1 M NaC1 solution at the same pH. Clearly, the NaHCO 3/Na2 CO 3 buffer raised the critical pitting potential by at least 250 mV and produced a more clearly defined E cp . The addition of buffer did not change the pH of the bulk solution, but it controlled the local pH near the specimen surface so that local acidification did not occur as readily as in the unbuffered solution. The NaHCO 3/Na2CO 3 buffer may also have contributed to the extension of the passive zone of zinc to a lower pH region by the formation of ZnCO 3 . This will be discussed in later sections (see E-pH diagram for H 2 0-0O 3 2- -Zn system in Section 6.2, and Section 7.6). The effect of different halide species on the pitting of the Zn (0001) surface was investigated by substituting 0.1 M NaBr for 0.1 M NaCl, while maintaining the pH at 9.2. The polarization results in these two solutions are compared in Figure 26. The critical pitting potential in the bromide solution is more clearly defined and about 100 mV higher than in the chloride solution, indicating that chloride ions are more aggressive than bromide ions in promoting the pitting corrosion of zinc.  -0.8  -0.85  in 0.1 M NaCI + 0.1 M carbonate buffer ^ in 0.1 M NaCI in 0.1 M NaBr  -0.9  -0.95  -1.15  -1.2  -1.25  0  2  4  6  8  LOG(i), nA/cm 2 Figure 26 Polarization curves of the zinc (0001) face in 0.1 M unbuffered and Na2CO 3/NaHCO 3 buffered halide solutions at pH 9.2  RESULTS^  77  5.1.4 Effect of Halides on Pitting of Al Single Crystals The Al { 001 } faces were chosen for the investigation of the effect of three different halide species (F, C1 and Br ) on the pitting corrosion of Al single crystals. -  Pitting scan tests were conducted in 0.1 solutions of HF and NaC1 at pH = 3.4 (Figure 27), and in 0.1 M solutions of NaF, NaC1 and NaBr at pH 6.0 respectively (Figure 28). Comparisons were made on the corrosion behavior of Al based on these polarization results. In Figure 27, the polarization curve in 0.1 M fluoride solution at pH 3.4 shows the active dissolution of Al, corresponding to general corrosion. The generally corroded Al (001) surface in the fluoride solution at pH 3.4 exhibited crystallographic features, and roughened faces were exposed which appeared to be the { 111 } planes (Figure 29). The polarization curve in 0.1 M NaC1 at pH 3.4 exhibits a defined passive region below Ecp and passivity breakdown at Ecp . Pitting corrosion occurred on the Al { 100} in the chloride solution at pH 3.4. In Figure 28, pitting corrosion occurred on the Al {001} surfaces in both 0.1 M NaC1 and 0.1 M NaBr at pH = 6.0, and pits appeared to be crystallographic, and bounded by pit walls of Al {100} planes (Figure 30). The critical pitting potential in the bromide solution is more than 50 mV higher than in the chloride solution. The polarization curve of Al {001} showed a large active dissolution nose near -1.2 V (SCE) in 0.1 M NaF at pH 6.0, followed by an anodic current decrease as a result of thick salt film formation on the surface (Figure 31). No pitting was observed in the fluoride solution at pH 6.0. The thick salt film was analyzed by the energy dispersive X-ray (EDX) method. Enrichment of Al and F was detected, and the salt film is identified to be an aluminum fluoride salt  RESULTS^  78  film. By potentiodynamically polarizing the Al { 100} to -0.6 V (SCE) from the corrosion potential, salt crystals were initially formed on the Al surface, and showed dendritic growth (Figure 32) The above results indicated that the corrosion behavior of aluminum is dependent on the specific type of halide species in the solution. The stability of the passive oxide film on aluminum varies with the type of halides in the solution, and the passive film is less stable in the chloride solution than in the bromide solution. The formation of a salt film also alters the corrosion behavior of Al.  -0.2 -0.3 -0.4 -0.5 -0.6  W 0  -0.7 -0.8 -0.9  ui  -1  -1.2 -1.3 -1.4 -1.5  0  ^ ^ 2  4  ^  6  ^  8  LOG(i), nA/cm 2  Figure 27 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 3.4  0.6 0.4 0.2  0 -0.2 -0.4 C/) -0.6 -0.8  -1 -1.2 -1.4 -1.6 -1.8  0  2  4  LOG(i), nA/cm  6  8  2  Figure 28 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 6.0  RESULTS  ^  Figure 29 Exposed (111) faces on the Al (100) surface corroded in 0.1 M HF at pH 3.4  81  RESULTS  ^  Figure 30 A crystallographic pit formed on the Al (100) in 0.1 M NaC1 at pH 6.0  82  RESULTS  ^  Figure 31 A thick salt film formed on the Al (100) in 0.1 M NaF at pH 6.0  83  RESULTS  ^  Figure 32 Salt crystals formed on the surface of Al in 0.1 M NaF at pH 6.0  84  RESULTS  ^  85  5.1.5 Overall Summary of Corrosion Behavior of Single Crystals Table 11 summarizes the results obtained from the pitting tests on tin, zinc and aluminum single crystals in halide-containing solutions.  Table 11 Summary of corrosion behavior of single crystals Metal  Solution No.  Sn  A-1  Zn  Corrosion Behavior Pitting  A-2  Ecp lowest on (111) Pitting, Ecp  A-3  (1010)>(1120 )>(0001)  Zn  A-4  Zn  A-5  Pitting Ecp (buf) > E cp (unbuf) Pitting  A-6  Ecp (Br) > Ecr,(C1 ) Pitting  A-7 A-8, pH=3.4 A-8, pH=6.0  Ecp(Br-) > E cp (Cr) Active Active/Passive  Pit Morphology Crystallographic Crystallographic Crystallographic Crystallographic  -  Al Al Al  Notes:^Ecp(buf) - critical pitting potential in the buffered solution. Ecp (unbuf) - critical pitting potential in the unbuffered solution. Ecp (Br) - critical pitting potential in the bromide solution. Ecp (C1 ) - critical pitting potential in the chloride solution. -  Crystallographic (Exposed { 111}) (Thick salt film)  RESULTS^  86  5.2 Pitting Corrosion Behavior of Polycrystalline Nickel  5.2.1 Polarization Behavior in Nitrate and Sulfate Solutions The polarization behavior of nickel in 1.0 M NaNO 3 and 1.0 M Na2 SO 4 at pH = 10.5 is shown in Figure 33, along with the polarization curve of Ni in 1.0 M NaC1 at the same pH. The results show passive behavior of Ni in 1.0 M NaNO 3 and 1.0 M Na2 SO4 at pH 10.5. The increases in anodic current density at higher potentials (above 0.6 V (SCE)) in these solutions were caused by the evolution of oxygen at the surface of the nickel specimen. The equilibrium potential for oxygen evolution at 25 °C can be calculated using the following equations: 2H 2 0 = 0 2 + 4H + + 4e -^(5.1) E (021H2 0) = 1.228 -0.0591pH^(V vs. SHE)^(5.2)  For the solution at pH 10.5, E (021H 2 0) = 0.608 V (SHE) = 0.367 V (SCE). The potential (about 0.6 V (SCE)) where the anodic current increases in Figure 33 is higher than E (0 21H20) (0.367 V (SCE)). Furthermore, unlike the polarization curve in the chloride solution, no anodic current hysteresis loop occurred on the polarization curves in 1.0 M NaNO 3 or 1.0 M Na2 SO4 . Therefore, the passive film on nickel was stable in 1.0 M NaNO 3 and 1.0 M Na2 SO4 at pH 10.5, without undergoing localized breakdown. In 1.0 M NaC1, the anodic current exhibited a sudden increase at the potential of about +0.3 V (SCE), and exhibited a large current hysteresis loop during the reverse potential scan on the polarization curve, indicating the localized breakdown of the passive film and the initiation of pitting.  1.6 1.4 1.2 1 0.8 L.1.1^0.6 0 U)^0.4 > LLI^0.2 0 -0.2 -0.4 -0.6 -0.8  0  ^  2  ^ ^ ^ 4 8 6  LOG(i), nA/cm 2  Figure 33 Potentiodynamic polarization curves of Ni in 1.0 M NaNO 3 , 1.0 M Na2 SO4 and 1.0 M NaC1 at pH = 10.5  1.6 1.4 1.2 1 0.8 w 0.6 0 CO^0.4 >  ul^0.2 0 -0.2 -0.4 -0.6 -0.8  0  2  4  6  8  LOG(i), nA/cm 2  Figure 34 Potentiodynamic polarization curves of Ni in 1.0 M Na2SO 4 and 1.0 M NaC1 at pH = 2.5  RESULTS^  89  Figure 34 presents polarization curves of Ni in 1.0 M Na2 SO4 and 1.0 M NaC1 at pH 2.5. There was an active/passive behavior in 1.0 M Na2 SO4 . The passive region extended to as high as the oxygen evolution potential (E (021H 2 0) = 0.840 V (SCE) at pH 2.5, calculated from Equation 5.2)., and no current hysteresis occurred on the polarization curve. No localized film breakdown was observed. However, in 1.0 M NaC1 at pH 2.5, the passivity was fully destroyed by the chloride species, and only active behavior was observed. The results suggest that sulfate and nitrate are not aggressive ions and do not induce pitting corrosion of nickel. These types of anions, especially nitrate, are considered to be inhibitors to pitting corrosion. On the other hand, the chloride ions are considered to be aggressive and are needed for the breakdown of the passive film and initiation of the pitting corrosion on nickel.  5.2.2 pH Effect on Pitting of Nickel in 0.1 M NaC1 In order to study the effect of the bulk solution pH on pitting corrosion behavior, pitting scan tests were carried out in several 1.0 M NaC1 solutions spanning the pH range from 2.5 to 14. An anodic polarization curve shows that Ni exhibited active behavior in 1.0 M NaC1 at pH 2.5 (Figure 35), indicating that Ni underwent general corrosion. An active/passive transition occurred in the solution at pH 4.5, followed by pitting at a higher potential (Figure 36). Pitting occurred on Ni in the solutions in the pH range from 6.5 to 12.5, but no active/passive transition behavior was observed (Figure 36 and 37). The polarization curve in the solution at pH 14 (Figure 38) shows that no passivity breakdown was detected up to the potential of 0.60 V (SCE), higher than the oxygen evolution potential (E (0 21H 2 0) = 0.160 V (SCE) at pH 14, calculated from Equation  RESULTS^  90  (5.2). The results showed a transition in the polarization behavior of nickel from full activation (Figure 35) to localized film breakdown (Figures 36 and 37), and then to fully stable passivation (Figure 38) as the bulk solution pH changed from 2.5 to 14. The effect of bulk solution pH on the critical pitting potentials is summarized in Figure 39. It is evident that the critical pitting potential is not influenced by pH in the region of 4.5 - 10.5, but there is an increase in the critical pitting potential at pH 12.5 when the solution becomes strongly alkaline. Pitting corrosion of nickel is completely retarded up to the potential of +0.60 V (SCE) when the bulk solution pH increases to 14.  5.2.3 Effect of Nitrate Inhibitors in Chloride Solutions Nitrate is usually considered to be a pitting corrosion inhibitor. Therefore, the addition of nitrate species to the chloride solution might be expected to raise the critical pitting potential. Pitting scan tests on nickel in 1.0 M NaC1, pH 10.5, containing additions of 0.001 M, 0.01 M and 0.1 M NaNO 3 are shown in Figure 40, and the resulting critical pitting potentials are plotted in Figure 41. Comparison of Figures 40 and 41 with the results obtained at the same pH in the absence of nitrate addition (see Figures 36 and 39) shows that nitrate additions have no effect on E cp when present in smaller concentrations of 0.001 M and 0.01 M. However, the addition of 0.1 M nitrate raised the critical pitting potential by 90 mV.  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 ru^0.4 0 (/)^0.3 > 0.2  u.i  0.1 0  -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0  ^  2  ^  4  ^  6  ^  LOG(i), nA/cm 2  Figure 35 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 2.5  8  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6 0  ^  2  ^  4  ^  6  ^  8  LOG(i), nA/cm 2  Figure 36 Potentiodynamic polarization results for Ni in 1.0 M NaC1 at pH = 4.5, 6.5, 8.5 and 10.5  1.2 1.1 1 0.9 0.8 0.7 0.6 W ^0.5  0 CI)^0.4 >^0.3 Lli^0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0^ 2^ 4^ 6^ 8  LOG(i), nA/cm  2  Figure 37 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 12.5  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 W 0.4 C.) u) 0.3 > 0.2 LII^0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6 -0.7 -0.8  0  ^  2  ^  4  ^  6  ^  LOG(i), nA/cm 2  Figure 38 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 14.0  8  0.8  no pitting  0.7  4  A  0.6  I  Ca 0.5 (...) C/) Ci. U W  0.4  1 i^i^A^ A  0.3  0.2  0.1  0  I^I^I^I^I^I^I^I^I^I^I  4  ^ ^ ^ ^ ^ 6 8 10 12 14  pH  Figure 39 Effect of bulk solution pH on the critical pitting potential of Ni in unbuffered 1.0 M NaC1 solution  ▪  1.2 1.1 1 — — 1.0 M NaCI + 0.1 M NaNO3  0.9  --- 1.0 M NaCI + 0.01 M NaNO3  0.8  1.0 M NaCI + 0.001 M NaNO3  0.7 0.6 •  0.5  •  co  0.4  •  0.3  ■■■  No N. •  r-—  Lir^0.2 0.1 0 -0.1  I  r.  -0.2 -0.3  m.■^  -0.4  go*  -0.5 -0.6  ■••••• "".  ■■=.^  ••••• "  ••••.'  0  2  4  6  8  LOG (i), nA/cm 2  Figure 40 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaNO 3 at pH = 10.5  0.7  0.6  0.5  0.3  0.2  0.1  -3  -2  -1  LOG[C] , M  Figure 41 Effect of nitrate concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5  RESULTS^  98  5.2.4 Effect of Buffers in Chloride Solutions The effect of buffer concentration on the pitting behavior of nickel in 1.0 M NaC1 solution was investigated at pH 10.5. Two types of buffer solutions were used based on (1) NaHCO 3/Na2 CO 3 additions and (2) Na2HPO4/Na3 PO4 additions. They will be referred to as carbonate and phosphate buffers respectively. The polarization behavior in the carbonate buffer is shown in Figure 42 for total NaHCO 3 /Na2 CO 3 concentrations ranging from 0.001 M through 0.01 M to 0.1 M. The corresponding critical pitting potentials are plotted as a function of buffer concentration in Figure 43. The results clearly show that the increase in the buffer concentration, while maintaining a constant pH of 10.5, raised the critical pitting potential. Comparison with Figure 39 at pH 10.5 shows that even the smallest buffer addition of 0.001 M raised E cp by about 30 mV, and the highest addition of 0.1 M increased Ecp by about 200 mV. Polarization curves in the phosphate buffered solutions are shown in Figure 44 for total Na 2HPO 4/Na3 PO4 additions of 0.001 M, 0.01 M and 0.1 M. The corresponding E SP potentials are plotted as a function of buffer concentration in Figure 45. It is clear that the phosphate buffer behaves in a very similar manner to the carbonate buffer, the E  cp  increased with increasing buffer concentration. Some comparative tests were conducted with a 1.0 M NaBr solution, pH 10.5, in the absence and presence of phosphate buffer having a total Na 2HPO4/Na3 PO4 concentration of 0.1 M. The polarization results are shown in Figure 46. Again, in a similar manner to the chloride solution, the critical pitting potential was raised about 250 mV in the buffered bromide solution.  1.2 1.1 In 1.0 M NaCI + 0.1 M NaHCO3 /Na 2 CO 3 — In 1.0 M NaCI + 0.01 M NaHCO3 /Na 2 CO 3  1 0.9  — - — - • In 1.0 M NaCI + 0.001 M NaHCO3 /Na 2 CO3  0.8  pH = 10.5  0.7 0.6 0.5 UT 0.4 0 co^0.3 >^0.2 Li 0.1  . ■ '.^'' .^. .■ . ■ . ... .^ 0,00 . . .m.... m.....  .^. •••• ''.  0  ''''  .7  .i .'"  •••• •  ......•  m... l =  -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  1  1^$1......_ I  0  2  4  ^  6  ^  8  LOG(i), nA/cm 2  Figure 42 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaHCO3/Na2 CO 3 buffer at pH = 10.5  0.7  0.6  In 1.0 M NaCI +x M NaHCO 3 /Na 2 CO3 pH = 10.5  0.5  0.4  0.3  0.2  0.1  -3  -2  -1  LOG[C] , M  Figure 43 Effect of carbonate buffer concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5  ^  1.2 1.1  ^ In 1.0 M NaCI + 0.1 M Na 2 HPO4 /Na3 PO4 In 1.0 M NaCI + 0.01 M Na 2 HPO 4 /Na3 PO4  1  In 1.0 M NaCI + 0.001 M Na 2 HPO4 /Na3 PO4  0.9  0.8 —^pH = 10.5 0.7 0.6 0.5 0 0.4 0.3  LLI^0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0  2  4  6  8  LOG(i), nA/cm 2  Figure 44 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M Na2HPO 4/Na 3PO4 buffer at pH = 10.5  0.7  0.6  In 1.0 NaCI + x M Na 2HPO4/Na 3 PO4 pH = 10.5  ■ ■  0.5  um  0.3  0.2  0.1  IIILIJ_I_LI_LIIIIIIIIIII  -3  ^  -2  ^  -1  LOG[C] , M  Figure 45 Effect of phosphate buffer concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5.  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 CCI 0 0.4 U)  > W-  0.3 0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0  ^  2  ^  4  ^  6  ^  8  LOG(i), nA/cm 2  Figure 46 Potentiodynamic polarization results for Ni in unbuffered 1.0 M NaBr and Na2HPO 4/Na3 PO4 buffered 1.0 M NaBr at pH = 10.5  RESULTS^  104  5.2.5 Comparison of Inhibitor with Buffers The common feature of carbonate and phosphate buffers, and nitrate inhibitor in solutions, when present in sufficient concentrations, is that they raise the critical pitting potentials. Their relative effects on the increase of the critical pitting potential, AE cp , in 1.0 M NaC1 at pH 10.5 are compared as a function of their concentrations in Figure 47 and listed in Table 12. The 6E cp is defined as Ecp - ecp , where ecp is the critical pitting potential in the simple unbuffered and uninhibited 1.0 M NaC1 solution at pH 10.5. Consequently, Ak p accounts for the net effect of the additives to 1.0 M NaCl solution at pH =10.5. It is evident from Figure 47 and Table 12 that increasing concentrations of additives raise AE cp , and that the buffers are much more effective in raising the critical pitting potential than nitrate additions. Furthermore, the phosphate buffer tends to produce the highest values. The results indicate that the control of pH by buffering the solution is an important means of controlling pitting, more important than the use of nitrate inhibitors.  RESULTS  ^  105  Table 12. Critical pitting potentials (E, p) and AE, p , for polycrystalline nickel in chloride solutions at pH 10.5.  Type  No.  Additives to solution  Eep, V (SCE)  Type II  B-7  -  0.295 (Ecp )  B-10  0.1 M NaNO 3  0.400, 0.370  0.105, 0.075  Type III B-11  0.01 M NaNO 3  0.310, 0.295  0.015, 0.000  B-12  0.001 M NaNO 3  0. 295, 0.230  0.000, 0.005  B-13  0.1 MNaHCO 3/Na2 CO 3  0.530, 0.500, 0.480  0.235, 0.205, 0.185  B-14  0.01 M NaHCO 3 /Na2 CO3  0.385, 0.380, 0.400  0.090, 0.085, 0.105  B-15 0.001 M NaHCO 3/Na2 CO3  0.345, 0.340  0.050, 0.045  B-16  0.1 M Na2 HPO4/Na3 PO4  0.565, 0.540  0.270, 0.245  B-17  0.01 MNa2 HPO4/Na3 PO4  0.440, 0.390  0.145, 0.095  °  AE,p, V -  Type IV  B-18 0.001 M Na 2 HPO4/Na3PO4 0.375, 0.355, 0.34  0.080, 0.060, 0.045  0.4  0.35  A 1.0 M NaCI + x M Na 2 HPO 4 /Na3 PO4 * 1.0 M NaCI + x M NaHCO3/Na2CO3  0.3  • 1.0 M NaCI + x M NaNO 3 A  6.1^0.25  A 0  0 U)  a 0 w  0  0.2  0 0.15  A A 0  0.1 A 0.05  0  g  ■ ■  A  a ■ ^ I^I^1^i^I^I^l_^I^ii^1^I^lit^tit^I^1 I^ 1 ^ ^  -3  -2  -1  LOG[C] , M  Figure 47 Comparison of the increment in the critical pitting potential CAE& with the presence of NaNO 3 , NaHCO 3INa2CO 3 and Na2HPO 4/Na3 PO 4 in 1.0 M NaC1 at pH 10.5  RESULTS^  107  5.2.6 Halide Ion Effect Some pitting scan tests were performed on nickel in different types of halide solutions to compare the relative influence of halides on passivity breakdown. Figure 48 shows the polarization curve in 1.0 NaBr, pH 10.5, along with the polarization curve in 1.0 M NaC1 at the same pH as a comparison. Pitting corrosion occurred in 1.0 M NaBr solution at pH = 10.5, but the critical pitting potential was about 100 mV higher than that in 1.0 M NaCl solution at the same pH. A similar result showing about a 50 mV difference in the critical pitting potential was obtained in the phosphate buffered bromide solution as compared with the phosphate buffered chloride solution (Figure 49). These results show that the passive oxide film on nickel is more stable in the bromide solution than in the chloride solution. The polarization results of Ni in 1.0 M NaF solution at pH = 10.5 and 6.0 are shown in Figures 50 and 51 respectively. The passive film is stable in 1.0 M NaF at pH 10.5, and no pitting occurred, as compared with Ni in 1.0 M NaC1 at the same pH. At pH 6.0, an extensive region of active behavior was observed in 1.0 M NaF, followed by the onset of passivation near +0.5 V (SCE). This passive film was stable at higher potentials and pitting was not observed. For comparison, at pH 4.5, passivation occurred at a lower potential near -0.2 V (SCE) in 1.0 M NaC1, but film breakdown and pitting occurred at a potential of +0.3 V (SCE). Figure 52 presents a polarization result for nickel in 1.0 M HF at pH 3.1. The result shows that Ni exhibited an active dissolution behavior in a similar manner to Ni in 1.0 M NaCl at pH 2.5 (Figure 35).  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0  ^  2  ^  4  ^  6  ^  LOG(i), nA/cm 2  Figure 48 Polarization results for Ni in 1.0 M NaC1 and 1.0 M NaBr at pH = 10.5  8  •  1.2 1.1  ^  1.0 M NaCI + 0.1 M Na2HPO4/Na 3 PO4 1.0 M NaBr + 0.1 M Na2 HPO 4 /Na 3 PO4  0.9 0.8 0.7  w  0.6 0.5 0.4 0.3  •  0.2  LLI^  0.1 0  -0.1 -0.2 -0.3 -0.4 -0.5 -0.6 0  2  4  6  8  LOG(i), nA/cm 2  Figure 49 Polarization results for Ni in Na 2HPO 4/Na3 PO4 buffered 1.0 M NaC1 and 1.0 M NaBr at pH 10.5  1.6 1.4 1.2 1 0.8 0.6 0 0.4 1.11  0.2 0 -0.2 -0.4 -0.6 -0.8  0  ^  2  ^  4  ^  LOG(i), nA/cm  6  ^  2  Figure 50 Polarization curves for Ni in 1.0 M NaF and 1.0 M NaC1 at pH 10.5  8  1.6 1.4  in 1.0 M NaF at pH 6.0  1.2  ---------------^in 1.0 M NaCI at pH 4.5  1 0.8 w 0.6 C.) 0.4  LLl^0.2 0 -0.2 -0.4 -0.6 -0.8  0  2  4  LOG(i), nA/cm  6  8  2  Figure 51 Polarization curves for Ni in 1.0 M NaF at pH 6.0 and 1.0 M NaC1 at pH 4.5  1.2 1.1 1 0.9 0.8 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0 -0.1 -0.2 -0.3 -0.4 -0.5 -0.6  0  2  4  6  LOG(i), nA/cm 2  Figure 52 A polarization result for Ni in 1.0 M HF at pH 3.1  8  RESULTS^  113  5.2.7 Summary 1. Pitting corrosion of nickel occurs in specific solutions, usually in halide - containing solutions. Therefore, halides play a very important role in the breakdown of passivity and the initiation of pits. Without the presence of aggressive ions in the solution, nickel would stay passive as observed in NaNO 3 and Na2 SO4 solutions. 2. Nickel exhibited active dissolution behavior in 1.0 M NaC1 at pH 2.5. Pitting occurred on Ni in 1.0 M NaC1 in the pH range of 4.5 - 12.5. The critical pitting potential, E cp , was independent of the bulk solution pH in the range from 4.5 to 10.5, but E cp increased at pH 12.5. No pitting was observed in the solution with pH 14. It is worth noting that Ni exhibited a transition in polarization behavior from active dissolution at pH 2.5, through active-passive/localized breakdown at pH 4.5 and passive/localized breakdown at pH 6.5 - 12.5, to fully passive behavior at pH 14 in 1.0 M NaCl. 3. The presence of the Na2 CO 3/NaHCO 3 and Na 2 HPO 4/Na3 PO4 buffers raised the critical pitting potential in 1. 0 M NaCl. The E el, increased with increasing buffer concentration, while maintaining the same bulk solution pH at 10.5. The additions of carbonate and phosphate buffers were more effective in raising the critical pitting potential than the addition of the nitrate inhibitor. 4. Nickel exhibited fully active, active/passive or fully passive behavior in F solutions, depending on pH. No localized film breakdown (pitting) was observed in F solutions. This is in contrast to Cl and Br solutions where pitting was observed on Ni. Bromide -  -  ions were shown to be less aggressive than chloride ions.  E-pH and X-pH DIAGRAMS  ^  114  6 E-pH and X-pH DIAGRAMS  6.1 Construction of E-pH and X-pH Diagrams  6.1.1 Chemical Equilibrium For a general chemical reaction at equilibrium that does not involve any electron transfer and any change in the oxidation states of reactants and products, we may write: aA + bB = cC + dD^  (6.1)  An example being: fr + OH - = H2 O  (6.2)  The equilibrium constant K for Equation (6.1) can be expressed as:  K  [CT [Di d —  [A  r [Br  (6.3)  where the square brackets represent the activity of the species. The standard free energy change, AG°, is the algebraic sum of standard chemical potentials, g,° of individual species. AG° = ciLe c + dp,' D - al.C A - bp.° B^(6.4) At the chemical equilibrium, we have the relation between AG ° and K:  E-pH and X-pH DIAGRAMS  ^  AG ° = —RT1nK  115  ^  (6.5)  6.1.2 Electrochemical Equilibrium Electron transfer is involved in an electrochemical reaction. A single electrode reaction, or a half cell reaction, is usually used to represent the electrochemical equilibrium between the oxidized and reduced forms of species. A simple example is the reduction of hydrogen ions. 211+ + 2e = 112^(6.6) For a general electrochemical equilibrium between oxidized and reduced species: 1L + jH + ne - = pP + qQ^  (6.7)  Similar to the chemical equilibrium, the standard free energy change for the electrochemical reaction (6.7) becomes: AG° = pjf + qi.t° Q -^- ji_t° H  (6.8)  At the electrochemical equilibrium, we have the Nernst equation:  E = E°  RT  ln  [P]P [Q]q  (6.9)  nF [L] l [HP  Where, E is the equilibrium electrode potential, F is the Faraday constant, and E° is the standard electrode potential, which is governed by the following equation: AG° = -nFE°^  (6.10)  E-pH and X-pH DIAGRAMS^  116  6.1.3 E-pH Diagrams An E-pH diagram shows the effects of potential and pH on the stability of individual species (metals, ions, oxides etc.). Each line in the diagram represents a chemical or an electrochemical equilibrium. The regions between the lines show the domain of stability of each species. The diagrams are plotted with pH as the abscissa and E as the ordinate. In an E-pH diagram there are three types of lines which represent  three types of equilibria. (1) Chemical equilibrium involving fr ions: For a general chemical equilibrium: aA + jH+ = cC + dD^  (6.11)  We have: AG ° = —RT1n  (6.12)  r c r [D i d [A]a [HI  and (6.13)  AG° pH = — 71 (Log rCic [Did [A]a^2.303RT  Therefore, dpH 0 dE  (6.14)  E-pH and X-pH DIAGRAMS^  117  The equilibrium (6.11) is independent of potential, but dependent on pH, and is represented by a vertical line in the E-pH diagram. (2) Electrochemical equilibrium without 11 ÷ ions For a general electrochemical reaction: IL + ne = pP + qQ^  (6.15)  According to the Nernst Equation (6.9): E = E° —  RT [Q l 4 In nF^[L,11  (6.16)  and dE/dpH = 0^  (6.17)  The equilibrium (6.15) is independent of pH, but dependent on potential, and is represented by a horizontal line in the diagram. (3) Electrochemical equilibrium involving fr ions For a general electrochemical reaction: IL + j1I+ + ne = pP + qQ^  (6.18)  we have: E = E°  RT  nF  in  WV [QV  ^  (6.19)  E-pH and X-pH DIAGRAMS^  118  and dE^2.303jRT _ dpH^nF  (6.20)  —  Therefore, the equilibrium is dependent on both potential and pH , and is represented by a sloping line in the diagram. When 1-1 + appears on the left hand side of the equilibrium equation, the line has a negative slope. The line has a positive slope when fr appears on the right hand side.  6.1.4 X-pH Diagrams The E-pH diagram is a very concise presentation of multiple equilibria and shows the thermodynamic stability of individual species for H 2 O - metal systems. However, for more complicated systems, such as solutions with halides, variables other than pH and electrode potential will be involved. The conventional E-pH diagrams have some limitations in the presentation of these complicated systems. For the purpose of this study on passivity and its breakdown, the X-pH diagram has been introduced , for the first time, to show the thermodynamic stability of the specific oxide' in the system where the formation of halide - metal complexes is involved. The X-pH diagram presents the effects of solution pH and halide (X) activity on the equilibrium between a specific oxide and halide complexes. Since the formation of a stable oxide film is influenced by the halide complexes, X-pH diagrams are anticipated to be particularly useful in a study of the passivity breakdown of metals in halide environments. The diagrams are constructed  1 The term "oxide" in the treatment including hydrated oxides.  E-pH and X-pH DIAGRAMS ^  119  with pH as the abscissa and X as the ordinate. Each line in the X-pH diagram represents an equilibrium between the oxide and a soluble species, where the soluble species may be a halide complex. If the equilibrium between oxide and the complex is a chemical one, the X-pH diagram will be independent of the electrode potential, which is the case for Zn(OH) 2 , Ni(OH) 2 and Al 2 0 3 .3H 2 0. If the equilibrium between the oxide and complex is an electrochemical one, the X-pH diagram should be constructed at a defined potential (such as in the case of Ni 2 03 ). This is the limitation in the application of X-pH diagrams. There are also three types of lines in X-pH diagrams: (1) Vertical: an equilibrium without the involvement of halide complexes. M(OH) n + n1-1 + = M' + nH 2 O or M(OH) n = MO' + nfr  (2) Horizontal: an equilibrium without the involvement of fr ions mn+ + mx- = mx in(m-n)-  (3) Sloping: an equilibrium involving both 11 + and halide complexes M(OH) n + nH + + rriX - = MX,n (m-n )- + nH 2 O A limited number of X-pH diagrams have been constructed for the hydrated oxides Zn(OH) 2 , Ni(OH) 2 and Al 2 0 3 .3H 2 0 systems in the interest of this study. For details, refer to Section 6.2.  E-pH and X-pH DIAGRAMS^  120  6.1.5 Thermodynamic Data and Assumption of Activity of Species In the construction of the diagrams, we have to know the standard chemical potential, [C, for each individual species. Most thermodynamic data used in the following section (Section 6.2) were obtained from Pourbaix's Atlas of Electrochemical Equilibria in Aqueous Solutions  [201  , and others are from NBS publications [1391 and  Critical Stability Constants [140] . Data from Pourbaix's Atlas  [20]  are not noted with a  reference number in the following Tables 13 - 17. Data from other sources are noted with their reference number in Tables 13 - 17. Some thermodynamic data of halide complexes are not available in the form of standard chemical potentials, but they are given in terms of a stability constant. Therefore, the standard chemical potential has been calculated from the stability constant based on Equation (6.5). For example, K1 is known for the following chemical equilibrium: F + Ni2+ = Nir , where logKi = 1.1 [140] Therefore, the standard chemical potential for NT + is calculated: g° = -330,996 J/mol. The activities of solid substances are assigned a value of unity. The fugacity of gaseous substances is taken as 1.0. The activities of halides are assumed to be 10 -2 , 10 -1 , 10° and 10 1 . Activities of other dissolved species are arbitrarily set at 10 -6 , 10-4 and 10-2.  E-pH and X-pH DIAGRAMS^  121  6.2 Diagrams for H 2 0-Metal and H2 O-Halide-Metal Oxide Systems Since aqueous solutions are involved in the systems, the stability of water is shown in every E-pH diagram. The stability of water is determined by two electrochemical reactions (a) and (b). Standard chemical potentials of species considered are listed in Table 13. 2H + + 26 = H2^  (a)  0 2 + 411- + 4e = 2H2 o^  (b)  -  Table 13. Standard chemical potentials of substances at 25 °C Substance  ^  State^ji°, KJ/mol.  H+^aq^0 OH -^aq^-157.297 H2 O^1^-237.180 H2^g^ 0 02^g^ 0  Equilibria (a) and (b) are shown in every E-pH diagram as lines (a) and (b) respectively. In the cases of Al, Zn and Ni, hydrated oxides are thermodynamically more stable than anhydrous oxides based on the thermodynamic data reported in the Pourbaix's Atlas  [201  .  Therefore, the hydrated forms of these oxides are considered 2 . However, for Sn, the  and by Bard 1142j , anhydrous oxides of Zn and Ni (ZnO and NiO) are slightly more stable than hydrated oxides (Zn(OH) 2 and Ni(OH) 2). The reported data are: p°(ZnO) = -318.32 KJ/mol, [t°(Zn(OH) 2) = -555.13 KJ/mol, µ°(NiO) = -216.0 KJ/mol and if(Ni(OH) 2 ) = -453.1 KJ/mol. The free energy changes are very small, thus the hydrated oxides are still considered. 2 However, according to the thermodynamic data reported by NBS  rmi  E-pH and X-pH DIAGRAMS^  122  anhydrous oxides are most stable and should be considered in preference to the hydrated oxides. Nevertheless, in order to be consistent with the Al, Zn and Ni diagrams, hydrated Sn-oxides are considered in this Section, and the anhydrous oxides are presented in Appendix II. New E-pH and X-pH diagrams, involving the formation of metal halide complexes, are constructed and presented together with the conventional Pourbaix E-pH diagrams for H2 O - metal systems. The inclusion of halide complexes in E-pH and X-pH diagrams and their applications to pitting are a new and unique contribution of the present research.  6.2.1 Sn The formation of stannous hydroxide Sn(OH) 2 and stannic hydroxide Sn(OH) 4 on the tin surface is considered in the aqueous solution, and their stability determines the passivity of tin. In the chloride solution, Sn - chloride complexes exist in the form of SnC1 4 2- , SnC1 6 2- and SnC1 3 - . The standard chemical potentials of the species considered in the Sn-Cl_-H2 0 system are listed in Table 14.  Two E-pH diagrams have been constructed for H 2 O - Sn and H 2 O - Cl - - Sn systems respectively (Figures 53 and 54). Shaded areas in the diagrams are the passive zones - the domain of hydrated tin oxides. Comparing the two E-pH diagrams, the passive zone is narrowed by the chloride complex formation when chloride ions are present in the solution, indicating the instability of passivity in the C1 - -containing solution.  E-pH and X-pH DIAGRAMS^  Table 14. Standard chemical potentials for Sn systems at 25 °C Substance  ^  State^KJ/mol.  Sn^s^0 Sn(OH)2^s^-492.040 Sn(OH)4^s^-951.850 Sn2+^aq^-26.255 Sn4+^aq^2.720 HSn02^aq^-410.030 Sn2 03^aq^-590.280 SnC142-^aq^-560.910E1401 SnC1 62-^aq^-785.800E'4°J SnC13^aq^-431.495E1401 Cl -^aq^-131.168  123  Tin - water system  1.4 1.2 1 0.8 0.6 0.4 0.2 0 -0.2 -0.4 -0.6 -0.8 -1 -1.2 -1.4 0^2^4^6  8  10  12  14  pH  Figure 53 E-pH diagram for H 2 O - Sn system; activities of all solute species at 10  -2  , 10 -4 , 10-6  Sn - chloride - water system 1.4 1.2 1 0.8 0.6 0.4 0.2 0 -0.2 -0.4 -0.6 -0.8 -1 -1.2 -1.4  0  ^ ^  2  6  ^ ^ ^ ^ 12 14 8 10 pH  Figure 54 E-pH diagram for H 2O - Cl - - Sn system; activities of chloride at 10 ° and 10 1 ; activity of other solute species at 10-6  E-pH and X-pH DIAGRAMS^  126  6.2.2 Zn The passivity domain of zinc normally corresponds to the formation of Zn(OH) 2 . However, insoluble ZnCO 3 may contribute to the passivation of zinc when CO 3 - is present in the solution. Four types of zinc - chloride and bromide complexes have been reported in chloride and bromide solutions respectively (ZnX. (m -2)- , where m=1,2,3 and 4). For complex formation, the reaction may be written: zn2+ + mX =  znxm( n-2)-  and the corresponding stability constant, K.= [ZnX.(m -2) ]1[Zn21[X]m . It has been reported that logK, = 0.49, logK2 = 0.61, logK3 = 0.5 and logK4 = 0.2 for ZnClm(m-2)- [ 140], and that logKI = 0.22, logK2 = -0.10, logK3 = -0.74 and logK4 = -1.00 for ZnBrn.,(m-2)- [ 142]. Standard chemical potentials of involved substances are listed in Table 15. Three E-pH diagrams have been plotted for H2 O - Zn, H 2 O - CO 3 - - Zn and H2 O Cl - - Zn systems (Figures 55 - 57), and one X-pH diagram for H 2 O - Cl - Zn(OH) 2 -  system (Figure 58). The formation of insoluble ZnCO 3 extends the passive zone of zinc to the low pH region in the CO 3 - -containing solution. In the chloride system (Figure 57), the formation of ZnC1 4 2- diminishes the passive zone of zinc. Therefore, the passive film of zinc is less stable when chloride ions are present in the solution. Other forms of Zn chloride complexes (such as ZnCr, ZnC1 2 and ZnC1 3 - ) are also formed, depending on the chloride concentration, but it is difficult to introduce more than one chloride complexes in the E-pH diagram. The effect of chlorides on zinc passivity can be clearly illustrated in the X-pH diagram (Figure 58). The passive zone becomes narrower with increasing chloride activity.  E-pH and X-pH DIAGRAMS^  Table 15. Standard chemical potentials of substances for zinc systems at 25 °C Substance^State^if , KJ/mol. Zn^s^0 Zn(OH)2 ^s^-559.108 Zn 2+^aq^-147.210 HZn0 2 -^aq^-464.005 Zn0 2 2^aq^-389.238 ZnCr^aq^-280.830E14°1 ZnC1 2^aq^-413.028E14°1 ZnC1 3 -^aq^-543.570E1401 ZnC14 2-^aq^-673.026E14°1 Cl -^aq^-131.168 ZnC 03^ s^-731.363[139]  127  Zinc-water system 0.8 0.6 0.4 0.2  -1.2 -1.4 -1.6 -1.8 0  2  4  6^8  10  12  14  pH  Figure 55 E-pH diagram for H 2 O - Zn system; activities of all solute species at 10-2 , 10-4 , 10-6  Zinc - water - Carbonate system 0.8 0.6 0.4 0.2 0 III  1  -0.2 U)^. > -0.4 LL ^  -0.6  -0.8 -1 -1.2 -1.4 -1.6 -1.8  0  ^ ^ ^ ^ ^ ^ 2 4 6^8 10 12 14  pH  Figure 56 E-pH diagram for H 2 O - CO 3 2- - Zn system; Activity of H 2 CO 31HCO 3 -/CO 3 2- at 10-1 ; activities of other solute species at 10 -2 , 10 -4 , 10-6  Zinc-chloride-water system  1 0.8 0.6 0.4  [cr]^10. 3 _  0.2  ZnCI 4 (1 06 )  10  1 2-^_ HZnO 2 (1 06 )  0 -0.2  „de-- Zn(OH) 2 > -0.4 -0.6 -0.8 -1 -1.2 -1.4  Zn  -1.6 -1.8 0  ^ ^ ^ ^ ^ ^ 2 4 6^8 10 12 14  pH  Figure 57 E-pH diagram for H 2 O -^- Zn system; activities of chloride at 10 °3 and 10 1 ; activity of other solute species at 10-6  Zn(OH)2 - chloride - water system  2  1.5  1  0.5  O  0  -0.5  -1  -1.5  -2  0  ^ ^ ^ ^ ^ ^ 2 4 6^8 10 12 14  pH  Figure 58 X-pH diagram for H2 O - Cl - - Zn(OH) 2 system; activity of all solute species at 10 -2 , 10-4 , 10-6  E-pH and X-pH DIAGRAMS^  132  6.2.3 Al Standard chemical potentials of all species considered are listed in Table 16 for Al systems. The hydrated Al oxide (Al 2 0 3 .3H 2 0) is considered to be the oxide formed on aluminum in aqueous solutions. Complexes have been reported in all types of halide solutions. Al - fluoride complexes with from one to six F ions have been reported [mi. But for Al - chloride and Al - bromide complexes, thermodynamic data are only available for A1C1 3 and AlBr 3  "  391 .  Figures 59 and 60 show the constructed E-pH diagrams for H 2 O- Al and H2 O - C1 - Al systems respectively. The formation of chloride complexes destroys the stability of the passivity of aluminum in the low pH region. The aggressiveness of halides can be seen based on the X-pH diagrams for H 2 O - F - Al 2 0 3 .3H2 0 and H 2 O - CF - Al 2 0 3 .3H2 0 systems (Figures 61 and 62). The formation of a very stable fluoride complex almost totally destroys the passivity of aluminum. The passive zone in the chloride solution is narrowed with increasing chloride activity.  E-pH and X-pH DIAGRAMS ^  Table 16. Standard chemical potentials of substances for Al systems at 25 °C Substance  ^  State^1.1,°, KJ/mol.  Al^s^0 Al 2 0 3 .3H 2 0^s^-2320.450 A1 3-'^aq^-481.160 A102^aq^-839.770 A1C1 3^aq^-878.640E1393 Cl-^aq^-131.168 A1F2-1.^aq^-792.490E14°3 A1F2+^aq^-1097.550E14°3 A1F3^aq^-1396.163E1403 A1F4^aq^-1689.755E14°3 A1F52-^aq^-1974.220[1403 A1F63-^aq^-2252.984E14°3 A1Br 3^aq^-794.960E139] F^ aq^-276.480  Br-^aq^-103.970  133  Al - Water system  0^2^4^6  8  10  12  14  pH  Figure 59 E-pH diagram for H 2 O - Al system; activities of all solute species at 10 -2  , 10 -4 and 10-6  Al - Chloride - water system  1  0.5  0  -------_  p i] = 10 0  -0.5 AICI 3 (10 -6 )  -1.5  -2 —  -2.5 —  -3  Al  I^I^I^i^1^I^I^I^1^I^I^I^1^I^I  0^2^4^6^8^10^12^14  pH  Figure 60 E-pH diagram for H 2 O - cr - Al system; activities of chloride at 10 ° and 10'; activity of other solute species at 10-6  Al 2 03 • 3H 2 0 - flouride - water system 2  1.5  1  0.5  -0.5  -1  -1.5  0  2  4  6^8  10  12  14  pH  Figure 61 X-pH diagram for H 2 O -F - Al 2 0 3 .3H 2 0 system; activity of all solute species at 10 -2 , 10-4 , 10-6  Al oxide - Chloride - water system  2  1.5  0.5  0 0  -0.5  -1  -1.5  -2  0  ^ ^ ^ ^ ^ ^ 2 6^8 4 10 12 14  pH  Figure 62 X-pH diagram for H 2 O -Cr - Al2 0 3 .3H2 0 system; activity of all solute species at 10 -2 , 10 -4 , 10-6  E-pH and X-pH DIAGRAMS^  138  6.2.4 Ni For the construction of nickel diagrams, the following standard chemical potentials have been used (Table 17). Several stable oxides (Ni(OH) 2 , Ni 2 0 3 and Ni 3 04 ) are formed on nickel in aqueous solutions. Halide complexes have been found in the form of NiF+ , NiC1 4 , NiC1 2 and NiBe.  Table 17. Standard chemical potentials used for nickel systems at 25 °C Substance  ^  State^jr, KJ/mol.  Ni^s^0 Ni(OH) 2^s^-453.127 Ni3 0 4^  s^-711.910  Ni 2 0 3^s^-469.738 Ni0 2^s^-215.140 Ni 2÷^aq^-48.242 HNi02^aq^-349.197 NiCr^aq^-181.293E1411 NiC1 2^aq^-314.992E1411 Cl -^aq^-131.168 NiF4-^aq^-330.996E14°1 NiBe^aq^-151.530E14°E F+^aq^-276.480  E-pH and X-pH DIAGRAMS^  139  Two E-pH diagrams have been constructed for H 2 O - Ni and H 2 O - Cl - Ni -  systems respectively (Figures 63 and 64). According to the E-pH diagram for the H 2 O Cl - Ni system (Figure 64), high activity of chlorides will narrow the passive zone of -  nickel relative to the E-pH diagram for the H 2 O - Ni system (Figure 63). In the X-pH diagrams for H2 O - F - Ni(OH) 2 and H 2 O - Cl - Ni(OH) 2 (Figures 65 and 66), it is shown -  that the passive zones are diminished with increasing halide activity.  Ni - H 2 0 system  0.8  11.1 0.6 2  Cn >  al  0.4 0.2  pH  Figure 63 E-pH diagram for H 2 O - Ni system; activities of all solute species at 10-2 , 10 -4 and 10-6  Ni - H 2 0 - 0I system -  1.8 1.6 1.4 1.2  1 0.8 IT U I 0.6 CO > 0.4 Lir^0.2  0 -0.2 -0.4 -0.6 -0.8  -1  0  ^ ^ ^ ^ ^ ^ ^ 2 4 6 8 10 12 14  pH  Figure 64 E-pH diagram for H 2 O - Cr - Ni system; activities of chloride at 10° and 10 1 ; activity of other solute species at 10-6  Ni(OH) 2 - chloride -water system  2  1.5  1  0.5  -1  -1.5  -2 0  2^4  6  8  10  12  14  pH  Figure 65 X-pH diagram for H 2 O -F - NiO system; activities of all solute species at 10 -2 , 10  -4  , 10-6  Ni hydrated oxide - chloride - water system  '—' U_ 0 —J  0  2  4  6  8  10  12  14  pH  Figure 66 X-pH diagram for  H2O -cr - NiO system; activities of all solute species at 10 -2 , 10-4 , 10-6  DISCUSSION^  144  7 DISCUSSION This section will discuss several important aspects, including issues related to the experimental results, the thermodynamic E-pH and X-pH diagrams, and other available information on pitting corrosion. It will focus on factors controlling the pit initiation process, such as the local solution chemistry and halide complex formation. A theory is proposed to attempt to explain the effects of pH, buffers, halides, crystallographic orientations, electrode potential and other factors on pitting corrosion.  7.1 Pit Initiation Stages and Governing Factors Pit initiation may be considered to consist of several stages, as shown in Figure 67. At the first stage, the metal is in a passive state with a full coverage of passive film on its surface. The film is locally broken down when it proceeds to stage 2. Following the initial breakdown of the passive film, changes in the local solution chemistry occur and certain conditions are established in the local region at stage 3. Stable pits are formed at stage 4, which then grow continuously. Stage 4 signifies the transition from pit initiation to pit propagation. There is still considerable debate over which stage critically controls the pit initiation process. Kruger [143] believes that stage 2 is the critical stage which must take place before stages 3 and 4 can proceed. The evidence to support his argument is that pitting resistance is strongly dependant on the properties of the passive film. Metals with passive films which are resistant to the initial breakdown (stage 2) do not undergo pitting corrosion, so stage 2 may control the pit initiation process. The importance of stage 2 is also emphasized by several proposed theories, such as competitive adsorption, ion penetration,  Passive Film  Metal (1)  Breakdown  Metal (2)  Changes in the local chemistry  Metal (3)  Figure 67 Several stages occurring during the pit initiation process  Pit Growth  Metal (4)  DISCUSSION^  146  chemical-mechanical breakdown, and transitional complex formation. All these theories deal with the film breakdown process which is occurring during the transition from stage 1 to stage 2. Therefore, pits will be initiated when stage 2 is reached. However, the initial breakdown of a passive film does not automatically lead to pit initiation, because the breakdown site may be repassivated to reproduce the full coverage of the passive film (back to stage 1). Therefore, both the repassivation process and the breakdown process should be taken into consideration during pit initiation. Consequently, if repassivation can occur at stage 2, it may be argued that stage 3 controls the pit initiation process. Videm's theory  [121]  suggests that pit initiation is determined when the breakdown rate exceeds the repassivation rate in a kinetic breakdown/repassivation process occurring at stage 3. The importance of changes in the local solution chemistry has been stressed in Galvele's local acidification theory  "  191 .  According to Boehni  [1441  , metastable pits are formed at stage 3, which either are  repassivated or proceed to the stable pit stage (stage 4) depending on the development of local chemical environments. Beyond stage 4 pits enter their growth stage. Some of the problems relating to pit initiation may be resolved more clearly by considering the structure of the interface between the metal substrate and the bulk solution. This interface is shown schematically in Figure 68. It consists of an oxide film layer, an electrical double layer, and a diffusion layer. Therefore, recognizing that pit initiation is a localized process occurring in the interfacial region, pit initiation may result from a local inhomogeneity in the oxide film and/or a local chemical inhomogeneity in the diffusion layer. Furthermore, the models of pit initiation must include a role for halide ions, because these are necessary for pitting, as shown in the present study. Defects in the oxide film are assumed to be a governing factor in several models of pit initiation. The assumption of pre-existing defects in the film, so-called priori assumption by  Interfacial Region of Metal/Solution  Figure 68 Structure of the interface between metal substrate and bulk solution  DISCUSSION^ Okada [145] , provides the basis for Wood's flaw model  148 [11]  and Lin's point defect model [1291 •  It is also the rationale for modelling the early development of local solution chemistry at stage 3 without considering the film breakdown process occurring from stage 1 to stage 2 (Figure 67) in pit initiation. However, defects alone present the difficulty of explaining why pitting does not take place in halide-free solutions. Huesler [143] asked how the local instability could happen in a homogeneous oxide film. The inhomogeneity does not pre-exist, being created as a result of the interaction between the oxide film and aggressive environments. Based on the posteriori assumption, so-called by Okada [1451 , a local interaction across the interface, such as the critical chloride nucleation [581  and transitional halide complex formation [135] , may give rise to the critical passive film  breakdown process. In summary, it seems reasonable to conclude that those factors which control interactions between the oxide film and the adjacent local solution are the factors which govern the pit initiation process. Thus, three considerations appear to be relevant to the pit initiation process: (1) electrode kinetics, the potential - current relationship. (2) nature of the passive film. (3) local solution chemistry.  DISCUSSION^  149  7.2 Electrode Kinetics When an anodic polarizing potential is applied to an electrode to move it away from the equilibrium potential, a kinetic process occurs across the double layer of the electrode. For activation controlled dissolution, Tafel' s law applies: =a  +  blog(i)  (7.1a)  where rl is the electrode overpotential, a = -(2.303RT/azF)log(i0), b = 2.303RT/azF is the Tafel slope (it, is exchange current density, a is the electron transfer coefficient, and z is the electric charge of the activated species). When the electrode process is controlled by mass transport (diffusion) in the liquid phase, then we have: 2.303RT zF  —^  log 1 --.-  (7.1b)  where i1 is limiting current density, is the concentration overpotential. However, when a passive film is formed on the surface, it changes the electrode kinetics. The passive film blocks the dissolution of the metal substrate. The anode kinetics are no longer controlled by the charge transfer in the double layer but are controlled by the processes occurring across the passive film. In the passive state, the passivation of the metal is maintained by an extremely small current, called the passive current, of the order of a few RA/cm2 . There is still little information available on the kinetic processes taking place across the oxide film, but they involve the transport of ions through the film. In general, the question remains as to whether the passive current is an ionic flux due to the migration of anions or cations through the oxide layer, and whether the passive current is distributed evenly over the passivated surface.  DISCUSSION^  150  When the passive film has a local breach, it is uncertain whether the anodic process at the breach is activation-controlled and follows the Tafel behavior. Experimental evidence has shown that the anodic current density inside initiated pits is very high, in the order of A/cm2 [77], [146], [147] Kaesche [146] reported that the pitting current inside Al pits is about 0.8 A/cm 2 . Sato et al. [77] found that the current density at the pit mouth is as high as 8 A/cm 2 for 18Cr-8Ni stainless steel. According to Tousek [1471 , the current density inside pits was 5.8 A/cm2 one second after pit initiation, decreasing to 0.5 A/cm 2 after 300 seconds in 0.5 M NaC1 at pH 8.4. Suzuki et a/. [1°41 claim that a minimum current density in the pit (in the order of 10mA/cm 2 ) exists, below which pits cannot develop. The current density in the pit varies with applied potential according to the Tafel law (Equation 7.1), and the Tafel slopes were reported to be 0.150 V for aluminum  [641  , 0.050 V [147] and 0.087 V [148] for  Fe-18Cr-10Ni stainless steel.  7.3 Dynamic Nature of the Passive Film The dynamic features of the passive film are easier to understand if the passive film is considered to be an adsorbed layer. For example, following Uhlig, the surface may be considered to be passivated by the adsorption of passivating ions ( 0 2- and OH- ). Consequently, there exists a dynamic equilibrium between adsorbing and desorbing passivating ions. Therefore, since desorption of passivating ions will produce a local depassivation, it follows that there will be a dynamic passivation/depassivation process on the metal surface. For an oxide layer, there is an analogous dynamic nature due to the film breakdown/reformation processes occurring on the surface. There are several reasons for the passive film to be broken down, or ruptured, then re-formed. First, the volume of oxide per  151  DISCUSSION^  mole of metal (Vox) is different from the molar volume of the pure metal (V M ). Hence, when the oxide is formed, the ratio of Vo/Vm is not equal to 1. Therefore, stresses will be developed in the film when it is formed, and the film can be ruptured when the stress becomes sufficiently large. The Vox/Vm can be calculated according to the following equation: voivm = ( wox pry)4n, WM pox)  ^  (7.2)  where W is the molecular weight, p is density and 'Is is the stoichiometric number of metal atoms in the oxide. For example, VNi(oH)/VNi = 3.39 , V Al2 03 3H2 01 V Al = 3.22 and V Zn (OH)2 IV Zn =  2.7 g/cm3 ,  3.56, where 0 Ni(OH)2= 4.15g1cm3 , p m = 8.90 g/cm3 ,  Zn(OH)2=  A/20331120 =  2.24g1cm3 , p m =  3.05g1cm3 and p z . = 7.14 g/cm 3 r 1491 . Secondly, if the oxide film is  crystalline with a different crystal structure from that of the metal substrate, the film is subject to breakdown when the epitaxy of the film is changed during film growth. According to Sato's chemical-mechanical breakdown model, the film could be ruptured by the electrostriction pressure due to the high electric field across the film. The dynamic feature of the passive film was also recognized by Baroux in the point defect model  [1501  .  The film breakdown/reformation process has been detected by the technique of electrochemical noise analysis [1521,  ^The noise exists in both non-aggressive (halide-free)  and aggressive (chloride) solutions. However, a "burst" type of noise was found in chloride solutions, which could be associated with the "birth/death" process of metastable pits.  7.4 Local Solution Chemistry during the Film Breakdown/Reformation Process The film breakdown/reformation process, as mentioned in Section 7.3, causes local  DISCUSSION^  152  fluctuations in the anodic current density and solution chemistry. In this section, the resulting local changes in pH and halide concentration are modelled for two simple situations: (1) changes due to one oxide film breakdown/reformation cycle and (2) changes due to a continuous breakdown of the oxide film. The changes in local solution chemistry are considered to be critical to the formation of stable halide complexes and pit initiation. (1) Local pH and halide concentration changes due to one film breakdown/reformation cycle. Assuming that a hydrated oxide film, with a thickness of  h,  breaks down and produces  a circular active site with a radius of r (Figure 69), the metal substrate will undergo an anodic dissolution: M = M" + ne-^(7.3a) then the active site will be re-covered by the hydrated oxide film during the film reformation according to the following reaction 3 : M" + nH 2 O = M(OH)" + al +^(7.3b) and the equilibrium constant for Equation (7.3b) is given: [H  r  + n  " [Mn  ]  (7.3c)  —  Therefore, the volume of the hydrated oxide (V ox) to be reformed at the active site is:  3  For the formation of an anhydrous oxide (M0 ( ",2)), we have the following reaction: + (n/2)H2 0 = MO ("il2) + al+  Metal  ^  Oxide Film  ^  Diffusion Layer  ^  Bulk Solution  The region containing generated hydrogen ions  Figure 69 An active site with a radius of r during oxide film breakdown.  DISCUSSION^  154  ^Vox = nr2h^  (7.4)  and nr2 hp oxlW„ moles of oxide will be generated for the coverage of the active site, where pox is the density of the hydrated oxide, and Wox is the hydrated oxide molecular weight. According the mass balance in reaction (7.3b), the number of moles of the generated W . , Mol(H+ ), will be: Mol(W) = nicAp oxlWox^(7.5)  Assuming that the film reformation process is fast (a pulse of current), then the fr ions generated by Equation (7.3b) will not have moved far from the electrode surface when repassivation is completed. Therefore, allowing the generated H + to be confined in a cylindrical region with the radius, r + d, and height, d, as shown in Figure 69, the volume of solution, 17, 01 „, containing the generated H + ions is given by: ^Vsoh, = m(r+d) 2d^  (7.6)  And the concentration of H + can be calculated by the combination of Equation (7.5) and (7.6):  [H1 —  nr2hpo,  (7.7)  (r + d) 2dWox  There is a difficulty in estimating the value of d in Figure 69. However, the volume of solution containing generated H + can be roughly calculated based on the film reformation rate, i. e., the time, t, required for the reformation of the passive film at the active site. If the reaction (7.3a) proceeds at an anodic current density, i, then the charge, Q, required for the reformation of nr2 hp 0/147„x moles of oxide is given by:  DISCUSSION^ Q = nr2 it = nicr2 hpo,FYI'Vox^  155 (7.8)  from which t may be determined: (7.9)  nhp ox F t=^ ilKx  where F is the Faraday constant. Ignoring electrical migration effects, the distance, x, travelled by I-I ÷ during the time of t can be estimated from solutions to Fick's second law, where the diffusion coefficient, D, is assumed to be constant: x = 2 (Dt) 1/2^(7.10)  then d approximates to the distance x travelled by W. Therefore, we have:  d—  (4nhpoxFD )112  (7.11)  i wox  Therefore, in neutral or near neutral solution where the initial fr concentration is negligibly small, the local concentration of I-I + ions can be estimated by combination of Equations (7.7) and (7.11).  [H1=  nr 2hp ox (4nhpox FD )1121 2 (4nhpox FD )112  {r +  ^Wox  iwox^iwox  (7.12)  DISCUSSION^  156  As we see in Equation (7.12), for a given oxide film, the local concentration of fr ions changes with two variables, r and i, without considering the neutralization by the bulk solution. If r and h are in units of cm, p ox is in g/cm3 , D is in cm2/s, and i is in A/cm2 , then [fil ] is given as moles/cm 3 . Consequently, under these circumstances 0] may be converted to molarity (moles/liter, M) by multiply Equation (7.12) by 1x10 3 , from which pH = -log(10 3 [HI). Figure 70 shows the variation of local pH with the size of active site, r, and the anodic current density, i, for reformation of Al 2 0 3 .3H2 0 on the aluminum surface, where h = 1x10 -7 cm, p ox = 2.42 g/cm 3 , D = 1x10 -5 cm2/s [151], n =6, Wox = 156 g and F = 96500 C. From Figure 70, we can see that the local pH decreases with the increase in the size of active sites and with increasing anodic current density. For an active site larger than 1 gm (10 -4 cm) radius and with i = 1.0 A/cm 2 , the local pH is lowered to less than 2. When an anodic current, i, flows from the metal electrode, the current is transported by the migration of ions in the electrolyte, either cations move away from the electrode or anions move into the interfacial region. Considering a solution containing only sodium halide without other supporting electrolytes, then the migration of halide X - carries a fraction of the anodic current, which is determined by the transport number, nx- . By knowing the transport number, n, the local current carried by the migration of X - ions is icr2 itix- and the amount of charge carried by migration of X - ion is nr2 inx- t during the breakdown/repassivation cycle, where t is expressed in Equation (7.9).  13 12 11 10 9 8  I  CL 7  Tt 1 0  o  J  6 5 4 3 2 1 0 1 0 "8  -2 10  1 o -6  r, cm  Figure 70 Change in local pH with variation of r and i  DISCUSSION^  158  Then assuming that halide ions migrate from the bulk solution into the same region as the IT- ion confining region shown in Figure 69, we have a similar Equation to (7.12) for the local increase in halide concentration, [X],„ e , in moles/cm 3 , and which can be converted to molarity (M) by : M = 10 3 moles/cm 3 .  =  (4nhp ox FD )112} 2 4nhp„FD )112  {r +  and  (7.13)  nnx—^r 2 -•n p ox ^ox  - - ox^ iwox  (7.14)  nnx—r 2 hp ox  [X] = [X1 0 +  (4nhpox FD )112} 2 (4nhp ox FD )112  {r +  ^Wax  iwox^ iwox  where [X - ],o, = [X] - [X] 0 , [X- ] is the local concentration of halide ions, and [K] 0 is the bulk concentration of halide ions. Figures 71 and 72 show changes of the local chloride concentration with the size of the active site, r, for formation of Al 2 0 3 .3H2 0 on aluminum, nx = 0.61 [151] , h = 1x10 -7 cm, p ox = -  2.42 g/cm 3 , D = 1x10 -5 cm 2/s, n = 6, Wo, = 156 g and F = 96,500 C. Figure 71 was calculated at different anodic current density for a given bulk solution concentration of 10 -2 M. In Figure 71, the local halide concentration is little affected and equal to the bulk concentration for an active site having radius less than 10 -6 cm. At r = 10-4 cm (1 gm), the local halide concentration rises with increasing anodic current density. In Figure 72, the results were calculated for different bulk halide concentrations at a given anodic current density of 1.0 A/cm2 . For halide solutions with bulk concentrations of 0.001 and 00.1 M, the local concentration increases as the active site becomes larger than 10 4 cm (1 gm). There is little change in the local concentration (equal to bulk concentration) for the halide solutions of 0.1 M and 1.0 M.  -1  -1.2 -1.3 -1.4  -1.7 -1.8 -1.9 -2 -2.1 10-8  10-6  1 02  r, cm  Figure 71 Change in local halide concentration with variation of r and i at a given bulk halide concentration of 0.01 M  0.2 0  [X] 0 =1.0  -0.2 -0.4 -0.6 -0.8 -1  ......  ...................  -1.2 -1.4 -1.6 -1.8  ..  .  ..........  [K] o =am  -2 -2.2 -2.4 -2.6 -2.8  [X] o  =o.00i  -3 -3.2  1 0 -8^1 0-6  ^  i o-4  ^  16  2  r, cm  Figure 72 Variation of local halide concentration with bulk halide concentration at a given anodic current density of 1 A/cm  2  Cr;  DISCUSSION^  161  (2) Local pH and halide concentration changes due to a continuous breakdown process at a local (active) site. Assuming that a passive film on the surface is continuously broken away at an active site with a steady anodic current, i, the metal substrate will undergo the anodic dissolution according to Equation (7.3a). The flux of generated M' ions is given: J = i/nF^  (7.15)  where F is the Faraday constant (96500 C), n is the charge on the metal ions, M'. The generated M' will diffuse out into the bulk solution through a diffusion layer, as shown in Figure 73. Therefore, the M' concentration can be modelled approximately via a steady state diffusion process. At the steady state, the M' concentration at the interface can be estimated by Fick's first law: [M n_  J  (7.16)  = nF = D^8  where [Mil ; is the concentration of M' at the interface, [M1 0 is the bulk concentration of Mn + , and 8 is the thickness of the diffusion layer. Then we have: [M n  ]  t —  [Mn ]o—  i8 nFD  (7.17a)  the concentration of M' in the bulk solution can be neglected, then we have:  Metal  ^  Film^Diffusion Layer  ^  Bulk Solution  [ Mill  [ re1 0  a  ^  x  Figure 73 Local solution chemistry profile due to a diffusion process across the diffusion layer  DISCUSSION^  163  i8 ^(7.17b)  = nFD  Assuming that an equilibrium is established between M' and H + according to Equation (7.3b), the concentration of H + at the interface can be calculated from Equations (7.3c) and (7.17b). (7.18a)  [HT i K„ nFD  and (7.18b)  i8K =( ") nFD  Units used in Equation (7.18b) are: i - A/cm2 , 8 - cm and D - cm2/s. Thus the [Hl i is in units of moles/cm 3 . The [H + ] ; is then converted to molarity (M) by multiplying by 10 3 , from which pH is calculated (pH = -log(10 3 [H+ ] ; ])). According to Equation (7.18b), the local H + concentration at the interface (yr j i) depends on the K„ for the specific oxide and increases -  with increasing i and 8. Any increase in anodic current density and diffusion layer thickness will increase^and lower the local pH at the interface. Figure 74 shows the variation of local pH with the anodic current density at a given diffusion layer thickness of 1x10 -2 cm for Ni, Zn, Al and Sn oxides, where D = 10-5 cm2/s, logK„ (Ni(OH) 2 ) = -12.18, logK„ (Zn(OH) 2 ) = -10.96, logK„ (Al 2 0 3 .3H 2 0) = -5.7 and logkx (Sn(OH) 2 ) = -1.5. The value of 8 was reported to be 1x10-2 - 5x10 -2 cm (100 - 500 gm) under quiescent conditions  "54], "55] .  As discussed above, the continuous breakdown process gives rise to a steady anodic current flow at the active site. The flux of halide ions migrating to the local site on the  DISCUSSION^  164  electrode surface will be in x%F. When a steady state is established, the flux of halide ion migration into the interfacial region is equal to the flux of halide ions diffusing back into the bulk solution through a diffusion layer of thickness 8. The concentration of halide ions at the interface at steady state then becomes: [Xl i —[X1 0 F^8  n;i  =D  [X ] = [X-13+ -  ;  n;i8  (7.19)  (7.20)  FD  where [X] ; is the halide concentration at the interface, and [X- ] 0 is the bulk halide concentration. Figure 75 shows the relationship between the local halide concentration ([X] i) and bulk concentration (fX1 0) at different anodic currents, calculated from Equation (7.20) for a sodium chloride solution without other supporting electrolyte, where n; = 0.61, D = 1 x10-5 cm2/s, and 8 = 1 x10-2 cm. The change in local chloride concentration is very  significant when the bulk concentration [X] 0 < 10 -2 M. The change becomes less pronounced when the bulk solution becomes more concentrated with chloride.  6  _ -------------  5  Ni  ---------------------------------  4  I  0_ 3  cis 0 O 2 _I  Al  1  0  Sn  -1  -2  1  10  I^i  - 2^  1  1  I^i  10  1  1  -1^  1  1^i  10 0  1  ^  1  1  10  1  i A /c m 2 ,  Figure 74 Effect of i on the local pH for Ni, Zn, Sn and Al in a continuous film breakdown process at 8 = 10-2 cm  0  -1  -2  -  [X l 0 = [X 1 i  3  -4 i = 1.0 A/cm 2 1= 0.1 A/cm 2  -5  i = 0.01 A/cm -6  -6  -4  -2  0  2  LOGTX 10, M  Figure 75 Variation of local halide concentration with bulk halide concentration in a continuous film breakdown process at 8 = 10-2 cm  DISCUSSION^  167  7.5 Halide Complex Formation and Pit Initiation Theory The consequences of changes in local solution chemistry are dealt with in this section. The changes include a lower pH and a higher concentration of halides at an active site (see Section 7.4). According to the E-pH and X-pH diagrams constructed for tin, zinc, aluminum and nickel (see Section 6.2), when the pH moves to lower values (to the left side of diagrams), the passive oxide is no longer thermodynamically stable. Figure 76 presents a schematic X-pH diagram for M(OH) n in a halide solution. The bulk solution pH and halide concentration are presented as point A on the diagram, indicating that the metal surface is passive. With decreasing pH the metal is subjected to a transition from passivity to dissolution via halide complex formation. An increase in halide concentration, [Xi, narrows the passive zone, and also causes the instability of passive films. Therefore, a combination of a lower pH with a higher concentration of halides (point B in Figure 76) leads to a trend towards increasingly stable complex formation and a destablization of the passive films. Based on the above considerations, a pit initiation theory is proposed, which emphasizes changes in local solution chemistry and the generation of critical conditions for the formation of stable halide complexes. In this theory, pit initiation consists of a series of stages as shown in Figure 77. Stage 1 is a dynamic process where the repetitive film breakdown/reformation process occurs rapidly and randomly over the surface, irrespective of the presence or absence of halides. In halide-containing solutions, the film breakdown/reformation process leads to changes in local solution chemistry (Stage 2) characterized by a lower pH and a higher concentration of halide ions (as discussed in Section 7.4). The local solution chemistry change can disappear very shortly after repassivation and only cause a short-lived  X 0  -1  0  _1  pH  Figure 76 A generalized X-pH diagram  DISCUSSION^  169  perturbation at the oxide/solution interface. If the fluctuation in local solution chemistry is significant enough, with an adequate lifetime, a further breakdown/reformation process may occur at the same site due to an autocatalytic process in the local region, and a continuous breakdown process can be established (Stage 3). At stage 3 metastable pits are formed, they may be repassivated due to any change in the kinetics of the film breakdown process. Continued changes in local solution chemistry allow stage 4 to be reached, where the critical conditions are met for stable complex formation. At this stage (stage 4), any oxide film will be unstable relative to halide complexes, and it is no longer possible for the surface to be repassivated. Therefore, stable pits are initiated. Consequently, stage 4 is the critical stage in pit initiation, where critical conditions are established for stable halide complex formation. From a thermodynamic point of view, pit initiation is determined by the formation of stable halide complexes. According to the X-pH diagrams for several metals in Section 6, the conditions for their formation are defined by the lines representing chemical equilibrium between the hydrated oxides and the halide complexes, according to the reaction in Equation (7.21) 4 : M(OH) n + mX - + al+ = MX. ("1- " )- + nH 2O^(7.21) and conditions for formation of MX,n (m - n)- : [MX,n(m-n)- 1 = KxGHTPCl m)  ^  4 For an anhydrous oxide, Equation (7.21) may be written:  M0 0/2) + mX - + nil+ = MX,, (' )- + n/2H 2 0  (7.22a)  ^  X^X  MXm (m-n)-  4k,  HI-^ H ' r-N, H+^  Oxide film  High [Xi Low pH  x  (2)^ (3) Local Change in pH and [X" ]^  Metal  (4) Critical stage for stable pit initiation  Continuous breakdown process (1) Breakdown/reformation process  Figure 77 A schematic diagram showing processes in pit initiation  DISCUSSION^  171  where the equilibrium constant is given by IC = [MX,n(m -11/(fHTPCP). Assuming that a critical concentration of halide complexes [MX. (m-n) ]cnt is required for film breakdown and formation of a stable pit, then we have a critical product of (VIT[X] m )crit, i.e. uva„, m n (  )  c rit  -  = KAHTPcncrit ^  (7.22b)  When the product of local pH and local halide concentration is greater than the value of alti m iXT)crit, MX„,1 (m-n)- formation is thermodynamically stable, which will lead to pit initiation. From a kinetics point of view, two reactions compete with each other at the critical stage. These are formation of the oxide film and formation of complexes:. M + nH2 O --> M(OH)„ + nH+ + ne^ (7.23) and^M + mX- MX.(m* + ne  ^  (7.24)  Their reaction rates, R, are determined by: RM(OH)„  1  ^  (7.25)  '  and^Rmxm(„,_„ _ )  rx f  0c L  ^  -  where 13 and 7 are the reaction orders of  [111 and  (7.26)  [K] respectively. Thus, we can see that  high [ff ] (low pH) slows down the repassivation process and that high halide concentration -  accelerates the formation of complexes. Pit initiation occurs when:  DISCUSSION^ R^RAmion^  172 (7.27)  7.6 Evidence from Experiments on Environmental Effects The proposed pit initiation theory emphasizes the importance of changes in the local solution chemistry and the formation of stable halide complexes. The critical stage in pit initiation is determined by the critical value of local pH and halide concentration (Equation 7.22b). Therefore, any factors that influence the local solution chemistry will also affect the pit initiation process. The following environmental effects in this study can be clearly explained by the proposed theory: (1) Complex-forming ions are needed for the establishment of the critical stage for pit initiation, otherwise pitting corrosion does not occur. The so-called aggressive ions are the complex-forming ions. The polarization results have indicated that there was no pitting of nickel in 1.0 NaNO 3 and 1.0 M Na2 SO4 at pH = 10.5, not even at pH = 2.5 in Na2 SO4 solution (Figure 33 and 34). The passive film formed on nickel is very stable in these solutions. Therefore, nitrate and sulfate are considered to be non-aggressive ions. However, halides are aggressive. They can form various types of complexes with nickel, aluminum, tin and zinc (such as NiCr, NiC1 2 , A1C1 3 , A1Br3 , ZnC1 4 2- , ZnBr4 2- , SnC14 2- and Sn 6 2- ), so that pitting corrosion of these metals is predicted if halides are present in the solution. (2) Some halide complexes are more stable than others; the greater the stability of complexes, the more readily the pit is initiated. In order to compare their stabilities, the minimum halide concentration, [X - ]„„„, above which metal complexes become the dominant soluble species is calculated from the stability constant Km for Equation (7.28a):  DISCUSSION^ M°+ + mx =^ -  173 (7.28a)  and the stability constant, Km = [MX. (111-1111[Ml[X] "^  (7.28b)  When [Mil = [MX,, (11'1, we have [X] = [X ]mi. and  ^  hIm [X ]min= Km ^(7.29)  The formation of the halide complex, MX, n (m- " )- , is favored when the halide ion concentration is greater than PC - Lin . Table 18 lists the minimum halide concentrations for nickel, aluminum and zinc, which may be used as a measure of the aggressiveness of the halides. For example, during the pit initiation process, and the accompanying change in the local solution chemistry, it is easier to meet a smaller minimum halide concentration requirement. Consequently, the halide requiring smaller minimum concentration is more aggressive towards pit initiation. Therefore, according to Table 18, bromide ion is less aggressive than chloride ion for all three metals. Figures 26, 28 and 48 show the pitting tests for nickel, aluminum and zinc in different types of halide solutions. From these figures, critical pitting potentials show a trend as follows: Ecp in NaCl < E cp in NaBr  Therefore, if aggressiveness is equated with a lower pitting potential, then the trend is consistent with the predictions based on the minimum halide concentration for complex formation (Table 18).  DISCUSSION^  174  Table 18 Minimum halide concentration for complex formation  Metal  Complex  Min [CF], M  Min [Bri, M  Ni  Nir  0.47  1.32  NiX 2  0.41  -  A1X3  0.585  0.776  Znr  0.32  0.60  ZnX2  0.50  1.12  ZnX3 -  0.68  1.76  ZnX42-  0.89  1.78  Al  Zn  (3) From the X-pH diagram for the H 2O - F - Al 2 0 3 3H2 0 system (Figure 61), it is known that Al-fluoride complexes are very stable, so a very low critical pitting potential would be expected in fluoride solution. Actually, the fluoride complexes are so stable that total passivity of the aluminum surface is destroyed, which results in an active dissolution behavior in 0.1 M fluoride solution at pH 3.4 (Figure 27). For a bulk solution with [F] = 0.1 M and pH = 3.4, it is clearly evident that the bulk solution conditions shown in Figure 61 are  located in the complex formation zone (A1F 5 2- ), far away from the passive zone. In this case, the bulk solution chemistry meets the conditions for stable complex (A1F 5 2- ) formation so that general corrosion occurs on aluminum, instead of pitting corrosion. A similar situation is seen for nickel in 1.0 M chloride and fluoride solutions at pH = 2.5 and 3.1 respectively  DISCUSSION^  175  (Figures 35 and 52). Bulk solution chemistry conditions also show that Ni is located in the stable complex formation zones in Figures 65 and 66, forming NiC1 2 and Nir respectively, and resulting in total depassivation of nickel. General corrosion, resulting from total depassivation, will be expected if the bulk solution chemistry meets the conditions for stable halide complex formation. According to the E-pH and X-pH diagrams, the total depassivation requires that the bulk solution chemistry exhibits a low pH and a high halide concentration. However, in most cases such as for Ni in 1.0 M NaC1 at pH 10.5 and for Al in 0.1 M NaCl at pH 6.0, the bulk solution chemistry favors passivation and not the formation of halide complexes. Therefore, only local regions on the surface meet the requirement for complex formation, due to the changes in local solution chemistry discussed in Section (7.4). Consequently, only localized breakdown of passivity is encountered, while the rest of the surface remains passive. (4). In the proposed theory, the local pH is emphasized instead of the pH of the bulk solution. For example, bulk solution chemistry indicates that nickel is in the passive zone on the X-pH diagram for 1.0 M NaC1 at pH = 10.5 (Figure 66), but the local pH may be lowered towards the stable complex zone. Hence, pit initiation depends on the local pH and any effect on local acidification will influence the pit initiation process. The results clearly show that there is little effect of bulk solution pH on E cp of Ni in the pH region from 4.5 to 10.5. The critical pitting potential is independent of bulk solution pH until the solution becomes strongly alkaline (pH > 12) in 1.0 M NaC1 (Figures 36 and 37). According to the theoretical model in section (7.4), the generation of hydrogen ions is determined by Equation (7.18b), provided that there is a continuous breakdown of the  DISCUSSION^  176  passive film. Taking into consideration the transport of OH through the diffusion layer -  from the unbuffered bulk solution, the local solution pH for Ni is calculated (see Appendix I) and plotted in Figure 78. Figure 78 shows that the local pH is independent of bulk solution pH in the acidic, neutral and slightly alkaline region, since there are insignificant amounts of OH available to react with and neutralize the generated H+ . ^the generated 11+ ions almost totally -  control the local solution pH. However, in strong alkaline solutions the large amount of available OH ions consumes the generated fr ions , so that the local pH change is retarded. -  The local pH becomes indistinguishable from the bulk solution pH in very strong alkaline solution (pH >12). The local pH change in Figure 78 correlates well with the variation of the critical pitting potential for Ni with bulk solution pH (Figure 39). Recalling the observed corrosion behavior of nickel in 1.0 M NaCl with increasing bulk solution pH (Section 5.2.2), the transition from uniform corrosion to pitting corrosion and then to complete passivity can now be well understood. At pH = 2.5 the bulk solution chemistry meets the conditions for complex formation, resulting in general corrosion of nickel. In the pH region from 4.5 to 12.5, chloride complexes are only formed locally due to the change in local solution chemistry, and the local pH is independent of the bulk solution pH in the pH range from 4.5 to 10.5. Consequently, localized breakdown occurs, with the critical pitting potential independent of the bulk solution pH. At higher pH values, the local pH is controlled by the bulk solution pH, and no local acidification occurs at pH = 14, so that nickel remains passive.  14 13 12 11 10  I^I^I^I^I^I^I^I^I^1  2  ^  4^6  ^  8^10  ^  I^I^I  12  ^  14  Bulk Solution pH  Figure 78 Change of the local solution pH with the bulk solution pH calculated for Ni using an one-dimensional diffusion model  DISCUSSION^  178  (5) According to the proposed pitting theory, the effect of a buffer on pitting corrosion behavior may be predicted. When a buffer is introduced into the solution, the generated fr ions will be consumed by a buffering reaction, which counteracts local changes in pH, as shown below for carbonate and phosphate buffers: CO 3 2- + fr = HCO 3 -^(7.30) and^PO43- + fr = HPO 4 2-^(7.31) The local pH control by the presence of a buffer increases the critical pitting potential. This has been confirmed for nickel in buffered chloride and bromide solutions at pH = 10.5 (Figure 42 - 46). Therefore, it is concluded that the CO 3 2 1HCO3 and PO 4 3 1HPO4 2- buffers do not act simply by competing with halide ions for adsorption sites on the surface, as suggested by others  [661, [67]  Their primary role is the retardation of local pH change. The  increase in critical pitting potential was also found on zinc in CO 32 1HCO3 - buffered solution (Figures 26). However, in this case, the increase should also have resulted from the broadening of the passive zone to include lower pH regions because of the formation of insoluble ZnCO 3 (Figure 56). (6) Aluminum should have undergone general corrosion in 0.1 M NaF at pH 6.0 according to the prediction arising from the X-pH diagram (Figure 61). Aluminum did corrode in the active region. However, the anodic current density dropped dramatically to a very low value above a potential of -1.1 V (SCE). This transition is caused by the formation of insoluble A1F 3 .3H2 0, which precipitates on the surface of aluminum, as the Al 3+ concentration at the surface increases with increasing anodic current density. A thick salt film was noted on aluminum in 0.1 M NaF at pH 6.0 (Figure 31). Chemical analysis showed that it was an aluminum fluoride film (see Section 5.1.4).  DISCUSSION^  179  At first, it is surprising that nickel does not suffer pitting corrosion in 1.0 M NaF solution at pH = 6.0 and 10.5 (Figures 50 and 51), because pitting corrosion is predicted by the X-pH diagram (Figure 65) on the basis that the fluoride complex, NiF, is formed. However, taking into account the fact that HF is a weak acid, however, a F/HF buffer reaction occurs: F + H ÷ = HF^  (7.32)  and reaction (7.32) consumes not only the generated II+ ions, but also fluoride ions. Therefore, the change in local solution chemistry is well retarded. This suggests that the resistance to the change in local pH and fluoride concentration prevents pitting corrosion from occurring.  7.7 An Explanation of the Pitting Dependence on Crystallographic Orientations Orientation-dependent pitting behavior was observed on single crystals of zinc and tin in the present study, and reported on aluminum and nickel single crystals by others  [80]-[82]  So far, there is no generally accepted theory to explain the anisotropy in pitting corrosion. One may suggest that the surface atomic densities of metals play a role in orientation-dependent pitting corrosion. The surface atomic density changes with orientation. The higher atomic densities in surfaces composed of the more closely packed planes will produce higher numbers of oxygen cations per unit area of surface monolayer that is oxidized. This may cause higher strains in oxide films forming on these surfaces, leading to easier breakdown of oxide films. This possibility may be related to Zn because the surface atomic density increases in the order:  DISCUSSION^  180  (1010) < (1120) < (0001) and the observed critical pitting potential in Figure 12 increases in the reverse order of the surface atomic density, as predicted. Ecp : (0001) < ( 11 20) < (1010)  If the predicted relationship between surface atomic density and the critical pitting potential is applicable to other systems, then the atomic density and the critical pitting potential for F.C.C. metals should be: Density:^{110} < {100} < {111} and predicted E cp :^{111} < {100} < {110} In fact, the observed critical pitting potentials on Al  [80]  increase in the order { 111 } < {110}  < {100}, which is not entirely consistent with the prediction. Furthermore, the pitting behavior of Ni suggests that there is no generally applicable relationship, because E el, has been reported to be the highest on { 111 }  [821  , instead of the predicted {100} surface.  In the proposed theory (see Section 7.5), the transition from stage 3 to stage 4 is considered to be the critical process in pitting initiation. Therefore, the possible explanation of orientation-dependent pitting may be related to the effects of crystallographic orientation on changes in the local solution chemistry. These changes depend on the local anodic current density (Section 7.4). Higher anodic currents promote a lower pH and higher concentration of halides at the local regions (see Figures 70, 71, 74 and 75), so that the  DISCUSSION^  181  critical condition for pit initiation is more readily established. Therefore, we expect a lower critical pitting potential on a surface which has a higher local anodic current density (a higher active dissolution rate). Data reported in the literature show that the anodic dissolution rates Ra under active conditions in acid solutions for aluminum {130} and nickel [132} are dependent on surface orientation: aluminum. R,{001} < R, {011} < Ra{111} and nickel: Ra{ 111)^ R 1 11} —0)111^ Ra{001} Pit morphologies also indicate that the lowest dissolution rates occur on Al {001} and Ni { 111 }, which findings are consistent with the observed critical pitting potentials for Al and Ni (see Section 2.9). The active anisotropic dissolution rates of zinc and tin are unknown, but the relative rates of dissolution of differently oriented surfaces may be ranked on the basis of the observed pit morphologies. For example, pit walls of crystallographic pits are composed of the most slowly dissolving surfaces. Hence, the formation of crystallographic pits is evidence that dissolution behavior is anisotropic. It was observed that pit walls on zinc were close to {10l0} hexagonal prismatic planes, suggesting that dissolution rates are lowest on these surfaces, consistent with the highest critical pitting potential for { 1010} oriented crystal surfaces (Figure 12). Also, walls of pits formed on Sn crystals were found to be close to {100} tetragonal prismatic planes and {011} tetragonal bipyramidal planes, suggesting that these are slowly dissolving surfaces, consistent with critical pitting potentials  DISCUSSION^  182  observed on (100)- and (011)-oriented crystal surfaces (Figure 10). Therefore, the different anodic dissolution rates on the differently oriented crystal surfaces cause the orientation-dependent pitting behavior of single crystals.  7.8 Critical Pitting Potential The critical pitting potential is the most important criterion for defining susceptibility to pitting corrosion. The explanation of the meaning of the critical pitting potential is a key issue in the approach to understand the pit initiation mechanism. Any successful theory should explain the critical pitting potential, as suggested by Smialowska  [45]  and Kruger [211 .  In the present theory, it is clear that the critical pitting potential is directly associated with the critical stage in pit initiation. The critical pitting potential is defined as the potential at which the local solution chemistry leads to the critical condition for the formation of stable halide complexes. The relationship between local solution chemistry and electrode potential can be determined through Equations (7.18b) and (7.20), provided that the anodic current density, i, at the active site follows the Tafel relationship (Equation 7.1a). Thus, from Equation (7.18b) OK ^[H1 1^exp = k1 ( °x^azFE nFD^nRT  ^where k 1 = is exp ^ nRT  and from Equation 7.20  (7.33)  ^  DISCUSSION^  183 (7.34)  k2n;s5^czzFE ^ exp i —[X1 0 = ED X] RT  [  where, k2 = io exp  -azFE0 RT  .  When the condition^(from Equation 7.22b) is met at  the interface, we have the critical pitting potential, E cp . Therefore, oczF Ecp 8K ) 1c2n 8^azF E, exp RT^ ([X10 ±^eXp RT P )in^(7.35) ([ H1 in [Xi i)m^ crit = -  The relationship between E cp and bulk halide concentration [X] 0 can be derived from Equation (7.35), if the term  kfcs  exp  azT Ecp RT  is significantly small relative to [X] 0 . This  situation is most likely when the halide concentration in the bulk solution is high and/or there is a high concentration of supporting electrolytes ( nx - is very small). Therefore, neglecting this term, we have: E  ')exp azFRT cP ([Xlo)m =^n [Hi crit = constant FD 14(-nsK  (7.36)  and rearranging Equation (7.36) Ecp = A  2.303mRT ocz,F  log[X lo  (7.37)  ^  DISCUSSION^  184  RT^FD( v17 [111 1  :  where A is a constant given by A = azF ln^Int , and n and m are the stoichiometric sic; —  numbers for 11 + and X - in Equation (7.21), respectively. Equation. (7.37) gives a semi - logarithmic law between E cp and the bulk halide concentration [K] o , which has been well established for many metal - halide systems (see Section 2.2.2) 5 . However, the relation will deviate from the semi - logarithmic law if the k2 n;s^az,F'E exp ^ term cannot be neglected. Then there is no simple relationship between the FDRTcp critical pitting potential and the bulk halide concentration.  7.9 Extension of the Proposed Theory to Other Aspects of Pitting The proposed pit initiation theory has been shown to explain some of the environmental and crystallographic aspects of pitting. There are still other effects on pitting corrosion, such as temperature and metallurgical effects (alloy compositions and inclusions), which were not investigated in the present study. We shall attempt to extend the theory to these aspects.  5 Equation (7.37) is derived by assuming that the relationship between the potential and anodic 2. 30 3mR T  current density follows Tafel's law. Thus the slope of ,,,zF is obtained. In situations where the anodic current density is high (in the order of A/cm 2), the relationship between the potential and the anodic current density does not fully obey Tafel's law, and the electrode kinetics are partially controlled by the mass transport process across the diffusion layer. Consequently, the 2.303mRT apparent slope may be larger than the predicted slope of azF in Equation (7.37).  DISCUSSION^  185  Some alloying elements greatly increase the pitting resistance of alloys, for example, the addition of Mo and nitrogen into austenitic stainless steels (see Section 2.2.8). The increase in pitting resistance by alloying can be explained in three ways according to the theory: (1) broadening the passive zone; (2) decreasing the anodic dissolution rate at the active site; (3) buffering local pH changes. The film formed on stainless steels is primarily a chromium-rich oxide film which extends the passive zone into a lower pH region than iron oxide if E-pH diagrams for the Fe-H2 0 and Cr-H 2 0 systems are compared. Hence, chromium increases the pitting corrosion resistance. Stainless steels alloyed with Mo and N have improved pitting resistance. The role of Mo is still disputed, but some results have shown that the Mo addition decreases the anodic dissolution at active sites  [156]  A smaller anodic current  density tends to slow down the change in local solution chemistry and to prevent the local region from establishing the critical conditions for pit initiation. It has been speculated that the beneficial effect of nitrogen is due to the formation of NH 4+ by reduction of N within the active site, which consumes II+ ions and buffers the local pH change  [157]  .  The detrimental effect of inclusions has been reported in several studies (see Section 2.2.9). The existence of inclusions creates inhomogeneities in the passive film, which act as active sites in halide solutions. The size of the larger inclusions is in the order of gm. When an active site is formed at these inclusions, the changes in local pH and halide concentration will be significant (Figure 70 and 71), and the inclusions become the preferential sites for pit  DISCUSSION^  186  initiation. Temperature is another factor influencing pitting corrosion. The critical pitting potential decreases with the increase in temperature (see Section 2.2.4). It is not too difficult to understand why temperature promotes pit initiation kinetically. According to the Arrhenius rate law, activation-controlling dissolution will be accelerated with an increase in temperature. So a higher local anodic current density will give rise to a lower pitting resistance. Thermodynamically, if the passive zone is narrowed at higher temperature due to the enlargement of the complex formation zone in the X-pH diagram, then it is expected that pit initiation will occur more easily. Fluid flow affects the diffusion layer thickness 6, which influences mass transport controlled changes in the local solution chemistry (see Section 7.4). The changes in local solution chemistry will be retarded with a decrease in the diffusion layer thickness. Increasing the fluid flow velocity decreases the diffusion layer thickness and should produce an increase in the critical pitting potential. This is consistent with the observed behavior reported by others [73]4761.  CONCLUSIONS^  187  8 CONCLUSIONS (1) Pitting corrosion of zinc and tin single crystals is crystallographic-orientation dependent. The critical pitting potential of zinc, E ci, varies with the crystallographic orientation in the order of E cp (1010) > Ecp (1120) > k p (0001), and the lowest pitting potential for Sn crystals has been found on the (111) oriented surface. (2) The pits formed on the surfaces of zinc and tin single crystals are crystallographic. Some pit walls on Sn crystals are identified as {011} tetragonal bipyramidal planes and {100} tetragonal prismatic planes, and pit walls on Zn crystals are { 1010} hexagonal prismatic planes. (3) The pit morphology reveals the crystallographic planes with the lowest dissolution rates, and single crystal surfaces of Zn and Sn with these orientations exhibit higher critical pitting potentials. There is a correlation between the local anodic dissolution rates and the critical pitting potentials. Therefore, the dependence of pitting corrosion on the crystallographic orientations is attributed to orientation-dependent changes in local solution chemistry caused by orientation-dependent local anodic dissolution rates. Lower local anodic current densities at the active site produce higher critical pitting potentials. (4) E-pH and X-pH diagrams have been constructed for Sn, Zn, Al and Ni in association with the formation of halide complexes. These diagrams contribute significantly to the understanding of pit mechanisms. The role of halide ions is emphasized in the pit initiation process. The formation of halide complexes is a necessary step to destablize the passive film and induce pit initiation. (5) The critical pitting potentials for nickel, aluminum and zinc vary with the specific halide species. They increase in the tested halide solutions in the order: Ecp in NaCl < E ci, in  CONCLUSIONS^  188  NaBr. The minimum halide concentrations for complex formation, ['g rain , and X-pH diagrams indicate that the formation of more stable chloride complexes of nickel, aluminum and zinc leads to lower critical pitting potentials than the less stable bromide complexes. (6) The pit initiation process is governed by the local solution pH, instead of the pH of the bulk solution. The critical pitting potential of nickel has been shown to be independent of the bulk solution pH in the range of 4.5 - 10.5. The local solution pH is controlled by the bulk solution pH in very strong alkaline solutions (pH > 12.5), where a large amount of 011" is available to consume the II+ generated by the film breakdown/reformation process. The addition of carbonate and phosphate buffers to the halide solutions prevents changes in the local solution chemistry, thereby increasing the critical pitting potential. (7) The metal undergoes general corrosion if the bulk solution chemistry meets the conditions for the formation of stable halide complexes, such as the situation for nickel in 1.0 M NaCl at pH 2.5. If the bulk solution chemistry does not meet these conditions, the critical conditions are only established in the local regions and pitting corrosion is encountered. (8) A theory for halide-induced pit initiation is proposed, whereby the local solution chemistry governs the pit initiation process. The local solution chemistry is modelled for two simple situations, and it is shown that the local region is associated with low pH and high concentration of halide. Pit initiation occurs at a critical stage when the conditions are met for the formation of stable halide complexes in the local region. The theory explains the effects of solution pH, the presence of buffers, halide species and crystallographic orientations, all of which were investigated in this study. 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G. Akano, Y. C. Lu, R. Guo, S. C. Srivastava, Corros. Sci., Vol. 31, p. 367 (1990). [157] K. Osozawa, N. Okato, Passivity and Its Breakdown on Iron and Iron Base Alloys, Etd. by R. W. Staehle, H. Okada, NACE, p. 135, 1976.  APPENDIX  ^  197  APPENDIX I Effect of Bulk pH on Local Solution pH in Unbuffered Solution The change in local solution pH with bulk solution pH can be estimated by a simple one-dimensional diffusion model discussed in Section 7.4, without considering the ion electrical migration effect (which is valid for large concentrations of supporting electrolyte). The Local fr concentration is given by Equation (7.18b). The flux of generated 11 + is given: iSKox  (A-1)  [H  D nFD  HH+ +^H+^8  and the OH- flux from bulk solution into the local region is given: JOH  (A-2)  figH1 0 — [OH  = DOH^8  where, [OH] 0 and [OH] ; are bulk and local concentrations of Off respectively. Hence, the net flux of fr is the sum of Equations (A-1) and (A-2): •1 net = H +— "IOW  therefore:  (A-3)  ^  1^ ( i61Cox yi [H+10 [Hl i — [H10^ nFD ^[01110— [011 1 =D  DH^ ,^  (A-4)  ^D —  8^8^OW^8  Assuming that differences in the diffusion coefficients of H + and OH - are sufficiently small to be neglected [151], then D OH = D H+ = D, and [H+1^ioKox  nFD^  (A-5) ([0H10— [011 1 1  )  and equilibrium is established between II + and OH, and [111[01-1] = 10 -14 , then we have: ^ (A-6) 10-14^iSICT 10-14 [HI 1 nFD ) [H 1 0 -  APPENDIX^  198  Therefore, the variation of local solution pH with bulk pH can be calculated form Equation (A-6). Table A-1 lists the calculated data of the local solution pH for Ni at an anodic current density of 1.0 and 10 A/cm 2 respectively, where n = 2, log(K0 ) = -12.18 ( for Ni(OH) 2), D = 1x10 -5 cm 2/s and 8 = 10 -1 cm.  Table A-1. Calculated the local solution pH vs. the bulk solution pH for Ni  [1/10,^M^Local pH^Bulk pH  i= 1.0 A/c m2 0.0001^0.0001584999^3.799971^4 0.00001^0.000068499^4.164315^5 0.000001^0.00005949^4.225556^6 0.0000001^0.0000585^4.232844^7 0.00000001^0.00005751^4.240256^8 0.000000001^0.000048501^4.314249^9 1.00000000E-10^0.0000000002^9.618048^10 1.00000000E-11^1.06213489E-11^10.97382^11 1.00000000E-12^1.00588442E-12^11.99745^12 1.00000000E-13^1.00058534E-13^12.99974^13 1.00000000E-14^1.00005850E-14^13.95^14  = 10 A/cm 2 0.0001^0.0006849999^3.164309^4 0.00001^0.000594999^3.225483^5 0.000001^0.00058599^3.232109^6 0.0000001^0.000585^3.232844^7 0.00000001^0.00058401^3.233579^8 0.000000001^0.000575001^3.240331^9 1.00000000E-10^0.0004850001^3.314258^10 1.00000000E-11^2.40963855E-11^10.61804^11 1.00000000E-12^1.06213489E-12^11.97382^12 1.00000000E-13^1.00588442E-13^12.99745^13 1.00000000E-14^1.00058534E-14^13.99974^14  APPENDIX  ^  199  APPENDIX II  E-pH Diagrams for Sn Systems  E-pH diagrams in respect to the formation of Sn anhydrous oxides are shown in Figures A-1 and A-2 for Sn - H 2 O and Sn - Cl - H 2O systems, respectively. -  APPENDIX ^  200  Tin - water system  1.4 1.2 1 0.8 0.6  10 -6  0.4 — 0.2 cn  w  2SnO3  0 -0.2 -0.4 -0.6 -0.8 HSnO 2  -1  –  -1.2 -1.4 0  2  4  6  8  10  12  14  pH  Figure A-1 E-pH diagram for H 2 O - Sn system; activities of all solute species at 10 -6  APPENDIX  ^  201  Sn - chloride - water system 1.4 1.2  0.8 0.6 0.4  w  0.2 0  w  -0.2 -0.4 -0.6 -0.8  -1 -1.2 -1.4  0  ^ ^ ^ ^ ^ ^ ^ 2 10 12 14 4 6 8 pH  Figure A-2 E-pH diagram for H2O - Cr - Sn system; activities of chloride at 10 ° and 10'; activities of other solute species at 10-6  APPENDIX^  202  APPENDIX III  Pit Morphologies of Cold Rolled and Annealed Polycrystalline Nickel Pits formed on the surfaces of all cold rolled nickel specimens were found to be non-crystallographic. Figure A-3 shows a hemispherical pit formed on the surface of cold rolled Ni tested in 1.0 M NaC1 at pH 10.5.  Some tests were conducted on the annealed nickel specimens, as a comparison with the cold rolled nickel. Nickel was annealed in a furnace at 800 °C for 4 hours. The polarization curves obtained from the annealed nickel in 1.0 M NaCl at pH 10.5 are similar to that obtained from the cold rolled Ni in the same solution (see Figure 36 in Section 5.2), and no difference in their critical pitting potentials (E cp ) was detected. However, the pits formed on the surface of the annealed Ni were found to be crystallographic, as shown in Figure A-4.  APPENDIX^  203  Figure A-3. A non-crystallographic pit formed on the cold rolled Ni in 1.0 M NaCl at pH 10.5.  Figure A-4. A crystallographic pit formed on the annealed Ni in 1.0 M NaC1 at pH 10.5.  APPENDIX  ^  204  APPENDIX IV  Some Suggestions of Further Work  (1) The local solution chemistry has been modelled in two simple situations (see Section 7.4) to show the trend of changes in terms of the local solution pH and the local concentration of halides. However, more work is needed to model the processes occurring in the local region in details. A more precise model should takes into consideration: (a) activity coefficients of species, (b) equilibria between the species involved in the local region and (c) migrations of charged species.  (2) Further work is suggested to measure the anodic dissolution rates on the differently oriented crystal surfaces of Sn and Zn under active dissolution conditions in acidic halide solutions, in order to determine the dependence of anodic dissolution rates on orientations and to confirm the orientation-dependent pitting behavior of single crystals.  

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