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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 EffectsbyRUIJIN GUOB.Sc., Dalian University of Technology, China, 1982M.Sc., Dalian University of Technology, China, 1984A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THEREQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHYinTHE FACULTY OF GRADUATE STUDIES(Department of Metals and Materials Engineering)We accept this thesis as conforming to the required standardTHE UNIVERSITY OF BRITISH COLUMBIASeptember 1993© Ruijin Guo, 1993In presenting this thesis in partial fulfilment of the requirements for an advanceddegree at the University of British Columbia, I agree that the Library shall make itfreely available for reference and study. I further agree that permission for extensivecopying of this thesis for scholarly purposes may be granted by the head of mydepartment or by his or her representatives. It is understood that copying orpublication of this thesis for financial gain shall not be allowed without my writtenpermission.(Signature)Department of meMc a—t fitated04 crf.ApAy7„The University of British ColumbiaVancouver, CanadaDate ^-e/Y`t/ler^ t DE-6 (2/88)11ABSTRACTThe susceptibility to pitting corrosion of tin and zinc was found to be dependent on thecrystallographic 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 thefollowing order: (1010) > (1120) > (0001). The critical pitting potentials for zinc andaluminum 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 thepitting of polycrystalline nickel were investigated in 1.0 M halide solutions. It was found thatthe critical pitting potential (Ecp) was independent of the pH of unbuffered solutions in the pHregion of 4.5 to 10.5. However, Ecp was greatly increased at pH 12.5, and pitting corrosion wastotally retarded at pH 14 in 1.0 M NaC1 solution. The addition of Na2CO3/NaHCO3 orNa3PO4/Na2HPO4 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) andhalide concentration - pH (X-pH) diagrams were constructed for H 20-halide-metal (Sn, Zn, Aland Ni) systems. With increasing halide concentration, the formation of halide complexes isthermodynamically favoured and the zone of passivity is diminished. These new diagrams areunique and useful in the understanding of pit initiation.A pitting theory is proposed, which emphasizes the change in local solution chemistryand the formation of halide complexes during the pit initiation process. Pits will be nucleated111when conditions are met for the formation of stable halide complexes in the local region. Thistheory successfully accounts for the effects of the solution pH, buffers, halide ions and thecrystallographic orientations of single crystals on pitting behavior. It may also explain otheraspects such as the effects of temperature, alloying elements, and fluid flow on the pittingcorrosion of metals.Table of ContentsABSTRACT ^  iiLIST OF TABLES  viiLIST OF FIGURES ^  viiiLIST OF SYMBOLS AND ABBREVIATIONS ^  xiiiACKNOWLEDGEMENTS ^  xv1 INTRODUCTION  12 LITERATURE REVIEW^  22.1 Passivity ^  22.1.1 Thermodynamics and Kinetics of Passivation ^ 22.1.2 Properties of the Passive Film on Metals  62.2 Pitting Corrosion - Localized Film Breakdown  102.2.1 Pitting Corrosion Phenomena and Criteria ^  112.2.2 Effect of Halide Environments ^  132.2.3 Effect of pH and Buffers  162.2.4 Effect of Temperature  172.2.5 Effect of Solution Flow ^  182.2.6 Pitting of Metal Single Crystals  182.2.7 Pit Morphology  202.2.8 Alloying Element Effects ^  212.2.9 Inclusions, Grain Boundaries and Dislocations ^ 222.2.10 Solution Chemistry inside Pits  242.2.11 Application of E-pH Diagrams to Pitting Corrosion ^ 252.2.12 Theories of Pit Initiation ^  282.3 Remaining Problems^  322.3.1 The Pitting Dependence on Crystallographic Orientations ^ 322.3.2 Solution pH and Local Solution pH Control ^ 332.3.3 Halide Complexes and Diagrams Dealing with Halide Complexes ^ 332.3.4 Pit Initiation Mechanisms^  343 OBJECTIVE^  35ivV4 EXPERIMENTAL^ 374.1 Materials  374.2 Single Crystal Specimen Preparation ^  384.3 Polycrystalline Nickel Specimen Preparation  424.4 Test Temperature ^  424.5 Test Solutions  424.6 Electrochemical Test Cell and System ^  454.7 Pitting Scan Test Technique ^  484.8 Test Procedure ^  484.9 Pit Morphology Examination  495 RESULTS^  505.1 Pitting Corrosion Behavior of Sn, Zn and Al Single Crystals ^ 505.1.1 Pitting Potential Dependence on Crystallographic Orientations ^ 515.1.2 Pit Morphology ^  595.1.3 Effect of pH Buffer and Halides on Pitting of Zn ^ 755.1.4 Effect of Halides on Pitting of Al Single Crystals  775.1.5 Overall Summary of Corrosion Behavior of Single Crystals ^ 855.2 Pitting Corrosion Behavior of Polycrystalline Nickel ^ 865.2.1 Polarization Behavior in Nitrate and Sulfate Solutions^ 865.2.2 pH Effect on Pitting of Nickel in 0.1 M NaC1  895.2.3 Effect of Nitrate Inhibitors in Chloride Solutions ^ 905.2.4 Effect of Buffers in Chloride Solutions ^  985.2.5 Comparison of Inhibitor with Buffers  1045.2.6 Halide Ion Effect ^  1075.2.7 Summary  1136 E-pH and X-pH DIAGRAMS^  1146.1 Construction of E-pH and X-pH Diagrams ^  1146.1.1 Chemical Equilibrium  1146.1.2 Electrochemical Equilibrium^  1156.1.3 E-pH Diagrams ^  1166.1.4 X-pH Diagrams  1186.1.5 Thermodynamic Data and Assumption of Activity of Species ^ 120vi6.2 Diagrams for H20-Metal and H20-Halide-Metal Oxide Systems ^ 1216.2.1 Sn ^  1226.2.2 Zn  1266.2.3 Al  1326.2.4 Ni ^  1387 DISCUSSION  1447.1 Pit Initiation Stages and Governing Factors ^  1447.2 Electrode Kinetics ^  1497.3 Dynamic Nature of the Passive Film  1507.4 Local Solution Chemistry during the Film Breakdown/ReformationProcess ^  1517.5 Halide Complex Formation and Pit Initiation Theory ^ 1677.6 Evidence from Experiments on Environmental Effects  1727.7 An Explanation of the Pitting Dependence on CrystallographicOrientations ^  1797.8 Critical Pitting Potential ^  1827.9 Extension of the Proposed Theory to Other Aspects of Pitting ^ 1848 CONCLUSIONS ^  1879 REFERENCES  189APPENDIX I ^  197APPENDIX II  199APPENDIX III ^  202APPENDIX IV  204List of TablesTable 1. Aggressive ions in pitting corrosion of metals^ 10Table 2. Minimum halide concentration for pitting on various metals ^ 15Table 3. Values of slope B in Equation (2.1) for various metals ^ 16Table 4. Structure, melting points and growing orientations of single crystals ^ 37Table 5. Electrochemical polishing parameters for Zn, Sn and Al ^ 41Table 6. Test solutions for pitting corrosion of single crystals  43Table 7. Test solutions for pitting corrosion of nickel ^  44Table 8. Experimental summary for single crystals  50Table 9. Critical pitting potentials obtained from five oriented faces of tin ^ 51Table 10. Measured pit wall angles from the pits formed on the tin (001)surface ^  59Table 11. Summary of corrosion behavior of single crystals ^ 85Table 12. Critical pitting potentials and AE, p, for polycrystalline nickel inchloride solutions at pH 10.5 ^  105Table 13. Standard chemical potentials of substances at 25 °C ^ 121Table 14. Standard chemical potentials for Sn systems at 25 °C  123Table 15. Standard chemical potentials of substances for Zn systems at 25 °C ^ 127Table 16. Standard chemical potentials of substances for Al systems at 25 °C ^ 133Table 17. Standard chemical potentials of substances for Ni systems at 25 °C ^ 138Table 18. Minimum halide concentration for the complex formation ^ 174viiList of FiguresFigure 1 A Pourbaix E-pH diagram for Al - H 2O system with corrosion,passivation and immunity zones ^  3Figure 2 A typical polarization curve for the passive metal in a non-aggressivesolution^  5Figure 3 A typical cyclic polarization curve for the pitting corrosion of thepassive metal in an aggressive solution (solid line) ^  12Figure 4 An experimental E-pH diagram for pitting corrosion of iron inchloride solution [1111 ^  27Figure 5 A X-ray back reflection Laue pattern from the tin (001) surface ^ 39Figure 6 A schematic diagram of the specimen used in pitting tests ^ 40Figure 7 A schematic diagram of the cell used in pitting tests ^ 46Figure 8 A schematic diagram of the corrosion measurement system used inpitting tests ^  47Figure 9 Potentiodynamic polarization results obtained on tin (011) and (111)faces in 0.1 M NaC1 + 0.5 M NaNO3 at pH 6.0 ^  54Figure 10 Variation of the critical pitting potential with crystallographicorientation of tin in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 6.0^ 55Figure 11 Potentiodynamic polarization curves obtained on three differentlyoriented surfaces of zinc 0.1 M NaC1 at pH 9.2^  56Figure 12 Variation of the critical pitting potential with crystallographicorientation of zinc in 0.1 M NaC1 at pH 9.2  57Figure 13 Variation of critical pitting potential with crystallographicorientation of zinc in 0.1 M NaC1 + 0.5 M NaNO3 at pH 9.2 ^ 58Figure 14 Pits formed on the tin (001) face (a, b, and c), and crystallographicfacets (pit walls) in the pit (d) ^  60Figure 15 A light interference photograph taken on the pitted (001) face of tin61Figure 16 A schematic view of the pit on the tin (001) face and thedetermination of the pit wall angle ^  62viiiFigure 17 Relationship between the apparent crystallographic angle and thesteps formed on pit walls ^  66Figure 18 Pit morphology of the tin (111) face (a, b); crystallographic facetsin the pit (c) ^  67Figure 19 Pit morphology of the tin (011) face (a, b); crystallographic facetsin the pit (c)  68Figure 20 Pit morphology of the tin (110) face (a); crystallographic facets inthe pit (b) ^  69Figure 21 Pit morphology of the tin (100) face ^  70Figure 22 Pit morphology of the zinc (0001) face  71Figure 23 Pit morphology of the zinc (1010) face ^  72Figure 24 Pit morphology of the zinc (1120) face  73Figure 25 Pits formed on the cleaved basal plane (0001) of zinc ^ 74Figure 26 Polarization curves of the zinc (0001) face in 0.1 M unbuffered andNa2CO3/NaHCO3 buffered halide solutions at pH 9.2 ^ 76Figure 27 Potentiodynamic polarization results obtained on the Al (100) facein halide solutions at pH = 3.4 ^  79Figure 28 Potentiodynamic polarization results obtained on the Al (100) facein halide solutions at pH = 6.0  80Figure 29 Exposed (111) faces on the Al (100) surface corroded in 0.1 M HFat pH 3.4^  81Figure 30 A crystallographic pit formed on the Al (100) in 0.1 M NaC1 at pH6.0  82Figure 31 A thick salt film formed on the Al (100) in 0.1 M NaF at pH 6.0 ^ 83Figure 32 Salt crystals formed on the surface of Al in 0.1 M NaF at pH 6.0 ^ 84Figure 33 Potentiodynamic polarization curves of Ni in 1.0 M NaNO3 , 1.0 MNa2SO4 and 1.0 M NaC1 at pH = 10.5 ^  87Figure 34 Potentiodynamic polarization curves of Ni in 1.0 M Na2SO4 and 1.0M NaC1 at pH = 2.5 ^  88Figure 35 A potentiodynamic polarization result for Ni in 1.0 M NaCI at pH =2.5 ^  91ixFigure 36 Potentiodynamic polarization results for Ni in 1.0 M NaC1 at pH =4.5, 6.5, 8.5 and 10.5 ^  92Figure 37 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH =12.5 ^  93Figure 38 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH =14.0  94Figure 39 Effect of bulk solution pH on the critical pitting potential of Ni inunbuffered 1.0 M NaC1 solution ^  95Figure 40 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with0.001, 0.01 and 0.1 M NaNO3 at pH = 10.5 ^  96Figure 41 Effect of nitrate concentration on the critical pitting potential of Niin 1.0 M NaC1 at pH = 10.5^  97Figure 42 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with0.001, 0.01 and 0.1 M NaHCO3/Na2CO3 buffer at pH = 10.5^ 99Figure 43 Effect of carbonate buffer concentration on the critical pittingpotential of Ni in 1.0 M NaC1 at pH = 10.5 ^  100Figure 44 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with0.001, 0.01 and 0.1 M Na2HPO4/Na3PO4 buffer at pH = 10.5 ^ 101Figure 45 Effect of phosphate buffer concentration on the critical pittingpotential of Ni in 1.0 M NaCl at pH = 10.5. ^  102Figure 46 Potentiodynamic polarization results for Ni in unbuffered 1.0 MNaBr and Na2HPO4/Na3PO4 buffered 1.0 M NaBr at pH = 10.5 ^ 103Figure 47 Comparison of the increment in the critical pitting potential (AEcp)with the presence of NaNO3 , NaHCO3/Na2CO3 and Na2HPO4/Na3PO4 in 1.0 MNaC1 at pH 10.5 ^  106Figure 48 Polarization results for Ni in 1.0 M NaC1 and 1.0 M NaBr at pH =10.5 ^  108Figure 49 Polarization results for Ni in Na 2HPO4/Na3PO4 buffered 1.0 M NaC1and 1.0 M NaBr at pH 10.5 ^  109Figure 50 Polarization curves for Ni in 1.0 M NaF and 1.0 M NaC1 at pH 10.5^  110Figure 51 Polarization curves for Ni in 1.0 M NaF at pH 6.0 and 1.0 M NaC1at pH 4.5 ^  111Figure 52 A polarization result for Ni in 1.0 M HF at pH 3.1 ^ 112Figure 66 X-pH diagram for H2O -C1 -species at 10-`, 10-4, 10 .6- NiO system; activities of all solute^  143Figure 53 E-pH diagram for H 2O - Sn system; activities of all solute species at10 2, 10 -4, 10 ^  124Figure 54 E-pH diagram for H 2O - C1 - Sn system; activities of chloride at 10°and 10 1 ; activity of other solute species at 10 -6 ^  125Figure 55 E-pH diagram for H 2O - Zn system; activities of all solute species at1072, 10-4, 10 ^  128Figure 56 E-pH diagram for H 2O - CO32- - Zn system; Activity ofH2CO3/HCO 3 -/CO 32- at 10-1 ; activities of other solute species at 10 -2, 10-4, 10.6 ^ 129Figure 57 E-pH diagram for H 2O - C1 - Zn system; activities of chloride at10" .3 and 10 1 ; activity of other solute species at 10-6 ^  130xiFigure 58 X-pH diagram for H2Ospecies at 10', 10 -4 , 10-6 ^- Cl- - Zn(OH)2 system; activity of all solute131 Figure 59 E-pH diagram for H 2O -10', 10-4 and 10-6Figure 60 E-pH diagram for H 2O - Cr - Al system; activities of chloride at 10°and 10 1 ; activity of other solute species at 10 -6 ^  135Al system; activities of all solute species at 134Figure 61 X-pH diagram for H2O -F -solute species at 10', 10 -4 , 10-6 ^Figure 62 X-pH diasram for H2O -Crsolute species at 10', 10 -4, 10-6Al203 .3H20 system; activity of all136- Al203 .3H20 system; activity of all 137Figure 63 E-pH diagram for H 2O - Ni system; activities of all solute species at10 - , 10-4 and 10 -6 ^  140Figure 64 E-pH diagram for H 2O - Cr - Ni system; activities of chloride at 10 °and 10 1 ; activity of other solute species at 10 -6 ^  141Figure 65 X-pH diagram for H 2O -F -species at 10', 10-4, 10-6 ^NiO system; activities of all solute142Figure 67 Several stages occurring during the pit initiation process ^ 145Figure 68 Structure of the interface between metal substrate and bulk solution^  147Figure 69 An active site with a radius of r during oxide film breakdown. ^ 153Figure 70 Change in local pH with variation of r and i ^ 157Figure 71 Change in local halide concentration with variation of r and i at agiven bulk halide concentration of 0.01 M ^  159Figure 72 Variation of local halide concentration with bulk halideconcentration at a given anodic current density of 1 A/cm 2 ^  160Figure 73 Local solution chemistry profile due to a diffusion process acrossthe diffusion layer ^  162Figure 74 Effect of i on the local pH for Ni, Zn, Sn and Al in a continuousfilm breakdown process at 8 = 10' cm ^  165Figure 75 Variation of local halide concentration with bulk halideconcentration in a continuous film breakdown process at 8 = 10-2 cm ^ 166Figure 76 A generalized X-pH diagram ^  168Figure 77 A schematic diagram showing processes in pit initiation ^ 170Figure 78 Change of the local solution pH with the bulk solution pHcalculated for Ni using an one-dimensional diffusion model ^ 177xiiLIST OF SYMBOLS AND ABBREVIATIONSa^ electron transfer coefficient3, Y^reaction orders5 diffusion layer thicknesselectrode overpotentialwavelength of thallium lightstandard chemical potentialp^ densityD diffusion coefficientE electrode potentialE°^standard electrode potentialEpp^critical pitting potentialEr^repassivation potentialEDX energy dispersive X-ray analysisF^ Faraday constantAG° standard free energy changeh^ thickness of the oxide filmi current densityio^ exchange current densityi1^limiting current densityip^passive current densityJ diffusion fluxK equilibrium constantKm (m = 1,2, ...)^ stability constantxivM^ molarity (mole/liter)M^subscript for metalsox^ subscript for oxides and hydrated oxidesQ electric charger^ radius of an active siteR reaction rateSCE^ saturated calomel electrodeSHE standard hydrogen electrodeSEM^ scanning electron microscopyState:aq^ aqueousg gaseous1 liquidsolidV^ volumew molecular weight[xLin^ minimum halide concentration for theformation of dominant soluble complexesz electric charge of speciesACKNOWLEDGEMENTSI would like to express my sincere thanks to my supervisors, Dr. Desmond Tromansand Dr. Fred Weinberg for their advice and guidance throughout this research. My thanksare also extended to other faculty and staff members, and fellow graduate students in thecorrosion group for their help.Special thanks go to my wife and son for their support and understanding during thiswork.The author is also grateful for the financial support provided by the UniversityGraduate Fellowships, and by the Research and Teaching Assistance of this Department.XVINTRODUCTION^ 11 INTRODUCTIONThe passivity of metals is important in commercial applications. Many corrosion resistantalloys, including stainless steels, Ni-base alloys, aluminum alloys and titanium alloys, have beendeveloped to provide materials which have good corrosion resistance based on the formation ofprotective (passive) films. However, in some cases, these materials can exhibit severe localizedcorrosion which can sharply restrict their use under certain service conditions. Several types oflocalized corrosion behavior include: (1) pitting corrosion, (2) crevice corrosion (3)intergranular corrosion and (4) stress corrosion cracking and corrosion fatigue. Localizedcorrosion 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 theyobserved can be divided into the following categories: (1) general corrosion 15%, (2)stress-induced corrosion cracking 39%, (3) pitting corrosion 8%, (4) other localized corrosion19%, (5) high temperature corrosion 6% and (6) others 13%. LaQue [2] , on the basis of his 44years of experience, considered that localized corrosion was responsible for about 90% ofmetals which failed by corrosion.Extensive studies have been carried out on passivity, breakdown of passive films andlocalized corrosion phenomena in efforts to find better alloy/environment combinations toreduce the frequency of localized corrosion. Many investigations in the past 50 years E3H5I havebeen undertaken to examine the fundamentals of passivity and its localized breakdown usingchemical, electrochemical and physical methods. The reported results are very extensive, toolarge to be referred to comprehensively in this literature survey. The present review is confinedto pit initiation on passivated metal surfaces, caused by localized electrochemical breakdown ofpassive films.LITERATURE REVIEW^ 22 LITERATURE REVIEW2.1 PassivityPassivity of a metal surface was initially found on iron in concentrated nitric acid in1790, according to Uhlig 161 . The passive oxide film was first isolated from the surface ofiron, 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 byelectrochemical and physical techniques such as electrochemical reduction [81 [91 , AC,impedance techniques [101-1121 optical elliopsometrytechniques of Auger Electron Spectroscopy (AES), X-ray Photoelectron Spectroscopy (XPS)and Secondary Ion Mass Spectroscopy (SIMS) [1514191 . The thermodynamics and kinetics ofpassivation have become more clearly understood, and many passive film properties havebecome known.2.1.1 Thermodynamics and Kinetics of PassivationThe thermodynamic approach to passivity is mainly the result of contributions byPourbaix 1201 , who proposed that a necessary condition for passivation is the formation ofa stable oxide, or hydroxide film. The most useful guide to passivation is via Pourbaix'spotential-pH diagrams (E-pH diagrams) for metal-water systems calculated fromaqueous thermodynamic equilibrium data. An E-pH diagram shows the stable regions forindividual species, so that corrosion (stable soluble species), immunity (stable metal) andpassivation (stable oxides) zones can be identified (Figure 1). However, E-pH diagrams11314141 , and the surface analysis10.50-2-2.5-3-1-0.5-1.5Al - Water system0^2^4^6^8^10^12^14pHFigure 1 A Pourbaix E-pH diagram for Al - H 2O system with corrosion, passivation and immunity zonesLITERATURE REVIEW^ 4alone are only a guide to possible conditions for passivation. Other factors may beequally important, such as the presence of imperfections and defects in the film, and theadherence of the film to the metal.Passivation is better defined from a consideration of the corrosion kinetics. Ametal is passive if it substantially resists corrosion in a given environment despite amarked thermodynamic tendency to react 16j. The passivation phenomenon can bedescribed kinetically by an anodic polarization curve (Figure 2) obtained on a metal in anaqueous solution without aggressive species (such as halides ions). Three differentpotential regions are identified as active, passive and transpassive regions on thepolarization curve. In the active region metal undergoes general corrosion until thepassivation potential (Er) is reached. After this point, the anodic current densitymarkedly drops to an extremely low value (passive current density, ip, in the order ofgA/cm2) in the passive region. It is in the passive region where many metals and alloysare 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 solublespecies with a higher oxidation state [21], [22]. Sometimes, if the transpassive potential ishigher than the potential for the decomposition of water, the transpassive behavior ismarked by large anodic current increases corresponding to the evolution of oxygen.Transpassive-------------------------LITERATURE REVIEWi PCurrent Density (i)Figure 2 A typical polarization curve for the passive metal in a non-aggressive solution5LITERATURE REVIEW^ 62.1.2 Properties of the Passive Film on MetalsIt 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 asolid-phase oxide layer (or hydrated oxide layer).The adsorption layer model [6] assumes that adsorbed oxygen atoms form a stabletwo-dimensional structure of mixed 02- anions with metallic ions on the surface of ametal. At the optimum ratio of metal to oxygen ions, Uhlig [23] suggested that theadsorption film is more stable than the 3-dimensional metal oxide film. Although such amono-layer adsorption may lead to the subsequent growth of the oxide film, Uhligproposed that passivity was very much dependent on the adsorption layer.However, the available experimental evidence suggests that the passive film ismore likely to be a solid-phase oxide layer formed on the metal surface [21], [24]. Strongevidence 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 andmeasured the composition of the oxide film, its oxidation state and the thickness. Thereis now general agreement that the passive layer on the metal is composed of oxide layerswith 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 ironsurface has two oxide sub-layers, with an inner Fe 304 layer and a y-Fe2O3 outer layer.The XPS results obtained by Strehblow [28] also confirmed this two-layer structure of thepassive film on iron. Chromium enrichment has been found in passive films on stainlessLITERATURE REVIEW^ 7steels [15]-(18], and Cr-rich mixed oxides are the main components of these films. Thethickness of the passive film on iron and its alloys, such as stainless steels, varies fromlnm to lOnm [281. The hydrated Ni (II) oxide was found on nickel passivated in both acidand alkaline solutions in the form of NiO•H 20 or NiOOH. The thickness of the oxidelayer was found to be 2 - 8 nm [291E301 . At more positive potentials, other types of oxideswith higher oxidation states, such as Ni 203 and Ni02, were assumed to be present.The different modifications of Al(III) oxide and hydroxide are thermodynamicallystable in neutral solutions [3114331 . Passive film oxides of Al20 3 , Al203 •H20 andAl203 .3H20 have been reported on the surface of aluminum by many authors, the degreeof hydration being dependent on the prevailing conditions. Under gaseous oxidationconditions, Al203 is formed with a largely amorphous structure and a thickness up to 15nm [311 . Altenpohl and Post [331 reported that a film of Al203 •H20 was formed above 75°C, and Al203 .3H20 below 75 °C, in double distilled water. The films consisted of twolayers. 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 tobe the most protective. The oxide film on aluminum under anodizing conditions couldgrow up to 100 gm, however, only the inner layer, 3-5nm thick, was responsible for thepassivation behavior. The outer layer, composed mostly of primary dissolution-precipitation products, was non-protective [34] .Oxides films formed on zinc and tin have been less studied. In the near neutral pHrange, ZnO and Zn(OH) 2 were reported and found to be relatively insoluble andprotective [381 , and Zn(OH)2 is the most stable form according to Pourbaix [201. In alkalinesolutions, thick porous oxide or hydroxide films of zinc were formed [361 , which couldgrow to a visible thickness. Little information appears to have been reported on theLITERATURE REVIEW^ 8passive film of Sn. The film formed on a tin surface is complex because of the existenceof two oxidation states, Sn(II) and Sn(IV). The Sn oxides, SnO and Sn0 2, arethermodynamically 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 moreprotective, less susceptible to breakdown by aggressive ions, exhibit greaterrepassivation rates, and are more ductile than crystalline films. Hoar [381 also claimedthat the passivity of a non-crystalline film is superior to that of the crystalline film. Thepassive film formed on iron has been found to be crystalline, but the passive films onFe-Cr alloys are reported to be amorphous [391 . It is believed that Cr in the Fe-Cr alloypromotes the formation of a non-crystalline film [391 . It has been claimed that the oxidefilm formed on nickel is crystalline [40], [41]. The passive films formed on aluminum havebeen reported to be non-crystalline [31], [421 Hydration of the oxide film (either as boundwater or as Off bonds) facilitated the formation of the non-crystalline structure in Al andstainless steels [19], ^Lawless [43] has reported the formation of crystalline oxide filmson Cu, Ni and Zn under low temperature gaseous oxidation conditions, but the structureof 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 bysome authors to be n-type semiconductors [21], [26], having a higher electronicconductivity. The higher electronic conductivity facilitates the kinetics of oxygenevolution at higher anodic potentials. However, the passive films formed on aluminum,zirconium and titanium have very low electronic conductivities [211 , which hinders theLITERATURE REVIEW^ 9oxygen evolution kinetics. Consequently, a high anodic potential results in the growth ofoxide films on these metals. The passive oxide films on metals such as Zn and Sn havean 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 anddepletion, secondary phase particles and non-metallic inclusions in the metal substrateproduce imperfections in the grown passive films 14614511 . These imperfections areassociated directly with the positions of the defects existing in the substrate, and are themain nucleation sites for localized corrosion of commercial alloys. Wood [521 suggestedthat defects in the passive films (film flaws) were produced when the oxide film formedover mechanical surface defects (such as scratches), and voids formed by vacancycoalescence in the substrate. Flaws may also occur in the growing film at substrate grainboundaries. According to Chao et al. [531 as-grown point defects are present in crystallineoxide films, and they suggest that these cation and anion vacancies in the passive oxidefilm play an important role in the growth and breakdown of the film.LITERATURE REVIEW^ 102.2 Pitting Corrosion - Localized Film BreakdownPitting corrosion is a result of the localized breakdown of the passive film. Pittingcorrosion occurs on the passivated metal surface in an environment containing aggressiveions such as Cl - , Br-, I- , C104 and SCN - . Of these, chloride ions are considered to be themost widely encountered and most aggressive anions. Table 1 lists the most common anion- metal systems where the localized breakdown of passivity occurs. Pitting corrosion isreviewed by Smialowska [45] in her comprehensive reference book titled "Pitting Corrosionof Metals ".Table 1. Aggressive ions in pitting corrosion of metals [451Iron C1, Br- , 1-,C104Nickel Cl-, Br",rStainless Steels Cl-, Bf, SCN-Aluminum Cl-, Br- , r, C104, SCN-Titanium Cl-, Br-, rZinc CL, Br- , r, C104, Br03 -Tin Cl"Cadmium Cr, Br-Zirconium Cl-, Br- , TLITERATURE REVIEW^ 112.2.1 Pitting Corrosion Phenomena and CriteriaIt has been well recognized that a certain anodic electrode potential has to bereached in order for pits to initiate on a passive metal surface [45] ' [54] . A typicalsingle-cycle anodic polarization curve (solid line) in the aggressive solution is givenschematically in Figure 3. There are two characteristic potentials in this figure: (1) thecritical pitting potential (E cp) and (2) the repassivation potential (E r). These twocharacteristic potentials divide the polarization curve into three regions. Above E cp , pitswill 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 berepassivated. 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 localizedbreakdown in the aggressive solution, which gives rise to a sudden increase in anodiccurrent and an anodic current hysteresis loop during the backward potential scan. Thecritical pitting potential, E cp, is considered to be a major parameter in the study of pitinitiation. 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 ofpitting corrosion in a metal/environment system can be determined from the criticalpitting potential.2. To evaluate the pitting susceptibility of metals. The effects of environmentaland metallurgical factors, such as temperature, pH, solution concentration, and alloycomposition, can be determined from the measurement of the critical pitting potentials.Eq.),E r...0-..-0-...........-00....^Transpassive---...........--Pit Initiation Current LoopIRepassivationActiveLITERATURE REVIEW^ 12i PCurrent Density (i)Figure 3 A typical cyclic polarization curve for the pitting corrosion of the passive metal inan aggressive solution (solid line)LITERATURE REVIEW^ 133. To develop pit initiation mechanisms. An understanding of the factorsgoverning the pit initiation process and critical pitting potential may lead to moreeffective development of pitting resistant alloy-environment systems.Besides the potential criterion, there are other useful criteria, such as the criticaltemperature 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 EnvironmentsMost metals are subject to pitting corrosion in halide solutions. There is aminimum concentration of halide for pit initiation of a given metal, below which pittingdoes not occur. These minimum halide concentrations are listed in Table 2.The aggressiveness of the halide species varies with different metals [451 . Foraluminum, iron, stainless steel and nickel, chloride ions are most aggressive, followed bybromide and iodide ions:Aggressiveness: Cl- > Br > I -However, for titanium and tantalum, bromide and iodide are more aggressive thanchloride:Aggressiveness: Br - > I- > Cl-LITERATURE REVIEW^ 14The critical pitting potential (E, p) is dependent on the concentration of halides inthe solution. An increase in the halide concentration will give rise to a decrease in thecritical pitting potential. A relationship between Ecp and halide concentration [X] hasbeen 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 electrolytecompositions.LITERATURE REVIEW^ 15Table 2. Minimum halide concentration for pitting on various metals *Metal Halide Ion Minimum Concentration, MIron Cl- (a) 0.0003Iron Cr (b) 0.0005Iron CF (c) 0.003Fe-5.6Cr Cr (a) 0.017Fe-11.6Cr Cl- (a) 0.069Fe-20Cr Cr (a) 0.1Fe-24.5Cr Cr (a) 1.0Fe-29.4Cr Cr (a) 1.0Fe-18.6Cr-9.9Ni Cl" (a) 0.1Nickel CF (a) 0.001Titanium Br- (d) 0.002* Source: Z. S. Smialowska, Pitting Corrosion of Metals, NACE Publication, 1986(a)H2SO4 + NaC1 solution(b)Phthalate buffer + NaC1, pH = 5(c)Borate buffer + NaC1, pH = 8.4(d)KBr solutionLITERATURE REVIEW^ 16Table 3. Values of slope B in Equation (2.1) for various metals *Metals B, VoltsIron 0.06 - 0.2Iron-base alloys 0.04 - 0.068Nickel and its alloys 0.071 - 0.078Al-base alloys 0.05 - 0.13Cadmium 0.03 - 0.18Titanium 0.11Zirconium 0.06 - 0.065* Source: Z. S Smialowska, Pitting Corrosion of Metals, NACE Publication, 19862.2.3 Effect of pH and BuffersThere are many studies dealing with the effect of bulk solution pH on pittingcorrosion [6°]-4653 . In general, the critical pitting potential is not affected, or hardlyaffected, in the pH range from acidic to slightly alkaline values. The critical pittingpotential 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 Ecp wasunchanged on iron in chloride solutions of pH 7 - 10, but a higher critical pittingpotential was obtained above pH 10. Similar results were reported for nickel, stainlesssteels 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. ItLITERATURE REVIEW^ 17was found that buffers affected the pitting corrosion behavior, and the presence of thebuffer in the solution increased the critical pitting potential [661, [67] Heusler and Fischer[66] reported that the critical pitting potential of iron increased with the concentration ofbuffers in 0.01 M NaC1 solution. The increase in the critical pitting potential alsodepends on the types of buffers. The critical pitting potential was raised about 150 mVin borate buffer, but was raised only 30 mV in phthalate buffer, when the concentrationof the buffers changed from 0.1 to 1.0 M in 0.05 M NaCl. Drogowska et al. [67] foundthat bicarbonate and phosphate buffers raised the critical pitting potentials of tin by a fewhundred mV's in chloride solutions.2.2.4 Effect of TemperatureAccording to the Arrhenius rate equation, reactions proceed more rapidly at highertemperature. Consequently, it would be expected that pit initiation and pit growth occurmore readily with increasing temperature. It was found that the critical pitting potentialdecreased 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 stainlesssteel in 0.5 M NaC1, and found that the critical pitting potential (Ecp) was a linearfunction of reciprocal absolute temperature (1/T), with a slope of 500 VK° in thetemperature range of 6 - 40 C°. Above 40 C° the effect of temperature on E cp was lesspronounced. Smialowska [68] reportedthat a linear function (E ci, = a - bT) existedbetween Ecp and temperature (T) for ferritic and austenitic stainless steels with a slope ofabout 3 mV/°C, but the linear relation was no longer valid for Mo-containing stainlesssteel above 70 C°, where the critical pitting potential showed little change.LITERATURE REVIEW^ 182.2.5 Effect of Solution FlowThe effect of solution flow on pitting corrosion has been investigated using arotating electrode in an electrolytic cell. Powers and Wilfore [73] reported that the criticalpitting potential of Ti-6A1-4V alloy was shifted from 1.8 V to 4.3 V (SCE) when therotational speed was changed from 0 to 5000 rpm. Riskin and Turkovakaia 1741 found thatEcr, for 18Cr - 8Ni stainless steel was shifted to a more positive value when the electrodewas rotated at 3000 rpm in the chloride solution. Increases in Eci, were also reported foraluminum, where the critical pitting potential was raised about 50 mV when therotational speed was changed from 0 to 2000 rpm 1753 . A 50 mV increase in the criticalpitting 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 inother investigations [77] ' 1781 .2.2.6 Pitting of Metal Single CrystalsKruger [791 studied the effect of crystallographic orientation on the film breakdowntendency for single crystals of iron. His study showed that the resistance to pittingincreases as the surface approaches the (100) orientation, with the lowest pittingresistance on the closest packed planes { 110}. It was also found that the pit densityvaried 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 topitting attack.A more detailed study of crystallographic orientation effects has been conductedby Yasuda et al. [W on aluminum single crystals in chloride-containing solutions, usingLITERATURE REVIEW^ 19potentiodynamic and galvanostatic methods. The results from potentiodynamicmeasurements showed that the critical pitting potential increased with thecrystallographic orientation in the following order: { 111 } < {110} < { 1001. Based ongalvanostatic testing results, the pit density and pit area were found to decrease with thecrystallographic orientation in the pattern: { 111 } > { 1101 > { 100 }. This indicated thatthe close packed planes {111} had the lowest pitting resistance when evaluated either bythe critical pitting potential or by pit density and pit area. However, the critical pittingpotentials were reported to be independent of crystallographic orientations in Al-Cualloy single crystals r801 .The dependence of passivity on the crystal planes was reported for nickel singlecrystals by Latianision et al. [811 . The passive current densities on nickel single crystalsurfaces decreased in 0.5 M H2SO4 solution in the following order: { 100} > {110} >11111, indicating that the passive film formed on the { 111 } faces was most protective.Garz et al.E821 reported that the pitting resistance of nickel crystals is dependent on thecrystallographic orientation of the tested surface. It has been found that the criticalpitting potential in 0.5 M NiC12 solution increased in the order: {100} < {110} < {111} .Lei et al. [831studied the breakdown of the passive film on single crystals and polycrystalsof 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. Theinduction time for the pitting of Ni { 100} surfaces was reported to be slightly longer thanthat 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 dataavailable on {110} and {111} oriented surfaces from their investigation to compare with{ 100} orientations.LITERATURE REVIEW^ 20It is worth noting that the pitting resistance of the {111}, { 110} and { 100} faces ofNi is in the reverse order from that reported for Al, although both metals have the F.C.C.crystal structure.2.2.7 Pit MorphologyCorrosion pits may be either crystallographic or non-crystallographic. Acrystallographic pit exhibits pit walls, whose surface traces are consistent with theintersection of specific crystal planes with the specimen surface. Consequently, the pitmorphology is related to the crystal structure. The walls of crystallographic - type pitsare composed of crystalline planes, but the wall morphology may become quite complexif composed of multiple steps, where each step face has a different variation of thecrystalline facet. Non-crystallographic pits are usually hemispherically shaped, anindication of an isotropic dissolution process within pits, with no evidence ofcrystallographic facets.Hemispherical pits were observed on iron in 0.5 M SO42- + 0.1 M Cl - solution byVetter and Strehbolw [601 . Herbsleb and Engell [84] found a layer of sulfate on the bottomof hemispherical pits in iron, which they considered controls the isotropic dissolution ina diffusion-controlled process. They suggested that the sulfate anions enhance theformation of a viscous layer inside a pit, which controls the dissolution rate by a masstransfer process. Pickering and Frankenthal [851 claimed that in a H 2SO4 + NaCl solutionpits were polished into hemispherical pits in a later growth stage. Hemispherical pitswere also observed on stainless steels and aluminum 186],[87].LITERATURE REVIEW^ 21Crystallographic pits have been found on iron in NaC1 solutions [601 , and in aHC1O4 + NaC1 solution [851 . Brauns and Schwenk [88] observed crystallographic pits onstainless steel in a chloride solution at a potential slightly above the critical pittingpotential (Ecp). Latanision [811 , Garz et al. [82] and MacDougall [301 reported that theappearance of crystallographic pits on nickel varies with the crystallographic orientationof 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 showedthat pits formed on Al {100}, {110} and { 111 } faces were all crystallographic pits andthe pit facets were {100}. Crystallographic pits, bounded by (1010) planes, were alsofound in zinc [891 .The transition from crystallographic pits to hemispherical pits depends on the alloycomposition, 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 + 0.01M chloride solution [601 . Scully [901 found that the morphology of pits on Al depends onthe kind of inhibitor added to the chloride solution. Crystallographic pits were found atthe potential close to the critical pitting potential in stainless steels [883 , but at higherpotential the pits tended to be non-crystallographic. The addition of alloying elementstended to change crystallographic pits into non-crystallographic pits on Al and Ni [801 ' [91] .2.2.8 Alloying Element EffectsAlloying additions to an alloy can change the pitting corrosion behavior of thematerial. Adding Mo and N to austenitic stainless steels dramatically increases theresistance to pitting of the steels [92], [93] . The addition of Cu to Al also raises the criticalLITERATURE REVIEW^ 22pitting potential [80] . Alloying elements can also affect the pitting behavior of metalsingle 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 crystallographicorientations of single phase Al-Cu single crystals, indicating that the addition of Curesults in the pitting independence of crystallographically oriented surfaces. Pits, whichwere initially crystallographic in pure metals, were reported to change tonon-crystallographic pits when Cu (>0.5%) was added to Al 1801 , and Mo (>1%) added toNi [9 ' ] .2.2.9 Inclusions, Grain Boundaries and DislocationsThe detrimental effect of sulfide inclusions on pitting resistance has been reportedfor carbon steels and stainless steels. In a study of pit nucleation conducted bySmialowski et al. [94] , it was found that pits nucleated preferentially at mixed manganeseand iron sulfide inclusions which were present either in the form of separated particles oras 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 stainlesssteels [50], [51], [94], ^Pits were also seen to nucleate at other kinds of inclusions such asoxides, silicates and carbides 1961. SEM observation of inclusions before and after pitinitiation demonstrated that pits could be nucleated directly at sulfide or oxide inclusionsGrain boundaries have been shown to be a preferred sites for pitting corrosion [46],[47], [80], [97]. In sensitized austenitic stainless steels chromium carbide precipitatespreferentially along grain boundaries and diminishes the content of Cr in the adjacentmatrix, which makes it more susceptible to pitting corrosion [46] . Segregation of P atLITERATURE REVIEW^ 23grain boundaries is also considered to cause preferential localized attack in stainlesssteels [471. A study on aged Al-Cu bi-crystals found that the precipitation of Al 2Cu alongthe grain boundaries leads to pit nucleation and growth in the Cu-depleted region at grainboundaries [981 . However, not all grain boundaries are sensitive to pitting corrosion. Itwas reported that there was no preferential pitting corrosion occurring at grainboundaries in solution-treated austenitic stainless steels in which the Cr-depletion zonesalong grain boundaries are eliminated during the solution treatment [991. No preferentialpitting at grain boundaries was detected on aluminum bi-crystals [981 . In contrast to agedAl-Cu bi-crystals, it was found that grain boundaries were not the preferential locationsfor pitting to occur in the solution-treated Al-Cu bi-crystals, where the Al 2Cu precipitatesand Cu depletion zones were eliminated at the grain boundaries [981 . Therefore, thesusceptibility of grain boundaries to pitting is due primarily to local chemicalinhomogeneities such as precipitates, solute segregation and depletion at grainboundaries. It appears unlikely that alloys undergo a preferential pitting attack at grainboundaries when these chemical inhomogeneities are eliminated by proper heattreatments.The effect of dislocations on pit initiation in commercial alloys is less significantthan inclusions and grain boundaries. Even in single crystals the effect of dislocationson pit initiation is not convincing. Wyon et al. E10°1 found that the number of pits formedon pure aluminum is much less than the dislocation density, whose observation showedthat pit formation depended on the breakdown sites in the passive film rather than theemergent dislocations at the metal/passive film interface. Edeleanu et al. [Ion also foundno correlation between pit initiation sites and dislocations in their study on aluminum.Kruger [791 , using transmission electron microscopy (TEM), showed that the breakdownLITERATURE REVIEW^ 24sites on iron single crystal surfaces are not related to dislocations. In a study of Fe-Cralloy single crystals in MgCl 2 solution, Ahlers and Riecke [1021 found that the pit densityin plastically deformed crystal specimens with a high dislocation density was the same asthe unstrained crystal with a much lower dislocation density. A recent study on Albi-crystals, which contained low angle tilt grain boundaries [981 that may be considered asa two-dimensional array of edge dislocations with their Burger's vectors parallel to thebi-crystal surface, showed no preferential boundary attack. Thus, the experimentalevidence mostly indicates that dislocations do not directly act as preferential sites forpitting corrosion. However, Haruyama et al. [1031 studied the anodic behavior of ironwhiskers substantially free of dislocations, and twisted whiskers, in a chloride containingsolution. They claimed that the breakdown of the passive film could occur anywhere inthe film, but pitting corrosion could occur only when the breakdown was located at anemergent dislocation. Hence, there is no general agreement on the effect of dislocationson pit initiation.2.2.10 Solution Chemistry inside PitsSuzuki et al. [1041 showed that for three stainless steels, 304L, 316L, and18Cr-16Ni-5Mo, the pH inside the pits dropped to 0.6-0.8, 0.06 -0.17 and -0.03 - 0.08respectively in a neutral chloride solution. Peterson et al. [1051 found that the pHdecreased to 1.2 -2.0 within pits for a stainless steel in seawater with a pH of 8.0. Butleret al. [1061 measured local pH changes during pit growth in a bulk chloride solution with apH 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 observedLITERATURE REVIEW^ 25within 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 thebulk solution pH of 8) [107], [109].Halide ions can accumulate within growing pits, and their concentration can reachas high as 10 M. Suzuki et al. ""observed a considerable increase in the chlorideconcentration (to 6 M) in artificial pits. Mankowski and Smialowska [110] measured theconcentration of chlorides within naturally grown pits on a stainless steel in 0.5 M NaC1+ 0.1 N H2SO4, and found that the chloride concentration depended on the stage of pitdevelopment. The maximum chloride concentration rose to more than 10 M at a pitdiameter of about 0.5 mm, and then decreased to a steady value of 2.5 M as the pits grewlarger.In general, the solution inside the pit can be significantly different from the bulksolution. The solution chemistry inside the pit exhibits a much lower pH and higherhalide concentration than the bulk solution, which conditions facilitate the further growthof pits2.2.11 Application of E-pH Diagrams to Pitting CorrosionTheoretical E-pH diagrams predict possible passivity conditions in metal - watersystems. However, their application to aqueous solutions containing aggressive ions islimited when the passive film is subject to localized breakdown. Some attempts havebeen made to use E-pH diagrams to study localized corrosion in halide solutions. Figure4 shows an experimental E-pH diagram for iron in the presence of chloride ionsconstructed by Pourbaix r ill] . The pitting corrosion zone, stable passive (perfectLITERATURE REVIEW^ 26passivity) zone, and the conditions in active pits are indicated in Figure 4. These zoneswere 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 pittingcorrosion of 18Cr-10Ni-Ti stainless steel at different temperatures. So far, only a fewexperimental E-pH diagrams have been constructed for application to localizedcorrosion, 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 explainpitting corrosion phenomena if there is the possibility of forming complexes with theaggressive anions..• _qr^ 41.^..^..^.^.^e^5^.IIP\ •^.^.^.,.^. ....,.^..^.^.^:---Pitting,  • .^ ..^.Imperfect pcsattlieity _ 0 _aprotection _ct pAssiilleyimmunity2^4-4^6^8^IC) 12 19pH-2LITERATURE REVIEW0 2 y 6 8 10 12 1927Figure 4 An experimental E-pH diagram for pitting corrosion of iron in chloride solution 11113LITERATURE REVIEW^ 282.2.12 Theories of Pit InitiationAccording to Kruger [211 , a successful theory for pit initiation should explain thefollowing:1. A critical pitting potential (Ecd must be exceeded for pit initiation.2. Aggressive ions are needed for the breakdown of the passive film on the metalsurface.3. Breakdown occurs at highly localized sites.Many theories have been proposed to account for pit initiation on the passivatedmetal surface. They are summarized as follows:1. Adsorbed Ion Displacement Theory: This mechanism was suggested byKolotyrkin [113] and Uhlig [541 . The passive film is considered to be an adsorbed film ofoxygen. Passivity breakdown occurs when chloride ions are more readily adsorbed thanpassivating ions (0 2- , OH-). The chloride ions displace the passivating ions, resulting inpit initiation. Therefore, the critical pitting potential is the potential at which theadsorption of aggressive ions occurs, and pit initiation sites are the weakest points in thepassive 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 asufficiently small diameter to penetrate through the pores in the film. Breakdown of thefilm occurs when the Cl - ions reach the metal surface.LITERATURE REVIEW^ 29Hoar et al. [114] proposed that pit initiation is caused by the penetration of anionsthrough the film under the influence of an electrostatic field across the film. Small ionsmore 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 isaccompanied by the exchange of a passivating ion (e.g. OH-), which occurs at siteswhere the metal-oxygen bond is weakest.3. Chemical-Mechanical Breakdown Theory: In 1967, Hoar E ll6iproposed amechanical breakdown model for pit initiation. The adsorbing anions were postulated toreplace adsorbed water on the film and reduce the interfacial tension or interfacial energyof 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 chargeoccurs. The cracks or splits thus produced in the film result in the breakdown ofpassivity. Another chemical-mechanical breakdown model was developed by Sato [117] .He suggested that a high potential field could lead to mechanical rupture of the passivefilm 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 pressureexceeds the critical compressive strength of the film. The role of chloride ions is toretard the repassivation process after the breakdown of the passive film. They play a lessimportant role in the rupture of the passive film.4. Localized Acidification Theory: In 1937, Hoar [118] suggested that pits developbecause hydrolysis of corrosion products in the pits causes acidification. Galvele [119]developed a model, based on the assumption that metal cations hydrolyze insidemicropits which already exist on the surface, and that the movement of cations out of pitsLITERATURE REVIEW^ 30is controlled by diffusion processes. Pit initiation is presumed to occur when a criticalpH value is reached by the local acidification within micropits. In his model, Galveleemphasized the influence of acidification more than the specific role of chloride ions.Another model that considers the change in the local chloride concentration insidepits is due to Hisamatsu 11201 . He assumed that there was a critical chloride concentrationin the pit electrolyte, C* , above which, the pit repassivation process is hindered andnucleated pits grow.5. Depassivation - Repassivation Theory: This theory emphasizes the concept thatpassivity breakdown is a dynamic process. Videm [1211 presumed the existence ofdynamic breakdown-repair events in the passive film. In the absence of aggressiveanions, or below the critical pitting potential, film breakdown is followed by rapidhealing, whereas in the presence of aggressive ions, and at the potential above Ecp , therepassivation process is blocked. The role of chloride ions is to retard the repassivationprocess. Pitting is thus expected to occur when the rate of passivity breakdown is greaterthan that of repassivation.Wood's theory [521' ^assumes the presence of pre-existing flaws in the passivefilm. His theory may be classified with the ion penetration model, presuming thataggressive ions penetrate through the flaws. However, it is more likely that a dynamiccrack-heal process occurs at the base of flaws. The flaws will be repassivated if thecritical pitting potential is not met for pit initiation.6. Chloride Nucleus Formation Theory: A critical chloride nucleus model has beendeveloped based on the localization of chloride ions on the surface of the passive film.LITERATURE REVIEW^ 31Such localization was found experimentally on the passive surface of iron and stainlesssteels by Janik-Czachor et al. [123], [124] anda was supported by the perturbation theorysuggested by Okada " 251 .A chloride cluster model was first proposed by Hoar [1261 , who assumed that severalchloride ions are adsorbed together on the passive film to form a localized chloridetransitional complex with the cations in the passive film. Heusler and Fischer "271suggested that the formation of a two-dimensional chloride nucleus causes the thinningand eventual breakdown of the passive film. The chloride nucleus theory [581 " 281 suggeststhat the formation of chloride nuclei is a dynamic process occurring randomly on thepassive surface. When a pitting criterion (such as critical pitting potential) is met, stablecritical chloride nuclei are formed and penetrate the oxide film layer, resulting in thebreakdown of the passive film and pit initiation.7. Point Defect Theory: This theory has been developed by Chao et al. [53] and Linet al. [1291. Point defects in the crystalline oxide film tend to accumulate to form a void atthe metal/oxide film interface. When the void grows to a critical size, the passive filmcollapses locally , resulting in the localized breakdown of the passive film. Based on thismodel, they have successfully derived a relationship between the critical pitting potentialand halide concentration, and between electrode potential and the induction time forpitting.LITERATURE REVIEW^ 322.3 Remaining ProblemsSeveral remaining problems are outlined below, such as the pitting dependence ofsingle crystals on crystallographic orientations, halide complex formation, and the change inlocal solution chemistry, which are the subjects of the research in the present thesis.2.3.1 The Pitting Dependence on Crystallographic OrientationsA 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 beenless systematically studied on single crystals of other metals. Therefore, it is still notclear whether orientation-dependent pitting is a common phenomenon on passivatedmetals.There is still no answer to the question why single crystals have different pittingresistances on the differently oriented surfaces. Pitting studies on iron and aluminumsingle crystals [79],[79] have suggested that the most closely packed surfaces ({ 110 } for ironand { 111 } for aluminum) exhibit the lowest pitting resistance, and these surfaces alsohave the highest dissolution rates [1301 ,[131] ^the other hand, the pitting of nickel singlecrystals indicated that the most closely packed { 111 } faces have the highest criticalpitting 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^ 332.3.2 Solution pH and Local Solution pH ControlThe critical pitting potential has been reported to be independent of bulk solutionpH in an acidic to weakly alkaline pH region. However, there is a pronounced increasein the critical pitting potential when the solution becomes strongly alkaline (see Section2.2.3). The local pH within growing pits has been found to be significantly differentfrom the bulk solution pH, thus the local pH rather than the bulk solution pH is importantin the pit initiation process (see Section 2.2.10). There is a lack of experimentalevidence on local pH changes during pit initiation, possibly due to difficulties faced inmonitoring the local surface pH. The significance of a local pH change during pitinitiation is still unknown.Pitting corrosion behavior is anticipated to be different if a local pH is controlledby buffering the solution. Previous work has shown that the addition of buffersincreased the critical pitting potential (see Section 2.2.3). However, the effect of thebuffers was explained on the basis that the buffers acted as competitive adsorptioninhibitors without attention paid to the local pH control. A recent study by Tromans andSun [134] has suggested that the presence of buffers controls the local pH and prevents thebreakdown of the film on Cu, rather than the buffer being competitively adsorbed on thesurface.2.3.3 Halide Complexes and Diagrams Dealing with Halide ComplexesThe interaction of halides with the oxide film had been long ignored until theformation of halide complexes was proposed by Hoar [135] for iron, and by Foley [136] foraluminum. Halide complex formation emphasizes the role of aggressive ions in theLITERATURE REVIEW^ 34localized breakdown of passivity, and it can be presented in a diagram calculatedthermodynamically in the same manner as Pourbaix's E-pH diagrams. Pourbaix's E-pHdiagrams on simple H 20-metal systems give useful information on the possibleconditions for passivity of metals, but their application is limited when they are appliedto the localized breakdown of passivity in halide solutions. There are few theoreticalE-pH studies which include halide complexes, though the thermodynamic data areavailable for the formation of halide - metal complexes. It is possible to introduce halidecomplexes into conventional E-pH type diagrams, or introduce them into diagrams thatuse halide concentration (X) and pH as the variations on the axes (termed X - pHdiagrams). Such diagrams of water - halide - metal systems, based on thermodynamicconsiderations, could prove to be useful to an understanding of the possible role ofcomplex halide ions in the pit initiation process.2.3.4 Pit Initiation MechanismsThere 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. Thelocal acidification model and chemical - mechanical breakdown model are faced withthe difficulty of explaining the aggressive role of halides during film breakdown.Surface analysis techniques have not detected the penetration of chloride ions into thefilm. The point defect theory cannot be applied to amorphous films. The chloridenucleus theory fails to give information on the specific chloride complexes which lead topit initiation. No generally accepted theory has been developed to account forcrystallographic orientation - dependent pitting phenomena.OBJECTIVE^ 353 OBJECTIVEThe purpose of the present investigation is to extend the fundamental knowledge ofcrystallographic orientation - dependent pitting phenomena and to enhance the mechanisticunderstanding of the pit initiation process in halide solutions. These goals are pursued byconducting experiments on pure metal single crystals of Sn, Zn and Al, and polycrystalline Ni inhalide solutions, and by constructing thermodynamic diagrams to represent equilibria in severaldifferent water-halide-metal systems. The general sequence of procedures is outlined below.1. Pitting Corrosion of Single CrystalsThe dependence of the critical pitting potential on the crystallographic orientation of Snand Zn single crystals is investigated in aqueous chloride solutions. The pit morphology isexamined using scanning electron microscopy (SEM).2. Effects of Solution pH and HalidesBuffers are used to help control the local surface pH and assess the significance of a localpH change during the pit initiation process on polycrystalline Ni. Pitting corrosion isinvestigated in F, Cr and Br solutions to correlate the aggressiveness of the halide speciesaccording to their tendency to form metal halide complexes on Zn and Al single crystals, andpolycrystalline Ni.3. E-pH Diagrams for Water-Halide-Metal SystemsNew E-pH diagrams are constructed for Sn, Zn, Al and Ni in halide environments, basedon 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^ 364. Pit Initiation MechanismAn approach to the mechanism of pit initiation is made by taking into consideration thedynamic nature of the oxide film formation, local solution chemistry and halide - metalcomplexes. A theory is proposed to explain the effect of crystallographic orientations, solutionpH, buffers and halide ions on pitting behavior.EXPERIMENTAL^ 374 EXPERIMENTAL4.1 MaterialsSingle crystals of zinc, tin and aluminum, with a purity of 99.995%, were used in thestudy of the dependence of pitting corrosion resistance and pit morphologies oncrystallographic orientation. The single crystals were grown by Professor Fred Weinbergusing the Bridgman technique with pre-oriented seeds in a horizontal furnace. The size ofthe grown crystals was about 8 x 6 x 150 mm. The crystal structures of zinc, tin andaluminum, their melting points and crystal growth directions are listed in Table 4.Table 4. Structure, melting points and growth directions of single crystalsMetal Structure Melting point, °C Growth DirectionZn 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 environmentaleffects 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 nickelspecimens.EXPERIMENTAL^ 384.2 Single Crystal Specimen Preparation(a) Crystallographic Orientation Determination:The crystallographic orientations of Sn, Zn and Al single crystals were determined bythe X-ray back-reflection Laue technique [137j . A Philips Model PW 1830 X-ray generatorwas used at a working voltage of 35 KV and a current of 20 mA. The distance between thesingle 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 alongthe basal plane after being cooled in liquid nitrogen. Due to the difficulty in distinguishingbetween the (1010) and (1120) Laue patterns, the deformation twinning method was used todetermine these orientations. The Zn single crystal was twinned along the well-definedtwinning (1012) plane by applying a tensile stress parallel to the basal plane r 1381 .(b) Sampling and MountingOnce the orientations were determined, the crystals were cut into small pieces (about10 mm long) along the desired crystallographic angles to obtain specifically orientedsurfaces using spark machining or a jewelry saw (for the zinc single crystal, spark machiningcauses cleavage along the basal plane). Five differently oriented surfaces, (100), (110),(111), (011) and (001) were prepared for tin, while (0001), (1010) and (1120) orientedsurfaces were prepared for zinc single crystals. The deviation of the oriented surface fromthe desired crystallographic orientation was ± 5°. A conducting wire (copper wire) wassoldered to the sampled crystal, then mounted in cold-curing epoxy resin (Lecoset). Amounted specimen is shown schematically in Figure 6.EXPERIMENTAL^ 39[110][100]Figure 5 A X-ray back reflection Laue pattern from the tin (001) surfaceFigure 6 A schematic diagram of the specimen used in pitting testsEXPERIMENTAL^ 41(c) Mechanical and Electrochemical PolishingThe specimen was mechanically polished up to grit 600 paper, and then to 1.0 gmusing a water base diamond suspension (BUEHLER METADI). The surface was thenpolished electrochemically in order to remove the deformed layer resulting from mechanicalpolishing. Table 5 lists the electrolytes and conditions for the electrochemical polishing ofzinc, tin, and aluminum single crystals.Table 5. Electrochemical polishing parameters for Zn, Sn and AlCrystal Electrolyte Cathode Current, A/cm2 Temperature, °C Time, sHC1O4 200 mL(CH3C0)20 800 mLPt 0.15 30 360Zn * KOH 250 gH2O 750 mLPt 0.22 20 900.Al **HC1O4 62 mLEthanol 700 mL2-butoxyethanol 100 mLH2O^137 mLPt 3.85 20 20* P. V. Shigelev, Electrolytic and Chemical Polishing of Metals, Freund Publishing House, 1974** Handbook of Metals, ASM, Vol. 9, P. 353, 1985The electrochemically polished specimen was cleaned with distilled water and acetoneand dried in cold air. In order to prevent preferential corrosion attack at the specimen/epoxyinterface, the interface was coated with a cellulose acetate lacquer (red lacquer), leaving anexposed surface area varying from 0.2 -0.6 cm2 for different samples.EXPERIMENTAL^ 424.3 Polycrystalline Nickel Specimen PreparationNickel specimens (15x15 mm) were cut from 3 mm thick Ni sheet, then mounted andmechanically polished in the same manner as the single crystal specimen shown in Figure 6,without further electrochemical polishing. For nickel, the specimen/epoxy interfaces wereprotected by a GLYPTOL coating (an alkaloid enamel paint) instead of the red lacquer usedfor the single crystals. The exposed surface area of the test specimen was about 1 cm 2 .4.4 Test TemperatureAll experiments were conducted at room temperature in the temperature range20-23°C. Temperature changes within this range had no apparent effect on the experimentalresults.4.5 Test SolutionsHalide - containing solutions were used in the pitting corrosion tests. All the solutionswere prepared using reagent grade chemicals and distilled water. The pH of the solutionwas 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 pHmeter. The halide - containing solutions which were used in tests for single crystals arelisted in Table 6.EXPERIMENTAL^ 43Table 6. Test solutions for pitting corrosion of single crystals.Crystal Solution No. Test Solution pHSn A-1 0.1 M NaC1 + 0.5 M NaNO3 6.0A-2 0.1 M NaC1A-3 0.1 M NaC1 + 0.5 M NaNO3Zn A-4 0.1 M NaC1 + 0.1 M NaHCO3/Na2CO3 9.2A-5 0.1 M NaBrA-6 0.1 M NaC1 3.4Al A-7 0.1 M NaBrA-8 0.1 MNaF 6.0In the study of the environmental effects of halides, pH, and buffers on the pittingcorrosion of nickel, several types of solutions were prepared for experiments. Type Isolutions, consisting of NaNO3 or Na2SO4 , were used as a basis for comparison with halidesolutions. Their pH values were adjusted by the addition of NaOH (for B-1 and 2 solutions),or H2SO4 (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 halideions. Type III were chloride plus nitrate solutions. The pH values of Type II and IIIsolutions were adjusted by the addition of NaOH or HC1. Type IV were halide solutionsbuffered by Na2CO3/NaHCO 3 or Na3PO4/Na2HPO4 , which were used to investigate thesignificance of pH control in pitting corrosion. All the solutions used for tests on nickel arelisted in Table 7.EXPERIMENTAL^ 44Table 7. Test solutions for pitting corrosion of nickelType Solution No. Composition pHB-1 1.0 M NaNO3 10.5Type I B-2 1.0 M Na2SO4 10.5B-3 1.0 M Na2S°4 2.5B-4 1.0 M NaC1 2.5B-5 1.0 M NaC1 4.5B-6 1.0 M NaC1 6.5B-7 1.0 M NaC1 8.5B-8 1.0 M NaC1 10.5Type II B-9 1.0 M NaC1 12.5B-10 1.0 M NaC1 14.0B-11 1.0 M NaBr 10.5B-12 1.0 M NaF 6.0B-13 1.0 M NaF 10.5B-14 1.0 MHF 3.1B-15 1.0 M NaC1 + 0.1 M NaNO3 10.5Type III B-16 1.0 M NaC1 + 0.01 M NaNO3 10.5B-17 1.0 M NaC1 + 0.001 M NaNO3 10.5B-18 1.0 M NaC1 + 9.5x10-2 M Na2CO3 + 5.0x10 -3 10.5M NaHCO3B-19 1.0 M NaC1 + 9.5x10-3 M Na2CO3 + 5.0x10 -4 10.5M NaHCO3B-20 1.0 M NaC1 + 9.5x10-4 M Na2CO3 + 5.0x10 -5 10.5M NaHCO3Type IV B-21 1.0 M NaCl + 8.5x10 -2 M Na2HPO4 + 10.51.5x10-2 M Na3PO4B-22 1.0 M NaC1 + 8.5x10-3 M Na2HPO4 + 10.51.5x10-3 M Na3PO4B-23 1.0 M NaC1 + 8.5x10-4 M Na2HPO4 + 10.51.5x10-4 M Na3PO4B-24 1.0 M NaBr + 8.5x10-2 M Na2HPO4 + 10.51.5x10-2 M Na3PO4EXPERIMENTAL^ 454.6 Electrochemical Test Cell and SystemThe electrolyte test cell was a glass cell (EG&G C, K47 cell), Figure 7, containing fiveinlets for a working electrode (specimen), two platinum counter electrodes, a saturatedcalomel reference electrode (SCE) and a nitrogen gas purger. For fluoride solutions, aTEFLON (polytetrafluoroethylene) cell was used with the same configuration as the glasscell. A KC1 saturated agar-agar gel salt bridge was connected to the glass calomel referenceelectrode to prevent contact with the fluoride - containing solution.A microprocessor controlled potentiodynamic polarization system (EG&G Model350A) 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 computerthrough a serial port (RS232C) and saved on floppy discs. Figure 8 shows schematically theoverall system used for the electrochemical measurements.WE - Working Electrode (Specimen)CE - Counter Electrode (Pt)SCE - Saturated Calomel Reference ElectrodeFigure 7 A schematic diagram of the cell used in pitting testsNitrogen Gas CylinderFigure 8 A schematic diagram of the corrosion measurement system used in pitting testsEXPERIMENTAL^ 484.7 Pitting Scan Test TechniqueA cyclic potentiodynamic polarization technique, called the pitting scan test, was usedin the present study. The test started at the corrosion potential (open - circuit potential), thenthe potential was increased in a positive (anodic) direction at a scan rate of 0.5 mV/s until apredetermined current threshold (5 mA/cm 2) was reached. Then the potential scan wasreversed to the negative (cathodic) direction. The test was completed when the potentialscanned 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 localizedcorrosion on the passivated surface. The schematic appearance of the resulting polarizationcurve is similar to that shown in Figure 3 (solid line). The critical pitting potential (E cp) isthus determined from the polarization curve. The measured value of the critical pittingpotential is reported with reference to the Saturated Calomel Electrode (SCE). Theconversion of potential values between the SCE and the Standard Hydrogen Electrode(SHE) is:E (SCE) = E (SHE) - 0.241 (Volts)4.8 Test ProcedureDissolved oxygen was removed from the test solution by deaerating with N2 for onehour before the specimen was placed into the cell. Subsequently, the specimen was put inthe 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 atthe corrosion potential. The nitrogen purge operated continuously during the test. AfterEXPERIMENTAL^ 49each test, the specimen was washed with distilled water and ethanol, then dried and stored ina desiccator until required for pit morphology examination. Most of the tests were repeatedtwo or three times to confirm the reproducibility and validity of the results.4.9 Pit Morphology ExaminationPit morphology was examined by scanning electron microscopy, using an HitachiS-2300 SEM. Secondary electron imaging was employed, using an excitation energy of 20KeV.In addition, a Zeiss Interference Microscope (Model 3001) was used to examinesurface topography. This technique uses a monochromatic light source (thallium lightsource) to produce interference bands across the specimen surface. The interference patternsfrom a pitted surface enable the orientation of the facets in the pits to be determined inrelation to the top surface of the specimen.RESULTS^ 505 RESULTS5.1 Pitting Corrosion Behavior of Sn, Zn and Al Single CrystalsTable 8 summarizes the pitting tests conducted on tin, zinc and aluminum singlecrystals. The differently oriented surfaces of the tin and zinc single crystals were tested inchloride solutions to investigate the dependence of the critical pitting potential on thecrystallographic orientations. The effect of the Na 2CO3/NaHCO 3 buffer was studied only onthe zinc (0001) surface. Several halide solutions were used in the investigation of the pittingcorrosion of the zinc (0001) and aluminum { 1001 surfaces in order to compare theaggressiveness of the different halide ions.Table 8. Experimental summary for single crystalsSingleCrystalSurfaceOrientationEffect Test Solution(001)(011) CrystallographicSn (111) Orientation A-1(110)(100)(0001) Crystallographic A-2Zn (1610) Orientation A-3(1120)Zn (0001) Buffer A-4Zn (0001) Halide A-5A-6Al { 1001 Halide A-7A-8RESULTS^ 515.1.1 Pitting Potential Dependence on Crystallographic OrientationsFigure 9 shows the potentiodynamic polarization curves obtained on the Sn (011)and (111) oriented faces in 0.1 M NaC1 + 0.5 M NaNO3 solution at pH 6.0. The resultsclearly demonstrate the difference in their critical pitting potentials. The critical pittingpotential of the Sn (011) face is about 25 mV higher than that of the (111) face. Fivedifferently oriented surfaces of Sn single crystals have been tested in 0.1 M NaC1 + 0.5M NaNO3 solution at pH 6.0. The results of these pitting tests, with respect to effects onthe critical pitting potentials, are listed in Table 9 and plotted in Figure 10.Table 9. Critical pitting potentials obtained from five oriented faces of tinTin 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 -0.310The critical pitting potential for each Sn surface was measured three times, thescatter for each surface orientation being ± 5 mV. It is clear from Table 9 and Figure 10that the critical pitting potentials for the (001), (110), (011) and (100) surfaces of Sn arevery close and lie between -0.282 V (SCE) to -0.290 V (SCE). However, the (111)RESULTS^ 52surface of Sn has a critical pitting potential that is about 22-25 mV lower than the othersurface orientations. Therefore, the (111) face of Sn has the lowest pitting corrosionresistance among the five tested surface orientations. The critical pitting potentials of theSn 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 singlecrystals in chloride solutions. The potentiodynamic polarization curves measured in 0.1M NaC1 at pH 9.2 are shown in Figure 11. In contrast to the polarization curves for tinsingle crystals (Figure 9), there is no sharp potential point (no well-defined E cp), where asudden increase in anodic current occurs, on the polarization curve of zinc singlecrystals. Consequently, an extrapolation technique is used to determine the criticalpitting potential on each orientated surface, as shown in Figure 11.The dependence of Ecp on the crystallographic orientation of the zinc surface isshown in Figure 12. In these experiments only two repeated measurements were madefor each surface and the reproducibility was ± 5 mV. The critical pitting potential isobserved to progressively decrease as the crystallographic orientations are changed from(1010) —> (1120) —> (0001). The critical pitting potentials were also measured on thesurfaces of zinc single crystals in 0.1 M NaC1 + 0.5 M NaNO3 at pH 9.2 (Figure 13). Theresults show that the presence of NaNO 3 raised the Ecp to significantly higher values thanthose observed in 0.1 M NaCl at pH 9.2. However, the order of the dependence of Ecp onRESULTS^ 53crystallographic orientations in 0.1 M NaC1 + 0.5 M NaNO 3 exhibits the same trend as in0.1 M NaCl. Therefore, in both 0.1 M NaC1 and 0.1 M NaC1 + 0.5 M NaNO3 solutions atpH 9.2, the critical pitting potentials of Zn single crystals may be summarized:Ecp for zinc; (10-10) > (11 -20) > (0001)0 2 4 6 8-0.2-0.3-0.4-0.6-0.7-0.8Log(i), nA/cm 2Figure 9 Potentiodynamic polarization results obtained on tin (011) and (111) faces in 0.1 M NaCl + 0.5 M NaNO3 at pH 6.0I •A•i•A••I1-0.26-0.27-0.28IA0Cr)^-0.29..._...,'-0.3a)50...^-0.31C:Y)CIP4-,.CL^-0.32-0.33-0.34(001)^(110)^(011)^(100)OrientationFigure 10 Variation of the critical pitting potential with crystallographic orientation of tin in 0.1 M NaC1 + 0.5 M NaNO 3 at pH6.0LOG(i), nA/cm 21 3 5 7-0.98-1Figure 11 Potentiodynamic polarization curves obtained on three differently oriented surfaces of zinc 0.1 M NaC1 at pH 9.2••••I^ 1 (10 10) (1120)^ (0001)OrientationGIO -1.05cr)> -1.06Tz -1.0711'^-1.08o^-1.090-CY)^-1.1CIP+ a^-1.11"am-1.12-1.13-1.14-1.02-1.03-1.04Figure 12 Variation of the critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 at pH 9.2Orientation(1120)(1010) (0001)•••-0.72-0.73-0.744E' -0.750ta_D) -0.7617 4-/a--0.77-0.78Figure 13 Variation of critical pitting potential with crystallographic orientation of zinc in 0.1 M NaC1 + 0.5 M NaNO 3 at pH 9.2RESULTS^ 595.1.2 Pit MorphologyThe shape of the pits on Sn crystals varied with the change of the crystallographicorientation of the surface. Figure 14 shows the crystallographic pits formed on the tin(001) face in 0.1 M NaC1 + 0.5 M NaNO3 at pH 6.0. These square pits show a 4-foldsymmetry and pit edges are reasonably parallel to [100] or [010] directions. There aremultiple 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. Figure15 shows an interference micrograph which was taken from the tin (001) face. Bymeasuring the number of interference bands on the pit wall, the depth of the pit wascalculated according to the equation: t = nX/2, where t is the depth of the pit, n is thenumber 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, where1 is the half width of the pit (Figure 16). Table 10 gives some pit wall angles measuredfrom 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) surfacen t, gm 1, gm a, °10 2.70 5.1 27.815 4.05 8.7 25.018 4.86 9.7 26.6(c)[010](d)RESULTS^ 60(a)^(b)Figure 14 Pits formed on the tin (001) face (a, b, and c), and crystallographic facets (pitwalls) in the pit (d)RESULTS^ 61Figure 15 A light interference photograph taken on the pitted (001) face of tin0146.0,..._.[010]top view1)A,side viewInterference BandsFigure 16 A schematic view of the pit on the tin (001) face and the determination of the pit wall angleRESULTS^ 63The 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 measuredpit 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 boundingthe 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 becomessmaller, as shown schematically in Figure 17. This was confirmed by the pit wall anglemeasurements performed on the larger pits with steps on the pit walls, where the anglevaried from 15° to 20°.Unfortunately, no clear interference patterns could be obtained from the pitsformed on other tin surfaces. Some calculations were made based on the interceptiontraces of the pit walls with the tin free surface (the pitting tested specimen surface) inorder 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 atthe 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 speculatedto be (010) and (100) planes, referred to as { 1001 tetragonal prismatic planes in thefollowing Sections. At the lower part of the pit in Figure 18.b, the angle between thetwo 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 beRESULTS^ 64parallel to the [211] direction, and the interception of (101) with the (111) surface will beparallel 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 theSn (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 wallsand a pit base as shown in Figure 19. Assuming that the pit base is (011) and three of pitwalls (excluding the right side pit wall) are (101), (101) and (011) planes as shown inFigure 19.c, the interception traces of the (101) and (101) on the tin (011) surface wouldbe the [111] and [111] directions respectively. The measured angle between the traceson 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 interceptiontrace 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 tiltingthe specimen while observing the surface in the SEM, and the surface interception tracesare 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^ 65Figures 22 - 24 show the pit morphology on the (0001), (1010) and (1120) faces ofzinc single crystals respectively in 0.1 M NaC1 at pH = 9.2. The shapes of the pits aredifferent 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 werefound on the zinc surface cleaved along the (0001) plane (basal plane) after pitting in 0.1M NaCl solution at pH = 9.2 (Figure 25).The hexagonally shaped pits formed on the cleaved zinc (0001) surface areconsistent with the hexagonal crystal structure of Zn (Figure 25). The pit edges areparallel to the interception traces of (1010), (1010), (0110), (6110), (1100), (1100) withthe 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 exhibitedcrystallographic features to some extent (Figure 23 and 24). However, it is difficult tospeculate on the possible orientation of the pit walls because no well-defined surfacetraces of the pit walls were observed.less than 28.6Figure 17 Relationship between the apparent crystallographic angle and the steps formed on pit walls(a)Mau 4,0(b)N134.7RESULTS^ 67Figure 18 Pit morphology of the tin (111) face (a, b); crystallographic facets in the pit (c)RESULTS^ 68a-0Figure 19 Pit morphology of the tin (011) face (a, b); crystallographic facets in the pit (c)RESULTS^ 69(a)(b)[00 1]Figure 20 Pit morphology of the tin (110) face (a); crystallographic facets in the pit (b)77110.,..11111.1101*MOW^,...411nraT4 gumRESULTS^ 704,4010101219119%.^::-.0.141.4^-77nis1,--r• - 7., r■_111114.4461...• --wo,^....orPIPIP ...up,............t, _ ,1773111-10 "opa. "79110111rFigure 21 Pit morphology of the tin (100) faceRESULTS^ 71Figure 22 Pit morphology of the zinc (0001) faceRESULTS^ 72Figure 23 Pit morphology of the zinc (1010) faceRESULTS^ 73Figure 24 Pit morphology of the zinc (1120) faceRESULTS^ 74Figure 25 Pits formed on the cleaved basal plane (0001) of zincRESULTS^ 755.1.3 Effect of pH Buffer and Halides on Pitting of ZnThe effect of the NaHCO 3/Na2CO3 buffer on the pitting of Zn single crystals wasinvestigated on the (0001) surface in the 0.1 M NaC1 + 0.1 M NaHCO3/Na2CO3 solutionat pH 9.2. The potentiodynamic polarization curves are presented in Figure 26, wherethey are compared with the unbuffered 0.1 M NaC1 solution at the same pH. Clearly, theNaHCO3/Na2CO3 buffer raised the critical pitting potential by at least 250 mV andproduced a more clearly defined E cp. The addition of buffer did not change the pH of thebulk solution, but it controlled the local pH near the specimen surface so that localacidification did not occur as readily as in the unbuffered solution. TheNaHCO3/Na2CO3 buffer may also have contributed to the extension of the passive zoneof zinc to a lower pH region by the formation of ZnCO 3 . This will be discussed in latersections (see E-pH diagram for H 20-0O32--Zn system in Section 6.2, and Section 7.6).The effect of different halide species on the pitting of the Zn (0001) surface wasinvestigated 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 criticalpitting potential in the bromide solution is more clearly defined and about 100 mVhigher than in the chloride solution, indicating that chloride ions are more aggressivethan bromide ions in promoting the pitting corrosion of zinc.6 8in 0.1 M NaCI + 0.1 M carbonate buffer^ in 0.1 M NaCIin 0.1 M NaBr-0.85-0.8-0.9-0.95-1.15-1.2-1.250 4LOG(i), nA/cm 22Figure 26 Polarization curves of the zinc (0001) face in 0.1 M unbuffered and Na2CO3/NaHCO 3 buffered halide solutions at pH9.2RESULTS^ 775.1.4 Effect of Halides on Pitting of Al Single CrystalsThe Al { 001 } faces were chosen for the investigation of the effect of threedifferent 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 (Figure27), 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 polarizationresults.In Figure 27, the polarization curve in 0.1 M fluoride solution at pH 3.4 shows theactive dissolution of Al, corresponding to general corrosion. The generally corroded Al(001) surface in the fluoride solution at pH 3.4 exhibited crystallographic features, androughened faces were exposed which appeared to be the { 111 } planes (Figure 29). Thepolarization curve in 0.1 M NaC1 at pH 3.4 exhibits a defined passive region below Ecpand passivity breakdown at Ecp. Pitting corrosion occurred on the Al { 100} in thechloride solution at pH 3.4.In Figure 28, pitting corrosion occurred on the Al {001} surfaces in both 0.1 MNaC1 and 0.1 M NaBr at pH = 6.0, and pits appeared to be crystallographic, and boundedby pit walls of Al {100} planes (Figure 30). The critical pitting potential in the bromidesolution is more than 50 mV higher than in the chloride solution. The polarization curveof Al {001} showed a large active dissolution nose near -1.2 V (SCE) in 0.1 M NaF atpH 6.0, followed by an anodic current decrease as a result of thick salt film formation onthe surface (Figure 31). No pitting was observed in the fluoride solution at pH 6.0. Thethick salt film was analyzed by the energy dispersive X-ray (EDX) method. Enrichmentof Al and F was detected, and the salt film is identified to be an aluminum fluoride saltRESULTS^ 78film. By potentiodynamically polarizing the Al { 100} to -0.6 V (SCE) from thecorrosion potential, salt crystals were initially formed on the Al surface, and showeddendritic growth (Figure 32)The above results indicated that the corrosion behavior of aluminum is dependenton the specific type of halide species in the solution. The stability of the passive oxidefilm on aluminum varies with the type of halides in the solution, and the passive film isless stable in the chloride solution than in the bromide solution. The formation of a saltfilm also alters the corrosion behavior of Al.-1W -0.70-0.8-0.9ui-0.2-0.3-0.4-0.5-0.6-1.2-1.3-1.4-1.50^2^ 4^ 6^ 8LOG(i), nA/cm 2Figure 27 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 3.4Figure 28 Potentiodynamic polarization results obtained on the Al (100) face in halide solutions at pH = 6.00 2 4 6 80.60.40.20-0.2-0.4C/)-0.6-0.8-1-1.2-1.4-1.6-1.8LOG(i), nA/cm 2RESULTS^ 81Figure 29 Exposed (111) faces on the Al (100) surface corroded in 0.1 M HF at pH 3.4RESULTS^ 82Figure 30 A crystallographic pit formed on the Al (100) in 0.1 M NaC1 at pH 6.0RESULTS^ 83Figure 31 A thick salt film formed on the Al (100) in 0.1 M NaF at pH 6.0RESULTS^ 84Figure 32 Salt crystals formed on the surface of Al in 0.1 M NaF at pH 6.0RESULTS^ 855.1.5 Overall Summary of Corrosion Behavior of Single CrystalsTable 11 summarizes the results obtained from the pitting tests on tin, zinc andaluminum single crystals in halide-containing solutions.Table 11 Summary of corrosion behavior of single crystalsMetal Solution No. Corrosion Behavior Pit MorphologySn A-1PittingEcp lowest on (111)CrystallographicZnA-2A-3Pitting, Ecp(1010)>(1120)>(0001)CrystallographicZn A-4PittingEcp(buf) > Ecp(unbuf)CrystallographicZn A-5PittingEcp(Br) > Ecr,(C1 -)CrystallographicAlA-6A-7PittingEcp(Br-) > Ecp(Cr)CrystallographicAl A-8, pH=3.4 Active (Exposed { 111})Al A-8, pH=6.0 Active/Passive (Thick salt film)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.RESULTS^ 865.2 Pitting Corrosion Behavior of Polycrystalline Nickel5.2.1 Polarization Behavior in Nitrate and Sulfate SolutionsThe polarization behavior of nickel in 1.0 M NaNO3 and 1.0 M Na2SO4 at pH =10.5 is shown in Figure 33, along with the polarization curve of Ni in 1.0 M NaC1 at thesame pH. The results show passive behavior of Ni in 1.0 M NaNO3 and 1.0 M Na2SO4 atpH 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 thenickel specimen. The equilibrium potential for oxygen evolution at 25 °C can becalculated using the following equations:2H20 = 02 + 4H+ + 4e-^(5.1)E (021H20) = 1.228 -0.0591pH^(V vs. SHE)^(5.2)For the solution at pH 10.5, E (021H20) = 0.608 V (SHE) = 0.367 V (SCE). Thepotential (about 0.6 V (SCE)) where the anodic current increases in Figure 33 is higherthan E (021H20) (0.367 V (SCE)). Furthermore, unlike the polarization curve in thechloride solution, no anodic current hysteresis loop occurred on the polarization curvesin 1.0 M NaNO3 or 1.0 M Na2SO4 . Therefore, the passive film on nickel was stable in1.0 M NaNO3 and 1.0 M Na2SO4 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 potentialscan on the polarization curve, indicating the localized breakdown of the passive filmand the initiation of pitting.1.61.41.210.8L.1.1^0.60U)^0.4>LLI^0.20-0.2-0.4-0.6-0.80^2^4^6^8LOG(i), nA/cm 2Figure 33 Potentiodynamic polarization curves of Ni in 1.0 M NaNO3 , 1.0 M Na2SO4 and 1.0 M NaC1 at pH = 10.5Figure 34 Potentiodynamic polarization curves of Ni in 1.0 M Na2SO4 and 1.0 M NaC1 at pH = 2.50 2 4 6 81.61.41.210.8w 0.60CO^0.4>ul^0.20-0.2-0.4-0.6-0.8LOG(i), nA/cm 2RESULTS^ 89Figure 34 presents polarization curves of Ni in 1.0 M Na2SO4 and 1.0 M NaC1 atpH 2.5. There was an active/passive behavior in 1.0 M Na2SO4 . The passive regionextended to as high as the oxygen evolution potential (E (021H20) = 0.840 V (SCE) atpH 2.5, calculated from Equation 5.2)., and no current hysteresis occurred on thepolarization curve. No localized film breakdown was observed. However, in 1.0 MNaC1 at pH 2.5, the passivity was fully destroyed by the chloride species, and only activebehavior was observed.The results suggest that sulfate and nitrate are not aggressive ions and do notinduce pitting corrosion of nickel. These types of anions, especially nitrate, areconsidered to be inhibitors to pitting corrosion. On the other hand, the chloride ions areconsidered to be aggressive and are needed for the breakdown of the passive film andinitiation of the pitting corrosion on nickel.5.2.2 pH Effect on Pitting of Nickel in 0.1 M NaC1In 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 rangefrom 2.5 to 14. An anodic polarization curve shows that Ni exhibited active behavior in1.0 M NaC1 at pH 2.5 (Figure 35), indicating that Ni underwent general corrosion. Anactive/passive transition occurred in the solution at pH 4.5, followed by pitting at ahigher potential (Figure 36). Pitting occurred on Ni in the solutions in the pH range from6.5 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 passivitybreakdown was detected up to the potential of 0.60 V (SCE), higher than the oxygenevolution potential (E (021H20) = 0.160 V (SCE) at pH 14, calculated from EquationRESULTS^ 90(5.2). The results showed a transition in the polarization behavior of nickel from fullactivation (Figure 35) to localized film breakdown (Figures 36 and 37), and then to fullystable 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 inFigure 39. It is evident that the critical pitting potential is not influenced by pH in theregion of 4.5 - 10.5, but there is an increase in the critical pitting potential at pH 12.5when the solution becomes strongly alkaline. Pitting corrosion of nickel is completelyretarded 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 SolutionsNitrate is usually considered to be a pitting corrosion inhibitor. Therefore, theaddition of nitrate species to the chloride solution might be expected to raise the criticalpitting potential. Pitting scan tests on nickel in 1.0 M NaC1, pH 10.5, containingadditions of 0.001 M, 0.01 M and 0.1 M NaNO3 are shown in Figure 40, and the resultingcritical pitting potentials are plotted in Figure 41. Comparison of Figures 40 and 41 withthe results obtained at the same pH in the absence of nitrate addition (see Figures 36 and39) shows that nitrate additions have no effect on E cp when present in smallerconcentrations of 0.001 M and 0.01 M. However, the addition of 0.1 M nitrate raised thecritical pitting potential by 90 mV.1.21.110.90.80.70.60.5ru^0.40(/)^0.3> 0.2u.i0.10-0.1-0.2-0.3-0.4-0.5-0.60^2^4^6^8LOG(i), nA/cm2Figure 35 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 2.51.21.110.90.80.70.60.50.40.30.20.10-0.1-0.2-0.3-0.4-0.5-0.60^2^4^6^8LOG(i), nA/cm 2Figure 36 Potentiodynamic polarization results for Ni in 1.0 M NaC1 at pH = 4.5, 6.5, 8.5 and 10.51.21.110.90.80.70.6W^0.50CI)^0.4>^0.3Lli^0.20.10-0.1-0.2-0.3-0.4-0.5-0.60^ 2^ 4^ 6^ 8LOG(i), nA/cm 2Figure 37 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 12.51.21.110.90.80.70.60.5W 0.4C.)u) 0.3> 0.2LII^0.10-0.1-0.2-0.3-0.4-0.5-0.6-0.7-0.80^ 2^ 4^ 6^ 8LOG(i), nA/cm 2Figure 38 A potentiodynamic polarization result for Ni in 1.0 M NaC1 at pH = 14.00.80.70.6Ca 0.5(...)C/)0.4Ci.UW0.30.20.10no pitting4AIi^i^A^1AI^I^I^I^I^I^I^I^I^I^I4 6 8 10 12 14pHFigure 39 Effect of bulk solution pH on the critical pitting potential of Ni in unbuffered 1.0 M NaC1 solutionr - —■■■No N. •0 2 4LOG (i), nA/cm 26 8— — 1.0 M NaCI + 0.1 M NaNO3---  1.0 M NaCI + 0.01 M NaNO31.0 M NaCI + 0.001 M NaNO3Ir.m.■^ ■••••• "".go*■■=.^ ••••• "•••▪ .'1.21.110.90.80.70.6• 0.5• 0.4co• 0.3Lir^0.20.10-0.1-0.2-0.3-0.4-0.5-0.6Figure 40 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaNO3 at pH = 10.5-3 -1-2LOG[C] , M0.70.60.50.30.20.1Figure 41 Effect of nitrate concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5RESULTS^ 985.2.4 Effect of Buffers in Chloride SolutionsThe effect of buffer concentration on the pitting behavior of nickel in 1.0 M NaC1solution was investigated at pH 10.5. Two types of buffer solutions were used based on(1) NaHCO3/Na2CO3 additions and (2) Na2HPO4/Na3PO4 additions. They will bereferred to as carbonate and phosphate buffers respectively.The polarization behavior in the carbonate buffer is shown in Figure 42 for totalNaHCO3/Na2CO3 concentrations ranging from 0.001 M through 0.01 M to 0.1 M. Thecorresponding critical pitting potentials are plotted as a function of buffer concentrationin Figure 43. The results clearly show that the increase in the buffer concentration, whilemaintaining a constant pH of 10.5, raised the critical pitting potential. Comparison withFigure 39 at pH 10.5 shows that even the smallest buffer addition of 0.001 M raised Ecpby 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 fortotal Na2HPO4/Na3PO4 additions of 0.001 M, 0.01 M and 0.1 M. The corresponding E SPpotentials are plotted as a function of buffer concentration in Figure 45. It is clear thatthe phosphate buffer behaves in a very similar manner to the carbonate buffer, the E cpincreased with increasing buffer concentration.Some comparative tests were conducted with a 1.0 M NaBr solution, pH 10.5, inthe absence and presence of phosphate buffer having a total Na 2HPO4/Na3PO4concentration of 0.1 M. The polarization results are shown in Figure 46. Again, in asimilar manner to the chloride solution, the critical pitting potential was raised about 250mV in the buffered bromide solution.0 2 4^6^8LOG(i), nA/cm 2In 1.0 M NaCI + 0.1 M NaHCO3 /Na 2 CO 3— In 1.0 M NaCI + 0.01 M NaHCO3 /Na 2 CO3— - — - • In 1.0 M NaCI + 0.001 M NaHCO3 /Na 2 CO3pH = 10.5.^. .■ . ■ . ... .^. ■ '.^'' '''' .i.7.'"0,00 . . .m.... m..... •••• •......•m... l =  1^$1......_ I.^. •••• ''.11.21.110.90.80.70.60.5UT 0.40co^0.3>^0.2Li0.10-0.1-0.2-0.3-0.4-0.5-0.6Figure 42 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M NaHCO3/Na2CO3 buffer at pH= 10.5-1-3 -2LOG[C] , M0.7In 1.0 M NaCI +x M NaHCO 3/Na2 CO3pH = 10.50.60.20.10.50.40.3Figure 43 Effect of carbonate buffer concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5^ In 1.0 M NaCI + 0.1 M Na2 HPO4 /Na3 PO4^ In 1.0 M NaCI + 0.01 M Na 2 HPO 4 /Na3 PO4 In 1.0 M NaCI + 0.001 M Na 2 HPO4 /Na3 PO4—^pH = 10.54LOG(i), nA/cm 2620 81.21.110.90.80.70.60.50 0.40.3LLI^0.20.10-0.1-0.2-0.3-0.4-0.5-0.6Figure 44 Potentiodynamic polarization results for Ni in 1.0 M NaC1 with 0.001, 0.01 and 0.1 M Na2HPO4/Na3PO4 buffer at pH= 10.50.70.60.50.3In 1.0 NaCI + x M Na 2HPO4/Na3PO4pH = 10.5 ■■umIIILIJ_I_LI_LIIIIIIIIIII0.20.1-3^ -2^ -1LOG[C] , MFigure 45 Effect of phosphate buffer concentration on the critical pitting potential of Ni in 1.0 M NaC1 at pH = 10.5.1.21.110.90.80.70.60.5CCI0 0.4U)0.3>W- 0.20.10-0.1-0.2-0.3-0.4-0.5-0.60^2^4^6^8LOG(i), nA/cm 2Figure 46 Potentiodynamic polarization results for Ni in unbuffered 1.0 M NaBr and Na2HPO4/Na3PO4 buffered 1.0 M NaBr atpH = 10.5RESULTS^ 1045.2.5 Comparison of Inhibitor with BuffersThe common feature of carbonate and phosphate buffers, and nitrate inhibitor insolutions, when present in sufficient concentrations, is that they raise the critical pittingpotentials. Their relative effects on the increase of the critical pitting potential, AE cp , in1.0 M NaC1 at pH 10.5 are compared as a function of their concentrations in Figure 47and listed in Table 12. The 6Ecp is defined as Ecp - ecp, where ecp is the critical pittingpotential in the simple unbuffered and uninhibited 1.0 M NaC1 solution at pH 10.5.Consequently, Akp accounts for the net effect of the additives to 1.0 M NaCl solution atpH =10.5.It is evident from Figure 47 and Table 12 that increasing concentrations ofadditives raise AEcp , and that the buffers are much more effective in raising the criticalpitting potential than nitrate additions. Furthermore, the phosphate buffer tends toproduce the highest values. The results indicate that the control of pH by buffering thesolution is an important means of controlling pitting, more important than the use ofnitrate inhibitors.RESULTS^105Table 12. Critical pitting potentials (E, p) and AE,p, for polycrystalline nickel in chloridesolutions at pH 10.5.Type No. Additives to solution Eep, V (SCE) AE,p, VType II B-7 - 0.295 (Ecp°) -B-10 0.1 M NaNO3 0.400, 0.370 0.105, 0.075Type III B-11 0.01 M NaNO3 0.310, 0.295 0.015, 0.000B-12 0.001 M NaNO3 0. 295, 0.230 0.000, 0.005B-13 0.1 MNaHCO3/Na2CO3 0.530, 0.500, 0.480 0.235, 0.205, 0.185B-14 0.01 M NaHCO 3/Na2CO3 0.385, 0.380, 0.400 0.090, 0.085, 0.105B-15 0.001 M NaHCO3/Na2CO3 0.345, 0.340 0.050, 0.045Type IVB-16 0.1 M Na2HPO4/Na3PO4 0.565, 0.540 0.270, 0.245B-17 0.01 MNa2HPO4/Na3PO4 0.440, 0.390 0.145, 0.095B-18 0.001 M Na2HPO4/Na3PO4 0.375, 0.355, 0.34 0.080, 0.060, 0.0450.40.350.306.1^0.25U)0.2a0w0.150.10.050A 1.0 M NaCI + x M Na 2 HPO4 /Na3PO4* 1.0 M NaCI + x M NaHCO3/Na2CO3• 1.0 M NaCI + x M NaNO 3AA000AAaAA0g■■■I^1^I^I^1^i^I^I^l_^I^ii^1^I^lit^tit^I^1-3 -2 -1LOG[C] , MFigure 47 Comparison of the increment in the critical pitting potential CAE& with the presence of NaNO 3 , NaHCO3INa2CO3 andNa2HPO4/Na3PO4 in 1.0 M NaC1 at pH 10.5RESULTS^ 1075.2.6 Halide Ion EffectSome pitting scan tests were performed on nickel in different types of halidesolutions to compare the relative influence of halides on passivity breakdown. Figure 48shows the polarization curve in 1.0 NaBr, pH 10.5, along with the polarization curve in1.0 M NaC1 at the same pH as a comparison. Pitting corrosion occurred in 1.0 M NaBrsolution at pH = 10.5, but the critical pitting potential was about 100 mV higher than thatin 1.0 M NaCl solution at the same pH. A similar result showing about a 50 mVdifference in the critical pitting potential was obtained in the phosphate buffered bromidesolution as compared with the phosphate buffered chloride solution (Figure 49). Theseresults show that the passive oxide film on nickel is more stable in the bromide solutionthan in the chloride solution.The polarization results of Ni in 1.0 M NaF solution at pH = 10.5 and 6.0 areshown in Figures 50 and 51 respectively. The passive film is stable in 1.0 M NaF at pH10.5, and no pitting occurred, as compared with Ni in 1.0 M NaC1 at the same pH. At pH6.0, an extensive region of active behavior was observed in 1.0 M NaF, followed by theonset of passivation near +0.5 V (SCE). This passive film was stable at higher potentialsand pitting was not observed. For comparison, at pH 4.5, passivation occurred at a lowerpotential near -0.2 V (SCE) in 1.0 M NaC1, but film breakdown and pitting occurred at apotential of +0.3 V (SCE). Figure 52 presents a polarization result for nickel in 1.0 MHF at pH 3.1. The result shows that Ni exhibited an active dissolution behavior in asimilar manner to Ni in 1.0 M NaCl at pH 2.5 (Figure 35).1.21.110.90.80.70.60.50.40.30.20.10-0.1-0.2-0.3-0.4-0.5-0.60^2^4^6^8LOG(i), nA/cm 2Figure 48 Polarization results for Ni in 1.0 M NaC1 and 1.0 M NaBr at pH = 10.5Figure 49 Polarization results for Ni in Na 2HPO4/Na3PO4 buffered 1.0 M NaC1 and 1.0 M NaBr at pH 10.50 2 4 6 81.21.10.90.80.70.60.5w• 0.40.3• 0.2LLI^0.10-0.1-0.2-0.3-0.4-0.5-0.6LOG(i), nA/cm 2^ 1.0 M NaCI + 0.1 M Na2HPO4/Na 3PO4 1.0 M NaBr + 0.1 M Na2HPO4 /Na3PO41.61.41.210.80.600.40.21.11 0-0.2-0.4-0.6-0.80^2^4^6^8LOG(i), nA/cm 2Figure 50 Polarization curves for Ni in 1.0 M NaF and 1.0 M NaC1 at pH 10.50 2 4 6 8in 1.0 M NaF at pH 6.0---------------^in 1.0 M NaCI at pH 4.51.61.41.210.8w 0.6C.)0.4LLl^0.20-0.2-0.4-0.6-0.8LOG(i), nA/cm 2Figure 51 Polarization curves for Ni in 1.0 M NaF at pH 6.0 and 1.0 M NaC1 at pH 4.50 2 4LOG(i), nA/cm 26 81.21.110.90.80.70.60.50.40.30.20.10-0.1-0.2-0.3-0.4-0.5-0.6Figure 52 A polarization result for Ni in 1.0 M HF at pH 3.1RESULTS^ 1135.2.7 Summary1. Pitting corrosion of nickel occurs in specific solutions, usually in halide - containingsolutions. Therefore, halides play a very important role in the breakdown of passivityand the initiation of pits. Without the presence of aggressive ions in the solution, nickelwould stay passive as observed in NaNO 3 and Na2SO4 solutions.2. Nickel exhibited active dissolution behavior in 1.0 M NaC1 at pH 2.5. Pitting occurredon Ni in 1.0 M NaC1 in the pH range of 4.5 - 12.5. The critical pitting potential, E cp, wasindependent of the bulk solution pH in the range from 4.5 to 10.5, but E cp increased at pH12.5. No pitting was observed in the solution with pH 14. It is worth noting that Niexhibited a transition in polarization behavior from active dissolution at pH 2.5, throughactive-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 Na2CO3/NaHCO3 and Na2HPO4/Na3PO4 buffers raised the criticalpitting 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 andphosphate buffers were more effective in raising the critical pitting potential than theaddition 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. Bromideions were shown to be less aggressive than chloride ions.E-pH and X-pH DIAGRAMS^ 1146 E-pH and X-pH DIAGRAMS6.1 Construction of E-pH and X-pH Diagrams6.1.1 Chemical EquilibriumFor a general chemical reaction at equilibrium that does not involve any electrontransfer 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 - = H2OThe equilibrium constant K for Equation (6.1) can be expressed as:K — [CT [Di d[A r [Brwhere the square brackets represent the activity of the species. The standard free energychange, AG°, is the algebraic sum of standard chemical potentials, g,° of individualspecies.AG° = ciLec + dp,'D - al.C A - bp.°B^(6.4)(6.2)(6.3)At the chemical equilibrium, we have the relation between AG ° and K:E-pH and X-pH DIAGRAMS^ 115AG ° = —RT1nK^ (6.5)6.1.2 Electrochemical EquilibriumElectron transfer is involved in an electrochemical reaction. A single electrodereaction, or a half cell reaction, is usually used to represent the electrochemicalequilibrium between the oxidized and reduced forms of species. A simple example is thereduction 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 theelectrochemical reaction (6.7) becomes:AG° = pjf + qi.t° Q -^- ji_t°HAt the electrochemical equilibrium, we have the Nernst equation:E = E° RT ln [P]P [Q]qnF [L] l [HPWhere, E is the equilibrium electrode potential, F is the Faraday constant, and E° is thestandard electrode potential, which is governed by the following equation:(6.8)(6.9)AG° = -nFE°^ (6.10)E-pH and X-pH DIAGRAMS^ 1166.1.3 E-pH DiagramsAn E-pH diagram shows the effects of potential and pH on the stability ofindividual species (metals, ions, oxides etc.). Each line in the diagram represents achemical or an electrochemical equilibrium. The regions between the lines show thedomain of stability of each species. The diagrams are plotted with pH as the abscissa andE as the ordinate. In an E-pH diagram there are three types of lines which representthree types of equilibria.(1) Chemical equilibrium involving fr ions:For a general chemical equilibrium:aA + jH+ = cC + dD^ (6.11)We have:AG ° = —RT1n rcr [D id[A]a [HI(6.12)andAG° pH = — 71 (Log rCic [Did[A]a^2.303RT(6.13)Therefore,dpH 0dE(6.14)E-pH and X-pH DIAGRAMS^ 117The equilibrium (6.11) is independent of potential, but dependent on pH, and isrepresented by a vertical line in the E-pH diagram.(2) Electrochemical equilibrium without 11 ÷ ionsFor a general electrochemical reaction:IL + ne = pP + qQ^ (6.15)According to the Nernst Equation (6.9):E = E° — RT In  [Q l4nF^[L,11(6.16)anddE/dpH = 0^ (6.17)The equilibrium (6.15) is independent of pH, but dependent on potential, and isrepresented by a horizontal line in the diagram.(3) Electrochemical equilibrium involving fr ionsFor a general electrochemical reaction:IL + j1I+ + ne = pP + qQ^ (6.18)we have:E = E° RT in WV [QV^ (6.19)nFE-pH and X-pH DIAGRAMS^ 118anddE^2.303jRT_ —dpH^nF(6.20)Therefore, the equilibrium is dependent on both potential and pH , and is represented bya sloping line in the diagram. When 1-1 + appears on the left hand side of the equilibriumequation, the line has a negative slope. The line has a positive slope when fr appears onthe right hand side.6.1.4 X-pH DiagramsThe E-pH diagram is a very concise presentation of multiple equilibria and showsthe thermodynamic stability of individual species for H2O - metal systems. However, formore complicated systems, such as solutions with halides, variables other than pH andelectrode potential will be involved. The conventional E-pH diagrams have somelimitations in the presentation of these complicated systems. For the purpose of thisstudy on passivity and its breakdown, the X-pH diagram has been introduced , for thefirst time, to show the thermodynamic stability of the specific oxide' in the system wherethe formation of halide - metal complexes is involved. The X-pH diagram presents theeffects of solution pH and halide (X) activity on the equilibrium between a specific oxideand halide complexes. Since the formation of a stable oxide film is influenced by thehalide complexes, X-pH diagrams are anticipated to be particularly useful in a study ofthe passivity breakdown of metals in halide environments. The diagrams are constructed1 The term "oxide" in the treatment including hydrated oxides.E-pH and X-pH DIAGRAMS^ 119with pH as the abscissa and X as the ordinate. Each line in the X-pH diagram representsan equilibrium between the oxide and a soluble species, where the soluble species maybe a halide complex. If the equilibrium between oxide and the complex is a chemicalone, the X-pH diagram will be independent of the electrode potential, which is the casefor Zn(OH) 2 , Ni(OH)2 and Al203 .3H20. If the equilibrium between the oxide andcomplex is an electrochemical one, the X-pH diagram should be constructed at a definedpotential (such as in the case of Ni203). This is the limitation in the application of X-pHdiagrams.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 2O or M(OH)n = MO' + nfr(2) Horizontal: an equilibrium without the involvement of fr ionsmn+ + mx- = mx in(m-n)-(3) Sloping: an equilibrium involving both 11 + and halide complexesM(OH)n + nH+ + rriX - = MX,n(m-n)- + nH2OA limited number of X-pH diagrams have been constructed for the hydrated oxidesZn(OH) 2 , Ni(OH)2 and Al203 .3H20 systems in the interest of this study. For details,refer to Section 6.2.E-pH and X-pH DIAGRAMS^ 1206.1.5 Thermodynamic Data and Assumption of Activity of SpeciesIn the construction of the diagrams, we have to know the standard chemicalpotential, [C, for each individual species. Most thermodynamic data used in thefollowing section (Section 6.2) were obtained from Pourbaix's Atlas of ElectrochemicalEquilibria in Aqueous Solutions [201 , and others are from NBS publications [1391 andCritical Stability Constants [140] . Data from Pourbaix's Atlas [20] are not noted with areference number in the following Tables 13 - 17. Data from other sources are notedwith their reference number in Tables 13 - 17. Some thermodynamic data of halidecomplexes are not available in the form of standard chemical potentials, but they aregiven in terms of a stability constant. Therefore, the standard chemical potential hasbeen calculated from the stability constant based on Equation (6.5). For example, K1 isknown 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 ofgaseous 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^ 1216.2 Diagrams for H20-Metal and H2O-Halide-Metal Oxide SystemsSince aqueous solutions are involved in the systems, the stability of water is shown inevery 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)02 + 411- + 4e- = 2H2o^ (b)Table 13. Standard chemical potentials of substances at 25 °CSubstance^State^ji°, KJ/mol.H+^aq^0OH -^aq -157.297H2O 1^-237.180H2^g 002 g^ 0Equilibria (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 thananhydrous 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, the2 However, according to the thermodynamic data reported by NBS rmi 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 energychanges are very small, thus the hydrated oxides are still considered.E-pH and X-pH DIAGRAMS^ 122anhydrous oxides are most stable and should be considered in preference to the hydratedoxides. Nevertheless, in order to be consistent with the Al, Zn and Ni diagrams, hydratedSn-oxides are considered in this Section, and the anhydrous oxides are presented inAppendix II. New E-pH and X-pH diagrams, involving the formation of metal halidecomplexes, are constructed and presented together with the conventional Pourbaix E-pHdiagrams for H2O - metal systems. The inclusion of halide complexes in E-pH and X-pHdiagrams and their applications to pitting are a new and unique contribution of the presentresearch.6.2.1 SnThe formation of stannous hydroxide Sn(OH)2 and stannic hydroxide Sn(OH) 4 onthe tin surface is considered in the aqueous solution, and their stability determines thepassivity of tin. In the chloride solution, Sn - chloride complexes exist in the form ofSnC142- , SnC162- and SnC13 - . The standard chemical potentials of the species considered inthe Sn-Cl_-H20 system are listed in Table 14.Two E-pH diagrams have been constructed for H2O - Sn and H2O - Cl - - Snsystems respectively (Figures 53 and 54). Shaded areas in the diagrams are the passivezones - the domain of hydrated tin oxides. Comparing the two E-pH diagrams, thepassive zone is narrowed by the chloride complex formation when chloride ions arepresent in the solution, indicating the instability of passivity in the C1 --containingsolution.E-pH and X-pH DIAGRAMS^ 123Table 14. Standard chemical potentials for Sn systems at 25 °CSubstance^State^KJ/mol.Sn^s^0Sn(OH)2 s -492.040Sn(OH)4^s^-951.850Sn2+^aq -26.255Sn4+^aq^2.720HSn02 aq -410.030Sn203^aq^-590.280SnC142-^aq -560.910E1401SnC162-^aq^-785.800E'4°JSnC13 aq -431.495E1401Cl -^aq^-131.1680^2^4^6 8 10 12 14Tin - water system1.41.210.80.60.40.20-0.2-0.4-0.6-0.8-1-1.2-1.4pHFigure 53 E-pH diagram for H 2O - Sn system; activities of all solute species at 10 -2, 10-4 , 10-61.41.210.80.60.40.20-0.2-0.4-0.6-0.8-1-1.2-1.4Sn - chloride - water system 0^2^6^8^10^12^14pHFigure 54 E-pH diagram for H 2O - Cl- - Sn system; activities of chloride at 10° and 10 1 ; activity of other solute species at 10-6E-pH and X-pH DIAGRAMS^ 1266.2.2 ZnThe 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 - ispresent in the solution. Four types of zinc - chloride and bromide complexes have beenreported in chloride and bromide solutions respectively (ZnX. (m -2)-, where m=1,2,3 and4). 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 beenreported 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 H2O - Zn, H2O - CO3 - - Zn and H2O -Cl- - Zn systems (Figures 55 - 57), and one X-pH diagram for H 2O - Cl- - Zn(OH)2system (Figure 58). The formation of insoluble ZnCO 3 extends the passive zone of zincto the low pH region in the CO 3 --containing solution. In the chloride system (Figure57), the formation of ZnC142- diminishes the passive zone of zinc. Therefore, the passivefilm of zinc is less stable when chloride ions are present in the solution. Other forms ofZn chloride complexes (such as ZnCr, ZnC12 and ZnC13 -) are also formed, depending onthe chloride concentration, but it is difficult to introduce more than one chloridecomplexes in the E-pH diagram. The effect of chlorides on zinc passivity can be clearlyillustrated in the X-pH diagram (Figure 58). The passive zone becomes narrower withincreasing chloride activity.E-pH and X-pH DIAGRAMS^127Table 15. Standard chemical potentials of substances for zinc systems at 25 °CSubstance^State^if , KJ/mol.Zn^s^0Zn(OH)2 s -559.108Zn2+^aq^-147.210HZn02 -^aq -464.005Zn022^aq^-389.238ZnCr aq^-280.830E14°1ZnC12^aq^-413.028E14°1ZnC13 -^aq^-543.570E1401ZnC142-^aq^-673.026E14°1Cl-^aq -131.168ZnC 03 s^-731.363[139]Zinc-water system0 2 4 6^8pH10 12 140.80.60.40.2-1.2-1.4-1.6-1.8Figure 55 E-pH diagram for H 2O - Zn system; activities of all solute species at 10-2, 10-4 , 10-60.80.60.40.20III1 -0.2U)^.> -0.4LL ^-0.6-0.8-1-1.2-1.4-1.6-1.8Zinc - water - Carbonate system0^2^4^6^8^10^12^14pHFigure 56 E-pH diagram for H 2O - CO32- - Zn system; Activity of H 2CO31HCO 3-/CO 32- at 10-1 ; activities of other solute species at10-2, 10-4 , 10-610.80.60.40.20-0.2> -0.4-0.6-0.8-1-1.2-1.4-1.6-1.81012-^_HZnO 2  (1 06 ))„de-- Zn(OH) 2ZnZinc-chloride-water system[cr]^10.3_ZnCI 4  (1 06 )0^2^4^6^8^10^12^14pHFigure 57 E-pH diagram for H 2O -^- Zn system; activities of chloride at 10 °3 and 10 1 ; activity of other solute species at 10-621.510.50O-0.5-1-1.5-2Zn(OH)2 - chloride - water system0^2^4^6^8^10^12^14pHFigure 58 X-pH diagram for H2O - Cl- - Zn(OH) 2 system; activity of all solute species at 10 -2 , 10-4 , 10-6E-pH and X-pH DIAGRAMS^ 1326.2.3 AlStandard chemical potentials of all species considered are listed in Table 16 for Alsystems. The hydrated Al oxide (Al203 .3H20) is considered to be the oxide formed onaluminum in aqueous solutions. Complexes have been reported in all types of halidesolutions. 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 onlyavailable for A1C13 and AlBr3 " 391 .Figures 59 and 60 show the constructed E-pH diagrams for H 2O- Al and H2O - C1 -- Al systems respectively. The formation of chloride complexes destroys the stability ofthe passivity of aluminum in the low pH region. The aggressiveness of halides can beseen based on the X-pH diagrams for H2O - F - Al203 .3H20 and H2O - CF - Al203 .3H20systems (Figures 61 and 62). The formation of a very stable fluoride complex almosttotally destroys the passivity of aluminum. The passive zone in the chloride solution isnarrowed with increasing chloride activity.E-pH and X-pH DIAGRAMS^ 133Table 16. Standard chemical potentials of substances for Al systems at 25 °CSubstance^State^1.1,°, KJ/mol.Al^s^0Al203 .3H20 s -2320.450A13-'^aq^-481.160A102 aq -839.770A1C13^aq^-878.640E1393Cl-^aq -131.168A1F2-1.^aq^-792.490E14°3A1F2+^aq^-1097.550E14°3A1F3^aq^-1396.163E1403A1F4 aq^-1689.755E14°3A1F52-^aq^-1974.220[1403A1F63-^aq^-2252.984E14°3A1Br3^aq^-794.960E139]F aq -276.480Br-^aq^-103.9700^2^4^6 8 1 0 12 14Al - Water systempHFigure 59 E-pH diagram for H2O - Al system; activities of all solute species at 10 -2, 10-4 and 10-6-0.5-1.5pi] = 10 0-------_AICI 3 (10 -6)10.50AlI^I^I^i^1^I^I^I^1^I^I^I^1^I^I 0 2 4 6 8 10 12 14pHAl - Chloride - water system-2 —-2.5 —-3Figure 60 E-pH diagram for H2O - cr - Al system; activities of chloride at 10 ° and 10'; activity of other solute species at 10-6Al 203 • 3H 2 0 - flouride - water system140 2 4 10 126^8pH21.510.5-0.5-1-1.5Figure 61 X-pH diagram for H2O -F - Al203 .3H20 system; activity of all solute species at 10 -2, 10-4, 10-621.50.50 0-0.5-1-1.5-2Al oxide - Chloride - water system0^2^4^6^8^10^12^14pHFigure 62 X-pH diagram for H 2O -Cr - Al203 .3H20 system; activity of all solute species at 10-2 , 10-4 , 10-6E-pH and X-pH DIAGRAMS^ 1386.2.4 NiFor the construction of nickel diagrams, the following standard chemical potentialshave been used (Table 17). Several stable oxides (Ni(OH) 2 , Ni203 and Ni304) are formedon nickel in aqueous solutions. Halide complexes have been found in the form of NiF+ ,NiC14 , NiC12 and NiBe.Table 17. Standard chemical potentials used for nickel systems at 25 °CSubstance^State^jr, KJ/mol.Ni^s^0Ni(OH)2^s -453.127Ni304 s^-711.910Ni203^s -469.738Ni02^s^-215.140Ni2÷^aq -48.242HNi02 aq^-349.197NiCr^aq -181.293E1411NiC12^aq^-314.992E1411Cl-^aq -131.168NiF4-^aq^-330.996E14°1NiBe aq -151.530E14°EF+^aq^-276.480E-pH and X-pH DIAGRAMS^ 139Two E-pH diagrams have been constructed for H 2O - Ni and H2O - Cl - - Nisystems respectively (Figures 63 and 64). According to the E-pH diagram for the H 2O -Cl- - Ni system (Figure 64), high activity of chlorides will narrow the passive zone ofnickel relative to the E-pH diagram for the H 2O - Ni system (Figure 63). In the X-pHdiagrams for H2O - F - Ni(OH) 2 and H2O - Cl - - Ni(OH)2 (Figures 65 and 66), it is shownthat the passive zones are diminished with increasing halide activity.Ni - H 20 system0.8211.1 0.6Cn0.4>al 0.2pHFigure 63 E-pH diagram for H2O - Ni system; activities of all solute species at 10-2 , 10-4 and 10-61.81.61.41.210.8ITUI 0.6CO> 0.4Lir^0.20-0.2-0.4-0.6-0.8-1Ni - H 2 0 - 0I - system0^2^4^6^8^1 0^12^14pHFigure 64 E-pH diagram for H 2O - Cr - Ni system; activities of chloride at 10° and 10 1 ; activity of other solute species at 10-66 8pH10 12 14Ni(OH) 2 - chloride -water system21.510.5-1-1.5-20 2^4Figure 65 X-pH diagram for H 2O -F - NiO system; activities of all solute species at 10 -2, 10-4 , 10-60 2 4 6 8 10 12 14Ni hydrated oxide - chloride - water system'—'U_0—JpHFigure 66 X-pH diagram for H2O -cr - NiO system; activities of all solute species at 10 -2 , 10-4 , 10-6DISCUSSION^ 1447 DISCUSSIONThis section will discuss several important aspects, including issues related to theexperimental results, the thermodynamic E-pH and X-pH diagrams, and other availableinformation 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 toattempt to explain the effects of pH, buffers, halides, crystallographic orientations, electrodepotential and other factors on pitting corrosion.7.1 Pit Initiation Stages and Governing FactorsPit initiation may be considered to consist of several stages, as shown in Figure 67. Atthe first stage, the metal is in a passive state with a full coverage of passive film on itssurface. The film is locally broken down when it proceeds to stage 2. Following the initialbreakdown of the passive film, changes in the local solution chemistry occur and certainconditions 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 pitpropagation.There is still considerable debate over which stage critically controls the pit initiationprocess. Kruger [143] believes that stage 2 is the critical stage which must take place beforestages 3 and 4 can proceed. The evidence to support his argument is that pitting resistance isstrongly dependant on the properties of the passive film. Metals with passive films whichare resistant to the initial breakdown (stage 2) do not undergo pitting corrosion, so stage 2may control the pit initiation process. The importance of stage 2 is also emphasized byseveral proposed theories, such as competitive adsorption, ion penetration,Pit GrowthBreakdownPassive Film Changes in the local chemistryFigure 67 Several stages occurring during the pit initiation processMetal(1)Metal(2)Metal(3)Metal(4)DISCUSSION^ 146chemical-mechanical breakdown, and transitional complex formation. All these theoriesdeal with the film breakdown process which is occurring during the transition from stage 1to stage 2. Therefore, pits will be initiated when stage 2 is reached. However, the initialbreakdown of a passive film does not automatically lead to pit initiation, because thebreakdown site may be repassivated to reproduce the full coverage of the passive film (backto stage 1). Therefore, both the repassivation process and the breakdown process should betaken into consideration during pit initiation. Consequently, if repassivation can occur atstage 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 repassivationrate in a kinetic breakdown/repassivation process occurring at stage 3. The importance ofchanges in the local solution chemistry has been stressed in Galvele's local acidificationtheory " 191. According to Boehni [1441 , metastable pits are formed at stage 3, which either arerepassivated or proceed to the stable pit stage (stage 4) depending on the development oflocal chemical environments. Beyond stage 4 pits enter their growth stage.Some of the problems relating to pit initiation may be resolved more clearly byconsidering 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, anelectrical double layer, and a diffusion layer. Therefore, recognizing that pit initiation is alocalized process occurring in the interfacial region, pit initiation may result from a localinhomogeneity in the oxide film and/or a local chemical inhomogeneity in the diffusionlayer. Furthermore, the models of pit initiation must include a role for halide ions, becausethese 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 pitinitiation. The assumption of pre-existing defects in the film, so-called priori assumption byInterfacial Region of Metal/SolutionFigure 68 Structure of the interface between metal substrate and bulk solutionDISCUSSION^ 148Okada [145] , provides the basis for Wood's flaw model [11] and Lin's point defect model [1291 •It is also the rationale for modelling the early development of local solution chemistry atstage 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 whypitting 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 betweenthe oxide film and aggressive environments. Based on the posteriori assumption, so-calledby 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 filmbreakdown process.In summary, it seems reasonable to conclude that those factors which controlinteractions between the oxide film and the adjacent local solution are the factors whichgovern the pit initiation process. Thus, three considerations appear to be relevant to the pitinitiation process:(1) electrode kinetics, the potential - current relationship.(2) nature of the passive film.(3) local solution chemistry.DISCUSSION^ 1497.2 Electrode KineticsWhen an anodic polarizing potential is applied to an electrode to move it away fromthe 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 theTafel slope (it, is exchange current density, a is the electron transfer coefficient, and z is theelectric charge of the activated species). When the electrode process is controlled by masstransport (diffusion) in the liquid phase, then we have:2.303RT—^log 1 --.-zF(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 electrodekinetics. The passive film blocks the dissolution of the metal substrate. The anode kineticsare no longer controlled by the charge transfer in the double layer but are controlled by theprocesses occurring across the passive film. In the passive state, the passivation of the metalis maintained by an extremely small current, called the passive current, of the order of a fewRA/cm2 . There is still little information available on the kinetic processes taking placeacross the oxide film, but they involve the transport of ions through the film. In general, thequestion remains as to whether the passive current is an ionic flux due to the migration ofanions or cations through the oxide layer, and whether the passive current is distributedevenly over the passivated surface.DISCUSSION^ 150When the passive film has a local breach, it is uncertain whether the anodic process atthe breach is activation-controlled and follows the Tafel behavior. Experimental evidencehas shown that the anodic current density inside initiated pits is very high, in the order ofA/cm2 [77], [146], [147] Kaesche [146] reported that the pitting current inside Al pits is about 0.8A/cm2 . Sato et al. [77] found that the current density at the pit mouth is as high as 8 A/cm2 for18Cr-8Ni stainless steel. According to Tousek [1471 , the current density inside pits was 5.8A/cm2 one second after pit initiation, decreasing to 0.5 A/cm 2 after 300 seconds in 0.5 MNaC1 at pH 8.4. Suzuki et a/. [1°41 claim that a minimum current density in the pit (in theorder of 10mA/cm2) exists, below which pits cannot develop. The current density in the pitvaries with applied potential according to the Tafel law (Equation 7.1), and the Tafel slopeswere reported to be 0.150 V for aluminum [641 , 0.050 V [147] and 0.087 V [148] forFe-18Cr-10Ni stainless steel.7.3 Dynamic Nature of the Passive FilmThe dynamic features of the passive film are easier to understand if the passive film isconsidered to be an adsorbed layer. For example, following Uhlig, the surface may beconsidered to be passivated by the adsorption of passivating ions ( 0 2- and OH-).Consequently, there exists a dynamic equilibrium between adsorbing and desorbingpassivating ions. Therefore, since desorption of passivating ions will produce a localdepassivation, it follows that there will be a dynamic passivation/depassivation process onthe metal surface.For an oxide layer, there is an analogous dynamic nature due to the filmbreakdown/reformation processes occurring on the surface. There are several reasons for thepassive film to be broken down, or ruptured, then re-formed. First, the volume of oxide perDISCUSSION^ 151mole of metal (Vox) is different from the molar volume of the pure metal (V M). Hence, whenthe oxide is formed, the ratio of Vo/Vm is not equal to 1. Therefore, stresses will bedeveloped in the film when it is formed, and the film can be ruptured when the stressbecomes sufficiently large. The Vox/Vm can be calculated according to the followingequation:voivm = (wox pry)4n, WM pox)^(7.2)where W is the molecular weight, p is density and 'Is is the stoichiometric number of metalatoms in the oxide. For example, VNi(oH)/VNi = 3.39 , V Al2033H201 V Al = 3.22 andV Zn (OH)2IV Zn = 3.56, where 0 Ni(OH)2= 4.15g1cm3 , pm = 8.90 g/cm3, A/20331120 = 2.24g1cm3 , pm =2.7 g/cm3, Zn(OH)2= 3.05g1cm3 and pz. = 7.14 g/cm3 r 1491 . Secondly, if the oxide film iscrystalline with a different crystal structure from that of the metal substrate, the film issubject 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 bythe electrostriction pressure due to the high electric field across the film. The dynamicfeature 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 ofelectrochemical 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 chloridesolutions, which could be associated with the "birth/death" process of metastable pits.7.4 Local Solution Chemistry during the Film Breakdown/ReformationProcessThe film breakdown/reformation process, as mentioned in Section 7.3, causes localDISCUSSION^ 152fluctuations in the anodic current density and solution chemistry. In this section, theresulting local changes in pH and halide concentration are modelled for two simplesituations: (1) changes due to one oxide film breakdown/reformation cycle and (2) changesdue to a continuous breakdown of the oxide film. The changes in local solution chemistryare 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/reformationcycle.Assuming that a hydrated oxide film, with a thickness of h, breaks down and producesa circular active site with a radius of r (Figure 69), the metal substrate will undergo ananodic dissolution:M = M" + ne-^(7.3a)then the active site will be re-covered by the hydrated oxide film during the film reformationaccording to the following reaction 3 :M" + nH2O = M(OH)" + al+^(7.3b)and the equilibrium constant for Equation (7.3b) is given:[H+] n —r " [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)H20 = MO ("il2) + al+Metal^Oxide Film^Diffusion Layer^ Bulk SolutionThe region containing generated hydrogen ionsFigure 69 An active site with a radius of r during oxide film breakdown.DISCUSSION^ 154^Vox = nr2h^ (7.4)and nr2hpoxlW„ moles of oxide will be generated for the coverage of the active site, wherepox 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) = nicApoxlWox^(7.5)Assuming that the film reformation process is fast (a pulse of current), then the fr ionsgenerated by Equation (7.3b) will not have moved far from the electrode surface whenrepassivation is completed. Therefore, allowing the generated H+ to be confined in acylindrical region with the radius, r + d, and height, d, as shown in Figure 69, the volume ofsolution, 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,(r + d)2dWox(7.7)There is a difficulty in estimating the value of d in Figure 69. However, the volume ofsolution containing generated H + can be roughly calculated based on the film reformationrate, i. e., the time, t, required for the reformation of the passive film at the active site. If thereaction (7.3a) proceeds at an anodic current density, i, then the charge, Q, required for thereformation of nr2hp 0/147„x moles of oxide is given by:d —i wox(4nhpoxFD )112 (7.11)DISCUSSION^ 155Q = nr2 it = nicr2hpo,FYI'Vox^ (7.8)from which t may be determined:nhpoxFt =^ilKx(7.9)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:Therefore, in neutral or near neutral solution where the initial fr concentration isnegligibly small, the local concentration of I-I + ions can be estimated by combination ofEquations (7.7) and (7.11).[H1=nr 2hp ox (7.12)(4nhpoxFD )1121 2 (4nhpoxFD )112{r +  ^Woxiwox iwox DISCUSSION^ 156As we see in Equation (7.12), for a given oxide film, the local concentration of fr ionschanges with two variables, r and i, without considering the neutralization by the bulksolution. 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 convertedto molarity (moles/liter, M) by multiply Equation (7.12) by 1x10 3 , from which pH =-log(103 [HI).Figure 70 shows the variation of local pH with the size of active site, r, and the anodiccurrent density, i, for reformation of Al203 .3H20 on the aluminum surface, where h = 1x10-7cm, pox = 2.42 g/cm 3 , D = 1x10-5 cm2/s [151], n =6, Wox = 156 g and F = 96500 C. FromFigure 70, we can see that the local pH decreases with the increase in the size of active sitesand 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 bythe migration of ions in the electrolyte, either cations move away from the electrode oranions move into the interfacial region. Considering a solution containing only sodiumhalide without other supporting electrolytes, then the migration of halide X - carries a fractionof the anodic current, which is determined by the transport number, nx- . By knowing thetransport number, n, the local current carried by the migration of X - ions is icr2 itix- and theamount of charge carried by migration of X - ion is nr2inx-t during thebreakdown/repassivation cycle, where t is expressed in Equation (7.9).Figure 70 Change in local pH with variation of r and i1 0 "8 1 o - 6 -21 01312111098ICL 7Tt 10o 6J543210r, cmDISCUSSION^ 158Then assuming that halide ions migrate from the bulk solution into the same region asthe IT- ion confining region shown in Figure 69, we have a similar Equation to (7.12) for thelocal increase in halide concentration, [X],„ e, in moles/cm3 , and which can be converted tomolarity (M) by : M = 103 moles/cm3 .=—^ -•nnx r2 n pox (7.13)(4nhpoxFD )112} 2 4nhp„FD )112{r +  ^ox- - ox iwox and[X] = [X10 +—nnx r2 hpox (7.14)(4nhpoxFD )112} 2 (4nhpoxFD )112{r +  ^Waxiwox iwoxwhere [X -],o, = [X] - [X] 0, [X-] is the local concentration of halide ions, and [K] 0 is the bulkconcentration of halide ions.Figures 71 and 72 show changes of the local chloride concentration with the size of theactive site, r, for formation of Al203 .3H20 on aluminum, nx- = 0.61 [151] , h = 1x10-7 cm, pox =2.42 g/cm3 , D = 1x10-5 cm2/s, n = 6, Wo, = 156 g and F = 96,500 C. Figure 71 was calculatedat different anodic current density for a given bulk solution concentration of 10 -2 M. InFigure 71, the local halide concentration is little affected and equal to the bulk concentrationfor an active site having radius less than 10 -6 cm. At r = 10-4 cm (1 gm), the local halideconcentration rises with increasing anodic current density. In Figure 72, the results werecalculated for different bulk halide concentrations at a given anodic current density of 1.0A/cm2 . For halide solutions with bulk concentrations of 0.001 and 00.1 M, the localconcentration increases as the active site becomes larger than 10 4 cm (1 gm). There is littlechange in the local concentration (equal to bulk concentration) for the halide solutions of 0.1M and 1.0 M.10-8 10-6 1 02-1-1.2-1.3-1.4-1.7-1.8-1.9-2-2.1r, cmFigure 71 Change in local halide concentration with variation of r and i at a given bulk halide concentration of 0.01 M[X] 0 =1.0............................[K] o =am..........[X] o =o.00i0.20-0.2-0.4-0.6-0.8-1-1.2-1.4-1.6-1.8-2-2.2-2.4-2.6-2.8-3-3.21 0 -8^1 0-6^i o-4^162r, cmFigure 72 Variation of local halide concentration with bulk halide concentration at a given anodic current density of 1 A/cm 2Cr;DISCUSSION^ 161(2) Local pH and halide concentration changes due to a continuous breakdown processat a local (active) site.Assuming that a passive film on the surface is continuously broken away at an activesite with a steady anodic current, i, the metal substrate will undergo the anodic dissolutionaccording 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 shownin Figure 73. Therefore, the M' concentration can be modelled approximately via a steadystate diffusion process. At the steady state, the M' concentration at the interface can beestimated by Fick's first law:[Mn_J = nF = D^8(7.16)where [Mil ; is the concentration of M' at the interface, [M10 is the bulk concentration ofMn+ , and 8 is the thickness of the diffusion layer. Then we have:i 8[Mn ] t — [Mn ]o— 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 0a^ xFigure 73 Local solution chemistry profile due to a diffusion process across the diffusion layerDISCUSSION^ 163i8 ^(7.17b)= nFDAssuming 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). [HT i K„ nFD(7.18a)andi8K =( ")nFD(7.18b)Units used in Equation (7.18b) are: i - A/cm2 , 8 - cm and D - cm2/s. Thus the [Hl i is in unitsof moles/cm3 . The [H+] ; is then converted to molarity (M) by multiplying by 10 3 , fromwhich 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 increaseswith increasing i and 8. Any increase in anodic current density and diffusion layer thicknesswill increase^and lower the local pH at the interface. Figure 74 shows the variation oflocal pH with the anodic current density at a given diffusion layer thickness of 1x10 -2 cm forNi, Zn, Al and Sn oxides, where D = 10-5 cm2/s, logK„ (Ni(OH)2) = -12.18, logK„(Zn(OH)2) = -10.96, logK„ (Al203 .3H20) = -5.7 and logkx (Sn(OH) 2) = -1.5. The value of 8was 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 anodiccurrent flow at the active site. The flux of halide ions migrating to the local site on theDISCUSSION^ 164electrode surface will be in x%F. When a steady state is established, the flux of halide ionmigration into the interfacial region is equal to the flux of halide ions diffusing back into thebulk solution through a diffusion layer of thickness 8. The concentration of halide ions atthe interface at steady state then becomes:n;i =D[Xl i —[X10F^8n;i8[X-] ; = [X-13+ FD(7.19)(7.20)where [X] ; is the halide concentration at the interface, and [X-]0 is the bulk halideconcentration. 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 verysignificant when the bulk concentration [X] 0 < 10-2 M. The change becomes lesspronounced when the bulk solution becomes more concentrated with chloride._ Ni----------------------------------------------AlSn1 1 1 1 1 1 1 11I^i I^i 1^i654I0_ 3cis0O 2_I10-1-210 -2 10 -1 100 101i, A / c m 2Figure 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-6 -4 -2 0 20-1-2-3-4-5-6LOGTX 10, Mi = 1.0 A/cm 21= 0.1 A/cm 2Figure 75 Variation of local halide concentration with bulk halide concentration in a continuous film breakdown process at 8 =10-2 cm[X l 0 = [X 1 ii = 0.01 A/cmDISCUSSION^ 1677.5 Halide Complex Formation and Pit Initiation TheoryThe 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 (seeSection 7.4). According to the E-pH and X-pH diagrams constructed for tin, zinc, aluminumand nickel (see Section 6.2), when the pH moves to lower values (to the left side ofdiagrams), the passive oxide is no longer thermodynamically stable. Figure 76 presents aschematic X-pH diagram for M(OH) n in a halide solution. The bulk solution pH and halideconcentration are presented as point A on the diagram, indicating that the metal surface ispassive. With decreasing pH the metal is subjected to a transition from passivity todissolution via halide complex formation. An increase in halide concentration, [Xi, narrowsthe passive zone, and also causes the instability of passive films. Therefore, a combinationof a lower pH with a higher concentration of halides (point B in Figure 76) leads to a trendtowards increasingly stable complex formation and a destablization of the passive films.Based on the above considerations, a pit initiation theory is proposed, whichemphasizes changes in local solution chemistry and the generation of critical conditions forthe formation of stable halide complexes. In this theory, pit initiation consists of a series ofstages as shown in Figure 77.Stage 1 is a dynamic process where the repetitive film breakdown/reformation processoccurs rapidly and randomly over the surface, irrespective of the presence or absence ofhalides. In halide-containing solutions, the film breakdown/reformation process leads tochanges in local solution chemistry (Stage 2) characterized by a lower pH and a higherconcentration of halide ions (as discussed in Section 7.4). The local solution chemistrychange can disappear very shortly after repassivation and only cause a short-livedpHX-100_1Figure 76 A generalized X-pH diagramDISCUSSION^ 169perturbation at the oxide/solution interface. If the fluctuation in local solution chemistry issignificant enough, with an adequate lifetime, a further breakdown/reformation process mayoccur at the same site due to an autocatalytic process in the local region, and a continuousbreakdown process can be established (Stage 3). At stage 3 metastable pits are formed, theymay 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 criticalconditions are met for stable complex formation. At this stage (stage 4), any oxide film willbe unstable relative to halide complexes, and it is no longer possible for the surface to berepassivated. Therefore, stable pits are initiated. Consequently, stage 4 is the critical stagein pit initiation, where critical conditions are established for stable halide complexformation.From a thermodynamic point of view, pit initiation is determined by the formation ofstable 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 equilibriumbetween the hydrated oxides and the halide complexes, according to the reaction in Equation(7.21) 4 :M(OH)n + mX- + al+ = MX. ("1-")- + nH2O^(7.21)and conditions for formation of MX,n(m -n)- :[MX,n(m-n)-1 = KxGHTPCl m)^(7.22a)4 For an anhydrous oxide, Equation (7.21) may be written:M00/2) + mX - + nil+ = MX,, (')- + n/2H20MXm (m-n)-Oxide filmMetalx(1)X^X4k,HI-^High [Xi^H ' r-N, H+ Low pH(2)^ (4)(3)Local Change in pH and [X" ]^ Critical stage for stable pit initiationContinuous breakdown processBreakdown/reformation processFigure 77 A schematic diagram showing processes in pit initiationDISCUSSION^ 171where the equilibrium constant is given by IC = [MX,n(m -11/(fHTPCP). Assuming that acritical concentration of halide complexes [MX. (m-n) ]cnt is required for film breakdown andformation of a stable pit, then we have a critical product of (VIT[X] m)crit, i.e.uva„, (m-n) crit = KAHTPcncrit^ (7.22b)When the product of local pH and local halide concentration is greater than the value ofaltim iXT)crit, MX„,1 (m-n)- formation is thermodynamically stable, which will lead to pitinitiation.From a kinetics point of view, two reactions compete with each other at the criticalstage. These are formation of the oxide film and formation of complexes:.M + nH2O --> 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(„,_„ ) _0c Lrx -f^ (7.26)where 13 and 7 are the reaction orders of [111 and [K] respectively. Thus, we can see thathigh [ff-] (low pH) slows down the repassivation process and that high halide concentrationaccelerates the formation of complexes. Pit initiation occurs when:DISCUSSION^ 172R^RAmion^ (7.27)7.6 Evidence from Experiments on Environmental EffectsThe proposed pit initiation theory emphasizes the importance of changes in the localsolution chemistry and the formation of stable halide complexes. The critical stage in pitinitiation is determined by the critical value of local pH and halide concentration (Equation7.22b). Therefore, any factors that influence the local solution chemistry will also affect thepit initiation process. The following environmental effects in this study can be clearlyexplained by the proposed theory:(1) Complex-forming ions are needed for the establishment of the critical stage for pitinitiation, otherwise pitting corrosion does not occur. The so-called aggressive ions are thecomplex-forming ions. The polarization results have indicated that there was no pitting ofnickel in 1.0 NaNO3 and 1.0 M Na2SO4 at pH = 10.5, not even at pH = 2.5 in Na2SO4solution (Figure 33 and 34). The passive film formed on nickel is very stable in thesesolutions. 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, NiC12 , A1C13 , A1Br3 , ZnC142-, ZnBr42-, SnC142- and Sn62-), so thatpitting 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 ofcomplexes, the more readily the pit is initiated. In order to compare their stabilities, theminimum halide concentration, [X-]„„„, above which metal complexes become the dominantsoluble species is calculated from the stability constant Km for Equation (7.28a):DISCUSSION^ 173M°+ + mx - =^ (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^[X ]min= Km hIm^(7.29)The formation of the halide complex, MX,n(m-")-, is favored when the halide ion concentrationis 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 localsolution chemistry, it is easier to meet a smaller minimum halide concentration requirement.Consequently, the halide requiring smaller minimum concentration is more aggressivetowards pit initiation. Therefore, according to Table 18, bromide ion is less aggressive thanchloride 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 pittingpotentials show a trend as follows:Ecp in NaCl < E cp in NaBrTherefore, if aggressiveness is equated with a lower pitting potential, then the trend isconsistent with the predictions based on the minimum halide concentration for complexformation (Table 18).DISCUSSION^ 174Table 18 Minimum halide concentration for complex formationMetal Complex Min [CF], M Min [Bri, MNi Nir 0.47 1.32NiX2 0.41 -Al A1X3 0.585 0.776ZnZnr 0.32 0.60ZnX2 0.50 1.12ZnX3 - 0.68 1.76ZnX42- 0.89 1.78(3) From the X-pH diagram for the H2O - F - Al 2033H20 system (Figure 61), it isknown that Al-fluoride complexes are very stable, so a very low critical pitting potentialwould be expected in fluoride solution. Actually, the fluoride complexes are so stable thattotal passivity of the aluminum surface is destroyed, which results in an active dissolutionbehavior in 0.1 M fluoride solution at pH 3.4 (Figure 27). For a bulk solution with [F] = 0.1M and pH = 3.4, it is clearly evident that the bulk solution conditions shown in Figure 61 arelocated in the complex formation zone (A1F52-), far away from the passive zone. In this case,the bulk solution chemistry meets the conditions for stable complex (A1F52-) formation sothat general corrosion occurs on aluminum, instead of pitting corrosion. A similar situationis seen for nickel in 1.0 M chloride and fluoride solutions at pH = 2.5 and 3.1 respectivelyDISCUSSION^ 175(Figures 35 and 52). Bulk solution chemistry conditions also show that Ni is located in thestable 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 bulksolution chemistry meets the conditions for stable halide complex formation. According tothe E-pH and X-pH diagrams, the total depassivation requires that the bulk solutionchemistry exhibits a low pH and a high halide concentration. However, in most cases suchas 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 solutionchemistry favors passivation and not the formation of halide complexes. Therefore, onlylocal regions on the surface meet the requirement for complex formation, due to the changesin local solution chemistry discussed in Section (7.4). Consequently, only localizedbreakdown 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 bulksolution. For example, bulk solution chemistry indicates that nickel is in the passive zone onthe X-pH diagram for 1.0 M NaC1 at pH = 10.5 (Figure 66), but the local pH may be loweredtowards the stable complex zone. Hence, pit initiation depends on the local pH and anyeffect 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 thepH region from 4.5 to 10.5. The critical pitting potential is independent of bulk solution pHuntil 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 isdetermined by Equation (7.18b), provided that there is a continuous breakdown of theDISCUSSION^ 176passive film. Taking into consideration the transport of OH - through the diffusion layerfrom the unbuffered bulk solution, the local solution pH for Ni is calculated (see AppendixI) 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 toreact with and neutralize the generated H+ . ^the generated 11+ ions almost totallycontrol the local solution pH. However, in strong alkaline solutions the large amount ofavailable 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 alkalinesolution (pH >12). The local pH change in Figure 78 correlates well with the variation ofthe critical pitting potential for Ni with bulk solution pH (Figure 39).Recalling the observed corrosion behavior of nickel in 1.0 M NaCl with increasingbulk solution pH (Section 5.2.2), the transition from uniform corrosion to pitting corrosionand then to complete passivity can now be well understood. At pH = 2.5 the bulk solutionchemistry meets the conditions for complex formation, resulting in general corrosion ofnickel. In the pH region from 4.5 to 12.5, chloride complexes are only formed locally due tothe change in local solution chemistry, and the local pH is independent of the bulk solutionpH in the pH range from 4.5 to 10.5. Consequently, localized breakdown occurs, with thecritical pitting potential independent of the bulk solution pH. At higher pH values, the localpH is controlled by the bulk solution pH, and no local acidification occurs at pH = 14, sothat nickel remains passive.I^I^II^I^I^I^I^I^I^I^I^114131211102 4 6 8 10 12 14Bulk Solution pHFigure 78 Change of the local solution pH with the bulk solution pH calculated for Ni using an one-dimensional diffusion modelDISCUSSION^ 178(5) According to the proposed pitting theory, the effect of a buffer on pitting corrosionbehavior may be predicted. When a buffer is introduced into the solution, the generated frions will be consumed by a buffering reaction, which counteracts local changes in pH, asshown below for carbonate and phosphate buffers:CO32- + fr = HCO3 -^(7.30)and^PO43- + fr = HPO42-^(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 321HCO3 and PO431HPO42- buffersdo not act simply by competing with halide ions for adsorption sites on the surface, assuggested by others [661, [67] Their primary role is the retardation of local pH change. Theincrease in critical pitting potential was also found on zinc in CO321HCO3 - buffered solution(Figures 26). However, in this case, the increase should also have resulted from thebroadening of the passive zone to include lower pH regions because of the formation ofinsoluble ZnCO3 (Figure 56).(6) Aluminum should have undergone general corrosion in 0.1 M NaF at pH 6.0according to the prediction arising from the X-pH diagram (Figure 61). Aluminum didcorrode in the active region. However, the anodic current density dropped dramatically to avery low value above a potential of -1.1 V (SCE). This transition is caused by the formationof insoluble A1F3 .3H20, which precipitates on the surface of aluminum, as the Al 3+concentration at the surface increases with increasing anodic current density. A thick saltfilm was noted on aluminum in 0.1 M NaF at pH 6.0 (Figure 31). Chemical analysis showedthat it was an aluminum fluoride film (see Section 5.1.4).DISCUSSION^ 179At first, it is surprising that nickel does not suffer pitting corrosion in 1.0 M NaFsolution at pH = 6.0 and 10.5 (Figures 50 and 51), because pitting corrosion is predicted bythe 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 bufferreaction 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 theresistance to the change in local pH and fluoride concentration prevents pitting corrosionfrom occurring.7.7 An Explanation of the Pitting Dependence on CrystallographicOrientationsOrientation-dependent pitting behavior was observed on single crystals of zinc and tinin 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 inorientation-dependent pitting corrosion. The surface atomic density changes withorientation. The higher atomic densities in surfaces composed of the more closely packedplanes will produce higher numbers of oxygen cations per unit area of surface monolayerthat 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 becausethe 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 thesurface atomic density, as predicted.Ecp: (0001) < ( 1120) < (1010)If the predicted relationship between surface atomic density and the critical pittingpotential is applicable to other systems, then the atomic density and the critical pittingpotential for F.C.C. metals should be:Density:^{110} < {100} < {111}and predicted Ecp :^{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 pittingbehavior of Ni suggests that there is no generally applicable relationship, because E el, hasbeen 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 isconsidered to be the critical process in pitting initiation. Therefore, the possible explanationof orientation-dependent pitting may be related to the effects of crystallographic orientationon changes in the local solution chemistry. These changes depend on the local anodiccurrent density (Section 7.4). Higher anodic currents promote a lower pH and higherconcentration of halides at the local regions (see Figures 70, 71, 74 and 75), so that theDISCUSSION^ 181critical condition for pit initiation is more readily established. Therefore, we expect a lowercritical pitting potential on a surface which has a higher local anodic current density (ahigher active dissolution rate).Data reported in the literature show that the anodic dissolution rates Ra under activeconditions in acid solutions for aluminum {130} and nickel [132} are dependent on surfaceorientation:aluminum. R,{001} < R,{011} < Ra{111}and nickel: Ra{ 1 11} R1)^—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 andNi (see Section 2.9).The active anisotropic dissolution rates of zinc and tin are unknown, but the relativerates of dissolution of differently oriented surfaces may be ranked on the basis of theobserved pit morphologies. For example, pit walls of crystallographic pits are composed ofthe most slowly dissolving surfaces. Hence, the formation of crystallographic pits isevidence that dissolution behavior is anisotropic. It was observed that pit walls on zinc wereclose to {10l0} hexagonal prismatic planes, suggesting that dissolution rates are lowest onthese surfaces, consistent with the highest critical pitting potential for { 1010} orientedcrystal surfaces (Figure 12). Also, walls of pits formed on Sn crystals were found to beclose to {100} tetragonal prismatic planes and {011} tetragonal bipyramidal planes,suggesting that these are slowly dissolving surfaces, consistent with critical pitting potentialsDISCUSSION^ 182observed on (100)- and (011)-oriented crystal surfaces (Figure 10). Therefore, the differentanodic dissolution rates on the differently oriented crystal surfaces cause theorientation-dependent pitting behavior of single crystals.7.8 Critical Pitting PotentialThe critical pitting potential is the most important criterion for defining susceptibilityto pitting corrosion. The explanation of the meaning of the critical pitting potential is a keyissue in the approach to understand the pit initiation mechanism. Any successful theoryshould 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 associatedwith the critical stage in pit initiation. The critical pitting potential is defined as the potentialat which the local solution chemistry leads to the critical condition for the formation ofstable halide complexes.The relationship between local solution chemistry and electrode potential can bedetermined 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)^[H1 1^exp= k 1 (  OK°x^azFEnFD nRT^where k 1 = is exp n^RT(7.33)and from Equation 7.20DISCUSSION^ 183k2n;s5^czzFE[X] i —[X1 0 = ED^ exp RT(7.34)-azFE0where, k2 = io exp RT . When the condition^(from Equation 7.22b) is met atthe interface, we have the critical pitting potential, E cp . Therefore,8K ) exp oczF EcpH1 in [Xi i) crit =m^1c2n -8^azF E,^RT^ ([X10 ±^eXp RT P )in([ ^(7.35)The relationship between E cp and bulk halide concentration [X]0 can be derived fromazT EEquation (7.35), if the term kfcs exp RTcp is significantly small relative to [X] 0 . Thissituation is most likely when the halide concentration in the bulk solution is high and/orthere is a high concentration of supporting electrolytes ( nx- is very small). Therefore,neglecting this term, we have:14(--nsKFD')exp azFRTEcP ([Xlo)m =^n [Hi crit = constant (7.36)and rearranging Equation (7.36)Ecp = A 2.303mRT log[X loocz,F(7.37)^DISCUSSION^ 184RT^FD(v17 [111 1:where A is a constant given by A = —azF ln Int , and n and m are the stoichiometricsic;numbers for 11+ and X- in Equation (7.21), respectively.Equation. (7.37) gives a semi - logarithmic law between E cp and the bulk halideconcentration [K] o, which has been well established for many metal - halide systems (seeSection 2.2.2) 5 . However, the relation will deviate from the semi - logarithmic law if thek2n;s^az,F'ERTcpexp^ term cannot be neglected. Then there is no simple relationship between theFD critical pitting potential and the bulk halide concentration.7.9 Extension of the Proposed Theory to Other Aspects of PittingThe proposed pit initiation theory has been shown to explain some of theenvironmental and crystallographic aspects of pitting. There are still other effects on pittingcorrosion, 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 tothese aspects.5 Equation (7.37) is derived by assuming that the relationship between the potential and anodic30 3mR Tcurrent density follows Tafel's law. Thus the slope of 2. ,,,zF is obtained. In situations wherethe anodic current density is high (in the order of A/cm 2), the relationship between the potentialand the anodic current density does not fully obey Tafel's law, and the electrode kinetics arepartially controlled by the mass transport process across the diffusion layer. Consequently, the2.303mRTapparent slope may be larger than the predicted slope of azF in Equation (7.37).DISCUSSION^ 185Some 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). Theincrease in pitting resistance by alloying can be explained in three ways according to thetheory:(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 whichextends the passive zone into a lower pH region than iron oxide if E-pH diagrams for theFe-H20 and Cr-H 20 systems are compared. Hence, chromium increases the pittingcorrosion resistance. Stainless steels alloyed with Mo and N have improved pittingresistance. The role of Mo is still disputed, but some results have shown that the Moaddition decreases the anodic dissolution at active sites [156]  A smaller anodic currentdensity tends to slow down the change in local solution chemistry and to prevent the localregion from establishing the critical conditions for pit initiation. It has been speculated thatthe beneficial effect of nitrogen is due to the formation of NH4+ by reduction of N within theactive site, which consumes II+ ions and buffers the local pH change [157] .The detrimental effect of inclusions has been reported in several studies (see Section2.2.9). The existence of inclusions creates inhomogeneities in the passive film, which act asactive sites in halide solutions. The size of the larger inclusions is in the order of gm. Whenan active site is formed at these inclusions, the changes in local pH and halide concentrationwill be significant (Figure 70 and 71), and the inclusions become the preferential sites for pitDISCUSSION^ 186initiation.Temperature is another factor influencing pitting corrosion. The critical pittingpotential decreases with the increase in temperature (see Section 2.2.4). It is not too difficultto understand why temperature promotes pit initiation kinetically. According to theArrhenius rate law, activation-controlling dissolution will be accelerated with an increase intemperature. So a higher local anodic current density will give rise to a lower pittingresistance. Thermodynamically, if the passive zone is narrowed at higher temperature dueto the enlargement of the complex formation zone in the X-pH diagram, then it is expectedthat pit initiation will occur more easily.Fluid flow affects the diffusion layer thickness 6, which influences mass transportcontrolled changes in the local solution chemistry (see Section 7.4). The changes in localsolution chemistry will be retarded with a decrease in the diffusion layer thickness.Increasing the fluid flow velocity decreases the diffusion layer thickness and should producean increase in the critical pitting potential. This is consistent with the observed behaviorreported by others [73]4761.CONCLUSIONS^ 1878 CONCLUSIONS(1) Pitting corrosion of zinc and tin single crystals is crystallographic-orientationdependent. The critical pitting potential of zinc, E ci, varies with the crystallographic orientationin the order of Ecp(1010) > Ecp(1120) > kp(0001), and the lowest pitting potential for Sn crystalshas 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 dissolutionrates, and single crystal surfaces of Zn and Sn with these orientations exhibit higher criticalpitting potentials. There is a correlation between the local anodic dissolution rates and thecritical pitting potentials. Therefore, the dependence of pitting corrosion on the crystallographicorientations is attributed to orientation-dependent changes in local solution chemistry caused byorientation-dependent local anodic dissolution rates. Lower local anodic current densities at theactive site produce higher critical pitting potentials.(4) E-pH and X-pH diagrams have been constructed for Sn, Zn, Al and Ni in associationwith the formation of halide complexes. These diagrams contribute significantly to theunderstanding of pit mechanisms. The role of halide ions is emphasized in the pit initiationprocess. The formation of halide complexes is a necessary step to destablize the passive filmand induce pit initiation.(5) The critical pitting potentials for nickel, aluminum and zinc vary with the specifichalide species. They increase in the tested halide solutions in the order: Ecp in NaCl < Eci, inCONCLUSIONS^ 188NaBr. The minimum halide concentrations for complex formation, ['g rain , and X-pH diagramsindicate that the formation of more stable chloride complexes of nickel, aluminum and zincleads 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 thebulk solution. The critical pitting potential of nickel has been shown to be independent of thebulk solution pH in the range of 4.5 - 10.5. The local solution pH is controlled by the bulksolution pH in very strong alkaline solutions (pH > 12.5), where a large amount of 011" isavailable to consume the II+ generated by the film breakdown/reformation process. Theaddition of carbonate and phosphate buffers to the halide solutions prevents changes in the localsolution chemistry, thereby increasing the critical pitting potential.(7) The metal undergoes general corrosion if the bulk solution chemistry meets theconditions for the formation of stable halide complexes, such as the situation for nickel in 1.0 MNaCl at pH 2.5. If the bulk solution chemistry does not meet these conditions, the criticalconditions 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 solutionchemistry governs the pit initiation process. The local solution chemistry is modelled for twosimple situations, and it is shown that the local region is associated with low pH and highconcentration of halide. Pit initiation occurs at a critical stage when the conditions are met forthe formation of stable halide complexes in the local region. 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Schwenk, Corrosion, Vol. 20, p. 129t (1969).[149] Handbook of Chemistry and Physics, Etd. by R. C. Weast, 70th Edition, CRC Press,Boca Raton, Florida, 1990.REFERENCES^ 196[150] B. Baroux, Proc. Int. Symp. Corros. Sci. and Eng., Etd. by R. A. Rapp, N. A. Gokcer,A Pourbaix, CEBELCOR, Vol. 1, p. 315, 1989.[151] Electrochemical Data, B. E. Conway, Elsevier Publishing Co., London, 1952.[152] G. Okamoto, K. Tachibana, S. Nishiyama, T. Sugita, Passivity and Its Breakdown onIron and Iron Base Alloys, Eds. R. W. Staehle, H. Okada, p. 106, NACE, 1976.[153] Y. Miyata, T. Handa, H. Takazawa, Corros. Sci., Vol. 31, p. 465 (1990).[154] K. J. Vetter, Electrochemical Kinetics - theoretical and experimental aspects,Academic press, N. Y., 1967.[155] J. O'M. Bockris, A. K. N. Reddy, Modern Electrochemistry, Plenum Press, N. Y.,1970.[156] M. B. Ives, U. 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^ 197APPENDIX IEffect of Bulk pH on Local Solution pH in Unbuffered SolutionThe change in local solution pH with bulk solution pH can be estimated by a simpleone-dimensional diffusion model discussed in Section 7.4, without considering the ionelectrical 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 [HH+ D  nFDH+^H+^8(A-1)and the OH- flux from bulk solution into the local region is given:figH10 — [OH JOH = DOH^8where, [OH] 0 and [OH] ; are bulk and local concentrations of Off respectively. Hence, the netflux of fr is the sum of Equations (A-1) and (A-2):•1 net = H+— "IOW^ (A-3)therefore:1^ (A-4)( i61Cox yi — [H+10[Hl i — [H10^nFD ^[01110— [011 1 D ,^ = D^DH 8 8^OW^8Assuming that differences in the diffusion coefficients of H + and OH- are sufficiently small tobe neglected [151], then DOH = D H+ = D, and[H+1^ioKoxnFD^([0H10— [0111 1 )(A-5)and equilibrium is established between II + and OH, and [111[01-1] = 10-14, then we have:(A-2)10-14^iSICT 10-14^ (A-6)[HI 1 nFD ) [H -10APPENDIX^ 198Therefore, 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 currentdensity of 1.0 and 10 A/cm2 respectively, where n = 2, log(K0 ) = -12.18 ( for Ni(OH) 2), D =1x10 -5 cm2/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 pHi= 1.0 A/cm20.0001^0.0001584999^3.799971^40.00001 0.000068499 4.164315 50.000001^0.00005949^4.225556^60.0000001 0.0000585 4.232844 70.00000001^0.00005751^4.240256^80.000000001 0.000048501 4.314249 91.00000000E-10^0.0000000002^9.618048^101.00000000E-11^1.06213489E-11^10.97382 111.00000000E-12^1.00588442E-12^11.99745^121.00000000E-13^1.00058534E-13^12.99974 131.00000000E-14^1.00005850E-14^13.95^14= 10 A/cm20.0001^0.0006849999^3.164309^40.00001 0.000594999 3.225483 50.000001^0.00058599^3.232109^60.0000001 0.000585 3.232844 70.00000001^0.00058401^3.233579^80.000000001 0.000575001 3.240331 91.00000000E-10^0.0004850001^3.314258^101.00000000E-11^2.40963855E-11^10.61804 111.00000000E-12^1.06213489E-12^11.97382^121.00000000E-13^1.00588442E-13^12.99745 131.00000000E-14^1.00058534E-14^13.99974^14APPENDIX^ 199APPENDIX IIE-pH Diagrams for Sn SystemsE-pH diagrams in respect to the formation of Sn anhydrous oxides are shown in FiguresA-1 and A-2 for Sn - H 2O and Sn - Cl - - H2O systems, respectively.1.4 Tin - water system1.210.80.60.410 -6— 0.22 64 80 141210cn02-SnO3w-0.2-0.4-0.6-0.8–HSnO2-1-1.2-1.4APPENDIX^ 200pHFigure A-1 E-pH diagram for H 2O - Sn system; activities of all solute species at 10 -6APPENDIX^ 201Sn - chloride - water systemww1.41.20.80.60.40.20-0.2-0.4-0.6-0.8-1-1.2-1.40^2^4^6^8^10^12^14pHFigure A-2 E-pH diagram for H2O - Cr - Sn system; activities of chloride at 10 ° and 10';activities of other solute species at 10 -6APPENDIX^ 202APPENDIX IIIPit Morphologies of Cold Rolled and Annealed Polycrystalline NickelPits formed on the surfaces of all cold rolled nickel specimens were found to benon-crystallographic. Figure A-3 shows a hemispherical pit formed on the surface of coldrolled Ni tested in 1.0 M NaC1 at pH 10.5.Some tests were conducted on the annealed nickel specimens, as a comparison with thecold rolled nickel. Nickel was annealed in a furnace at 800 °C for 4 hours. The polarizationcurves obtained from the annealed nickel in 1.0 M NaCl at pH 10.5 are similar to that obtainedfrom the cold rolled Ni in the same solution (see Figure 36 in Section 5.2), and no difference intheir critical pitting potentials (E cp) was detected. However, the pits formed on the surface ofthe annealed Ni were found to be crystallographic, as shown in Figure A-4.APPENDIX^ 203Figure 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^ 204APPENDIX IVSome Suggestions of Further Work(1) The local solution chemistry has been modelled in two simple situations (see Section7.4) to show the trend of changes in terms of the local solution pH and the local concentrationof halides. However, more work is needed to model the processes occurring in the local regionin details. A more precise model should takes into consideration: (a) activity coefficients ofspecies, (b) equilibria between the species involved in the local region and (c) migrations ofcharged species.(2) Further work is suggested to measure the anodic dissolution rates on the differentlyoriented crystal surfaces of Sn and Zn under active dissolution conditions in acidic halidesolutions, in order to determine the dependence of anodic dissolution rates on orientations andto confirm the orientation-dependent pitting behavior of single crystals.

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