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The high temperature electrochemical behavior of carbon steel in alkaline sulfide solutions Crowe, David Charles 1985

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THE HIGH TEMPERATURE ELECTROCHEMICAL BEHAVIOR OF CARBON STEEL IN ALKALINE SULFIDE SOLUTIONS by DAVID CHARLES CROWE M.A.Sc. (Metallurgical Engineering) The University of British Columbia, 1982 B.Sc. (Mechanical Engineering) The University of Manitoba, 1977 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY xn THE FACULTY OF GRADUATE STUDIES Department of Metallurgical Engineering We accept this thesis as conforming to the required standard THE UNIVERSITY OF BRITISH COLUMBIA February 1985 © David Charles Crowe, 1985 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make i t freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the head of my department or by his or her representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission. Department of ? ? t o l c ^ ^ ^ 2 ^ l ^ a^l^/x^yU/x^--The University of British Columbia 1956 Main Mall Vancouver, Canada V6T 1Y3 Date ABSTRACT The high temperature, high pressure electrochemical behavior of A516 Gr. 70 carbon steel in aqueous alkaline sulfide solutions was studied by means of polarization tests and c y c l i c volt-ammetry. The effects of variation of temperature (90-150°C), s u l f i d e concentration (0-3 m), scan rate (1-50 mV/s) and scan range, and the effects of s t i r r i n g and polarization at the switching potentials between scans were investigated. Passiv-ation was consistent with formation of a protective Fe2C>3 film. An understanding of the electrochemical behavior of iron in the alkaline sulfide solutions was facilitated by the construc-tion of E-pH diagrams for S-H20 and Fe-S-H^O systems at 25, 100 and 150°C. Sulfide, S^-, currently considered to be stable only at extremely high pH, was excluded from the E-pH diagrams. Reference electrodes, compatible with sulfide solution, were designed for use with an autoclave. The response of the Ag/Ag2S electrode (SSSE) to variation of temperature, sulfide and hydroxide concentration, and chloride addition was studied. Liquid junction potential and thermal liquid junct.ion potential corrections were applied. The SSSE was not reliably predictable thermodynamically, but was stable and was proven effective in the polarization study. The electrochemical potentials of inert metal reference electrodes, Hg and Pt, were found to be con-sistent with the mixed potential between polysulfide, , and thiosulfate, S2O3 * i i The Tafel slopes from anodic polarization curves of Pt in alkaline sulfide solutions were consistent with oxidation of HS~ to S 2 0 3 2 - at the mixed pote n t i a l . At more noble potentials, oxidation to polysulfide occurred. Reaction path mechanisms were proposed. Understanding of su l f i d e oxidation aided in i n t e r -pretation of the electrochemical behavior of steel in the alkaline sulfide solutions. i i i TABLE OF CONTENTS Page Abstract i i Table of Contents iv List of Tables v i i List of Figures v i i i List of Symbols and Abbreviations xi Acknowledgement x i i i 1. . INTRODUCTION 1 1.1 Corrosion Studies of Steel in Highly Alkaline Sulfide Solutions 1 1.2 Alkaline Sulfide Chemistry 11 2. E-pH DIAGRAMS 18 2.1 Introduction 18 2.2 The S-H20 System 20 2.3 The Fe-S-H20 System 39 3. REFERENCE ELECTRODES 54 3.1 Introduction 54 3.2 Experimental 66 3.2.1 Ag/Ag2S Electrode 66 3.2.2 Polysulfide Electrode 71 3.3 Results and Discussion 72 3.3.1 Silver-Silver Sulfide Electrode 72 3.3.2 Polysulfide Electrode 92 i v TABLE OF CONTENTS (Cont) Page 4. POLARIZATION STUDY 104 4.1 Introduction 104 4.2 Experimental 112 4.3 Results and Discussion 119 4.3.1 Cyclic Voltammetry in Plain NaOH Solution 119 4.3.2 Sulfide Oxidation 132 4.3.3 Polari z a t i o n Behavior of A516 Steel in Alkaline Solutions 141 4.3.3.1 Effect of a Low Concentration of Sulfide 141 4.3.3.2 Anodic P o l a r i z a t i o n with a Medium Concentration of Sulfide 151 4.3.3.3 Cyc l i c Voltammetry with a Medium Concentration of Sulfide 160 4.3.3.4 Effect of a High Concentration of Sulfide 168 4.3.3.5 Effect of Scan Range 172 4.3.3.6 Scan Rate 177 4.3.3.7 Po l a r i z a t i o n at the Switching Potentials Between Scans 179 4.3.3.8 Effect of S t i r r i n g 183 4.4 Relevance to P r a c t i c a l Corrosion Problems 186 5. SUMMARY 188 BIBLIOGRAPHY 190 v TABLE OF CONTENTS (Cont) APPENDIX I APPENDIX II APPENDIX III APPENDIX IV APPENDIX V APPENDIX VI The Criss-Cobble Extrapolation Method Calculation of AG° and S° 2 Q 8 of S2 -Standard E l e c t r o d e P o t e n t i a l s of Ele c t r o c h e m i c a l E q u i l i b r i a and Free Energy Changes of Chemical E q u i l i b r i a for the S-H20 and Fe-S-H20 Systems Relative Thermodynamic S t a b i l i t y of Ag20, AgCl and Ag2S Estimation of Oxide Film Thickness Oxidation of Sulfide to Thiosulfate and Polysulfide Page 198 201 203 209 211 213 vi LIST OF TABLES Page I Thermodynamic Data f o r S u l f i d e and P o l y s u l f i d e s 24 II Thermodynamic Data f o r P o l y s u l f i d e s from Reference 37 27 I I I E l e c t r o c h e m i c a l E q u i l i b r i a f o r the S-H20 System 28 IV E l e c t r o c h e m i c a l E q u i l i b r i a f o r the Fe-S-H20 System 44 V Comparison of C a l c u l a t e d P o t e n t i a l s f o r E q u i l i b r i a I n v o l v i n g HFe02~ and Fe(0H)3- i n 3m NaOH 53 VI Standard E l e c t r o d e P o t e n t i a l of Ag/Ag2S 59 VII Standard E l e c t r o d e P o t e n t i a l s at Temperature f o r E l e c t r o d e Reaction 14 80 VIII P o t e n t i a l D i f f e r e n c e , AE, between Ag/Ag2S I n t e r n a l E l e c t r o d e s at Temperature and Ag/Ag2S E x t e r n a l Reference E l e c t r o d e at 25°C 89 IX P o t e n t i a l s of C y c l i c Voltammogram Current Peaks and P o t e n t i a l s of Relevant E q u i l i b r i a at 95, 120 and 150°C i n 3m NaOH. 121 X C o r r o s i o n P o t e n t i a l of Platinum and C a l c u l a t e d E q u i l i b r i u m P o t e n t i a l s at 90°C. 134 XI T a f e l Slopes on Platinum at 90°C. 136 XII P o t e n t i a l s of C y c l i c Voltammogram Current Peaks and Relevant E q u i l i b r i u m P o t e n t i a l s i n 2.5m NaOH +0.05m Na2S at 95, 120 and 1 5 0° C 144 XIII P o t e n t i a l s of P o l a r i z a t i o n Current Peaks and P o t e n t i a l s of Relevant E q u i l i b r i a at 95, 125 and 150°C i n 2.5m NaOH + 0.5m Na2S. 154 XIV P o t e n t i a l s of C y c l i c Voltammogram Current Peaks and P o t e n t i a l s of Relevant E q u i l i b r i a at 95, 120 and 150°C i n 2.5m NaOH + 0.5m Na2S. 162 XV P o t e n t i a l s of C y c l i c Voltammogram Current Peaks and P o t e n t i a l s of Relevant E q u i l i b r i a i n 3m Na2S at 95, 120 and 150°C. 170 v i i LIST OF FIGURES Pag_e 1. Schematic Anodic Polarization Curve. 2. S-H2O at 25°C. Potential-pH diagram considering thio-s u l f a t e as the only oxidized form of s u l f u r . E q u i l i b r i a 1.2,1.3. 29 3. S-H9O a t 25°C. Potential-pH diagram considering poly-sulfides and elemental sulfur as the only oxidized forms. Equilibria 1.2, 1.6, 1.9-1.16. 30 4. S-H2O at 100°C. Potential-pH diagram considering thiosulfate as the only oxidized form of su l f u r . Equilibria 1.2, 1.3. 31 5. S-H2O a t 100°C. Potential-pH diagram considering polysulfides and elemental sulfur as the only oxidized fms. Equilibria 1.2, 1.6, 1.9-1.16. 32 6. S-H2O at 150°C. Potential-pH diagram considering thiosulfate as the only oxidized form of sul f u r . Equilibria 1.2, 1.3. 33 7. S-H2O at 150°C. Potential-pH diagram considering polysulfides and elemental sulfur as the only oxidized forms. Equilibria 1.2, 1.6, 1.9-1.16. 34 8. S-H20 at 25°C. Potential-pH diagram based on Giggenbach's thermodynamic data for polysulfides, considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.9, 1.11, 1.12, 1.21, 1.22. 35 9. S-H20 at 100°C. Potential-pH diagram based on Giggenbach's thermodynamic data for polysulfides, considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.10, 1.12, 1.21, 1.22. 36 10. S-H20 at 150°C Potential-pH diagram based on Giggenbach's thermodynamic data for polysulfides, considering polysulfides and sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.10-1.12, 1.15, 1.16, 1.21. 37 11. Fe-S-H20 at 25°C. E q u i l i b r i a 2.2, 2.3, 2.10, 2.11, 2.17, 2.24, 2.27, 2.30, 2.31, 2.40, 2.47, 2.52, 2.53, 2.58, 2.60, 2.63, 2.81. 48 v i i i LIST OF FIGURES (Cont) Page 12. Fe-S-HoO at 100°C. E q u i l i b r i a 2.2, 2.3, 2.10, 2.11, 2.17, 2.24, 2.27, 2.30, 2.31, 2.40, 2.47, 2.52, 2.53, 2.60, 2.63, 2.81. 49 13. Fe-S-H2o at 150°C E q u i l i b r i a 2.2, 2.3, 2.10, 2.11, 2.17, 2.24, 2.28, 2.30, 2.31, 2.41, 2.48, 2.52, 2.60, 2.63, 2.81. 50 14. External Reference Electrode and Bridge. 68 15. Detail Cross-Section of External Reference Electrode. 69 16. Potential as a function of temperature for SCE/KC1 (sat.) //2.5m NaOH + 0.423m Na2S/Ag2S/Ag. 73 17. Potential as a function of temperature for SCE/KC1 (sat.) //2.5m NaOH + 0.5m Na2S/Ag2S/Ag. 74 18. Potential as a function of temperature for SCE/KC1 (sat.) //2.5m NaOH + 1.0m Na2s/Ag2S/Ag. 75 19. Potential as a function of temperature for SCE/KC1 (sat.) //2.5m NaOH + 0.05m Na2S/Ag2S/Ag. 76 20. Potential as a function of temperature for SCE/KC1 (sat.) //12m NaOH + 0.423m Na2s/Ag2S/Ag. 77 21. Potential as a function of temperature for SCE/KC1 (sat.) //2.5m NaOH + 0.5m Na2S + 1.0m NaCl/Ag2S/Ag. 78 22. Potential as a function of temperature for Ag/Ag2S (298°K)/2.5m NaOH + 0.5m Na2S/Ag2S/Ag(T). 87 23. P t / Sx2 - potential as a function of temperature for SCE/KC1 (sat.) // 2.5m NaOH + 0.423m Na2S/Pt. Open symbols are calculations of E gx2 -/ n s _ for different Sx^"concentrations. 93 24. Hg/Sx2~ potential as a function of temperature for SCE/KC1 (sat.) // 2.5m NaOH + 0.423m Na2S/Hg. Open symbols are calculations of ESx2-/HS- f o r different Sx concentrations. 94 25. Evans diagram of sulfide oxidation in alkaline sulfide. 97 26. Schematic c y c l i c voltammogram of steel in alkaline sulfide solution. 105 27. Schematic of Teflon test c e l l for polarization studies. 113 ix LIST OF FIGURES (Cont) Pa&e 28. Schematic diagram of external reference electrode for autoclave tests ( o r i g i n a l design). 115 29. C y c l i c Voltammogram of A516 S t e e l i n 3m NaOH(50 mV/s). 120 30. Anodic P o l a r i z a t i o n Curves f o r Platinum i n NaOH + NaoS S o l u t i o n s with Varying Na2S20o Concentration (1 mV/s). 133 31. C y c l i c V o 11 a m m o gr a m s f o r A516 S t e e l i n 2.5m NaOH + 0.05 m Na2S (50 mV/s) . 143 32. P o l a r i z a t i o n of A516 S t e e l i n 2.5m NaOH + 0.5m Na2S (1 mV/s). 152 33. C y c l i c Voltammograms f o r A516 S t e e l i n 2.5m NaOH + 0.5m Na2S (50 mV/s). 161 34. C y c l i c Voltammograms for A516 Steel i n 3m Na2S. 169 35. Variation of Cathodic Switching Potential (labeled at s t a r t of each c y c l e ) , 2.5m NaOH + 0.5m Na2S, 90°C, 50 mV/s. 173 36. Variation of Anodic Switching Potential (labeled at midpoint of each cycle), 2.5m NaOH + 0.5m Na2S, 90°C, 50 mV/s. 176 37. Effect of Scan Rate, 2.5m NaOH + 0.5m Na2S, 90°C 178 38. Effect of Scan Rate, 2.5m NaOH + 0.5m Na2S, 120°C. 180 39. Effect of Scan Rate, 2.5m NaOH + 0.5m Na2S, 150°C 181 40. E f f e c t of Cathodic Cathodic Switching NaOH + 0.5m Na2S, 50 P o l a r i z a t i o n f o r 300 s at the P o t e n t i a l between C y c l e s . 2.5m mV/s. 182 41. Effect of Anodic Polar i z a t i o n for 300 s at the Anodic Switching P o t e n t i a l before the Reverse Scans. 2.5m NaOH + 0.5m Na2S, 50 mV/s. 184 42. E f f e c t of S t i r r i n g , 50 mV/s. , 2.5m NaOH + 0.5m Na2S, 90°C, 185 x LIST OF SYMBOLS AND ABBREVIATIONS Tafel slopes, anodic and cathodic, mV/decade current. symmetry factor. molar heat capacity at constant pressure, J/mol °K. molar heat capacity at 1 atm. at standard state. average molar heat capacity between T^ and T 2 . film thickness. potential difference, V. standard electrode potential, V. Faraday, 96500 coulombs. free energy of formation, kj/mol at 298°K. activity coefficient. enthalpy, J/mol. current density, J/mol. exchange current density. equilibrium constant. second dissociation constant of H2S. dissociation constant of H2O. liquid junction potential, V. transmission coefficient. 2 — 1 equivalent conductance, cm A" . molal. molar . molecular weight. xi LIST OF SYMBOLS AND ABBREVIATIONS (Cont) MMOE mercury-mercuric oxide electrode, n oxidation state, valency. overpotential, V. Q number of equivalents. R gas constant 8.31 J/mol °K. S° entropy, J/mol °K. S ° "absolute" entropy, J/mol °K at 298°K. SCE saturated calomel electrode. SHE standard hydrogen electrode at 25°C. SSSE silver-silver sulfide electrode. sp. gr. specific gravity. T temperature. TLJP thermal liquid junction potential. ionic mobility, cm2 o,-l. z ionic valency. approximately., x i i ACKNOWLEDGEMENT I am indebted to Professor Des Tromans for his guidance and suggestions throughout this work, and for pushing me that extra distance to improve my experimental interpretations. I am very grateful for the understanding and support of my wife Tracy, which have eased the effort of the work. For two years I received a MacMillan Family Fellowship. It is gratefully acknowledged, with thanks to the people who endowed the fellowship, and those who awarded i t to me. Fi n a l l y I must express my appreciation to family, friends and colleagues who encouraged and assisted me. x i i i 1. INTRODUCTION 1.1 Corrosion Studies of Steel in Highly Alkaline Sulfide  Solutions The kraft wood pulping process u t i l i z e s s u l f i d e aqueous solutions to digest wood chips. Large digester vessels, of batch or continuous design, contain the process liquor at high temp-erature and pressure, up to 170°C and 1.03 MPa. These vessels and ancillary equipment now are constructed typically of A285 or A516 steel* and are subject to severe corrosion. General corrosion may cause unacceptable thinning of vessel walls. Stress corrosion and corrosion fatigue may result in catastrophic f a i l u r e2. The alkaline sulfide pulping liquor varies in composition from plant to plant and changes throughout the digestion process. Fresh pulping solution, termed "white liquor", contains the main pulping ingredients, hydroxide and s u l f i d e , and several impur-i t i e s including chloride, carbonate and oxidized s u l f i d e s . During the cook, the liquor is heated to approximately 170°C and hydroxide is consumed as the lignin in the wood chips is prefer-e n t i a l l y dissolved, leaving the cellulose fibers behind as a pulp. It is inevitable that corrosion of carbon steel w i l l occur in the presence of the hot cooking liquor but unfortunately the corrosion kinetics are poorly understood. The understanding developed in the past 35 years provides a useful foundation for further study. 1 The Swedish kraft pulp industry became concerned with corrosion and established a committee in 1951 to study the prob-lem. Increasing concentrations of Na2S, NaOH, Na2S203 and Na2Sx stimulated corrosion, as did higher silicon and carbon content in the steel-^. During "cooking", corrosion stopped when a temp-erature of about 120°C was attained. Further, they claimed that the amount of corrosion was related to the amount of oxygen present in the (batch) digester^. Apparently, American pulpmakers encountered problems around the same time. The account of a lab study indicated that poor reproducibility of test results was related to the surface con-d i t i o n of the s t e e l , that i s whether i t was active or passive^. Polysulfide solutions prepared by adding sulfur dependably s t a b i l i z e d the surface although very low concentrations were perhaps detrimental. Solutions containing sulfur as a mono-sulfide (HS~, S2-) and thiosulfate were much more corrosive. The electrochemical potential of passive surfaces was found to be the same in mono- or polysulfide solutions, posing a seeming dilemma to the investigators trying to explain the advantages of poly-s u l f i d e . They suggested that the polysulfide benefited the passive f i l m by depolarizing hydrogen before i t could cause damage to the film. Haegland an d Roald6, surveying the previous investigations, suggested t h a t t h e sulfur compounds of intermediate oxidation states stimulated corrosion through their cathodic reduction or by removing nascent hydrogen. In their own measurements, they could detect no hydrogen evolution from digester steels immersed in white liquor; either i t doesn't form, or i t is a l l consumed in s u l f i d e reduction. A reduction of cathodic polarization which resulted when sulfur was added was attributed to the presence of po1ysu1fides. Thiosulfate increased polysulfide (Sx^-) con-centration and s u l f i t e (SO32-) decreased i t , apparently via the equilibrium: s 2 o 3 2 ~ + S2~ ^ s o 3 2 ~ + S22~ . . . (1) It was concluded that the corrosion reaction consists of anodic dissolution of ferrous ions which precipitate in the solution as ferrous sulfide, with cathodic reduction of polysulfide ions to s u l f i d e . Mass transfer of polysulfide to the surface through a developing layer of ferrous sulfide was considered to control the corrosion rate. The importance of high concentrations of poly-s u l f i d e in maintaining p a s s i v i t y i n d i c a t e d to them that polysulfide played a role in conserving oxygen type passivity rather than producing a sulfide type passivation. In a later paper, Roald^ developed an equation to measure the corrosivity of a solution, based on the rate of diffusion of Sx2 - to the electrode surface. Mueller^, at the Pulp and Paper Research Institute of Canada, implemented the use of anodic polarization curves to 3 explain the activity or passivity of the steel within a framework of electrochemistry. A large anodic current peak was observed over a narrow potential range between the active and the passive potential ranges, as illustrated schematically in Figure 1. The natural corrosion potential of the steel was above or below this peak, that is passive or active respectively. Higher concen-trations of oxidizing substances c o u l d s h i f t the corrosion potential in the positive d i r e c t i o n . During a batch cook, the potential shifted into the passive range after a t i m e ^ . U n f o r -tunately, in work with mixtures of black (spent) and white (fresh) liquor, corrosion potentials were not reported. Changes in corrosion potential might have accounted reasonably for some of the observations regarding activity and passivity. The Institute of Paper Chemistry adopted and improved the use of polarization curves in a study of the effect of liquor additives-^. Carbonate, chloride and sulfate slightly depressed the maximum anodic current density, but corrosion rate was i n -creased only i f sodium sulfite concentration was less than 1 gpL. The additives had a negligible or slightly inhibitive effect when sulfite concentration was higher. Corrosion potentials were not discussed, but may have been influenced by the s u l f i t e content which is in equilibrium with other sulfide species such as poly-sulfide, via equation 1. Landmark and Roaldll determined that increasing concen-t r a t i o n s of p o l y s u l f i d e s d e c r e a s e d the c r i t i c a l a p p l i e d a n o d i c current density (active/passive peak) to passivate the material. LO to c Ol o 0) x» o o o> LJ Passive Active/Passive Peak Active Logarithm of Current Density . i Figure 1: Schematic Anodic Polarization Curve The c r i t i c a l current density was decreased by the amount of the current supplied through the reduction of po1ysu1fides. The corrosion potential was shifted upward with addition of poly-sulfide . Muellerl2 measured polarization curves to 135°C i n a mixture of white and non-oxidized black liquor. Current peaks decreased in s i z e at 130 and 135°C during a laboratory cook, probably due to consumption of hydroxide by the process. Anodic protection of digesters was shown to be feasible. The possibility of anodic protection of pulp digesters was evaluated further by means of polarization studies of 1020 steel up to 177°C.13 The investigators suggested that, as the temp-erature increased, the cathodic hydrogen reaction current over-shadowed the iron anodic dissolution current. At 177°C, the steel would passivate spontaneously. Their polarization curves indicate that the corrosion potential was s h i f t i n g during the polarization scan, obscuring the effect of temperature on the shape of the anodic polarization curve. In conjunction with a polarization study, T r o i a n s ^ perform-ed analyses of corrosion products formed on steel in an alkaline s u l f i d e solution at 92°C. These results were related to react-ions expected at various potentials according to the polarization curves and thermodynamic c a l c u l a t i o n s . P a s s i v a t i o n was i n h i b i t e d by the i n c o r p o r a t i o n of sulfide into the Fe^O^ f i l m . During 6 p o l a r i z a t i o n scanning, p a s s i v a t i o n could occur only when s u f f i c i e n t l y noble potentials were achieved to oxidize the i n -corporated sulfide ions to S2032 -« Wensley and Charltonl5} in a study of polarization behaviour in kraft white liquor, agreed that sulfide impaired passivation, as did t h i o s u l f a t e . Sulfate and s u l f i t e had no e f f e c t . The corrosion potential was controlled by polysulfide concentration. Ahlersl6 examined the effect of polysulfide on corrosion of carbon steel. More than 1 gpL of polysulfide ensured immediate passivation. They did not mention whether corrosion potential was controlled during weight-loss experiments. The appearance of corrosion products depended on the amount of elemental sulfur added, a r e s u l t probably r e l a t e d to unmeasured change of corrosion potential. Stress corrosion of mild steel in alkaline solutions was investigated in the classic paper by Singbeil and Tromans^. In slow strain rate tests, cracking occurred within a range of potentials between the active and passive ranges. They warned that, i f the corrosion potential in kraft digesters lay within the cracking range, stress corrosion would occur. Failure of a continuous digester at Pine H i l l in 1980 confirmed their spec-ulation2. Sulfur additions affected the corrosion potential of steel 7 in NaOH + Na2S solution. Polysulfide would be formed via the equilibrium: (x-1) S + OH- + HS" =^  Sx2- + H20 ... (2) Polysulfides, which are r e l a t i v e l y strong oxidants, raised the corrosion potentia 117. In another a r t i c l e , Singbeil and Tromans described the effect of sulfide on caustic cracking^. Sulfide ions increased the size of the active-passive current peak and displaced i t to more noble potentials. The potential range for cracking was thus displaced to higher potentials, too. Cracking rates from frac-ture mechanics tests were consistent with a dissolution process of cracking. Hydrogen embrittlement was shown by thermodynamic calculation to be unlikely in hot alkaline solutions at the potentials of study. The same authors published results of further study of cracking rates using fracture mechanics specimens^. In elab-orating the mechanism of cracking, they considered that cracking was slower in solutions containing sulfide than in straight NaOH because of faster repassivation kinetics and the effect of sul-fide on iron dissolution rates. As part of the response to the f a i l u r e of the digester at Pine H i l l , work at the Institute of Paper Chemistry was funded through the Technical Association of the Pulp and Paper Industry (TAPPI)20. A survey of pulp mills indicated that digesters with large rated capacities or incomplete stress r e l i e f were more susceptible, as were those pulping southern softwoods or mixed hardwood/softwood species. Possible remedies for cracking were discussed. In t h e i r f i n a l r e p o r t2^ they claimed that no correlations could be found to relate s u s c e p t i b i l i t y to major, minor or trace constituents of cooking liquors. A group at Battelle Pacific Northwest Laboratory determined that a kraft digester they investigated suffered cracking where the corrosion potential was a c t i v e2 2. Auger electron spectro-scopic measurements suggested that S, CI and C were di f f u s i n g into the material ahead of the crack tip. Slow strain rate and fracture mechanics tests were performed at the Pulp and Paper Research Institute of Canada with the support o f TAPPll. No cracking was detected under simulated acid cleaning conditions, although pitting was observed in inhibited HC1 solutions. Polarization behavior was unaffected by additions of Na2C03, NaCl or Na2S203 but was affected by the presence of NaOH and Sx2~. Slow strain rate experiments were performed on digester base metal (A516 Grade 70) and weldment (7018 f i l l e r ) in solutions simulating impregnation and cooking zone liquors of a kraft continuous digester. Cracking occurred in a l l solutions at a l l temperatures tested. In slow strain rate tests, the 7018 weld-9 merit appeared to crack more readily than the base metal. Long term fracture mechanics tests were carried out in simulated impregnation zone liquor in a p i l o t plant digester. Anodic protection was recommended to prevent cracking of digesters23. Although considerable progress has been made on understand-ing corrosion of steel in highly alkaline sulfide solutions more questions remain. This present work seeks to answer some of the questions. In particular, the effects of higher temperature on polarization behavior wil l be explored. The few high temperature studies to date have not explained the polarization behavior in thermodynamic terms13* 2 3. The corrosion process i t s e l f should be better defined by thermodynamics so that i t becomes less mysterious to the corrosion engineer. Reference electrodes for use in sulfide solutions are poorly developed. These are the problems addressed by this work. High temperature, high pressure polarization tests in an autoclave w i l l be described. A variety of temperatures and solution compositions have been investigated. Reference elec-trodes for use in alkaline s u l f i d e solutions have been studied and then employed in the polarization study. E-pH diagrams for S-H2O and Fe-S-H2° systems have been constructed to aid in inter-preting the thermodynamics of the polarization behavior. Some analyses of corrosion products w i l l be interpreted with regard to the E-pH diagrams and the polarization curves. The resultant improved understanding of the corrosion process w i l l be discussed. 10 1.2 Alkaline Sulfide Chemistry The chemistry of aqueous alkaline sulfide solutions is com-plicated and contentious. This poses obstacles to corrosion studies which should be f u l l y described before they are encountered in the work. As well, this w i l l provide the ground-work for reviews on the thermodynamics and electrochemistry of sulfides later in the thesis. Sulfur, in i t s most reduced form, i s present in aqueous solutions as H2S, HS~ or S2 - depending on the pH of the solution. The equilibrium between H+, H2S and HS~ via reaction 3: (H2S)ag 5* H+ + HS~ . . . (3) has an equilibrium constant pKj = 7 and this value i s widely acce pted2 4. The equilibrium between HS~ and S2 - is the subject of continuing disagreement. HS- ^ H+ + S2" K 2 - fH+1 r s 2 - ] [HS~] ... (A) . . . ( 4 A ) K2, the so-called second dissociation constant of hydrogen sulfide is in dispute. Knox2^ obtained a value of 14.92 in 1908. Values of 12.88 and 13.90 were published later.26, 27 in 1 9 5 9 , 11 pK2 was measured as 14 at 20°C by E l l i s and Golding^ . By the 1960's this value was accepted by many2^- 3 0. Macdonald and Hyne3! expressed some concern that the pK2 value tabulated in the Handbook of Thermodynamic Data3 2 was too low (12.92 at 25°C). In 1971, Giggenbach3 3» 3^ published results of a new deter-mination of K2 obtained using an improved technique to exclude oxygen. A value of pK2 — 17.1 at 24°C was advanced. Giggenbach suggested that previous lower values resulted from slight oxid-ation of s u l f i d e . Although there have been some doubts about this value3^, recent experiments have confirmed i t3^ . The higher value of K2 means that S2 - is not present in significant amounts u n t i l the pH i s close to 17. Thus HS~ i s the important form of sulfide in solutions of pH 7-17. Employing spectrophotometry, Giggenbach obtained upper limits for K2 to 80°C. Stephens and Cobble2^ developed an equation for K2 up to 250°C based on the lower values reported in the l i t e r a t u r e for lower temperatures. Because of this assumption, the K2 values d i f f e r from those by Giggenbach but the variation with temp-erature is similar. Sulfide may be oxidized to a large number of other species in aqueous solution. These include thiosulfate (S2032 -)» poly-sulfides (Sx2~), sulfite (S032 _) and sulfate ( S O 42 -) . Thiosulfate, polysulfides, and sulfite are metastable forms of sulfide, that is they may be oxidized further. Metastable forms predominate in 12 many sulfide solutions. Sulfite and Sulfate Sulfite and sulfate were not formed in significant amounts during disproportionation of polysulfide over a wide range of pH and temperature37. 38. Oxidation to sulfate is very slow2 8. Thiosulfate Avrahami and Golding2 8 studied the formation of S2032 - as an _ O -intermediate oxidation product in the oxidation of HS to SO^ in the pH 11-13 range. The oxidation rate increased rapidly with 2-increasing pH. They claimed that the rate of oxidation of S2O3 to SO^2- was very slow. The presence of SO32 - led rapidly to formation of S2O32- ions. Giggenbach37 revealed that thiosulfate was the predominant oxidized species present above 100°C in a l k a l i n e s u l f i d e solutions. The thiosulfate was formed through decomposition of polysulfides, as described below. Polysulfide Pol y s ul f i de s , Sx^-~, are present in solution always with a range of ionic sizes (e.g. S22", S32-, Sz.2-) in equilibrium with each other and in turn in equilibrium with other sulfide species in solution. The presence of polysulfides is indicated by their 13 yellow color in s o l u t i o n . The complexity of p o l y s u l f i d e solutions has inhibited study of them. Moreover, they are susceptible to oxidation, thus complicating analysis. Research-ers have directed their e f f o r t s towards determining ranges of conditions in which polysulfides are s t a b l e ^ . The e q u i l i b r i a between polysulfides and other species, especially sulfur, have been investigated3^> 3^» 40". Early measurements^1 based on electrode potential measure-ments suggested that a l l polysulfide ions, Sx2 - (x = 2-5), were formed in the pH 10-13 range. Cloke^2 investigated the d i s t r i b u t i o n of polysulfides via electrode potential. His model included S2 -, S^~, S 5 2 - , S £2~ and HS-. Much of the work on polysulfide solutions has been done in connection with polysulfide pulping. These solutions have been formed by the addition of sulfur to alkaline sulfide solutions. Gustafsson and Teder3^ studied the thermal decomposition of aqueous polysulfide solutions (to S2032 - and HS-) via: nS2032 - + 2(n+2) HS" ==* 4SnS2" + 4(n-l) OH- + (4-n) H20 ... (5) The limit between stable and unstable polysulfide solutions was described by an equation in [0H~] and [ H S-] . As the s t a b i l i t y increasing effect of [ H S-] was considerably larger than the 14 counteracting effect of [OH-], concentrated polysulfide solutions as a rule were more stable than dilute ones. The smaller the polysulfide, the higher was the activation energy for decompos-i t i o n . Thus, in highly a l k a l i n e s o l u t i o n s , with smaller polysulfides, decomposition would be hindered. In the very con-centrated p o l y s u l f i d e s o l u t i o n s which Teder s t u d i e d , decomposition r a t e s of polysulfide were independent of t h i o -sulfate concentration. Teder^O obtained equilibrium constants between each poly-su l f i d e (S x 2~, x = 3-6) and pentasulfide (Ss 2 -) by means of spectroscopy at 25°C and 80°C. His results revealed that tetra-and pentasulfide ions predominated at pH 9-14. The polysulfides equilibrated very rapidly after adjustments to OH- or HS-. Giggenbach38 determined the equilibrium d i s t r i b u t i o n of polysulfide ions in aqueous solution (pH 6.8-17.5) at 20°C. Di-, t r i - , t etra- and pentasulfide ions were present. S 5 was not detected. Equilibrium constants were derived as a function of [OH-] [SH-]/S(o) where S(o) i s the t o t a l c o n c e n t r a t i o n of z e r o -valent sulfur in solution, n[S-S2 -]. Other constants were developed to relate the concentration of one polysulfide to another. The tetrasulfide ion, S^2-} predominated in solutions of intermediate a l k a l i n i t y , followed at higher pH by the t r i -sulfide ion, S 3 2 - , ^he disulfide ion, S 2 2 - , was the predominant species only at extremely high a 1ka1inities. The pentasulfide i o n , S 5 2 -f was formed i n s i g n i f i c a n t amounts o n l y i n n e a r l y neutral solution. 15 Subsequently, Giggenbach37 found that the equilibrium con-stants between the polysulfides varied l i t t l e with temperature up to 240°C. Disproportionation into sulfide and t h i o s u l f a t e and dissociation of polysulfide ions into" radicals S2- or S3- became s i g n i f i c a n t above 150°C and was accompanied by decoloration of the solutions. At pH >8, polysulfide ions became metastable even at room temperature. Addition of s u l f i t e or thiosulfate i n -creased p o l y s u l f i d e concentration under some c o n d i t i o n s , presumably via: nS032- + SnS2~ + H20 ^ nS2032- + HS~ + OH" ... (6) There was evidence that only thiosulfate reacted with monosulfide to form zerovalent sulfur. At pH 14-16 polysulfide was found to be present as t r i - and disulfide and was very stable up to 100°C. Disproportionation to thiosulfate was rapid only above 100°C.3 7 A kinetic study^3 explained the high kinetic stability of polysulfide solutions at high pH in terms of the pentasulfide ion being the main reactive polysulfide species. Smaller polysulfides would react only by f i r s t equilibrating to S^2-. Metastable species are the rule rather than the exception in s u l f i d e s o l u t i o n s . Their presence must be considered in corrosion studies. In experimentation, precautions against 16 p a r t i a l o x i d a t i o n i n c l u d e c a r e f u l s e a l i n g of c h e m i c a l s i n b o t t l e s and s t o r i n g s o l u t i o n s under n i t r o g e n . In a n a l y s i s of r e s u l t s , the r o l e of p o l y s u l f i d e s i n the c o r r o s i o n p r o c e s s must be a p p r a i s e d . 17 2 . E-pH DIAGRAMS 2 . 1 Introduction E-pH diagrams have become an important part of corrosion science since their development by Pourbaix^. They provide, at a glance, a means of determining which ionic species or s o l i d compound i s thermodynamica 11y stable in aqueous solution at a given pH and electrochemical potential. Metallic dissolution is usually associated with formation of some soluble metallic species and passivation often results from formation of an oxide layer. By means of measurement of pH and p o t e n t i a l , and com-parison with potential-pH diagrams, corrosion processes can be analyzed in terms of electrochemical reactions. The boundaries of stability of each species are determined by Nernst equations for the electrochemical reactions or by chemical equilibria. A large number of E-pH diagrams have been constructed over the years, some for single metals (e.g. Fe-R^O) a nd others for more complicated systems (e.g. Fe-S-H 20). The a c t i v i t i e s of soluble species have been set arbitrarily for use in the Nernst equations used to construct the diagrams or have been chosen to correspond to a particular system being studied. E-pH diagrams have been constructed in the present work, with special attention devoted to making them r e f l e c t the true composition of the corrosive solutions. F i r s t , S-H2O diagrams were developed to aid in understanding sulfide reactions. Second, Fe-S-H^O diagrams were constructed to elucidate the 18 mechanisms of corrosion of steel. Relevant E-pH diagrams already available and sources of thermodynamic data are referenced in the section on each system ( S - H2Q and Fe-S-H20). The methods of construction of E-pH dia-grams are well establishe d*4. Criss-C obble extrapolation was used to develop the diagrams at elevated temperatures^' -^*» The Criss-Cobble method is described in Appendix I. Wherever thermodynamic data were estimated, the method of estimation has been explained. The convention adopted in the present work has been that the standard hydrogen reference electrode potential is a function of temperature and i s zero only at 298°K. Duby^7 has c r i t i c i z e d this method because i t introduces an additional error when est-imating the temperature c o e f f i c i e n t of the hydrogen electrode. On the other hand, i t is more convenient for corrosion engineers to convert measurements to the standard conventional electrode because in many situations we place the reference electrode outside the test chamber. Errors in calculated potentials have not been estimated here. The uncertainties in thermodynamic data are not published or, for that matter, known. Extrapolation methods such as the Criss-Cobble theory^8 are inherently inexact to an unknown extent but in the absence of experimental data extrapolations provide a u s e f u l approximation. This i m p r e c i s i o n d i c t a t e s that 19 potential-pH diagrams developed in this way can serve only to indicate "landmarks" in the thermodynamic realm. The accuracy of Criss-Cobble extrapolations is considered satisfactory to 200°C, beyond the temperature considered here. 2 . 2 The S-H2Q System Potential -pH diagrams of the sulfur-water system have been developed and.refined over the past four decades. Some of the f i r s t diagrams were presented by V a l e n s i ^ a nj Pourbaix^. Dia-grams for the system at elevated temperatures were constructed by Biernat and Robins^ using estimates of high temperature heat capacities of ions in solution. P e t e r s ^ published a S - H 2 O diagram for metastable species, neglecting sulfate, SO^2-. Most recently, Murray and Cubicciotti-*1 tabulated thermodynamic values of sulfur species to 300°C and included diagrams at 300°C for diff e r e n t metastable forms of oxidized sulfur ( s u l f i t e , thio-sulfate and sulfur). Metastable forms of oxidized sulfur predominate in many s u l f i d e s o l u t i o n s . A comprehensive study of s u l f i d e s o l u t i o n s3 7 , 38 showed that at pH >8 polysulfides were metastable to ~100°C and then slowly equilibrated with thiosulfate at higher temperatures until thiosulfate predominated above ~130°C. These results indicate that polysulfide is a significant component of moderately concentrated sulfide solutions at high pH. The significance of polysulfide in wood pulping solutions has long 2 0 been r e a l i z e d3^ , 40, 52. gy their presence, polysulfide and thiosulfate must be considered in studies of corrosion electro-chemistry. Already, polysulfide has been shown to affect the corrosion potential of steels°» 17. As an aid to further study, two sets of potential-pH dia-grams have been constructed in the present study at 25, 100 and 150°C. In one set, thiosulfate has been taken as the most highly oxidized form with s u l f i t e and sulfate disregarded for kinetic reasons. In the second set polysulfide and elemental sulfur have been considered to be the only oxidized forms. E-pH diagrams for thiosulfate or polysulfide were not available previously. They have been constructed because they were considered essential to understanding the behavior of aqueous s u l f i d e s o l u t i o n s encountered in this study. Furthermore, polysulfide may be the only oxidized sulfide present in significant concentrations. Thermodynamic Data Obtaining thermodynamic values for aqueous sulfide species has been a major obstacle to the construction of E-pH diagrams. Values are poorly established, partly as the result of dis-agreement on the second dissociation constant of hydrogen sulfide, K2. The value of pK2 = 17.1 at 25°C3^ accepted here, is a much lower value than that used previously in calculating the thermodynamic data for aqueous sulfide species26, 38, 30-32^ With pK2 = 17.1, HS- could be considered as the predominant form of reduced sulfide throughout the pH range considered (pH 8-16). 21 The lower K2 value necessitated recalculation of thermodynamic values for aqueous sulfide species. The AG0 and S° of S2 - were calculated to be 109.7 kJ/mol and -94.6 J/mol °K (see Appendix II). Similar calculations were apparently done by others using the old K 2 data to obtain the d i f f e r e n t values given by them26- 3 0. Thermodynamic data for polysulfides were d i f f i c u l t to obtain and poorly established. Bichowsky and Rossini^3 listed AH values for S 2 2 - > S3 2"" and S4 2"" which they calculated from data obtained in 1891 by Sabatier. Their values were used by Latimer^, along with estimated values for S° to calculate A G° values. Maronny and Valensi^S, 56 provided data for S 5 2 - and S^2" obtained from electrochemical measurements and in 1957 systematized the values for Sx2 - (x = 1-5)-*7. Barnes and Kullerud^8 seem to have been unaware of this l a t t e r work and l i s t e d the same values as Latimer. Cloke^2 recalculated AG° values but assumed the 2_ presence of S§ The NBS30 l i s t e d d i f f e r e n t thermodynamic values for the polysulfides without explanation or reference. The values are close to those of Maronny and V a l e n s i ^ with the An va iues showing the most difference. Giggenbach37 published values for AG0 and S° of Sx2_(x=2-5), based on equilibrium constants of sulfur dissolution. He assumed that A H and A Cp| 2 9 8 w e r e equal to zero throughout the 2 2 range investigated. These assumptions are inadequate with res-pect to previous investigations. No values for S 2 ~ or S were provided; this raises doubts about how to relate values available from other sources. F i n a l l y , Giggenbach3 7 admitted that his values were subject to "considerable modification". Naumov et a l3 2 tabulated AH, A G° and S° values for poly-sulfides which are the same as those of the NBS30. They also included Cp values for Sx2 - (x = 1-3) calculated by an "empirical correlation between S 2 0 and C p2° " ' Murray and Cubicciotti^l apparently were not confident about the A G° and S° values tabulated. Their reservations may have been related to the general uncertainty about thermodynamic data on aqueous sulfides or to their acceptance of the smaller value for K2. They listed AG and AS°values for the reactions: S + Sx2" = Sx + 12~ (x = 1-4) ... (7) and these values were the same as those calculated using the NBS data3 0. values tabulated by Murray and Cubicciotti^1 calculate A G° and S° values for the poly-values are listed in Table I. The AG and AS° were used here to sulfides, and these 23 TABLE I Thermodynamic Data for Sulfide and Polysulfides AG° kj/mol So* J/mol°K Cp° Cp 373 298 Cp 423 298 J/mol°K s2- 109.7 -94.6 -188 -242 -253 S22- 103.4 -51.5 -519 -599 -573 S32- 97.5 -13.8 -460 -519 -498 92.9 23.4 -381 -440 -427 S52- 89.6 60.7 -310 -360 -351 *Note: These are not absolute entropies. 24 Heat capacity values for the polysulfides were especially d i f f i c u l t to obtain. The Handbook of Thermodynamic Data3 2 did not include values for S^2- a n (j S 5 2 - . Macdonald and Hyne31 used the values in the Handbook for Sx2 _ (x = 1 _ 3) from which to extrapolate values for S^2~ and S52 _. They used the extrapolated values to plot concentrations of HS~, S2~ and Sx2 - at selected temperatures as a function of pH. The heat capacity values for Sx2 - have been recalculated here. The Criss-Cobble metho d48 was employed, considering S 2 - to be a simple anion and Sx2~ (x = 2-5) to be in the oxyanion class. This approach obtains values of the same order of magnitude for S 2 2 - and S32 -, but a much smaller value for S 2 - as compared to the Handbook of Thermodynamic Data3 2. The calculated values are in Table I. These heat capacity values have the opposite trend from Macdonald and Hyne's values3 1. A G° and S° values for S2032~, HS~, H+, H20 and S032~ were obtained from the Handbook of Thermodynamic Data3 2. Heat cap-acities of the ions were calculated by the Criss-Cobble method. AG° and S° values for S were obtained from reference 30 and were considered to be consistent. Heat capacity values of S and H 2 0 at 298° were obtained from this latter source. High temperature values of heat capacity of S and H20 were taken from Biernat and Robins*6. Note that sulfur is liquid at 150°C. A second set of E-pH diagrams of polysulfides at 25 and 100°C were constructed using Giggenbach's AG° and S° values for 25 p o l y s u l f i d e s3 7. Again heat capacity values were calculated by the Criss-Cobble method. The thermodynamic values provided by Giggenbach and the Cp values calculated from them are listed in Table II. A l l other data remained the same as those used in the f i r s t s e t . Equilibrium Equations High temperature e q u i l i b r i a were c a l c u l a t e d f o r t h e reactions listed in Table III using the method described by Criss and Cobble^ to obtain the change in free energy of the reactions at elevated temperatures. Activities of reactants were approx-imated and s u b s t i t u t e d into the e q u i l i b r i u m equations. A c t i v i t i e s used were 10--6 fo r S 2 O 3 2 - and S032-> 10-3 for Sx (x = 1-5), 10_1 for HS- and 1 for S unless otherwise s p e c i f i e d . These activ i t i e s approximate to concentrations anticipated in the solutions used in the experimental study except S which was assumed to have unit a c t i v i t y i f present in the so l i d phase. Standard electrode potentials at 25, 100 and 150°C have been listed in Appendix III. Electrochemical potentials calculated from the Nernst equilibrium equations were plotted versus the standard hydrogen electrode at 25°C as a function of pH at 25, 100 and 150°C in Figures 2-10. 26 TABLE II Thermodynamic Data for Polysulfides from Reference 37 AG° kJ/mol • So* J/mol°K Cpo Cp"l 373 | 298 J/mol°K S22- 85 10 -402 -468 S32- 77 72 -301 -337 66 105 -226 -267 S52- 67 102 -226 -272 ote: These are not absolute entropies. 27 TABLE III Electrochemical Equilibria for the S-H20 System 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 1.10 1.11 1.12 1.13 1.14 1.15 1.16 1.17 1.18 1.19 1.20 1.21 1.22 02 + 4H+ + 4e~ 2H20 2 H+ + 2e~ ^  H2 S2032 - + 8H+ + 8e~ ^ = 2HS~ + 3H20 2 S032 _ + 6H+ + 4e~ ^  S2032 _ + 3H20 S2- + H+ =s=- HS" S + H+ + 2e" ^ = HS~ S O 3 2 - + 6H+ + 4e~ =?= S + 3H20 s2°32~ + 6 H + + 4 e~ ^ = 2S + 3 H2° S52~ + 5H+ + 8e~ 5 HS" S42 _ + 4H+ + 6e~ 4 HS" S32- + 3H+ + 4e" 3 HS~ S22" + 2H+ + 2e~ ^  2 HS-5 S + 2e" ^  S52" 4 S52 - + 2e" ^  5 S42" 3 S42~ + 2e" ^  4 S32~ 2 S32 _ + 2e" ^ » 3 S22 _ S22 _ + 2e~ ^  2 S2-5 S2032" + 30 H+ + 24e~ 2 S52- + 15 H20 2 S2032- + 12 H+ + 10e~ ^ » S42 _ + 6 H20 3 S2032 - + 18 H+ + 16e~ ^  S32" + 9 H20 4 S + 2e" =-= S4' S42~ + 2e" ; 2-2 S22 _ 28 'e 9 10 11 12 13 V, 15 16 pH Figure 2: S-H2O at 25°C. Potential-pH diagram considering thio s u l f a t e as the only oxidized form of s u l f u r . E q u i l i b r i a 1.2, 1.3. 29 S-H2O at 25°C Potential-pH diagram considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.6, 1.9-1.16. 30 -0.31 1 1 i 1 1 1 r Figure A : S-H2O at 100°C. Potential-pH diagram considering thiosulfate as the only oxidized form ofsulfur. E q u i l i b r i a 1.2, 1.3. 31 -0.2 r~ 1 1 "—1 1 1 r Figure 5: S-H^ O at 100°C. Potential-pH diagram considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.6, 1.9-1.16. 32 S-H2O at 150°C. Potential-pH diagram considering thiosulfate as the only oxidized form of s u l f u r . E q u i l i b r i a 1.2, 1.3. 33 S-H2O at 150°C. Potential-pH diagram considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.6, 1.9-1.16. 34 Figure 8: S-H2O at 25°C. Potential-pH diagram based on Giggenbach's thermodynamic data for po l y s u l f i d e s and considering polysulfides and elemental sulfur as the most oxidized forms. E q u i l i b r i a 1.2, 1.9, 1.11, 1.12, 1.21, 1.22. The shaded region i s uncertain. 35 Figure 9: S-H2O at 100°C. Potential-pH diagram based on Giggenbach's thermodynamic data for polysulfides, considering polysulfides and elemental sulfur as the only oxidized forms. E q u i l i b r i a 1.2, 1.10, 1.12, 1.21, 1.22. 36 F i g u r e 10: S - H 2 0 a t 1 5 0 ° C . P o t e n t i a l - p H d i a g r a m b a s e d o n G i g g e n b a c h ' s t h e r m o d y n a m i c d a t a f o r p o l y s u l f i d e s , c o n s i d e r i n g p o l y s u l f i d e s a n d s u l f u r a s t h e o n l y o x i d i z e d f o r m s . E q u i l i b r i a 1 . 2 , 1 . 1 0 - 1 . 1 2 , 1 . 1 5 , 1 . 1 6 , 1 . 2 1 . 37 General Observations A problem was encountered in constructing the high temp-erature diagrams. The transitions from HS- to S2 -, as calculated from the thermodynamic data, did not agree with those extra-polated from lower temperature experimental d a t a3 4. The extrapolated experimental data were considered to be correct, that i s the extrapolated K2 values were accepted. The diagram was restricted to a pH below the HS-/S2 - equilibrium to avoid the discrepancies. There i s a s i g n i f i c a n t difference in the potential pH dia-grams for the thermodynamic values of polysulfides as calculated here versus those provided by Giggenbach3 7. Polysulfide form-ation would occur at a lower potential according to Giggenbach's values. s32 - had no region of s t a b i l i t y at 25°C or 100°C. A region of uncertainty in Figure 8 resulted from the instability of S32"". None of the diagrams drawn in the present work reflect the observed coexistence of a range of polysulfides in a s o l -ution. If thiosulfate is the major species present at 150°C, as experiment shows37, then the potential pH diagrams of polysulfide at that temperature may be useful only i f polysulfide persists for kinetic reasons. 38 2.3 The Fe-S-H?o System The Fe-S-H^O system i s very important to c o r r o s i o n s t u d i e s because i r o n i s f r e q u e n t l y i n c o n t a c t w i t h s u l f i d e - c o n t a i n i n g s o l u t i o n s , e.g. K r a f t p u l p i n g , G.S. heavy w a t e r . The e v o l u t i o n of E-pH diagrams f o r t h i s system has been dependent on improved understanding of the S-H20 and Fe-H20 systems. The development of the S-H20 diagram has alre a d y been d i s c u s s e d . G e o l o g i s t s were probably the f i r s t to c o n s t r u c t diagrams of t h e s t a b i l i t y o f i r o n and s u l f u r c o m p o u n d s . B a r n e s and K u l l e r u d5 8 p u b l i s h e d p 02_ pH d i a g r a m s f o r the anhydrous Fe-S-O system. G a r r e l s and C h r i s t6 0 o f f e r e d an E-pH diagram of i r o n o x i d e s , s u l f i d e s and carbonate at 25°C. Using r e s u l t s p u blished by Pourbaix and V a l e n s i f o r Fe-H^O and S-H20 r e s p e c t i v e l y , and i n f o r m a t i o n on i r o n s u l f i d e s , Bouet and B r e n e t6! c o n s t r u c t e d d i a g r a m s f o r the Fe-S-H^O syste m a t 25°C. Diagrams f o r both i r o n oxides and hydroxides were i n c l u d e d with FeS and FeS2 or with FeS and Fe2S3. Pourbaix's diagrams of S-H20 and F e - H2° were made a v a i l a b l e i n E n g l i s h i n 1966 with the t r a n s l a t i o n of h i s A t l a s 4 4 . Townsend6 2 grams of the Fe inc l u d e d e s t i m a t i o n s of e r r o r s i n h i s -H20 system at 60,- 100, 150 and 2 0 0°C 39 E-pH d i a -D i f f e r e n t f o r m s o f i r o n o x i d e s o r h y d r o x i d e s a r e t h e r m o -d y n a m i c a l l y s t a b l e i n v a r i o u s c o n d i t i o n s . The c h o i c e o f t h e a p p r o p r i a t e f o r m f o r use on t h e d i a g r a m has been an a r e a of u n c e r t a i n t y . L e w i s8 3 c l a i m e d t h a t Fe(0H)2 seemed t o be s t a b l e o n l y b e l o w 60°C, i f a t a l l . F e ( 0 H )2 w o u l d be o x i d i z e d t o Fe3C>4 v i a the a p p a r e n t l y i r r e v e r s i b l e S c h i k o r r r e a c t i o n : 3 F e ( 0 H )2 _ ^ Fe304 + 2H2O + H2 . . . (8) Fe30*4 seemed t o be t h e o n l y l o w e r o x i d e t h a t i s t r u l y s t a b l e b e t w e e n 25 and 350°C. I n L e w i s ' s a c c o u n t , Fe(0H)2+ was t h e most p r e v a l e n t Fe ( I I I ) i o n . M a c d o n a l d , S h i e r m a n and B u t l e r ^ u s e d C r i s s - C o b b l e e x t r a -p o l a t i o n t o c a l c u l a t e thermodynamic f u n c t i o n s f o r Fe-H^O t o 300°C but d i d not c o n s t r u c t p o t e n t i a l - p H d i a g r a m s . B i e r n a t and Robins4-* c o n s t r u c t e d a diagram f o r Fe-S-H20 at 25°C showing r e g i o n s of s t a b i l i t y d e c r e a s i n g i n n o b i l i t y i n the o r d e r F e ( 0 H )3, F e ( O H )2 and F e S2. R e g a r d i n g t h e Fe ( I I I ) s t a t e , F e ( 0 H )3 i s m e t a s t a b l e up to 90°C and FeOOH i s m e t a s t a b l e t o 130°C a f t e r w h i c h i t decomposes to F e203. Fe(OH)2 decomposes to Fe304 above 150°C v i a r e a c t i o n (8). R i c k a r d65 found t h a t FeS ( f e r r o u s s u l f i d e ) and m a c k i n a w i t e were f o r m e d by s u l f i d a t i o n of FeOOH ( g o e t h i t e ) . These o b s e r v -40 ations indicate that mackinawite and FeOOH should be included on the diagram at least at low temperatures, although the inclusion of mackinawite in the present work was precluded by the absence thermodynamic data. Peters5 0 presented an Fe-S-H20 diagram at 25°C showing FeS and FeS2 stability regions. Macdonald and Hyne31 studied the Fe-S-R^O system and con-structed E-pH diagrams by computer. Mackinawite and FeS2 were shown. They considered thermodynamic values for the complexed iron/hydrosulfide species (Fe(HS)2 and Fe(HS)3~ to be inaccurate and so these species could not be included. Taylor and Shoesmith6° have described a colloidal iron sul-fide species, NaFeS2, which was formed at pH 11-13 probably by HS- and Fe(0H)4~ ions. The solution had a dark green-black color. The NaFeS2 decomposed at 80°C, even in the presence of HS-, and p o l y s u l f i d e was formed. In contrast to t h i s , Raudsepp67 observed NaFeS2 at 100 and 200°C, and speculated that i t might be involved in a leaching process. Macdonald and Syrettes1 drew their E-pH diagram for Fe-S-H20 in terms of FeS ( t r o i l i t e ) , although they recognized that F"e(l + x)S and FeS(i + x) , pyrrhotite, exist and that Fe(i+ x)S might form at lower potentials than FeS. Conversion of FeS to FeS2 w a s considered to involve HS-. On the other hand, e q u i l i b r i a with Fe30£ and Fe203 were thought to involve sulfur oxyanions, SO^ . 41 New thermodynamic data were used by Silverman^ to construct a new Fe-H20 diagram at 25°C. A region of s o l i d Fe(0H)2 w a s found to l i e between Fe and Fe304. The stability region of Fe304 was much larger than that of Fe(0H)2- T h e present treatment neglects F e (0 H) 2 at 25°C for the sake of s i m p l i c i t y , bearing in mind that high temperature behavior i s of greater interest in this work. The inclusion of Fe(0H)3- was favored by Silverman instead of HFe02-. Tuff r e y7^ experienced d i f f i c u l t y in producing FeO-42- and found that i t is less stable at higher temperatures. For this reason i t has been omitted in the diagrams drawn here. The diagrams at 25, 100 and 150°C have been constructed here to reflect changes in the S-H20 diagrams plus new thermodynamic d a t a3 2, 69. The rmodynamic Data The principal source of thermodynamic data for the Fe-S-H20 diagram was the Handbook of Thermodynamic Data^. Biernat and Robins4-* and Silverman^ were the sources of the balance of the data. The heat capacity of 3-FeS at 150°C was assumed to be the same as for a-FeS at the same temperature. Heat capacities of the ions at 100 and 150°C were estimated by the Criss-Cobble method, as described in Appendix I. 42 Equilibrium Equations High temperature e q u i l i b r i a were c a l c u l a t e d for the reactions l i s t e d in Table IV. The a c t i v i t i e s of reactants were approximated and substituted into the equilibrium equations. A c t i v i t i e s of most ions were taken as 10~6. A c t i v i t i e s of 10- 3 and 10-1 were used for Sx2- and HS- respectively, after section 2.2. Standard electrode potentials of electrochemical equilibria and free energy changes of chemical equilibria at 25, 100, and 150°C have been tabulated in Appendix III. Electrochemical potentials calculated from the Nernst equil-ibrium equations were plotted versus the standard hydrogen electrode at 25°C (SHE) as a function of pH at 25, 100 and 150°C in Figures 11, 12 and 13. The presence of two lines relating FeS2 to Fe304 results because one is with respect to HS- and the other is with respect to S 2 O 3 2 - (i.e. e q u i l i b r i a 2.30 and 2.31 respectively). 43 TABLE IV Electrochemical Equilibria for the Fe-S-H2Q System 2.1 02 + 4 H+ + 4e" ^  2 H20 2.2 2 H+ + 2e" ^  H2 2.3 FeS + H+ + 2e" Fe + HS~ 2.4 2 Fe + S2032" + 6 H+ + 4e" ^ 2 FeS + 3 H20 2.5 Fe304 + 8 H+ + 8e" ^ 3Fe + 4 H2° 2.6 Fe(0H)2sol + '2 H+ + 2e_ ^ Fe + 2 H20 2.7 Fe2 + + 2e~ ^ Fe 2.8 FeOH+ +H+ + 2e" ^  Fe + H20 2.9 Fe(OH)2dis + 2 H+ + 2e" ^  Fe + 2 H20 2.10 Fe(0H)3- + 3 H+ + 2e" ^  Fe + 3 H20 2.11 FeS2 + H+ + 2e" FeS + HS" 2.12 2 FeS + S2032" + 6 H+ + 4e" ^  2 FeS2 + 3 H20 2.13 FeOOH + HS" + 2 H+ + e~ ^ FeS + 2 H20 2.14 2 FeOOH + S2032" + 12 H+ + 10e~ ^ 2 FeS + 7 H20 2.15 Fe203 + S2032~ + 12 H+ + 10e~ 2 FeS + 6 H20 2.16 Fe203 + 2 HS~ + 4 H+ + 2e" 2 FeS + 3 H20 2.17 Fe304 + 3 HS" + 5 H+ + 2e" ^  3 FeS + 4 H2° 2.18 Fe(0H)2 sol + HS" + H+ =^ FeS + 2 H20 2.19 Fe2 + + HS~ ?=*FeS + H+ 2.20 2 Fe2+ + S2032~ + 6 H+ + 6e~ 2 FeS + 3 H20 2.21 FeOH+ + HS" ^ FeS + H20 2.22 2FeOH+ + S2032- + 8- H+ + 8e" ^  2 FeS + 5 H20 2.23 Fe(OH)2 dis + HS~ + H+ ^ FeS + 2 H20 2.24 Fe(0H) 3" + HS" + 2 H+ ^ FeS + 3 H20 2.25 2 Fe(OH)4" + S2032" + 11 H+ + 7e~ ^ 2 FeS + 11 H20 44 TABLE IV (Cont.) 2.26 FeS2 '+ 2 H20 + e- FeOOH + 2 HS~ + H+ 2.27 FeOOH + S 2 0 3 2 - + 9 H+ + 7e" ^  FeS2 + 5 H20 2.28 Fe203 + 2 S 2 0 3 2 - + 18 H+ + I4e~ ^ »'2FeS2 + 9 H2° 2.29 2 FeS2 + 3 H2O + 2e" *9 Fe203 + 4 HS~ + 2 H+ 2.30 3 FeS2 + 4 H20 + 4e" -a Fe304 + 6 HS~ + 2 H+ 2.31 Fe3o4 + 3 S2032 _ + 26 H+ + 20e" ^ » 3 FeS2 + 13 H20 2.32 FeS2 + 2 H2O + 2e~ Fe (0H) 2 s o 1 + 2 H S ~ 2.33 FeS2 + 2 H+ + 2e~ Fe2 + + 2 HS" 2.34 Fe2 + + S 2 O 3 2 - + 6 H+ + 6e~ FeS2 + 3 H 2 0 2.35 FeS2 + H2O + H+ + 2e~ Fe0H+ + 2 HS" 2.36 FeOH+ + S 2 O 3 2 - + 7 H+ + 6e~ ^  FeS2 + 4 H 2 0 2.37 FeS2 + 2 H2O + 2e~ ^» Fe (0H) 2 dis + 2 HS-2.38 FeS2 + 3 H20 + 2e" Fe(OH)3;_ + 2 HS" + H+ 2.39 Fe(OH)4- + S 2 0 3 2 ~ + 10 H+ + 7e~ ^  FeS2 + 7 H20 2.40 3 FeOOH + H+ + e~ Fe304 + 2 H 2 0 2.41 3 Fe 2 0 3 + 2 H+ + 2e" ^ 2 Fe304 + H 2 ° 2.42 FeOOH + 3 H+ + e~ ^ * Fe2+ + 2 H20 2.43 Fe 2 0 3 + 6 H+ + 2e" ^ » 2 Fe2+ + 3 H20 2.44 FeOOH + 2 H+ + e~ ^ » FeOH+ + H20 2.45 Fe 2 0 3 + 4 H+ + 2 e - ^  2 FeOH+ + H20 2.46 FeOOH + H+ + e~ ««» Fe ( 0 H ) 2 d i s 2.47 Fe ( 0 H ) 4 - + H+ FeOOH + 2H 2 0 2.48 2 Fe (0H)4 _ + 2 H+ ^ * Fe 2 0 3 + 5 H20 2.49 Fe 3 0 4 + 2 H20 + 2 H+ + 2 e ' ^ 3 Fe (0H) 2 sol 2.50 Fe304 + 8 H+ + 2e" ^ ft 3 Fe2+ + 4 H2° 2.51 Fe3o4 + 2 H20 + 2 H+ + 2e~ ^  3 Fe(OH)2 dis 2.52 Fe304 + 5 H20 + 2e~ 3 Fe(OH)3~ + H+ 45 TABLE IV (Cont.) 2.53 3 Fe(0H)4~ + 4 H+ + e" Fe304 + 8 H 2 0 2.54 F e 2 + + 2 H20 ^  Fe(OH)2 sol + 2 H + 2.55 FeOH+ + H2O Fe(0H)2 s o 1 + H + 2.56 Fe(0H)2 d i s Fe(0H)2 sol 2.57 Fe ( 0 H ) 3 _ + H+ ^ Fe(OH)2 sol + H20 2.58 FeOH+ + H+ ^ = Fe2+ + H2o 2.59 Fe(0H)2 d i s + H + ^ FeOH+ + H 2 ° 2.60 Fe(0H)4" + 3 H+ + e~ ^  Fe0H+ + 3 H20 2.61 Fe ( 0 H ) 3 " + H+ ^= Fe (0H) 2 dis + H20 2.62 Fe ( 0 H ) 4 " + 2 H+ + e" ^  Fe ( 0 H ) 2 dis + 2 H20 2.63 Fe ( 0 H ) 4 - + H+ + e~ ^ * Fe ( 0 H ) 3 _ + H20 2.64 Fe(HS) 2 + 2e~ ^  Fe + 2 HS" 2.65 FeOOH + 2 HS" + 3 H+ + e~ ^  Fe(HS) 2 + 2 H20 2.66 FeS + HS~ + H+ Fe(HS) 2 2.67 Fe(HS)3" + 2e~ Fe + 3 HS" 2.68 FeOOH + 3 HS" + 3 H+ + e~ ^  Fe(HS) 3- + 2 H20 2.69 FeS + 2 HS" + H+ ^  Fe(HS) 3~ 2.70 Fe(HS) 2 + HS~ ^  Fe(HS) 3-2.71 FeOOH + H 2 0 + e " ^ = Fe ( 0 H ) 3 ~ 2.72 Fe 20 3 + 3 H20 + 2 e _ ^  2 Fe ( 0 H ) 3 -2.73 2 Fe 30 4 + 3 S 4 2 - + 16 H+ + 10e" ^  6 FeS 2 + 8 H20 2.74 2 FeOOH + S 4 2" + 6 H+ + 4e" ^ * 2 FeS 2 + 4 H 2 0 2.75 Fe 20 3 + S 4 2 _ + 6 H+ + 4e" ^  2 FeS 2 + 3 H20 2.76 3 FeOOH + 2S 3 2" + 9 H+ + 5e" =^  3 FeS 2 + 6 H20 2.77 3 F e 2 0 3 + 4 S 3 2 _ + 18 H+ + 10e" 6 FeS2 + 9 H2O 2.78 FeOOH + S22~ + 3 H+ + e~ ^  FeS2 + 2 H20 2.79 Fe 20 3 + 2 S 2 2 ~ + 6 H+ + 2e" ^  2 FeS 2 + 3 H 2 0 46 TABLE IV (Contd) 2.80 FeS2 + 4 H20 + e" ^ Fe(0H)4~ + 2 HS" + 2 H+ 2.81 Fe(0H)3- + 2 H+ ^ FeOH+ + 2 H20 2.82 F e ( 0 H )4" + 4 H+ + e" ^ F e2 + + 4 H20 2.83 F e ( 0 H )4" + 2S + 4 H+ + 3e~ ^ FeS2•+ 4 H20 47 48 F i g u r e 12: F e - S - H 2 0 a t 1 0 0 ° C E q u i l i b r i a 2 .2 , 2 . 3 , 2 .10 , 2 . 1 1 , 2.17* 2 .24 , 2 .27 , 2 .30 , 2 . 3 1 , 2 . 4 0 , 2 .47 , 2 . 5 2 , 2 .53 , 2 .60 , 2 .63 , 2 .81 . 49 50 General Observations The Figures 11-13 for the Fe-S-H20 system differ s i g n i f i c -antly from the figures constructed by others*5* 68. There was no stability region for S2 -, as explained in section 2.2. Fe304 and FeOOH were considered at 25°C instead of Fe(0H)2 and Fe(0H)3» An interesting feature of the present diagrams i s that the Fe3o4 stability region embraces that of FeS2 a t H-13. The present diagrams include Fe(0H)4~ and HFe02~ i s replaced with F e ( 0 H ) 3-«D^ Table V summarizes the calculated potentials for e q u i l i b r i a of Fe and Fe304 with both HFe02~ a n d Fe(0H)3-. They were calculated according to the Criss-Cobble theory, assuming for the calculation of Cp that both ions were in the "acid oxyanion" class. The difference in potentials of e q u i l i b r i a with HFe02- and Fe(0H)3- i s due to their different free energies, AG0. In the calculation, the activities of HFe02-and Fe(0H)3_ were approximated as IO- 6. The pH of 3m [OH-] solution at each temperature71 was employed in the calculation. The effect of considering Fe(0H)3_ instead of HFe02~ becomes more important as temperature is increased and is most apparent in the equilibrium with Fe304« The gradients of the equilibrium lines are the same for the two ions. E q u i l i b r i a 2.10 and 2.52 were used to calculate potentials for reactions involving Fe(0H)3~. The equilibria involving HFe02- were: HFe02_ + 3 H+ + 2e" ^  Fe + H20 ... (9 ) Fe304 + 2 H20 + 2e~ ^  3 HFe02~ + H+ ... (10) 51 Townsend's"2 AG0 and S° values for HFeC>2- were used to calculate the standard potentials. Tromans7 2 has reviewed the entropy values for HFeO^- and found general agreement on the value used ji e r e62-64> bu t considerable discrepancy with the value used by Biernat and Robins4-* in constructing their E-pH diagrams. The difference, A, between the e q u i l i b r i a involving Fe and Fe3C>4 (Table V) provides a r e l a t i v e comparison of the size of the stability range of the dissolved iron in 3m NaOH. The stability range of HFe02~ increases with temperature, but that of Fe(0H)3 increases and then decreases. The difference in the s t a b i l i t y range for these two species arises from the difference in thermo-dynamic values for the two species. 52 TABLE V Comparison of Calculated Potentials for Equilibria Involving HFe09 * and Fe(OH)^ * in 3m NaOH Equilibrium 25°C E A 100OC E A 150OC E A Fe(OH)3-/Fe -0.994 0.179 -1.101 0.262 -1.146 0.214 Fe304/Fe(0H)3- -0.815 -0.839 -0.932 HFe02~/Fe -0.971 0.086 -1.108 0.291 -1.175 0.336 Fe304/HFe02~ -0.885 -0.817 -0.839 * - based on 1 0-° m soluble iron species A = EFe304/HFe02- " EHFe02-/Fe or A= EFe304/Fe(0H)3_ " EFe(0H)3"/Fe 53 3. REFERENCE ELECTRODES 3.1 Introduction A r e f e r e n c e e l e c t r o d e s e r v e s a s a s t a b l e r e f e r e n c e p o i n t w i t h r e s p e c t to w h i c h the p o t e n t i a l of a w o r k i n g e l e c t r o d e may be m e a s u r e d . The r e f e r e n c e e l e c t r o d e may be l o c a t e d a t t h e same t e m p e r a t u r e and p r e s s u r e a s t h e w o r k i n g e l e c t r o d e ; s u c h an a r r a n g e m e n t i s t e r m e d an i n t e r n a l r e f e r e n c e e l e c t r o d e . A l t e r n a t -ively, a reference e l e c t r o d e situated at a different temperature and pressure i s an external reference electrode. A conduction path must be preserved between the reference and working elec-trodes in both cases. The solution surrounding the external or internal reference electrodes may be different from that in which the working elec-trode i s located. If the solution i s d i f f e r e n t , a l i q u i d junction potential i s created at the interface between the two elect r o l y t e s due to the differences in mobil i t i e s of different ions on either side of t h e interface. Invariably the differences between positive and negative ions do not cancel and therefore a small current should flow. In the absence of current flow, a potential difference develops to equalize the positive and neg-ative charge flow. This potential difference can be estimated by the Henderson equation7 3: (11) 54 where U j ±s ionic mobility, Cj is concentration in equivalents, and zj 1 s the ionic valency. R, T and F have their usual mean-ings. The integers 1 and 2 refer to the two sides of the junction. Equivalent conductance, A, cm2£2_ 1, may be substituted into equation 11 in place of u7^. The sign of the equation is consistent with the IUPAC convention for electrochemical values. A reduction potential results. When an external reference electrode i s used, another potential drop arises in the bridge between the working and reference electrodes due to any temperature difference between them, which influences ionic m o b i l i t i e s to d i f f e r i n g extents. This i s termed a thermal l i q u i d junction potential (TLJP). The cold junction assumes the charge of the fastest i o n 7 5. in some solutions, KC1 for example, the mobilities of the K+ and C l- ions are closely matched at a l l temperatures and for small temperature differences the potential difference is sraa 1175, 76 but when the difference in temperature is large the TLJP becomes very s i g n i f i -cant7 7. Over an extended period of time, the temperature difference can lead to the formation of a concentration gradient along the bridge. This i s the Soret e f f e c t7 8. There i s evidence that i t is unimportant in experiments of duration similar to those per-formed in the present study78* 79. Another potential difference may arise when there i s a 55 pressure drop between the internal and external reference elec-80 trode compartments. This is a streaming potential. Cellulose or asbestos plugs8 1* 8 2 have been used to reduce the flow due to a pressure difference. Some experimentalists minimize the stream-ing potential and flow by applying a back-pressure to the reference electrode compartment83. To avoid streaming potential altogether the reference electrode may be at room temperature but at the same pressure as the working electrode; this has been done in the present investigation. Measured potentials should (LJP), thermal l i q u i d junction according to the demands of the be corrected for liquid junction (TLJP) and streaming potential experimental set-up. Reference electrodes should be reproducible and stable, that is they should have a constant half-cell potential with respect to a thermodynamic scale such as the hydrogen electrode. It i s desirable that reference electrodes be reversible, that is they should not be polarized by the very small currents involved in measuring the potential of another electrode. This implies a high exchange current density on the electrode. If the electrode is to be used for thermodynamic purposes, i t should behave in a thermodynamic manner. It should obey Nernst's equation, and the electrode potential should be related to thermodynamic chemical potentials7 5. If the electrode is not thermodynamic i t may s t i l l be useful for practical engineering studies provided i t is cheap and rugged and maintains a stable potential. 56 This present investigation demanded reference electrodes which were compatible with highly alkaline and sulfide-containing solutions. A number of r e f e r e n c e e l e c t r o d e s have been developed f o r use i n c a u s t i c s o l u t i o n s including the hydrogen e l e c t r o d e ^ and a gold -2% palladium electrode8 5' 8 o. The most popular choice f o r work i n c a u s t i c solutions has been the Hg/HgO electrode8 5* 87-91# Every and Banks89 found the EMF of the electrode with respect to the calomel electrode at 25°C was unchanged by temperature in 20% NaOH, and decreased only s l i g h t l y as NaOH was increased from 5 to 40%. Case and Bigno l d90, 91 found the Hg/HgO electrode to be r e l i a b l e and predictable up to 250°C in solutions of 5 M NaOH. The electrode was poisoned i f Na2S was added to the NaOH solution8 8, presumably by the formation of mercuric sulfide. The popular silver-silver chloride electrode, Ag/AgCl, be-haved in NaOH + NaCl s o l u t i o n as though i t was an Ag/AgO electrode, but the AgO decomposed at high temperature making i t unacceptable for use in high temperature alkaline s o l u t i o n s9 0. Others used the electrode but located i t in another electrolyte with a liquid junction92, 93# D a n i e l s o n7 9 d e s i g n e d a v e r y r u g g e d A g / A g C l electrode for use in geothermal fluids containing H2S r e f e r e n c e and H2- A n 57 impermeable barrier of zirconia protected the reference electrode from poisoning. There was a porous-zirconia plug at the tip of the electrode. A more desirable approach to the problem of sulfide poison-ing would be to develop an electrode which would not be poisoned. Sulfide electrodes are the most obvious candidates. Silver sulfide was a l i k e l y choice for further investig-« ation. It was considered by Ive s7 4 to be one of the most promising sulfide electrodes. The Ag/Ag?S Electrode Due to the p r a c t i c a l l y complete h y d r o l y s i s of S2 - in alkaline solutions (section 1.2) the silver-silver sulfide elec-trode reaction is expressed as: Ag2S + H20 + 2e" ^  2 Ag + OH" + HS- ... (12) where the electrode potential at 25°C, based on the best thermo-dynamic data available, i s : E = -0.684 - 0.0295 log [0H-][HS-], VS H E ... (12A) Standard potentials available in the literature for this reaction are listed in Table VI (column 3). 58 TABLE VI Standard Electrode Potential of Ag/Af^S Reported Standard Potential Reference Standard Potential Standard Potential E°> VSHE E ° 1 2 ' VSHE E°14' VSHE -0.7125 94 -0.683 -0.266 -0.7135 95 -0.681 -0.264 -0.691 97 -0.691 -0.274 -0.598* 98 -0.629 -0.212 -0.648 99 -0.648 -0.231 -0.660 100 -0.660 -0.243 -0.684 101 -0.684 -0.267 -0.692 102 -0.692 -0.275 •Measured electrode potential in 1 M NaOH + 0.10 M Na2S. 59 for E a r l i e r workers^" the reaction: quoted standard electrode potentials Ag2s + 2e~ # 2 Ag + S2~ ... (13) They interpreted their measured values in terms of the sulfide species, HS- or S2-, which were expected to predominate at the experimental pH. I f they considered HS- to be the predominant species, their quoted values must be related to the standard potential values for reaction 12 by means of the K2 values they used to calculate their quoted values. The values l i s t e d in Table VI from references 94 and 95 have been corrected in this manner. The reported values from references 98 and 100, l i s t e d in Table VI, must be considered consistent with reaction 12. It should be noted that scatter is reduced when potentials listed in Table VI are calculated on the basis of the same equilibrium equation. Reaction 12 may be rewritten alternately as: Ag2S + H+ + 2e" ^  2 Ag + HS" ... (14) where E = -0.267 - 0.0295 log [HS"]/[H+], Vg H E ... (HA) 60 Equations (12A) stant for water, and (14A) Kw via: are related by the dissociation con-E °1 2' - 2.303RT log Kw = E°i4 ... (15) 2F Published values of standard p o t e n t i a l have been converted to this equilibrium, and then l i s t e d in column 4 of Table VI. Ag/A g2S forms a reversible electrode of the second kin that is a porous oxidized phase (Ag2s) on a metal substrate (Ag), which i s generally accepted as behaving in accordance with the Nernst equation. Golding97f however, claimed that the electrode was not reversible. Ag2S i s thermodynamically more stable than Ag20 or AgCl as discussed more f u l l y in Appendix IV, where i t i s shown that the concentration of sulfur species must be many orders of magnitude lower than that of [CI-] or [OH-] in order for AgCl or Ag20 to form. This fact should preclude formation of these species on the silver electrode in the aqueous sulfide solutions encountered in the present work. Hseu and Rechnitz9 6 calibrated the silver sulfide electrode with sulfide concentration as a function of ionic strength, and with pH as a function of Na2S concentration of constant ionic strength. They found that a variety of anions including CI-, SO^2-, and CrO^2 - did not affect the measured potential at pH 12 and 0.1 M ionic strength. 61 In a study relevant to wood pulping, Swartz and Light98 observed that the quality of t i t r a t i o n curves obtained with a silver sulfide ion-selective electrode deteriorated in proportion to the quantity of organosulfur compounds in the solution. They also observed errors caused by polysulfides and aromatic poly-hydroxy compounds. Teder1 0 3 observed that an Ag/Ag2S electrode could not be used to determine the a c t i v i t y of free sulfi d e ions in polysulfide solutions. In polysulfide-free solutions, the electrode behaved as expected from theory, but diverged from theory when poly-sulfid e was added. The malfunction was ascribed to probable formation of silver polysulfides. The Hg/HgS Electrode Ives'^ suggested that the mercury-mercuric s u l f i d e , Hg/HgS electrode was worthy of further consideration. Using data of Goates et a l - ^4, the standard potential may be calculated as - 0 . 6 6 1 Vg^E f o r t h e r e a c ti o n HgS + H2o + 2e" Hg + HS" + 0H~ ... (16) Unfortunately, HgS is soluble in moderate sulfid e concen-trations^) 54. it forms a soluble sulfide complex, HgS22 -, at 62 pH >9.1°5» 1°° Somewhat in contrast, Armstrong et a l1 0 7 measured the potential of HgS-black as -0.692 V"sHE a t P H 9*5 , In a 1 M Na2S solution, anodic dissolution was identified. Their results indicated that HgS was deposited on the electrode as two mono-molecular layers. These findings seemed to support the idea of i n o an H g / H g S reference electrode. S i m i l a r l y , Szebenyi et al performed polarization tests at pH 10.5 and obtained results consistent with mercuric sulfide formation. They claimed H g / H g S / H S ~ predominated at the free corrosion potential at pH 10.5. Preliminary experiments confirmed that a black sulfide formed at weakly alkaline pH and disappeared at higher pH (>11). Further investigation at high pH was attempted but on the assumption that the electrode was no longer Hg/HgS. The Polysulfide Electrode The electrode potential of inert electrodes in sulfide sol-utions i s consistent with r e a c t i o n s between s u l f u r and polysulfides in solution*1* 109. (17) (18) (19) (20) S + 2e" ^ S2-s22" + 2 e" ^ 2S2~ • • • S32- + 2e- ^ S2" + s22- • • • S42- + 2e~ ^ S2- + S32- • • • 63 Allen and Hickling*0 9 published some of the earliest work on the sulfur-sulfide electrode. Reproducible potentials resulted when inert electrodes including platinum, platinized platinum, graphite, nickel and gold were immersed in alkaline aqueous solutions of sulfur and sodium sulfide at 18°C. A subsequent a r t i c l e1^ indicated that anodic polarization of the electrodes caused the formation of polysulfide in solution. These workers added sulfur to sulfide solutions and then related equilibrium potential with sulfur/sulfide ratio via: E = Eo + RT in _ _ I Z S J _ . . . (21) 2F [ZNa 2 S] 2 where E° = -0.522 VS H E ZS i s moles of sulfur added per l i t e r of sodium sulfide solution. Their results suggested that total dissolved sulfur, sulfide and bisulfide control the electrode behavior in dilute solutions. The potential in more concentrated solutions may be given by: E = Eo + RT In ([ZS] -2 [S32~] -3 [S42-]) ... (22) 2F ([ZNa2S] - [ES] + [S32-] +2 [ S42" ] )2 Maronny27* 5 7 used platinum electrodes to obtain elec-trode potentials for polysulfide equilibria. 64 Tederi U J also used platinum to obtain redox potentials of polysulfide solutions. The potentials depended on the ratio, Xg» of polysulfide excess sulfur to sulfide s u l f u r , but decreased only slightly with increasing ionic strength. Polysulfide excess s u l f u r and s u l f i d e s u l f u r are r e l a t e d by: (x-l)S + S2" ^ Sx2 _ ... (23) where ( x - l ) S i s p o l y s u l f i d e excess s u l f u r S2" is sulfide sulfur The redox potential decreased linearly with increasing logarithms of the OH- concentration and excess sulfur concentration. Temp-erature had only a moderate effect on the redox p o t e n t i a l , with a greater effect at higher pH. An empirical equation for the redox potential, based on the above variables, was offered. 65 3.2 Experimental 3.2.1 Ag/Ag9S Electrode The object of the experiments was to determine the su i t -a b i l i t y of the electrode as a reference electrode in alkaline sulfide solutions. The potentials of the electrode as a function of temperature and solution composition were determined and the electrode behaviour interpreted. Silver sulfide electrodes were produced by anodizing silver wires at 0.5V versus the saturated calomel electrode (SCE) for 30 s in a 1 M Na2S solution. They were stored in the same sol-ution. The p r i n c i p a l test solution was 2.5m NaOH + 0.423 m Na2S, a composition that has been used to simulate kraft white liquor. Also, NaOH and Na2S concentrations were varied: A 2.5m NaOH + 0.5m Na2S + lm NaCl solution was chosen for study of the effect of the upper l i m i t of chloride contamination anticipated in a coastal pulp m i l l . Test solutions were prepared by boiling d i s t i l l e d water, nitrogen purging, and adding desired quantities of analytical reagent grade NaOH, Na2S»9H20 and NaCl. The f i r s t series of experiments, to ~95°C, was performed in a Teflon c e l l . Nitrogen purging of the solution was maintained throughout each test. The potential of the Ag/Ag2s electrode was measured with respect to a room temperature SCE via a Teflon 66 Luggin capillary and saturated KC1 salt bridge. A cotton thread passing through the bridge minimized the effect of vapor bubble formation. Measured results were recorded for at least two separate runs through the temperature range. The typical rate of temperature increase was 4°C/h but slower rates were used where the electrode was slow to stabilize. The reported potentials were referred to the standard hydrogen electrode at 25°C (SHE) where: VSHE = VSCE + °«2416V ••• (2*) The second series of experiments was conducted in an auto-clave (1 l i t e r PARR, Monel with Teflon l i n e r ) to 170°C. Potential measurements were made between an internal Ag/Ag2S electrode at test temperature and an external Ag/Ag2S electrode at 25°C in a common test solution. The design of the external reference electrode is similar to one described by Danielson7 9, and is illustrated in Figure 14 and 15. The silver electrode was cast from 99.99% pure material and the silver sulfide was formed as previously described. The external electrode compartment was connected to the autoclave by a 1/4 inch (6.3 mm) stainless steel tube through which a Teflon bridge tube passed. A cotton thread was passed through the inside of the Teflon bridge tube to pre-serve an e l e c t r i c a l path i f bubbles formed. The cotton thread was stable throughout the period of the tests. A porous zirconia plug was f i t t e d onto the end of the Teflon bridge tube inside the 67 y—Conax F i t t ing P 6 2 - 2 5 0 - A T U /—Autoclave C o n n e c t i o n C 3 -Te f lon Tube Porous Z i rconia P lug E x t e r n a l R e f e r e n c e E l e c t r o d e Figure 14: External Reference Electrode and Bridge 68 Tube 304 SS 1/4" x 0.035" Gyroloc Fitting 1/4"T-1/4"NPT Teflon Seal Conax Fitting 1/4*" NPT E6T-125A Teflon Tube TT250 sizelO Silver Rod Figure 15: Detail Cross-section of External Reference Electrode 69 autoclave. The plug was boiled and cooled several times to expel trapped air and then stored in the solution***. The bridge and electrode compartment were f i l l e d with s o l -ution by means of a syringe and Teflon intramedic tubing. The external reference electrode operated at the same pressure as the autoclave. This technique eliminated isothermal liquid junction potentials (LJP) and streaming potentials and allowed an assess-ment of the re l a t i v e magnitude of thermal l i q u i d junction potentials (TLJP). The internal electrode was a length of silver wire attached to a Teflon sheathed 316 stainless steel conductor rod with the connection mounted in epoxy. Sufficient epoxy was used to ensure a seal even though some surface softening occurred after periods at high temperature. The internal electrode conductor, the bridge to the external reference electrode, and a graphite counter electrode penetrated the autoclave head through Conax seals. Potential measurements were obtained with a Corning pH meter (model 125) or via a Princeton Applied Research potentiostat (model 173). Sulfide analyses were performed by t i t r a t i o n with 0.1 N AgN03j using a commercial Ag2S electrode. 70 3.2.2 Polysulfide Electrode Although the polysulfide i s a mixed electrode and thus cannot be considered "thermodynamic", i t may serve as a very useful electrode for practical systems. An attempt to relate i t s behavior to electrode reactions has been made here to confirm that the electrode behavior of a noble metal (inert) electrode is due to polysulfides. Platinum was used as an inert base for the electrode. It was in the form of wires, 0.5 mm diameter. Mercury was tested to ascertain whether i t behaved as a noble metal electrode. The electrode was a mercury reservoir with an insulated platinum wire lead. The test solution had a composition 2.5 m NaOH + 0.423 m Na28. It was prepared as described in section 3.2.1. Experiments were conducted only at atmospheric pressure to ~95°C, in a Teflon c e l l , with nitrogen purging. The potential of the Hg or Pt/Sx2 -> po l y s u l f i d e , electrode was measured with respect to a room temperature saturated calomel reference elec-trode (SCE) via a Teflon Luggin capillary and salt bridge f i l l e d with saturated KC1 solution. Potential measurements were obtained with a Corning pH meter (model 125). 71 3.3 Results and Discussion 3.3.1 Silver-Silver Sulfide Electrode The measured potential values, E, of the s i l v e r - s i l v e r sul-fide electrode, SSSE, were plotted for a range of temperatures in a variety of solutions (Figures 16-21). Corrections for l i q u i d junction potentials (LJP) between the test solution and the saturated KC1 bridge were calculated by means of Henderson's equation (eq. 11, section 3.1). Thermal l i q u i d junction poten-t i a l s (TLJP) for the saturated KC1 bridge were approximated from the results of Macdonald et a l ^ > 7<3. These corrections are included in the figures. The Soret effect was considered to be negligible. For comparison, the theoretical electrode potentials were calculated based on reaction [14], and plotted for comparison. E° values for reaction 14 at temperature T, E ° ^ , were obtained using the method described by Criss and Cobble^. The thermo-dynamic data for Ag and Ag2S were taken from Barin and Knacke1 1 2. AG0 and S°298 of H+ were obtained from Criss and Cobble-5^, a n cj CpT was calculated using a formula given by Taylor**3. The A G° and S° values of HS- have been given by Biernat and Robins.4 o The Cp|^298 values were approximated by the Criss-Cobble method^0,, 72 •••• •• E E*UP E*LJP*TUP 20 30 40 50 60 70 80 90 100 TEMPERATURE ,°C SSSE potential as a function of temperature for SCE/KC1(sat.)// 2.5mNaOH+ 0.423m Na2S/Ag2s/A8 73 -0.650 -0.660 -0.670 x g -0.680 •0.690 o CL -0.700 -0.710 -0.720 -0.730 • . . . • * . • ••••» • ** • .. ... • •• • • • « .. ... . . . . . . i , . ....... «•.. !• • • • .... ......... • J ». v.v E — E « U P — E*LJP*TUP — ET 20 30 40 50 60 70 80 90 100 TEMPERATURE ,°C Figure 17. SSSE potential as a function of temperature for SCE/KC1 (sat.) // 2.5 m NaOH + 0.5m Na2S/Ag2S/Ag. 74 Ul I C O > U J L U I— o •0.670 -0.680 -0.690 -0.700 0.710 0.720 -0.730 -0.740 -0.750 1 • •• • • • •• • • • • •• E - E*LJP "E*LJP*TLJP - E T 20 30 40 50 60 70 80 90 100 TEMPERATURE ,°C Figure 18. SSSE potential as a function of temperature for SCE/KC1 (sat.) // 2.5 m NaOH + lm Na2S/Ag2s/A8-75 0.610 -0.620 •0.630 LU -0.6AO L J •a650 o CL LU -0.660 -0.670 -0.680 -0.690 V . V £ — E*LJP —•E*LJP*TLJP — ET 20 30 AO 50 60 70 80 90 100 TEMPERATURE ,°C F i g u r e 1 9 . S S S E p o t e n t i a l a s a f u n c t i o n o f t e m p e r a t u r e f o r S C E / K C 1 / / 2.5 m NaOH + 0 .05m N a 2 S / A g 2 S / A g . 76 -0.690 -0.700 •0.710 5 - ° -7 2 0 > UJ ~. -0.730 LU r— o Q. •••• «... -0.740 -0.750 0.760 -"0.770F ^ E + L J p •— E*LJP*TLJP 0.780 h — £ T 40 50 60 70 80 90 TEMPERATURE ,°C 100 F i g u r e 2 0 . S S S E p o t e n t i a l a s a f u n c t i o n o f t e m p e r a t u r e f o r S C E / K C 1 / / 12m NaOH + 0 .423m N a 2 S / A g 2 s / A 8 -77 T 1 1 -r 1 1 ' ' r -0.660 -0.670 LO -0.680 LJ 5 -0.690 5 -0.700 -0.710 -0.720 -0.730 o OL * * * • • • E — E*LJP — E+LJP*TUP — ET - I L . _ l I 1 L. 20 30 AO 50 60 70 80 90 100 TEMPERATURE ,°C F i g u r e 2 1 . S S S E p o t e n t i a l a s a f u n c t i o n o f t e m p e r a t u r e f o r S C E / K C 1 / / 2 . 5 m N a O H + 0 . 5 m N a 2 s + 1 m N a C l / A g 2 S / A g . , 78 The calculated values of E°rp are listed in Table VII. These values were substituted in the Nernst equation for reaction 14 at temperature T, to give E-j.; ET = E °T - 2.303 RT pH -2.303 RT log [HS-] ... (25) h~F RT Calculated values of pH and [HS-] were required to obtain E.p f0r each solution and temperature. The HS- concentration was calculated as follows. Values of K2 at temperature for reaction 3 were extrapolated from data given by Giggenbach3 4. These values were corrected for ionic strength according to Hseu and Rechnitz^0 who converted the ionization constant, K2, to an ionization constant, K21, i n terms of concentrations of any given ionic strength, I, by using the extended Debye-Huckel equation: pK2I = pK2 _ vT . .. (26) 1+/I Following their method, [HS-] was related to total concen-tration of sulfide, [ S2~ ] f v i a [HS"] = [ S2- ]t ... (27) 1 + K 2 I / [ H . + . ] 79 TABLE VII Standard Electrode Potentials at Temperature  for Electrode Reaction (14) T E ° T 298 -0.2673 3 0 3 - 0 .2 6 6 2 313 -0.2644 323 -0.2633 333 -0.2627 343 -0.2629 353 -0.2637 363 -0.2652 373 -0.2674 383 -0.2701 393 -0.2737 403 -0.2780 413 -0.2826 423 -0.2884 433 -0.2947 443 -0.3018 453 -0.3093 80 The pH was obtained from Macdonald and McKubre's tabulation of pH versus [OH-].7 1 The [OH-] was estimated to be the sum of hydroxyl concentration from NaOH additions and the complete dissociation of S2 _ to HS-, equation 4. The concentration of H+ for use in equation 27 was approximated to 10~P^. The calculation of [HS-] confirmed that sulfide was present mainly as HS". The values of [HS-] and pH were substituted into equation 25 to obtain for each solution and temperature. The E-j values have been plotted together with the measured potent i a l , E, and the measured potential corrected for LJP and TLJP, E + LJP + TLJP, in Figures 16-21. The experimental data obtained in 2.5m NaOH + 0.423m Na2S, Figure 16, showed r e l a t i v e l y good agreement between the theoretical electrode potentials and the LJP + TLJP corrected measurements. For example, the difference i s 12 mV at 25°C and 10 mV at 100°C. The observed difference may be attributed prin-c i p a l l y to uncertainty in thermodynamic data on which the calculations were based and the assumption that the activity of HS- was the same as the concentration. Activity c o e f f i c i e n t s , ^ , l e s s than unity for HS- would produce better agreement. K i e l l a n d1 1* has l i s t e d a value of ^ HS- = °*3 8 i n °«2m ionic strength solution. No values for "tf jjg_ could be found for more concentrated solutions. 81 Figure 17 illustrates the results for Ag/Ag2S in 2.5m NaOH + 0.5m Na2^« The direction of change of potential with temp-erature differs from Figure 16. The difference in concentrations of Na28 cannot account for the differeace in behavior as temp-erature i s increased; the results should be very close. The difference illustrates the major problem with the silver sulfide electrode results . The discrepancy, which i s d i f f i c u l t to account for, was not discovered until near the end of the invest-igation. Early tests were performed with silver wire which was pre-anodized for 30 s in Na2^ solution and then placed in the test solution; a very small surface area was exposed to the solution. These early results were consistent with Figure 16 from test to test. The later tests used a rod electrode which had been pre-anodized but stored overnight in the Na2$ solution between tests. Figure 17 was one of the later tests. This change of r e s u l t s suggested that there was some difference in the electrodes. However, when both electrodes were placed together in the same c e l l , there was no difference in the measured potential. It was observed that, in the tests with the rod electrode, the potential increased over the duration of the test i f the temperature was maintained at a fixed point. This suggested a change in the electrode with time, but a freshly anodized elec-trode introduced into the c e l l had the same potential. It was suspected that the solution sulfide composition changed over the duration of the test, despite continuous nitrogen purging, but analysis f a i l e d to confirm t h i s . Furthermore, potentials 82 measured mixed the in solutions stored overnight were the same as those day of the test. The p o s s i b i l i t y of dissolution from the electrode i t s e l f remained. To investigate this, two tests were conducted at the same time under identical conditions at 70°C. The solutions were mixed and stored overnight at temperature. One c e l l had the Ag/Ag2S rod electrode in i t ; the other had no electrodes. In the morning, the potential of the electrode was measured in each ce l l using the same Luggin capillary and millivoltmeter. When the rod was placed in the c e l l which had had the rod in i t overnight the measured potential was high, consistent with Figure 17. The same rod, in the c e l l in which there had been no electrodes overnight, had a lower potential, consistent with Figure 16. After confirm-ing this difference several times, the rod was placed in the c e l l in which i t had not been located overnight. Three hours l a t e r there was no difference in the electrode potentials measured in the two c e l l s . These results suggest that the s i l v e r rod dissolved into the solution. There is no information in the l i t e r a t u r e on the formation of a dissolved s i l v e r species, but there has been l i t t l e work at this high pH at elevated temp-eratures. The difference in measured potentials is consistent with dissolution of Ag2s a t t n e beginning of the tests in new solutions. This would cause more Ag2S to be formed from Ag. The smaller wire electrodes used to obtain the data in Figure 16 would have dissolved less than the rod electrodes used for the data in Figure 17. 83 In several t e s t s , the electrode was maintained at a high temperature for a period of time. The potential approached the values consistent with Figure 17. Then, i f the temperature was increased fa i r l y rapidly, the potential measurement increased or i f the temperature was decreased, the potential measurement decreased, in proportion to the temperature at a rate consistent with Figure 16, but with a different base pot e n t i a l . Thus, i t seems reasonable to suggest that the electrode is thermodynamic, but is affected by something in solution. Although mercury forms a complex at high pH1 0 5, no information on s i m i l a r s i l v e r com-plexes could be found. In results for early tests comparable to those in Figure 16, the sulfide concentration was increased to lm Na2S as shown in Figure 18. Agreement between measured and theoretical data is excellent at a l l temperatures. On the other hand, in tests comparable to those in Figure 17', decrease of sulfide concen-tra t i o n to 0.05m resulted in behavior extremely different from that predicted, Figure 19. The discrepancy would be consistent with oxidation of the s u l f i d e , however analysis by t i t r a t i o n after the test indicated this was not the reason. Possibly, there i s i n s u f f i c i e n t s u l f i d e in solution to maintain the Ag2S surface. Alternately Ag2S dissolved quickly into solution and changed the chemistry as in 2.5m NaOH + 0.5m Na2S. Figure 20 i l l u s t r a t e s the effect of extremely high NaOH concentration. The measurements were not repeatable within 84 ~15mV, perhaps due to varying amounts of Ag2S dissolution from the wire electrode. There is a large difference between the measured and theoretical values which may be related to the very high concentration used. A high activity coefficient for OH" would be expected in these nearly saturated solutions. Yagil has found that in concentrated hydroxide solutions, most water molecules are bound to ions, thus reducing the a c t i v i t y of water and i n -creasing the a c t i v i t i e s of e l e c t r o l y t e s ^ ^ , This factor may account for the low measured values. Furthermore, heat capacities of high ionic strength at high temperature may be much different from those estimated and used here to calculate theoretical response1^. Alternatively dissolution of Ag2S may occur, thus affecting the solution chemistry. This explanation could account for the poor repeatability, which would depend on the amount of dissolution. Chloride caused no deviation from measured potential values in 2.5m NaOH + 0.5m Na2S as shown in Figure 21 compared with Figure 17. Chloride is unlikely to form an Ag/AgCl electrode in these solutions (Appendix IV). Hseu and Rechnitz9 0 found that CI- did not affect the Ag/Ag2S electrode potential when the ionic strength was maintained constant in the solution at 25°C. The results of the present work confirm their findings at 25°C and extend them to higher temperatures. Thus, the SSSE has not been found to be t he r mod y na m ica 11 y reliable. While the electrode may be stable under a given set of 85 conditions, the electrode behavior i s not predictable based on thermodynamic data. Unless improvements or better understanding can be realized the electrode i s not useful as a thermodynamic reference electrode. It remains useful for corrosion engineering s t u d i e s because of i t s r e l a t i v e s t a b i l i t y i n s u l f i d e s o l u t i o n s . I n t e r n a l v e r s u s E x t e r n a l E l e c t r o d e s a t P r e s s u r e To d e t e r m i n e the s u i t a b i l i t y of the SSSE at h i g h t e m p e r a t u r e (to 170°C) i t was necessary to measure i t s potential at temp-erature (SSSE(T)) versus a known reference p o i n t . A high temperature reference was considered to be unsuitable because of the d i f f i c u l t y of finding a high temperature electrode compatible with this solution. An external reference electrode was con-sidered satisfactory when corrected for junction potential. The reference point used could have been the SCE as was utilized earlier in this work. This would have involved building an external SCE for the autoclave. Measurements of the SSSE(T) obtained in this way would have compared easily with the values already obtained for the SSSE to 95°C. However there i s the possibility of introducing KC1 by leakage from the KC1 bridge. An alternative external reference electrode was the SSSE (25°C). It was selected because i t was easy to build, resistant to poisoning, useful for subsequent experiments and required no c o r r e c t i o n s f o r LJP. T h i s a r r a n g e m e n t r e q u i r e d e s t i m a t i o n o f TLJP for the caustic solution. It allowed the estimation of TLJP 86 T— i 1 1 1 1 1 1 — i 1 1 — i 1 1 r • i i i i I I i i i i 1 — i 1 1 1— 20 30 40 50 60 70 80 90 100 HO 120 130 140 150 160 170 TEMPERATURE , ° C Figure 22. Potential as a function of temperature for Ag/Ag2S (25°C)/2.5m NaOH + 0.5m Na2S/Ag2S/Ag (T). 87 to be appraised and SSSE p o t e n t i a l s to be measured to 170°C. The p o t e n t i a l d i f f e r e n c e between the SSSE(T) and SSSE (25°C) has been p l o t t e d i n F i g u r e 22. No c o r r e c t i o n s f o r LJP were n e c e s s a r y i n t h i s s i t u a t i o n . C o r r e c t i o n has been made f o r the s m a l l d i f f e r e n c e i n p o t e n t i a l of the two e l e c t r o d e s at 25°C. The s c a t t e r i n values observed at low temperature may be r e l a t e d to the problems d i s c u s s e d with regard to r e p e a t a b i l i t y ( F i g u r e s 16 and 17) and the p o s s i b i l i t y of changes i n the s o l u t i o n . Thermal l i q u i d j u n c t i o n p o t e n t i a l s were c a l c u l a t e d a c c o r d i n g to Macdonald7^. He noted t h a t , f o r c a u s t i c s o l u t i o n s : TLJP 0.5 (T°K - 298) mV . . . (28) A more recent s t u d y7 7 i n d i c a t e d that t h i s equation might be true only f o r very s m a l l temperature d i f f e r e n c e s . Macdonald et a l7 7 o b s e r v e d t h a t f o r a l a r g e t e m p e r a t u r e d i f f e r e n c e , TLJP i n KC1 inc r e a s e d i n approximately p a r a b o l i c manner with the temperature d i f f e r e n c e between the two h a l f - c e l l s . In the absence of pub-l i s h e d i n f o r m a t i o n , i t has been assumed f o r the p r e s e n t s t u d y that TLJP f o r NaOH + Na2S s o l u t i o n i n c r e a s e s i n a s i m i l a r manner. Thus, the TLJP has been assumed to be the sum of the j u n c t i o n p o t e n t i a l c a l c u l a t e d f o r s m a l l t e m p e r a t u r e d i f f e r e n c e s by e q u a t i o n 28 p l u s the t h e r m a l j u n c t i o n p o t e n t i a l measured by Macdonald et a l7 7 f o r KC1. T h i s summed TLJP i s t a b u l a t e d f o r a range of t e m p e r a t u r e s i n T a b l e V I I I (column 3). These c o r r e c -88 TABLE VIII Potential Difference, AE, Between Ag/Ag2S Internal Electrode at Temperature and Ag/Ag9S External Reference Electrode at 25°C T AE T L J P ........ AE+ E X ( M ) = ET (M) ° C mV mV T L J P mV E M G A P + L J P M VS H E - E 2 5 (M) 25 0 0 0 -674 0 30 1.5 -3.5 -2 -676 -2 40 3.5 -10.5 -7.6 -680 -6 50 4.5 -17.5 -13.5 -685 -11 60 6.5 -24.5 -18 -691 -17 70 8.5 -31.5 -23 -695 -21 80 11 -38.5 -27.5 -701 -27 90 13 -45.5 -32.5 -706 -32 100 14.5 -54.5 -40 -711 -37 110 15.5 -65.5 -50 -45" 120 17.5 -75.5 -58 -53 130 19.5 -84.5 -65 -60 140 21 -95.5 -74.5 - 68l r est. 150 23.5 -107.5 -84 -78 160 25 -119.5 -94.5 -88 170 27.5 -131.5 -104 1 -98, 89 tions were applied to the measured potential differences between SSSE (25°) and SSSE(T) to obtain the corrected values l i s t e d (column 4). The values in column 4 should be the true difference in potential of the half-cells at temperature and 25°C. The values in column 5 are those measured versus the SCE, referred to the SHE and corrected for LJP and TLJP as presented in Figure 16. Column 6 is the calculated difference between the SSSE (T) and SSSE (25<>C). Values from 110 to 170OC are extra-polated from data below 110°C. The values in column 6 should be the true difference in potentials of the half cells SSSE (T) and SSSE (25°C), and i f the estimated TLJP and LJP are correct c o l -umns 4 and 6 should be in reasonable agreement. The results in Table VIII indicate that in 2.5m NaOH + 0.5m Na2S, the electrode behavior up to 170°C i s consistent with extrapolations from lower temperature data. The potential difference between the internal SSSE (T) and SSSE (25°C), column 4, i s very close to the values obtained at low temperature versus the SCE, column 6. Considering the approximations, the agreement is very good. This confirms the re l a t i v e accuracy of the TLJP correction and the values of the electrode potential extrapolated to 170°C. Electrode Durability It was observed that the internal SSSE(T) became b r i t t l e and fragile and spalling of the Ag2S surface occurred. This would be consistent with rapid s u l f i d a t i o n of the electrode via rapid 90 d i f f u s i o n , and g e n e r a t i o n of e l a s t i c s t r a i n s and s u b s e q u e n t s p a l l i n g . The f r a g i l i t y of the e l e c t r o d e makes i t u n s u i t a b l e f o r use i n process streams. The e x t e r n a l S S S E i s u s e f u l i n l a b o r a t o r y s t u d i e s even though i t i s not r e l i a b l y p r e d i c t a b l e thermodynamica 11y . The e x t e r n a l e l e c t r o d e i s favored over an i n t e r n a l design i n t h i s work because no epoxy mounting was required and E°-p did not have to be c a l c u l a t e d . Furthermore, the e x t e r n a l e l e c t r o d e was not i n f l u e n c e d by metastable s u l f i d e s which may have formed by o x i d -a t i o n of the bulk t e s t s o l u t i o n . Thus, the e x t e r n a l SSSE (25°C) i s p r e f e r r e d because i t i s s t a b l e under a l l c o n d i t i o n s . 91 3.3.2 Polysulfide Electrode The measured potential values on platinum and mercury elec-trodes were plotted for a range of temperatures in 2.5m NaOH +• 0.423m N a 2 S ( F i g u r e s 23 and 24). L i q u i d j u n c t i o n p o t e n t i a l s were c a l c u l a t e d a s d e s c r i b e d in s e c t i o n 3.3.1. The m e a s u r e d potentials were closest to t h e polysulfide equilibrium p o t e n -t i a l s , f o r example a t 90° the p o t e n t i a l on p l a t i n u m was measured t o be - 0 . 5 6 8 V g H E # T h i s was c l o s e t o - 0 . 5 6 3 VsHE c a l c u l a t e d f o r equilibrium 1.11, HS-/S32", a t ts32~ 3 = °«001> b u t m u c n differ-ent from -0.714 VgH E calculated for equilibrium 1.3, HS"/S2032-, at [S2O32-] = 0.001. These concentrations of oxidized sulfur species are of the order of magnitude expected in these s o l -utions, based on the impurity level in the Na2S»9H20 chemical used (as described more fully on p.96). In order for equilibrium 1.3 to s a t i s f y the measured value, a r i d i c u l o u s l y high value of [S2O32 -] '"' 10l3 would be required. In general, i f true equilibrium exists between S^2 - and S2O32 - species in the as-prepared solution, then the equilibrium potentials for 29 and 30 w i l l be identical. (x-1) S2032- + 8 (x-1) H+ + 8 (x-1) e" ^ 2 (x-1) HS" + 3 (x-1) H20 ... (29) 4 Sx2 - + 4x H+ + 8 (x-1) e~ ^ 4x HS" ... (30) 92 -0/50 -0.460 -0.470 -0.480 ID -0.490 i </) > ^ Q -0.500 < -a5io L J £ -a520 -a530 -a540 -0.550 h -0.560 f — E E'UP'TLJP ET« ^ x(mol.S x*") x'2  v—0.156 m(5gplS) A 4.8 x 10* m O—2.4x10"3m 1.2x10"3m 20 30 40 50 60 70 80 TEMPERATURE,°C F i g u r e 2 3 : P t / S x 2 p o t e n t i a l a s a f u n c t i o n o f t e m p e r a t u r e f o r S C E / K C 1 ( s a t . ) / / 2 . 5 m N a O H + 0 . 4 2 3 m N a 2 S / P t . O p e n s y m b o l s a r e c a l c u l a t i o n s o f ^ S x ^ H S - ^ o r d i f f e r e n t S x ^ ~ c o n c e n t r a t i o n s . 93 -0.470 -0.480 -0.490 -0.500 ui I > -0,510 < -a520 -0.530 -0.540 -0.550 -a560 -0.570 h z LU t— O O-T r - - - E * L J P * T L J P ET©Zx(mot.Sx2") x=2  v—0.156m(5gplS) A—4^x10'3m O—2.4x10'3m •—1.2 x 10'3 m 20 30 40 50 60 70 80 90 T E M P E R A T U R E , ° C 100 F i g u r e 2 4 : H g / S _ 2 - p o t e n t i a l a s a f u n c t i o n o f t e m p e r a t u r e f o r S C E / K C 1 ( s a t . ) / / 2 . 5 m N a O H + 0 . 4 2 3 m N a 2 S / H g . O p e n s y m b o l s a r e c a l c u l a t i o n s o f S x 2 - / H S - ^ o r d i f f e r e n t SX2~ c o n c e n t r a t i o n s . 94 For one For s i m p l i c i t y , the case of the polysulfides. this total reaction: x=3 has been considered, that i s only Equations 29 and 30 have been added. AG = -16F ( E °3 0 - E°2 9) ... (31) and ln K = i i i ( E °3 0 - E02q) ... (32) RT where K = [H+ ] 4 [ S 2 Q 3 2 - ] 2 [ H S ~ ] 8 ... (33) [ S32" ]4 [H20]6 From thermodynamic data, the standard potentials were calculated for reactions 29 and 30, and K was calculated by equilibrium 32 to be 2 X 101 3. The activities of H+, HS- and H20 were considered for 2.5m NaOH + 0.5m Na2S + 0.1m Na2S203 solution at 90°C and were 10-12'92, Q-.5 and 1.0 respectively. The concen-trations of S2032 - and S32 - would form by equilibration from the S2032 - added (0.1m). Thus the sulfur balances via: 3 [S32-] + 2[S2032"] = 2[0.1] ... (34) Solving 33 with 34 to obtain [ S32~ ] and [ S2032~ ] , the concentration of [S32 _] was calculated to be <10~^ m. Due to this low [ S32 -] the e q u i l i b r i u m p o t e n t i a l should be v i r t u a l l y i d e n t i c a l with the potential for HS~/S2032_. The measured p o t e n t i a l s were not co n s i s t e n t with HS~/S2032 -, and so equilibrium must have not been attained. The measured potential must be a mixed potential and l i e s n e a r but below the H S ~ / S X 2 -equilibrium, reflecting the known3? metastability of Sx2 -. The 95 s t a b i l i t y of S^2~ has been v i s u a l l y confirmed in the present study. Sulfur additions to s u l f i d e solution form po l y s u l f i d e , g i v i n g the s o l u t i o n a bright yellow c o l o r . T h i o s u l f a t e solutions, on the other hand, are c o l o r l e s s . Polysulfide solutions maintained their yellow c o l o r f o r extended periods of time, even with thermal cycling, demonstrating the raetastability of the polysulfide. Thiosulfate solutions remained colorless for extended periods of time during thermal cycling, confirming their s t a b i l i t y . Giggenbach37, 43 found that, below 150°C, polysulfide 9 solutions are stable with respect to dissociation to S2O3 • Thus, polysulfide and thiosulfate cannot reach equilibrium, and the measured potential was a mixed potential due to polysulfides. Figure 25 i s an Evans diagram i l l u s t r a t i n g a possible and consistent explanation for the observed potential. The exchange current density for oxidation to S2032~, ±QI, may be low enough, or the Tafel slope high enough that the mixed potential l i e s close to the HS~/S^2~ equilibrium. The polysulfide equilibrium potential is determined by the equilibria between the various polysulfides in solution, and in turn on their concentrations. The concentrations have not been determined a n a l y t i c a l l y in this work. Instead, the concen-trations were calculated using equilibrium constants for the polysulfides-^8. The total amount of polysulfide was estimated to be the same as the level of s u l f i t e and thiosulfate impurity i n sodium sulfid e reagent, N a2S « 9 H 2 0 , that i s no oxidized sulfide 96 LU I— O Q_ • 01 'oz 'm L06 CURRENT DENSITY Figure 25: Evans Diagram of Sulfide Oxidation in Al k a l i n e Sulfide. 97 species other than polysulfides. The crystals of Na2S*9H20 had a s l i g h t l y yellow color when received, which was in d i c a t i v e of polysulfide. This is the justification for the assumption that polysulfide was present in the c r y s t a l s . The assumption makes t h i s c a l c u l a t e d Sx2~ e q u i l i b r i u m p o t e n t i a l an a p p r o x i m a t i o n . The t heoretical equilibrium potential of the polysulfide e l e c t r o d e was based on the same thermodynamic data f o r the poly-s u l f i d e i n s e c t i o n 2 . 2 . Two e q u i l i b r i a were c o n s i d e r e d to be important in these highly alkaline solutions: S32 _ + H+ + 2e~ ==s S22~ + HS" ... (35) S42" + H+ + 2e~ =52= S 32- + HS" . . . (36) According to Giggenbach38t and confirmed by p r e l i m i n a r y calculations, HS- would be the dominant reduced sulfide. Hence no equations were written with respect to s 2 - . E° was calculated for e q u i l i b r i a (35) and (36) as a function of temperature to 100°C and the resulting Nernst equations were obtained: E = E °3 5 - 2.303 RT log [S22~] [HS"] - 2.303 RT pH ... (37) 2F [S32 _] 2F E = E °3 6 _ 2.303 RT log [S32~] [HS~] - 2.303 RT pH ... (38) 2F [ S42 _] 2F 98 The pH and [HS-] were calculated as described in section 3.3.1. The Nernst equations for the reactions were equated to each other because they are in equilibrium. An equation in terms of the three unknown concentrations of polysulfides was obtained. - 2.303 RT log [ S 9 2~ ] [HS~]= E ° ^ - 2.303 RTlog [S^2~] [HS-] 2F [S32 _] [H+] 2F [S42 _] [H+] ... (39) Another equation in terms of the unknown concentrations of polysulfide was written by equating the t o t a l moles of the polysulfide sulfur to the moles of sulfur impurity in the solution, 4.8 X 10-3m. 4 I x (molality of Sx2 -) = 2 (molality of Na2S203 impurity) x = 2 ... (40) The impurity sulfur was contained in the Na2S»9H20 reagent and was quoted in terms of Na2S203« This impurity polysulfide was taken to be the only source of polysulfide in the nitrogen-purged solutions used in this work. Sulfur was not added to generate polysulfide. A third equation relating the concentrations of polysulfides to the equilibrium constant, Kj/2 for S22 - and S3 2"" was obtained. log K1 / 2 + log [HS-] + log [OH-] = 2 log [S22~] - log [S32"] ... (41) 99 A value of log Ki/2 = -3.56 was obtained from Giggenbach38. Equations 39, 40 and 41 were combined to yield an equation in terms of the concentration of S 3 2 " " . This was solved by iteration to obtain a value for S32 - which was then used to solve for S 4 2 ~ and § 2 ^ ~ ' The equilibrium potential was calculated using the Nernst equations for e q u i l i b r i a 35 or 36 and the p o l y s u l f i d e c o ncentrations. This c a l c u l a t e d t h e o r e t i c a l equilibrium potential was then plotted for comparison with the measured value, E, and the measured potential corrected for LJP and TLJP, E + LJP + TLJP, in Figures 23 and 24. Potentials were referred to the SHE scale. Theoretical equilibrium potentials were also calculated for one-half and twice the concentration of impurity giving different polysulfide concentrations and are plotted in Figures 23 and 24 for comparison. To appraise the effect of very high concentration of polysulfide, potentials were calculated for the case of 5 gpl sulfur addition, assuming an HS" concentration the same as for the previous cases, after the sulfur addition. The calculated values were then plotted in Figures 23 and 24. The potential measurements of the platinum electrode were unstable, and each data point has error bars indicating the oscillation of a few millivolts either side of the plotted mean value. The readings of the mercury electrode were more stable, and no error bars are shown. 100 The agreement between the measured and t h e o r e t i c a l values was very good considering the approximations. A theoretical value of electrode potential at the impurity l e v e l used in the experiments results in values which are within 25 mV of the corrected measured value. The results suggest that the potential of inert electrodes is due primarily to polysulfide reactions on the electrode surface. This is an important observation because engineers have employed inert electrodes as reference without understanding the basis of the potential on that reference e l e c t r o d e . C a r e f u l measurements and expert a n a l y s i s of polysulfide solutions at elevated temperatures would be necessary to provide conclusive interpretation of the measured potentials. Such an investigation was beyond the scope of the present project. The agreement between measured and theoretical potentials was improved at high temperature i f a lower concentration of polysulfide was assumed in the calculation of the theoretical equilibrium potential. A lower measured potential at high temp-erature would be consistent with lower [Sx2 -] than assumed and/or the S 2 O 3 2 - equilibrium and would be due to formation of a mixed potential due to oxidized sulfur species, e.g. S 2 O 3 2 - , as shown in Figure 25. The equilibrium 1.3, (i.e. HS- /S2O3 2 -) at [S2032 _] = 0.001 had a calculated value of -0.714 Vg H E at 90°C and would shift the mixed potential downward. Referring to Figure 25, i t may be seen that i f iQ^ for HS-/S2032- increases with temperature faster than iQ2 f o r HS~/SX2_, the mixed potential would be 101 shifted downward, as was observed. Both io values at 90°C may be small enough that insignificant amounts of S2O32- were generated (i.e. [Sx2 -] and [S2032-] remained the same throughout the test). According to the calculations, the addition of 5 gpL sulfur would be expected to change the polysulfide electrode potential by only ~15mV. Thus the potential would be probably relatively insensitive to large changes in polysulfide concentration. The electrode may provide a reference potential which is adequately stable for use in a corrosion control application where a ~25mV variation in potential in polysulfide solutions is tolerable. At high temperatures, the potential of the inert electrode may be strongly influenced by the presence of S 2 O 3 2 - because we then have a mixed potential and the temperature may strongly influence the kinetics of electrochemical oxidation (i.e. io) which deter-mine the mixed potential. Thus, the potential may be very dependent on the l o c a l temperature. In a continuous digester where temperature varies considerably along the length, the potential may vary s i g n i f i c a n t l y and w i l l make operation of an anodic protection system with an inert reference electrode com-plicated. The polysulfide electrode would have limited usefulness in corrosion studies due to i t s dependence on solution chemistry which is d i f f i c u l t to analyze and because of shifts in the mixed S x 2 - / S 2 0 3 2 - potential. On the other hand, with careful mainten-ance and awareness of l i m i t a t i o n s , the v a r i a t i o n s may be 102 t o l e r a b l e w i t h i n the l i m i t s r e q u i r e d f o r anodic p r o t e c t i o n or i n -s i t u c o r r o s i o n s t u d i e s . 103 4. POLARIZATION STUDY 4.1 Introduction The literature of polarization behavior of iron in caustic solutions i s substantial. Instead of treating the subject exhaustively, this present review places emphasis on the alkaline sulfide solutions and some of the most recent studies in straight caustic solutions. Many investigations of polarization behavior have involved cyclic voltammetry7*» 92, 117-119f a technique of polarizing the electrode by scanning l i n e a r l y through the potential range of interest whilst measuring the current density, and then reversing the scan direction to return to the i n i t i a l p o t e n t i a l . This d i f f e r s from a polarization curve which is for only one run through the potential range. A schematic cyclic voltammogram is illustrated in Figure 26. The changes in the current peaks with repeated cycling are of interest because they provide inform-ation on the kinetics of the electrochemical reactions. The speed of the potential scanning, in mV/s, is termed the scan rate. At the end of the scan range, the point at which the scan reverses direction has been denoted as the switching potential. Where there is a peak on the forward scan (oxidation) and a peak on the reverse scan (reduction) which result apparently from the same electrochemical reaction they have been termed conjugate peaks. V a r i a t i o n of the s w i t c h i n g potential i s u s e f u l i n determ-ining which peaks are conjugate. Interpretation of c y c l i c 104 i : i i 1 1 1 1— -1.50 -1.25 -1.00 -0.75 -0.50 -0.25 0.0 POTENTIAL, VSHE Figure 26: Schematic Cyclic Voltammogram of Steel in Alkaline Sulfide Solution. 105 voltammograms can be complicated because peaks i a t e d with thermodynamic p o t e n t i a l s due displacement of peaks by kinetic factors. may to not be assoc-skewing and Anodic polarization tests of steel have been performed in a l k a l i n e s u l f i d e pulpi ng l i q u o r s6. 8-10, 12, 13, 120 with increasing attention devoted to identifying corrosion reactions occurring at potential. Interpretation i s often complex, and most investigators did not interpret their results in terms of electrochemistry. Macdonald and Owen*!7 investigated the electrochemistry of iron in LiOH solution at 22°C and 200°C using cyclic voltammetry. At 22°C the anodic scan of the voltammogram had three peaks. They attributed the f i r s t peak to Fe(0H>2 formation. At 200°C, the anodic scan had only one peak located at approximately the same potential as the f i r s t peak at 22°C. The peak current of both this peak and i t s conjugate cathodic peak were proportional to the square root of the sweep rate, which they interpreted as due to a charge transfer process controlled by the mass transport of 0H~ to the metal surface. The d i f f e r -ence in potential between the conjugate oxidation/reduction peaks increased significantly with scan rate, indicating to them that the charge transfer processes occurring on the electrode surface were not very reversible. At 200°C, anodic peaks were observed on the reverse (cathodic) scan coincident with the f i r s t peak on 106 the forward scan. To the investigators, this behavior was indic-ative of reactivation processes involving charge transfer reactions which produce ionic products. Dihypoferrite, HFe02~» seemed to be implicated in this behavior. 0 j e f o r s l 2 1 s t u d i e d t h e polarization o f i r o n electrodes i n 5 N KOH. He s u g g e s t e d t h a t H F e 0 2 ~ was an i n t e r m e d i a t e i n t h e f o r m a t i o n o f F e ( 0 H ) 2 and FeOOH. F e r r i t e i o n , Fe02~> w a s a n intermediate i n t h e f o r m a t i o n o f FeOOH. The f i r s t peak on t h e anodic polarization curve was attributed to Fe(0H)2 formation, and the second and third peaks to FeOOH formation from Fe and Fe(0H)2 respectively. The electrochemical behavior of iron in alkaline s u l f i d e solutions of pH >12 at 25°C was investigated by Shoesmith et al92. In that study, the f i r s t peak on the anodic polarization scan was attributed to Fe(0H)2 or Fe304 formation. It almost disappeared as the hydroxide and b i s u l f i d e concentrations i n -creased. The second peak was attributed to Fe203« This peak and i t s conjugate almost disappeared i f the b i s u l f i d e and hydroxide concentrations were very large. The second peak did not increase in size with continuing cycling as i t did in sulfide-free solutions. The growth of oxide apparently was in-hibited in solutions containing sulfides. At s u f f i c i e n t l y anodic potentials, Shoesmith et a l9 2 con-c l u d e d t h a t t he o x i d e was a t t a c k e d and NaFeS2 was f o r m e d . T h i s formation of NaFeS2 caused the largest anodic peak. The peak 107 shifted to lower potential when the HS~ concentration increased r e l a t i v e to the OH- c o n c e n t r a t i o n . The NaFeS2 continued to be formed above the peak p o t e n t i a l . The charge passed during the peak depended on the HS~ concentration, indicating a k i n e t i c a l l y l i m i t e d r e a c t i o n which depended on the H S - c o n c e n t r a t i o n . P i t formation was observed at the peak. It was claimed that oxid-ation of the su l f i d e ion passivated the electrode by a deposit of elemental s u l f u r at pores i n the oxide l a y e r , but i t was not conclusively proven that the sulfur was not associated with iron in these deposits. Repassivation was claimed to be impaired by s u l f i d e because s u l f i d e would d i s s o l v e the p a s s i v a t i n g s u l f u r deposit to form p o l y s u l f i d e s . There was evidence that the Fe(OH) 2/Fe304/'6'-Fe203 layers were reduced prior to the reduction of the S deposits. A subsequent p u b l i c a t i o n1 2 2 dealt with the electrochemical formation of Mackinawite, F e S i _Xf in alkaline s u l f i d e solutions of pH 9-12. Mackinawite nucleated at f a u l t s i t e s i n the pre-cursor Fe3C"4 f i l m . The s o l u t i o n adjacent to the patch was considered to be very acid due to the reaction: Fe + SH" FeS + H+ + 2e~ ... (42) or by S formation. This a c i d i t y caused l o c a l attack and NaFeS2 fo r m a t i o n . S t i r r i n g i n h i b i t e d growth of Mackinawite patches, demonstrating that growth of patches was controlled by formation of some active soluble species at the edge of the growing patch. 108 M a c k i n a w i t e f i l m s were h i g h l y a l l o w i n g f u r t h e r d i s s o l u t i o n . s t r e s s e d and c r a c k e d r e a d i l y , Guzman, V i l c h e and A r v i a * * S i n v e s t i g a t e d the p o l a r i z a t i o n b e h a v i o r of i r o n i n c a u s t i c s o l u t i o n s v i a c y c l i c v o l t a m m e t r y . They a t t r i b u t e d the f i r s t two anodic peaks to Fe(0H)2 f o r m a t i o n . The t h i r d peak o b s e r v e d on the a n o d i c scan i n c r e a s e d i n h e i g h t w i t h each c y c l e and was a t t r i b u t e d to the t r a n s f o r m a t i o n of Fe(0H)2 t o FeOOH. T h i s peak d e c r e a s e d i n s i z e i f the c a t h o d i c p o t e n t i a l on the reverse scan was more p o s i t i v e than -1.2 Vgjjg. T h i s study provided the b a s i s f o r one by S a l v a r e z z a , V i d e l a and A r v i a * 1 9 which d e a l t w i t h the b e h a v i o r of m i l d s t e e l i n a l k a l i n e s u l f i d e s o l u t i o n s , pH ~11.5. They e x p l a i n e d t h e i r r e s u l t s i n terms of c o m p e t i t i v e a d s o r p t i o n of OH- and HS-. H y d r o x i d e s t i m u l a t e d the f o r m a t i o n of an o x i d e p a s s i v e l a y e r . The b i s u l f i d e , HS-, caused e i t h e r the f o r m a t i o n of mackinawite or a s o l u b l e f e r r o u s s p e c i e s . I t a l s o caused the FeOOH peak to d e c r e a s e i n s i z e w i t h c y c l i n g . P i t t i n g was i n h i b i t e d at pH 12. In c o n t r a s t to e a r l i e r s t u d i e s ^ , s u l f u r f o r m a t i o n was not ob-s e r v e d . An i n t e r e s t i n g c h e m i c a l t e s t was p e r f o r m e d as p a r t of t h i s s t u d y . I r o n powder i n 0.01M Na2S s o l u t i o n y i e l d e d Mackinawite and a s o l u b l e green product, NaFeS2* S i m i l a r r e s u l t s were obtained by r e a c t i n g FeS powder, but no r e a c t i o n took place w i t h Fe20*3. I t was considered t h a t , due to the l a c k of evidence t h a t s u l f i d e s p e c i e s c o u l d d i s s o l v e F e20 3 , i t was more l i k e l y t h a t FeS d i s s o l v e d i n the p r e s e n c e of H S ~ to y i e l d Fe2+ (Note: Fe2+ i s u n s t a b l e at pH 12. They may mean F e ( I I ) ) . D i s s o l u t i o n 109 of the metal was attributed to FeSH+ formation, which would then be s t a b i l i z e d in solution to form a hydrosulfide complex. The cathodic peak which Shoesmith et a l9 2 had attributed to S reduc-tion was ascribed instead to Fe(0H)2 reduction. Macdonald and McKubre7* conducted cyclic voltammetry experi-ments over a range of temperatures for iron in NaOH solutions. The peaks did not s h i f t much with temperature. The peak which they identified with FeOOH formation disappeared by 80°C in 28% NaOH. McKubre and Madou^2-^ interpreted similar cyclic voltammetric curves. At temperatures greater than 50°C, Fe203 was formed on electrodes from Fe(OH)2 or by precipitation of Fe(0H)4~. Accord-ing to their interpretation, at high potentials this Fe203 was converted to FeOOH. Tromans^ provided a thorough electrochemical interpretation of his data on the anodic polarization of mild steel in hot alkaline sulfide solutions. He i d e n t i f i e d the active-passive peak current with HFe02_ formation. Sulfide ions were incorp-orated in the Fe^O^ passive layer making i t nonprotective. The fi l m became protective only when the potential reached* that at which the incorporated sulfide was oxid ized to thiosulfate. The formation of FeS prior to the peak was attributed to a high l o c a l concentration of HFe02~, causing the solubility product of FeS to be exceeded. 110 Wensley and C h a r l t o n ^ observed more peaks than did T r o i a n s ^ but did not attempt to id e n t i f y them with calculated equilibrium potentials. Wiekowski, Ghali and Le*2 4 observed a correlation between changes in the cyclic voltamraogram of steel in caustic solution with the onset of susceptibility to c a u s t i c c r a c k i n g . Inspection of t h e i r voltaramograms and comparison with others suggests that the stress corrosion susceptibility correlates with the decrease of the FeOOH formation peak (thus poorer f i l m repair) and i n -crease of the Fe(0H)2 formation peaks which i s mediated by the HFe02- species. Most investigations described in the l i t e r a t u r e have been concerned with lower pH, lower temperature, s u l f i d e - f r e e solutions. There i s scant information in the l i t e r a t u r e on electrochemical behavior of steel in alkaline sulfide solutions at high temperature. This present investigation studies the electrode behavior of iron in highly alkaline sulfide solutions at high temperatures and pressures using cyclic voltammetry. I l l A.2 Experimental Anodic polarization curves and c y c l i c voltammograms were obtained for A516 Grade 70 carbon steel in alkaline sulfide solutions up to 1 8 0 o C. Some of the tests at temperatures below the boiling point of the solution were conducted in a Teflon corrosion c e l l . The higher temperature tests were performed in a pressurized autoclave. Polarization tests of platinum electrodes also were performed. The Teflon corrosion c e l l consisted of a 600 ml beaker f i t t e d with a Teflon l i d . It i s i l l u s t r a t e d in Figure 27. The c e l l was placed in an e l e c t r i c heating mantle and current supplied through a variable transformer. A Yellowsprings temp-erature controller (model 71) with a Teflon coated thermistor probe regulated the current supply to the heater to control temperature to within +/- 1°C. The thermistor probe, a Teflon-coated glass thermometer and a nitrogen purge l i n e entered the c e l l through the l i d . The c e l l was fitted also with a condenser. The working electrode consisted of a c y l i n d r i c a l specimen 6.2 mm in diameter and 11 mm in length. It was threaded onto a 316 stainless steel welding rod which was tightly fitted into a Teflon tube. The electrode and the threaded end of the conductor rod were potted in "Quickmount" epoxy so that the connection was insulated, and only the c i r c u l a r end face of the specimen was exposed. Immediately before testing,this end face was polished 112 Figure 27: Schematic of Teflon Test C e l l for P o l a r i z a t i o n Studies a) test electrode b) Luggin capillary c) counter electrode d) temperature probe e) nitrogen purge f) l i d g) beaker 113 to 600 g r i t on a clean paper with d i s t i l l e d water and rinsed. The counter electrode was a spectroscopic grade graphite rod, 6.2 mm diameter. The Teflon Luggin c a p i l l a r y , with a cotton thread passing through i t , was f i l l e d with saturated KC1 and connected to a calomel reference electrode v i a a g l a s s v a l v e . The autoclave (IL, Monel), used for high temperature polar-iza t i o n studies, was supplied by P a r r Instrument Co. A Teflon liner purchased with the autoclave was fitted with a Teflon l i d to minimize condensation inside the autoclave. The working elec-trode was of i d e n t i c a l design to that used in the Teflon c e l l . The graphite counter electrode was mounted on a stainless steel conductor rod and sealed with epoxy. The internal silver-silver sulfide reference electrode has been described in section 3.2.1. Two designs of external reference electrode were used. The f i r s t design i s i l l u s t r a t e d in Figure 28. It was designed so that both mercury-mercuric oxide and silver-silver sulfide refer-ence electrodes could be used. The design provided satisfactory service with both electrodes. There was concern, however, that this design could be affected by stray currents or ground loops and so a second, simpler, design was built. It is illustrated in Figure 14 and 15, section 3.2.1. The simpler design i s approp-riate only for solid reference electrodes; i t would be unsuitable for mercury. No difference in performance of the silver-silver sulfide reference electrode was observed between the two designs. 114 TO AUTOCLAVE CONAX PACKING GLAND T E F L O N W A S H E R S STAINLESS S T E E L BODY CAP & BOLT 1/8" T E F L O N BRIDGE T U B E . FLARED E N D . 1/A" STAINLESS ST. BRIDGE T U B E 1/A"ALUMINA T U B E COTTON T H R E A D & T E F L O N FILLING T U B E A g / A g 2 S T E F L O N S E A L S Ag WIRE IN T E F L O N T U B E hematic diagram of e x t e r n a l r e f e r e n c e e l e c t r a u t o c l a v e t e s t s ( o r i g i n a l d e s i g n ) . 115 The mercury-mercuric oxide reference electrode was used for polarization in sulfide-free solutions. It has been demonstrated to be a dependable reference electrode in straight caustic s o l u t i o n s ^ ' 87-91# The Hg/HgO (3.35 m NaOH) reference electrode was calibrated with respect to the saturated calomel electrode, VgCE* This measurement was corrected for liquid junction potential and con-verted to the standard hydrogen s c a l e , V g ^E. Electrode measurements versus the mercury-mercuric oxide, Hg/HgO, electrode in 3.35 m NaOH, VM MQE» were related to VgHE via: VSHE = VM M 0 E + 0.070 V ... (43) The electrodes entered the top of the autoclave through Conax fitting s with Teflon seals. A 316 stainless steel pressure gauge(6.9 MPa max.) was f i t t e d onto the gauge block on top of the autoclave. The gauge block also had a gold-faced Inconel rupture disc (7.185 MPa), a pressure hose connection, and a valve. There was a thermal well for the thermistor probe and thermocouple wires. The thermal well was covered with a Teflon sheath inside the autoclave to prevent ground loops with the electric heater. The autoclave heater was obtained from Parr Instrument Co. It was an insulated resistance heater, 120V. Power was supplied to i t through a variable transformer, 22A max. which was manually set, according to the desired temperature range. The temperature 116 range was regulated by a specially calibrated Yellowsprings con-t r o l l e r with a high temperature range thermistor probe (model 632). The temperature was measured also by a chromel-alumel thermocouple with Keithley microvoltmeter (model 153). When the autoclave temperature f e l l below the set point the heater c i r c u i t was closed and current was supplied to the heater. When the autoclave temperature reached the set point the heater c i r c u i t opened and a cooling blower turned on. The blower forced a i r i n t o the bottom of the autoclave heater, past the autoclave and out the top of the heater. This cooled the heater and autoclave to prevent overshooting the set temperature. The speed of the blower was regulated by a variable transformer. A r e s i s t o r in parallel with the heating switch allowed some current supply to the heater even when the blower was on, so that the heater would be warmed up to supply heat more quickly when the temperature f e l l below the set point. Temperature could be maintained at the setpoint for about 10 minutes, providing ample time for polariz-ation tests at the potential scan rates employed. The ASME A516 Grade 70 hot-rolled carbon steel u t i l i z e d in this investigation had a chemical position, in weight percent, C-0.18, Si-0.22, Mn-1.12, P-0.021 and S-0.007. The yield stress was ~340 MPa. Solutions for the tests were prepared as previously des-cribed in s e c t i o n 3.2.1. During the t e s t s at atmospheric pressure, nitrogen purging was maintained throughout the test. 117 Prior to the tests in the autoclave, a nitrogen overpressure of ~0.345 MPa was applied. The working electrode was cathodically polarized usually at -1.25 VgsSE ^o r 300 s p r i o r to the p o l a r i z a t i o n or c y c l i c v o l t -araraetric t e s t s . P o t e n t i a l was scanned during the t e s t s by using a P r i n c e t o n A p p l i e d R e s e a r c h p o t e n t i o s t a t , model 173, e q u i p p e d w i t h an electrometer probe, model 178, logarithmic current converter, model 376, and programmer, model 175. The potential and current were recorded on a Houston X-Y recorder, model 2000. Corrections for TLJP have been described in section 3.3.1. 118 A.3 Results and Discussion 4.3.1 Cyclic Voltammetry in Plain NaOH Solution The b e h a v i o r o f A516 s t e e l i n p l a i n NaOH s o l u t i o n s was i n v e s t i g a t e d to a i d i n the i n t e r p r e t a t i o n of b e h a v i o r of s t e e l i n the more complex a l k a l i n e s u l f i d e s o l u t i o n s . The c y c l i c v o l t a m m o g r a m s i n 3m NaOH a t 9 5 , 120 and 150°C are i l l u s t r a t e d in Figure 2 9 . The p o l a r i z a t i o n curves a r e un-corrected for LJP and TLJP. The current peaks were increasing in size during the f i r s t two cycles, but after cycle 3 the behavior stabilized, producing superimposable traces. The third cycle at each temperature is illustrated. The peak potentials are listed in Table IX. The measured potential of peak I did not change with temperature, but the potential corrected for TLJP decreased with increasing temperature. The decrease of potential is con-sistent with displacement downward of lines in E-pH diagram at higher temperatures. The potential of peak II decreased at higher temperature, i t broadened and appeared as a shoulder on peak I at high temperature. The peak current increased slightly with cycling, and decreased at higher temperature. The increas-ing size of peaks during the f i r s t three cycles may have been due to increasing surface area during cyclic voltammetry. The increase in size of the peaks with increasing temp-e r a t u r e was c o n s i s t e n t w i t h i m p r o v e d d i f f u s i o n a t h i g h e r temperature, supplying and removing soluble reactants from the 119 r o i o X E — < > CO cr on O •1.25 -1.00 -0.75 -0.50 -0.25 POTENTIAL , V S H E 0.0 Figure 29: Cyclic Voltammogram of A516 Steel in 3m mV/s) . NaOH (50 120 TABLE IX Potentials of Cyclic Voltammogram Current Peaks and Potentials of Relevant Equilibria at 95, 120 and 150°C in 3m NaOH  Peak 95 120 150 E T -0.740 -0.740 -0.740 ± E + TLJP -0.790 -0.816 -0.848 E II E + TLJP -0.555 -0.605 -0.615 -0.691 -0.660 -0.768 E III E + TLJP -0.851 -0.901 -0.895 -0.971 -0.895 -1.003 E IV E + TLJP -1.055 -1.105 -1.085 -1.161 -1.085 -1.193 Equilibrium 2.5 -1.030 -1.049 -1.093 Fe/Fe304 2.52 -0.841 -0.872 -0.930 Fe(0H)3-/Fe304 2.71/2.72 -0.779 -0.799 -0.859 Fe(0H)3~/Fe00H or Fe2o3 2.40/2.41 -0.643 -0.661 -0.715 Fe304/Fe00H or Fe203 2.63 -0.721 -0.721 -0.736 Fe(0H)3"/ Fe(0H)4" 2.10 -1.083 -1.107 -1.149 Fe/Fe(0H)3" 46 -1.019 -1.033 -1.082 Fe/Fe(0H)2 47 -1.037 -1.065 -1.133 Fe(0H)2/Fe304 121 electrode. During a polarization scan the peaks were displaced in the direction of the scan. T h e r i s i n g current corresponding to d i s s o l u t i o n d e c r e a s e s when a f i l m f o r m s on t h e s u r f a c e . The process of film formation (oxidation) requires time for nu-cleation and growth, and i f the potential scan is proceeding during t h i s t i m e , a higher o v e r p o t e n t i a 1 f o r dissolution w i l l have been reached b e f o r e f i l m f o r m a t i o n p roduces f u l l c o v e r a g e of the surface. Thus the peak w i l l occur at a higher potential. As the scan rate is increased, the potential w i l l reach increasingly higher values before s i g n i f i c a n t f i l m coverage occurs. On the reverse scan, reduction takes place and the peak is moved in the direction of the scan (toward lower potentials) because time is required for the film to be reduced. Thus the oxidation and reduction conjugate peaks for the same process would be displaced in opposite directions. They would not be displaced by equal amounts because the reaction paths of film formation and reduc-tion processes could be different. Peak IV appeared to be the conjugate of peak I because of their similarity in size. Due to scan rate effects, the equilib-rium potential must be located somewhere between the two conjugate peak potentials. One method of estimating this e q u i l i b r i u m p o t e n t i a l i s to assume i t corresponds to the potential at which a line connecting the two conjugate peak maxima intercepts the zero current axis. By this method, the 122 "intercept average" potential for the conjugate pair I and IV at 95, 120 and 150°C i s -0.925, -0.963, and -0.973 VSHE' corrected for TLJP. These p o t e n t i a l s are close to the c a l c u l a t e d potentials for e q u i l i b r i a 2.5 arrd 2.52,- Fe ^ Fe304 and Fe(0H)3 ^ F e304, l i s t e d in Table IX. They are closer to 2.52 than to 2.5, in agreement with the E-pH diagram which indicates that Fe(0H)3- -» Fe304 should be expected. The r i s i n g current of peak I was probably due to oxidation of iron to Fe(0H)3~ via equilibrium 2.10. This would have pro-duced a concentration of Fe(0H)3_ which was estimated by considering the charge passed during a scan. F i r s t , i t was assumed that this charge was partly due to the oxidation of iron to surface oxides and partly to soluble iron species, and that the charge passed in the cathodic scan was due to reduction of surface oxides. The d i f f e r e n c e i n charge passed in the anodic and cathodic scans represented an estimation of the quantity of soluble iron species formed. The charge passed was measured by plotting the current versus time and determining the area under the curves (i.e. to t a l coulombs) for the anodic and cathodic scans. The difference in these measurements provided an estimate of the number of equivalents, Q, of soluble species formed. The number of moles of a particular soluble species was then c a l -culated from Faraday's laws via: Q/Fn = No. of moles ... (44) where F = Faraday (96500 coul./mole) 123 and n = oxi Molality was in the c e l l , dation state of the soluble m e t a l l i c approximated by dividing by the volume of v. s p e c i e s , solution . Q/Fnv = molality ... (45) Assuming Fe(0H)3~ was formed (n=2), and diluted into the f u l l volume of the test c e l l , v, then concentrations were at 95oC: 1.9 x 10-7m, 120OC: 3.9 x 10-7ra and at 150oC: 6.7 x 10-7m. Even i f i t were assumed that a l l current up to the maximum of peak I formed Fe(0H)3~t concentrations were only marginally increased, e.g. at 95°C: 2.86 x 10~7m, 120°C: 5.94 x 10-7m, and at 150°C: 8.13 x 10~^m. These concentrations are lower than 10-D taken as the activity potentials in Table IX. The use of higher values is quite reasonable for the calculation though, because mass trans-fer would be too slow to cause f u l l dilution in the 3 seconds i t takes to move from the beginning of Peak I to i t s maximum, and so the local concentration near the electrode would be much higher. In s i m i l a r tests in which a v i s i b l e coloration formed at the electrode, the color lingered at the electrode surface, occupying a volume approximately one hundredth the volume of the test c e l l . About a minute was required for this to dissipate completely. This could account for the better agreement i f 10-^ i s used for the Fe(0H)3 activity in the equilibrium equations. The shift of peak I with temperature is less than predict-ed for equilibria involving Fe(0H)2 but greater than predicted 124 using HFe02~ instead (see Table IX this section and Table V, section 2.3). The value for EHFEQ _/pe Q at 150°C was too high 2 3 4 to account for the observed potential. Thus the thermodynamic values of Fe(0H)3~ and HFe02- were not s u f f i c i e n t l y r e l i a b l e to distinguish between the processes. More accurate Cp values might clarify the situation. If a l l the charge passed in peak I were due to Fe^Q^ formation, with a l l Fe(0H)3~ being converted to Fe304, then the resulting thickness of Fe304 fi l m present may be estimated as described in Appendix V. Accordingly, the thicknesses were at 95°C: 0.075 jum, at 120°C: 0.160 jim and at 150°C: 0.258 j i m . The film thickness is overestimated because some of the charge passed in peak I would be to form Fe(0H)3~ which is dissipated into the solution. These are f a i r l y thick films in r e l a t i o n to the several monolayer thick films normally associated with passivity. The possibility that peaks II and III were conjugates was considered. Peak III was smaller than peak II but was displaced in potential consistent with being a conjugate of peak II. If they are conjugates, the difference in size may be due to ir r e v e r s i b i l i t y of the reaction. The intercept average potential for the conjugate pair was measured using a line drawn from peak to peak, as described above. Corrected for TLJP, i t was -0.838, -0.926 and -0.951 Vg H E at 95, 120 and 150°C respectively. These intercept average potentials correlated with the calculated values for e q u i l i b r i a 2.71 and 2.72, Fe(0H)3" FeOOH or Fe304. The FeOOH is stable to 130°C and Fe203 i s stable above 130°C. 125 The observed decrease in potential of peak II with increasing temperature was consistent with the lower potentials for equilib-ria 2.71 and 2.72 at higher temperature. Peaks II and III d i f f e r in size, and so the p o s s i b i l i t y that they are not conjugate must be considered. Thus, some other reaction must account for peak II or II I . If the alternative F e ^0 4 s#= FeOOH and Fe304 ssg= Fe203 e q u i l i b r i a were considered, then peaks II and III are unlikely to be conjugates because their intercept average potential was not consistent with the c a l -culated potentials for e q u i l i b r i a 2.40 and 2.41. Peak III was displaced in the active direction from e q u i l i b r i a 2.40 and 2.41 consistent with the anticipated scan rate e f f e c t . If FeOOH or Fe203 is formed at peak II via equilibria 2.71 or 2.72 the charge passed w i l l be 1 charge per mole of iron (n=l). If peak III i s due to reduction of that FeOOH or Fe203 to Fe304, then only 1/3 charge per mole of iron w i l l be passed (n = l/3). This difference in charge passed could account for the difference in size of the two peaks. Assuming peak II i s due to FeOOH or Fe203 formation from Fe(0H)3~, the thickness of Fe00H-Fe203 film formed was estimated, as described in Appendix V. The estimated thickness of film was at 95°C: 0.127 jim at 120°C: 0.084 um and at 150°C: 0.048 |im. These numbers probably overestimate the actual thickness because the measurements of charge passed in peak II may include some charge from the t a i l of peak I, the size of which could only be 126 estimated. This layer of Fe203 o r FeOOH may be deposited on top of the FegOA fi l m already formed at peak I. The Fe203 may com-pletely cover the Fe304, preventing oxidation of the Fe304 t 0  F e203. On the reverse scan, peak III would represent reduction of FeOOH or Fe203 to Fe(0H)3~ v i a equilibrium 2.71 or 2.72 or the reduction of FeOOH or F e2 0 3 t 0 Fe304 via equilibrium 2.40 or 2.41. Assuming peak III was entirely due to reduction of Fe203 to Fe304, the tota l charge passed during the peak was used to calculate the thickness of FeOOH or Fe304 fi l m which could be reduced. The thickness was at 95°C: 0.054 ^i m, at 120°C: 0.016 ^im, and at 150°C: 0.012 j i m . These numbers are comparable with those for film formation at peak II. This is good evidence that the suggested reactions are occurring. The film formed at high temperature might have been more protective than FeOOH formed at lower temperature, and thus may be thinner. These thicknesses of Fe203/Fe00H being reduced at peak III were lower, but comparable, with the estimated thicknesses of film formed at peak I I . The thickness formed at peak II may have been over-estimated, for the reason already described. The Fe203 or FeOOH being reduced may have been present as a layer on top of the Fe304' Alternatively, the entire thickness of film may have been FeOOH or Fe203. This latter possibility w i l l be examined below. If peaks I and IV are conjugates, then Fe304 must be present after peak III. Peak IV should have resulted from the reduction process 127 F e 3 0 4 > F e ( 0 H ) 3 - a c c o r d i n g t o e q u i l i b r i u m 2 . 5 2 i f i t w a s t h e c o n j u g a t e o f peak I . The F e ( 0 H ) 3 - may t h e n h a v e b e e n r e d u c e d t o F e v i a e q u i l i b r i u m 2 . 1 0 . T h e c a l c u l a t e d p o t e n t i a l s i n T a b l e IX s u p p o r t t h i s p r o p o s e d s e q u e n c e . T h e t h i c k n e s s o f o x i d e f i l m r e d u c e d was e s t i m a t e d f r o m t h e c h a r g e p a s s e d d u r i n g peak I V , as d e s c r i b e d i n A p p e n d i x V. A s s u m i n g i t was e n t i r e l y F e 3 0 4 r e d u c e d t o F e , t h e t h i c k n e s s was a t 9 5 ° C : 0 . 0 6 9 jim , 1 2 0 ° C : 0 . 1 1 2 ^ m a n d a t 1 5 0 o C : 0.164 pm. T h e s e t h i c k n e s s e s a r e c o m p a r a b l e w i t h t h o s e c a l c u l a t e d f o r t h e f i l m f o r m a t i o n s e q u e n c e . The t h i c k n e s s e s a r e c g r e a t e r t h a n t h e t h i c k n e s s e s e s t i m a t e d t o f o r m a t p e a k I I I , i n d i c a t i n g t h a t t h e F e 2 0 3 " F e O O H l a y e r w a s o n l y p a r t o f t h e f i l m t h i c k n e s s , p r o b a b l y o n t o p o f a n F e 3 0 4 l a y e r . T h i s F e 3 0 4 u n d e r l a y e r w o u l d be r e d u c e d a t p e a k I V , a l o n g w i t h t h e F e 3 0 4 l a y e r f r o m r e d u c e d F e 2 0 3 o r F e O O H . T h e t h i c k n e s s e s a r e l e s s t h a n t h o s e c a l c u l a t e d f o r F e 3 0 4 f o r m a t i o n a t p e a k I , b u t t h e t h i c k -n e s s e s c a l c u l a t e d f r o m peak I w e r e o v e r e s t i m a t e d . As n o t e d a t t h e b e g i n n i n g o f t h i s s e c t i o n , p e a k s i n c r e a s e d i n s i z e d u r i n g t h e f i r s t 3 c y c l e s , a p p a r e n t l y d u e t o i n c r e a s i n g e l e c t r o d e a r e a . T h i s i n c r e a s e i n a r e a w o u l d a r i s e f r o m t h e f o r m a t i o n a n d r e d u c t i o n o f o x i d e f i l m s w h i c h w o u l d r o u g h e n t h e s u r f a c e . M a c d o n a l d a n d O w e n * ! 7 o b s e r v e d a l a r g e p e a k , s i m i l a r t o p e a k I , o n i r o n i n 1 M L i O H a t 2 0 0 ° C . T h e y a t t r i b u t e d t h e p e a k t o F e 3 0 4 f o r m a t i o n f r o m i r o n a n d H F e 0 2 _ (now b e l i e v e d t o be F e ( 0 H ) 3 - ) > i n a g r e e m e n t w i t h t h e i n t e r p r e t a t i o n o f t h e p r e s e n t 128 r e s u l t s . Shoesmith et al92 a nd Guzman et a l l l° \ in studies at lo wer pH and temperature, suggested that Fe(0H)2 formed instead of Fe304« Tromans1^ ascribed a peak at -0.77 VgHE t 0 F e304 formation. McKubre and Madoul23 attributed peak I at ~75°C to F e 3 0 4 f o r m a t i o n d i r e c t l y f r o m i r o n o r f r o m H F e 0 2 - v i a F e ( 0 H ) 2 p r e c i p i t a t i o n . The c a l c u l a t i o n of the e q u i l i b r i a : Fe(0H)2 + 2 e _ =^ Fe + 20H- . . . (46) F e3o4 + 2H20 + 2H+ + 2e~ =^  3 F e ( 0 H ) 2 . . . (47) indicate that Fe(0H)2 forms at a potential above that for Fe304 formation, as summarized in Table IX. For this reason, Fe(0H)2 formation has been eliminated in the present work. McKubre and Macdona ld71 linked peak IV, which was smaller in their study in plain caustic solutions to a highly reversible surface process. They found that the peak size decreased as the temperature was increased, in contrast to the present r e s u l t s . McKubre and Madoul23 attributed peak IV to the reduction of FeOOH to HFe02- and its precipitation as Fe(0H)2- Their interpretation indicated that peaks I and IV were not conjugates. They linked peak I with Fe304 formation and peak IV with FeOOH reduction. Also, peak IV was attributed to FeOOH or Fe203 reduction rather than Fe304 reduction and peak III to Fe203 transforming to FeOOH. If peaks II and IV are considered to be conjugate and assuming Fe(0H)3~ concentrations about 10~7m, the "intercept average" p o t e n t i a l w o u l d be i n a g r e e m e n t w i t h e q u i l i b r i a 2.71 o r 2 . 7 2 , Fe(0H)3~ / Fe203 o r FeOOH, lending some support to their inter-129 p r e t a t i o n . However, t h e i r i n t e r p r e t a t i o n did not e x p l a i n a l l of the present r e s u l t s adequately, f o r example the observed s i m i l a r -i t y i n s i z e of peaks I and IV. The d i f f e r e n c e s may a r i s e from t h e i r use of a r o t a t i n g e l e c t r o d e i n c o n t r a s t to the s t a t i c e l e c t r o d e used h e r e . D i f f e r e n c e s i n the c y c l i c voltammogram would be expected under these circumstances-'-2^. I f the cathodic p o l a r i z a t i o n i s performed at a p o t e n t i a l which i s not low enough, peak I i s s m a l l e r H S , s u g g e s t i n g t h a t p o t e n t i a l s much l o w e r than peak IV, i n the hydrogen e v o l u t i o n r e g i o n are r e q u i r e d to reduce F e3o4 f u l l y . In the present work, p o t e n t i a l s were i n s u f f i c i e n t l y a c t i v e to s a t i s f y t h i s c o n d i t i o n . Shoesmith et al92 observed c u r r e n t peaks f o r i r o n i n NaOH s o l u t i o n w hich they a s c r i b e d to F e 2 03 f o r m a t i o n and r e d u c t i o n . O t h e r s7 1' 1 1 8 o b s e r v e d a peak at -0.400 VS H E due to FeOOH f o r m -a t i o n . I t s c o n j u g a t e r e d u c t i o n peak at -0.850 VsHE w a s s a i d to be c o n t r o l l e d by f i l m t h i c k n e s s , and d e c l i n e d i n s i z e at temp-e r a t u r e s above 4 0° C . T h i s t r e n d c o n t i n u e d a t t h e h i g h e r t e m p e r a t u r e s u s e d h e r e . T r o m a n s1^ a t t r i b u t e d a peak at -0.66 VS H E to the t r a n s f o r m a t i o n of F e304 to Fe203- When compar-ing the present peak p o t e n t i a l s to other work,1 4' 7 1» 92, 1 1 8 i t must be noted t h a t no c o r r e c t i o n f o r LJP and TLJP was i n c l u d e d by the o t h e r s . McKubre and Madou1 2 3 observed a peak II f o r i r o n i n 30% KOH at ~ 7 5 ° C and s u g g e s t e d t h a t i t was due to F e 2 03 f o r m a t i o n from Fe(0H)2 p r e c i p i t a t e d from HFe02~- Peak I I I was s a i d to be due to r e d u c t i o n of F e 2 03 to F e304. A l l of t h e s e s o u r c e s agree 130 in general that peaks II and III represent oxidation and reduc-tion involving FeOOH or Fe203 with Fe(0H>2 or Fe304. According to McKubre and Madou,^3 the proposed sequence of reduction reactions at peak III i s not the reverse of the oxidation reac-t i o n s a t peak I I , i n a g r e e m e n t w i t h t h e p r e s e n t i n t e r p r e t a t i o n . E x c e p t f o r t h e f i r s t t h r e e c y c l e s , p e a k s II and III w e r e no t o b s e r v e d to i n c r e a s e i n s i z e w i t h c y c l i n g , as some i n v e s t i g a t o r s o b s e r v e d a t 2 5 0 C . 9 2 , 118 0 t h e r s l 2 4 m n o t o b s e r v e peak II a t a l l i n 1 M NaOH at 95°C. The p r e s e n c e of the peak i n the p r e s e n t w o r k , and i t s s l i g h t i n c r e a s e i n s i z e w i t h c y c l i n g may have resulted from the use of a higher 0H~ concentration. In summary, i t i s concluded that the present study i s consistent with: Peak Reaction Equation No. I Fe —» Fe(0H)3~ —• Fe304 2.5, 2.10, II Fe(0H)3- -> FeOOH or Fe203 2.71 , 2.72 Above II Fe304 FeOOH or Fe203 2.40, 2.41 III FeOOH or Fe203 Fe304 2.40, 2.41 IV Fe304 -» Fe(0H)3" —> Fe 2.52, 2. 10 Below IV remaining Fe304 —» Fe 2.5 , 131 A.3.2 Sulfide Oxidation The anodic behavior of platinum in alkaline solutions was investigated to provide information on sul f i d e oxidation. The polarization behavior was due to sulfide oxidation only because platinum behaves as an inert electrode. In this way, the features of the polarization behavior of iron in caustic sulfide solutions, due to sulfide alone, could be distinguished from features involving iron oxidation interactions with sulfide. Figure 30 illustrates the 2.5m NaOH + 0.5m Na2S with add Na2S203. The thiosulfate was a was the anticipated f i n a l metas and so should have affected the the reversible equilibrium pot equation: behavior of platinum of 90°C in it i o n s of 0.01m, 0.1m and 1.0m dded to the solution because i t table product of HS~ oxidation, electrode behavior by displacing e n t i a l according to the Nernst E = E° - 2.303 RT log fHS~]2 - 2.303 RT (pH) ... (A8) 8F [S2032-] F The equilibrium electrode potentials based on equation 48 are l i s t e d in Table X. The corrosion potentials were more con-sistent with the HS- / Sx2 - equilibrium potentials (equation 38 and 39) than the HS- / S2032 - equilibrium (A8), as discussed in section 3.3.2. The corrosion potentials represented mixed p o t e n t i a l s . The s l i g h t amount of p o l y s u l f i d e i n s o l u t i o n would 132 Figure 30: Anodic Polarization Curves for Platinum of NaOH + Na2S Solutions with Varying Na2 S2O3 Concentration (lmV/s) TABLE X Corrosion Potential of Platinum and Calculated Equilibrium Potentials at 90°C Solution EC0RR EC0RR + TLJP +LJP E Rx. 1.3 NaOH Na2S Na2S203 2.5 0.5 0.01 -0.570 -0.599 -0.705 2.5 0.5 0.10 -0.600 -0.629 -0.696 2.5 0.5 1.00 -0.575 -0.604 -0.687 134 derive from the polysulfide impurity in the Na2S*9H20 c r y s t a l used. During anodic polarization of this electrode, i n i t i a l oxidation to S 2 O 3 2 - would occur above the mixed potential. Oxid-ation to Sx2 - would occur at slightly higher potentials above the equilibrium potential for HS- / S x 2 - (- 0.4 V ) , as shown in Figure VI-1 in Appendix VI. This oxidation to Sx2~ was v i s u a l l y confirmed by coloration changes. A platinum electrode was main-tained at -0.260 Vgjjg in alkaline sulfide solution for a day, during which time the solution turned yellow. This yellow color resulted from polysulfides formed in the solution. The Tafel slopes varied as shown in Table XI. There were two linear sections on the anodic polarization curves: above and below -0.4 Vgjjg. The section above -0.4 Vgjjg had a lower slope. Appendix VI demonstrates how two slopes may arise when the Tafel behavior for two concurrent reactions d i f f e r , these two oxidation reactions resulting in S2032 - and Sx2~ respectively. The relative amounts of each oxidation product are also indicated. At the lowest S2032 - concentration, 0.01m, the Tafel slope was approximately 145mV/decade, and was attributed mainly to oxidation of HS~ to S2032 -. The transmission c o e f f i c i e n t was calculated via: X = 2.303 RT ... (43) baSF where the transmission coefficient, A, is the charge transferred 135 TABLE XI Tafel Slopes on Platinum at 9QQC Solution Composition bg mV/decade NaOH Na2s Na2S203 < -o.A v S H E > "0.4 VS H E 2 . 5m 0.5m 0.01m 150-137 2. 5m 0. 5m 0. lm 142-140 75 2. 5m 0.5m 1.0m 125 75-87 136 through the double layer each time the rate determining step occurs, ba i s the Tafel slope and 8 i s the symmetry factor, assumed equal to ~0.5. The symmetry factor i s the f r a c t i o n a l distance of the activated complex across the double layer. According to the calculation, a single charge was transferred each time the rate determining step occurred. Below -0.4 Vg^g. the Tafel slope was almost the same in the solution with 0.1m Na2S203« This slope should not vary i f the reaction is occurring under activation control. The higher Tafel slope obtained in 0.01m Na2S203 at > -0.4 Vgjjg i s suggestive of some di f f u s i o n control, with the higher concentration of S 2 O 3 2 - polarizing the 9 oxidation to S2O3 • A reaction path for the oxidation of HS- to S 2 O 3 2 - con-sistent with X =1 for the rate controlling step i s : 2HS- + OH- HS2~ (ads.) + H20 + 2e~ ... (49) HS2_ (ads.) + 60H- HS203- + 3H20 + 6e~ ... (50) H S2°3~ + 0 H~ -*• s2°32~ + H2 ° ••• <51) The adsorbed species has been introduced here because Latimer^ has indicated that H S 2 - is a possible form of oxidized s u l f i d e . Equation 50 i s considered to be the rate determining step. The net reaction for steps 49, 50 and 51 i s : 2HS- + 80H- + S2032- + 5H20 + 8e" ... (52) 137 which i s equation 1.3. Reaction 50 may be s p l i t into two successive reactions: HS2" (ads.) + OH- ->- 2S(ads.) + H20 + 2e" ... (53) and 2S (ads.) + 50H- -+ HS203_ + 2H2o + 4e~ ... (54) Reaction 54 then becomes the rate determining step. Equations 53 and 54 are considered because they include sulfur passing through the zero oxidation state, making the oxidation more gradual: HS" (-II) + HS2- (-1) •*• S(0) HS203- (ID Formation of S22 - from adsorbed sulfur would be a competing reaction which may occur above EHS-/S2032-. Either reaction 50 or 54 is thought to be rate controlling because of the number of species involved in the reaction and because they result in a singly charged species. In agreement with the calculated transmission coefficient, the singly charged H S 2 O 3 - ion would be transferred across the double layer, where i t would react with OH- to form S2032 -. The existence of HS203_ a t the surface is enabled by the consumption of OH- ions which may l o c a l l y depress the pH. This suggested reaction path through reaction 53 is also consistent with the observation of Moscardo-Levelut and Plichonl26, 127 that sulfur 138 is an intermediate in the oxidation of Na2S in equimolar NaOH-H^ O melts at 75-150°C. The HS2O3- species is unlikely to be detected in the bulk solution because of i t s anticipated short lifetime in highly caustic solutions. The passage of H S 2 - through the double layer after reac-tion 49 would be a less satisfactory explanation to account for a unit transmission coefficient because the H S2- may react via: HS2_ + OH" + S22 _ + H20 ... (55) yielding the kinetically stable S2 » which would not be expected to react to form $ 2®3^~'^^ which is the predominant species produced below -0.4 VgHE' (a s s h o w n i n Appendix VI). At higher potential, above -0.4 VgH E, the smaller Tafel slopes observed in solution containing 0.1 and 1.0m Na2S203 yielded a tr ansmission c o e f f i c i e n t , X, equal to ~2, indicating that a double charged species i s passing through the double layer. This would be consistent with production of a poly-2— sulfide, for example S2 » via: 2HS- + 20H- S22 _ + 2H20 + 2e~ ... (56) This reaction is in agreement with the yellow coloration due to polysulfide produced in solution by anodic polarization of plat-inum above -0.260 VsHE' Appendix VI describes how current due to 139 oxidation to polysulfide exceeds that due to thiosulfate in this potential region. Previous investigators92 Qf polarization behavior in a l k a l i n e s u l f i d e s o l u t i o n s a t t r i b u t e d a n o d i c c u r r e n t s above -0.158 Vgjjg to o x i d a t i o n of s u l f i d e to p o l y s u l f i d e , which they confirmed s p e c t r o p h o t o r a e t r i c a l l y . The present r e s u l t s are con-s i s t e n t with o x i d a t i o n to p o l y s u l f i d e , but i n d i c a t e that there i s a n o t h e r major o x i d a t i o n p r o d u c t , i . e . t h i o s u l f a t e . T r o m a n s ^ observed no change in color due to polysulfides when a m i l d steel electrode was polarized above -0.5 VgHE' a n d concluded that sul-fide was oxidized to th i o s u l f a t e . According to the present findings, thiosulfate would be the pr i n c i p a l oxidized sulfur species produced at -0.5 VgH E. Higher potentials would be required to form significantly more polysulfide. 140 4.3.3 Polarization Behavior of A516 Steel in Alkaline Sulfide Solutions 4.3.3.1 Effect of a Low Concentration of Sulfide A small concentration of sulfide was added to NaOH solution to observe its effect. Cyclic vo1tammograms obtained at 50 mV/s in 2.5 m NaOH + 0.05 m Na2S are i l l u s t r a t e d in Figure 31. The stable cycle following several cyclic scans at each temperature has been shown, i.e. the diagrams for which further cycling produced no consistent change in behavior. At 95°C, peak I was composed of two peaks in which the second appeared as a shoulder on the f i r s t , and was larger than in the plain 3 m NaOH, Figure 29. Based on the E-pH diagram for Fe-H20, a smaller peak was expected in the sulfide-containing solution because the hydroxide concentration, 2.55 m, and there-fore the pH, was lower. Consequently, the sulfide must be responsible for the increased current. With respect to Figure 29, peak II was more than doubled in size, and increased in size with cycling, perhaps due to an increasing electrode area. Peaks III and IV were enlarged, also. At 120°C, Peak I current density was very high, and the maximum current varied ±0.5A/m2 from cycle to cycle in no apparent pattern. The maximum of peak II occurred at a slightly lower potential at 120°C than at 95°C, consistent with behavior in plain NaOH solution. Both peaks II and III grew with cycling. 141 Peak IV varied in the same manner as i t s conjugate, peak I. At 150°C, peak II disappeared. Apparently i t moved to a lower potential and was absorbed into peak I, as in the plain NaOH s o l u t i o n . Peak I I I d e c r e a s e d i n s i z e . P e a k s I and IV v a r i e d i n magnitude w i t h c y c l i n g as o b s e r v e d at 120°C. The p e a k s I, I I , I I I and IV w e r e l o c a t e d a t c l o s e t o t h e c o r r e s p o n d i n g peak p o t e n t i a l s i n the s u l f i d e - f r e e s o l u t i o n s and behaved similarly. It is reasonable to interpret them in terms of the same e q u i l i b r i a described in section A.3.1, recognizing that surface films in the sulfide solutions may be less protect-ive , thus accounting for the s h i f t in potential or increase in current density. Table XII summarizes the measured potentials and potentials calculated for various relevant e q u i l i b r i a for comparison purposes. Measured potentials were corrected for TLJP and for LJP because different solution (2.5 m NaOH + 0.5 m Na2^) was used in the bridge and external reference electrode. For the calculated potentials, the a c t i v i t i e s of Fe(0H)3 _ and Fe(0H)4 were approximated to 10-^m. The activity of S2O3 2"" was approx-imated to lO-^m, based on the impurity in the Na2S used. Peak I, in plain NaOH, was attributed to Fe304 formation, consistent with the E-pH diagram at high pH and temperature. However, in the sulfide solution, Fe(0H)3_ may also form FeS due to mass transfer effects, and this would be consistent with the o b s e r v e d d a r k e n i n g of the e l e c t r o d e . Thus a t peak I: 142 143 TABLE XII Potential of Cyclic Voltammetric Current Peaks and Relevant Equilibrium Potentials in 2.5 m NaOH + 0..05 m_Na2S a t 9 5» 1 2 0 a n d 1 5 0°C Peak 95 120 150 I E E+TLJP+LJP -0.762 -0.807 -0.725 -0.796 -0.675 -0.778 II E E+TLJP+LJP -0.530 -0.575 -0.575 -0.646 -VI E E+TLJP+LJP - ~-0.175 "-0.246 -0.238 -0.341 VII E E+TLJP+LJP - -0.025 -0.046 0.012 -0.091 III E E+TLJP+LJP -0.870 -0.915 -0.912 -0.983 -0.895 -0.998 IV E E+TLJP+LJP -1.037 -1.082 -1.125 -1.196 -1.162 -1.265 Equilibrium 2.10 2.5 2.52 2.13/2.16 -1.072 -1.010 -0.843 -0.894 -1.096 -1.041 -0.876 -0.939 -1.136 -1.084 -0.935 -0.914 Fe/Fe(0H)3~ Fe/Fe304 Fe(0H)3-/ Fe 3o 4 FeS/FeOOH or Fe 2o 3 144 TABLE XII (Cont) Equilibrium 95 120 150 2 . 7 1 / 2 . 7 2 - 0 . 7 7 9 - 0 . 7 9 9 - 0 . 8 5 9 F e ( 0 H )3~ / FeOOH or Fe203 2.11 -0.779 - 0 . 7 9 5 -0.777 FeS+HS~/FeS2 1.3 - 0 . 6 9 8 - 0 . 7 2 8 - 0 . 7 7 0 H S - / S90 o2-2.63 -0.714 -0.714 -0.727 Fe(0H)3-/ Fe(0H)A-2.27/2.28 -0.704 -0.739 -0.789 FeS2/FeOOH or F.go,+S203*-2.26/2.29 -0.664 -0.651 -0.747 FeOOH or Fe203+HS_/ FeS2 2.25 -0.632 -0.631 -0.660 FeS/J/e(0H)47+ S2032" (10- 4 ) 2.AO/2.41 -0.636 -0.653 -0.705 Fe304/FeOOH or Fe203 2.53 -0.431 -0.387 -0.310 Fe304/ FetOH)4-2.83 -0.402 -0.465 -0.463 FeS2/Fe(0H)4~ + S 2.76/2.77 -0.233 -0.244 -0.294 FeS 2/F e 0 0 H o r F e203 + S32" 1A5 Fe Fe(OH)3_ (2.52^ Fe304 (2.24) FeS The f o r m a t i o n o f FeS may have i n t e r f e r e d w i t h o x i d a t i o n o f F e ( 0 H )3- to F e 3 0 4 , m a k i n g Fe3C>4 defective. The " i n t e r c e p t a v e r a g e s " f o r the conjugate p e a k s I and IV we re - 0 . 9 7 5 , - 0 . 9 8 1 and -1.015 VgHE a t ^5, 120 and 150°C, respectively, corrected for TLJP and LJP. These were more active than in 3 m NaOH, consist-ent with the calculated values for equilibrium 2.52, which indicated the potential should have shifted just slightly to more active values in the lower pH sulfid e solution. The higher currents in sulfide solution would have been due to increased dissolution to Fe(0H)3~ through poorly formed Fe304' The Fe(0H)g- formed on the r i s i n g side of peak I may have reacted to form FeS via equation 2.24 i f log [Fe(0H)3-] exceeded -5.34, -5.10 and -6.34 at 95, 120 and 150°C respectively. These concentrations could have been present at the electrode surface due to slow diffusion of Fe(0H)3- away from the electrode. If the difference in total current passed at peaks I and II in 3 m NaOH (section 4.3.1) and in 2.5 m NaOH + 0.5 m Na2S was due to FeS formation, then the thickness of FeS formed would have b e e n a t 9 5 ° C: 0 . 1 9 7 jim, a t 120°C: 0 . 4 7 4 ^ m a n d a t 150°C: 0.571 urn, calculated as described in Appendix V. These are 146 thicker layers than calculated for oxide films in plain NaOH, but are in accord with the heavy deposit observed on electrodes. Probably the thickness was less than calculated because poorly adherent FeS may have f a l l e n off the electrode surface. There was a considerable difference in charge passed during the forward and reverse cycles, indicating that much of the oxidation product was removed from the surface. The peak and shoulder on peak I at 95°C, and also observed on the f i r s t cycle at 120°C and 150°C was indicative of two reactions occurring. The smaller, more noble peak may have resulted from the formation of an hydrated oxide. This poss-i b i l i t y could not be evaluated for lack of thermodynamic data. A l t e r n a t i v e l y , this peak, which was not observed in plain NaOH, was due to oxidation of FeS to FeOOH or Fe203 via equation 2.13 or 2.16 or to FeS2 via equation 2.11. This shoulder w i l l be referred to as peak IA in future. The variation in size of peak I from cycle to cycle at 120 and 150°C could be explained as due to random changes in mass transfer of Fe(0H)3- away from the electrode and thus changes in the potential at which FeS formed, or at which passivation took place. Peak IV would have been affected s i m i l a r l y due to v a r i -ation in the thickness of the surface f i l m which must be reduced. The increase in size of peak I compared with the peak in plain NaOH, Figure 29, was in contrast with the decrease observed by Shoesmith et al92 in alkaline s u l f i d e solutions of pH ~12 at 147 25°C. Perhaps, because Fe(0H)3~ would not be formed significant-ly in the lower pH solutions according to the E-pH diagram, their results would not reflect the mass transport controlled reaction to form FeS. On the other hand, Salvarezza et a l ^ ^ observed peak I to increase in size when sulfide was added to solutions of similar pH, in agreement with the present results. Peak II was considered to be the same as in plain NaOH, due to Fe(0H)3- -> FeOOH or Fe203 (eq. 2.71 and 2.71). The s h i f t of the peak potential to lower values as temperature was increased was consistent with behavior in plain NaOH as predicted for reaction 2.71 and 2.72. Previous i n v e s t i g a t i o n s9 2* 1 1 9 at lower pH showed peak II only in sulfide-containing solutions, but not in plain NaOH solutions. The concentration of Fe(0H)3- may not have been high enough for detectable current resulting from reaction 2.71 or 2.72. In the present work, higher currents were observed also in the sulfide solution compared with the plain NaOH. The increased size of peak II in the sulfide solution may have been due to higher rates of dissolution through a defective passive f i l m c o n taining s u l f i d e . Oxidation of Fe(OH)3~ to Fe(OH)4~ (eq. 2.63), and HS- to S2032~ would contribute to the current, also. FeS2 may have been oxidized to FeOOH or Fe203 (eq. 2.27 or 2.28). This oxidation could not occur until S203 was stable. Shoesmith et al92 i d e n t i f i e d peak II with deposition of 148 s u l f u r o n t h e s u r f a c e . W h i l e t h i s may h a v e b e e n p l a u s i b l e a t t h e l o w e r p H , i t i s l e s s l i k e l y a t h i g h e r p H a t t h i s p o t e n t i a l , a n d i n d e e d t h e i r r e s u l t s d o n o t p r o v e t h a t s u l f u r i s n o t a s s o c -i a t e d w i t h i r o n a s a s u l f i d e . M o r e o v e r , c a l c u l a t i o n o f t h e p o t e n t i a l f o r e q u a t i o n 2.83, FeS2 •+ F e ( 0 H ) 4 _ + S, i n d i c a t e d t h a t s u l f u r w o u l d n o t b e d e p o s i t e d u n t i l t h e p o t e n t i a l i s m u c h m o r e n o b l e t h a n p e a k I I . T h e d e p o s i t e d s u l f u r w o u l d be e x p e c t e d t o d i s s o l v e t o f o r m p o l y s u l f i d e s . S a l v a r e z z a e t a l 1 ^ w e r e u n a b l e t o d e t e c t s u l f u r o n t h e i r e l e c t r o d e s u r f a c e a n d a l s o e x p r e s s e d r e s e r v a t i o n s a b o u t t h e r o l e o f d e p o s i t e d s u l f u r i n t h e f i l m a t t h e s e p o t e n t i a l s . I t may h a v e b e e n t h a t , a s T r o m a n s ^ p r o p o s e d , s u l f i d e w a s i n c o r p o r a t e d i n t o t h e f i l m a t t h i s p o t e n t i a l , a n d w e a k e n e d i t , t h u s e n h a n c i n g d i s s o l u t i o n . T h e i m p a i r m e n t o f p a s s i v a t i o n by s u l f i d e s h a s b e e n r a t i o n a l -i z e d t h u s f a r b y c o n s i d e r i n g t h e i n t e r m e d i a t e F e ( 0 H ) 3 ~ t o f o r m F e S i n s t e a d o f F e 3 0 4 o r F e 2 0 3 « S h o e s m i t h e t a l 9 2 h a v e p r o p o s e d a d i f f e r e n t e x p l a n a t i o n . T h e y o b t a i n e d e v i d e n c e s u p p o r t i n g t h e t h e o r y t h a t o x i d e f i l m g r o w t h o c c u r s by t h e t r a n s f e r o f o x y g e n f r o m 0 H ~ t o t h e o u t e r m o s t l a t t i c e l a y e r o f m e t a l o x i d e . T h e r a t e o f o x i d e g r o w t h w o u l d t h e n b e p r o p o r t i o n a l t o t h e n u m b e r o f s u r f a c e s i t e s o c c u p i e d by a d s o r b e d 0 H ~ . T h e r a t e o f o x i d e t h i c k -e n i n g was r e d u c e d by t h e p r e s e n c e o f H S - , s u g g e s t i n g t h a t i t was s t r o n g l y a d s o r b e d o n t h e s u r f a c e , i n c o m p e t i t i o n w i t h t h e 0 H ~ . T h e c u r r e n t d e n s i t y a t p o t e n t i a l s n o b l e t o p e a k I I w a s h i g h e r t h a n i n t h e p l a i n N a O H , a n d may h a v e r e s u l t e d f r o m c o n -t i n u i n g d i s s o l u t i o n t h r o u g h a p o o r l y f o r m e d p a s s i v e f i l m 149 containing s u l f i d e . It was l i k e l y that these currents were related to some reaction involving both iron and sulfide because similar currents were not observed for steel in plain NaOH (iron alone) or for platinum in sulfide solution (sulfide alone). Reactions 2 .83 (FeS2 + F e ( 0 H ) 4 ~ +S) and 2 .25 (FeS -»• F e ( 0 H ) 4 ~ ) would have resulted in the soluble F e ( 0 H ) 4 _ species. A resultant soluble species would have been consistent with the absence of a conjugate peak on the reverse scan. The current due to oxidation of HS- to $2®3^~ a n d ^x^~ c a n be seen as the anodic current observed within the same range during the reverse scan. Peak VI corresponded well to oxidation of FeS2 to S 32 - + FeOOH or Fe203 via equations 2.76 and 2.77, but the peak could not be identified positively. Peak VII could not be identified, although oxidation of an iron su l f i d e to FeOOH or Fe203 plus a polysulfide was possible. Neither peak VI nor VII have con-jugates which would be consistent with production of a soluble product, e.g. S 2-. On the reverse scan, anodic current due to H S - oxidation continued until the potential was decreased below -0.645, -0.670, and -0.800 at 95, 120 and 150°C respectively, closely correspond-ing to the equilibrium potential of reaction 1.3. Peak III could be attributed to Fe2Q3 -+ Fe304, equation 2.40 150 or 2.A1, as in p l a i n NaOH, and l i k e w i s e peak IV could be attributed to Fe203.->- Fe(OH)3~ •*• Fe, equations 2.52 and 2.10. Increased currents may have been due to reduction of possibly thicker oxide films formed in the sulfide-containing solution, plus reduction of FeS film (equation 2.3). In summary, the behavior of steel in NaOH solutions with a low concentration of sulfide was very similar to the behavior in plain NaOH, but with increased currents apparently due to inter-ference, by sulfide, with passivation, thus enhancing dissolution and the formation of iron sulfides. A.3.3.2 Anodic Polarization with a Medium  Concentration of Sulfide Anodic polarization curves (1 mV/s) were obtained for A516 steel in 2.5 m NaOH + 0.5 m Na2S, for comparison with previously published anodic polarization curves1' 17, 18, 23 a n (j a s a basis for the c y c l i c voltammetry study in the same solution. Figure 32 illustrates the results at 95, 125 and 150°C. At very active potentials the currents were cathodic, consistent with hydrogen evolution at the electrode. Peaks I and IA were observed, in agreement with the observations in low sulfide NaOH solutions. Much higher current densities were attained at peak IA in the medium concentration sulfide solution. A new peak, V, was observed. At 95°C, current densities at peaks 151 152 V and I were very low. In other tests this minor peak and inflection became more prominent i f the electrode was cathodical-ly p o l a r i z e d for 0.5 h p r i o r to the t e s t . This cathodic polarization may have reduced surface oxide film and exposed more surface, thus producing higher currents. These peaks were more pronounced at 125 and 150°C. Table X I I I l i s t s the peak potentials at each temperature, and the potentials corrected for TLJP. No corrections for LJP were necessary because the same solution was used in the external reference electrode and autoclave. The theoretical potentials were calculated for reactions occurring in the potential range of the observed peak using the concentrations of soluble species indicated. These calculated potentials have been tabulated for comparison with the measured peaks. Peak V was consistent with oxidation of Fe to FeS, via equation 2.6. The slow scan rate, 1 mV/s, employed here would have allowed more time for reduction of surface f i l m , thus enhancing subsequent oxidation. Extended periods of cathodic polarization prior to the test would have ensured complete reduc-tion of surface oxides and enhanced subsequent formation of FeS at peak V, as was observed at 95°C. 153 TABLE XIII Potentials of Polarization Current Peaks and Potentials of Relevant Equilibria at 95, 125 and 150°C in 2.5 m NaOH + 0.5 m Na2S Peak 95 • 125 150 V E -0.885 -0.885 -0.900 E+TLJP+LJP -0.935 -0.946 -1.001 T E -0.760 -0.785 -0.790 I E+TLJP+LJP -0.810 -0.846 -0.891 IA E E+TLJP+LJP -0.660 -0.710 -0.670 -0.731 -0.685 -0.786 Equilibrium 2.3 -1.120 -1.076 -1.153 Fe/FeS 2.10 -1.093 -1.119 -1.149 Fe/Fe(0H)3-2.5 -1.030 -1.055 -1.093 Fe/Fe304 2.52 -0.952 -0.996 -1.056 Fe(0H)o-(io-5)7 Fe304 2.52 -0.841 -0.879 -0.930 Fe(0H)3-(io-6)7 Fe304 2 .13 or 2.16 -0.857 -0.883 -0.848 FeS/FeOOH or Fe 2o 3 2 .71 or 2.72 -0.853 -0.884 -0.943 FeOOHor Fe203 2 .71 or 2.72 -0.779 -0.805 -0.859 Fe(0H)o-(10-6)7 ! FeOOH or Fe203 154 TABLE XIII (Cont) Equilibrium 95 125 150 2.11 -0.819 -0.845 -0.823 FeS/FeS2 i 2.27/2.28 -0.703 -0.738 -0.789 FeS2/Fe00H2or (10-3) J | 2.26/2.29 -0.803 -0.799 -0.821 FeOOH or j Fe203 + HS-/ FeS2 1.3 -0.714 -0.745 -0.790 HS - /S2032" (10-3) 2.63 -0.721 -0.723 -0.736 Fe(0H}3" (10-o)/FeOH--( i o - 6 ) 2.25 -0.664 -0.672 -0.699 FeS/Fe(0H)4~ 2.40/2.41 -0.643 -0.668 -0.715 Fe304/FeOOH or Fe203 2.83 -0.411 -0.475 -0.475 FeS2/Fe(0H)4~ + S 2.40 or 2.41 -0.643 -0.668 -0.715 Fe304/FeOOH or Fe 2 o 3 2.83 -0.411 -0.475 -0.475 Fe(0H)4_ + S/ FeS2 155 The absence of peak V i n the 3 m NaOH and 2.5 m NaOH + 0.05 m N a 2 S w o u l d have been due t o t h e a b s e n c e or low c o n c e n t r a t i o n of s u l f i d e i n those s o l u t i o n s . At lower temperatures, peak V has been a t t r i b u t e d to Fe(0H)2 f o r m a t i o n ,7 2' 1 1 7 but F e ( 0 H ) 2 may no t be s t a b l e a t t h e s e temp-e r a t u r e s ^ and was n o t o b s e r v e d i n the p r e s e n t study o f c y c l i c v o l t a m m e t r y i n p l a i n NaOH. O t h e r s ^ ' 131 d i d not o b s e r v e t h i s peak, perhaps due t o a reduced p e r i o d o f c a t h o d i c p o l a r i z a t i o n p r i o r to the t e s t . Wensley and C h a r l t o n l ^ a t t r i b u t e d the peak to o x i d a t i o n of s u l f i d e s p e c i e s on the e l e c t r o d e s u r f a c e . Peak I c o r r e s p o n d e d w i t h e q u a t i o n s 2.5 ( F e / F e ^ O ^ ) , 2.10 ( F e / F e ( 0 H )3-) and 2.52 ( F e ( 0 H )3" / F e30 4 ) as i n p l a i n NaOH. The l o n g e r t i m e s s p e n t at p o t e n t i a l s a c t i v e to the peak maximum i n t h e s e low scan r a t e p o l a r i z a t i o n t e s t s may have r e s u l t e d i n h i g h e r F e ( 0 H )3~ c o n c e n t r a t i o n , c a u s i n g o x i d a t i o n v i a e q u a t i o n 2.52 to o c c u r at l o w e r p o t e n t i a l . Peak I was not a double peak as i n voltammograms i n the low s u l f i d e c o n c e n t r a t i o n s o l u t i o n because peak IA was f u l l y s e p a r a t e d from peak I. C u r r e n t s f o r r e a c t i o n of FeS to FeOOH or F e203 ( e q u a t i o n 2 .13 or 2 .16) or to FeS2 (equation 2 .11) , now at peak IA, would have been i n c l u d e d i n peak I i n the low s u l f i d e c o n c e n t r a t i o n s o l u t i o n . Thus, the higher peak I p o t e n t i a l s measured i n those low s u l f i d e s o l u t i o n s may have r e s u l t e d because peak I i n c l u d e d peak IA. At temperatures 95 , 125 and 150°C, FeS may have formed from F e ( 0 H )3- , e q u a t i o n 2.24, i f l o g [ F e ( 0 H )3~ ] exceeded - 5 . 1 4 , -5.06 156 and -6.76 r e s p e c t i v e l y . These c o n c e n t r a t i o n s c o u l d have been p r e s e n t n e a r t h e s u r f a c e due t o l i m i t e d mass t r a n s f e r . D i s s o l u t i o n t o F e ( O H )3- w o u l d have been e x p e c t e d as t h e c u r r e n t i n c r e a s e d t o t h e peak, and FeS f o r m a t i o n w o u l d have o c c u r r e d as soon as the l o c a l c o n c e n t r a t i o n exceeded the c a l c u l a t e d amounts. Above t h e peak p o t e n t i a l , t h e p r o d u c t i o n of Fe(0H)3"~ may have been i n h i b i t e d by an improved p a s s i v e l a y e r formed v i a o x i d a t i o n of FeS t o FeOOH or Fe203 v i a e q u a t i o n 2.13 or 2.16. P e a k IA a g r e e d f a i r l y w e l l w i t h FeS •+• FeOOH o r F e 2 0 3 ( e q u a t i o n 2.13 or 2.16). In the medium s u l f i d e s o l u t i o n , peak IA i n c r e a s e d i n s i z e s u b s t a n t i a l l y c ompared w i t h l o w s u l f i d e s o l u t i o n , s u g g e s t i n g t h a t f o r m a t i o n of FeOOH or Fe203 might have been i n h i b i t e d by c o m p e t i t i v e a d s o r p t i o n of s u l f i d e ( t o produce F e S ) , as a l r e a d y d e s c r i b e d f o r t h e b e h a v i o r i n l o w s u l f i d e s o l u t i o n , t h u s a l l o w i n g i n c r e a s e d d i s s o l u t i o n t h r o u g h a non-p r o t e c t i v e f i l m . Some of t h e c u r r e n t a t peak IA may a l s o have been due t o o x i d a t i o n o f FeS t o FeS2 v i a e q . 2.11. O x i d a t i o n of F e S2 t o FeOOH or Fe203 v i a e q u a t i o n 2.17 or 2.28 may have been o c c u r r i n g i n a d d i t i o n to eq. 2.13/2.16, (FeS -»• FeOOH or Fe203) t o c a u s e p a s s i v a t i o n a t t h e peak IA maximum. T h i s r e a c t i o n w o u l d o c c u r when S 2 O 32 - i s s t a b l e , i . e . above t h e e q u i l i b r i u m p o t e n t i a l of r e a c t i o n 1.3. Peak IA was c o n s i s t e n t at a l l t e m p e r a t u r e s w i t h o x i d a t i o n of HS- t o S 2 0 32 -, e q u a t i o n 1.3. At peak I A , t h e o x i d a t i o n of HS~ 157 apparently caused the surface f i l m to become more protective, apparently by allowing oxidation of FeS2 to FeOOH or Fe203 via eq. 2.27 or 2.28. Some oxidation of FeS, remaining for kinetic reasons, may also occur via eq. 2.13 or 2.16. Tromans^ determined that sulfide was incorporated in the Fe304 film, and he suggested that the film would become protect-ive only at more noble potential, where the sulfide content was decreased by Fe203 formation and the sulfide was oxidized (to ^2^3^" a n d Sx2 -) . The present results are consistent with this theory. The oxidation of the sulfide would have enhanced the film's protectiveness by reacting with the FeS2 to form FeOOH or Fe203i thus reducing further dissolution. Shoesmith et a 192 observed NaFeS2 (i.e. FeS2 ) formation as the current increased in peak IA. They detected no sulfur in films formed at potentials active to the beginning of peak IA, and claimed that sulfur or sulfide was incorporated into the film at peak IA. It i s possible that sulfur was present in the f i l m as FeS rather than as S (which they suggested). Their theory that FegO^ was attacked by H+ and NaFeS2 was ultimately formed, is untenable at the high pH of the present solutions. It has been suggested already that FeS may form FeS2 at peak IA via eq. 2.11. If this FeS is oxidized only to FeS2~ via: FeS + HS" + OH" -c FeS2~ + H20 + e- ... (57) 9 2 then the observation of FeS2~ formation by Shoesmith et al 158 could be accounted for. Reaction 57 could be decreased dramatic-a l l y when the potential i s raised above that for oxidation of HS~ to S20 32 -» which would remove the reactant HS- from the electrode surface i f the oxidation of HS- to S 2 0 32 - were much faster than oxidation of FeS to FeS2_* At high temperatures, FeS2~ is considered to be unstable^2, decomposing to ferrous iron and polysulfide. On the other hand, Raudsepp67 detected FeS2~ at 100 and 200°C in alkaline sul f i d e solutions and thought i t might be involved in a leaching process. This FeS2~ caused solutions to be green in color, as observed in pulp mill "green liquor". Precipitated iron sulfides, observed in the autoclave, may have resulted from the unstable FeS2-' I t s instability does not preclude i t s formation and subsequent disproportionation. The thermodynamics of FeS2~~ production could not be evaluated because no thermodynamic data were available for this species. The only comparable reaction was 2.11, for formation of FeS2 fr o m FeS. In the sulfide solution, either OH- or HS~ could adsorb on surface reaction sites. If the former adsorbed, then oxidation of Fe304 to Fe20"3 could occur. If HS- adsorbed, iron su l f i d e could be formed. As HS- concentration was increased, HS- adsorp-tion would be increased causing greater amounts of iron sulfide to be formed and making the fi l m less protective. Increased dissolution through a poorly protective f i l m would account for 159 the increased size of peak IA. The r i s i n g current above -0.5 VgH E was consistent with oxidation to S 2 O 3 2 - and Sx2 -' The polarization behavior in sulfide solution i s compli-cated. In addition to iron-water reactions, there are iron-sulfide and su1fide-water reactions. This complexity makes interpretation d i f f i c u l t and uncertain. Passivation apparently occurs when the iron sulfides are oxidized to FeOOH or ^ 2 0 3 . 4.3.3.3 Cyclic Voltammetry with a Medium Concentration of Sulfide Cyclic voltammograms (50 mV/s) in 2.5 M NaOH + 0.5 m Na2S at 95, 120 and 150°C are i l l u s t r a t e d in Figure 33, uncorrected for TLJP. The large number of peaks appearing at faster scan rate i l l u s t r a t e the complexity of the i r o n - s u l f i d e interactions and the strong influence of kinetics on them. Table XIV summarizes the peak potentials measured at each temperature and corrected for TLJP. The calculated values for relevant e q u i l i b r i a have been tabulated for comparison with the peaks. Peaks V, I and IA were shifted from the values measured in the polarization study in the direction of the scan, as expected for the higher scan rates used in the cyclic voltammetry study. 160 -1.25 -1.00 -0.75 -0.50 POTENTIAL , V S H E F i g u r e 33. C y c l i c voltammograms f o r A516 s t e e l i n 2.5 m NaOH + 0.5 m Na 2S. 161 TABLE XIV Potentials of Cyclic Voltammogram Current Peaks and Potentials of Relevant Equilibria at 95, 120 and 150°C in 2.5 m NaOH + 0.5 m Na2S Peak 95 120 • 150 V E E+TLJP ' i -0.798 -0.874 -0.810 ' i -0.918 i I E E+TLJP -0.774 -0.824 -0.726 -0.802 -0.702 -0.810 IA E E+TLJP -0.637 -0.687 -0.637 -0.713 -0.637 -0.745 II E E+TLJP -0.510 -0.560 -0.563 -0.639 -0.550 -0.658 VIII E E+TLJP -0.690 -0.740 -0.702 -0.778 -0.690 -0.798 III E E+TLJP -0.964 -1.014 -0.911 -0.987 -0.940 -1.048 Equilibrium 2.3 -1.120 -1.090 -1.153 Fe/FeS 2.10 -1.093 -1.107 -1.149 Fe/Fe(0H)3-2.5 -1.030 -1.049 -1.093 Fe/Fe304 2.52 -0.841 -0.872 162 -0.930 Fe(0H)o-(10-6)7 F e3o4 TABLE XIV (Cont) Equilibrium 95 120 150 2.13/2.16 -0.857 I I -0.877 -0.848 FeS/FeOOH or Fe203 2.71/2.72 ! -0.779 I i i -0.799 -0.859 FeOOHor 1 Fe 2o 3 2.11 -0.819 -0.838 -0.823 FeS + HS"/ FeS2 ; 2.38 -0.788 -0.801 -0.827 Fe(0H)3_/FeS2 2.26/2.29 -0.803 -0.799 -0.821 FeOOH or F e ? 0 i + HS-/ F e s V 2.27/2.28 -0.703 -0.738 -0.789 FeS2/FeOOH or F e203 + , S2o32" (lO"3) 1.3 -0.714 -0.733 -0.790 HS-/S2032" ( 1 0 - 3 J 2.63 -0.721 -0.721 -0.736 Fe(0H)3-/ Fe(0H)4-2.25 -0.664 -0.665 -0.699 FeS/Fe(0H)4~ 2. AO/2.41 -0.643 -0.661 -0.715 Fe304/Fe00H or Fe203 2.53 -0.460 -0.418 -0.347 Fe304/ Fer0H) 4-2.83 -0.411 -0.A75 -0.475 FeS2/Fe(0H)4_ + S 163 Peak V was not observed at 95°C, and may have merged with peak I. This would account for the lower value observed for peak I at 95°C compared to 120 and 150°C. P e a k I A , w h i c h was p a r t o f peak I a t l o w s u l f i d e c o n c e n -tration, was shifted to higher potential, separate from peak I. P e a k II was a t approximately t h e same potentials as i n t h e weak s u l f i d e s o l u t i o n . Some i n d i s t i n c t p e a k s o c c u r r e d at potentials noble to peak II. Peak VIII, an anodic peak, appeared on the reverse (cathodic) scan. It decreased in size at 150°C, compared to that at lower temperature. This anodic peak was in addition to the anodic current for sulfide oxidation (HS- ->- S2O32 - and Sx2 -) . Peak V has been attributed already to Fe ->- FeS, and peak I has been ascribed to Fe ->- Fe (0H)3~ + Fe304. Peak IA was dis-cussed as part of peak I in the low sulfide concentration, and was attributed to FeS being oxidized to FeOOH or Fe203 (equation 2.13, 2.16). In the polarization study at medium sulfide concen-t r a t i o n , peak IA was also correlated w i t h oxidation o f F e S 2 to FeOOH or Fe203 (equation 2.27, 2.28) with some current due to oxidation of FeS to FeS2 (eq. 2.11). The peak potentials were more noble than those observed in the plain NaOH or l o w sulfide solution, perhaps due t o some interference by the sulfide in the formation of the passive film. 164 The s h i f t of potential could not be pH related, because the pH was the same in the 3 m NaOH and the 2.5 m NaOH + 0.5 m Na2S' Peak VIII, on the reverse cycle, appeared to result from the same anodic process as peak IA. The occurrence of an anodic peak on the reverse scan has been described as indicative of a react-ivation process involving charge transfer reactions which produce ionic products.H7 jn this case the reactivation process would have reactivated dissolution to Fe(0H)3~- Reactivation may have occurred when the potential decreased below that of reaction 2.27/2.28 (-0.703 VS H E) and FeOOH or Fe203 would have been re-duced to FeS2. This reduction would have occurred for a short time (~11 mVgH E) u n t i l the p o t e n t i a l of r e a c t i o n 1.3 (-0.714 VS H E) was reached, and S2032 - was reduced to HS-. Then oxidation of any FeOOH or Fe203 would occur via reactions 2.26 or 2.29 to yield FeS2. Whether the conversion of FeOOH to FeS2 i s an o x i d a t i o n r e a c t i o n (2.26/2.29) or a reduction r e a c t i o n (2.27/2.28) depends on the sulfide species which is stable at the potential of interest. Formation of FeS2 would weaken the film and allow dissolution. Reactions 2.26 and 2.29 would have been oxidation reactions between --0.714 and -0.803 Vg^E, in agreement with the observed anodic currents at peak VIII, however most of the current was probably due to dissolution through the non-protective FeS2 film. The decreased size of peak VIII at 150°C suggested that dissolution was being prevented by improved passive film at the 165 higher temperature, perhaps an Fe2C>3 f i l m instead of the FeOOH expected at lower temperature, or simply that a thicker f i l m formed (consistent with the high currents on the forward scan at high temperature). The Fe203 may have been converted to FeS2 less easily than FeOOH. Rickard°5 determined that the rate of s u l f i d a t i o n of FeOOH to FeS was controlled by the rate of dis-solution of FeOOH. Sulfidation to FeS2 may be controlled similarly. Perhaps dissolution of Fe 203 i s more d i f f i c u l t than FeOOH. If dissolution i s rate-controlling and Fe 203 dissolves slowly, then the passive film would remain somewhat protective after reactivation at 150°C, thereby lessening subsequent dis-solution currents. Peak III has been attributed to the reduction of FeOOH or Fe203 to Fe304 (section 4.3.1). It was much smaller at 150°C than at 120 or 95°C. A similar trend was observed in the low concentration sulfide solution. The differe n t behavior may be related to the change of stable ferric oxide/hydroxide from FeOOH to Fe203 between 120 and 150°C. The Fe203 may not undergo the same reduction at peak I I I , and may react less e a s i l y , as already suggested to account for the decreased size of peak VIII at high temperature. The s h i f t of peak III with temperature could not be rationalized. Possibly, some of the current at peak III resulted from reduction of iron sulfide via reaction 2.3. Peak IV was not observed, and must have been occurring in the hydrogen evolution region. The hydrogen evolution currents may have increased because iron sulfides are known to catalyze 166 the hydrogen peak IV and sulfur f i l m , The reduced size of peaks at 120°C compared with those at 95°C indicates that the material should be expected to passivate more easily at 120°C than at 95°C because less charge must be passed to move the corrosion potential from the active to the passive region. This is in accord with the observation that pulp digesters, which contain a similar solution, passivate easily at ~130°C.5 1 The onset of passivity in sulfi d e solutions appears to depend on Fe203 or FeOOH formation. Films formed in the Fe304 stability range contain iron sulfides which make them unprotect-ive. At higher temperature, behavior approaches that in sulfide-free solution, i.e. the Fe304 seems to be more protective. At 150°C, after Fe304 formed at peak I, some degree of passivation occurred. At lower temperature, FeOOH or Fe203 formation seemed to be necessary. evolution reaction. 1^ 2 Shoesmith et a l9 2 observed a attributed i t to cathodic stripping of a surface which could have been present at lower pH. 167 4 . 3 . 3 . 4 Effect of a High Concentration of Sulfide Figure 34 i l l u s t r a t e s the effect of a high s u l f i d e concen-t r a t i o n , 3 m Na2S, on the polarization behavior of A516 s t e e l . The potentials are uncorrected for LJP and TLJP. No NaOH was added because 3 m OH- concentration was produced in solution by the dissociation of S 2 - to HS- and OH-. Thus, the OH- concen-tration and pH were the same as in the 3 m NaOH and 2.5 m NaOH + 0.5 m Na2S solutions. The measured potentials of the current peaks, corrected for TLJP and LJP, are l i s t e d in Table XV for comparison with calculated equilibrium potentials. For the cal-culations, the activity of S 2 O 3 2 - was approximated to 6 x lO-^ m, the anticipated concentration based on the impurity in the Na2S used. A very large peak was observed, with a potential close to that of peak IA or II measured in the 2.5 m NaOH + 0.5 m Na2S solution. The current density was nearly 3 times that of peak I in the 2.5 m NaOH +. 0.5 m Na2S. At 9 5 Q C , the peak appeared to include peak IA, but at 120 and 1 5 0 ° C i t appeared to include peak II, also. Peak V increased in size with cycling, as was observed at lower HS~ concentrations. Peak VIII was a sharp spike which was especially prominent at 120 and 150°C, perhaps related to the larger peak IA/II in the forward scan at these temperatures. It i s uncertain whether IX i s a peak because i t was observed only at 150°C; i t may be only the appearance of a peak, caused by the adjacent anodic peak VIII. 168 I.IA.II •1.25 -1.00 -0.75 -O.50 -0.25 POTENTIAL ,V S H E Figure 34. Cyclic voltammograms for A516 steel in 3 m Na2S 169 TABLE XV Potentials of Cyclic Voltammogram Current Peaks and Potentials of Relevant Equilibria in 3 m Na?S at 95, 120 and 150°C  Peak 95 120 150 V E E+LJP+TLJP -0.887 -0.941 -0.870 -0.955 -0.850 -0.967 I.IA.II E E+LJP+TLJP -0.650 -0.704 -0.530 -0.615 -0.475 -0.592 VIII E E+LJP+TLJP -0.760 -0.814 -0.720 -0.805 -0.755 -0.872 IX E E+LJP+TLJP - --1.000 -1.117 III E E+LJP+TLJP -1.087 -1.141 -1.080 -1.165 -1.050 -1.167 Equilibrium 2.3 -1.067 -1.121 -1.186 Fe/FeS 2 .5 -1.165 -1.049 -1.093 Fe/Fe 30 4 2 .52 -0.942 -0.989 -1.056 Fe(0H).3" (io-*)/ Fe 30 4 2 .52 -0.831 -0.872 -0.930 Fe(0H) 3-(10-5)7 Fe 30 4 2 .11 -0.851 -0.853 -0.856 FeS + HS"/ FeS2 170 TABLE XV (Cont) Equilibrium 95 120 150 2.71/2.72 -0.842 -0.877 -0.943 FeOOHor Fe 2o 3 j 2.71/2.72 I -0.768 -0.799 -0.859 FeOOHor Fe203 i 1.3 -0.716 -0.753 -0.798 2.26/2.29 -0.777 -0.772 -0.807 FeOOHor Fe203 + HS-/ FeS2 2.27/2.28 -0.695 -0.729 -0.780 FeS2/Fe00H or Fe20a + s2°3 (6 x 10-3) 2.40/2.41 -0.640 -0.661 -0.715 Fe304/Fe00H or Fe2Q3 171 The major effect of a high concentration of sulfide was to increase greatly the size of peak I/IA/II. The higher concen-tratio n of sulfide may have interfered with passivation to a greater extent than at the lower concentrations. The Fe304 film produced at low potential would have contained large amounts of iron sulfide, FeS or FeS2» making i t nonprotective. Dissolution to Fe(0H)3- or Fe(0H)4~ through the non-protective f i l m would account for the very large currents. Passivation occurred at a potential noble to e q u i l i b r i a 1.3 and 2.27/2.28 where FeOOH or Fe203 would be stable. In conclusion, i t i s noted that sulfide interferes with Fe304 passivity in proportion to i t s concen-t r a t i o n , and the formation of FeOOH or Fe203 i s necessary for effective passivation to occur. 4.3.3.5 Effect of Scan Range In the cyclic voltammograms illustrated in section 4.3.3.3, the potential was cycled between an anodic switching potential of 0.25 V g g g £ ( 2 5 ) and a c a t h o d i c s w i t c h i n g p o t e n t i a l of -1.25 VgggE(25)' Changes in the voltammograms resulted when the anodic and cathodic switching potentials were changed to vary the scanning range. The maximum range of c y c l i n g was from -1.25 Vggg£^25) to 0.75 VgggE(25)' The effect of scan range was investigated in 2.5 m NaOH + 0.5 m Na2s t o 9 0°c o n ly » a n d similar behavior at higher temperature was assumed. The results are i l l u s t r a t e d of i n varying the cathodic endpoints of Figure 35. Peak V did not appear 172 the cycles u n t i l the Figure 35. Variation of Cathodic Switching Potential (labeled at start of each cycle) 2.5 m NaOH + 0.5 m Na2S, 90°C, 50 mV/s. 173 second cycle and then increased in size in the third cycle. This increasing peak size was observed also in the low concentration s u l f i d e solution, and was attributed to increase of electrode area during cycling. Peak V then decreased in cycle 4 because the electrode apparently was not sufficiently cathodically polar-ized after cycle 3 to reduce the surface passive f i l m . This decrease in peak V current continued in cycle 5 following a further increase of the cathodic switching potential. Peak IA was not affected by variation of the cathodic end-point of the previous scan. This indicated that the anodic process was not a film formation process involving film reduction and reformation with each cycle. Thus, the current of IA was probably due to formation of a soluble species, consistent with dissolution to FeCOH^- or Fe(0H)4- being responsible for most of the current. Peak IA was not observed in cycle 6 because cycle 5 was terminated at 0 VSSSE(25)> before S 2 O 3 2 - could be reduced to FeS2 (e cl* 2.26), reactivating dissolution. If FeOOH was being formed at peak IA, i t was apparently not being reduced f u l l y at peak I I I , otherwise peak IA would have been smaller in cycle 5, after cycle 4 terminated above peak III. This was consistent with peak IA being due almost e n t i r e l y to dissolution to soluble iron species through a poorly protective film containing iron sulfide. The size of peak VIII diminished as the cathodic endpoint 174 was increased, perhaps because passive films were not reduced f u l l y during the previous reverse scan and were better formed with each additional forward scan. Figure 36 illustrates the changes in the cyclic voltammogram when the anodic endpoint of the scans was varied. Peak V i n -creased in size with each cycle. Evidently, there was sufficient cathodic polarization to reactivate the electrode before each cycle, thus increasing FeS formation (eq. 2 .3 ) . Decreasing time was spent at anodic potentials where FeOOH was formed (eq. 2.40) and so less passive f i l m would have had to be reduced, thereby ensuring more complete reduction of surface films during cathodic polarization. The decreasing amount of anodization with each cycle caused peak VIII to increase in size through cycle 4. Surface film may have been weaker due to the decreased period of anodic p o l a r i z -a t i o n . Subsequent to reduction of t h i o s u l f a t e (eq. 1.3), breakdown of the f i l m (via equation 2.26) may have been fa s t e r . Peak VIII was unaffected by ending cycle 5 at the middle of peak IA, although dissolution did not recommence until a more active potential was reached on the reverse scan. This suggested that some reduction, perhaps of S2O32"" had to take place before oxid-ation of FeOOH to FeS2 (e1« 2.26) and subsequent dissolution through a weakened passive film. Peak III was disappeared during unaffected in cycles 1 through 4 cycle 5, apparently because the 175 , but almost passive film ro O > E < to z UJ o or rr z> o -1.25 •1.00 -0.75 -0.50 -0.25 POTENTIAL , V S H E 0.0 Figure 36. Variation of Anodic Switching Potential (labeled at midpoint of each c y c l e ) . 2.5 m NaOH + 0.5 m Na2S, 90°C, 50 mV/s. 176 which would have formed above -0.66 VgHE present results support the theory that peak III reduction and indicate that peaks III and V were not form. The was due to FeOOH not conjugates. The effects of variation of the anodic and cathodic switch-ing potentials were consistent with the interpretation of the polarization behavior described in previous sections. 4.3.3.6 Scan Rate The effect of varying the scan rate was investigated at 90, 120 and 150oC in 2.5 m NaOH + 0.5 m Na2S-Figure 37 illustrates the effect of varying the scan rate at 90°C. The current peaks were shifted in the direction of the scan by an amount proportional to the scan rate, but the s h i f t was almost negligible below 20 mV/s. Slower scan rate would allow more time for nucleation and growth of a f i l m , so passivation could occur at a lower potential. Except for peak VIII, peak current densities were lower at the slower scan rates. Peak VIII became larger and sharper as the scan rate was decreased. In some tests at 2 mV/s (not shown here), peak VIII attained a higher anodic current density than peak IA. The high current then decreased to zero after a decrease of potential of <10 mV. This was consistent with current at peak VIII being entirely due to oxidation of FeOOH to FeS2 via equation 2.26, or 177 178 with production of some passive film which stopped dissolution. During the reverse cycle, at slow scan rate, black solids were visually observed to be produced briefly at the electrode during peak VIII, confirming the production of iron sulfides at peak VIII, either via equation 2.26, or by sulfidation of Fe(0H)3-. Figure 38 i l l u s t r a t e s the effect of scan rate at 120OC. Again, peak VIII was sharper at lower scan rates. The effect of varying scan rate at 150°C i s i l l u s t r a t e d in Figure 39. Again, peaks were shifted s l i g h t l y in the direction of the scan in proportion to the scan rates. Peak VIII was not observed; passivation at this temperature may have been so effective that dissolution could not recommence s i g n i f i c a n t l y . Peak III almost disappeared, as was observed previously (section 4.3.3.3). The results of the investigation of scan rate were consis-tent with the interpretations discussed previously. 4.3.3.7 Polarization at the Switching Potentials Between Scans Cathodic polarization for 300S at -1.900 VsHE' before each forward scan improved the repeatability of the results as shown in Figure 40. Apparently, cathodic polarization was effective in restoring the electrode surface to the same state before each cycle by reducing any surface deposits. 179 -1.25 -1.00 -0.75 -0.50 -0.25 0.0 POTENTIAL,V S H E Figure 38. Effect of Scan Rate, 2.5 m NaOH + 0.5 m N a2s» 120°C. 180 T 1 1 1 1 r -1.25 -1.00 -0.75 -0.50 -0.25 0.0 POTENTIAL, VS H E Figure 39. Effect of Scan Rate, 2.5 m NaOH + 0.5 m Na2s» 150°C. 181 Figure AO. E f f e c t of Cathodic P o l a r i z a t i o n f o r 300S at the Cathodic Switching P o t e n t i a l Between Cycles. 2.5 m NaOH + 0.5 m N a 2S» 5 0 m V / S . 182 Figure 41 i l l u s t r a t e s the effect of anodically polarizing the electrode for 300s at 0.090 VgHE» t n e a n o d i c switching poten-t i a l , before each reverse scan. The location and size of peak III (FeOOH reduction) varied considerably after this p o l a r i z -a t i o n . T h i s e f f e c t may have been due to v a r i a t i o n i n the thickness or composition of FeOOH f i l m formed at the anodic switching potential. 4.3.3.8 Effect of Stirring In the cyclic voltammetry conducted in the autoclave, sol-utions were not s t i r r e d . At lower temperature, in t e s t s conducted in a Teflon c e l l instead of the autoclave, some stirring was effected by nitrogen purging. The influence, on the cyclic voltammogram, of vigorous stirring with a magnetic s t i r bar i s i l l u s t r a t e d in Figure 42. S t i r r i n g depressed peak V and increased peaks I, IA and VIII. Improvement of mass transfer by stirring reduced the amount of FeS formed at peak V, probably by causing some passivation by reaction 2.5 (Fe ->- Fe304) at lower potential. Improved mass transfer must have stimulated oxidation or dissolution at peak I, IA by increasing the supply of HS- or 0H_ for reactions 2.10, 2.1 1, 2.63, 2.25 or 2.26. The effect of stirring may account for much of the observed variation in behavior from test to test or the differences between tests conducted in the autoclave versus those in the Teflon c e l l with nitrogen purging. 183 POTENTIAL, V S H E Figure 41. Effect of Anodic P o l a r i z a t i o n for 300s at the Anodic Switching Potential Before the Reverse Scan. 2.5 m NaOH + 0.5 m Na2S, 50 mV/s. 184 POTENTIAL,VSHE Effect of St i r r i n g . 2.5 m NaOH + 0.5 m Na2S, 9 0 ° , 50 mV/ 185 4.4 Relevance to Practical Corrosion Problems The cyclic voltammetric behavior indicated that, while some degree of passive protection was afforded by Fe3<")4 in plain NaOH solutions, the formation of Fe203 was necessary for passivation in sulfide-containing solutions. Dissolution, which increased with sulfide concentration, was associated witli formation of iron sulfides which led to a less protective film composed of a mix-ture of iron oxides and iron s u l f i d e s . Although temperature increased the currents, i t had a lesser effect than sulfide. At higher temperatures, the behavior tended to be similar to that at lower sulfide concentrations. Passivation of pulp digesters should occur more readily at lower sulfide concentrations. This present work w i l l aid in design and operation of anodic protec-tion systems for digesters by providing improved information concerning the potential required for passivation of the steel, and an improved understanding of electrochemical reactions in alkaline sulfide solutions. In addition, knowledge of the potential ranges for film breakdown and dissolution may assist in predicting ranges of SCC susceptibility.'''7 Repeated fast scanning gave fa i r l y reproducible results for a particular set of experimental conditions. Reproducibility from test to test at slow scan rates was poorer. The high scan rates used in this study may be more revealing of the non-equilibrium, kinetically controlled, film formation and breakdown at stress corrosion crack tips, where corrosion films are being ruptured by deformation processes and alternate cycles of f i l m 186 rupture and repassivation occur. Slower scan rate tests are closer to equilibrium conditions and may be less useful to the understanding of SCC behavior. The f i l m formed at high temperature seemed more r e s i s t a n t to breakdown, as i n d i c a t e d by the s m a l l e r peak VIII, ( s e c t i o n 4.3.3.3); t h i s may reduce s u s c e p t i b i l i t y to s t r e s s c o r r o s i o n c r a c k i n g , a l t h o u g h t h e r e i s a s c a r c i t y of q u a n t i t a t i v e d a t a at elevated temperatures. The high temperature polarization tests performed in the autoclave without s t i r r i n g may more closely simulate the poor mixing anticipated inside a stress corrosion crack. It is worth noting that the size of the active/passive peak (I/IA/II) was smaller in the unstirred than in the stirred condition, suggest-ing less dissolution inside a crack where mass transport i s poorer. 187 5. SUMMARY S e v e r a l a s p e c t s of the e l e c t r o c h e m i c a l b e h a v i o r of s t e e l were s t u d i e d : 1. E-pH diagrams f o r S-H20 and Fe-S-H20 systems at 25, 100 and 150°C and pH >8 were c o n s t r u c t e d w i t h u p - t o - d a t e thermo-dynamic data. They did not i n c l u d e S2~ because HS- was the s t a b l e s p e c i e s throughout the pH range. HFe02~ w a s r ep la c e d by FeCOH)^- because t h i s was c o n s i d e r e d more a p p r o p r i a t e . These d i f f e r e n c e s p r oduced s u b t l e changes i n the shape of s t a b i l i t y r e g i o n s f o r d i f f e r e n t s p e c i e s c o m p a r e d t o p r e v i o u s l y p u b l i s h e d E-pH d i a g r a m s . Diagrams f o r S - H 2 ^ » with p o l y s u l f i d e s and s u l f u r as the most o x i d i z e d forms of s u l f i d e , i n d i c a t e d r e g i o n s o f k i n e t i c s t a b i l i t y of Sx a t noble p o t e n t i a l s . 2. The response of the Ag/Ag2S e l e c t r o d e , SSSE, to v a r i a t i o n of t e m p e r a t u r e , s u l f i d e and h y d r o x i d e c o n c e n t r a t i o n , and c h l o r i d e a d d i t i o n was s t u d i e d . The SSSE was not r e l i a b l y p r e d i c t a b l e thermodynamically, but was s t a b l e and was proven e f f e c t i v e i n the p o l a r i z a t i o n s t u d i e s . The thermal l i q u i d j u n c t i o n p o t e n t i a l t h r o u g h an NaOH s o l u t i o n b r i d g e to an e x t e r n a l r e f e r e n c e e l c t r o d e was eva l u a t e d . 3. I n e r t e l e c t r o d e s , Hg and P t , were i n v e s t i g a t e d f o r use as ref e r e n c e e l e c t r o d e s . T h e i r e l e c t r o d e p o t e n t i a l s were found to be c o n s i s t e n t w i t h the mixed p o t e n t i a l between p o l y -188 s u l f i d e , Sx^- and t h i o s u l f a t e , S 2 O 3 . Below 60°C, the potential was close to that for H S-/ Sx2 -» but at higher temperature, the electrode potential was strongly influenced by the HS-/S2032_ equilibrium. Tafel slopes were obtained from anodic polarization curves of sulfide on Pt in alkaline sulfide solutions. Two Tafel slopes were measured on each curve and i d e n t i f i e d with oxidation to S 2 O 3 2 - at lower potential and to Sx2 _ at more noble potential. Reaction path mechanisms were proposed to account for the slopes. The high tempertaure electrochemical behavior of A516 carbon steel in alkaline sulfide solutions was studied by means of polarization tests and cyclic voltammetry. Dissolution and film formation increased with temperature and sulfide con-centration. The effects of variation of scan rate and scan range, and the effects of s t i r r i n g and polarization at the switching potentials between scans were investigated. Passivation in alkaline sulfide solutions occurred at a more noble potential than in plain NaOH and was consistent with formation of a protective Fe203 film associated with oxid-ation of HS- to S 2 O 3 2 - . The electrochemical behavior at high scan rates in unstirred solutions was considered to involve mass transfer effects. 189 BIBLIOGRAPHY 1. D. Singbeil and A. Garner, F i r s t Semi Annual Progress  Report of the Research Program to Investigate Cracking of  Continuous Digester, Pulp and Paper Research Institute of Canada, Montreal, July 1983. 2. K. Smith, Pulp and Paper 55 (10) 66-69 ( 1 9 8 1 ) . 3. L. Ruus and L. Stockman, Svensk Papperstidn 56 857 (1953). 4. L. Stockman and L. Ruus, Svensk Papperstidn 57 831 (1954). 5 . J. W. Hassler, TAPPI 38 265-274 ( 1 9 5 5 ) . 6 . B. Haegland and B. Roald, Norsk S k o g i n d 9 (10) 3 5 1 - 3 6 4 (1955). 7. B. Roald, Norsk Skogind 10 (8) 285-289 (1956). 8. W. A. Mueller, Can. J . Tech. 34 162-181 (1956). 9. W. A. 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A s t l e , e d s , CRC Handbook o f C h e m -i s t r y and P h y s i c s CRC P r e s s , Boca R a t o n , F l a . , 1982. 1 3 1 . D. L. S i n g b e i l , A S t u d y of the Stress C o r r o s i o n Cracking o f  M i l d Steel i n Alkaline and Alkaline Sulphide Solutions, M.Sc. thesis, University of B r i t i s h Columbia, July 1981. 132. R. L . M a r t i n and R. R. Annand, C o r r o s i o n 37 297-301 ( 1 9 8 1 ) . 197 APPENDIX I  The C r i s s - C o b b l e E x t r a p o l a t i o n Method The s t a n d a r d G i b b s e n e r g y of f o r m a t i o n o f a compound a t T2 may be w r i t t e n i n t e r m s of t h e f r e e e n e r g y of f o r m a t i o n , A G ° , and e n t r o p y of f o r m a t i o n , S ° , a t T^ and t h e h e a t c a p a c i t y of t h e compound, Cp, over the t e m p e r a t u r e i n t e r v a l T j to T2. T 2 A G °T 2 = A G °T i - S ^ ^ ( T2- T i ) - T2 \ Cp dT Tj_ T T2 + \ Cp dT ... ( 1 - 1 ) Tl The change i n Gibbs energy f o r a r e a c t i o n i s g i v e n by A G °T = E -?p A G ° - Z - J R A G ° . . . ( 1 - 2 ) 2 P T2 R T2 where ^ and -y^R are the s t o i c h i o m e t r i c c o e f f i c i e n t s Heat c a p a c i t i e s m easured o v e r a r a n g e o f t e m p e r a t u r e s a r e no t a v a i l a b l e , so some a p p r o x i m a t i o n i s n e c e s s a r y . C r i s s and C o b b l e1^ p r o p o s e d t h a t an a v e r a g e v a l u e of h e a t c a p a c i t y be ACp be used f o r t h e r e a c t i o n i n t h e t e m p e r a t u r e r a n g e o f i n t e r e s t . U s i n g t h i s and s u b s t i t u t i n g e q u a t i o n ( 1 - 1 ) i n t o ( 1 - 2 ) g i v e s : A G° = A G° - A S° AT + ACp AT - T2 A Cp l n (T2/T!) T2 T]_ Tx . . . ( 1 - 3 ) 1 9 8 For nonionic substances, ACp can be calculated directly from tabulated equations for the heat capacity. For ionic species, the average heat capacity may be given by: T2 -o -o T1 2 Al ST - ST . . . ( 1 - 4 ) In (T2/Ti) where S~o is the "absolute" entropy. S ° at 298°K has been defined as: S298 = S298 - 20.93Z J°K-1 mol"1 ... (1-5) where z i s the i o n i c charge. C r i s s and Cobble used the "absolute" entropy to estimate S ^ via their "correspondence principle" o o ST = aT + bT S 2 9 8 • • • (1-6) They tabulated values of a and b to be used in this equation.^8 These constants were determined for different classes of ions: simple anions and OH-, simple cations, oxyanions and acid oxy-anions. The calculation of equation 1-3 was performed with a pro-~ ,T2 graramable calculator. Values of T, AG°2Qg, S°2ag, Cp | , and 298 stoichiometric coefficients were input. Note that "conventional" 199 entropies are used in equation 1-3 rather than "absolute" entropies. The program also used the calculated AG value to evaluate the standard electrode potential, E0, at temperature. 200 APPENDIX II  Calculation of AG°20g and S°?Qft of S2~ A. Free Energy of Formation, AG°29g Using Giggenbach's37 value of K2 at 298°K, in the reacti on K2 HS" U H+ + S2" . . . ( I I - l ) AG = -2.303 RT log K2 ... (H-2) O O O AGH+ + AGS2_ - AGRS_ = - 2 . 3 0 3 RT log K2 (I I _ 3 ) substituting A G J J+ = 0 , AGjjg_ = 2.88 from NBS30 Note 2 7 0 - 3 and solving: AG§2- = 2 6 . 1 9 9 kcal mol B. Entropy, S°9Q« Cobblei Z O noted that AS*"* may be calculated for a reaction by either 3AG° AS° = [ ] ••• (II-*) 9T p or AS° = AH° - AG° ••• (I1"5) T where II-5 was preferred because i t involves an independent free 201 energy and enthalpy, while II-4 involves free energy differences over a necessarily small temperature interval which is subject to more error. Using values from Note 270-3 but substituting o the newly calculated value of AGg2- into ( A H S 2 _ + AHH20 " A H H S - " A H 0 H - ) = ( A G ° S 2 _ + A G ° H 2 0 " A G ° H S - - A G ° 0 H - ) + - 2 9 8 (S°S2- + S°H20 " S ° H S - " S ° 0 H - > and solving: S° = -22.638 cal/mol °K, S2 _ 202 APPENDIX III Standard Electrode Potentials of Electrochemical Equilibria and Free Energy Changes of Chemical Equilibria for the S-H2Q and Fe-S-H?0 Systems  Standard Electrode Potentials, E ° , for Electrochemical Equilibria for S-H2p_ Equilibrium ^°298 E°373 ^°423 1 . 3 0.225 0. 244 0.245 1.4 0.653 0.715 0.754 1.6 -0.062 -0.068 -0.092 1.7 0.583 0.636 0.668 1.8 0.513 0.557 0.581 1.9 0.038 0.045 0.030 1.10 0.077 0.091 0.080 1.11 0.159 0.185 0.183 1.12 0.411 0.473 0.499 1.13 -0.464 -0.524 -0.581 1.14 -0.551 -0.634 -0.718 1.15 -0.577 -0.660 -0.744 1.16 -0.596 -0.680 -0.765 1 .17 -0.601 -0.649 -0.682 1. 18 0.350 0.377 0.388 1.19 0.314 0.336 0.343 1.20 0.258 0.274 0.275 203 base don Standard Electrode Potentials, E ° , for Electrochemical Equilibria for S-H2O Giggenbach's Thermodynamic Values for Polysulfides Equilibrium ^°298 ^°373 ^°423 . 9 0 . 0 0 9 0 . 0 1 1 - 0 . 0 0 9 . 10 0 . 0 3 1 0 . 0 3 1 0 . 0 1 0 .11 0 . 1 0 6 0 . 1 1 1 0 . 0 9 2 .12 0 . 3 1 5 0 . 3 4 8 0 . 3 4 9 . 13 -0.347 -0.387 -0.435 .14 -0.325 -0.291 -0.279 .15 -0.568 -0.604 -0.649 .16 -0.523 -0.600 -0.680 .17 -0.696 -0.774 -0.832 . 18 0.369 0.400 0.413 .19 0.342 0.372 0.386 .20 0.285 0.311 0.321 .21 -0.343 -0.368 -0.404 .22 -0.538 -0.601 -0.670 204 Standard Electrode Potentials, E°V, of Electrochemical Equilibria for Fe-S-HpQ Equilibrium E°298 E°373 E°423 2 .3 - 0 . 5 8 3 - 0 . 5 9 0 -0.661 2.4 1.033 1 .078 1. 150 2 .5 -0.092 -0.081 -0.084 2.6 -0.096 - -2.7 -0.477 -0.436 -0.421 2.8 -0.213 -0.175 -0.163 2.9 0.138 - -2.10 0.470 0.553 0.618 2.11 -0.373 -0.363 -0.331 2.12 0.823 0.851 0.820 2.13 1.009 1.064 -2.14 0.382 0.408 -2.15 - - 0.435 2.16 - - 1.196 2.17 1.381 1.448 1.646 2. 20 0.371 0.428 0.486 2.22 0.410 0.452 0.494 2.25 0.904 1.003 1.103 2.26 -1.754 -1.789 -2.27 0.508 0.535 -2.28 - - 0.545 2.29 - - -1.857 2.30 -1.249 -1.268 -1.319 205 Equilibrium E°298 E°373 E°A23 2.31 0.520 0.547 0.557 2.32 -0.860 - -2. 33 -0.478 -0.516 -0.570 2.34 0.460 0.498 0.516 2.35 -0.743 -0.777 -0.829 2.36 0.548 0.585 0.602 2.37 -1.094 - -2.38 -1.425 -1.506 -1.610 2.39 0.687 0.742 0.779 2.40 0.264 0.297 -2.41 - - 0.295 2.42 0.797 0.757 -2.43 - - 0.717 2.44 0.268 0.234 -2.45 - - 0.199 2.46 -0.434 - -2.49 -0.080 - -2.50 1.064 0.987 0.928 2.51 -0.782 - -2.52 -1.777 -1.982 -2.192 2.53 4.034 4.649 5.203 2.60 1.525 1.685 1.835 2.62 0.823 - -2.63 0.160 0.228 0. 273 206 Equilibrium E°298 E°373 E°423 2 . 6 4 2 . 6 5 2 . 6 7 2 . 6 8 2 .71 2 . 72 2.73 2.74 2.75 2. 76 2.77 2.78 2.79 2.80 2.82 2.83 - 0 . 7 4 3 1 . 3 2 9 - 0 . 8 0 3 1 . 4 4 9 - 1 . 0 9 7 1.138 0.993 1.307 2.575 -3.010 2.054 0.920 - 1 . 2 2 3 1.177 1.031 1.369 2.734 -3.240 2.208 0.989 -1 .363 1.199 1.048 1.407 2.854 Not calc. Not calc. 1.039 207 Free Energy Change, AG ca l . , of Chemical Equilibria for Fe-S-H9,0 Equilibrium A G 2 9 8 A G3 7 3 A G4 2 3 2.18 -22460 - -2. 19 -4880 -7084 -11045 2.21 -17070 -19129 -22978 2 . 23 -33260 - -2. 24 -48550 -52724 -58994 2.47 -57960 -33457 -2.48 - - -75448 2.54 17580 - -2.55 5390 - -2.56 -10800 - -2.57 -26090 - -2. 58 -12190 -12046 -11933 2.59 -16190 - -2.61 -15290 - -2.66 -7380 - -2.69 -10160 - -2. 70 -2780 - -2 . 81 -31480 -33594 -36016 208 APPENDIX IV Relative Thermodynamic Stability of Ag?Q, AgCl and PLRQS DeBethune and Loudl29 have l i s t e d the following standard electrode potentials for reactions IV-1, IV-2 and IV-3: Ag 2S + H+ + 2e~ 2 Ag + HS", 2 AgCl + 2e~ 2 Ag + 2 C l " , Ag20 + H20 + 2e" ^  2 Ag + 2 OH", E° = - 0.66 VS H E ... (IV-1) Eo = 0.222 VS H E ... (iv-2) E° = 0.345 VS H E .•• (IV-3) To illustrate the stability of the Ag/Ag2S electrode in the presence of chlorides, consider the potential at which reactions IV-1 and IV-2 are in equilibrium. Equating the Nernst equations: 0.222 + 2.303 RT l o g 1 = _ 0.66 + 2.303 RT l o g laH+1 2F ( aCi - )2 2F ( aH S- ) . . .'(IV-4) Simplifying: (aH +) (aci")2 = (aHS") 6'8 x l°2 9 (IV~5) Thus, the product of the activities of hydrogen and chloride ions must be many orders of magnitude larger than than of bisulfide, in order to influence the electrode potential. Similarly, to illustrate the stability of the sulfide elec-trode in caustic solutions, the Nernst equations for reactions IV-1 and IV-2 are equated as: 209 0.345 + 2.303 RT io g 1 = -0.66 + 2.303 RT io g (aH+^ 2F (a0H") 2F (aHS") ... (IV-6) which i s : (aH+) (aOH-)2 = (aHS~) 9.9 x 1 03 3 ... (I V - 7 ) A g a i n , the pro d u c t of the a c t i v i t i e s of the hydrogen and h y d r o x y l i o n s must be many o r d e r s o f m a g n i t u d e l a r g e r t h a n t h a t o f t h e b i s u l f i d e i n or d e r t o i n f l u e n c e the p o t e n t i a l . 210 APPENDIX V Estimation of Oxide Film Thickness The thickness of oxide films formed on the electrode surface was estimated based on the charge required to form or reduce them. For example, the charge to reduce F e 3 0 4 i n peak I V could be measured and used to calculate the amount of Fe^O^ which was present on the surface prior to reduction. F i r s t , the current, i , measured on the cyclic voltammogram was plotted versus time, t, which was determined from the scan rate. This plot was done for the complete range of peak IV. The area under the curve, ^ i . d t , was equal to the charge passed, Q, in coulombs or A-sec. The area was measured graphically. The number of moles, M, of FegO^ reduced was: Q/Fn = M . . . ( V - l ) where F i s the Faraday number ( 9 6 5 0 0 coul./mole), and n i s the oxidation state of the F e 3 0 4 t 2 . 6 6 7 . The Fe-jC^ would have been deposited fa i r l y uniformly on the electrode surface of area, a. Thickness, d, was calculated by: M(M.W.) = d . . . ( V - 2 ) a (sp.gr.) 2 1 1 were M.W. i s the molecular weight and sp.gr. i s the s p e c i f i c gravity (x 106). Values of M.W. and sp.gr. for FeOOH, Fe203 a n d Fe304 were obtained from the Handbook for Chemistry and Physics.1 3 0 212 APPENDIX VI Oxidation of Sulfide to Thiosulfate and Polysulfide The mixed potential for an inert electrode, (Pt or Hg) i s l o c a t e d at a p o t e n t i a l s l i g h t l y below the e q u i l i b r i u m p o t e n t i a l for the polysulfides, i.e. HS_/SX2-. This mixed potential i s d e t e r m i n e d by the polysulfides and thiosulfate i n solution as d e s c r i b e d i n s e c t i o n 3.3.2. D u r i n g a n o d i c o x i d a t i o n of the s u l f i d e s o l u t i o n on a p l a t i n u m e l e c t r o d e , the p o t e n t i a l i s scanned upward slowly from the mixed potential and SH- i s oxidized to S 2 O 3 2 - . Referring to Figure VI-1, i t is noted that oxidation to Sx2 - w i l l occur when the potential is raised above the HS-/SX2 - e q u i l i b r i u m p o t e n t i a l , mV from the mixed potent i a l , and both Sx2 - and S2O3 2 "" w i l l be produced. It i s of interest to know the r e l a t i v e amounts of the two products. The proportions may be estimated using measured overpotentials and Tafel slopes, and considering the relationships between the elec-trode k i n e t i c behavior for oxid a t i o n to p o l y s u l f i d e and thiosulfate. The measured Tafel slope was ~145 mV/decade below -0.4 Vc^g. In the f o l l o w i n g arguments, i t w i l l be assumed that anodic and cathodic Tafel slopes are equal, i.e. t>a = bc. The Tafel slope for oxidation is S 2 O 3 2 - cannot be < 145 mV/decade, otherwise the Tafel slope measured above the mixed potential would have to be less than the measured slope, 145 mV/decade, and due to oxidation to Sx2 - or S2032 _ can't be > 145 mV/decade because the charge 213 transferred through the double layer per rate determining step, X would be < 1. Ions with fractional charges are unknown. CASE I Tafel Slopes for Oxidation and Reduction to Sx2~ and ^2^3^~ a r e Equal, b-^  = b2  This case is i l l u s t r a t e d in Figure V I - 1 . Ea nd E 2 are the e q u i l i b r i u m p o t e n t i a l s for oxi d a t i o n to Sx2 _ and S20 32 -respectively. The graphical construction shows that above the HS_/SX2 _ equilibrium potential, E]_, the current of oxidation to Sx2 _. At the equilibrium potential, E^, which is 7[ mV above the mixed potential, \7[ I = I bc l I log ( i M / io i ) ••• (V1-1) and = I ba 2 I log (is2032~/:LM) ••• (VI-2) If I bg 2 I = I bci I = b, then these equations may be summed: 2 \7[ I = b log ( iS 20 3 2 - / i M ) + b log ( i M / io i ) ••• (VI-3) 2 | ^ | = b log ( iS 2o32- / i0l ) (VI-4) The ratio of S2032 - to Sx2 - at Ej was estimated. The Tafel slope, b, was the measured value of 145 mV/decade. For example, the overpotential, ^  , in 2.5m NaOH + 0.5m Na2S + 0.01m Na2S203 at 90°C was estimated to be the difference between the measured potenti a l , -0.599 Vg H E (corrected for LJP and TLJP) and the 214 theoretical polysulfide equilibrium potential -0.535 Vgjjg, from section 3.3.2, i.e. 0.064 V. The ratio 1 S2032~/io = 7 .6 was calculated via equation VI-4. Thus, most of the current i s due to oxidation to th i o s u l f a t e . This relationship w i l l be main-tained at high overpotentia 1 because the Tafel slopes are par a l l e l . CASE II Tafel Slopes for Sx2- Oxidation and Reduction are One-Half those for S2032 _, b_[ = 0.5 b2  If the rate determining step for oxidation of Sx2 - involves the movement of a doubly charged species through the double layer, the Tafel slope w i l l be halved, i.e. 72 mV/decade. This case i s illustrated in Figure VI-1. The section MX in the figure would have the slope of the measured Tafel slope at <-0.4Vgjjg, 145 mV/decade. Now: log (VI-5) log (VI-6) But b2 may be substituted in VI-5 and the equation multi-plied by 2 to obtain: 2 (VI-7) Summing VI-6 and VI-7 to eliminate ij^ : 215 3 | ^ | = b log ( iS 2032 - / i0 1) ... (VI-8) If 7^= 0.064 V and b =0.145 V/decade, then i s20 32~ /i0 1 = 2 1-Thus, at the polysulfide equilibrium p o t e n t i a l , the current due to oxidation of su l f i d e to thiosulfate w i l l be 21 times that due to p o l y s u l f i d e . However, at higher o v e r p o t e n t i a 1s, above X, oxidation to polysulfides w i l l overtake current for oxidation to t h i o s u l f a t e . To find the over po t e n t i a l , ^x> a t which the oxid-ation to S x 2 - becomes predominant, consider: ^ x = ba i log ( ix/ i0l ) = 0.5 ba 2 log ( ix/ i Q l ) (V I~9) and = ba 2 log ( ix/ iS 02-) ••• (VI-10) M u l t i p l y i n g VI-9 by 2 and s u b t r a c t i n g VI-10 to e l i m i n a t e ix y i e l d s : 7\ = ba 2 log ( is 02 - / i0 1) ... (VI-11) L 2 3 Equations VI-8 and VI-11 may be equated, and thus i n 2.5m NaOH + 0.5m Na2S + 0.01m Na2S203: 7[ x = 3?[ = 3(0.064) = 0.192V ... (VI-12) The current for o x i d a t i o n to Sx2 _ should equal the current for oxidation to S2032 - at point x, which i s : 216 = ^ x + ?l = 0.192 + 0.064 = 0.256V above the mixed p o t e n t i a l . Above t h i s p o t e n t i a l , e l e c t r o c h e m i c a l o x i d a t i o n to S>x^~ would be p r e d o m i n a n t . In the 2.5m NaOH + 0.5m N a 2 S + 0.01m N a 2 S 2 0 3 , t h i s p o i n t x s h o u l d be a t EMIXED + ^7 T = - 0 . 5 9 9 + 0 .256 = - 0 . 3 4 3 V g g g . T h e r e was no o b s e r v e d change i n the s l o p e i n t h i s s o l u t i o n , F i g u r e 29. S i m i l a r c a l c u l a t i o n s f o r t h e s o l u t i o n s c o n t a i n i n g 0.1 and 1.0m Na2S203 y i e l d e d Ex of -0.253 and -0.328 VgHg r e s p e c t i v e l y , c o r r e s p o n d i n g t o t h e e x p e r i m e n t a l l y m e a s u r e d p o i n t s -0.379 and -0.399 ^SHE r e s p e c t i v e l y . The agreement of measured and t h e o r e t i c a l v a l u e s f o r E„ i s s u f f i c i e n t l y good to c o n c l u d e t h a t the o x i d a t i o n t o HS conforms to the proposed r e a c t i o n s . The l a c k of s l o p e change i n t h e 0.01m Na2S203 s o l u t i o n , and t h e r e l a t i v e shape of t h e c u r v e i s c o n s i s t e n t w i t h mass t r a n s p o r t e f f e c t s , i . e . s m a l l e r a n o d i c c u r r e n t s . 217 LOG CURRENT DENSITY 2 1 8 

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