• TABLE OF CONTENTS
HIDE
 Title Page
 Dedication
 Acknowledgement
 Table of Contents
 List of Tables
 List of Figures
 Abstract
 Introduction
 Literature survey
 Experimental procedure
 Results and discussion
 Conclusions
 Recommendations for further...
 Appendices
 Bibliography
 Biographical sketch














Group Title: significance of the "protection potential" for Fe-Cr alloys at room temperature
Title: The significance of the "protection potential" for Fe-Cr alloys at room temperature
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Permanent Link: http://ufdc.ufl.edu/UF00098192/00001
 Material Information
Title: The significance of the "protection potential" for Fe-Cr alloys at room temperature
Physical Description: xxxiii, 357 leaves. : illus. ; 28 cm.
Language: English
Creator: Starr, Kenneth Kirch, 1944-
Publication Date: 1973
Copyright Date: 1973
 Subjects
Subject: Chromium-iron alloys   ( lcsh )
Chromium alloys -- Corrosion   ( lcsh )
Metallurgical and Materials Engineering thesis Ph. D
Dissertations, Academic -- Metallurgical and Materials Engineering -- UF
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
 Notes
Thesis: Thesis--University of Florida.
Bibliography: Bibliography: leaves 342-356.
General Note: Typescript.
General Note: Vita.
 Record Information
Bibliographic ID: UF00098192
Volume ID: VID00001
Source Institution: University of Florida
Holding Location: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: alephbibnum - 000585195
oclc - 14198788
notis - ADB3827

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Table of Contents
    Title Page
        Page i
        Page ii
    Dedication
        Page iii
    Acknowledgement
        Page iv
    Table of Contents
        Page v
        Page vi
        Page vii
        Page viii
        Page ix
        Page x
    List of Tables
        Page xi
        Page xii
        Page xiii
        Page xiv
    List of Figures
        Page xv
        Page xvi
        Page xvii
        Page xviii
        Page xix
        Page xx
        Page xxi
        Page xxii
        Page xxiii
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        Page xxv
        Page xxvi
        Page xxvii
        Page xxviii
        Page xxix
        Page xxx
        Page xxxi
    Abstract
        Page xxxii
        Page xxxiii
    Introduction
        Page 1
        Page 2
        Page 3
        Page 4
    Literature survey
        Page 5
        Page 6
        Page 7
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    Experimental procedure
        Page 67
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        Page 73
        Page 74
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        Page 78
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    Results and discussion
        Page 80
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    Conclusions
        Page 176
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    Recommendations for further research
        Page 181
    Appendices
        Page 182
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    Bibliography
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    Biographical sketch
        Page 357
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Full Text

















THE SIGNIFICANCE OF THE "PROTECTION POTENTIAL"
FOR Fe-Cr ALLOYS AT ROOM TEMPERATURE








By



KENNETH KIRCH STARR


A DISSERTATION PRESENTED TO THE GRADUATE
COUNCIL OF THE UNIVERSITY OF FLORIDA IN PARTIAL
FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY


UNIVERSITY OF FLORIDA
1973
















































) 1974

KENNETH KIRCH STARR



ALL RIGHTS RESERVED




























Dedicated to My Wife,

Barbara Jo











ACKNOWLEDGMENTS


I would like to express my thanks to the chairman of my

supervisory committee, Dr. Ellis D. Verink, Jr., for his guidance,

inspiration and patience. Thanks are also extended to the other

members of the committee: Dr. G. M. Schmid, Dr. R. T. DeHoff and

Dr. J. J. Hren. Special thanks are extended to Dr. M. Pourbaix for

his inspiration and helpful suggestions at several stages of this

investigation.

Technical assistance was provided by Mr. P. D. Kalb, Mr. C. J.

Minier, Mr. C. M. Simmons, Mr. E. C. Logsdon and Mr. H. W. Willis.

I wish to acknowledge the financial support supplied by the

National Science Foundation in the form of an N.S.F. Traineeship. I

would also like to acknowledge the financial support supplied by the

Advanced Research Project Agency under Contract Number N00014-68-A-0173-

0003, administered by the Office of Naval Research, Washington, D.C.,

and the financial support supplied by the Office of Naval Research

under Contract Number N00014-68-A-0173-0015.

The Fe 12% Cr alloy was donated by the Naval Research Laboratory

and the other Fe-Cr alloys were donated by the United States Steel

Corporation. These contributions are gratefully acknowledged.

The author is also grateful to his parents, for their help

and encouragement, and to his wife, who supported him in many ways.













TABLE OF CONTENTS

Page

ACKNOWLEDGMENTS............................................. iv

LIST OF TABLES .............................................. xi

LIST OF FIGURES............................................. xv

ABSTRACT .................................................... xxxii

CHAPTERS:

1. INTRODUCTION....................................... 1

2. LITERATURE SURVEY .................................. 5

2.1 Methods of Investigation ...................... 6

2.1.1 Immersion Tests......................... 7

2.1.2 Galvanostatic Tests.................... 7

2.1.3 Potentiostatic Tests................... 8

2.1.4 Potentiokinetic Tests.................. 8

2.2 Passivity...................................... 9

2.3 Aggressive Anions ............................. 12

2.4 Pitting Corrosion.............................. 13

2.4.1 Pit Initiation.......................... 14

2.4.1.1 Breakdown of passivity:
adsorption theory............. 14

2.4.1.2 Breakdown of passivity:
oxide-film theory............. 14

2.4.1.3 Breakdown of passivity:
combined theory................ 15

2.4.1.4 Passive film breakdown
parameters..................... 16

2.4.1.5 Critical pitting potentials... 17

2.4.1.6 Critical chloride ion
concentrations................ 19











TABLE OF CONTENTS (Continued)


2.4.1.7 Effect of pH..................

2.4.1.8 Effect of alloy composition...

2.4.1.9 Induction periods.............

2.4.2 Propagation of Pitting.................

2.4.2.1 Geometry of pitting...........

2.4.2.2 Current distribution in and
around pits...................

2.4.2.3 Potential distribution in
and around pits...............

2.4.2.4 Electrolyte composition
inside pits...................

2.4.2.5 pH changes in pits............

2.4.2.6 Chemical analysis of artifi-
cial pits.....................

2.4.2.7 Hydrolysis in pits...........

2.4.2.8 Electromigration and
diffusion.....................

2.4.2.9 Stability of pitting..........

2.5 Crevice Corrosion............................. .

2.5.1 Initiation of Crevice Corrosion.........

2.5.1.1 Metal ion cells...............

2.5.1.2 Differential aeration cells...

2.5.1.3 Passive-active cells..........

2.5.1.4 Critical parameter concept....

2.5.1.5 Induction periods..............

2.5.2 Propagation of Crevice Corrosion.......

2.5.2.1 Geometry of crevice
corrosion.....................


Page

20

21

22

22

23


25


27


28

29


30

31


32

34

36

37

37

39

39

40

42

42


43











TABLE OF CONTENTS (Continued)


Page

2.5.2.2 Current distribution in and
around crevices.............. 44

2.5.2.3 Potential distribution in
and around crevices............ 45

2.5.2.4 pH changes in crevices........ 48

2.5.2.5 Chemical analysis of crevices. 49

2.5.2.6 Hydrolysis in crevices........ 51

2.6 Limitations of Laboratory Tests............... 52

2.7 The "Protection Potential".................... 53

2.7.1 Definition and Verification............ 53

2.7.2 Experimental Methods ................... 54

2.7.3 Interpretation of the "Protection
Potential" for Cu-Rich Alloys.......... 55

2.7.4 Effects of Experimental and Environ-
mental Variables on the "Protection
Potential" ............................. 56

2.7.4.1 Effect of alloy composition
on Ep ......................... 56

2.7.4.2 Effect of pH on Ep.... ....... 58

2.7.4.3 Effect of [Cl ] on Ep......... 58

2.7.4.4 Effect of temperature on E ... 60

2.7.4.5 Effect of extent of propaga-
tion on E .................... 60

2.7.5 Effect of Geometry on E ... ............ 60

2.7.6 Material Vs. Environmental Property.... 63

2.7.7 "Repassivation" Vs. "Deactivation"..... 64

2.7.8 Experimental Correlation.............. 65












TABLE OF CONTENTS (Continued)


Page

3. EXPERIMENTAL PROCEDURE....... ..................... 67

3.1 Potentiostatic Polarization Experiments....... 67

3.2 Potentiokinetic Polarization Experiments...... 67

3.3 Galvanostatic Polarization Experiments........ 68

3.4 Artificial Occluded Cell Experiments.......... 71

3.4.1 First Type.............................. 71

3.4.2 Second Type............................. 74

4. RESULTS AND DISCUSSION .............................. 80

4.1 Potentiokinetic Polarization Results........... 80

4.1.1 Effect of Alloy Composition............ 89

4.1.2 Effect of Bulk Electrolyte pH.......... 89

4.1.3 Effect of Bulk Electrolyte Chloride
Ion Concentration...................... 97

4.1.4 Effect of Extent of Propagation......... 104

4.2 Galvanostatic Polarization Results............ 113

4.3 Artificial Occluded Cell Results.............. 119

4.3.1 First Type.............................. 119

4.3.2 Second Type............................. 122

4.4 Thermodynamic Calculations.................... 134

4.4.1 Potential Vs. pH (Pourbaix) Diagrams... 134

4.4.2 Hydrolysis Calculations................ 141

4.5 Metal Chloride Solution Scans................. 157

4.6 Potentiostatic Titration Results.............. 169

4.6.1 pH Titrations........................... 169

4.6.2 Metal Chloride Titrations.............. 170












TABLE OF CONTENTS (Continued)


Page

5. CONCLUSIONS........................................ 176

6. RECOMMENDATIONS FOR FURTHER RESEARCH ............... 181

APPENDICES................ .................................. 182

1. ALLOY CHARACTERIZATION................................ 183

2. STANDARD ELECTROCHEMICAL POLARIZATION CELL.......... 188

3. SAMPLE PREPARATION AND MOUNTING.................... 190

4. ELECTROLYTES....................................... 192

5. DEAERATION OF ELECTROLYTES................. ........ 195

6. LIST OF EQUIPMENT.................................. 196

7. ZERO-IMPEDANCE AMMETER............................ 198

8. EXPERIMENTAL POTENTIAL VS. pH DIAGRAMS AND
CHARACTERISTIC VALUES FROM CYCLIC POTENTIOKINETIC
POLARIZATION CURVES FOR Fe-Cr ALLOYS IN 0.1 M C1,
H2-SATURATED (OR 02-SATURATED) SOLUTIONS............ 200

9. SPECIAL CYCLIC POTENTIOKINETIC POLARIZATION SCANS
FOR THE Fe 12% Cr ALLOY IN 0.1 M Cl-, H2-
SATURATED SOLUTIONS................ ............ 228

10. SPECIAL CYCLIC POTENTIOKINETIC POLARIZATION SCANS
FOR THE Fe 12% Cr ALLOY IN 0.1 M CI-, H2-
SATURATED SOLUTIONS, USING PAINTED-OFF SAMPLES
APPROXIMATELY 1 cm2 IN AREA ........................ 245

11. EXPERIMENTAL POTENTIAL VS. LOG CHLORIDE ION
MOLARITY DIAGRAMS AND CHARACTERISTIC VALUES FROM
CYCLIC POTENTIOKINETIC POLARIZATION CURVES FOR
Fe-Cr ALLOYS IN H2-SATURATED SOLUTIONS OF NOMINAL

pH 5.4, 8.8 AND 10.8............................... 262

12. GALVANOSTATIC POLARIZATION EXPERIMENT DATA FOR THE
Fe 12% IN A 0.1 M Cl H2-SATURATED SOLUTIONS OF
NOMINAL pH 10.1 .................... 300

13. GALVANOSTATIC SINGLE-PIT POLARIZATION EXPERIMENT
DATA FOR THE Fe 12% Cr ALLOY IN 0.1 M Cl-, H -
SATURATED SOLUTIONS OF NOMINAL pH 10.2............. 303












TABLE OF CONTENTS (Continued)

Page

14. ARTIFICIAL OCCLUDED CELL DATA FOR THE Fe 12% Cr
ALLOY AND ARMCO Fe IN SOLUTIONS INITIALLY 0.1 M Cl 307

15. ANALYSES OF FINAL ARTIFICIAL OCCLUDED CELL
SOLUTIONS FOR THE Fe 12% Cr ALLOY AND ARMCO Fe
IN SOLUTIONS INITIALLY 0.1 M Cl-.................... 311

16. CARBIDE CALCULATIONS .............................. .. 313

17. CALCULATED EQUILIBRIUM POTENTIAL VS. pH (POURBAIX)
DIAGRAMS, THERMODYNAMIC DATA AND PERTINENT EQUATIONS 316

18. HYDROLYSIS CALCULATIONS AND RESULTS................. 331

19. CYCLIC POTENTIOKINETIC POLARIZATION SCAN DATA FOR
THE Fe 12% Cr ALLOY IN H2-SATURATED METAL
CHLORIDE SOLUTIONS ................................. 338

BIBLIOGRAPHY ................................................ 342

BIOGRAPHICAL SKETCH .................. .......................... 357











LIST OF TABLES


Table Page

1 Chemical Compositions of Commercial Stainless
Steels..................... ......................... 2

2 Details of Investigations of the "Protection
Potential" for Ferrous Alloys...................... 57

3 Extent of Propagation Data for the Fe 12% Cr
Alloy in 0.1 M Cl H2-Saturated Solutions of pH
Approximately 10.1................................. 109

4 Single-Pit Extent-of-Propagation Data for the Fe -
12% Cr Alloy in 0.1 M Cl-, H2-Saturated Solutions
of pH Approximately 10.2........................... 112

5 Solubility of Metal Chlorides...................... 145

6 Thermodynamic Data for Complex Ion Hydrolysis
Reactions......................................... 146

7 pH's of H2-Deaerated Metal Chloride Solutions...... 152

8 pH Titration Results: Fe-12% Cr Alloy.............. 171

9 Metal Chloride Titration Results: Fe-12% Cr Alloy,
0.1 M NaC1, H2-Saturated Initial Solutions.......... 173

10 Metal Chloride Titration Results: Fe 5.0% Cr,
Fe 12% Cr, Fe 16.9% Cr and Fe 24.9% Cr
Alloys 0.1 M NaC1, H2-Saturated Initial
Solutions, Holding Potential -0.250 v S(0.000
SHE) .................................................. 175

11 Compositions of Binary Iron-Chromium Alloys Under
Investigation.......... ........................... 184

12 0.1 M Cl Electrolytes ............................ . 193

13 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-0.5% Cr
Alloy in 0.1 M Cl-, H2-Saturated Solutions........ 213

14 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-2.0% Cr
Alloy in 0.1 M Cl-, H2-Saturated Solutions......... 215

15 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-5.0% Cr
Alloy in 0.1 M Cl-, H2-Saturated Solutions......... 217











LIST OF TABLES (Continued)


Table Page

16 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-12% Cr
Alloy in 0.1 M Cl-, H -Saturated Solutions........ 219

17 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-12% Cr
Alloy in 0.1 M Cl-, 02-Saturated Solutions........ 222

18 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-16.9% Cr
Alloy in 0.1 M Cl H2-Saturated Solutions........ 224

19 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-24.9% Cr
Alloy in 0.1 M Cl-, H2-Saturated Solutions........ 226

20 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-0.5% Cr
Alloy in H2-Saturated Solutions................... 287

21 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-2.0% Cr
Alloy in H2-Saturated Solutions................... 290

22 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-12% Cr
Alloy in H2-Saturated Solutions................... 293

23 Characteristic Values from Cyclic Potentio-
kinetic Polarization Curves for the Fe-16.9% Cr
Alloy in H2-Saturated Solutions................... 297

24 Galvanostatic Polarization Experiment Data for
the Fe-12% Cr Alloy in a 0.1 M Cl-, H2-Saturated
Solution of pH 10.1: Run No. KS-145 (Sample
Area 0.95 cm2)............................ ....... 301

25 Galvanostatic Polarization Experiment Data for
the Fe-12% Cr Alloy in a 0.1 M Cl-, H2-Saturated
Solution of pH 10.0: Run No. KS-146 (Sample
Area = 1.13 cm2) .................................. 301

26 Galvanostatic Polarization Experiment Data for the
Fe-12% Cr Alloy in a 0.1 M Cl-, H2-Saturated
Solution of pH 10.1: Run No. KS-148 (Sample
Area = 1.33 cm2) .................................. 302


xii











LIST OF TABLES (Continued)


Table Page

27 Galvanostatic Polarization Experiment Data for
the Fe-12% Cr Alloy in a 0.1 M Cl-, H2-Saturated
Solution of pH 10.1: Run No. KS-149 (Sample
Area = 1.33 cm2).................................. 302

28 Single-Pit Galvanostatic Polarization Experiment
Data for the Fe-12% Cr Alloy in a 0.1 M CI-, H -
Saturated Solution of pH 10.2: Run No. KS-185
(Sample Area 0.000883 cm2)...................... 304

29 Single-Pit Galvanostatic Polarization Experiment
Data for the Fe-12% Cr Alloy in a 0.1 M Cl-, H2
Saturated Solution of pH 12.1: Run No. KS-186
(Sample Area = 0.000883 cm )...................... 305

30 Single-Pit Galvanostatic Polarization Experiment
Data for the Fe-12% Cr Alloy in a 0.1 M Cl-, H2
Saturated Solution of pH 10.2: Run No. KS-187
(Sample Area 0 0.000883 cm2)................. ..... 306

31 Artificial Occluded Cell Data for the Fe-12% Cr
Alloy and Armco Fe in Solutions Initially 0.1 M C1- 308

32 Analyses of Final Artificial Occluded Cell
Solutions for the Fe-12% Cr Alloy and Armco Fe
in Solutions Initially 0.1 M CI"................... 312

33 Free Energy of Formation Data for Fe-Containing
Species Considered................................. 323

34 Free Energy of Formation Data for Cr-Containing
Species Considered................................. 324

35 Pertinent Nernst Equations for the Fe-H 0 System
Potential Vs. pH (Pourbaix) Diagram (Activities
of All Dissolved Species Taken to Be 10 )........ 325

36 Pertinent Nernst Equations for the Fe-C1 -H 0
System Potential Vs. pH (Pourbaix) Diagram
(Activity of C1- Taken to Be 0.1 and Activities
of All Other Dissolved Species Taken to Be 10-6... 326

37 Pertinent Nernst Equations for the Special Fe-C1--
H20 System Potential Vs. pH (Pourbaix) Diagram
(Activity of Cl- Taken to Be 0.373 and
Activities of All Other Dissolved Species Taken
to Be 0.156)...................................... 327


xiii












LIST OF TABLES (Continued)


Table Page

38 Pertinent Nernst Equations for the Fe-H20
System Potential Vs. pH (Pourbaix) Diagram
(Activities of all Dissolved Species Taken to Be
-6
106 )......................................... 328

39 Pertinent Nernst Equations for the Fe-C1 -H20
System Potential Vs. pH (Pourbaix) Diagram
(Activity of Cl- Taken to Be 0.1 and Activitieg
of All Other Dissolved Species Taken to Be 10 .. 329

40 Pertinent Nernst Equations for the Special Cr-Cl -
H20 System Potential Vs. pH (Pourbaix) Diagram
(Activity of C1 Taken to Be 0.373 and Activities
of all Other Dissolved Species Taken to Be 0.02).. 330

41 Data Used for HYD1E, HYD2E, HYD3A and HYD4A
Computer Programs.................................. 334

42 Hydrolysis Calculation Results.................... 335

43 Cyclic Potentiokinetic Polarization Scan Data
for the Fe 12% Cr Alloy in H2-Saturated Metal
Chloride Solutions ("Teflon Gasket" Samples)...... 339

44 Cyclic Potentiokinetic Polarization Scan Data for
the Fe 12% Cr Alloy in H2-Saturated Metal
Chloride Solutions ("Microshield" Samples)........ 341












LIST OF FIGURES


Figure Page

1 Hypothetical cathodic and anodic polarization
plots for a passive metallic electrode............ 11

2 Proposed potential vs. concentration ratio plot
for an active-passive metal showing regions of
active dissolution (and etching), passivity,
brightening and transpassivity. Seven possible
kinds of reactions are indicated by numbers....... 24

3 Effect of sodium chloride molarity on the
"protection potential" for Types 410, 304 and
316 stainless steels in 0.1 M NaHCO N -
saturated solutions at 20 C ....................... 58

4 Effect of chloride ion normality on the
"protection potential" for Type 316L stainless
steel in stagnant chloride solutions at 70C ...... 59

5 Effect of temperature on the "protection
potential" for Types 410, 304 and 316 stainless
steels in 0.1 M NaC1 + 0.1 M NaHCO N -
saturated solutions............................... 61

6 Effect of temperature on the "protection potential"
for Type 316L stainless steel in stagnant 0.88 N
NaC1 solutions.................................... 62

7 Effect of extent of propagation on the "protection
potential" for Types 430 and 304 stainless steels
in 1 M NaC1, N2-saturated solutions at 250C.
2
Initial specimen area: 5 cm ..................... 63

8 Correlation between the "difference potential"
(ER Ep) and corrosion weight loss of stainless
steel alloys exposed in sea water for 4.25 years.
The "difference potentials" were determined in
3.5 wt. % NaC1, air-saturated solutions at 25C... 66

9 Schematic diagram of equipment used to
potentiokinetically polarize specimen and record
current density as a function of potential........ 69

10 Effect of scan rate on potentiokinetic polariza-
tion curves for the Fe-12% Cr alloy in 0.1 M
NaCI + 0.1 M NaHC03, H2-saturated solutions
(pH 8.7) at 250C.................................. 70


XV











LIST OF FIGURES (Continued)


Figure Page

11 Schematic diagram of equipment used to
galvanostatically polarize sample and record
resulting potential as a function of time......... 72

12 Modified sample and sample holder for first type
of artificial occluded cell used during this
investigation...................................... 73

13 Schematic diagram of equipment used for first
type of artificial occluded cell.................. 75

14 Second type of artificial occluded cell used
during this investigation......................... 76

15 Electrical schematic diagram for the second
type of artificial occluded cell.................. 78

16 Schematic diagram of equipment used for second
type of artificial occluded cell. S1 and S2 are
the bulk and artificial occluded cell samples,
respectively...................................... 79

17 a. Schematic diagram for the development of an
experimental potential vs. pH diagram for iron
in chloride-free, N2-saturated solutions. The
"a" and "b" dashed lines represent the equili-
brium hydrogen and oxygen electrodes,
respectively................................... 81

b. Schematic diagram for the development of an
experimental potential vs. pH diagram for iron
in 0.01 M Cl-, N2-saturated solutions. The
"a" and "b" lines represent the equilibrium
hydrogen and oxygen electrodes, respectively.. 81

18 Experimental potential vs. pH diagrams for the
series of Fe-Cr alloys under investigation in
0.1 M Cl-, H2-saturated solutions, as reported by
Cusumano. The cross-hatched areas represent
regions of general corrosion...................... 82

19 Some typical potentiokinetic polarization curves
for the series of Fe-Cr alloys under investigation
in 0.1 M Cl H2-saturated solutions, according to
Cusumano. Epp: primary passivation potential,
ER: rupture potential, Ep: "protection potential". 83











LIST OF FIGURES (Continued)


Figure Page

20 Experimental potential vs. pH diagrams for the
Fe 12% Cr alloy in 0.1 M CI-, H2-saturated
and 02-saturated solutions....................... 84

21 Schematic experimental potential vs. pH diagrams
for the Fe 12% Cr alloy in 0.1 M C-l, H2
saturated and 02-saturated solutions............. 85

22 Schematic representation of two types of cyclic
potentiokinetic polarization curves observed for
the Fe 12% Cr alloy in C1--containing, H 2-
saturated solutions. E : zero current potential,
Epp: primary passivation potential, ER: rupture
potential, Ep: "protection potential"............. 86

23 Experimental potential vs. pH diagrams for the
series of Fe-Cr alloys in 0.1 M C1-, H2-saturated
solutions, including zero current potentials,
primary passivation potentials and rupture
potentials........................................ 88

24 "Protection potential" (Ep) data, superimposed
on the zero current potential, primary passiva-
tion potential and rupture potential lines taken
from Figure 23.................................... 90

25 Results of a series of special cyclic potentio-
kinetic polarization scans for the Fe 12% Cr
alloy in 0.1 M Cl-, H,-saturated solutions of pH
3.2, 4.4, 5.5 and 7.1.............................. 93

26 Results of a series of special cyclic potentio-
kinetic polarization scans for the Fe 12% Cr
alloy in 0.1 M Cl-, H -saturated solutions of pH
8.9, 9.7, 10.9 and 12.0........................... 94

27 Results of a series of special cyclic potentio-
kinetic polarization scans for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solutions of
pH 3.2, 4.4, 5.5 and 7.2, using painted-off
samples (- 1 cm2 in area)......................... 95











LIST OF FIGURES (Continued)


Figure Page

28 Results of a series of special cyclic potentio-
kinetic polarization scans for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solutions of
pH 8.9, 9.7, 10.9 and 12.0, using painted-off
samples (u 1 cm2 in area)......................... 96

29 Experimental potential vs. log chloride ion
concentration diagrams for four Fe-Cr alloys in
H2-saturated solutions of pH 5.4 to 5.5, including
zero current potentials, primary passivation
potentials and rupture potentials................. 98

30 Experimental potential vs. log chloride ion
concentration diagrams for four Fe-Cr alloys in
H2-saturated solutions of pH 8.7 to 8.8, including
zero current potentials, primary passivation
potentials and rupture potentials.................. 99

31 Experimental potential vs. log chloride ion
concentration diagrams for four Fe-Cr alloys in
H2-saturated solutions of pH 10.8 to 11.0,
including zero current potentials, primary passiva-
tion potentials and rupture potentials............. 100

32 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 29......................................... 101

33 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 30......................................... 102

34 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 31......................................... 103

35 Plot of "protection potential" (Ep) data vs. total
charge passed per unit sample area, Q/A (plotted
logarithmically), for the Fe 12% Cr alloy in
0.1 M Cl-, H2-saturated solutions of indicated
pH 's .............................................. 105


xviii











LIST OF FIGURES (Continued)


Figure Page

36 a. Plot of "protection potential" (E ) data vs.
total charge passed per unit (original)sample
area, Q/A (plotted logarithmically) for the
Fe 12% Cr alloy in 0.1 M C1-, H -saturated
solutions of pH approximately 0.I........... 108

b. Plot of "protection potential" (Ep) data vs.
pit density (number of pits divided by
original sample area)......................... 108

c. Plot of "protection potential" (Ep) data vs.
active area fraction, AA (pitted plus
creviced area divided by original sample area).
Points connected by a line represent a
single sample.................................. 108

37 Plot of single-pit "protection potential" (E )
data vs. total charge passed per unit sample area,
Q/A (plotted logarithmically), for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solutions of pH
approximately 10.2. Points connected by a line
represent a single sample.......................... 111

38 Instantaneous mixed potential vs. total charge
passed, Q (plotted logarithmically) for the
Fe 12% Cr alloy in 0.1 M Cl-, H2-saturated
solutions of pH approximately 10.1. Sample areas
b 1 cm2. Each diagram presents results for a sin-
gle sample........................................ 115

39 Superposition of data from Figures 36a and 38d.
A superposition of potentiokinetic and galvano-
static results.................................... 116

40 Single-pit instantaneous mixed potential vs.
total charge passed per unit sample area, Q/A
(plotted logarithmically), for the Fe 12% Cr
alloy in 0.1 M ClI, H2-saturated solutions of pH
approximately 10.2. Each diagram presents results
for a single sample................................ 118












LIST OF FIGURES (Continued)


Figure Page

41 Potentiokinetic plot of applied potential vs.
Ag-AgCl electrode potential for the first type
of Fe 12% Cr artificial occluded cell in a
0.1 M Cl-, H2-saturated solution of pH 5.4. The
zero current potential, rupture potential and
"protection potential," observed with an ammeter
incorporated into the potentiostat, are indicated
with arrows....................................... 120

42 Potentiokinetic plots of applied current vs.
Ag-AgCl electrode potential for the first type
of Fe 12% Cr artificial occluded cell in 0.1 M
Cl-, H2-saturated solutions of pH 8.8............. 121

43 a. Progress of artificial occluded cell
experiments for the Fe 12% Cr alloy in
solutions initially 0.1 M NaC1 + 0.01 M
NaHCO3, with H2-saturated bulk solutions...... 124

b. Progress of artificial occluded cell experi-
ments for the Fe 12% Cr alloy in solutions
initially 0.1 M NaC1 + 0.01 M NaHCO3, with
02-saturated bulk solutions. The results are
superimposed on the experimental potential vs.
pH diagram (0.1 M Cl-, H -saturated solutions)
for the same alloy............................ 124

44 Progress of artificial occluded cell experiments
for the Fe 12% Cr alloy in solutions initially
0.1 M NaC1, adjusted to approximately pH 10 with
0.1 M NaOH. The bulk solutions were 02-saturated.
The results are superimposed on the experimental
potential vs. pH diagram (0.1 M Cl-, H2-saturated
solutions) for the same alloy..................... 125

45 Progress of two artificial occluded cell experi-
ments conducted with Armco Fe in solutions
initially 0.1 M NaCI, adjusted to approximately
pH 10 with 0.1 M NaOH. The bulk solutions were
02-saturated. The results are superimposed on
the experimental potential vs. pH diagram for
the Fe 12% Cr alloy (0.1 M Cl-, H2-saturated
solutions) for purposes of comparison............. 128











LIST OF FIGURES (Continued)


Figure page

46 Minimum pH and corresponding potential attained in
the Fe 12% Cr artificial occluded cell experi-
ments presented in Figures 43a, 43b and 44,
superimposed on the experimental potential vs. pH
diagram for the same alloy in 0.1 M Cl-, H -
saturated solutions............................... 130

47 Return zero current potentials and their
corresponding pH's measured for some of the
artificial occluded cell samples, superimposed on
the results of Figure 46. Run numbers are
included for reference............................ 131

48 Results of atomic absorption spectrophotometric
analyses of the final artificial occluded cell
solutions for the Fe 12% Cr alloy. The solid
and dashed lines were calculated assuming no
macroscopic dealloying occurred................... 133

49 Plot of chloride concentrations (determined by
AgNO3 titrations) vs. linear combinations of the
data shown in Figure 48. The solid line was
calculated assuming the predominant corrosion
products to be FeC 2.xH20 and CrCl yH 20.......... 135

50 Comparison of calculated equilibrium potential vs.
pH (Pourbaix) diagrams for the Fe-H20 system and
experimentally determined potential vs. pH
diagrams for the Fe 12% Cr alloy in 0.1 M Cl-
solutions......................................... 136

51 Calculated equilibrium potential vs. pH (Pourbaix)
diagram for the Fe-FeCl2'4H20-H20 system, assuming
a solution saturated in FeC12'4H20................ 137

52 Calculated equilibrium potential vs. pH diagrams
for the Cr-H20 system. Cr(OH)3'nH20 is assumed to
form in the presence of chloride solutions......... 139

53 a. Calculated equilibrium potential vs. pH
(Pourbaix) diagrams for the Fe-H20 and
Fe-C1--H20 systems: The activity of Cl was
taken to be 0.1 and the activities of gll other
dissolved species were taken to be 10 ....... 140


xxi











LIST OF FIGURES (Continued)


Figure Page

53 b. Calculated equilibrium potential vs. pH
(Pourbaix) diagrams for the Cr-H2 0 and
Cr-C1--H20 systems............................ 140

54 Specially calculated equilibrium potential vs. pH
(Pourbaix) diagrams for the Fe-Cl--H 0 and
Cr-C1--H 0 systems................................ 142
-H-
55 Results of hydrolysis calculations for Fe,
Fe Cr and Ni ............................. 147

56 a. Potential vs. pH curves calculated by combin-
ing hydrolysis equations and Nernst Equations
for pure Fe................................... 150

b. Potential vs. pH curves calculated by
combining hydrolysis equations and Nernst
Equations for pure Cr and for pure Ni......... 150

c. Potential vs. pH curves calculated by
combining hydrolysis equations and Nernst
Equations for the Fe 12% Cr alloy........... 150

57 Superposition of Fe 12% Cr artificial occluded
cell results and hydrolysis calculations for
aquo-complex ions involving Cr.................... 154

58 a. Potential vs. log chromic ion molarity
diagrams for the Fe 12% Cr alloy in 0.1 M
NaC1, H2-saturated solutions with CrC3l 6H20
and FeCl2 4H20 added........................... 159

b. "Protection potentials" (Ep), return peak
potentials, return zero current potentials
and second feature potentials (all determined
from return scans), superimposed on the zero
current potential, primary passivation
potential and rupture potential lines taken
from Figure 58a............................... 159

59 Some of the metal chloride solution scan data for
the "Teflon Gasket" samples, superimposed on the
experimental potential vs. pH diagram for the Fe -
12% Cr alloy in 0.1 M Cl-, H2-saturated solutions. 162











LIST OF FIGURES (Continued)


Figure Page

60 Log passivation current density (a/cm2) vs. log
chromic ion molarity diagrams constructed from
the metal chloride solution scan data for
"Teflon Gasket" samples and for "Microshield"
samples of the Fe 12% Cr alloy ..................... 163

61 Log current density (a/cm2) vs. log chromic ion
molarity diagrams constructed from the metal chlo-
ride solution scan data for "Teflon Gasket"
samples and for "Microshield" samples of the Fe -
12% Cr alloy...................................... 164

62 Log passivation current density (a/cm2) vs. log
chloride ion molarity diagrams constructed from
potentiokinetic polarization data for the Fe -
12% Cr alloy in H2-saturated solutions of pH
5.4, 8.8 and 10.8................................. 166

63 Log passivation current density (a/cm2) vs. pH
diagram constructed from potentiokinetic
polarization data for the Fe 12% Cr alloy
in 0.1 M (molar) Cl-, H2-saturated solutions...... 167

64 Metal chloride solution ([Fe]/[Cr] = 7.82) scan
data superimposed on the log passivation current
density (a/cm2) line for the Fe 12% Cr alloy
in 0.1 M (Molar) CI-, H -saturated solutions
(from Figure 63) .................................. 168

65 Phase diagram for the Fe-Cr system, according to
Hansen............................................. 185

66 Photomicrograph of the as-forged Fe 12% Cr
alloy showing delta ferrite in a martensite
matrix (magnification 500 X)...................... 186

67 Photomicrograph of the annealed Fe 12% Cr
alloy showing a matrix of equiaxed ferrite grains,
with randomly dispersed particles of chromium
carbide (magnification 500 X)..................... 186

68 Standard electrochemical polarization cell used
during this investigation........................... ....... 189

69 Assembled and exploded views of sample holder..... 191

70 Electrical schematic diagram of zero-impedance
ammeter used to eliminate internal resistance
errors during current measurements................ 199


xxiii










LIST OF FIGURES (Continued)


Figure Page

71 Experimental potential vs. pH diagram for the
Fe 0.5% Cr alloy in 0.1 M C-l, H2-saturated
solutions......................................... 201

72 "Protection potential" (Ep) data for the Fe -
0.5% Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 71...... 202

73 Experimental potential vs. pH diagram for the
Fe 2.0% Cr alloy in 0.1 M C-, H 2-saturated
solutions......................................... 203

74 "Protection potential" (Ep) data for the Fe -
2.0% Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 73...... 204

75 Experimental potential vs. pH diagram for the Fe -
5.0% Cr alloy in 0.1 M Cl-, H2-saturated solutions 205

76 "Protection potential" (E ) data for the Fe 5.0%
Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 75...... 206

77 Experimental potential vs. pH diagram for the
Fe 12% Cr alloy in 0.1 M, H -saturated
solutions ........................................ 207

78 "Protection potential" (Ep) data for the Fe 12%
Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 77...... 208

79 Experimental potential vs. pH diagram for the
Fe 16.9% Cr alloy in 0.1 M Cl-, H2-saturated
solutions......................................... 209

80 "Protection potential" (Ep) data for the Fe -
16.9% Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 79...... 210

81 Experimental potential vs. pH diagram for the
Fe 24.9% Cr alloy in 0.1 M Cl H2-saturated
solutions......................................... 211


xxiv












LIST OF FIGURES (Continued)


Figure Page

82 "Protection potential" (Ep) data for the Fe -
24.9% Cr alloy, superimposed on the zero current
potential, primary passivation potential and
rupture potential lines taken from Figure 81...... 212

83 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H -saturated solution of pH
3 .2 ............................................... 229

84 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl H -saturated solution of pH
3 .2 ............................................... 230

85 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M C1-, H -saturated solution of
pH 4.4 ............................................ 231

86 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12%
Cr alloy in 0.1 M Cl-, H2-saturated solution of
pH 4.4 ............................................ 232

87 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M C-l, H2-saturated solution of
pH 5.5............................................ 233

88 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H -saturated solution of
pH 5.5 ............................................ 234

89 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M C-l, H2-saturated solution of
pH 7.1............................................ 235

90 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
7.1 .................................... .......... 236

91 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M CI-, H2-saturated solution of pH
8 .9 ............................................... 237


xxv











LIST OF FIGURES (Continued)


Figure Page

92 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M C1-, H2-saturated solution of
pH 8.9 ........................................ 238

93 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of
pH 9.7 ...... ..... ...... ........... 239

94 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H -saturated solution of
pH 9.7 ............................................ 240

95 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of
pH 10.9.................... .................... 241

96 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of
pH 10.9........................................... 242

97 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of
pH 12.0 ........................................... 243

98 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H -saturated solution of
pH 12.0........................................... 244

99 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
3.2, using painted-off sample approximately
2
1 cm in area..................................... 246

100 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of
pH 3.2, using painted-off sample approximately
2
1 cm in area..................................... 247


xxvi












LIST OF FIGURES (Continued)


Figure Page

101 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
4.4, using painted-off sample approximately
2
1 cm in area..................................... 248

102 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M CI-, H -saturated solution of
pH 4.4, using painted-off sample approximately
2
1 cm in area...................................... 249

103 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
5.5, using painted-off sample approximately 1 cm
in area........................................... 250

104 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
5.5, using painted-off sample approximately 1 cm
in area........................................... 251

105 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
7.2, using painted-off sample approximately 1 cm
in area........................................... 252

106 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
7.2, using painted-off sample approximately 1 cm
in area........................................... 253

107 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M CI-, H2-saturated solution of pH
2
8.9, using painted-off sample approximately 1 cm
in area........................................... 254


xxvii











LIST OF FIGURES (Continued)


Figure Page

108 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
8.9, using painted-off sample approximately 1 cm
in area........................................... 255

109 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M C1 H2-saturated solution of pH
2
9.7, using painted-off sample approximately 1 cm
in area........................................... 256

110 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl H2-saturated solution of pH
2
9.7, using painted-off sample approximately 1 cm
in area........................................... 257

111 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
10.9, using painted-off sample approximately 1 cm
in area........................................... 258

112 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl H2-saturated solution of pH
2
10.9, using painted-off sample approximately 1 cm
in area........................................... 259

113 Out-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M CI-, H2-saturated solution of pH
2
12.0, using painted-off sample approximately 1 cm
in area........................................... 260

114 Return-scan portion of special cyclic potentio-
kinetic polarization scan for the Fe 12% Cr
alloy in 0.1 M Cl-, H2-saturated solution of pH
2
12.0, using painted-off sample approximately 1 cm
in area........................................... 261

115 Potential vs. log chloride ion molarity diagram
for the Fe 0.5% Cr alloy in H -saturated solu-
tions of pH 5.5................................... 263


xxviii











LIST OF FIGURES (Continued)


Figure Page

116 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 115......................................... 264

117 Potential vs. log chloride ion molarity diagram for
the Fe 0.5% Cr alloy in H2-saturated solutions of
pH 8.7 ............................................. 265

118 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 117......................................... 266

119 Potential vs. log chloride ion molarity diagram for
the Fe 0.5% Cr alloy in H2-saturated solutions of
pH 10.8 ............................................ 267

120 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 119......................................... 268

121 Potential vs. log chloride ion molarity diagram for
the Fe 2.0% Cr alloy in H2-saturated solutions of
pH 5.5............................................. 269

122 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 121......................................... 270

123 Potential vs. log chloride ion molarity diagram for
the Fe 2.0% Cr alloy in H2-saturated solutions of
pH 8.7 ............................................. 271

124 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 123......................................... 272

125 Potential vs. log chloride ion molarity diagram for
the Fe 2.0% Cr alloy in H2-saturated solutions of
pH 11.0 ............................................ 273


xxix











LIST OF FIGURES (Continued)


Figure Page

126 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 125......................................... 274

127 Potential vs. log chloride ion molarity diagram for
the Fe 12% Cr alloy in H2-saturated solutions of
pH 5 .4............................................. 275

128 "Protection potential" (Ep) data, superimposed on

the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 127......................................... 276

129 Potential vs. log chloride ion molarity diagram for
the Fe 12% Cr alloy in H -saturated solutions of
pH 8.8 ............................................. 277

130 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 129......................................... 278

131 Potential vs. log chloride ion molarity diagram for
the Fe 12% Cr alloy in H2 -saturated solutions of
pH 10.8............................................ 279

132 "Protection potential" (E ) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 131......................................... 280

133 Potential vs. log chloride ion molarity diagram for
the Fe 16.9% Cr alloy in H2-saturated solutions
of pH 5.4.......................................... 281

134 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 133......................................... 282

135 Potential vs. log chloride ion molarity diagram for
the Fe 16.9% Cr alloy in H2-saturated solutions
of pH 8.8.......................................... 283


xxx











LIST OF FIGURES (Continued)


Figure Page

136 "Protection potential" (Ep) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 135.......................................... 284

137 Potential vs. log chloride ion molarity diagram for
the Fe 16.9% Cr alloy in H2-saturated solutions
of pH 10.8.......................................... 285

138 "Protection potential" (E ) data, superimposed on
the zero current potential, primary passivation
potential and rupture potential lines taken from
Figure 137.......................................... 286

139 Calculated equilibrium potential vs. pH (Pourbaix)
diagram for the Fe-H20 system. The activities of
-6
all dissolved species were taken to be 10 ......... 317

140 Calculated equilibrium potential vs. pH (Pourbaix)
diagram for the Fe-C1--H20 system. The activity
of Cl was taken to be 0.1 and the activities of
all other dissolved species were taken to be 10-6... 318

141 Specially calculated equilibrium potential vs. pH
(Pourbaix) diagram for the Fe-Cl--H20 system. The
activity of Cl was taken to be 0.373 and the
activities of all other dissolved species were taken
to be 0.156 ......................................... 319

142 Calculated equilibrium potential vs. pH (Pourbaix)
diagram for the Cr-H20 system. The activities of
-6
all dissolved species were taken to be 106 ........ 320

143 Calculated equilibrium potential vs. pH (Pourbaix)
diagram for the Cr-CI--H20 system. The activity
of Cl was taken to be 0.1 and the activities of
-6
all other dissolved species were taken to be 10-.. 321

144 Specially calculated equilibrium potential vs. pH
(Pourbaix) diagram for the Cr-Cl--H20 system. The
activity of Cl was taken to be 0.373 and the
activities of all other dissolved species were
taken to be 0.02................................... 322













Abstract of Dissertation Presented to the
Graduate Council of the University of Florida in Partial Fulfillment
of the Requirements for the Degree of Doctor of Philosophy



THE SIGNIFICANCE OF THE "PROTECTION POTENTIAL"
FOR Fe-Cr ALLOYS AT ROOM TEMPERATURE


By


Kenneth Kirch Starr

March, 1973


Chairman: Dr. Ellis D. Verink, Jr.
Major Department: Materials Science and Engineering


The significance of the "protection potential," a potential

below which active pits cease to propagate (a critical potential for

pit propagation), was investigated using electrochemical hysteresis

(cyclic potentiokinetic polarization), artificial occluded cell, atomic

absorption analysis and wet chemical analysis techniques.

The effects of certain material, environmental and

experimental variables on the "protection potential" were studied:

alloy composition, pH, chloride ion concentration and extent of

propagation. Results from artificial occluded cells were compared

with thermodynamic calculations. Preliminary studies of the passive-

active behavior of certain Fe-Cr alloys in the presence of metal

chloride solutions were undertaken.

The "protection potential" appeared to be influenced by both

alloy composition and environment. Both "repassivation" and

"deactivation" mechanisms appeared to be operative for the Fe-Cr


xxxii












alloys, depending upon the particular alloy/environment system.

Evidence was obtained which suggested that an active occluded cell of

at least one alloy (Fe 12% Cr) could propagate to an extent precluding

any possibility of a "repassivation" mechanism unless there was dilution

of the contents of the occluded cell. Hydrolysis calculations indicated

that the CrOH+H complex ion may be one of the hydrolysis products

primarily responsible for localized acidification of active occluded

cells of Fe-Cr alloys.


xxxiii













CHAPTER 1

INTRODUCTION


The great utility of stainless steel alloys is common knowledge.

There are many commercially available stainless steel alloys, each

designed to have the mechanical properties and corrosion resistance

necessary for particular applications. Table 1 lists the nominal

compositions of a number of these alloys.

The martensitic chromium steels, Group I, can be hardened by

heat treatment. The ferritic steels, Group II, cannot be hardened by

heat treatment but possess better corrosion resistance. Groups I and

II are basically Fe-Cr alloys.

The austenitic stainless steels, Group III, are basically

Fe-Cr-Ni alloys, with a few notable exceptions. These alloys also

cannot be hardened by heat treatment but possess better corrosion

resistance than either Group I or II.

The age-hardenable steels, Group IV, possess both reasonably

good corrosion resistance and the ability to be hardened by heat

treatment.

The alloys studied in this investigation are Fe-Cr alloys of

0.5% Cr, 2.0% Cr, 5.0% Cr, 12% Cr, 16.9% Cr and 24.9% Cr. The 0.5% Cr,

2.0% Cr and 5.0% Cr alloys are called chromium steels. The 12% Cr

alloy is a martensitic stainless steel, and the 16.9% Cr and 24.9% Cr

alloys are ferritic stainless steels. The detailed compositions and

characterization of these alloys are given in Appendix 1.

Fe-Cr and stainless steel alloys may suffer pitting and/or










TABLE 1


CHEMICAL COMPOSITIONS OF COMMERCIAL STAINLESS STEELS [1, p. 164]


AISI
type % C % Cr % Ni % Other Elements


Group I: Martensitic Cr Steels
410 0.15 max 11.5-13.5
416 0.15 max 12-14 Se, Mo, or Zr
420 0.35-0.45 12-14
431 0.2 max 15-17 1.25-2.5
440A 0.60-0.75 16-18

Group II: Ferritic Nonhardenable Steels
405 0.08 max 11.5-14.5 0.5 max 0.1-0.3 Al
430 0.12 max 14-18 0.5 max
442 0.25 max 18-23 0.5 max
446 0.20 max 23-27 0.5 max 0.25 N max

Group III: Austenitic Cr-Ni Steels
201 0.15 max 16-18 3.5-5.5 5.0-7.5 Mn 0.25 N max
302 0.15 max 17-19 4-6 7.5-10 Mn 0.25 N max
301 0.15 max 16-18 6-8 2 Mn max
302 0.15 max 17-19 8-10 2 Mn max
302B 0.15 max 17-19 8-10 2-3 Si
304 0.08 max 18-20 8-12 1 Si max
304L 0.03 max 18-20 8-12 1 Si max
308 0.08 max 19-21 10-12 1 Si max
309 0.2 max 22-24 12-15 1 Si max
309S 0.08 max 22-24 12-15 1 Si max
310 0.25 max 24-26 19-22 1.5 Si max
310S 0.08 max 24-26 19-22 1.5 Si max










TABLE 1 (Continued)


AISI
type % C % Cr 7 Ni % Other Elements

314 0.25 max 23-26 19-22 1.5-3.0 Si
316 0.10 max 16-18 10-14 2-3 Mo
316L 0.03 max 16-18 10-14 2-3 Mo
317 0.08 max 18-20 11-14 3-4 Mo
321 0.08 max 17-19 8-11 Ti 4 x C (min)
347 0.08 max 17-19 9-13 Cb + Ta 10 x C (min)
Alloy 20 0.07 max 29 20 3.25 Cu, 2.25 Mo

Group IV: Age-Hardenable Steels
322 0.07 17 7 0.07 Ti, 0.2 Al
17-7PH 0.07 17 7 1.0 Al
17-4PH 0.05 16.5 4.25 4.0 Cu
A-8MoPH 0.05 max 14 8.5 2.5 Mo, 1 Al
AM350 0.10 16.5 4.3 2.75 Mo
CD4MCu 0.03 25 5 3.0 Cu, 2.0 Mo












crevice corrosion in particular environments such as sea water.

Interest has arisen in the possibility of protecting these alloys

against localized attack by altering or controlling their electrode

potentials. The Fe-Cr alloys are of special interest for three

reasons:

1. they contain only one basic alloying element and thus

have only one basic composition variable,

2. the martensitic grades can be hardened by heat treatment

and

3. they are the least expensive of the stainless steels.

The "protection potential," a potential below (more active than)

which active pits cease to propagate (i.e., a critical potential for

pit propagation), has been the sujbect of considerable debate among

corrosion experts in recent years.

The purpose of this research was to characterize the material,

environmental and experimental aspects of the "protection potential,"

and to distinguish between the possibilities of "repassivation" and

"deactivation" mechanisms.

The approach has been to study the effects of alloy composition,

bulk electrolyte pH, bulk electrolyte chloride ion concentration and

extent of propagation on the "protection potential." Results from

artificial occluded cells were compared with thermodynamic calculations.

These results primarily concern the "repassivation" vs. "deactivation"

question. In addition, preliminary work was carried out involving the

passive-active behavior of alloys exposed to artificial pit or crevice

solutions.












CHAPTER 2

LITERATURE SURVEY


Two phrases need to be clarified: "protection potential" and

"occluded cell." The phrase "occluded corrosion cell" was used by

B. F. Brown [2] to tie together in a general way a number of forms of

localized corrosion. These included crevice corrosion, pitting, stress

corrosion cracking, turberculation, intergranular corrosion, filiform

corrosion and exfoliation. Brown concluded that the unifying feature

of these different forms of corrosion was localized acidification by

hydrolysis and that this localized acidification was largely responsible

for their self-perpetuating character.

In this study, an occluded cell will be considered to be an

electrochemical system involving a solid metallic electrode in contact

with an electrolyte of restricted volume. The restriction may result

from any combination of the following conditions:

1. limited total quantity of electrolyte,

2. confining geometry of cell [2],

3. presence of corrosion product [2],

4. stagnation.

The phrase "protection potential" was used by M. Pourbaix and

coworkers [3] in reference to the experimental observation that, for

stainless steels in chloride solutions, the propagation of active pits

could be stopped by making the sample potential more negative (active)

than a critical value. This critical potential was considered to be

the "protection potential" against pitting.













This survey will be focused on the discussion of crevice

corrosion and pitting of stainless steel alloys, although mention will

be made of stress corrosion cracking and other types of occluded cell

corrosion. Emphasis will be placed on the significance of the "protec-

tion potential" concept with respect to crevice corrosion and pitting.

Some of the terms used are defined as follows:

1. Corrosion The deterioration of a substance,
usually a metal, because of a reaction with its
environment.

2. Passivation The process or processes by means of
which a metal becomes inert to a given environment
or environments.

3. Crevice Corrosion Localized corrosion as a result
of the formation of a crevice between a metal and
a nonmetal, or between two metal surfaces.

4. Pitting Localized corrosion taking the form of
cavities at the surface.

5. Stress Corrosion Cracking Spontaneous cracking
produced by the combined action of corrosion and
static stress (residual or applied). [4, pp. 267t-
268t)


2.1 Methods of Investigation


Experimental techniques which have been employed in the study

of crevice corrosion and pitting phenomena include immersion tests,

galvanostatic polarization techniques, potentiostatic polarization

techniques, potentiokinetic polarization techniques, the use of

artificial occluded cells, indicator techniques, autoradiography,

chemical analysis, X-ray analysis, electron probe microanalysis, optical

microscopy, electron microscopy and scanning electron microscopy.

Greene and Fontana [5] discussed advantages and disadvantages













of a number of these techniques. The first four techniques listed

will be briefly described here. Other methods of investigation will

be described later.


2.1.1 Immersion Tests


Immersion tests involve the placing of a sample or samples

within a suitable environment. The environment may consist of any

combination of gaseous, liquid or solid phases. Atmospheric exposure

is an important type of immersion test involving gaseous, liquid and

solid phases. Frequently, immersion tests involve samples totally

immersed in liquid electrolytes at ambient temperatures. Dissolved

gases and solid objects may also be present.

Information obtained may result from visual observations,

metallographic examinations, weight-loss measurements and potential

measurements. Such potentials are "free-corrosion" or "mixed"

potentials [6, 7]. They are determined only by the metal-

environment system.

In the case of stainless steel alloys, there exists an ASTM

Standard Method for total immersion corrosion tests, ASTM Designation:

A279-63 [8]. This standard discusses general considerations, leaving

specifics to the individual investigator.


2.1.2 Galvanostatic Tests


"Galvanostatic" implies a technique of applying a constant

current to a sample immersed in an electrolyte [9]. Current is the

independent variable and usually potential is the dependent variable.

Visual observations may also be made.












In normal practice the current is varied, either stepwise or

continuously (galvanokinetically), and the potential is monitored.

Brennert [10] employed such a technique to measure his "break-through"

potentials, denoting the initiation of pitting of stainless steel

samples.


2.1.3 Potentiostatic Tests


"Potentiostatic" implies the potential of an electrode being

held constant during some experimental procedure, such as measuring

the current over a time interval [9]. Potential is the independent

variable and usually current is the dependent variable. Visual

observations may be made and weight-loss data may be obtained.

In normal practice the potential is varied in a stepwise manner

and the current is monitored. A brief review concerning "the

classical potentiostat" and some simplified circuitry was given by

Greene [11]. He also showed that such an instrument was particularly

useful in the study of passivity.


2.1.4 Potentiokinetic Tests


"Potentiokinetic" (also "potentiodynamic") implies a technique

of varying the potential of an electrode in a continuous manner at a

preset rate, frequently used to prepare polarization plots [9].

Potential is the independent variable and usually current is the

dependent variable.

The rate of potential scan (scan rate) may be a significant

variable [12, 13]. ASTM Standard Number G5-69 [14] presents a












reference method for conducting both potentiostatic and potentiokinetic

anodic polarization measurements. Greene [15] has given a detailed

treatment of experimental aspects of electrode kinetics.


2.2 Passivity


Fontana and Greene stated that


...the nature of the passive film, and consequently
the basic nature of passivity, still remains an
unsolved problem. [1, p. 321]


However, the fact that many passivable metals are particularly

susceptible to pitting [16, 17, 18] and/or crevice corrosion [19] under

particular conditions makes it imperative that some thought be given

to the passivation process itself.

Today there are two basic theories of passivity. The first

theory, sometimes referred to as the "oxide-film theory," holds that

passivity is due to a diffusion barrier of reaction products. These

reaction products are usually considered to be oxides of some sort,

although they may not necessarily correspond to bulk oxides [20, 21,

22, 23, 24].

The second theory attributes passivity to the chemical

adsorption of atoms, ions, or molecules on the surface of a metal.

The adsorbed layer may be only a monolayer thick and probably contains

chemically adsorbed oxygen ions and/or hydroxide ions in many cases

[20, 22, 23].

Tomashov and Chernova [22] attributed the first description

of passivity in metals to Lomonosov in 1738. Tomashov and Chernova

[22] and Uhlig and Wulff [25] attributed the original "oxide-film












theory" to Faraday in 1836.

Uhlig [26], in a paper tying together the concepts of chemisorp-

tion and the "adsorption theory" of passivity, attributed the original

suggestion to Langmuir in 1916. Uhlig and Wulff [25] developed a

theoretical basis for the "adsorption theory" of passivity in alloys

according to an electron-sharing scheme.

It appears that the "adsorption theory" of passivity, in its

strictest sense, is presently being abandoned or modified [18].

However, Tomashov and Chernova have stated that:


It must be pointed out that the oxide-film and
adsorption theories do not contradict, but rather
supplement one another. As the adsorbed film in the
process of thickening gradually passes into an oxide
film, the retardation of the anodic process promoted
by change in the double-layer structure will also be
supplemented by the greater difficulty encountered by
ions passing directly through the protective film. Thus,
one may speak of a combined oxide-film-adsorption theory
of passivity. [22, p. 13]


Figure 1 shows schematic apparent anodic and cathodic

polarization curves for a passive metal in an aqueous electrolyte

containing aggressive anions, representing steady-state conditions.

More will be said later about "aggressive anions." The plot in Figure

1 is according to the conventions outlined in ASTM Standard Designation

G3-68 [9], with slight modifications [27]. The critical points on

that curve may be defined as follows:

1. E corrosion potential The potential of a
corr.
corroding electrode in a stated environ-
ment. Often used to mean the open circuit
(also called "steady state," freely corroding
or "rest") potential in an electrochemical
cell; that is, the potential without any
external current flowing. [9, pp. 546-548]






11





_j
ll-


-J
E 0
R >

Imperfect /
Passivity


EP O
Perfect aC
Passivity ul

Active Region Epp 0


~corr. C _j



Immunity Region


LOG CURRENT DENSITY (ma/cm2)


Figure 1. Hypothetical cathodic and anodic polarization
plots for a passive metallic electrode [9, 27].



2. E primary passive potential The potential
corresponding to the maximum active current
density in an electrode which exhibits
active-passive corrosion behavior. [Also
called primary passivation potential.]
[9, pp. 546-548]

3. iC critical anodic current density The maximum
anodic current density observed in the active
region with a metal or alloy electrode which
exhibits active-passive behavior in the
environment studied. [9, pp. 546-548]

4. ER rupture potential The potential above which
the passivating film becomes locally non-
protective leading to pitting. [Also called
"breakthrough potential," pitting potential,
critical pitting potential or critical
potential for pit initiation.] [27, p. 431]











5. E protection potential The potential below
which formed pits will not grow and thus
become harmless. [Also called critical
potential for pit propagation (or growth).]
[27, p. 431]


2.3 Aggressive Anions


The terms "aggressive anions" and "unaggressive anions" are

misleading if taken too literally, but are useful for general

classification. Kolotyrkin [17] used the former term with reference

to specifically acting anions, particular Cl Br or I He used the

latter term with reference to SO4 C103, CO3 NO3 and Cr0 with

C104 falling into both categories.

Borgmann [28] found that the initial corrosion rates of iron

in salt solutions were similar for chloride, iodide and sulfate

solutions. Foley [29] reviewed the literature on the role of the

chloride ion in the corrosion of iron. Although there was no general

agreement in the literature, it appeared that many ions could be

either aggressive or inhibitive, depending upon their concentrations

[30]. The situation was even more complicated for solutions containing

two or more anions (other than OH ).

Leckie and Uhlig [31] studied the effects of additions of

SO C104, OH and NO3 to chloride solutions on pitting of an

18% Cr 8% Ni stainless steel. They found, for each anion, a critical

activity ratio of anion to chloride ion necessary to inhibit pitting

corrosion.

It becomes clear that the aggressiveness of any particular

anion depends upon several factors, including the nature of the












metal, whether or not it is passivated, the concentration of the anion,

and the presence and concentration of other anions.

Greene and Fontana [5] noted, however, that the halides and

halogen-containing anions were almost always associated with the pitting

of metals. In particular, chlorides, bromides and hypochlorites were

usually the most aggressive of the halogen-containing anions. It is

in this context that the term "aggressive anion" will be used in

this discussion.


2.4 Pitting Corrosion


Reviews of the literature of pitting corrosion have been

published by Greene and Fontana [5], Kolotyrkin [17], and Szklarska-

Smialowska [32]. It has been found meaningful and convenient to

divide pitting into two steps [33, 34, 5, 32]: pit initiation or

surface breakdown and pit growth in depth and volume [34], or more

simply, initiation and propagation [33]. Any complete theory of pitting

must take both steps into account in a consistent manner.

Szklarska-Smialowska [32] expressed the opinion that none

of the existing theories was completely successful. A brief look

will be taken at some of the theories concerning pit initiation and a

more thorough look taken at the theories and experiments concerning

pit growth.

Any theory of pit initiation must apply to at least one

theory of passivity. In fact, the two basic theories of passivation

have attendant theories of pit initiation. Aggressive anions, usually

halide ions, will be assumed present in the electrolytes.












2.4.1 Pit Initiation


2.4.1.1 Breakdown of passivity: adsorption theory

According to the adsorption theory, chloride ions [35, 5, 20,

32] or other aggressive anions [17] competitively or preferentially

adsorb on a passive metal surface, displacing oxygen [35, 5, 17, 20,

32] or hydroxide ions [20]. Substitution of adsorbed oxygen by

chloride ions has also been suggested [36]. Pitting is then assumed

to initiate at sites or areas where the passivating species has been

successfully displaced or replaced by aggressive anions.


2.4.1.2 Breakdown of passivity: oxide-film theory


Numerous theories of the breakdown of passive oxide films have

been proposed. Since these theories are well described in the

previously mentioned review articles on pitting corrosion [5, 17, 18],

only a few of the more important historical and current theories will

be mentioned here.

The penetration theory was suggested by Evans [37] in 1927.

According to this theory, the small diameter of the chloride ion

allowed it to penetrate the protective oxide layer on passive iron.

The acid theory was proposed by Hoar [38] in 1947. According

to Hoar, who then assumed the passive oxide film on iron to be Fe(OH)2,

transference of chloride ions to the anode surface allowed the

electrolyte next to the anode to become acidified. Eventually this

localized acidification led to destruction and undermining of the

passive film. In a later work Hoar [39] assumed the existence of

incipient cracks or pores in the passive oxide film.












Thermodynamic theories of pit initiation have been given in

the literature [40, 41, 42]. These theories suggest that at some

critical anodic potential,conversion of a protective oxide film into

a nonprotective phase (such as a metal chloride) becomes energetically

possible. Unfortunately, the thermodynamic theories offer only a

qualitative description of the pit initiation process.


2.4.1.3 Breakdown of passivity: combined theory


Assuming a modern combined theory of passivation to be

operative, Hoar and Jacob [43] proposed a plausible mechanism for

the breakdown of passivity of stainless steels by halide ions. Their

theory proceeded as follows:


Three or four halide ions jointly adsorbb" on the
oxide film surface around a lattice cation--one next to
a surface anion for preference. The transitional
complex thus formed will be of high energy and the
probability of its formation at any instant will be very
small. But, once formed, the complex can readily and
immediately separate from the oxide ions in the lattice,
the cation dissolving in the solution, very much
more readily than the non- or aquo-complexed cations
present in the film surface in the absence of halide
ion. Under the anodic field, a further cation comes up
through the film to replace the dissolved cation--the
field at constant anode potential increases at the
"thinned" point of the film; but arriving at the film/
solution interface, it finds, not stabilizing oxide ion
formed from water (nor, in de-aerated solution, oxygen
molecules), but several halide ions, so that the
catalytic process, once begun, has a strong probability
of repeating itself, and of accelerating because of the
increasing electrostatic field. Thus, once localized
breakdown starts with the initial transitional complex,
it accelerates "explosively." [43, p. 1300]

However, for the case of iron, Hoar suggested that

...a process of migration through the passivating
film of one chloride ion at a time may determine the
breakdown rate. [18, p. 20C]












One apparent weakness of the existing theories of the breakdown

of passive films by aggressive anions is that they are unable to predict

or even explain observed experimental values of critical pitting

potentials.

Some of the experimental methods available for studying the

breakdown of passive films by aggressive anions will now be discussed.


2.4.1.4 Passive film breakdown parameters


Reinoehl and Beck [44] noted that for an active metal to

become passive there existed a number of critical parameters,

including a critical concentration of oxidizing agent, a critical

potential, a critical anodic current density, a critical temperature,

a critical pressure and a critical pH. It seems reasonable that many

of these parameters should also be important in the breakdown of

passivity. In pitting, an additional parameter is a critical concen-

tration of aggressive anion [45, 40].

The critical parameters most often used in experimental

studies of pit initiation are potential, pH, and aggressive anion

concentration (usually chloride ion concentration). Oxidizing agents

(particularly oxidizing metal chlorides) have been used to promote

pitting of passive metals [46, 47, 34], but as a first approxima-

tion the action of oxidizing agents should be describable in terms of

potential, pH, and aggressive anion concentration.

Uhlig [46] proposed that only solutions capable of forming

sufficiently noble oxidation-reduction potentials would cause

pitting of stainless steels. Kolotyrkin later suggested a general

rule:












At present, the following general rule can possibly
be formulated: a metal is subject to pitting corrosion
only if the redox potential of the solution containing
aggressive anions lies at values more positive than V
c
[critical pitting potential]. In other words, the redox
system can stimulate the development of pits only on those
metals for which the V lies at more negative values than
the redox potential of the system. [17, p. 263t]


Streicher [34] found that the pitting intensity of Type 304

stainless steel at 250C in nine different chloride salts was largely

a function of the pH of the solutions (due to hydrolysis). Oxidizing

agents may also affect pitting by contributing an aggressive anion to

the electrolyte. The concentration of that particular anion would

depend upon the concentration of the oxidizing salt.

Szklarska-Smialowska [32] noted that no systematic work had

been done on temperature effects in pitting. The work that had been

done indicated a greater tendency for pitting at higher temperatures

[3, 32]. Little or no work has been done on pressure effects.


2.4.1.5 Critical pitting potentials


Electrical stimulation of pit initiation may be accomplished

galvanostatically, galvanokinetically (continuously increasing

applied current), potentiostatically, potentiokinetically or with

simple applied voltages.

Uhlig and Wulff [25] used a battery and slide wire resistances

to apply a voltage across a cell consisting of a stainless steel

anode and a silver-silver chloride cathode. As long as the anode

remained passive and little current flowed, the silver-silver chloride

electrode should not have been appreciably polarized. Under those












conditions the anode potential would have been nearly equal to the

applied potential and the apparatus should have operated as a potentio-

stat. However, the relatively large currents associated with pitting

would have polarized the silver-silver chloride cathode and potential

control would then have become uncertain.

Uhlig and Wulff [25] noted that Donker and Dengg had used a

similar apparatus to make film breakdown measurements on iron in 1927.

The procedure was to increase the applied voltage (equal to the

potential) until a large increase in current was measured by an ammeter,

corresponding to the initiation of pitting. The applied potentials at

the moment of film rupture were called "threshold potentials."

In 1937, just prior to the work of Uhlig and Wulff, Brennert

[10] had developed a slightly different method of evaluating the

pitting resistance of stainless steels. Brennert used an auxiliary

cathode and a reference electrode instead of a combined cathode-

reference electrode. He applied an increasing voltage between the

anode and auxiliary cathode until the passive film ruptured, at

that point measuring his "breakthrough potentials."

Mahla and Nielsen [48] used a similar apparatus to study the

effects of prepassivation and inhibitors. Streicher [34] also used a

similar apparatus in his studies of pit initiation. Streicher, however,

maintained a breakthrough voltage for a fixed time and used the

resulting density of pits as a measure of pit initiation resistance.

Pourbaix et al. [49, 3] employed a potentiokinetic method

to determine "breakdown" and "protective tensions" for AISI Types 410,

304, and 316 stainless steels in bicarbonate solutions at temperatures

of 200, 40' and 80'C. Such terms as "breakthrough potential," "threshold












potential," and "breakdown potential" are somewhat confusing. In

addition, the terms "critical pitting potential" and "critical potential

for pit initiation" have been used in the literature. All of these

potentials are related to the rupture potential. Szklarska-Smialowska

and Janik-Czachor [50] presented a critical analysis of the experimental

methods available for the determination of characteristic potentials

of pitting corrosion. Bond [51] has also reviewed testing methods for

pitting corrosion and their interpretation.


2.4.1.6 Critical chloride ion concentrations


Some early work to determine the critical minimum chloride

ion concentration necessary for the initiation of pitting was published

by Fenwick [45] in 1935, describing an "electrometric titration"

method. This method involved the titration of a standard chloride

solution into an electrolyte containing a carbon steel or stainless steel

sample and a bridge tube leading to a silver-silver chloride reference

electrode.

Usually, a sudden change in the electrode potential of up to

one volt occurred at some critical chloride ion concentration.

Fenwick took this critical concentration to be a measure of the relative

resistance to corrosion (pitting) of a ferrous alloy in the passive

state.

Pourbaix et al. [52, 3] used a modification of this technique

for a similar purpose. Their method consisted of a continuous

chloride titration, under potentiostatic conditions, while monitoring

the applied current. A sharp increase of applied current was noted at











a critical value of chloride ion concentration. The anode material in

this case was an AISI Type 304 stainless steel rod.

Nobe and Tobias [53] used a similar technique as part of their

study of the effects of chloride ions on the potentiostatic anodic

polarization of iron.

Pourbaix et al. [52, 3] used rupture potentials, determined

potentiokinetically for stainless steels in solutions of known chloride

ion concentrations, to study the effect of chloride ion concentration

on the initiation of pitting.

Leckie and Uhlig [31] used a potentiostatic method for studies

of an 18% Cr 8% Ni stainless steel. Verink [54] employed a

potentiokinetic method to study the effect of chloride ion concentration

on rupture potentials for a series of binary Fe-Cr alloys.

The results of the studies mentioned indicate that the break-

down of passivity is strongly influenced by the chloride ion concentration

in solution. Higher chloride ion concentrations result in lower

rupture potentials and an increased tendency for the initiation of

pitting corrosion.


2.4.1.7 Effect of pH


Szklarska-Smialowska [32] indicated that most of the recent

literature suggested the rupture potential to be rather insensitive

to pH. Streicher [34] found pitting intensity of Type 304 stainless

steel to decrease with increasing pH. Pourbaix et al. [3] found the

addition of NaHCO3 to chloride solutions (thus buffering the pH

at around 8.3) to raise the rupture potential as much as several

hundred millivolts.












Leckie and Uhlig [31] noted a small increase in the critical

pitting potential with pH (18% Cr 8% Ni stainless steel in 0.1 N NaC1)

up to a pH of about 8, where a large increase occurred. Pourbaix [27]

presented experimental potential vs. pH diagrams for iron and stainless

steels showing a considerable increase (as much as several hundred

millivolts) in the rupture potential with pH, depending upon the

chloride ion concentration.

Verink [54] and Cusumano [55] found both dependent and

relatively independent behavior for a series of binary Fe-Cr alloys,

depending upon the particular alloy.

Both Pourbaix's [27] and Cusumano's [55] work indicated that

at some critical pH value the primary passivation and rupture

potentials should become equal, and that below this pH value passivation

should not occur at all.

Leckie [13] determined critical pitting potentials for Type

304 stainless steel as a function of pH in 0.01 M, 0.1 M, and 1.0 M

NaC1. His values were nearly constant up to a pH of 10 to 12,

depending on chloride concentration, where a marked shift to more

noble potentials occurred.


2.4.1.8 Effect of alloy composition


A wide variation in rupture potentials has been found between

different metals and/or alloys. Hoar [18] cited a few typical

examples for a given electrolyte ranging from a few tenths of a volt

to approximately thirty volts.

Kolotyrkin [17] studied the effect of Cr in Fe-Cr alloys and












Ni in Fe-Ni alloys on the critical pitting potential (0.1 N chloride

solutions of pH % 2). An abrupt increase (more noble) was observed

at 28-30% Cr for the Fe-Cr alloys. Steigerwald [56] also found

28-30% Cr to be critical for stabilizing the passive state in both

ferric chloride and 0.1 N NaC1 solutions.


2.4.1.9 Induction periods


Induction periods for the initiation of pitting corrosion

have been noted in the literature [57, 40, 58, 59, 43, 60, 32].

Engell and Stolica [40] found the induction time and the chloride ion

concentration to be inversely proportional for carbon steels. This

relationship was confirmed for stainless steel samples by Schwenk [58].

Hoar and Jacob [43], however, found 1/T for an 18% Cr 8% Ni

stainless steel in chloride solutions (T = induction time) to be

proportional to the nth power of the chloride ion concentration, where

2.5 < n < 4.5. For bromide solutions the exponent was between 4 and 4.5.

The authors calculated an apparent energy of activation of 60 kcal/mole.

They used this information as a basis for their theory of the break-

down of passivity, cited in Section 2.4.1.3.

It has also been established that the induction time depends

upon potential [59, 43] and temperature [43]. In general, it appears

that the induction time decreases with increasing potential, temperature

or chloride ion concentration.


2.4.2 Propagation of Pitting


The second stage of pitting, propagation, contains several

important aspects: the geometry of pitting (the size, shape, structure,












and distribution of pits); the current distribution in and around pits;

the potential distribution and resistance effects in and around pits;

the electrolyte composition within pits; mass transport and diffusion

in pits; and stability and reversibility considerations in pitting

corrosion.


2.4.2.1 Geometry of pitting


Champion [61] characterized the possibilities of variations

in intensity and distribution of metallic corrosion. He suggested

a scheme of classification ranging from general corrosion to cracking

and graduations in pit number and size. The tools of investigation

were visual observation, optical microscopy and radiography. More

recently, transmission electron microscopy [62] and scanning electron

microscopy [63, 64] have been used to study pit morphology.

Szklarska-Smialowska [32] reviewed the recent literature on

the subject of pit morphology. Her conclusion was that the shape

of a pit depended on both the environmental conditions within the

pit and the composition, properties and structure of the particular

metal.

Schwenk [58] found that at low potentials and current densities

regularly etched pits (mostly hexagons and squares) were sometimes

formed. High potentials and current densities lead to isotropic

(hemispherical) pits. Schwenk noted that these conditions were

similar to those present in electrolytic polishing, although the pit

areas might look either dull or polished.

Hoar [65] proposed a schematic diagram of potential vs.

anion/water concentration ratio, shown in Figure 2, on which he





























-J

5

z passive 5

O
- B
0




7 7' 6

A active





ANION/ WATER CONCENTRATION RATIO

Figure 2. Proposed potential vs. concentration ratio plot for
an active-passive metal showing regions of active
dissolution (and etching), passivity, brightening
and transpassivity. Seven possible kinds of
reactions are indicated by numbers [65, p. 354].












delineated areas of active dissolution (and etching), passivity,

brightening and transpassivity.

In addition to pit shape, pit size and distribution can be

important. Agar and Hoar [66] noted that the e.m.f. of an active

cell had to overcome electrolyte resistance, electrode polarization

and external circuit resistance. Thus, very large cells (referring

to current path distance) are frequently limited by electrolyte

resistance and very small cells are often limited by electrode

polarization.

Waber [67] treated the problem similarly, but in more depth.

He referred to the above-mentioned limiting cases as macroscopicc"

and "microscopic" cells. For coplanar electrodes, Waber found the

potential of the "composite electrode" (the potential measured at

large enough distances to be a negligible function of position) to be

proportional to the anode area fraction.

Artificial pits have been employed from time to time in the

study of pitting corrosion, particularly the second stage. Greene and

Fontana [5] cited as one criticism of this technique the fact that

most artificial pit designs were not geometrically similar to actual

corrosion pits.


2.4.2.2 Current distribution in and around pits

Three types of current measurements have been made with

respect to pitting corrosion: total applied current measurements

(normally under potentiostatic control), localized current measurements

in the vicinity of pits and current measurements associated with

artificial pits.












Szklarska-Smialowska [32] also reviewed the recent literature

pertaining to pit growth kinetics. A semiempirical current equation

for potentiostatic control of the form i = kt where i = applied

current, k = constant depending upon chloride ion concentration, b =

constant (sometimes 2 or 3) and t = time, has been found to hold in

some cases.

Rosenfeld and Danilov [36] employed a twin probe method to

measure the field strength in the solution adjacent to an active

stainless steel pit. They were then able to calculate the current

distribution adjacent to the pit. The effect of time on this distri-

bution was also studied.

Rosenfeld and Danilov discovered, as they traversed an active

pit, that the anodic current density was a maximum at the center of

the pit and the cathodic current density was a maximum near the edge

of the pit. The apparent current density of a pit decreased with time

-1/2
according to the equation i = kt where k = constant and t = time.
a

Artificial pits have been employed for current measurements,

the determination of electrochemical control and the effects of

additions of cathodic depolarizers.

Brown and Mears [68-70] developed their well-known "scratch

technique" and applied it to electrochemical studies of Al [68],

stainless steel (18% Cr 8% Ni) [69], and steel and Mg [70]. This

method involved the short-circuited coupling, through a zero-resistance

ammeter, of a waxed anode with one or more scratches present and a

waxed cathode of definite exposed area. The potential was also

monitored. The technique gave an indication of anodic, cathodic or

mixed electrochemical control.












Another type of artificial pit was developed by Greene and

Fontana [71] and used for a similar purpose [72]. Their artificial

pit consisted of a sheet cathode with a hole in it, through which

passed a small diameter anode wire, electrically insulated. It was

hoped that such an arrangement would approximate the geometry of an

actual pitted metal surface, neglecting pit interaction. Murray [73]

demonstrated the use of a somewhat similar artificial pit, using a

drilled plate cathode backed by an anode plate, electrically insulated.

The general result of Brown and Mears' work [69] and Greene

and Fontana's work [72], for stainless steels, seems to have been

that electrochemical control could change, depending upon such factors

as time, presence of depolarizers and extent of propagation. As

pointed out by Greene and Fontana [5], the presence of local-action

currents on any artificial pit could lead to questionable results.


2.4.2.3 Potential distribution in and around pits


Herbsleb and Engell [74] used a fine Haber-Luggin capillary

(0.08 mm) to measure the potential inside pits under potentiostatic

control. They found that for an iron electrode in a 1 N H2S04 +

0.003 M Cl solution, potentiostated to +1.230 vSHE, the potential

inside an active pit was as low as +0.100 vSHE. Rosenfeld and Danilov

[36] measured a substantial variation in electric field strength as a

function of position in the electrolyte just outside of an active pit.

Pickering and Frankenthal [75], also using a fine Haber-Luggin capillary

probe, measured a potential drop corresponding to their calculated

electrolyte IR drop ( 100 my) as they lowered the probe into an












active iron pit. However, within 50 vm of the bottom of the pit, they

observed larger potential drops and sharp transients that were 1.2 v

more negative (active) than the applied potential. They also observed

H2 gas bubbles eminating from iron and stainless steel pits under

similar conditions of applied potential.

Akimov [76] discussed the potential drop due to current flow

through a solution for a pair of electrochemically coupled electrodes.

This is similar to the potential drop between the auxiliary and working

(specimen) electrodes in a potentiostatic circuit. Piontelli [77]

showed that "additional resistance" and "screen" effects, due to

positioning of the Haber-Luggin capillary and current flow, could

introduce significant experimental errors into measured potentials.


2.4.2.4 Electrolyte composition inside pits


Pryor [78] used X-ray diffraction to identify the corrosion

products of steel after immersion in several electrolytes exposed

to air. He found a and y-FeO.OH to be a common constituent.

Fontana and Greene [1 ] noted that "corrosion tubes" could

form on ferrous alloys due to the action of dissolved oxygen in the

electrolyte. Here X-ray diffraction showed the presence of several

layers or "rings" of iron in various oxidation states.

Heyn and Bauer [79], and later Foroulis and Uhlig [80],

identified iron carbide (Fe3C) as a corrosion residue. Herbsleb and

Engell [81] identified Fe C in pits of mild steel, but no solid corrosion

products in pits of spectroscopically pure iron. Staehle et al. [82]

observed the stability of iron carbide structures to be a function of

potential, pH and anion.












Strehblow and Vetter [83] developed a technique for transferring

a sample with active pits into an electron probe microanalyzer. The

pits were grown potentiostatically on iron samples in chloride and

chloride/sulfate solutions. In order to prevent chloride precipitation

inside the pits, it was found necessary to limit the diameters of the

pits to between 5 and 10 pm.

Enrichment in chloride was measured inside previously active

(5-10 pm) pits, but no solid corrosion products were observed for

chloride solutions. A thick, porous sulfate layer was observed over

the chloride layer for chloride/sulfate solutions.


2.4.2.5 pH changes in pits


Ideally, it would be best to analyze for the chemical composition

of actual "in-service" pits. In 1926, Baylis [84] carried out such

an analysis of the tubercles on the inside of cast-iron pipe exposed

to a public water system. He noted pH's of about 6 and the presence

of Fe, chlorides, sulfates and other constituents.

The difficulty of making in situ analyses on actual corrosion

pits has led many investigators to the use of artificial pit devices.

These devices, although varying in size and design, usually possessed

a common feature: an anode compartment of restricted volume separated

from the bulk or cathode compartment by some sort of diffusion barrier

(porous ceramic material, parchment or paper, asbestos fibers or

fritted glass). Porous debris [85], corrosion products [86, 85] and

migrational screening effects [ 5 ] have been viewed as possible

diffusion barriers in real pits.













2.4.2.6 Chemical analysis of artificial pits


Evans [85] used his "parchment cell apparatus" to study

differential aeration conditions. He noted a much larger accumulation

of iron in the anode unaeratedd) compartment. He did not determine

the final pH's of the compartment electrolytes.

Uhlig [46] employed a porous cup, a piece of 18% Cr 8% Ni

stainless steel pipe and a graphite cathode to simulate a natural pit.

He passed a constant direct current through a 5% NaC1 aqueous solution,

with the stainless steel anode inside the porous cup. After the

experiment, the anode was found to have pitted over its entire

surface.

Analysis of the anolyte solution revealed no preferential

dissolution of the alloy constituents (Fe, Cr, Ni). Chloride ion

increased in concentration within the porous cup and decreased

outside. The catholyte reached a pH of about 11 and the anolyte

attained of pH of less than 2.

Parsons et al. [87] constructed artificial pits using Fe anodes

and Cu cathodes. Their starting solutions contained 9 parts of 0.1 N NaCI

and 1 part of 0.1 N NaOH, giving an initial pH of about 12. A yellowish

membrane formed between the electrodes. In those cases where the

membrane remained intact, it acted as a diffusion barrier.
t1-
Analyses of cells forming a single membrane yielded Fe

concentrations of 0.085 M to 0.276 M, Cl concentrations of 0.154 M to

0.324 M and pH's of 5.8 to 6.3. Analyses of a few multiple-

membrane pits gave slightly lower pH's and nearly double the Fe-

and Cl concentrations of single-film pits.












Pryor [78] employed an artificial pit device, making use of

a porous alundum pot, to study the behavior of steel anodes. He

applied a constant anodic current for a period of 24 hours and then

analyzed the anolyte. The resulting pH's were found to vary from 3.4

to 6.7. In those solutions containing chlorides, 230 to 290 percent

increases in chloride concentration were observed.

More recently, Wilde and Williams [88] devised an artificial

pit anode consisting of an AISI Type 304 stainless steel wire embedded

in epoxy. The sample was potentiostated at various potentials in

1 M NaC1 at 250C. Each sample, after having been maintained at a given

potential for an appropriate time, was quickly frozen in liquid nitrogen.

The frozen plug of corrosion products was then melted and analyzed

for pH and Fe /Fe+ concentration ratio.

For applied potentials ranging from +0.500 to -0.200 vSCE, they

observed pH's ranging from 3.6 to 0.0 and Fe /Fe+ concentration

ratios ranging from 1.59 to 9.8. An unexpected result was that more

noble applied potentials resulted in higher pit pH's (lower H

concentrations).


2.4.2.7 Hydrolysis in pits


In a general sense, hydrolysis has been defined as the

reaction of any substance with water [89]. In a more restricted

sense, it has been defined as the reaction of an ion with water to

form an associated species plus H or OH [89]. Cation hydrolysis
+
can result in an increase in H concentration and thus a reduction

in pH.












Whitman et al. [90] attributed the nearly constant corrosion

rate of steel in natural waters of pH 4 to 10 as being due to the

formation of a constant pH of 9.5 at the surface of the steel (relating

to the solubility product of Fe(OH)2).

A number of authors [28, 46, 47, 27, 88] have noted that iron

and chromium chloride solutions tended to be acidic. There appears to

be no general agreement as to what hydrolysis products form from iron

and stainless steel reaction products in chloride solutions.

Uhlig [46] found the corrosion products of an 18% Cr 8% Ni

stainless steel (driven anodes in NaCI solutions) to be Fe Cr

and Ni chlorides. Stolica [60] suggested that, according to his data,

Cr was formed as a primary reaction product and was oxidized to

Cr by the electrolyte.

Pourbaix [27] attributed hydrolysis of iron to the formation

of Fe304 and Fe(OH)3. Wilde and Williams [88] tabulated a number of
44 . 4 +++
possible hydrolysis reactions involving Fe Fe Cr Cr

and Ni but were unable to decide which reaction or reactions

predominated.


2.4.2.8 Electromigration and diffusion


The observed tendency for chloride ions to concentrate in

active pits may be accounted for in two ways:

1. Chloride ions react directly or indirectly with metal

inside pits to form a corrosion product or associated

species, resulting in a drop in chloride ion activity

in the pits. Inward diffusion of chloride ions follows.

2. The rapid production of positively charged metal ions












inside pits tends to cause a local charge imbalance.

This must be offset by the outward movement of cations

or the inward movement of anions.

The first possibility would seem difficult to prove, due to

the known high solubilities of many metal chlorides in water [91, 92,

93]. In addition, if operative, such a process should not result in

any increase in chloride ion activities in pits.

The second possibility can be dealt with using the concept of

transference or transport numbers. Assuming a current passing through

an electrolyte, transference numbers can be defined as follows [94]:

t = fraction of current carried by cation,

t = fraction of current carried by anion,

where t + t = 1.

This definition could be extended to include transference numbers

for each individual cation and anion present in the electrolyte.

The anions in solution tend to migrate toward the anode under

the influence of an applied electric field and the cations tend to migrate

toward the cathode. If particular ions are discharged at the appro-

priate electrode, a buildup in concentration may not occur. If

discharge does not occur, there may be a buildup in concentration

near the appropriate electrode.

A situation can be envisioned where the transference

number of a particular anion, such as the chloride ion, might be high.

At potentials below that required for chlorine evolution, chloride

ions should concentrate themselves in the vicinity of the anode.

Local electrical neutrality (down to molecular dimensions) must be












maintained in an electrolyte [94]. Thus, a large concentration

of metal ions in a pit must be accompanied by enough anions to assure

an equal number of positive and negative charges.

The ionic mobility is defined as the average velocity with

which an ion moves toward an electrode under the influence of a

potential of 1 v applied across a 1 cm cell [94]. H+ and OH- ions

have the highest ionic mobilities at infinite dilution in aqueous

solutions, with other ions possessing lower but about equal mobilities

[94]. It would appear that the transference number for a given ion

would depend on both availability (concentration) and mobility. This

would imply that an increase of chloride ion concentration in a pit

would be due to the requirement for electrical neutrality and to ready

availability, rather than to an inherently high ionic mobility.

A number of authors have suggested that migration [46, 47, 78, 95]

and diffusion [47, 78, 17] effects were instrumental in the changes

of electrolyte composition occurring in active pits. Other authors

[96, 97, 98, 5] discussed an ion-screening mechanism with reference

to the diffusion of oxygen into a pit. Their view was that corrosion

products reacted with oxygen and hindered or prevented it from entering

the interior of a pit. This concept might be applied to other species

in solution, such as the OH ion for pits on ferrous alloys (due to

the formation of Fe(OH)2).


2.4.2.9 Stability of pitting


Franck [99] showed that there should be two stable stationary

states (active and passive) for an open system involving an active-












passive metal or alloy. Franck also showed (mathematically) that

only one stationary state should be stable at a time. This implied

that pitting corrosion should be an unstable process, since both

active and passive areas existed simultaneously.

Realizing that pitting corrosion did exist, Franck made further

assumptions. Using an analog computer simulation of pitting, involving

two variables, he was able to show mathematically that there were

cases where pitting corrosion was stable, unstable, or oscillatory.

The presence of a very high-resistance polishing film was postulated,

whose rate of dissolution determined the corrosion current. The true

potential of a pit was viewed as lying in the active corrosion region.

Brauns and Schwenk [100] suggested that perhaps both the

resistance effect and the electrolyte composition change (concentration)

effect were required for stable pitting. Schwenk [101] later showed

that the concentration effect could lead to stable pitting under

galvanostatic conditions (only at the pitting potential) and that

the resistance effect could lead to stable pitting under potentiostatic

conditions at any potential within the pitting range. In this [101]

and another work [58] Schwenk again noted that both the resistance

and concentration effects might be necessary for stable pitting

corrosion.

Rosenfeld and Danilov [36] helped to confirm this idea by

showing that pits could be deactivated by destroying the protecting

layer over the pits, thus allowing access to the bulk electrolyte.

Pickering and Frankenthal [75] presented evidence that the

potential in a localized corrosion cell (pit or crevice) was in the












active part of the anodic polarization curve of the metal, even when

the specimen surface was potentiostated to more noble potentials.

They attributed this effect to the presence of high-resistance paths

in the form of H2 gas bubbles within pits or crevices.

Hoar et al. [102] held the opinion that the anode potential

within pits could not be more than a few my different from that

measured and controlled (potentiostatically) outside. They based

their conclusion partly on the fact that crystallographic etching (often

found with active dissolution) was suppressed, with anodic brightening

type of attack occurring instead (characteristic of higher anode

potentials). The authors suggested the presence of a "contaminated

oxide" compact solid film of high cation conductivity on the pit anode.

At about the same time Vetter [42] expressed similar doubts

about the validity of the resistance effect. He also developed a

thermodynamic theory for the stability of pitting, postulating the

coexistence of passivating oxide layers and nonpassivating salt layers

(Cl-, Br I ).


2.5 Crevice Corrosion


As in the case of pitting corrosion, two stages or steps in

crevice corrosion are generally recognized: initiation and

propagation [103]. Bombara [104] noted that initiation and propagation

were two well-distinguished stages in stress corrosion cracking

and in many forms of localized attack.

A number of authors [105, 34, 1, 106, 107, 75] have noted

similarities between the propagation stages of pitting and crevice













corrosion. Schafer et al. [103] stated that they were identical

processes. Bombara [104] went so far as to suggest that a strict

analogy might exist between the electrochemical mechanisms of localized

attack and stress corrosion cracking in both the initiation and pro-

pagation stages. For this reason, some results from the literature

pertaining to stress corrosion cracking will be included here.


2.5.1 Initiation of Crevice Corrosion


France [106] recently reviewed the literature of crevice

corrosion (Rosenfeld and Marshakov [108] had done so earlier). He

categorized the currently popular mechanisms under a general heading

of "concentration cells." The categories included metal-ion cells,

differential aeration cells, active-passive cells and other concen-

tration cells. The final category included hydrogen-ion, neutral

salt and inhibitor cells. Primary emphasis will be placed on crevice

corrosion of previously passive metals.


2.5.1.1 Metal ion cells


The metal ion cell is a special case of the electrolyte

concentration cell. A concentration cell is an electrolytic cell,

the e.m.f. of which is due to differences in composition of the

electrolyte at anode and cathode areas [4]. McKay [109] studied such

a cell in 1922 by measuring the potential between a pair of Monel metal

electrodes (70% Ni 30% Cu) in sulfuric acid solutions of different

concentrations of cupric sulfate, separated by a porous cup.

By the use of current and weight-loss measurements, McKay


Trade name of International Nickel Co.
Trade name of International Nickel Co.












verified that the Monel electrode immersed in the solution of lower

copper ion concentration was the anode. He correlated these results

with those predicted by the Nernst equation for pure Cu. In addition,

McKay demonstrated that corrosion could be produced by cells due to

differences in the concentration of acids, concentration of dissolved

oxygen or hydrogen, or concentration of dissolved oxidizing and reducing

salts.

Evans [85] made some calculations on the possible effects of

stirring, using the Nernst equation. Assuming fixed electrode potentials

of +0.1, +0.2, and +0.3 vSHE, he calculated the equilibrium concen-

tration of metal ions (assuming the Nernst equation to be valid) for

several metals. The results showed that equilibrium concentrations

could be attained for Ag and Cu, but that equilibrium concentrations

for Pb, Cd, Fe and Zn (according to the Nernst equation) were clearly

unattainable.

In a later paper, McKay extended an existing definition of

concentration cell to

...cells whose e.m.f. is set up by two electrodes of the
same material in different electrolytes. [110, p. 23]

LaQue et al. [19] used the metal ion cell concept to explain a type

of crevice corrosion observed on Cu alloys, where the attack was

limited to the area just outside a crevice. Schafer and Foster [111]

argued that the metal ion cell, in order to operate, would have to

violate the laws of chemical diffusion. They suggested that the type

of attack referred to by LaQue et al. was in fact a special case of

differential aeration.











2.5.1.2 Differential aeration cells


Differential aeration cell has been defined as an electrolytic

cell, the e.m.f. of which is due to a difference in oxygen concentra-

tion at two otherwise similar electrodes [4]. McKay [109, 110]

recognized the existence of such a cell in the early 1920's.

Evans [85, 112, 96, 113] treated the subject in a number of

his papers. It was established that, in general, the electrode

immersed in the solution of lower dissolved oxygen concentration

became the anode and that the electrode immersed in the solution of

higher dissolved oxygen concentration became the cathode [85, 110, 96,

113]. More recently, Korovin and Ulanovskii [114] have studied the

effects of dissolved oxygen concentration on the electrode potentials

of stainless steels. Schafer et al. [103] noted that the differential

aeration cell mechanism should be operative primarily only in the

initiation stage of crevice corrosion.


2.5.1.3 Passive-active cells


Passive-active cell (or active-passive cell) has been defined

as a cell, the e.m.f. of which is due to the potential difference

between a metal in an active state and the same metal in a passive

state [4]. Uhlig [20] attributed the formation of passive-active

cells in chloride solutions to the localized breakdown of passivity

by chloride ions. France [106] discussed the mechanism in terms of
+2
a localized deficiency in oxidizers (e.g., dissolved oxygen, Cu+ or

Fe 3).

Although the concept of localized cells being responsible for

the initiation of crevice corrosion is well established, another












possible way of viewing the subject for the case of previously passive

metals in chloride solutions (or solutions containing other aggressive

anions) presents itself. This is by way of the concept of critical

parameters for the breakdown of passivity.


2.5.1.4 Critical parameter concept


The concept of critical parameters for the breakdown of passivity

was discussed in Section 2.4.1.4. It was noted then that in solutions

containing chloride or other aggressive anions, breakdown of passivity

could be attained by reaching a critical potential (rupture potential),

a critical chloride concentration (or aggressive anion concentration)

or a critical pH.

Thus, the initiation of crevice corrosion on a previously

passive metal might be affected by raising the potential of the metal

above (more noble than) the critical pitting or rupture potential in

an existing crevice. The potential might be maintained electronically,

through the use of a potentiostat, or chemically, by the use of oxidizing

agents. The role of dissolved oxygen might then be viewed as twofold:

to maintain the potential (in the absence of any applied current) and

to alter the rupture potential.

The data of Wilde and Williams [115] indicated a higher (more

noble) rupture potential for aerated solutions than for deaerated

solutions, although a more exhaustive study might be required to

verify the point. Accepting their data, one could envision the

rupture potential in an existing inactive crevice to be lower (due to

a lower concentration of dissolved oxygen) than that outside the crevice












in solutions containing dissolved oxygen. If the potential were

maintained between the two rupture potentials (more noble than the

crevice rupture potential), breakdown of passivity could be

expected to occur inside the crevice.

If the chloride ion concentration inside an existing inactive

crevice were increased while the potential was maintained at a constant

value, a critical concentration necessary for the breakdown of passivity

might be reached. Such a concentration increase could conceivably

result from the electromigration of chloride ions accompanying the

small corrosion current present in the passive state. This type of

mechanism would depend upon the relative transference of the various

ions present in solution.

If the pH inside an existing inactive crevice were lowered

while the potential was maintained at a constant value, a critical

pH necessary for the breakdown of passivity might be attained. Such

a pH decrease might result from the hydrolysis of metal ions, slowly

formed by the dissolution process in the passive state. This type of

mechanism would depend upon the relative diffusion rates of the

various ions present in solution. In reality all three processes, or

any combination thereof, might operate simultaneously.

Some aspects of this approach follow closely the reasoning

put forth by Fontana and Greene [ 1, 106] in a unified mechanism of

crevice corrosion (e.g., chloride ion migration and metal-ion hydrolysis),

but the additional possibility of critical parameters may be useful for

better understanding of the initiation of crevice corrosion on previously

passive metals in solutions containing aggressive anions.













2.5.1.5 Induction periods


Another aspect of the initiation stage of crevice corrosion

should be mentioned: the existence of an induction period for the

initiation of crevice corrosion. France [106] attributed the existence

of an incubation period to the time required for consuming the oxygen

in a shielded crevice area. Indirect experimental verification of

this process was obtained by Karlberg and Wranglen [116] through the

use of an artificial crevice.

The critical parameter concept would also be amenable to the

existence of an induction period for crevice corrosion initiation.

Wilde [117] claimed the induction period for crevice corrosion to be

shorter than that of pitting. Fontana and Greene stated that pitting


...is a self-initiating form of crevice corrosion.
...it does not require a crevice--it creates its own.
[1, p. 54]


The existing geometry at a crevice helps to rationalize a longer

induction period for the initiation of pitting than for the initiation

of crevice corrosion.


2.5.2 Propagation of Crevice Corrosion


As was mentioned at the beginning of Section 2.5, the propaga-

tion stages of pitting and crevice corrosion (and perhaps other forms

of localized corrosion) appear to proceed by similar mechanisms.

For this reason a subject order similar to that used in the section

on the propagation stage of pitting will be followed.

Rosenfeld [118] reviewed the literature of crevice corrosion,













with particular emphasis on the propagation stage. One observation

that he made was that it was not clear just which hydrolysis process

was responsible for acidification of the electrolyte in a crevice.


2.5.2.1 Geometry of crevice corrosion


The name given to the type of attack presently under considera-

tion implies the presence of a small interstice or crevice. In

practice, it can be very difficult to eliminate unwanted crevices.

It has been shown that some of the materials commonly used for mounting

or partial incasement of electrochemical samples can form crevices

at material/sample interfaces [119, 120].

Crevice corrosion has been known to occur under a number of materials,

including fabrics, plastics, wood, porcelain, glass, rubber, wax, mica,

asbestos, packing (especially when graphite was present) and marine

growths [19]. Rosenfeld and Marshakov [121] studied the nature of attack

produced under a series of dielectrics on the surface of iron and between

two homogeneous iron samples, and observed no differences. Some crevice-

forming substances may contain reactive constituents such as

sulfur in sulfur-bearing rubbers [19].

The geometry, or shape and size, of a crevice can be important

in characterizing crevice corrosion. This is due partly to the effect

of geometry on diffusion and electromigration processes and partly to

the effect of geometry (relative surface distribution of anodic and

cathodic areas) on electrochemical polarization.

Ellis and LaQue [122] studied area effects on crevice

corrosion of an Fe 17% Cr stainless steel exposed to fresh sea












water. Using machined samples of known crevice dimensions, they found

both the average weight loss and the average maximum pit depth to be

directly proportional to the exposed area of sample outside the crevice.

Other authors have shown that intensity of attack within a

crevice may vary with width or depth. Rosenfeld and Marshakov [108]

noted that a maximum of corrosion intensity occurred at an intermediate

crevice width. Bombara et al. [123] showed that a maximum rate of

crevice attack occurred at a depth which depended upon the externally

applied potential and the crevice width. Vermilyea and Tedmon [124]

made calculations based on experimental work which predicted how the

metal-ion concentration and potential should vary with depth for a

simple type of crevice. Some of the mathematical treatments of crevice

corrosion are quite involved and fraught with simplifying assumptions.

Nevertheless, they can be instructive and useful by pointing out the

danger of intuitive reasoning.


2.5.2.2 Current distribution in and around crevices


The several types of current measurements that conceivably

might be made include the total net sample current (including crevices

and other areas), the total net crevice current, and net local currents

inside and outside crevices.

Measurements of the total net sample (working electrode)

current can be made using standard potentiostatic or potentiokinetic

apparati. These were the kindsof measurements made by Greene et al.

[119], and Lizlovs and Bond [120]. The presence of crevices resulted

in higher measured applied currents in passive potential ranges.












Lizlovs [125] designed a more reproducible crevice assembly,

in which the crevice could be introduced or removed at will. Large

current increases were frequently observed when the artificial crevice

was introduced.

The "scratch technique" of Brown and Mears [68-70] was briefly

discussed in Section 2.4.2.2. Their scratches (in the wax) might be

considered to be shallow artificial crevices. Thus, their conclusions

regarding the observation that all types of electrochemical control

could be observed in pitting corrosion probably also apply to crevice

corrosion. If this is true they were, in fact, measuring the total

crevice currents of artificial crevices.

Rosenfeld and Marshakov [108] presented a plot of current

vs. depth in crevice (for an Fe 13% Cr alloy in 0.5 N NaC1) [There

is some question as to the alloy composition.] which showed the crevice

edges to be functioning as cathodes and the interior of the crevice to

be functioning as anode. The authors also employed a sectional artificial

crevice to measure current and potential as a function of depth inside

the crevice.


2.5.2.3 Potential distribution in and around crevices


The potential, as measured experimentally, of an electrode

possessing an active crevice, may depend upon the position of the

measuring apparatus (i.e., the position of the Haber-Luggin capillary)

[77, 76, 67]. The measured potential may be that of a "composite

electrode" [67] (a sort of average potential) or that of a local

electrode process (a local potential). The measured potential

conceivably may also fall between these two extremes.












Bombara et al. [126, 104] determined potentiokinetic polarization

curves for several austenitic stainless steels at 40C in solutions of

0.75 M H2SO4, with and without 0.01 M NaC1. They then maintained

samples potentiostatically for appropriate times and later examined

them metallographically. Combining the data, they delineated potential

ranges according to attack morphology.

Larin and lofa [127] expressed an opinion that, for the case

of the crevice corrosion of iron at three-phase boundaries, the potential

of the anodic area could not be very different from that of the

cathodic area, due to their close proximity. This statement of

opinion was similar to that expressed by Hoar et al. [102] for the

case of pitting corrosion.

Localized anodic and cathodic polarization curves (potential

vs. current density) were determined for iron in 0.5 N NaCl by Rosenfeld

and Marshakov [108]. A sectional artificial crevice was employed and

curves were presented for several depths inside the crevice. They found

the anodic process to be greatly accelerated and the cathodic process

to be retarded inside a narrow crevice.

Somewhat related experiments were conducted by France and

Greene [128] in 1 N H2SO for Type 304 and 18% Cr 12% Ni stainless

steels. They used an electrode assembly fabricated from a plastic

block, with internal probe openings spaced at 0.5-in. intervals. They

maintained the potential of the external crevice face potentiostatically.

Potential differences of up to approximately 1 v were measured in

some cases.

Another type of artificial crevice was developed by Vermilyea











and Tedmon [124]. Their (calculated) potential differences for iron

were smaller than those observed by France and Greene [128], of the

order of 100 mv. Similar results were obtained directly by Karlberg

and Wranglen [116] for stainless steels, using yet another type of

artificial crevice. Pourbaix [129, 27] presented data relating

potential and pH inside an artificial crevice to the potential and

pH outside the crevice for an ordinary carbon steel. Efird [130] used

an improved artificial crevice design to obtain data for a Cu-Ni alloy.

A. Pourbaix [131, 132] used a modified version of the artificial

crevice employed by M. Pourbaix [129, 27] to determine simultaneous

polarization curves for duplex (active and passive segments) steel

electrodes in 10 M NaOH and either 10-3 or 10-0.3 M NaC1 solutions.

The device consisted of a large "bulk" (cathode) compartment and a

small "crevice" (anode) compartment, separated by a diffusion barrier

in the form of a hole filled with asbestos fibers.

Pickering and Frankenthal [75] observed H2 gas bubbles

emanating from pits and crevices on iron and stainless steel samples,

even with the external applied potential as noble as +1.4 vSHE. Similar

observations were made by Rhodes [133] for the case of stress corrosion

cracking. B. F. Brown et al. observed potential differences between

stress corrosion cracks and unaffected material for freely corroding

[134, 135] and potentiostatically controlled [135] samples. They

studied 4340 [135] and other alloy steels [134], and found the potential

of the stress corrosion crack in every case to be below that of the

equilibrium hydrogen electrode.












2.5.2.4 pH changes in crevices


A number of authors have reported pH changes (from the bulk

solution pH) in crevices [136, 137, 27, 116] and stress corrosion

cracks [138, 134, 135]. Most of the work concerning crevice corrosion

was done with the aid of artificial crevices, but the work concerning

stress corrosion cracks was performed in situ on real cracks.

Ulanovskii and Korovin [136] reported attempts to make direct

pH measurements on crevices, using both a potentiometric method and

an indicator method. Details were not given, but a minimum pH of

about 3 was noted for an Fe 13% Cr stainless steel. They then

polarized specimens in a cell of limited volume (about 5 ml of electro-

lyte) and monitored pH. The pH was found to decrease towards a steady-

state value, both with increasing applied current (for a constant time)

and with increasing time (for a constant current). A pH of about 2

was attained after a period of 24 hours with an applied current density

of +0.20 ma/cm2

Peterson et al. [137] used a thymol blue indicator to make

direct pH measurements of active crevices on a Type 304 stainless steel

exposed to sea water. They estimated the pH to range from 1.2 to 2.0.

The authors also found that cathodic protection yielded alkaline pH's,

above 10 for specimens coupled to zinc anodes (probably due to H2 gas

evolution).

Pourbaix [27] showed, on potential vs. pH diagrams, a dependence

of the pH of an artificial crevice on the externally applied potential.

He reported a steady-state crevice pH of about 4.7 for ordinary

carbon steel with an applied potential (due to oxygen reduction) of about












-0.1 vSHE. Pourbaix observed that the crevice pH could become either

more acidic than or more basic than the bulk solution, depending on

the externally applied potential. Acidification was attributed to

metal-ion hydrolysis and alkalization was attributed to H2 gas evolution.

Karlberg and Wranglen [116] measured the pH of a freely

corroding artificial crevice (after up to 14 days) on an Fe 13% Cr

stainless steel. The external potential was about -0.15 vSHE and

the crevice potential was about -0.3 vSHE. The measured pH was 4.0,

although the authors noted that the pH may have been lower locally.

Brown et al. employed chemical indicator [138, 135], indicator

electrode [135] and colorimetric methods [134] in their studies of

chemistry changes within actual stress corrosion cracks. Early work

[138] reported pH values at crack fronts of 3.5, 3.8 and 1.7 for an

Al alloy, a carbon steel and a Ti alloy, respectively, in solutions of

3-1/2% NaC1 by weight.

Later work [135] on an AISI 4340 steel showed the pH at crack

tips to be virtually independent of the bulk pH and that electro-

chemical conditions at the crack tips were favorable for H2 gas evolution

regardless of the externally applied potential. Work [134] performed

on a series of alloy steels (including an Fe 12% Cr alloy) in

solutions of 3-1/2% NaC1 by weight, under freely corroding conditions,

yielded pH values at the crack tips of about 3.7 for all steels.


2.5.2.5 Chemical analysis of crevices

Some of the authors cited in the previous section also attempted

chemical analyses of solutions within crevices [137, 130, 131, 132] and












stress corrosion cracks [138, 134]. Brown et al. employed chemical

indicator [138, 134] and colorimetric [134] methods to study the

solution chemistry of actual stress corrosion cracks. Their chemical

indicator work identified Fe but not Fe in the crack solution

of a 0.45% C steel exposed to a 3-1/2% NaC1 solution by weight [138].

Similar results were obtained in later work [134] done on several

alloy steels. In addition, the alloying elements in the steels were

found in the solution near the crack tip (analyzed colorimetrically)

in approximately the same proportions as in the original steel

compositions.

Peterson et al. [137] used a similar chemical indicator technique

to analyze qualitatively for the presence of Fe++ and Fe+ in crevice

solutions of a Type 304 stainless steel exposed to sea water. They

noted a strong indication for Fe and only a trace for Fe

Any buildup of ions in crevices (or other occluded cells) may

result in indirect effects. One of these is a decrease in oxygen

solubility with increasing concentrations of salt solutions [139, 140].

This may tend to decrease the rate of transfer of oxygen into an

active occluded cell.

The topics of hydrolysis, electromigration and diffusion, and

stability (of pitting) were discussed in the section on the propagation

of pitting (2.4.2). The pertinent remarks made there should also

apply, at least qualitatively, to the propagation stage of crevice

corrosion. Consideration should be made of the different geometries

involved.












2.5.2.6 Hydrolysis in crevices


It was implied in Section 2.5.1.3 that the concentration of

A-H-
oxidizing agents (e.g., dissolved oxygen, Cu or Fe ) in an active

crevice or occluded cell should be lower than that in the bulk solution.

Whether or not oxidizing conditions are maintained would depend on the

relative rates of oxidizer reduction inside the crevice and mass

transport into the crevice.

The reason for concern as to whether or not oxidizing conditions

exist in a given occluded cell is that hydrolysis reactions depend on

the oxidation state of any ionic species involved. The only general

statement that can be made is that many authors [108, 106, 118, 75]

consider the dissolved oxygen concentration inside an active crevice

or occluded cell to be low.

As was done in the case of pitting corrosion, the acidity

frequently found in crevice solutions has been attributed to metal or

metal-ion hydrolysis [108, 137, 27, 116]. Similar conclusions were

reached for the case of stress corrosion cracking [138, 133, 134, 135].

There has, however, been no general agreement as to the responsible

hydrolysis products.

Several authors [108, 133, 27, 116] considered only solid

reaction products. Other authors [138, 134, 135] considered only

soluble hydrolysis products. Finally, Peterson et al. [137] did

consider a number of possible solid and soluble hydrolysis products.

In addition, these authors suggested that the most likely reaction

under their particular conditions (Type 304 stainless steel in sea

water) was the hydrolysis of Cr












Brown et al. [134] measured the acidity of freshly prepared

aqueous solutions of 1 M FeC12 4H20 (Parsons et al. [87] may have done

a similar experiment earlier). These solutions were prepared both

with boiled (to lower dissolved oxygen content) and unboiled

distilled water. The initial pH's of both solutions were 4.0. The

pH of the solution prepared with boiled distilled water slowly drifted

to a value of about 3.2. The unboiledd" solution attained about the

same pH, but much more rapidly. The authors then interpreted their

pH results on stress corrosion cracks of alloy steels as being

due to the hydrolysis of partially oxidized (by dissolved oxygen)

ferrous ions.


2.6 Limitations of Laboratory Tests


One criticism of relatively rapid laboratory corrosion tests

(both freely corroding and with externally applied current) has been

that results might not correlate with in-service tests. Induction

periods have been observed for the initiation of occluded cell

corrosion, leading some authors [141, 18] to question the validity

of short-term laboratory corrosion tests. Other authors [13, 107] have

attempted to show that short-term data could be applicable to longer

exposures.

France and Greene [95] questioned the correspondence between

chemically and electrolytically induced pitting corrosion. Brigham

[142], and Wilde and Williams [143] appear to have countered their

arguments.

The extrapolation of data obtained for salt solutions or

synthetic sea water solutions to an actual sea water environment is












generally not encouraged. This attitude arises from the known fact

that salt solutions and synthetic sea water solutions may be different

in their "corrosiveness" from fresh, live sea water [144, 145].


2.7 The "Protection Potential"


"Tension de protection," translated as "protective tension"

[3] or "protection potential," [27] was a term coined by M. Pourbaix

et al. [52] in 1962. The original work was done for Types 410, 304

and 316 stainless steels [52, 146, 49, 147, 3]. Work also was done

involving the "protection potential" for copper [148, 149, 150, 27].


2.7.1 Definition and Verification


The "protective tension" (later the "protection potential")

was defined in the following manner:


If after polarizing the steel at an increasing
tension beyond this breakdown tension it is then
polarized at a decreasing tension, a tension
[potential]/current curve is obtained which does not
coincide with the curve for increasing tension, but
lies considerably below it. The curve meets the
ordinate axis (with a practically zero current
density) at an electrode tension representing a second
well-defined critical value, which we have called the
"protective tension" Ep, below which the steel stops
corroding. [3, p. 245]


Three electrode potential regions have been recognized

[151, 143]:

(1) Above (more noble than) the rupture potential, new pits

can initiate and propagate and existing pits can continue

to propagate.

(2) Between the rupture potential and the "protection












potential," existing pits can continue to propagate but

no new pits can initiate.

(3) Below (more active than) the "protection potential,"

pits can neither initiate nor propagate.

Pourbaix et al. [52, 3] conducted a series of chronoamperometric

experiments (monitored current as a function of time at a fixed

potential) for a Type 410 stainless steel. They observed that the

applied current for actively pitting specimens (whose potentials had

previously been raised above (more noble than) the rupture potential)

approached low values when the potential was lowered below (more active

than) the "protection potential." Physical examination of the specimens

after conclusion of the experiments confirmed the current measurements.

Verink et al. [152] employed a diamond scribe to further

verify the "protection potential" concept for an Fe 16.9% Cr alloy.

It was observed that scratches made on the surface of the sample, in

solution, would not propagate as active pits at potentials below

(more active than) the rupture potential for the alloy. It was also

observed that scratches, activated at potentials above (more noble

than) the rupture potential, would cease to propagate when the

potential was lowered below (more active than) the "protection

potential."

It should be noted that completely meaningful "protection

potential" data can be obtained only in solutions containing negligible

quantities of dissolved oxidizing species [3, 153, 154].


An electrochemical hysteresis technique [153, 154] (a cyclic


2.7.2 Experimental Methods












potentiokinetic method), developed by Pourbaix et al. [52, 49, 3] has

been the usual method used for determining "protection potentials."

Details of this method will be given in Chapter 3.

Szklarska-Smialowska and Janik-Czachor [50] discussed various

potentiostatic, potentiokinetic, galvanostatic and galvanokinetic

methods for the determination of characteristic pitting potentials.

The characteristic pitting potentials were a critical potential for

pit initiation (rupture potential) and a critical potential for pit

propagation ("protection potential").

Up to the present time the experimental work has primarily

been done on iron-rich alloys (Fe [129, 27], Fe-Cr [54, 55, 153, 154]

and stainless steel [52, 49, 3, 88, 155]), copper-rich alloys (Cu

[149, 150, 156], Cu-Ni [157, 156] and Cu-Zn [158, 159, 160]) and

nickel-rich alloys (Ni-Cu [156]).


2.7.3 Interpretation of the "Protection Potential" for Cu-Rich Alloys


One result of the research mentioned above was that the

"protection potential" for copper and copper-rich alloys was observed

to be relatively pH-independent [157, 158, 159, 160, 156] and approximately

equal to +0.200 vSHE [149, 150, 157, 158, 159, 160, 156]. This

potential appeared to correspond to the potential for the reaction

Cu + Cl = CuCl + e [158, 159, 160] (assuming an activity for Cl of

0.1), as given by Pourbaix et al. [161,149].

The "protection potential" of copper and copper-rich alloys

may also correspond to the potential for the reaction


Cu = Cu++ + 2e


[149, 129, 156] .












Lee [156] investigated this possibility for a Cu 10% Ni alloy in

0.1 M chloride solutions and obtained good agreement.


2.7.4 Effects of Experimental and Environmental Variables on
the "Protection Potential"


The effects of a number of experimental and environmental

variables on the "protection potential" have been investigated for iron-

rich alloys [52, 49, 54, 27, 88, 155]. These are described

chronologically in Table 2.

Pourbaix et al. [52] observed "protection potential"

results for Type 410 stainless steel to be nearly the same for

mechanically and electrolytically polished specimens. The authors

also observed lower (more active) "protection potentials" for

samples which had received a sensitizing heat treatment (10 minutes

at 1000-1050C).


2.7.4.1 Effect of alloy composition on E


The results of Pourbaix et al. [52], as given in Figure 3, showed

the "protection potential" to increase (become more noble) with

increasing chromium or nickel content (in the order Types 410, 304 and

316). Cusumano concluded that the "protection potential"


...rises to more noble values with increasing
chromium content. [55, p. 106]

The data of Wilde and Williams [88] also showed a similar trend, becoming

more noble in the order Type 410, USS 100, Types 430, 446, 304, 316,

and Hastelloy C.















Year Autl

1962 Pourbaix



1962 Pourbaix



1970 Verink [5'


1970 Pourbaix

1971 Wilde and
[88]










1972 Suzuki an
Kitamura


DETAILS OF INVESTIGATIONS

hors Alloy

et al. [52] 410
304
316

et al. [49] 410
304
316

4] Fe-0.5% Cr
Fe-2.0% Cr
Fe-5.0% Cr
Fe-12% Cr
Fe-16.9% Cr
Fe-24.9% Cr
[27] iron

Williams USS 100

410
430
446
304
316
Hastelloy C
430
304

d 316L
[155]


TABLE 2

OF THE "PROTECTION

Electrolyte Gas

0.1 M NaHCO3 N2


0.1 M NaHCO3



buffers






buffers

3.5 wt. %
NaC1







1 M NaCI


POTENTIAL" FOR FERROUS ALLOYS

pH [Cl] Prop. Temp. S.P.

X X
X
X


H X
2

X
X
X
X
X
N- X


chlorides stag- X
nant
air


H.T.

X
X












E
SCE

+1.0 -
8 I
+.4
+ -- -(-



0-
+.4 :^ --. -\^- --------- -s


4 r. 304

-.6 I
6 8 12 2 4 6 8 1 2 4 6 NaC
10 IO M NaCI


Figure 3. Effect of sodium chloride molarity on the
"protection potential" for Types 410, 304
and 316 stainless steels in 0.1 M NaHCO3,
N2-saturated solutions at 200C [52].




2.7.4.2 Effect of pH on Ep


Previous authors have reported the "protection potential" to

be either independent of [155] or only slightly dependent on [27, 54,

55, 153, 154] the bulk solution p1H.


2.7.4.3 Effect of [Cl ] on Ep


Pourbaix et al. [52] observed the "protection potential" for

some stainless steel alloys to become more active with increasing

bulk solution chloride concentration, as shown in Figure 3. Suzuki and

Kitamura [155] obtained qualitatively similar results, shown in Figure

4. Other workers [27, 54, 153, 154] have reported the "protection

potential" to be relatively independent of chloride concentration.




















E
SCE





-25-



316 L

-.30 o


.01 .03 .1 .3 1 3
CHLORIDE ION NORMALITY



Figure 4. Effect of chloride ion normality on the
"protection potential" for Type 316L
stainless steel in stagnant chloride
solutions at 70C [155, p. 5].











2.7.4.4 Effect of Temperature on E


Pourbaix et al. [49] presented a plot of "protection potential"

vs. temperature, reproduced in Figure 5, showing a decrease with increasing

temperature for Types 316 and 304 stainless steels, and a decrease

followed by an increase with increasing temperature for Type 410 stainless

steel. Suzuki and Kitamura [155] observed a decrease of the "protection

potential" with increasing temperature for Type 316L stainless steel,

as shown in Figure 6.


2.7.4.5 Effect of extent of propagation on Ep


Wilde and Williams [88] studied the effect of extent of

propagation on the "protection potential" for Types 430 and 304 stainless

steels. They observed that the "protection potential" decreased

(became more active) linearly with the increasing logarithm of the

charge passed (obtained by integration of cyclic potentiokinetic

polarization scans). Their results are shown in Figure 7.


2.7.5 Effect of Geometry on Ep


Suzuki and Kitamura [155] employed artificial occluded cells

to study the "protection potential" for pits, crevices and stress

corrosion cracks. They observed the "protection potential" for

pitting to be more noble than the "protection potential" for crevice

corrosion, and that the latter was approximately equal to that for

stress corrosion cracking.

Wilde concluded [117] that the breakdown of passivity on

stainless alloys occurred by either pit initiation or crevice initiation,

possessing differing kinetics. He also concluded that the initiation











E
SCE


-X 316

304

\ _4-10
e - -


-.6 I I I I I I 1 1 1 1 ;1
0 20 40 60 80 Temperature (C)



Figure 5. Effect of temperature on the "protection
potential" for Types 410, 304 and 316
stainless steels in 0.1 M NaCI + 0.1 M
NaHCO3, N2-saturated solutions [49].


t8



















E
SCE
0

-05

-.10

-15



316L
-25
"':2 5 316L


-:30
0 10 20 30 40 50 60 70
TEMPERATURE (C)




Figure 6. Effect of temperature on the "protection
potential" for Type 316L stainless steel
in stagnant 0.88 N NaC1 solutions [155,
p. 5].



















-7


o 304




02 16( 100 101 102 103 104 105
AMOUNT OF PROPAGATION COULOMBSS)

Figure 7. Effect of extent of propagation on
the "protection potential" for Types
430 and 304 stainless steels in
1 M NaC1, N2-saturated solutions at
250C. Initial specimen area:
5 cm2 [88, p. 1980].




kinetics of crevice corrosion appeared to be more rapid than those of

pit initiation.


2.7.6 Material Vs. Environmental Property


The "protection potential" has apparently not been established

as either a material property or as an environmental property. Some

authors [3, 27, 54, 153, 154, 155], although admitting the "protection

potential" might be somewhat dependent on both material and environment,

considered it to be potentially more useful than the rupture potential












(which was strongly dependent on environmental variables) in predicting

"safe" electrochemical potential ranges for ferrous alloys and stainless

steels in chloride solutions.

Other authors [31, 162, 117] considered the "protection potential"

to be unreliable. Leckie and Uhlig expressed the opinion that the

"repassivation" of active pits would depend


...on accidental favorable geometry of the pit cavities
and a stirring rate high enough to mix electrolyte
inside the pits with electrolyte outside. [31, p. 1267]


Pessall and Lui [163] tried to show that a true critical pitting potential

fell between the rupture potential and the "protection potential."

Pickering [162] considered the "protection potential" to be

unpredictable due to the constant egress of H2 gas bubbles from

active pits and crevices. Wilde went a step further and concluded

that the "protection potential" was not a material property and that it

varied


...in magnitude with the experimental conditions
used to evaluate it. [117, p. 290]


2.7.7 "Repassivation" Vs. "Deactivation"

A number of authors [31, 36, 125, 163, 164, 117] considered

the "protection potential" to involve some sort of "repassivation"

process. Other authors [27, 153, 154, 131, 132, 162] considered a

"deactivation" process to be operative. By "deactivation" it is

meant that the sample potential is lowered (made more active) enough

to cause it to become thermodynamically immune to corrosion (i.e., a

form of cathodic protection).












Pourbaix has interpreted the "protection potential" as being


...approximately equal to (or better slightly higher
than) the potential...inside these pits. [27, p. 432]


This appears to be equivalent to saying that the "protection potential"

is approximately equal to an instantaneous mixed potential involving

active metal dissolution and H2 gas evolution inside active pits

(implying a "deactivation" mechanism).


2.7.8 Experimental Correlation


Wilde and Williams [88] were able to obtain a linear correlation

between a "difference potential" (rupture potential minus "protection

potential") as measured in the laboratory and weight loss of stainless

steel alloys exposed to sea water. Their results are given in Figure 8,

although no definite theoretical basis is available at the present time.



















_L

-j


.2 410
S430 USo100


Z 304 446
oj
w

WUb- J 316
U-
U-

HostelloyC
10 20
WEIGHT LOSS (mg/cm2)


Figure 8. Correlation between the "difference
potential" (ER Ep) and corrosion

weight loss of stainless steel alloys
exposed in sea water for 4.25 years.
The "difference potentials" were
determined in 3.5 wt. % NaC1, air-
saturated solutions at 25C [88, p. 1978].












CHAPTER 3

EXPERIMENTAL PROCEDURE


The standard electrochemical polarization cell, described in

Appendix 2, was modified (by drilling an additional hole in the

Teflon lid) to permit the insertion of a combination glass/Ag-AgCl

pH electrode into the cell when desirable. Sample preparation and

mounting techniques are discussed in Appendix 3. Electrolyte composi-

tions are given in Appendix 4. The vacuum deaeration of electrolytes

is discussed in Appendix 5. A list of equipment is given in Appendix 6.


3.1 Potentiostatic Polarization Experiments


Potentiostatic polarization experiments were conducted in the

standard corrosion cell described in Appendix 2. Applied potentials

were measured with an electrometer or with an electrometer/digital

voltmeter combination. Applied currents were measured with an ammeter

incorporated into the potentiostat or with an electrometer (ammeter mode).

When desired a burette was put in place of the thermometer in

the cell lid so that solutions could be introduced into the cell

during the course of an experiment. The duration of the experiments

ranged from a few minutes to a few days.


3.2 Potentiokinetic Polarization Experiments


Potentiokinetic polarization experiments were conducted in




Trade name of E. I. DuPont de Nemours Co., Inc.




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