CHROMATE-FREE CORROSION INHIBITION OF ALUMINUM
ALLOYS: VANADATES AND ANIONIC EXCHANGE CLAY PIGMENTS
A Dissertation
Presented in Partial Fulfillment of the Requirements for
the Degree Doctor of Philosophy in the
Graduate School of The Ohio State University
By
Kevin Douglas Ralston, M.S.
*****
The Ohio State University
2008
Dissertation Examination Committee: Approved by
Professor Rudy Buchheit, Advisor
Professor Gerald Frankel ______
Professor Michael Mills Advisor Graduate Program in Materials Professor R. Allen Miller Science and Engineering
ABSTRACT
In this study, aqueous vanadates and vanadate pigments were studied for possible
use as chromate replacements to inhibit corrosion of Al 2024-T3. Vanadate inhibition on
Al 2024-T3 was characterized as a function of pH and concentration using anodic and
cathodic polarization experiments and scanning electron microscopy (SEM). The results
showed a strong correlation between inhibition and the availability of tetrahedrally
coordinated vanadate species in test solutions. In particular, solutions containing predominately tetrahedrally coordinated vanadates were observed to act as modest anodic inhibitors and to reduce cathodic kinetics through the suppression of oxygen reduction kinetics. Further, the effect of these tetrahedral vanadates on individual intermetallic particles commonly found in Al 2024-T3 was characterized using a microcapillary electrode. Tetrahedral vanadates were generally found to increase breakdown potentials and decrease cathodic kinetics on all tested materials. Open circuit potential (OCP) was observed to shift in the active direction as a result of decreased cathodic kinetics, just below the observed breakdown potential of Al2CuMg; a phase that plays a critical role in
corrosion susceptibility of Al 2024-T3. OCP measurements, SEM images, and
potentiostatic hold experiments were used to show suppressed Al2CuMg dissolution and
damage accumulation in vanadate solutions.
ii Furthermore, synthetic hydrotalcite anion exchange clay pigments were
synthesized with vanadates and other possible inhibitor anions. Hydrotalcites allow the
use of inhibitors that are too soluble for direct use in organic coatings without leading to
coating blistering. A number of hydrotalcite pigments were synthesized and compared to
a SrCrO4 standard using electrochemical impedance spectroscopy and salt spray exposure
of scribed organically coated panels. Typically, vanadate pigmented PVB coatings were
7 observed to have total impedance within an order of magnitude of SrCrO4 (2 x 10 ohms.cm2). Further, all vanadate coatings were observed to provide some scribe
protection during 750 hours of salt spray exposure, however, these coatings also had a
tendency to blister. Release from vanadate hydrotalcites was characterized using neutron
activation analysis. Interestingly, vanadate hydrotalcite pigments that released relatively
small total concentrations of vanadium resulted in the best performance. Low
concentrations of vanadium may promote the formation of tetrahedrally coordinated
species which were shown to act as inhibitors earlier in this study.
iii
Dedicated to Robin
iv
ACKNOWLEDGMENTS
I would like to thank my advisor, Dr. Rudy Buchheit, whose encouragement, approach to advising, and patience have allowed me the time to mature and ultimately complete a dissertation, which I would not have believed possible 5 years ago. I also would like to thank Dr. Jerry Frankel who first introduced me to the world of corrosion while I completed my senior project under his supervision. Further, I would like to thank
Dr. Mike Mills who provided me with my first experience with materials research when I worked in his lab as an undergraduate.
I also wish to acknowledge those who have sponsored portions of my work, namely, the Air Force Office of Scientific Research, Concurrent Technologies, and Luna
Innovations. Further, I acknowledge the Journal of the Electrochemistry Society for permission to reproduce significant portions (Chapter 3) of J. Electrochem. Soc., 155,
C350 (2008). Copyright 2008, The Electrochemical Society.
I would like to thank Dr. T.L. Young for assistance with NMR measurements and
Mr. Joe Talnagi for help and guidance in conducting neutron activation analysis.
Finally, I wish to thank and acknowledge a number of people who have contributed advice, expertise, and assistance to this work, namely Dr. Nick Birbilis, Dr.
Nikki Padgett, Dr. Santi Chrisanti, Dr. Marianno Iannuzzi, Dr. Belinda Hurley, Dr.
Sudhaker Mahajanam, Dr. Girdhari Kumar, Hong Guan, Shuyan Qui, and Casey Grimez.
v
VITA
2003…………………………………B.S. (Materials Science and Engineering),
The Ohio State University, Columbus, OH
2006…………………………………M.S.
The Ohio State University, Columbus, OH
2003 - Present……………………….Ph.D.
The Ohio State University, Columbus, OH
PUBLICATIONS
1. K. D. Ralston, S. Chrisanti, T. L. Young, and R. G. Buchheit (2008) Corrosion Inhibition of Aluminum Alloy 2024-T3 by Aqueous Vanadium Species, Journal of the Electrochemical Society 155 (7): C350-C359.
2. P. J. Wurm, P. Kumar, K. D. Ralston, M. J. Mills, and K. H. Sandhage (2002) Fabrication of Light Weight Oxide/Intermetallic Composites at 1000°C by the Displacive Compensation of Porosity (DCP) Method, Ceramic Transactions, 93-101.
3. P. J. Wurm, P. Kumar, K.D. Ralston, M. J. Mills, and K. H. Sandhage (2001) Fabrication of Dense, Lightweight, Oxide-Rich Oxide/Aluminide Composites at 1000oC by the Displacive Compensation of Porosity (DCP) Process, Powder Materials: Current Research and Industrial Practices (2001 TMS Fall Meeting), November 4-8, 129-139.
vi FIELDS OF STUDY
Major Fields: Materials Science and Engineering
vii
TABLE OF CONTENTS
Page
ABSTRACT ii
DEDICATION iv
ACKNOWLEDGMENTS v
VITA vi
LIST OF TABLES xiv
LIST OF FIGURES xv
1. INTRODUCTION 1
2. LITERATURE REVIEW 6
2.1 Introduction 6
2.2 Aluminum 2024-T3 Metallurgy, Microstructure, and Corrosion
Susceptibility 8
2.2.1 Aluminum Alloy 2024-T3 Metallurgy and Microstructure 8
2.2.2 Aluminum 2024-T3 Susceptibility to Localized Corrosion 11
2.2.3 Importance of Al2CuMg (S Phase) 13
2.3 Chromate Background and Toxicity 14
2.3.1 Aqueous Chemistry of Chromate 14
2.3.2 Chromate as Anodic and Cathodic Inhibitors 15
2.3.3 Functionalizing Chromates as Films and Pigments 17
viii 2.4 Vanadate Inhibitors as Possible Alternatives to Chromate 20
2.4.1 Vanadate Background and Toxicity 20
2.4.2 Vanadates as Inhibitors of Corrosion 21
2.4.3 Aqueous Speciation of Vanadate 22
2.4.4 Current Understanding of the Mechanism of Vanadate
Inhibition 25
2.4.5 Vanadates Functionalized in Films and Coatings 27
2.5 Hydrotalcite Pigments 29
2.5.1 Structure of Hydrotalcite and Hydrotalcite-like Materials 29
2.5.2 Synthesis of Hydrotalcite-like Materials 32
2.5.3 Anion Exchange and Selectivity 33
2.5.4 Hydrotalcites in Coatings 34
2.6 Critical Issues 35
3. CORROSION INHIBITION OF ALUMINUM ALLOY 2024-T3 BY
AQUEOUS VANADIUM SPECIES 53
3.1 Introduction 53
3.2 Experimental Procedures 55
3.2.1 Materials and Chemicals 55
3.2.2 Sample Preparation 56
3.2.3 Nuclear Magnetic Resonance (NMR) 56
3.2.4 Potentiodynamic Polarization 58
3.2.5 Exposure Experiments 59
3.3 Results 60
ix 3.3.1 Changes in Vanadate Speciation with pH Adjustment, Time,
and Exposure to Aluminum 60
3.3.2 Soluble Inhibitor Release from Pigments 63
3.3.3 Aluminum 2024-T3 Aerated Polarization in NaCl Solutions 64
3.3.4 Aluminum 2024-T3 Deaerated Polarization in NaCl Solution 66
3.3.5 Corrosion Morphology of Aluminum in Vanadate Solution 67
3.4 Discussion 68
3.4.1 Speciation versus Corrosion Inhibition 68
3.4.2 Inhibition and Oxygen Dependence 70
3.4.3 Action of Vanadates on Al Alloy Surfaces 71
3.4.4 Vanadate Speciation and Vanadates in Hydrotalcite Pigments 72
3.5 Conclusions 73
4. ELECTROCHEMICAL EVALUATION OF CONSTITUENT
INTERMETALLICS IN ALUMINUM ALLOY 2024-T3 EXPOSED TO
AQUEOUS VANADATE INHIBITORS 98
4.1 Introduction 98
4.2 Experimental Procedures 103
4.2.1 Solution Preparation 103
4.2.2 Nuclear Magnetic Resonance (NMR) 104
4.2.3 Potentiodynamic Polarization Using the Microcapillary
Electrode 105
4.2.4 Electrochemical Experiments on Bulk Al 2024-T3 Sheet 106
4.3 Results 108
x 4.3.1 Inhibition from Tetrahedral Vanadate Species vs. Octahedral
Species 108
4.3.2 Tetrahedral Vanadate Species in Alkaline Electrolytes 109
4.3.3 Polarization of Intermetallics in Tetrahedral Vanadate
Solutions 109
4.3.4 OCP and SEM Images of Al 2024 Exposed to Tetrahedral
Vanadate Solution 113
4.3.5 Suppressed Al2CuMg Dissolution in Tetrahedral Vanadate
Solutions 116
4.4 Discussion 117
4.4.1 Cathodic Inhibition from Tetrahedrally Coordinated Vanadate
Species 117
4.4.2 Suppression of Al2CuMg Breakdown 118
4.4.3 Variation in Anodic Behavior of Al2CuMg 119
4.4.4 Vanadate Buffering and Circumferential Attack 120
4.5 Conclusions 122
5. HYDROTALCITE PIGMENTS FOR CORROSION INHIBITION OF
ALUMINUM ALLOY 2024-T3 168
5.1 Introduction 168
5.2 Experimental Procedures 173
5.2.1 Materials and Chemicals 173
5.2.2 Hydrotalcite Pigment Synthesis 174
5.2.3 XRD Structure Confirmation 176
xi 5.2.4 Organic Coating Preparation and Application 176
5.2.5 Coating Evaluation by EIS on Panels Exposed to Static
NaCl Solutions 177
5.2.6 Coating Evaluation by Salt Spray Exposure of Scribed
Panels 177
5.2.7 Inhibitor Release Characterization Using NAA 178
5.3 Results 180
5.3.1 Structural Confirmation of Synthesized Pigments 180
5.3.2 EIS of PVB Coatings Pigmented with Vanadate
Hydrotalcites 182
5.3.3 Salt Spray Exposure of Vanadate Hydrotalcite Pigmented
Coatings 187
5.3.4 Inhibitor Release from Vanadate-Bearing Hydrotalcite
Pigments 188
5.3.5 EIS of PVB Coatings Pigmented with Various Non
Vanadate Hydrotalcites 190
5.3.6 Salt Spray Exposure of Non-Vanadate Hydrotalcite
Pigmented Coatings and Controls 192
5.4 Discussion 192
5.4.1 Influence of Vanadate Speciation and Concentration on
Vanadate Hydrotalcite Pigment Performance 192
5.4.2 Possible Influence of Cation Inhibitors Released into
Solution 194
xii 5.4.3 Effect of Solubility on Vanadate Pigment Performance 195
5.4.4 Non-Vanadate Hydrotalcite Pigments Generally Do Not
Provide Inhibition 196
5.5 Conclusions 196
6. CONCLUSIONS AND FUTURE WORK 224
6.1 Conclusions 224
6.2 Future Work 226
BIBLIOGRAPHY 228
xiii
LIST OF TABLES
Page
3.1. List of probable vanadium species in solutions of varied NaVO3 concentration and pH. The dashed line delineates vanadate solutions that demonstrated inhibition, on the right, from solutions in which inhibition was not observed, on the left. Species listed in red (grey in black and white copies) are likely present in relatively small concentrations compared to species listed in black 95
4.1. Averaged electrochemical data for intermetallics tested in approximately pH 9.17 0.5 M NaCl solutions with and without 10 mM NaVO3. A “NaCl” heading indicates data from NaCl-only solutions while a “NaCl + V” heading indicates data collected in NaCl solution that contained 10 mM NaVO3. 164
5.1. Synthesis details and salt spray exposure results for PVB coated panels pigmented with hydrotalcites and various controls. For rankings of blistering and scribe protection a “1” indicates the best performance. 220
xiv
LIST OF FIGURES
Page
2.1. Corroded aluminum surface showing two different kinds of intermetallic particles commonly found in Al 2024, small round Al2CuMg and Al2Cu intermetallics, and large irregular shaped Al-Cu-Fe-Mn intermetallics. 38
2.2. Aqueous speciation of Cr6+ species as a function of concentration and pH. 39
2.3. Sequence of events of Cr3+ film formation. 40
2.4. Equilibrium predominance diagram for V5+ aqueous species. 41
2.5. Illustrations of different vanadate structure as a function of pH and concentrations ranging from 10-3 to 10-1 mol L-1. 42
2.6. Schematic representation of selective inhibitor release and exchange for contacting aggressive anions. 43
2.7. Ideal structure for a hydrotalcite containing carbonate in the interlayers. 44
2.8. X-ray diffraction pattern for Al-Zn hydrotalcites containing decavanadate and chloride. 45
2.9. Titration curves for different cation combinations showing the ranges of pH where co-precipitation and precipitation of cation-hydroxides occur. 46
3.1. Equilibrium predominance diagram for VV-OH- species as a function of concentration and pH. Adapted from The Hydrolysis of Cations and Heteropoly and Isopoly Oxometalates. The dotted line indicates delineation between octahedral coordinated and tetrahedral coordinated vanadate species. Approximate pH and concentration of test solutions are indicated with an “x” 75
3.2. Change in pH as a function of time of a pH 8.76 100 mM NaVO3 solution initially acidified to pH 4.12 with HNO3 and then adjusted to pH 7.58 with NaOH. 76
xv 3.3. NMR spectrum as a function of time of a 100 mM NaVO3 solution initially acidified to pH 4.12 with HNO3 and then adjusted to pH 7.58 with NaOH 77
3.4. NaVO3 solution after serial additions of 10 N NaOH: A) initial orange 100 mM NaVO3 solution at pH 6.08, B) yellow pH 6.63 solution 27 hours after the addition of a second drop of 10 N NaOH immediately prior to a third NaOH addition (spectrum collected at 52 hrs total), C) yellow pH 9.15 solution, NMR sample taken and pH measured approximately 30 minutes after the addition of a third drop of 10 N NaOH (spectrum collected at 52 hrs total). 78
3.5. NMR spectra of 100 mM NaVO3 solutions in contact with a pure Al wire A) initial orange decavanadate-metavanadate pH 6.08 solution prior to exposure, B) dark emerald green solution with pH near 6 after 52 hours of contact, C) initial pale yellow decavanadate-free pH 8.10 solution prior to contact. D) Clear decavanadate-free solution after 52 hours of contact. 79
3.6. NMR spectrum of filtrate from 4.0 grams of an Al-Zn-V hydrotalcite pigment soaked in 40 mL of 0.1 M NaCl for 20 hours, solution pH 6. 80
3.7. Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant pH 10 adjusted solution and varied NaVO3 concentration. 81
3.8. Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant 0.0032 M NaVO3 solution and varied pH. 82
3.9. Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant pH 8 solution and varied NaVO3 concentration. 83
3.10. Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant 0.0032 M NaVO3 concentration and varied pH. 84
3.11. Condensed data from aerated polarization experiments, corrosion potential as a function of pH. 85
3.12. Condensed data from aerated polarization experiments, corrosion current as a function of pH. 86
3.13. Condensed data from aerated polarization experiments, pitting potential as a function of pH. 87
3.14. Condensed data from aerated polarization experiments, cathodic current at -1.25VSCE as a function of pH. 88
xvi
3.15. Condensed data from deaerated polarization experiments, A) corrosion potential as a function of pH. 89
3.16. Condensed data from deaerated polarization experiments, corrosion current as a function of pH. 90
3.17. Condensed data from deaerated polarization experiments pitting potential as a function of pH. 91
3.18. Condensed data from deaerated polarization experiments, cathodic current at -1.25VSCE as a function of pH. 92
3.19. Sample chemical maps collected in aerated 50 mM NaCl showing observed suppression of S-phase dissolution in alkaline NaVO3 solutions compared to vanadate-free solutions. A) pH 10 0.0 M NaVO3 B) pH 10 0.0032 M NaVO3. 93
3.20. Sample chemical maps collected in aerated 50 mM NaCl showing different manifestations of vanadium on the surface of Al 2024-T3. A) pH 5 0.32 M NaVO3, B) pH 8 0.0032 M NaVO3 C) pH 8 0.32 M NaVO3. 94
4.1. NMR spectra of pH 5.1 0.5 M NaCl solutions with a) 0.0025 M NaVO3 and b) 0.25 M NaVO3. The dilute NaVO3 solution has a greater proportion of tetrahedrally coordinated species (V1 and V4) relative to octahedrally coordinated species (V10) compared to the more concentrated NaVO3 solution which contains mostly octahedrally coordinated species. 124
4.2. Cathodic polarization curves on Al 2024-T3 in aerated pH 5.1 0.5 M NaCl with 0.25 M NaVO3, 0.0025 M NaVO3, and without NaVO3. These experiments show an inverse relationship between NaVO3 concentration and inhibition of cathodic kinetics at pH 5.1, which correlates well with a transition from solutions dominated by octahedrally coordinated vanadates to tetrahedral vanadates. 125
4.3. NMR spectra showing the presence of tetrahedrally coordinated vanadates (V1, V2, V4, and V5) in pH 9.17 0.5 M NaCl + 10 mM NaVO3 solution used for microcapillary electrochemical experiments a) immediately prior to experimentation and b) 14 days later after completion of experiments. 126
4.4. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Al. 127
xvii 4.5. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Cu. 128
4.6. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. 129
4.7. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. 130
4.8. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg. 131
4.9. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. 132
4.10. Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. 133
4.11. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al. The “V” in the legend indicates curves collected in solutions containing NaVO3. 134
4.12. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 pure Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 135
4.13. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 136
4.14. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 137
xviii 4.15. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg. The “V” in the legend indicates curves collected in solutions containing NaVO3. 138
4.16. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. The “V” in the legend indicates curves collected in solutions containing NaVO3. 139
4.17. Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. The “V” in the legend indicates curves collected in solutions containing NaVO3. 140
4.18. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al. 141
4.19. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Cu. 142
4.20. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. 143
4.21. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. 144
4.22. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg. 145
4.23. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. 146
4.24. Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. 147
xix 4.25. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al. The “V” in the legend indicates curves collected in solutions containing NaVO3. 148
4.26. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 149
4.27. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 150
4.28. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3. 151
4.29. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for, Al2CuMg. The “V” in the legend indicates curves collected in solutions containing NaVO3. 152
4.30. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. The “V” in the legend indicates curves collected in solutions containing NaVO3. 153
4.31. Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. The “V” in the legend indicates curves collected in solutions containing NaVO3. 154
4.32. OCP of Al 2024-T3 over 4 h in aerated and deaerated pH 9.17 0.5 M NaCl with and without 10 mM NaVO3. Solutions were not intentionally buffered and duplicate curves are shown. 155
4.33. SEM images of Al 2024-T3 after 4 h of OCP measurement in aerated 0.5 M NaCl solutions with an initial approximate pH 9.17. 156
4.34. SEM images of Al 2024-T3 after 4 h of OCP measurement in aerated 0.5 M NaCl solutions with an initial approximate pH 9.17 with 10 mM NaVO3. 157
xx
4.35. SEM images of Al 2024-T3 after 4 h of OCP measurement in deaerated 0.5 M NaCl solutions with an initial approximate pH 9.17. 158
4.36. SEM images of Al 2024-T3 after 4 h of OCP measurement in deaerated 0.5 M NaCl solutions with an initial approximate pH 9.17 and with 10 mM NaVO3. 159
4.37. Duplicate anodic polarization curves of bulk Al 2024-T3 sheet in aerated 0.5 M NaCl + 10 mM NaVO3 solution at pH 9.17. 160
4.38. Current response of potentiostatic hold experiments on Al 2024-T3 in pH 9.17 0.5 M NaCl with 10 mM NaVO3. 161
4.39. Current response of potentiostatic hold experiments on Al 2024-T3 in pH 9.17 0.5 M NaCl without NaVO3. 162
4.40. Total charged passed as a function of potential hold for Figures 4.11.1 and 4.11.2. 163
5.1. XRD patterns of hydrotalcites synthesized with interlayer vanadates. 198
5.2. XRD patterns of hydrotalcites synthesized to contain anions other than vanadate. 199
5.3. An explicit equivalent circuit model of the coating system (a) and the simplified model used to fit data from PVB coated Al 2024-T3 (b). 200
5.4. Bode magnitude and phase angle plots of HTV3 pigmented PVB coatings subject to static exposure in 0.5 M NaCl. The area exposed was approximately 8.55 cm2. 201
5.5. Bode magnitude and phase angle plots of SrCrO4 pigmented PVB coatings subject to static exposure in 0.5 M NaCl. The area exposed was approximately 8.55 cm2. 202
5.6. Bode magnitude and phase angle plots of HTV3 pigmented PVB coating after 4 days exposure in 0.5 M NaCl and the corresponding data fit produced using the simplified equivalent circuit. The area exposed was approximately 8.55 cm2. 203
5.7. Nyquist plot of HTV3 pigmented PVB coating after 4 days exposure in 0.5 M NaCl and the corresponding data fit produced using the simplified equivalent circuit. The area exposed was approximately 8.55 cm2. 204
xxi 5.8. Total impedance of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl. 205
5.9. Pore resistances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl. 206
5.10. Defect capacitances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl. 207
5.11. Oxide capacitances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl. 208
5.12. Capacitances associated with the PVB layer of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl. 209
5.13. Al 2024 panels coated with vanadate hydrotalcite pigmented PVB. Panels were scribed and exposed to salt spray for 750 h. Relative blister (B) and scribe protection (SP) performance are indicated at the bottom of each panel with “1” indicating strong performance. 210
5.14. Close up image of SrCrO4 and HTV3 panels after 750 hours of salt spray exposure. 211
5.15. Vanadate release from vanadate hydrotalcite pigments exposed to 0.5 M NaCl compared to SrCrO4 release using NAA. 212
5.16. Release of Zn and Ni from vanadate hydrotalcite pigments exposed to 0.5 M NaCl collected using NAA with long counting times or long sample decay times. 213
5.17. Total impedance at low frequency of PVB coatings pigmented with non- vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl. 214
5.18. Pore resistances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl. 215
5.19. Defect capacitances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl. 216
5.20 Oxide capacitances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl. 217
xxii 5.21. Capacitances associated with the PVB layer in a PVB coating pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl. 218
5.22. PVB coated Al 2024 pigmented with non-vanadate hydrotalcites. Panels were scribed and exposed to salt spray for 750 h. Relative panel blister (B) and scribe protection (SP) performance is indicated at the bottom of each panel with “1” indicating strong performance. 219
xxiii
CHAPTER 1
INTRODUCTION
Aluminum 2024-T3 is a high strength aluminum alloy commonly used in the
aerospace industry (1). The alloy is generally designed, through alloy additions and heat-treatments, to maximized mechanical properties without consideration for possible implications for corrosion resistance; although the alloy can be supplied in a T8 temper to
increase resistance to stress corrosion cracking (1, 2). The microstructure of Al 2024 is
heterogeneous, containing numerous second phase particles with different
electrochemistry than the matrix (1-4). This heterogeneity makes the alloy inherently
susceptible to localized corrosion (4). Historically, chromate-based coatings and pigments have been used successfully to increase corrosion resistance of many aluminum alloys (5, 6). However, chromate is a known carcinogen, and as such, there is a strong desire to find a more “green” and less hazardous replacement (7). Recently, vanadates have garnered attention as non-carcinogenic chromate replacements and have shown promise as corrosion inhibitors on aluminum in pigments and conversion coatings (8-10).
Although generally accepted as inhibitors, vanadates have a complex aqueous chemistry,
1 dependent on both pH and concentration, which prevents an easy explanation of which vanadate species are involved and the mechanism of inhibition (11). There is a body of evidence that suggests that vanadates are inhibitors of corrosion, and it seems likely that inhibition depends on speciation. However, this relationship has not been fully characterized. Additionally, the presence of intermetallic particles, and in particular Cu- bearing intermetallics, plays a large role in the corrosion processes of Al 2024. An understanding of how inhibiting vanadates interact with these particles and the matrix may help in elucidating the mechanism of inhibition.
Unfortunately, the use of vanadates as a drop in replacement for chromates is not possible. In addition to strong corrosion inhibition, a convenience of many chromate- based pigments is that they have appropriate solubility and inhibitor efficiency to provide inhibition without leading to blistering of the coating (12). Vanadates typically do not possess the same inhibitor efficiency as chromates, and generally vanadate salts are too soluble for direct use as pigments in an organic coating. However, studies using anion- exchange clays, known generically as hydrotalcites, have provided a pathway by which highly soluble inhibitor salts may be incorporated directly into organic coatings (13).
Previous work has shown vanadate-bearing hydrotalcite pigments to be inhibitors on organically coated Al 2024 (8, 14). However, these pigments have yet to be optimized, and the possibility of using hydrotalcites to deliver other inhibitors that would otherwise be too soluble has not been explored. The objectives of this dissertation are as follows:
• To systematically characterize vanadate speciation and inhibition over a
range of pH and vanadate concentration using electrochemical techniques
2 to further the understanding of how various vanadate species effect
inhibition.
• To determine how inhibiting vanadates interact with the matrix and
constituent particles of Al 2024 and to develop an understanding of why
vanadates suppress Al2CuMg dissolution and the implications for damage
accumulation on the alloy.
• To develop a framework by which hydrotalcite pigments can be tailored to
contain a number of different inhibitor anions. Further, to synthesize,
characterize, and evaluate performance of a number of different
hydrotalcite pigments relative to a SrCrO4 standard.
This dissertation consists of six chapters. Chapter 2 is literature review that
provides pertinent details about Al 2024, chromates, vanadates, and hydrotalcite
pigments. In Chapter 3 the effects of vanadate speciation on corrosion inhibition of Al
2024 are characterized. This work shows that inhibition is likely correlated to the
presence of tetrahedrally coordinated vanadates. Chapter 4 explores how inhibiting
tetrahedral vanadates interact with specific intermetallic particles commonly found in Al
2024. Generally, tetrahedrally coordinated vanadates were found to decrease cathodic
kinetics on intermetallic particles and the matrix; this has strong implications for the
dissolution of Al2CuMg intermetallics. Chapter 5 focuses on developing an
understanding of issues critical to the synthesis of a number of different hydrotalcite pigments. Using this understanding, a variety of different hydrotalcites were synthesized
3 and evaluated relative to SrCrO4 standards, with a particular focus on optimization of
vanadate hydrotalcites. Chapter 6 is a summary of key conclusions of this dissertation and suggestions for future work.
4 REFERENCES
1. I. J. Polmear, Light Alloys Metallurgy of the Light Metals, Arnold, London (1995).
2. T. H. Muster, A. E. Hughes and G. E. Thompson, in Corrosion Research Trends, I. S. Wang Editor, p. 35, Nova Science Publishers, Inc. (2007).
3. R. G. Buchheit, Journal of the Electrochemical Society, 142, 3994 (1995).
4. R. G. Buchheit, R. P. Grant, P. F. Hlava, B. Mckenzie and G. L. Zender, Journal of the Electrochemical Society, 144, 2620 (1997).
5. R. G. Buchheit and A. E. Hughes, in ASM Handbook, S. D. Cramer and B. S. C. Jr Editors, p. 720 (2003).
6. M. W. Kendig and R. G. Buchheit, Corrosion, 59, 379 (2003).
7. P. O'Brien and A. Kortenkamp, Transition Metal Chemistry, 20, 636 (1995).
8. S. P. V. Mahajanam and R. G. Buchheit, Corrosion, 64, 230 (2008).
9. H. Guan and R. G. Buchheit, Corrosion, 60, 284 (2004).
10. M. Iannuzzi, T. Young and G. S. Frankel, Journal of the Electrochemical Society, 153, B533 (2006).
11. J. Charles F. Baes and R. E. Mesmer, The Hydrolysis of Cations, Robert E. Krieger Publishing Company, Inc., Malabar, Fl (1986).
12. J. Sinko, Progress in Organic Coatings, 42, 267 (2001).
13. F. Deflorian and I. Felhosi, Corrosion, 59, 112 (2003).
14. R. G. Buchheit, H. Guan, S. Mahajanam and F. Wong, Progress in Organic Coatings, 47, 174 (2003).
5
CHAPTER 2
LITERATURE REVIEW
2.1 Introduction
Aluminum alloy 2024-T3 is a high strength precipitation-age-hardened alloy widely used in aerospace due to its low density and superior mechanical properties.
However, the same alloying elements that provide Al 2024-T3 with advantageous mechanical properties also make the alloy inherently susceptible to localized corrosive attack. Historically, the use of chromate-based coatings and pigments have been used to protect aluminum used in the aircraft industry (1, 2). However, chromate is a known carcinogen, and there is a strong desire to find more environmentally-friendly and healthy alternatives (1-3). To fully appreciate the challenge associated with finding a chromate replacement and work towards a solution, a fundamental understanding of how Al 2024-
T3 corrodes and of the mechanisms by which chromate is able to suppress corrosion is required. Chromate is a unique inhibitor that is easily applied as a conversion coating or dispersed as a pigment in organic coatings, and is effective by dramatically slowing cathodic reaction kinetics in a wide range of service environments (4, 5). Chromate-based
6 coatings have also been shown to “self-heal,” meaning damaged areas of a coating can be repaired by aqueous chromate species leached into solutions from the coating itself (4, 5).
A number of possible chromate alternatives have previously been evaluated, but none possesses the same combination of properties that chromates provide. As a result, a piecewise solution that incorporates the properties of multiple inhibitors and inhibitor systems will likely be necessary to achieve similar levels of inhibition attained with chromate.
Vanadium-based coatings and pigments have shown promise as corrosion inhibitors of aluminum alloys, providing protection without the same degree of adverse health effects associated with chromate (6-9). Although vanadate is generally accepted as an inhibitor, the complexities of aqueous vanadium chemistry prevent an easy explanation of the exact mechanism of inhibition, necessitating a more in-depth exploration. Anionic exchange clay pigments, broadly known as hydrotalcites, provide a novel way of storing and releasing vanadates and other potential inhibitor anions that are not easily incorporated into organic coatings (8, 9). Hydrotalcites can be tailored to store a number of different inhibitor anions, which could be used either individually or in concert to provide inhibition over a similar range of environmental conditions as chromates.
The objective of this chapter is to provide a literature-based background of the challenges of chromate replacement and to justify research on the use of vanadates, in conjunction with hydrotalcites, to replace the use of chromate-based inhibitors on copper containing aluminum alloys, particularly Al 2024-T3. This will be accomplished through a description of 1) the metallurgy and microstructure of Al 2024-T3 and why the alloy is
7 susceptible to localized corrosion, 2) the mechanisms by which chromates prevent
corrosion, 3) vanadates and vanadate inhibition as is currently understood, and 4)
hydrotalcites and the unique properties that might be exploited to develop a chromate- free corrosion inhibitor for Al 2024.
2.2 Aluminum 2024-T3 Metallurgy, Microstructure, and Corrosion Susceptibility
Aluminum 2024-T3 is a high strength age-hardened aluminum alloy commonly
used in the aerospace industry. A combination of alloy additions, heat-treatments, and aging allows substantial gains in mechanical properties to be achieved, with reported tensile strengths near 500 MPa, as well as good damage tolerance and resistance to fatigue (10, 11). Additionally, microstructural variability from the presence of intermetallics, namely Cu containing-intermetallics, leads to susceptibility to localized corrosion necessitating the use of additional corrosion protection to prevent property degradation.
2.2.1 Aluminum Alloy 2024-T3 Metallurgy and Microstructure
Al 2024-T3 is a solution-heat-treated and cold-worked alloy that contains by weight percent 3.8-4.9 Cu, 1.2-1.8 Mg, 0.3-0.9 Mn, and small quantities of Si, Fe, Zn, Cr, and Ti (10). Alloy additions result in both superior mechanical properties and a heterogeneous microstructure, which renders the alloy susceptible to localized corrosion
(10). Appreciable quantities of copper, magnesium, and manganese are added as strengtheners and are reported to have maximum solubility in weight percent at elevated temperatures of 5.65, 17.4, and 1.82, respectively (10). Manganese is an efficient
8 strengthener. However, due to relatively low solubility only 0.2-0.3% remains in solid solution while the rest precipitates as second phase particles (10). Other strengtheners, such as Cu, are more soluble than manganese, but also form micro-constituents that decrease the corrosion resistance of the alloy (10). Magnesium additionally acts as a strengthener, but not as efficiently on an atomic basis as either copper or manganese (10).
Other alloy additions such as Cr and Ti play a role in grain refinement and recrystallization (11, 12).
Al 2024-T3 is an alloy that is heat treated to maximize mechanical properties in the final product. The heat treating schedule is generically described by a solution heat treatment in a single phase region, rapid quenching, cold-work, and solid solution decomposition through aging (10, 11). During the initial solution heat treatment the alloying additions form a equilibrium solid solution with Al (10, 11). This solid solution is then quenched, which results in the formation of a supersaturated solid solution and the precipitation of insoluble intermetallics (11). Appreciable quantities of copper, magnesium, and manganese remain in solid solution, however, and a fine dispersion of
Cu and Mg clusters begins to form and subsequently grow into coherent Guinier-Preston
(GP) zones (11). GP zones are small clusters of atoms that contain high concentrations of solute atoms such as Cu and Mg (10). Cold working increases dislocation density, facilitating the formation of semi-coherent and ultimately incoherent phases to occur heterogeneously at lattice incongruencies such as GP zones or dislocations (10).
Hardening particles, such as Al2Cu and Al2CuMg, are largely responsible for increased strengths and can range in sizes up to nearly a micrometer (11, 12). Optimal hardness is achieved by obtaining a critical dispersion of hardening particles and GP zones through
9 an appropriate heat-treatment and aging schedule (10). This maximizes the amount of strain around precipitates and GP zones in the alloy and in turn maximizes resistance to dislocation motion (10).
Two other types of particles that are found in the alloy are dispersoids and constituent particles. Dispersoids form as the result of reactions with Cr, Mn, and Ti and play a role in grain refinement, however, dispersoids are not believed to contribute to localized corrosion due to their relative inertness (0.05 to 0.5 μm) (11, 12). Constituent particles, formed as the initial ingot solidifies, are large with sizes up to 10 μm and may include the following: Al7Cu2Fe, Al12(Fe, Mn)3Si, Al2CuMg, Al2Cu, and Al6(Cu, Fe) (11,
12). The total number of particles in Al 2024 is large and estimated to have a density between 300,000 and 530,000/cm2 (11).
Among intermetallic particles, Al2CuMg is of particular interest because
Al2CuMg is reported to be one of the most abundant intermetallics found in Al 2024 (13,
14). Specifically, Al2CuMg is believed to account for approximately 60% of particles
larger than 0.5 μm, while most of the remaining particles are some combination of Al-
Cu-Fe-Mn intermetallics (11, 14). In general, Al2CuMg and Al2Cu intermetallics appear
as small round particles, in contrast to Al-Cu-Fe-Mn intermetallics which often have
blocky irregular shapes and sizes typically larger than 5 μm (13-15). Figure 2.1 illustrates the typical difference in shape and size between round (Al2CuMg and Al2Cu) and blocky particles (Al-Cu-Fe-Mn) intermetallic particles (15).
10 From a mechanical properties perspective, the previously discussed alloy
additions result in an alloy with superior properties. However, property gains come at a
cost because the heterogeneous microstructure makes the alloy inherently susceptible to
localized corrosion (10).
2.2.2 Aluminum 2024-T3 Susceptibility to Localized Corrosion
The effects of intermetallic particles on corrosion of aluminum alloys has been studied widely (11-21). In general, localized attack is rooted in compositional differences between the matrix and intermetallic particles, leading to differences in electrochemical behavior. A convenient and simplified view of Al 2024 is as series of galvanic relationships, where the intentionally heterogeneous microstructure of the alloy in contact with an electrolyte leads to the formation of numerous galvanic couples. A galvanic couple forms when two different metals or alloys are in electrical contact with each other and ionic contact through an aqueous electrolyte. This results in the distinct separation of anodic and cathodic reactions where the surfaces of each metal or alloy are polarized, causing one to become an anode and preferentially dissolve while the other becomes a cathode and supports a reduction reaction. Studies of aluminum alloys have been conducted to classify whether particular intermetallic particles are active or noble relative to the surrounding matrix (12, 17, 21, 22). Once the relative nobility or activity is
known, judgments about its expected behavior to either undergo anodic dissolution or
support a reduction reaction, which drives corrosion in the surrounding matrix, can be made. Most intermetallics are expected to be noble to the matrix and support reduction
reactions, however those containing primarily Zn, Mg, and Si are expected to be active to
11 the matrix and preferentially corroded (12). For Al 2024-T3 it has been found that
intermetallics containing Cu are typically noble or will become noble to the surrounding aluminum matrix during exposure to many electrolytes and these particles are capable of supporting rapid cathodic kinetics (11, 13, 14, 16, 19, 23). Other noble intermetallics, containing Fe or Mn for example, are capable of supporting oxygen reduction, but at efficiencies much lower than intermetallics containing Cu (11, 12). Noble intermetallic particles support oxygen reduction in acidic (Eqn. 1) and neutral or alkaline aqueous environments (Eqn 2) by the following equations (24, 25):
+ - O2 + 4H + 4e → 2H2O Eqn. 1
- - O2 + 2H2O + 4e → 4OH Eqn. 2
Such cathodic particles drive corrosion in the surrounding matrix leading to
pitting and trenching attack morphologies (19, 26). This attack has been attributed to the
previously mentioned galvanic interaction and the formation of local galvanic cells
between the matrix and intermetallic particles (11, 13-15, 26). Alternately, these
morphologies have also been attributed to development of increased alkalinity near
cathodic intermetallics associated with oxygen reduction (26-29). The increase in
alkalinity from oxygen reduction dissolves nearby aluminum from the matrix creating
grooves that may subsequently grow by an acid-pitting mechanism (27, 28). Either way,
intermetallic particles make the alloy susceptible to localized corrosion.
12 Although it is useful to view and explain corrosion processes of Al 2024 as
galvanic interactions between the matrix and constituent intermetallic particles, this may
be an over simplification. While rankings of corrosion potentials will give an idea of
whether an intermetallic will support a cathodic reaction or be selectively dissolved,
reaction rates will vary widely depending on the composition and size of the intermetallic
(12). Further, it has been argued that the idea of local corrosion cells is “dangerous” as
each pit requires current from many intermetallics supporting oxygen reduction to
support stable pit growth (27).
2.2.3 Importance of Al2CuMg (S Phase)
As previously discussed, Al2CuMg is one of the most abundant intermetallic particles found in Al 2024 and in large part has been found to be responsible for the
susceptibility of Al 2024 to localized corrosion (13, 14). The corrosion of Al2CuMg is
complex; under free corrosion conditions the intermetallic is initially anodically polarized
by the matrix, leading to selective dissolution of Mg from the intermetallic and non- faradaic liberation of Cu, which can then be oxidized to form ions that can be reduced on the surrounding matrix (13, 14). Often what remains of the particle is an enriched Cu remnant, which acts as a local cathode that is capable of supporting rapid oxygen reduction and corrosion in the surrounding matrix (13, 14). Prevention of Mg dissolution from Al2CuMg, and as a result the subsequent formation of local Cu cathodes capable of
supporting rapid oxygen reduction, could be an effective way to increase the resistance of
Al 2024 to localized corrosion (30).
13 Ultimately, how damage is accumulated on the alloy surface and the relationship
between intermetallic particles and the matrix is complex. What is agreed on is a need to
slow or prevent corrosion damage on the alloy in order to prevent degradation of mechanical properties.
2.3 Chromate Background and Toxicity
Chromate has been used in pigments for paints and primers and in conversion
coatings to inhibit corrosion of aluminum alloys since the beginning of the 20th century
(1, 2, 5). However, chromate is a known carcinogen; specifically, the reduction of Cr6+ to
Cr3+ and the resultant reaction debris are believed to cause mutations and damage to
DNA (1, 3-5). As a result, a strong desire exists to find healthier, more environmentally-
friendly, alternatives to chromate-based protection schemes for aluminum alloys (1, 5).
An understanding of the mechanism of chromate action on Cu-containing aluminum
alloys may give valuable insight into the development of chromate-free alternatives.
2.3.1 Aqueous Chemistry of Chromate
The aqueous chemistry of chromium is relatively simple and is an important
consideration in developing an understanding of chromate inhibition. Depending on pH
and concentration, soluble Cr6+ will hydrolyze into 4 different species as illustrated in
- 2- Figure 2.2 (5, 31). HCrO4 (bichromate) and Cr2O7 (dichromate) are predominant in acidic solutions with relatively low and high concentrations of Cr6+, respectively (5, 31).
2- H2CrO4 is present in strongly acidic solutions and CrO4 (chromate) predominates in
neutral and basic solutions (5, 31). Despite the variable speciation, all charged Cr6+
14 species are believed to play a role in corrosion inhibition of aluminum alloys, and as a result, inhibition from chromates is observed over a wide range of pH values (4). It has been noted that the electronic properties of Cr6+ and Cr3+ are well suited for corrosion inhibition of aluminum alloys (4). Cr6+ species are tetrahedrally coordinated and generally soluble in aqueous solutions, where they can readily be transported and adsorb onto oxide surfaces (4). In contrast, Cr3+ species are octahedrally coordinated and form inert oxides and films (4). The presence of reducible high-valence Cr6+ species across a wide range of concentrations is largely responsible for the protective properties of chromates, and ideally a chromate replacement would be able to deliver similar performance across the same range of solution pH.
2.3.2 Chromate as Anodic and Cathodic Inhibitors
Chromates have been observed to act as modest anodic inhibitors on aluminum, primarily by increasing the pitting potential (4, 5, 22, 32). Pits are a form of attack caused by local breakdown of a thin protective film, in the case of aluminum, an oxide.
This breakdown leads to an increased local pH at the initiation site, as well as an increase in anodic current density, causing rapid localized anodic dissolution. The potential at which the film becomes unstable and breakdown occurs can be measured and used as a comparative gauge of corrosion resistance: an increased or more noble pitting potential implies increased resistance to pitting attack (25).
An increased pitting potential from chromates has been attributed to their ability to decrease the metastable pit nucleation rate and pit growth rates (4). Although the exact mechanism of chromate action in pits is not clear, it is believed that the release of small
15 amounts of soluble Cr6+ species from conversion coatings or pigments migrates or
3+ 6+ diffuses to active pit sites where it adsorbs to Al(OH)x forming a Al /Cr mixed oxide
(33). It is unclear how this oxide slows corrosion, but Cr6+ may neutralize charge on the
surface, preventing Cl- adsorption or increase the pH in the pit through consumption of
H+ during Cr6+ reduction (33).
Despite modest anodic inhibition from chromates, the primary mechanism of
inhibition is through suppression of cathodic reduction reactions and the ability of
chromate-based protection schemes to “self heal” defects that may occur in a coating.
The effects of Cr6+ on cathodic reduction kinetics have been studied widely (4, 5, 16, 30,
34, 35). For instance, sodium chromate at 0.01 M concentration decreases the net rate of
cathodic reaction on Al 2024 (30). Further, chromate has been observed to decrease
cathodic kinetics by passivating Al2CuMg, which prevented the formation of a Cu-
enriched cluster capable of efficiently supporting oxygen reduction (30).
In particular, the reduction of Cr6+ to Cr3+ and the resultant formation of a thin
adherent hydroxide layer on the surface of an alloy plays a critical role in slowing
reduction reactions (4, 35). An illustration of the process is shown in Figure 2.3 (36).
The layer has monolayer or near monolayer thickness and subsequently inhibits electron
transfer, in particular the reduction of oxygen and Cr6+ (4, 34, 35). On Al-Cu alloys, Cr6+ is rapidly adsorbed on sites capable of supporting reduction, in particular noble Cu- containing intermetallics, and once the Cr6+ is reduced it irreversibly occupies the site
preventing future O2 reduction at that site (4, 34, 35, 37). In this way the overall ability
of the alloy to support oxygen reduction and accumulate damage is substantially
decreased.
16 2.3.3 Functionalizing Chromates as Films and Pigments
Chromates have been used as conversion coatings and pigments to provide
inhibition to aluminum alloys (4, 5). Chromate conversion coatings (CCC) applied under
proper conditions, are spontaneously formed 0.01 to 3 μm thick films (5). Coatings are typically formed in acidic baths of pH 1.2 to 3.0, with sufficient concentration of Cr6+ to
2- 6+ ensure Cr2O7 is the predominant species in solution; 50 mM Cr is commonly used in conversion baths (4, 5). The formation of CCCs occurs in two stages. The first stage involves a short period of electrochemical activity associated with the adsorption of Cr6+
species in solution on the surface and their reduction, accompanied by surface oxidation,
to insoluble Cr3+ species (36, 38). This process results in the formation of a barrier film
consisting of Al2O3 and Cr(OH)3 which protect against corrosion (39). Although the
initial monolayer thick film substantially reduces electron transfer and oxygen reduction,
Cr(OH)3 octahedral units, sharing corners and edges, polymerize on the surface, further
inhibiting oxygen and Cr6+ reduction (4, 35, 36, 40).
After the formation of a thin Cr3+ hydroxide film, the second stage of CCC
formation occurs when Cr6+ species bind to the surface of the Cr3+ monolayer film,
forming a Cr3+-O-Cr6+ bond (36, 38). Cr6+ is incorporated into the film resulting in additional resistance and coating thickness (38, 40). More importantly, Cr6+ incorporated
into the oxide film participates in a reversible Cr3+-O-Cr6+ bond with the insoluble Cr3+
oxide on the alloy surface resulting in a mixed oxide (40). In this sense, the CCC acts as
a reservoir of Cr6+ that may be released if the covalent bond is broken when in contact
with an aqueous solution (40-42). The release of Cr6+ from CCC is governed by the
following reaction (42):
17
3+ 6+ 2- + 3+ 6+ - Cr - OH (solid) + Cr O4 (aq) + H (aq) ↔ Cr -O-Cr O3 (solid) + H2O Eqn. 3
This equilibrium indicates the release of Cr6+ is pH dependent, with acidic solutions
favoring the formation of oxide and alkaline solutions favoring Cr6+ release (42). This
allows CCC to “self heal” small defects. Increased alkalinity from oxygen reduction near
a defect would trigger release of Cr6+, which then could migrate to actively corroding
sites where it would be reduced to form a Cr3+ barrier.
The presence of surface activators and accelerators in coating baths is necessary
for quick and thick conversion coating formation (5). NaF, typically in concentrations of
30 mM to 40 mM, is used as an activator to destabilize the native aluminum oxide film
and promote the oxidation of the aluminum surface by Cr6+ (5, 38). Additionally,
additives like K3Fe(CN)6 in concentrations of 2 to 5 mM are used as accelerators because
the reduction kinetics of Cr6+ to Cr3+ and the oxidation of the aluminum surface are
3- 4- sluggish (5, 43). Fe(CN)6 is reduced to Fe(CN)6 as the aluminum surface is oxidized,
4- 6+ 3+ Fe(CN)6 can then be oxidized by Cr species resulting in the formation of a Cr layer
(43).
Chromates may also be used directly as pigments in paints and primers, providing a reservoir of inhibitor anion available for inhibition (44). Paints and primers exposed to an aqueous electrolyte will take up water. As water saturates a pigmented coating, the pigment will provide inhibition to the underlying substrate (8). A usable pigment must be soluble enough to be sufficiently concentrated to provide passivity to the underlying alloy when exposed, but not so soluble as to promote osmotic blistering (44). Osmotic
18 blistering is a phenomenon caused by osmotic pressure resulting from differences
between the water activity at the substrate-coating and coating-water interfaces (44). To
avoid osmotic blistering, it has been reported that the solubility of a pigment must be less
than 2 grams per 100 mL of water (44). There are few inorganic species that satisfy the
overall solubility criteria. However, a number of chromates, and in particular SrCrO4, have appropriate solubility and inhibition efficiency for use as inhibitor pigments. For example, SrCrO4 has a saturation concentration of approximately 5 mM in near neutral
solutions, but the critical concentration necessary for inhibition is between 0.1 and 1 mM
(44). The consequence is that SrCrO4 pigments are soluble enough to be present in concentrations well above the critical concentration necessary for inhibition (44).
6+ Additionally, Cr species from SrCrO4 pigments may be leached into solution and
transported to active defect sites where Cr6+ species are reduced and corrosion is suppressed (8). In this respect, SrCrO4 pigmented primers are similar to CCCs in that
both demonstrate ability to “self heal”. However, there is an important distinction in the
6+ mechanism of Cr release from CCC and SrCrO4 pigments dispersed in a primer; the
prior is a reversible pH dependent reaction while the former will leach into solution until a finite solubility is reached (5 mM) (42). Once released from either CCC or a pigmented
paint, the interaction between Cr6+ species and an actively corroding site is likely similar.
Soluble Cr6+ species, available across a wide range of pH values, inhibit reduction
reactions through formation of a thin Cr3+ layer. A desire to identify an alternative inhibitor that supplies the same level of inhibition in similar environments, but without
the carcinogenic risks, exists.
19 2.4 Vanadate Inhibitors as Possible Alternatives to Chromate
Vanadates have shown promise as alternatives to chromates for corrosion inhibition of aluminum alloys such as Al 2024-T3 (1, 2, 6-9, 45-48). In particular, vanadates in NaCl solutions have been shown, under specific conditions, to be modest anodic inhibitors and potent cathodic inhibitors for Al 2024-T3 (6-9, 45-48). There are ways to functionalize soluble vanadates through the formation of films and conversion coatings, and by incorporation into organic coatings (7-9, 49). Additionally, a number of vanadate surface treatments and pigmented organic coatings applied to zinc and aluminum substrates have been observed to “self-heal” in a fashion resembling that shown by chromate-based coatings (7, 49). Although evidence exists that vanadates in general are inhibitors, a complicated aqueous speciation has prevented an understanding of the precise mechanism of inhibition, and vanadate use is not completely benign.
2.4.1 Vanadate Background and Toxicity
Vanadium is a trace element widely distributed throughout nature, comprising
0.014% of the earth’s crust and in concentrations of 20-35 nM in the earth’s oceans (50,
51). The effect and essentiality of vanadium for biological systems has had a long, contentious, and yet unresolved history (51, 52). There is some evidence that low concentrations of vanadates, which have similar chemistry to phosphates, may interact with biological molecules in a similar fashion (53-56). Additionally, vanadates may influence cardiovascular activity, hormones, enzymes, and have been shown to mimic the behavior of insulin (55-57). It is generally accepted that high levels of vanadium are toxic and evidence exists that large doses are reproductive and developmental toxins (51,
20 52, 55, 56). Overexposure to vanadium through diet is unlikely as animal and human studies have shown that ingested vanadium is not readily absorbed in the digestive track.
However, exposure to airborne vanadium is a risk (51, 52). The toxicity of vanadate has
been suggested to be tied to valance, with V5+ being toxic, while V4+ is benign with
possible uses in medicine (58). Despite the risks associated with vanadates, they have
been shown to be promising inhibitors and their use relative to chromates appears to be
less dangerous because there is no suggestion that vanadates are carcinogens.
2.4.2 Vanadates as Inhibitors of Corrosion
The focus on vanadates as possible alternatives to chromates is based on previous
work which has shown vanadates to be corrosion inhibitors in multiple systems. For
example, Bienstock and Field observed significant inhibition on carbon steels exposed to
boiling potassium carbonate solutions with sodium metavanadate and vanadium
pentoxide (59). Hinton has shown decreased corrosion rates on Al 2014 in 0.001 M NaCl
with 0.002 M NaVO3 (1, 2). Further, vanadium alloy additions to 18% chromium ferritic
stainless steel have been observed to increase the pitting potential, albeit by a different
mechanism (60).
Additionally, previous work suggests that whether soluble vanadates provide
inhibition to Al 2024-T3 is strongly dependent on the solution pH. Guan et al. and
Buchheit et al. reported an approximately 100 mV increase in pitting potential for a 0.124
M NaCl solution at pH 6 with 0.1 M NaVO3 and a small inhibiting effect on the rate of
oxygen reduction (7, 8). Cook observed a two-order-of-magnitude decrease in corrosion
current from polarization measurements in 0.6 M NaCl solution at pH 7 with 3.4 mM
21 NaVO3 on Al 2024-T3 compared to a control; similar to the decrease observed for
sodium chromate additions (6). Additionally, electrochemical impedance spectroscopy
(EIS) measurements on Al 2024-T3 samples exposed to 0.6 M NaCl solutions with 3.4
mM NaVO3 carried out over 10 days showed aqueous NaVO3 was poorly protective at
pH 3, but was inhibiting at pH 7 and 10 (6). Iannuzzi observed significant decreases in
cathodic kinetics in 0.5 M NaCl solutions containing 10 mM NaVO3 at pH values
ranging from 7.8 to 9.3 (48). A similar reduction in cathodic kinetics was observed in pH
8.71, 0.5 M NaCl solution with 5 mM monovanadate (45). There is a body of evidence
that suggests that vanadates are inhibitors of corrosion depending on pH. Further,
speciation of vanadates is also dependent on pH (31). Therefore, inhibition is likely
dependent on speciation as has been suggested by others (45-48).
2.4.3 Aqueous Speciation of Vanadate
Compared to Cr6+, V5+ has a complex aqueous chemistry, which depends on pH,
concentration, and ionic strength (31, 55, 61, 62). The equilibrium speciation diagram in
Figure 2.4 shows speciation as a function of pH and concentration (31, 61). Most of the
species in the predominance diagram fall into two broad categories: tetrahedrally
coordinated species, which dominate in alkaline and dilute solutions, and octahedrally coordinated species, which dominate in concentrated acidic solutions (31, 55, 61-63).
Although a specific species will be dominant in a solution of a given pH, concentration, and ionic strength, a number of other species are present in solution in smaller, but appreciable concentrations, creating a complex mixture of oligomers in equilibrium (31,
55, 56, 61-66). These oligomers may include monomers (V1), dimers (V2), trimers (V3),
22 tetramers (V4), pentamers (V5), hexamers (V6), and decamers (V10) (55, 56, 61, 62, 64,
65). Additionally, a number of these oligomers may exist in either linear or cyclic
configurations or only as transients or in solutions of high ionic strength such as the case for V3 and V6 (55, 61, 62, 64, 66). Vanadates that have a cyclic structure are more stable
and, as a result, less reactive than non-cyclic vanadates (67). Additionally, V2 is reported
to be the most reactive species (67). From a corrosion inhibition perspective, the variety
of possible species in solution and the complex equilibria among them are daunting,
making identification of the species responsible for inhibition and to what extent difficult.
There is some disagreement in the literature as to the exact pH at which particular
species or oligomers dominate and multiple species are often grouped together for
discussion in the literature. The following is a summary with nomenclature, solution
color, coordination, structure and approximate pH range of various species broadly
grouped as vanadates, pyrovanadates, metavanadates, decavanadates, and
monovanadates. Figure 2.5 shows some general structures of vanadate species to be
discussed in the following (68).
3- Vanadate or orthovanadate (VO4 ) is a single tetrahedral unit, which forms a colorless
solution and dominates at a pH values greater than 13-14 (31, 55, 61, 63).
4- 3- 2- Pyrovanadates (V2O7 , HV2O7 , and VO3(OH) ) are tetrahedrally coordinated,
predominate at pH values between 9 and 12 and form colorless solutions (31, 61, 63).
Pyrovanadates contain two vanadium atoms and have a structure that consists of two corner-sharing tetrahedral vanadate units (55, 61, 62).
3- 4- 5- 6- - Metavanadates (V3O9 , V4O12 , V5O15 , V6O18 , VO(OH)3, and VO2(OH)2 ) are
tetrahedrally coordinated, form colorless or yellow solutions depending on concentration,
23 and predominate at pH values between 6 and 9 (31, 61, 63, 65). The structure of metavanadates consists of corner-sharing tetrahedral vanadate units that form rings or strings (55, 61, 62).
6- 5- 4- Decavanadates (V10O28 , V10O27(OH) , V10O26(OH)2 ) consist of a combination of 10 edge-sharing VO6 octahedral coordinated vanadium units and predominate at pH between
2 and 6, forming orange or red solutions (31, 61-63). Six edge-sharing octahedral vanadate units are arranged in a rectangle with two added edge-sharing octahedral units positioned above and below the center of this rectangle plane (31, 61-63).
3- - 2- 2+ Monovandates (VO4 , VO(OH)3, VO2(OH)2 , VO3(OH) , and VO ), which can alternatively be classified as vanadate, pyrovanadate, and metavanadates as previously discussed, predominate at dilute concentrations and are typically tetrahedrally
+ coordinated (31, 61). The pervanadyl cation (VO2 ) is yellow and dominates in strongly acidic solutions (31, 61, 63).
Equilibria between different tetrahedral vanadate species with each other has generally been reported to occur quickly, on the millisecond time scale (62, 66, 67).
Likewise, the formation of decavanadate from the acidification of metavanadate or other solutions containing tetrahedrally coordinated species occurs relatively quickly (69). In contrast, the kinetics of decavanadate decomposition into metavanadate or tetrahedral vanadates is slow and likely involve the formation of intermediate species (31, 56, 61, 65,
69). The half-life of decavanadate decomposition, when thermodynamically favorable, has been reported to be approximately 5 to 15 h or days depending on exact solution conditions (56, 61, 62). Slow decomposition kinetics are possibly attributed to the compact structure of decavanadate or the fact that decavanadate has a different
24 coordination than the tetrahedrally coordinated products (69). As will be discussed in
later chapters, the kinetics of decavanadate decomposition have direct bearing on
synthesis and observed inhibition from vanadate-based synthetic anionic exchange clay pigments (8, 9).
2.4.4 Current Understanding of the Mechanism of Vanadate Inhibition
As previously discussed, vanadates are inhibitors of corrosion and inhibition seems to be linked closely with speciation. Specifically, there is a body of evidence that suggest vanadates are modest anodic inhibitors, particularly in near neutral solutions but mainly act through the suppression of cathodic kinetics. For example, anodic polarization experiments on Al 2024-T3 samples exposed to 0.124 M NaCl solution at pH 6 with and without 0.1 M NaVO3 showed the development of a region of imperfect
passivity which occurred up to 100 mV above the OCP (7, 8). Cathodic polarization in
similar solutions showed that vanadates in neutral solutions inhibit rates of oxygen
reduction (7, 8). Analogous results were found on Al 2024-T3 coated with a vanadate
conversion coating (VCC) in 0.5 M NaCl compared to the uncoated alloy; an increase in
the pitting potential was observed from anodic polarization experiments and suppressed
oxygen reduction rates (7). Further, Mahajanam found that, in pH 6.5 NaCl solutions
6- with 0.1 M NaVO3 and 0.1 M [V10O28] , increased pitting potential and decreased
cathodic kinetics were observed (9). Additionally, 0.1 M NaVO3 at pH 6.5 was observed
to increase the pitting potential more than solutions at pH 6.0 (9). In a study of pit depth
25 and frequency, 3.4 Mm NaVO3 addition to 0.6 M NaCl adjusted to pH 7 were observed to decrease pitting damage significantly on Al 2024 (6). The same study noted that the presence of oxygen was necessary for inhibition from NaVO3 (6).
Reductions in cathodic kinetics have been observed in alkaline NaCl solutions
containing metavanadate. Iannuzzi et al. found very modest inhibition of oxygen
reduction decavanadate solutions adjusted to alkaline pH values, but observed significant
cathodic inhibition in alkaline metavanadate solutions (46, 48). It was suggested that
higher molecular weight vanadates, such as V4 and V10, do not contribute to reduced
oxygen reduction rates (48). Further, anodic polarization curves in both decavanadate
and metavanadate solutions indicated reduction or elimination of transient dissolution of
Al2CuMg (45, 46). As Al 2024 is anodically polarized, the alloy will have two different
breakdown potentials; the first is associated with the dissolution of Al2CuMg and the second is from intergranular and matrix attack (45, 70). Previous observations in vanadate solutions are inline with work using insitu atomic force microscopy scratching, which showed additions as small as 0.1 mM metavanadate to 0.5 M NaCl suppressed attack of Al2CuMg particles (47). The mechanism of suppression of transient dissolution
was not clear, but it was speculated that monovanadates on the matrix prevent or displace
Cl- adsorption on the surface, which hinders subsequent oxide film breakdown (46). In
regards to cathodic inhibition, it was found that reduction of monovanadates occurs
slowly suggesting that protection from vanadates is from adsorbed species on
intermetallics, which block cathodes, rather than reduction of species to form a protective
film (45, 46).
26 Although S-phase dissolution was suppressed and modest reductions in oxygen
reduction were observed in dilute decavanadate solutions, in comparison to metavanadate
solutions decavanadates are poor inhibitors (45, 48). Decavanadates were found to be easily reduced in both aerated and deaerated solutions and decavanadate reduction
contributes to cathodic current (46).
2.4.5 Vanadates Functionalized in Films and Coatings
There are a number of different ways that vanadates can be incorporated into
coatings, including as conversion coatings and pigments (7-9). VCC can be formed using
similar procedures and bath chemistries as for CCCs (7). The formation of VCCs occurs
by a sol-gel process that involves hydrolysis, condensation, and polymerization of
octahedrally coordinated vanadate species (7). These VCCs have been observed to
increase corrosion resistance by both increasing pitting potential and decreasing the rate
of oxygen reduction in NaCl solutions (7).
Although soluble vanadates in solution have been shown to have promise as
inhibitors of corrosion on Al 2024-T3, vanadates are generally too soluble for direct use
in organic coatings (1, 2, 6-9, 45-48). A useful pigment must be soluble enough to
provide inhibition to the substrate but so not soluble as to cause blistering of the coating
(44). The use of relatively less soluble vanadate compounds as pigments, such as
. Sr(VO3)2 H2O, which has approximately one fifth the solubility of SrCrO4, has been
attempted and shown increased corrosion resistance on zinc substrates (49, 71).
. Although the leaching kinetics of Sr(VO3)2 H2O were found to be similar to SrCrO4, the
vanadate pigment could not supply high enough concentrations of vanadate to satisfy
27 demand for inhibitor on the substrate (49, 71). In this respect, leaching rate was found to govern the amount of inhibitor available for inhibition rather than saturated solubility of the pigment (49, 72). This possibly indicates that vanadates have lower inhibitor efficiency than chromates, and that for successful use of vanadates as pigments higher concentration must be made available in solution. Also, this presents an obstacle to elevation of the concentration of leached vanadate in solution without leading to blistering.
Vanadates incorporated into organic coatings and VCC have been observed to have “self healing” characteristics on both aluminum and zinc substrates (7, 49, 71).
Increased corrosion resistance has been observed on unprotected metal panels in close proximity, through an electrolyte, to a vanadate-pigmented or conversion coated panel (7,
71).
Additionally, vanadates incorporated into anionic-exchange pigments (to be discussed in detail later) have been observed to inhibit the corrosion of Al 2024 (8, 9).
When these pigments come into contact with an aggressive electrolyte inhibitor, anions stored within the pigment are selectively exchanged for aggressive anions, releasing inhibitor into solution. Successful inhibition has been observed from decavanadate anions released from anionic-exchange pigments (8, 9). These results are at odds with previous work has shown decavanadate to be a poor inhibitor of corrosion on Al 2024
(48). It is possible that under alkaline solution conditions, such as near sites of oxygen reduction, decavanadates quickly begin to decompose into vanadate species that have been observed to inhibit corrosion.
28 2.5 Hydrotalcite Pigments
Hydrotalcite is a naturally occurring Mg-Al hydroxycarbonate discovered in
Sweden in 1842 (73). However, the term “hydrotalcite” is often generically used to
describe other layered double hydroxides (73, 74). Hydrotalcite materials are anionic
exchange clays that have been studied widely for use in catalysis, filtration, and scavenging, but have recently been shown to be a novel way to store and deliver corrosion inhibitors (8, 9, 73, 75, 76). The main advantage of ion exchange clay pigments is that they allow for storage and release of inhibitors that otherwise would be too soluble for direct use in organic coatings without causing osmotic blistering and extreme leaching (77, 78). When an aggressive electrolyte comes into contact with a hydrotalcite-pigmented coating, inhibitor anions incorporated into the pigment will be selectively exchanged with aggressive solution anions. This process is shown schematically in Figure 2.6. The unique properties and ease of synthesis allow for the development of custom made, synthetic hydrotalcite-pigments using a wide variety of cation hosts and inhibitor guests that otherwise would not be appropriate for use in organic coatings.
2.5.1 Structure of Hydrotalcite and Hydrotalcite-like Materials
Hydrotalcites are mixed metal layered double hydroxide anionic clays and may be represented generically as follows:
x+ n-. x- [Mg1-xAlx(OH)2] [Ax/n mH2O] Eqn 4.
29 where, x is between 0 and 0.33, and An- is the exchangeable anion (79). The structure of
an ideal hydrotalcite, seen in Figure 2.7, consists of alternating negatively-charged anion
layers and positively charged cation layers (73). The basic building blocks of
hydrotalcite structures are based on brucite, Mg(OH)2 cationic sheets (74, 80-82).
Brucite consist of Mg2+ cations with 6-fold coordination to OH-; these ocatahedral units
form sheets by sharing edges (73, 74). The cationic sheets are separated by inter-layers,
containing H2O and anions, and can be stacked to produce either rhombohedral or
hexagonal symmetries (73, 74, 80). The ability to store anions in interlayers comes from
a netpositive charge developed in the cationic layers from substitution of approximately
one-third of trivalent cations for divalent cations (73, 74, 80, 82, 83). For substitution to
occur, the trivalent cations must have similar size radii as the divalent cations (74, 80,
82). Within the cationic-brucite-like layers, the trivalent ions tend to distance themselves
from each other due to repulsion of like charges resulting in a degree of ordering (74).
The net positive charge developed in these layers is balanced by anions, which can be
incorporated with water into the inter-layer regions of the structure (73, 74, 80, 82). The
inter-layers are highly disordered and are capable of incorporating a number of different
metal cations that can subsequently be exchanged for other cations (73, 74, 82).
The previously presented formula can be expanded to include another variant of the hydrotalcite structure, which can be synthesized with a combination of trivalent and monovalent cations (74, 84):
z+ 3+ A+ m- . [M 1-xM x(OH)2] X A/m nH2O Eqn. 5
30 where A=x for z=2 and A=2x-1 for z=1. These structures can only be synthesized using
Li+ because Li+ is the only monovalent cation with an ionic radius of appropriate size to form hydrotalcite-like materials (74). In this structure, the aluminum cations are arranged in sheets similarly to gibbsite, Al(OH)3, with lithium ions producing a net positive charge by occupying the vacant octahedral sites; gibbsite has the same structure as brucite except one-third of the octahedral sites are vacant (74, 84).
The exact structure and dimensions of a hydrotalcite depend on a number of factors, including the type and ratio of cations, the type of anions, and the degree of hydration. In discussions about hydrotalcite, ‘a’ is the cation-cation distance within a positively charged brucite-type layer (83). This parameter is largely determined by the composition and ratio of cations in the brucite-type layer; as smaller radius trivalent ions are substituted into the brucite skeleton, the ‘a’ dimension shrinks (73, 74). Another parameter, ‘d’, refers to the thickness of an anion interlayer, also known as a gallery, plus the thickness of one cation sheet (73). A unit cell consists of 3 repeat d spacings (73).
The layered structure of hydrotalcites lends itself to easy characterization using x-ray diffraction (XRD).
XRD is a quick and easy method available for structural confirmation of hydrotalcite-like materials. Figure 2.8 shows a sample XRD pattern for Al-Zn hydrotalcites containing decavanadate or Cl- (8). The figure shows features typical of hydrotalcites. Reflections resulting from the cationic layers occur at low 2θ and are given (003), (006), and (009) indices (8). The structure within cationic sheets results in
31 reflections at large 2θ (8, 74). Additionally, gallery height can be used to identify what
anion is intercalated in the interlayer and changes in height can be used to sense anion
exchange (8).
2.5.2 Synthesis of Hydrotalcite-like Materials
Many methods are available to produce hydrotalcites intercalated with different
anions (85). Direct synthesis is an easy and attractive method that can be accomplished
without the complications of other synthesis methods (86). However, many issues need
to be taken into account to successfully synthesize a “clean” hydrotalcite, including
cation and anion ratios, co-precipitation pH, anion stability, and possible complications
from atmospheric CO2.
Cavani et al. offer two generic formulas to produce pure hydrotalcites based on
divalent and trivalent cation skeletons (74):
M (III) 0.2 ≤ ≤ 0.4 Eqn. 6 [M (II) + M (III)]
1 An− ≤ ≤ 1 Eqn. 7 n M (III) where M(III) and M(II) are concentrations of trivalent and divalent cations, respectively,
and n and A are anion charge and concentration, respectively. These two equations help
ensure that the correct ratio of cations is available for synthesis and that there is a
sufficient anionic charge to populate the interlayers with a desired anion. The importance
of these ratios is that pure hydrotalcites are only formed in a narrow window of cation
32 ratios; deficiencies of the divalent cation lead to the formation of the trivalent cations hydroxide and excess of the divalent cation leads the formation of the divalent cations
hydroxide (87). There does not appear to be a similar published set of rules for synthesis
using monovalent and trivalent cations; however, the theoretical ratio for the formation of
ordered cation sheets for monovalent and trivalent cation is 2 to 1 (84). Additionally, Li+ is the only monovalent cation with an ionic radius of appropriate size to form hydrotalcite-like materials (74).
A further consideration is that different cation combinations will co-precipitate at different pH values. The range of co-precipitation can be determined using titration curves which are produced when a solution containing a pair of cation salts is titrated with alkali. The resultant curves have distinct plateaus, which indicate pH regions of precipitation; the intermediate plateau occurs at the pH of co-precipitation and plateaus at pH extremes occur from precipitation of metal hydroxides. Figure 2.9 shows sample titration curves for different cation combinations (74). For direct synthesis, the desired anion to be incorporated into the hydrotalcite must be stable and predominant in solution at the same pH range of cation co-precipitation. The stability and predominance of particular anions can be determined by referencing predominance diagrams and speciation studies.
2.5.3 Anion Exchange and Selectivity
When hydrotalcite-like pigments come into contact with an electrolyte, the inhibitor anions intercalated in the pigment may be selectively exchanged for aggressive anions in solution. High selectivity for aggressive anions such as Cl- is advantageous,
33 providing inhibition while simultaneously neutralizing the environment through removal
of aggressive species (75, 76). However, this exchange process is also an important
consideration during the synthesis of hydrotalcite-like materials as it is possible for
contaminant anions to be incorporated into galleries of the pigment. Both anion size and
charge density affect the selectivity of an anion (79, 88). In general, the selectivity for an anion increases as the diameter of the anion decreases; smaller anions are more easily incorporated into interlayer space than larger anions (79, 88). Further, species with higher charge densities have higher selectivity (88). For example, for exchange into an
3- 2- - Al-Li compound, the selectivity for phosphates is as follows: PO4 > HPO4 > H2PO4
(88).
- 2- OH and CO3 anions are both noted to have particularly high selectivity among
anions of similar valence (79). During synthesis of hydrotalcites requiring co-
2- precipitation at pH values higher than approximately 6.5, the availability of CO3 in solutions is a critical issue and its elimination requires purging the synthesis bath of CO2
(74, 82, 86).
2- The CO3 comes from dissolved atmospheric CO2. CO2 forms H2CO3 which has
- 2- equilibria involving HCO3 and CO3 , both of which have increased solubility in alkaline
solutions (89).
2.5.4 Hydrotalcites in Coatings
There has been a limited amount of work on the use of hydrotalcites as inhibitor
pigments and coatings, but hydrotalcites are suggested to have potential to be strong inhibitors. The formation of a continuous protective hydrotalcite-like coating on
34 3+ aluminum alloys had been achieved in Li2CO3 solutions doped with excess Al ions
using a similar procedure as used for other conversion coating processes (89).
Hydrotalcites have also been used as pigments in organic coatings and have been shown
to provide increased resistance to corrosion; specifically, Al-Zn-decavanadate hydrotalcites were observed to increase corrosion resistance on Al 2024-T3 (8, 9).
Protection was attributed to the release of decavanadates and Zn2+ into solution and the
subsequent uptake of Cl- ions (8, 9). Also, it is suspected that hydrotalcites used in these
studies were not “clean”, possibly causing an understatement of corrosion resistance.
- 2- 2- NO3 , CO3 , and CrO4 hydrotalcite pigments in organic coatings have been shown to
slow the propagation of filiform corrosion and coating delamination on Al 2024-T3 (75,
76). These hydrotalcite pigments have been shown to slow attack by exchanging
inhibitor anions for aggressive chloride anions and by neutralizing the pH by
consumption of hydrogen ions at sites of anodic activity (75, 76). More impressively, the
2- HT-CrO4 coatings were able to provide comparable corrosion resistance to regular
2- SrCrO4 pigments using 30 times less CrO4 on a molar basis (75, 76). This may suggest
that exchange and neutralization properties make hydrotalcites potent inhibitors
regardless of what inhibitor anion is intercalated.
2.6 Critical Issues
There is a body of evidence that suggests that vanadates are inhibitors of
corrosion and that vanadates may be functionalized into coatings to provide inhibition to
aluminum alloys (7-9). Further, inhibition from vanadates has been observed to vary
depending on pH and depend strongly on speciation; however, this relationship has not
35 been fully characterized. Additionally, the microstructure of Al 2024-T3, which results
in superior mechanical properties, also makes the alloy inherently susceptible to localized corrosion. For a complete understanding of inhibitor action from vanadates, inhibition must be viewed through the prism of electrochemical differences between constituent intermetallic particles and the matrix. Further, inhibition has been observed from Al-Zn decavanadate-bearing hydrotalcites and anion exchange pigments may offer a way to incorporate a wide variety of highly soluble inhibitors into organic coatings allowing for tailored application of specific inhibitors.
First, inhibition from vanadates must be systematically characterized over a range of pH values and vanadate concentrations. This would help in identifying what pH values and concentrations result in the most potent inhibition. Ultimately this characterization will further the understanding of inhibition mechanisms and the roles that various vanadate species have in those mechanisms.
Second, the microstructure of Al 2024 is complex and plays a direct role in corrosion and property degradation of the alloy. Vanadate inhibition must be characterized in terms of interactions with the matrix and constituent particles of Al
2024-T3. Specifically, an understanding of vanadate interaction with Al2CuMg intermetallics in inhibiting vanadate solutions is necessary.
Finally, it is known that vanadate-bearing hydrotalcite pigments provide inhibition to organically coated aluminum alloys. The refinement of vanadate-bearing hydrotalcites and development of hydrotalcites with other inhibitor anions is desired.
The hope is that hydrotalcites tailored and maximized for specific environments may be able to provide comparable inhibition as chromates over narrow environmental windows.
36 A framework by which hydrotalcites could be designed to incorporate a number of different inhibitor anions into various cation skeletons is necessary. Additionally, these pigments must be characterized in terms of performance and inhibitor release relative to
SrCrO4 pigments.
37 FIGURES
Figure 2.1: Corroded aluminum surface showing two different kinds of intermetallic particles commonly found in Al 2024, small round Al2CuMg and Al2Cu intermetallics, and large irregular shaped Al-Cu-Fe-Mn intermetallics. Reprinted with permission from Elsevier (15).
38
Figure 2.2: Aqueous speciation of Cr6+ species as a function of concentration and pH. Reprinted with permission of ASM International® (5). All rights reserved. www.asminternational.org
39
Figure 2.3: Sequence of events of Cr3+ film formation (36). Reproduced with permission from J. Electrochem. Soc., 150, B367 (2003). Copyright 2003, The Electrochemical Society.
40
Figure 2.4: Equilibrium predominance diagram for V5+ aqueous species (90). Reproduced with permission from J. Electrochem. Soc., 155, C350 (2008). Copyright 2008, The Electrochemical Society.
41
Figure 2.5: Illustrations of different vanadate structure as a function of pH and concentrations ranging from 10-3 to 10-1 mol L-1. Reprinted with permission from (68). Copyright 1991 American Chemical Society.
42
Figure 2.6: Schematic representation of selective inhibitor release and exchange for contacting aggressive anions.
43
Figure 2.7: Ideal structure for a hydrotalcite containing carbonate in the interlayers. Reprinted with permission from Elsevier (73).
44
Figure 2.8: X-ray diffraction pattern for Al-Zn hydrotalcites containing decavanadate and chloride. Reprinted with permission from Elsevier (8).
45
Figure 2.9: Titration curves for different cation combinations showing the ranges of pH where co-precipitation and precipitation of cation-hydroxides occur. Reprinted with permission from Elsevier (74).
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52
CHAPTER 3
CORROSION INHIBITION OF ALUMINUM ALLOY 2024-T3 BY AQUEOUS
VANADIUM SPECIES
3.1 Introduction
Vanadates have shown promise as alternatives to chromates for corrosion
inhibition of aluminum alloys such as 2024-T3 (1-7). In particular, vanadates in NaCl
solutions have been shown, under specific conditions, to be modest anodic inhibitors and
potent cathodic inhibitors for 2024 (3-8). Further, there are ways to functionalize soluble vanadates through the formation of films and conversion coatings, and by incorporation into organic coatings (7-12). Additionally, a number of vanadate surface treatments and pigmented organic coatings applied to zinc and aluminum substrates have been observed to “self-heal” in a fashion resembling that shown by chromate-based coatings (7, 11, 12).
Whether soluble vanadates provide inhibition of Al 2024-T3 is strongly dependent on the solution pH. Guan et al. and Buchheit et al. reported an approximately
100 mV increase in pitting potential for a 0.124 M NaCl solution at pH 6 with 0.1 M
NaVO3 with little effect on the rate of oxygen reduction (7, 8). Cook observed a two-
53 order of magnitude decrease in corrosion current from polarization measurements in 0.6
M NaCl solution at pH 7 with 3.4 mM NaVO3 on Al 2024-T3; similar to the decrease
observed for sodium chromate additions (6). Further, electrochemical impedance
spectroscopy (EIS) measurements on Al 2024-T3 samples exposed to 0.6 M NaCl
solutions with 3.4 mM NaVO3 carried out over 10 days showed aqueous NaVO3 was
poorly protective at pH 3, but was inhibiting at pH 7 and 10 (6). Iannuzzi observed significant decreases in cathodic kinetics in 0.5 M NaCl solutions containing 10 mM
NaVO3 at pH values ranging from 7.8 to 9.3 (3).
V5+ has a complex aqueous chemistry that depends on pH and concentration as
seen in the equilibrium speciation diagram in Figure 3.1 (13, 14). Although there is some
disagreement in the literature as to the exact pH at which particular species dominate, the
following is a summary with nomenclature, solution color, and pH range of the various
3- species under consideration. Orthovanadates (VO4 ) are colorless and dominate at a
solution pH greater than 13 (13, 15). As highly basic solutions are acidified to pH values
between 9 and 12, orthovanadate tetrahedral units can combine to form pyrovanadates
4- 3- 2- (V2O7 , HV2O7 , and VO3(OH) ), which are also colorless (13, 15). Continued
acidification to pH values between 6 and 9 leads to the formation of colorless or yellow
3- 4- 5- - metavanadates (V3O9 , V4O12 , V5O15 , VO(OH)3, and VO2(OH)2 ) (13-16). Additional
acidification to pH between 2 and 6 causes the coordination of vanadium to change from
tetrahedral to octahedral. A combination of 10 octahedral units forms the decavanadate
6- 5- 4- ion (V10O28 , V10O27(OH) , V10O26(OH)2 ), which are orange or red solutions (13-15).
+ Pervanadyl (VO2 ) is yellow and dominates at dilute concentrations in strongly acidic
solutions (13-15). As seen in Figure 3.1, small concentrations of vanadate exist
54 exclusively as monovanadate species, while higher vanadate concentrations result in
larger species consisting of multiple vanadium atoms (13). Although a specific species
will be dominant at a certain pH and concentration combination, a number of other species are present in solution in smaller, but appreciable proportions (13). As a result, complex equilibria may mask which species are the causes of inhibition and to what extent. Equilibria between different vanadate species have generally been reported to occur quickly except for the equilibrium between metavanadate and decavanadate; when approached from the acidic side, equilibrium is established only after several hours and possibly involves the formation of intermediate species (13-15).
There is a body of evidence that suggests that vanadates are inhibitors of corrosion and the extent of inhibition depends on speciation, however, this relationship has not been fully characterized. The objective of this work is two-fold: 1) to systematically characterize vanadate speciation and inhibition over a range of pH and vanadate concentration, and 2) to further the understanding of inhibition mechanisms and the roles that various vanadate species have in those mechanisms.
3.2 Experimental Procedures
3.2.1 Materials and Chemicals
2.0 mm thick Al 2024-T3 sheet was used in all electrochemical and exposure experiments. Aluminum 2024-T3 is a solution heat-treated and cold-worked alloy whose nominal composition is Al, 3.8-4.9wt% Cu, 1.2-1.8wt% Mg, 0.3-0.9wt% Mn, 0.5wt% Fe,
0.5wt% Si, 0.25wt% Zr, 0.1wt% Cr, and 0.1wt% Ti (17). Solutions for potentiodynamic polarization and exposure experiments were prepared using reagent grade NaCl, 10 N
55 NaOH solution, and HCl purchased from Fisher Scientific, and NaVO3 (assay ≥ 98%)
purchased from Fluka Chemika. The NaVO3 used for NMR was obtained from Alfa
Aesar (assay 96%). 18.2 MΩ.cm deionized water was used for preparation of all
solutions.
3.2.2 Sample Preparation
Samples for potentiodynamic polarization experiments were prepared by abrading
Al 2024-T3 coupons with a Scotch-Brite© pad for 15-20 seconds. The samples were then
rinsed with tap water to remove debris followed by a deionized water rinse. The coupons
were dried using an air hose and experiments were carried out within minutes of sample
preparation. This method allowed for quick and easy sample preparation and
reproducible results. For solution exposure experiments, Al 2024-T3 sheets were cut into
coupons, 2 cm by 2 cm, and mounted in epoxy to avoid any possible galvanic coupling
associated with conductive Bakelite mounting media. Coupons for the exposure
experiments were polished with SiC paper and finished with 1μm diamond paste. Ethyl alcohol was used for the last SiC polishing step to help avoid onset of corrosion prior to experimentation. Samples were washed using an ultrasonic ethyl alcohol bath prior to experiments to remove any residual polishing media.
3.2.3 Nuclear Magnetic Resonance (NMR)
A Bruker DPX 400 MHz superconducting magnet was used to collect high resolution 51V (105.2 MHz) NMR spectra. An indirect detection probe was used with a
90o pulse duration of 10.38 μs. Spectra were collected using 1024 transients, a spectral 56 window of 73,529 Hz, a 0.051 s acquisition time, and a 0.20 s relaxation delay. Each
spectrum had the subsequent process parameters applied: 1.00 Hz line broadening, zero
filling (25 K points), and baseline correction. A solution consisting of 20% v/v VOCl3 in
51 51 C6D6 (δ V = 0ppm) was used as an external standard to reference the V chemical
shifts.
Four different NMR experiments were performed. In the first experiment, a 100
mM as dissolved NaVO3 solution, initial pH 8.76, was acidified to pH 4.12 with HNO3, followed by a pH adjustment to 7.58 with NaOH to observe the kinetics of decavanadate decomposition. In the initial solution, once the NaVO3 was dissolved the solution pH
remained stable. NMR spectra of the initial pH adjusted and readjusted solution were
conducted within a few minutes of pH modification. NMR spectra and pH measurements were taken with time over a period of 1468 hours. A second experiment was conducted to observe the effects of multiple pH adjustments on a NaVO3 solution. Using a pH 6.08
NaVO3 solution produced after 1468 hours of equilibration in the previously mentioned
experiment, the pH was adjusted by one-drop additions of NaOH at 0, 24, and 51 hours.
To characterize the solution throughout the experiment, NMR spectra were collected and
pH was measured for samples of the initial solution, a solution after 2 NaOH additions
collected prior to the third NaOH addition (spectrum at 52 hours of total time), and a
solution after 3 NaOH additions (spectrum at 52 hours total time). A third experiment
was conducted to observe the effects of NaVO3 solution exposure to aluminum. A 2.0
mm diameter aluminum wire (99.999%) obtained from Alfa Aesar was abraded with 600
grit SiC paper, cleaned with alcohol, and placed inside an NMR tube in contact with the
final aged pH 6.08 NaVO3 solution produced after 1468 hours of equilibration from the
57 previously mentioned first NMR experiment. Another experiment was run in parallel
where an aluminum wire was placed in contact with an equilibrated pH 8.10 NaVO3 solution that contained primarily tetrahedral vanadate species. The wires were removed and subsequent NMR spectra were collected at 52 hours. A final NMR experiment was used to determine which vanadate species were released from a vanadate hydrotalcite pigment exposed to NaCl solution. Hydrotalcite pigments are being evaluated for possible use as inhibitor pigments (8, 18-20). The pigments were synthesized by the co- precipitation of zinc and aluminum chlorides in an orange decavanadate solution. 4.0 grams of the resultant decavanadate-hydrotalcite pigment were soaked in 40 mL of 0.1 M
NaCl for 20 hours at which time a sample of this solution was taken for NMR analysis.
3.2.4 Potentiodynamic Polarization
To characterize the inhibitive nature of vanadate, anodic and cathodic polarization curves were collected for Al 2024-T3 coupons in aerated and deaerated 50 mM NaCl solutions with NaVO3 additions resulting in 0.0 M, 0.0032 M, and 0.32 M NaVO3 and pH
adjustments of 3, 5, 8, and 10. This approach created a matrix of test solutions that, when
considered relative to the equilibrium predominance diagram for vanadate, allowed
characterization of the inhibitive effects of a number of different vanadate species. It
should be noted that NaVO3 is a buffer. At pH 5, the 0.32M NaVO3 solution required
enough HCl to cause the overall Cl- concentration to be nearly 4x that of other test
solutions. For this reason, a 0.32 M NaVO3 solution at pH 3 was not used in
experiments. To differentiate the effects of NaVO3 inhibition from pH-induced changes,
58 measurements were made in NaVO3-free solutions and the solution pH was measured
before and after every experiment to ensure the solution pH remained near the intended
pH.
Polarization curves were replicated a minimum of four times in aerated solutions
and twice in deaerated solutions. All polarization experiments were carried out in a flat
cell with a teflon o-ring using a standard three-electrode setup consisting of saturated
calomel reference electrode (SCE), a platinum counter electrode mesh, and an Al 2024-
T3 working electrode. Measurements were conducted using Princeton Applied Research
potentiostat/galvanostat models 263A or 273A in conjunction with Corrware© data acquisition software. The exposed sample area of each Al 2024-T3 coupon was 1cm2.
All polarization curve measurements were preceded by a 30-minute measurement of the open circuit potential. Anodic polarization curves were initiated –30 mV versus the OCP and reversed at 3.93x10-4A/cm2 finishing at –30 mV versus OCP. A scan rate of 0.5 mV/sec was used in all experiments. Cathodic polarization curves were initiated +30 mV versus the OCP and finished at –2.0 V vs. OCP. Most cathodic polarization curves were stopped before –2.0 V versus OCP was reached, but well after hydrogen evolution had begun.
3.2.5 Exposure Experiments
Al 2024-T3 coupons were exposed to different vanadate solutions to explore differences in corrosion damage accumulation as a function of pH and NaVO3 concentration. Exposure experiments were conducted in actively aerated 200 mL solutions of the same pH-concentration combinations used in the polarization
59 experiments. For comparison, a parallel set of samples were exposed to chloride solutions that did not contain NaVO3. Exposure experiment solutions were allowed to
sit for 1 hour prior to use to allow time for solution equilibration. All samples were
placed face up in beakers containing test solutions for 2.5 hours as air was actively
bubbled. The samples were carefully removed, gently rinsed with ethyl alcohol, and
allowed to air dry. Samples were then examined using a Quanta 200 scanning electron
microscope equipped with energy dispersive spectrometry (EDS) capabilities.
3.3 Results
3.3.1 Changes in Vanadate Speciation with pH Adjustment, Time, and Exposure to
Aluminum
The decomposition of decavanadate to metavanadate in an initially acidic vanadate solution whose pH was adjusted into the basic regime was characterized by
NMR. The goal of these experiments was to observe how quickly octahedrally coordinated decavanadate formed tetrahedrally coordinated species after an increase in solution pH. Figure 3.2 shows changes in pH of a 100 mM NaVO3 solution initially pH
8.76 then acidified to pH 4.12 with HNO3, followed by pH adjustment to pH 7.58 with
NaOH. These measurements were made to record pH changes in solutions on which
NMR measurements were made. The initial NaVO3 solution was clear and turned a
brilliant deep orange color upon acidification. After pH adjustment from pH 4.12 to 7.58
the solution remained a brilliant orange color with no observable color change. The pH
was observed to drop to 6.2 within 27 hours and remained between 6.05 and 6.2 over the
next 1441 hours. Figure 3.3 is a collection of NMR spectra that show changes in the
60 vanadate speciation as a function of time. Identification of vanadate species from NMR
measurements was made by comparison to literature (16, 21-23). As seen in Figure 3.3,
the initial as-dissolved vanadate solution contained monovanadates, VO(OH)3,
- 2- 4- 3- VO2(OH)2 , and VO3(OH) (V1) , divanadates V2O7 and HV2O7 (V2), cyclic
4- 5- tetravanadates, V4O12 (V4), and cyclic pentavanadates, V5O15 (V5) species (3, 13-15,
21). The cyclic species are typically combinations of tetrahedrally coordinated
monovanadates. After the pH was adjusted to 4.12, tetrahedrally coordinated
metavanadate and pyrovanadate fully transform to decavanadate within the detection
6- 5- 4- limits of the experiment, V10O28 , V10O27(OH) , V10O26(OH)2 (V10) (13, 14, 16). The solution pH was then readjusted to 7.58 and formation of V1, V2, and V4 were observed
within minutes. The relative amounts of metavanadate and pyrovanadate in comparison
to decavanadate continued to increase with time despite of a decrease in solution pH.
Although the initial formation of metavanadate and pyrovanadate from decavanadate is
observed to occur rapidly, the solution continues to equilibrate for hundreds of hours.
These experiments show that relatively modest increases in pH trigger rapid formation of
tetrahedrally coordinated metavanadate and pyrovanadate species, which persist even in
mildly acidic solutions.
Figure 3.4 shows the effect of serial NaOH additions on speciation in 100 mM
NaVO3 solution. The aim of this experiment was to observe vanadate speciation in the
presence of a “continuous” source of alkalinity, as might occur in the vicinity of a cathode in a localized corrosion cell. The initial solution was orange with a pH of 6.08 and contained decavanadate, metavanadate, and pyrovanadate species. A one-drop addition of 10 N NaOH resulted in a pH increase to 9.21 with no observable color
61 change. After 24 hours the pH had drifted down to 6.34. Another drop of 10 N NaOH
was added and the pH increased to 8.97. After 27 hours the pH had drifted to 6.63 and
the color was observed to have changed to yellow. NMR spectra on the pH 6.63 solution showed the presence of decavanadate, metavanadate, and pyrovanadate, but the relative amount of metavanadate and pyrovanadate increased dramatically compared to the initial solution. Addition of a third drop of 10 N NaOH caused the pH to increase to 9.16 with no observable change in color. NMR showed this solution to contain V1 and an increase
in V1 and V2 concentration compared to other forms of vanadate. This experiment
demonstrated that vanadates tend to form simple tetrahedral units on alkalization.
Figure 3.5 shows NMR spectra from 100 mM NaVO3 solutions exposed to a pure aluminum wire. This experiment was carried out to understand how contact with a reducing metal might affect speciation of vanadate in solution. Spectra A and B are from a solution initially containing decavanadate. When an aluminum wire was placed into a
pH 6.08 100 mM NaVO3 orange solution (spectra A), the amount of metavanadate
increased relative to decavanadate (spectra B). Within 4 hours, the solution color was
observed to change from orange to an emerald green, and after 52 hours the solution had
turned dark green. Over the course of these color changes the pH remained near 6. The
green solution may be an indication of the formation of a vanadous ion (V3+, VOH2+, or
VO+) from the reduction of V5+ (24).
Spectra C and D in Figure 3.5 show behavior in a solution that was initially free
of decavanadate. Spectrum C shows the initial solution dominated by tetrahedrally
coordinated metavandates. After 4 hours of contact by an Al wire, the solution color had
changed from a pale yellow to clear, but the NMR signature was largely unchanged. The
62 only difference noted was a small shift in the V1 position, perhaps associated with an increase in pH. The VX species detected is likely a peroxovanadium complex resulting from solution treatment with H2O2 prior to contact with Al (23). The results of this experiment suggest that strongly reducing metals like aluminum accelerate decomposition of decavanadates by reduction, while tetrahedral vanadate species remain stable against reduction.
3.3.2 Soluble Inhibitor Release from Pigments
Synthetic hydrotalcites are anionic exchange clays that consist of alternating positively charged cation layers and negatively charged, exchangeable anion layers (20,
25). Hydrotalcites may be used as corrosion inhibiting pigments in organic coatings to deliver inhibiting anions by exchange with anions in the attacking solution. Hydrotalcites synthesized with exchangeable decavanadate have been observed to provide significant corrosion protection to aluminum alloys (8, 20). However, other work has shown decavanadate to be a poor inhibitor of corrosion on aluminum (3). Therefore, the aim of the present experiment was to determine what species are released from a decavanadate- bearing hydrotalcite. Figure 3.6 is the NMR spectrum of filtrate solution from 4.0 grams of a hydrotalcite pigment containing exchangeable decavanadate (HT-V) soaked in 40 mL of 0.1 M NaCl at pH 6.0 for 20 hours. The filtrate solution from a HT-V pigment was observed to contain both octahedrally coordinated decavanadates and tetrahedrally coordinated metavanadate and pyrovanadate species. The relative proportions of the tretrahedrally coordinated forms to octahedrally coordinated forms is greater than might
63 be expected based on the data presented in Figure 3.3. The following sections illustrate
the correlation between the presence of tetrahedral forms of vanadate in solution and inhibition.
3.3.3 Aluminum 2024-T3 Aerated Polarization in NaCl Solutions
Figures 3.7-3.10 show anodic and cathodic polarization curves measured for Al
2024-T3 in 50 mM NaCl solution. In these experiments, the solution pH and NaVO3 concentration were systematically varied from experiment to experiment. Figure 3.7 shows anodic polarization curves at pH 10 for several different NaVO3 concentrations.
In basic solutions, the corrosion potential was observed to shift to more active potentials,
corrosion current decreased in the presence of vanadate. In the case of the 0.0032 M
NaVO3 addition, an increase in pitting potential was observed. Figure 3.8 shows anodic polarization curves in solutions with a constant NaVO3 concentration (0.0032 M) and
varied pH. Compared to control experiments in chloride-only solutions at pH 3, 5, 8, and
10 (not shown), the corrosion potential became more active, corrosion current decreased,
and the pitting potential increased, but only in alkaline solutions (pH 8 and 10). Figure
3.9 shows cathodic polarization curves collected in pH 8 solutions with varied NaVO3 concentration. Here, addition of NaVO3 resulted in more active corrosion potentials, and an overall decrease in cathodic kinetics. Figure 3.10 shows cathodic polarization curves with constant 0.0032 M NaVO3 concentration with varied pH. As solution pH increased,
reduction kinetics were observed to slow compared to those measured in chloride-only
solutions at pH 3, 5, 8, and 10 (not shown). The observation that anodic and cathodic
64 inhibition were more evident in alkaline solutions than in acidic ones is consistent with
the idea that inhibition is associated with the presence of tetrahedrally coordinated
metavanadate and pyrovanadate species.
The effects of pH and vanadate concentration on the potentiodynamic polarization
response and corrosion inhibition are summarized in Figures 3.11-3.14. Corrosion
potential and current were obtained from polarization curves by extrapolation carried out
with the aid of Corrview© software. Figure 3.11 shows changes in corrosion potential as
a function of pH and NaVO3 concentration. Each data point represents an average of
eight experiments. The figure shows that NaVO3 additions have no effect at pH 3 and 5,
but lead to a considerable decrease in corrosion potential at pH 8 and 10. Figure 3.12
shows changes in corrosion current density as a function of pH with NaVO3 addition.
Approximately an order of magnitude reduction in corrosion current was observed at pH
8, and 10 in solutions with NaVO3. At pH 5, high concentrations of NaVO3 resulted in a
modest reduction of corrosion current with greater inhibition occurring at lower vanadate
concentrations. Figure 3.13 shows changes in pitting potential. In the 0.32 M NaVO3 solutions, breakdown was not observed at the point of scan reversal and was not included in the plot. Generally, the pitting potential was observed to increase in alkaline solutions containing NaVO3. Figure 3.14 shows a comparison of the cathodic current density at -
1.25 VSCE in NaVO3 solutions as a function of pH. The current density at -1.25 VSCE was used as a point of comparison between different experiments because this is a point below the lowest corrosion potential observed and allows direct comparison of reduction kinetics from all the collected data. The data shows an order of magnitude or greater
65 decrease in current density at -1.25 VSCE for alkaline solutions containing NaVO3 compared to those without. Around pH 5, the more dilute vanadate solutions slow cathodic kinetics more than concentrated ones.
3.3.4 Aluminum 2024-T3 Deaerated Polarization in NaCl Solution
Figures 3.15-3.18 show variation in polarization characteristics as a function of pH and NaVO3 concentration in deaerated 50 mM NaCl solutions. All experiments were
run in duplicate at a minimum. Figure 3.15 shows corrosion potential as a function of pH
in deaerated solutions with and without NaVO3. NaVO3 ennobles the corrosion potential
in acidic deaerated solutions, but has little effect in alkaline solutions. Figure 3.16 shows
the variation in corrosion current density. NaVO3 in acidic deaerated solutions was
observed to increase corrosion current; a decrease in corrosion current was observed in
alkaline solutions, particularly at pH 10. Figure 3.17 shows the variation in pitting
potential. The data is similar to the data collected in aerated solutions showing an
increase in pitting potential in alkaline NaVO3 solutions. Suppression of pitting by
vanadates in alkaline solutions appears to be relatively independent of the extent of
solution aeration. Figure 3.18 shows current density at -1.25 VSCE. There is no
significant reduction in the cathodic kinetics associated with the presence of NaVO3 except at pH 10. This suggests that cathodic inhibition by vanadate is largely through the suppression of oxygen reduction with little effect on hydrogen evolution.
66 3.3.5 Corrosion Morphology of Aluminum in Vanadate Solution
Figure 3.19 shows images and EDS chemical maps collected in the vicinity of
constituent particle clusters in polished Al 2024-T3 surfaces after immersion in aerated pH 10 50 mM NaCl with and without NaVO3. A description of constituent particles in
aluminum alloys and Al2CuMg dissolution can be found in the literature (17, 26, 27).
There was no evidence of Mg at particle locations in chemical maps on samples exposed
in NaVO3-free solutions at any pH tested (3, 5, 8, and 10). This indicates selective
dissolution of Mg from Al2CuMg particles. In contrast, in alkaline solutions containing
NaVO3, chemical maps showed both Mg and Cu present at particle locations. Even in
mildly acid solutions with NaVO3 small amounts of Mg were detected. These findings indicate that vanadate additions suppress dealloying of Cu and Mg-bearing particles.
Figure 3.20 is a collection of chemical maps obtained from Al 2024-T3 exposed
to aerated 50 mM NaCl showing different manifestations of vanadium deposition on the
surface of Al 2024-T3 at different pH and NaVO3 concentration. Typically, in acidic
solutions, where polarization results indicated no significant inhibition, high
concentrations of vanadium were observed in areas of localized corrosion, such as near
pits and filiform tracks. There is some evidence that the vanadium is only present near
pits which that copper. Vanadium was not observed near Fe-Mn or Fe-Mn-Cu
intermetallic particles. Despite the fact that mildly alkaline solutions provided the
strongest inhibition according to polarization results, no vanadium was detected near areas of attack or on specific intermetallic particles. In strongly alkaline solutions and those with high concentrations of NaVO3, in which polarization scans also indicated
corrosion inhibition, no significant vanadium segregation was observed in areas of attack;
67 however, patches or a dusting of vanadium was observed over large portions of the
surface. This dusting effect may be the result of rinsing the samples with ethyl alcohol
after exposure; vanadate complexes with ethyl alcohol (28). Further, an association
between Al2CuMg particles and vanadium was sometimes observed in strongly alkaline
solutions with NaVO3. In summary, the solutions that demonstrated the strongest
inhibition typically resulting in post-exposure surfaces characterized by intact Al2CuMg particles and vanadium deposition that was usually below EDS detection limits.
3.4 Discussion
3.4.1 Speciation versus Corrosion Inhibition
The formation of decavanadate in solution results in a distinct orange colored solution. The formation of metavanadate and pyrovanadate from an orange decavanadate solution is not immediately accompanied by a visually detectable color change. Given enough time and alkali dose, the solution will turn yellow or clear with an appropriate
NaVO3 concentration, and tetrahedral vanadates will dominate the vanadate solution
chemistry.
In these experiments, the best overall inhibition was observed in alkaline 0.0032
M NaVO3 solutions. However, good inhibition was observed in alkaline solutions with
0.32 M NaVO3 and even reasonable inhibition was seen in 0.0032 M NaVO3 solutions at
pH 5. The equilibrium predominance diagram in Figure 3.1 gives a general idea of the
types of species present in a solution of given pH and concentration, however, a number
of other species are also likely to be present. Interpretation of experimental results and
conclusions about speciation and resultant inhibition must be made with that fact in mind.
68 Table 1 is a summary of possible species that might be found at different combinations of pH and concentration at which experiments were performed; species are listed in order of abundance (13). A comparison of Table 1 with experimental results shows that inhibition is strongly correlated with the presence of tetrahedrally coordinated species in solution.
This includes the metavanadates and pyrovanadates, with the strongest inhibition possibly correlated to single tetrahedrally coordinated vanadate as previously suggested
(3).
Polarization data collected from 2024-T3 exposed to a 0.32 M NaVO3 solution at pH 5 and a 0.0032 M NaVO3 at pH 3 show no evidence of inhibition. Octahedrally coordinated decavanadates prevail under these conditions and do not appear to directly contribute inhibition. In fact, the ennobled corrosion potentials and increased corrosion current densities in deaerated acidic conditions (Figs. 3.15 and 3.16) suggest that decavanadates are oxidizing agents (5). Reduced forms of vanadium noted in the presence of pure Al may arise from reduction of decavanadates.
It is interesting to note that at pH 5 and 8, the polarization response characteristics suggest that 0.0032 M NaVO3 additions are more inhibiting than 0.32 M additions
(Figures 3.12 and 3.14). This is consistent with inhibition dominated by tetrahedrally coordinated forms of vanadate. Figure 3.1 shows that at pH 5, tetrahedral vanadates predominate at lower concentrations and octahedral forms predominate at higher concentrations. The differences in the extent of inhibition in pH 5 solutions at the two concentration ranges used in this study clearly show that there can be an inverse relationship between vanadate concentration and effectiveness of inhibition in mildly
69 acidic, neutral and perhaps mildly alkaline solutions. It is also of interest that vanadates operate as cathodic inhibitors over the widest pH range when present in dilute
concentrations.
Under strongly alkaline conditions where vanadate is speciated almost exclusively
in tetrahedrally coordinated forms, vanadate inhibition exhibits a more regular
dependence on NaVO3 concentration. Anodic and cathodic inhibition are demonstrated.
4- Under these conditions, a combination of the following species is expected: V4O12 ,
3- 2- 4- 3- - V3O9 , VO3(OH) , V2O7 , V2O6(OH) , and VO2(OH)2 . Among these, the strongest
2- inhibition was observed in the presence of tetrahedral pyrovanadate VO3(OH) (3).
However, this does not exclude other metavanadate and pyrovanadate species from
providing inhibition either directly or indirectly through respeciation into V1. For instance, strong inhibition was observed in pH 8 0.32 M NaVO3 solutions which likely
2- 4- 3- contain relatively small concentrations of VO3(OH) relative to V4O12 and V3O9 . It is
2- not clear whether sufficient concentrations of VO3(OH) initially exist in solution to
4- 3- account for observed inhibition or if V4O12 and V3O9 directly inhibit or rapidly
speciate into inhibiting species.
3.4.2 Inhibition and Oxygen Dependence
Vanadates inhibit both anodic and cathodic reactions. Anodic inhibition, as
assessed by the effect on the pitting potential, seems to be independent of oxygen. In this
regard, vanadates are similar to chromates in that both may act as anodic inhibitors
70 regardless of solution oxygen content (29). This is an advantage over other anodic
inhibitors such as phosphates, molybdates, and silicates, which function best in
oxygenated environments (29).
Tetrahedrally coordinated vanadates are inhibitors of oxygen reduction on Al
2024-T3. The trends in oxygen reduction closely mirror those that favor speciation of vanadates in their tetrahedral form. Inhibition of oxygen reduction shifts the corrosion potential in the negative direction. Because tetrahedral vanadates simultaneously increase the pitting potential, the tendency for localized corrosion under free corrosion conditions is significantly decreased.
It should be noted that apparent acceleration of cathodic kinetics in both aerated and deaerated environment occurs under conditions where octahedrally coordinated vanadates are expected to be present. In these cases, increased cathodic kinetics is associated with reduction of decavanadates (5).
3.4.3 Action of Vanadates on Al Alloy Surfaces
Copper-containing intermetallic particles have been reported to play a central role in localized corrosion of Al 2024-T3, essentially acting as local cathodes leading to
localized cathodic corrosion (30). In particular, Al-Cu-Mg intermetallic particles have
been reported to account for a large percentage of all particles in number and area
fraction (26). Al2CuMg, S-phase, under free corrosion conditions is initially active to the
surrounding matrix due to constituent magnesium (27). Strong anodic polarization of the
intermetallic causes dealloying and magnesium dissolution, the remaining copper will
function as a strong local cathode which supports oxygen reduction (27). In light of
71 observed anodic and cathodic inhibition in alkaline solutions, it is interesting that chemical maps show Al-Cu-Mg intermetallic particles remain relatively intact after exposure experiments to pH 8 and 10 NaVO3 solutions. In contrast, chemical maps in
strongly acidic solutions and those with no NaVO3 showed no evidence of magnesium. It is possible that vanadates in alkaline solutions play a role in preventing the rapid dissolution of magnesium from Al-Cu-Mg particles. As a result, these particles may not become strong cathodes that contribute to localized corrosion damage accumulation. A similar observation was made by Iannuzzi, et al., where metavanadate solutions were observed to suppress S-phase dissolution (4, 31).
The chemical maps collected in pH 8 NaVO3 solutions are possible evidence of
localized film formation over the matrix (Fig. 3.11). There is no detectable segregation
of vanadium to any intermetallic particles although the particles appear to be intact.
Chemical mapping lacks the sensitivity to detect very thin layers and deposits and a
complete interpretation of vanadate interaction with corroding surfaces is needed to fully
describe these interactions.
3.4.4 Vanadate Speciation and Vanadates in Hydrotalcite Pigments
NMR solution experiments show that vanadate solutions are very dynamic, with
changes in pH resulting in rapid changes in speciation. This is true even of the
dissociation of decavanadates to metavanadates. Solutions that are dominated by
tetrahedrally coordinated vanadates will speciate into octahedrally coordinated vanadates
72 and reach equilibrium quickly after an acidic pH adjustment. Octahedral vanadate
solutions will begin to speciate into tetrahedral vanadates quickly on alkaline pH
adjustment, but will not reach equilibrium for long periods of time.
This latter point is significant in understanding the behavior of corrosion
inhibiting hydrotalcite pigments. Previous work has shown hydrotalcites containing vanadates inhibit corrosion of Al 2024-T3 when dispersed into an organic resin and
applied as a coating (8, 20). However, hydrotalcite synthesis can only be carried out with
octahedrally coordinated decavanadate. Speciation of decavanadate to tetrahedrally
coordinated forms is essential for corrosion inhibition to be imparted. Indeed, the
presence of tetrahedrally coordinated forms of vanadate in solutions in contact with
hydrotalcite pigments shows that inhibiting forms of vanadate develop from this pigment.
In coatings on metallic substrates, the generation of tetrahedral forms of vanadate is
expected to be further stimulated by local increases in pH associated with sites supporting
oxygen reduction, and the low vanadate concentrations in solution.
3.5 Conclusions
• Inhibition of Al 2024-T3 in NaCl solutions by vanadates is associated with
tetrahedrally coordinated forms of vanadate. Octahedrally coordinated vanadates
do not appear to provide inhibition and may accelerate corrosion under deaerated
conditions.
• On acidic pH adjustment of alkaline solutions, speciation and equilibration of
octahedral forms of vanadate from tetrahedral forms is fast. On alkaline pH
adjustment of acid solutions, speciation of tetrahedral forms of vanadate from
73 octahedral forms begins within minutes of pH adjustment, however, equilibration
is slow.
• NaVO3 solutions inhibit cathodic kinetics of Al 2024-T3 corrosion in alkaline
aerated NaCl solutions. NaVO3 solutions were not observed to inhibit cathodic
kinetics in deaerated solutions. The main inhibiting action of vanadates on
cathodic reactions is suppression of oxygen reduction.
• An increase in pitting potential was observed in both aerated and deaerated NaCl
solutions where tetrahedrally coordinated vanadates were present. Anodic
inhibition from tetrahedrally coordinated vanadates in alkaline solutions is
independent of solution aeration.
• Chemical maps showed intermetallic particles containing magnesium to be
largely intact after exposure to NaCl solution containing tetrahedrally coordinated
vanadates.
• A filtrate solution produced by contact with a decavanadate-bearing hydrotalcite
inhibitor pigment was shown to contain both metavanadate and decavanadate
species. Vanadates in solution will speciate quickly in response to the prevailing
pH.
• Vanadate solution color change is not always a reliable indicator of species
present in vanadate solutions.
74 FIGURES
Figure 3.1: Equilibrium predominance diagram for VV-OH- species as a function of concentration and pH. Adapted from The Hydrolysis of Cations and Heteropoly and Isopoly Oxometalates. The dotted line indicates delineation between octahedral coordinated and tetrahedral coordinated vanadate species. Approximate pH and concentration of test solutions are indicated with an “x”.
75
Figure 3.2: Change in pH as a function of time of a pH 8.76 100 mM NaVO3 solution initially acidified to pH 4.12 with HNO3 and then adjusted to pH 7.58 with NaOH.
76
Figure 3.3: NMR spectrum as a function of time of a 100 mM NaVO3 solution initially acidified to pH 4.12 with HNO3 and then adjusted to pH 7.58 with NaOH.
77
Figure 3.4: NaVO3 solution after serial additions of 10 N NaOH: A) initial orange 100 mM NaVO3 solution at pH 6.08, B) yellow pH 6.63 solution 27 hours after the addition of a second drop of 10 N NaOH immediately prior to a third NaOH addition (spectrum collected at 52 hrs total), C) yellow pH 9.15 solution, NMR sample taken and pH measured approximately 30 minutes after the addition of a third drop of 10 N NaOH (spectrum collected at 52 hrs total).
78
Figure 3.5: NMR spectra of 100 mM NaVO3 solutions in contact with a pure Al wire A) initial orange decavanadate-metavanadate pH 6.08 solution prior to exposure, B) dark emerald green solution with pH near 6 after 52 hours of contact, C) initial pale yellow decavanadate-free pH 8.10 solution prior to contact. D) Clear decavanadate-free solution after 52 hours of contact.
79
Figure 3.6: NMR spectrum of filtrate from 4.0 grams of an Al-Zn-V hydrotalcite pigment soaked in 40 mL of 0.1 M NaCl for 20 hours, solution pH 6.
80
Figure 3.7: Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant pH 10 adjusted solution and varied NaVO3 concentration.
81
Figure 3.8: Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant 0.0032 M NaVO3 solution and varied pH.
82
Figure 3.9: Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant pH 8 solution and varied NaVO3 concentration.
83
Figure 3.10: Anodic and cathodic polarization curves for Al 2024-T3 in 50 mM NaCl solution with constant 0.0032 M NaVO3 concentration and varied pH.
84
Figure 3.11: Condensed data from aerated polarization experiments, corrosion potential as a function of pH.
85
Figure 3.12: Condensed data from aerated polarization experiments, corrosion current as a function of pH.
86
Figure 3.13: Condensed data from aerated polarization experiments, pitting potential as a function of pH.
87
Figure 3.14: Condensed data from aerated polarization experiments, cathodic current at - 1.25VSCE as a function of pH.
88
Figure 3.15: Condensed data from deaerated polarization experiments, A) corrosion potential as a function of pH.
89
Figure 3.16: Condensed data from deaerated polarization experiments, corrosion current as a function of pH.
90
Figure 3.17: Condensed data from deaerated polarization experiments pitting potential as a function of pH.
91
Figure 3.18: Condensed data from deaerated polarization experiments, cathodic current at -1.25VSCE as a function of pH.
92
Figure 3.19: Sample chemical maps collected in aerated 50 mM NaCl showing observed suppression of S-phase dissolution in alkaline NaVO3 solutions compared to vanadate- free solutions. A) pH 10 0.0 M NaVO3 B) pH 10 0.0032 M NaVO3.
93
Figure 3.20: Sample chemical maps collected in aerated 50 mM NaCl showing different manifestations of vanadium on the surface of Al 2024-T3. A) pH 5 0.32 M NaVO3 B) pH 8 0.0032 M NaVO3 C) pH 8 0.32 M NaVO3.
94 TABLES
pH 3 pH 5 pH 8 pH 10 4- 5- 4- 4- 0.32 M NaVO3 V10O26(OH)2 V10O27(OH) V4O12 V2O7 5- 3- 2- V10O27(OH) V10O28 V3O9 VO3(OH) 4- 2- 3- V10O26(OH)2 VO3(OH) V2O6(OH) 3- V2O6(OH) 4- 5- 3- 2- 0.0032 M NaVO3 V10O26(OH)2 V10O27(OH) V3O9 VO3(OH) 5- 3- 2- 4- V10O27(OH) V3O9 VO3(OH) V2O7 + - - VO2 VO2(OH)2 VO2(OH)2 6- 4- VO(OH)3 V10O28 V4O12 - 4- 3- VO2(OH)2 V4O12 V2O6(OH) 4- V10O26(OH)2
Table 3.1: List of probable vanadium species in solutions of varied NaVO3 concentration and pH. The dashed line delineates vanadate solutions that demonstrated inhibition, on the right, from solutions in which inhibition was not observed, on the left. Species listed in red (grey in black and white copies) are likely present in relatively small concentrations compared to species listed in black.
95 REFERENCES
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97
CHAPTER 4
ELECTROCHEMICAL EVALUATION OF CONSTITUENT
INTERMETALLICS IN ALUMINUM ALLOY 2024-T3 EXPOSED TO
AQUEOUS VANADATE INHIBITORS
4.1 Introduction
Aluminum 2024-T3 is high-strength age-hardened aluminum alloy commonly
used in the aerospace industry. Al 2024 contains by weight percent between 3.8-4.9 Cu,
1.2-1.8 Mg, 0.3-0.9 Mn, and small quantities of Si, Fe, Zn, Cr, and Ti (1). Alloy
additions result in both superior mechanical properties and a heterogeneous
microstructure which renders the alloy susceptible to localized corrosion (1).
Appreciable quantities of copper, magnesium and manganese, added as strengtheners,
remain in solid solution. However, through heat treatments and natural aging, a
dispersion of fine Cu and Mg particles and insoluble intermetallic precipitates form
within the matrix phase. The main constituent particles in Al 2024-T3 include Al2CuMg,
Al7Cu2Fe, Al2Cu, and Al20Cu2Mn3. The effect of intermetallic particles on corrosion of
aluminum alloys has been studied widely (2-11). For Al 2024-T3 it has been found that
98 intermetallics containing Cu are typically noble or will become noble to the surrounding aluminum matrix during exposure to many electrolytes and these particles are capable of supporting rapid cathodic kinetics (4, 5, 8, 9, 12). Such cathodic particles drive corrosion in the surrounding matrix, leading to pitting and trenching attack morphologies (10, 11,
13). Al2CuMg (S phase) intermetallic particles are of particular interest because
Al2CuMg is one of the most abundant intermetallic particles found in Al 2024 and in large part has been found to be responsible for susceptibility of Al 2024 to localized corrosion (4, 5). The corrosion of Al2CuMg is complex; under free corrosion conditions
the intermetallic is initially anodically polarized by the matrix, leading to selective
dissolution of Mg from the intermetallic and non-faradaic liberation of Cu, which can
then be oxidized to form ions that can be reduced on the surrounding matrix (4, 5). Often
what remains of the particle is an enriched Cu remnant, which acts as a local cathode,
supporting rapid oxygen reduction and corrosion in the surrounding matrix (4, 5).
Prevention of Mg dissolution from Al2CuMg, and as a result, the subsequent formation of
local Cu cathodes capable of supporting rapid oxygen reduction could be an effective
way to increase the resistance of Al 2024 to localized corrosion (14).
Historically, chromate-based pigments and coatings have been used successfully
prevent to corrosion of aluminum alloys (15). However, due to environmental and
carcinogenic risks associated with chromate use, more “green” alternative inhibitors and
coatings have recently received attention. In particular, soluble vanadates, vanadate-
based coatings, and inhibitor pigments have been observed to inhibit the corrosion of
aluminum alloys and have shown promise as chromate replacements (16-23). However,
unlike chromates, aqueous vanadates have a relatively complex aqueous chemistry,
99 dependent on pH, concentration and ionic strength, which convolute a straight forward understanding of inhibition (24-26). In a simplified description of aqueous vanadate speciation, tetrahedrally coordinated species, metavanadates and pyrovanadates, predominate in alkaline solutions, octahedral coordinated species, decavanadates, predominate in acid solutions, and single tetrahedral species exist over a wide pH range at low concentrations. Previous work has shown that the extent of inhibition depends strongly on vanadate speciation, with the greatest inhibition from tetrahedrally coordinated species, which predominant in alkaline solutions (18, 20-22, 27).
Tetrahedrally coordinated vanadates have been shown to act primarily through decreased oxygen reduction, however, vanadates have also been observed to be modest anodic inhibitors independent of aeration (18, 22). Decavanadate ions, which are combinations of ten octahedrally coordinated vanadate units, predominate in acidic solutions of appropriate vanadium concentration and have been shown to be poor inhibitors of oxygen reduction (18, 20-22, 24, 25, 27). There is evidence that decavanadate increases the cathodic kinetics in acidic NaCl solutions (18). However, small increases in pH, as found near sites supporting oxygen reduction, can trigger the decomposition of non-inhibiting octahedrally coordinated species into inhibiting tetrahedral species, which helps explain corrosion protection observed from pigments containing decavanadate (16, 18, 28, 29).
Evidence exists that inhibiting tetrahedrally coordinated vanadates suppress the dissolution of Al2CuMg intermetallics (18, 21, 27). Ralston et al. noted suppressed Mg dissolution from Cu-Mg particles exposed to alkaline and mildly acidic 50 mM NaCl solutions with NaVO3 compared to particles exposed to NaVO3-free solutions (18).
Iannuzzi and Frankel used in-situ atomic force microscopy scratching to observe
100 additions as small as 0.1 mM metavanadate to 0.5 M NaCl suppressed attack of Al2CuMg particles, while, corrosion in the surrounding matrix was still observed (27). Iannuzzi further found that 5 mM metavanadate prevented transient Al2CuMg dissolution,
resulting in increased corrosion resistance at OCP (20, 21). The mechanism of
suppression of transient dissolution was not clear, but it was speculated that
monovanadates on the matrix prevent or displace Cl- adsorption on the surface, which
hinders subsequent oxide film breakdown (20). Generally, it is not certain whether
vanadates slow corrosion of Al 2024 through acting exclusively on Al2CuMg or if suppressed Al2CuMg dissolution is a consequence of overall corrosion inhibition.
Although previous work has established that solutions containing predominately
tetrahedrally coordinated vanadates prevent or slow Mg dissolution from Al2CuMg, and
in turn the formation of Cu-rich cathodes capable of supporting rapid oxygen reduction,
the precise relationship between tetrahedral vanadates, Al2CuMg, and the matrix is not
currently understood.
Vanadate is a known buffer, and observed inhibition must be rationalized in the
context of the effects that buffers have on corrosion. The presence of low concentrations
of buffer can have pronounced effects on the corrosion of Al-Cu alloys (30, 31). In
unbuffered systems, oxygen reduction results in an increase in alkalinity near cathodic
sites, which dissolves the surrounding Al matrix, leading to grooving and trenching
around intermetallics (30, 31). Dissolution of the Al matrix leads to self-corrosion and
cathodic current is consumed through driving widespread cathodic corrosion across the
matrix. In contrast, buffers prevent alkalinization and the resultant matrix dissolution,
however, larger net cathodic currents are achieved and focused exclusively on driving
101 localized corrosion (30, 31). The production of reducible H+ at sites of localized corrosion contributes to the larger net cathodic currents. This leads to deep and discrete pits, but, comparatively less mass loss than observed in unbuffered systems (30).
Owing to the small size of constituent intermetallic particles present in Al 2024, direct electrochemical testing of different intermetallics in the matrix is not feasible.
However, previous work on aluminum alloys using a microcapillary electrode and synthesized “bulk” intermetallics demonstrates that intermetallic-specific electrochemical data can be obtained (2, 3, 32). The microcell is a modified standard three electrode setup that uses a thin glass silicon coated capillary connected to an electrolyte reservoir, containing both a reference electrode and counter electrode to contact and allow electrochemical experiments on micrometer scale diameter working electrodes (33, 34).
General details of the microcell setup and a specific description of the microcell used for this work can be found in the literature (3, 33, 34). By choosing bulk intermetallics that are representative of constituents in Al 2024, pure Al and Cu, and an Al 4% Cu solid solution used as a matrix analog, electrochemical characteristics of specific intermetallic phases can be catalogued and used to rationalize observed behavior of the bulk alloy.
The objective of this work is to determine how inhibiting vanadates interact with the matrix and constituent particles of Al 2024-T3. In addition, this work aims to develop a deeper understanding of the suppression of Mg dissolution from Al2CuMg intermetallics in inhibiting vanadate solutions.
102 4.2 Experimental Procedures
4.2.1 Solution Preparation
Solutions for all experiments were prepared using reagent grade chemicals. The
NaVO3 for solution preparation was purchased from Fluka Chemika with an assay ≥98%.
Cathodic polarization experiments were conducted in 0.5 M NaCl solutions adjusted to pH 5.1 using HCl with 0.25 M and 0.0025 M NaVO3 to show the effect that tetrahedral
vanadates have on inhibition compared to octahedral vanadates. However, most
experiments were conducted in alkaline 0.5 M NaCl solutions with and without 10 mM
NaVO3 to characterize the inhibitive effect tetrahedrally coordinated vanadates have on
different constituent intermetallics in Al 2024-T3. The initial as-dissolved 0.5 M NaCl +
10 mM NaVO3 solution was yellow and had a pH of 6.37. As previously mentioned,
vanadates have been shown to provide strongest inhibition when coordinated
tetrahedrally, which occurs in alkaline solutions. As a result, the pH of the master test
solution was adjusted with drop wise additions of 10 N NaOH until the pH reached 9.18.
Before experimentation, the test solution was allowed to equilibrate for more than two weeks, during which a few additional drops of NaOH were used to maintain the pH above 9. Once the pH was stable, nuclear magnetic resonance (NMR) was used to characterize the solution. To help ensure the vanadate species in solution were not evolving with time, the solution pH was monitored daily over the course of microcell experimentation and NMR spectra were collected prior to the first experiment and after completion of the last experiment. The same 0.5 M NaCl + 10 mM NaVO3 solution was
used for subsequent experiments after completion of the microcell polarization work.
Since the solution appeared stable over the duration of microcell experiments, pH was
103 used as a sufficient measure of solution and species stability. Although the NaVO3 solution remained stable, the pH of NaCl solutions adjusted to approximately pH 9.2 would decrease with time. This is likely the result of H2CO3 formation from dissolved
- atmospheric CO2 and the subsequent proton formation from equilibria involving HCO3
2- and CO3 , which both have increased solubility in alkaline solutions (35). As a result, care was taken to monitor and measure the pH immediately prior and during each testing session. Unless specifically stated, the pH of NaCl-only solutions was between 9.05 and
9.20. Additionally, the reservoir, capillary, and tubing of the microcell were frequently flushed with fresh solution throughout experimentation.
4.2.2 Nuclear Magnetic Resonance (NMR)
Vanadates have a complex aqueous speciation depending on both concentration and pH and, as a result, small changes in pH can have dramatic effects on the type and concentration of specific species in solution. NMR was used to characterize solutions used for cathodic polarization experiments on Al 2024-T3 sheet in pH 5.1 NaVO3 solutions and microcell polarization experiments in alkaline 10 mM NaVO3 solution which were expected to take a number of weeks to complete and for which the possibility of solution evolution with time was a concern. NMR spectra were collected immediately prior to microcell work in vanadate solutions and 10 days later after completion of experimentation. NMR spectra for cathodic polarization experiments in pH 5.1 solutions were collected immediately after the pH of the vanadate solutions were adjusted. A
Bruker DPX 400 MHz superconducting magnet was used to collect high resolution 51V
(105.2 MHz) NMR spectra. An indirect detection probe was used with a 90o pulse
104 duration of 10.38 μs. Spectra were collected using 8192 transients, a spectral window of
73,529 Hz, a 0.051 s acquisition time, and a 0.20 s relaxation delay. Each spectrum had
the subsequent process parameters applied: 10.0 Hz line broadening, zero filling (25 K
51 points), and baseline correction. A solution consisting of 20% v/v VOCl3 in C6D6 (δ V
= 0ppm) was used as an external standard to reference the 51V chemical shifts. Peaks
were identified by comparison to literature (26, 29, 36).
4.2.3 Potentiodynamic Polarization Using the Microcapillary Electrode
To characterize the inhibitive effects of tetrahedrally coordinated vanadates,
anodic and cathodic polarization curves were collected on representative bulk versions of
intermetallics found in Al 2024 in alkaline 10 mM NaVO3 and NaVO3-free 0.5 M NaCl
solutions. Samples used for potentiodynamic polarization experiments using the
microcell were sourced from previous work and commercial suppliers. The 99.999% Al
and 99.9% Cu samples were obtained from Alfa Aesar©. The intermetallic samples used
in this study were prepared and studied previously: Al7Cu2Fe and Al20Cu2Mn (32), Al
4%Cu (3), Al2CuMg (37), and Al2Cu (38). Samples were ground in 200 proof ethyl alcohol to 1 μm using SiC papers, followed by polishing using 6 μm and 1 μm diamond pastes. All electrochemical experiments presented in this paper were made using an
Autolab PGSTAT 100 potentiostat in conjunction with General Purpose Electrochemical
Systems© (GPES) data acquisition software. Both anodic and cathodic polarization
experiments were preceded by 30 seconds of OCP measurement and were carried out in
aerated solutions using a 0.01 V/s scan rate. Anodic polarization curves were initiated at
-0.03 V vs. OCP and reversed at either 0.0 V vs. OCP or manually at approximately 0.05 105 V above any observed breakdown. Cathodic polarization experiments were initiated at
0.03 V vs. OCP and terminated at -2.0 V vs. SCE, although capillary tip leaking often resulted in an early termination of the experiment. The contact area used for area normalization of data was estimated from digital images taken after individual experiments.
4.2.4 Electrochemical Experiments on Bulk Al 2024-T3 Sheet
Experiments on bulk Al 2024 electrodes in alkaline 0.5 M NaCl solutions with and without 10 mM NaVO3 were used to obtain electrochemical data from the actual
alloy for comparison to microcell results. Experiments in pH 5.1 NaVO3 solutions were
used to show the inhibiting effect that tetrahedral coordinated vanadates have on cathodic
kinetics compared to octahedrally coordinated vanadates. Further, work on bulk Al 2024
was used to gain insight into Al2CuMg dissolution. Al 2024-T3 sheet was used for four
different sets of experiments; cathodic polarization, OCP measurement, anodic
polarization, and potentiostatic experiments. All samples were polished by hand in 200
proof ethyl alcohol to at least to 1 μm in a similar fashion as discussed for microcell
sample preparation except samples for OCP measurements, which were polished to ¼ μm
using diamond paste, the samples for the potentiostatic experiments, which were polished
to 1 μm using a diamond suspension and an automatic polisher rather than diamond paste
by hand, and samples used for cathodic polarization which were polished to 1200 grit
under ethyl alcohol. Experiments on bulk Al 2024-T3 sheet were carried out using a
standard three electrode setup, which included a SCE reference, a platinum counter
electrode mesh, and 1 cm2 exposed working electrode. Cathodic polarization 106 experiments in actively aerated pH 5.1 NaVO3 solutions were preceded by a 30 minute measurement of OCP. The scan was initiated at 0.03 V vs. OCP and a scan rate of 0.5 mV/s was used. OCP was measured for 4 hours in actively aerated and deaerated 0.5 M
NaCl solutions at approximately pH 9.2 with and without 10 mM NaVO3 to determine
the effect of tetrahedrally coordinated vanadates on OCP with time. For deaerated OCP
measurements, solutions were deaerted for 1 hour before the electrolyte came in contact
with the sample. Anodic polarization curves on Al 2024-T3 sheet in 0.5 M NaCl solution
with 10 mM NaVO3 solution at approximately pH 9.17 were used to make comparisons
between microcell data and data collected from Al 2024-T3 sheet. Anodic polarization
experiments were preceded by a 30 minute measurement of OCP. The scan was initiated
at -0.03 V vs. OCP, a scan rate of 0.5 mV/s was used with scan reversal at -0.25 VSCE.
Potentiostatic hold experiments were conducted in 0.5 M NaCl solutions between pH 9.1 and 9.23 with and without 10 mM NaVO3. These experiments were used to determine if
tetrahedral vanadates have an effect on the repassivation of the surface once activated and
to show the suppression of Mg dissolution from Al2CuMg intermetallics. Each 100 mL
test solution was deaerated for 30 minutes prior to experimentation and sample exposure.
The samples were held at a conditioning potential of 1 VSCE for 1 s and then held at a
specific potential for the next 120 s; potential holds at -1.2, -0.9, -0.8, -0.7 -0.6, and -0.5
VSCE were used.
107 4.3 Results
4.3.1 Inhibition from Tetrahedral Vanadate Species vs. Octahedral Species
The effect that tetrahedrally coordinated vanadates have on cathodic kinetics
compared to octahedrally coordinated species can be observed through experiments in
mildly acidic NaVO3 solutions. Figure 4.1 shows the NMR spectra from two different
pH 5.1 0.5 M NaCl solutions with a) 0.0025 M NaVO3 and b) 0.25 M NaVO3. The subscript of the peak labels in the Figure describes the number of vanadium atoms in each oligomer. For example, V1 indicates single tetrahedrally coordinated vanadium
3- 2- - 4- (VO4 , VO3(OH) , VO2(OH)2 , VO(OH)3) while V2 indicates dimeric vanadate (V2O7 ,
3- 4- 6- 7- 5- V2O6(OH) ), V4 is tetrameric (V4O12 , V4O13 ), V5 is pentameric (V5O16 , V5O15 ) and
6- 4- 5- V10 represents decameric vanadate species (V10O28 , V10O26(OH)2 , V10O27(OH) ) (25,
26, 29, 36). V1, V2, V4, and V5 are tetrahedrally coordinated species and V10 is octahedrally coordinated. Also, it should be noted that the scales of the two spectra in
Figure 4.1 are not the same and have been adjusted so peaks from both concentrations can be observed in the same figure. The 0.0025 M NaVO3 solution has a greater
proportion of tetrahedral species relative to octahedral species compared to the 0.25 M
NaVO3 solution, which contains significantly more octahedral species than tetrahedral
species.
Figure 4.2 shows cathodic polarization curves on Al 2024-T3 in aerated pH 5.1
0.5 M NaCl with 0.25 M NaVO3, 0.0025 M NaVO3, and without NaVO3. These
experiments show the inverse relationship between NaVO3 concentration and inhibition
of cathodic kinetics at pH 5.1, where the dilute solutions containing relatively more
108 tetrahedrally coordinated vanadate to octahedral vanadates have a larger reduction in
cathodic kinetics than more concentrated solutions with relatively more octahedrally coordinated decavanadate.
4.3.2 Tetrahedral Vanadate Species in Alkaline Electrolytes
Small changes in solution pH can have a large effect on vanadate speciation.
Concerns that the vanadate test solutions would change over the course of microcell
experimentation, with consequences for inhibitor behavior were addressed using NMR.
Figure 4.3 shows the spectra from two samples of pH 9.17 0.5 M NaCl + 10 mM NaVO3 test solution taken immediately before microcell experimentation and two weeks later after the conclusion of experimentation in vanadate solutions. The solution did remain stable over the course of experimentation and was found to contain a number of different tetrahedral vanadate species as expected from previous work on vanadate inhibition (18,
22). It should be noted that the assignment of V4 and V5 are not certain as two different
standards available in the literature leave room for speculative interpretation (26, 36).
However, definitive assignment of these species is not critical for this work. The solution used for this work contained tetrahedrally coordinated species and predominately V1; no
octahedrally coordinated vanadates were detected.
4.3.3 Polarization of Intermetallics in Tetrahedral Vanadate Solutions
Figures 4.4-4.10 show sample anodic polarization curves for pure Al, pure Cu, Al
4%Cu, Al2Cu, Al2CuMg, Al7Cu2Fe, and Al20Cu2Mn3, respectivley. The objective of
these experiments was to determine the effect of tetrahedral vanadates on the anodic
109 behavior of different constituent intermetallic particles in Al 2024. Detailed anodic
polarization results are discussed below in conjunction with cathodic polarization results
for each intermetallic.
Figures 4.11-4.17 are a collection of cumulative distribution plots showing Ecorr,
Epit, and Erp for various intermetallic compounds. Ecorr is the corrosion potential, Epit is the pitting potential and Erp is the reversible potential on the reverse scan, which was
defined as the potential at the smallest observed current on the reverse scan. Cumulative
probability plots are of value because there can be significant variation in the measured
characteristic potentials of intermetallic compounds. Knowing this variation is important
in the interpretation of corrosion processes. The statistical variation, which would be lost
by simply taking averages, can be shown fully in cumulative distribution plots. The
corrosion rate was estimated from both anodic and cathodic polarization curves (to be
presented later) by extrapolation of linear passive regions and regions of oxygen
reduction, respectively, to the intersection with corrosion potential.
Figures 4.18-4.24 show sample cathodic polarization curves for pure Al, pure Cu,
Al 4%Cu, Al2Cu, Al2CuMg, Al7Cu2Fe, and Al20Cu2Mn3, respectively. These
experiments were used to determine the effect of vanadates on cathodic kinetics of
intermetallic particles in Al 2024.
Figures 4.25-4.31 are a collection of cumulative distribution plots showing corrosion current density (icorr), passivation current density (ipass), and current density at -
1.3 VSCE for tested intermetallics and metals. Passivation current density was defined as
the current density immediately before breakdown. The current density at -1.3 VSCE was
used as a comparative measure of cathodic reduction kinetics because the potential at this
110 point was below the most active Ecorr and this point of the curve allowed a direct comparison of reduction kinetics at potentials where oxygen reduction reactions contributed to the cathodic response. It should be noted that considerable hydrogen evolution might also be present at this potential. The following presents results for each tested material.
Pure Al- The addition of 10 mM NaVO3 to alkaline 0.5 M NaCl results in a shift of
corrosion potential to more active potentials relative to NaVO3-free solutions on pure Al.
However, NaVO3 appears to have little effect on corrosion current density. It seems the small decrease in cathodic kinetics, as seen by a decrease in current density at -1.3 VSCE in vanadate solutions, is largely offset by an increase in anodic kinetics, as seen by an increase in ipass observed in NaVO3 solutions. NaVO3 increases the breakdown potential
of pure Al and does not appear to have an effect on the repassivation potential on the
reverse scan.
Pure Cu- NaVO3 was observed to have little effect on the corrosion potential, corrosion
current density, or repassivation potential on the reverse scan of pure Cu. Pure Cu did
not demonstrate passive behavior with an observable characteristic breakdown.
However, NaVO3 was observed to have an effect on dissolution kinetics. Comparison of
current density at -1.3 VSCE shows that NaVO3 solutions slowed cathodic kinetics.
Al 4%Cu- Al 4%Cu was used as an analog material for the matrix of Al 2024-T3.
NaVO3 was observed to shift the corrosion potential to more active potentials and the
corrosion current density was observed to decrease as a result of a decrease in cathodic
kinetics. A relatively small increase in Epit was observed in NaVO3 solutions compared
111 to increases observed on other intermetallic compounds. Vanadate increased the current
density observed in the passive region of the anodic polarization scans.
Al2Cu, Al7Cu2Fe, and Al20Cu2Mn- These three intermetallics were observed to behave
similarly in NaVO3 solution. All three intermetallics showed a large shift in corrosion
potential to more active potentials in NaVO3 solutions. Al7Cu2Fe and Al2Cu showed a
decrease in corrosion current density in NaVO3 solution, while vanadate had little effect on the corrosion current density of Al20Cu2Mn3. These intermetallics showed an order of
magnitude or more decrease in current density at -1.3 VSCE indicating an overall decrease
in cathodic kinetics at this potential. Further, the breakdown potentials of all three
intermetallics was observed to increase in NaVO3, however, NaVO3 had little effect on
the current magnitude observed in the passive region of the anodic polarization curves.
Al2CuMg- A large degree of variability was observed in the anodic polarization curves of
Al2CuMg. This variability is most pronounced in the breakdown potential distribution.
Breakdown in Al2CuMg is believed to be associated with loss in protection conferred by
a Cu-enriched layer that forms on the intermetallic surface as the intermetallic dissolves
Mg and, to a lesser extent, Al. Two different reproducible behaviors were observed on
Al2CuMg in 0.5 M NaCl solutions, which are shown in the Figure 4.8. Most anodic polarization curves in NaCl-only solutions showed passive behavior up to and above -0.4
VSCE. In some cases the rapid scan rate resulted in polarization of sample intermetallics
well beyond potentials necessary for breakdown as illustrated by one of the NaCl-only
curves in Figure 4.8. In these cases, the potential of the plateau after breakdown was
used for Epit values shown in cumulative probability plots. A small number of polarization curves in NaCl-only solution were observed to have corrosion potentials
112 more noble and breakdowns more active relative to the previously discussed Al2CuMg anodic polarization curves. Generally, the addition of NaVO3 resulted in much more
reproducible behavior with less data scatter. For example, Al2CuMg in NaCl solutions
had breakdown potentials scattered from approximately -0.8 to 0.2 VSCE. NaVO3 caused the data distribution to become much tighter, ranging from approximately -0.7 to -0.6
VSCE, which on average is a shift in breakdown potentials to more active potentials.
Interestingly, the breakdown potential distribution is shifted to lower potentials when
vanadate is present in solution. This behavior is different than that observed for other
intermetallic compounds where breakdown potentials are ennobled by vanadate
additions. Additionally, on average NaVO3 caused a small decrease in the corrosion
potential and corrosion current density, however, larger passive current densities were
observed in NaVO3 solutions. NaVO3 was observed to decrease the current density at -
1.3 VSCE, an indication of suppressed cathodic kinetics
In summary, all tested materials showed passive behavior and a breakdown with
the exception of pure Cu. NaVO3 in approximately pH 9 test solutions typically shifted
the corrosion potential to more active potentials, increased the pitting potential, except for
Al2CuMg, and decreased the cathodic kinetics. A summary of averaged characteristic
potentials and current densities from polarization experiments is found in Table 1.
4.3.4 OCP and SEM Images of Al 2024 Exposed to Tetrahedral Vanadate Solutions
Figure 4.32 is a plot of the OCP of Al 2024-T3 in aerated and deaerated 0.5 M
NaCl solutions with and without 10 mM NaVO3. The solutions were not intentionally
buffered in any way and as a result the solution pH at initial sample exposure and during
113 OCP measurement varied. Typically the pH of actively aerated solutions would become more acidic during the measurement. For example, the pH was observed to drop from
9.17 to 6.72 and from 9.19 to 8.85 from the beginning of OCP measurement until completion 4 hours later for aerated solutions without NaVO3 and with NaVO3, respectively. In contrast, deaerated solutions would become more alkaline during the hour of deaeration prior to experimentation and then pH would change little during the 4 hours of OCP measurement. Similar to aerated solutions without NaVO3, deaerated
solutions without NaVO3 typically were observed to have larger changes in pH compared
to solutions with NaVO3. Vanadates can act as buffers and this may be important with
regards to the mechanism of vanadate solution inhibition. In aerated solutions without
NaVO3 the OCP was observed to initially trend towards more noble potentials while
cyclically fluctuating between approximately -0.89 VSCE and -0.64 VSCE for the first 20-
25 minutes of measurement before stabilizing to -0.62 VSCE. Although pH was not
measured during the experiment, it is suspected the pH was becoming more acidic with
2- - time as dissolved CO2 dissociated to and equilibrated with CO3 and HCO3 . In aerated
solutions, the addition of NaVO3 maintained the OCP at values below -0.8 VSCE after 4 hours and the pH remained alkaline. In contrast, in deaerated solutions the addition of
NaVO3 was observed to increase the OCP from approximately -1.100 to -0.85 VSCE after
4 hours.
Figures 4.33-4.36 are SEM images of Al 2024-T3 in aerated and deaerated 0.5 M
NaCl solutions with and without 10 mM NaVO3 corresponding to the samples used in
Figure 4.32 after 4 hours of OCP measurement. A comparison of samples in aerated solutions with and without vanadate show that in aerated solutions the presence of
114 vanadate greatly decreased attack on and around intermetallic particles, Figures 4.33 and
4.34. It should be noted that the comparison is not completely direct as the pH of the
NaVO3-free solution drifted over experimentation to below pH 7. Regardless, the effect of vanadate even if only acting as a buffer is apparent. In particular, in aerated solutions without vanadate nearly every intermetallic demonstrates attack and a large number of intermetallics appear to have undergone complete dissolution or undercutting. In addition, a number of small bright particles, possibly redistributed Cu, are observed on the surface near pits. Previous studies have detailed the release and redistribution of metallic Cu, observed around some pits, from Al2CuMg intermetallics (5, 39). Although there is some attack at the periphery of intermetallic particles in aerated solutions with vanadate, for the most part intermetallics remain intact. As seen in Figures 4.35 and 4.36, vanadates also appear to decrease the amount of attack around the periphery of intermetallic particles in deaerated solutions. Circumferential attack is seen around round intermetallics in deaerated solutions without vanadate to a larger degree than in deaerated solutions with vanadate. Although no direct chemical analysis was performed, these small round intermetallics are likely Al2CuMg or Al2Cu (4, 5, 40). In both aerated and deaerated solutions, NaVO3 appears to greatly decrease the amount of trenching attack observed around intermetallic particles.
Figure 4.37 shows duplicate anodic polarization curves of Al 2024-T3 sheet in 0.5
M NaCl with 10 mM NaVO3 at pH 9.17. The purpose of this experiment was to show a correlation between electrochemical features of the real alloy with observed behavior from the microcell. The curves show an inflection between -0.75 and -0.8 VSCE, which corresponds well to the breakdown of Al2CuMg observed in NaVO3 solution from
115 microcell results and is likely the onset of Mg dissolution. Further, breakdown is
observed at approximately -0.575 VSCE which corresponds to the breakdown of pure Al.
The unusual behavior observed at the beginning of each scan is the result of starting the
polarization immediately after the sample was exposed to deaerated solution which
allowed little time for OCP stabilization.
4.3.5 Suppressed Al2CuMg Dissolution in Tetrahedral Vanadate Solutions
Figures 4.38 and 4.39 are plots of the resultant current responses after
potentiostatic hold experiments on Al 2024-T3 sheets in 0.5 M NaCl at pH 9.17 with and
without 10 mM NaVO3. Figure 4.40 shows total charge passed from Figures 4.38 and
4.39. Polished Al 2024-T3 samples were activated for 1 s at 1 VSCE and then held at
different potentials spanning a range of expected intermetallic breakdown for 120 s to
observe intermetallic and surface repassivation once activated. For samples exposed to
the NaCl-only solutions the sample held at -0.9 VSCE repassivated after activation,
however, potential holds at more noble potentials produce a distinct transient. This is
consistent with the observed breakdown of Al2CuMg observed in Figure 4.15 for NaCl-
only solutions, although pure Al is also expected to breakdown in this range of potentials.
The transient is believed to be associated with the dissolution of Mg from Al2CuMg
intermetallic particles which have shown dealloying at similar potentials (4). The plot
shows that, regardless of potential hold, Al2CuMg dissolution is nearly complete within
30 s. Similar behavior was observed in samples exposed to vanadate solutions where,
once activated, samples held at potentials at or more positive than -0.8 VSCE exhibited
dissolution transients. However, for each curve obtained in NaVO3 solution the current
116 response had a lower magnitude than the current response at corresponding potentials
obtained in NaCl-only solutions. A comparison of total charge passed in Figure 4.40
shows that at all potentials except -0.6 VSCE vanadate reduced total charge passed.
4.4 Discussion
4.4.1 Cathodic Inhibition from Tetrahedrally Coordinated Vanadate Species
Figure 4.2 shows that an inverse relationship exists between NaVO3 concentration
and inhibition of cathodic kinetics at pH 5.1, where dilute solutions perform better than
more concentrated solutions. The pH and concentration of these experiments was chosen
purposely to be near the transition from solutions dominated by octahedral species (0.25
M NaVO3) to those dominated by tetrahedral species (0.0025 M NaVO3). The dilute
NaVO3 solution contains relatively more tetrahedral vanadate as compared to the
concentrated NaVO3 solution, which contains relatively little tetrahedral species compared to octahedral species. At this pH, concentrated NaVO3 solutions are expected
to predominately have decavanadate species which are octahedrally coordinated. As the
concentration of NaVO3 in solution decreases, tetrahedral species become more
prevalent. The increased presence of tetrahedral species correlates with reduced cathodic
kinetics observed in cathodic polarization experiments. This is in agreement with other
work which has suggested that tetrahedrally coordinated vanadates inhibit the corrosion
of Al 2024 (18, 20-22, 27). However, the more concentrated NaVO3 solution containing
predominately octahedral decavanadate appears to be a modest cathodic inhibitor at this pH in some contrast to the findings of previous work which suggest that
117 decavanadates are reduced and possibly increase corrosion (18, 20). If decavanadate is
not responsible for the observed inhibition then the small concentrations of tetrahedral
species in solution are likely responsible.
4.4.2 Suppression of Al2CuMg Breakdown
In alkaline NaCl solutions a portion of the total Al2CuMg intermetallic particle
population would be expected to breakdown at potentials as low as -0.8 VSCE. At the
OCP in aerated NaCl solutions, which fluctuates approximately from -0.9 VSCE and -0.75
VSCE before the solution begins to become acidic from dissolved atmospheric CO2 with eventual OCP stabilization near -0.65 VSCE, Al2CuMg breakdown leads to selective Mg dissolution and formation of Cu enriched local cathodes. In contrast, Al2CuMg in
alkaline aerated NaCl with 10 mM NaVO3 is expected to breakdown at potentials only
above -0.7 VSCE. However, the OCP in NaVO3 solutions after 4 hours was observed to
remain below approximately -0.79 VSCE. The shift in alloy OCP to more active potentials
is a result of a decrease in cathodic kinetics observed on the matrix and all intermetallic
phases as seen in Figure 4.18-4.24. Iannuzzi et al. argue that decreased kinetics might be
the result of the adsorption of monovanadate to the surface, blocking reactive sites on
intermetallic particles and displacing Cl- from the matrix surface (20). The consequence
of decreased cathodic kinetics and the resultant shift in OCP combined with a small
increase in Al2CuMg breakdown potential observed in NaVO3 solutions is that Al2CuMg intermetallics remain “inefficient” cathodes. If Cu-enriched particles resulting from
118 Al2CuMg dissolution are assumed to act like pure Cu (Figure 4.19) then the rate at which
cathodic reactions are supported is greater than compared to rates supported by intact
Al2CuMg in NaVO3 solutions (Figures 4.22).
4.4.3 Variation in Anodic Behavior of Al2CuMg
As seen in Figure 4.8 two different anodic polarization behaviors were observed
in NaCl solutions. In a minority of cases Ecorr was observed to be -0.8 VSCE or more
noble with breakdown at potentials near -0.6 VSCE or more active. Yoon and Buchheit
reported that Al2CuMg has an initial OCP in 0.5 M NaCl near -1.2 VSCE, which drifts and
becomes stable between –0.86 and –0.76 VSCE after approximately 1000 s (12). It was
claimed that Al2CuMg experienced transient dissolution followed by passivation and
stable low-rate dissolution with an OCP near –0.8 VSCE (12). Possibly anodic
polarization curves of Al2CuMg that showed corrosion potentials at approximately -0.8
VSCE had already experienced transient dissolution compared to experiments that showed
more active corrosion potentials near the -1.2 VSCE OCP reported by Yoon. Generally,
the wide variably in Ecorr and Epit data obtained on Al2CuMg in NaCl could be attributed to samples being at various stages of dissolution prior to polarization, perhaps a result of extreme sensitivity to variation in polishing conditions. In contrast, the behavior of
Al2CuMg in NaVO3 solutions showed little variability in electrochemical response,
which is possibly evidence that vanadates control the process of dealloying and
redistribution of Cu on the surface, resulting in overall suppression of Al2CuMg dissolution. It is interesting to note that the breakdown potential distribution measured in
vanadate solutions is shifted to more active potentials compared to vanadate-free
119 solutions. This is consistent with the idea that vanadates suppress Mg dissolution from
the phase. In this case, suppression of Mg dissolution would result in a thinner, less
protective Cu-enriched layer on the surface of the phase that breaks down readily at lower
potentials. Even though breakdown of Al2CuMg is shifted to lower potentials when
vanadate is present in solution, the corrosion potential of the alloy remains below the
Al2CuMg breakdown potential distribution and corrosion susceptibility is decreased.
4.4.4 Vanadate Buffering and Circumferential Attack
In this work and in previous work it has been observed that vanadate solutions act
as buffers requiring more alkali or acid to adjust the pH than necessary for equivalent
vanadate-free solutions (18). Particularly, for this work once vanadate solutions were adjusted to pH values near 9 the solution pH remained stable for months. In contrast, equivalent NaCl solutions adjusted to pH values near 9 would quickly become more acidic as carbonic acid formed from dissolved atmospheric CO2. Leclère et al. have
argued that the presence of a buffer or weak acid can suppress alkaline dissolution near
cathodic sites resultant from hydroxyl produced by oxygen reduction, which results in
large cathodic currents that drive pitting and localized attack more than observed in
unbuffered solutions (30, 31). The production of reducible H+ at sites of local attack may
contribute to the increased cathodic currents. A rough estimate of buffering capacity, β,
for both carbonate and vanadate at pH 9.17 is easily calculated using the following simplified equation:
+ K aCt [H ] β = 2.303 + 2 Equation 1 (K a + [H ])
120 where, Ka is the acid dissociation constant and Ct is the total concentration of acid and base or total buffer concentration (41). At pH 9.2 the dominant species from atmospheric
- 2- CO2 in solution are HCO3 and CO3 which have a pKa of 10.33 and a total concentration of approximately 0.025 M which gives a buffering capacity of 0.0035 M/pH unit (35, 42).
If pH 9.17 10 mM NaVO3 is assumed to be dominated by monomeric vanadate, a pKa of
8.2 is appropriate (other tetrahedral vanadates have pKa values ranging from 7.92 to 9.1) and will give a buffering capacity of 0.002 M/pH unit (43, 44). At this pH the contribution of vanadate as a buffer is small compared to carbonate, while both seem to be relatively ineffective buffers. Figure 4.33 shows an Al 2024 sample exposed to 0.5 M
NaCl without vanadate, which is analogous to a nearly unbuffered solution. The observation of circumferential attack and white corrosion product around intermetallics is consistent with initial alkaline attack from increased pH near cathodic sites supporting oxygen reduction followed by aggressive pitting and crevice corrosion around intermetallics. In vanadate-free unbuffered deaerated NaCl solutions (Figure 4.35), despite the lack of oxygen, attack and trenches around intermetallic particles are still observed. In contrast, the degree of attack in alkaline deaerated (Figure 4.36), and in particular aerated (Figure 4.34), NaCl solutions that contain NaVO3 is significantly less than observed on corresponding NaVO3-free samples. There is some attack near intermetallic particles, but complete circumferential attack of intermetallics is not observed. If vanadate were assumed to have an effect as a buffer, this would be consistent with previous observations of attack on Al-Cu alloys in buffered and unbuffered solution, where more trenching is observed around intermetallics in unbuffered solutions as a result of alkalinazation while buffers decrease trenching by
121 neutralizing local cathodic alkalinization (30, 31). However, it has also been previously
observed that total cathodic current is greater in buffered solutions because current is not
used up for alkalinization and cathodic attack (30, 31). This is in contrast with current microcell results which show decreases in cathodic kinetics in vanadate solutions and
Figure 4.40. As a result, the action of vanadate as a cathodic inhibitor seems to dominate any effect from buffering. It may be possible that vanadate acting as a buffer does prevent trenching as observed in Figure 4.34, but there is not a corresponding increase in cathodic reactions that would be expected if vanadates were acting strictly as buffers rather than inhibitors. It is of interest to note that from a purely subjective analysis of the
Figures 4.34 and 4.36 the sample exposed to aerated solution shows less attack. The sample exposed to deaerated solution even shows some white particulate on and near intermetallics, possibly redistributed Cu. Vanadates have been shown to act primarily through the suppression of oxygen reduction kinetics and as a result are not nearly as effective in deaerated solutions, although vanadates are modest anodic inhibitors in both aerated and deaerated solutions (18).
4.5 Conclusions
• An inverse relationship exists between reduction in cathodic kinetics and NaVO3
concentration in mildly acidic solutions, where dilute solutions with tetrahedrally
coordinated vanadates inhibit cathodic kinetics better than concentrated solutions
with octahedrally coordinated vanadates.
122 • Tetrahedrally coordinated vanadates inhibit the ability of intermetallics particles
and the matrix of Al 2024-T3 to support cathodic reactions in alkaline NaCl
solutions.
• Tetrahedrally coordinated vanadates decreased corrosion potential and generally
increased pitting potential.
• The overall reduction of cathodic kinetics on intermetallics and the matrix of Al
2024- T3 in alkaline vanadate solutions shifts the OCP below the potential
required to cause Al2CuMg breakdown. This may prevent selective Mg
dissolution from Al2CuMg and subsequent formation of Cu-enriched clusters
capable of supporting rapid cathodic reaction kinetics.
• Tetrahedral vanadates decrease the incidence of circumferential attack around
intermetallic particles.
• The inhibition of cathodic reactions provided by tetrahedral vanadates dominates
any effect from buffering.
123 FIGURES
Figure 4.1: NMR spectra of pH 5.1 0.5 M NaCl solutions with a) 0.0025 M NaVO3 and b) 0.25 M NaVO3. The dilute NaVO3 solution has a greater proportion of tetrahedrally coordinated species (V1 and V4) relative to octahedrally coordinated species (V10) compared to the more concentrated NaVO3 solution which contains mostly octahedrally coordinated species.
124
-0.6
0.5 M NaCl -0.8 0.5 M NaCl + 0.25 M NaVO -1 0.5 M NaCl + 3 ) 0.0025 M NaVO 3 SCE
E (V -1.2
-1.4
-1.6
10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2
i (amps/cm2)
Figure 4.2: Cathodic polarization curves on Al 2024-T3 in aerated pH 5.1 0.5 M NaCl with 0.25 M NaVO3, 0.0025 M NaVO3, and without NaVO3. These experiments show an inverse relationship between NaVO3 concentration and inhibition of cathodic kinetics at pH 5.1, which correlates well with a transition from solutions dominated by octahedrally coordinated vanadates to tetrahedral vanadates.
125
Figure 4.3: NMR spectra showing the presence of tetrahedrally coordinated vanadates (V1, V2, V4, and V5) in pH 9.17 0.5 M NaCl + 10 mM NaVO3 solution used for microcapillary electrochemical experiments a) immediately prior to experimentation and b) 14 days later after completion of experiments.
126
-0.2
-0.4 NaCl NaCl + NaVO -0.6 3 )
SCE -0.8 E (V
-1
-1.2
-1.4 10-9 10-7 10-5 10-3 10-1
i (Amps/cm2)
Figure 4.4: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Al.
127
0.1
NaCl 0 NaCl + NaVO 3
) -0.1 SCE E (V -0.2
-0.3
-0.4 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.5: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Cu.
128
-0.2
-0.4 NaCl NaCl + NaVO 3 -0.6 )
SCE -0.8 E (V
-1
-1.2
-1.4 10-10 10-8 10-6 10-4 10-2 100
i (Amps/cm2)
Figure 4.6: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu.
129
0
-0.2 NaCl
NaCl + NaVO 3 -0.4 ) SCE
E (V -0.6
-0.8
-1 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.7: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu.
130
0
-0.2 NaCl NaCl NaCl + NaVO -0.4 3 ) SCE -0.6 E (V
-0.8
-1
-1.2 10-10 10-8 10-6 10-4 10-2 100
i (Amps/cm2)
Figure 4.8: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg.
131
0
-0.2 NaCl
NaCl + NaVO 3 -0.4 ) SCE
E (V -0.6
-0.8
-1
10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.9: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe.
132
0
-0.2 NaCl
NaCl + NaVO -0.4 3 )
SCE -0.6 E (V
-0.8
-1
-1.2 10-9 10-7 10-5 10-3 10-1
i (Amps/cm2)
Figure 4.10: Sample anodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3.
133
0.4
E E E 0 corr pit rp E V E V E V corr pit rp
-0.4 ) SCE E (V -0.8
-1.2
-1.6 020406080100
Percent Below
Figure 4.11: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al. The “V” in the legend indicates curves collected in solutions containing NaVO3.
134
0.4
0
-0.4 ) SCE E (V -0.8
-1.2 E E corr rp E V E V corr rp -1.6 020406080100
Percent Below
Figure 4.12: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 pure Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
135
0.4
E E E corr pit rp 0 E V E V E V corr pit rp
-0.4 ) SCE E (V -0.8
-1.2
-1.6 020406080100
Percent Below
Figure 4.13: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
136
0.4
0
-0.4 ) SCE E (V -0.8
-1.2 E E E corr pit rp E V E V E V corr pit rp -1.6 020406080100
Percent Below
Figure 4.14: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
137
0.4
0
-0.4 ) SCE E (V -0.8
-1.2 E E E corr pit rp E V E V E V corr pit rp -1.6 020406080100
Percent Below
Figure 4.15: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg. The “V” in the legend indicates curves collected in solutions containing NaVO3.
138
0.4
0
-0.4 ) SCE E (V -0.8
-1.2 E E E corr pit rp E V E V E V corr pit rp -1.6 020406080100
Percent Below
Figure 4.16: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. The “V” in the legend indicates curves collected in solutions containing NaVO3.
139
0.4
0
-0.4 ) SCE E (V -0.8
-1.2 E E E corr pit rp E V E V E V corr pit rp -1.6 020406080100
Percent Below
Figure 4.17: Cumulative probability plot displaying Ecorr, Epit, and Erp of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. The “V” in the legend indicates curves collected in solutions containing NaVO3.
140
-0.8
-1
-1.2
) -1.4 NaCl SCE NaCl + NaVO 3 E (V -1.6
-1.8
-2
-2.2 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2
i (Amps/cm2)
Figure 4.18: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al.
141
0
-0.4
NaCl NaCl + NaVO 3 ) -0.8 SCE E (V
-1.2
-1.6
10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.19: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Cu.
142
-0.8
-1
-1.2
) -1.4 SCE E (V -1.6 NaCl NaCl + NaVO 3 -1.8
-2
10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.20: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu.
143
-0.4
-0.8
NaCl ) -1.2
SCE NaCl + NaVO 3 E (V
-1.6
-2
10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.21: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu.
144
-0.6
-0.8
-1
-1.2 NaCl
) NaCl + NaVO SCE -1.4 3 E (V -1.6
-1.8
-2
-2.2 10-10 10-8 10-6 10-4 10-2
i (Amps/cm2)
Figure 4.22: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2CuMg.
145
-0.4
-0.8 )
SCE -1.2 NaCl
E (V NaCl + NaVO 3
-1.6
-2 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.23: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe.
146
-0.4
-0.8
) -1.2 SCE NaCl E (V NaCl + NaVO 3 -1.6
-2
10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1
i (Amps/cm2)
Figure 4.24: Sample cathodic polarization curves obtained using the microcapillary electrode in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3.
147
10-1
10-2
10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
-7 i i i 10 corr pass -1.3 V SCE i V i V i V corr pass -1.3 V SCE 10-8 0 20406080100
Percent Below
Figure 4.25: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for pure Al. The “V” in the legend indicates curves collected in solutions containing NaVO3.
148
10-1
10-2
10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
-7 i i 10 corr -1.3 V SCE i V i V corr -1.3 V SCE 10-8 0 20406080100
Percent Below
Figure 4.26: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Pure Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
149
10-1
-2 10 i i i corr pass -1.3 V SCE i V i V i V corr pass -1.3 V SCE 10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
10-7
10-8 0 20406080100
Percent Below
Figure 4.27: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al 4%Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
150
10-1
10-2
10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
-7 i i i 10 corr pass -1.3 V SCE i V i V i V corr pass -1.3 V SCE 10-8 0 20406080100
Percent Below
Figure 4.28: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al2Cu. The “V” in the legend indicates curves collected in solutions containing NaVO3.
151
10-3
10-4
-5 )
2 10
-6
i (Amps/cm 10
10-7 i i i corr pass -1.3 VSCE i V i V i V corr pass -1.3 VSCE 10-8 0 20406080100
Percent Below
Figure 4.29: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for, Al2CuMg. The “V” in the legend indicates curves collected in solutions containing NaVO3.
152
10-1
10-2
10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
-7 i i i V 10 corr pass -1.3 V SCE i V i V i V corr pass -1.3 V SCE 10-8 0 20406080100
Percent Below
Figure 4.30: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al7Cu2Fe. The “V” in the legend indicates curves collected in solutions containing NaVO3.
153
10-1
10-2
10-3 ) 2 10-4
10-5 i (Amps/cm
10-6
-7 i i i 10 corr pass -1.3 V SCE i V i V i V corr pass -1.3 V SCE 10-8 0 20406080100
Percent Below
Figure 4.31: Cumulative probability plot of icorr, ipass, and i at -1.3 VSCE of tested materials in pH 9.17 0.5 M NaCl solution with and without 10 mM NaVO3 for Al20Cu2Mn3. The “V” in the legend indicates curves collected in solutions containing NaVO3.
154
-0.5
-0.75 ) SCE -1 OCP (V OCP
-1.25 Air No Air
Air + NaVO No Air + NaVO 3 3
-1.5 0 3600 7200 10800 14400
Time (s)
Figure 4.32: OCP of Al 2024-T3 over 4 h in aerated and deaerated pH 9.17 0.5 M NaCl with and without 10 mM NaVO3. Solutions were not intentionally buffered and duplicate curves are shown.
155
Figure 4.33: SEM images of Al 2024-T3 after 4 h of OCP measurement in aerated 0.5 M NaCl solutions with an initial approximate pH 9.17.
156
Figure 4.34: SEM images of Al 2024-T3 after 4 h of OCP measurement in aerated 0.5 M NaCl solutions with an initial approximate pH 9.17 with 10 mM NaVO3.
157
Figure 4.35: SEM images of Al 2024-T3 after 4 h of OCP measurement in deaerated 0.5 M NaCl solutions with an initial approximate pH 9.17.
158
Figure 4.36: SEM images of Al 2024-T3 after 4 h of OCP measurement in deaerated 0.5 M NaCl solutions with an initial approximate pH 9.17 and with 10 mM NaVO3.
159
-0.2
-0.4
) -0.6 SCE
-0.8 Potential (V
-1
-1.2 10-8 10-7 10-6 10-5 10-4 10-3 10-2
i (Amps/cm2)
Figure 4.37: Duplicate anodic polarization curves of bulk Al 2024-T3 sheet in aerated 0.5 M NaCl + 10 mM NaVO3 solution at pH 9.17.
160
0.008 -1.2 V -0.7 V SCE SCE -0.9 V -0.6 V 0.006 SCE SCE -0.8 V -0.5 V SCE SCE 0.004 ) 2
0.002 i (Amps/cm 0
-0.002
-0.004 0 20406080100120140
Time (s)
Figure 4.38: Current response of potentiostatic hold experiments on Al 2024-T3 in pH 9.17 0.5 M NaCl with 10 mM NaVO3.
161
0.008
0.006
0.004 ) 2
0.002 i (Amps/cm 0
-1.2 V -0.7 V SCE SCE -0.002 -0.9 V -0.6 V SCE SCE -0.8 V -0.5 V SCE SCE -0.004 0 20406080100120140
Time (s)
Figure 4.39: Current response of potentiostatic hold experiments on Al 2024-T3 in pH 9.17 0.5 M NaCl without NaVO3.
162
1000
NaCl NaCl + NaVO 3
100 Charge PassedCharge (mC)
10 -1 -0.9 -0.8 -0.7 -0.6 -0.5 -0.4
E (V ) SCE
Figure 4.40: Total charged passed as a function of potential hold for Figures 4.38 and 4.39.
163
Table 4.1. Averaged electrochemical data for intermetallics tested in approximately pH 9.17 0.5 M NaCl solutions with and without 10 mM NaVO3. A “NaCl” heading indicates data from NaCl-only solutions while a “NaCl + V” heading indicates data collected in NaCl solution that contained 10 mM NaVO3.
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167
CHAPTER 5
HYDROTALCITE PIGMENTS FOR CORROSION INHIBITION OF
ALUMINUM ALLOY 2024-T3
5.1 Introduction
Aluminum alloy 2024-T3 is a high strength precipitation-age-hardened alloy
commonly used in aerospace applications because it has low density and superior
mechanical properties (1). Alloy additions of Cu, Mg, and Mn are largely responsible for
the high strengths achieved in Al 2024-T3. However, these alloy additions result in the formation of constituent intermetallic particles that are electrochemically different from the surrounding matrix (1-6). These electrochemical differences make the alloy inherently susceptible to localized corrosion, necessitating the use of supplemental corrosion inhibitors to protect in-service parts (2, 5, 7). Historically, chromate-based inhibitors and protection schemes have been used to prevent corrosion of Al 2024-T3, mainly resulting from the reduction of Cr6+ to Cr3+ on the alloy surface and the
subsequent suppression of oxygen reduction kinetics from an irreversibly adsorbed thin
Cr3+ hydroxide film (8-14). Additionally, chromate-based coatings and pigments have
168 demonstrated an ability to “self-heal”, where soluble Cr6+ species can migrate to
unprotected sites and be reduced to form the protective Cr3+ hydroxide film (9, 15, 16).
However, desires to find more “green” methods of corrosion protection have emerged because chromates are known carcinogens and environmental hazards (8-11, 17). A number of previous studies have focused on screening possible replacements, including various vanadates, molybdates, tungstates, phosphates, borates, and silicates among others (10, 11, 18-20). Although there are promising candidates, vanadates in particular, finding an inhibitor that provides the same level of protection as chromates has been elusive (10, 11, 18, 21-23).
In addition to finding an inhibitor pigment that provides sufficient inhibition, the number of possible replacement candidates is severely limited by solubility requirements
(19). An inhibitor must be soluble enough to be available in sufficient concentrations in a contacting electrolyte to passivate the substrate, but not so soluble as to cause blistering of the coating (18, 19). SrCrO4 is one such inorganic pigment that has solubility above what is necessary for inhibition, but well below concentrations that will lead to blistering
2- 3- (19). Other potential pigment grade inhibitor anions for Fe include: MoO4 , PO4 ,
2- 5- 2- - 2- 2- HPO3 , P3O10 , (SiO3 )n,n>1, BO2 , and NCH as identified by Sinko (19). MoO4 and
- BO2 are additionally recognized as inhibitors on Al (19). Identifying an inhibitor
pigment as potent as SrCrO4, which also has an appropriate solubility for direct use in organic coatings, has yet to be accomplished.
However, ion exchange pigments, which have been the focus of recent corrosion studies, may greatly increase the number of candidate inhibitor ions available for use in organic coatings by allowing the use of highly soluble inhibitors while minimizing the
169 risk of blistering (22-27). In particular, anionic exchange clays and hydrotalcites, which
structurally consist of alternately charged positive cation layers and negatively charged
inter-layers, have received attention (22, 23, 26-28). The structure of hydrotalcites can
generically be described by:
z+ 3+ A+ m- . [M 1-xM x(OH)2] X A/m nH2O Eqn. 1
where A=x for z=2 and A=2x-1 for z=1 (29). The cation sheets have a structure similar
to brucite, Mg(OH)2, or gibbsite, Al(OH)3, and develop a net positive charge from
random substitution of trivalent cations for divalent cations or vacancy occupation by
monovalent cations, respectively (28-32). The net positive charge developed in the
cation layers is balanced by anions in the inter-layers, which are also known as galleries
(28-32). This presents the possibility of intercalating a wide variety of inhibitor anions
into the inter-layers of a hydrotalcites. These inhibitor anions may be subsequently
released into solution by selective exchange for aggressive anions in a contacting
electrolyte (23). Anions with large charge densities and small size, such as Cl-, will selectively exchange for intercalated anions with lower charge densities and large size
(32-34). X-Ray diffraction (XRD) is commonly used for structural verification of hydrotalcites and may be used to track exchange-induced structure changes in the inter- layer spacing of hydrotalcite pigments exposed to NaCl (23).
There are a number of methods available to synthesize hydrotalcites (35).
However, direct synthesis is an easy and attractive method that may be accomplished without many of the complications associated with other synthesis methods (36). There are a number of issues that need to be taken into account to successfully synthesize a
170 “clean” hydrotalcite, including cation and anion ratios, co-precipitation pH and anion
stability, and possible complications from atmospheric CO2. “Clean” refers to a final
product that is primarily hydrotalcite with galleries filled by desired anions. This implies
that impurity anions have not been intercalated and the product is free of either of the
cation hydroxides that may form.
Cavani et al. offer two generic formulas to produce pure hydrotalcites based on
divalent and trivalent cation skeletons (28):
M (III) 0.2 ≤ ≤ 0.4 Eqn. 2 [M (II) + M (III)]
1 An− ≤ ≤ 1 Eqn. 3 n M (III)
where M(III) and M(II) are moles of trivalent and divalent cations, respectively, and n
and A are anion charge and concentration, respectively. Equation 2 ensures that
sufficient and appropriate concentrations of trivalent cations are available to substitute for
approximately one third of divalent cations in brucite-type structure. This is necessary to
develop a net positive charge needed to incorporate anions. In addition, Equation 3
guarantees at a minimum that there is a sufficient anionic charge in the synthesis bath to
populate the inter-layers with a desired anion. Pure hydrotalcites are only formed in a
narrow window of cation ratios; deficiencies of the divalent cation lead to the formation
of the trivalent cation hydroxide and excess of the divalent cation leads the formation of
the divalent cation hydroxide (37). There does not appear to be a similar published set of
rules for synthesis using monovalent and trivalent cations. However, the theoretical ratio 171 for the formation of ordered cation sheets for monovalent and trivalent cations is 2 to 1
(29). Additionally, Li+ is the only monovalent cation with an ionic radius of appropriate
size to form hydrotalcite-like materials (28).
A further consideration for hydrotalcite synthesis is that different combinations of
cations will co-precipitate at different pHs. The range of co-precipitation can be
determined using titration curves that are produced when a solution containing a pair of
cation salts is titrated with alkali. The resultant curves have distinct plateaus, which indicate pH regions of precipitation; an intermediate plateau occurs at the pH of co-
precipitation, while plateaus at pH extremes occur from precipitation of metal
hydroxides. For direct synthesis, the desired anion to be incorporated into a hydrotalcite must be stable and predominant in solution at the same pH range of cation co- precipitation. Anion stability can be determined with aid from a number of different literature sources including predominance diagrams, speciation studies, and Pourbaix diagrams.
A potential complication in hydrotalcite synthesis is the possible incorporation of
2- CO3 anions, which are easily intercalated into hydrotalcite interlayers as a result of the
2- ion’s small size and large charge density. CO3 anions come from dissolved atmospheric
- 2- CO2, which forms H2CO3. H2CO3 has equilibria involving HCO3 and CO3 , which both
have increased solubility in aqueous solutions of pH greater than approximately 6.5 (38).
Therefore, if co-precipitation of a pair of cations occurs above pH 6.5, then synthesis must be conducted under deaerated conditions to prevent competition for inter-layer space.
172 Previous studies suggest that hydrotalcite pigments can be used as corrosion inhibitors and possibility used to deliver inhibitor anions that are too soluble for direct
use in organic coatings (22, 23, 26, 27). The objectives of this work are to create and
develop a framework by which hydrotalcite pigments can be tailored to incorporate a
number of different anion inhibitors into various cation host skeletons. This process will
include identification of issues critical to synthesizing “clean” hydrotalcite powder
pigments for use in organic coatings. Further, this work aims to synthesize and evaluate
the performance of a number of hydrotalcite pigments relative to a SrCrO4 standard, with
a particular focus on vanadate-bearing hydrotalcites. Performance will be compared
using electrochemical impedance spectroscopy (EIS) of coated panels exposed to static
NaCl solutions and salt spray exposure tests of coated and scribed panels. Finally, this work aims to characterize inhibitor release for a number of vanadate hydrotalcite
pigments and SrCrO4 using neutron activation analysis (NAA).
5.2 Experimental Procedures
5.2.1 Materials and Chemicals
Commercially obtained Al 2024-T3 sheet, which contains Al, 3.8-4.9 Cu, 1.2-1.8
Mg, 0.3-0.9 Mn, 0.5 Si, 0.5 Fe, 0.25 Zn, 0.1 Cr, and 0.15 Ti by weight, was coated with a pigmented polyvinyl butyral (PVB) coating and used for EIS and Salt Spray exposure experiments (1). 18.2 MΩ.cm deionized water and reagent grade chemicals were used for
synthesis of hydrotalcite pigments, organic coating preparation, long term static EIS
experiments, and NAA sample preparation.
173 5.2.2 Hydrotalcite Pigment Synthesis
A variety of anion exchange pigments were synthesized to contain inhibitor
anions previously reported to be corrosion inhibitors, including vanadates, molybdates, tungstates, phosphates, silicates, and borates (10, 11, 19-21). Details of synthesized pigments are summarized in Table I. A shorthand notation was used in the table and throughout this chapter for different pigments to indicate whether or not the pigment was a hydrotalcite and to specify the primary element of the intercalated inhibitor anion.
HTV1 for example is a hydrotalcite containing vanadate, while NaVO3 is a reagent grade
salt used as a control. Additionally, for each hydrotalcite pigment the table shows the
cations used for synthesis and the relative quantity of inhibitor anion used for synthesis,
in terms of the ratio of moles Al to moles inhibitor anion. In most cases the primary
elemental component of the inhibitor anion is used to represent generically the anion in
the table. This was done because the precise predominant anion used for a specific
synthesis depends on a number of factors including pH and concentration and an exact
assignment would only be speculative. Other details about pH of co-precipitation, pH of
post-synthesis exchange, whether synthesis was deaerated or not deaerated, and what acid
was used for pH adjustment are also included in the table. Pigments requiring synthesis above approximately pH 6.5 were conducted in deaerated solutions to avoid complications from atmospheric CO2, while syntheses in acidic solutions were performed
open to air. Deaerated synthesis was performed in a special cell designed to bubble Ar
simultaneously through the cation, anion, and NaOH solutions prior to and during
synthesis. The cell was also equipped with a pH meter and heating element so that both
synthesis and aging could be accomplished under an Ar atmosphere. Additionally, H2O2
174 and HNO3 were used in some instances to adjust the solution pH rather than HCl to
achieve a decavanadate-free and Cl--free synthesis, respectively. Synthesis parameters
and inhibitor anions were varied to observe and draw conclusions from differences in
observed inhibition. The table also includes information on controls against which
- pigments were compared. These include SrCrO4, HTCl , NaVO3, HTCO3, and neat PVB.
References were consulted to determine an appropriate pH of anion stability at which co-
precipitation should occur for vanadates (39), molybdates (39), tungstates (39),
phosphates (40), silicates (41), and borates (42).
The exact synthesis procedure followed for each hydrotalcite pigment was different; however, all syntheses followed the same general pattern. First a solution of
dissolved anion salt was adjusted to the pH at which a particular inhibitor anion of
interest was predominant. Then a solution containing both dissolved cation components,
in an appropriate ratio, was added to the anion solution while maintaining the pH in the
range of co-precipitation for that particular pair of cations using NaOH. Following
synthesis, the pigment slurry was aged for 24 hours at 55-60oC to improve crystallinity.
After aging, the slurry was washed and filtered with DI water and methyl alcohol in
conjunction with a vacuum aspirator. The resultant powders were air dried and ground
by hand using a mortar and pestle to produce a fine powder. This powder was then
exchanged for 24 h in 100 mL of an equivalent anion solution that was used for that
pigments’ respective synthesis. The goal of the post-synthesis exchange was to achieve a
2- - higher purity pigment by forcing unwanted contaminant anions such as CO3 and Cl to be exchanged for the desired anion of a particular hydrotalcite. After the post-synthesis exchange, the pigments were once again washed and filtered using DI water and finally
175 air dried. The resultant powder was ball milled for 1 hour in methyl alcohol and then
dried in a furnace at 40oC overnight before being added to a PVB resin and applied to test
panels.
5.2.3 XRD Structure Confirmation
A Scintag XDS-2000 diffractometer with a CuKα source (λ=1.542 Å) was used to
obtain XRD patterns of synthesized hydrotalcite pigments. Patterns were collected from
2-theta values ranging from 3o to 68o with a scan rate of 1o per minute using a 45 kV
accelerating voltage and 20 mA of current. EVATM analysis software was used to
remove the background, strip Kα, and smooth all data.
5.2.4 Organic Coating Preparation and Application
Al 2024-T3 sheet was cut into 3 by 4 inch panels and degreased for 2 minutes in
o 1.0 L solution of DI water containing 48.0 g Na2CO3 and 32.4 g Na2SiO3 held at 65 C.
The panels were rinsed in DI water and then deoxidized for 3 minutes in 1.0 L of DI
o water containing 30.0 g Sanchem™ and 72.0 mL HNO3 at a temperature of 55 C. Panels were coated with a resin consisting of 15 wt% PVB dissolved into methyl alcohol over several days. The edges of the panels were masked using electrical tape to help ensure consistency in coating thickness and prevent crevice corrosion during salt spray exposure.
Volumes of PVB were loaded with 5 wt% pigment, which previous in-house coatings
work showed to be an appropriate loading level to produce uniform coatings without
pigment clumping. The PVB and pigment mixtures were agitated overnight and then ball
milled for 1 hour prior to application. Coatings were applied within 24 h of panel
176 degreasing and deoxidizing using a single pass of a drawdown bar. After application, the
coatings were allowed to cure for 24 h before further contact. Two sets of panels were
prepared for each pigment; one for static solution EIS measurements and a second for
ASTM B 117 salt spray exposure. ASTM B 117 is a standardized test which consists of
exposure of coated panels to a salt spray, 5 wt% NaCl, in an enclosed chamber held at
35oC.
5.2.5 Coating Evaluation by EIS on Panels Exposed to Static NaCl Solutions
Cells for EIS experiments were prepared by gluing a 1 5/16 inch inner diameter polyvinyl chloride (PVC) coupling to the surface of the coated panels with a silicone
sealant. This left an exposed area of approximately 8.55 cm2. The couplings were filled
with approximately 40 mL of 0.5 M NaCl. EIS experiments were carried out using a
Gamry Reference 600TM system with a saturated calomel electrode (SCE) reference
electrode and a graphite rod counter electrode. Spectra were collected using 10 mV AC
amplitude with an initial frequency of 105 Hz, a final frequency of 0.01 Hz, and
collection of 7 points per decade. EIS spectra were collected at 1, 4, 10, 20, and 35 days
TM and later modeled with the assistance of Zview data analysis software.
5.2.6 Coating Evaluation by Salt Spray Exposure of Scribed Panels
Salt spray exposure testing was carried out on one set of coated panels for 750 h.
Prior to exposure, coated panels were scribed using two passes of an electrical scribe to
thoroughly perforate and expose bare metal under the PVB coating. Blistering and scribe protection performance were subjectively ranked by optical inspection using a
177 magnifying lens. These rankings are found in Table 5.1 with a ranking of “1” indicative of a strong performance relative to other tested coatings. For example, a scribe protection ranking of 1 indicates bare metal was observed over the entire scribe after exposure, 2 indicates bare metal observed over most of the scribe with some corrosion product, 3 indicates some shiny metal observed in the scribe but lots of corrosion product, and a 4 indicates that the scribe is completely choked with corrosion product. A similar ranking system, from 1 to 5, was used to compare coating blistering.
5.2.7 Inhibitor Release Characterization Using NAA
1 mL samples of 0.5 M NaCl exposed to different vanadate hydrotalcites and
SrCrO4 pigments, were dehydrated and analyzed using NAA for the presence of Na, Cl,
Al, Zn, Ni, V, and Cr. 2 grams of pigment were exposed to 40 mL of 0.5 M NaCl in 50 mL centrifuge tubes. 1 mL samples were collected at 2, 7.25, 28.5, 101.25, 169, and 290 hours using a micropipette and placed into 2 mL polypropylene microvials. Immediately prior to taking samples, the solutions were centrifuged at 1000 rpms for 3 minutes to drive any particulate matter to the bottom of the tube and help ensure that only the solution was sampled. After each sampling, the centrifuge tubes were thoroughly shaken to remix pigment with solution. The 1 mL samples were placed under a heat lamp to evaporate off water and, once dehydrated, the sample were subsequently stored in a desiccator. The mass of the dehydrated samples were later recorded.
NAA was conducted at The Ohio State University Research Reactor (OSURR) using procedures for similar work conducted in this lab as an outline (22, 43). Trace element analysis from samples is possible by activation and subsequent detection of
178 gamma radiation, which is distinctive for each element (22, 43). This technique allows
elemental identification and quantification; up to part per billion resolutions if necessary.
A high resolution gamma ray spectrometer, based on a hyper pure germanium gamma ray detector was used for elemental characterization. An Aptec Nuclear Model 5008 MCArd was used in conjunction with OSQ Plus/Professional software from Aptec Nuclear for
data collection and analysis. Each sample was irradiated and counted twice, once for elements with short half-lives and once for elements with long half-lives. Samples containing short-lived elements (Na, Cl, Mg, Al, and V) were placed on the edge of the reactor core for an irradiation time of 1 min using a pneumatic transfer system or
“rabbit”. These samples were run at 50 KW or 10% of full power, which resulted in an approximate thermal neutron flux of 2.3 x 1011 neutrons.cm-2.s-1. Short-lived samples
were counted for 60 s. Long-lived elements were placed in the Central Irradiation
Facility (CIF) located in the center of the core where higher fluxes help improve sensitivity of the analysis. Long-lived samples (Cr, Ni, and Zn) were irradiated for 60 min at 300 KW or 60% of full power which gives an approximate thermal neutron flux of
7.2 x 1012 neutrons.cm-2.s-1. 600 s counting times were used for elements with long half- lives. Additionally, a limited number of samples were counted for times up to approximately 20 h to get more accurate counts.
Standards for detection of Na, Cl, V, Zn, Ni, Al, Mg, and Cr using NAA were
made. Appropriate quantities of NaCl (assay 100.5%), NaVO3 (assay 98%), ZnCl2 (assay
. . 97%), NiCl2 (assay 98%), AlCl3 6H2O (assay 99.9%), and MgCl2 6H2O were dissolved in
100 mL of DI water to create 1 M solutions of each chemical, respectively. Using serial
dilutions, 0.1 M and 0.01 M solutions were made from the initial 1 M solution. Samples
179 of each solution were taken using a micropipette. Ideally this would allow for a three point calibration to be made. However, in some instances 0.1 mL samples of 0.01 M solution did not produce signals significant enough to be used. The volumes of solution
for all standard preparation were 0.1 mL except for the samples taken for the NaCl standard, which were 1 mL. Standards were evaporated under a heat lamp and the
masses were recorded. Cr2O3 (assay 98%) powder samples of approximately 0.1, 0.01,
and 0.001 grams were used as the standard for Cr.
5.3 Results
5.3.1 Structural Confirmation of Synthesized Pigments
Figure 5.1 shows XRD patterns of synthetic vanadate hydrotalcite pigments.
These pigments were fabricated to ascertain whether intercalation of vanadates other than decavanadates or the use of cations such as Ni, Mg, and Li rather than Zn in the host structure improved corrosion resistance. XRD is a quick and easy method available for structural confirmation of hydrotalcite-like materials. Reflections resulting from the cationic layers of hydrotalcites occur at low 2θ and are assigned (003), (006), and (009) reflections (23). Additionally, changes in inter-layer spacing may be used to detect anion exchange and generically may be used to help identify what type of anion is intercalated in the inter-layer (23). In Figure 5.1 most pigments show reflections at approximately
15o and 23o which correspond to the (006) and (009) reflections for decavanadate,
respectively (23, 36). The peak at approximately 9o is likely from the formation of Al or
Zn polyoxovanadate and may overlap with the (003) decavanadate reflection at slightly
lower 2θ (36). Both are resolved in the pattern of HTV4. HTV3 does not have the
180 reflections associated with decavanadate as this pigment was synthesized using H2O2 rather than a strong acid to adjust the pH of the original NaVO3. The use of H2O2 allows acidification of vanadate solutions without the formation of decavandate (44). HTV3 was also observed to have much sharper peaks than other vanadate hydrotalcites which may indicate a higher degree of crystallinity. Reflections at large 2θ result from the structure within cationic sheets (23, 28).
Figure 5.2 shows the XRD patterns for hydrotalcite pigments synthesized to contain molybdates, phosphates, silicates, tungstates, borates, chlorides, and carbonates.
These pigments were synthesized to observe whether anions, additional to vanadates, would function as inhibitors when intercalated into hydrotalcites. Additionally, both
- 2- HTCl and HTCO3 where synthesized and used as negative controls because they often appear as contaminants while synthesizing other hydrotalcites. Attempted synthesis
using phosphate anions does not appear to have been successful in producing a layered
- 2- structure. In contrast, synthesis using Cl and CO3 resulted in highly crystalline
materials. Cl- hydrotalcites are expected to have strong peaks at approximately 11o and
o 2- 22 for (003) and (006) reflections, respectively (23, 45, 46). CO3 hydrotalcites have
reflections at similar 2θ values with reflections from (003) and (006) at approximately
11o and 23o, respectively (47). Additionally, the peaks at approximately 53o and 57o in
2- the pattern for HTCO3 are likely from impurity Al2O3 and do not appear in any other
patterns (47). Based on comparisons of XRD patterns it does not seem that hydrotalcites
- 2- with either Cl or CO3 were present in large quantities as contaminants in other
synthesized pigments.
181 5.3.2 EIS of PVB Coatings Pigmented with Vanadate Hydrotalcites
EIS was used to monitor corrosion with time of Al 2024 panels coated with pigmented PVB exposed to static 0.5 M NaCl. The Nyquist plots of most tested
pigments typically were observed to have two resolved arcs and a partially resolved third arc, indicating three time constant behavior. SrCrO4 and neat PVB were exceptions.
Three time constant behavior is believed to be associated with two layers, PVB and
oxide, and pores through those layers as seen schematically in Figure 5.3. Figure 5.4
shows sample Bode magnitude and phase angle plots as a function of time for an HTV3
pigmented PVB coating. This set of data was chosen as a sample because it displayed
behavior similar to HTV2, HTV4, HTV5, and HTV7, and it shared some similarity with
data of HTV1 and HTV6. At the highest frequencies, the solution resistance in the Bode
magnitude plot is not resolved and behavior is dominated by capacitance of the PVB
layer, which does appear to change slightly with time. At frequencies between
approximately 1 and 100 Hz, behavior is initially dominated by a capacitance associated
with the oxide layer. With time, this capacitive region shifts to lower frequencies and
intermediate frequencies increasingly show resistive behavior as a resistance associated
with the PVB layer in the model becomes resolved. The change in oxide film
capacitance is possibly the result of hydration. Through defects in the coating, the
electrolyte may have direct access to the interface between the PVB and oxide layers and hydration of the oxide layer may initiate at this interface. Behavior is dominated by pore resistance at yet lower frequencies and is initially observed to decrease with time before
stabilizing at approximately 1 x 105 ohms. At the lowest frequencies capacitance
behavior associated with defects becomes apparent. The total impedance at low
182 frequencies also decreases with time. As a check the theoretical capacitance of the PVB coating can be compared to that obtained from the data. The theoretical capacitance of a
PVB coating can be calculated using the relation, t = εεoA/C, where t is thickness, εo is
8.85 x 10-12 F/m, ε is the dielectric constant, A is area, and C is capacitance. The thickness of the PVB coatings was measured to be approximately 24 μm, the dielectric
constant is 2.69, and the area of the electrode was 8.55 cm2 (48). Using these parameters
the capacitance of the PVB coating is approximately 8.5 x 10-10 F which is within two
orders of magnitude of the capacitance obtained from extrapolation of the capacitive
region associated with the PVB layer in Figure 5.4 to 0.16 Hz. Further, the resistance of
the PVB coating can be calculated using the relation, R = ρl/A, where ρ is resistivity, l is
thickness or length, and A is area. If the resistivity is assumed to be approximately 1 x
1011 Ω.cm the PVB coating should have a resistance equal to approximately 2.8 x 107 Ω
(49).
Figure 5.5 shows an analogous Bode magnitude and phase angle plot for PVB
coatings pigmented with SrCrO4. The behavior is largely capacitive over the all
frequencies. Relative to the HTV3, and other vanadate hydrotalcites, the total impedance
at low frequencies of the SrCrO4 is larger. Additionally, the response of SrCrO4 pigments in PVB was not observed to change with time except at low frequencies were the beginnings of a resistive plateau associated with pore resistance decreased with time until becoming stable at approximately 1.9 x 106 ohms after 10 days.
EIS spectra were modeled using the simplified equivalent circuit found in Figure
TM 5.3 in conjunction with Zview data analysis software. By using circuit elements to represent physical phenomena occurring in the coating, characteristic capacitances and
183 resistances may be extracted and used for comparisons of performance and to develop an understanding of coating degradation. Both an explicit (a) and simplified (b) equivalent circuit for PVB coated Al 2024-T3 are shown in Figure 5.3. These models were designed to replicate a coating system that consists of a porous PVB coating on an oxide layer.
The explicit model was developed because attempts to model spectra with three time constants using more conventional equivalent circuits were not successful. In particular, it was possible to model the bode plots well with a number of more simple and conventional models used for similar coatings work, but none of these were able to achieve a remotely accurate fit for the third arc often observed in the Nyquist plots (43).
In an attempt to use the simplest circuit that provided acceptable data fitting, the oxide resistances were omitted from the simplified circuit because they are too large to detect and essentially act as an open in the circuit. Further, the constant phase element (CPE) and charge transfer resistance associated with the defect interface were simplified to a single CPE which fit the data better and allowed for the fact that it is not known whether charge transfer of diffusion leads to this response in the data. Generally, the equivalent circuit developed for three time constant behavior was able to fit the data well as seen in
Figures 5.6 and Figures 5.7. Figure 5.6 shows the Bode magnitude and phase angle plots of HTV3 after 4 days overlaid with the fit produced using the simplified equivalent circuit. Figure 5.7 shows the Nyquist plot and fit for HTV3 after 4 days of exposure.
The only instance where variation in fitting was observed was at low frequencies in the
Nyquist plot as seen in Figure 5.7. However, even in these instances fitting was acceptable. The simplified equivalent circuit model was used to extract total impedance
(Ztotal), pore resistance (Rpore), defect capacitance (Cdefect), and capacitances of the PVB
184 (CPVB) and oxide layers (Coxide). To obtain better fits, CPEs were used in place of capacitors and the following equation was used to extract true capacitance (Ctrue) (50):
1 1−φ φ φ Ctrue = Y R Eqn. 4
where, Y is the CPE magnitude, R is resistance, and φ is a value ranging from 0 to 1 that
is a measure of the capacitive behavior of the CPE. A φ value of 1 would indicate purely
capacitive behavior.
Figure 5.8 shows the total impedance at low frequency for vanadate pigmented
coatings relative to a SrCrO4 standard. None of the vanadate pigments performed as well
as the SrCrO4. However, most vanadate pigments did have total impedance within approximately one order of magnitude of the SrCrO4 standard. Interestingly, the HTV2
pigment, which was synthesized with an abundance of vanadate, performed the poorest of
all vanadate hydrotalcites. The largest total impedance observed after 840 h was from
HTV4 and HTV5, which were synthesized without the presence of chloride and with Mg
rather than Zn in the cation host structure, respectively. The total impedance of all
vanadate hydrotalcite pigments showed an overall slight decrease with time. In
particular, most of the decrease in total impedance occurred between 24 and 96 h at
which point the total impedance of all pigments except HTV1 appeared to stabilize.
Figure 5.9 shows the pore resistance over 840 h of PVB coatings pigmented with
vanadate hydrotalcites relative to a SrCrO4 standard. Pore resistances were determined at frequencies between approximately 10 and 0.05 Hz. Large pore resistances are characteristic of good coatings. The pore resistances of all vanadate hydrotalcites
185 decreased between the first measurement at 24 and 96 h. This was in contrast to the pore
resistance of a SrCrO4 pigmented coating, which was observed to increase after 24 h
before becoming stable at approximately 1 x 107 ohms.cm2. After 96 h the pore
resistance of most vanadate hydrotalcites remained stable up to 840 h except for HTV1
which decreased after 480 h. Additionally, the pore resistances of all vanadate
hydrotalcites were distributed between approximately 2 x 105 and 2 x 106 ohms.cm2 after
840 h. None of the vanadate pigmented coatings performed as well as the SrCrO4, similar to the results for total impedance.
Figure 5.10 shows the defect capacitance of PVB coatings containing vanadate hydrotalcite pigments. Defect capacitances were determined from frequencies lower than
0.1 Hz. All defect capacitances except for HTV2 increased between 24 and 96 h before generally becoming stable. HTV6 showed a sharp increase in defect capacitance after
480 h. The SrCrO4 pigment was observed to have a smaller defect capacitance compared
to vanadate hydrotalcites, although most vanadate pigments had defect capacitances
within an order of magnitude of the SrCrO4.
Figure 5.11 shows the oxide capacitances for PVB coatings pigmented with vanadate hydrotalcites and SrCrO4. All tested vanadate hydrotalcites had oxide
capacitances within approximately an order of magnitude of SrCrO4 up to 480 h.
However, the oxide capacitance of the SrCrO4 pigmented PVB coating remained
relatively stable over 840 h, while the capacitances of vanadate hydrotalcites trended
upward with time. The same trend with time is evident in Figure 5.4 in the capacitive regions at intermediate frequencies and is possibly associated with hydration of the
coating.
186 Figure 5.12 shows the capacitances associated with the PVB layer pigmented with vanadate hydrotalcites and a SrCrO4 standard. In general, capacitance of most of the
PVB coatings decreased between 24 and 96 h before becoming stable. Additionally, the capacitance of all PVB coatings fell within an order of magnitude, between approximately 2 x 10-10 F.cm-2 and 2 x 10-9 F.cm-2. The capacitance of this layer is expected to fall within a tight distribution and may help confirm that the equivalent circuit model is correct. However, there is incongruence in interpretation as the generally stable capacitance associated with the PVB layer does not suggest water uptake, as would be expected, while the gently increasing capacitances of the oxide layer with time in
Figure 5.11 may suggest hydration of the oxide near defects.
5.3.3 Salt Spray Exposure of Vanadate Hydrotalcite Pigmented Coatings
Figure 5.13 shows scribed PVB coated panels after 750 hours of exposure in a salt spray chamber. The numbers superimposed on the bottom of each panel are subjective rankings of the degree of blistering (B) and scribe protection (SC). A ranking of “1” indicates the best performance while rankings of “4” and “5” indicate the poorest performance for scribe protection and blistering, respectively. A complete listing of the ranking of both degree of blistering and scribe protection is found in Table 5.1. The
SrCrO4 panel displayed the best overall performance with no observed blistering and complete protection of the scribe. All other tested panels displayed some blistering, attack in the scribe, or some combination of the two. However, the HTV3 and HTV6, synthesized with a peroxovanadium complex rather than decavanadate and with Ni rather than Zn cations, respectively, showed a combination of scribe protection and blister
187 resistance that was similar to SrCrO4. Figure 5.14 shows a close up of the SrCrO4 and
HTV3 panels after 750 hours of exposure in salt spray for comparison. In general, all
hydrotalcite pigments containing vanadates showed moderate to strong scribe protection,
but were observed to blister. In the case of HTV2, blistering was severe. This was
somewhat perplexing as a similar formulated pigment without the post-synthesis
exchange had previously been observed to have superior performance during prior tests
in this lab. During the post-synthesis exchange for HTV2 a number of blocky orange
precipitates were observed to form throughout the pigment. No effort was made to
remove these precipitates from the pigment before it was filtered, ground, and milled. It
seems likely that these precipitates were some form of a highly soluble vanadate salt that
could be the cause of extensively observed blistering on the coated panel.
5.3.4 Inhibitor Release from Vanadate-Bearing Hydrotalcite Pigments
NAA analysis was used to measure release of Na, Cl, Al, Zn, Ni, V, and Cr from
vanadate hydrotalcite pigments and SrCrO4 into 0.5 M NaCl solutions. Figure 5.15
shows the release kinetics of vanadium from vanadate hydrotalcites and Cr from SrCrO4 pigment exposed to 0.5 M NaCl. Release of vanadium from most vanadate hydrotalcites appears to be a diffusion controlled process in agreement with previous work (22, 43).
Depending on the exact vanadate hydrotalcite, the concentration of total vanadium in solution after 840 h ranged from 0.00023 M to 0.0154 M, while the concentration of total
Cr in solution was 0.0066 M. The pH of release experiment solutions was not recorded meaning that a determination of the exact quantity and type of vanadate species in solution is not possible. However, if solutions are assumed to be approximately pH 6, as
188 reported for a similar exposure of vanadate hydrotalcite to NaCl solution, then
tetrahedrally coordinated vanadates will predominate (21). Further, the most dilute
vanadate solutions, 0.00023 M total vanadium, would likely form solutions that are
predominately monovanadate (21). Similarly, the pH of the SrCrO4 solution was not measured, but solutions with concentrations near 0.0066 M Cr6+ would speciate as either
- 2- HCrO4 or CrO4 depending on if the pH were acidic or alkaline, respectively (9). The
solubility of SrCrO4 has been reported to be 0.005 M in the literature (19). Figure 5.16
shows results of total Zn and Ni release, presumably as Zn2+ and Ni2+, from three
different vanadate hydrotalcite pigments. The Zn data was obtained using long counting
times and the Ni data was obtained after long decay times because accurate results were
not obtained during the original count of these samples. Ni was present in concentrations
of approximately 0.0022 M and the Zn was detected in 0.0045 M concentrations. Both
the pigments used for Zn counting, HTV1 and HTV2, were observed to blister
extensively while the pigment used for Ni counting blistered only slightly. Without counting Zn release from other pigments a definitive conclusion can not be made, but Zn containing hydrotalcites may blister more than Ni containing hydrotacites. Similarly, release of Mg from HTV5 was not accurately determined as a result of the short half-life
of Mg and the counting times necessary to decrease error to acceptable levels. The
release of small concentrations of Al was detected but could not be accurately assigned
concentrations.
189 5.3.5 EIS of PVB Coatings Pigmented with Various Non-Vanadate Hydrotalcites
Figure 5.17 shows the total impedance for PVB coatings pigmented with non-
vanadate hydrotalcites and control pigments. The plot has been scaled, as have all the
following plots, to allow easy comparison to previously presented SrCrO4 and vanadate
hydrotalcite data. Among these pigments, neat PVB had the largest total impedance
2- followed by HTCO3 and HTW, which had similar total impedance as SrCrO4 in Figure
5.8. The high total impedance of the neat PVB may be associated strong barrier properties resultant from the PVB coating resin not being disrupted by the addition of a pigment. The total impedance of neat PVB does begin to decrease sharply at 480 h which could be an indication that the barrier properties of the coating are beginning to degrade. The disparate performance of neat PVB relative to pigmented PVB implies that apparent total impedance is a function of both the corrosion resistance and the barrier properties of the coating. The bode magnitude plots for neat PVB show strongly capacitive behavior until low frequencies at which point a resistive plateau occurs at approximately 5 x 108 ohms.cm2. The behavior changes little with time until after 480 h
when the resistive plateau drops to 2 x 107 ohms.cm2 and a second, and perhaps third,
time constant appear. The Nyquist plot shows single time constant behavior up to 480 h
after which the plot develops three time constants. This may indicate that the barrier
properties of the neat PVB have failed and corrosion had initiated. The ability of PVB to
wet and form a coherent barrier with various pigments may affect total impedance at
earlier times and apparently, in turn, observed total impedance. The HTB, HTP, and
NaVO3 had similar total impedance as most HTV pigments over 840 h, while the HTMo
190 and HTSi had low total impedance at low frequency. Additionally, it is of interest to note
that the impedance of the NaVO3 increased with time.
Figure 5.18 shows the pore resistances of non-vanadate hydrotalcites and controls
with time. Neat PVB was observed to have the largest pore resistance of any tested
pigment. The pore resistance of HTP, HTB, and NaVO3 were observed to trend upward
2- with time. Additionally, the HTW, HTCO3 , and neat PVB had comparable pore
resistances to SrCrO4.
Figure 5.19 shows the defect capacitances of non-vanadate hydrotalcites and
controls used as pigments in a PVB coating. Neat PVB was observed to have the
smallest defect capacitance over the course of the experiment. However, after 480 h the
defect capacitance of PVB was observed to increase dramatically. It is also possible that
the data for PVB was not correctly modeled in this case as the data showed single time constant behavior up to 480 h but the data was fit with a model for three time constants.
Figure 5.20 shows oxide capacitances for non-vanadate hydrotalcite pigments and controls dispersed in PVB coatings. The oxide capacitances of all pigments were observed to increase and generally to drift upward over the course of experimentation, except for neat PVB. A similar trend was observed in Figure 5.11 and may be associated with hydration of the oxide layer.
Figure 5.21 shows the capacitances for the PVB layer of coatings pigmented with non-vanadate hydrotalcites and controls. All PVB layers had capacitances within an
order of magnitude of each other, 2 x 10-10 F.cm-2 to 2 x 10-9 F.cm-2, and had similar
magnitude as observed for vanadate hydrotalcites in Figure 5.12. These results do not
suggest water uptake or that PVB is highly permeable.
191 5.3.6 Salt Spray Exposure of Non-Vanadate Hydrotalcite Pigmented Coatings and
Controls
Figure 5.22 shows Al 2024 panels coated with PVB pigmented with non-vanadate hydrotalcites and controls. Panels were scribed and exposed to salt spray for 750 h. The same subjective ranking system, as has been previously discussed, for blister and scribe protection was used for comparison. A number of tested pigments including HTMo,
HTW, HTB, and the PVB were not observed to blister in contrast to vanadate hydrotalcite pigments. However, none of the tested pigments or controls was observed to provide any protection to the scribed area. Additionally, NaVO3 arguably displayed the
worst overall performance in salt spray, with extensive blistering and no protection of the
scribe. Furthermore a number of panels demonstrated a halo effect around the scribe which is clearly seen on the HTMo, HTCl-, and PVB coated panels in Figure 5.22. The
coating appears to be intact but discolored in these regions. Further, smaller areas of the
same panels show similar discoloration on what would have been the top edges of the
panels during salt spray exposure, near where the coating meets the electrical tape used
for masking. This effect may result from direct electrolyte contact with the PVB oxide
interface as would occur in the scribe.
5.4 Discussion
5.4.1 Influence of Vanadate Speciation and Concentration on Vanadate Hydrotalcite
Pigment Performance
The best combination of scribe protection without observed blistering from panels
exposed salt spray was observed from HTV3, a decavanadate-free synthesis, and HTV6,
192 which was synthesized with Ni rather than Zn. HTV3 was synthesized to test whether the
type of vanadate species intercalated into a hydrotalcite had an effect on inhibition. It is
known that decavanadate hydrotalcites provide inhibition to aluminum alloys (22, 23,
43). However, it has also been reported that decavanadates are not particularly good
corrosion inhibitors, although given enough time and correct solution conditions they will
speciate into tetrahedral forms, which have been shown to be strong inhibitors of oxygen
reduction (21, 51). The performance of HTV3 compared to decavanadate hydrotalcites in
this study may be an indication that inhibition is improved and perhaps more rapid when
inhibition is not dependent on the decomposition of decavanadates into inhibiting
tetrahedral coordinated species.
Further, both HTV3 and HTV6 were shown to release relatively low
concentrations of V compared to other vanadate hydrotalcite pigments. These two pigments released between 0.0023 M and 0.0053 M total V at 290 h and had the peculiar
behavior of slightly decreasing solution concentrations after approximately 30 h. It is unclear whether vanadate is being readsorbed into the hydrotalcite or this is the result of experimental error associated with sampling and massing samples for NAA. Low concentrations of vanadate in solution promote the speciation and availability of tetrahedral species, and in particular, simple tetrahedral species, while high concentrations promote the formation of larger more complex vanadate molecules (39).
3- 2- - Species such as VO4 , VO(OH) , VO2(OH)2 , and VO(OH)3 , which have been linked to
corrosion inhibition, become increasingly predominant and are available over a wider
range of pH values (approximately pH 3.5 to 14) in dilute concentrations (>0.0001 M)
compared to more concentrated solutions (21, 39, 51). It may be the case that, as long as
193 vanadates are available in solution above some critical concentration, they are more
effective in lower concentrations than higher concentrations because low concentrations
promote the formation of tetrahedrally coordinated species capable of delivering
inhibition. The critical concentration for inhibition has been reported to be 0.05 mM
monovanadate, with significant inhibition starting at concentrations as low as 0.007 mM
(52). Further, there appears to be a loose correlation between the degrees of blistering
observed in Figure 5.13 and the concentration of vanadate released in Figure 5.15.
Pigments that released high concentrations, above approximately 0.001 M total
vanadium, were typically observed to blister more than pigments that released lower
concentrations of vanadium.
5.4.2 Possible Influence of Cation Inhibitors Released into Solution
Another factor that has been reported to affect overall inhibition from hydrotalcite
pigments is the release of cations like Zn2+ into solutions. Zn2+ in concentrations from
0.05 M to 0.15 M has been shown to decrease oxygen reduction kinetics on Al 2024-T3
in NaCl solutions (22). In this work, significant concentrations of both Ni and Zn were
observed to be released from some hydrotalcite pigments, 0.0022 M and 0.0045 M,
respectively. In particular, during salt spray exposure, HTV6, which was synthesized
with Ni rather than Zn, blistered less than all pigments synthesized with Zn, except
HTV3, while providing similar or better scribe protection. This may indicate that hydrotalcites that contain Ni provide additional inhibition compared to hydrotalcites with
Zn. Or this may simply mean that Ni was less soluble which resulted in less blistering.
194 Conversely, vanadate hydrotalcite pigments synthesized with Li and Mg cations were
observed to perform similarly to pigments containing Zn.
Synthesis using chloride salts was not observed to have an adverse effect on
inhibitor performance. A number of vanadate hydrotalcite pigments were observed to
perform as well or better than HTV4 which was synthesized using nitrate salts and nitric
acid to avoid exposure to chlorides during synthesis.
5.4.3 Effect of Solubility on Vanadate Pigment Performance
Solubility issues appear to play a major role in observed inhibition from vanadate- bearing hydrotcites. NaVO3 dissolved into NaCl solutions has been shown to be able
inhibit oxygen reduction on Al alloys (21). However, NaVO3 dispersed directly into
PVB resulted in blisters and was not observed to provide scribe protection. Further,
HTV2, which possibly contained highly soluble vanadate salts in addition to hydrotalcite,
was observed to blister excessively, while providing minimal scribe protection. The best
performing pigments were those that released relatively low concentrations of V into
solution. The availability of vanadate in a coating is not enough to ensure inhibition. If a
pigment is too soluble not only will the coating blister, but released solution vanadates
will not be favorably speciated for inhibition. To be effective pigments, it appears that
vanadate must be released in low concentrations to prevent blistering and promote
tetrahedrally coordinated species.
195 5.4.4 Non-Vanadate Hydrotalcite Pigments Generally Do Not Provide Inhibition
One of the goals of this work was to survey a number of other possible inhibitor anions intercalated into hydrotalcite pigments to determine if they could be used to
inhibit corrosion of Al. Although it is possible to synthesize pigments with a variety of
anions, and many that do not appear to cause blistering, none provide similar levels of
scribe protection as observed from the vanadate hydrotalcites. The total impedance from
hydrotalcites containing carbonate was observed to approach values for SrCrO4 pigmented coatings.
5.5 Conclusions
1. A framework was developed by which synthetic hydrotalcites could be tailored to
incorporate different cations into host skeletons and anion into hydrotalcite
interlayers. By following a few simple rules and staying within certain
boundaries it is possible to create hydrotalcites with numerous cation and anion
permutations.
2. Generally vanadate hydrotalcites performed well, but not compared to a SrCrO4
standard. Vanadate hydrotalcites were observed to have total impedances within
one order of magnitude of SrCrO4 and provide at a minimum some scribe
protection, although these pigments had a tendency to blister.
3. Vanadate hydrotalcites synthesized without decavanadate were shown to perform
as well as or better than hydrotalcite pigments intercalated with decavanadate.
There is possible evidence that intercalation of vanadates other than decavanadate
into hydrotalcites may result in additional corrosion resistance. Further, the best
196 performing pigments were those that released relatively low concentrations of
vanadate into solution. Low solution concentrations promote vanadate speciation
into tetrahedral species which previous studies have noted to inhibit oxygen
reduction.
4. Non-vanadate hydrotalcite pigments were observed to have good resistance to
blistering, but showed little ability to protect scribed areas on Al panels exposed
to salt spray.
197 FIGURES
Figure 5.1: XRD patterns of hydrotalcites synthesized with interlayer vanadates.
198
Figure 5.2: XRD patterns of hydrotalcites synthesized to contain anions other than vanadate.
199
(a)
(b)
Figure 5.3: An explicit equivalent circuit model of the coating system (a) and the simplified model used to fit data from PVB coated Al 2024-T3 (b).
200
Figure 5.4: Bode magnitude and phase angle plots of HTV3 pigmented PVB coatings subject to static exposure in 0.5 M NaCl. The area exposed was approximately 8.55 cm2.
201
Figure 5.5: Bode magnitude and phase angle plots of SrCrO4 pigmented PVB coatings subject to static exposure in 0.5 M NaCl. The area exposed was approximately 8.55 cm2.
202
Figure 5.6: Bode magnitude and phase angle plots of HTV3 pigmented PVB coating after 4 days exposure in 0.5 M NaCl and the corresponding data fit produced using the simplified equivalent circuit. The area exposed was approximately 8.55 cm2.
203
Figure 5.7: Nyquist plot of HTV3 pigmented PVB coating after 4 days exposure in 0.5 M NaCl and the corresponding data fit produced using the simplified equivalent circuit. The area exposed was approximately 8.55 cm2.
204
1011
HTV1 1010 HTV2 HTV3 HTV4 9 HTV5 ) 2 10 HTV6 HTV7 8 SrCrO 10 4
107
106
5 Total Impedance (Ohms.cm Impedance Total 10
104
103 0 200 400 600 800 1000
Time (h)
Figure 5.8: Total impedance of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl.
205
1010
109
8
) 10 2
107
106
105 HTV1 HTV2 4 HTV3 Pore Resistance (Ohms.cm Resistance Pore 10 HTV4 HTV5 3 HTV6 10 HTV7 SrCrO 4 102 0 200 400 600 800 1000
Time (h)
Figure 5.9: Pore resistances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl.
206
10-3
10-4
10-5 ) 2
10-6
-7 HTV1 10 HTV2 HTV3 -8 HTV4 10 HTV5 HTV6 HTV7 Defect Capacitance (F/cm Capacitance Defect -9 SrCrO 10 4
10-10
10-11 0 200 400 600 800 1000
Time (h)
Figure 5.10: Defect capacitances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl.
207
10-4
10-5
-6 )
2 10
10-7
10-8 HTV1 HTV2 -9 HTV3
Oxide Capacitance Oxide(F/cm Capacitance 10 HTV4 HTV5 HTV6 -10 HTV7 10 SrCrO 4
10-11 0 200 400 600 800 1000
Time (h)
Figure 5.11: Oxide capacitances of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl.
208
10-8
HTV1 HTV2 HTV3 HTV4 HTV5 HTV6 ) 2 HTV7 SrCrO 4
10-9 PVB Capacitance (F/cm
10-10 0 200 400 600 800 1000
Time (h)
Figure 5.12: Capacitances associated with the PVB layer of PVB coatings pigmented with vanadate hydrotalcites and SrCrO4 subject to static exposure in 0.5 M NaCl.
209
Figure 5.13: Al 2024 panels coated with vanadate hydrotalcite pigmented PVB. Panels were scribed and exposed to salt spray for 750 h. Relative blister (B) and scribe protection (SP) performance are indicated at the bottom of each panel with “1” indicating strong performance.
210
Figure 5.14: Close up image of SrCrO4 and HTV3 panels after 750 hours of salt spray exposure.
211
10-1
10-2
10-3
Concentration (M) HTV1 -4 10 HTV2 HTV3 HTV4 HTV5 HTV6 HTV7 -5 SrCrO 10 4
0 50 100 150 200 250 300
Time (h)
Figure 5.15: Vanadate release from vanadate hydrotalcite pigments exposed to 0.5 M NaCl compared to SrCrO4 release using NAA.
212
10-2
-3 10 HTV1-Zn HTV2-Zn HTV6-Ni Concentration (M)
10-4 0 50 100 150 200 250 300
Time (h)
Figure 5.16: Release of Zn and Ni from vanadate hydrotalcite pigments exposed to 0.5 M NaCl collected using NAA with long counting times or long sample decay times.
213
1011
1010
9 )
2 10
108
107
106
5 Total Impedance (Ohms.cm Impedance Total 10 HTB HTMo HTW 4 - HTP NaVO 10 HTCl 3 2- HTSi PVB HTCO 3 103 0 200 400 600 800 1000
Time (h)
Figure 5.17: Total impedance at low frequency of PVB coatings pigmented with non- vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl.
214
1010
109
8
) 10 2
107
106
105
4 Pore Resistance (Ohms.cm Resistance Pore 10 HTB HTMo HTW
- HTP NaVO 103 HTCl 3 2- HTSi PVB HTCO 3 102 0 200 400 600 800 1000
Time (h)
Figure 5.18: Pore resistances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl.
215
10-3
10-4
10-5 ) 2
10-6 HTB -7 - 10 HTCl HTCO 2- 3 10-8 HTMo HTP HTSi
Defect Capacitance (F/cm Capacitance Defect -9 HTW 10 NaVO 3 PVB 10-10
10-11 0 200 400 600 800 1000
Time (h)
Figure 5.19: Defect capacitances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl.
216
10-4
10-5
-6 )
2 10
10-7
10-8
-9
Oxide Capacitance (F/cm Oxide Capacitance 10 HTB HTMo HTW
-10 - HTP NaVO 10 HTCl 3 2- HTSi PVB HTCO 3 10-11 0 200 400 600 800 1000
Time (h)
Figure 5.20: Oxide capacitances of PVB coatings pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl.
217
10-8 HTB HTMo HTW
- HTP NaVO HTCl 3 HTSi PVB HTCO 2- 3 ) 2
10-9 PVB Capacitance (F/cm
10-10 0 200 400 600 800 1000
Time (h)
Figure 5.21: Capacitances associated with the PVB layer in a PVB coating pigmented with non-vanadate hydrotalcites and controls subject to static exposure in 0.5 M NaCl.
218
Figure 5.22: PVB coated Al 2024 pigmented with non-vanadate hydrotalcites. Panels were scribed and exposed to salt spray for 750 h. Relative panel blister (B) and scribe protection (SP) performance is indicated at the bottom of each panel with “1” indicating strong performance.
219 TABLES
Table 5.1: Synthesis details and salt spray exposure results for PVB coated panels pigmented with hydrotalcites and various controls. For rankings of blistering and scribe protection a “1” indicates the best performance. 220 REFERENCES
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223
CHAPTER 6
CONCLUSIONS AND FUTURE WORK
6.1 Conclusions
In this study, inhibition from aqueous vanadates was characterized as a function of vanadate concentration and pH to identify which species result in inhibition. Further,
the effect of vanadates on specific intermetallic particles commonly found in Al 2024
was characterized. A framework was developed to synthesize hydrotalcites containing a
number of different inhibitor anions within a variety of different cation sheets. The
performance of these pigments in organic coatings was compared to that of a SrCrO4
standard. The following is a summary of important findings:
1. Inhibition of Al 2024-T3 in NaCl solutions by vanadates is associated with
tetrahedrally coordinated forms of vanadate which predominate in alkaline and
dilute mildly acidic solutions. Octahedrally coordinated vanadates do not appear
to provide inhibition and may accelerate corrosion under deaerated conditions.
224 2. Vanadates result in a modest in increase in pitting potential independent of
aeration. However, the primary action of vanadate inhibitors is through the
suppression of oxygen reduction kinetics.
3. Specifically, tetrahedrally coordinated vanadates inhibit the ability of
intermetallic particles and the matrix of Al 2024-T3 to support cathodic reactions
in alkaline NaCl solutions. Additionally, tetrahedrally coordinated vanadates
decreased corrosion potential and generally increased pitting potential.
4. The overall reduction of cathodic kinetics on intermetallics and the matrix of Al
2024-T3 in alkaline vanadate solutions shifts the OCP below the potential
required to cause Al2CuMg breakdown. This may prevent selective Mg
dissolution from Al2CuMg and subsequent formation of Cu-enriched clusters
capable of supporting rapid cathodic reaction kinetics. SEM-EDS showed that
intermetallic particles containing magnesium to be largely intact after exposure to
NaCl solutions containing tetrahedrally coordinated vanadates.
5. Synthetic hydrotalcite pigments can be synthesized to contain a number of
different inhibitor anions with various cationic skeletons.
6. Vanadate hydrotalcites were shown to be inhibitors of corrosion. In particular,
vanadate hydrotalcite pigments synthesized without decavanadate and those that
resulted in release of relatively low concentrations of vanadium in solution were
the best inhibitors. Low concentrations of vanadium in solution may promote the
presence of tetrahedrally coordinated vanadate species which were shown to
inhibit corrosion.
225 6.2 Future Work
This work has helped further develop an understanding of vanadate inhibition and create a platform by which tailored hydrotalcite pigments may be synthesized. However, there are a number of issues that yet remain unresolved and new questions that have arisen over the course of this work.
1. It has been suggested that the decrease in oxygen reduction from vandadates is the
result of adsorbed species rather than reduction of soluble species as occurs with
chromates. Surface analysis of aluminum exposed to vanadate solutions could
reveal which specific tetrahedral vanadate species bind to the surface to reduce
rates of oxygen reduction. Raman spectroscopy has been used to characterize the
surfaces of aluminum exposed to Cr6+ solutions. Following a similar procedure,
the specific vanadate species that are absorbed to the surface to stifle oxygen
reduction might be identified.
2. Work with hydrotalcites has shown that they may be used as chloride sensors
when incorporated in a polymeric matrix. It may be possible that hydrotalcite
inhibitor pigments, dispersed in an organic coating, could be used as sensors to
provide rapid in-situ concentration measurements of contacting electrolytes.
Preliminary work in this direction is promising; however, this technology has not
currently been fully developed.
3. Vanadate hydrotalcites were shown to be inhibitors of Al corrosion. However, all
tested pigments were observed to lead to blistering of coated panels exposed to
salt spray. Future work to decrease the level of observed blistering in HTV
pigmented coatings would likely improve performance. Additionally,
226 hydrotalcite pigments have not been tested in commercial prototype coating
formulations and evaluation in these coatings is a logical next step.
4. Most studies using hydrotalcites as inhibitors have been focused on inhibition of
aluminum alloys. This work could be expanded to evaluate the performance of
various hydrotalcites on other substrate metals. Perhaps pigments containing
inhibitors such as molybdates, phosphates, and borates would prove to be more
effective corrosion inhibitors on Fe.
227
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