Materials Corrosion and Protection 9783110310054, 9783110309874

Citation preview

Yongchang Huang and Jianqi Zhang (Eds.) Materials Corrosion and Protection

Also of interest Series: Advanced Composites. J. Paulo Davim (Ed.) ISSN 2192-8983 Published titles in this series: Vol. 6: Wood Composites (2017) Ed. by Aguilera, Alfredo/Davim, J. Paulo Vol. 5: Ceramic Matrix Composites (2016) Ed. by Davim, J. Paulo Vol. 4: Machinability of Fibre-Reinforced Plastics (2015) Ed. by Davim, J. Paulo Vol. 3: Metal Matrix Composites (2014) Ed. by Davim, J. Paulo Vol. 2: Biomedical Composites (2013) Ed. by Davim, J. Paulo Vol. 1: Nanocomposites (2013) Ed. by Davim, J. Paulo/Charitidis, Constantinos A. Metals and Alloys Industrial Applications Benvenuto 2016 ISBN 978-3-11-040784-6, e-ISBN 978-3-11-044185-7

Corrosion for Engineers Zander, Dietzel, Atrens 2018 ISBN 978-3-11-028131-6, e-ISBN 978-3-11-028132-3

Thermophysical Properties of Multicomponent Liquid Alloys Brillo 2016 ISBN 978-3-11-046684-3, e-ISBN 978-3-11-046899-1

Materials Corrosion and Protection Edited by Yongchang Huang and Jianqi Zhang

Editors Prof. Yongchang Huang, Shanghai Jiao Tong University, No. 800 Dongchuan Road, Shanghai, China [email protected] Prof. Jianqi Zhang, Inner Mongolia University of Science&Technology, No. 7 Inner Mongolia Arding Street, Baotou, Inner Mongolia, China [email protected] Translated by Prof. Yongchang Huang (Chapter 1,2,3,8,9,11), Prof. Shaozong Zhang (Chapter 5,6), Prof. Guangfu Li (Chapter 7,10) Prof. Yuanwei Huang (Chapter 4), and Dr Haiyan Yang (Chapter 10).

ISBN 978-3-11-030987-4 e-ISBN (PDF) 978-3-11-031005-4 e-ISBN (EPUB) 978-3-11-038295-2 Library of Congress Cataloging-in-Publication Data A CIP catalog record for this book has been applied for at the Library of Congress. Bibliographic information published by the Deutsche Nationalbibliothek The Deutsche Nationalbibliothek lists this publication in the Deutsche Nationalbibliografie; detailed bibliographic data are available on the Internet at http://dnb.dnb.de. © 2018 Shanghai Jiao Tong University Press, Shanghai and Walter de Gruyter GmbH, Berlin/Boston Typesetting: Compuscript Ltd., Shannon, Ireland Printing and binding: CPI books GmbH, Leck Cover image: Kittikorn / iStock / Getty Images Plus ∞ Printed on acid-free paper Printed in Germany www.degruyter.com

Preface The destruction or deterioration of a metal and its properties through chemical or electrochemical action to the surrounding environment is defined as metal corrosion. Since humans began using metal, this phenomenon has been found. Consequently, people took action to prevent metal corrosion consciously. For example, Herodotus, an ancient Greek historian, and Plinius, an ancient Roman natural scientist, have used tin to prevent the corrosion of iron. In China, Sn-bronze has appeared as early as the Shang dynasty. The excavated swords of kings Wu and Yue, which were cast in Sn-bronze, are as good as new ones after having been buried under the earth for 2600 years. Since the 1950s, with the wide use of nonmetal composite materials has received more attention as the corrosion of these materials increases quite a lot. Corrosion of nonmetal material includes aging, checking, degradation, and breaking of nonmetal material due to the effect of chemical and mechanical reaction and physical force in the environment. This book is an attempt to mainly present metal corrosion and protection and to give a brief introduction to the corrosion of nonmetal materials and at the same time make our readers well informed of material corrosion and protection.

https://doi.org/10.1515/9783110310054-202

Contents Preface 

 v

List of contributing authors 

 xv

Yongchang Huang 1 Introduction 1 1.1 Severity and damage of corrosion 1 1.2 Aim of studying corrosion and protection measures 3 1.3 Definition 3 1.3.1 Corrosion 3 1.3.2 Corrosion subject 4 1.4 Classification of corrosion and destructive types of corrosion 1.4.1 Classification of corrosion 5 1.4.2 Electrochemical corrosion 6 1.4.3 Chemical corrosion 8 1.4.4 Destructive forms of metal corrosion 10 1.5 Metal corrosion level expression 14 1.5.1 Corrosion rate expression 14 1.5.2 Evaluation standard of metal and alloy corrosion resistance 1.6 Protection of metal corrosion 17 1.6.1 Isolation control method 17 1.6.2 Thermodynamics control method 18 1.6.3 Kinetics control method 19 1.7 Development and prospect 20 References 23

5

16

Yongchang Huang 2 Thermodynamics of materials corrosion 25 2.1 Thermodynamic criteria of corrosion 25 2.2 Electrode potential 26 2.2.1 Electric double-layer structure and electrode potential 26 2.2.2 Equilibrium electrode potential and Nernst equation 28 2.2.3 Standard electrode potential and electromotive force series 29 2.2.4 Metal’s thermodynamic stability of in aqueous solution 31 2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution 32 + 2.3.1 E-pH diagram of the ​​H​2O ​​ -​​H​2-​​ O ​​ ​2-​​ H ​​ ​ -​​ O ​​ H​−​​ system 33 2.3.2 E-pH diagram of a metal-​​H​2O ​​ system 34 2.3.3 Applications of E-pH diagram in the field of metal corrosion 37

viii 

 Contents

2.4 Constitution and classification of corrosion cell 2.4.1 Constitution of corrosion cell 41 2.4.2 Types of corrosion cells 46 2.5 Electrode process of corrosion cell 52 2.5.1 Anodic process 52 2.5.2 Cathodic process 55 References 57

40

Yongchang Huang 3 Electrochemical corrosion kinetics 59 3.1 Corrosion potential and its polarization diagram 59 3.1.1 Corrosion potential 59 3.1.2 Polarization effect and corrosion polarization diagram 62 3.2 Polarization curves of the corroded metal’s electrodes 70 3.2.1 Polarization curves of the corroded metal’s electrodes in the double-electrode system 70 3.2.2 Polarization curves for the corroded metal electrode in a multielectrode system 74 3.3 Corrosion process controlled by the activation polarization 78 3.3.1 Electrochemical polarization equation 78 3.3.2 Hydrogen depolarization corrosion 86 3.4 Corrosion process controlled by concentration polarization 93 3.4.1 Steady-state diffusion equation 93 3.4.2 Oxygen depolarization corrosion 98 References 107 Yuanwei Huang 4 Oxidation and hot corrosion of metals and alloys 109 4.1 Introduction 109 4.2 High-temperature oxidation 110 4.2.1 Thermodynamics of metal oxidation reaction 110 4.2.2 Drawing and application of the figure of oxide standard resultant free energy versus temperature 112 4.2.3 Metal oxidation dynamics and mechanism 115 4.2.4 Oxidation of alloy 125 4.2.5 Oxide film and its basic properties 132 4.3 High-temperature alloy materials 140 4.3.1 Strengthening of alloy, high-temperature resistance principle, and classification of high-temperature materials 140 4.3.2 Nickel base high-temperature alloy 151 4.3.3 Cobalt base high-temperature alloy 160

Contents 

Iron base high-temperature alloy 4.3.4 Other high-temperature materials 4.3.5 References 178

168 171

Jianqi Zhang Uniform corrosion and passivation in metallic materials 5 181 5.1 Introduction 181 Uniform corrosion 5.2 181 5.2.1 Characteristics of uniform corrosion 181 5.2.2 Measurement of uniform corrosion rate 182 5.3 Passivation 192 5.3.1 Metal passivation and passive film 192 5.3.2 Active-passive corrosion behavior 195 5.3.3 Determination and measurement of passivation 202 5.3.4 Influencing factors on passivation 203 5.3.5 Properties of passive film 209 References 217 Jianqi Zhang 6 Localized corrosion in metallic materials 219 6.1 Introduction 219 6.1.1 Microcorrosion cell – electrochemical inhomogeneity 219 6.1.2 Material inhomogeneity 220 6.1.3 Environmental inhomogeneity 224 6.2 Galvanic and concentration cell corrosion 227 6.2.1 Galvanic couple and galvanic series 227 6.2.2 Galvanic effects 229 6.2.3 Determination and evaluation of galvanic corrosion rate 231 6.2.4 Concentration cell 232 6.3 Pitting corrosion 233 6.3.1 Pitting corrosion mechanisms 234 6.3.2 Influencing factors in pitting corrosion 236 6.3.3 Determination of pitting corrosion 238 6.3.4 Avoiding pitting corrosion 240 6.4 Crevice corrosion 240 6.4.1 Crevice corrosion mechanism 240 6.4.2 Influencing factors in crevice corrosion 242 6.4.3 Determination of crevice corrosion 245 6.4.4 Avoiding crevice corrosion 246 6.5 Selective corrosion 246 6.5.1 Selective corrosion principles 247

 ix

x 

 Contents

6.5.2 Selective corrosion systems 247 Influencing factors in selective corrosion 6.5.3 248 Determination of selective corrosion 6.5.4 248 Protection of selective corrosion 6.5.5 248 Intergranular corrosion 6.6 249 Intergranular corrosion mechanism 6.6.1 249 Intergranular corrosion system 6.6.2 249 Influencing factors in selective corrosion 6.6.3 251 6.6.4 Determination of intergranular corrosion 252 6.6.5 Avoiding intergranular corrosion 255 6.7 Exfoliation corrosion 255 6.7.1 Exfoliation corrosion behaviors and reasons 255 6.7.2 Avoiding exfoliation corrosion 255 6.8 Microbial corrosion 256 6.8.1 Spices and functions of microorganisms 256 6.8.2 Microbial corrosion in stainless steel 257 6.8.3 Microbial corrosion in other metallic materials 257 6.8.4 Determination and control of microbial corrosion 257 References 258 Guangfu Li 7 Stress corrosion cracking of metals 259 7.1 Introduction 259 7.2 Characteristics of SCC 261 7.2.1 Mechanical characteristics 262 7.2.2 Environmental characters 264 7.2.3 Material characters 265 7.3 Mechanisms of SCC 267 7.3.1 SCC mechanism of anodic dissolution 268 7.3.2 SCC mechanism of hydrogen-induced cracking 7.4 Test/Evaluation of SCC 272 7.5 Control of SCC 275 References 276

270

Yongchang Huang 8 Corrosion and protection of metalsin the natural environment and chemical medium 277 8.1 Corrosion in aqueous environment 277 8.1.1 Mechanism of corrosion in aqueous solutions 277 8.1.2 Characteristics of seawater corrosion 279 8.1.3 Factors impacting seawater corrosion 282 8.1.4 Control of the seawater corrosion 287

Contents 

 xi

288 Atmospheric corrosion 8.2 Classification of atmospheric corrosion and its mechanism 8.2.1 288 Corrosion product film 8.2.2 291 Factors affecting atmospheric corrosion 8.2.3 293 Protection measures 8.2.4 298 Soil corrosion 8.3 299 Mechanisms of soil corrosion 8.3.1 300 Characteristics and impacting factors of soil corrosion 8.3.2 301 8.3.3 Stray current corrosion 306 8.3.4 Soil corrosion controlling methods 310 8.4 Chemical engineering medium corrosion 311 8.4.1 Types of chemical engineering medium 311 8.4.2 Corrosion of sulfur and sulfide 313 8.4.3 Corrosion of chlorine and chloride 315 8.4.4 Corrosion of nitric acid 316 8.4.5 Corrosion of sodium hydroxide and alkali solution 317 8.4.6 Corrosion of urea solution 318 8.4.7 Corrosion of refining oil and petrochemical media 319 References 321 Hongliang Pan 9 Corrosion and protection of nonmetal material 323 9.1 Introduction 323 9.2 Corrosion of fireproof material and ceramic material 323 9.2.1 Corrosion of fireproof material 324 9.2.2 Corrosion of constructional ceramic material 326 9.2.3 Corrosion protection of fireproof material and ceramic material 9.3 Corrosion of glass 329 9.3.1 Characteristics of the glass structure 329 9.3.2 Corrosion mechanism of glass 330 9.3.3 Factors affecting glass corrosion 332 9.4 Corrosion of macromolecule material 334 9.4.1 Structure characteristics of macromolecule material 334 9.4.2 Corrosion type and mechanism 335 9.4.3 Protection of corrosion of macromolecule material 338 9.5 Corrosion of concrete 341 9.5.1 Characteristics of concrete structure 341 9.5.2 Corrosion modality and mechanism 342 9.5.3 Factors affecting corrosion 344 9.5.4 Corrosion of steel bar in concrete 345 9.6 Corrosion of timber 346 9.6.1 Structure feature of timber 346

328

xii 

 Contents

9.6.2 Corrosion modality of timber 347 Chemicals preventing timber degradation 9.6.3 References 350

348

Jianqi Zhang and Guangfu Li Corrosion of materials in high-technology field 10 351 10.1 Introduction 351 Aerospace and aviation field 10.2 351 10.2.1 Corrosion and protection of aircraft 351 10.2.2 Corrosion and protection of space shuttles 374 10.2.3 Corrosion and protection of satellites 379 10.3 Corrosion in nuclear power systems 383 10.3.1 Introduction 383 10.3.2 Statistics of cost and failure due to corrosion in nuclear power plants 387 10.3.3 Main corrosion problems and countermeasures in PWR plants 390 10.3.4 Corrosion and countermeasures of structural materials in BWR plants 399 10.4 Energy field 400 10.4.1 Corrosion of lead-acid battery 401 10.4.2 Corrosion of NI-MH battery (nickel-metal hydride battery) 403 10.4.3 Corrosion of lithium ion battery 405 10.4.4 Corrosion of fuel cells 406 10.5 Biological medicine field 414 10.5.1 Corrosion of implanted materials and corrected instruments 415 10.5.2 Corrosion of dental alloy 419 References 421 Yongchang Huang 11 Protection technology of material corrosion 425 11.1 Introduction 425 11.2 Corrosion preventive design 425 11.2.1 Selecting suitable material 426 11.2.2 Anticorrosion design of engineering structures 432 11.3 Protective coating 433 11.3.1 Characteristics and classification of protective coating 433 11.3.2 Metal plating and its technology 433 11.3.3 Inorganic nonmetallic coating and its technology 440 11.3.4 Nonmetallic organic coating and its technology 445 11.4 Electrochemical protection 451 11.4.1 Overview 451 11.4.2 Cathodic protection 452

Contents 

471 Anodic protection 11.4.3 Corrosion inhibitor protection 11.5 474 11.5.1 Introduction 474 Classification of inhibitor 11.5.2 476 Influencing factors 11.5.3 478 Application of inhibitor 11.5.4 479 References 482 Subject index 

 483

 xiii

List of contributing authors Prof. Yongchang Huang, Shanghai Jiao Tong University, No. 800 Dongchuan Road, Shanghai, China [email protected] Chapter 1, 2, 3, 8, 11 Prof. Jianqi Zhang, Inner Mongolia University of Science&Technology, No. 7 Inner Mongolia Arding Street, Baotou, Inner Mongolia, China [email protected] Chapter 5, 6, 10 Prof. Yuanwei Huang Shanghai Institute of Microsystem and Information Technology, Chinese Academy of Sciences, No. 865 Changning Road, Shanghai, China Chapter 4

Prof. Hongliang Pan East China University of Science and Technology, No. 130 Meilong Road, Shanghai, China Chapter 9 Prof Guangfu Li Shanghai Research Institute of Materials, No. 99 Handan Road, Shanghai, China Chapter 7, 10

Yongchang Huang

1 Introduction 1.1 Severity and damage of corrosion The phenomenon of metal corrosion is quite common. From the perspective of thermodynamics, the majority of metal has a tendency to change into metallic compounds gradually, except several noble metals (Au, Pt, etc.) due to the exposure to the environment. Metal corrosion is a process of natural deterioration, that is to say, an inverse process of metallurgy. Most accidents, such as rusting of steel and tarnishing of bronze when exposed to air, corrosion of ship hulls in seawater, perforation of gas and oil piping, destruction of boilers in thermoelectric power stations and fracture of bridges, explosion of petrochemical plants, and aviation accidents, are due to metal corrosion. The whole country pays a high price for material corrosion. The corrosion happens in numerous situations where metal and nonmetal materials are involved. There are various kinds of corrosion problems in transportation, architecture, machinery, electric power industry, petrochemical industry, metallurgy, and nation defense, as well as in daily life. Statistics show that the quantity of the metal equipments and materials scrapped annually by corrosion constitutes 1/3 of annual production. Assuming that 2/3 of metals can remelted, then the remaining metal or about 1/10 of metal material cannot be recycled because of corrosion and chalking. It is clear that metal corrosion is a great waste of limited natural resources. On one hand, corrosion wastes metal resources; on the other hand, it wastes energy and water consumed in the processing of the metal structure. In addition, the destruction of metal equipment and engineering structure caused by material corrosion and their scrappiness ahead of time gives rise to economic losses, which are worth much more than the price of material. According to the statistics of industrialized countries, the annual cost of corrosion is about 2–4% of gross domestic product (GDP). Sectors of the UK industry have conducted a special research on national corrosion loss as early as the 1960s and published a famous Hoar report in 1971 [1]. Just as the report pointed out, the annual cost of corrosion was £1.36 billion, about 3.5% of GDP. In the last several decades, corrosion loss in industrialized countries has been rising year after year. Taking the United States as an example, a corrosion scholar, Uhlig, estimated that the annual cost of corrosion in United States was about $5.5 billion in 1949, $10 billion in 1966, $75 billion in 1975, then up to $126 billion in 1982, and reached a peak of $300 billion in 1995, which was approximately 4.2% of its gross national product [2]. Using effective protection techniques then, the increasing rate got under control and the annual cost of corrosion was reduced to $275.7 billion in 1998. Table 1.1 shows the annual cost of corrosion in industrially advanced countries. https://doi.org/10.1515/9783110310054-001

2 

 1 Introduction

Tab. 1.1: Corrosion loss of several industrialized countries. Country

Year Corrosion loss

Proportion of GDP (%)

Country

Year

Corrosion loss

USA

1995 $300 billion

4.2

Japan

1997

UK

3.5 3.5 2.5

Sweden

1986

Germany

1970 £1.365 billion 1985 $17 billion 1982 £45 billion

Union of Soviet Socialist Republics

1985

¥3937.7 billion Kr35 billion ($5 billion) Rbl40 billion

India

1986 $4.82 billion

1987

$100 billion

Proportion of GDP (%) 3.0 3.5 2.0

In 1999, the Chinese Academy of Engineering officially launched the first consulting project of Investigation and Countermeasures of China’s Industrial and Natural Environmental Corrosion problems, which was organized by the Institute of Metal, Research, Chinese Academic of Science and under academician Ke Wei’s charge to conduct detailed research on material corrosion loss. The research results showed that the direct corrosion loss was 228.8 billion yuan, with indirect loss of 500 billion yuan, amounting to $62.5 billion, which accounted for 5% of the GDP in 1999 [3]. Besides, it was reported that in 2004, the corrosion loss was up to 819 billion yuan, increasing by 64% within just 5 years. The increasing rate is so surprising. It is obvious that corrosion loss exerts an extremely negative influence on China’s economy. Economic loss caused by material corrosion can be divided into direct loss and indirect loss. Direct loss includes (a) the replacement costs of the corroded equipment and structures and costs of better corrosion-resistance metals; (b) material and labor expenditures on the protective measures of painting, applying cathodic protection, coatings, and use of inhibitors; and (c) metal equipment drying and storage expenses. All the above direct losses can be calculated easily. Indirect loss includes (a) the idle time cost, (b) product loss resulting from corrosion leakage, (c) product contamination caused by corrosion, (d) loss of efficiency because of the accumulation of corrosion products and partial damage, and (e) an extra cost resulting from corrosion allowance in design, an additional thickness of structural parts. Apart from serious economic loss, corrosion accidents usually endanger human security, leading to loss of life and properties. For example, the failures of pressure vessels, boilers, aircraft components, bridges, nuclear power plant containers, and steam turbine blade can give rise to accidents. In 1967, with the collapse of Silver Bridge across the Ohio River in the United States, 31 out of 37 cars on the bridge fell down to the river, which led to 75 deaths and 9 serious injuries. In 1970, a gas explosion caused by pipelines’ corrosion break in Osaka, Japan, caused 75 deaths on the

1.2 Aim of studying corrosion and protection measures 

 3

spot. We have frequently heard about fatal airplane crashes caused by corrosion in the recent years. In 1985, a total of 500 people died in a crash of a Japan Boeing 747 because of stress corrosion on the airplane engine [4]. Now, people pay much more attention to the environmental pollution problems resulting from material corrosion. Various kinds of industrial wastewater and residues and lots of hazardous substance emitted to the atmosphere caused by corrosion do cause harm to humans’ health and break the ecological balance.

1.2 Aim of studying corrosion and protection measures In order to eliminate the large losses in economy and resources and the potential safety hazards, researching the causes of material corrosion and its countermeasures is the aim of studying corrosion and protection. The science of corrosion and protection is a comprehensive technology science that has been developed in the 1930s. And now it is an independent discipline and will be in constant development. Corrosion and protection are important disciplines involving many fields that play critical roles in a developing national economy. Choosing the proper skills and methods of corrosion control widely and correctly can prevent or slow down corrosion damage and minimize the economy loss and the negative social impact. It is generally believed that as long as the existing corrosion protection techniques are taken full advantage of, corrosion loss can be reduced by 25–30%. Correct corrosion control and prevention measures can protect public safety, prevent damage and failure of industrial equipment, protect the environment, save resources and energy, and avoid billions or even hundreds of billions loss. The discipline of material corrosion and protection consists of two researching aspects: one is to study the basic laws and mechanisms in the process of corrosion and the other is to develop corrosion controlling technology and its application.

1.3 Definition 1.3.1 Corrosion The word “corrode” is derived from the Latin word Corrodere, which means “to rot.” In China’s Ci Hai (published in 1979), the definition of corrosion is the degradation of materials occurring on the surface due to chemical or electrochemical action [5]. At that time, this definition mainly meant metal corrosion. Foreign scholars gave different definitions of corrosion. Evans regarded that metal corrosion was chemical or electrochemical transformation of metal elements from a form of single element to a combination. Moreover, Fontana claimed that metal corrosion was the reverse process of metallurgy. Uhlig considered that substance or material corrosion was

4 

 1 Introduction

damage under the effect of chemical and electrochemical action in the environment. Toma Shchev took metal corrosion as destruction of metal under the chemical or electrochemical action of surrounding environment. Cao Chunan described that metal corrosion was damage caused by condition changes into a new phase with the effect of environment. Now, material corrosion consists of metal and nonmetal corrosion, which can be defined as damage, deterioration, and degradation of material due to the effect of chemical and electrochemical action in environments.

1.3.2 Corrosion subject Xiao Jimei, China’s famous corrosion scholar, pointed out that corrosion includes microcorrosion and macrocorrosion subjects [6]. The microcorrosion subject focuses on the microanalysis of corrosion to establish corrosion theories and develop protection techniques under the guidance of these theories, that is, material corrosion and protection. However, the macrocorrosion subject aims to analyze corrosion from the perspective of overall level, namely, studying the corrosion as a whole and investigating the interaction between the corrosion and social environment, as well as the corrosion impact on economy and the society. The establishment of the corrosion subject has experienced a long process. From the mid-18th century to the early 20th century, people started to recognize material corrosion, which was a transition process from experimental stage to theories researching stage. After the 1900s, research theories on corrosion appeared and representatives include the following researches. In 1903, Whitney pointed out the essence of electrochemistry by an experiment on iron corroding in water. In 1920, Tamman proposed the parabolic law of oxidation kinetics by studying the oxidation laws of Ag, Fe, Pb, and Ni, which set up the foundation of the theories of metal oxidation [7]. In 1929, Evans established the polarization diagrams of metal corrosion, which promoted the studying of the essence of electrochemistry [8]. In 1938, Wagner and Trand established the theory of mixed potential of metal corrosion to lay the foundation of kinetics of corrosion science [9]. In the same year, Pouxbaix drew the pH-potential diagram, which laid the foundation of the thermodynamics of corrosion [10]. In 1950, Uhlig put forward the autocatalytic model of pitting ­corrosion, which promoted the development of the theories of corrosion electrochemistry [2]. In 1957, Stern and Geary presented the linear polarization technique, fostering the development of corrosion electrochemistry [11]. In 1970, Epellbain used electronic impedance spectroscopy for the first time to provide a new method of corrosion electrochemical studying [2]. In the recent two decades, corrosion theories have developed gradually, and at the same time, protection techniques have developed fully. The macrocorrosion subject is a cross-discipline subject of natural science and social science that emphasizes on the economic and social benefits of the

1.4 Classification of corrosion and destructive types of corrosion 

 5

corrosion subject. It includes corrosion education, corrosion economy (the core of macrocorrosion subject), corrosion management, etc. Spreading knowledge of corrosion science, promoting modern protection techniques, strengthening corrosion management to reduce the economic loss caused by corrosion, prolonging the service life of the materials and equipment, saving resources and energy, and protecting the environment are the most urgent priorities.

1.4 Classification of corrosion and destructive types of corrosion 1.4.1 Classification of corrosion Material includes metal and nonmetal material. Consequently, material corrosion consists of metal corrosion and nonmetal corrosion. 1.4.1.1 Metal corrosion Metal corrosion can be divided into different ways, mainly based on corrosion mechanism, corrosion environment, and the morphology of corrosion. (1) Based on corrosion mechanism, corrosion can be divided into four categories, that is, chemical corrosion, electrochemical corrosion, biological corrosion, and physical corrosion (Fig. 1.1).

Fig. 1.1: Classification of metal corrosion [4].

6 

 1 Introduction

(2) Based on corrosion environment, corrosion consists of dry corrosion, wet corrosion, anhydrous organic liquid and gas corrosion, corrosion in the molten metal, and corrosion in slag and molten salt. (3) Based on the morphology of corrosion, corrosion can be divided into general corrosion, nonstress localized corrosion, and stress localized corrosion. 1.4.1.2 Nonmetal material corrosion Compared with metal corrosion, nonmetal corrosion has rich variety and less regularity. However, in principle, it can be classified by the morphology of corrosion and mechanism of corrosion. (1) Based on the morphology of corrosion, nonmetal corrosion consists of three categories: mild corrosion, moderate corrosion, and destructive corrosion. Mild corrosion refers to bubbling, swelling, microcracking, tarnishing, and thinning. Moderate corrosion means debonding and delaminating, checking, embrittlement, hardening, and deformation. Destructive corrosion refers to dehiscence, chalking, strength loss decomposition, degradation, oxidation, and peeling. (2) According to nonmetallic corrosion mechanism, nonmetal corrosion can be divided into physical corrosion, chemical corrosion, weathering, and environmental stress cracking. Physical corrosion is the most common type of nonmetallic corrosion, mainly resulting from the interaction between osmotic diffusion of the environmental medium and the stress. Chemical corrosion is the decomposition caused by the chemical action between certain functional groups and the environmental medium. Atmospheric aging means that polymeric materials degrade under the continual effect of ultraviolet radiation, thermal radiation, air temperature, humidity, moisture, oxygen, and other factors. Environmental stress cracking refers to the cracking of organic nonmetallic materials occurring when the material fracture stress is lower than the normal level under the joint action of stress and some kinds of active medium.

1.4.2 Electrochemical corrosion Electrochemical corrosion is the corrosion caused by electrochemical reaction between metal and electrolyte solution (mostly aqueous solution). Its characteristic is that there exists two kinds of independent reaction simultaneously, anodic reaction and cathodic reaction, which generates electric current. Metal is dissolved in anodic reaction, whereas depolarizer is deoxidized in cathodic reaction. These two electrodes exist independently in time and space, which form a circuit through the solution. Corrosion products usually generate in the area between anode and cathode, which cannot cover the whole corrosion area, so they do not have the function of protection.

1.4 Classification of corrosion and destructive types of corrosion 

 7

Electrochemical corrosion is the most common corrosion. Corrosion of metal constructions in natural condition such as atmosphere, water, soil, and most mediums in the chemical industry and metallurgic process has the properties of ­electrochemical corrosion. 1.4.2.1 Atmospheric corrosion Atmospheric corrosion generally refers to corrosion of raw metal materials and their products, as well as metal parts, in the process of production, transportation, storing, and using, due to the action of atmosphere. The water in the atmosphere forms a thin conductive water film with dissolving of oxygen and impurities on the surface of metal, which is the condition for electrochemical corrosion. Atmospheric corrosion is a commonplace in daily life and in industrial and agricultural production. Steel get rusted and chalked after stacking in open air for a long time, and after household aluminum products and copper products have been used for a long time, the surface then loses the original metallic luster. Even landmark buildings like the Shanghai Exhibition Center, with its gold-filled spire, have changed their color due to long-time exposure to the atmosphere. In some cases, atmospheric corrosion can damage metal structures seriously. For example, crane frames in sulfuric acid pickling workshops have corroded to a depth of 0.5 mm each year and the paint film on it peeled off after being in use for 3 months. Atmospheric corrosion can obviously change with the changes in temperature and humidity; therefore, the rates of atmospheric corrosions around the world are quite different. According to the action of the atmosphere on metals, we can divide atmosphere into industrial, rural, and marine atmosphere and so on. 1.4.2.2 Water environment corrosion Water environment corrosion means that metal structures and components corrode in environments such as seawater, fresh water, brackish water, salt lake water, etc. Research will focus on seawater corrosion in the long-term. About 7/10 of the earth’s surface is covered by ocean, which is closely connected to humans’ lives. The number of ships sailing in the ocean is increasing, and the production of offshore oil and gas and other marine development projects and facilities have been building largely. Submarine cables, optic cable, and undersea oil pipelines are put in use. In recent years, in order not to deplete the terrestrial energy, deep-sea exploitation has been on the agenda. Seawater is natural electrolyte with high corrosivity. Regulations of the shipbuilding industry claim that the ships must return to the shipyards after sailing for 2 years to repaint and reinstall the sacrificial anodes. Besides, the steel pipes on coastal wharf have a certain service life, 30 years on average, with the auxiliary effect of coatings and cathodic protection. Although freshwater is less corrosive than seawater, we must pay enough attention as it can also result in electrochemical corrosion. Brackish seawater is the water near

8 

 1 Introduction

estuaries. Every year, seawater flows back for a period of time, raising the amount of salt in the water. The wharfs of steel mills and power plants in different regions of Shanghai have been threatened, and thus, anticorrosion measures should be taken correspondingly. 1.4.2.3 Soil corrosion The corrosion of buried metal structures and components such as large storage tanks, oil pipelines, urban underground pipelines, cables, etc., is called soil corrosion. Soil corrosion leads to the leak of oil, gas, and water, as well as communication failure, which does great harm to the people’s living and industrial production. Moreover, as the hydraulic constructions have been buried under the earth for a long time, the reinforced concrete structures can also be corroded. The soil is a special kind of electrolyte which is composed of various particulate minerals, water, gas, and microbe, and as a result, soil properties vary in different areas. So there are large differences in soil corrosion rate. In China, pipeline construction has developed very quickly in recent three decades. For example, the gas pipeline length in the West-East natural gas transmission project is 4122 km, and the pipelines go through lake marshes, saline soil areas, Gobi desert, and gravel areas. Soil environment is quite complicated, and its properties are very different. Consequently, studying soil corrosion and its anticorrosion techniques has its significance in the national economy. 1.4.2.4 Chemical medium corrosion Chemical medium corrosion is corrosion of metal components in acid, alkali, and salt solution, including corrosion in the process of oil extracting and refining. Various kinds of chemical medium are strong electrolyte with high corrosive power. As a result, when the metal components are exposed to the harsh environment, serious corrosion happens. For the different media, it is necessary to choose proper metallic material and take appropriate preventive measures accordingly.

1.4.3 Chemical corrosion Chemical corrosion is the result of oxidation-reduction reactions, which take place without generating current when metal contacts corrosive media directly. In the process of corrosion, the changing of metal atoms into the ionic state and the reduction reactions of the medium oxidants occur at the same location at the same time. Corrosion products cover the entire reaction surface and form protective films, whose properties will determine the rate of the chemical corrosion. If the integrity, strength, and plasticity of the films are great and the expansion coefficients of the metal and films are close, with a strong affinity between metals and films, metals can get

1.4 Classification of corrosion and destructive types of corrosion 

 9

well protected and the corrosion rate can slow down. The most important types of chemical corrosion are the drying gas corrosion and the high-temperature gas corrosion, such as metal oxidation and high-temperature corrosion. Furthermore, chemical corrosion includes the corrosion of metal in nonelectrolyte solutions, such as aluminum corrodes in carbon tetrachloride, chloroform, or ethanol and magnesium and titanium corrode in methanol. 1.4.3.1 Oxidation of metal and alloy Metal oxidation has two meanings, in a narrow and broad sense. Oxidation narrowly refers to the reaction of metals and oxygen to generate the corresponding oxides. However, in a broad sense, oxidation can be described as loss of electrons and elevation of metal valence. The products of oxidation include not only oxides but also sulfides, nitrides, carbides, halides, etc. In this section, we will focus on oxidation in a narrow sense. Almost all of reaction free energy changes (∆G)T,P in metal oxidation reaction exposed to the atmosphere are less than zero, which indicates that the oxidation reaction can freely perform. At high temperatures, metal oxidation is much easier. The alloy oxidation is more complex than the oxidation of the pure metal, as when an alloy contains two or more metals, one of the metals is usually preferentially oxidized. Whether it can generate a new phase or not depends on the relative concentration of each component and oxide phases. For example, in the process of oxidation of the Fe-Cr alloy, the protective film Cr2O3 comes first on the alloy surface, thereby preventing further oxidation of the alloy. In addition to selective oxidation, if the two components of the alloy have the same oxygen affinity and the oxygen partial pressure in the environment is larger than the partial pressure of the two oxides, two components of the alloy can be simultaneously oxidized. 1.4.3.2 High-temperature corrosion High-temperature corrosion is also known as hot corrosion or gas corrosion. It mainly refers to the accelerated corrosion of the alloy surface at high temperature caused by being covered with molten salt (Na2SO4) or liquid alloy oxides, or a mixture of these two substances, which usually occurs on gas turbine blades. High-temperature corrosion often happens in the range of 700 to 1000°C, which is more serious than pure oxidation at the same temperature. When the temperature is higher than 884°C (the melting point of Na2SO4), as the salts deposited on the metal surface are in a molten state, the sulfurs of salts penetrate the protective oxide film (Al2O3 or Cr2O3) to generate sulfide as a result. At the same time, the oxides begin to dissolve into the molten salt and the growth stress generates in the oxide film, which leads to the destruction of the oxide integrity. As a result, the film becomes porous and the corrosion accelerates further.

10 

 1 Introduction

1.4.3.3 Low-temperature hot corrosion Low-temperature hot corrosion is the corrosion happening at temperatures below the sodium sulfate melting point (884°C). It has been proved that the most serious low-temperature hot corrosion occurs between 700 and 750°C, when liquid nickel or cobalt sulfate generates, thereby destroying the protective oxide film. In the atmosphere of O2-SO2-SO3, CoSO4-NaSO4 or eutectic salts can form. Hot corrosion occurs in the ambient atmosphere with the presence of sulfur.

1.4.4 Destructive forms of metal corrosion According to the destructive forms of metal corrosion, we can classify corrosion into two categories, general corrosion and local corrosion. General corrosion refers to corrosion throughout the entire surface of the metal components and linked together. Local corrosion refers to the corrosion occurring only on a small part of the metal surface, while most of other regions don’t suffer from corroding substantially. Considering the metal structure, the damage of local corrosion is much more serious than general corrosion. The forms of general corrosion and several typical local corrosions are shown in Fig. 1.2.

Fig. 1.2: Destructive forms of metal corrosion [4]. 1 – Uniform corrosion. 2 – Nonuniform corrosion. 3 – Selective corrosion. 4 – Porphyritic corrosion. 5 – Ulcerative corrosion. 6 – Pitting corrosion. 7 – Intergranular corrosion. 8 – Transgranular corrosion. 9 – Internal corrosion.

1.4.4.1 General corrosion General corrosion includes two types of corrosion, uniform corrosion and nonuniform corrosion. Uniform corrosion occurs uniformly throughout the entire metal surface

1.4 Classification of corrosion and destructive types of corrosion 

 11

and the depths of corrosion on different parts are substantially the same. For example, steel corrodes in the atmosphere and silver changes its color; the high-temperature oxidation of metal belongs to this type of corrosion as well. Compared with other corrosion, uniform corrosion is least dangerous. We can control it by increasing the thickness of metal parts. However, in daily life and industrial production, nonuniform corrosion and local corrosion occur more easily. For example, after going sailing for 4 to 5 years, we can observe that the entire underwater parts of the surface of the hull have corroded unevenly. Usually, corrosion occurs on the coating defects or the smaller surface areas with mechanical abrasions. The corrosion area continues to expand and eventually forms into a potholed surface with the passage of time.

1.4.4.2 Local corrosion in stress-free conditions Local corrosion can be divided into two types: stress-free and stress corrosion. The types of local corrosion under stress-free conditions are as follows. (1) Galvanic corrosion Galvanic corrosion occurs when two different metals or alloys contact with each other directly. When they are immersed in the corrosive or conductive medium, the corrosion cells form because of the potential difference between the two coupled metals or alloys whose corrosion rates change largely. Metals or alloys with lower potential have higher corrosion rates, inducing the corrosion in contacting parts, while those metals and alloys with higher potential have lower corrosion rates, and therefore, they don’t corrode obviously or even don’t corrode at all. Galvanic corrosion often results from the improper material selection of the composition of the metal structure. For example, in a luxury yacht with a beautiful Monel (Ni70Cu20) shell and steel rivets, galvanic corrosion occurred after launching. With quick rivet corrosion, the yacht had to return to the dock to get repaired. (2) Pitting corrosion Pitting corrosion, also known as hole corrosion, refers to a local corrosion with tiny holes generated on the surface of metals. Holes can be big or small. In most cases, the diameters of the holes are only several tens of micrometers, but their depths are equal to or much larger than the diameters. Therefore, as the majority pits are covered with corrosion products, they are not easily found at an early stage. Pitting corrosion is one type of corrosion with the strongest destructive effect and hidden risk, resulting in pipe perforation and leakage of water, gas, and products. Serious accidents may lead to a plant to stop working and even turn it down. A plant in Shanghai Jinshan Petrochemical Company lost millions a day because of pitting corrosion. Stainless steel, aluminum, and aluminum alloys often have pitting corrosion in the solution containing chlorine ions, which needs much more attention.

12 

 1 Introduction

(3) Crevice corrosion Crevice corrosion usually occurs at the connection points of the metal parts, as whatever the connection method is, such as riveting, welding, bolts, etc., gaps exist in connection places. Serious local corrosion occurs at the connection once the corrosive medium gets into the crevice where the corrosive mediums are stagnant. As the oxygen of the solution is consumed, the inside area of the crevice becomes an oxygen-depleted region, which is an anode with lower potential regions. And the outside area of the crevice contains sufficient oxygen, which is a cathode with higher potential regions. In addition to the structure problems on the crevice, the presence of the other substances on the metal surface, such as sand, dust, corrosion products, and other solids, can cause crevice corrosion as well. Crevice corrosion results in the decline of metal strength, deterioration of the connection between parts, and production of additional stress. (4) Selective corrosion Selective corrosion is in close relation to two factors, components and structures. Component selective corrosion means preferential corrosion of one component in a multicomponent alloy under the effect of corrosive medium with another component concentrating on the surface of the alloy. The most common example is dezincification of brass that zinc is preferentially dissolved so that the original surface yellow color of brass becomes red with the enrichment of copper. The gray cast iron corrosion, that is, graphitizing corrosion, is the organizationally selective corrosion. The graphite in gray cast iron distributes in the ferrite with a network-like morphology, and selective dissolution of ferrite leaves the cast iron a porous graphite skeleton. (5) Intergranular corrosion Intergranular corrosion is defined as corrosion along the metal grain boundary, resulting in the loss of link and binding force between the grains. The shape of metal changes a little while mechanical properties deteriorate seriously. Usually, the atoms in the grain boundaries which are concentrated with impurity atoms are arranged loosely and disorderly, resulting in adsorption and precipitation in grain boundaries as well, that is, the grain boundary precipitation. Therefore, compared with the grain, the grain boundary that is more active may preferentially become the anode and is dissolved in the corrosive medium. Many metals and alloys have the tendency to corrode intergranularly. Intergranular corrosion is a prominent phenomenon for stainless steels and aluminum alloys in practice.

1.4.4.3 Local corrosion in stress conditions Local corrosion sometimes occurs under the combined effect of stress and corrosive medium. Usually, there are two stress sources. First is residual stress produced in

1.4 Classification of corrosion and destructive types of corrosion 

 13

the metal melting, processing, and metal parts assembling or the thermal stress generated by temperature difference and the second residual stress caused by deformation of the solid phases. Second is the stress applied externally in the different parts of the structure during the process of using. The types of local corrosion under the stress are as follows. (1) Stress corrosion cracking Stress corrosion cracking (SCC) refers to the brittle fracture of metals or alloys under the cooperative effect of a specific corrosive medium and fixed tensile stress. This means that metallic materials can break in a short time without any obvious signs in SCC. Thus, SCC is one of most dangerous local corrosions. Plane crashes, bridge collapses, explosion of buried oil and gas pipes, and the sinking of offshore oil platforms are all related to SCC. Austenitic stainless steel, copper alloy, titanium alloy, high-strength steel, and high-strength aluminum alloy are very sensitive to SCC. Once these materials get in touch with the corresponding corrosion medium, local corrosion occurs under the stress effect. (2) Corrosion fatigue Corrosion fatigue, a special type of mechanical fatigue, refers to cracking under the combined effect of repeated stress and corrosive environment. There are various forms of repeated stress, among which alternately changed stress and compressive stress are the most common types. Corrosion fatigue is the most frequent reason for the sudden destruction of safety-designed structures. For example, due to the corrosive effect of salt water in oil production wells, steel piston rods used to pump oil from underground a have very short service life. Thus, the US oil industry losses millions of dollars a year for this reason. Marine ship propellers, cable wires at mines, automobile springs, combustion engine blades, turbine rotors, and turn tables suffer from corrosion fatigue. Of course, corrosion fatigue also exerts a negative influence on the nuclear power industry and aerospace industry. (3) Hydrogen damage The corrosion damage resulting from the presence of hydrogen and the reactions between metal and hydrogen is known as hydrogen damage. The types of hydrogen damage include hydrogen embrittlement, hydrogen blistering, hydrogeninduced cracking, etc. Hydrogen embrittlement and hydrogen blistering occur as hydrogen penetrates into metals, causing a decrease in materials ductility and tensile strength or local deformation of materials. Not only the acidic condition and the welding process are contributory factors of hydrogen damage, but also the contact with petrochemical fluid can also produce it. Cracking induced by hydrogen is referred to as hydrogen cracking, that is, the brittle fracture of materials under the effect of hydrogen; oxygen-containing copper embrittlement is a representative example.

14 

 1 Introduction

(4) Erosion corrosion Erosion corrosion is an accelerated destruction caused by a relative movement between the corrosive medium and the metal surface. There is a wide variety of mediums resulting in erosion corrosion, such as gas, aqueous solution, organic system, liquid metal, and liquid containing solid particles and bubbles, in which the solid particles suspended in liquid are extremely harmful. The appearance and characteristics of erosion corrosion present like slots, grooves, corrugations, circular holes, and valleys with directionality on the smooth surface, and the direction of movement cutting into the metal surface layer is along with the fluid. Most metals and alloys are subjected to erosion corrosion, and alloys with lower hardness such as lead and copper are more susceptible to such corrosion. Industrial production equipment and components, such as ship propellers, slurry pump impeller producing phosphate fertilizer, and the inlet elbow and bend of heat exchanger and all curve systems, may suffer from different degrees of erosion corrosion during the period of working. The range of erosion corrosion is quite wide, which can be divided into impingement corrosion (turbulent corrosion), cavitation corrosion, friction vibration corrosion (fretting corrosion), friction and erosion corrosion, etc.

1.5 Metal corrosion level expression 1.5.1 Corrosion rate expression The type of methods for expressing the degree of metal corrosion depends on corrosion types. Generally, for general corrosion, it can be expressed by the average corrosion rate. 1.5.1.1 Evaluation based on weight change The degree of metal corrosion can be expressed as the weight change before and after corrosion, whether it increases or decreases. If the corrosion rate is measured as weight loss, it can be calculated from the following formula:

K− =

g0 − g1 , S0 ⋅ t

g − where ​​K​  ​​is corrosion rate, ​​ __    ​​ · h; ​​g0​ ​​ is the weight before corrosion, g; g​ ​​ 1​​ is the ​m​ 2​ weight after corrosion, g; ​​s​0​​is the surface area of specimen, m2; and t is the corrosion time, h. Obviously, this method can be used in the uniform corrosion only, when the corrosion products on the metal surface can be removed clearly. In some cases, if the corrosion products on the metal surface are all adhered to it or fall off but can be

1.5 Metal corrosion level expression 

 15

collected completely, corrosion rate can be calculated based on the specimen’s weight gain. Thus, corrosion rate can be represented by the following formula: K+ =



g2 − g0 , S0 ⋅ t

g + where ​​K​  ​​is corrosion rate, ​​ __     ​​· h; and ​​g​2​​is the weight of alloy containing corrosion ​m​ 2​ products, which should be converted to the alloy’s weight according to the components of products, g. The unit of corrosion rate calculated by mass can also be mg g expressed as ​​ ___    ​​ · day (mdd), ​​ ___     ​​· h, etc. ​cm​ 2​ ​dm​ 2​

1.5.1.2 Evaluation based on corrosion depth Weight-loss methods have a serious drawback in common; i.e., they fail to indicate the corrosion depth in the rate formula expressed. From the viewpoint of engineering application, corrosion depth or the structure’s thickness change will directly affect the service life of the components to be tested, and therefore, it has more practical significance. The formula converting weight loss of parts into corrosion depth is

𝜋=

k− 24 × 365 k− × = 8.67 × , d 1000 d

___   ;​​  ​​K​ −​​is the corrosion rate where π is the corrosion rate expressed as penetration, ​​ mm a g g expressed as weight loss, ___ ​​ ​m​ 2 ​h ;​​  and d is the density of metal, ​​ ___     .​​ ​cm​ 3​ The widely used international unit of the corrosion depth rate is mils/year (abbreviation mpy) (1 mil = 1/1000 inches); in fact, the corrosion rate of corrosion-resistant material is between 1 and 200 mpy. In this way, the corrosion rate can be expressed by an integer, avoiding the use of decimals. The corrosion rate of carbon steel and ironbased materials is approximately 0.1–1.0 mm/a, and the rate of corrosion-resistant alloy is less than 0.0025 nmm/a. In addition to the above two conventional methods, the corrosion rate of metal can also be measured by the electrochemical corrosion current density of the formula as follows:

K=

A ⋅ i corr , nF

g where A is the relative atomic mass of 1 mol metal, ​​ ___     ;​​ n is the valence number of mol mA metal cation; F is the faraday constant; and i​​ ​corr​​ is the current density of corrosion,  ​​___    .​​ ​cm​ 2​ It must be noted that the corrosion rate of metal generally varies with time. For example, the corrosion rate of carbon steel seawater decreases a little with the passage of time due to the protective effect of the corrosion products formed on its surface.

16 

 1 Introduction

This must be kept in mind when evaluating the service life of carbon-steel structure in seawater. In corrosion tests, in order to select the proper total test time to obtain a stable rate value, the relationship between the corrosion rate and time should be determined. Local corrosion rate cannot be expressed by the measure used for the rate of general corrosion. The speed of pitting corrosion is represented by pitting factor, which is the ratio of the deepest pit depth and the average corrosion depth. Local corrosion will inevitably cause a reduction in the effective cross-section of the metal structure and the stress concentration and a decline in mechanical properties. Therefore, we can adopt the mechanical properties such as tension and torsion test to assess the extent of corrosion destruction indirectly. Recently, research has found that changes on corrosion resistance could be applied to evaluate the corrosion rate of the most local corrosion.

1.5.2 Evaluation standard of metal and alloy corrosion resistance Table 1.2 lists ten-level standards to evaluate the corrosion resistance of the general corrosion of metals and alloys, and the three-level evaluation standards listed in Tab. 1.3 are also used. Tab. 1.2: Common standard of corrosion resistance for metals [12]. Classification of corrosion resistance Level of corrosion

Corrosion rate (mm/y)

I

Excellent corrosion resistance

1

II

Good corrosion resistance

2 3

0.001–0.005 0.005–0.01

III

Corrosion resistance

4 5

0.01–0.05 0.05–0.1

IV

Ordinary corrosion resistance

6 7

0.1–0.5 0.5–1.0

V

Poor corrosion resistance

8 9

1.0–5.0 5.0–10.0

VI

No corrosion-resistance

10

>10.0

Tab. 1.3: Three-level standards of corrosion resistance for metals [12].

0. Given that ​​E​0​​is the equilibrium potential of an electrochemical reaction, the equilibrium condition of an electrochemical reaction can be further written as

� 𝜐i 𝜇i = � 𝜐i 𝜇i − nF EO = 0. i

i

From chemical thermodynamics, the relationship between the chemical potential of solution components with its activity is as follows:

𝜇i = 𝜇i0 + RTlnai ,

where ​​𝜇​i0 ​ ​​is standard chemical potential whose value is equal to the chemical potential per unit activity in a given temperature and pressure. Substituting this relationship into the formula of equilibrium condition of electrochemical action can identify the relationship between the equilibrium potential and the activities of each component participating in the electrochemical actions.

2.2 Electrode potential 



 29

� 𝜐i 𝜇i0 + � 𝜐i RT lna i − nFE0 = 0 i

i

E0 =



1 RT � 𝜐i 𝜇i0 + � 𝜐 lna i nF i nF i i

E ​​ ​0o ​ ​​ is the standard electrode potential; that is, all substances involved in the reaction are at the electrode potential under the standard state (e.g., 1 mole of solute, gas is one standard atmospheric pressure). Substituting ​​E​0o ​ ​​ into the above equation, we get

E0 = E0o +

RT � 𝜐 lna i . nF i i

This is the famous Nernst equation. When substituting R = 8.314 J​​K​ −1 ​​​​ mol​ −1​​, F = 965,000 C, and T = 298K ​into the above equation and then modifying the natural logarithm into the logarithm based on 10, it is the simplified Nernst equation:

E0 = E0o +

0.05915 � 𝜐i lga i . n i

When electrochemical corrosion occurs, the actual electrode potential established in the interface of a metal and solution is E. Comparing E and E​ ​​ 0​​, it can be determined whether the electrochemical corrosion process can occur spontaneously. There are three cases: � 𝜐i 𝜇i = � 𝜐i 𝜇i − nF EO < 0 (1) E > E0 i i The electrochemical corrosion reaction can occur spontaneously. � 𝜐i 𝜇i = � 𝜐i 𝜇i − nF EO > 0 (2) E < E0 i i The electrochemical corrosion reaction cannot occur spontaneously.

(3) E = E0

� 𝜐i 𝜇i = � 𝜐i 𝜇i − nF EO = 0 i

i

The electrochemical corrosion reaction is at equilibrium. The above-mentioned equilibrium potential of electrochemical reaction, ​​E​0,​​ can be calculated by the Nernst equation, while the actual electrode potential in a corrosion reaction can only rely on measuring the electrode potential. Therefore, in order to determine the corrosion tendency of a metal in a given medium, its electrode potential must actually be measured. 2.2.3 Standard electrode potential and electromotive force series According to the discussion of the previous section, standard electrode potential refers to the electrode potential of all substances involved in the electrochemical reactions

30 

 2 Thermodynamics of materials corrosion

being at the standard state. The hydrogen standard electrode potential is defined artificially as zero. The electrode to be tested constitutes a primary cell with a SHE, that is: (−)M|Mn+ (aMn+ = 1)|H+ �a󸀠+ H = 1�, H2 (PH2 = 1)|Pt(+).



The electromotive force of this primary cell is the value of the standard electrode potential of the electrode to be tested. The symbol of electrode potential uses the following notation: when any given electrode and a SHE constitute a primary cell, if the electrode reaction is a reduction reaction, the potential is positive (for example, copper electrode), and if the electrode reaction is an oxidation reaction, the potential is negative (for example, zinc electrode). Tab. 2.1: Standard electrode potentials ​​E​eo​ ​​ for metals at 25°C [2]. Metals

Electrode reactions

Li Rb K Cs Ra Ba Sr

Li → L​​i​ ​​ + e + Rb → R​​b​ ​​ + e + K → ​​K​ ​​ + e + Cs → C​​s​ ​​ + e 2+ Ra → R​​a​ ​​ + 2e 2+ Ba → B​​a​ ​​ + 2e 2+ Sr → S​​r​ ​​ + 2e

Ca Na La Ce Y Mg Am Sc Pu Th Np Be U Hf Al Ti Zr U Ti V Mn Nb Cr

+

2+

Ca → C​​a​ ​​ + 2e + Na → N​​a​ ​​ + e 3+ La → L​​a​ ​​ + 3e 3+ Ce → C​​e​ ​​ + 3e 3+ Y → ​​Y​ ​​ + 3e 2+ Mg → M​​g​ ​​ + 2e 3+ Am → A​​m​ ​​ + 3e 3+ Sc → S​​c​ ​​ + 3e 3+ Pu → P​​u​ ​​ + 3e 4+ Th → T​​h​ ​​ + 4e 3+ Np → N​​p​ ​​ + 3e 2+ Be → B​​e​ ​​ + 2e 3+ U→U ​​ ​ ​​ + 3e  4+ Hf → H​​f​ ​​ + 4e 3+ Al → A​​l​ ​​ + 3e 2+ Ti → T​​i​ ​​ + 2e 3+ Zr → Z​​r​ ​​ + 3e 4+ U → ​​U​ ​​ + 4e 3+ Ti → T​​i​ ​​ + 3e 2+ V → ​​V​ ​​ + 2e 2+ Mn → M​​n​ ​​ + 2e 3+ Nb → N​​b​ ​​ + 3e 3+ Cr → C​​r​ ​​ + 3e

​​E​eo​ (​​ V ) −3.405 −2.925 −2.924 −2.923 −2.916 −2.906 −2.890

−2.866 −2.714 −2.522 −2.480 −2.372 −2.363 −2.320 −2.080 −2.070 −1.900 −1.860 −1.847 −1.800 −1.700 −1.662 −1.628 −1.529 −1.500 −1.210 −1.186 −1.180 −1.100 −0.913

Metals

Electrode reactions

Cd In Tl Mn Co Ni Mo

Cd → C​​d​ ​​ + e + In → I​​n​ ​​ + e + Tl → T​​l​ ​​ + e 3+ Mn → M​​n​ ​​ + 3e 2+ Co → C​​o​ ​​ + 2e 2+ Ni → N​​i​ ​​ + 2e 3+ Mo → M​​o​ ​​ + 3e

Ge Sn Pb W Fe Sn Ge Bi Sb Re As Cu Tc Co Cu Rh Tl Pb Hg Ag Rh Hg Pd

+

4+

Ge → G​​e​ ​​ + 4e 2+ Sn → S​​n​ ​​ + 2e 2+ Pb → P​​b​ ​​ + 2e 3+ W → ​​W​ ​​ + 3e 3+ Fe → F​​e​ ​​ + 3e 4+ Th → T​​h​ ​​ + 4e 2+ Ge → G​​e​ ​​ + 2e 3+ Bi → B​​i​ ​​ + 3e 3+ Sb → S​​b​ ​​ + 3e 3+ Re → R​​e​ ​​ + 3e 3+ As → A​​s​ ​​ + 3e 2+ Cu → C​​u​ ​​ + 2e 2+ Tc → T​​c​ ​​ + 2e 3+ Co → C​​o​ ​​ + 3e 2+ Cu → C​​u​ ​​ + 2e 2+ Rh → R​​h​ ​​ + 2e 3+ Tl → T​​l​ ​​ + 3e 4+ Pb → P​​b​ ​​ + 4e + Hg → H​​g​ ​​ + e + Ag → A​​g​ ​​ + e 3+ Rh → R​​h​ ​​ + 3e 2+ Hg → H​​g​ ​​ + 2e 2+ Pd → P​​d​ ​​ + 2e

​​Ee​o​ (​​ V ) −0.402 −0.342 −0.336 −0.283 −0.277 −0.250 −0.200

−0.156 −0.136 −0.126 −0.110 −0.037 +0.007 +0.010 +0.216 +0.240 +0.300 +0.300 +0.337 +0.400 +0.418 +0.521 +0.600 +0.723 +0.784 +0.789 +0.799 +0.800 +0.854 +0.987

2.2 Electrode potential 

 31

Tab. 2.1 (continued) Metals

Electrode reactions

V Ta Zn Cr Ga Te Fe

V → ​​V​ ​​ + 3e 5+ Ta → T​​a​ ​​ + 5e 2+ Zn → Z​​n​ ​​ + 2e 3+ Cr → C​​r​ ​​ + 3e 3+ Ga → G​​a​ ​​ + 3e 2+ Te → T​​e​ ​​ + 2e 2+ Fe → F​​e​ ​​ + 2e

3+

​​E​eo​ (​​ V ) −0.876 −0.810 −0.763 −0.744 −0.529 −0.510 −0.440

Metals

Electrode reactions

Ir Pt Au Au

Ir → I​​r​ ​​ + 3e 2+ Pt → P​​t​ ​​ + 2e 3+ Au → A​​u​ ​​ + 3e + Au → A​​u​ ​​ + e

3+

​​Ee​o​ (​​ V ) +1.000 +1.190 +1.498 +1.691

Arranging the values of standard potential E of the electrode reactions where all metal anodes dissolve into corresponding metal ion from small to large can get an electromotive force series that can clearly demonstrate the tendencies of all kinds of metal transforming into oxidation state and the corrosion. The metals in the front of the electromotive series are generally easier to lose electrons to be oxidized and thus are prone to corrode. The electromotive series, also known as the standard redox potential table, is shown in Tab. 2.1. But it must be noted that in many real corrosion systems, since it is impossible to meet the above equilibrium conditions, using the standard electrode potential will often bring the wrong judgment. This will be discussed in the next chapter on the electrochemical corrosion kinetics.

2.2.4 Metal’s thermodynamic stability of in aqueous solution Metal corrosion mostly occurs in aqueous solutions. The thermodynamic stability of the metal in an aqueous solution can be determined approximately by using the datum of standard electrode potential. According to the equilibrium potentials of hydrogen electrode at pH = 7 (neutral solution) and pH = 0 (acidic solution) being −0.414 V and 0 V, respectively, and the equilibrium potential of the oxygen electrode being +0.815 V and +1.23 V, respectively, all metals are classified into five groups with different thermodynamic stabilities. Group 1: quite thermodynamically unstable metal (active metal): ​​E​eo ​ ​​ < −0.414 V

The standard electrode potential of metal is less than the equilibrium potential of the hydrogen electrode, i.e., 1.23V

In all metals, Au is the most resistant to corrosion and does not corrode even in an acid solution with oxygen. However, Au can dissolve into the complexing agents with an existing oxidant, such as 2Au + 4KCN + _​​ 21 O ​​​​  ​2​​ + ​​H2​ O ​​ → KAu​​(CN)​2​​ + 2KOH. The variao tion of the Gibbs free energy for the reaction is negative, Δ​​G​ ​​ = −45KJ. As it can be seen from the above discussions, under the natural conditions, only a handful of metals, such as noble metals in groups 4 and 5, can be considered completely stable. In neutral aqueous solutions, even without the presence of dissolved oxygen, the vast majority of metals are thermodynamically unstable, with a tendency to corrode spontaneously. It should also be noted that thermodynamic stability depends not only on the nature of a metal but also the corrosive agent. Thermodynamically stable metals (group 4) may corrode in the acidic solution containing oxygen and strong oxidant. Even the completely thermodynamically stable metal Au may become a thermodynamically unstable metal in an oxidizing solution containing complexing agents as its electrode potential becomes negative. In contrast, in the liquid saturated hydrocarbon without moisture and oxygen, the most active metals become completely stable.

2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution Most of the electrochemical corrosion of metals occurs when the metal is in contact + with an aqueous solution. In addition to other ions, there are always ​​H​ ​​ and ​​OH​−​​ ions

2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution 

 33

staying in the aqueous solution. The contents of these two ions can be expressed by pH value. Therefore, the stability of metal in an aqueous solution has a relationship not only with its electrode potential but also with the pH value of the aqueous solution. When plotting the relationship between the electrode potential of the electrochemical reaction system and the pH value of the aqueous solution into a curve, we can determine directly from the curve the possibility of corrosion reaction in the given condition. This relationship curve is called a potential-pH diagram (or E-pH diagram), which was first proposed by Pourbaix, a Belgian scholar, in 1938 and is also known as Pourbaix diagram. Before discussing the E-pH diagram of metal-​​H​2​​O system, let us firstly evaluate the relationship between the stability and the potential of water.

2.3.1 E-pH diagram of the ​​H​2O ​​ -​​H2​ -​​ O ​​ 2​ -​​ H ​​ ​ -​​ O ​​ H​−​​ system +

As described above, aqueous corrosion has the essence of electrochemical reaction. + The electrochemical reactions involving ​​H​ ​​ and ​​OH​−​​are the two electrode reactions that are mostly close to metal corrosion. (1) Hydrogen electrode reaction + ​​ ​2​​ ⇔ 2​​H​ ​​ + 2e. The standard electrode potential of this electrode reaction is zero, i.e.,​​ H E​0​​ = 0, and therefore, its equilibrium potential can be expressed as

E0 �

2 𝛼 + H2 RT 𝛼H+ � ln = 0.0591 lg H . = 2F pH2 pH2 H+

Because pH = −​​lga​​H​ ​​,​ when the atmospheric pressure P​ ​​ ​H​ ​​​ is equal to 1 standard atmospheric pressure, the above equation can be rewritten as +

2



E0 �

H2 � = −0.0591pH. H+

(2) The oxygen electrode reaction 4​​OH​−​​ ⇔ ​​O2​ ​​ + 2​​H​2O ​​ + 4e. In dilute solution, water activity ​​α​H​ ​ ​O​​can be considered a fixed value. Therefore, under a certain temperature and pressure, the chemical potential of water ​𝜇​H​ ​ ​O​​in a dilute aqueous solution is also a fixed value, which can be classified together with the standard chemical potential of other substances into the standard electrode potential ​​E​0​​. Thus, the equilibrium potential of the oxygen electrode reactions can be expressed as 2

2



E0 �

OH− RT PO2 � = E o OH− + ln 4 . 0� O � O2 4F 𝛼OH − 2

34 

 2 Thermodynamics of materials corrosion

The standard electrode potential of the oxygen electrode reaction is equal to 0.401V under the condition of 25°C and ​​P​​O​ ​​​ is equal to 1 standard atmospheric pressure. Taking ​​lgα​​OH​  ​​​ = pH − 14 into consideration, the relationship between the equilibrium potential of the oxygen electrode reactions and the pH of the aqueous ­solutions is 2

−



E0 �

OH− � = 1.23 − 0.0591 pH. O2

Through the above discussions, the potential-pH diagram of H ​​ 2​ O ​​ -​​H2​ -​​ ​​O2​ -​​ H ​​ ​ -​​ O​​H−​ ​​ system can be plotted, shown in Fig. 2.1. The hydrogen electrode reaction corresponds to line a on the diagram, and the oxygen electrode reaction, to the zone above line b in the diagram. Below line a, i.e., E( H2 ) < E0( H2 ) , the hydrogen electrode reaction +

H+

H+

proceeds in the reduction direction, and H​ ​​ 2​​becomes stable. While above line b, i.e., E( H2 ) > E0( H2 ), the oxygen electrode oxidation reaction proceeds in the oxidation H+

H+

direction and ​​O​2​​ becomes stable. In other words, the region between line a and line b in the potential-pH diagram is a water stable zone. According to the equilibrium potential of metal and metal ions E0( M ) and the relative position of the stable zone Mm+

of water, the stability of metal in water can be determined. There are two cases: 1. The value of E0( M ) falls in the water stable zone where the metal does not react Mm+

with water. For example, the equilibrium potential of copper in an aqueous solution is above line a, which means copper does not corrode in aqueous solution in absence of oxygen. 2. The value of E0( M ) falls outside the water stable zone, where the metal will Mm+

react with water and the corrosion is accompanied by the evolution of hydrogen or oxygen. For example, the equilibrium potential of iron in an aqueous solution is below line a, which indicates that the iron in an aqueous without oxygen would corrode and release hydrogen.

2.3.2 E-pH diagram of a metal-​​H​2O ​​ system Some of the chemical reactions involved in the corrosion process of metals in aqueous solutions are related to the electrode potential, but not to the pH value of the solution, while some of them are related only with the pH of the solution but independent of the electrode potential, and some of them are dependent on both the electrode potential and the pH value of solution. Type I reactions are in the horizontal lines, which are parallel to the axis of pH on the E-pH diagram, and the activity variation of these reaction substances will change with the height of these horizontal lines. Type  II reactions are in vertical lines, which are parallel to the E axis on the E-pH diagram, and the activity of reaction substances will make the vertical equilibrium line move

2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution 

 35

Fig. 2.2: Potential-pH diagram of hydrogen electrode and oxygen electrode [1].

laterally. Type III reactions are in the oblique lines, which, in the E-pH diagram, are parallel to neither the pH axis nor the E axis. The equilibrium relationship of these reactions is actually the Nernst electrode relationship of the electrode. Take the E-pH diagram of the Fe-H2O system (Fig. 2.2) as an example for the discussion below. There are nine basic chemical and electrochemical reactions between the relevant substances in the Fe-​​H​2​​O system. These reactions and their equilibrium relationships are shown in Tab. 2.2. As described above, for a reaction, it is possible to make countless equilibrium lines that correspond to different activities of reaction substances. However, for clarity’s sake, we only draw the equilibrium lines corresponding to a few typical values. Therefore, in Fig. 2.3, numbers such as 0, −2, −4, and −6 marked on some lines respectively indicate that if the reaction is a heterogeneous reaction between the solution and the solid phase, the activities of the reaction substances in 0 −2 −4 −6 the corresponding solution are the equilibrium lines of 10​ ​​ ,​​ ​​10​ ,​​ ​​10​ ,​​ and ​​10​ ​​ mol, and if the reaction is a homogeneous reaction in the solution, the ratio of the activities 0 −2 −4 −6 corresponding to the two reactions are the equilibrium line of ​​10​ ,​​ ​​10​ ,​​ ​​10​ ,​​ and ​​10​ .​​ An E-pH diagram can get further simplified. When stipulating the activity of −6 metal ions or metal complexing ions in a solution with ​​10​ ​​mol as a critical condition, ​​ [1]. Tab. 2.2: Some chemical and electrochemical equilibrium reactions of Fe-​​H​2O 2+

3+

3+

​​ ⇄ ​​Fe​2O ​​​​ ​3​​ + 6​​H​ ​​ 2 2​​Fe​ ​​ + 3​​H2​ O 2+ 3 Fe ⇄ ​​Fe​ ​​ + 2e 4 5 6 7 8 9

​α​F​ e​ ​​ ​​E​0​​ = 0.771 + 0.591lg​​ ____ ​α​F​ e​   ​​​​ pH = 0.24 − __ ​​ 1 ​​lg​​α​F​ e​ ​​​ 3 ​​E0​ ​​ = −0.440 + 0.0295lg​​α​F​ e​ ​​​ ​​E0​ ​​ = 0.50 + 0.0295lg​​αHFe​ ​ O​ ​ ​​​ − 0.0886pH ​​E​0​​ = 0.728 − 0.177pH − 0.0591lg​​α​F​ e​ ​​​ ​​E​0​​ = 0.947 − 0.236pH − 0.89lg​​α​F​ e​ ​​​ ​​E0​ ​​ = −1.878 − 0.089lg​​αHFe​ ​ O​ ​ ​​​ − 0.0295pH ​​E​0​​ = −0.093 − 0.059pH ​​E​0​​ = 0.29 − 0.059pH 3+

1 ​​Fe​ ​​ ⇄ ​​Fe​ ​​ + e

+

Fe + 2​​H2​ O ​​ ⇄ HF​​O2​ −​ ​​ + 3​​H​ ​​ + 2e 2+ + 2​​Fe​ ​​ + 3​​H2​ O ​​ ⇄ ​​Fe​2O ​​​​ ​3​​ + 6​​H​ ​​ + 2e 2+ + 3​​Fe​ ​​ + 4​​H2​ O ​​ ⇄ ​​Fe​3O ​​​​ ​4​​ + 8​​H​ ​​ + 2e + HF​​O2​ ​​ + 4​​H​ ​​ ⇄ ​​Fe​3O ​​​​ ​4​​ + 2​​H2​ O ​​ + 2e + 3Fe + 2​​H2​ O ​​ ⇄ ​​Fe​3O ​​​​ ​4​​ + 8​​H​ ​​ + 8e + 2​​Fe​3O ​​​​ ​4​​ + ​​H2​ O ​​ ⇄ 3​​Fe​2O ​​​​ ​3​​ + 2​​H​ ​​ + 2e +

2+

3+

2+

− 2

2+

2+

− 2

36 

 2 Thermodynamics of materials corrosion

Fig. 2.3: Potential-pH diagram of the Fe-​​H​2​​O system [3].

we would get a number of different zones constituted by some “boundary lines” in the E-pH diagram of the Fe-​​H​2​​O system. On one side of the “boundary lines,” the activity of the metal ion or metal complexing ion in the solution is higher than ​​10​−6​​mol, and the solid phase can be considered to be dissolved. On the other side of the “boundary lines,” the activity of the metal ion or metal complexing ion in the solution is lower than ​​10​−6​​mol, and the corresponding solid phase can be considered to be stable. The solid phases here include the metal itself and the insoluble compounds with protective effect. For the zone between the two different solids, since the activity does not change, only one straight line exists as a boundary line of two phases. The simplified E-pH diagram of the Fe-​​H​2​​O system is shown in Fig. 2.4. The E-pH diagram of the Fe-​​H​2​​O system can be divided into three basic zones: iron stable zones, iron corrosion zone, and iron passivation zone. (1) Fe stable zone (also known as immunity zone) When E < −0.62 V, in the most range of pH, Fe stays at a thermodynamically stable condition; that is, Fe does not corrode in water.

(2) Fe corrosion zone 2+ 3+ In a wide range of E > −0.62V, pH < 9, there is a stable region of ​​Fe​ ​​ and ​​Fe​ ​​ ions 2+ 3+ where Fe corrodes in water. When E > 0.771 V, ​​Fe​ ​​ions are oxidized into Fe​ ​​ ​​ ions. At a higher potential, ferric complexing ion with high valence may form. At the bottom right of the E-pH diagram, there is a small piece of iron corrosion zone, indicating that in the highly alkaline conditions (pH > 13), due to the formation of the soluble HFe​​O​2 ​ −​​ion, Fe can dissolve in water.

2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution 

 37

Fig. 2.4: Simplified potential-pH diagram of Fe-​​H​2O ​​ [3].

(3) Fe passivation zone From pH > 1.76 onward, in a large circle range of potential, the formation of an insolubly protective oxide film on the surface of Fe to separate Fe from the aqueous agent inhibits the corrosion reaction. Of course, the corrosion behavior of metal in the passivation zone differs from the thermodynamically stable region. Under the condition of passivation, the corrosion of Fe has a great relationship with the properties of oxide film formed on its surface. The more complete and dense the oxide film is, the lower the corrosion rate of iron in passivation zone is. From the above discussion, the E-pH diagram is an important tool to analyze and study the metal corrosion process in an aqueous solution. It has been widely used to explain a variety of corrosion phenomenon and study its mechanism, and it also provides the theoretical basis for the application of inhibitor and electrochemical protection. Of course, the E-pH diagram also has its limitations. First, the E-pH diagram is built on the basis of thermodynamic data, which can be able to determine only whether the metal corrosion may occur in the given conditions (potential, pH value) but is unable to answer how much the actual corrosion rate is. Second, the E-pH diagram is an electrochemical equilibrium diagram, but most of actual corrosion systems are far from equilibrium state, and moreover, the pH value of the local solution on the metal surface is often not the same as the pH value of the overall solution. Thus, sometimes, using the theoretical E-pH diagram may bring errors to judgment. To overcome the above defects, combined with research of electrochemical kinetic, the E-pH diagram, through experiments, can play a greater role in corrosion research and its control. Simplified potential-pH diagrams of 43 metal-​​H​2​​O systems are shown in Fig. 2.5. 2.3.3 Applications of E-pH diagram in the field of metal corrosion Potential-pH diagrams show the datum of different metals and their compounds when they stay in equilibrium in water. Therefore, an E-pH diagram has its wide applications

38 

 2 Thermodynamics of materials corrosion

in the field of electrochemistry, hydrometallurgy, element separation, analytical chemistry, and geological science and has become an important research tool. As mentioned above, most metals in natural conditions are thermodynamically unstable. The extent of thermodynamic instability of a metal is dependent not only on its nature but also on the corrosive agents and the external conditions. Thus, the extent of thermodynamic instability of metals can be compared qualitatively by using the equilibrium electrode potentials of metals. The thermodynamic stability datum of various metals in different aqueous solutions can be obtained by analyzing the E-pH diagrams. First, we can predict whether a metal corrodes or not in a given potential and a pH value and what the corrosion product is. Although sometimes this prediction is slightly different from that in the real condition, it would bring much convenience to the engineering design and construction, as well as the study of the corrosion science. For example, after pickling a boiler, alkali is often used to neutralize it, which may bring an “alkali crisp” phenomenon due to improper treatment. This problem can easily be explained by the E-pH diagram of the Fe-​​H​2​​O system. Fe would redissolve into the strongly alkaline solution whose pH value is higher than 13 and generate HF​​O​2− ​ ​​ ions. Such agent, coupled with the effect of stresses, makes the device crack partially. Therefore, it is necessary to control pH value not to be large, generally between 10 and 10.6, at the time of neutralization after pickling boiler. In this range, Fe may be subject to the passivation protection in the oxygenated conditions. As a neutralizing agent,

2.3 Potential–pH diagrams of electrochemical equilibrium in aqueous solution 

 39

Fig. 2.5: Simplified potential-pH diagrams of 43 metals [4].

trisodium phosphate is much safer than caustic soda. Furthermore, the stress corrosion cracking of brass can refer to the potential-pH diagram of Cu-​​H​2O ​​ -N​​H​3​​ ternary system to find out the reason. This diagram indicates that when pH < 7.3, the tendency to generate ​​Cu​2​​O, which promotes the cracking, increases. When pH = 7.1–7.3, the destruction of brass gets accelerated, and its electrode potential increases from 0.15 V up to 0.25 V when Cu is oxidized into C​​u​2​​O. While when the electrode potential is lower than −0.05 V, it is unable to generate C​​u​2O ​​ , and brass cracking will not occur at this time. An E-pH diagram can also teach us how to take protective measures to prevent the metal from corrosion. Figure 2.6 shows the E-pH diagram of the Fe-​​H​2O ​​ system as an example for discussion, assuming that Fe corrodes at the given potential and pH value, indicated by point Ο in Fig. 2.5. Making point Ο shift into the immunity zone or into the passive zone has the following paths: (1) Adjusting the pH value Figure 2.6 indicates that increasing the pH value of the solution to the range of 9 to 13 can significantly reduce or even completely prevent the corrosion of Fe. At this time, point Ο enters the passivation zone and a protective oxide film forms on the Fe surface. For some reasons (such as the solution containing a certain amount of Cl− ions), the surface cannot generate an oxide film, and the iron will continue to corrode, not get passivated. (2) Cathodic protection When the pH value is less than 9 and also not allowed to increase, it can artificially reduce the electrode potential of Fe to less than −0.62 V when point Ο enters the

40 

 2 Thermodynamics of materials corrosion

Fig. 2.6: Application of the E-pH diagram of Fe-​​H​2O ​​ system [5].

immunity zone and Fe gets fully protected. The specific method is to connect a metal with a lower electrode potential than Fe (such as Al alloy) with which to form a galvanic corrosion or link the negative terminal of an external power supply to Fe and the positive terminal to the auxiliary anode to form a current loop. People call the method that refers the corrosion system as the cathode of cell and shifts its potential in the negative direction as the cathodic protection. (3) Metal passivation Using a passivating agent to improve the oxidability of solution or using anodic polarization can enhance the potential of Fe, shifting point Ο into the passivation zone. The passivating agents commonly used include nitric acid, chromate, etc. When the anode gets polarized, the positive terminal of the external power get connected to Fe, and the negative terminal, to the auxiliary electrode, which is called anodic protection. It must be noted that the third method applies only to metals that can be passivated. Sometimes, due to insufficient amount of passivating agent or the improper control of the anode protection parameters, the protective film on metal surface is not complete, which will cause severe localized corrosion. When there is ​​Cl​​​ ions existing in the solution, there is the need to pay attention to prevent the emergence of localized corrosion. −

2.4 Constitution and classification of corrosion cell Metal corrosion in the electrolyte solution is an electrochemical corrosion process. Steels stacked in an open storage are exposed to sun and rain, and with the passage of time, the steel surface becomes rusty and a large sheet of rust peels by a

2.4 Constitution and classification of corrosion cell 

 41

gentle rub after being subjected to atmospheric corrosion. When the steel is exposed to air, an invisible layer of water film whose thickness is only about 100 to 1 μm will ​​ 2,​​ ​​SO​2,​​ ​​NH​3,​​ ​​H2​ ​​S, and some solid salts or dust in the condense on its surface. Since CO​ air can dissolve into the water film, it becomes a conductive electrolyte solution. The anode part of steel dissolves and the dissolved oxygen in the water film is reduced when electrochemical corrosion occurs. Soil has a porous particulate structure, which contains air, water, mineral salts, etc. Therefore, it is also a conductive medium. All metal pipelines buried in the ground suffer from different degrees of electrochemical corrosion. In addition, cable and water conservancy facilities are also buried in the soil for long time. Thus, soil corrosion is a common type of corrosion. Seawater is a good conductive medium, and ships at sea, various types of marine engineering structures, submarine cables, and cooling water equipment in coastal plants are subject to the seawater corrosion. As to the extent of the corrosion, these corrosions are much more serious than atmospheric corrosion and soil corrosion. The average corrosion rate in seawater is equal to 0.10–0.13 mm/y, according to the test results of 80 different carbon steels and low alloy steels. The electrochemical corrosion mentioned above is inseparable from the anodic dissolution occurring on the interface of metal and electrolyte, at the same time accompanied by the reductions of some substances in solution, which occurs on the metal surface. Therefore, the cause of the electrochemical corrosion of metal is a result of a spontaneous and short-circuited primary cell.

2.4.1 Constitution of corrosion cell Dipping the zinc slice and copper slice into a dilute sulfuric acid solution, and then connecting them together (shown in Fig. 2.7) [6], because the potential of zinc is lower and the potential of copper higher. The equilibriums of the electrode process established on their respective interfaces of the electrodes and solution are destroyed and the two electrodes respectively have the following electrode reactions. On the zinc electrode, the metal loses electrons and is oxidized

Zn → Zn2+ + 2e

On the copper electrode, the hydrogen ions in the sulfuric acid solution obtain electrons and are reduced: 2H+ + 2e → H2

The whole electrode reaction is

Zn + 2H+ → Zn2+ + H2 ↑ .

42 

 2 Thermodynamics of materials corrosion

It can be seen that after the copper and zinc cell are connected because of the dissolution of zinc slice, electrons flow along the wire to copper slice, and the current direction is from copper to zinc.

Fig. 2.7: Copper-zinc cell [1].

The cell mentioned above is essentially different  from Daniel cell that it neither is a reversible cell nor can convert the chemical energy into the electrical energy. The chemical energy released by the oxidation-reduction reaction in such cells emits in the type of heat energy and cannot be converted into electric power. If we connect the copper with zinc directly and then immerse them into sulfuric acid, the same change as the galvanic cell mentioned above will happen, as shown in Fig. 2.8. In this case, the zinc and copper can still be two electrodes of the cell, and the zinc is anode and the copper is cathode. The remanent electrons provided by the zinc dissolution flow into the copper, which contacted with zinc directly and is accepted by the hydrogen ion of solution on the copper surface. Lastly, hydrogen molecules form on the copper surface and then escape. As long as the cathode reduction reaction of hydrogen ions exists, zinc will continue to dissolve. Similar cells are called corrosion galvanic cell or corrosion cell. Therefore, the working characteristic of corrosion cell is that it is a short-circuited galvanic cell that only can lead to the destruction of metal while it does have useful work for outside. H2 bubbles

H+

H+ – Cu

H H+ H+ Zn2+

H+

–

H2SO4→2H++SO4

H+ H+ –

H

+

Zn2+ Zn+

– – – – –

–

– –

Zn2+ – – – Zn

Zn2+

Fig. 2.8: Sketch map of Zn dissolution in Cu-Zn cell [1].

2.4 Constitution and classification of corrosion cell 

 43

In the above-mentioned copper-zinc cell, replacing Cu with other metals such as Sn, Sb, or Fe also can play a role in promoting the corrosion of zinc, shown in Fig. 2.9. In the diagram, the corrosion rate of zinc is indirectly represented by the quantity of hydrogen that evolved on the Cu slice. The higher the potential difference between zinc anode and the cathode metal, the larger the corrosion rate of zinc is. In addition, the corrosion rate of zinc is also related to the hydrogen evolution potential of cathode. Industrial pure metal that is not in contact with another metal can produce a similar corrosion cell as the above-mentioned Cu-Zn cell after immersing in an electrolyte solution. For example, industrial zinc often contains small amounts of impurities, such as Fe​​Zn​7​​, whose potentials are generally higher than the potential of zinc, and therefore, the corrosion cell of the zinc as an anode and the impurities as cathode is formed. The working results of this cell is the same as that of the Cu-Zn cell, and the only difference is that there are many such minor corrosion cells on a piece of zinc. Figure 2.10 shows the dissolving of the industrial zinc in the sulfuric acid. In addition to impurities, the differences of the processing extent on the metal surface, the microstructure, or the stress, as well as the existence of grain boundaries and dislocation defects, and even the different energy state of the metal atoms are likely to produce an electrochemical inhomogeneity and therefore result into the micro-­ corrosion cell. We take the corrosion of Fe sheet in the dilute hydrochloric acid as a further example.

Fig. 2.9: Influence of different impurities on the corrosion rate of Zinc in 0.5 N sulfuric acid [7].

44 

 2 Thermodynamics of materials corrosion

Fig. 2.10: Illustration of dissolution of industrial Zinc with impurities into ​​H​2S ​​​​ O​4​​ [1].

Immersing the iron slice into the dilute hydrochloric acid, which is deoxidized, we can observe an intense chemical reaction between iron and hydrochloric acid, whose the reaction formula is as follows:

Fe + 2HCl → FeCl + H2 ↑ .

In this case, Fe dissolves continuously while it has a large quantity of hydrogen bubbles generated on the surface (Fig. 2.11), which is a typical case of metal corrosion in acidic medium. According to the theory of electrochemical corrosion, the above corrosion reactions may be decomposed into two different electrode reactions: Anodic reaction: Fe ⇐⇒ Fe2+ + 2e Cathode reaction: 2H+ + 2e ⇐⇒ H2 Due to the presence of the above-mentioned electrochemical heterogeneity on the surface of iron, which makes its surface covered with numerous micro-anode and micro-cathode, in dilute hydrochloric acid a large number of microscopic or sub­ microscopic corrosion cells form. The latter anode and cathode are very small, and even under the microscope, they are difficult to be distinguished. Thus, as the effective action area of a single cell is very small, and there is little potential drop between the anode and cathode, the entire metal surface may be considered as equipotential surface. Or it can be regarded that the boundary of anode and cathode does not exist in the microcell. Moreover, with the corrosion process going on, the positions of the micro-cathode and micro-anode alternate continuously, which causes the metal to dissolve uniformly. At the macroscopic level, the whole surface of iron corrodes in dilute hydrochloric acid with a uniform rate. Thus, when this uniform corrosion occurs, the entire surface of the iron slice may be conditionally considered as the anode and cathode, which work simultaneously [8]. In summary, the electrochemical corrosion process is inseparable from the migration of the electron on the interface of metal and electrolyte. An electron shifts from a metal with a lower potential (or a part of the same metal) to the metal with a

2.4 Constitution and classification of corrosion cell 

 45

Fig. 2.11: Electrochemical corrosion of Fe [1].

Fig. 2.12: Illustration of electrochemical corrosion process [2]. D – depolarizer.

higher potential (or another part of the same metal) and finally moves to the oxygen agent of electrolyte solution. Figure 2.12 is an illustration of a metal corroding electrochemically in the electrolyte solution, where the upper part is an anode reaction and the lower part is the cathodic reaction. Thus, it can be seen that a corrosion cell must include four integral parts: an anode, a cathode, an electrolyte solution, and a circuit, which constitute three steps of the corrosion cell. (1) Anodic process. This is a process in which metal undergoes anodic dissolution and moves into the solution as the form of the metal ions or hydration ions and, at the same time, the same quantities of electrons are left on metal. (2) Cathodic process. In this process, the electrons flowing from the anode are absorbed by the substances and the electrolyte that is capable of accepting electrons is oxidizing the substances. In metal corrosion, the electron acceptor in the solution is called depolarizer. (3) Electron flowing. The current flow in the metal relies on the electrons flowing from the anode to the cathode, while in the solution, it relies on the moving of ions, that is, the cations moving from the cathode to the anode and anions from the anode to the cathode so that the current in the electrode system constitutes a circuit. The three steps of the corrosion cell are independent of each other and are closely linked with each other. Just one step stops working, the entire corrosion process

46 

 2 Thermodynamics of materials corrosion

will stop. It should be noted that the space separation of the anode and the cathode caused by the electrochemical heterogeneity in the corrosion system is not the only reason to have electrochemical corrosion. In fact, in many instances of corrosion, the anodic process and cathodic process can occur alternately on the same metal surface with time. The only reason for metal having electrochemical corrosion is the existence of an oxidant (depolarizer) that is able to accept the excess electrons on the metal surface. If these substances are absent, even though there are different anodes and cathodes on the metal surface, electrochemical corrosion will not occur. Of course, the presence and distribution of micro-corrosion cells on metal surfaces can affect the rate and the morphology of electrochemical corrosion. In addition, under certain conditions, there is the possibility of secondary corrosion occurring where the products of anodic and cathodic process interact with each other to produce insoluble compounds. In most cases, the produced precipitates of the metal hydroxide compounds will play a role in inhibiting the corrosion process to some extent. Of course, the protective properties of these deposition films are much worse than the oxide film produced directly by the chemical reactions of metals and oxygen.

2.4.2 Types of corrosion cells Based on the sizes of the electrodes in the corrosion cell and the stability of the distribution of the anode and cathode changing with time, coupled with the influence of factors that promote the formation of corrosion cell and corrosion characteristics, corrosion cells are generally divided into three types: macro-corrosion cell, micro-corrosion cell, and submicro-corrosion cell. 2.4.2.1 Macro-corrosion cell The anode and cathode in such corrosion can be visually distinguishable, and the cathode and anode maintain long-term stability, causing the apparent localized corrosion. (1) Bimetallic cell and dissimilar metal cell This is corrosion in which two or more metals contact with each other in the electrolyte solution, also known as the corrosion couple. In real metal structures, there are usually different metals in contact with each other, and it can be observed that the metals with more negative potential (anode) corrode faster while the metals with more positive potential (cathode) corrode slowly or even are protected. The higher the potential difference of the metals constituting the bimetallic cell is, the more serious the localized corrosion is. Such corrosion damage is called as the galvanic corrosion. For example, a copper device has an iron rivet on it (Fig. 2.13). When it is exposed to moist air, the moisture in the air condenses on its surface, forming a thin layer of water film, shown below the dotted line of Fig. 2.13. Since the water film is dissolved

2.4 Constitution and classification of corrosion cell 

 47

with CO2, O2, etc., in the air, a conductive electrolyte solution is formed. In this case, copper and iron will form a corrosion cell, which can be expressed as 󵄨 󵄨 Fe󵄨󵄨󵄨H2 O(CO2 )󵄨󵄨󵄨Cu.



Fig. 2.13: Connecting corrosion of Fe and Cu [1].

Because the potential of Fe is lower than the potential of Cu, in this corrosion cell, Fe is anode and Cu is cathode. The electrons can be transferred from the iron rivet to the copper sheet. The following reactions on both electrodes will occur: Anode (iron rivet): 2Fe → 2Fe2+ + 4e Cathode (copper sheet): O2 + 2H2 O + 4e → 4OH− Electrode reaction: 2Fe + O2 + 2H2 O → 2Fe2+ + 4OH− 2+

When the solution pH is higher than 5.5, F ​​ e​ ​​ ions and ​​OH​−​​ ions generate precipitate of​​ Fe(OH)​2​​, followed by the oxygen in the air oxidating it into the ferric hydroxide:

4Fe(OH)2 + O2 + 2H2 O → 4Fe(OH)3 .

Thus, it can be seen that under the effect of bimetallic cell constituted by copper and iron, the iron rivets dissolve gradually and their initial shapes and sizes have changed, while the copper ones do not change at all. Such combination of the dissimilar metals is very dangerous, and after a certain time, the copper sheet would fracture at the iron rivets. Another example is saltwater corrosion such as the corrosion and protection of the stern section that has a variety of metals in contact with others, which constitute a multielectrode cell, shown in Fig. 2.14. The corrosion potentials of these metals in seawater are ranked as the following order:

⋅

Zinc alloys (or aluminium alloys), mild steel, hardened steel, bronze → Potential increasing

Fig. 2.14: Illustration of the tail part of a seagoing ship [1]. 1 – Axis (hardened steel). 2 – Propellern (bronze). 3 – Rudder (mild steel). 4 – Ship hull (mild steel). 5 – Sacrificial anode (aluminum alloy). 6 – Skeel (hardened steel).

48 

 2 Thermodynamics of materials corrosion

Zinc (or aluminum) having minimum potential is the anode of this multielectrode system. Due to their prioritized solution, they protect the other three metals. (2) Concentration cell This is a corrosion cell constituted by one metal contacting with electrolyte solutions composed of different components, which can be further divided into the differential aeration cell and the salt concentration cell. 1. The differential aeration cell is a ubiquitous and harmful phenomenon that causes corrosion damages. Generally, if the cell is formed after the metal contacts with agents containing different oxygen contents, it is called the oxygen concentration cell. An example is putting two iron sheets into the dilute brine with the same concentration that one of them contacts with the aerated solution filled with air and another of them contacts with the solution filled with nitrogen (Fig. 2.15). The negative charge densities on the two sheets should be same; while the oxygen contents in the two solutions are different, the extents of the following reactions are different:

O2 + 2H2 O + 4e → 4OH− .

At the air aerated electrode, the reaction of oxygen reduction is rapid, and thus, the electron on the surface is neutralized so that the negative charge density decreases and the potential become more positive. After being aerated with nitrogen gas, in the electrode containing less air, the oxygen-reduction reaction is difficult to occur, leading to the negative charge density being unchanged and the potential being more negative. Thus, the potential difference caused by the different oxygen concentrations in the solution generates corrosion current. In the above-mentioned cell, it can be observed that the iron sheet in contact with the nitrogen is dissolved, while the iron in contact with air does not. Of course, the above discussion is relatively superficial. The essence of the oxygen concentration cell will be discussed in detail in the following chapter on the dynamics of corrosion process in detail.

Fig. 2.15: Oxygen concentration cell [1].

2.4 Constitution and classification of corrosion cell 

 49

In fact, the oxygen concentration cell we often see is due to the various parts of metal being in contact with solutions with different oxygen contents. For example, the most vulnerable part of an iron bucket that is filled with water is the area below the waterline. The metal above the waterline is directly contact with air, whose surface has a higher oxygen content in the water film. However, as the oxygen in the solution below the waterline has a certain solubility, coupled with the slow diffusion rate, the oxygen content in the surface of the iron below the waterline is much lower. This phenomenon produces a similar oxygen concentration cell as the one mentioned above, and therefore, iron below the waterline is an anode and is consequently corroded. Due to the impact of solution resistance, the part suffering from severe corrosion is usually not far away from the waterline, and so it is called waterline corrosion. This type of corrosion occurs frequently in the bridge steel piles, steel pile wharfs, and offshore oil platforms. Metal pipelines and equipment buried in the soil with different densities or depths also suffer from oxygen concentration cell corrosion due to the nonuniformly aerating oxygen to the soil. Engineering components commonly use rivets, welding, and screws to connect. The deep gaps formed at the connecting points are subject to severe corrosion, due to the difficulty of aerating the oxygen. So, the crevice corrosion is also caused by the oxygen concentration cell. 2. The salt concentration cell is another form of concentration cell. If one end of a long copper rod contacts with the dilute copper sulfate solution and another with the concentrated copper sulfate solution, according to the Nernst equation, the potential of the part contacting with the dilute copper sulfate is more negative as anode, while the part contacting concentrated copper sulfate is more positive as cathode. The two parts constitute copper salt concentration cell. Cu2+ ions precipitate on the cathode, while the copper as the anode will be dissolved. The dissolution reaction is carried out until the concentrations of copper sulfate in the two parts of the copper rod are same when the electrode potentials of the two copper electrodes are equal and the electromotive force of the cell is zero so that no corrosion current is generated. (3) Differential temperature cell The metal immersed in an electrolyte solution may form a differential temperature cell because of the different temperatures of each metal part. This often occurs in devices of heat exchangers, boilers, etc. For example, when checking a heat exchanger that is made of ​​ carbon steel, it is found that the part with the higher temperature corrodes more seriously than the part with low temperature because the higher-temperature part become the anode of corrosion cell, due to its lower potential than that of the carbon steel with lower temperature. However, the electrode processes at different temperatures of copper and aluminum in a certain solution are opposite to the one of carbon steel, and in the copper sulfate solution, the copper plate with lower temperature is an anode while the higher temperature side is cathode. Examples of the differential temperature cell needs to be further concluded, and few theoretical research works about it are currently being carried out.

50 

 2 Thermodynamics of materials corrosion

2.4.2.2 Micro-corrosion cell A lot of micro-electrodes formed on the same metal surface due to the electrochemical nonuniformity generate a corrosion cell, which is called micro-corrosion cell. In such a cell, the size of the electrodes whose polarity can stay for a long time in the corrosion process could be equal to the size of the metal crystal structure (from 1 mm to 0.1 μm). Of course, this can cause the localized corrosion in the microrange (such as the pitting corrosion and the intergranular corrosion). But from the macro aspect, it can be considered as the uniform corrosion. The reasons of the electrochemical heterogeneity on the metal surface are various, which can be summarized to the following situations. (1) The nonuniformity for the chemical compositions of the metal Generally, industrial pure metals often contain various impurities, and when they come into contact with the electrolyte solution, the impurities as a micro-electrode will form many short-circuited cell with the base metal. If the impurities are cathode, they would accelerate the corrosion of the base metal. For example, Fe​​Zn​7​​ in industrial pure zinc,​​ Fe​3​​C in carbon steel, graphite in cast iron, Al-Fe and Al-Fe-Si in industrial pure aluminum, etc., are the cathodes in the microcells, which accelerate the dissolution of the base metal in the electrolyte solution. The segregation of alloy components can also cause electrochemical heterogeneity. For example, α-Cu-Zn alloy has the segregation phenomenon during the process of crystallization that the parts that crystallize first contain more copper to become a cathode because of its higher potential, while the parts crystallizing later contain less copper so that they become the anode due to their lower potential. (2) Inhomogeneity of metal structure. Most of the metallic materials used currently are polycrystalline, so there are countless grain boundaries on their surfaces. The grain boundaries of metal are the areas where the arrangement of atoms is relatively loose and disordered and where it is easy to gather impure atoms and produce grain boundary precipitates. In addition, the densities of crystal defects such as dislocations, holes, and distortions are high. Thus, the potential of the grain boundary is usually lower than that inside the grain, which makes the grain’s boundary in the corrosive agent more active and the dissolution occur preferentially. For example, the potential difference of the grain and grain boundary in the industrial pure aluminum, according to measurements, is about 0.091 V and the grains are the cathodes of the microcells and the grain boundary is the anode. In another example, during the tempering process of austenitic stainless steel, due to the precipitation of the chromium-rich phase ​​Cr​23​​​​C​6​​along the grain boundary, the grain boundary lack of chromium becomes the anode of the microcell, causing the well-known intergranular corrosion of stainless steel. The potentials of the different phases in the multiphase alloys are often different, which is also the reason for forming the micro-corrosion cell. The phase selective corrosion in duplex stainless steel is an example.

2.4 Constitution and classification of corrosion cell 

 51

(3) Inhomogeneity of metal’s physical state The machining process of metal or the assembly of metal members often causes the nonuniform deformation and nonuniform internal stress in all parts of metal. Generally speaking, the part with a larger deformation is anode. For example, the bend part of iron sheet and the rivet head are prone to corrosion, and that is the reason. During the process of using a metal, the material bears various loads. Experience has shown that the stressed parts are usually anode and corrode preferentially, such as boiler corrosion and the corrosion for the stress concentrates of ships and bridges that are caused by the stress inhomogeneity. (4) Incompleteness of the film (the coating) on the metal surface If the film and coating on the metal surface are not complete or have gaps and get damaged, a potential difference between the film (or coating) and the base metal below the gaps and damages is produced to form a micro-corrosion cell. In most cases, the potentials of the metal in the gaps and damages are lower than that of the surface film, and therefore, they become the anode of a microcell. In the medium containing chloride ions, due to the strong destructive effect of chloride ions on the passivation film of the stainless steel, the weak points of the film can more easily accelerate the destruction of the film, causing pitting corrosion. 2.4.2.3 Submicroscopic corrosion cell This usually refers to the corrosion cell whose cathode and anode are difficult to be distinguished by the naked eye and an ordinary microscope. The electrode surface area of this cell, which is statistically distributed with no order on the whole surface of metal, is very small (10–100 Å) and its cathode and anode are constantly changing with time. The macro-performance of the results caused by the submicro-corrosion cell is a uniform corrosion. So we can regard the entire surface of the metal undergoing the corrosion as an electrode where the anode and cathode reactions are occurring simultaneously. That is, the entire surface is not only an anode but also a cathode. The reason for the formation of the submicroscopic corrosion cell is not only the nonuniformity of the metal surface structure but also many other reasons such as the differences of the components or the individual crystal orientation, the existence of the grain boundaries or the presence of heterogeneous inclusions, and the ultramicroscopic nonuniformity of metal surface such as the incomplete lattice and crystal lattices containing dislocations, heterogeneous atoms and the atoms at the different energy state, and the thermal expansion between the metal atoms and the solvents ions. When the presence of submicroscopically electrochemical nonuniformity exists, coupled with the mechanical stress in the metal, the directional movement of dislocations caused by the stress deformation or the constant of the anode polarity generated by the stacking dislocations may cause stress corrosion cracking or corrosion fatigue. In this case, although the size of the anode area that works at each moment is small, as the continuous directional movement of the anode, the total anode region (corrosion

52 

 2 Thermodynamics of materials corrosion

cracking) can be extended sufficiently long enough and up to the length of parts when the cracking occurs. It must be noted that in the actual metal corrosion, the effects of macro- and micro-corrosion cell often exist simultaneously. For example, the corrosion of Zn sheet alone immersed in dilute sulfuric acid is caused by the microscopic corrosion cell, while after being linked together with copper, the corrosion of zinc sheet in the dilute sulfuric acid is accelerated by not only the effect of many microcells on the surface of zinc itself but also the macrocells. Additionally, in the cathodic protection, the dissolution of the aluminum sacrificial anode is mainly the result of the macrocell constituted by the linked body to be protected, such as the steel hull and the anode. However, we should not ignore the dissolution caused by the microcell of the aluminum itself. Therefore, the current efficiency of the sacrificial anode cannot reach 100%.

2.5 Electrode process of corrosion cell 2.5.1 Anodic process An anodic process of corrosion cell refers to the process of electrochemical dissolution process of metal as an anode. Anodic dissolution reaction in aqueous solution is

Mn+ ⋅ ne + mH2 O ⇐⇒ Mn+ ⋅ mH2 O + ne.

If the metal ions and anions in the solution generate the complexing ions, the overall reaction is

Mn+ ⋅ ne + xA− + yH2 O ⇐⇒ (MAx )n−x ⋅ yH2 O + ne.

In fact, the anodic dissolution of metal at least includes the following successive steps of

M (lattice) → M (adsorption) (2.1) M (adsorption) → (Mn+ ) ∗ +ne(2.2) (Mn+ )∗ → Mn+  (2.3)

Step (2.1) indicates that the metal atoms leave the crystal lattice and convert into surface-­adsorbed atoms; Step (2.2) indicates that the surface-adsorbed atoms get across the electric double layer and convert into ions; and Step (2.3) indicates that metal ions (or metal hydrated ions) move from the electric double-layer solution to the deep solution.

2.5 Electrode process of corrosion cell 

 53

As we all know, structural metals (including alloys) used in the industry are all polycrystalline solid metal. Metal surface is constituted by a number of crystal faces and is filled with narrow bands – the grain boundaries that distort the crystal structure. The basic unit of grain is crystal cell. Metal atoms are mainly located on the lattice nodes by the effect of the electrostatic forces of metal bond, ionic bond, covalent bond, and Van der Waals forces to bond the atoms together. Atoms have a variety of arrangement forms that are metallographically called space lattice. The most common metal structures include three kinds: fcc, hcp, and bcc (Fig. 2.16). The metal atoms are located on spots of the space lattice and vibrate only in the vicinity of the spots. The gap locations in the unit cells of different metals are not the same, which affects the embedding of heterogeneous atoms and the distortion of the crystal cells. In order to make the metal atoms leave the metal surface, the vibration energy in the perpendicular direction of the surface must be large enough so that it can overcome the electrostatic attraction. Based on calculations, the energy of a metal atom into vacuum is approximately 6000 kJ mol−. Therefore, even at high temperatures, most metals do not evaporate in the vacuum. However, if immersing the metal into aqueous solution, the situation is quite different. The solvent, water, in aqueous is polar molecule, whose charge distribution is not uniform. In a water molecule, oxygen termination is negatively charged and hydrogen termination is positively charged, which form the so-called dipole. The electrostatic attraction that the atoms at the top of the crystal lattice, which are most vulnerable

Fig. 2.16: Lattice structures of typical metals [9].

54 

 2 Thermodynamics of materials corrosion

to the action of water polar molecules, suffer from is minimal. In the body-centered cubic crystal, the surface atoms at the top of crystal lattice have only three atoms around them, and therefore, it is only necessary to overcome the gravities of three bonds to escape from lattice while the water dipolar molecules can approach them from three directions. Under the effect of water dipolar molecules, metal atoms move from the lattice to the surface to become the adsorption atoms. Compared with the atoms on the lattice points, they have a larger degree of freedom so that they can transform into metal ions or metal hydrated ions. The potential energy of metal ions in the lattice is higher, while the potential energy of its hydrated ions is lower, which means the hydration process is companied by the lowering of energy, that is, a spontaneous process. Figure 2.17 illustrates that a lattice atom transforms into an absorption atom that is partially hydrated, which then converts into complete hydration state on the edge of absorption layer, and lastly, the ions migrate to the deep solution. The above discussion applies to an intact crystal, as the anode of an intact crystal always starts from the top of edge of lattice. However, the fact is that industrially pure metals often contain uniformly dissolved heterogeneous inclusions, including metals and non-metals. These inclusions would cause the distortion of each individual crystal lattice, which thus makes the transition of the lattice atoms into the absorption layer easier. In addition, if the crystal has defects such as screw dislocation, its outcrop point can also accelerate the dissolution, and the solution components are easily absorbed onto such crystal defects, accelerating or inhibiting anodic dissolution. When the adsorbed solution components can generate complex compounds with metal ions, it can

Fig. 2.17: The process of lattice atom dissolving into solution [6].

2.5 Electrode process of corrosion cell 

 55

reduce the activation energy of the anode to facilitate the anodic process. Conversely, when the adsorbed solution compounds form an inhibition layer on the metal surface, it would prevent the metal ions from dissolving freely into the solution, which can suppress the anodic process. The anodic dissolution process of metal is always accompanied by an anodic polarization when the electrode potential of metal deviates from the primary equilibrium potential and shifts in the positive direction under the effect of anodic current. In general, the anode current increases with an increase in anode potential, that is, the increase of the rate of metal dissolution, which refers to the dissolution of metal in the active state. However, some metals may be passivated during the process of anodic dissolution. A passivation film is formed on the metal surface to inhibit the metal ion from dissolving into the solution, which transforms the metal into a passivated state. At this time, the anodic current declines sharply, and the dissolution process nearly stops. The relevant contents will be discussed in detail in the chapter of “metal passivation.”

2.5.2 Cathodic process A cathodic process of corrosion cell refers to the reaction of the combination of the oxidant in the electrolyte solution with the residual electrons released after metal anode dissolving. The above discussion has pointed out that if there is no presence of such an oxidant the in solution, even if there are many microcells on the metal surface, electrochemical corrosion cannot occur. Therefore, the basic condition that constitutes electrochemical corrosion is that the cathode process and anode process must occur simultaneously. There are two most important cathodic processes in an electrochemical corrosion: (1) Reduction reaction of hydrogen ion

H+ + e ⇐⇒ H H + H ⇐⇒ H2

This reaction occurs commonly on the corrosion of certain metals in acidic mediums. The electrode potentials of metal Fe, Zn, Al, etc., are lower than the electrode potential of hydrogen; therefore, this metal corrosion in acidic medium is accompanied by hydrogen evolution. The corrosion rates of metals are controlled by the cathodic process and are related with the overpotential of hydrogen evolution. (2) Reduction reaction of dissolving oxygen into the solution In a neutral or alkaline solution, oxygen is reduced into ​​OH​−:​​

1 O + 2H2 O + 2e ⇐⇒ 2OH− 2 2

56 

 2 Thermodynamics of materials corrosion

In an acidic aqueous solution, oxygen is reduced into water:

1 O + 2H+ + 2e ⇐⇒ 2H2 O 2 2

This cathodic process is the most commonly in metal corrosion. Most metal corrosion in the atmosphere, soil, water, and neutral salt solution is inseparable from the oxygen reduction reaction, whose rate is controlled by the oxygen depolarization process. Metal corrosion in the acidic medium containing oxygen may have two reduction reactions occurring simultaneously. At this time, the corrosion rate is controlled by both depolarization processes. Questions about the dynamics will be discussed in the next chapter. In addition to the two cathodic processes of the reduction reaction of hydrogen and oxygen, there may be other processes occurring. To sum up, the reactions are as follows: (1) Reduction reaction of high valence ions in solution such as the reduction reaction of trivalent iron ion in rust.

Fe3 O4 + H2 O + 2e ⇔ 3FeO + 2OH− or Fe(OH)3 + e ⇔ Fe(OH)2 + OH−

(2) Reduction of the noble metal ions in the solution, such as the reaction where divalent copper ion is reduced into metallic state.

Cu2+ + 2e ⇐⇒ Cu

This reaction may occur in the selective dissolution of brass when the dissolved copper ions may redeposit on the surface of brass to form a layer of loose red copper sponge, which plays a role as an additional cathode in the cell, further accelerating the dissolution of zinc in its alloy. (3) Reduction of oxidizing acids and some anions

NO−3 + 2H+ + 2e ⇐⇒ NO−2 + H2 O

(4) Reduction of certain organic compounds

RO + 4e + 4H+ ⇐⇒ RH2 + H2 O



R + 2e + 2H+ ⇐⇒ RH2

In summary, the cathodic processes in the corrosion of various metals are not a single reaction, which are related with both the nature of metals and the composition of agents. Thus, for a particular corrosion process, a specific analysis

References 

 57

should be made. Because the reduction reaction of hydrogen ions and the dissolved oxygen are the two most important cathodic processes, the metal corrosion caused by the reduction reaction of hydrogen ions at the cathode is called hydrogen evolution corrosion, while the metal corrosion caused by the reduction reaction of the dissolved oxygen at the cathode is called the oxygen reduction corrosion.

References [1] Huang, Y. C. The Fundamentals of Metals Corrosion and Protection. Shanghai Jiao Tong Univ. Press: Shanghai, 1989. [2] Tomashow, N. D. The Theory of Metals Corrosion and Protection (B. D. Hua, Trans.). China Machine Press: Peking, 1965. [3] Scully, J. C. The Fundamentals of Corrosion, 2nd edition (Q. Z. Li, Trans.). China Water and Electricity Press: Peking, 1984. [4] Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions, 2nd English edition. National Association of Corrosion Engineers (NACE): Houston, 1974. [5] Zook, N. D.:A Course of Metals Corrosion and Protection. Metallurgy Press: Moscow, 1976. [6] Wei, B. M. The Theory of Metals Corrosion and Application. Chemical Industry Press: Peking, 1979. [7] Cao, C. N. The Fundamentals of Corrosion's Electrochemistry, 6th edition. Chemical Industry Press: Peking, 2004. [8] Kerser, H. Corrosion of Metals (Wu Yin Shun Trans.). Chemical Industry Press: Perking, 1984. [9] Xu, Z. Y.; Xing, L. P. An Introduction of Materials Science. Shanghai Science Technology Press: Shanghai, 1986.

Yongchang Huang

3 Electrochemical corrosion kinetics 3.1 Corrosion potential and its polarization diagram 3.1.1 Corrosion potential When metals suffer from electrochemical corrosion in aqueous solutions, their surfaces have at least two different electrode reactions at the same time, one of which is the metal oxidation reaction (anodic process) and another which is the reduction reaction of the oxidants in the solution (cathodic process). The rates of these two reactions are equal so that there is no accumulation of charges on the metal surface and the net current is zero. The potential of the corroded metal stays constant after metal immersion in an agent for a long time. The absolute potential cannot be measured currently, while its relative value can be measured by the reference electrode. The electromotive force of the galvanic cell constituted by the reference electrode and the electrode to be measured is the potential of the electrode to be measured. Now, let us continue to discuss the meaning of the electrode potential when there are two electrode reactions on a metal at the same time. As the net current on the metal is zero, it can be recognized as an isolate electrode. Below, we will analyze the problem by an imaginary short-circuited primary cell. In Fig. 3.1(a), the anode and cathode are made up by the same metal M, and the two electrodes are linked by a wire whose resistance is nearly zero, which means the anodic reaction and cathodic reaction are likely to occur on one metal surface, shown in Fig. 3.1(b).

Fig. 3.1: Short-circuited primary cells constituted by the same metal [1]. (a) Primary cell whose electrode materials are made up by one metal. (b) Relationship between the primary cell constituted by the electrodes made up by same metal and the coupled system of the electrode reactions.

https://doi.org/10.1515/9783110310054-003

60 

 3 Electrochemical corrosion kinetics

The two electrode reactions in the above-mentioned primary cell are as follows: − Anodic reaction: X 󳨀→ X + e − Cathodic reaction: Y + e 󳨀→ Y

So the whole reaction is as follow: − − Redox reaction: X + Y 󳨀→ X + Y

Therefore, when there are two electrode reactions occurring on an isolated electrode at the same time, the two electrode reactions are likely to occur in a short-circuited primary cell. That is to say: (1) The electrode reaction with higher equilibrium potential proceeds in the direction of cathodic reaction, and the one with lower equilibrium potential proceeds in the direction of anode reaction. (2) The overall result of the two electrode reactions is redox reaction, whose propelling force comes from the equilibrium potential difference of the two electrodes. (3) The two electrode reactions, one of which is in the anode reaction and another of which is in the cathode reaction, have the same rate. The current from the electrode to the solution in the anode reaction is just the opposite current in the cathode reaction. In addition, the chemical energy released from the redox reaction constituted by the two electrodes diffuses in the form of heat energy without any working capacity. The phenomenon that on an isolated electrode there are an anode reaction and a cathode reaction proceeding at the same rate is called the coupling of the electrode reaction. The two coupling reactions are also called the conjugated reaction. The meaning of the conjugation is to rely on each other. When the two electrode reactions conjugate with each other, if reaction 1 is anode reaction whose equilibrium potential is E01 and reaction 2 is cathode reaction whose equilibrium potential is E02, then the result of conjugation is getting a new electrode potential E. The coupled electrode reactions proceed under the condition of the nonequilibrium potential E. Consequently, the electrode potential E is the nonequilibrium potential of not only the anode reaction but also the cathode reaction. The value of E is among the equilibrium potentials of the two electrode reactions is as follows: E02 > E > E01

Therefore, we call the electrode potential E as the mixed potential or the compromising potential. If in the coupled electrode reactions, the anode reaction is the anodic dissolution reaction of metal M: M 󳨀→ Mn+ + ne

3.1 Corrosion potential and its polarization diagram 

 61

While the cathode reaction is the reduction reaction of the oxidant in the solution: Y + e 󳨀→ Y−

Then the inevitable result of the electrode reactions is to cause the destruction of the metal M. The coupling of the two electrode reactions, which causes the corrosion, is referred to as the corrosion couple, whose mixed potential is corrosion potential. The whole redox reaction is the corrosion reaction, which is an electrochemical corrosion process. According to the electrochemical thermodynamics, the equilibrium potential of the anode reaction is E01 = E�

𝜇 n+ − 𝜇M M �= M . nF Mn+

And the equilibrium potential of the cathode reaction is E02 =

𝜇Y − 𝜇Y− . F

The condition of the electrochemical corrosion is △G = −nFE < 0 △G = nF�

𝜇Mn+ − 𝜇M 𝜇Y − 𝜇Y− � 0.

Therefore, the reason that a metal suffers from electrochemical corrosion in the aqueous solution is that there exists substances whose equilibrium potential of reduction reaction must be higher than the equilibrium potential of oxidation reaction of this metal. The discussion above can be illustrated by an example of the corrosion of metal Fe in the 0.1 N HCl. There are two electrode reactions occurring simultaneously on the surface of Fe, which immerses in the dilute hydrochloric acid.

Fe 󴀘󴀯 Fe2+ + 2e (3.1)



H2 󴀘󴀯 2H+ + 2e(3.2)

62 

 3 Electrochemical corrosion kinetics

At this time, the Fe sheet plays the roles of both an iron electrode and a hydrogen electrode, and therefore, it can be considered as the double electrode to distinguish from the single reversible electrode. The reactions on the surface of the Fe sheet are called the coupled reaction because if there is no reaction (3.2), the continuation of the dissolution reaction of Fe is impossible. Given that the concentration of Fe2+ in the solution is 10−2 M, according to the calculating result of the Nernst formula, the equilibrium potential of Fe electrode is about −0.50 V (vs. standard hydrogen electrode [SHE]). Now, as there is another reaction that destroys the equilibrium of reaction (3.1), the dissolution reaction continues in the oxidation direction, which corresponds to an existing anode current in the doubled electrode Fe|HCl, Ia > 0. Similarly, as reaction (3.2) provides residual electrons to reaction (3.1), it can continue to proceed in the reduction direction, which corresponds to an existing cathode current in the doubled electrode Fe|HCl, |Ic| < 0. The above-mentioned Fe|HCl system does not connect with the external circuit, whose net current I is equal to zero. Consequently, at a stable corrosion state, the anode dissolution current of Fe is equal to the cathode reduction current of hydrogen ion and with the opposite direction. IFe = −I H

or  I = IFe + IH = 0 Fe electrode and hydrogen electrode are polarized under the effect of each anode current. The result of polarization is to make the equilibrium potential of iron dissolution shift positively and the equilibrium potential of hydrogen reduction shift negatively and finally stay in stable values. At this moment, the potential of the Fe sheet, which differs from both the equilibrium potential of the Fe electrode (−0.50 V) and the equilibrium potential of the hydrogen electrode (0.06 V in 0.1 N HCl), is one that can maintain the above current equations. The real measured potential of the Fe sheet is −0.25 (vs. SHE), which is the stable potential of Fe in 0.1 N HCl. The potential is called natural corrosion potential or corrosion potential, indicating the potential of metal staying in the nature corrosion state, expressed as Ik or Icorr. 3.1.2 Polarization effect and corrosion polarization diagram 3.1.2.1 Polarization of galvanic cell In the last section, it was introduced that when the two electrode reactions occur in one insolated electrode simultaneously, it can obtain a certain stable potential whose value is not equal to the equilibrium potential of the anode reaction and cathode reaction but among the two potentials. Such phenomenon is caused by the polarization of electrode.

3.1 Corrosion potential and its polarization diagram 

 63

  Fig. 3.2: Corrosion cell of Cu-Zn [1]. It is supposed that a corrosion cell is constituted by Zn|3% NaCl|Cu, where Zn is anode and Cu is cathode, whose surface areas are equal. Zn and Cu were immersed in the same container with 3% NaCl solution and linked with the wire of ampere meter A with switch K, shown in Fig. 3.2. The measured potentials of Zn and Cu in the 3% NaCl solution are, respectively, E​ ​​  oZn   ​​​= −0.76 V, E​ ​​  oCu  ​​​ = 0.04 V, and the whole resistance of the primary cell is 230 Ω When it is open circuit, as the resistance R ⟶ ∞, I o 󳨀→ 0. At the moment the circuit is shorted, the electrode surface is too late to change, and according to Ohm’s law, the current in the cell should be E0Cu − E0Zn R 0.04 − (−0.76) = 3.5 (mA). = 230

Istart =

However, with the electrifying time extending to several seconds, it is found that the cell current decreases slowly and finally reaches a stable current of 0.2 mA, which is 1/17.5 of the starting current. So what is the reason for this? According to Ohm’s law, the factors influencing cell current includes two aspects, one of which is cell resistance and another which is the potential difference of the two electrodes. Under the above condition, the cell resistance does not change much, and the decreasing of the primary cell current certainly is caused by the potentials of the anode and cathode and its potential difference. Theoretically, before the cell gets switched on, the determined potential of the structure of the electric double layer on the Zn electrode surface is −0.76 V, and the determined potential of the structure of the electric double layer on the Cu electrode surface is 0.04 V when the two electrodes reach the state of charge equilibrium, without the flowing of net current. Once the cell is connected, the equilibrium of the

64 

 3 Electrochemical corrosion kinetics

two electrodes is immediately destroyed, and the electrons transfer from the Zn anode to the Cu cathode. With the change in the charge density on the electrode surface, the potentials of the electrodes also get varied. By measuring the two electrode potentials, this is proved to be correct. Figure 3.3 shows the changing of the electrode potentials after connecting the cells. From the diagram, it can be seen that after the cell has been connected for several seconds or several minutes, the cathode potential shifts in the negative direction and the anode potential difference shifts in the positive direction, which causes a decrease in the potential between the anode and cathode and, therefore, a quick decrease in the stable current in the Cu-Zn cell.

Fig. 3.3: Diagram of polarization curves

  before connecting the circuit [1].

The phenomenon wherein the current intensity decreases because of the current passing through the primary cell, lowering the potential difference of the two electrodes, is called polarization of cell. The essence of the polarization of the cell is the result of the two electrode polarization. The anode potential in the primary cell shifts in the positive direction after switching on the current. Such phenomenon is called the anodic polarization. The difference between the anodic polarization potential and the starting potential is the anode overpotential, expressed by 𝜂a. Anodic polarization caused by the following aspects: (1) When the rate of the metal ions entering into the solution is less than that of the electrons getting through the wire from anode to cathode, there is excessive accumulation of positive charges on the anode, which causes a decline in the negative charges on the electric double layer and, therefore, the positive shift of the anode potential in the positive direction.

3.1 Corrosion potential and its polarization diagram 

 65

(2) At the moment of metal anode dissolution, the diffusion of metal ions in the solution is not fast enough, which leads to an increase in the concentration of the metal ions near the anode, causing the so-called concentration polarization. (3) When there is an anode current flowing through some metals, it is easy for their surfaces to form a dense oxygen film, which inhibits the transformation of the metal ions. At this time, the electrode potential shifts intensively in the positive direction. Due to the formation of the surface film, the resistance of the cell system increases as well, which causes the polarization. The potential change caused by such reason is called resistance polarization. The phenomenon that the cathodic potential in the primary cell shifts in the negative direction after connecting the current is called cathodic polarization. The difference between the cathodic polarization potential and the starting potential is called cathode overpotential, expressed by 𝜂c. The reason for cathodic polarization is that the substances in the solution near the cathode that can accept the electrons combine with electrons in a slow rate because of some reasons so that the electrons transferred from the anode cannot be consumed and there is accumulation of negative charges on the cathode, which leads to the negative shift of the cathode potential. Except for the easy-passivated metals, the corrosion process of most metals is controlled by the cathode polarization. Therefore, the following sections will introduce this in detail. Now, let us continue to discuss the polarization of Cu-Zn corrosion cell and the changing regulation of the two electrode potentials after switching on the current. If connecting variable resistance to the external circuit in Fig. 3.2 and a voltmeter in parallel, shown in Fig. 3.4, the electrode potential in the different currents can be measured by controlling the current.

Fig. 3.4: Illustration of measurement instrument for polariza-

  tion behavior of corrosion cell [2].

66 

 3 Electrochemical corrosion kinetics

Ee d

e

/ V

f c

b Ea a

I

Ik /A

Fig. 3.5: Relationship between cathode and anode

  potentials and current density [2].

When the variable resistor R is large and a small current gets through the external circuit, the values of ammeter A and voltmeter V are recorded​​. Then, decreasing the variable resistor to increase the current of external circuit, it is observed that the voltage value in the voltmeter decreases. When the resistance of the external circuit decreases to the minimum, close to the battery short-circuit condition, the current is largest and the voltage on voltmeter is smallest, which indicates that the larger the currents through the two electrodes are, the more serious their polarizations are, and so the voltage drop between the two electrodes is smaller. In order to study the polarization of the cathode and anode separately, their potential changes are determined in the above-described experiment at the same time, and a curve is drawn with the potentials of the cathode and the anode change with increases in their currents, shown in Fig. 3.5, where E​ ​​ c​​ and E​ ​​ a​​ are, respectively, the open-circuit potentials of Cu and Zn electrodes in 3% NaCl (the potentials without current flowing in the cell). As the current increases, the potential of the zinc electrode changes along the abc line, the potential of the copper electrode along def line changes. It can be seen that the anode potential is gradually increases and the cathode potential gradually reduces, which causes the anode and cathode potential values to get close to each other. When the current in the ammeter is I1, the potential of the zinc electrode is polarized to point b, and the potential of the copper electrode, to point e. The polarization potential difference between the two electrodes is equal to (e-b). When the cell is short-circuit, the current is biggest, expressed by Imax, when the external resistance can be ignored and the potential difference between Cu and Zn electrodes is ( f−c). The corroded metals are similar to the short-circuiting galvanic cell, Imax is the true corrosion current ​​I​k​​ when the two electrodes of the cell have been replaced by the electrodes on one metal, of course. Therefore, the above discussion of cell

3.1 Corrosion potential and its polarization diagram 

 67

polarization can be used to explain some of the kinetic laws of the steady-state corrosion process of metal occurring in the electrolyte solution. 3.1.2.2 Diagram of corrosion polarization and its application The establishment of the steady-state corrosion process of metal in the electrolyte solution can be expressed by a diagram of corrosion polarization, which is a potential-current diagram, that is, plotting the anodic and cathodic polarization curves that characterize the corrosion cells on a same graph. In order to simplify the discussion, we can ignore the details of the potential changing with the current and draw polarization curves by using straight lines. Such diagram of corrosion polarization is called an Evans map (Fig. 3.6). The initial potentials of the cathode and anode are the equilibrium potential of the cathodic and anodic reactions, expressed by ​​E ​ c0 ​ ​​ and ​​E ​ a0 ​ ​​, respectively. When ignoring the ohmic resistance of the solution, the simplified polarization curves intersect at one point. The corresponding potential at the intersection is the conjugate mixed potential, that is, the corrosion potential. The current corresponding to the corrosion potential is corrosion current. Generally, the areas of anode and cathode in a corrosion cell are not equal so that using an E-I polarization diagram is more convenient. However, in the case of uniform corrosion, the metal surfaces act as both a cathode and an anode, and therefore, the E-I polarization diagram can be adopted. When the anode and cathode reactions are controlled by the electrochemical polarization, corrosion polarization expressed by the E-lgi diagram can be clearer. The next section will mention that, under the condition of corrosion controlled by electrochemistry, the relationship of the logarithm of current density and potential changes is linear.

  Fig. 3.6: Polarization map (Evans map) [3].

68 

 3 Electrochemical corrosion kinetics

  Fig. 3.7: Polarization map [3]. The anodic and cathodic polarization processes in the corrosion are not the same, and the degree of polarizations can be expressed by the slopes of the polarization curves in the diagram, respectively, i.e., the value of polarizability P, shown in Fig. 3.7. Pc =

E0c − E K 𝜂 = c = tan 𝛽 IK IK

Pa =

E K − E0a 𝜂a = = tan 𝛼 IK IK

Polarizability can be viewed as resistance in the cathodic process. Its value means that as the polarization potential rises, the current changes a little. According to the definition of polarizability, I K =

E0c − E K E − E0a = K . Pc Pa

After arranging mathematically, one can obtain I K =

E0c − E0a ΔE = . Pc + Pa Pc + Pa

From the above equation, in fact, corrosion current depends not only on the initial potential difference between the anode and cathode of the cell but also on the value of the polarization, which can be seen very clearly on the corrosion polarization diagram that when the polarizabilities of the anode and cathode are large, the corresponding corrosion current on the intersection of the two polarization curves is very small, and on the contrary, the Ik is high. In addition, when in the two electrode processes, one polarizability is high and another is small, the corrosion reaction is controlled by the process with high polarizability. We will respectively discuss these in the following.

3.1 Corrosion potential and its polarization diagram 

 69

(1) Corrosion process controlled by cathode If Pc ≫ Pa, that is, the cathodic polarization curves on the diagram are very steep (shown in Fig. 3.8), the corrosion potential Ek is close to the equilibrium potential of the anode reaction E 0a and the corrosion current ik is also mainly determined by the value of Pc. Under the condition of cathode control, any factor increasing the cathodic polarizability Pc will not significantly reduce the corrosion rate of the cathode, and any factor affecting the anodic reaction will not make a significant change in the corrosion rate. Therefore, under such conditions, the corrosion rate can be controlled by changing the slope of the cathodic polarization curves. For example, the corrosion of the metals Fe and Cu in aqueous solutions is associated with the cathodic process in the oxygen reduction reaction. By using the deoxygenation method, the oxygen concentration in the solution is reduced to increase the cathodic polarization resistance, which can achieve a significant inhibition effect. Additionally, the corrosion of many metals in acidic mediums is related to the cathode process of the hydrogen reduction reaction. By increasing the activation overpotential of the hydrogen reduction reaction, the dissolution of these metals can weaken.

  Fig. 3.8: Corrosion polarization map of cathodic control [3].

  Fig. 3.9: Corrosion polarization map of anodic control [3].

70 

 3 Electrochemical corrosion kinetics

(2) Corrosion process controlled by anode If Pa ≫ Pc, that is, the slope of the anodic polarization curve in the diagram is very steep (shown in Fig. 3.9), the corrosion potential Ek is close to the equilibrium potential of cathodic reaction E 0c and the corrosion current ik is mainly determined by Pa. As shown in Fig. 3.9, any factor that would affect the cathodic reaction will not change the corrosion rate significantly, and any factor that increases the anodic polarizability Pa will significantly reduce the corrosion rate. For example, the metals and alloys that can form passive state in the solution are the typical examples of the corrosion controlled by anode. Adding a small amount of inhibitor to promote the passivation of metal or alloy can greatly reduce the corrosion rate. In addition to the corrosion processes controlled by the cathode and the anode, in the polarization curve diagram, the corrosion processes controlled by the ohmic resistance and the mixed process can also be distinguished. In practice, all of these controlling factors are likely to participate in the controlling process to determine the corrosion rate, and only their respective controlling degrees are different. In order to calculate the degree of each controlling factor, the stable potential of a metal or an alloy in such corrosion condition should first be measured and then the equilibrium potentials of metal, E 0c and E 0a, under the given conditions are then calculated. Based on these data, we can determine the degrees of the cathode polarization control and the anode polarization control. A corrosion polarization diagram is an important tool for the study of electrochemical corrosion, which is widely used and, apart from comparing the controlling degree of each resistance to determine the main controlling factors, is also used to explain the phenomenon of corrosion, analyze the nature and impact factors of corrosion, determine the mechanism of corrosion additives, and graphically calculate the corrosion problem of the multielectrode system. Determination of cathode and anode polarization curves is the basis for drawing the corrosion polarization diagrams.

3.2 Polarization curves of the corroded metal’s electrodes 3.2.1 Polarization curves of the corroded metal’s electrodes in the double-electrode system The metal electrodes with corrosion galvanic couples are called corroded metal electrodes. In the study of the electrochemical behavior of the corroded metal electrodes, one must take into account the case that the isolated electrode has two or more electrode reaction couplings. Let us start by discussing the simplest corrosion process, which consists of two conjugate electrode reactions. Therefore, the polarization curve of the corroded metal electrode is made of an anodic polarization curve and a cathodic polarization curve.

3.2 Polarization curves of the corroded metal’s electrodes 

 71

The initial potentials of the cathodic and anodic polarization curves are respectively the equilibrium potentials of the cathodic reduction reaction and anodic oxidation reaction. And ignoring other secondary factors affecting the polarization process, this polarization curve is called ideal polarization curve, which can be obtained only in ideal electrodes, which means that anode only has the anodic process and the cathode only has the cathode process. In other words, the electrode system has only a single reaction. Obviously, this is an ideal situation. In the previous section, for the purposes of simplifying the discussion, this ideal polarization curve was used. The setup potential of the anode polarization curve in the diagram is the equilibrium potential of the metal dissolution reaction, and the initial potential of the cathode polarization curve is the equilibrium potential of the depolarizing agents in the solution under this condition. As conjugate reactions occur on the two-electrode couple and the metal electrodes suffer the anodic dissolution, the potential deviates from the equilibrium potential in the positive direction, resulting in the presence of an anodic polarization curve, while the depolarizing agents react on the metal electrode, which leads to the negative deviation of the potential and the presence of the cathodic polarization curve. In the case of ignoring solution resistance, the two polarization curves intersect at one point, obtaining the corrosion potential of the metal electrode (Fig. 3.10). Therefore, once the metal immerses in an electrolyte solution, due to the inhomogeneity of the surface electrochemistry, plenty of local microscopic corrosion cells will form and become the polarized electrodes. In the absence of impressed current, the two electrode reactions perform in the local anode area and local cathode area at the same time. The two electrode reactions will not maintain their equilibrium states and the polarization results in the establishment of a common stable potential (corrosion potential). The ideal polarization curve generally cannot be measured directly.

  Fig. 3.10: Ideal polarization curve [4].

72 

 3 Electrochemical corrosion kinetics

  Fig. 3.11: Measured polarization curves [4]. When impressing the polarization into the corroded metal, the measured polarization curves are called apparent polarization curves. The starting points of the measured anodic and cathodic polarization curves are all the corrosion potential EK, shown in Fig. 3.11. E 0s SD and E 0a SB in the diagram respectively stand for the envisaged ideal anode and cathode polarization curves. In the early stage of the polarization, when the impressed current is very small, the ideal polarization curves and the apparent polarization curve have essential differences. However, under the effect of a large impressed current, they increasingly approach each other. When the impressed current is large enough and the polarization potential is higher than the initial equilibrium potential, the two curves coincide. As can be seen from Fig. 3.11, the apparent anodic polarization curve is the adduct of the ideal cathode and anode polarization curves. According to this property, on the basis of the measured apparent polarization curves, the ideal polarization curves can be plotted by using the graphing method, which is the synthesis and degradation methods for the polarization curves that are used frequently in the electrochemical corrosion. Another expression for the polarization curve of the corroded metal electrode is shown in Fig. 3.12. The origin O in the diagram corresponds to the natural corrosion state without connecting the outer circuit, which is expressed by Ek. This corrosion potential corresponds to a corrosion current ik whose value is equal to the value of the anodic oxidation current and equal to the cathodic reduction current. Thus, at this on, the net current on the corroded metal electrode is zero.​​ E​ 0a ​​​ QTV is an ideal anodic polarization curve and E​ ​​  c0 ​ S ​​ PU is an ideal cathodic polarization curve. When the anodic polarization starts, it is supposed that the electrode potential moves upward from the origin point O to M, and at this time, the oxidation current ia increases, corresponding to the line MT in the diagram and the reduction current decreases, corresponding to the line MS. The impressed anodic current Ia flowing through the corroded metal electrode should satisfy the following condition: I a = i a − |i c | = i a + i c .

3.2 Polarization curves of the corroded metal’s electrodes 

 73

  Fig. 3.12: Polarization curve of corroding metal [4]. Accordingly, the values of the impressed current are equal to the line ME. With the potential of the corroded metal electrode shifting positively, the value of ia gets higher and the value of ic gradually gets smaller. When the electrode is polarized to a state that the potential is equal to the equilibrium potential E 0c of the reduction reaction occurring in the partial cathode area, the oxidation current becomes the only current on the electrode and the reduction current disappears. At this moment, Ia = ia, ic = 0. The impressed anode currents flowing through the metal electrode are all for metal oxidation, and the reduction reactions in the local cathode completely stop. Here, we continue to discuss the cathodic polarization curves. When the electrode potential moves from origin O down to point N, the reduction current ic increases, corresponding to line NP, and the oxidation current decreases, equivalent to line NQ. The impressed cathodic current Ic flowing through the corroded metal electrode should satisfy the following condition:

I c = −(|i c | − i a ) = i c + i a .

Therefore, the impressed current value is equivalent to line NF. With the negative shifting for the potential of the corroded metal electrode, the ic value gets bigger and the ia value gets smaller. When the electrode polarization potential is equal to the equilibrium potential E​ ​​  0c ​ ​​of the oxidation reaction in the local anode area, the reduction current becomes only current on the electrode and the oxidation current disappears. At this moment, Ic = ic, ia = 0. In this case, the dissolution of the corroded metal stops, which is the theoretical basis for the electrochemical cathodic protection of metal.

74 

 3 Electrochemical corrosion kinetics

By utilizing the above principle, it is possible under certain conditions that the ideal polarization curve can be obtained from the decomposition of the apparent polarization curve. (1) When the hydrogen ions act completely as the only depolarizing agents in the corrosion process, by using the volumetric method, the hydrogen evolution per unit of time can be accurately measured at the first impressed potential, and then the value is converted into the corresponding current value ic, and according to Faraday’s law, we can draw the ideal anodic polarization curves by using I c = −|i c | − i a . (2) When the cathode of corrosion process only has the oxygen depolarization reaction, the anode polarization curve of metal anodic dissolution can be measured directly in the case of complete absence of oxygen. And the ideal cathodic polarization curves can be plotted by adopting I a = i a − |i c |. Sometimes, in the process of metal corrosion, it is possible that the reduction reactions of two polarization agents can occur simultaneously on a cathode. For example, in an acidic solution containing oxygen, magnesium has the reduction reactions of the dissolved oxygen and the hydrogen ion at a relatively negative potential. In this case, the apparent polarization curves are merged by three ideal polarization curves.

3.2.2 Polarization curves for the corroded metal electrode in a multielectrode system The above-mentioned corroded metal electrodes are seen as short-circuit cells with two electrodes, which is sufficient for explaining the general electrochemical corrosion behavior of metals. However, for nonuniform alloys and the multimetal components widely used in the industry, such simplification is inappropriate. For example, when two nonuniform metals contact in a solution, each of the metals at least can be seen as a double electrode, and therefore, it should be a simple four-electrode short-circuit primary cell in the simple condition. In the previous section, it was mentioned that there are two electrode reactions that occur on the electrode surfaces of corroded metal in the double-electrode system. The absolute value of the current of two conjugate reactions is equal; i.e. i a = |i c |.

When the metal M1 with higher corrosion potential and M2 with lower potential consist of galvanic couples and their areas are the same, there are i aM1 + i aM2 = |i cM1 | + |i cM2 |.

3.2 Polarization curves of the corroded metal’s electrodes 

 75

Therefore, for the galvanic couple as a whole, the total current of the anodic reactions on the electrode remains equal to the total current value of the cathodic reactions. But in two different metals, the reaction rates on the anode and the cathode are no longer equal. In M1, the anode current is greater than the cathode current, which is the anode component of galvanic couples; the cathodic current is greater than the anode current value in M2, which is the cathode component of galvanic couples. Extending this to a case where a galvanic couple is composed of N metals or alloys, the corrosion current of this system is 󵄨󵄨 N 󵄨󵄨 󵄨󵄨 󵄨󵄨 I = � Iaj =󵄨󵄨󵄨󵄨� Icj 󵄨󵄨󵄨󵄨, 󵄨 󵄨󵄨 j=1 󵄨󵄨󵄨 j=1 N

where Iaj and Icj are the anode current and the cathode current of the jth electrode reaction, respectively. The current intensity is used here because the areas of each metal or alloy are not necessarily equal. In the above-mentioned multielectrode coupled system, there must be a part of the electrode reaction as an anode reaction, and another, a cathode reaction. Since the N electrode reactions are carried out at the same mixed potential, the polarization potential of the anode reaction is equal to the equilibrium potential plus the overpotential, and the polarization potential of the cathode reaction is the equilibrium potential minus the overpotential. For convenience, the mixed potential of the coupling system is E = E e1 + 𝜂1 = E e2 + 𝜂2 = ⋅ ⋅ ⋅ = E eN + 𝜂N .

When summing the equations of the N electrode potentials, the following equation is obtained: N

N

j=1

j=1

N E = � E ej + � 𝜂j

E=

N N 1 ��j=1 E ej + �j=1 𝜂j � N

The mixed potential of this multielectrode system has the following characteristics: (1) Mixed potential E is always among the highest and lowest electrode potentials. (2) In a multielectrode corrosion system, a metal or an alloy whose electrode potential is lower than the mixed potential is the anodic component, and a metal or an alloy whose electrode potential is higher than the mixed potential is the cathodic component.

76 

 3 Electrochemical corrosion kinetics

If there are N electrode reactions on the electrode in the electrode corrosion system, the rates (i) of each electrode reaction has the following relationship with the electrode potential (E): i1 = f1 (E) i2 = f2 (E) ⋅⋅⋅ ⋅⋅⋅ i N = f N (E)

When the electrode potential is equal to E, the impressed current intensity of the whole electrode is I = i1 + i2 + · · · + i N = f1 (E) + f2 (E) + · · · + f N (E) = f (E).

Therefore, if the relationships between the current intensities I of each electrode reaction and the electrode potential are known, the polarization curves of the whole electrodes can be obtained. However, the relationships between E and I for every electrode reaction actually are often not simple functions, and the impressed current intensity I and potential E of the whole multielectrode system are complex functions. Therefore, getting the polarization curves for the corroded metal electrodes in the entire multielectrode system from the E-I relationships of each electrode reactions generally always uses the mapping methods.

Fig. 3.13: E-I polarization map of five-electrode

  systems [5].

3.2 Polarization curves of the corroded metal’s electrodes 

 77

According to the size of each electrode, change the potential-current density curve measured when the electrode exists alone into the potential-current intensity curve. Fig. 3.13 is an E-I polarization diagram of a five-electrode system where E1, E2, E3, E4, and E5 are the corrosion potentials of each electrode system, and E5 > E4 > E3 > E2 > E1. As mentioned before, a multielectrode system is the same as the double-electrode corrosion system. At such corrosion potential, the sum of whole cathode currents of the system is equal to the sum of whole anode currents. That is, Ek in Fig. 3.13 is the intersection point q of the adduct curve mnpr of E-I for all cathodic processes and adduct curve stqu of E-I for all anodic processes (shown by a dotted line) and the adduct profile mnpr E-I curves of all the anodic processes (shown by a solid line). The adduct here means the sum of currents at the same potentials. When adding the cathodic polarization curves, there is only one cathodic polarization curve 5’ between E5 and E4, that is, between m and n, while between E4 and E3, that is, n and p, there are two cathode polarization curves of 5ʹ and 4ʹ, which, when added together, will give the total cathodic polarization curve, the np line. From E3, three cathodic polarization curves, 5ʹ, 4ʹ, and 3ʹ, must be added together. And likely, we can continue adding to obtain a total cathodic polarization curve, the pr line. But as can be seen from the diagram, the existing mnpr curve has intersected with all anodic polarization curves so that adding more cathodic polarization curves is not necessary. Using the same method, the total anodic polarization curve stuqu can be obtained. The total cathodic polarization curve and the total anodic polarization curve in the diagram intersect at point q, whose coordinates are the mixed potential Ek and the total current Ik of the five-electrode corrosion system. From the diagram, we can determine the polarity of each electrode and the current value of the system. Since the cathodic polarization curves of 3ʹ, 4ʹ, and 5ʹ respectively intersect with horizontal line Ek-q at b, d, e, the electrodes in the electrode system are cathode, the current values of which depend on the length of lines ab, ad, and ae. Similarly, the anodic polarization curves of 1 and 2 intersect with the horizontal line Ek-q at c, f so that they are the anode whose current values depend on the length of lines ac and af. After further analyzing Fig. 3.13, it can be seen that the smaller the polarizability of the electrode reaction is, that is, the more smooth the polarization curve is, the greater the role that the electrode plays on the current adding and on the distribution of the polarities of other electrodes. For example, when decreasing the polarizability of the most effective cathode, a cathode can be converted to an anode, whereas decreasing the polarizability of the most effective anodic then promotes the conversion of the anode into the cathode. The above discussion can be explained by the experimental data. Let five metals of zinc, aluminum, cadmium, lead, and platinum make up the multielectrode corrosion cell in 3% NaCl solution, where the areas of each electrodes in the cell are equal and all 4 cm2. In this case, only the zinc is anode, but after adding the hydrogen peroxide solution into the NaCl solution, the polarity of the electrodes has changed.

78 

 3 Electrochemical corrosion kinetics

Since H2O2 will greatly reduce the polarizability of the effective cathode Pt, even Al and Cd may also become the anode. The polarity of the electrodes can be expressed by the positive and negative currents in Tab. 3.1, where the positive current is the cathode and the anode current is negative. Tab. 3.1: Influence of the addition of H2O2 on each electrode partial current of Zn-Al-Cd-Pb-Pt corrosion couple in 3% NaCl [5]. Content of H2O2 (%) 0

Partial current of each electrode (mA) Zn

Al

−0.35

+0.03

Cd +0.06

Pb

Pt

+0.07

+0.19

0.05

−8.15

+0.16

+2.28

+2.88

+2.84

0.10

−14.00

+0.04

+3.54

+5.50

+4.92

0.30

−36.00

−5.40

+5.60

+9.60

+27.00

0.50

−40.00

−17.00

+3.00

+4.40

+50.00

1.0

−47.00

+33.20

−8.60

+2.40

+87.00

1.50

−44.00

−34.60

−19.00

+1.60

+96.00

Data listed in the table show that for the 3% NaCl solution without H2O2, only Zn is the anode and Al, Cd, Pd, and Pt are the cathodes. When the H2O2 content is 0.3%, in addition to Zn, Al has become the anode. When H2O2 content continues to increase to 1%, Cd also becomes the anode. Further, if increasing the surface area of the effective cathode Pt, the intermediate electrodes can become the anodes. The larger the area of ​​Pt is, the faster the corrosion rates of Zn, Al, and Cd are.

3.3 Corrosion process controlled by the activation polarization 3.3.1 Electrochemical polarization equation Metal corrosion in aqueous solutions occurs at the corrosion potential (nonequilibrium potential). According to the electrode process kinetics, the whole electrode process in a nonequilibrium potential is regarded as a heterogeneous reaction process that occurs at the interface between the electrode and the solution. In general, the electrode process involves three basic steps: (1) The reaction materials transmit from the interior solution to the liquid layer near the electrode surface, and the transmission power is the gravity of the electric field and the concentration difference for the reaction materials between the interior solution and the electrode surface.

3.3 Corrosion process controlled by the activation polarization 

 79

(2) The reaction materials gain and lose electrons on the interface between the electrode and the solution. (3) The reaction materials transfer to a stable state or transmit from the solution near the electrode surface to the interior solution. Under a certain condition, the total rate of the electrode process lies on the one of the above three steps, which has slowest rate. The step with the slowest rate is called the control step of the electrode process. It can be contemplated that the three steps of the electrode process can be counted as the three resistances of the net current flowing on the electrode, which is the reason for the electrode polarization, as well as the current decrease. The polarization caused by the resistance that results from the slow transmission of the reaction substances or reaction products in the solution is called concentration difference polarization. The polarization resulting from the resistance caused by the slow rate of the electrochemical reaction is called electrochemical polarization. The polarization caused by the resistance that is due to oxidation or other substances, with high resistance, was produced when the current flows through the electrode surface is called resistance polarization. The corrosion whose rate is controlled by the electrochemical polarization is called corrosion controlled by activation polarization. For example, when a metal corrodes in the nonoxidizing acid solution without oxygen and other depolarizing agents and if a passivation film does not generate its surface, the corrosion process will generally be controlled by the activation polarization. In this case, the only depolarizing agent is the hydrogen ions, and the reduction reaction of hydrogen ions and the dissolution reaction of anode are controlled by the activation polarization. Before discussing the electrochemical equation of corroded metal electrode, let us review the electrochemical polarization equation of the single electrode reaction first. The general equation of the electrochemical reaction is as follows: R 󴀘󴀯 O x + ne

According to chemical kinetics equation, the rates of the forward and reverse reactions are linked with the reaction’s activation energy. The activation energy of positive reaction (oxidation reaction) and the reverse reaction (reduction reaction) are expressed by W+ and W−, respectively. For the unit electrode surface, the reaction rate is W+

V Ox = k1 · C R · e− RT

and W−

V R = k2 ⋅ C OX ⋅ e− RT .

80 

 3 Electrochemical corrosion kinetics

k1 and k2 in the above two formulas are the proportionality constants and CR and COX are the concentrations of reducing agent and oxidizing agent, respectively. If the reaction rates are expressed by the current density, the two equations above become 󳨀→ = nF · k1 · C R · e−

−W+ RT

←󳨀 = nF ⋅ k2 ⋅ C OX ⋅ e−

−W− RT

i Ox

and .

iR

When there is no current flowing through the electrode, the electrode is in the equilibrium state, whose electrode potential is Eo, corresponding to an exchanging current io. 󵄨 󳨀 󵄨󵄨← → 󵄨󵄨i󳨀R 󵄨󵄨󵄨 < 󳨀 󵄨󵄨 󵄨󵄨 i Ox 󳨀→ ←󳨀 When it is at equilibrium, iOx and i R are in the opposite directions, with equal values. They are all equal to the exchange current at equilibrium potential. Therefore, 󳨀→ 󵄨󵄨󵄨 󵄨󵄨󵄨 i o = iOx = 󵄨󵄨󵄨←󳨀󵄨󵄨󵄨 󵄨󵄨 i R 󵄨󵄨 i o = nFk1 · C R · e

o −W + RT

= nF · k2 · C OX · e

o −W + RT

.

For the charged particles, the change in electrode potential will change their energy, thus changing the activation process. For example, for the oxidation reaction, when the electrode potential shifts more positively, the atoms in the lattice with higher energy can easily leave the surface to the solution. Accordingly, the activation energy of oxidation reaction decreases and the reaction rate speeds up. When the electrode potential is higher than the equilibrium potential for 󳵻E , the activation energy of metal dissolution on the electrode decreases by 󳵻E . For the reduction reaction, it is the opposite; that is, when the electrode potential shifts negatively for 󳵻E , the activation energy of charged ions increases by 𝛼nF󳵻E. The above discussions are shown in Fig. 3.14, which can be expressed by the following equations: W+ = W+o − 𝛽nF ⋅ 󳵻E

3.3 Corrosion process controlled by the activation polarization 

 81

and W− = W−o + 𝛼nF ⋅ 󳵻E,

where α and β are dynamic coefficient, α + β = 1. Typically, α and β preferably are​​ 0.5.

Fig. 3.14: Diagram of activation

  energy in the electrode [6].

When, in the reaction rate equation, considering the impact of electrode potential changes on the activation energy, we can further get the following equations: o −𝛽𝜂F△E) −(W+ 󳨀 󳨀 → RT i Ox = nFk1 · C R · e

= nFk1 ⋅ C R ⋅e = io ⋅ e

o −W− RT

⋅e

𝛽𝜂F△E RT

𝛽𝜂F△E RT

−(W− +𝛼𝜂F△E) ← 󳨀 − RT = nFk ⋅ C ⋅ e i R 2 OX    o

 = nFk2 ⋅ COX ⋅e o   =i ⋅ e

𝛼𝜂F△E RT

−

o −W− RT

⋅e

𝛼𝜂F△E RT

.

Obviously, when the electrode potential does not change, i.e., there is no net current 󳨀→ 󵄨󵄨 󵄨󵄨 flowing through the electrode, △E = 0 , iOx = 󵄨󵄨󵄨󵄨←󳨀󵄨󵄨󵄨󵄨. After the electrode is polarized, 󵄨󵄨 i R 󵄨󵄨

82 

 3 Electrochemical corrosion kinetics

that is, the electrode potential deviates from the equilibrium potential, the current densities for the positive and negative directions are no longer equal. When the anode is polarized, as the potential shifts to the positive direction, the activation energy of the oxidation is less than that of the reduction process, which indicates that the oxidation process can go on more easily. So, 󵄨 󳨀 󵄨󵄨← → 󵄨󵄨i󳨀R 󵄨󵄨󵄨 < 󳨀 󵄨󵄨 󵄨󵄨 i Ox .

The difference between them is the actual measured current density ia. 󳨀󵄨󵄨 󳨀 󳨀 → 󵄨󵄨← i a = i Ox − 󵄨󵄨󵄨i R 󵄨󵄨󵄨 󵄨 󵄨 = i o �exp

𝛽𝜂F△E RT

− exp

−𝛼𝜂F△E RT

�

When the cathode gets polarized, since the potential shifts in the negative direction, the activation energy of the reduction process is lower than that of the oxidation process, indicating that the reduction process is easier to perform, and then 󵄨 󳨀 󵄨󵄨← → 󵄨󵄨i󳨀R 󵄨󵄨󵄨 > 󳨀 󵄨󵄨 󵄨󵄨 i Ox

The difference between them is the actually measured current density ic. 󵄨󵄨← 󳨀 → 󳨀󵄨󵄨 󳨀 i c = 󵄨󵄨󵄨i R 󵄨󵄨󵄨 − i Ox 󵄨 󵄨 = i o �exp

−𝛼𝜂F△E RT

− exp

𝛽𝜂F△E RT

�

Citing the Tafel slope: ba =

2.3RT 𝛽nF

bc =

2.3RT 𝛼nF

The relationship between the above-mentioned current density of anode, cathode, and the potential changes can be rewritten as i a = i o �exp

2.3ΔE ba

− exp

−2.3ΔE bc

�

3.3 Corrosion process controlled by the activation polarization 

 83

and i c = i o �exp

−2.3ΔE bc

− exp

2.3ΔE ba

�.

These are the mathematical expressions of the electrochemical polarization of the single electrode reaction. Now, we apply the above discussion into the corrosion system. When a metal corrodes, there is a pair of conjugate reactions occurring, due to the dissolution of metal and the reduction of depolarizing agents, and the electrode no longer stays in equilibrium. The results of electrochemical polarization are to obtain corrosion potential and corrosion current, between which the relationship can also be expressed through an electrochemical polarization equation. We take the example of the corrosion of Fe in a 0.1 N HCl solution without oxygen. The total reaction of the corrosion process is Fe + 2HCl 󴀘󴀯 FeCl2 + H2 .

According to the mixed potential theory, the above corrosion reaction can be decomposed into a pair of conjugate reactions. At the corrosion potential, the two reactions occur on the Fe electrodes simultaneously: Fe 󴀕󴀬 Fe2+ + 2e

and H2 󴀕󴀬 2H+ + 2e.

Iron sheet dissolves in the hydrochloric acid as Fe electrode, whose dissolution current is 󳨀 → ←󳨀󳨀󳨀 iFe = iFe − iFe2+ .

Hydrogen ions are reduced on the iron sheet, which has H2 electrode in the hydrochloric acid solution, and its reduction current is ← 󳨀 󳨀 󳨀󳨀→ iH2 = iH+ − iH2 .

Under the effect of anode current, the electrode potential of the Fe oxidation reaction moves in the positive direction, while under the effect of cathode current, the equilibrium potential H2 reduction reaction shifts in the negative direction. From the

84 

 3 Electrochemical corrosion kinetics

E-lgi corrosion polarization diagram (shown in Fig. 3.15), it can be found that the relationship between potential and current can be expressed by lines EF and FE. Line AB o o intersects with line CD at point M, which corresponds to E FeFe2+ and lgi FeFe2+ , and line FE o o intersects with line GH at point N, which corresponds to E HH+2 and lgi HH+2 . At the corrosion potential, the Fe electrode doesn’t have net current flowing, and therefore, the dissolution current iFe must be equal to the reduction current ​​i​H ​, which 2 is indicated in the polarization diagram that line AB intersects with line GH at point O, corresponding to the corrosion potential Ek and the corrosion current lgik of Fe in 0.1 N HCl. In other words, the charge reaches balance at point O, and the total current of the oxidation is equal to that of the reduction (shown in Fig. 3.16).

Fig. 3.15: Diagram of the relationship between E-lgi of

  Fe in the HCl solution [1].

Fig. 3.16: Illustration of electric charge balance under

  corrosion potential [1].

In the following, we substitute the single-electrode electrochemical equation into the corrosion system, then we get

iFe =

o �exp iFe/Fe 2+

�

2.3(EFe −E o ) Fe/Fe2+ b a(Fe/Fe2+ )

− exp

−2.3(EFe −E o ) Fe/Fe2+ b c(Fe/Fe2+ )

� �

3.3 Corrosion process controlled by the activation polarization 

i H2 =

� iHo 2 /H+ �exp

−2.3(EH −E o ) 2 H2 /H+ b cH2 /H+

− exp

−2.3(EH −E o ) 2 H2 /H+ b a(H2 /H+ )

�

 85

� �. �

If the Fe electrode is not connected with an external circuit, then � i = iFe + iH2 = 0.

Or iFe = −iH2 = ik. That is, when the external current is zero, the anode dissolution rate of metal Fe is equal to the escaping rate of hydrogen at Fe, which is the rate of metal corrosion. Supposing the surface of Fe sheet is uniform and its and the solution’s resistance is negligibly small, all points at the Fe surface stay at the same potential, when the relationship between the corrosion potential and corrosion current is

iK =

=

o �exp iFe/Fe 2+

2.3(E K −E o ) Fe/Fe2+ b 2+ a(Fe/Fe )

�

� −iHo 2 /H+ �exp

−2.3(E K −E o ) H2 /H+ b cH2 /H+

− exp

− exp

−2.3(E K −E o ) Fe/Fe2+ b 2+ c(Fe/Fe )

� �

2.3(E K −E o ) H2 /H+ b a(H2 /H+ )

�

� �. �

This corrosion potential is a compromise potential whose value is between the equilibrium of the cathode and the anode reactions. When the corrosion potential deviates from ____ their own equilibrium, more than 50 mV (higher than 2.303 ​​  b   ​​ ), the second term of the above formulas, ic and ia, can be ignored. This is because at this time, the Fe electrode mainly has the oxidation reaction, and for H2 electrode, it’s mainly the reduction reaction. Thus, the relationship between corrosion potential and corrosion current can be simplified as

iK =

o iFe/Fe 2+

⋅ exp

� 2.3�EFe −E o Fe/Fe2+ b a Fe/Fe2+

(

)

=

−iHo 2 /H+

⋅ exp

−2.3(E K −E o ) H2 /H+ . b + cH2 /H

Extending the special case of Fe corrosion in 0.1 N HCl to the general, we can get the following formula: i K = i oa exp�

2.3�E oc −E K � 2.3�E K − E oa � � = −i oc exp . baa bcc

where baa and bcc are the Tafel slopes of the anode reaction and the cathode reaction, respectively.

86 

 3 Electrochemical corrosion kinetics

When using the activation overpotential 𝜂 to express the difference between corrosion potential and equilibrium potential and sorting out the formula, we have: 2.3𝜂a ik = exp , where 𝜂a =E K − E oa i oa b aa

and 2.3𝜂c ik = exp , where 𝜂c = E oc − E K . i oc b cc

Taking a logarithm on both sides, we get the following: 𝜂a = −baa lgi oa + baa lgi k 𝜂c = −bcc lgi oc + bcc lgi k

−blgio in the above formula is a constant, expressed by a, and then we obtain the final relationship between the activation potential and corrosion current density: 𝜂 = a + blgi.

This is the famous Tafel equation, which indicates that the relationship between activation overpotential and current density is a semilogarithmic one.

3.3.2 Hydrogen depolarization corrosion The typical corrosion process controlled by the activation polarization is one whose cathodic process is the reduction reaction of hydrogen ions, which is called hydrogen depolarization corrosion or hydrogen evolution corrosion. A necessary condition for the hydrogen depolarization corrosion is that in the corrosion cell, the potential of the anodic metal is more negative than the equilibrium potential of hydrogen electrode, that is, Em < EH. The corrosion of negatively charged metals, such as Fe and Zn, in an oxidizing acid without oxygen and metals with much lower potentials, such as Mg, in neutral or alkaline solutions belongs the hydrogen depolarization corrosion. However, for some oxidizing metals, such as Ti and Cr, despite their potential meeting the conditions of hydrogen depolarization corrosion, as the passivation film increases the anode resistance, as long as the passivation film can stay stable in the acid solution of low concentration, corrosion will not occur.

3.3 Corrosion process controlled by the activation polarization 

 87

In general, the relationship between the overpotential of the anodic dissolution of a metal in the active state and the current density follows the Tafel equation, indicating that the anodic dissolution process of metal goes on essentially, according to the mechanism of electrochemical controlling. Moreover, the slope of an anode polarization curve of a metal is lower than that of a polarization curve of hydrogen electrode, indicating that the hydrogen depolarizing corrosion is controlled by part cathode reactions. The mechanism of hydrogen evolution reaction in aqueous solutions has been described much in the kinetics of electrode process. Hydrogen depolarization reaction is mainly constituted by the following successive steps: (1) Hydrogen ions or hydrated hydrogen ions transfer from the internal solution to the electrode surface and get dehydrated on the electrode surface: H+ (H2 O)S 󳨀→ H+ ⋅ (H2 O)M → H+ + H2 O.

(2) Hydrogen ions bind the electrons which are on the electrode (M) surface and adsorb on the electrode surface: H+ + e → (M − H)ad .

(3) Compositing desorption of hydrogen ions: (M − H)ad + (M − H)ad → H2 + 2M

or electrochemical desorption: (M − H)ad + H+ + e → H2 + M.

(4) Hydrogen molecules form bubbles and escape. In these steps, if there is a step that is proceeding slowly, the other steps’ proceeding smoothly will be affected, leaving the hydrogen depolarization being hampered, which results in the shift of the electrode potential in the negative direction and a certain overpotential. Tafel, from experimental facts, has found that the relationship between hydrogen overpotential and current density conforms to the following relationship: 𝜂H = a + blgi.

Descendants’ work showed that for the hydrogen evolution reaction, when the current density is between 10−9 and 100 A/cm2, the Tafel relationship can be well conformed.

88 

 3 Electrochemical corrosion kinetics

We know that the Tafel relationship reflects the basic characteristics of the electrochemical polarization. Thus, the controlling step of hydrogen depolarization reaction is not step (1) but only step (2) or (3). The slow discharge theory considers the second step is slowest, but the slow compositing theory emphasizes that the third step is the controlling step. According to the slow discharge theory, b in the equation is determined as 118 mv (25°C), which is substantially the same with b values measured in most electrodes, and therefore, it has a more common value. But there are few metals such as platinum whose hydrogen evolution overpotential can be explained by the slow compositing theory. Hydrogen ions in the acidic medium has a large migration rate, and the released hydrogen gas bubbles can stir so that the concentration polarization can often be ignored. In the neutral medium, pH values are not stable as in acidic medium, so in the case of the hydrogen depolarization, the impact of concentration polarization must be considered.

Fig. 3.17: Cathodic polarization curve of

  hydrogen depolarization reaction [7].

The cathodic polarization curve of hydrogen depolarization reaction is shown in Fig. 3.17. At the equilibrium potential E​ ​​  oH ​​ ​, there is no hydrogen that has evolved from the electrode, and the current is zero. With an increase in the current density, the polarization gets intensified. The polarization curve can be divided into two parts: when the current density is very small, namely i < 10−4–10−5 A/cm2, the I-E curve is linear, and when the current density is within a large range, the I-E curve is logarithmic. At a certain current density, the difference between the hydrogen equilibrium potential ​​E​ 0H ​​ ​ and hydrogen evolution potential ​​E​H​​ is the hydrogen overpotential at this current density. 𝜂H = −�E H − E oH � = E0H − E H

In hydrogen depolarization corrosion, the larger the hydrogen evolution overpotential ​​𝜂H​ ​​ is, the more serious the cathodic polarization is and the smaller the corrosion

3.3 Corrosion process controlled by the activation polarization 

 89

rate is. When metal or alloy is subjected to uniform corrosion, if the impurity phases ​​ H,​​ the corrosion rate increases, or alloy phases as a cathode have a relatively lower 𝜂​ and on the contrary, it will reduce corrosion. There are many factors affecting the value of hydrogen potential, mainly the current density, the electrode material, the surface state of electrode, the solution compositions, concentration, temperature, etc. (1) Current density. The Tafel equation mentioned above points out that the relationship between the hydrogen overpotential and the current density is 𝜂H = a + blgi c ,

where a =

−2.3RT 2.3RT − lgi c , b = . 𝛼nF 𝛼nF

For a given electrode at certain compositions and temperature of a solution, a and b are constants. Constant a is related to the condition of the electrode surface, the nature of electrode material, the solution composition, and its temperature, and its value is equal to the overpotential at a unit current density. The larger the value of a is, the higher the overpotential at a given current density is. Constant b is independent regardless of the electrode material. Values of b in hydrogen evolution reactions for a variety of metal cathodes are substantially the same, about between 0.11 and 0.12 V.

Fig. 3.18: Illustration of 𝜂 – lgi of deferent

  metals [7].

The hydrogen overpotentials of different metals have a linear relationship with the logarithmic value of current density, shown in Fig. 3.18. Since the values of b in the Tafel relationship expression are close, these straight lines are substantially parallel.

90 

 3 Electrochemical corrosion kinetics

(2) Electrode materials. The impact of different metal electrodes on the hydrogen evolution overpotential is mainly reflected on the constant a in the Tafel equation, as shown in Tab. 3.2.

Tab. 3.2: Values of a and b in Tafel equations for cathodic hydrogen evolution of different metals (t = 20°C) [8]. Metals

Ag Al Au Bi Cd Co Fe Mg Mn Mo Nb Ni Pb Pt Pd Sn Sb Be Ge Ti Tl W Zn Cu

Acid solution

Alkaline solution

a

b

a

b

0.95 1.00 0.40 0.84 1.40 0.62 0.70 1.41 0.80 0.66 0.80 0.63 1.56 0.10 0.24 1.20 1.00 1.08 0.97 0.82 1.55 0.43 1.24 0.87

0.10 0.10 0.12 0.12 0.12 0.14 0.12 0.114 0.10 0.08 0.10 0.11 0.11 0.03 0.03 0.13 0.11 0.12 0.12 0.14 0.14 0.10 0.12 0.12

0.73 0.64 – – 1.05 0.60 0.76 1.54 0.90 0.67 – 0.65 1.36 0.31 0.53 1.28 – – – 0.83 – – 1.02 0.96

0.12 0.14 – – 0.16 0.14 0.11 0.11 0.12 0.14 – 0.10 0.25 0.10 0.13 0.23 – – – 0.14 – – 0.12 0.12

The data in Tab. 3.2 can divide metallic materials into three categories, based on the hydrogen overpotentials: (1) High hydrogen overpotential metals: a value is between 1 and 1.6 V, such as Pb, Cd, Hg, Tl, Zn, Bi, Sn, etc. (2) Mediate hydrogen overpotential metals: a value is between 0.5 and 1 V, such as Fe, Co, Ni, Cu, Ag, etc. (3) Low hydrogen overpotential metals: a value is between 0.1 and 0.5 V, such as Au, Pt, Pd, W, etc.

3.3 Corrosion process controlled by the activation polarization 

 91

It is believed that the controlling steps of the hydrogen depolarization reactions for these three different metal electrodes are not the same. A metal with low hydrogen overpotential exerts great catalytic activity to the hydrogen ion discharge. Thus, the exchange current densities of the hydrogen evolution reactions for these metals are large (shown in Tab. 3.3). However, their ability to adsorb hydrogen atoms is strong. Therefore, the slowest step in the hydrogen depolarization process of these metal electrodes is the compositing desorption of the absorbed hydrogen atoms. For the metal with mediate hydrogen overpotential such as Fe, Ni, Cu, and other metals, the slowest step in the hydrogen depolarization process is the electrochemical desorption. High hydrogen overpotential metal has the weakest catalytic effect on the discharge reaction, and therefore, the exchange current density of hydrogen evolution reaction is also minimal. The slow discharge of hydrogen ions on these metal electrodes constitutes the controlling step of the hydrogen depolarization. Tab. 3.3: Exchanging current densities of hydrogen evolution of different metals [1] (except as noted, others are the data in 1 MH2SO4). Metals

lg​​i​o​​ (A/cm2)

Metals

lg​​io​ ​​ (A/cm2)

Pd Pt Rh Ir Ni Au Fe W Cu

−3.0 −3.1 −3.6 −3.7 −5.2 −5.4 −5.8 (1.0 N HCl) −5.9 −6.7 (0.1 N H2SO4)

Nb Ti Zn Cd Mn Tl Pb Hg

−6.8 −8.2 −10.3 (/NH2SO4) −10.8 −10.9 −11.0 −12.0 −13.0

The surface state of electrode also has an influence on the hydrogen overpotential. For the same metal, the hydrogen overpotential on the roughened surface is smaller than the one on the smooth surface, as the area of roughened surface is larger than that of the smooth one. (3) Solution composition and its temperature. When there is the presence of positively charged metal ions in the solution, they will be reduced at the electrode surface, which would exert different effects on the hydrogen evolution overpotential. If Pt ions exist in the solution, they will precipitate on the surface of Fe electrodes or Zn electrode. The hydrogen overpotential of Pt is much lower than that of Fe and Zn, and if Pt acts as an additional cathode, it will significantly improve the corrosion rate of Fe and Zn in the acid solution (shown in Fig. 3.19). However, if some solution contains heavy metal salts, such as AsCl, As will precipitate on the electrode surface. At this time, the hydrogen overpotential of metal electrode increases, and the corrosion rate decreases.

92 

 3 Electrochemical corrosion kinetics

Fig. 3.19: Influence of zinc and iron corrosion

  in the acid solution by adding platinum [9].

Additionally, when the solution contains the surface active substances, they are adsorbed on the surface of the metal electrode, hindering the hydrogen evolution reaction, thereby increasing the hydrogen overpotential. So, adding them as an additive agent can play a role as inhibitor. The role of solution pH value on the hydrogen evolution overpotential is that in an acidic solution, the hydrogen overpotential increases with increasing pH value, while in an alkali solution, the hydrogen overpotential decreases with increasing pH. An increase in solution temperature leads to a decrease in the hydrogen overpotential. Generally, if temperature is increased by 1°C, the hydrogen overpotential decreases by about 2 mV. In hydrogen depolarization corrosion, people tend to care about that: in the corrosion process, will hydrogen embrittlement happen? Hydrogen embrittlement is caused by the adsorption of hydrogen into the surface of metal. After the hydrogen atoms are absorbed into the metal, depending on the type of metal materials, the metal surface may experience hydrogen blistering or a metal gets brittle and then fractures. Hydrogen embrittlement generally occurs in ferritic and martensitic steels in sulfide mediums, and so it is often associated with petrochemical production. High-strength steels are particularly prone to experiencing hydrogen embrittlement, and thus, attention should be paid to the pickling process and the welding operation. In the acidic agent, for some certain steels, the cathodic polarization may lead to serious consequences. It is a must to strictly control the protect potential, which is generally not less than −1.0 V (vs. Ag/AgCl reference electrode) when applying the cathodic protection to the engineering structures. Otherwise, it may cause danger of hydrogen embrittlement.

3.4 Corrosion process controlled by concentration polarization 

 93

3.4 Corrosion process controlled by concentration polarization 3.4.1 Steady-state diffusion equation Electrode process kinetics tells us that material movement in a solution can be caused by convection, diffusion, and electromigration. For oxygen concentration polarization, as oxygen is an uncharged molecule, in a static solution, it still mainly relies on the diffusional mass transfer process. For simplicity, here, we discuss only one-dimensional diffusion; i.e., in the three-dimensional orthogonal coordinate, expressing the space position, the concentration only has one coordinate direction. For example, the concentration changes only in the x-axis direction in Fig. 3.20. For each x value, the concentrations in the y-axis direction and z-axis direction are same, which constitutes an equivalent concentration surface (e.g., surface A in Fig. 3.20). The diffusion process is only along the x-axis direction, across through numerous equivalent concentration surfaces.

Fig. 3.20: Illustration of diffusion direction [1].

The expression of Fick’s first law is � = − Di � xi

dC i �, dx

where the negative sign on the right indicates that diffusion direction is opposite to the concentration increasing direction. ∏ xi- – diffusing speed of substance i, mol/ m2·s;   Di – diffusing coefficient of substance i, m2/s; dC i - – concentration gradient of substance i, (mol/m3)/m. dx

94 

 3 Electrochemical corrosion kinetics

In the derivation of Fick’s first law, it is assumed that particle concentration gradient in the solution does not change with time and space. Thus, the law applies only to the steady-state diffusional mass transfer. In the early stage of electrochemical reaction on the electrode surface, due to the consumption of the reaction particles, these particles’ concentrations in the solution near the electrode surface change greatly, leading to the concentration differences. However, when the concentration differences of liquid near the electrode surface appear, the diffusional mass transfer process slowing down the concentration begins. Therefore, when the reaction stays at steady state, the number of particles transferring from the surface liquid layer to the electrode surface can compensate the particles consumed in the electrode reactions, which makes the materials achieve a certain balance. Although at this time the concentration difference of the reactive ion on the surface liquid still exists, it no longer continues to develop and then sets up a so-called steady-state diffusion process. There exists a diffusion layer with a thickness of 𝛿, close to the electrode surface, and the internal liquid mainly depends on the diffusional mass transfer. The thickness of diffusion layer 𝛿 decreases with the strengthening of the solution convection motion. In the case of natural convection, the thickness of diffusion layer is about 0.1 mm, whose schematic diagram is shown in Fig. 3.21. In the case of steady diffusion, the concentration gradient of diffusion layer is equal to the difference of the concentration of the deep solution outside the diffusion layer (that is, the concentration of the whole solution) ​​C​ 0i ​ ​​and the concentration of electrode surface ​​C​ Si ​ ​divided by the thickness of the diffusion layer 𝛿: C o − C si dC i = i . dx 𝛿

Another mathematical expression of Fick’s first law is � = − Di xi

C oi − C si . 𝛿

  Fig. 3.21: Diagram of diffusion layer [1].

3.4 Corrosion process controlled by concentration polarization 

 95

In the corrosion process, the cathodic reduction reaction of the depolarizing agent is often related to the diffusion process, and therefore, here, we discuss only the diffusion process of cathodic reaction. The diffusion process of anodic reaction is basically similar. Now, i means the current density of the cathodic reduction reaction and m represents the moles of the substances reduced after diffusing to the electrode surface through a per unit area. Since the consumed electric quantity in reducing one mole substance is nF therefore: �x =

dm i =− , dt nF

where the negative sign denotes that the direction of cathodic current is usually negative. Arranging the above different mathematical expression of Fick’s first law, we get i = nFD i

C oi − C si . 𝛿

With the cathodic reduction current increasing, the concentration of particle i on the electrode quickly becomes small, eventually approaching zero. Since then, the current no longer increases and reaches a certain limit iL. Once the reduced depolarizer transfers to the surface electrode, it is reduced immediately, and thus, ​​C​ Si ​ ​​ = 0. The above equation can be rewritten as i = nFD i

C oi . 𝛿

iL is called as the limiting diffusion current density. Now, we consider two cases of electrode reaction: (1) The exchange current density of the electrode reaction is large, and even if the external cathodic current flows through the cathode, the electrode reaction can still be recognized as an equilibrium state. This electrode reaction is reversible, whose electrochemical overpotential is zero. According to the Nernst equation, it can be determined that the diffusion overpotential depends only on the ratio of the concentration of the whole solution and the electrode surface concentration. 𝜂d =

RT C s ln nF C o

After rearranging the above equation, we obtain C s = C o ⋅ exp�

nF �𝜂 . RT d

96 

 3 Electrochemical corrosion kinetics

Then substituting C S into i = nFD i

C oi − C si , we finally get 𝛿

i = nFD i [C o − C o exp�

= nFD i C o [1 − exp�

=i L �1 − exp�

nF �𝜂 ] RT d

nF �𝜂 ] RT d

nF � 𝜂 �. RT d

The formula is plotted graphically in Fig. 3.22.

Fig. 3.22: Relationship curve of concentration

  polarization overpotential and current density [10]. For the cathodic polarization, since 𝜂d is negative, 𝜂d is larger and the exponential term on the right side of the expression can be ignored. Then the external current is close to the limit diffusion current density: i ≈ iL .

Further, from two equations: i = nFD

Co − Cs 𝛿

i L = nFD

Co . 𝛿

it can be derived that C s = C �1 − o

i �. iL

3.4 Corrosion process controlled by concentration polarization 

 97

Substituting CS/C 0 into the initial diffusion overpotential relational expression, we get 𝜂d =

RT i ln�1 − �. nF iL

For cathodic polarization, in order to make ​​𝜂​d​​ positive, the above expression can i RT ln� L � . be rewritten as 𝜂d = nF iL − i When i → i L , 𝜂d 󳨀→ ∞ , which indicates that the limit diffusion density that does not change with the electrode potential will appear. Fig. 3.22 clearly illustrates the conclusion. However, it virtually impossible for 𝜂​ ​​ d​​ to become infinite because when ​𝜂​ increases to a certain value, the new electrode reaction will begin. In aqueous solutions, when the cathode potential becomes sufficiently negative, the reduction reaction of hydrogen ions will appear. (2) The exchange current density of the electrode reaction is very small and the balance is easy to get destroyed under the effect of external current. This electrode reaction is irreversible, which can be divided into two cases. 1. If the discharge process through a double electric layer is the sole controlling step, from the previous section, it is already known that the relationship between the absolute value of the cathodic reduction current density and cathodic overpotential 𝜂 is i = i o · exp�−

2.3𝜂 �. bc

The exchange current density i0 is proportional to the concentration of the reactant in the liquid layer near the electrode surface, and the cathode current is related to C0 as well. (In this case, CS and C0 have no significant difference, since the rate of the reactant approaching the electrode surface is much faster than that of the electrode reaction). 2. If the diffusion process is also considered as one of the steps affecting the rate of the entire electrode reaction, the concentration of the reactant close to the liquid layer on the electrode surface is reduced from C0 to CS, and the cathode current at this time should be as follows: i = io ·

Substituting

2.3𝜂 Cs �. · exp�− o C bc

i Cs = 1 − into the above equation, then o C iL i = i o ⋅ �1 −

2.3𝜂 i � ⋅ exp�− �. iL bc

98 

 3 Electrochemical corrosion kinetics

After arrangement, i=

i o · exp�− 1+

io iL

2.3𝜂 � bc

· exp�−

2.3𝜂 � bc

.

Taking the logarithm of both sides, the above equation can be rewritten as 𝜂 = b c lg�1 −

i i � − b c lg o . iL i

Thus, it can be seen that in cases where the electrode reactions are reversible or irreversible, the results are different. In the corrosion process, the dissolved oxygen is an important depolarizer. The corrosion potential of the general metals used industrially is far from the equilibrium potential of oxygen reduction reaction, and the exchange current density of the oxygen reduction reaction is very small, and so this depolarization reaction occurs irreversibly. In this case, the latter derived formulas should be used to handle the overpotential problems caused by diffusion.

3.4.2 Oxygen depolarization corrosion Corrosion where the oxygen reduction reaction is the cathodic process is called oxygen depolarization corrosion or oxygen evolution corrosion. A necessary condition for oxygen depolarization corrosion is that the potential of the anode metal in the corrosion cell is more negative than the equilibrium potential of the oxygen electrode, namely, EM < EO2 . The oxygen reduction reaction in neutral and alkaline medium is O2 + 2H2 O + 4e 󳨀→ 4OH− .

The oxygen reduction reaction in an acidic medium is O2 + 4H− + 4e 󳨀→ 2H2 O.

The relationship between the equilibrium potential of the oxygen reduction reaction and the solution acidity is EOo 2 = 1.23 − 0.059 pH.

3.4 Corrosion process controlled by concentration polarization 

 99

Comparing oxygen depolarization corrosion and hydrogen depolarization corrosion, it is easy to see that oxygen depolarization corrosion has more general significance. Since the potential of the oxygen electrode is much more positive than the potential of the hydrogen electrode, except for a few noble metals, most metals in the solution containing oxygen would corrode in this way. The corrosive agents include humid atmosphere medium, natural water, seawater, and soil, all of which contain a certain amount of oxygen. Therefore, the oxygen depolarization process and metal corrosion under natural conditions are closely linked. The corrosion of most metals in neutral and alkaline solutions, as well as corrosion of a small number of electropositive metals in a solution with the presence of dissolved oxygen and a weak acid solution, belongs to oxygen depolarization corrosion. When metal undergoes oxygen depolarization corrosion, in most cases, the anode process is the active dissolution of metal and the corrosion is controlled by the cathodic process. The rate of oxygen depolarization corrosion depends on the following two factors: 1. The rate of oxygen transferring to the electrode surface 2. The discharge rate of oxygen on the electrode surface Oxygen depolarization corrosion can be roughly divided into three cases: (1) If the corroded metal has a more positive potential in the solution and the transfer rate of oxygen in the corrosion process is large, the corrosion rate is mainly determined by the discharge rate of oxygen in the electrode. At this time, the anodic polarization curves and cathodic polarization curves intersect at the electrochemically controlled area in the oxygen reduction reaction (shown in Fig. 3.23). An example is the corrosion of copper in an open container stirring vigorously.

Fig. 3.23: Diagram of oxygen depolarization corrosion process [11]. 1 – Electrochemical control. 2 – Diffusion   control. 3 – Mixed control.

(2) If the potential of the corroded metal in a solution that stays at an actively dissolved state is negative and the oxygen transfer rate is limited, the corrosion rate

100 

 3 Electrochemical corrosion kinetics

of a metal is controlled by the extreme diffusion current density of oxygen. From the polarization diagram (Fig. 3.23), it can be seen that the anodic polarization curves and the cathodic polarization curves intersect at the oxygen-diffusion controlling areas. Therefore, within a certain potential range, the corrosion current is not affected by the slope and position of the anodic polarization curves, indicating that the corrosion rate of a metal has nothing to do with the nature of this metal. For example, in the sea, there is no significant difference between the corrosion rates of plain carbon steel and low-alloy steel. (3) If the potential of the corroded metal in the solution is quite negative, regardless of the value of the oxygen transferring rate, the cathodic process is constituted by two reactions of oxygen: depolarization and hydrogen depolarization, which are shown in Fig. 3.23. The corrosion current is higher than the extreme oxygen diffusion current density. An example is magnesium corrosion in neutral solution. Oxygen depolarization corrosion can be divided into two basic processes: oxygen transfer and reduction of oxygen molecular at the cathode. Oxygen transfer process includes the following: 1. Oxygen in the air transfers into the solution through the solution interface. 2. Dissolved oxygen diffuses uniformly and is distributed in the solution by convection. 3. Due to the effect of diffusion, oxygen gets through the diffusion layer close to the cathode surface to reach the electrode surface and forms the adsorbed oxygen. The oxygen reduction process is quite complex, for which people do not have a clear understanding of how the hydrogen reduction process goes, since the oxygen electrode reactions are involved with four electrons, whose mechanisms are complex. In one reaction step, there may be generation of a layer of oxide or appearance of particles with intermediate valence, which brings great difficulty to the research. In addition, the degree of reversibility of the oxygen electrode reaction is very low, even if on the metal surface, such as Pt, Pd, Ag, etc., the exchange current densities are no more than 10−9–10−10 A/cm2. Therefore, it is prone to have the polarization under the effect of a tiny external current, causing a change in the structure of the electrode surface. The oxygen reduction process may be carried out in several steps, whose reaction mechanisms are different in acidic solutions and neutral or alkaline solutions. In an acidic solution, the overall reaction is O2 + 4H+ + 4e 󳨀→ 2H2 O.

First step: O2 + 2H+ + 2e 󳨀→ H2 O2 This step is constituted by the following several consecutive steps: O2 + e 󳨀→ O−2

3.4 Corrosion process controlled by concentration polarization 

 101

O−2 + H+ 󳨀→ HO2 HO2 + e 󳨀→ HO−2 HO−2 + H+ 󳨀→ H2 O2

The second step: H2 O2 + 2H+ + 2e 󳨀→ 2H2 O This step also consists of two steps: H2 O2 + H+ + e 󳨀→ 2H2 O + OH OH + H+ + e 󳨀→ H2 O

In neutral and alkaline solution, the total reaction is O2 + 2H2 O + 4e 󳨀→ 4OH− .

The reduction reaction can be divided into a two steps as well: O2 + H2 O + 2e 󳨀→ HO−2 + OH− HO−2 + H2 O + 2e 󳨀→ 3OH−

In the above basic steps, the first step in the acidic solution (i.e., the reaction generating the oxygen ion with half valence) tends to be regarded as the controlling step, and the first step in the alkaline solution (i.e., the reaction forming H​​O​ –2 ​ ​​) is the controlling step. To get a full understanding of the two basic reduction processes, in most cases, the rate of the first basic process determines the total rate of the reduction process, which is relatively complex, including oxygen passing through the interface of the air and the electrolyte solution, and then through the internal solution by mechanical force or thermal convection, and finally through a thin layer closely attached to the metal electrode surface and being adsorbed on the electrode surface. As long as one step proceeds slowly, the oxygen concentration near the cathode would decrease, resulting in concentration polarization, hindering the oxygen depolarization process. Particularly, the closed system does not flow, where the rate of oxygen transfer to the cathode surface is much lower than the rate of electron transfer from the anode so that the oxygen concentration decreases to a very low level, leading to severe concentration polarization. However, if the solution is vigorously stirred, or the diffusion layer on the metal electrode surface is very thin, in a well-inflated condition, it is easy to reach the cathode for oxygen. In addition, if the current density in a cathodic process is very small, the oxygen replenished in time can also keep up with its consumption. Only in the above two cases is the rate of oxygen reduction process determined by the second basic step.

102 

 3 Electrochemical corrosion kinetics

Fig. 3.24: Cathodic polarization curve in the

  oxygen depolarization process [5].

The above discussion can be represented graphically by using a cathode polarization curve of oxygen depolarization process (shown in Fig. 3.24). There are three different areas on this polarization curve: (1) When the cathode current density is small and the electrode surface is supplied fully with oxygen, the overpotential of oxygen reduction reaction and the current density meet the Tafel relationship: 𝜂O2 = a󸀠 + b󸀠 lgi c .

The rate of cathodic reaction is determined by the slow step of the oxygen ionization process. Apart from the current density, the overpotential of oxygen ionization is related to the cathode material, electrode surface state, composition, and temperature of the solution and other factors. Tab. 3.4 lists some overpotentials of oxygen ionization. Tab. 3.4: Overpotentials of oxygen ionization of different metals [5]. Metals Pt Au Ag Cu Fe Ni Carbon Stainless steel

​​η0​ /​​ V (ica = 1 mA/cm2) 2

0.70 0.85 0.97 1.05 1.07 1.09 1.17 1.18

Metals Cr Sn Co Fe3O4 Pb Hg Zn Mg

​​η​0 /​​ V (ica = 1 mA/c​​m​ 2)​​ 2

1.20 1.21 1.25 1.26 1.44 1.62 1.75 ~2.55

When the current density is small, the oxygen overpotential and current density have a linear relationship.

3.4 Corrosion process controlled by concentration polarization 

 103

If the oxygen supply is always sufficient, then the cathodic polarization curves along the ​​Eo​ 0 ​​ ​ curve toward BPC t over a wide range of current density, shown in Fig. 3.24. 2 But in fact, it is very difficult to do this. When ​​i​c​​ > 1/2​​i​L,​​ as the appearance of the concentration polarization, the curve will be different from E​ ​​ o 0 ​​, ​ BPC. Accordingly, unlike 2 the cathodic process of hydrogen ion depolarization, the overpotential curve of the oxygen ionization in the oxygen depolarization is not a complete cathodic polarization curve. (2) When the cathode current density continues to increase, the cathode process is controlled by the mixed effect of the oxygen ionizing reaction and the oxygen diffusion process. In the case of i​​​L/​​ 2 < i < i​​​L​​ , the concentration polarization is caused by the limited rate of oxygen diffusing to the electrode, leading to a rise in polarization curves along PF. At this time, the relationship between the oxygen overpotential and current density is not in accordance with the simple Tafel relationship, which should also include the concentration overpotential in addition to the ionization overpotentials. The mathematical expression is 𝜂O2 = �a󸀠 + b󸀠 lgi c � + b󸀠 lg�1 −

ic �. iL

If ​​ic​ ​​ ≪ ​​iL​ ,​​ the last item of the above right side is close to zero, then return to the Tafel equation. The Tafel equation was proposed in the condition of ignoring the concentration polarization. ​i​c​ According to the assumption that ​​ __  ​​  < 1, the last item of the above formula ​iL​ ​ is negative, which shows that when the concentration polarization appears, the shifting value in the negative direction of the cathode potential would be higher than 1 when there is no value of concentration polarization. (3) The cathode current density is further increased. When ​​i​c​​ ≈ ​​iL​ ,​​ the potential increases and the current does not change to reach a certain limit value, as shown in lines FSN of Fig. 3.24. As the electrode potential moves drastically in the negative direction, the reduction reaction of oxygen is significantly activated. At this moment, as long as oxygen reaches the surface, it immediately combines with electrons. Therefore, compared with the oxygen diffusion step, the electrochemical reaction steps are not control factors. In this concentration polarization region completely diffusion controlled, the relationship of concentration overpotential and current density accords to the derived formula in the last section: 𝜂d =

i RT ln� L � nF iL − ic

= b󸀠 lg�

iL �. iL − ic

104 

 3 Electrochemical corrosion kinetics

(4) The negative move of potential in the process of oxygen depolarization cannot go on indefinitely along the FSN direction because when the cathode potential is sufficiently negative, the hydrogen reduction reaction may occur in the aqueous solution when the cathode process consists of oxygen depolarization and hydrogen depolarization. As shown in Fig. 3.24, when the cathode potential reaches the equilibrium potential of the hydrogen electrode ​​E​ oH ​​ ​, the depolarization process of the hydrogen ions (curve E​ ​​  oH ​​ ​ M in Fig. 3.24) begins to combine with the oxygen depolarization process, appearing as the SQG segment on the polarization curve. The total current density on a cathode electrode is the sum of the current densities of the oxygen depolarization process and the hydrogen depolarization process, namely: i c = −iO2 + iH2 .

The ratio of the total cathode current densities ​​i0​ ​​ and ​​iH​ ​​depends on the nature of 2 2 the metal electrode and the pH value of the aqueous solution. Through the above discussion, we can know that in most cases, the rate of oxygen depolarization corrosion is determined by the value of the diffusion current density iL. All factors affecting the value of iL will affect the rate of oxygen depolarization corrosion. Therefore, increasing the oxygen diffusion coefficient D and the concentration of dissolved oxygen in the solution​​ C​ o​​and reducing the thickness of the diffusion layer 𝛿 can accelerate the rate of oxygen depolarization corrosion. Accordingly, the factors affecting oxygen depolarization corrosion include the following. (1) Influence of the dissolved oxygen concentration. When the dissolved oxygen concentration increases, the limit of current density for the oxygen diffusion will increase and the rate of oxygen ionization reaction will increase as well, and thus, the rate of the oxygen depolarization corrosion increases. Fig. 3.25 illustrates that when the oxygen concentration increases (for example, blowing oxygen into the solution), the starting potential on the cathodic polarization curve appropriately shifts in the positive direction, increasing the limiting diffusion current density accordingly and raising the corrosion potential from ​​E​K ​​ to​​ 1 E​K ​​, and the corrosion current from i​​​K ​​ to ​​i​K ​​. However, if the metal has the char2 1 2 acteristics of corrosion passivation, when the dissolved oxygen is increased to a certain content, because iL has reached the critical current density of passivation to turn the metal from the active state to the passive one, the rate of oxygen depolarization corrosion gets significantly reduced, as shown in Fig. 3.26, with the corrosion current decreasing from iL to iP. Thus, the dissolved oxygen tends to have an opposite double impact on metal corrosion, which has important implications for researching the corrosion behavior of metals such Fe and Zn in neutrally aerated solutions.

3.4 Corrosion process controlled by concentration polarization 

 105

Fig. 3.25: Diagram of the influence of oxygen concentration

  on the corrosion process controlled by diffusion [1].

Fig. 3.26: Influence of oxygen concentration on the

  corrosion rate of passivated metals [1].

Fig. 3.27: Influence of flow velocity on oxygen depolarization corrosion [12].

106 

 3 Electrochemical corrosion kinetics

Fig. 3.28: Influence of NaCl concentration on the rate of iron corrosion [10].

(2) Influence of the solution flow rate and mechanical stirring. Under the condition of the oxygen concentration being unchanged, the limiting diffusion current density is directly proportional to the thickness of the diffusion layer 𝛿. The higher the solution flow rate is or the more intense the mechanical stirring is, the smaller the thickness of the diffusion layer is, and the higher the limiting oxygen diffusion current density is, the larger the corrosion rate is. It can be found in Fig. 3.27 that when stirring the solution or making the solution flow, the corrosion current is increased from ​​i​0​​at the static condition to ​​i​1​​when the solution rate is v​​ ​1​​ or ​​i​2​​when the solution rate is ​​v​2​​. When the limiting diffusion current density of oxygen increases (e.g., ​​v​3)​​ with an increase in flow rate, to a certain extent, the cathode and the anode polarization curves no longer interact in the oxygen diffusion region when the corrosion rate does not increase with the increase in flow rate anymore. When the flow rate further increases to a large extent, under the effect of the liquid with high flow rate, the metal will suffer from cavitation corrosion. Similarly, for a metal or alloy having passivation tendency, when they have not entered a passive state, increasing the flow rate of the solution or enhancing the stirring would make the limiting diffusion current density reach or exceed the critical current density of passivation to promote the formation of a passive state and lower the corrosion rate. (3) Influence of NaCl concentration. With an increase in the concentration of the NaCl solution, because of the increase in solution conductivity, the corrosion rate will increase. For example, when the neutral solution contains 3% NaCl (corresponding to the NaCl content in seawater), the corrosion rate of iron reaches its maximum value. After increasing the salt concentration further, due to the significant decrease in oxygen solubility, iron corrosion rate decreases instead. Fig. 3.28 shows the effect of NaCl concentration on the corrosion rate of iron in a neutral solution. (4) Influence of the solution temperature. Increasing the solution temperature can improve the diffusion rate of oxygen and the electrode reaction rate, and therefore, if the temperature is within a certain range, the corrosion rate would

References 

 107

increase with an increase in temperature. However, the temperature rise causes a decrease in the solubility of oxygen in water. Shown in Fig. 2.29, the maximum corrosion rate appears in the relationship curve of the solution temperature and corrosion rate, which decreases when the temperature exceeds a certain value. In a closed system, the temperature rise increases the gas phase partial pressure of oxygen, which increases the solubility of oxygen in the solution, which offsets the temperature effect of reducing the solubility of oxygen; thus, the corrosion rate always increases with an increase in temperature.

Fig. 3.29: Relationship curve between the limit diffusion current density of oxygen reduction of steel in the 0.5 N NaCl and temperature [13]. 1, 2, 3, 4, 5 – Different rotation rates of rotating   disc steel electrode.

References [1] Huang, Y. C. The Fundamentals of Metals Corrosion and Protection. Shanghai Jiao Tong Univ. Press: Shanghai, 1989. [2] Wei, B. M. The Theory of Metals Corrosion and Application. Chemical Industry Press: Peking, 1979. [3] Evans, U. R. An Introduction to Metallic Corrosion. Edward Arnold: London, UK, 1948. [4] Uhlig, H. H. Corrosion and Corrosion Control, 3rd edition. John Wiley and Sons: New York, 1985. [5] Tomashow, N. D. The Theory of Metals Corrosion and Protection. China Machine Press: Peking, 1965. [6] Cao, C. N. The Fundamentals of Corrosion’s Electrochemistry, 6th edition. Chemical Industry Press: Peking, 2004. [7] Martesohn, E. The Foundation of Corrosion (Huang, J. Z., Zhong, J. L., Trans.). Metallurgical Industry Press: Peking, 1990. [8] Antropor, R. E. The Theoretical Electrochemistry. Higher Education Press: Moscow, 1984. [9] Liu, B. J. Materials Corrosion and Protection. Peking Aviation and Spaceflight Univ. Press: Peking, 1989. [10] Zeng, R. C., Han, E. H. Materials Corrosion and Protection. Chemical Industry Press: Peking, 2006. [11] Li, X. G. Corrosion and Protection of Materials. Zhong Nan Univ. Press: Changsha, 2009. [12] Scully, J. C. The Fundamentals of Corrosion, 2nd edition (Li, Q. Z, Trans.). China Water and Electricity Press: Peking, 1984. [13] Masamichi, K. Metals Corrosion Destroy and Protection Technology (Yuan, B. L., Trans.). Chemical Industry Press: Peking, 1988.

Yuanwei Huang

4 Oxidation and hot corrosion of metals and alloys 4.1 Introduction The oxidation and hot corrosion of metals and alloys can be classified into the category of high-temperature corrosion. The high-temperature corrosion of metals is the deteriorative or destructive process of metals caused by the chemical or electrochemical reaction of metals, under high temperature, with oxygen, sulfur, chlorine, carbon, etc., in the atmosphere. It includes narrow sense oxidation, broad sense oxidation, and hot corrosion. Narrow sense oxidation is the process of oxide formation caused by the reaction between metal and oxygen or an oxidizing medium. Broad sense oxidation extensively indicates the reaction of metal to lose electrons and increase its positive atomic valence. Hence, broad sense oxidation includes the process of formation of oxides, sulfides, chlorides, carbides, and other chemical compounds caused by high-temperature chemical reaction between metal and oxygen, sulfur, chlorine, carbon, etc., in the atmosphere, namely, oxidation, sulfuration, chlorination, carburetion reaction, etc. Metal hot corrosion is an accelerated high-temperature corrosion form caused by metal salt deposited on the metallic surface formed by the reaction of metal and an environmental medium, a simultaneous action under oxygen, sulfur, and other corrosive gases. Metallic oxidation, sulfuration, carburetion, etc., belong to chemical corrosion, while metallic hot corrosion includes acid alkali melt-corrosion mechanism and electrochemical corrosion mechanism. Metallic high-temperature corrosion frequently leads to the deterioration or destruction of metal materials or the descent of material properties; hence, research on high-temperature corrosion and the development of metallic alloys and coatings resisting high-temperature corrosion have an important position in the development of modern science and technology and engineering. It is a key science and technology and engineering problem that must be solved in the industries of energy resources, petrochemistry, power, chemical engineering, etc., and the high-­ technology industrial development of aerospace, nuclear energy, etc. It has important scientific significance and application value.

https://doi.org/10.1515/9783110310054-004

110 

 4 Oxidation and hot corrosion of metals and alloys

4.2 High-temperature oxidation Narrow sense oxidation is the process of oxide formation caused by the reaction between metal and oxygen or an oxidizing medium. It can be expressed by the following reaction equation:

M+

n O 󴀔󴀭 MOn . 2 2

In the reaction, the metal atom M loses an electron and turns into metallic ion to increase its positive atomic valence. And the oxygen atom obtains an electron and becomes oxygen ion. The metal ion and oxygen ion are combined into metal oxide. It is the most simple and basic chemical corrosion reaction, with high-temperature oxidation leading to the damage of metal material properties and the destruction of structure. For metals applied in the industries of mechanical engineering, chemical engineering, power, aerospace, etc., high-temperature oxidation resistance is a key property having the same important significance as the high-temperature mechanical property.

4.2.1 Thermodynamics of metal oxidation reaction Can the metal material spontaneously carry on oxidation reaction under high temperature? How about the stability of the oxide? All these can be analyzed and judged by means of basic knowledge of chemical thermodynamics, and it is generally judged by the change in free energy. The oxidation reaction of metal can be expressed as

M+

n O 󴀔󴀭 MOn .​ 2 2 ​

(4.1)

According to classic thermodynamics, the relation between the reaction free energy (∆G) at a given temperature (T) and the component activity (a) is

ΔG = ΔG0 + RT ln

aM On aM an/2 O

.(4.2)

2

When ∆G = 0,



ΔG0 = −RT ln

aM On aM an/2 O

.(4.3)

2

Δ​​G​ 0​​is all free energy change in the substance taking part in the reaction, under standard state (for gaseous reactant and resultant, it is the state in which its

4.2 High-temperature oxidation 

 111

partial pressure is one atmospheric pressure, and for liquid and solid state, its pure state at one atmospheric pressure is adopted as standard state), namely, the difference between the standard resultant free energy of the substance taking part in the reaction and the standard resultant free energy of the reaction product.

So, ΔG0 = ΔG0MOn − ΔG0M −

n 0 ΔG (4.4) 2 O2

0 And because ∆​​G​M  ​​ ​ ​​​∆G​​O0 ​​ , ​ the resultant free energy of pure elementary substance 2​ ​ is 0,

0 0 so ΔGMOn = ΔG = −RT ln

aM O N aM an/2 O

.(4.5)

2

The activity of pure substance equals to 1. a​ ​​ M​​ = ​​aMOn ​ ​​ = 1, and the activity of oxygen (​​a​​O​2​​​) is approximately equals to its partial pressure P​​o​2​​under low pressure and high temperature; under the balance state, it is, namely, the decomposition pressure of oxide (​P​O​0 ​​ ) ​​ . 2

So ΔG0MOn = −RT ln



1 (P0O )n/2 2

=

n RT lnP0O2 .(4.6) 2

This is the relation equation of the decomposition pressure of oxide and the standard resultant free energy.

P0O2 n n n n 0 Yet ΔG = ΔG − RT lnPO2 = RT lnPO2 − RT lnPO2 = RT ln .(4.7) 2 2 2 2 P O2 0

Here, ​​P​​O0 ​​​  ​​ is the decomposition pressure of MOn at a given temperature (T), and ​P​​O​2​​ is 2 the partial pressure of oxygen at gaseous phase. According to ∆G value, we can judge the oxidation reaction: When ​P​​O​2​​ >​​ P​​O0 ​​​  ​​ , namely, ∆G < 0, the reaction goes on toward the direction of pro2 ducing MOn. When ​ P​​O​2​​ = ​​P​O​0 ​​​ , ​​ namely, ∆G = 0, the reaction of metal oxidation reaches 2 balance. When ​P​​O​2​​ < ​​P​​O0 ​​​  ​​ , namely, ∆G > 0, the reaction goes on toward the direction of MOn 2 decomposition; Therefore, by means of calculating the ∆G value or comparing the oxygen partial pressure ​P​O​ 2​ ​​ at practical atmosphere with the decomposition pressure of oxide at this temperature, the possibility of oxidation reaction and the direction of oxidation reaction can be determined.

112 

 4 Oxidation and hot corrosion of metals and alloys

4.2.2 Drawing and application of the figure of oxide standard resultant free energy versus temperature The thermodynamical calculation of oxide standard resultant free energy is relatively difficult and complicated. For the convenience of application, in 1944, Ellinggham [1] drew a ∆​​G​ 0​​-T figure of the oxide standard resultant free energy ∆​​G​ 0​​ versus temperature T. The oxide standard resultant free energy at T temperature can be read by means of a diagram, and accordingly, the possibility and the direction of oxidation reaction can be judged. In 1948, Richardson and Jeffes [2] added balance oxygen pressure (​P​O​ 2​ )​​ and (CO/​​CO​2,​​ ​​H​2/​​ H ​​ ​2​​O) three types of gaseous partial pressure supplementary coordinates to the ∆​​G​ 0​​-T figure, which constituted the RichardsonJeffes figure (see Fig. 4.1). This made the content of the ∆​​G​ 0​​-T figure richer, and thus, it was more widely used. From the ∆​​G​ 0​​-T figure (shown in Fig. 4.1), we can immediately read the change value (∆​​G​ 0​​) of the standard resultant free energy of metal oxidation reaction at any temperature. The method of reading the standard resultant free energy of metal oxide by diagram method is, as per Fig. 4.1, finding out this metal oxidation reaction line and making a vertical line passing a given temperature at a temperature horizontal coordinate, which intersects the reaction line. We can thus read the change value of the standard resultant free energy (∆​​G​ 0​​) of the metal oxide at a given temperature from the vertical coordinate corresponding to the point of intersection. The more negative ∆​​G​ 0​​is, the more stable the metal oxide is, the stronger the power of this metal seizing oxygen by reduction is, and thereby, the stability of metal oxide at standard state can be judged, and the possibility of a metal reducing another metal oxide can be anticipated. The affinity to oxygen of any metal at the lower part of the ∆​​G​ 0​​-T figure is greater than that of any metal at the upper part of the ∆​​G​ 0​​-T figure. Therefore, any metal at the lower part can reduce the oxide of the upper metal. For example, for aluminum and iron at 600 °C standard state, all can be oxidized, and the reaction occurs as follows:

4/3Al + O2 → 2/3Al2 O3 ; ΔG0Al2 O3 = −928 KJ(4.8)



2Fe + O2 → 2FeO; ΔG0FeO = −417 KJ (4.9)

0 Because ∆​​G​A​ 0l​  O​​ ​ ​​​ ​is more negative than ​​∆G​FeO   ​​,​  the oxide of aluminum is more stable, the 2 3 oxidation tendency of aluminum is stronger than that of iron, and the oxide of iron can be reduced by aluminum. The reaction is



2FeO + 4/3Al = 2/3Al2 O3 + 2Fe; ΔG0600 ∘ C = −511 KJ.(4.10)

4.2 High-temperature oxidation 

 113

Fig. 4.1: ∆​​G ​ 0​​-T figure of metal oxidation reaction [3].

0 ​​∆G​600   °C​​ ​​is a negative value, so at 600 °C, it is possible to reduce FeO by aluminum. This is why iron oxide can be reduced by aluminum as deoxidizer in steelmaking [4]. Three types of supplementary coordinates, ​​P​​O​2,​​​ ​​P​H​ 2​ /​​​ P ​​ ​​H​2​O,​​ and ​​P​CO/​​ P ​​ ​​CO​2​​​ are respectively listed in the rightmost side (including the uppermost part and the lowermost part) of Fig. 4.2. The O point, H point, and C point marked in the line on the leftmost side of the ∆​​G​ 0​​-T figure are respectively the corresponding original points when seeking for ​​P​​O​2,​​​ ​​P​​H2​ /​​​ P ​​ ​H​ 2​ ​O,​​ and ​​P​CO/​​ P ​​ ​C​ O​2​​​under the balance state by means of the figure line. For

114 

 4 Oxidation and hot corrosion of metals and alloys

Fig. 4.2: Graphic schematic diagram defining the decomposition pressure of metal oxide (​​Al​2O ​​​​ ​3)​​ and the ratio of balance gas components by ∆​​G​ 0-​​ T figure [5].

example, corresponding to a reaction of metal with oxygen, ∆G = RT ln ​​P​​O​2;​​​ when T = 0, ∆​​G​ 0​​ = 0; namely, the 0 point is located at (0,0) coordinate point. The balance oxygen partial pressure (oxide decomposition pressure) at a given temperature can be sought, and the direction and possibility of oxidation reaction under actual oxygen pressure can be judged by means of P​ ​​ ​O​2​​​ coordinate system of the ∆​​G​ 0-​​ T figure. The figure method seeking for the decomposition pressure of metal oxide is as follows: Make a vertical line passing a given temperature at temperature horizontal coordinates, which intersects the metal oxidation reaction line at S point. Link the 0 point in the leftmost side in a straight line with the S point and extend it to intersect the P​ ​​ ​O2​ ​​​ coordinate and ​​P​​O​2​​​at the intersection point, namely, balance oxygen pressure of MOn at this temperature (or the oxide decomposition pressure) [4, 5] (shown in Fig. 4.2). Take A​​l​2O ​​​​ ​3​​as an example; it can be obtained using the figure that its balance partial pressure at 1600  °C equals to 4 × ​​10​ −16​​ atm. Based on the ∆​​G​ 0-​​ T figure and ∆​​G​ 0-​​ P ​​ ​O​ ​2​​​ coordinate system, the stability of metal oxide under a given temperature and oxygen partial pressure can also be judged easily. The judgment method is as follows: first, under the coordinate system ∆​​G​ 0-​​ P ​​ ​O​ 2​ ​​​, make a straight line passing the original point 0 and the point of known oxygen partial pressure at the ​​P​​O2​ ​​​ coordinate (constant

4.2 High-temperature oxidation 

 115

pressure line), then make a vertical line passing a given temperature at the ∆​​G​ 0-​​ T coordinate system, and the vertical line will respectively intersect the oxidation reaction line and the constant pressure line; the stability of oxide can simply be judged from the relative position of these two focus points. If the intersection point at the oxidation reaction line is above the intersection point at the constant pressure line, the oxide will decompose. Conversely, the oxide will be stable. The reaction of metal with water vapor or ​​CO​2​​to generate metal oxide is also a common oxidation reaction. If the reaction relates to C and CO​ ​​ 2​​ or ​​H2​ ​​ and ​​H2​ O ​​ , the ∆​​G​ 0-​​ T figure and ​​P​CO/​​ P ​​ ​C​ O​2​​​ and ​​P​H​ ​2/​​​ P ​​ ​H​ 2​ ​O​​coordinate system can also be used respectively to seek the balance ratio of P​ ​​ CO/​​ P ​​ ​​CO​2​​​ and ​​P​​H2​ /​​​ P ​​ ​​H2​ ​O​​. The concrete method is as follows: the C point and H point at the left vertical line can be linked respectively with the S point at the reaction line and extended to intersect the ​​P​CO/​​ P ​​ ​C​ O​2​​​ or ​​P​H​ ​2/​​​ P ​​ ​H​ ​2​O​​ coordinate, and the component ratio of balance gaseous phase in these two mixed gases (​​P​CO/​​ P ​​ ​​CO​2​​​ and ​​P​​H​2/​​​ P ​​ ​​H​2​O​​) can be sought out. The stability of metal oxide at a given temperature and given P​ ​​ CO/​​ P ​​ ​​CO​2​​​ or ​​P​H​ 2​ /​​​ P ​​ ​​H2​ ​O​​can also be sought out by means of the same method as the above-mentioned ∆​​G​ 0-​​ P ​​ ​O​ ​2​​​ coordinate system. 4.2.3 Metal oxidation dynamics and mechanism 4.2.3.1 Metal oxidation dynamics The rule of metal oxidation dynamics is generally expressed by the mathematical formula of the change in oxidation weight gain or oxide film with time and determining the constant temperature dynamics curve of the oxidation process is the basic method of studying the oxidation process dynamics. The reaction of metal with oxygen depletes metal and oxygen and forms oxide on the metallic surface. If the oxide film is complete and stable, the mass of the metallic sample will increase. The oxidation velocity of the metal can generally be expressed by the weight gain of unit area and sometimes by the thickness of the oxide film, the thickness reduction of metal sample, the oxygen partial pressure in the system, or the absorbing quantity of oxygen on unit area. The rule of the metal oxidation dynamics is generally expressed by the metallic constant temperature oxidation dynamics curve (namely, the ∆m-t curve). It can not only provide the information related to the oxidation mechanism, such as the velocity control factors of the oxidation process, the protectiveness of the film, the reaction velocity constant, the reaction activation energy, etc., but also be generally used as the basis of engineering design. The rule of metal oxidation dynamics is related to the metal element, oxidation temperature, and time. Different metals follow different oxidation rules. The same metal may follow different rules at different temperatures. Even at the same temperature, with the lengthened oxidation time and increased thickness of the oxidation film, its dynamics rule may also transform from one type into another

116 

 4 Oxidation and hot corrosion of metals and alloys

type. To sum up these oxidation rules, we can classify these dynamics curves into five types: straight line, parabola, cube, logarithm, and inverse logarithm. A schematic diagram of these five metal constant temperature oxidation dynamics curves is shown in Fig. 4.3 [3].

Fig. 4.3: Schematic diagram of five types of metal constant temperature oxidation dynamics curves.

(1) Straight line rule: when metal oxidizes, if there is no resulting protective oxide film or the gaseous phase or liquid phase products depart from the surface due to reaction, then the oxidation velocity is directly determined by the reaction velocity. Hence, the oxidation weight gain or the thickness of oxide film is in direct proportion to the oxidation time, namely,

dy = k or y = kt + C.(4.11) dt

In this equation, y is the mass change of the unit area of sample or the thickness of the oxide film, k is the oxidation velocity constant, and C is the integral constant (relating to the original metal state). When the metal oxidizes according to the straight line rule, its oxidation velocity is constant and these metals do not have antioxidation property. Magnesium and alkaline-earth metals, as well as tungsten, molybdenum, and vanadium, oxidize, all accord with the straight line rule. (2) Parabola rule: when the metal oxidizes, the square of its oxidation weight gain or the thickness of oxide film is in direct proportion to time, which is the parabola rule. Its expressive equation is



Y2 = 2kt + c; after diferentiating, it can be obtained that

dy k = .(4.12) dt y

4.2 High-temperature oxidation 

 117

In the equation, y is the oxidation weight gain or the thickness of oxide film; k is the parabola velocity constant, which is an important parameter content; and c is the integral content, which represents deviation from the parabola at the initial reaction stage; for example, at the beginning stage, the temperature does not reach constant or the sample is overheated by the reaction heat. It is known from the equation that oxidation velocity is in inverse proportion to the weight gain or film thickness. The parabola rule of the oxidation reaction indicates that the oxide film has protectiveness, and with the increment of film thickness, the oxidation velocity decreases. The oxidation of many metals at a relatively wide temperature range adheres to the parabola rule, such as the common metals of iron, nickel, and copper, at most ranges of temperature. (3) Cube rule: the cube of the oxidation weight gain or the thickness of oxide film is in direct proportion to time, and the oxidation velocity is in inverse proportion to the square of the oxidation weight gain or film thickness. Namely:

y3 = 3kt + C, or

dy k = 2 .(4.13) dt y

Here, k is the velocity constant and c is the integral constant. As compared with the parabola rule, when the metal oxidizes, following the cube rule, its oxidation velocity decreases at a faster velocity; namely, this type of metals has a better antioxidation property. For zirconium at 600–900 °C and copper at 800–1000 °C, their oxidation in air adheres to the cube rule. The oxidation of some metals at low temperature in thin oxide film is also in accord with the cube rule. (4) Logarithm rule and inverse logarithm rule: when the metal oxidizes at a low temperature (such as 10% Cr 4–5% Al 2–3% Al

Produce selective oxidation Selective oxidation temperature (°C) generated oxide 1100 C​​r2​ O ​​​​ ​3​​ 1100 A ​​​ l​2​O3​ ​​ 700 ​​​Al​2​O3​ ​​ ​ 3​ ​​ 1000 ​​​Al​2O

(2) Producing better protective, structure stable composite oxide film of spinel type. At a high temperature, an alloy with suitable components can form spinel- or iron-olivine-type multielement composite oxide film. Its complicated compact structure leads to a large increase in the activation energy needed by the ion movement in this structure, a decrease in the movement velocity, and thereby an increase in the oxidation resistance of the alloy. Up until now, the diffusion mechanism of ion in the lattice is still not clear, but its excellent oxidation resistance is obvious. For example, for Fe-Cr alloy containing about 10% Cr, although its Cr content has not yet reached the critical concentration to produce selective oxidation, it cannot form a single Cr​ ​​ 2O ​​​​ ​3​​protective film but it can form an FeO-​​Cr​2O ​​​​ ​3​​spinel-type composite oxide film; and for Ni-Cr alloy, it can form a NiO-​​Cr​2O ​​​​ ​3​​film and thereby increase the oxidation resistance. The composite oxide produced in Fe-Si alloy is not spinel type, but 2FeO • ​​SiO​2​​ composite oxide of iron-olivine type. It is an amorphous composite oxide composed of ​​SiO​2​​and FeO. The amorphous ​​SiO​2​​under a relatively wide range of temperature is stable, and therefore, adding Si into Fe will increase its oxidation resistance. The formation of composite oxide film depends on appropriate alloying components and heating at the temperature to produce the composite oxide. However, in order to increase oxidation resistance, the composite oxide film is also required to have 1. a high melting point; 2. a low steam pressure; and 3. a low diffusion velocity of ion among it. For example, the steam pressure of Al​ ​​ 2O ​​​​ ​3​​is lower than that of C​​r​2O ​​​​ ​3​​, and its diffusion velocity is lower. Thus, adding simple Al into Fe and adding Cr and Al into Fe are both very effective for resisting high-temperature oxidation. (3) Decrease the lattice defects of the oxide film, and reduce the diffusion velocity of ion. Most of the metal oxides are electronic semiconductors or mixed conductors of electrons and ions. Now, take as an example the oxide of electronic semiconductor to explain that by adding different alloying elements and through the

4.2 High-temperature oxidation 

 131

control of the lattice defects of oxide, the oxidation resistance of the alloy will be enhanced. When the oxide is metal excess-type semiconductor, namely, n-type semi­ conductor, which contains excess metal components, such as ZnO, CaO, BeO, ​​V​2O ​​​​ ​5,​​ ​​P​bO ​​​​ ​2,​​ Mo​​O​3​​, etc.: In the lattice interstice of these oxides exist excess metal ions and electrons, and the metal oxidation is mainly determined by the diffusion of metal ions passing the interstice; if at this time, the metal capable of forming higher valence metal ions is added, the oxidation resistance of the metal will be increased. If zinc is oxidized to form ZnO and ​​Zn​ 2+​​interstitial cations will be produced. If aluminum is added into zinc, during oxidation, ZnO will be mixed with​​ Al​2O ​​​​ ​3.​​ Two ​​Al​ 3+​​displace two ​​Zn​ 2+,​​ which will decrease one ​​Zn​ 2+,​​ thus decreasing the interstitial zinc ions in ZnO and leading to the reduction in the electric conductivity of cation and then the decrease in the growth velocity of ZnO. When the oxide is metal insufficient-type semiconductor, namely, p-type semiconductor, such as NiO, etc.: In this type of oxide, excess oxygen exists; thus, in the normal lattice of metal ion, vacancy exists. If in this type of metal, the metal capable of forming metal ions with lower atomic valence is added, the ­oxidation resistance of the metal will be increased. For example, adding low+ valence metal ion ​​Li​  ​​ into NiO will lead to a decrement of metal ion vacancies in NiO, a reduction in the growth velocity of NiO, and thus, an increase in its oxidation resistance. (4) Enhance the combining and adhesive power of oxide films to the matrix, which is the active element effect. That a Cr​ ​​ 2O ​​​​ ​3​​ or ​​Al​2O ​​​​ ​3​​film formed by the selective oxidation on alloy surface makes the alloy have certain oxidation resistance. During temperature change, the peeling off phenomenon of the oxide film often occurs. Hence, enhancing the combining and adhesive power of oxide film to the matrix to increase the anti-peeling off property of the oxide film is the main aspect of improving the oxidation resistance of alloy. Adding active trace elements into the alloy, such as the rare-earth elements, titanium, zirconium, hafnium, etc., can obviously decrease the oxidation velocity of the alloy and enhance the adhesion of the oxide film. This is the so-called active element effect. For example, adding rare-earth elements into heat-resistant steel and alloy will obviously increase their high-temperature oxidation resistance. Adding rareearth elements Ce, La, Y, etc., into Fe-Cr-Al alloy will obviously increase their service temperature and life. Although the action mechanism of active elements such as rare-earth elements has not been completely clarified, it has been verified by research results time and again. It is generally recognized that 1. Active elements decrease the content of the critical aluminum or chromium forming the ​​Al​2O ​​​​ ​3​​ or ​​Cr​2O ​​​​ ​3​​ protective film. 2. Active elements obviously reduce the growth velocity of ​​Cr​2O ​​​​ ​3​​and change the growth mechanism of ​​Cr​2O ​​ ​3​​ film.

132 

 4 Oxidation and hot corrosion of metals and alloys

3. Active elements obviously increase the adhesion of oxide film and enhance the combining and adhesive power of oxide film to matrix. After adding rare-earth elements into the alloy, it is usually observed that the localized oxide film penetrates in branch shape deeply into the alloy inside along the grain boundary at the alloy/oxide film interface, and its bulging parts assume the shape of a needle, flake, or post. This morphology can increase the actual contact area of oxide film and metal, lengthen the spread distance of crack along the interface, and thereby increase the adhesive power of oxide film to the matrix, taking the action of “pinning.” This is exactly the pinning effect of rare-earth elements. It enhances the adhesive power of oxide film to the matrix, decreases the cracking trend of oxide film, and promotes the formation of homogeneous and compact protective film. Hence, adding active trace elements, enhancing the combining and adhesive power of oxide film to the matrix, is also one of the effective ways to increase the oxidation resistance of alloy.

4.2.5 Oxide film and its basic properties 4.2.5.1 Growth process and control factors of metal oxide film The growth process of metal oxide film includes oxygen and other gas molecules reacting with surface metal, finally forming a layer of continuous and compact oxide film, and isolating the metal from the gas environment; the oxide film continuously grows and thickens through the diffusion of metal ions or oxygen, links of the oxidation course. In these aspects, a lot of research work has been carried out, many achievements have been obtained, and basic understanding has been established, but the details are still to be further researched. Once the oxide film is formed, the progress of the oxidation process to be continued will be determined by two factors: (1) Interface reaction velocity, including the reaction at two interfaces: metal/oxide and oxide/gas. (2) Diffusion velocity of the substance taking part in the reaction through the oxide film, including the diffusion caused by concentration gradient chemical potential and the migrate diffusion caused by the potential difference of electric potential gradient. These two factors in fact control the whole process of the continued oxidation and thus also control the velocity of further oxidation. They both may become the actual control factor. When the reaction on the metal surface with oxygen forms an extremely thin oxide film, the interface reaction plays a leading role; namely, at this time, the control

4.2 High-temperature oxidation 

 133

factor is the interface reaction. But along with the growth and thickening of the oxide film, the diffusion process is gradually playing a more and more important role and becomes the control factor of the continued oxidation. The diffusion of metal ions and oxygen through the oxide film generally may occur in three ways: (1) The one-way outward diffusion of metal ions, and thus, the reaction is carried out at the oxide/gas interface, such as the oxidation process of copper (Fig. 4.10(a)). (2) The inward diffusion of oxygen, and thus, the reaction is carried out at the metal/ oxide interface, such as the metallic oxidation of Ti, Zr, etc. (Fig. 4.10(b)). (3) Diffusion in two directions, namely, the outward diffusion of metal ions and the inward diffusion of oxygen are being carried out at the same time, both of them meeting each other in the oxide film and the reaction goes on, such as the oxidation of Co (Fig. 4.10(c)). Different metals have different ways of growing oxide films.

Fig. 4.10: Schematic diagram of the diffusion of metal ions or oxygen in the metal oxidation process [4]. (a) Outward diffusion of metal ions. (b) Inward diffusion of oxygen. (c) Two-way diffusion.

4.2.5.2 Protectiveness of the oxide film When metal is oxidized, oxide film is formed on the metal surface. The oxide film is generally in solid state, but under higher temperatures, the oxide of some metals is in liquid state (e.g., the melting point of V ​​ 2​ O ​​​​ ​5​​is 674 °C), and the oxides of some metals are highly volatile (e.g., ​​M​0O ​​​​ ​3​​begins volatilization above 450 °C). If the oxide is in liquid state or volatile, the metal surface is continuously exposed to oxidation medium, and the oxidation will go on rapidly. If the oxide is in solid state, the oxide film formed on the metal surface will hinder the substance transmission between metal and medium to a certain degree. Thereby, the oxide film has constant protectiveness. The protectiveness of the metal oxide film is determined by many factors such as the completeness and compactness of the film, change in surface layer volume resulting from the oxide film formation, heat stability of the oxide film and its structure and thickness,

134 

 4 Oxidation and hot corrosion of metals and alloys

relative thermal expansion coefficient of the film and metal and combining strength, stress in the film, etc. It is necessary to make specific analysis according to different metals and different situations. 4.2.5.3 Completeness of the oxide film The completeness of the oxide film is the necessary condition for the protectiveness of the oxide film, and the necessary condition for maintaining the film’s completeness is that the volume of oxide must be more than that of the metal consumed by oxidation, which is the so-called Pilling-Bedworth principle, or PB ratio for short. They indicate that the necessary condition for the completeness of the oxide film is that the volume of the metal oxide films resulting from the oxidation (​​V​OX​​) must be more than that of the metal consumed by the formation of these oxide films (​​V​M)​​ , namely: VOX /VM > 1(4.32)



𝜐=

VOX Md Md M/D = = > 1.(4.33) = VM nA/d nAD mD

In this equation: M – molecular weight of metal oxide N – metal reactive atomic valence in metal oxide A – metal atomic weight m – metallic weight consumed by oxidation, nA d – density of metal D – density of metal oxide υ – PB ratio Obviously, only when υ > 1 can the metal oxide film completely cover the surface and thus have protectiveness. When υ < 1, the formed oxide film is not complete and cannot completely cover the surface to form a loose and porous oxide film and thus has no protectiveness. For example, the alkali metal and alkaline-earth metal oxides MgO and CaO are just like this. But υ > 1 is only a necessary condition for the protectiveness of the oxide film, not its sufficient condition. The film may produce stress in the growth process, result in film cracking, and thereby decrease or even lose its protectiveness. When υ » 1, this situation will appear. For example, the PB ratio of tungsten oxide film is 3.4, which has a poor protectiveness. The PB ratio of molybdenum oxide film is 3.3, but it will produce volatilization under high temperature, so its protectiveness is also not good. Practice shows that it is better to take the PB ratio of an oxide film with better protectiveness as slightly greater than 1. For example, the PB ratio of Al oxide film is 1.28, which leads to a very good protectiveness. The PB ratios of some metal oxide films are listed in Tab. 4.4.

4.2 High-temperature oxidation 

 135

Tab. 4.4: PB ratios of some metal oxide films [6].

Oxide ​​K​2O ​​ C​​s​2O ​​ ​​​Rb​2​O3​ ​​ ​​Li​2O ​​ ​​Na​2O ​​ CaO SrO BaO MgO BaO

PBR < 1

PBR ≈ 1

PBR

Oxide

PBR

Oxide

0.45 0.47 0.50 0.56 0.57 0.57 0.64 0.65 0.69 0.80

​​La​2O ​​​​ ​3​​ ​​​Y​2​O3​ ​​ ​​Nd​2O ​​​​ ​3​​ ​​Ce​2O ​​​​ ​3​​ ​​CeO​2​​ ​​Er​2O ​​​​ ​3​​ TiO ​​In​2O ​​​​ ​3​​

1.10 1.13 1.13 1.15 1.17 1.20 1.22 1.23

𝛼-​​Al​2O ​​​​ ​3​​ PbO ​​SnO​2​​ 𝛾-​​Al​2O ​​​​ ​3​​ ​​ThO​2​​ ​​Ti​2O ​​​​ ​3​​ PtO Zr​​O​2​​ ZnO PdO

PBR ≫ 1

PBR

Oxide

PBR

Oxide

PBR

1.28 1.28 1.31 1.31 1.35 1.47 1.56 1.57 1.58 1.59

​​ u​2O C ​​ NiO BeO Si​​O​2​​ CuO CoO TiO W​​O​2​​ ​​Co​3O ​​​​ ​4​​ FeO

1.67 1.70 1.70 1.72 1.72 1.74 1.76 1.87 1.98 1.78

​​Cr​2O ​​​​ ​3​​ 𝛼-​​Fe​2O ​​​​ ​3​​ ​​ReO​2​​ I​​rO​2​​ ​​Ta​2O ​​​​ ​5​​ ​​Nb​2O ​​​​ ​5​​ Mo​​O​3​​ Os​​O2​ ​​

2.02 2.15 2.16 2.23 2.47 2.74 3.27 3.42

4.2.5.4 Structure types of metal oxide The oxide formed on metal or alloy may be amorphous (noncrystalline), but also in a crystal state. As a general rule, a very thin oxide film ( 1250 °C), Mo​​O​2,​​ Fe​​S​2,​​ (Os​​S2​ )​​ , (Ir​​O2​ )​​ , Ru​​O2​ ,​​ PbS

4.2.5.5 Structure of matrix metal and directional adaptability of the growth of oxide film In the initial stage of oxidation, namely, the thin film growth stage, oxide film and matrix may have the directional adaptability of structure orientation, which is beneficial to the growth of oxide film; this mainly refers to the fact that the atom distribution in the lattice of metal and oxide film may present similarity, and the initial growth of oxidation may continuously develop on the metal matrix in a certain crystallographic orientation, namely, the so-called “extensional growth” of oxidation. The crystallographic orientation of the extensional growth of oxide film and that of the matrix metal have a certain corresponding relation, and with the obvious change in the lattice parameter of oxide film, pseudo-crystal oxide is formed. The thickness of this pseudo-crystal oxide is very thin (generally 1, press stress is formed in oxide films; when υ < 1, pulling stress is formed. (3) The recrystallization of oxide film in the thickening caused by high-temperature growth may produce or change the stress state. In fact, this and stress interact as both cause and effect. (4) The chemical components near the oxide film/metal interface have changed, such as the segregation of rare-earth trace element yttrium at the metal/oxide film interface, which will change the stress state of the interface. (5) The movement of lattice defect vacancy in the oxidation process. An example is the metal movement as a result of vacancy poured into the vicinity of the metal/ oxide film interface or its disappearance in metal, which will cause or change the stress state of the interface oxide film. (6) The production of new oxide phase in the oxide film will cause stress. For example, when copper is oxidized, the new oxide phase may form in the cracking of oxide film, thus producing press stress. Someone thinks: when nickel is oxidized at high temperature, oxygen may diffuse fast into the oxide film along the grain boundary, and new oxide will form in the film, thereby producing press stress. (7) Stress is produced as a result of the difference in the heat expansion coefficients between metal and oxide. For example, in the cooling process of cyclic oxidation, due to the difference in the linear expansion coefficients between metal matrix

4.2 High-temperature oxidation 

 139

and oxide film, pulling stress is produced in the oxide film. The stress in the oxide film may obtain complete or partial slacking or release through the plastic deformation of the film itself, the plastic deformation of the matrix metal, the separation of oxide film and matrix metal, the destruction of film, etc. The dislocation slip and climbing in the oxide film are generally considered the main mechanisms of plastic deformation of oxide film.

4.2.5.7 Hole, cracking, and destruction of oxide film Many irregular holes are frequently seen in the oxide section, which are usually distributed near the oxide film/metal interface. For example, when nickel is oxidized at high temperature, its inner layer is often equiaxed crystal layer containing many holes (Fig. 4.13). According to the oxide film completeness condition when PB ratio υ < 1, it is sure to form incomplete oxide film with holes or cracks; when PB ratio υ > 1, complete oxide film will be formed. In the complete oxide film, the formation of hole is related to the mechanism of vacancy injection. Near the metal/oxide film interface, the collection of supersaturated vacancies will form a hole. The mechanical destruction of cracking of oxide film is the result of stress release after the stress in the oxidation reaches a certain degree. The common destruction forms are rise, bulge and cracking, microvesicle forming, peeling, shearing, and edge and corner cracking of the oxide film (Fig. 4.14). In most cases, these destructions are related to the lateral press stress in the film. The destruction form of the oxide film is related to the factors of intensity of the lateral stress, strength, and plasticity of the film and combining power of the oxide film/metal.

Fig. 4.13: Holes in the oxide film inner layer when nickel is oxidized at high temperature [13].

140 

 4 Oxidation and hot corrosion of metals and alloys

Fig. 4.14: Various types of destructions in the growth of oxide film [3]. (a). Hole of film without cracking. (b) Cracking hole. (c) Gas impermeable microvesicle. (d) Peeling. (e) Cracking of notch. (f) Cracking of edge and corner.

The stress concentration of the oxide film is the largest at the sharp corner and edge side or the sharp turnings. Moreover, when the metal changes into the oxide, its volume will suddenly change; therefore, the oxide film is most likely to crack and peel at the edge and corner.

4.3 High-temperature alloy materials 4.3.1 Strengthening of alloy, high-temperature resistance principle, and classification of high-temperature materials The metal materials applied at high temperature generally are called high-­temperature alloy. At present, the high-temperature alloys applied in engineering mainly are iron base alloy, nickel base alloy, and cobalt base alloy. To be applied safely in long-term at high temperature, these high-temperature alloys must meet three property requirements: first is having better high temperature mechanical property, second is having excellent high temperature corrosion resistance, and third is having high temperature structure stability. In a certain sense, the high temperature mechanical property is the most basic requirement, and even more important in practice. These high temperature mechanical properties mainly include high temperature creep property, high temperature endurance property, high temperature fatigue property, high temperature slackness property, high temperature endurance plasticity and may be called heat-strengthening requirement, relating to the heat strengthening of alloy with high temperature resistance. 4.3.1.1 The strengthening theory and strengthening of high-temperature alloy The typical structures of iron-base, nickel-base, and cobalt-base high-temperature alloy basically are alloying austenite structures and strengthening phase dispersed

4.3 High-temperature alloy materials 

 141

and distributed in austenite, such as various carbide phases or intermetallic compound phases, some carbides or intermetallic compounds strengthening the grain boundary, and suitable content of trace elements partially gathered nearby grain boundary to further strengthen or purify the grain boundary. According to the analysis of the effect of alloy structure on heat strengthening, high-temperature strengthening requirements are as follows: (1) increase the resistance of dislocation movement on slip plane, in order to increase the deformation resistance of slip or deformation mechanism; (2) decrease the dislocation diffusion type movement process, in order to inhibit the diffusion-type deformation; and (3) improve the grain boundary structure, increase the strengthening effect of grain boundary, and decrease the slip deformation of grain boundary [13]. Figure 4.15 shows the location of iron, nickel, and cobalt in the element periodic table and the acting relation of those elements with other elements [13]. According to their relation with the location situation of iron, nickel, and cobalt in the periodic table, other elements can be divided into four types, namely, the elements nearby iron, nickel, and cobalt (I type), whose atomic size, crystal structure, and electron layer structure are similar, able to form continuous solid solution; the slightly far away elements (II type) compose definite solid solution; some more far away elements form the ionic compound (III type); and those without interaction (IV type). According to this, we can divide the applicable alloy elements into the following: (1) Elements that can form continuous or limited solid solution with bigger dissolubility: manganese, copper, and noble metals such as rhodium, platinum, etc.; (2) elements that can form limited solid solution with certain dissolubility: ­chromium,

Fig. 4.15: Periodic table of alloy elements to form solid solution with iron, nickel, and cobalt.

142 

 4 Oxidation and hot corrosion of metals and alloys

molybdenum, tungsten, zirconium, hafnium, beryllium, etc. They and iron, nickel, and cobalt are the nearby family. When its content exceeds the limit value of dissolubility, it will form exactly the second phase. These second phases may have a strengthening effect or may be harmful; (3) nonmetal elements whose differences in atomic sizes are extremely big but can form interval solid solution with less dissolubility: carbon, boron, nitrogen, etc.; and (4) elements that can have only extremely little or even no dissolubility: magnesium, lanthanum, cerium, calcium, barium, etc. (metal elements). These elements generally gather partially in the grain boundary. In addition, among the high-temperature alloys, lead, tin, arsenic, antimony, bismuth, etc., are harmful impurity elements with low melting point, which should be decreased as much as possible or be completely eliminated; silicon, phosphorus, sulfur are common impurity elements, whose dissolubilities are small, and they generally gather partially on the grain boundary and are not favorable for strengthening of the grain boundary. As shown from the above analysis, the most common way of strengthening ironbase, nickel-base, and cobalt-base high-temperature alloy may be solid solution strengthening, precipitated phase strengthening, and grain boundary strengthening. (1) Solid solution strengthening Increasing heat strengthening through alloy elements’ solid dissolution into the matrix is called solid solution strengthening. It is a matrix strengthening method widely applied in high-temperature alloys. For example, in Ni-Cr base alloy, we generally add high-melting-point metals like W, Mo, Ta, Nb, etc. (V, VI family), to commonly act with Cr. A larger amount of this type of metal dissolves into γ-Ni, which will decrease the stacking-fault energy (SFE), increase the width of diffusion dislocation, increase the activation energy of diffusion, and retard the slip and intersection slip of dislocation in high temperature creep condition, thereby increasing the heat strengthening of heat-resisting alloy. The heat strengthening enhancement of alloy through solid solution strengthening is mainly realized by increasing the atom binding force and lattice divergence change of alloy element. When temperature ​​T ≤ 0.6T​melting point​​ (absolute temperature), the increase in atom binding force and divergence change in lattice increase the slip resistance in the solid solution, making it difficult to produce slip deformation, thereby reaching the objective of strengthening. When temperature T ​​ > 0.6T​melting point,​​ the effect of increasing atom binding force is especially important; the addition of alloy elements decreases the diffusion power of the elements in the solid solution and increases the recrystallization temperature, thereby retarding the progress of diffusion deformation. Figure 4.16 indicates the effects of solid solution strengthening and electron vacancy number caused by lattice constant change in binary nickel base alloy. As shown in Fig. 4.16, the greater the difference in atomic valences between the solute and the solvent is, the greater the difference of lattice strains is, and the more obvious

4.3 High-temperature alloy materials 

 143

Fig. 4.16: In binary nickel base alloy, the effect of (a) solid solution strengthening and (b) electron vacancy number (Nv) caused by lattice constant change [15].

the strengthening effect. Add Cu, Co, Fe, Mo, W, Cr, and Ti into Ni to change the lattice constant, thereby causing the increment of yield strength. And in the same change of lattice constant, the strengthening effect caused by Ti with the biggest atomic valence difference is the biggest; the strengthening effect caused by Fe, Co, Cu with less atomic valence differences is the least. The increment of electron vacancy number (​​N​v​​) increases the yield stress, as shown in Fig. 4.16(b). Besides, the effects of different elements on SFE are also different; they are caused by the concentration difference of solute atom in stacking-fault from the mother phase. The dislocation is required to pass this type of partial gathering area of solute atom and must overcome more supplementary energy barrier, just to realize the strengthening objective. When temperature ​​T ≥ 0.6T​melting point​​, the diffusion deformation mechanism takes great effect in high-temperature creep, and the effect of realizing solid solution strengthening through lattice divergence change is limited; thus, the more important thing is to decrease diffusivity in order to retard diffusion-type deformation. The relation between the self-diffusion coefficient of different metals and tempera​​  −13– ture is shown in Fig. 4.17 [15]. As shown in this figure, taking D at 10​ ​​ ​​10​ −14​​ ​​cm​ 2/​​ s as requirement, the order of the realized temperatures of different metals and their heat strengthening is tungsten, molybdenum, tantalum, niobium, chromium, nickel, cobalt, iron, titanium, and aluminum. At high temperature, such as the solid solution strengthening of nickel base alloy at 1000 °C, the diffusion factor appears to be especially outstanding. The strengthening effect of tungsten is more excellent than that of molybdenum, and cobalt has a large quantity of solid solution, thus effectively decreasing its SFE, which is extremely favorable to high-temperature strengthening. Polybasic alloying can increase the

144 

 4 Oxidation and hot corrosion of metals and alloys

Fig. 4.17: Self-diffusion coefficients of different metals.

heat strengthening of high-temperature alloy by decreasing the diffusion process factor. According to the dislocation theory of high-temperature strengthening, polybasic alloying is favorable to the formation of inhomogeneously distributed group of solute; this will increase the resistance of dislocation slip, to retard the progress of diffusion deformation. An example is retarding the dislocation climb, which may produce a better effect of solid solution strengthening. Hence, in practice, it is generally applied to increase the heat strengthening of high-temperature alloy through polybasic alloying. This point is more effective at temperature T > ​​0.6T​melting point​​. Table 4.6 lists the data of polybasic nickel base alloy and diffusion feature. As shown in Tab. 4.6: polybasic alloying is very effective in increasing heat strengthening of nickel base alloy. Certainly, in Tab. 4.6, besides relating to solid strengthening, polybasic alloying also has the effect of precipitated phase [12]. In addition, the effect of solid solution strengthening is not only related to the unit strengthening effect of the added element itself but also directly related to the adding quantity of element. Within the dissolubility range, we can add solid solution element as much as possible, to fully take the effect of solid solution strengthening. (2) Precipitated phase strengthening (or called precipitation hardening) The strengthening caused by dedissolution from supersaturated solid solution and the precipitation dispersed second phase is called precipitation phase strengthening. It is a main method used to increase heat strengthening of heat-resisting steel and high-temperature alloy. The dispersed carbide phase VC in heat-resisting steel, the dispersed γ’ phase [​​Ni​3​​(AlTi)] in iron base heat-resisting alloy, and the γ’ dispersed phase in nickel base heat-resisting alloy with a high volume fraction of 30–60% are all examples of precipitated phase strengthening.

Diffusion feature

Heat strengthening

​​D​0​​ 1.6 ×

​​ 0​ −9​​ 1

216.1

Q/ (kJ/mol)

/(​​cm​ 2/​​ s)

400 (0.39Tm)

Single Ni

​​ 100 hour σ ​ = ​​ 14.7 ​​MP​a​​ ​​t​O0 ​​​ 

Feature factor

–

286.7

~ 700 0.56Tm

Binary Ni-Ti

3×

​​ 0​ −13​​ 1

351.1

800 0.63Tm

Three part Ni-Cr-Ti

3.8 ×

​​ 0​ −13​​ 1

366.2

950 0.71Tm

Five part Ni-Cr-Ti-W-Al

Alloy system

6.2 ×

​​ 0​ −13​​ 1

381.6

970 0.72Tm

Six part

Tab. 4.6: Effect of polybasic alloying on diffusion feature and heat strengthening of nickel base alloy.

5.6 ×

​​ 0​ −13​​ 1

410.5

980 0.73Tm

Seven part

8.1 × ​​10​ −13​​

467.36

1,020 0.75Tm

Eight part

4.3 High-temperature alloy materials   145

146 

 4 Oxidation and hot corrosion of metals and alloys

To realize precipitated phase strengthening, the following conditions must be met: first, considering the phase figure, it should have a curve of change in solid solution dissolubility with temperature; at the same time, the composition of the alloy is in a two-phase area; second, the precipitated phase itself should have good enough high temperature stability and good heat strengthening, while the distribution of the precipitated phase should disperse homogeneously, and in long-term high-temperature application condition, it is not easy to gather and grow. Thus, it can be seen that the strengthening effect of age precipitated phase and the precipitated phase type in alloy, the crystal structure, chemical composition, and size and quantity are closely related to stability. From the point of view of strength theory, the strengthening effect of the precipitated phase is related to the interaction of dislocation and the precipitated phase. As shown in Fig. 4.18, the dislocation forms a stress field obstacle through the precipitation of second phase. Especially in the second phase coherent precipitation, owing to the existing high elastic stress field, flow stress greatly increases; when the dislocation passes the second phase particle through climb, it should overcome the barrier; when the dislocation passes in a curve through the precipitated phase particles or cuts across the precipitated phase particle, it all should overcome the barrier. These barriers of dislocation movement can all increase the strain strength of alloy [13]. Another effect similar to the strengthening principle of precipitated phase is the so-called dispersion strengthening. It realizes the strengthening effect by use of additional high-dispersion, high-hardness particles. The difference between dispersion strengthening and precipitated phase strengthening lies in the following: the dispersed particle is not produced from dedissolution of matrix but is additional. These particles mostly are high-melting-point oxides, whose heat

Fig. 4.18: Schematic diagram of moving dislocation and second phase interaction and some possibilities to overcome the barrier.

4.3 High-temperature alloy materials 

 147

­stability is high, firmly combined into the matrix, and does not react with the matrix. The dispersion strengthening alloys are mostly produced by using the powder metallurgy method. (3) Grain boundary strengthening The crystal structure of the grain boundary is nonregular, the atomic arrange is complicated and disordered, the lattice is twisted, and at the same time, there also exist various crystal defects (such as dislocation, vacancy, etc.). Therefore, the grain boundary at high temperature is a weak link. The diffusion of grain boundary is faster than that of the inside crystal, therefore helping to lead to creep crack diffusion along the grain boundary, to form grain boundary cracking vertical to the received force direction. Therefore, strengthening the grain boundary is an important link to increase high temperature strength of high-temperature alloy. Strengthening the grain boundary can be done by two important ways: one is to decrease as much as possible or to avoid that the grain boundary contains harmful impurity elements to weaken the grain boundary; the second is to add a suitable amount of microelements to effectively strengthen the grain boundary. The general content of impurity elements in alloy is very small, and the dissolubility in alloy is also very low, some of which have a very low melting point or can produce low-melting-point eutectic. They generally partially separate out at the grain boundary, with a higher concentration at the grain boundary. The distribution of these impurities’ partial gathering at the grain boundary will seriously weaken the grain boundary, decreasing its high temperature strength. Therefore, in high-temperature alloy production, the control of harmful impurity elements is very strict. Take nickel base alloy as an example: American Aeronautical Materials Specification AMS 2280 specifies: the content of bismuth, thallium, tellurium, lead, selenium (fire types of elements) is controlled within 0.5 ppm–5 ppm. At the same time, 15 types of impurity elements (antimony, arsenic, cadmium, gallium, germanium, gold, indium, mercury, potassium, sodium, thorium, silver, tin, uranium, and zinc) are controlled to be under 50 ppm. Their general content does not exceed 400 ppm. Sulfur is another type of harmful element among the high-temperature alloys. When containing sulfur, it is easy to form low-melting-point eutectic at the grain boundary, seriously deteriorating the hot working and heat-strengthening of alloy. Therefore, in high-temperature alloy, its sulfur content generally should be controlled to be under 0.015%; if the requirement for high-temperature alloy is high, its sulfur content should be controlled within 0.015 and 0.007%. Regarding gaseous elements, such as oxygen, nitrogen, etc., their effect on the properties of high-temperature alloy is also very big. They mainly appear in alloy as impurity, which easily reunites at the grain boundary. Purging and strengthening the grain boundary by eliminating harmful impurities and gases are realized generally by applying the smelting technology of vacuum smelting and adding some trace alloying elements. The other way is to suitably add elements to effectively strengthen the

148 

 4 Oxidation and hot corrosion of metals and alloys

grain boundary. The most common elements used to strengthen the grain boundary are boron, zirconium, magnesium, and calcium. Boron is a common trace element applied in strengthening the grain boundary among high-temperature alloys. Its content in alloy is very small; but as a result of the localized alloying caused by its partial gathering at the grain boundary, which intensely changes the grain boundary state and decreases the diffusion of element at the grain boundary, the grain boundary is therefore strengthened. For the precipitation process of iron, a nickel base high-temperature alloy, boron mainly affects the precipitation of grain boundary carbide or some other intermetallic compounds. And particularly, boron can inhibit a harmful layer like “cellular precipitation” of the precipitated phase and improve the concentrated nonhomogeneous distribution state of the precipitated phase, thereby being favorable to heat strengthening. Zirconium and boron have a similar effect, but the former is weaker. Moreover, zirconium has a bigger dissolubility in nickel and can also form Ni​ ​​ 3​​Zr intermetalic compound with nickel and get into the γ’ phase to replace partial aluminum and titanium atoms, thereby promoting the strengthening of the γ’ phase. Therefore, generally, at the same time, we add boron and zirconium into iron, nickel base high-temperature alloys to realize better strengthening effect. The alkali-earth metals of magnesium, calcium, etc. (with content less than 0.01%), in iron, nickel base high-temperature alloys can improve its endurance strength and plasticity by strengthening the grain boundary. In recent years, they have obtained widespread application in various high-­ temperature alloys [13]. Alkali-earth metals like magnesium mainly partially gather at the grain boundary to produce inner-adsorption and affect the morphology and distribution of the grain boundary precipitated phase, fill up vacancies, retard the diffusion of grain boundary in creep process, prevent fast development of crack along the grain boundary, and thereby being favorable to increasing heat strengthening of alloy.

4.3.1.2 Strengthening high-temperature corrosion resistance A high-temperature material is required to have not only good heat strengthening but also excellent high temperature corrosion resistance. Two possible ways to increase the oxidation resistance of alloy have been discussed in Section 4.2. To increase the high temperature corrosion resistance of high-temperature alloys, the most basic thing is to realize the corrosion-resisting alloying of alloys; another important way is exactly to apply suitable high-temperature protective coating on the surface of high-temperature alloy. With regard to corrosion-resisting alloying, for iron, nickel, cobalt base alloy, the main alloying elements used to increase their high temperature corrosion resistance are chromium, aluminum, silicon, etc. These alloys generally contain a substantial amount of chromium, aluminum, and silicon (elements); through selective oxidation to form ​​Cr​2​​O​3​​, ​​Al​2​​​​O​3​​, etc. (protective oxide film), on alloy surface, basic high temperature corrosion resistance of alloy can be ensured. At the same time,

4.3 High-temperature alloy materials 

 149

we generally add titanium, zirconium, hafnium, rare-earth elements, etc. (active elements), to produce the effect of active elements, for further strengthening the protectiveness of oxide film and the binding force of oxide film with the matrix, promoting peeling off resistance property, therefore obviously decreasing high temperature corrosion velocity of alloy. Certainly, the design of high temperature corrosion-resisting alloying and heatresisting alloy is a complex theory and practice problem. It considers not only temperature and atmosphere (environmental conditions) but also type of oxide film, structure, thermodynamic stability and growth kinetics, surface morphology of the oxide film, microstructure, growth velocity, self-repair property, etc. (many problems). It has many complicated influencing factors. Therefore, a concrete alloy/ environment system needs concrete analysis to specifically solve relevant related high-temperature corrosion problems. 4.3.1.3 Structure stability of high-temperature alloy As for long-term application of high-temperature alloy at high temperatures, we should consider not only its heat strengthening, high temperature corrosion resistance but also its structure stability at high temperature. If alloy is applied long-term at high temperatures, a brittle phase will appear, leading to a sharp decrease in its strength, plasticity, and shock toughness; we need to manage to give prevention and treatment. For example, some iron-chromium alloys, iron-chromium-nickel alloys, and nickel base alloys at high temperature can produce σ phase, which is a hard and brittle intermetallic compound phase. The formation of σ phase may greatly decrease the strength, plasticity, and shock toughness of alloy, even leading to cracking. This is extremely harmful to some high-temperature alloys. Figure 4.19 presents a binary phase figure of a Cr-Fe system. As shown in Fig. 4.19, at 820–870 °C, when iron content is 49–56%, an intermetallic compound FeCr is formed, namely, σ phase [15].

Fig. 4.19: Binary phase figure of Cr-Fe.

150 

 4 Oxidation and hot corrosion of metals and alloys

Fig. 4.20: 870 °C isothermic section of Fe-Cr-Ni phase figure.

Figure 4.20 is an 800 °C isothermic section of an Fe-Cr-Ni phase figure. It indicates that at 800 °C, within certain a Fe, Cr, and Ni content range, σ phase exists [16]. The σ phase is an intermetallic compound of tetragonal lattice; its unit cell has 30 atoms, it has a coordination number of 15, and it is a topologic close-packed structure belonging to electron compound. Its formation is controlled by electron factor and is related to electron vacancy number. However, the formation of the σ phase generally has a pregnant period: it often begins to grow in alloy only after application at high temperature for a certain period of time. Figure 4.21 indicates that after long application at about 650 °C, in Incoloy 800 alloy (Fe-20Cr-32Ni), an σ phase appears. The σ phase generally appears as needle-like, flat-like lump or grain structure. Due to its hard-brittleness, cracks are produced in the alloy, leading to cracking [17]. Therefore, in order to prevent the formation of the σ phase to maintain the stability of high-temperature alloy structure, we may control the electron vacancy number of the residual solid solution in the design of alloy, to make it to be less than a certain critical value and thus avoiding the growth of the σ phase. For example, for denatured nickel base alloy, we can make its electron vacancy __ number of residual solid solution N​​ ​​ v  < 2.45 through the design of alloy; for casting nickel base alloy, we can make its electron vacancy number of residual solid solu__ tion ​​N​​v  < 2.32. However, generally, it can eliminate the hidden trouble of hightemperature σ phase precipitation.

4.3 High-temperature alloy materials 

 151

Fig. 4.21: Appearance of σ phase in Incoloy 800 alloy and its crack produced after 2000 h at a high temperature of 650 °C.

(4) Development and classification of high-temperature alloy With the development of aviation, aerospace, energy resources, chemical industry, petrochemical, etc. (industries), according to the requirements of application environment and property, metallurgic scientific and technological circles work hard to research and develop various high-temperature alloys, to basically satisfy the needs. Figure 4.22 lists various high-temperature alloys researched and produced along with age development [13]. High-temperature alloys can be classified according to different methods. According to manufacturing processes, they can be mainly classified into casting high-­ temperature alloy, deforming high-temperature alloy, powder metallurgy and mecha­ nical alloying high-temperature alloy, directional solidification high-temperature alloy, and single-crystal high-temperature alloy. According to composition classification, they can be mainly divided into iron base high-temperature alloy, nickel base high-temperature alloy, and cobalt base high-temperature alloy. Along with the increase in applied temperature and requirements for various high temperature comprehensive properties, high-temperature alloys are also under continuous development.

4.3.2 Nickel base high-temperature alloy Nickel has a face center cube crystal structure, with a melting point of 1454 °C, and its structure is very stable and there is no allotropic transformation. These features

152 

 4 Oxidation and hot corrosion of metals and alloys

Fig. 4.22: Change in applied temperature of heat-resisting steel and high-temperature alloys with age.

are very suitable for application as high-temperature material. Therefore, it has an important position among high-temperature alloy materials. Suppose we take 150 Mpa, 100 h endurance strength as standard, at present, the highest temperature that nickel base high-temperature alloy can bear is about 1100 °C, that of cobalt base alloy is about 950 °C, and that of iron base alloy is under 850 °C, and that of nickel base alloy is correspondingly higher by about 150–250 °C. Nickel has a very good alloying capability, which seldom forms a harmful phase with other alloying elements, and this exactly provides the potential possibility to improve various properties of nickel. Thus, nickel base alloy is broadly applied in aero-engine manufacturing, the hottest end part of various gas turbines, such as the working blade of turbine part, guide blade, turbine plate, combustion chamber, etc.

4.3 High-temperature alloy materials 

 153

4.3.2.1 Alloying of nickel base high-temperature alloy According to the strengthening mechanism of high-temperature alloy, nickel base alloy can respectively realize solid solution strengthening, precipitated phase strengthening, and grain boundary strengthening through alloying. Moreover, some elements also have the effect of resisting high-temperature corrosion. Among them, some can take both the effect of high-temperature strengthening and that of resisting high-temperature corrosion. Nickel base alloy generally contains ten-odd types of alloying elements, whose effects, respectively, are the following: Alloying elements resisting high-temperature corrosion: chromium, aluminum, tantalum, niobium, etc. Solid solution strengthening elements: tungsten, molybdenum, chromium, cobalt, vanadium, tantalum, etc. Precipitated phase strengthening elements: titanium, aluminum, niobium, zirconium, etc. Grain boundary strengthening elements: boron, zirconium, magnesium, cerium, carbon, etc. (1) Effect of chromium One of the important effects of chromium in nickel base alloy is that it increases its high temperature corrosion resistance. Figure 4.23 is the binary phase figure of nickel-chromium [13]. As shown in the figure, chromium can dissolve into nickel, but the dissolvability of chromium in nickel at room temperature is not high. When the chromium content in alloy is greater

Fig. 4.23: Binary phase figure of Ni-Cr.

154 

 4 Oxidation and hot corrosion of metals and alloys

than 10%, a compact layer of ​​Cr​2O ​​​​ ​3​​film can form on its surface, which has very good corrosion resistance and oxidation resistance. Hence, the chromium content in nickel base alloy is generally between 10% and 20%. From the point of view of strengthening effect, chromium in nickel base alloy mainly takes the effect of solid solution strengthening, only to form a few amounts of ​​Cr​23C ​​​​ ​6​​-type carbide; chromium solid solution strengthens the matrix grain and precipitates carbide (​​Cr​23C ​​​​ ​6)​​ , therefore affecting the strength of alloy. And the ​​Cr​23C ​​​​ ​6​​grains formed at the grain boundary and their distribution state have a great effect on the high temperature endurance strength of alloy. Practice indicates that the chromium content in nickel base alloy should be maintained above 12%, to reach the effect of resisting high-temperature corrosion and strengthening. (2) Effect of titanium and aluminum on the structure and phase composition of nickel base alloy Titanium and aluminum are the main elements of precipitated phase strengthening of nickel base alloy. Aluminum and titanium are the main elements in forming strengthening γ phase (​​Ni​3​​Al), and titanium is the main element in forming strengthening η phase (​​Ni​3​​Ti). As shown from its structure, aluminum is in phase balance between the γ phase and γ’ phase in solid solutions, titanium is in phase balance between the γ phase and ŋ phase in solid solutions, while the γ’ phase or ŋ phase undergoes changes in dissolubility with temperature. Hence, we can produce dispersion strengthening by means of heat treatment. The strengthening phases that are also possible besides γ’ and ŋ phase are carbide TiC, Ti(CN), and Cr​ ​​ 23C ​​​​ ​6​​. In view of this, aluminum and titanium are the most important alloying elements among the nickel base alloys. But the increment of aluminum content in alloy (>3–4%) will make alloy difficult to have deformation processing. The development of vacuum smelting technique can increase the aluminum content of alloy to 6–7% and the general content of aluminum and titanium to 7–8%, and thus to fully take the effect of aluminum and titanium. (3) Effect of molybdenum, tungsten, vanadium, and niobium Many nickel base alloys contain different quantities of molybdenum, tungsten, vanadium, niobium, etc., alloying elements, which all increase the high-temperature strength of alloy in various degrees. These elements all belong to difficult-to-melt metal, with a high melting point and high density. Research indicates that when carrying out the comprehensive alloying of these elements, its strengthening effect is bigger than that of single addition. Molybdenum and tungsten are solid solution strengthening elements and are also forming elements of carbide. When the carbon content in alloy is very low, molybdenum and tungsten do not dissolve into the strengthening phase but take effect by means of increasing the stability of solid solution, thus increasing its recrystallization temperature. When alloy contains carbon (or boron), it can form carbide (or boride), thus taking the effect of strengthening phase.

4.3 High-temperature alloy materials 

 155

Some nickel base alloys have a low vanadium content ( 5%; or when Cr + W + Mo = 20, Al > 6%; or when Cr + W + Mo = 17, Al > 7%, an α phase area will appear. So, the adding quantity of ferrite formed elements into the nickel base alloy should be controlled appropriately. (4) Effect of boron, zirconium, magnesium, cerium These are all grain boundary strengthening elements. They gather at the grain boundary, retard the progress of diffusion process, and increase heat strengthening. Adding of rare earth elements is also helpful to enhancing the adhesive force of oxide film and improving the alloy’s property of high temperature oxidation resistance.

4.3.2.2 Heat treatment of nickel base alloy The heat treatments of nickel base alloy are mainly solid solution and aging. But we should select different heat treatment processes according to different requirements for the composition, structure, and properties of a specific alloy. Sometimes, its heat treatment process is very complicated. (1) Diffusion annealing The aim of diffusion annealing is to have homogeneous composition of alloy and decrease the segregation degree. Generally, we heat the alloy to 1150–1200  °C at a suitable heating velocity, then after heat preservation for a long time, rapidly cool it down (air cooling or water cooling). For some nickel base alloys, we may also use this process to eliminate Laves phase in casting ingot. This type of heat treatment is also called homogeneous annealing. (2) Solid solution treatment. The basic requirement to meet solid solution treatment is to obtain supersaturated solid solution and appropriate grain size. Before the solid solution treatment, the alloy is in deformation state, containing one order γ’ phase, second order γ’ phase, various carbides and boride, etc. In the solid solution treatment, we generally heat the alloy to solid solution temperature, then after a short period of heat preservation

156 

 4 Oxidation and hot corrosion of metals and alloys

(generally 10 minutes), cool it down in a certain way. The basic change in heating in solid solution treatment is the dissolution of the above-mentioned various phases. And the dissolution temperature of various phases is related to the saturation degree of alloy and the properties of precipitated phase. The growth of grain is related to the dissolution of the second phase. Temperature increase and the extension of heat preservation time accelerate the dissolution of the second phase. The more the second phase dissolves, the faster the grain will grow. Therefore, selecting the solid solution temperature should mainly be according to the requirement for grain size. For some alloys, we need to carry out solid solution treatment twice; in the second treatment, the solid solution temperature may be lower than that of the first treatment, and its aim is to control the precipitation of carbide at grain boundary and to regulate the size of the γ’ phase. The cooling velocity after solid solution treatment should be defined according to the requirement for maintaining enough supersaturated solid solution. It can be judged according to TTT curve (Fig. 4.24) [13]. Curve 1 in Fig. 4.24 represents the beginning of phase precipitation. Curve 2 represents the end of phase precipitation. ​​υ​1– ​​ υ ​​ 5​ ​​represent different cooling velocities, and ​​υ​4​​is critical cooling velocity. When υ > ​​υ4​ ​​, we can obtain supersaturated solid solution.

Fig. 4.24: Isotherm decomposition lines of nickel base alloy solid solution.

(3) Aging After solid solution treatment, we can carry out aging treatment of alloy. The object of aging is to produce a supersaturated solid solution to precipitate a suitable amount of homogeneously distributed and dispersed precipitated phase, to strengthen the alloy. Because alloy can precipitate the γ’ phase, various different carbide and boride phases in the aging process, the aging temperature, and the cooling velocity of alloy can be defined according to the precipitation TTT curve of every type of precipitated phase. For some alloys, we need to carry out aging two or more times. To ensure the structure stability of alloy, we generally should select the highest aging temperature, which is higher than the application temperature. The heat preservation time of aging should ensure enough precipitation of strengthening phase, but without overaging occurring. The feature of nickel base alloy aging is to ensure an enough amount of precipitation of γ’ phase, and the γ’ phase has the feature of coherent precipitation

4.3 High-temperature alloy materials 

 157

and the returned dissolution and reprecipitation phenomenon; it is exactly the reversibility of the precipitation hardening of nickel base alloy. It makes repeated aging possible many times. What is commonly used is, first, high-temperature and, later, low-temperature aging, which can obtain the best coordination of strength and plasticity. (4) Special curved grain boundary heat treatment The grain boundary obtained by general heat treatment is straight. If we can obtain curved grain boundary, resistance to the extension of grain boundary crack can be produced, thereby not only increasing the creep and endurance properties of alloy but also increasing plasticity at the same time. If we want to obtain saw tooth curved grain boundary, special curved grain boundary heat treatment must be carried out. Its principle is, first, precipitate second phase grains along the grain boundary; then, let second phase grains nail part of the grain boundary when grain boundary movement at high temperature occurs, to make it immovable. However, at the grain boundary without precipitated matter between the second phase grains, grain boundary movement occurs, which creates curved grain boundary. The heat treatments of curved grain boundary mainly have the following three types, as shown in Fig. 4.25 [13].

Fig. 4.25: Schematic diagram of three types of typical curved grain boundary heat treatments (aging treatment not indicated all). (a) Cooling control treatment, (b) returned dissolution treatment, and (c) isothermic treatment.

Cooling control treatment: heat the alloy to solid solution temperature; after heat preservation, make the grains grow to an ideal degree, and the precipitated phase is also fully dissolved; then cool it down to a certain temperature at the specified velocity, and carry out the aging treatment. Control the cooling velocity, and coarse precipitating phase, plus the effect of orientation and interface energy, to produce a curved grain boundary. After this kind of treatment, although the plasticity of the alloy has improved, the properties of creep, endurance, fatigue, etc., decrease, and although the grain boundary is strengthened by being curved, the grain inside is weakened by “overaging.” Therefore, it is difficult to obtain better comprehensive properties.

158 

 4 Oxidation and hot corrosion of metals and alloys

Returned dissolution treatment: first, dissolve most of the thick second phase inside the matrix, then carry out air cooling or isothermic aging at a lower temperature, then again, precipitate the second phase as dispersion. The returned dissolution treatment refines the grain and also maintains the curved grain in a certain degree, thereby improving the coordination of grain inside strength with grain boundary strength, enhancing its strength and plasticity. Isothermic treatment: heat the alloy to solid solution temperature and preserve it for a certain period of time, and later carry out air cooling to a certain temperature, conduct heat preservation again, and finally cool it down to room temperature. During air cooling after solid solution, γ’ phase precipitates fine grain to be distributed in dispersion. The curved grain is caused by precipitated second phase (mostly carbide) at the grain boundary in isothermic heat preservation. The temperature of the formed curved grain by isothermic treatment is lower, and the structure is refined. Thus, the strength and plasticity can have a better coordination at a higher level. 4.3.2.3 Oxidation-resisting nickel base alloy The main property requirements of this type of nickel base alloy are high temperature oxidation resistance and high temperature gas corrosion resistance. They are applied for manufacturing part alloy bearing with no outside force or relatively small outside force, such as parts of electric heating element, heating furnace component, engine combustion chamber, etc. As viewed from the alloy composition, it mainly contains enough amount of the alloying element of chromium resisting high-temperature corrosion; some also contain a small amount of strengthening elements, such as tungsten, molybdenum, titanium, aluminum, etc. This type of alloys has better processing forming property and can be processed to form a line, belt, plate, tube material; its representative alloys are Cr20Ni80 and Cr20Ni78TiAl. Table 4.7 lists the composition and main usage of this type of alloys, most of which are line, plate, tube, and belt material and can be welded. 4.3.2.4 Heat strengthening nickel base high-temperature alloy This type of alloy mainly is used in parts working at a high temperature, such as blade, turbine plate, etc. They need to bear high-temperature gas corrosion and additional load; the alloy is required to have both high temperature corrosion resistance and high temperature comprehensive mechanical properties. The composition of this type of nickel base high-temperature alloy is relatively complicated; besides the adding of element chromium resisting high-temperature oxidation and high-temperature corrosion, it also contains molybdenum, tungsten, niobium, etc. (difficult-to-melt metals); titanium, aluminum, etc. (elements forming intermetallic compound precipitated phase); and boron, cerium, magnesium, etc. (grain boundary strengthening elements).