(Materials Science Monographs 59) Janusz Flis (Eds.)-Corrosion of Metals and Hydrogen-Related Phenomena_ Selected Topics-Elsevier Science Ltd (1991)

Contributing Authors

J. FLIS

J. FORYST

A. F R A C K I E W I C Z

E. HAJEWSKA

A. J A N K O

E. t U N A R S K A - B O R O W I E C K A

W . P A L C Z E W S K A

J. P IELASZEK

W . RACZYKlSKI

A. S A D K O W S K I

j . ε ξϋζ ΐΜ^

L SUSKI

Z . S Z K L A R S K A - S M I A t O W S K A

J. W I L D

T. Z A K R O C Z Y M S K I

Z . Z E M B U R A

A. ZIELlKlSKI

P. Z O L T O W S K I

Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland

Institute of Fundamental Problems of Metallurgy, Polish Academy of Sciences, Cracow, Poland

Institute of Physical Chemistry, Polish Academy

of Sciences, Warsaw, Poland

Institute of Nuclear Energy, Warsaw, Poland

Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland





Institute of Non-Ferrous Metallurgy, Technical University of Mining and Metallurgy, Cracow, Poland

Institute of Physical Chemistry, Polish Academy of Sciences, Cracow, Poland

Department of Metallurgical Engineering, Ohio

State University, Columbus, Ohio

Institute for Building Technics, Warsaw, Poland

Institute of Physical Chemistry, Polish Academy of Sciences, Warsaw, Poland Institute of Non-Ferrous Metallurgy, Technical University of Mining and Metallurgy, Cracow, Poland

Marine Higher Academy, Gdynia


MATERIALS SCIENCE MONOGRAPHS, 59

CORROSION OF METALS AND HYDROGEN-RELATED PHENOMENA SELECTED TOPICS

Edited by

JANUSZ FLIS

Institute of Physical Chemistry,

Polish Academy of Sciences,

Warsaw, Poland

ELSEVIER Amsterdam - Oxford - New York - Tokyo

PWN — POLISH SCIENTIFIC PUBLISHERS Warszawa 1991

Revised and enlarged translation from the Polish original Wodorowe i korozyjne niszczenie metali

published in 1979 by Panstwowe Wydawnictwo Naukowe, Warszawa

Translated by Andrzej Skup

Distribution of this book is being handled by the following publishers: for the USA and Canada ELSEVIER SCIENCE PUBLISHING CO., INC. 655 Avenue of the Americas, New York, NY 10010

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for all remaining areas ELSEVIER SCIENCE PUBLISHERS Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands

Library of Congress C atalogi η g-in-Publi cation Data

Wodorowe i korozyjne niszczenie metali. English. Corrosion of metals and hydrogen-related phenomena: selected topics /edited by Janusz Flis; [translated by Andrzej Skup].

p. cm. — (Materials science monographs; 59) "Revised and enlarged translation from the Polish original

Wodorowe i korozyjne niszczenie metali" — T.p. verso. Includes bibliographical references. ISBN 0-444-98793-2 1. Corrosion and anti-corrosives. 2. Metals—Hydrogen

embrittlement. I. Flis, Janusz. II. Title. III. Series. TA467. W6213 1990 620. Γ623—dc20 89-23684

CIP ISBN 0-444-98793-2 (Vol. 59) ISBN 0-444-41685-4 (Series)

Copyright (Q) by ΡWN—Polish Scientific Publishers—Warszawa 1991

.All rights reserved

N o part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise, without the prior written permission of the copyright owner.

Printed in Poland by D.R.P.

Dedicated to Professor Michal Smialowski

by his coworkers

Preface

The cost of corrosion in industrialized countries amounts annually up to about 4% of the Gross National Product. According to the estimates of the National Bureau of Standards, in the USA about 40% of the annual production of metals is used to repair or replace materials damaged by corrosion. Corrosion causes waste of the natural material and energy resources, it creates serious materials problems for many technologies and adversely affects almost every area of en-gineering. Metals are used in various aggressive environments and this results in an extremely wide diversity of corrosion problems.

This book presents a collection of review articles on selected topics of me-tallic corrosion and on some aspects of interaction of hydrogen with metals.

Part I, entitled "Corrosion of Metals", covers the problems of pitting corro-sion and stress corrosion cracking in aqueous solutions, high temperature corro-sion in liquid metals and molten salts, displacement (cementation) of metals, and the role of diffusion in the kinetics of corrosion processes.

Part II, entitled "Interaction of Hydrogen with Metals and Their Resultant Deterioration", presents the problems of the entry of hydrogen into a metal from the gas phase and from aqueous solutions, problems of hydrogen behav-iour within the metal phase, and of the destructive action of hydrogen. In parti-cular, hydrogen decreases the mechanical strength and ductility of metals, causing the effect known as hydrogen embrittlement. Apart from the destructive action, the desired effects of hydrogen are also presented, such as the possibility of the modification of catalysts for chemical reactions and the use of hydrogen in fuel cells. In connection with the latter topic, the electrochemical oxidation reactions of hydrogen are described.

The authors have been the close collaborators of Professor Michat Smia-lowski who pioneered modern corrosion research in Poland in the early thirties. Professor Smialowski initiated in Poland the research in the field of hydrogen embrittlement, corrosion inhibition, stress corrosion cracking and other forms of localized corrosion. His work on the problems of hydrogen in metals and

2 PREFACE

hydrogen embrittlement has been widely recognized as an essential contribu-tion to the knowledge in this area of the environmental deterioration.

The research promoted by Professor £mialowski has been carried on and developed by many of his students and coworkers. Some areas of this research are reflected in the content of this book. By this work the authors would like to pay a tribute to Professor Smialowski for his great scientific achievements and his creative teaching.

The scope of the presented topics is fairly broad, and it is therefore hoped that the book will be of interest to corrosion scientists of various specialities.

/. Flis

Major Symbols and Units Used in the Book

Symbol Name SI Unit Relations between units

1 2 3 4

A Activation energy J/mol 1 J/mol = 0.2388 cal/mol 1 cal/mol = 4.1868 J/mol

a, B, b9 η Coefficients — a Chemical activity mol/1 a Crystal lattice parameter nm 1 nm = 10 Â b Tafel slope, dE/dlogi — — C Capacitance F — c Concentration mol/1 D Diffusivity m

2/s —

d Diameter m . Ε Electrode potential V

SCE Saturated Calomel Electrode —_

e Electron — F Faraday constant c 1 F = 96520 C/g-equiv. f Frequency s"

1

G Gibbs free energy J/mol 1 J/mol = 0.2388 cal/mol g Surface energy inhomo-

geneity coefficient — H Enthalpy J/mol — H Hardness Pa 1 MPa = 0.102 kG/mm

2

h Height, depth m I Current intensity A i Current density A/m

2 1 A/m

2 = 0.1 mA/cm

2

j Flux of substance kg/s j Imaginary number j / — 1 —

MAJOR SYMBOLS AND UNITS USED IN THE BOOK

1 2 3 4

κ Equilibrium constant .— —

k Reaction rate constant various — I Length, thickness m — M Molarity mol/1 — m Mass kg — m Speed of rotation of the

disk electrode r.p.s. Ν Number — —

ΝΗΕ Normal Hydrogen Elec-trode _

η Number of electrons in an electrode reaction — —

RHE Reversible Hydrogen Electrode (in a given electrolyte)

Ρ Force, load Ν 1 Ν = 0.102 kG Ρ Hydrogen permeation

rate various _ Ρ Pressure Pa 1 Pa = 1.02xl0"

5 kG/cm

2

Ρ Reaction order — — Q Quantity of heat J 1 J = 0.2388 cal

q Electronic charge C 1 C = 1 A · s

R Molar gas constant J/mol · Κ R = 8.31696 J/mol · Κ = 1.986 cal/mol-Κ

R Electric resistance Ω — Ultimate tensile strength, Pa 1 MPa = 0.102 kG/mm

2

r Radius m —-S Surface area m

2

— S Solubility product various — S Solubility various — S Dimensional factor — — s Potential sweep rate V/s — Τ Absolute temperature Κ 0°C = 273.15 Κ t Time s — V Volume m

3 —

VNHEJ VSCE Volts vs. NHE or SCE — —

2

MAJOR SYMBOLS A N D UNITS USED IN THE BOOK 5

ν Process rate various — W Energy J 1 J = 1 Ν · m = 10

7 erg

= 0.102 k G - m Ζ Cold work % — Ζ Impedance Ω — ζ Number of electrons in a

charge transfer reaction — — α Transfer coefficient — — β Symmetry factor — — a Thickness of diffusion layer m — ε Free charge of an electron C — ε Strain — — η Overpotential V — λ Thickness of deposit m — ν Kinematic viscosity m

2/s 1 m

2/s = 10

4 St

— Dynamic viscosity Ν · s/m2 1 Ν · s/m

2 = 10 Ρ

ρ Density kg/m3 1 kg/m

3 = 10"

3 g/cm

3

a Stress Pa 1 MPa = 0.102 kG/mm2

χ Ionic fraction — ω Angular velocity of the

disk electrode rad/s 1 rad/s = 0.159 r.p.s. — Specific heat J/kg · deg 1 J/(kg · deg) = 2.388 χ 10"

4

cal/(g · deg) — Volume specific heat J/(m

3 · deg) 1 J/(m

3 · deg) = 2.388 χ 10"

7

cal/(cm3 · deg)

— Thermal conductivity J/(m3 -s -deg) 1 J/(m

3 · s · deg) =2.388 χ ΙΟ""

7

cal/(m3 · s · deg)

1. Nucleation and Development of Pitting Corrosion in Iron and Steel

1.1 Introduction

One of the most frequently encountered varieties of localized corrosion is pitting. In common with other kinds of localized corrosion damage, such as crevice corrosion or stress corrosion cracking, pitting occurs primarily in chloride-con-taining solutions. Metals and alloys covered with passivating oxide films are susceptible to this mode of corrosion.

The effective prevention of pitting requires not only a knowledge of the fac-tors and conditions which cause its occurrence in a given material but also a thorough understanding of the mechanism of nucleation and development of pitting.

The considerable experimental data accumulated and published in the past thirty years have provided much valuable information on the pitting of various metals and alloys. In addition, methods have been developed for the determina-tion of the susceptibility of metals and alloys to pitting. A comprehensive basic research effort has been focussed on the elucidation of the initial stage of the process, i.e. the nucleation of pits, as well as their subsequent growth.

However, despite a huge research effort our knowledge in this field remains incomplete. As regards the initial stage of corrosion process in particular, a com-plete understanding of the reasons for pit initiation at specific spots on a metal surface continues to be lacking, and individual researchers differ significantly among themselves in their views.

Until recently a large number of specialists in the field were of the opinion that pits develop randomly, that is, there are no preferred nucleation sites on the metal surface. However, in the light of results of the latest studies carried out with the aid of more precise methods, this theory has been discarded. The re-sults demonstrated the surface inhomogeneities to be particularly liable to pitting.

10 NUCLEATION A N D DEVELOPMENT OF CORROSION IN IRON AND STEEL

This is clearly of prime significance for both the manufacturer and the user, since they need to know which surface inhomogeneities are particularly undesirable and harmful from the standpoint of corrosion sensitivity.

1.2 Sites of pit nucleation

The pit nucleation sites are undoubtedly related to the quality of the passivat-ing film. It has been known that those alloying elements which favourably affect the passivity of steel, such as chromium or molybdenum, also increase its resistance to pitting corrosion. Additionally, it is generally agreed that passive films of superior protective properties are produced on physically and chemically homo-geneous surfaces; it has often been demonstrated that the resistance of a given metal to pitting increases as its surface homogeneity becomes greater. If a metal surface layer has some physical defects or chemical inhomogeneities, then the pits nucleate at these weak sites on the surface.

As should be expected the passive properties of the film are more strongly affected by inhomogeneity of chemical composition of the surface than by physi-cal defects. It has long been known that sulphide inclusions are preferable sites of intense corrosion of carbon steels [1, 2] and stainless steels [3, 4].

Forchhammer and Engell [5] established that the number of corrosion-sen-sitive sites depends on the sulphide distribution on the surface and this in turn is dependent on the thermal treatment employed. If the treatment involves coagulation of the sulphides, fewer pits are formed and they are relatively large in size.

As shown by microscopic studies and electron microprobe analysis the major sources of pit nucleation in carbon steels are precipitates of complex manga-nese and iron sulphides (Mn, Fe) S [6]. In steels containing manganese in suffi-ciently high concentrations, manganese-rich sulphides are formed, whereas in Mn-deficient steels iron-rich sulphides are normally produced. Occasionally, however, Mn-rich sulphides are observed even in materials containing manga-nese in relatively low quantity. It is these complex manganese-rich sulphides that were found to be the principal pit nuclei [7].

The behaviour of the mixed iron and manganese sulphides depends on the Mn/Fe concentration ratio. The Mn-rich sulphides are readily attacked by an acid corrosive medium (Fig. 1.1), whereas the Fe-rich sulphides are not (Fig. 1.2). In the corrosion process, micro-crevices are produced around the Mn-rich sul-phide inclusions and both the inclusions and the metal matrix are attacked.

SITES OF PIT NUCLEATION 11

Fig. 1.1. Μη-rich sulphide inclusions within a pit formed in steel containing 0.22% S and 0.16% Mn. Corrosion in 0.11 M H 3B 0 3 + 0.023 M N a 2B 40 7 + 0.1 M KC1 (pH 8.4) at 175 mV SCE after [7]). Micrograph obtained with secondary electrons using a scanning electron microscope; magnification 5000 χ .

Fig. 1.2. The initial stage of pit formation on an Fe-rich sulphide inclusion in steel containing 0.21% S and 0.16% Mn. Corrosion in 0.11 Μ H 3B 0 3 + 0.023 M N a 2B 40 7 + 0.1 M KC1 (pH 8.4) at —175 mVsce. Micrograph obtained with secondary electrons of a scanning elec-tron microscope; magnification 500 χ after [7]).

Conversely, for Fe-rich sulphide inclusions, the metal around an inclusion under-goes corrosion, whereas the inclusion itself is resistant.

Gainer and Wallwork [8], using a scanning microprobe potentiometer, found the potential distribution on the surface of mild steel, in an NaCl solution, to

12 NUCLEATION A N D DEVELOPMENT OF CORROSION IN IRON A N D STEEL

be highly localized and to be more anodic at the sulphide inclusions than in the surrounding matrix.

Comparative studies on mild steels and steels containing 1% Cr, Al, Ni, or Cu led these authors [9] to believe that Cr and Al tend to increase both the pitting and general corrosion, whereas Cu and Ni are beneficial. The authors suggest that in the presence of Cr and Al inclusions of complex sulphides of the type (Mn, Cr)S + C r 2S 3 and MnS + A1203 are formed, which results in an ap-preciable potential difference between the inclusions and the matrix. The effec-tiveness of metallurgical heterogeneities as pitting promoters increases in the order [8]:

oxides < ξ f**"1 b o u n d a r

y < ( Mn Cu)S < MnS < A1203 < MnS + t n p l e p 0 i n tS

+ Cr2S3(Mn,Cr)S

Electron microprobe analyses have also shown that on the surface of stainless steels in contact with an aggressive medium, pits nucleate in the vicinity of mixed iron and manganese sulphides, which are present as separate precipi-tates or as envelopes around the oxides [10, 11]. Pits are not formed, on the other hand, on precipitates composed of oxides alone.

Not all sulphide inclusions are equally capable of pit nucleation. Depending on their chemical composition and shape, some of the inclusions may be prefer-able pit nuclei to others. In the presence of both kinds of inclusions, the active and the less active ones, pits already present may protect from further attack the adjacent parts of steel surface which contain the less active inclusions [12].

Furthermore, not all sulphides are the sites of pit nucleation. Cihal et al. [13] in particular, from their study on pit nucleation in Crl8Ni9Ti steel by means of the electron microprobe, found that titanium sulphides exhibit a higher sta-bility than titanium carbide or manganese sulphide and fail to induce pitting. This fact may be explained by a low rate of dissolution of titanium sulphide in acids.

A number of authors have noted that the shape of inclusions also influences pitting [14-17]. Spherical sulphide inclusions are less corrosion sensitive than elongated, plastically deformed ones. It has also been found that susceptibility to pitting depends upon the direction of grinding. Transverse grinding, which cuts through inclusions, increases the effective interfacial area and leads to a more negative pitting potential, compared with longitudinal grinding [16, 17].

If manganese sulphide or iron sulphide inclusions are absent in steel, corros-ive pit nuclei may alternatively be induced by carbide precipitates, grain bounda-ries, and other heterogeneities.

SITES OF PIT NUCLEATION 13

The electron microprobe detected the presence of Al and Si [18] on the edges of some pits on the surface of very high purity iron. There was no certainty however as to whether these elements had come from the metal or from impur-ities introduced during the mechanical polishing of specimens.

In stainless steel, Brauns and Schwenk [19], as well as Streicher [4], noted the formation of pits at the grain boundaries where carbide precipitates oc-curred.

In 18Cr5Ni steels subjected to solution heat treatment, which contained from 5.7 to 15% Mn, the sites most prone to pit nucleation proved to be the bound-aries between the γ phase and carbide precipitates (probably of the Me2aC6 type). In hot-forged steels, pits formed at the interface between the γ and θ phases; both these phases were subject to corrosion attack [20].

The most preferred sites of pit nucleation in a 18Cr5Nil0Mn steel containing from 0.07 to 0.35% nitrogen were the interfaces between the austenitic matrix and inclusions (most probably carbides [21].

Pit nucleation often occurs at grain boundaries, which are commonly the sites of chemical composition heterogeneity. Specifically, Tomashov et al. [22] report that in an 18Crl4Ni steel with additions of Mo, V, and Si, pits are formed at the grain boundaries, probably on account of the presence of mixed carbides, which are responsible for a depletion of the grain boundaries in chromium and other alloying components.

In an iron alloy containing 13% chromium, pit nucleation was again observed to occur at the grain boundaries. In the initial corrosion stages pits grew on different grains at different rates dependent on the crystallographic orientation of the grains relative to the specimen surface [23],

In pure iron, Vetter and Strehblow [24] observed pit inititation on surface scratches.

In summary, the preferred sites for pit nucleation on the surface of iron and its alloys may be not only the sulphide inclusions with particular emphasis on the manganese-rich mixed compounds of the (Mn, Fe)S type, but also:

— interfaces between the metal matrix and either non-metallic inclusions or precipitates that are capable of binding the alloy components which are re-sponsible for its passive condition; certain carbides seem to be particularly active in this respect;

— grain boundaries along which impurities segregate; — surface defects. It is significant that manganese sulphide has a relatively large thermal expan-

sion coefficient. Hence, during steel cooling, the sulphide inclusions shrink


more than the ferrite or martensite and, as a result, microcracks develop around MnS which, if brought into contact with an aggressive medium, may promote pit initiation. In addition, manganese sulphide inclusions display considerable reactivity. Hydrogen sulphide, which increases the rate of both the anodic and cathodic process, is formed in MnS reactions [25-27]. Furthermore, at anodic potentials, manganese sulphide may undergo oxidative solution to yield SOU" and M n

2+ + H

+ [28]. Thus, both the sulphide dissolution and oxidation prod-

ucts may speed up the corrosion of iron.

1.3 The effect of thickness and structure of the passivating oxide films on pitting

It is well established that the more contaminated is a metal, the less perfect is the passive film produced on its surface, and hence the worse the protective properties of the film. Unfortunately, there are as yet no data to permit a quanti-tative estimation of how the concentration and composition of impurities affect the quality of a passive film; only sparse and indirect information is available.

10

273 373 473 523 573 Temperature, Κ

673

Fig. 1.3. Thickness of oxide film vs. oxidation temperature for various nickel types (after [29D: •—pure nickel, O—cold-rolled nickel, Δ —annealed in vacuum for 1 h at 973 Κ (the latter two were technical grade nickel specimens).

THE EFFECT OF THICKNESS 15

It has however been established that the oxidation rate of metals and their alloys in contact with air at high temperatures is greater for contaminated materials, and that the oxide films produced on them are thicker than on higher-purity materials.

In the case of nickel, for example, Fig. 1.3 shows the dependence of the oxide film thickness on the oxidation temperature (373 to 573 K) for three sets of nickel specimens oxidized in the air: pure nickel (99.98%) and technical grade nickel (95.50%), either cold-rolled or annealed in vacuum for 1 h a 973 K. As can be seen, at the same temperature the thinnest oxide films are formed on pure nickel (squares), whereas the thickest films grow on cold-rolled, lower-purity nickel (circles) [29].

Analogous data on the effects of anodic oxidation are lacking, and studies aimed at correlating the thickness and the structure of oxide films with suscep-tibility to pitting are scarce.

Bianchi et al [30] investigated the relation between pitting corrosion of 19Crl0Ni steel and oxidation temperature in air over the range 423 to 673 K. They consider that the number of pits produced per unit surface is independent of the film thickness but depends on the semiconducting properties of the oxide. The conductivity type was established on the basis of the polarization curves deter-mined in a redox solution of 0.5AfNa2SO4+0.1AfK3Fe(CN)6+0.1MK4Fe(CN)6. Steel oxidized at 573 Κ proved to be the most sensitive to pitting which, in the author's opinion [30], was due to «-type conductivity of the oxide film formed under these conditions.

Ellipsometric and electrochemical studies were carried out on oxide films grown on carbon steel oxidized in the air at various temperatures [31, 32]. The resistance to local breakthrough of the film produced at 473 Κ (magnetite), when in contact with an aggressive medium, increased as the film thickness increased (up to 20 nm). The films produced at temperatures from 523 to 623 Κ were less resistant; they consisted of two layers, magnetite and hematite. Hence, the susceptibility to pitting increased with the film thickness in this instance.

The kinetics of film growth, for a temperature range of 373 to 473 K, follows a logarithmic relationship, whereas in a temperature range from 523 to 623 Κ a parabolic dependence of film thickness vs. time is observed. A difference was also found in the activation energies for the formation of magnetite and hema-tite, these being 33.5 and 100.5 kJ/mol, respectively.

The sensitivity of the passive film to breakthrough was estimated from the number of pits formed per unit surface area of specimens anodically polarized in a buffer solution of pH 8.4 with addition of O.lMNaCl, at 293 Κ [32]. As


600

x 373 Κ

Δ 423 Κ

ο 473 Κ

10 15 20 Film thickness, nm

25 30

Fig. 1.4. Relation between the number of pits per cm

2 and thickness of the magnetite layer

formed on steel at 373, 423 and 473 K. The number of pits was determined after polarization of the steel specimen at Ε = 145 n ï V W for 20 min in 0.11 Μ H 3B 0 3 + 0.023 M N a 2B 40 7 at pH 8.4 (after [32]).

2000

Ι ι i l i ι I d 1 1—„—ι 1 0 10 20 30 40 50 60 70 80

Thickness, η m


2 and thickness of the oxide film formed

on steel at 523, 573 and 623 K. Conditions of the experiment as in Fig. 1.1 (after [32]). Roman numerals denote oxide films differing in optical properties.

seen from Fig. 1.4, the density of pit occurrence (only pits over 7 μιη were con-sidered) decreases with the thickness of the magnetite film formed. As soon as hematite starts forming on the magnetite layer at 523 K, the number of pits rapidly increases (Fig. 1.5). The number rises with oxidation time and with the thickness of the layer. In the temperature range from 573 to 673 K, the pit den-


sity decreases with temperature and time and pit formation becomes replaced by corrosion processes. If, however, polarization is off after 1-2 min, then multitudes of pits form (a few thousands per cm

2) and gradually coalesce. For

this reason, the experimental points in Fig. 1.5 corresponding to the number of pits found in ranges II to IV are connected by a broken line.

Results of the ellipsometric studies show that the oxide films of thickness greater than a certain critical value which depends on the growing conditions and type of oxide, etc., become increasingly porous. This results in a gradual rise in the density of pit distribution so that ultimately general corrosion super-venes.

On the basis of electrochemical measurements in a ferri-/ferrocyanide redox system it was established [31] that the films formed on carbon steel exhibit «-type conductivity irrespective of their thickness. This proves that the observed differ-ences in sensitivity of these films to breakthrough are caused by differences in thickness, rather than by changed electronic properties.

In the case of nickel oxidized in the air at temperatures from 373 to 573 K, X-ray analysis showed the presence of just one oxide, NiO. Results of the ellip-sometric studies, on the other hand, lead to the conclusion that in temperature range from 373 to 523 Κ a single oxide film is formed on the metal, and at higher temperatures this film is overgrown by another film which, as shown by the ellipsometric evidence, differs in the value of the complex index of refraction [29].

The film growth kinetics study [29] demonstrated the film to grow via a log-arithmic relationship up to 523 K, whereas at 573 K, from a thickness of 3 nm the relationship becomes parabolic at higher thickness values.

A clear correlation was found to exist between the optical properties and the film growth kinetics on the one hand, and sensitivity to pitting of oxidized nickel on the other [33]. The least sensitive to pitting were 99.98% nickel and 99.5% annealed nickel specimens oxidized at 373 to 473 K, and also 99.5% cold rolled nickel specimens oxidized at 373 K. The films exhibited a relatively high complex index of refraction, and grew logarithmically. Microscopic investi-gation showed that the films were more compact than those formed at higher temperatures. With increasing thickness of these films, the susceptibility to pitting decreased (Fig. 1.6).

A change in the nickel oxidation kinetics, observed at above 523 K, consisting in the passage of a logarithmic relation of the oxide film thickness vs. time into a parabolic one, resulted in a film composed of two layers. A second and more porous layer grew on top of the primary compact layer, as demonstrated by a lower value of the complex index of refraction.


850 m V

11 ι l ι I ι I ι I ι I 0 2 4 6 8 10

Thickness, nm


2 and thickness of the film obtained at

various temperatures for three Ni specimens. The number of pits was determined after a 5-min polarization in a solution of 0.05 M N a 2S 0 4 + 0.013 M NaCl (pH 5.8) at 750 and 850 HIVNHE (after [33]).

The oxidation of nickel in the air at temperatures of 573 or 673 Κ enhances susceptibility to pitting, in spite of the fact that, under these conditions, far thicker oxide films are produced than at 373 to 473 K. This evidently results from the higher porosity of thick layers, in the presence of which multiple pits are formed during the corrosion process.

Under the same experimental conditions, considerably more pits were invari-ably formed on contaminated nickel (99.5%) than on pure nickel (99.98%).

Interestingly, on cold rolled 99.5% nickel specimens, oxidized either at lower or at higher temperatures, the corrosion pits appear immediately on the appli-cation of a potential of 750 m V N HE to a specimen immersed in a solution of 0.05MNa2SO4 with 0.013MNaCl. Conversely, on 99.98% nickel, as well as on the annealed 99.5% nickel specimens, following oxidation at 373 and 473 K, an induction period is required to initiate pitting under the same conditions.


Thicker oxide films result in longer induction periods. It seems that the induc-tion period occurs for film growing logarithmically; in contrast, no induction time is necessary to break through the films showing a parabolic growth pattern (specimens oxidized at 573 K). The results of the relevant study are reported in Table 1.1.

TABLE 1.1

THE INDUCTION TIME FOR PITTING OF NICKEL IN 0.05 M N a 2S 0 4 -f 0.013 M NaCl AT A POTENTIAL OF 750 m V N HE

Oxidation temperature,

Κ

Thickness of films, nm

Induction time, min Oxidation

temperature, Κ 99.50%

annealed Ni 99.98%

Ni 99.50%

annealed Ni 99.98%

Ni

373 1.0±0.1 6 ±0 .2 1 3

473 1.6±0.2 9 ±0 .2 1.7 3.5

573 8.9 ±0.3 6.7 ±0 .4 — —

From the experimental data, both for mild steel and nickel, a clear effect is seen of the oxidation conditions in air on the pitting behaviour of these two metals, namely:

1. In both cases, with increasing films thickness, the number of pits formed per unit area of specimen initially diminishes, and then rises dramatically after a certain critical thickness is reached;

2. The sensitivity of the film to breakthrough depends on the kinetics of its growth; in the case of a logarithmic dependence of film thickness on oxida-tion time, less pits are formed than for parabolic dependence.

It seems that in the case of steel, the variations in susceptibility of the film to breakthrough are independent of the presence of magnetite or hematite in the film but, as for nickel, such variations depend on the film porosity.

Only one paper was found in the available literature, dealing with the corre-lation of the susceptibility to pitting with thickness of the anodic passive film on iron. McBee and Kruger [34] studied films of thickness ranging from 2 to 6 nm, produced in a boric acid buffer solution. They found that, in the presence of chloride ions, the thicker films took longer for a breakthrough to occur.

20 NUCLEATION AND DEVELOPMENT OF CORROSION IN IRON A N D STEEL

1.4 Role of chloride ions in the pit nucleation process

The numerous studies of oxide films on metal surfaces concerned with the vari-ous aspects—identification of potential pit nuclei sites, structure, growth kin-etics, sensitivity to breakthrough, etc.—have resolved a number of difficulties and have led to the rejection of some erroneous ideas. However, a complete elucidation of the mechanism of pit nucleation has not yet been achieved. In particular, the role of chloride ions in pit nucleation is somewhat obscure, al-though valuable information has come from more recent investigations of the composition of surface films formed on metals and alloys at the moment immedi-ately precedent to pit initiation, as well as during pitting.

A number of authors have established the adsorption of chloride ions on metals. The adsorption is not uniform but takes place at certain active sites of the metal surface [35-38]. Bevilacqua and Brown [35] showed an inhomogeneous adsorption of

3 6C1 on type 347 stainless steel. Bergen [36] found that

3 8C1 ions

migrate towards areas of maximum stress. On rough ground steel specimens, chloride ions were found to agglomerate primarily in grooves [37]. Investiga-tions made with an electron microprobe at potentials somewhat lower (i.e. more negative) than the potential at which pit nucleation takes place in a given electro-lyte (this potential will be referred to as the critical pit nucleation potential, and will be denoted as Epn) demonstrated localized accumulation of chloride ions at specific areas of the metal surface [39, 40]. For 18Cr9Ni steel such accu-mulations were found at the boundaries between the alloy matrix and non-me-tallic inclusions [40], whereas for pure iron containing only 27 ppm of impuri-ties, some enrichment in Cl~ was detected in certain unspecified sites of the metal surface [39].

A typical example of impurity distribution at a site of possible pit nucleation on surface of mild steel is shown in Fig. 1.7. The results illustrated were obtained by scanning electron microprobe along a straight line cutting through a sul-phide inclusion on a carbon steel specimen polarized to a potential lower than ^pn [40]. The characteristic radiation curve for Cl(â) is shifted in relation to the analogous curve for S which suggests that accumulation of CI" occurs at the interface between a sulphide inclusion and the matrix.

Approximate estimates of the thickness of the locally accumulated layer of the chloride salt, detectable by the electron microprobe, ranged from 6 nm, to an upper limit of the order of 80 nm [39]. The autoradiographic method furnished similar results [41].

There are some indications that accumulation of chloride ions at certain

ROLE OF CHLORIDE IONS IN PIT NUCLEATION PROCESS 21

Ci s (Κα) {Κα) 1 x 1 0 \ - 3 x 1 0

4c p m

5χ10'

2x10

4cpqL~

LlxIO^cpm

SxIO^cpm

-tex10

Jcpm

1x10

dcpm

10

Fig. 1.7. Distribution of the elements along a straight line running through sulphide inclusions after polarization of a carbon-steel specimen in a chloride-containing solution. The analysis was made using an electron microprobe (after [39]).

active sites of the metal surface is a forerunner of the pit nucleation process. Heusler and Fischer [42] claim, on the grounds of their electrochemical study with a ring electrode, that the formation of chloride islands on the metal surface precedes pitting. Other investigators have also noted a local acidification of steel surfaces to occur during the pit initiation process [38]. It is supposed that the chloride salt layer forms either as a result of an exchange between the chlor-ide and oxide ions, or as a consequence of migration of CI" and F e

2+ through

the passive film, which would result in accumulation of salts at weak spots of the film.

There is a considerable controversy in the literature over the incorporation of chloride ions into the oxide film.

Ambrose and Kruger [43] tried to demonstrate the occlusion of CI" in the passive film. They removed part of the passive film by cathodic reduction, and then reproduced it anodically in the presence of chloride ions. This gave rise


to a reduced time required to break through the film, which was thought to be evidence of chloride ions being incorporated in the film. Conversely, Abd-Rabbo et al. [44], using a sensitive method of secondary ion mass spectroscopy, failed to find Cl~ ions in the oxide film on aluminium specimens previously treated in chloride solutions.

The identification of individual elements in passive films of variable thickness, performed by Auger-electron spectroscopy, also seems to contradict the incor-poration of chloride ions into the passive film [45]. The experiments are carried out on pure iron specimens during polarization in a borate buffer with addition of 0.1 Âf NaCl, at potentials more negative than the pit nucleation potential. Under these conditions the passive films analysed by electron microprobe exhi-bited a localized accumulation of chloride on the metal [39]. On the other hand,

Li ^ E,eV

Fig. 1.8. Energy distribution of Auger electrons from an iron surface: (a) passivated in 0.11 M H3BO3 + 0.027 M N a 2B 40 7 + 0.1 M KC1 at Ε = 200 mVNHE for 60 min at 293 K; (b) the same specimen after a 105-min bombardment with A r

+ ions at a density of 5 χ 10

l fl

ions/m

2 (after [45]).

ROLE OF CHLORIDE IONS IN PIT NUCLEATION PROCESS 23

10

1

τ—ι ι ι ι ι ι

,19 Number of ions, ArVm

2

10

: ,20

Τ 1 — ι —I 1 I I I

The Ar

4" bombarding time, min

Fig. 1.9. Auger peak height vs. the time of bombarding with A r

+ ions and the quantity of Ar

+

used for bombardment of a passivated iron surface (after [45]).

the peaks characteristic of CI, as obtained by Auger-electron spectroscopy, were only slightly higher than the background, and they disappeared soon after bom-barding the film with argon ions. This shows that only minor quantities of chlor-ide ions were present in the oxide film, and that such ions were weakly bonded to the film. The results of the experiments are illustrated in Fig. 1.8. This shows the characteristic spectrum of Auger electrons for a specimen anodically polar-ized at 200 m VN HE for 60 min (Fig. 1.8a) and a spectrum of the same spot upon bombarding with Ar

+ ions for 105 min (Fig. 1.8b). Figure 1.9 shows the de-

pendence of the composition of the passive film, expressed in terms of the height of the characteristic peaks of individual elements, on the time of bombardment of the film with argon ions. As can be seen, the chloride ions must have been present only on the film surface, as their concentration diminishes rapidly under A r

+ bombardment, whereas boron that originated from the electrolyte must

have been occluded in the bulk oxide film, as its concentration did not decrease under these conditions.

These results indicate that pit nucleation is caused by the accumulation of chloride ions at certain defect sites in the passive film, rather than by their migra-tion through the film.


1.5 The critical potential for pit nucleation and the induction time for pitting

Phenomenologically, the critical potential for pit nucleation can be defined as the potential at which the first pit is noticed, and above which severe pitting pro-ceeds.

The term induction time is the time required for the first pit to appear. The value of the critical potential for pit nucleation depends not only on the

kind of metal, chloride concentration, electrolyte composition, temperature, etc., but also on the way this value is determined. The potentiostatic method is most frequently used; it entails plotting the anodic polarization curve and determin-ing the potential at which passive film is broken through. This effect manifests itself by a sudden strong rise in the current at a potential that is more negative than the oxygen evolution potential. A number of modifications of this pro-cedure has been detailed in the literature [46]. Another method for determination of the resistance of metals to pitting is the observation of the number of pits that are formed per unit surface area of metal. This technique, however, char-acterizes largely the degree of inhomogeneity of the passive film.

The length of induction time at a constant potential depends on chloride concentration [47-49] and at a constant Cl~ concentration it depends on the potential [19, 48, 50-52]. The higher the Cl~ concentration, and the higher the potential, the shorter is the induction time.

The induction time is usually determined at a constant potential by measuring the time required to produce the first pit, which manifests itself by a sudden rise in the anodic current.

It has repeatedly been found that induction time tind vs. chloride concentra-tion in solution may be expressed by the following formula:

log^nd = A + B\ogcCv

For mild carbon steel Engell and Stolica [47] found Β = — 1, and the same value was observed for austenitic chromium-nickel steel [53]. However, Β has later been found to depend on the conditions of the experiment. Ashworth et al. [54] for instance found various values of this coefficient at various potentials; for a potential of - 3 6 5 m V S C E, Β = -2 .2 , whereas for -165 mVS CE> Β = - 1 . 7 .

1.6 Pit nucleation theories

An early attempt to explain the mechanism of pitting was the theory of competi-tive ion adsorption, by Kolotyrkin [55, 56] and Uhlig [57, 58]. It assumes that pit formation is the result of a competitive adsorption of chloride ions and oxy-

PIT NUCLEATION THEORIES 25

gen. According to this concept, pits develop at the sites where the oxygen ad-sorbed on the metal surface is replaced by Cl~. Kolotyrkin believes the density of the anodic current of metal dissolution to be unequal over the whole metal surface on account of surface inhomogeneity : metal dissolves more rapidly where chlorides are more strongly adsorbed. According to the theory, the break-down potential is that minimal electrode potential at which aggressive anions are capable of replacing the passivating oxygen on the metal surface.

The adsorption theory however fails to account for all the facts observed, especially the occurrence of induction time, and it is inadequate in the light of the current views on the nature of passive films.

Penetration and migration theories have enjoyed a wider acceptance. Accord-ing to Rozenfeld [59, 60], at the sites of passive film where the metal-oxygen bonding is the weakest, oxygen may become replaced by chloride ions. Aggres-sive anions which interchange with oxygen penetrate into the passive film to result in a pit. In terms of this theory the induction time may be defined as the time required for an aggressive anion to penetrate the passive film and reach the metal surface. The assumption is made that the first stage of this penentration process involves adsorption of aggressive ions on the surface of the passive film. Consequently, the critical potential of pit nucleation would be the potential at which these ions adsorb on the passivated metal surface.

Hoar and co-authors [61] assume that the aggressive ions permeate the oxide film without any exchange. This would result in the formation of a "contamina-ted" path of a relative high ionic conduction in the film, which would facilitate the flow of a fairly high intensity current at certain spots, leading to localized corrosion.

Pryor et al. [62, 63] uphold the "contaminated film" concept, but they sug-gest that cationic vacancies arise, which reduce the ionic conduction and favour film breakdown.

McBee and Kruger [34] propose the following pattern for ionic exchange and migration :

CI" + D o - + 2 0 H - = Clfa t t i ce + 2ΠΟΗ- + 20Η"0ΐ

where the square • symbolizes vacancies in the crystal lattice of the film. Another interpretation of the pit nucleation process was put forward by

Vetter [64, 65]. He supposed that the oxide film on the metal surface remains in equilibrium with a thin chloride salt film, provided that films have the same chemi-cal potential and that an equilibrium exists in both films between M e

2+ cations

and electrons. Pitting is considered to start on the non-porous oxide film con-


taining salt nuclei. A tendency to form a salt film exists at potentials more posi-tive than the critical potential of pit nucleation, whereas at more negative po-tentials the more stable form is the oxide.

In their later study [24] Vetter and Strehblow assumed the pits to arise from mechanical breakdown of the passive film, accompanied by the formation of a non-porous salt or a chemisorptive film in the pit.

Yet another model for pit nucleation has been advanced by Hoar [66]. It is suggested that adsorption of aggressive anions on surface of the passive film lowers the surface tension at the solution interface and results in "peptization" of the film. If mutual repulsion of the adsorbed ions becomes strong enough, the passive film breaks and aggressive ions reach the uncovered metal surface.

Yahalom [67] also claims that physical adsorption of chloride ions on the passive film surface reduces the surface energy of the oxide, which may result in film breakdown and reduced adhesion of the film to the substrate. Penetra-tion of the electrolyte between the film and metal results in a closed-circuit cell, while the concentrated electrolyte from the operating local cell will leak out through the crack in the film, partly blocked by pieces of the film. Depending on the rate of formation of concentrated anolyte and its outflow through the film, either repassivation of the given spot, or pit nucleation and its consequent devel-opment may occur. The rate of anolyte formation is controlled by the local current density which, in turn, depends on the value of the potential applied, whereas the rate of anolyte outflow is dependent on the crack geometry and on the local hydrodynamic conditions.

Another viewpoint has been presented by Hoar and Jacob [68]. These in-vestigators carried out measurements of the induction time for pitting on 18Cr8Ni steel. The results obtained by them have led to the following conclu-sions:

(1) 1/înd is an approximate measure of the rate of film breakdown (tind is induction time);

(2) find depends on the potential and temperature; (3) the logfi nd vs. logger relation is a straight line with a slope Β ranging

from - 2 . 5 to - 4 . 5 ; (4) the activation energy for the induction process is 251 kJ/mol. The Β value corresponds, in the authors' opinion, to the amount of the Cl~

ions involved in the formation of a "transitional" complex which brings about a breakdown of the film. This is due to the fact that the complex becomes de-tached from the passive film and cations migrate in the electric field, through the thinned out region of the film, towards the outer surface, where they meet


chloride ions, form a complex with the latter, and pass into solution. The high value of the activation energy for the induction process indicates that the for-mation of the transitional complex mentioned occurs with considerable difficulty. The investigators [68] consider that the migration model is quite inappropriate for stainless steels, because of the high activation energy and the large amount of chloride ions required to form the complex.

The film thinning concept has also been put forward by Heusler and Fischer [42, 69].

An entirely different view is expressed by Richardson and Wood [70, 71]. They reject the concept of film breakthrough altogether, claiming that pores and defects are always present, thus permitting an immediate penetration of aggressive ions to the metal surface with consequent instantaneous pitting. Thus, the authors also entirely discount the involvement of any induction time process; the time lag observed in the initial pit development is ascribed by them to low values of the corrosion current that flows through the tiny, uncovered spots on the metal surface. According to this model, the critical potential of pit nuclea-tion is equivalent to the potential of the uncovered metal at the bottom of pores.

The mechano-chemical concept of pit nucleation has been developed by Sato [72]. He assumes that the high intensity of the electric field operating at the electrolyte-passive film-metal interfaces may cause the film to break, if the elec-trostrictive pressure set up in the film becomes higher than its compressive strength. The value of the electrostrictive pressure depends, among other fac-tors, on the surface tension of the film. As a result of adsorption, aggressive ions lower the surface tension of the oxide, thus increasing the electrostrictive pressure. The critical potential for pit nucleation for this model would be the potential above which the electrostrictive pressure exceeds the compressive strength of the film. Sato [72] claims that either a pit originates, or its growth is suppressed, depending on the rate with which the film is recovered at the site of the breakdown. This model, however, can hardly be invoked to explain the induction time for pitting and the specific role played by the aggressive ions in this process.

Videm [73] believes that the pit formation depends on a dynamic equilibrium between the processes of mechanical breakdown and repassivation. Competi-tiveness of these two processes determine whether or not the breakdown is per-manent and thus results in pitting. For pitting on aluminium it was demonstrated that within a certain narrow range of potentials a dynamic equilibrium condi-tion does occur between the film repassivation and breakdown. Below this range the passive film recovery predominates, whereas above this range film break-


down occurs. The pit nucleation potential should, in terms of this concept, be determined by the relative rates of the repassivation and breakdown processes.

An extended version of the dynamic breakdown-repassivation model has been advanced by Ambrose and Kruger [74]. They assumed that the aggressive ions, having penetrated to the metal surface, result in the formation of a local "pocket" at the metal-oxide interface. This pocket is filled with corrosion prod-ucts that have no protective properties (possibly y-FeOOH). The growth of these pockets sets up a localized concentration of stresses and results in the film breaking. The authors assume that the specific action of aggressive anions is to affect the rate of repassivation. In turn, the rate of the repassivation process, the mechanical properties of the film and the adsorption of aggressive ions, affect the critical passivation potential.

Hisamatsu et al. [75] believe that the critical potential of pit nucleation is the minimum potential sufficient to concentrate the chlorides inside the pit to a critical value that would render pit passivation impossible.

It is sometimes assumed that the pit nucleation potential is the potential above which a film or deposit is produced that exhibits no protective proper-ties [43].

Ambrose and Kruger [43] found that Epa for iron in a borate solution of pH 8.4 containing Cl~ ions coincided with the zero charge potential determined for iron by Lorenz and Fischer [76]. This potential also agreed with the equilib-rium potential of the formation of unprotective FeOOH that grew in successive stages [77]. It was assumed that film penetration by Cl~ opens the way for ionic conduction for F e

2 +, as a result of which, at potentials higher than Epn, a non-

protective y-FeOOH deposit is formed. Galvele [78, 80] and Galvele and de Wexler [79] suggest that the critical po-

tential for pit nucleation is the minimum potential at which a local acidity of the electrolyte in contact with metal can be maintained. Calculations performed by these authors showed that this acidity probably occurred at a very early stage of pit nucleation. This view, however, is contradictory to the opinion of Vetter and Strehblow [24] and Vetter et al. [91] who argue that the initiation of pitting is related neither to pH changes nor to variations in potential or ionic concentraion, as these variations are minute at the initial stage of pitting.

Lin et al. [81] assume that during the film growth caused by the diffusion of anions from the film-solution interface, diffusion of metal cations from metal-film to film-electrolyte interface occurs, producing metal vacancies at the metal-film interface. When the metal vacancy production rate is higher than the rate of vacancy submergence into the bulk, a void forms at the metal-film interface.


After reaching a critical size, the passive film collapses. In the presence of halide anions, this process occurs more easily, because these ions are incorporated into the outermost layer of the passive film by occupying anion vacancies in the film.

All these pitting theories share the view that the adsorption of Cl~ on the metal surface is the first stage of the process. The majority of the theories suggest that the weakest sites in the passive film are attacked, but individual investiga-tors differ in their views as to the nature of these sites.

The following are some of the .reasons which have been advanced for the occurrence of weak sites in the film:

(a) weak metal-oxygen bonding (Rozenfeld); (b) mechanical damage (Vetter, Videm); (c) breakthrough due to a local reduction of the surface tension at the film-

electrolyte interface as a result of Cl~ adsorption (Hoar, Yahalom); (d) breakthrough due to electrostrictive pressure (Sato); (e) pores and flaws in the film (Richardson, Wood); (f) sulphide inclusions and other chemical and physical inhomogeneities of

the metal, in the neighbourhood of which the film has a disturbed structure (Smialowska).

In the same way, there is a lack of consensus as to how chloride ions pass through the film to the metal surface. Advocates of the penetration theories suggest either migration or gradual penetration of Cl~", causing pit nucleation as soon as the chloride ions reach the bare metal surface. Mechanical theories, on the other hand, assume that electrolyte reaches the uncovered metal through the break developed in the passive film as a result of one of the reasons men-tioned above. Numerous researchers suppose that chloride anions cause dis-solution of the passive film.

The arrival of chloride ions at the metal surface, however, is not on its own sufficient to induce nucleation of pits that are capable of further sustained growth. This is indicated by the data related to pit nucleation and repassivation below the value of the critical potential for their nucleation, being evident from the oscillation of the current at potentials lower than the breakdown potential. Consequently, the breakdown potential can neither be expressed by the poten-tial at which aggressive anions are adsorbed on the passivated metal surface, nor by the potential at which the film starts breaking, but only by the potential cor-responding to a permanent occurrence of pits.

Currently the most acceptable views on the nature of the breakdown po-tential appear to be the following: that expressed by Videm [73] who claims the

30 NUCLEATION AND DEVELOPMENT OF CORROSION IN IRON AND STEEL

breakdown potential to be the potential above which the repassivation rate is lower than the dissolution rate: and the assumption made by Hisamatsu et al [75], who believe that Epn is the potential at which chloride ions become con-centrated to such an extent that pit repassivation becomes impossible.

On the basis of (a) the results of electron microprobe studies, which have demonstrated the presence of chloride salts on the metal surface in the form of localized islets, and (b) the results of profile analyses of the oxide films by Auger-electron spectroscopy, which showed Cl~ ions to be absent in the oxide film itself at a potential somewhat lower than the critical pit nucleation potential, the following mechanism can be advanced for this process as the most likely one.

At anodic potentials, the passive film undergoes a slow dissolution process. This process occurs faster at local defects on the metal surface. At these sites, cations are liberated and agglomerate, and chloride anions migrate to these sites from the solution. As a result, micro-islands of chloride salts are formed. Pit nucleation occurs at a potential somewhat higher than the potential at which critical agglomeration of chloride ions takes place on certain parts of metal surface, at "weak" sites of the passive film. As a result of this agglomeration, islets of iron chloride (or chromium and nickel chlorides in stainless steels) are formed in a layer of a thickness from several scores to a few hundreds of nano-metres. This salt is likely to hydrolyze, leading to an increased acidity and a re-sultant dissolution of the passive film, which had previously been impaired by a strong local emergence of F e

2+ ions. As a consequence, the aggressive solu-

tion reaches the metal surface. The permanent growth of pits depends on whether or not conditions exist for the repassivation of the exposed metal surface. If the anodic potential is higher than the critical pit nucleation potential, then passiva-tion of the nucleus will not be possible, since the supply of chloride ions will be sufficient to form an aggressive solution in the pit within which the wide range of anodic potentials will keep the surface in the active state and the precipi-tating corrosion products will not be insufficiently protective. As already empha-sized, sulphide inclusions appear to be particularly significant in the pit nuclea-tion process [82]. Even though it was found that accumulation of chloride ions and thus the formation of local salt islets, may occur not only at the MnS-matrix interfaces but also on the Mn-Al oxide inclusions, the oxide-metal boundaries are not attacked in the presence of sulphide inclusions. This suggests that the sulphides have a greater effect on the pit nucleation process than the oxides.

If it is assumed that the solution adjacent to the surface becomes locally acidic as a result of hydrolysis of the salt, the local attack on the manganese sulphide

KINETICS OF PIT GROWTH 31

occurring in the surface metal layer in the form of mixed sulphide-oxide inclu-sions can be essily explained. The sulphide dissolves in the acid, whereas the oxide is resistant to the acid; in consequence, a crevice develops around the partly dissolved sulphide. The exposed metal normally dissolves more rapidly than the sulphide. It is therefore probable that sulphide inclusions are not only dis-solving but that they are also leached or washed out into solution.

If a passive film and oxide inclusions are present on the stainless steel sur-face, then pit initiation is more difficult, because the oxides are resistant to dilute acids. Conversely, the occurrence of sulphide-oxide inclusions facilitates this process, which begins on the sulphides. Narrow and deep crevices formed at the sulphide-oxide boundary favour pit development, which may be a reason for the exceptionally high ability of mixed sulphide-oxide inclusions to initiate pitting.

This pit nucleation model, in common with nearly all other models, fails to account for the specific effect of chloride ions. In particular there has been no justification as to why chloride anions promote pitting, whereas sulphate anions, for instance, do not. Sulphate anions also become adsorbed on the metal sur-face, and may be assumed to agglomerate at the sites of F e

2+ emergence. Iron(II)

sulphate also undergoes hydrolysis in aqueous solutions; it can therefore render the surface acidic and hence destroy the passive film. The only difference in the behaviour of the two anions appears to lie in the difference in the protective properties of the corrosion products formed.

It seems that in the presence of Cl~ anions, the development of pits (es-pecially on austenitic steels) is due to the formation on the bottom of pits of a semiprotective layer of the chloride salt which prevents repassivation, without entirely suppressing the dissolution of metal. The presence of a sulphate salt layer on the bottom of pits would probably be more protective, facilitating repassivation.

1.7 Kinetics of pit growth

The majority of investigations on the kinetics of pit growth have been made under potentiostatic conditions. At a constant potential of the electrode, the time dependence is measured for such quantities as current intensity, the number of pits formed, their depth and diameter.

Engell and Stolica in 1959 [83] were the first to show that the following rela-tion holds between the current intensity / and pit development time t:

i = ktb (1.1)


These investigators have established that if the number of pits Ν per unit area of specimen is constant, then b = 2, whereas if pit density increases with time, b = 3. It has been assumed that the pits are hemispherical in shape and that the current density inside pit is constant; in line with this assumption the pit radius increases directly with time.

A large number of researchers have obtained results that fit the above rela-tion. The more recent studies, however, have given values of b both much lower than 2 and much higher than 3, depending on whether the number of pits pro-duced inreases directly with time or whether it follows a more complex function.

Relation (1.1) may be expressed in the form [84]:

(..2,

where d is the pit diameter and h is the depth. If d ~ tbd

9 h ~ tbh, and Ν ~ t

b"

where bd, bh, and bN are values of the exponents of the time dependence of the diameter, depth and number of pits, respectively, then b = 2bd + bh + bN—l.

Figure 1.10 illustrates the / vs. pit growth time relation for a given steel and

308 Κ

293 Κ

0.1 « 1 1 I ι ι m l ι ι ι ι ι m l ι ι ι

1 10 100 Time,- m in

Fig. 1.10. Current density vs. pit growth time for a 17Crl2Ni2MoTi type steel in 0.5 M NaCl + + 0.05 M H 2S 0 4 at 293, 308 and 323 Κ; Ε = 880 m V N HE (after [85]).

COMPOSITION OF THE ELECTROLYTE WITHIN PITS 33

TABLE 1.2

VALUES OF THE EXPONENTS b, bd, AND b„ FOR THE EXPERIMENTS ILLUSTRATED IN FIG. 1.10

Temperature, Κ Experimental Calculated

Temperature, Κ bd bh b b = 2bd + bh + bN-\

293 0.91 0.75 1.6 1.6

308 0.51 0.50 0.6 0.6

323 0.29 0.35 0.1 0.1

for a given solution at different temperatures [85]. The values of b are shown. Table 1.2 reports the b, bd, and bh values found experimentally and calculated from formula (1.2), showing a good agreement in the values of b.

1.8 Composition of the electrolyte within pits

Until recently the precise composition of the solution filling the growing pits has not been known. It is generally recognized that iron within pits passes into solution as the bivalent cation, that the solution is rich in chloride anions, and that it exhibits a low pH [86]. Studies of the electrolyte composition inside pits formed on surface of chromium-nickel stainless steels have shown that the concentration of individual components (Fe, Cr, Ni) in the resultant sol-ution is roughly the same as in the alloy [52, 58, 87]. Several researchers have established the presence of precipitates at the bottom of the pit [53, 70, 87].

Measurements of pH variations and changes in chloride concentration are of particular significance in the interpretation of the processes occurring inside a developing pit. The pH measurements inside a pit, crevice or fissure have been carried out by independent methods by several investigators. The determina-tion of chloride concentrations is, however, much more difficult. Suzuki et al [88] made such a determination with the aid of artificial pit models; he found a sub-stantial enrichment of the solution in CI" anions.

Recently [89] quantitative determinations of chloride content have been successfully made in natural pits growing in an austenitic stainless steel, type 18Crl2Ni2MoTi. Steel specimens were polarized anodically to a constant po-tential of 860 m F N HE in a 0.5 MNaCl solution made acidic with H 2S 0 4, at 293 K. The variations of current density and pit diameter with time were followed. Under specified conditions covered pits developed under the passive film,

TA

BL

E 1

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DIA

ME

TE

R d

, D

EP

TH

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F A

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Cl"

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N C

ON

CE

NT

RA

TIO

N

IN P

ITS

AF

TE

R

120,

180

AN

D

240-

min

CO

RR

OS

ION

d,

mm

1.

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1.22

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107

170

150

102

92

277

min

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6.3

5.9

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mm

0.

87

1.18

1.

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1.76

1.

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2.03

2.

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, m

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0.26

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COMPOSITION OF THE ELECTROLYTE WITHIN PITS 35

but the convexity of the pit-covering film allowed the estimation of the pit di-ameter.

The specimen was frozen in sol;d C 0 2 and the frozen pit content was analysed

microcolorimetrically by the Harczenko-Chohij-Lenarczyk method [90], which involves oxidizing chlorides to chlorine. The accuracy of the determinations was ±0.2 μg CI. The volume of the solution in the pit was evaluated from the diameter and the depth of the pit.

For every pit, a tiny hole was present in the middle of the covering film. Other holes appeared in the film some time later in the pit development process, through which electrolyte could emerge. In the initial stage of pit growth the pits developed uniformly in all directions, assuming an irregular shape, depend-ent on the electrode position, in later stages.

The CI" concentration inside pits, as determined at a definite stage of the corrosion process, becomes less as the pit volume increases (Table 1.3). The differences in pit development rate are probably due to unequal permeabilities of the film covering pits. The more porous the film, the higher the pit growth rate. ' To show the variation in Cl~ concentration inside pits with time, pits were classed into two groups: one, which included slow growing pits (pit diameter

classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which included slow growing pits (pit diameter

classed into two groups: on


from 0.48 to 0.65 mm after 120 min) and the other, which comprised pits devel-oping more rapidly (pit diameter from 0.65 to 1.4 mm after 120 min).

Figure 1.11 shows the chloride concentration inside pits vs. their diameter and growth time. In small pits the CI" concentration characteristically reaches 12 M, i.e. it exceeds many times the average concentration of these ions in the bulk electrolyte. As the pits increase in size, the Cl~ concentration declines to about 2.5 Af, due to loosening of the pit-covering film. This facilitates con-vection and diffusive exchange between the solution inside and outside the pit. Furthermore, it was observed that the initially hemispherical shape of the pit gradually becomes flattened. This indicates an uneven current density and Cl~ concentration within pit and also the variability with time of these parameters.

A knowledge of the chloride concentration within pits has contributed to an understanding of the very low pH values of the electrolyte in pits and crev-ices. On the basis of studies of numerous authors [85, 87, 92-94], in the so-called occluded cells, the pH was found to drop to much lower values than those re-sulting from the hydrolysis of ferrous salts. In the case of crevice corrosion of type 304 L stainless steel, Peterson et al. [92] established pH values ranging from 1.2 to 2.0, while Suzuki et al. [88] demonstrated that in model pits, the pH may drop to 0.6 or 0.8. In covered pits that developed in 18Crl2Ni2MoTi steel specimens immersed in a solution of an average pH of 6.5, the value was found to reach 1.3 [85].

The reason for such a high acidity of the solution in pits and crevices has been sought either in the formation of hydroxychloride complexes of cations, or in an increased activity of hydrogen ions in solutions of a high chloride concentra-tion [88, 95]. According to Rhodes [95] a decreased pH in cracks formed in stainless steel during stress corrosion in chloride solutions is the result of the formation of insoluble oxides in the anodic process:

2Cr + 3 H 20 C r 20 3 + 6 H+ + 6e

Studies carried out on model solutions that resemble the solutions in the pits led us to believe that the reason for the low pH of these solutions is the occurrence of high concentrations of chloride ions [89]. The results illustrated in Fig. 1.12 show how variations in FeCl2 and CrCl3 concentrations affect the pH of the solution. It follows from these data that in the presence of high CI" concentrations, the pH values obtained are of the same order as those measured in the occluded corrosion cells.

PIT MORPHOLOGY 37

Chloride concentration, M

Fig. 1.12. The concentration effect of FeCl2 · 3 H 20 and CrCl3 · 6 H 20 on the pH of the sol-ution (after [89]): Ο 30 min after solution preparation; — Ο— 24 h after solution preparation.

1.9 Pit morphology

Depending on the experimental conditions, corrosion pits may be of various shapes. Thus, we may see open or closed pits, hemispherical or spherical caps, regular polyhedral pits (i.e. lined with flat walls that reveal the crystalline struc-ture of the metal involved), or entirely irregular ones. It is not completely clear what is responsible for the shape of the developing pit. Among the likely factors are: composition of the electrolyte and conditions within the pit, polarization conditions, nature, composition and structure of the metal, as well as the compo-sition and structure of the passive film.

Herbsleb and Engell [53, 96] and Engell and Stolica [83] observed isotropic pits to form on iron in chloride solutions containing sulphates, whereas on pure iron subjected to corrosion in a buffered borate solution containing chlorides in various amounts, Janik-Czachor [18] obtained solely anisotropic pits. The pit walls were usually the {110} planes, which are the closest packed faces of α-iron crystals.

In sulphate solutions containing Cl~ Vetter and Strehblow [24] observed isotropic hemispherical pits on iron only when the SO4" concentration in the


solution exceeded 0.5 M. A low Cl~ concentration (10~3M) resulted in the forma-

tion of irregular pits. On the bottom of hemispherical pits formed on iron, Herbsleb and Engell

[53] found the presence of a sulphate layer which, in their opinion, is suggestive of participation of the SOl~ anion in the isotropic dissolution of the metal. These investigators suggested that SO4- anions facilitate the formation of an adherent polishing film which might favour the hemispherical shape assumed by the pit. A similar viewpoint is expressed by Frankenthal and Pickering [97]. They claim that for iron, the pit morphology is determined by the composition of the electrolyte. If a film formed inside the pit controls the localized corrosion rate (as for instance in the case of sulphate solutions), then the initially poly-hedral pit undergoes electropolishing and becomes round. If, on the other hand, such a film cannot be formed (e.g. in perchlorate solutions), the pit maintains its polyhedral shape.

In a study on austenitic stainless steel, Hisamatsu et al. [75] observed solely round pits over the entire range of the anodic polarization curve.

It has been proposed that the shape of growing pits depends on whether the metal dissolution rate is activation or diffusion controlled. If this assumption is correct, then with increasing potential and concentration of solution inside the pits, a change from the initial activation control to diffusion control should be observed, i.e. a change to a mechanism that would favour generation of hemi-spherical pits. The results of certain investigators support this concept. Brauns and Schwenk [19] for example observed pits of a square or hexagonal cross section for 18Crl9Ni steel only when the electrode was kept at a potential cor-responding to the pit nucleation potential. The generation of flat-walled pits in austenitic steels at potentials close to the critical potential of pit nucleation has been noted by numerous investigators.

In the light of recent findings [33], pit shape may also depend on the struc-ture of the oxide film. In the case of thin films formed on nickel during oxida-tion in the air at 373 and 473 K, the majority of pits generated in a corrosion experiment in a 0.1 AfNa2S04 + 0.013 M NaCl solution over a wide range of anodic potentials was of a square or hexagonal cross section and governed by the crystallographic orientation of the substrate. This was presumably due to the epitaxial structure of the oxide film produced under the oxidation condi-tions of the experiment. It follows from observations of the shape of pits formed on specimens oxidized at 473 Κ that in the case of cold-worked 99.5% nickel (which was the most susceptible to pitting), chloride ions attacked the {100} face. This is the least closely packed plane in a face-centered crystal lattice.

PIT MORPHOLOGY 39

On nickel specimens oxidized at 573 and 673 K, only spherical-cap shaped pits were observed. Presumably, oxidation at higher temperatures damages the metal surface, and results in the formation of pits of a shape unrelated to orienta-tion of metal grains, but determined by the diffusion-controlled transport of reactants and by the electrochemical conditions within pits.

The formation of the flat-walled pits is evidence for the hypothesis that under specific conditions the rate of metal dissolution is controlled neither by ohmic polarization nor by diffusion but by crystal anisotropy.

In the case of austenitic stainless steels, as a rule, covered pits are formed, with a tiny hole in the pit-covering film. The film is relatively thick, as it con-sists not only of the passive oxide film but also of the corrosion products ad-hering to it.

The generation of covered pits indicates that the passive film is virtually insoluble both in the surrounding electrolyte and in the solution within pits. On the other hand, in less resistant steels, such as carbon or low-alloy types, open pits are produced, as the oxide films are easily attacked by the aggressive solution.

The formation of hemispherical pits suggests that processes similar to those occurring during electropolishing take place in the interior of the pits. The experi-mentally found presence of CI" anions at very high concentrations in pits sup-ports this suggestion. At the pit bottom a salt layer may occur, which resembles an electropolishing layer. The change in pit shape from hemispherical to flat-tened spherical one is related to an uneven distribution of current intensity, caused by a salt layer whose thickness varies with time, within the pit.

Interestingly, the degree of accumulation of chloride ions in pits was found [98] to depend not merely on the permeability of the pit-covering film, but also on orientation of the metal surface in relation to the gravitational field. De-pending on the position of the attacked metal surface and the direction of pit growth, three major pit types can be observed, shown schematically in Fig. 1.13. Type 1 pits develop when the exposed surface of the electrode is horizon-tally arranged and pointing downwards, so that pits develop in the upward direction.

Type 2 features pits that are growing on specimens whose active surface is horizontal, and pointing upwards so that the pit grows downwards, in the direction of the force of gravity.

Type 3 represents the pit growing on a vertical electrode wall. In the initial stage of pit growth the shape is independent of the eletrode position. First, hemi-spherical pits are formed [98] which then assume a spherical cap shape. It is


Type 1

(a) (b)

Type 2

(a) (b)

Fig. 1.13. Pit types depending on the orientation of the exposed metal surface and the direction of pit growth (after [98]): (a) presumed distribution of corrosion products; (b) pit shape after prolonged corrosion.

only in the later stages that the pit shape changes in a manner dependent on the direction of their further development.

Type 1 pits, after a prolonged corrosion, assume the shape shown in Fig. 1.13b. The cavity at the centre of the pit bottom is due to a higher local rate of metal dissolution.

Type 2 pits gradually become flattened, as in Fig. 1.13d. Type 3 pits, at a later stage of corrosion process, grow at a somewhat higher

rate in the lower than in the higher part of the attacked zone. The ultimate shape of such pits is shown in Fig. 1.13f.

These various shapes may be accounted for as follows. As a result of metal dissolution in the pit and the restricted outward transport of corrosion products, a layer of solid matter is produced at the pit bottom, over which a saturated solution of the corrosion products is present. The thickness of the solid layer

DEVELOPMENT OF PITS 41

affects the pit shape as shown in Fig. 1.13a, c and e. In the case of type 1 pits the solution exits mainly through the centre hole in the pit-covering film. In the middle of the pit bottom, from which the outflow of the anodic product is the easiest, the thickness of the solid layer that hampers the dissolution process is less, i.e. the local current density is larger there than in the outer portions of the pit bottom, where the layer is thicker. A hemispherical cap is consequently formed.

In type 2 pits a relatively thick layer of solid corrosion products is formed, which uniformly covers the pit bottom. As a result, the anodic current also is uniformly distributed over the whole of the bottom, which accordingly assumes a flat shape. A thick layer hinders pit growth. Therefore, corrosion is largely operative at the pit edges, where mixing is more intense and concentration of the aggressive solution is lower. A consequence of an uneven coverage by the solid layer of the pit bottom and its edges is an uneven distribution of the anodic current within the pit; thus, lateral pit enlargement goes at a faster pace than deepening.

To sum up, the difference in shape of corrosion pits generated on differently oriented electrodes is due to the conditions under which the dissolution reaction products are removed outwards form the site where the reaction takes place.

1.10 Development of pits

The process of metal or alloy dissolution in a pit may proceed in two ways de-pending on the dissolution rate of the metal or alloy in acid solution. For a metal (or alloy) that exhibits a high passivation current, e.g. iron, the metal dissolution reaction is activation controlled at the initial stage of pit growth. The pit bottom potential corresponds to the potential of the metal in its active condition. Flat-walled pits would be expected to form in this case, at least in the initial stage of corrosion. If the dissolution process changes from activation control to diffusion control, the pits become hemispherical in shape. For instance, an increased dissolution rate at a higher anodic potential will cause originally flat-walled pits to become smooth and round. This is due to sup-pression of the exchange of the solution inside the pits with the surrounding elec-trolyte, which results in an increased concentration within pits. If, on the other hand, the metal or alloy (e.g. austenitic steel) readily undergoes passivation in acid solution, then pits will grow steadily only if the pH within pits is low and there is a salt-containing layer on the pit bottom. In this case, pit growth is diffusion cntrolled and spherical pits are formed,


Chloride ions play a double role in pit growth. On the one hand, in acid sol-ution they increase the hydrogen ion activity, and thus enhance the corrosion rate; on the other hand, once a sufficiently low pH is attained, chlorides cause a salt layer to form at the bottom of pits, which tends to hinder metal dissol-ution but prevents passivation.

High-alloy stainless steels which readily undergo passivation tend to form pits that are covered with a passive film. Here, the metal dissolution rate is diffu-sion controlled by the restricted transport of reactants and corrosion products to and from the reaction site, and also by the local formation of a salt layer. A limited diffusion of water into the pit may also affect its growth [99, 100].

A particular role in the development of covered pits is played by a low per-meability of the film, which hinders the exchange of solution with the surround-ings. As measurements of chloride concentration inside pit show, the highest dissolution rate is in pits where the Cl~ concentration is not too high (pits grow-ing upwards) and also at sites where chloride accumulation in the form of a salt layer is rendered more difficult.

100

11—ι ι ι ι ι ι I ι ι I 50 100 300

Time, min

Fig. 1.14. The effect on current density of breaking of the pit-covering film during the devel-opment of a type 3 pit (Fig. 1.13) on 18Crl2Ni2MoTi steel in 0.5 MNaCl + 0.05 M H2SO* at a potential of 860 mVNHE at 293 Κ (after [98]). Arrow indicates the moment at which the film was broken.

DEVELOPMENT OF PITS 43

A sudden increase in film permeability, e.g. by mechanical extension of the original hole or piercing of an adjoining one, can dramatically change the convection conditions in the operating pit and entail a change in the resulting current. Figure 1.14 illustrates an example of current rise as a function of time at a constant potential with only one covered pit present on the electrode sur-face. As soon as the film was pierced with a glass point, an increase of the anodic current was oberved. If, however, convection was made much easier, e.g. by a complete removal of the film, this would result in pit passivation, and hence the suppression of corrosion.

Polarization curves for a stainless steel grade 18Crl2Ni2MoTi in 1 M HC1 with variable amounts of FeCl2 added, shown in Fig. 1.15, indicate that there is a maximal concentration of chloride anions above which corrosion rate decreases.

10000

1000

e <

100

1.75M FeCl 2

1.5MFeCl2

IMFeCU

1MHCI +*MFeCt2

in 11 ι » ι ι ι ι ι—ι—ι—ι—ι—ι—ι—ι—ι—L 0 02 0.4 0.6 0.8 1.0 1.2 1.4

Potential, VN HE Fig. 1.15. Polarization curves for 18Crl2Ni2MoTi steel in 1 M H Q with various quantities of FeCl2 added (after [89]).


The highest current densities are obtained in solutions containing 1.75 M and 2 M ferrous chloride, whereas higher CI" concentrations result in a decreased cur-rent density. Pitting occurs in the solutions containing 1 M and 1.5 MFeCl2.

The results of this study [89] led to the conclusion that even though a high Cl~ concentration and low pH values are prerequisites for pit development in the initial stage, after prolonged corrosion the formation of a supersaturated solution of corrosion products at the bottom of pits results in suppression of the metal dissolution process.

1.11 Pitting of iron and steel in chloride-free solutions

Pitting is usually associated with the presence of chloride ions in solution [101]. This is understandable, as corrosion of this type in other environments is rare. Thus, all the processes described hitherto in this chapter are concerned with the corrosion of iron and steel in chloride solutions. Under certain conditions, however, pitting of iron and steel may occur in halogen-free solutions, and the mechanism of pitting in these systems may differ from that in chloride solu-tions.

Examples of such systems, where halide ions are absent but pitting does occur, are given below.

1. Iron in sulphate solutions in the active-passive region of potentials. Compe-tition between oxide film formation and iron sulphate film formation on a free metal surface is a cause of pitting. The formation of a salt prevents the passiva-tion of iron in the potential region close to the active-passive transition [102-105].

2. Iron in borate and silicate solutions in the transpassive region of potentials. It has been found by ellipsometry that in the transpassive potential region a sec-ond film, characterized by low optical parameters, is formed on top of the oxy-hydroxide film. In this case, breakdown of the film occurs at certain sites of the surface, where more soluble compounds of iron with H 3B 0 3 or H 2S i 0 3 are formed [106].

3. Carbon steel in water at 323 to 473 K. Pitting of carbon steel was found to occur in oxygenated pure water in the range 323 to 473 Κ [107-109]. The pits nucleate at manganese sulphide inclusions. The aggressive species derived from H2S come from dissolution of sulphide inclusions. It is known that HS~ (or H2S) catalytically increases the dissolution rate of iron. In addition, corrosion prod-ucts containing sulphur prevent repassivation of the neutral surface.

4. Stainless steel in methanol at anodic potentials. Pitting of stainless steels in methanol solutions at anodic potentials occurs as a result of film breakdown

PITTING AT ELEVATED TEMPERATURES 45

under high field strength and attack on the bare metal surface, due to the forma-tion of formic acid. In the presence of higher concentrations of water, pitting does not occur, because the water repassivates the pits [110].

1.12 Pitting at elevated temperatures

The preceding sections have mainly dealt with pitting at room temperature or close to it. In recent years, several studies have concentrated on pitting at higher temperatures. It has been shown that under these conditions the pitting poten-tial is especially strongly affected by temperature.

For example, Manning and Duquette [111], Lin et al. [112] and others found that with increasing temperature the pitting potential of various iron and nickel base alloys in chloride solutions shifts to much less positive values relative to that at 293 K, but starting from a certain temperature (about 423 K) does not change significantly.

In another study [113], the pitting potential of Alloy 600 in a deaerated 0.1 MNaCl solution was shown to change from 400 mVS HE at 333 Κ to —300 mVsHE at 553 K.

The above variations in the pitting potential with temperature are probably related to some changes in the composition and structure of the passive film.

Measurements of the susceptibility to pitting of Alloy 600 performed in de-aerated CuCl2 solutions of different concentration showed that at 333 Κ more than 2000 ppm of CuCl2 was needed to cause pitting, while above 473 Κ as little as 20 ppm CuCl2 sufficed to initiate pits in seven hours [114].

At high temperatures, the kinetics of pit growth obeys the same law as at ambient temperature; namely, the rate at which pits grow in depth is governed by

h = ktb

where h is the pit depth, t is time and b is a constant. In deaerated 100 ppm CuCl2 solution, under open circuit conditions, pits

were found [115] to grow in depth on Alloy 600 proportionally to i0 4

, whereas in deaerated 0.1 M NaCl at applied potentials 150 mV more positive than the respect-ive pitting potentials, the pit depth was proportional to i

0 3, t

0'5 and t

0'7 at

373, 423 and 523 K, respectively. The reason for those differences is not clear. After the same exposure time of Alloy 600 to the above two solutions, the

density of pits formed at 423 Κ was less, but their depths were greater than at higher temperatures [115].


1.13 References

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[10] Smialowski, M., Szklarska-Smialowska, Z., Rychcik, M., Szummer, Α., Corr. Set, 9, 123 (1969).

[11] Eklund, G., Jernkont. Ann., 155, 637 (1971). [12] Hoar, T. P., Mears, C. D., Rothwell, G. P., Corr. Sci., 5, 279 (1965). [13] Cihal, V., Kashova, J., Kubelka, J., Zashchita Met., 8, 19 (1972). [14] Poyet, P., Desestret, Α., Mom. Sci. Rev. Met., 72, 133 (1975). [15] Richter, B., Prause, W., Werkstoffe u. Korrosion, 27, 856 (1976). [16] Scotto, V., Ventura, G., Traverso, E., Corr. Sci., 19, 237 (1979). [17] Manning, P. E., Duquette, D . I., Savage, W. F., Corrosion, 35, 151 (1979). [18] Janik-Czachor, M., Brit. Corr. J., 6, 57 (1971). [19] Brauns, E., Schwenk, W., Arch. Eisenhuttenw., 32, 387 (1961). [20] Lunarska, E., Szklarska-Smialowska, Z., Janik-Czachor, M., Corrosion, 31, 231 (1975). [21] Janik-Czachor, M., Lunarska, E., Szklarska-Smialowska, Z., Corrosion, 31, 394 (1975). [22] Tomashov, N. D., Chernova, G. P., Markova, N., Corrosion, 20, 166t (1964). [23] Szklarska-Smialowska, Z., Janik-Czachor, M., Corr. ScL, 7, 65 (1967). [24] Vetter, K. J., Strehblow, Η. H., Ber. Bunsengesel. Phys. Chem., 74, 1024 (1970). [25] Kuznetsov, V. Α., Jofa, Ζ. Α., Zhurn. Fiz. Khim., 21, 201 (1947). [26] Jofa, Ζ. Α., Tomasheva, G. N., Zhurn. Fiz. Khim., 34, 1036 (1960). [27] Jofa, Ζ. Α., Batrakov, V. V., Ba Cho Ngok, Zashchita Met., 1, 55 (1965). [28] Syrret, B. C , Macdonald, D. D., Shih, R., Final Report to NSF (RANN)—ERPA

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[31] Szklarska-Smialowska, Z., Jurek, J., Corrosion, 32, 294 (1976). [32] Jurek, J., Szklarska-Smialowska, Z., Corrosion, 32, 309 (1976). [33] Oranowska, H., Szklarska-Smialowska, Z., Corr. ScL, 16, 363 (1976). [34] McBee, C. L., Kruger, J., Proc. Intern. Confer. Localized Corrosion, Williamsburg 1971,

Eds. R. W. Staehle, B. F. Brown, J. Kruger, A. Agrawal, NACE, 1974, p. 252. [35] Bevilacqua, F., Brown, G. M, GNEC-295, October 18, 1963.

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REFERENCES 47

[38] Seyes, Α. Α., van Haute, Α. Α., Brabers, M. J., Werkstoffe u. Korrosion, 25, 663 (1973). [39] Janik-Czachor, M., Szummer, Α., Szklarska-Smialowska, Z., Corr. Sci., 15, 775 (1975). [40] Szummer, Α., Janik-Czachor, M., Brit. Corr. J., 9, 216 (1974). [41] Overman, R. F., Corrosion, 22, 48 (1966). [42] Heusler, Κ. E., Fischer, L., Werkstoffe u. Korrosion, 27, 551 (1976). [43] Ambrose, J. R., Kruger, J., Proc. 4th Intern. Congress on Metallic Corrosion, NACE,

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[44] Abd-Rabbo, M. F., Wood, C. C , Richardson, J. Α., Jackson, C. K., Corr. ScL, 14, 645 (1974).

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USA-Japan Seminar, Honolulu 1975, Eds. R. W. Staehle, H. Okada, NACE, 1976, p. 91.

48 NUCLEATION AND DEVELOPMENT OF CORROSION IN IRON A N D STEEL

[75] Hisamatsu, Y., Yoshii, T., Matsumura, Y., Proc. Intern. Confer. Localized Corrosion, Williamsburg 1971, Eds. R. W. Staehle, B. F. Brown, J. Kruger, A. Agrawal, NACE, 1974, p. 427.

[76] Lorenz, W. J., Fischer, H., Electrochim. Acta, 11, 1597 (1966). [77] Hoar, T. P., J. Electrochem. Soc, 117, 17 C (1970). [78] Galvele, J. R., Passivity and its Beakdown on Iron and Iron Base Alloys, USA-Japan

Seminar, Honolulu 1975, Eds. R. W. Staehle, H. Okada, NACE, 1976, p. 118. [79] de Wexler, S. B., Galvele, J. R., / . Electrochem. Soc, 121, 1271 (1974). [80] Galvele, J. R., Publicaciones Tecnicas y Cientificas ECOMAR, ECOVAR 10021 Ρ 1975. [81] Lin, L. F., Chao, C. Y., Macdonald, D . D., / . Eoectrochem. Soc. 128, 1194 (1981). [82] Szklarska-Smialowska, Z., Sulfide Inclusions in Steel, American Society for Metals, Eds.

J. J. de Barbadillo, E. Snape, 1974, p. 380. [83] Engell, H., Stolica, N., Z. Physik. Chem. N. F., 20, 113 (1959). [84] Szklarska-Smialowska, Z., Werkstoffe u. Korrosion, 22, 780 (1971). [85] Szklarska-Smialowska, Z., Mankowski, J., Corr. Sci., 12, 925 (1972). [86] Szklarska-Smialowska, Z., Proc. Intern. Confer. Localized Corrosion, Williamsburg 1971,

Eds. R. W. Staehle, B. F. Brown, J. Kruger, A. Agrawal, NACE, 1974, p. 312. [87] Novakovskii, V. M., Sorokina, A. N., Corr. Sci., 6, 227 (1966). [88] Suzuki, T., Yamabe, M., Kitamura, Y., Corrosion, 29, 18 (1973). [89] Mankowski, J., Szklarska-Smialowska, Z., Corr. ScL, 15, 493 (1975). [90] Marczenko, Z., Choluj-Lenarczyk, L., Chem. Anal., 11, 1221 (1966). [91] Vetter, K. J., Strehblow, H. H., Willgalis, Α., Ber. Bunsengesel. Phys. Chem., 75, 822

(1971). [92] Peterson, M. H., Lennox, T. J. Jr., Grooves, R. E., Materials Protection, 9, 1, 23 (1970). [93] Brown, B. F., Fuji, C. T., Dahlberg, E. P., J. Electrochem. Soc, 116, 218 (1969). [94] Brown, B. F., Corrosion, 26, 249 (1970). [95] Rhodes, P. R., Corroison, 25, 462 (1969). [96] Herbsleb, G., Engell, H. J., Z. phys. Elektrochem., 65, 881 (1961). [97] Frankenthal, R. P., Pickering, H. W., / . Electrochem. Soc, 119, 1302 (1972). [98] Mankowski, J., Szklarska-Smialowska, Z„ Coorr, ScL, 17, 725 (1977). [99] Tomashov, N. D., Chernova, G. P., Markova, Ο. N., Zashchita Met., 7, 104 (1971).

[100] Faita, G., Mazza, F., Bianchi, G., Proc. Intern. Confer. Localized Corrosion, Williamsburg, 1971, Eds. R. W. Staehle, B. F. Brown, J. Kruger, A. Agraval, NACE, 1974, p. 34.

[101] Shreir, L. L., Corrosion (in Polish), Vol. I, WNT, Warszawa 1966, p. 147. [102] Matsuda, S., Uhlig, H. H., / . Electrochem. Soc, 111, 156 (1964). [103] Szklarska-Smialowska, Z., Corr. ScL, 18, 97 (1978). [104] Szklarska-Smialowska, Z., Mrowczynski, G., Brit. Corr. / . , 10, 187 (1975). [105] Tousek, J., Corr. ScL, 12, 15 (1972). [106] Kozlowski, W., Szklarska-Smialowska, Z., / . Electrochem. Soc, 131, 723 (1984). [107] Coriou, H., Grail, L., Peltraz, M., Perez, Α., Corrosion et Anticorrosion, 12, 358 (1964). [108] Choi, H., Beck, F. H., Szklarska-Smialowska, Ζ., Macdonald, D. D., Corrosion, 38, 125

(1982).

[109] Mizumo, T., Pednekar, S., Szklarska-Smialowska, Z., Macdonald, D . D. , Proceedings of the International Symposium on Environmental Degradation of Materials in Nuclear Power Systems—Water Reactors, NACE, 1984, p. 395.

REFERENCES 49

[110] Szklarska-âmialowska, Z., Mankowski, J., Corr. ScL, 22, 1105 (1982).

[ I l l ] Manning, P. E., Duquette, D. J., Corr. ScL, 20, 597 (1980). [112] Lin, L. F. Cragnolino, G, Szklarska-Smialowska, Z. Macdonald, D . D., Corrosion, 37,

616 (1981). [113] Park, J. R., Szklarska-Smialowska, Z., Corrosion, 41, 665 (1985). [114] Park, J. R., Szklarska-Smialowska, Z., Proc. 2nd Intern. Symp. on Environmental Degra-

dation of Materials in Nuclear Power Systems—Water Reactors, Monterey, CA, 1985, p. 456.

[115] Szklarska-Smiatowska, Z., Grimes, D., Park, J. R., Corr. ScL, 27, 859 (1987).

2. Stress Corrosion Cracking of Structural Steels in Nitrate Solutions

2.1 Introduction

Stress corrosion cracking (SCC) is one of the most dangerous kinds of corrosion damage occurring in metals. In the chemical industry, SCC-induced corrosion failures amount to 21-25% of all corrosion damage [1, 2], coming next to gen-eral corrosion (approx. 32%).

SCC occurs as a result of the combination of an aggressive environment and tensile stresses. It proceeds in a macroscopically brittle manner, despite a con-siderable plasticity of the materials under normal conditions. The cracking can occur under stresses much below the ultimate tensile strength in air. often below the yield stress.

The failures under service conditions are usually unexpected, because they occur with very weak general corrosion and with no obvious corrosion deposits. Moreover, the cracks are fine and hardly visible. In view of the problems associ-ated with early detection of impending SCC, the consequences may sometimes be disastrous.

The number of reported failures attributed to SCC is constantly increasing, and therefore much attention has been paid to this problem in the last two or three decades. The subject is treated in many monographs [3-7] and conference proceedings [2, 8-12],

SCC may occur in the presence of the following three factors: a specific en-vironment, a susceptible material, and tensile stresses.

The number of environments known to cause SCC is constantly increasing so that now it is difficult to consider them in terms of their strict specificity. At an appropriate electrode potential, pH and temperature, SCC may appear in many solutions which, until recently, were regarded as entirely safe.

Nearly all commercially important metals and alloys are susceptible to SCC. Cracking occurs in carbon steels and austenitic stainless steels, as well as in alu-

GENERAL CHARACTERISTICS OF CARBON STEELS 51

minium, titanium, magnesium, and nickel alloys. SCC of carbon steels is a par-ticularly serious problem because of their extensive use in industry. This problem is the subject of this chapter.

2.2 General characteristics of carbon steels

Carbon (unalloyed) steels contain carbon as the major addition. Other consti-tuents occur in minor quantities and are related to particular metallurgical pro-cesses; the main additives are Mn (up to 0.8 wt. %), Si (up to 0.4%), Cr (up to 0.3%), Ni (up to 0.3%) and Cu, W, and Co (each up to 0.2%). Sulphur and phos-phorus are impurities; they occur in amounts up to approx. 0.03% each.

Steels with purposefully introduced constituents such as Ni, Cr, Mn, W, Mo, V, Co, Si, Al, Cu, Ti, B, ets. are referred to as alloy steels.

As concerns the carbon content, steels are classed into low-carbon (up to 0.25% C), medium-carbon (0.25 to 0.65%) and high-carbon (0.65 to 2.0% C).

The following categories of use can be distinguished: (a) structural steels, used for the construction of machines, equipment and

civil engineering structures; (b) tool steels, used for the production of a wide variety of tools; (c) alloy steels, which exhibit high corrosion and heat resistance due to the

alloying additions. This chapter deals only with structural steels. Depending on the carbon con-

tent, composition and heat treatment, the ultimate tensile strength Rm of these steels may vary between approximately 300 and 2000 MPa. Low and medium strength steels have a ferritic or ferritic-pearlitic structure; high-strength steels may have a bainite structure or a tempered martensite structure (after water or oil quenching, followed by tempering at a temperature of 473 to 923 K). The tempered-martensite steels may contain approx. 0.25 to 0.7% C and addi-tionally Mn up to 1.8%, Si up to 2%, Ni up to 4.4%, W up to 1.2%, Mo up to 0.4% or V up to 0.3%.

Structural steels are widely used because of their good engineering properties and low price. They are used in large quantities in the construction engineering and in chemical industry for tanks, pipes, heat exchangers and other equipment. A serious limitation in their use is their low resistance to corrosion in a wide variety of media; nevertheless, their use in such media is often more economically sound than the use of alloy steels. In dilute acid or alkaline solutions these steels undergo a fairly uniform dissolution; in near-neutral solutions or in alkalis, on the other hand, these steels may undergo intergranular corrosion and SCC.

52 STRESS CORROSION CRACKING OF STRUCTURAL STEELS

2.3 Stress corrosion cracking of structural steels in practice

A classical case of SCC is the well known caustic embrittlement, also known as boiler embrittlement, of unalloyed ferritic steel which at one time was a very hazardous type of damage of steam boilers [13]. It appeared in the form of intergranular cracks in riveted joints where concentrated lye accumulated as a result of water evaporation. Evaporators for N H 4N 0 3 have also long been notorious for SCC damage [14].

Present-day SCC of ferritic steels is a serious problem primarily in the fol-lowing four areas of engineering and industrial practice:

(a) the fertilizer industry, particularly that manufacturing ammonium nitrate (NH4NO3) and calcium nitrate (Ca(N03)2) [15]. The most common places of damage are fertilizer containers and structures exposed to nitrate dust [16];

(b) soda lye tanks and alkaline-water steam boilers; (c) liquid-ammonia tanks and pipelines; (d) underground pipelines, especially those for long-distance gas or water

transport. A disastrous gas pipeline failure occurred in 1965 in the United States near a gas compression station. The cracks developed beneath the coating, prob-ably due to nitrates, or more likely, to the hydroxide and carbonate COf^/HCOg* ions that form under the conditions of cathodic protection of pipelines [17].

SCC damage has also been observed in the following cases [18]: (e) coke gas scrubbers in liquids containing N H 3 -f C 0 2 + H2S + HCN [18]; (f) pipelines in an H 20 / C O / C 0 2 environment [19]; (g) steam generator blades made of low-alloy steel (3% Cr, 0.5% Mo) in

high-purity steam, but containing e.g. traces of acetate anions [20]; (h) gas combustion equipment; (i) bridges and cables, possibly from nitrates in rain water. High-strength martensitic steels are used for high-reliability airfield or civil

engineering structures and in very large quantities as concrete reinforcement. Their cracking may occur even in pure water or moist air. Failures of these materials under service conditions happen seldom, but if they do, they may be disastrous.

2.4 Environments causing stress corrosion cracking of structural steels

It was believed for many years that only OH~ or NOj ions were responsible for SCC of ferritic steels. Consequently, the various SCC cases observed were ascribed to the presence of these ions in trace amounts. The list of currently

INVESTIGATION OF STRESS CORROSION CRACKING 53

known SCC-inducing environments is, however, much more extensive. The en-vironments and the conditions of SCC are given in Table 2.1.

Among the media that induce SCC of ferritic and martensitic steels, nitrates exert the strongest effect. For this reason nitrates are often used as test sub-stances. The results of nitrate tests allow a rapid evaluation of the general re-sistance of a steel to SCC, as well the choice of a steel to be used also in other solutions, especially alkalis. Much basic research has been carried out in ni-trates, and a number of general conclusions can be drawn from the results, since they also apply to other environments. Particular problems related to SCC of structural steels in nitrate solutions are presented in subsequent parts of this chapter.

2.5 Investigation of stress corrosion cracking

The major aspects of SCC testing are covered in reference [37]. Studies with low-carbon low-strength steels are, as a rule, carried out only on smooth speci-mens, whereas studies of high-strength steels are performed either on smooth specimens or on specimens with neck or a microcrack. Stress is imposed upon a specimen by a constant strain, a constant load, or by straining at a constant rate.

In the constant strain test specimens are stressed in grips and placed in a vessel containing an aggressive medium. The test is inexpensive but reproducibility is poor. Stress relaxation occurs during crack development, resulting in a longer time to failure.

In the constant load test, stress increases as cracks grow. The elongation of the specimen is measured, from which the induction time and crack development time can be determined.

The constant strain rate test [38] is well suited for comparative testing, but because of the severe stress conditions it is not appropriate for predicting the corrosion behaviour of a material under service conditions. SCC occurs only in a definite range of strain rates which depends on the particular system. For ferritic steels in nitrate solutions SCC appears clearly at strain rates of 10~

4 to

10"6 s"

1 [37-40]. The effect of strain rate on the ultimate tensile strength in

boiling 4 M C a ( N 0 3) 2 and on the Rm and Re values in paraffin oil is illustrated in Fig. 2.1.

The constant strain rate test is inherently positive, since it is carried out till the failure of the specimen. As a consequence of the severe conditions of the test, transgranular cracks are formed as well as intergranular cracks [17].

TA

BL

E 2

.1

ME

DIA

TH

AT

PR

OM

OT

E S

TR

ES

S C

OR

RO

SIO

N C

RA

CK

ING

OF

ST

RU

CT

UR

AL

ST

EE

LS

Typ

e of

st

eel

Med

ium

C

ond

itio

ns

for

SC

C

Cra

ckin

g p

ath

s R

efer

ence

s

Aq

ueo

us

nit

rate

s N

H4

NO

3,

Ca

(N0

3) 2, N

aN

03,

pH

3-8

, p

oten

-ti

als

app

rox.

-2

00

to 4

-160

0 m

VNH

E;

hig

h co

r-ro

sivi

ty

of c

one,

so

luti

ons

(20-

60%

) an

d h

ot

solu

tion

s (a

t b

oili

ng

poi

nt)

Inte

rgra

nu

lar;

fer

rite

gra

in b

oun

d-

arie

s, a

nd

in q

uen

ched

ste

els

the

form

er

bou

nd

arie

s of

au

sten

ite

grai

ns

3,4

2

1-2

4

Aq

ueo

us

Na

OH

C

once

ntr

atio

ns

abov

e 5%

at

373

K,

up

to 5

0%

at

323

K,

02

in m

inor

am

oun

ts,

pot

enti

als

app

rox.

—

850

to

—55

0 m

VNH

E

and

+5

00

to

700

mV

NH

E

As

abov

e,

and

tran

sgra

nu

lar

at

hig

h te

mp

erat

ure

s 6

,7

17,

22-2

6

Fer

riti

c st

eels

L

iqu

id N

H3

An

hyd

rou

s N

H3 w

ith

trac

e 0

2, p

oten

tial

s ab

ove

-45

0 m

VN

HE

, 0.

2-%

ad

dit

ion

of H

20 in

hib

its

SC

C

Inte

r- o

r tr

ansg

ran

ula

r 24

, 27

Aq

ueu

os

CO

|-/H

CO

r so

luti

ons

1 M

(NH

4) 2C0

3, 0.

5 N

N

a2C

03

+

0.5

Ν

Na

HC

03,

pH

8-1

0,

343-

368

K,

pot

enti

als

app

rox.

-5

00

to

-35

0m

VN

HE

Inte

r- o

r tr

ansg

ran

ula

r 17

, 24

, 28

H20

—C

O—

C0

2 T

emp

erat

ure

293

-373

K,

pH

6

Tra

nsg

ran

ula

r 19

, 29

Aq

ueo

us

ph

osp

hat

es

1 M

Na

H2P

04,

pH

4.8

; 29

5 or

373

K,

pot

en-

tial

s at

373

Κ

-40

0 to

-2

50

and

-13

0 to

—

80

mV

NH

E,

ten

sile

str

ess

app

lied

at

a ra

te o

f 1

0"

6 s"

1

Tra

nsg

ran

ula

r 30

Aq

ueo

us

acet

ates

Sa

tura

ted

stea

m o

n th

e lo

w-p

ress

ure

sid

e of

a

stea

m t

urbi

ne, 0

.1-1

.0 M

CH

3CO

ON

H4,

pH

8,

363

K,

pot

enti

als

from

-3

60

to -

20

0 m

VN

HE

Inte

rgra

nula

r 20

, 31

Aq

ueo

us

HC

N

Pur

e so

luti

ons

or

con

tain

ing

NH

3, H

2S,

CO

, C

02

(gas

liq

uor)

T

rans

gran

ular

32

Aq

ueo

us

FeC

l 3 ΙΟ

-MO

"

3 MF

eCl 3,

tem

p.

583

Κ

Inte

r- a

nd

tran

sgra

nula

r 33

Aq

ueo

us

eth

anol

amin

e so

luti

ons

15%

eth

anol

amin

e in

w

ater

co

nta

inin

g H

2S an

d C

02,

tem

p,

up

to 4

23 Κ

U

nre

por

ted

34

Aq

ueo

us

Na

3A1

03

Na

3A1

03

in

wat

er

wit

h b

auxi

te

and

lim

e, a

t 41

8 Κ

U

nre

por

ted

34

AIC

I3

+ S

bC

l 3 in

hyd

roca

rbon

s C

atal

yst—

10%

A1C

1 3 +

90%

Sb

Cl 3

in

hyd

ro-

carb

ons,

363

Κ

Inte

rgra

nula

r 35

Mar

tens

itic

st

eels

W

ater

, hu

mid

air

, aci

ds,

sa

lts,

alk

alis

3%

NaC

l, 3

% N

a2S

04,

20%

H2S

04,

room

ter

n-pe

ratu

re o

r ab

ove,

cat

hod

ic o

r an

odic

pol

ariz

a-ti

on;

hot

NH

4N0

3 or

Ca

(N0

3) 2 so

luti

ons

Pre

viou

s b

oun

dar

ies

of a

ust

enit

e gr

ains

or

inte

rgra

nula

r 36


m

100

0 10"

. -1

Fig. 2 .1 . Effect of strain rate of a 0.07% C steel on fracture stress, Rmax, in boiling 4 M C a ( N 0 3) 2 and on plastic yield Rc and UTS, Rm, in paraffin oil at the same temperature of 115°C (after [37]).

High-strength steels are frequently investigated by the fracture mechanics technique. Specimens are used which have an initial microcrack produced at the notch tip. The field stress at the tip of a growing crack is determined by the parameter K, which is called the stress intensity factor. For an infinite plate this factor is given by:

where a is the applied stress and a is crack length. In measurements by this method Kscc is usually determined, i.e. the Κ value in a given environment be-low which failure does not occur, and the dependence of crack growth rate on the Κ value. Application of fracture mechanics to SCC studies has been described by Brown [41].

The main information obtained in measurements with smooth specimens is the time to failure at a given stress. The simplest tests are carried out in boiling solutions without external polarization and without measurement of elonga-tion. The tests are most frequently performed under potentiostatic conditions, at a controlled temperature, in solutions of a specific composition and pH, with simultaneous measurement of strain or load and monitoring of the polarization current.

For a better description of the electrochemical conditions involved in SCC, surface films are studied (composition, structure, thickness, protective and physi-cal properties, and behaviour in the presence of stresses), and also corrosion cracks (composition and pH of solution, potential). Recent studies on the film

a

CORROSION OF IRON IN NITRATE SOLUTIONS 57

composition and thickness have been carried out by Auger-electron spectro-scopy, ESCA and ellipsometry.

Metallographic methods are used to study the crack nucleation sites and crack growth paths and rate. Transmission [39, 40, 42] and scanning electron micro-scopy [24] are employed in addition to optical microscopy.

In studies of SCC in nitrates, most frequently the N H 4N 0 3, Ca (N03)2 or N a N 0 3 solutions of the following concentrations are used:

(a) 4 M (28.6%) N H 4N 0 3 or 5 M (34.8% N H 4N 0 3, pH 5, b.p. 378 K; (b) 30, 55, or 60% C a ( N 0 3) 2; 55% C a ( N 0 3) 2, b.p. 388 K; (c) 60% Ca(N03)2 · 4 H 20 + 3% N H 4N 0 3, b.p. 383 K; (d) 4 M N a N 0 3. Equipment for the application of stress to specimens is described in reference

[43]. Some cells and experimental set-ups are presented in references [37, 44-47].

2.6 Corrosion of iron in nitrate solutions

During the spontaneous corrosion of iron in nitrate solutions the pH of the solutions increases [48-50], and in hot solutions the electrode potential of iron

298 Κ

J I I I I I KrA I 1 U

0 1 2 3 20 60180140 Corrosion time, h

Fig. 2.2. Spontaneous changes of potential and pH during corrosion of Armco iron in

5 MNH4NO3 at various temperatures, without stress (after [50]).


rises [49, 50] (Fig. 2.2). The potential rise is a consequence of the formation of oxide film, and the pH increase is a result of the reduction of nitrates and oxygen.

The overall chemical reaction of iron in nitrate solutions is given by Engell and Baumell [51] as follows:

lOFe + 6NOJ + 3 H 20 -> 5 F e 20 3 + 60H~ + 3 N 2 Smiatowski and Szklarska-Smialowska [48] give the following scheme for the main reaction:

2 H+ + Fe + N O j -* F e

2+ + H 20 + N O j

The F e2+

ions may be oxidized to F e3+

by N O j , NOj or oxygen from the air, in a process that is accompanied by the formation of OH" ions. Thus con-ditions ensue for the precipitation of iron hydroxides, the formation of a passive film and passivation of the metal. Reduction of N O j ions results in the formation of OH- , N 2 and N H 3 [48, 52]:

2NOJ + 6 H 20 + 10e -> N 2 + 120H"

N O j + 6 H 20 + 8e N H 3 + 90H~

The main products of the cathodic reaction are N O j , OH"" and N H 3, where-as the main products of the anodic reaction may be F e

2+ or F e

3+ ions, sol-

-0.1

-0.2 CM

Ο

-0.3

£ -0.1

-0.2

-0.3

\ 6 min Ρ L' ' \ ' '|3*βκ|

\ 80min _ ~

F e ++

F e 2°

3

/ \ F e 30 4

1 j j j •τ - I Κ J

.1323KJ - Χ ρ θΛ .

F e ++

f12h

|298K| \

_ 6 N\ F e 20 3

F e

++ Ν Γ ^ ^ / '

\ Fe304

\ Λ -1 7 . . i— l \ _j ι ι 1

|298K| \

_ 6 N\ F e 20 3

F e

++ Ν Γ ^ ^ / '

\ Fe304

\ Λ -1 7 . . i— l \

pH(at298K)

Fig. 2.3. The potential vs. pH curves for Armco iron in 5 M NH4NO3 at temperatures 298 to 373 Κ on the F e - H 20 diagrams. The arrowed figures denote the time for the attainment of the F e 30 3 region. Symbols (Δ , x, etc.) denote the values measured in SCC tests (after [50]).

CORROSION OF IRON IN NITRATE SOLUTIONS 59

uble complexes of F e2+

with N H 3 (probably Fe(NH3) l+ [53]) as well as iron

hydroxides, oxyhydroxides and oxides. The kinetics and the reaction mechanism of the electrode reactions on iron

in nitrate solutions were investigated by Moshtev and Hristova [54]. The type of corrosion products on iron depends to a considerable extent on

temperature. Protective oxide films of F e 30 4 (magnetite) or F e 20 3 form rapidly only in hot solutions [50, 52, 55], whereas at room temperature discontinuous F e 30 4 deposits are formed [50].

Môssbauer studies [56] have shown that y-FeOOH is the major product of hydrolytic precipitation at pH above 5.6, while at pH below 3.0 only a-FeOOH is formed. Magnetite is always present in the vicinity of the metal surface. a -Fe203 is the major product of hydrolytic precipitation at temperatures above 353 K.

The nature of the corrosion products formed, especially the oxides, may be deduced from the potential and pH values established in the system. In Fig. 2.3

10V

1QJ ι I ι I ι I I l ι l ι I l I ι I I I ι l ι I -0.4 -0.2 0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 1.8

Potential ,VN HE Fig. 2.4. Potentiodynamic curves of anodic polarization of pure decarburized iron in flowing 5 M NttjNOa at temperatures 298 to 373 K; dE/dt = 2 V/h (after [59]).


the values of average potential and pH derived from Fig. 2.2 are shown as solid lines for definite times of corrosion at temperatures of 298 to 373 K. The lines are drawn on the potential-pH equilibrium diagrams for the F e - H 20 system, in which regions of thermodynamic stability of F e

2 +, F e 30 4 and F e 20 3 are

marked. It may be assumed that if the potential-pH line passes through the area of stability of F e 20 3, it indicates that a passive film is present, essential compo-nent being F e 20 3 [57, 58]. The detection of a thin film of y - F e 20 3 in the pres-ence of F e 30 4 by X-ray diffraction is practically impossible.

The potential-pH lines in Fig. 2.3 indicate that F e 30 4 and especially F e 20 3 form spontaneously on iron surface only at elevated temperatures.

The anodic polarization curves for pure decarburized iron in flowing 5 Μ M H 4N 0 3 at temperatures of 298-373 Κ are shown in Fig. 2.4, whereas the cathodic polarization curves for this iron, for Fe-0.64% C and for Fe3C together with dissolution rates at 348 Κ are illustrated in Fig. 2.5. It follows from Fig.

-1.2 -1.0 -0.8 ).8 -0.6 Potential,VN HE

-0.2

Fig. 2.5. Potentiodynamic curves of cathodic polarization of pure decarburized iron, marten-site Fe-0.65% C, and cementite Fe3C and curves of dissolution rate of iron and martensite in a flowing 5 Μ N H 4N 0 3 at 348 Κ (after [59]).

FORMATION OF CORROSION CRACKS 61

2.5 that iron also undergoes dissolution under cathodic polarization, which was established earlier by Smialowski and Szklarska-Smialowska [48].

On increasing the pH from 5 to 8, the corrosion potential and passivation potential of iron in 5 Μ N H 4N 0 3 shift towards negative values, and the corro-sion current and critical passivation current decrease [48, 60].

Addition of chlorides has an insignificant effect on the dissolution rate of iron in the active state but it makes passivation difficult and it strongly increases anodic dissolution in the passive region [49, 61].

Studies made on iron single crystals showed significant differences in the electrochemical behaviour of the {100}, {110}, and {111} crystal planes of iron in 5 Μ N H 4N 0 3 at room temperature over a wide range of potentials [62, 63].

2.7 Formation of corrosion cracks

In slowly cooled steels of a ferritic structure, cracks develop along the bound-aries of ferrite grains, whereas in steels of a martenistic structure quickly cooled (quenched) from austenitization temperatures, cracks grow along the previously existing austenite grain boundaries and, in some cases, along the boundaries of martensite laths [24, 40, 64].

The formation and development of corrosion cracks is closely related to the preferential corrosion of grain boundaries, which occurs also in the absence of stresses and as a consequence of the increased reactitity of these boundaries. An increased reactivity may be due to their physical and chemical hetero-geneity.

Among other factors, grain boundaries differ from the grain bulk, by their increased free energy (for austenite it is approx. 0.85 J/m

2 at 1373 Κ [65]), and

possibly also by the energy of elastic strain at the precipitated phases or dislo-cation accumulations. They differ chemically from the grain bulk because of the segregation of impurities and the precipitation of second phases. Traces of the original austenite grain boundaries in quenched steels are not the grain bound-aries in the physical sense, since they represent only a chemical heterogeneity of those grain boundaries.

Investigations of high-purity iron indicate that pure grain boundaries may undergo preferential corrosion [39, 63]. The reactivity of grain boundaries re-sulting from physical heterogeneity is, however, much lower than the reactivity resulting from chemical heterogeneity. Figure 2.6 shows a transmission electron micrograph of iron quickly cooled from 1253 K, which was then subjected to


Fig. 2.6. Transmission electron microphotograph of rapidly cooled high-purity iron after corro-sion in 5 M NH4NO3 at 348 Κ for 30 s (after [38]): A—dissolution along pure grain boundary in α-Fe; Β—dissolution along prior grain boundary in y-Fe, at which there remained segre-gated impurities.

corrosion in 5 Μ N H 4N 0 3. The narrow corroded path coincides with a new grain boundary in a-Fe, while the broad path corresponds to a trace of the original grain boundary in y-Fe of which there remained only segregated impurities. It will be seen that preferential corrosion does occur along pure grain boundaries, but it is clearly less pronounced than that along segregated impurities.

In metallurgical products the role of the physical factor of grain boundaries is certainly negligible. Intergranular corrosion of these materials is determined by segregated components and impurities as well as by precipitated phases.

In ferritic steels the formation and development of corrosion cracks is pri-marily due to carbon and nitrogen accumulated at grain boundaries. In mar-tensitic steels, on the other hand, this process is most probably due to substitu-tional components and impurities, whose atoms take up positions of iron atoms in the crystal lattice.

The accumulation of carbon at the ferrite grain boundaries is due to the segre-gation of carbon atoms or precipitation of carbides at these boundaries. Some authors ascribe intergranular corrosion and SCC of ferritic steel to precipi-tates of spherical cementite Fe3C [66, 67]. The precipitates may be responsible

CRACK GROWTH RATE 63

for localized corrosion of grain boundaries in two ways: in the active state they cause preferential corrosion of the adjacent ferrite, while in the passive state they dissolve faster than the ferrite [59, 68].

Electrochemical conditions affect not only the action of carbides but also of other components. In ferritic steels, carbon exerts a detrimental effect in the passive state [59], whereas in the active state the harmful influence is primarily due to segregated substitutional impurities [40].

The problem of the increased reactivity of grain boundaries is one of the major research directions in the study of SCC in carbon steels; the solution of this problem will considerably advance the understanding of the mechanism of SCC in these steels.

2,8 Crack growth rate

During the SCC process, there is a distinction to be made between the induction period and the development period. In the former, macroscopic cracks start to develop, whereas in the latter, the cracks propagate continuously. These periods can be distinguished by the measurement of acoustic wave propagation, measure-ment of electrical resistance, determination of relaxation of the load under-

Time, h

Fig. 2.7. Maximum depth of cracks in an annealed 0.08% C steel in boilling 4 M C a ( N 0 3) 2 + + 0.1 kg NH4NO3/I as a function of time in a constant-strain test (after [38]).

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71


ENVIRONMENTAL FACTORS IN STRESS CORROSION CRACKING 65

constant strain [37] or, most frequently, by the measurement of elongation as a function of time under a constant load [46, 49, 50].

In carbon steels the crack growth rate in the development period is a straight-

forward function of stress [69] :

-~ = k[a(t)-a0]

where / is crack length, t is the time, a{t) is the true stress at time t, and aQ is the threshold stress of crack propagation.

The crack growth rate additionally depends on the electrode potential (it increases exponentially with potential [46]) and on a number of environmental and structural factors.

Figure 2.7 shows the experimental results from measurements of maximum crack depth in ferritic steel as a function of time in a test under a constant strain [38]. The crack growth rate was decreasing up to the moment when open cracks began to form. A reduced crack growth rate in this test is related to decreasing stress. In constant load tests the stress rises as the cracks propagate. As a result, the crack growth rate also increases [69].

The values of crack growth rates for various materials and conditions are presented in Table 2.2. The values range from 0.04 to 2.58 mm/h.

2.9 Environmental factors in stress corrosion cracking

2.9.1 Determination of the active-passive state

Stress corrosion cracking generally occurs under electrochemical conditions in which the active and passive states can coexist on metal surface. Under these circumstances the environment manifests a moderate ability to passivate the metal.

Passivation ability can be determined from measurements of the kinetics of growth of oxide films or from the kinetics of the decay of the anodic dissolu-tion current. The initial surface of specimens should be free of oxide films for these measurements.

A simple procedure for the evaluation of the active-passive state has been proposed by Humphries and Parkins [72]. It consists in taking potentiodynamic curves of anodic polarization at a high and a low potential sweep rate, for example at 1000 and 1 mV/min. At a high sweep rate the potential range of an increased anodic activity extends farther than at a low sweep rate, since the formation of a protective film is not fast enough. A large difference between the anodic cur-


rents for the two potential sweep rates indicates a possible active-passive state, and hence conditions for the occurrence of SCC. The potential at which SCC occurs in nitrates [72], carbonates [28], acetates [31] and alkalis [23] are all within the ranges of the potentials of the active-passive state determined by the above procedure.

Major environmental factors which are responsible for the aggressiveness of an environment as regards SCC are: composition of solution, temperature, potential and the types of surface layers.

2.9.2 Effect of the composition of the solution

Studies have been made on the effect of cations, solution concentration, pH value, and additives to the solution.

Investigations of the aggressiveness of nitrates with various cations [73-78] showed that the most active species in inducing SCC are NH4 ions. According to Parkins and Usher [77], the aggressiveness of cations decreases in the order: Ν Η ί > C a

2+ > L i

+ > K

+, N a

+. This is shown by the values of threshold

stresses presented in Table 2.3. The values also indicate that the aggressiveness of solutions rises as the concentration increases from 1 M to 8 M.

TABLE 2.3

THRESHOLD STRESS VALUES (MPa) FOR A 0.05% C STEEL IN BOILING NITRATE SOLUTIONS OF VARIOUS CONCENTRATIONS (after [7/])

Nitrate Concentration

Nitrate 8 Ν 4 Ν 2.5 Ν 1 Ν

N H 4N 0 3 C a ( N 03)2 L i N 0 3 K N 0 3 N a N 0 3

15.4 38.6 (4M) 38.6 46.3 61.8

23.2 54.0 (2M) 61.8 69.5

146.7

54.0 92.6 (1.25M)

146.7 (2M) 108.0 169.8

92.6 177.6 (0.5Μ) 177.6 185.3 201.0

This effect, however, changes at higher concentrations. Above a concentra-tion of approx. 7.5 M9 the aggressiveness of an ammonium nitrate solution decreases, especially for materials with a lower susceptibility to cracking (Fig. 2.8). It will be seen from Fig. 2.8 that in accordance with the sequence of cations


3 4 8 Κ

Decarburized Fe

Fe-0.009%C

2.5 5.0 7.5 10.0 Molarity of NH4N03 solutions

12.5

km

2.5M [Ca(N03)JCc

X X D •

1.5M

WCaCb

Fig. 2.8. The time to failure of iron of various carbon contents in a range of N H 4N 0 3 solutions and in C a ( N 0 3) 2 solutions with and without chlorides, initial pH = 5.0; temperature 348 K; a = 0.9R, (after [49]).

given above, a Ca (N03)2 solution is less aggressive than an N H 4N 0 3 solution of the same molarity.

The degree of aggressiveness of particular nitrates has been related to their pH values [77]. It seems, however, that in this case the pH value is not the de-termining factor, as differences in aggressiveness occur at equal pH values [46, 49, 79]. The cation effect is possibly associated with incorporation of the cation in the oxide films and, as a result, with a change in protective properties of these films.

It is not easy to determine the effect of pH unambiguously, because the pH changes during the corrosion of steel; in acid solutions the pH increases, whereas in alkaline solutions it decreases during the course of corrosion [23, 48-50, 80]. The magnitude of the pH change depends, among other factors, on the ratio of the volume of solution to the quantity of iron dissolved, and it may reach several pH units.

The time to failure vs. pH value before and after a test is shown for two ma-terials and two solutions in Figs. 2.9 and 2.10. These figures and the results


Fig. 2.9. Effect of pH on the time to failure of unkilied 0.13% C steel in 5 M N H 4N 0 3, at 369 Κ under stress of 0 .7Rm. Arrows indicate changes of pH during the test (after [52]).

Time to failure, h

Fig. 2.10. Effect of pH on the time to failure of killed 0.13% C steel in boiling 9% C a ( N 0 3) 2 under a stress of 0.7Ro (after [80]).


reported in [23, 77, 78] indicate that the time to failure increases substantially at a pH value above approx. 3. Figure 2.10 shows that the aggressiveness of the solution also decreases at a pH below approx. 5.5. A similar decrease was ob-served in a 2 Ca (N03)2 solution at pH 4.0-4.5 [23]. Z^bik and Krukiewicz [78] report that in a constant strain rate test, cracks develop readily in a boiling 4 M N a N 0 3 solution at an initial pH of 1-2, less readily at pH 5-8, and do not develop at all at a pH 3-4.

The pH values at the corroded surface, and especially inside cracks, may differ from those in the bulk of solution. Some authors [23, 77] suppose that the pH value within the cracks is markedly lower, but there is no experimental evidence for this.

The aggressiveness of nitrates is substantially affected by other substances. Oxidizing additives, such as K M n 0 4, MnS04, N a N 0 2 or K 2C r 20 7 promote cracking, whereas additives such as N a 2C 0 3, N a H C 0 3, N a 2H P 0 4, dilute H 3P 0 4 or CO(NH2)2, which favour the formation of insoluble products on the iron surface, inhibit cracking [58, 73, 77, 80, 81].

Addition of chlorides suppresses cracking [55, 76, 82-84] and changes the type of corrosion. Lees and Lockington [55] observed pitting corrosion in a chloride-containing Ca (N03)2 solution, and intense general corrosion in a chloride-containing N H 4N 0 3 solution.

Bromides extend the time to failure to a lesser degree than chlorides, and fluorides have no effect [82].

233 Effect of temperature

SCC of iron and carbon steels occurs most readily in boiling solutions. At room temperature under open circuit conditions (without external polarization) these metals show a considerable resistance [23, 74], nevertheless cracking is possible [50, 80, 85]. SCC can also readily occur at room temperature under conditions of anodic polarization [50, 86].

Over a certain temperature range the logarithm of the time to failure is a linear function of the reciprocal of the absolute temperature:

\ogtf = ki+B/T

where tf is the time to failure, and kx, Β are constants. If the cracking rate is expressed in terms of \/tf, the above relation can be

expressed in the form of the Arrhenius equation:


where A is the apparent activation energy of the cracking process and k2 is a constant.

373

Temperature, Κ 348 3 2 3 298

-10

1

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10

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10

Reciprocal temperature, 1000/Κ

Fig. 2.11 Effect of temperature on time to failure and induction time for a 0.009% Armco iron in 5 M NH4NO3 of initial pH 5.0 and stress 0.92?e (after [50]). The numbers denote the values of apparent activation energy (kJ/mol).

The temperature dependence of the time to failure under open circuit condi-tions is shown in Fig. 2.11. In the range from boiling point down to room tem-perature, the slope of the \ogtf vs. T"

1 line varies, primarily due to the forma-

tion of different oxide films [50]. At potentials corresponding to specific films, the log tf vs. 7

1 -1 is of different

character (Fig. 2.12). In the presence of a passive film the apparent activation energy for Armco iron is only 34.8 kJ/mol and it is less than one third of the acti-vation energy under open-circuit conditions at high temperatures (cf. Fig. 2.11).

Values of the apparent activation energy reported in the literature range approx. 35 to 134 kJ/mol [50, 52, 80, 83, 84, 87]. Values of this energy for vari-ous materials and conditions are given in Table 2.4.

The majority of the activation energy values are in the range 35-70 kJ/mol; this range comprises the activation energies of common chemical and corrosion processes. The lower values of activation energy of SCC are roughly equal to


Temperature, Κ

1 ι ι ι ι 1 1 1 •

2.7 2.8 2.9 3.0 3.1 3.2 3.3 3.4 Reciprocal temperature, 1000/K

Fig. 2.12. Effect of temperature on time to failure, t ti for a 0.009% Armco iron in 5 M N H 4N 0 3 at a pH of 7.0 under a stress of 0.9RC: (a) at the potential of minimum u in the presence of a passive film (Emmtt, F e 20 3) ; (b) at the potential of the second peak on the anodic polarization curve (largely connected with the presence of carbon) (Enmax a.p.); (c) at the potential of maximum U in the presence of F e 30 4 (EmAxU, F e 30 4) (after [50]).

the activation energy for the corrosion of iron in the active state, 38.8 kJ/mol, in a 5 Μ N H 4N 0 3 solution [59].

In the passive state the activation energy of the cracking process (Fig. 2.12) is close to the activation energy of corrosion in the active state. This indicates that crack development occurs as a result of electrochemical dissolution along the sites poorly protected by the passive film. The reason for a decreased pro-tective ability of the passive film along the ferrite grain boundaries is discussed in Section 2.10.1.

Mean values of activation energy (approx. 50-70 kJ/mol) may be connected with diffusion of the electrolyte components through oxides or other corrosion products, and high values could be associated with the formation of the passive film.

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ο ο ENVIRONMENTAL FACTORS IN STRESS CORROSION CRACKING 71


2.9.4 Effect of potential

Cathodic polarization inhibits SCC of structural steels in nitrate solutions, whereas anodic polarization promotes it [49-52, 67, 72, 84, 88-91].

The potential range for SCC in nitrates, compared with other environments, is wide. In hot N H 4N 0 3 solutions it may range from approx. —300 to approx. + 1700 m V N HE [49, 52, 72, 91]. At other potentials SCC is either not observed or it proceeds at a negligible rate. At potentials below approx. —400 m V N HE in hot NH4NO3 solutions the failure of metals is chiefly due to general corro-sion [52, 75, 49].

Figure 2.13 shows a typical dependence of the time to failure on potential

- 0 . 4 \rr 0 C U 0 8 U 1 6

Potential, VN HE Fig. 2.13. Effect of potential on time to failure of a 0.13% C steel in 5 MNH4NO3 at 348 Κ under stress of 0.7Rm (after [52]): Δ —intergranular failure, O—failure caused by thinning as a result of general corrosion.

in NH4NO3 (after Szklarska-Smialowska [52]). An analogous dependence has been observed by others [49, 50, 72], whereas Bombara [91] reports that in a boiling 50% N H 4N 0 3 solution over a potential range from -150 to +1700 m V N H E, there is a considerable variation in the susceptibility of steel to SCC. Bombara observed a substantial decrease in susceptibility to SCC in this range at poten-tials from +250 to +400 m V N H E-

As the temperature decreases, the time to failure at a given potential increases, but at the same time the range of potentials at which SCC occurs shifts towards negative values [49, 50, 84].


The replacement of N H 4 by C a2+

in nitrates results in a decrease of the potential range of SCC. In concentrated Ca (N02)2 solutions, cracking begins only when the potential exceeds approx. —100 m V N HE [84, 89],

The acceleration of the SCC process by anodic polarization is explained by Mazille and Uhlig [84] in terms of a more intense adsorption of aggressive anions at the tip of a corrosion crack, which would favour the mechanical rupture of the metal. More probably, however, it is associated primarily with the active-passive state that occurs at these potentials. The appearance of this state de-pends on the type of oxide films formed.

2.9.5 Role of the oxide films

The insoluble products of corrosion of iron in nitrate solutions consist mainly of iron oxides. Using X-ray analysis or electron diffraction, the following oxides were detected among the corrosion products in nitrate solutions: F e 20 3 · H 20 or F e 20 3 + F e 30 4 [52], F e 30 4 [59,72] and a -Fe203 [72]. The oxide y - F e 20 3, which is responsible for the passive state of iron, could not be identified by the diffraction method because of the great similarity of diffraction patterns of F e 30 4 and y - F e 20 3 and because of the small thickness of the passive film of y - F e 20 3 (approx. 2-3 nm [92]); nor can these oxides be distinguished by ellip-sometry.

Potential and pH measurements provide an indirect method for an assessment of the presence of any of these oxides (Section 2.6) [46, 49, 50].

The significance of oxide films in SCC of iron and steel in hot nitrate solu-tions has been emphasized by many investigators [46, 48-52, 69, 70, 72, 76, 79, 82, 88, 93, 94]. The oxide playing the major role in SCC is not usually specified. Humphries and Parkins [72] have reported that cracking occurs in nitrates in the presence of F e 30 4 and F e 20 3, and particularly in the presence of an a - F e 20 3 film. The authors suggest that the role of oxide films is limited mainly to blocking the cracks and suppressing their development, so that the SCC process is not dependent on the nature of the oxide.

On the other hand, other studies [46, 49, 50, 79] show that the rate of SCC strongly depends on the nature of the films produced. It has been shown that the rate may be very high in the presence of a ferric oxide passive film, but it is negli-gible in the presence of magnetite F e 30 4.

In hot solutions under open circuit conditions, cracking occurs only after the attainment of the potential and pH values corresponding to the formation


of an F e 20 3 film. At the potential and pH values corresponding to the formation of F e 30 4 the rate of SCC is very slow.

The results of measurements carried out at a controlled potential and pH, corresponding to the thermodynamic stability of F e

2 +, F e 30 4 and F e 20 3, are

shown in Fig. 2.14. They indicate that in the presence of an F e 20 3 film, the time

Potential,^

Fig. 2.14. The potential dependence of: 1) time to failure tt for a 0.009% C Fe in 5 M NH4NO3 at pH 7.0 at 348 Κ under a stress of 0.9 Rc; 2) electric charge of a total anodic dissolution of iron during test, and 3) ratio / l n d/ / f Omd is induction time) (after [49]).

to failure may be three orders of magnitude shorter than in the presence of an F e 30 4 film.

From the available data on the effect of the potential it follows that, as a rule, the shortest times to failure occur under conditions of stability of the F e 20 3 film, whereas these times increase near the potential of the formation of F e 30 4 [52, 72, 84, 89, 91].

The dependence of the SCC rate on major environmental factors (solution composition and concentration, temperature, and potential) is to a large extent determined by the effect of these factors on the nature of the oxide film formed. Cracking is promoted by those factors which favour the formation of a y - F e 20 3 passive film, i.e. elevated temperature, anodic polarization, or passivating addi-

METALLURGICAL FACTORS IN STRESS CORROSION CRACKING 77

tions. Cracking is slow or may not occur at all if the conditions favour the forma-tion of magnetite F e 30 4, i.e. lower temperatures, cathodic polarization in the region of stability of F e 30 4, and additions that make the passivation difficult (e.g. chlorides, in the presence of which the passive film is not formed, but F e 30 4 is produced in large amounts) [49].

A considerable tendency to cracking is shown by commercial iron in the pres-ence of the passive film, whereas pure decarburized iron remains resistant to SCC [50]. The susceptibility of steel to SCC in the presence of a passive film is therefore closely connected with the composition of the steel and, in particular, with the presence of carbon in the material.

2*10 Metallurgical factors in stress corrosion cracking

2.10.1 Effect of steel composition

Pure iron may undergo cracking and intergranular corrosion in 5 Μ N H 4N 0 3 [39, 49, 63, 79], but this has not been observed in Ca (N03) 2 solutions [66, 71, 95-99]. However, the addition of carbon or nitrogen makes iron susceptible to SCC in the latter solution also [71]. It is now generally recognized that the susceptibility of structural steels to SCC is determined in the first place by the presence of carbon or nitrogen.

2.10.1.1 Effect of carbon

Of all the various components and impurities of steel that affect its susceptibility to SCC in nitrates, carbon has been studied most extensively [40, 59, 64, 66, 67, 71, 79, 90, 95-104].

The effect of carbon on susceptibility of steel to SCC depends strongly on its content. Ferritic steels show the highest susceptibility at a carbon content of approx. 0.01 to 0.1 wt. %. Outside this range the susceptibility decreases both with increasing and decreasing carbon content [66, 71, 79, 95-97, 99-103]. A high or even complete resistance is shown by carbon-free pure iron, as well as by ferritic and ferritic-pearlitic steels containing over approx. 0.25% C [64,95,97-99, 102, 104].

The effect of carbon also depends on heat treatment and on the steel struc-ture. In annealed steels in the absence of stresses, the intergranular corrosion initially increases, then it decreases with the carbon content (Fig. 2.15). A similar tendency is observed for the dependence of the extent of SCC on carbon content [77, 99]. This type of relationship is the most common one, and it can therefore be considered typical.


OA 02 0.3 ° Caption content, %

Fig. 2.15. The depth of intergranular corrosion of unstressed steels vs. carbon content, after a 300 h corrosion test in boiling 57% C a ( N 0 3) 2 + 3% N H 4N 0 3 at 383 Κ (after [99]).

J— 1 1 1 1 ι I 0 0.04 0.08 0.12 0.16 0.20 0.24

Carbon content. %

Fig. 2.16. The threshold stress for cracking of annealed and quenched steels with various carbon contents in a boiling solution of 0.57 kg/1 C a ( N 0 3) 2 + 0.03 kg/1 N H 4N 0 3 (after [64]): · — annealed, Δ —quenched.

Figure 2.16 shows a somewhat different relation. In the annealed material the corrosion resistance showed a gradual increase with carbon content, where-as in the quenched material it initially increased and then decreased. The differ-ent behaviour of the annealed and quenched steels is probably associated with different types of crack growth; in annealed steels, cracks develop along the ferrite grain boundaries on which carbon readily accumulates in the form of cementite or segregated atoms, whereas in quenched steels of a martensitic structure the cracks develop along the original austenite grain boundaries, on


which chiefly substitutional components are segregated. The cracking of high-strength martensitic steels containing 0.3-0.8% C also takes place along the latter type of boundaries [36, 86].

There are two theories of crack propagation and the influence of carbon: the adsorption theory, whose major proponent is Uhlig [105], and the electro-chemical dissolution theory.

The adsorption theory proposes that the development of a crack during SCC is a consequence of a lowered surface energy of the metal as a result of the ad-sorption or chemisorption of specific ions [106]. In nitrate solutions these may be NO3 ions or corrosion products [96]. The adsorption of such ions at the crack tip in a stressed material may facilitate the breaking of the atomic bonds, thereby favouring crack development. The adsorption may occur on surface defects, such as grain boundaries or dislocations. According to Uhlig et al. [71, 84, 96, 104] the detrimental effect of carbon and nitrogen consists in rendering the de-fects more stable and in forming sites that favour adsorption. Unambiguous experimental evidence in support of this theory, however, is lacking.

The electrochemical dissolution theory explains the development of cracks in terms of the anodic dissolution of the material along grain boundaries. Car-bon may favour this process by enhancing the electrochemical reactivity of the material of grain boundaries or by lowering the protective properties of the oxide film at the grain boundaries.

An increased electrochemical reactivity of the material at grain boundaries with accumulated carbon manifests itself in the active state mainly on account of the presence of cementite precipitates. The ferrite undergoes preferential corrosion around the cementite precipitates [40], probably due to the cathodic effect of the carbide [107].

Many authors connect susceptibility to SCC with the precipitation of car-bides [23, 66-68, 95, 101], According to the carbide theory, the higher densities of spherical cementite precipitates at the grain boundaries lead to a higher sus-ceptibility to cracking. An increased resistance to cracking at a high carbon content is explained by the presence of pearlite, and thus by the lower number of cementite precipitates at the grain boundaries [23, 66]. This reasoning is supported by the increased susceptibility of steel to cracking on annealing below the eutectoid point (996 K), which brings about spheroidization of carbides and their accumulation at grain boundaries [64].

In the passive state cementite corrodes faster than ferrite [59, 68]. Under these conditions, preferential corrosion would be expected on the precipitates rather than on the adjacent ferrite.


The carbide theory fails to account, however, for all cases of cracking. It does not explain the preferential corrosion at carbide-free grain boundaries [42] or the increased resistance to SCC with carbon content in the cases where the amount of carbides at the grain boundaries also increases [79]. In such cases the detrimental effect will be exerted primarily by the carbon atoms segregated on grain boundaries. The segregated carbon perhaps enhances the electrochem-ical reactivity of grain boundaries, but its most likely effect is the modification of the protective properties of the oxide films [79],

Reference [59] shows that carbon affects the rate of overall dissolution of iron in nitrate solutions in three ways, leading to activation, hindrance, and depassivation:

(a) the activating effect of carbon is demonstrated by an increased dissol-ution rate of iron in the active state. The activation energy only slightly in-creases with the carbon content;

(b) the hindering effect involves a decreased dissolution rate as a result of an enhanced deposition of magnetite and improvement of its protective prop-erties. On the surface of Fe-C alloys in 5 Μ Ν Η 4Ν Ό 3 more magnetite is for-med than on the surface of pure iron [59, 79], resulting in a much reduced corro-sion rate. This is shown by the decrease in corrosion rate constants with carbon content. In a flowing 5 Μ N H 40 3 solution at 373 Κ the parabolic rate constants for pure iron, Fe-0.64% C martensite and 1.4% C martensite are: 11 .2x l0

1 0,

2.1 x l O1 0

, and 1.3 χ 1010 m g ^ m

4^ (1120, 210, and 130 m g ^ c m

4^ ) , respect-

ively; (c) the depassivating effect is demonstrated by hindrance of the formation

of a passive film accompanied by impairing the protective properties of the film. Figure 2.17 shows that passivation of Fe-C alloys occurs much less readily than passivation of pure iron.

The hindering and depassivating effects are most probably due to the amor-phous carbon accumulated on the corroding steel surface. Since carbon in ferritic steels is accumulated at the grain boundaries, its effect will manifest itself pri-marily at these boundaries. If a passive film is present on the metal surface the depassivating effect will be largely exerted on the carbon-enriched grain bound-aries, whereas the ferrite grains with negligible carbon content will be pass-ivated in the same way as pure iron. This will result in a preferential anodic dissol-ution of the grain boundaries between the passivated ferrite grains. It is suggested that this effect of carbon explains the intergranular corrosion and SCC of ferri-tic steels in the passive state.

The fact that SCC occurs chiefly in the presence of a passive film discussed


10

4

ε 10

3

<

10'

-0.4

5M NH4NO3 ,348 Κ

/F e - 0 . 1 0 % C M

Jr L j î î " ~u o / oC

>n/Fe-0.64%C M

Ξ" /ffy/ W-O.OOO%C

1 1 1 1 1 I 1 .1 .. TP—Ί T I L L 0.4 0.8 1.2 1.6 -2.0

Potential, VN HE Fig. 2.17. Potentiodynamic curves of anodic polarization of quenched Fe-C alloys in 5 M NH4NO3 at 348 Κ (after [79]).

in Section 2.9.5, can be explained by the susceptibility of this film to the depassi-vating effect of carbon and of other constituents segregated at grain boundaries. This will limit the corrosion to the very narrow regions along the grain bound-aries.

SCC is slow in the presence of magnetite, probably because the magnetite layer is thicker than the passive film and thus less sensitive to inhomogeneities of the metal substrate. Therefore this layer protects grains and grain boundaries alike. Carbon improves the protective properties and in this way it can enhance the resistance to SCC in the presence of magnetite.

The increase followed by a decrease in the susceptibility of ferritic steel to in-tergranular corrosion and SCC with increasing carbon content may be accounted for by overlapping of the opposing depassivating and hindering effects, the latter becoming stronger with increasing carbon content [59, 79]. At a low carbon content, SCC is enhanced as a result of the prevailing depassivating effect; at a high carbon content, on the other hand, SCC becomes suppressed as a result of the counteraction of the hindering effect. The dependence of time to failure


on carbon content is perhaps due to the combination of these two effects, as represented schematically in Fig. 2,18.

The adverse action of carbon in the passive state can be explained by its depassivating effect only for the case of ferritic steels, in which grain boundaries

ff(dep) \

Carbon content

Fig. 2.18. A schematic representation of the effect of carbon on time to failure under condi-

tions corresponding to the depassivating effect, /f(dep), and to the hindering effect, /f(hin),

as well as on the resultant time, tt(res).

are enriched in carbon. This explanation, however, is not tenable for martensitic steels that crack along the original austenite grain boundaries, on which the car-bon segregation is negligible. Substitutional components, which may increase the electrochemical reactivity of these boundaries and/or reduce the protective properties of the passive film, seem to be critically important in this case [40].

2Λ0Λ.2 Effect of nitrogen

After carbon, the second most important element in determining the suscepti-bility of steel to SCC is nitrogen. However, its effect on SCC has not been studied in such detail as in the case of carbon.

The detrimental effect of nitrogen has been demonstrated by many authors [71, 99, 101, 103, 108, 109]. In quenched steels it is detrimental at contents above 0.014% Ν [108, 109]. The susceptibility of pure steel containing 0.043% Ν varies with the degree of cold work and heat treatment, as with pure carbon steel.


Nitrogen in moderate quantities enhances the susceptibility to SCC, whereas in large quantities it reduces the susceptibility [66, 110]. This behaviour may indicate an analogy between the effects of carbon and nitrogen in the SCC pro-cess.

2.10.1.3 Effect of substitutional constituents

The main studies of the effects of substitutional constituents have been made mainly with alloying additions. Their effect on susceptibility of steel to SCC is usually explained by their influence on the detrimental components, carbon and nitrogen, but it may also result from the effect on mechanical properties of steel, on the kinetics of dissolution, the formation of protective films, or on the corrosion potential [23].

In keeping with the concept of the influence on carbon and nitrogen, a bene-ficial effect should be expected from the carbide- and nitride-forming compo-nents because they tie up carbon or nitrogen. The graphitizing components, on the other hand, which favour segregation of carbon at grain boundaries, will exhibit an adverse effect.

Long and Lockington [98] report [111] the following sequence of major alloying components according to their carbide-forming or graphitizing power:

Generally, the chemical affinity of carbon and nitrogen to transition elements is the higher, the more to the left from iron in the Periodic Table a given element lies. Long and Lockington therefore suggest that a beneficial effect can be ex-pected from the components lying on the left from iron in the above sequence, and an adverse effect from those on the right from iron. This suggestion is borne out by experience in many cases.

The effect of particular elements is as follows: Titanium—beneficial effect [34, 74, 98, 101]: effective in concentrations above

approx. 0.4% [98, 101]; Molybdenium—beneficial effect [102,104]; effective in concentrations of approx.

0.5-1.0%;

Chromium—beneficial effect [97, 98, 104, 112, 113]; in annealed materials it is effective at contents above approx. 2% [102, 104];

Niobium and tantalum—beneficial effect [114]; Manganese—beneficial effect [97, 98, 102];

Increasing <- carbide-forming power Ti, Mo, W, Cr, Mn |Fe|

Increasing graphitizing power -» Ni, Zr, Cu, Al, Si


Nickel—ambivalent effect: it enhances susceptibility to SCC at contents up to approx. 2% [98, 104] and reduces it at higher contents [104];

Copper—at concentrations from 0.3 to 1.0% its effect is unfavourable [112]; Aluminium—the results are often contradictory, but in the majority of cases

they indicate a beneficial effect especially at concentrations above approx. 0.5% [74, 98, 101, 112, 114, 115];

Vanadium—it is beneficial [102] in some cases (up to 0.17%), and unfavourable in others [68];

Silicon—ambivalent effect; it enhances susceptibility to SCC at contents up to 0.5%, but reduces the susceptibility at higher concentrations [104];

Germanium—unfavourable effect [104]; Uranium—ambivalent effect: it reduces the susceptibility to SCC at contents

up to 0.15%, and it enhances it at higher concentrations [116]. Phosphorus—a highly adverse effect [117-122]; phosphorus segregated at the

grain boundaries promotes a strong intergranular corrosion of steel, both under stress [117, 120, 122] and without stress [119, 121];

Kiipper et al. [119] found that Fe-P alloys passivate less readily than iron; an Fe-P (2.5 wt. %) alloy failed to undergo passivation over the entire range of potentials studied, from the active to the transpassive region. This suggests that enhanced corrosion of grain boundaries with segregated phosphorus is connected with a more difficult passivation of the material at these grain bound-aries. The mechanism of the detrimental effect of phosphorus therefore is likely to be the same as that of carbon [59, 79].

In potential regions of intense SCC, segregation of phosphorus at grain boundaries does not affect the sensitivity of carbon steels to cracking [122]. Decrease of phosphorus content to low values cannot therefore guarantee re-sistance against SCC [122].

Sulphur—no effect of sulphur on intergranual cracking has been observed [123];

Tin—a study of 3.5 Ni-1.7 Cr-0.060 Sn steel showed no effect of the segre-gated tin on intergranular cracking [117].

2.10.2 Effect of heat treatment and structure

Low-carbon steels are generally used hot- or cold-worked, annealed or normal-ized (cooling down in the air from the austenitizing temperatures; ferrite or ferrite/pearlite structure), whereas high-strength steels with medium and high-carbon content are used after heat treatment involving quenching followed by


tempering (a tempered martensite structure). The effects of the major heat treat-ment factors on the susceptibility of steel to SCC are as follows:

Austenitizing temperature. With increasing austenitizing temperature the susceptibility to SCC rises [74, 82]. This is associated with increasing grain size and consequently, with a reduced yield point [90], and with an enhanced segre-gation of harmful components at the grain boundaries.

Cooling rate from the austenitizing temperature. The higher the cooling rate, the greater susceptibility to SCC. This susceptibility decreases as the cooling rate decreases in the order: cooling in water, in oil, in the air, and in a furnace [71, 74, 75, 124]. This may indicate that cracking along the former austenite grain boundaries (with substitutional constituents) occurs much more readily than along the ferrite grain boundaries. This, however, is observed only for steels containing carbon above a certain level (Fig. 2.16).

Temperature and time of tempering. The effect of these factors depends to a large extent on the carbon content of the steel. On the basis of the relation shown in Fig. 2.16. Parkins et al. [64] have suggested that the tempering process, which causes the metal to approach a state analogous to that after slow cooling, will raise the susceptibility to SCC of low-carbon steels (up to approx. 0.1% C), but it will reduce it in the case of high-carbon steels. This is observed only for certain tempering conditions, e.g. 1 h at 973 K; at higher temperatures a rise in the susceptibility is noted for high-carbon steels also. An increase in temper-ing time causes the changes in susceptibility to SCC to occur at lower tempera-tures [64].

Despite a considerable diversity of the conditions of the experiments carried out by individual authors (using various materials and times), there is a general trend in the variations in the susceptibility to SCC during tempering, as shown in Fig. 2.19. The trend is as follows:

(a) temperatures up to approx. 473 K: no changes [71, 74, 124] or an en-hanced susceptibility of high-carbon steels [64];

(b) temperatures from approx. 473 to 973 K: increased resistance [64, 71, 74, 104, 124, 125] (in Fig. 2.19 a peak in resistance occurs at 773-823 K);

(c) temperatures above approx. 773-973 K; a drop in resistance [71, 74,124].

During the tempering process a transformation of martensite into ferrite takes place, accompanied by the following phenomena: precipitation of highly disperse carbides, dispersion of the impurities segregated at the original auste-nite grain boundaries, recrystallization and growth of ferrite grains, and pre-cipitation of carbides at the ferrite grain boundaries [40, 42, 64, 124].


Tempering temperature, Κ

Fig. 2.19. Time to failure of a 0.11% C steel quenched in water and tempered for 15, 30 and 60 min at various temperatures: tests carried out in a boilling 40% N H 4N 0 3 (after [124]).

A change in susceptibility to SCC as a result of tempering at various tem-peratures is probably due to the following transformations at the active paths :

(a) at temperatures up to approx. 473 K, cracking occurs along the former austenite grain boundaries [24, 40, 64]; the reactivity of these paths is primarily due to the segregation of substitutional constituents [40];

(b) at temperatures from approx. 473 to 973 K, cracking occurs in low-car-bon steels (approx. 0.05% C) along the ferrite grain boundaries formed upon decomposition of martensite; and in higher-carbon steels (approx. 0.25% C),

EFFECT OF COLD WORK 87

along the former austenite grain boundaries with crack branching along the bound-aries of the martensite lath. The higher the carbon content, the longer the former grain boundaries remain as sites of preferential corrosion [64]. The in-crease in the resistance to SCC is most probably caused by dispersion of detri-mental substitutional impurities from the former boundaries. Parkins et al [64] suggest that a maximum resistance is determined by the competitiveness of reactivity of the original austenite grain boundaries and the new ferrite bound-aries;

(c) temperatures above 773-973 K: cracking occurs largely along the bound-aries of ferrite grains. An enhanced susceptibility to SCC may result from an increased precipitation of spheroidal cementite at the grain boundaries [67]. Another possible reason for the higher susceptibility may be a lower concentra-tion of segregated carbon atoms at grain boundaries, giving rise to a correspond-ingly lower hindering effect (Fig. 2.18).

Ageing of pearlite. Pearlite forms as a result of the slow cooling of steel from the austenitizing temperature. The susceptibility of pearlite to SCC is low, which is explained by the precipitation of cementite within the grains instead of at the grain boundaries. Annealing of pearlite at temperatures below the eutectoid point (1000 K) increases the susceptibility of steel to SCC probably as a result of spheroidization of cementite and its precipitation at the ferrite grain bound-aries [64, 66].

Effect of welding. In welded constructions, cracks form primarily in the heat-affected zone adjacent to the welded joint. Residual stresses are the major reason for these cracks. Removal of these stresses by tempering at about 923 Κ effec-tively reduces the susceptibility of welds to SCC [126]. Another possible reason for the susceptibility of welds is that their structure resembles that of quenched material.

2.11 Effect of cold work

The effect of cold work on the susceptibility of steel to SCC depends on the carbon content and the initial susceptibility of material.

Low-carbon steels with a high initial susceptibility become more resistant with increasing degree of cold work (by rolling or drawing) [38, 71, 82, 104, 127].

The increase in resistance to SCC is explained by the dispersion of carbon or nitrogen from the grain boundaries and a disorder of the structure, as a re-sult of which the continuity and regularity of grain boundaries disappears [71].


Resistance to SCC also increases as a result of hardening of the steel surface. Often this effect is explained in terms of favourable compressive stresses set up in a material as a result of this operation, but this may well be also related to a disturbance in regularity of grain boundaries [128]. In the case of high-carbon steels of a low initial susceptibility, the effect of a heavy cold work is the same as for low-carbon steels, that is, resistance to SCC increases. The effect of mod-erate cold working, however, is in the opposite direction; the resistance may decline [66, 110, 127].

Cold work may also affect general corrosion. In dilute N H 4N 0 3 solutions (0.5 and 0.6 M)9 plastic deformation does not affect the corrosion rate of steel in the active state, but it strongly increases corrosion in the passive region. At room temperature the rise in corrosion in the latter region is observed only for a carbon content above approx. 0.01% [68] and the increase is proportional to the degree of cold work [68, 129]. Since the effect of cold work on the passive state occurs only in the presence of carbon, it is suspected that the changes ob-served are in fact due to the depassivating effect of carbon (see Section 2.10.1.1).

2.12 Role of mechanical stresses

With increasing tensile stresses the time to failure, t{, decreases. Figure 2.20 shows the dependence of tf on relative stresses <rr related to ultimate tensile strength, ar = a/Rm. The data in this figure can be expressed in an exponential form of tf vs. σΓ:

Fig. 2.20 Time to failure of steel with various amounts of carbon vs. relative stress at — a/Rm; tests were carried out in a boiling 60% C a ( N 0 3) 2 solution (after [101]).

0

!"

1 1 » 1 M l l l l l l 1 t i i ι 1111 , , , , ,,,,1

10

2 10

3 10

4 10

5

Time to. failure tf . min .10

ROLE OF MECHANICAL STRESSES 89

tf = Acxp(—αστ)

or in a power form as t{ vs. absolute stresses a [22]

tf = Ba~b

where A, B, a, b are constants dependent on steel composition, and particularly on carbon content and heat treatment.

The power relation describes the behaviour of many systems that are known to undergo SCC. Azhogin [36] reports that the majority of systems obey the equation:

(tf-crt h)if = Κ

or, alternatively

where ath is the threshold stress, below which cracking does not occur, and Κ is a constant. The cracking of high-strength steels in an ammonium nitrate sol-ution follows this relation [130].

In some cases there are deviations from this relation. Z^bik [131] demon-strated that the t{ vs. tensile stress relation may be non-monotonous near the yield stress and that t{ may increase with stress.

In structural steels in nitrate solutions, intergranular corrosion occurs also in the absence of stresses, and hence stresses are not indispensable for the in-itiation of cracks. Incipient cracks form as a result of the corrosion processes themselves. These processes are also responsible for an increase in the potential and pH value, and for the formation of a passive film [46] whose presence fa-vours a fast SCC (Section 2.9.5).

In the absence of stresses intergranular corrosion stops after some time, re-commencing only if stresses are applied. The growth is then a consequence of the rupture of oxide films and the electrochemical dissolution of uncovered, bare metal. This is borne out by the fact that the relation between dissolution current and potential following film rupture, and the relation between crack growth rate and potential in the presence of stresses both show the same expo-nential character [46].

Crack development in structural steels proceeds by an electrochemical mech-anism, that is, the crack develops as a result of the preferred anodic dissol-ution of metal at the crack tip [38, 46, 70, 93]. The action of the stresses con-sists of widening the cracks, thus facilitating the entry of electrolyte, and of breaking the oxide films and deposits in the crack. This leads to the exposure of metal at the crack tip and a subsequent rapid corrosion.


2.13 Mechanism of stress corrosion cracking

Stress corrosion cracking of carbon steels in nitrate solutions is inseparably connected with its susceptibility to intergranular corrosion. This susceptibility results mainly from a chemical heterogeneity of the material of grain bound-aries. In quenched steels cracking occurs along the original austenite grain bound-aries, on which mainly substitutional segregated constituents remain; in an-nealed steels, on the other hand, cracking occurs along the ferrite grain bound-aries on which both substitutional and interstitial constituents are accumulated. Of the various interstitial constituents that exert an adverse effect, carbon is of primary significance. It accumulates at grain boundaries as segregated atoms and precipitated carbides.

In the active state and in the presence of magnetite, intergranular corrosion results from an enhanced electrochemical reactivity of grain boundaries, which may largely be ascribed to the segregated substitutional components and car-bides. In the passive state intergranular corrosion results primarily from the depassivating effect of segregated constituents. At ferrite grain boundaries this effect is exerted mainly by carbon; at the original austenite grain boundaries, on the other hand, this effect is possibly due to substitutional constituents.

During the initial period of the contact with nitrate solutions (especially NH4NO3) the metal is in the active state. The intergranular corrosion that de-velops in this period is caused primarily by segregated substitutional impurities and carbides. As a result of iron dissolution, the pH of the solution increases; oxide films and deposits form and carbon accumulates, while the potential shifts spontaneously towards positive values. An essential part in the cracking process is played by two oxide phases which correspond to magnetite F e 30 4 and to the passivating oxide y - F e 20 3.

Magnetite in the form of a film or deposit renders the cracking more difficult (a favourable effect), whereas the passive film of y - F e 20 3 may promote cracking (an unfavourable effect). The passive film promotes preferential corrosion at the ferrite grain boundaries, as it is sensitive to the depassivating effect of carbon and possibly of other constituents or impurities segregated at grain boundaries. Moreover, the film speeds up the crack growth in the presence of stresses, as it involves a considerable shift of the potential towards more positive values.

Because of the sensitivity of the y - F e 20 3 film to the depassivating effect of segregated constituents, it is the material of grain boundaries that will undergo dissolution rather than the walls of the crack. The walls will become pass-ivated more readily as they are not enriched by depassivating components. This

PREVENTION OF STRESS CORROSION CRACKING 91

may be one of the major reasons for a strictly intergranular path of cracks and their penetration, with a negligible lateral growth.

The detrimental effect of the y - F e 20 3 film will decrease under conditions that favour rapid repassivation (a higher potential, a strongly oxidizing medium). Under these conditions the depassivating effect of constituents may not appear. This may explain the fact that a strong anodic polarization and oxidizing agents reduce the cracking rate.

Tensile stresses play an important role only in the process of rapid crack growth, but they are of minor significance in the crack nucleation process. A ra-pid crack growth occurs after a longer corrosion time, when intergranular pe-netration has reached depth great enough for a local flow of material at the crack tip to occur, and at the same time a passive film has been formed on the metal surface. The effect of stresses consists in opening the intergranular cracks and rupturing the oxide film at the crack tip. Crack development occurs by electrochemical dissolution at a rate that grows exponentially with potential.

A moderate anodic polarization accelerates cracking, since the intensity of the depassivating effect of carbon or other constituents increases, and the crack growth rate in the presence of stresses grows exponentially with potential.

In summary, carbon affects the liability to Stress Corrosion Cracking by increasing the electrochemical reactivity of grain boundaries and by influencing the protective properties of oxide films. Carbon also favours deposition of the protective layer of magnetite, and thus reduces the susceptibility to cracking. The dominant effect will depend on the carbon content, its distribution, and on the potential and temperature.

2.14 Prevention of stress corrosion cracking

The susceptibility of ferritic steels to stress corrosion cracking can be reduced by the following measures.

1. Reduction of stresses. Stresses should be reduced below the threshold value, if there is any. External stresses are reduced by increasing the thickness of struc-tural parts under load or by changing the design. Internal stresses (for example welding) can be relieved by tempering at 773-923 K.

2. Choice of an appropriate grade of steel and modification of its composition. In addition to corrosion resistance, the following factors need to be considered in the selection of a steel: adequate strength and physical properties, machin-ability and weldability, availability on the market, as well as economic factors.

In modifying the steel composition the data on the effect of particular el-


ements will serve as guidelines (cf. Section 2.10.1.3). Under mild corrosion con-ditions a satisfactory resistance can be expected from a steel containing approx. 0.2% C, 0.3% Cr, 1% Mn, 0.2% Al, and possibly also other alloying additions.

3. Surface treatment. Compressive stresses or a large degree of cold work decrease SCC. Hence such treatments as shot or sand blasting, or peening may prove advantageous.

4. Mechanical treatment. Cold rolling to above approx. 50% renders low-car-bon steel virtually completely resistant to SCC

5. Electrochemical protection. In many cases this may prove to be the most economically sound method. An efficient protection should be achieved by polarization at potentials that correspond to the thermodynamic stability of magnetite on the corroding surface.

6. Use of inhibitors. Additives to solutions that are known to reduce SCC of low-carbon steels are treated in Section 2.9.2.

2.15 References

[1] Fontana, M. G., Met. Trans., 1, 3251 (1970). [2] Conference: Stress Corrosion Cracking and Hydrogen Embrittlement in Iron Based Alloys,

NACE, Unieux-Firminy 1973. [3] Parkins, R. N., Met. Rev., 9, 201 (1964). [4] Logan, H. L., The Stress Corrosion of Metals, J. Wiley, New York 1966. [5] Karpenko, G. V., Vasilenko, 1.1., Stress Corrosion Cracking of Steels (in Russian),

Tekhnika, Kiev 1971. [6] Brown, B. F., Stress Corrosion Cracking Control Measures, US National Bureau of

Standards, Washington DC 1977. [7] Scully, J. C , "Stress Corrosion Cracking", in: Corrosion: Aqueous Processes and Passive

Films, Ed. J. C. Scully, Academic Press, New York 1983. [8] Conference: Fundamental Aspects of Stress Corrosion Cracking, NACE, Houston 1969. [9] Conference: The Theory of Stress Corrosion Cracking in Alloys, NATO, Brussels 1971.

[10] Conference: Stress Corrosion Cracking of Metals—A State of the Art, ASTM, Phila-delphia 1972.

[11] Conference: Mechanisms of Environment Sensitive Cracking of Materials, Surrey 1977. [12] Conference: Electrochemical Test Methods for Stress Corrosion Cracking, Firminy 1978,

published in Corr. ScL, 20, 1980. [13] Bach, C , Zeit Verein Deutsche Ingenieure, 50, 1 (1906). [14] Jones, J. Α., Trans. Faraday Soc, 17, 102 (1921).

[15] Shvarts, G. L., Kristal, M. M., Corrosion of Chemical Equipment (in Russian), Mashgiz,

Moskva 1958. [16] Malek, K., Ochrona przed korozjq, 17, 276 (1974). [17] Popperling, R., Schwenk, W., Werkstoffe u. Korrosion, 27, 81 (1976). [18] Week, R., Inst. Gas. Engineers, publ. No. 517, 1957.

REFERENCES 93

[19] Kowaka, M., Nagata, S., Corrosion, 24, 427 (1968). [20] Parker, J. G., Brit. Corr. J., 8, 124 (1973). [21] Bûhler, H. E., Robusch, G., Friehe, W., Stahl u. Eisen, 96, 333 (1976). [22] Parkins, R. N., in ref. [9], p. 361. [23] Parkins, R. N., in ref. [2]. [24] Poulson, B., Corr. ScL, 15, 469 (1975). [25] Bohnenkamp, K., Arch. Eisenhuttenw., 39, 361 (1968). [26] Grâfen, H., Werkstoffe u. Korrosion, 20, 305 (1069). [27] Class, I., Gering, K., Werkstoffe u. Korrosion, 25, 314 (1974). [28] Sutcliffe, J. M., Fessier, R. R., Boyd, W. K., Parkins, R. N., Corrosion, 28, 313 (1972). [29] Brown, Α., Harrison, J. T., Wilkins, R., Corr, ScL, 10, 547 (1970). [30] Middleton, W. R., Brit. Corr. J., 8, 62 (1973). [31] Parker, J. G., Pearce, W. G., Corrosion, 30, 18 (1974). [32] Huckholtz, H., Pusch, R., Stahl u. Eisen, 62, 21 (1942). [33] Strauss, M. B., Bloom, M. C , Corrosion, 16, 533t (1960). [34] Schmidt, H. W., Gerner, P. J., Heinemann, C , Pogacar, C. F., Wyche, Ε. H., Corrosion,

7, 295 (1951). [35] Tresede, R., Wachter, Α., Corrosion, 5, 383 (1949). [36] Azhogin, F. F., Stress Corrosion Cracking and Protection of High-Strength Steels (in

Russian), Metallurgiya, Moskva 1974. [37] Parkins, R. N., Mazza, F., Royuela, J. J., Scully, J. C, Brit. Corr. J., 7, 154 (1972). [38] Henthorne, M., Parkins, R. N., Corr. ScL, 6, 357 (1966). [39] Flis, J., Brit. Corr. J., 3, 182 (1968). [40] Flis, J., Scully, J. C , Corrosion, 24, 326 (1968). [41] Brown, B. F., Met. Rev., 13, 171 (1968). [42] Flis, J., Scully, J. C, Corr. ScL, 8, 235 (1968). [43] Cihal, V., Desestret, Α., Wagner, H. G., Métaux, Corrosion, Industrie, No, 534, 1970. [44] Baumel, Α., Engell, H. J., Arch. Eisenhuttenw., 32, 379 (1961). [45] Lee, H., Uhlig, H., / . Electrochem. Soc, 117, 18 (1970). [46] Flis, J., Brit. Corr. J., 8, 57 (1973). [47] Romanov, V. V., Research Methods in Corrosion of Metals (in Russian), Metallurgiya,

Moskva 1965. [48] Smialowski, M., Szklarska-Smialowska, Z., Corrosion, 18, It (1962). [49] Flis, J., Corr. ScL, 15, 553 (1975). [50] Flis, J., Brit. Corr. J., 10, 79 (1975). [51] Engell, H. J., Bâumel, Α., Physical Metallurgy of Stress Corrosion Fracture, Ed. T. N.

Rhodin, Int. Publ., New York 1959, p. 341. [52] Szklarska-Smialowska, Z., Corrosion, 20, 198t (1964). [53] Jones, D. , Hackerman, N., Corr. ScL, 8, 565 (1968). [54] Moshtev, R. V., Hristova, Ν. I., Corr. ScL, 7, 255 (1967). [55] Lees, D . J., Lockington, Ν. Α., Brit. Corr. J., 5, 167 (1970).

[56] Bhat, P. C , Sathyavathiamma, M. P., Puttaswamy, N. G., Corr. ScL, 23, 733 (1983).

[57] Nagayama, M., Cohen, Μ., / . Electrochem. Soc, 109, 781 (1962); 110, 670 (1963).

[58] Sato, N., Kudo, K., Nodo, T., Corr. ScL, 10, 785 (1970).

[59] Flis, J., Corr. ScL, 10, 745 (1970).


[60] Mieluch, J., Bull. Acad. Polon. Sci., sér. Sci. Chim. 13, 33 (1965). [61] Mieluch, J., ibid., 10, 493 (1962). [62] Mieluch, J., ibid., 6, 337 (1963); 9, 429 (1967), 12, 597 (1967). [63] Fus, J., Mieluch, J., Smialowski, M., Corrosion, 20, 184t (1964). [64] Parkins, R. N., Slattery, P. W., Middleton, W. R. Humphries, M. J., Brit. Corr. J., 8,

117 (1973). [65] Van Vlack, J. H., / . Metals, 3, 251 (1951). [66] Parkins, R. N., / . Iron Steel Inst., 172, 149 (1952). [67] Graf, J., Becker, H., Z. Metallk., 62, 685 (1971). [68] Tikkanen, M. H., Puranen, P., Proc. 1th Scandinavian Corrosion Congress, Trondheim

1975. [69] Engell, H. J., Bohnenkamp, K., Baumel, Α., Arch. Eisenhutenw., 33, 285 (1962). [70] Hoar, T. P., Galvele, J. R., Corr. Sci., 10, 211 (1970). [71] Uhlig, H. H., Sava, J., Trans. ASM. 56, 361 (1963). [72] Humphries, M. J., Parkins, R. N., ref. [9], p. 384. [73] Smialowski, M., Korrosion u. Metallschutz, 14, 111 (1938); 15, 81 (1939). [74] Houdremont, E., Bennek, H., Wentrup, H., Stahl u. Eisen, 60, 757, 791 (1940). [75] Herzog, E., Portevin, M., Métaux et Corr., 24, 40 (1949). [76] Smialowski, M., Proc. 1st Intern. Congress on Metallic Corrosion, Butterworths, London

1962, p. 296. [77] Parkins, R. N., Usher, R., ibid., p. 289. [78] Zabik, W., Krukiewicz, W., Ochrona przed korozjq, 15, 327 (1972). [79] Flis, J„ Corrosion, 29, 37 (1973). [80] Nurnberger, U., Arch. Eisenhiittenw., 44, 775 (1973). [81] Schroeder, W. E., Berk, Α. Α., O'Brien, R. Α., Met. and Alloys, 8, 320 (1937). [82] Râdeker, W., Grâfen, H., Stahl u. Eisen, 76, 1616 (1956). [83] Cubicciotti, D. , Boyer, W., Weld. Res. Suppl. to Weld. Journ., 29, 140 (1950). [84] Mazille, H., Uhlig, H. H., Corrosion, 28, 427 (1972). [85] Heeley, E. J., Symposium on Internal Stresses in Metals and Alloys, Inst, of Metals, Lon-

don 1948, p. 473. [86] Gilchrist, J. D., Narayan, R., Corr. Sci., 11, 281 (1971). [87] Winterstein, M. G., McDonald, H. J., Waber, J. T., Weld. Res. Suppl. to Weld Journ.,

723, Dec. 1947. [88] Logan, H. L., Physical Metallurgy of Stress Corrosion Fracture, Ed. T. N. Rhodin, Int.

Publ., New York 1959, p. 295, [89] Baumel, Α., Engell, H. J., Arch Eisenhiittenw., 32, 379 (1961). [90] Henthorne, M., Parkins, R. N., Brit. Corr. J., 5, 186 (1967). [91] Bombara, G., Corr. ScL, 9, 519 (1969). [92] Kruger, J., Colvert, J., / . Electrochem. Soc, 114, 43 (1967). [93] Logan, H. L., / . Research Nat. Bur. Stand., 48, 99 (1952). [94] Grâfen, H., Werkstoffe u. Korrosion, 17, 485 (1966).

[95] Parkins, R. N., Stress Corrosion Cracking and Embrittlement, Ed.„W. D. Robertson, Chapman and Hall, London 1956, p. 140.

[96] Long, L. M., Uhlig, Η. Η., / . Electrochem. Soc, 112, 964 (1965). [97] Bellot, J., Hugo, M., Rolin, M., Herzog, E., Mem. ScL Rev. Met., 65, 607 (1968).

REFERENCES 95

[98] Long, L. M., Lockington, Ν. Α., Corr. Sci., 11, 853 (1971). [99] Herzog, H., Corrosion Trait, Prot. Fin., 20, 187 (1972).

[100] Smialowski, M., Métaux et Corr., 14, 56 (1939). [101] Munster, R., Grâfen, H., Arch. Eisenhuttenw., 36, 277 (1965). [102] Bâdecker, W., Mishra, Β. N., Werkstoffe u. Korrosion, 17, 193 (1966). [103] Rehm, G., Nûrnberger, U., Stahl u. Eisen, 91, 689 (1971). [104] Uhlig, H. H., Perumal, K. B., Talerman, M., Corrosion, 30, 229 (1974). [105] Uhlig, H. H., ref. [9], p. 86. [106] Uhlig, H. H., Physical Metallurgy of Stress Corrosion Fracture, Ed. T. N. Rhodin, Int.

Publ., New York 1959, p. 1. [107] Green, J. A. S., Parkins, R. N., Corrosion, 24, 66 (1968). [108] Monier, G., Gobin, P., Mazille, H., C. R. Acad. Sci., Paris, 267, 1771 (1968). [109] Mazille, H., Monier, G., ibid., 272, 278 (1971). [110] Athavale, G., Eilender, W., Korrosion u. Metallschutz, 16, 127 (1940). [111] Sachs, G., Van Hoorn, R. R., Practical Metallurgy, ASM, 1940. [112] Parkins, R. Ν., Brown, Α., / . Iron Steel Inst., 193, 45 (1959). [113] Herzog, E., Métaux et Corr., 24, 29 (1949). [114] Baerlecken, E., Hirsch, W., Stahl u. Eisen, 73, 785 (1953). [115] Râdeker, W., Mishra, B., ibid., 17, 193 (1966). [116] McDonald, R. D. , Corrosion, 24, 214 (1968). [117] Moloznik, K. L., Briant, C. L., McMahon, C. J., Jr., Corrosion, 35, 331 (1979). [118] Lea, C , Metal Sci., 14, 107 (1980). [119] Kupper, J., Erhart, E., Grabke, H.-J., Corr. Sci., 21, 227 (1981). [120] Lea, C , Hondros, E. D. , Proc. R. Soc. London, A 377, 477 (1981). [121] Wei, W., Grabke, H. J., Corr. Sci., 26, 223 (1986).

[122] Krautschick, H. J., Bohnenkamp, K., Grabke, H. J., Werkstoffe u. Korrosion, 38, 103

(1987). [123] Tauber, G., Grabke, H.-J., Corr. Sci., 19, 793 (1979). [124] Smialowski, M., G^sior, E., Bieniosek, C , Bull. Acad. Polon. Sci. Lettres Suppl, 1,

163 (1950). [125] Nohl, F., Werner, W., Stahl u. Eisen, 59, 1155 (1939). [126] Parkins, R. N., Brit. Weld. J., 2, 495 (1955); 8, 24 (1961). [127] Humphries, M. J., Parkins, R. N., Proc. 4th Intern. Congress on Metallic Corrosion, Ex-

tended Abstracts, NACE, Houston 1969, p. 27. [128] Cocks, F. H., Bradspies, J., Corrosion, 28, 192 (1972). [129] France, W. D., Jr., ibid., 26, 189 (1970). [130] Azhogin, F. F., Intercrystalline Corrosion and Corrosion of Metals under Stress, Ed. I. A.

Levin, Consultants Bureau, New York 1962, p. 211. [131] Zabik, W., Ochrona przed korozjq, 12, 73 (1969).

Added in proof

[1] Lea, C , Corrosion, 40, 337 (1984). [2] Narayan, R., Kumar, Α., Singh, K. P., Corr. Sci., 25, 449 (1985). [3] Hunt, C. P., Corr. Sci., 28, 901 (1988). [4] Krautschick, H. J., Grabke, H. J., Diekmann, W., Corr. Sci., 28, 251 (1988).

3. Corrosion of Structural Materials in Metal-Cooled Nuclear Reactors

3.1 Introduction

The late 1950's witnessed the onset of a rapid growth in nuclear technology. In this period, research was begun into the design of fast breeder reactors. The operation of these reactors is associated with a release of large quantities of which must be removed from the system. The core temperatures of reactors of this type sometimes exceed 873 K. Conventional cooling media are unsuitable, and so the problem arose of finding adequate new coolants. Two cooling media have found practical application: gases under high pressures, and liquid me-tals.

This chapter will deal with some of the corrosion problems that are encoun-tered in metal-cooled nuclear reactors.

3.2 Liquid metals and structural materials used in fast breeder reactors

Liquid metals is an arbitrary term for metals whose melting point does not exceed 933 Κ (the melting point of aluminium) [1]. The major physical and nuclear properties of these metals are set out in Tables 3.1 and 3.2. Some of those metals, with particular thermal properties, can be used as coolimg media in nuclear practice. Such metals, in addition to having a fairly low melting point and a high boiling point, and adequate thermal properties including a good thermal con-ductivity and a high heat capacity, should not be corrosive towards structural materials. For applications in the radiation zone of reactor systems the liquid metal should additionally be resistant to neutron and gamma radiation, and should exhibit appropriate nuclear properties, such as a low cross section for neutron absorption, a short half-life of radioactive isotopes, etc. The price and availability are additional factors to be considered in selecting a coolant, as the cooling medium is generally required in substantial quantities in the circuit.

TA

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LIQUID METALS USED IN FAST BREEDER REACTORS 103

98 CORROSION OF STRUCTURAL MATERIALS IN REACTORS

TABLE 3.2

NUCLEAR PROPERTIES OF LIQUID METALS (after [2, 3])

Metal and its isotope

Isotope percentage

Thermal neutron ab-

sorption cross

section

Fission products

Half-life Energy, MeV, and the type of radia-

tion

Lithium 6Li 7Li

7.5 92.5

65 370

0.033 ! H e + f H

iLi 31 years 0.85 s

0.015/5 13.0/5; 1.4/5

Sodium 2 3N a

100 0.45 ÏÎNa 14.8 h 1.4/5; 1.38/9 2.76/5

Potasium 3 9 K 4 0K 4 1K

93.4 0.01 0.6

1.97 1.87

1.19

4 0 V-1 9 *^ 4 0A 4 2

* r 19JV

1.3 10~

9y

stable 12.4 h

1.46y; 1.36/5

3.6/5; 1.9/5 1.51y

Mercury 1 9 6H g 1 9 8H g

1 9 9H g

200Hg 2 0 1

H g 2 0 2H g

2 0 4Hg

0.15 10.1

17.0 23.3 13.2 29.6

6.7

430 3100

20

2500 60 60

2.4 0.34

'fo'Hg »88 Hg 2S8Hg 2 0 1 τ τσ a88Hg 2gaHg a88Hg

64 h 23 h stable stable stable stable 47 days

5.5 min

Ο.8Ο7 0.16y isomers 0.16y

0.46/5 0.3y 1.62/5

Lead 2 0 4 pb 2 0 6 pb 2 0 7 pb 2 0 8 pb

1.5 23.6 22.6 52.3

0.2 v. low v. low v. low v. low

radioactivity negligible

Bismuth 2 0 9Bi 100 0.015 2ioBi

ultimate 2ioPo

5 days 1.17/5; 5.3a

273 373 473 573 673 773 '873 973 1073 Κ

Fig. 3.1. Resistance of various structural materials to mercury at high temperatures (after [9]): I—good, II—reduced, III—poor, IV—unknown, V—data obtained from loop and commercial testing, VI—loop test data, VII—data from steady-state testing; 1—Ti, Mg, 2—low-carbon steel, 3—low-carbon steel containing 0.1-4% Al, 4—low-carbon steel containing 4% Cr, 5— 5% Cr steel, 6—low-carbon 0.5% Mo steel, 7—low-carbon 20% Mo steel, #—low-carbon 2-3% Si steel, 9—low-carbon 1-2% Ti steel, 10—low-carbon 2% Al + 2% Cr steel, //—Nitra-lloy (1.23% Al + 1.49% Cr), 12—low-carbon 5.7% Cr steel, 13—low-carbon 4.5% Mo steel, 14—low-carbon + 5.7% Co + 1 . 2 % W steel, 15—low-carbon 15-20% Mo + 3% Si steel, 16—low-carbon 8% Cr + 0.5% Al + 0.3% Mo, 17—Silchrome 5 S (5% Cr + 0.5% Mo + + 0.5% Si), 18—low-carbon 5.5% Cr + 6.4% Mo + 1.4% Si, 19—chromium-nickel type 304 and 310 steels, 24—high Ni and CrNi alloys, 21—chrome stainless steels, 22—tungsten, 23— molybdenum, 24—chromium, 25—beryllium, 26—-Ta, C, Nb, Si, Ti, V, 27—Ni, Cr and their alloys, 28—cobalt and stellites, 29—platinum, magnesium, zirconium, 30—Al, Bi, Cd, Ce, Au, Pb, Mg, Ag, Sn, Zn, 31—glass, 52—-ceramics, 33—graphite.

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19

88

975

375

2 90

0 U

02-P

u0

2 N

a 66

1/80

8 S

UP

ER

PH

ÉN

IX-I

I F

ran

ce

4 19

90

3700

15

00

U0

2-Pu

02

Na

CD

FR

U

K

1990

32

30

1320

6

600

U0

2-Pu

02

Na

643/

813

SN

R-2

F

RG

19

90

3420

13

00

12 0

00

U0

2-Pu

02

Na

663/

813

BN

-160

0 U

SS

R

1990

42

00

1600

8

800

U0

2-Pu

02

Na

623/

823

Dem

onst

rati

on

Japa

n 19

90

2400

10

00

80

00

U0

2-Pu

02

Na

1 In

coo

pera

tion

wit

h F

RG

and

Eur

atom

. 2

Th

0 2-coa

ted

3 In

coo

pera

tion

wit

h B

elgi

um a

nd t

he N

eth

erla

nd

s.

4 A

joi

nt p

roje

ct w

ith

Ital

y, F

RG

, B

elgi

um a

nd t

he N

eth

erla

nd

s.


TABLE 3.4

CHEMICAL COMPOSITION OF THE STRUCTURAL MATERIALS USED

Material type Grade

Content,

Material type Grade C max

Fe Cr Ni Mn

Stainless steels

304 0.08 matrix 18-20 8-12 max 2.0

Stainless steels

316 0.08 matrix 16-18 10-14 max 2.0

Stainless steels 321 0.08 matrix 17-19 9-12 max 2.0 Stainless steels

347 0.08 matrix 17-19 9-13 max 2.0

Stainless steels

Incoloy 800 0.04 46.0 20.5 32.0 0.75

Nickel-based alloys

Inconel 600 0.15 10.0 17.0 matrix 1.0

Nickel-based alloys

Inconel X750 0.08 9.0

18.5

17.0 matrix 1.0

Nickel-based alloys

Hastelloy X 0.1

9.0

18.5 22.0 matrix 1.0 Nickel-based alloys

Inconel 718 0.1 matrix 21.0 55.0 0.35

Nickel-based alloys

Inconel 625 0.1 5.0 23.0 matrix 0.5

Ferritic steels

10CD9, 10 2V4Cr-lMo

0.15 matrix 2.0-2.5 0.3-0.6

Ferritic steels

10CD Nb 9, 10

2

1/4Cr- lMo- lNb

0.10 matrix 2.0-2.5 0.3-0.8 0.4-0.8

Ferritic steels Z10CD9 9Cr-lMo

0.15 matrix 8.0-10.0 0.3-0.6 Ferritic steels

Z10CD NbV 9,2 9Cr-2Mo-NbV

0.15 matrix 8.5-10.5 0.8-1.3

The majority of the liquid metals listed in Tables 3.1 and 3.2 have found no practical use. Some data related to metal-cooled reactors are collected in Table 3.3. The listing is not complete, as information on the reactors in operation is frequently undisclosed and hence unavailable, and other data are publicized with


%by wt.

Mo Nb Al J Ti Si

max W Ρ

max S

max Cu V

1.0 0.045 0.030

2.0-3.0 1.0 0.045 0.030

5 Χ Cmin 1.0 0.045 0.030

10 Χ Cmin 1.0 0.045 0.030

0.35 0.30

0.5

1.2 1.0 2.75 0.5

9.0 0.75 0.6

3.3 5.5 0.8 1.15 0.35

10.0 4.15 0.4 0.4 0.5

0.9-1.1 0.10-0.50 0.03 0.03

0.9-1.1 10 χ Cmin 0.15-0.50 0.02 0.02

0.9-1.1 0.02 max

0.25-1.00 0.03 0.03

1.7-2.3 0.30-0.55 0.20-0.65 0.03 0.03 0.2-0.4

considerable delay. Certain of the reactors specified in Table 3.3 have long been removed from operation and were designed as research reactors. This was for instance the case of the American Clementine reactor or of the Soviet BR-1 and BR-2 reactors, which were mercury coolled.The use of mercury in the sys-

IN SODIUM-COOLED NUCLEAR REACTORS (after [10, 14])


Iron Alloys Armco iron Carbon steel Grey cast iron Steel SAE 52100 (1.45 Cr,1.0C Sicroma 5asteel (5Cr, 0.5Mo,1. Steels 2-9Cr,0.5-1Mo' Steel 18/8qrade Steel 310 (Z5Cr,20Ni) Ferritic stainless steel (27Cr) — Refractory steel (20Cr, S i , Mo, Cu) High-speed steel (18W, 4Cr, 1V) Invar (36 Ni) Fenico(28Ni,18Co) Alnico5(Ni,Co, Al, Cu)-Hadfietd steel Nitralloy G (nitriding steel) 80Fe/20Psotder Nickel and Nickel Alloys

Nickel -Inconel (13 Cr, 6.5 Fe>-Nichrome (15Cr, 25Fe)-Monel (30 Cu)-Hastelloy A, Β and C -Ni-Mn, Ni-Mo and N-P solders-Copper and Copper Alloys Copper -Electrolytic copper Aluminium bronze (5-8Al) Beryllium bronze (Z Be) — Nickel alloys (20 and 30 Ν Γ Brass (40 Zn) German (nickel) silver type alloys(17Zn,18Ni)-t Refractory Metals Niobium-Molybdenui Tantalum Titanium Tungsten Vanadium Zirconium Chromium Other Metals

Cobalt and high-cobalt alloys Aluminium, 2Sand 3S -Aluminium, 24S and 52! Eutectic alloy, Al - Si Beryllium Sb, Bi, Cd, Ca, Au, Pb, Se. Ag, S, Sn. Magneskji Platinum and silicon Non-metallic Materials

273 373 473 573 673 773 .873 973 1073 1173 Κ

AI2O3 BeO MgO Pyrex vycor etc. type glasses Darkoic* AsbesU Silicone rubber Teflon High-density graphite

Corrosion resistance rating:

duced poor unknown

Fig. 3.2. Resistance of structural materials to sodium at high temperatures (after [12]).


tern gave considerable trouble with the pumping, in view of the high density of mercury and its high electrical resistivity. Furthermore, mercury is corrosive towards many commonly used structural materials (Fig. 3.1). Accordingly, the reactors mentioned were operated for a short time and, as a result of studies in service, mercury has been ruled out as a cooling agent in nuclear reactors. Of the metal-cooled power breeder reactors the sodium-cooled ones have re-mained as the only variant of practical significance. The listing of Table 3.3 also includes the data on the fast breeder reactors under current service or for future construction.

Figure 3.2 illustrates diagrammatically the corrosion resistance of various structural materials to liquid sodium at high temperatures. A comparison of Figs. 3.1 and 3.2 reveals that the majority of structural materials at high tem-peratures is more resistant to liquid sodium than to mercury. This observation is also true for common materials used in nuclear practice such as stainless chromium-nickel steels type 18-10 and the ferritic chromium steels. In the con-struction of sodium-cooled fast breeder reactors, three main types of structural materials are used in practice: austenitic chromium-nickel steels, some nickel alloys and, increasingly, ferritic chromium steels. The compositions of the more important of these materials are reported in Table 3.4. It should be men-tioned that analogous steel and nickel alloy grades are used for the construction of sodium cooling circuits in the USSR [13].

The structural material to be used depends on the operating conditions of the cooling system and especially on the sodium temperature. If this is fairly low (e.g. 773-793 K) ferritic chromium steels can be used because of their fairly good mechanical properties and their adequate resistance to liquid sodium [14-16]. As far as corrosion resistance goes in these steels, an essential factor is the decarburization process, which influences the steel structure. It has been found that niobium added alone to stabilize the structure of ferritic steels has a minor effect; if, however, niobium is combined with vanadium as an addi-tion, the stabilizing effect is satisfactory [16]. In ferritic steels, complex oxides of the chromite type can form with sodium, or intergranular corrosion may occur. Such steels, despite their relatively inferior properties, are commonly used because of their low price.

The most popular structural materials in fast breeder reactors are auste-nitic chromium-nickel steels, both as general structural and cladding materials [10, 13, 16-23]. Molybdenum and high-nickel steels are generally used, in view of their improved mechanical properties at high temperatures and a better creep resistance. Incoloy type steel is employed in the core sections of reactors, as it exhibits a low tendency to swelling under large doses of neutron radiation [11].


3.3 Corrosion mechanisms in liquid metals

The results of corrosion of solid metals immersed in liquid metals are the same as those observed in other environments, namely: change in weight, surface oxidation, and intergranular corrosion. The mechanism of these effects, however, differs from that for conventional corrosion processes. Essentially two types of corrosion mechanism are generally distinguished:

—• corrosion that consists of the dissolution of solid metal or some of its constituents in the liquid metal, not involving chemical reactions, at least at the initial stage of the process;

— corrosion which is the result of chemical reactions between the liquid metal or some of its constituents or impurities and the solid metal or one of its constituents.

Obviously, in actual practice the two mechanisms never operate independ-ently, and in general the corrosion processes are highly complex, as the rate and character of the initial stage of dissolution corrosion changes into that of the chemical corrosion stage. Characteristically, no electrochemical effects have been observed in liquid metals.

3.3.1 Dissolution

As mentioned previously, a solid metal can dissolve in a liquid metal uniformly or selectively, if certain metal constituents alone pass into solution. These pro-cesses are closely related to diffusion in the liquid layers next to the solid metal surface and to the mass transfer caused by the flow of the coolant [13, 25, 26].

Individual metals exhibit different solubilities in different liquid metals.

TABLE 3.5

SOLUBILITY OF IRON, ppm, IN LIQUID METALS AT HIGH TEMPERA-TURES (after [2])

Temperature, Κ

Liquid metal Temperature, Κ Na Hg Pb Bi

673 0.0046 0.075 0.25 2.0

773 0.0084 0.23 0.97 12.3

873 0.013 0.55 2.71 50.1 973 0.018 1.08 6.72 155

1073 0.025 1.9 11.9 377

CORROSION MECHANISMS IN LIQUID METALS 107

For instance sodium readily dissolves cadmium, lead, tin, magnesium and silver, while iron, nickel, chromium and molybdenum are sparingly soluble. In practice, therefore, alloys containing readily soluble components are avoided, as such components are the site of accelerated corrosion attack.

Table 3.5 provides the solubility data for iron in a number of liquid metals [24]. The solubility varies considerably with the temperature in individual liquid metals, and similar variations occur in the solubilities of other metals.

Theoretically, the metal dissolution process under steady-state conditions should proceed until an equilibrium is established on reaching saturation of the liquid metal with a given component. At this point corrosion damage of the solid metal should be arrested, for example in the case of iron when its concentra-tion in sodium at 973 Κ attains a value of 0.018 ppm. In practice, however, this equilibrium may be disturbed either because of a temperature gradient or if another solid metal, capable of forming an intermetallic compound with iron, is present in the system. This, for instance, is the case of aluminium. In sodium, this metal displays an equally low solubility as iron. In a ternary Fe-Al-Na system, however, it dissolves rapidly to yield an intermetallic compound on the iron surface [27].

The major factor disturbing the equilibrium of the dissolution process in the mass transfer process. In view of the change of solubility with temperature of the dissolving metal components, a component dissolving in the hot zone of the system (i.e. in the higher-temperature region) may be carried along with the cooling agent to the cold zone, where it deposits, for instance, on the inner walls of the plant. Thus, in the hot zone the concentration of the dissolving compo-nent is constantly lower than that corresponding to the saturation condition, and the process continues, never reaching an equilibrium condition.

In steady-state systems the rate of corrosion due to dissolution is controlled by the diffusion of atoms of the component dissolving from the surface layer to the liquid metal [13]:

ν = — ^ — - (3.1)

where ν is the dissolution rate, cT is the equilibrium concentration at dissolution temperature Γ, c is instantaneous concentration, δ is the thickness of the diffu-sion layer of solvent, D is the diffusivity.

The instantaneous concentration cT in a steady-state system depends only on the temperature. The temperature dependence of the dissolution rate was found


to be affected to a considerable degree by the ratio of atomic radii of dissolving component A and solvent Β [28] :

S = rBrA (3.2)

X-ray evidence has shown that for atoms dissolved in a liquid metal, the coordination number is frequently between 8 and 12. For these cases the tem-perature coefficient of solubility can be evaluated from the equation:

167.60

Ε 146.65 ο ο en

®

12570

104.75

ο 83.80

Ç, 62.85

I 41.90

20.95

Sol- Dissolv-vent ing

metal 1 Cd Ni 2 GA BE 3 Sb U 4 Tl Th 5 TL Ρ 6 CO C 7 FE C 8 Mn c 9 Ni c

10 PB ZN 11 PB CO 12 PB CR 13 PB CU 14 PB FE 15 PB Mn 16 P B Ni 17 PB U 18 PB W 19 PB SR 20 NA TA 21 NO Ag 22 NA CA 23 NA Cd 24 NA FE 25 NA NI 26 NA PB 27 NA SB 30 Bi ZR 31 Bi C 32 Bi Cr 33 Bi FE 34 Bi Pu 35 Bi U 36 Bi BA 40 Hg SN 41 Hg BA 42 Hg FE 43 Hg RB 50 Mg U 51 Mg FE 52 Mg MN 53 Mg Ti

BA 54 Mg Ti BA

58 ZN Th 59 ZN CE 60 ZN τ j 61 ZN Bi 62 ZN Ni 63 ZN PB 64 ZN U 65 ZN MO 66 ZN BA 67 ZN LA 68 Zn NA 69 ZN ZR 70 SN TE 71 SN BE 72 SN LA 80 IN U 81 IN B E 90 Al B I 91 Al PB 92 Al U

^ M 0.6 0.8 1.0 1.2

r 1.4 1.6 1.8

Size factor S=T%-

2.0

Fig. 3.3. The temperature coefficient of solubility vs. size factor (after [28]).


d(RlnNAW) _ RT2 dNA(»

d(l/T) NAU dT

where NAm is the atomic fraction of the dissolved metal in the saturated sol-ution.

Figure 3.3 presents the dependence of temperature coefficient of solubility on the value of size factor S (for 0.6 < S < 2.0). The curve is composed of two branches which reach a maximum at S = 1.4.

If a dissolution process follows the van't Hoff equation for an equilibrium condition Αω ^ A (in liquid B), where A is pure solute in equilibrium with the solution of metal A in metal B, the temperature coefficient of solubility can be estimated from the equation :

d(RlnNAil}) dl/T

= àQH+AQAi0 (3.4)

where AQ£ is the heat of fusion of metal A, AQA(l) is the partial molar heat of dissolution of component A in dilute solutions.

In cases where an intermetallic compound is present in the composition of the solid metal, equation (3.4) must be modified. In the majority of cases the value of AQA can be evaluated from the following equation:

AQm

A = 2.273 (3.5)

where is the melting point, in K. The temperature coefficient of solubility therefore depends primarily on the

AQaw value, which represents the difference in the energy of interaction between atoms A surrounded by atoms of the same metal A and atoms A surrounded by atoms of metal B. This value can be found from atomic properties. The corre-lation of temperature coefficient of solubility with size factor is satisfactory. If the size factor equals unity, the energy of dissolution begins to rise, as occupa-tion of the sites in the metal Β lattice by metal A atoms is increasingly difficult. If S = 1.4, the coordination number changes, and atom A is then surrounded by only six Β atoms with which it interacts. This is considered to explain the decline in the temperature coefficient of solubility observed when S > 1.4. The squares in Figure 3.3 represent points obtained from calculations performed on the basis of this assumption. The legitimacy of the assumption is again con-firmed by calculations of carbon solubility in transition metals, which are known to have carbon in interstitial positions. Heavy-line symbols, on the other hand, refer to the points that correspond to systems for which equation (3.4) holds good.


In the case of dynamic systems with a moving coolant the dissolution process changes [28-30], A number of equations are available for evaluation of the dissolution rate of a solid metal in a liquid one. For instance, in reference [28] Epstein's equations are reported; these account for the mass transfer as a func-tion of the temperature gradient of solubility:

— the case of iron dissolving in mercury:

where νΎ is the corrosion (dissolution) rate, D is diffusivity, d is pipe diameter, vp is coolant velocity, ν is kinematic viscosity of liquid metal, S° is equilibrium solubility, AT is the temperature difference in the system, / is length of the sys-tem, α is a coefficient, characteristic of the system.

In the former case, dissolution rate is high and the mercury layer that re-mains in contact with iron is constantly saturated. The rate at which iron passes into the liquid is controlled by the rate of diffusion of iron in mercury. In the latter, the sodium layer in contact with iron is not saturated, and the rate with which iron passes into liquid is controlled by the dissolution rate.

The dissolution process is also affected by other factors. For instance,

(3.6)

— the case of iron dissolving in sodium:

(3.7)

s

ο on

JO

Time

Fig. 3.4. Effect of oxygen content on the solubility of iron in sodium (after [2]).

CORROSION MECHANISMS IN LIQUID METALS 1 1 1

in the case of the iron-liquid sodium system, the oxygen content in the liquid is of great importance. With increasing oxygen content in the sodium, the iron dissolution rate is higher, but the iron solubility at a given temperature remains the same (Fig. 3.4).

The selective dissolution frequently observed may also be divided into two categories: extraction which occurs when only one alloy component dissolves at a high rate, and intergranular attack, which involves penetration of the liquid metal into the bulk of the solid metal along grain boundaries. The latter occurs for instance if a number of readily soluble precipitates have accumu-lated at the grain boundaries. An example of this is a selective leaching out of phosphorus and silicon from the surface layers of soldered joints, observed during a corrosion study of such joints formed between chromium-nickel steel pieces using Nicrobraz type alloys, exposed to liquid sodium flowing at high rates and at high temperatures [17, 31]. The drop in phosphorus content was found to occur at greater depths.

3.3.2 Chemical reactions

The corrosion of structural metals in liquid metals is a complex phenomenon and, as has been stated, it does not involve only dissolution processes. A sig-nificant role is also played by chemical reactions of the solid metals or some of their constituents with liquid metals or impurities present in them. In the case of the liquid sodium-structural metal or mercury-structural metal systems, the reactions of solid metals or their constituents and impurities with the oxygen dissolved in liquid metal are the most important [2, 9, 12, 13, 26, 32].

In accordance with Henry's law, at a given temperature the activity of oxygen in a liquid metal is determined by the following ratio :

^ _ [ppm] oxygen present in liquid phase ^ ^ a° [ppm] oxygen dissolved at equilibrium

and the change in Helmholtz free energy of oxygen AF- in solution is

AF' = FT+RTloga0 (3.9)

where AFT is the change in the Helmholtz free energy of formation of oxides at a given temperature T.

Figure 3.5 shows the variations in the Helmholtz free energy accompanying the reactions of formation of various metal oxides. The data presented indicate, for example, that sodium can reduce iron oxides, as well as nickel and chromium oxides at temperatures below 973 K. The protective films on chromium-nickel


Temperature, Κ 1073 1473 1873 2273 2673

83.6

167.2

250.8

334.4!

418

501.6

° 585.2 "ο e ^668.8 χ.

^ 752.4 ι

836

919.6

1003.2

1086.8

1170.4

1254

Fig. 3.5. of metal

Zr02-U02 7 Ce02' 2Li20-2MgO

Th02 2BeO 2CaO"

Variation in the Helmholtz free energy accompanying the reactions of the formation

oxides (after [2]).

stainless steels thus undergo destruction at temperatures as low as approx. 773 Κ [13]. Often, however, films of complex oxides of the type 2 N a 20 - FeO or N a 20 · C r 20 3 are formed on steel surface, as confirmed by X-ray evidence gathered by Horsley [32]. Cavell and Nicholas found chromite N aCr 02 to form on surface of steel 316 in liquid sodium containing trace oxygen (a few ppm) [34],

The most significant of the chemical reactions involved in these systems is the reaction between solid metal and sodium oxide dissolved in the liquid phase:

(x/y)Me + N a 20 = (I/y)(Me)xOy + 2Na


The reactions of individual components of structural metals with sodium oxide are the subject of a large number of studies [12, 13, 35-39].

Weeks et al [37], in their study on the effect of oxygen on the dissolution of iron in sodium, found that iron oxide was formed in the hot leg of the test loop;* this was then carried over to the cold section where it was reduced to metallic iron as the result of another reaction with sodium. Thus, instead of iron oxide accumulating in the cold section, metallic iron is deposited as a porous layer on the loop walls.

Possible reactions of solid metals with other impurities present in the sodium were also considered. Borgstedt and Frees [38] studied the effect of calcium on the corrosion of vanadium-based alloys over a temperature range of 773-873 Κ in liquid sodium. They found calcium oxide CaO to be incapable of reacting with vanadium or with other metals in the same way as sodium oxide, and that it directly forms complex oxides, following the equation:

CaO + MeO = CaMe02

The authors concluded that vanadium-based alloys can be operated in so-dium that contains calcium below 10 ppm and oxygen below 10 ppm.

Pits were also observed to form as a result of a reaction of sodium with non-metallic inclusions occurring in steels [41].

There is lack of consensus as to the role played by oxygen dissolved in so-dium in corrosion processes. For instance Klueh [42], who studied the corrosion of niobium at a temperature of 873 K, found that oxygen, when present in mi-nute quantities, substantially accelerates the dissolution of niobium, and as soon as the solubility limit is exceeded a violent process of sodium penetration starts along the grain boundaries into the bulk of the solid metal. The latter effect is ascribed by the author [42] to the formation of complex oxides resulting from the reaction of sodium with the impurities present in the solid phase.

Opportunities for chemical reactions between the impurities contained in the liquid metal and structural materials are many and varied. Such reactions are perhaps best considered by analogy with metallurgical processes, where the commonly formed phases exhibit no stoichiometric composition and form sys-tems composed of mutually substituting constituents.

* Test loops are the facilities provided to carry out research under dynamic conditions with the cooling agent recirculating. They constitute a closed piping conduit whose individual sections, referred to as legs, are heated or cooled as desired. The loops are generally equipped with auxiliary devices as pumps, valves, and settling tanks known as traps, etc.


3.3.3 Intergranular corrosion

Intergranular corrosion and the liquid metal embrittlement has often been ob-served, in addition to other types of corrosion, in studies on the corrosion of structural metals in liquid metals [9, 12, 13, 35-38, 41, 43, 47, 48]. Figure 3.6

Fig. 3.6. Intergranular corrosion in liquid sodium. Steel 1H18N9T kept in sodium at 773 Κ

for 2000 h (after [38]).

illustrates an example of intergranular corrosion of steel caused by liquid so-dium. It is well established that easiest diffusion paths for a liquid agent into the bulk of a solid metal are the grain boundaries, because it is at the grain boundaries, which have a higher energy, that inclusions commonly accumulate; often this is also true of precipitates and defects. Occasionally, the liquid metal may occupy the voids formed in the solid metal as a result of a selective leaching-out of one or more of its components. This is, for example, the case of chro-mium-nickel steels immersed in liquid sodium, where chromium or nickel is selectively leached out from the steel. In a study [13] of the stainless steel-liquid sodium system, no effect of oxygen concentration on the intergranular corrosion rate and depth was noted. It was found, however, that specimens removed from the sodium and thoroughly washed in water and alcohol tarnish after some time to produce a white film. The film is water-soluble and its solution is alkaline; sodium was found to be one of its components.

Other explanations of the mechanism of intergranular corrosion have been put forward [13]. Some investigators believe that it is related to low melting


point eutectics that appear at the grain boundaries, for example Na2 · 2S i02 and S i 0 2, melting point 1066 K. Others claim [13] that in view of the fact that the atoms situated at the grain boundaries are thermodynamically less stable, they are thus capable of reacting more readily with, say, oxygen (a sodium impurity) to yield complex oxides of the ferrite type. In this case the oxygen content is a decisive factor and may also accelerate intergranular corrosion.

Generally, however, intergranular corrosion in sodium systems is considered fairly unharmful, but it is dangerous on the water side of heat exchangers.

Intergranular corrosion was also noted in steels immersed in mercury at high temperatures [9]. .This has been found not only from laboratory research but also during operational testing. The solubility of iron in mercury is minute (Fig. 3.7) and essentially this metal forms no amalgam. At elevated tempera-

tures, however, mercury was found to penetrate into the solid metal, during which process the steel undergoes both general and intergranular corrosion. These processes are markedly accelerated by a recirculating coolant, because of the mass transfer involved; the corrosion of steel in mercury then occurs at a much higher rate.

Fig. 3.7. Solubility of iron in mercury vs. temperature (after [9]).

1 1 6 CORROSION OF STRUCTURAL MATERIALS IN REACTORS

The corrosion of steel and other structural materials was also studied in saturated mercury vapours at high temperatures (673-973 K) in an attempt to find an adequate resistance of steel to this environment [2, 9].

The results of studies obtained by various authors are often discrepant. One reason for this, undoubtedly, is the failure to standardize matters such as ex-perimental conditions, use of materials of a similar composition but differing in structure, thermal treatment, and the quantity of impurities, etc. An equally common reason for the disagreement of results is the use of liquid metals of differing qualities.

The corrosion processes occurring in liquid metals are highly complex. No uniform theory has so far been developed, neither is there agreement as to which is the main factor in the process.

Much attention has been paid to the chemical reactions which result in the formation of oxide films on the surface of structural materials. The formation of such films may be accompanied by a substantial rise in volume, which gives rise to stresses that result in strain in the surface layers of the solid metal; slip planes formed as a result may ultimately become paths of facilitated diffusion [41].

3.4 Mass transfer, carburization and decarburization of steel

In systems with a recirculating cooling agent, mass transfer occurs, largely due to the motion of the liquid medium. Mass transfer accelerates corrosion processes, as it causes the dissolved components and products of chemical reactions to be continuously removed from the hot section to the cold section of the system. Thus equilibrium conditions for the chemical reactions or physical phenomena occurring at the solid-liquid interface and in the surface layers of a structural material cannot be attained.

Another consequence of the mass transfer related to the deposition of various products on the walls in the cold section of the system, is an impaired heat exchange and a hindered flow of coolant.

Examples are the deposition of iron as a porous layer [40], or the formation of oxide films [32, 33, 38, 40].

Among the corrosion products carried along with the cooling agent, there are also various radioactive isotopes produced in the core section of the reactor. The most important of these are:

5 4M n ,

5 8C o , and

6 0C o . The mass transfer

process may also include fuel components from damaged fuel elements (or the products of reaction of the fuel with sodium) [10, 49, 50]. Deposition of radio-

INTERACTION OF FUEL WITH FUEL-ELEMENT CLADDING MATERIAL 117

active particles on the surfaces of the cooling system results in a general rise in the activity.

Dissolved alloying components or impurities from structural materials and the products of their reaction with the impurities contained in the liquid phase are also carried by the coolant stream. In the case of steel the major results of this process are carburization and decarburization. These phenomena are given considerable attention, as, by affecting the structure, they can substantially change the mechanical properties of steel.

Diffusion of carbon into austenitic steel may occur through the grain surface or along the grain boundaries. The latter process, known as intergranular diffu-sion, was found to be about 1.5 times faster than the former, recognized as trans-granular diffusion [33].

The carbon source in liquid sodium systems are carbon- and low-alloy steels which undergo decarburization. This process, next in importance to diffusion and dissolution, is claimed by some authors [13, 51] to be a probable result of complex chemical reactions between liquid metal and cementite, releasing free carbon atoms.

The solubility of carbon in sodium is low. In the light of recent findings [52], the carbon concentration in liquid sodium at temperatures below 723 Κ does not exceed 0.005 ppm. In the absence of mass transfer, therefore, a perceptible decarburization of steel would be unlikely.

However, a clear carburization of chromium-nickel steels, containing car-bon in micro quantities, is actually observed. Thorley and Tyzack [17] found a six-fold increase in percentage of carbon in steel grades 316 and 316 L in their contact with flowing sodium at a temperature of 923 K. The carbon source in this system was a low-alloy steel containing 0.36% C. As a result of carburiza-tion, composite carbides formed both at the grain boundaries and inside grains.

As a result of the carburization of chromium-nickel steels, their plastic prop-erties deteriorate sharply [17, 38, 53], while the hardness of their surface layers increases [17]. In the carburized steel layers, largely carbides and the σ phase are formed [53, 54]. The carbides usually precipitate at the grain boundaries, as well as at the slip and twinning planes.

3.5 Interaction of fuel with the fuel-element cladding material

The heat source in nuclear reactors is the fuel, composed of the radioactive elements: uranium, thorium, and plutonium. Ceramic fuel is commonly used, consisting of sintered oxides or carbides of these elements. The fuel is introduced


into the reactor core in the form of fuel elements, composed of hermetically sealed fuel in a thinly clad can. The can is designed to protect the fuel against the corrosion and erosion effects of the coolant and to prevent radioactive fission products from entering the reactor cooling circuit. It should be resistant to the coolant and to the fuel at the operating temperature, and it should ensure an adequate heat sinking.

If the cladding becomes damaged, fuel may react with metallic sodium to give the oxides N a 3U 0 4 or N a 3P u 0 4 [12]. These compounds, being carried along with the coolant stream, may strongly contaminate the entire cooling system.

The fuel element surface in fast breeder reactors may reach temperatures above 873 K. The fuel cladding for these reactors is hence commonly made of stainless steels or of Inconel type nickel alloys.

At high temperatures the fuel may affect the cladding material by dissolution or diffusion. Frequently the fuel components form some low-melting eutectics with the cladding material, which quickly results in damage.

In fast breeder reactors uranium-plutonium fuel is used. An addition of plutonium to uranium or its alloys considerably lowers the melting point of the alloy, and the melting point of the eutectics of the fuel elements and cladding material. For example, the fuel alloy U-20 Pu-10 Mo forms a eutectic with stainless 304 grade steel at a temperature as low as approx. 873 Κ [23, 55-57]. The system is therefore unsuitable for fast breader reactors and requires inter-mediate layers to be provided to separate the fuel from the cladding. This solution has been adopted in the French reactor Rapsodie, for which the U-Pu-Mo alloy fuel elements were cast in thin-walled niobium tubes, which were then sealed in stainless steel cans. To ensure maximum heat exchange, the space between the niobium and the steel cladding was filled with pure sodium.

Recognized refractory materials suitable for operation in conjunction with the uranium-plutonium fuel are the Nb-1% Zr and Ta-0.1% W alloys. They are not, however, often used in fast breeder reactors in view of the fact that they unfavourably affect the performance of these reactors, and also are less resis-tant to liquid sodium than are stainless steels. The vanadium-based alloys, e.g. V-20 Ti are better in this respect [58, 59].

3.6 Effect of radiation on the properties of metals The main types of radiation in nuclear reactors are gamma rays and neutrons. Depending on their energy, the neutrons are classed into thermal neutrons (energy measured in eV), and fast neutron (energy measured in MeV) [60],

EFFECT OF RADIATION ON PROPERTIES OF METALS 119

The neutron radiation intensity is commonly expressed as the neutron flux den-sity, i.e. the number of neutrons that pass through a unit area in a unit time. The total radiation dose absorbed by a material is also frequently specified as the number of neutrons per cm

2 of surface.

TABLE 3.6

MECHANICAL PROPERTIES OF CHROMIUM-NICKEL STEELS UPON NEUTRON IRRADIATION TESTS AT 293 Κ (after [61, 62])

Irradia-tion Irradiation Tensile Yield Elongation

Steel tempera- dose, strength, strength, (A5), ture, η/cm

2 MPa MPa %

Κ

304, solution heat 592 235 63 treated at 1373 Κ 368 7 x l 0

19 714 518 58

and quenched in 343 1.8 x l O

20 696 — 47

water 353 2.43 χ 1 0

20 677 677 12

316, solution heat 657 294 70 treated at 1373 K, 1.0 x l O

20 834 657 45

water quenched 573 7 x l 0

18 775 353 57

and annealed at 743 1 χ 1 0

19 686 294 55

1123 Κ 873 1.1 χ 1 0

19 696 275 51

321, solution heat 637 235 66 treated at 1373 Κ 573 2.4 x l O

18 716 412 50

and quenched in 653 5.2 x l O

18 657 304 60

water 743 3.5 x l O

18 735 284 52

873 4.4 x l O

18 745 265 70

923 3.4 x l O

20 569 — 31

321, solution heat 923 3.4 x l O

20 618 — 25

treated at 383 Κ water quenched and annealed at 1123 Κ

H18N10T (GOST), 637 235 66 solution heat 373 3 χ 10

17 706 363 65

treated at 1373 Κ and quenched in 373 1.4 x l O

18 726 431 55

water


The main effects induced in metals by neutron radiation are radiation damage and fission processes. Radiation damage causes lattice point defects, disloca-tions, and vacancies to arise. Damage of this kind is accompanied by increased stress, swelling and deformation of the metal. A variety of products are formed in the fission reactions, including gases (He, H2) which may accumulate at the grain boundaries or in the voids.

These effects bring about changes in the mechanical and physical properties of materials, of which the most important from the viewpoint of a metal struc-ture is radiation-induced embrittlement. A drop in plastic properties is observed (percent elongation, necking) and a reduced impact strength with a simulta-neous rise in yield point, whereas the ultimate tensile strength of a metal does not always increase. Table 3.6 lists the effect of irradiation on the mechanical properties of chromium-nickel steels, based on the American and Russian data.

The two most important gaseous products of the fission reaction involved in the radiation-induced embrittlement effect are helium, produced by a nuclear transformation of boron and hydrogen, formed as a result of a nuclear transforma-tion of nickel. Accumulations of these elements are particularly hazardous if they occur at the grain boundaries, as they promote embrittlement of the metal. Radiation-induced embrittlement is generally an irreversible phenom-enon.

Once a dose of the order of 1022 n/cm

2 is reached in fast breeder reactors,

swelling of the materials becomes prevalent. This effect brings about deforma-tion and failure of the reactor structure [11]·

3.7 Corrosion control in sodium-cooled reactor systems

The corrosion rate of structural metals in liquid sodium was found to de-pend to a large extent on the purity of this metal coolant, and particularly on the concentration of oxygen contained in it. The major corrosion control measure, therefore, in sodium-cooled reactor systems is to ensure the maxi-mum possible purity of the cooling agent. The oxygen content in sodium can be reduceed (down to 15 ppm) by means of cold traps, using the variation of oxygen solubility in sodium with temperature (Fig. 3.8).

There are two kinds of such cold traps: one employs natural circulation (diffusion type traps), the other is based on forced circulation (by-pass type). Another technique for the removal of oxygen from the system, is the use of hot traps (oxygen is decreased to a level of 2-5 ppm) which incorporate getters, i.e. sodium-insoluble metals (zirconium, titanium) which readily form oxides

CORROSION CONTROL IN SODIUM-COOLED REACTOR SYSTEMS 121

— ,

1

>

CO

fr Ρ CD

. -i_

XL CO

ES 1 -4

73 K

i

ro c-* ro

1 1.4 1.6 1.8 2.0 2.2 2.4 2.6

10

3/r, K"

1

Fig. 3.8. Temperature dependence of solubility of oxygen in sodium (after [2]).

at high temperatures. Less frequently, oxygen-removing metals such as calcium, barium, or magnesium are added to the sodium; their oxides are subsequent-ly separated by filtration, decanting or by any other suitable mechanical pro-cedure.

A much more difficult method of corrosion control is the use of inhibitors. Substances suitable for this purpose should be capable of forming an insoluble film on the surface that separates the liquid from the solid metal. Thus a film of carbides or nitrides which are thermodynamically stable can be produced [2, 15, 16]; deposition of another metal is an alternative approach along these lines. The majority of corrosion control methods for structural materials in contact with liquid metals, involving the deposition or formation of protective films, are still at the laboratory testing stage.

Although coolant purity is a very important factor in corrosion control in sodium cooling systems, the structural materials also need to be critically con-sidered, especially as regards purity and smoothness of surface. To reduce the amount of non-metallic inclusions, vacuum-melted or electroslag-melted steels are used and the surfaces of finished steel products are thoroughly polished.


3.8 References

[1] Kummerer, K., Atomkernenergie, 9, Jg (1964). [2] Champaix, L., Energie Nucléaire, 8, 471 (1966). [3] Reactor Handbook—Materials, US Atomic Energie Commision, 1955. [4] Directory of Nuclear Reactors, Vol. X, Power and Research Reactors, IAEA, Vienna 1976. [5] Hausner, H. H., Modern Materials, Vol. 3, Academic Press, New York-London 1962. [6] Simnae, M. T., Fuel Element Experience in Nuclear Power—Reactors, Gordon and Breach

Science Publishers, New York-London-Paris 1976. [7] Palmer, R. G., Piatt, Α., Fast Reactors, Temple Press, London 1961. [8] Colle, R., Matériaux de Centrales Nucléaires, Technique Documentation, Paris 1972. [9] Gavze, M. N., Gudtsov, N. T., The Mercury Coolant Effect on Steel in the Power Systems

(in Russian), Izd. Akademii Nauk SSSR, Moskva 1963. [10] Walter, A. E., Reynolds, A. B., Fast Breeder Reactors, Pergamon Press, New York-Ox-

ford-Toronto-Sydney-Paris-Frankfurt 1982. [11] Weisz, M., Matériaux et Technique, 6, 1977. 112] Liquid Metals Handbook, AEC, Washington 1955. [13] Nevzorov, B. A. Corrosion of Structural Materials in Sodium (in Russian), Atomizdat,

Moskva 1968. tl4] Berge, P., Matériaux et Techniques, 9-10, 1977. £15] International Conference on Ferritic Steel for Fast Reactor Steam Generators, London UK,

30 May-2 June, 1977. [16] Mathern, G., Grolet, J. L., Séraphin, L. Tricot, R., Matériaux et Techniques, 9-10, 1977. [17] Thorley, A. W., Tyzack, T., Alkali Metals Coolants, IAEA, Vienna 1967. [18] Brush, E. G., Koenig, R. F., Materials and Methods, 42, 111 (1955). [19] Agapova, N. P. et al, Alkali Metal Coolants, IAEA, Vienna 1967. [20] Zebrowski, E. L., Young, R. S., Comprelli, F. Α., Dutina, D., ibid. [21] Champaix, L., Darras, R., Sannier, J., ibid. [22] Design, Construction and Operating Experience of Demonstration LMBRs, Proceedings

of a Symposium, Bologna, 10-14 April, 1978, IAEA, Vienna 1978. [23] Weisz, M., Berge, P., Rabbe, P., Matériaux et Techniques, 9-10, 1977. [24] Horsley, G. W., Reactor Technology, 1, 84 (1959). [25] Suzuki, T., Mutoh, L, Yagi, T., Ikenaga, Y., / . Nucl. Mater., 139, 2 (1986). {26] Champaix, L., Borgstedt, H. U., Eurocorr-87, Karlsruhe, FRG, 1987. [27] Vreeland, D. C , Hoffman, Ε. E., Manly, W. D. , Nucleonics, 11, 36 (1953). [28] Strauss, S. W., White, J. L., Brown, B. F., Acta Metallurgica, 6, 604 (1958). [29] Balandin, J. F., Gorinin, I. W., Zviezdin, J. I., Markov, W. G., Structural Materials for

Nuclear Power Stations (in Russian), Energoatomizdat, Moskva 1984. [30] Suzuki, T., Mutoh, I., / . Nucl. Mater., 140, 1 (1986). [31] Sheward, G. E., Welding J., 49, 7 (1970). [32] Horsley, G. W., Reactor Technology, 1, 84 (1959). [33] Zotov, W. W., Nevzorov, Β. Α., Umiashkin, E. W., Filin, M. P., Conf. Fast Reactor,

Dymitrovgrad 1972. [34] Cavell, I. W., Nicholas, M. G., / . Nucl. Mater., 95 ,1 -2 (1980). [35] Bagdasarov, J. E. et al., Technological Problems of Fast-Neutron Reactors (in Russian),

Atomizdat, Moskva 1969.

REFERENCES 123

[36] DiStefano, J. R., Litman, A. P., Corrosion, 20, 12 (1964). [37] Weeks, J. R., Klamut, C. J., Gurinsky, D. H., Alkali Metal Coolants, IAEA, Vienna 1967. [38] Borgstedt, H. U., Frees, G., Corrosion, 24, 7 (1968). [39] Chiotti, P., Wu, P. C , Fischer, R. W., / . Nucl. Mater., 38, 3 (1971). [40] Leipunskii, Α. I. et al.. Alkali Metal Coolants, IAEA, Vienna 1967. [41] Hajewska, E., Ph.D. Thesis, IBJ, Warszawa 1973. [42] Klueh, R. L., Corrosion, 27, 3 (1971). [43] Lashenko, W. S. et al., 2nd Intern. Confer on Peaceful Use of Atomic Energy, Geneva 1958. [44] Agapova, N. P. et al., Alkali Metal Coolants, IAEA, Vienna 1967. [45] Zebrovski, E. L., Young, R. S., Comprelli, F. Α., Dutina, D., ibid. [46] Champaix, L., Darras, R., Sannier, J., ibid. [47] Borgstedt, H. U., Dreschler, G., Frees, G., Z. Werkstofftech., 12, 7 (1981). [48] Borgstedt, H. IL, Grundmann, M., Eurocorr-87, Karlsruhe, FRG, 1987. [49] Garaud, J., Matériaux et Technique, 6, 1977. [50] Colburn, R. P., Meeting on Sodium Removal and Decontamination, IAEA, Richland,

14-16 Feb. 1978. [51] Beniere, F., Chemla, M., Aucouturiee, M., LaCombe, P., Corrosion, 24, 4 (1968). [52] Luner, C , Cosgarea, Α., Feder, Η. M., Alkali Metal Coolants, IAEA, Vienna 1967. [53] Valibus, L., Intern. Confer, on Sodium Techn. and Fast Reactors Design, Nov. 1968. [54] Watanabe, H., Kurusawa, T., Kikuchi, T., Furukawa, F., Nihei, I., / . Nucl. Mater., 40,

2 (1971). [55] Fast Reactor Technology, ANS, National Topical Meeting, Detroit, 26-28 April, 1965. [56] Palmer, R., Piatt, Α., Reactors for Fast Neutrons (in Russian), Gosatomizdat, Moskva

1963. [57] Kazachkovski, O. D. , Lytkin, Κ. B., Fast Power Reactors Atomic Energy Review, 3 ,4 (1965). [58] Greenberg, S., Ruther, W. E., Levin, Η. Α., Alkali Metal Coolants, IAEA, Vienna 1967. [59] Suzuki, T., Iwao, N., Noda, T., Ohno, K., Watanabe, R., / . Nucl. Mater., 79, 8 (1978). [60] Richards, C. W., La science des matériaux de Γ ingénieur, Dunod, Paris 1965. [61] Chyrkin, B .G, Thermophysical properties of Materials for Nuclear Engineering (in Russian),

Atomizdat, Moskva 1968. [62] Metals Handbook, New York 1961. [63] Barker, M. G., International Colloquium on Refractory Oxides for High Temperatures,

Energy Sources, Odeillo, France, 28 June-1 July 1977.

4. Corrosion in Molten Salts—A Thermodynamic Approach

4.1 Introduction

The study of electrochemical corrosion has been largely concerned with processes involving metallic materials in aqueous solutions of electrolytes, since water is far and away the most important liquid solvent which is stable in the temperature range of approx. 250-500 K. At higher temperatures, however, a large number of other inorganic compounds, such as oxides and salts, are encountered in this liquid state and, therefore, there are many liquid media in which metal corrosion processes may take place. Single- and multicomponent systems of molten salts and inorganic oxides have been used as media for vari-ous industrial processes for a considerable time. The knowledge and under-standing of their ability to corrode metals, however, is much less complete and developed than in the case of electrochemical corrosion in aqueous solutions.

Molten salt media are usually highly corrosive. Many modern technological processes pay the price for this, because although they find favourable kinetic conditions at high temperatures, on the other hand, the high electrical conduc-tivities of molten salts and high reaction rates of chemical processes at elevated temperatures also create good conditions for corrosion processes.

These electrochemical corrosion processes at high temperatures are many and varied and our knowledge about them is still developing.

4.2 Structural characteristics of molten salts: single- and multicomponent systems

An introduction to the study of electrochemical corrosion in aqueous electro-lytes usually starts with basic details concerning the physical properties of such media. In the same way, a description of metal corrosion processes in molten inorganic electrolytes must begin with an outline of the properties and structure of these liquids.

The term "molten salts" can be applied to the liquid phase of numerous inor-ganic as well as to some organic compounds. The properties of these substances

STRUCTURAL CHARACTERISTICS OF MOLTEN SALTS 125

in the liquid state are determined by the nature of interactions between the par-ticles of which they are composed. The real structure of these liquids may be referred to the following two limiting models :

1. A liquid composed of positive and negative ions, corresponding to a sol-ution of two liquid ionic plasmas with opposite electrical charges compensat-ing each other. The inter-ionic forces, both attractive and repulsive, are of a coulombic, long-range nature.

2. A liquid composed of molecules with relatively weak inter-molecular forces of the van der Waals type.

Real molten salts never correspond rigorously to either of these two extremes, since they contain elements of both types of interaction. Most of these liquids, however, are very similar to the first ionic plasma model. This group of molten salts is commonly called ionic liquids.

It is convenient to consider the structure of a real liquid in terms of either the gaseous state or the solid state. At temperatures near their critical points, liquids are often described in terms of a gas-like structural model, based on the concept of a random motion of molecules or other structural micro-species. Inter-molecular interactions in gases introduce an element of structural order-ing which increases, the lower the temperature and the higher the density. When, by a decrease of temperature, the substance passes its critical point, the struc-tural ordering becomes more and more distinct, but nevertheless, the nature of the liquid according to this model is best considered as the continuation of gas and vapour properties. The second model is formulated by analogy with crystal properties and compares the liquid structure to the ideal ordering of a perfect solid substance. In terms of this model, the liquid is considered as a crystal-like structure in which the ideal micro-species ordering has been strongly disturbed. It will be evident that this second model is more readily applicable to the case of liquids at temperatures near the triple points, i.e. near the melting temperatures of substances, than at temperatures in the critical point region.

In practice, the molten salts are used, as solvents and reaction media, at temperatures not much exceeding their melting points. The critical temperatures of many of them are not known at all. The imaginary model of the molten salt structure, therefore, should be based on the analogy with the solid phase struc-ture, rather than on a comparison with the respective gaseous phase properties.

The molten alkali metal halides most closely resemble the ideal ionic liquid model. As the electrical charge density of ions increases, the inter-ionic interac-tions become more purely coulombic. Hence molten lithium fluoride, where both

126 CORROSION IN MOLTEN SALTS

the cation and anion radii are the smallest of the alkali metal and halide ions, respectively, comes closest to the ideal.

Experimental studies of ionic structural configurations have been made on molten alkali metal halides. The existence of short range ordering has been demonstrated in these liquids. This means that the average distance between ions with opposite electrical charges is much less than the anion—anion and cation—cation distances and, therefore, that some character of residual ionic crystal lattice structure is preserved in the ionic liquid. Such quasi-lattice structures may also be found in other groups of molten salts which, in spite of lower ionic charge densities and higher polarizabilities of ions, may be also classified as ionic liquids. The ionic liquid model can be applied with good approximation to all the halides of alkali metals, alkali earth metals, and rare-earth metals. Similarly, the inorganic salts with alkali metal cations and oxy-gen-containing anions—carbonates, nitrates and sulphates—have a typically ionic chaarcter. Many physical properties of these salts, e.g. high values of the melting point and the heat of fusion, are evidence of a high contribution of electrostatic forces to the total potential energy of interaction between the struc-tural micro-species, and show that relaxation of the quasi-lattice structures in the fusion processes requires considerable energy.

The physical properties of typical ionic liquids show similarities to many of the respective properties of water, which is surprising when one considers the essential differences between ionic nature of the former and molecular nature of the latter. Viscosity coefficients, surface tensions and refractive indices are of the same order of magnitude in both water and ionic molten salts. As will be seen in Table 4.1, water resembles ionic liquids more closely than it resembles typical organic solvents. One physical property—the electrical ionic conductivity, de-termines the essential difference between the ionic liquids and the rest. In mol-ten salts it is higher by several orders of magnitude than even the conductivity of the concentrated water+electrolyte solutions. This feature is the deciding factor in the use of molten salts as media for high temperature electrochemical technologies involving either electrolysis or current generation processes. But at the same time, it enhances the electrode reactions in local electrochemical cells on the metal molten salt interfaces, and thereby makes the corrosion processes in these media very intense.

The properties and structures of single fused salts and oxides of metals of the various subgroups in the periodic table are different. In some cases these substances have a layer structure in the solid-phase, with a high contribution of van der Waals forces in the network configuration energy. The existence of

STRUCTURAL CHARACTERISTICS OF MOLTEN SALTS 127

TABLE 4.1

SELECTED PHYSICAL PROPERTIES OF SOME IONIC LIQUIDS, WATER AND BENZENE AT TEMPERATURE 25°C ABOVE MELTING POINTS

NaCl KC1 CdCl2 K N 0 3 N a 2S 0 4 Water Benzene

Fusion tempera-ture, + 25K 1098 1068 866 635 1187 298 279

Density, 10

3 kg/m

3 1.543 1.513 3.368 1.852 2.054 1.000 0.879

Refractive index 1.419 1.393 1.708 1.417 1.394 1.332 1.490

Viscosity coefficient, mPa · s 1.33 1.12 2.40 2.56 — 0.90 0.610

Surface tension, 10

3 N/m 112.2 99.2 84.1 110.1 193.3 72.0 29.0

Specific electrical conductivity, 10

3/ohm -m 3.654 2.224 1.93 0.706 2.354 6.7 10"

8 4 .010-

17

associates, clusters and "autocomplexes" has been revealed in these liquids, and it is very difficult to picture their precise configuration. Compared to typical ionic liquids the long-range coulombic forces give place to short-range, weak interactions between the structural associates and complexed; to different ex-tents however, in molten salts of different metals. These structural properties manifest themselves in macroscopic measurements. For example, the very high viscosity coefficient of molten zinc chloride is due to the very high degree of association in this salt's network up to 100° above its melting point. This coeffi-cient is higher than those of typical ionic liquids by three orders of magnitude. In molten cadmium chloride the network structure properties are not so pro-nounced, and the existence of a range of anion—anion distances in this liquid has been shown by Raman spectroscopy. The purely molecular structure of molten mercury—and beryllium halides has been shown by their very low elec-trical conductance. Numerous molten salts of other transition metals show high vapour pressures even at the fusion temperatures, revealing the molecular nature of these liquids. At normal pressure many of them are stable only over a very small range of temperatures, and some of these substances, e.g. aluminium chlor-ide, are not stable as liquids at all.


For technical purposes, the molten salts are used mainly in the form of multi-component systems. Distinct differences between the properties of these sys-tems and that of single salts should be emphasized. The molecular, associate-forming molten salts relax their network structures when mixed with ionic liquids. The ions with high charge densities increase the contribution of long-range coulombic interactions and hence increase the ionic character of the liquid. At high ionic molten salt concentrations, e.g. at more than 0.5 mole fraction of alkali metal halides in solution with a molten salt of the second type, structural rearrangements take place, consisting in the orientation of ligands around the ions of higher electrical charge density. New structural entities usually called complexes are formed. The stoichiometric compositions of these complexes sometimes correspond to the solid state compounds with congruent fusion points.

Consider for example the cadmium chloride/alkali metal chloride systems. Cadmium chloride in the liquid state shows a residual network structure and is supposed to contain a large fraction of the associate molecule—ions CdClj, CdCl

+, etc. Alkali ions when added in a small concentration, relax the structure,

resulting in an increase of surface tension in this range of concentrations. The change of this parameter is the greater, the higher is the electrical charge density of the alkali metal ion. In the region of high mole fractions of the alkali metal ions however, the relaxed C d

2+ ions become trapped by CI" ligands and complex

ions are formed, possibly CdClu-. Under these conditions the inter-ionic forces

are weaker and the surface tension thus decreases. This effect is greatest in case of caesium chloride: the low charge density around the large Cs

+ ions leaves

the Cl~ ions free to associate around the relatively small and double charged cadmium cations. This double influence of alkali metal ions on the structure of their solutions with associated molten salts may be observed in the dependence of other physical properties on concentration, mainly in the partial thermody-namic functions—partial free enthalpies and entropies. The absolute values of the excess functions are relatively high in the region of concentrations corre-sponding to complex formation. The partial entropies then depend on concen-tration according to complicated relationships.

Readers interested in the present-day knowledge of molten salt chemistry are directed to the following monographs and data reviews. Two monographs, one edited by Sundheim [1], the other by Blander [2], were published in the U.S.A. as early as 1964. In spite of some twenty years of further development in this field these books still represent important bibliographic positions. In the same year some results were reviewed in a book edited by Mamantov [3]. The editorial work of this author has been continued in four consecutive volumes

CLASSIFICATION OF METAL-MOLTEN SALT EQUILIBRIA 129

of a fundamental series entilted Advances in Molten Salt Chemistry [4-7]. The book by Bloom [8] is valuable as a student's handbook. A number of books are devoted to particular topics in molten salt chemistry. The thermodynamic prop-erties of molten salts and their solutions have been presented in a comprehen-sive monograph by Lumsden on the basis of the state of knowledge up to 1969 [9]. The electrochemistry of molten salts was the subject of a monograph in two volumes by Morand and Hladik in French [10], and four books in Russian [11-14]. Recent achievements in molten salt chemistry, especially in salts with oxy-anions are presented in Ionic Liquids, edited by Inman and Lovering [15]. Copeland [16] discusses the transport properties of molten salt media. A short introduction to the statistical thermodynamics of molten salts has been given by Braunstein [17]. The monograph by Volkhov and Yatsimirskii in Russian is devoted to the spectroscopy of molten salts [18]. Some properties of molten salts as solvents are reported by Chariot and Trémillon in French [19]. The publications containing tabulated data of molten salt physical properties by Janz et al. [20-30] are of particular value in the development of this field of chemistry. The contribution of Russian authors to the compilation of phase diagrams for molten salt double and multicomponent systems is most useful [31-36]. The phase diagrams of binary systems of salts with organic anions are reviewed by Franzosini et al. [37].

4.3 Classification of metal-molten salt equilibria

The foregoing short review of molten salt structure and properties shows that above all, these liquids are good electrical ionic conductors. Corrosion phenom-ena taking place on metal-molten salt interfaces are of an electrochemical na ture and hence they can be compared with metal corrosion processes in aqueous electrolyte solutions. The values of the reversible electrochemical potentials in these latter systems are classified with respect to the potential of the standard hydrogen electrode. Such an unified ordering of all the various metal/molten salt interface potential values is evidently impossible, the potential series being as numerous as the molten salt solvents. For single molten salts the potential series may be established on the basis of standard free enthalpies of their de-composition reaction. For example, for the halide of a given metal Me such a reaction is

where X denotes F, CI, Br or I, and ζ is the metal valency. Two reactions

MeXz -> Me + z/2X2 (4.1)

M ez+ + ze~ ^ Me (4.2)


and

2X- ^ X 2 + 2e- (4.3)

determine the electrochemical redox potentials of the metal and of the halide:

4e = E&9jM++ + 2.3(RT/zF)log[aM̂(aMer>) (4.4)

and

Ex2 = El,x- + 23(RT/zF)\og[aX2 (αχ-)~2] (4.5)

respectively, where E° denotes standard potentials, and a is the activity of the respective species. The difference between the two standard potential values is connected with the standard free enthalpy AG° of reaction (4.1) by the relation

AE° = Ε°2-ΕΪ = -AG°/zF (4.6)

The standard electrode potentials for the anion forming substances may be conventionally considered as equal to zero. In this way the standard electrode potential series can be defined on the basis of thermal decomposition reactions which are, for chlorides :

MeCl2 -+ Me + z/2 Cl2 (4.7)

for carbonates:

Me(C03) , /2 -+ Me + z /2(C02 + 1/2 0 2) (4.8)

and for nitrates:

Me(N03)z -> Me + z ( N 0 2 + 1/2 0 2) (4.9)

and so on. In the same way, the electrochemical potentials of metals is equilibrium

with ions dissolved in multicomponent solutions of molten electrolytes can be

classified. The concentration of ions may be expressed in mole fractions, or

according to the Temkin's definition, in ionic fractions. Using this definition,

the ionic fraction of cations formed by dissociation of the i-th solution compo-

nent of the molten salt system is

where nci is the number of moles of particular cations, and nc is the number of moles of all the cations in the system. Similarly, the ionic fractions of anions of the f-th molten salts component is

/=i

TABLE 4.2

STANDARD ELECTROCHEMICAL POTENTIALS OF METALLIC ELECTRODES IN WATER AND IN SOME MOLTEN CHLORIDE SYSTEMS (after [38])

-E° (experimental), mV

-E°, mV water 298 Κ

- £ ° , m V (calculated)

MeCl2 1073 Κ NaCl + KC1

1073 Κ

MgCl + + NaCl +

+ KC1

LiCl + + KC1 673 Κ

748 Κ

LiCl + + KC1 673 Κ

1 2 3 4 5 6

L i

+ 4405 3457 3626

K

+ 4285 3441 — — —

C a

2+ 4230 3323 — —

N a

+ 4074 3240 — —

M g

2+ 3730 2460 — 2722 2796

U

3+ 3160 2280

3

— 234 2469 A l

3+ 3020 — 1990 2013

Z r

4+ 2890 1270

3

2010

2

— 2076 M n

2+ 2540 1807 2010 1870 2065

C r

2+ 2270 1385 1560 1472 1641

Z n

2+ 2123 1476 1655 1626 1782

C r

3+ 2100 1113 1205 1235 1376

F e

2+ 1800 1118 1330 1259 1388

C d

2+ 1763 1193 1400 1388 1532

T l

+ 1696 1473 1485

2 — 1586

C o

2+ 1637 977 1120 1046 1207

N i

2+ 1610 865 0950

2 868 1011

S n

2+ 1496 1259 1175 1223 1298

P b

2+ 1486 1112 1195 1247 1317

F e

3+ 1396 — 956 1004

H 2 1358 1050 1050

1 1027

1 1016

C u

2+ 1023 400 640 595 619

C u

+ 839 970 1075 939 1067

A g

+ 560 826 820 756 853

H g

2+ 506 290 — 700

P d

2+ 373 331 281 430

P t

2+ 160 150 76 216

C l 2

4 0 0 0 0 0

A u

+ - 3 2 0 — — - 2 0 1 - 0 9 5

F e

3 +/ F e

2+ 588 — — 350 236

C u

2 +/ C u

+ 1207 - 1 7 0 205 251 171

C r

3 +/ C r

2+ 1760 570 495 761 847

1 Calculaced JG°,

2 extrapolated,

3 solid salt,

4 reference electrode.


The pure /-th component constitutes the reference standard state for its solution. In the reference state

Xi = Xci = Xai = 1

The standard potentials are determined by extrapolation of the dependence of electrode potential on ion concentration, to the standard state.

Ketelaar has compiled some of standard potential series in different molten chloride solvents [38]. These series are reproduced in Table 4.2. The chloride electrode, reversible with respect to chloride ion concentration, is recognized as the reference electrode for these systems. As can be seen, the positions of some metals are different in particular electrochemical series.

On the basis of possible chemical reactions taking place between metals and molten salts one can distinguish three groups of thermodynamic equilibria occur-ring in these systems:

Group L Equilibria between the given metal and the molten salt of another one, determined by the difference in the normal electrochemical potentials of these metals.

Group 2. Equilibria between the given metal and its own molten salt, e.g. halide, consisting in the formation of a solution of metals in molten salts.

Group 3. Equilibria between the given metal and gases dissolved in the molten salt solvent, as well as between the metal and complex anions of the given salt.

Let us classify reactions leading to equilibria in group 1 and to some equilibria in group 3, as "corrosion processes". A distinct solubility of metals in molten salts as in group 2 precludes any anticorrosion resistance of the metal. In prac-tice, only such metallic materials may be used in contact with molten salts whose standard potentials are much more noble than the same parameter of metals forming cations of the molten electrolytes. In other cases the dissolution of metals in these media is catastrophic, the rate greatly exceeding that of corrosion pro-cesses. The difference between the standard electrode potentials also determines the behaviour of the local electrochemical cells formed in heterogeneous metallic materials in contact with molten electrolytes.

4.4 Dissolution of metals in their molten salts

Although the system of a metal+its molten salt is not directly relevant to metal corrosion phenomena, it needs to be mentioned, however, in this review be-cause it represents one of the specific features of such systems and because the metal + salt solutions show an extremely high corrosion aggressivity against other

CORROSION EQUILIBRIA IN MOLTEN ALKALI METAL HALIDE 133

materials such as ceramics. The nature of these liquid solutions has been exten-sively investigated, but still is not completely clear. The free enthalpies of mixing of the metal+salt solution cover a wide range of values. In some cases, for example for cadmium forming a solution of concentration up to 20 mole per-cent of the metal dissolved in molten cadmium chloride, the energy of inter-action between the salt solvent and the metal solute is high, and a complex ion-molecule formation is claimed. In this solution the dissolved metal exists at a valency lower than those of the cations in the pure salt. This may be pre-sented as the equilibrium

Cd + C d2 +

^ C d i+ (4.10)

In other cases, for instance in systems of alkali metal halides, the interaction between the two components is much weaker and the nature of the solutions may be compared more to the "physical" atomic metal+molten salt solution. An electronic component of the electrical conductance appears in this latter group of liquid systems.

The effect of the solubility of metals in their salts influences the electrode potentials of metals in these media and should be taken into account in the formal thermodynamic calculations of their potential values.*

4.5 Corrosion equilibria in a pure molten alkali metal halide

The stability of a given molten salt, e.g. of a divalent metal chloride, is deter-mined by the free enthalpy of the thermal dissociation according to reaction (4.7). This equilibrium is usually rapidly established at high temperatures, determining the redox potential of the salt as a solvent for the metal and for chloride. In the equilibrium constant

<*Me * tfCl2A*MeCl2 (4.11)

one can put aMeCi2 = 1 with good approximation, because the salt decomposes to a very low extent. The activity of the metal in a condensed phase is small, and the activity of chlorine is determined by its partial pressure.

The electrochemical redox potentials of the metal and of the chlorine in the molten MeCl2 are determined by the redox equilibria (4.2) and (4.3), respect-ively, and may be expressed respectively by formulae (4.4) and (4.5).

* The interesting problem of the metal-fits salt solution structure and properties has a large bibliography. Chapters by Corbett, Bredig and Warren in monographs [1], [2] and [7] are the most comprehensive reviews of this topic.


The standard electrochemical potential in formula (4.5) equals zero, accord-ing to the convention considering the halogen electrode as the reference elec-trode. In the equilibrium state the redox potentials of any other substance dis-solved in molten chloride are equal. This is demonstrated on diagrams by Ede-leanu and Littlewood [39] as in Fig. 4.1. The diagram shows the redox electro-

chemical potentials of the metal + its molten chloride system for the case of fused potassium chloride as the solvent and medium of corrosion processes at 1073 K. These potentials have been calculated according to equation (4.4) and are represented as functions of log[aM ez

+/aM e] (scale A). Scale Β corresponds

to the concentration range of the dissolved metals, when aMe*+ = 1? a n

d scale C corresponds to the concentration of the pure metal phase in equilibrium with the metal + salt solution, when a Me = 1. Scale D represents the pressure over the system, when CI"" is the only anion in the melt: ac l- = 1. Scales Β and C are denoted by symbols p, defined as ρ = - l o g ^ , by analogy to the pH scale used for the aqueous solutions.

Edeleanu and Littlewood discuss the case of nickel in the pure, oxygen free potassium chloride. For the equilibrium state to be established the redox po-tentials of metallic nickel and potassium dissolved in molten KC1 should be-

classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which included slow growing pits (pit diameter classed into two

CORROSION EQUILIBRIA IN MOLTEN ALKALI HALIDE 135

come equal, therefore some concentration of N i2+

must be formed in the mol-ten KC1 by the reaction

The metallic potassium produced in this way dissolves in the melt. Its concentra tion is twice that of the N i

2+ ion concentration:

The values of the redox potential corresponding to the N i /N i2+ and K/K+

coupled equilibria, can be readily determined from Fig. 4.1 because the re-spective lines are separated by a distance equal to log 2 at this point. The line parallel to the concentration scale A intersects both Ni and Κ lines at a point distant by log 2, determining the N i

2+ concentration in equilibrium with nickel

in the KC1 bath under these conditions. We can read this concentration on the scale C and it turns out that the content of N i

2+ in KC1 is very small. Accord-

ing to the definition by Pourbaix [40], a metal is considered corrosion resistant when the concentration of its cations does not exceed 10"

6 mole fraction. We

can conclude, therefore, that nickel does not corrode in molten KC1 when this solvent is absolutely dry and oxygen-free. As can be seen in Fig. 4.1, chromium corrodes to a greater extent.

For all these systems, the range of electrochemical redox potential values is limited. The potential of the 2 C1~/C12 electrode equals zero and represents the highest value. The lowest value corresponds to the lowest Cl2 activity as calcu-lated from the dissociation equilibrium (4.11) for different vapour pressures of the metal dissolved in molten salt. The metal vapour pressure depends on the total pressure of the chloride vapour, therefore the negative potential boundary is a function of the volatility of the molten chloride.

4.6 Corrosion equilibria in molten alkali halide containing oxide ions

The preparation of molten alkali halide absolutely free from both water and from oxide ions is extremely difficult. Moreover, atmospheric oxygen and water are not the only sources of oxide ions which may be found in the molten solvent; these ions may also come from the reaction of the molten solvent with the re-fractory materials of the container.

Ni + 2K+ = N i2+

+ 2K (4.12)

[Ni2 +

] = 1/2 [K]

Assuming the ideality of this solution we have

p N i2 +

- / ? K = log2

(4.13)

(4.14)


The standard redox electrochemical potential of the oxide ions in molten potassium chloride, as determined by the equilibrium

%02+2e~ τ±02~ (4.15)

may be calculated from the difference between the respective standard potentials of the reactions

2 Κ + ^ 0 2^ 2 Κ+ + O

2

and (4.16)

(4.17) K ^ K+ + e~

The electrochemical redox potential determined by reaction (4.15) is

Eo2io>- = ~2.38 + 0.106 [p02~-%p02] (4.18)

at 1073 K, in volts. Figure 4.2 represents the dependence of this potential on

p 0 2

7 V r

I /

A

Ί y

1 / I / 1 /

y-Y-—

Ni

2' /

NiO /

"~-f—~

NiO / s /

/ / A / / s /

• / / /

' Ni y

10 20 30 p o

2\

- 2 C f / C l2 +40

^ , p N i

2 += 0

+60

^ p N i

2' = 6

K/K

A0

Fig. 4.2. Stability diagrams of nickel in molten potassium chloride containing 0

2~ ions, at

1073 Κ (after [38]): Ni/NiO stability boundary.

the pO*~ parameter. The various lines correspond to different concentrations of oxygen dissolved in the salt. The diagram covers potentials between the zero value corresponding to the standard chloride electrode and the potassium elec-trode standard potential.

The electrochemical potential of nickel in equilibrium with molten KCl must take the value of the redox potential of the solvent as determined by the dis-

CORROSION EQUILIBRIA IN MOLTEN ALKALI HALIDE 1 3 7

solved oxygen and oxide ions concentration. This concentration also determines the N i

2+ ions content according to the solubility product of nickel oxide in

molten KC1. According to the equilibrium

NiO( s) ^ Ni(

2

K

+

c l) + 0(

2KCI) (4.19)

(where the indices (s) and (KC1) denote the solid phase and species dissolved in the molten "potassium chloride respectively), the solubility products may be calculated from the difference of standard free enthalpies of the reagents in equilibria (4.2) and (4.15) for Ni and 0 2, respectively, as

InS = (AG&i0-AG&i2+-AG&-)lRT = ln(aNi*+ · a0>-) (4.20)

The value of the solubility product for NiO in molten KC1 is low, being of the order of 1 0 "

22 mole fraction. On the basis of this value one can determine

the region of stability of solid nickel oxide in this medium. This region is re-presented on the diagram in Fig. 4.2. According to Pourbaix's convention the region of the susceptibility of nickel to corrosion starts from N i

2+ concentra-

tions equal to 10"6 mole fraction and this value determines the boundary of the

nickel corrosion resistivity on the diagram. The dotted vertical line in Fig. 4.2 denotes the corresponding O

2" ions concentration. The line determining the

boundary of the Ni and NiO stability passes through the intersection points of the constant potential value (horizontal lines) with the vertical lines of the re-spective O

2"" concentration values. The equation for the Ni/NiO boundary has

been determined, when expressing the nickel ions concentration in the formula for the nickel electrode potential

Em = -0 .856-0.10 6/?Ni2+ (4.21)

by the oxide ion concentration and the nickel solubility product. The region above the Ni/NiO boundary corresponds to the stability of NiO. In the appropri-ate conditions the nickel immersed in the molten KC1 became oxidized and covered by a solid NiO layer. At concentrations of 0

2~ higher than those deter-

mined by the Ni/NiO line, undissolved NiO rather than dissolved N i2+

is the product of the metallic nickel corrosion. In other words the Ni/NiO boundary determines the oxygen pressure above the NiO at a given temperature. When the 0 2 pressure above the molten salt exceeds this dissociation pressure of the metal oxide, the oxidation of metal proceeds. This process may be interrupted by the formation on the metal of a continuous passivation layer of NiO, imper-meable to oxygen, and the system may achieve a quasi-stable state. The redox potential of the solution is then determined by the limiting O

2"" concentration

value and by the 0 2 pressure in the atmosphere, whereas the metal potential is


determined by the concentration of the saturated N i2+

solution. The difference existing between the two potentials corresponds to an unequilibrated state and this occurs only when the metal oxide layer on the metal has a distinctly passive character.

The aim of the above discussion is to show the important role of oxygen dissolved in molten salt in metal corrosion processes in this medium. The dissol-ution of this gas from the atmosphere may be relatively easily avoided by the application of an inert gas atmosphere. The content of oxide ions in the salts is however mainly due to oxides in the refractory container materials. The most complicated case is when the oxide ions in the salt are a product of the decompo-sition of water. This impurity is the most difficult one to avoid, and it consti-tutes the main source of corrosion processes in molten salt media.

Let us describe the influence of water on the redox potential of molten salt solvent by the ionization processes

The equilibrium constant of water ionization may be calculated in the same way as used in NiO solubility product determination. At 1073 Κ the thermal disso-ciation equilibrium constant of water, Kd, is 1 0 "

9 1 8. From the thermal disso-

ciation constant of HC1, we obtain the ionization constant for water dissolved in KC1 at this temperature

Using this value we can discuss the conditions of nickel corrosion in KC1 in the presence of water, using the triangular coordinate system, as in the paper by Edeleanu and Little wood [39].

In Fig. 4.3 a set of triangles with common sides is drawn. The particular

triangles represent the following equilibria:

Triangle /—ionization of water

pH20 = -21.5 + 2pU+ + p0

2~

Triangle 2—redox potential of the solution

^redox = -2 .38 + 0.106[/>O2~-iO2]

Triangle 3—dissociation of water

pR2Q = - 9 . 1 8 - p i J 2- i p 0 2

H 2 2 H+ + 2e~

and equation (4.15), and the thermal dissociation reaction

H 2O ^ H 2 + i 0 2

(4.22)

(4.23)

Κι = (αΗ+)2 - 0 6 - • Κ ο Γ 1 = 10

2 1·5

(4.24)

CORROSION EQUILIBRIA IN MOLTEN ALKALI HALIDE 139

p 0 2 pH20 10

P H 2 5 15

Fig. 4.3. Relations between redox electrochemical potential values and 0 2, H 2, H 20 , H

+ and

0

2~ activities in molten potassium chloride at 1073 Κ (after [39]).

Triangle 4—redox potential of the solution

i W = Ell2H+ + 0.106r/>H2-2/>H+]

Edeleanu and Littlewood used this method to solve the following problem by a graphical method: how to avoid the corrosion of nickel in molten KC1 containing dissolved water at a temperature of 1073 K. In other words, how can the concentration of metal ions be kept below 10~

6 mole fraction by using an

appropriate concentration of hydrogen in the system. Using Fig. 4.2, this con-centration of N i

2+ ions corresponds to the value of —1.53 V of the nickel elec-

trode potential. Therefore, the potential of the whole system should not exceed this value, denoted in Fig. 4.3, in triangles 1 and 3 by the lines A and B, respec-tively. In Fig. 4.2 it is also shown that the formation of NiO is possible only at p02 below 13.0. In Fig. 4.3 this value is denoted in triangles 2 and 3 by the lines C and D, respectively. Thus line Ε in triangle 2 determines the p0

2~ value

above which nickel cannot be oxidized. Further discussion depends on which parameter is established as constant. If, for example, the concentration of water in KC1 is determined to be 10""

6 mole fraction (lines F and G), the other data

would be deduced from the diagram by the lines H91, / , Κ and L. Line Η indi-cates the lowest hydrogen pressure one must use to protect nickel against cor-rosion and the pure molten KC1 solvent against the N i

2+ impurities. The higher


the hydrogen pressure, the lower is the criterion for the dehydration of the mol-ten KC1. The concentration of water, however, must not be higher than 1 0 "

14

(the dotted line in triangle 1) because above this value, the application of hydro-gen pressure cannot prevent the oxidation of nickel.*

This example shows the very stringent conditions for the dehydration of molten salt media, when a complete elimination of their corrosion aggressivity with respect to metals is required.

4.7 Corrosion equilibria in molten carbonates

Although the redox potentials of molten chloride media are mainly determined by the impurities dissolved in them, in the molten salts of containing oxy-anions —carbonates, nitrates and sulphates—the conditions enabling corrosion pro-cesses may exist even in the purest state of these liquids. Consider, for example the equilibrium of nickel with the molten lithium 4-sodium+potassium carbon-ates ternary system at 873 K. This system is important, for instance, as the electrolyte in modern electric power generating devices called Molten Carbon-ate Fuel Cells (MCFC), [41, 42]. In this type of generator the electric current is directly obtained by the conversion of chemical energy of oxidation of gas-eous fuels containing hydrogen and carbon monoxide as well as gaseous hydro-carbons. Such conversions involve the electrochemical anodic oxidation of fuel on a nickel electrode. The redox potential of the molten carbonates elec-trolyte is determined by the electrochemical equilibrium

CO?T ^ C 0 2 + i 0 2 + 2e~ (4.25)

The reference electrode for the molten carbonate media is usually defined as a noble metal electrode immersed in the molten carbonate bath, in equilibrium with the gas mixture of gases containing carbon dioxide and oxygen with the partial pressures stoichiometric ratio | C 0 2 to ^ 0 2 under 0.1 MPa total pressure.

The redox electrode potential of the molten carbonate solvent is

2.3(RT/2F) [pC02 + tp02] (4.26)

This potential equals zero when pC02 and p02 take values corresponding to the reference electrode gas mixture. Therefore, the normal potential is

£ & > r / o a. c o a = - 2 . 3 ( « r / 2 f ) l o g ( 2 / 3 ] / 3 ) (4.27)

* It should be emphasized that throughout this chapter all parameters are calculated as-

suming the ideality of solutions.

CORROSION EQUILIBRIA IN MOLTEN CARBONATES 141

The 02~ ion concentration in molten carbonates depends on the C 0 2 pressure

due to the acid-base equilibrium:

M e 2C 0 3 ( 1) ^ M e 20 ( 1) + C 0 2 ( g> (4.28)

with an equilibrium constant:

P^ca-b) = p O2- + PC02 (4.29)

In the (Li, N a ) 2C 0 3 eutectic solution, p#( a-.b) equals 7.72 at 873 Κ [43, 44]. Figure 4.4 presents the redox potential scale in this solvent, calculated by the formula (4.26). The upper scale denotes ρΌ

2~ values, the lower one, pC02.

p0< 10 15 1

κ

1 1 1 1

I ^y^>

! Ni

2* y

1 ' M ι

^ 1 NiO !

^ 1 NiO !

*s

* y y

^ S Ni S y'

y'

pNi

2 t=0

ρ Ni

2*=6

Να/Να

pco2

Fig. 4.4. Stability diagram of nickel in a molten lithium/sodium carbonate eutectic mixture at 873 K. Ni/NiO stability boundary, solid carbon deposition boundary.

The Ni stability area in molten carbonates is denoted in Fig. 4.4 in the same way as in Fig. 4.2. The nickel carbonate decomposition equilibrium is

N i C 0 3 (0 ^ Ni( s) + C02 ( g )(2 /3 MPa) + i 0 2 ( e )( l / 3 MPa) (4.30)

The standard free enthalpy of this reaction is 6.33 χ 104 J/mole at 873 Κ [45],

which gives the N i / N i2+ standard potential value of —290 mV. The nickel elec-

trode potential in molten carbonates is, therefore, Em = -0 .29-0 .087pNi

2+ (4.31)

The redox potential function denoting the NiO stability region at 873 Κ is determined by the pressure of oxygen resulting from the dissociation reaction

N i O ( s )^ N i ( s) + i 0 2 ( e) (4.32)


the standard free enthalpy of which equals 1.698 χ 105 J/mole [45]. In Figure 4.4

the diagonal lines correspond to the p02 values. The Ni/NiO boundary calcu-lated using the respective values in equation (4.26), corresponds to p02 equal

From equations (4.26) and (4.31) one obtains the relation for the concentra-tion of N i

2+ ions in equilibrium with nickel oxide as dependent on the partial

pressure of C 0 2 in the system

The corrosion agressivity of the molten carbonate solvent in respect to a given metal (nickel in this case) can hence be deduced from Fig. 4.4. A complete unsusceptibility of the metal to corrosion is conventionally defined as the re-sistance to forming an equilibrated solutions of ions with a concentration higher than 10"

6 mole fraction. In the range of concentrations 2 < pNi

2+ ^ 6 one

observes the electrochemical corrosion sensitivity of the metal. At /?Ni2+ < 2

the dissolution of metal in the molten medium is catastrophic, but in the case of nickel this occurs at very high C 0 2 pressure values.

The oxidation of nickel in molten carbonates can hardly be avoided. At pC02 = 1 NiO is formed at solvent redox potentials above - 9 0 0 mV at 873 Κ and for protecting the metal against the formation of a solid NiO layer, a large negative voltage should be applied. The range of accessible negative potentials, however, is rather narrow. The boundary of solid carbon deposition due to dissociation of carbon dioxide

is the limit of the molten carbonate solvent stability. This boundary is also shown in Fig. 4.4.

The discussion of the metal corrosion equilibria in molten alkali carbonates is not complete if it is limited only to the behaviour of these media in a com-pletely dehydrated form. Such a state of molten carbonates corresponds to an unrealistic and theoretical situation never encountered in technical practice. The use of a molten carbonates electrolyte in the high temperature fuel cells mentioned at the beginning of this paragraph, furnishes an important example of corrosion equilibria involving molten carbonate solvents containing water.

* Ketelaar [38] and Ingram and Janz [46] report other values of the N i / N i

2+ electrode

standard potential: —340 mV, and of the pNi

2+ in pure C 0 2 atmosphere: 5.26. These values

result from the AG

0 values used by these authors: 6.57χ 10

4 and 1.56-»10

5 J/mole for nickel

carbonate and oxide dissociations, respectively. Moreover, the small value of £cor /o2, co2, 36 mV at 873 K, has probably been neglected by these authors.

to 19.12.

/>Ni2+ = 5.77 + pC02* (4.33)

C 0 2 ( g) -> C(S) -f- 0 2 ( g) (4.34)


The MCFC power generators are supplied with gaseous fuel containing hydro-gen and carbon monoxide (this fuel is usually called the CHO gas), and water and carbon dioxide constitute the products of electrochemical oxidation of fuel in such a cell. Methane is the fifth CHO gas component, its concentration, however, is rather low and is neglected in this discussion. In the corrosion equi-libria taking place in the MCFC type generator, namely the corrosion of nickel, iron, cobalt, chromium, manganium, etc—the main materials for the fabrica-tion of porous MCFC electrodes as well as of the alloys used in the fuel cell housing—water does not take part as an impurity in the carbonate melt but rather as an essential component of the molten electrolyte. The equilibrium constant value for the reaction

M e 2C 0 3 ( 1, + H 20 ( g) ^ C 0 2 ( g) + 2MeOH( 1) (4.35)

pH20 •

-20 -15 -10 - 5 0 5 10 15

P H 2 Û

Fig. 4.5. Relations between redox electrochemical potential values and CO, C 0 2, H 20 OH", O

2- , H 2 and 0 2 activities in a molten lithium/sodium carbonate eutectic mixture at 873 K.

— Ni/NiO stability boundary, — solid carbon deposition boundary.


in the case of the (Li, N a ) 2C 0 3 eutectic mixture is 9.72 χ 10""5 at 873 Κ [43]. This

acid-base equilibrium obliges us to consider the corrosive medium of the MCFC electrolyte as a solution of carbonates + hydroxides -f oxides, rather than as a pure carbonate melt. The presentation of the corrosion equilibria in the Pour-baix type diagrams, as in Fig. 4.2 and 4.4, needs to be replaced by an appropriate modification of the diagrams by Edeleanu and Littlewood of the type presented in Fig. 4.3. Let us therefore consider the four relevant equilibria at 873 Κ and represent them on a set of four triangles (Fig. 4.5):

Triangle 1—hydrolysis of molten carbonates [43] ρϋ20 = - 4 . 0 + pÇ02 + 2pOR-

Triangle 2—oxidation of carbon monoxide [45] pCO = 1 2 . 0 - i p O 2 + pC02

Triangle 3—dissociation of water [45] pU20 = -12.0 + / > H 2- i p 0 2

Triangle 4—oxidation of hydrogen and carbon monoxide [41] pCO = 4.0 + /?H2-2/>OH-*

In triangle 2 the redox electrochemical potential scale corresponding to for-mula (4.26) is also presented. In all these equilibria the approximation pCOl~ = 0 is made.

TABLE 4.3

STANDARD ELECTROCHEMICAL POTENTIALS OF METALLIC ELECTRODES AND pMe

z+ VALUES AT pC02 = 0 FOR METAL

OXIDES, IN MOLTEN CARBONATES AT 873 Κ

System i ? M e / M «s + »

calculated [45]

/>Me

z +,/>C02 = 0

calculated [45]

Fe/FeO/Fe

2+ - 6 6 3 4.77

Ni /NiO/Ni

2+ - 2 9 4 5.77

Co/CoO/Co

2+ - 3 9 9 5.27

Zn/ZnO/Zn

2+ - 8 5 2 5.37

Mn/MnO/Mn

2+ - 1 3 6 2 1.37

Ag /Ag20 /Ag

+ + 532 1.65

1

1 Calculated from data given by Janz et al. [47].

* The values of the constants in these equations have been estimated using the data reported

in the cited works.


Let us draw in triangle 2 the metal/metal oxide and carbon deposition bound-aries. As in Fig. 4.3, one can estimate the minimum hydrogen partial pressure sufficient to protect metals against the formations of solid oxide, at a given con-centration of water or OH" ions in the melt. The data used to construct this dia-gram correspond to the Ni/NiO/Ni

2+ system. In Table 4.3 are given the relevant

data necessary to construct similar diagrams for other metal/metal oxide/metal ions equilibria in molten carbonates. The triangle 2 in Fig. 4.5 shows the same relationships as the one in Fig. 4.4, but in different coordinate system. We will take into consideration the triple point of NiO/Ni boundary with /?Ni

2+ = 6

line intersection. At this state of the system the metallic nickel may still exist, whereas the corrosion product N i

2+ ion exists at a permissible concentration.

Let the C 0 2 partial pressure be less than 0.1 MPa, the pC02 value is then a little higher than 0. If the oxygen partial pressure is very low, this state of the system may be reached when the CO partial pressure is of the order of —3. The electrochemical redox potential is then approx. —800 mV and one can consider metallic nickel to be protected against oxidation and dissolution in the melt. Higher oxygen and/or lower carbon monoxide partial pressures will re-sult in metal oxidation.

The equilibria represented in triangles / and 3 permit us to discuss the same problem in presence of water in the system. Values of pC02 and p02 are pro-jected to these triangles, whereas the pCO value is projected to triangle 4. Con-sidering some water vapour partial pressure (in Fig. 4.5 pH20 of approx. 3.5 is assumed), at the given electrochemical potential the system is determined and the hydrogen partial pressure may be graphically estimated. The protective hydrogen partial pressure must be the higher, the higher is water vapour content in the system.

To maintain the system in the state corresponding to the triple point we have to control 4 among the 5 following parameters: the electrochemical redox po-tential and /?(/), where i is C 0 2, CO, 0 2 and H 2.

The diagrams in this paper present the data using logarithmic scales. In case of Fig. 4.5, for example, the diagram is not precise enough to furnish data con-cerning the equilibria between the CHO gas, the molten electrolyte and the solid metal and metal oxide phases, as required for operational purposes for the over-all MCFC power generation system. The reader interested in this field will find the relevant data in an atlas of equilibrium diagrams [48]. Figure 4.5, however, enables a good estimation of corrosion condition range for metallic materials working in the C 0 2, H 20 , CO and H2-containing gas mixtures ("water gas",


"synthesis gas", etc.), in each case in the presence of molten alkali carbonates. Such compounds may be easily formed in the reactions of fossil fuel combustion products with minerals containing different alkali metal compounds.

4.8 Corrosion equilibria in molten sulphates. "Hot corrosion"

The system considered in this section, metal + molten alkali sulphate, is the most widespread in industrial practice and perhaps the most important from the practical point of view. Sulphur dioxide is formed in enormous amounts in the combustion of sulphur-containing fossil fuels and, unfortunately, is largely oxidized to trioxide on the catalytically active surfaces of metal oxides in the tubes and collectors of combustion devices. Vanadium pentoxide present in the ashes of many of these fuels plays a special catalytic role in the sulphur (IV) to sulphur (VI) oxidation. The ashes also contain some sodium chloride, which eventually leads to the formation of sodium sulphate in the outlet collectors of the industrial installations. Sodium sulphate, being liquid at temperatures of the overal processes, gives a paste-like product with the solid particles of the ash deposits. The aggressivity of these deposits to metallic materials is related to the electrochemical corrosion phenomena in molten salts media.

Due to the industrial importance of this type of corrosion, commonly called hot corrosion phenomena, this problem has been extensively investigated. The modern concepts concerning its mechanism are reviewed in the literature [49]. Three mechanistic hypotheses have been proposed. According to some authors [50, 51] the reduction of the liquid N a 2S 0 4 leads to the deposition of solid sulphide phases on the metallic surfaces. These phases destroy the oxide pro-tective layers on the refractory, stainless steel outlet gas collectors, rendering them unresistant to corrosion. Other authors [52-54] suppose that these pro-tective layers are directly dissolved in the molten sodium sulphate, and a third explanation of the corrosion processes in refractory steel tubes emphasizes the role of chloride ions present in the melts based on sodium sulphate [55].

We shall continue with the chemical-thermodynamic approach to the problem of these complicated corrosion processes and discuss the equilibria in the metal -f molten alkali sulphate system, in the same way as the molten halide and carbon-ate systems. The relevant diagram for the redox potentials in the (Li, K ) 2S 0 4 eutectic melt is shown in Fig. 4.6, according to Ketelaar [38]. The redox poten-tial scale is defined with respect to the appropriate reference electrode for these media, consisting of a noble metal electrode in equilibrium with a sulphur dioxide-f sulphur trioxide mixture with a concentration [ (1~«)S02 + « S 0 3- f

CORROSION EQUILIBRIA IN MOLTEN SULPHATES 147

Η - (1 -1«)02] / (2-£Λ) , where η varies between 0 and 1, depending on tempera-ture, under 0.1 MPa pressure [38]. At temperatures below 1150 K, in the "hot corrosion" temperature range, the equilibrium

S 0 2 + i02 ^ S 0 3 (4.36)

is strongly shifted to the formation of sulphur trioxide, so the redox potentials

of molten solutions may be expressed with good approximation as

£ r e d 0X = £ So r / o 2, s o 2- 2 . 3 ( i * r / 2 F ) (pS03-ip02) (4.37)

The standard potential at 873 Κ is 36 mV. Figure 4.6 represents the boundaries of Ni, NiO and NiS stability in the

(Li, K ) 2S 0 4 eutectic melt at 873 K. The redox potentials as functions of pS03 are conformable to formula (4.37). The diagonal lines correspond to the par-ticular p02 values. The p0

2~ scale is given on the basis of acid-base equilibrium

M e 2S O 4 0) = M e 20 ( 1) + S 0 3 ( 8) (4.38)

The value of the pKia_h) parameter is 19.5 at 873 K, as calculated from the free enthalpy of reaction (4.38). The range of Ni/NiO stability is limited by the sulphate decomposition boundary, below which the sulphate ions are reduced to sulphide:

SO2.- = S

2" + 2 0 2 (4.39)

Fig. 4.6. Stability diagram of nickel in molten lithium, potassium sulphate eutectic mixtureat 373 K (after [38]): — — — Ni/NiO stability boundary, — — — NiS stability boundary.


For this reaction pK = 40 at 873 K. At low oxygen pressures considerable amounts of NiS are formed which limits even further the range of metallic Ni stability.

A comparison of Figs. 4.4 and 4.6 makes clear both the similarities and differ-ences between the abilities of molten carbonates and suphates to corrode me-tals. In both cases the range of metal stability is determined by the partial press-ures of anion-forming gases, in other words, by the acid-base properties of the molten media. The values of the acid-base equilibrium constants of both molten media are low and, therefore, the concentrations of O

2" ions in both cases are

also very small. The data show, however, that active metallic Ni dissolution leading to N i

2+ ion formation takes place at lower S 0 3 partial pressures than

those of C 0 2. These equilibria are determined by the activities of the respective gases dissolved in both solvents. The solubility of carbon dioxide in molten car-bonates is of a "physical" nature, whereas sulphur trioxide reacts with molten sulphate solvent forming pyrosulphates as products

Me2S04 (i ) + S 03 ( 8) -» M e 2S 20 7 ( 1) (4.40)

The total solubility of S 0 3 in molten alkali sulphates is, therefore, much higher than the solubility of C 0 2 in carbonates. This property of the sulphur trioxide (molten sulphate system enhances the "hot corrosion" processes of metals, covered by a layer of molten alkali sulphates. Although this effect refers to kin-etic phenomena, nevertheless it is mentioned in this review because it results from thermodynamic equilibria characteristic of the molten alkali sulphate solvents. The corrosion of nickel by S 0 3 acting via its solution in molten sul-phates is much more intensive than the analogous process taking place in a "dry" gaseous S 0 3 atmosphere, in the absence of a molten sulphate medium.

4.9 Conclusion

Throughout this chapter the author has emphasized the limitation of the problem to the relevant aspects of the thermodynamic equilibria in the science of metal corrosion concerned with molten salt media, and it is hoped that the reader who has followed the discussion will appreciate the importance of such an approach to the problems in this field. The protection of metals against corrosion by means of passive films, thermodynamically unstable in these media, is relatively ineffi-cient, because a metal 4-molten salt systems cannot remain in a non-equilibrium state for a long time. The consideration of these systems in the equilibrium state is the only key to the limitation of corrosion losses in the melt 4-salt media.

REFERENCES 149

A rapid establishment of equilibrium over the whole system must be taken into account when considering the corrosion problems of the use of metals in con-tact with molten electrolytes; this is the main feature distinguishing this topic from the * corresponding corrosion science dealing with aqueous electrolytes at normal temperatures.

4.10 References

[1] Sundheim B. R. (Ed.), Fused Salts, McGraw-Hill, New York 1964. [2] Blander, M. (Ed.), Molten Salts Chemistry, J. Wiley and Sons, New York 1964. [3] Mamantov, G. (Ed.), Molten Salts. Characterization and Analysis, M. Dekker, New York

1964.

[4] Mamantov, G., Smith, G. P., Braunstein, J. (Eds.), Advances in Molten Salt Chemistry, Vol. 1, Plenum Press, New York 1971.

[5] Braunstein, J., Mamantov, G., Smith, G. P. (Eds.), Advances in Molten Salt Chemistry, Vol. 2, Plenum Press, New York 1973.

[6] Braunstein, J., Mamantov, G., Smith, G. P. (Eds.), Adavances in Molten Salt Chemistry, Vol. 3, Plenum Press, New York 1975.

[7] Mamantov, G., Braunstrein, J. (Eds.), Advances in Molten Salt Chemistry, Vol. 4, Plenum Press, New York 1981.

[8] Bloom, H., The Chemistry of Molten Salts, Benjamin Inc., New York 1967. [9] Lumsden, J., Thermodynamics of Molten Salt Mixtures, Academic Press, London 1966.

[10] Morand, G., Hladik, J., Eléctrochimie des sels fondus, Vols. 1 and 2, Masson et Cie., Paris 1969.

[11] Delimarskii, Yu. K., Markov, B. F., Electrochemistry of Molten Salts (in Russian), Gos. Ν. T. Isdat., Moskva 1960.

[12] Antipin, L. N., Vazhenin, S. F., Electrochemistry of Molten Salts (in Russian), Metallur-gizdat, Moskva 1964.

[13] Delimarskii, Yu. K., Electrochemistry of Fused Ionic Electrolytes (in Russian), Metal-lurgiya, Moskva 1978.

[14] Delimarskii, Yu. K., Tumanova, N. H., Shilina, G. V., Vartshuk, L. P., Polarography of Ionic Salts (in Russian), Naukova Dumka, Kiev 1978.

[15] Inman, D., Lovering, D . G. (Eds.), Ionic Liquids, Plenum Press, New York, 1981. [16] Copeland, J. L., Transport Properties of Ionic Liquids, Gordon and Bread Science Pub-

lishers, New York 1974. £17] Braunstein, J., "Statistical Thermodynamics of Molten Salts and Concentrated Aqueous

Electrolytes", in: Ionic Interactions, Ed. S. Petrucci, Academic Press, New York 1971, Vol. 1, p. 180.

{18] Volkhov, S. V., Yatsimirskii, Κ. B., Spectroscopy of Molten Salts (in Russian), Naukova Dumka, Kiev 1977.

[19] Chariot, G., Trémillon, B., Les réactions chimiques dans les solvants et les sels fondus, Gauthier-Villars, Paris 1963.

[20] Janz, G. J., Molten Salts Handbook, Academic Press, New York 1967. [21] Janz, G. J., Dampier, F. W., Lakshminarayanan, G. R., Lorenz, P. K., Tomkins, R. P. T.


Molten Salts, Vol. 1, Electrical Conductance, Density and Viscosity Data, Nat. Stand, Ref. Data System, National Bureau of Standards US 15, Washington, D . C , Oct. 1968.

[22] Janz, G. J., Lakshinarayanan, G. R., Tomkins, R. P. T., Wong, J., Molten Salts, Vol. 2, Section 2, Surface Tension Data, Nat. Stand. Ref. Data System, National Bureau of Stan-dards US 28, Washington, D. C , Aug. 1969.

[23] Janz, G. J., Krebs, U., Siegenthaler, H. F., Tomkins, R. P. T., Molten Salts, Vol. 3, Nitra-tes, Nitrates and Mixtures—Electrical Conductance, Density, Viscosity and Surface Ten-sion Data; J. Phys. Chem. Ref Data, 1, 581 (1972).

[24] Janz, G. J., Gardner, G. L., Krebs, U., Tomkins, R. P. T., Molten Salts, Vol. 4, Part 1, Fluorides and Mixtures—Electrical Conductance, Density, Viscosity and Surface Tension Data; J. Phys. Chem. Ref. Data, 3, 1 (1974).

[25] Janz, G. J., Tomkins, R. P. T., Allen, C. B., Downey, J. R., Jr., Gardner, G. L., Krebs, U., Singer, S. K., Molten Salts; Vol. 4, Part 2, Chlorides and Mixtures—Electrical Conductance, Density, Viscosity and Surface Tension Data; J. Phys. Chem. Ref. Data, 4, 871 (1975).

[26] Janz, G. J., Tomkins, R. P. T., Allen, C. B., Downey, J. R., Jr., Singer, S. K., Molten Salts, Vol. 4, Part 3, Bromides and Mixtures, Iodides and Mixtures—Electrical Conductance, Density, Viscosity and Surface Tension Data; J. Phys. Chem. Ref. Data, 6, 409 (1977).

[27] Janz, G. J., Allen, C. B., Downey, J. R., Jr, Tomkins, R. P. T., Physical Properties Data Compilations Relevant to Energy Storage, I. Molten Salts: Eutectic Data, Nat. Stand. Ref. Data System, National Bureau of Standards US 61, Part I, Washington, D. C , 1978.

[28] Janz, G. J., Allen, C. B., Bansal, N. P., Murphy, R. M., Tomkins, R. P. T., Physical Properties Data Compilations Relevant to Energy Storage, II. Molten Salts: Data on Single and Multicomponent Salt Systems, Nat. Stand. Ref. Data System, National Bureau of Standards US 61, Part II, Washington, D. C , 1979.

[29] Janz, G. J., Tomkins, R. P. T., Physical Properties Data Compilations Relevant to Energy Storage, IV. Molten Salts: Data on Additional Single and Multicomponent Salt Systems, Nat. Stand. Ref. Data System, National Bureau of Standards US 61, Part IV, Washington, D. C , 1981.

[30] Janz, G. J., / . Phys. Chem. Ref. Data, 9, 791 (1980). [31] Possipayko, W. I., Alekseeva, Ε. Α., Vassina, Ν. Α., Diagrams of Melting of Salt Systems

(in Russian), Vol. I., Binary Systems with Anions, Metallurgiya, Moskva 1977. [32] Possipayko, W. I., Alekseeva, Ε. Α., Vassina, Ν. Α., Diagrams of Melting of Salt Systems

(in Russian), Vol. II, Binary Systems with Anions, Metallurgiya, Moskva 19*77. [33] Possipayko, W. I., Alekseeva, Ε. Α., Vassina, Ν. Α., Diagrams of Melting of Salt Systems

(in Russian), Vol. I, Binary Systems with Cations, Metallurgiya, Moskva 1979. [34] Possipayko, W. I., Alekseeva, Ε. Α., Diagrams of Melting Salt Systems. Threefold Systems

(in Russian), Khimya, Moskva 1977. [35] Possipayko, W. I., Alekseeva, Ε. Α., Diagrams of Melting of Salt Systems. Multiple Sys-

tems (in Russian), Khimya, Moskva 1977.

[36] Possipayko, W. I., Alekseeva, Ε. Α., Diagrams of Melting of Salt Systems. Threefold Sys-tems (in Russian), Khimya, Moskva 1977.

[37] Franzosini, P., Ferloni, P., Spinolo, G., Molten Salts with Organic Anions—An Atlas of Phase Diagrams, Université di Pavia—Istituto di Chimica Fisica, Pavia 1973.

[38] Ketelaar, J. Α. Α., Chemie Ing. Techn., 45, 667 (1975). [39] Edeleanu, C , Littlewood, R., Electrochim. Acta, 3, 195 (1960).

REFERENCES 151

[40] Pourbaix, M., Atlas d'Equilibres Eléctrochimiques, Gauthier-Villars, Paris 1963. [41] Selman, J. R., Maru, H. C , "Physical Chemistry and Electrochemistry of Alkali Carbon-

ate Melts", in: [7], p. 159. [42] Shores, D. Α., Singh, P., "Hot Corrosion oin Molten Carbonate Fuel Cell", in: Proceedings

of the Symposium on Molten Carbonate Fuel Cell Technology, Eds. J. R., Selman, T. D. Claar, The Electrochemical Society, Inc., Proceedinds Vol. 84-13, Pennington, Ν 1984, p. 271.

[43] Andersen, Β. K., "Dissociation Constants and Hydrolysis Constants in Molten Alkali Carbonates", in: Troisième Journées Internationales d'Etude des Piles a Combustible. Com-ptes rendus, Presses Académiques Européennes, Bruxelles 1969, p. 87.

[44] Andersen, Β. K., Thermodynamic Properties of Molten Alkali Carbonates, Doctoral Thesis, Techn. Univ. of Denmark, Lyngby 1975.

[45] Barin, I., Knacke, O., Thermochemical Properties of Inorganic Substances, Springer-Ver-lag, Berlin 1973.

[46] Ingram, M. D., Janz, G. J., Electrochim. Acta, 10, 783 (1965). [47] Janz, G. J., Neuenschwander, E., Conte, Α., Corr. Sci. 3, 177 (1963). [48] Jewulski, J., Suski, L., Chemical Equilibrium Diagrams Relevant to the Molten Carbonate

Fuel Cell, Thermodynamical Data for Technology, Ser. B, Institute of Physical Chemistry, Warszawa 1985.

[49] Goebel, J. Α., Petit, F. S., Goward, G. W., "Hot Corrosion Mechanism in Stationary Gas Turbines", in: Deposition and Corrosion in Gas Turbines, Eds. A. B. Hart, A. J. B. Cutler Applied Science Publ., London 1973, p. 96.

[50] Simons, E. L., Browning, G. V, Liebhafsky, Η. Α., Corrosion, 11, 505 (1955). [51] Seybold, Α. V , Trans. AIME, 242, 1955 (1958). [52] Bornstein, N. S., DeCrescente, Μ. Α., Trans. AIME, 245, 1947 (1969). [53] Goebel, J. Α., Petit, F. S„ Met. Trans., 1, 1943 (1970). [54] Goebel, J. Α., Petit, F. S., Met. Trans., 1, 3421 (1970). [55] Conde, J. F. G., after [49].

5. Rotating Disk Studies in the Transport of Reactants during Corrosion of Copper, Iron and Zinc with Hydrogen and Oxygen Depolarization

5.1 Introduction

During the corrosion of a metal, accompanied by an effect known as hydrogen depolarization, the following overall anodic process occurs :

Me -» M ez+ + ze (5.1)

and the following overall cathodic process :

2H+ + 2e -> H 2 (5.2)

If oxygen is dissolved in the corrosive solution, another cathodic process appears which, for an acidic solution, can be written as:

0 2 + 4 H+ + 4e 2 H 20 (5.3)

It is generally believed that processes (5.2) and (5.3), and possibly* also (5.4), proceed independently in the sense that their rate constants are independent. The metal corrosion rate with concomitant oxygen and hydrogen depolariza-tion is hence the sum of the rates of processes (5.2) and (5.3), and possibly (5.4).

In the corrosion of a metal in sufficiently acid electrolyte solutions, with participation of the above two depolarizers, and if soluble products are formed, the anodic process is usually faster, while the cathodic processes are slow. Pro-cess (5.2) occurs most frequently in the kinetic control region, whereas process (5.3), or (5.4), occur in the diffusion control region. In this connection, mass transfer effects can play an essential role in metal corrosion processes, especially in corrosion with participation of the oxygen dissolved in solution.

* It is also very probable that in certain cases of corrosion with oxygen depolarization, the oxygen dissolved in the electrolyte is partly reduced to hydrogen peroxide:

02 + 2 H

+ + 2e 2 H 20 2 (5.4)

THE DIFFUSION-CONVECTION MASS TRANSEFR EQUATION 153

In extreme cases (diffusion control), transfer of the reactants is the rate-de-termining step in the overall metal corrosion process.

The rate of transfer of reactants to the reaction surface is affected by a num-ber of factors. In addition to such physical quantities as: concentration of the transferred substance, its diffusivity, viscosity of the medium and temperature, a further parameter can be introduced, often encountered in corrosion studies, and referred to as stirring intensity. The numerical value of this parameter is known for only a few simple systems. Fortunately one of these systems, which is of much assistance in corrosion studies, is the rotating disk system.

5.2 The diffusion-convection mass transfer equation and the rotating disk

Corrosion processes, like all heterogeneous reactions, proceed through a number of consecutive steps. Reactants reach the reaction surface or reaction zone to undergo a chemical transformation, then the reaction products leave the reac-tion zone or else the zone itself is displaced. The chemical reaction may be pre-ceded by adsorption of reactants, and it may comprise both consecutive and simultaneous elementary steps, combined with adsorption and desorption of the intermediate and final products.

As mentioned previously, one possibility is that the rate-determining step of a corrosion process is the transfer of reactants to or from the reaction surface or reaction zone. In the case of corrosion of solids in liquids with the formation of soluble products this transfer of reactants is in principle* accomplished in the liquid phase, brought about by molecular diffusion and accompanied by a natural or forced convection. In certain specific systems, for example during an electrolysis process, migration of ions in the electric field may overlap these two kinds of transport.

The convective diffusion equation for a solute in an incompressible liquid can be written thus [1-4]:

where c is concentration of solute, t is time, V is liquid velocity vector, D is the Fick's diffusion coefficient for the solute. Expression (5.5) is applicable to a system in which the concentration of a solute is low enough for D to be inde-pendent of concentration gradient.

* Even in this corrosion system it is conceivable that corrosion rate may be controlled by transport in the solid, for example for an alloy, semiconducting material or a chemical com-pound.

(5.5)

154 ROTATING DISK STUDIES IN THE TRANSPORT OF REACTANTS

For immobile liquids (v = 0) equation (5.5) converts to the following form:

ΐ =

D Ac <

5-

6>

Solutions of equation (5.6) for the case specified as semiinfinite unidimensional diffusion are used in electrochemical methods of corrosion studies. As far as the direct observation of a corrosion process in liquid media is concerned, equa-tion (5.6) is generally inapplicable, because such corrosion cases are invariably accompanied by convection.

In the case of diffusion-convection mass transfer of a solute in a liquid, a stationary condition is rapidly established, i.e. dcjdt = 0, and equation (5.5) can be re-written in the form:

Fgradc = DAc (5.7)

In expression (5.7) the velocity of the liquid is considered to be determined by the Navier-Stokes equation. The înitial and boundary conditions need to be specified: the latter are determined by the physical and chemical transforma-tions at phase boundaries.

Equation (5.7) can be roughly visualized physically in the following manner. In the bulk of a liquid, a solute is transported almost entirely by convection; the share of molecular diffusion is negligible. On approaching the solid surface, the liquid moves increasingly slower in relation to the surface, and the motion vanishes entirely at the interface. If the solute is removed as a result of a reac-tion occurring at the solid surface, then in the liquid layer adjacent to the solid the share of molecular diffusion in the reactant transport constantly increases. Thus in the immediate neighbourhood of the solid two layers are formed: one, with a decreasing velocity of liquid, termed the hydrodynamic boundary layer or Prandtl's layer, and the other, with a decreasing concentration, termed the diffusion boundary layer.

The hydrodynamic boundary layer, as well as the diffusion layer in the di-rection normal to the solid surface, extend to a distance that is very small in comparison to the linear dimensions of the interface. The use of this observa-tion [1, 2] especially in hydrodynamics, has allowed relatively simple solutions to the liquid motion equations. These results,* combined with the analogy

* Progress in studies of convective mass transfer especially in engineering applications, has also been achieved by the use of dimensional analysis and a semi-empirical description of phe-nomena in terms of the criterion numbers. This approach is suitable for models of mass transfer processes.

THE DIFFUSION-CONVECTION MASS TRANSFER EQUATION 155

between mass and heat transfer, have in turn led to solutions of equation (5.7) for systems of a particularly simple geometry and a straightforward mechanism of liquid motion (laminar flow).

Such a simple (in the meaning specified in the above paragraph) system is the rotating disk system. The solution of equation (5.7) provided by Levich [1, 2] for this case is :

j = D(dc/dz)z=0 = D^-c^/d (5.8)

δ = l . o l D ^ V / S f t r

1'

2 (5.9)

where j is the diffusion flux, D is the diffusion coefficient of the reactant, c° is its concentration in the bulk of solution, c

1 is concentration at the disk surface,

ν is the kinematic viscosity, and ω is the angular velocity of disk. The quantities appearing in equations (5.8) and (5.9) are usually expressed in the mol-cm-s system.

The rotating disk system is particularly helpful in kinetic studies of hetero-geneous reactions. This is due to the following features of this arrangement: the thickness of the diffusion layer is independent of the disk radius, so that the change in the type of control of a reaction is well defined, as it simultaneously appears over the entire disk surface; and the thickness of the diffusion layer is a simple function of the angular velocity of the disk.

Equation (5.8) is identical to the well known Nernst formula. However, in the latter the diffusion layer thickness δ is an empirical constant, whereas in the former, i.e. as provided by Levich [1], it is determined by simple physical quan-tities.

If the rate of a chemical change is much higher than the rate of reactants transfer, then c

1 = 0 and the reaction proceeds in the diffusion region. If the

rate of transport greatly exceeds the reaction rate, then on the disk surface c1 χ c° and the reaction occurs under activation control. In intermediate cases

where the reaction occurs in the mixed kinetic region, the discussion is much more involved. The mass-balance on the disk surface can be written thus

j = v = k(cxY (5.10)

where ν is the rate of the chemical reaction, is a formal rate constant, depend-ent on the electrode potential, and ρ is the reaction order with respect to sol-ute. The solution of the system of equations (5.8) and (5.10) allows the de-termination of the order and the rate of reaction. Corresponding practical ex* amples and procedures will be given later in this chapter.


Equations (5.8) and (5.9) have often been verified experimentally.* The differ-ence between the value of diffusion flux measured and that calculated is not greater than the error of measurement which in precision experiments may be below 3%.

The equation for the diffusion flux to the rotating disk surface was derived on the assumption of a laminar flow. For carefully made smooth disks the la-minar flow condition has been found to correspond to Reynolds numbers, Re = r

2a>/v « 10

4, where r is the disk radius. Since it was assumed that the disk

surface area is infinitely large, it must hold that r > δ. The boundary effect, however, manifests itself clearly at r values as low as r < 0.5 mm.

5.3 Diffusion control of the corrosion rate

When the rate of a chemical change at the interface is about one order of magni-tude larger than the rate of transfer of any of the reactants to the reaction surface [1, 2], then c

1 # 0 and diffusion control of the process is observed ex-

clusively in experiments. Then from equations (5.8) and (5.9) we obtain

7i,um = ΟΛΙΏ^ω^ν"1^ (5.11)

If ji is measured in mol/sm2, then on multiplying equation (5.11) by nF we will

obtain the corrosion rate expressed in terms of current density, A/m2 thus:

illm = 0.62«FA2 / 3

c?co1/2^

1/6 (5.12)

where η is the number of the electrons involved in the net reaction and F is Fa-

raday's constant. It follows from equation (5.11) that an experimental criterion of diffusion

control is a linear run of corrosion rate vs. the square root of the angular vel-ocity of the disk. Moreover, upon linear extrapolation the line should go through the origin of the coordinates. The use of equation (5.11) is possible only if the corrosion rate is constant with time. This corresponds to a linear relation of specimen weight loss with time.

The results [5] presented in Fig. 5.1 may be taken as a typical example of diffusion control of a corrosion process. The results were obtained by dissolving a disk of area 7.07 cm

2, made of a Cu-15 Zn brass, in an air-saturated 2 .0MH2SO4

solution. The alloy weight losses were determined from analysis of solution

* In high-precision measurements more precise expressions are used for δ than in equation (5.9) (cf. the treatment in [2]). The difference between such measurements and those from equation (5.9) does not exceed 3%, provided the value of the Schmidt number Sc = v\D > 5 x l 0

2.

DIFFUSION CONTROL OF THE CORROSION RATE 157

Fig. 5.1. The corrosion rate of Cu-15 Zn brass vs. square root of the disk speed, rps, at a tem-perature of 338 ± 0.2 Κ in air-saturated 2 . 0 M H 2S O 4 (after [5]): O—copper corrosion rate, •—zinc corrosion rate.

samples for copper and zinc. The weight loss vs. time relation is linear, and a slope is equal to the time-constant dissolution rate.

Moreover, in the same experimental setup [5] the limiting oxygen flux was measured in a series of measurements of the two-electron cathodic reduction of the oxygen dissolved in the air-saturated solution under study.

Brass (Cu-15 Zn) fails to dissolve in a 2.0 M H 2S 0 4 in the absence of oxygen [5], as was demonstrated by measurements [5] on an argon-saturated solution. This fact, in conjunction with the fact that the corrosion rate vs. square root of the disk speed relation reported in Fig. 5.1 is linear and that the straight line, upon extrapolation, goes through the origin, strongly suggest what is fairly ob-vious; the brass studied corrodes in the diffusion region only when accompanied by oxygen depolarization. The overall corrosion reaction may then be written thus:

(Cu-Zn) + 0 2 + 4 H+ -» C u

2+ + Zn

2+ + 2 H 20 (5.13)

Since the reaction is controlled by the transport of the oxygen dissolved in the acid, the following relation should hold iCu + / Zn = *o2,iim> and, as follows from reaction (5.13), the limiting current of oxygen is calculated for a four-elec-tron reduction. Indeed (cf. Fig. 5.1) at 338 K i Cu + *zn = (1.5±0.1)j/m A/m

2,

and io2,lim = (1.7±0.2)^m A/m2.

In the system described, the slow-diffusing substance has been identified and conclusions have been made as to the corrosion reaction. On many occasions this sort of discussion is not possible, as some highly complicating effects often


become operative. An example of this is the dissolution of copper metal in air-or oxygen-saturated ammonia solutions at concentrations above approx. 0.5 M. In particular, it has been found that even though the reaction take place in the diffusion region, the flux of dissolved copper is larger than that suggested by the stoichiometry of the reaction of the reduction of oxygen on the electrode sur-face.

The situation is made somewhat clearer by the data of Fig. 5.2. Copper dis-

li ο Γ

I τ

δ /

0 /

Ο /

Time, min

Fig. 5.2. The loss in weight of a rotating copper disk vs. time of corrosion in an air-saturated I . O M N H 3 solution; temperature 298±0.2 K, m = 25 rps (after [6]): O— first measurement series, Δ—second measurement series.

solves with oxygen depolarization following the overall reaction:

Cu + 4 N H 3 + H 20 + i02 Cu(NH3) i+ + 20H~

The Cu(II) complexes formed react with metallic copper:

Cu + Cu(NH3) l+ -> 2Cu(NH3)2

h

(5.14)

(5.15)


At the same time, in the presence of oxygen dissolved in the solution:

2Cu(NH3)i + i 0 2 + 4 N H 3 + H 20 = 2Cu(NH3)I+ + 20H~ (5.16)

Reaction (5.16) is fast compared with reaction (5.15). Initially, at a low Cu(II) concentration in solution, copper corrodes mainly as a result of reaction (5.14) but as time goes on, reaction (5.15) prevails. As mentioned previously, both reactions take place in the diffusion region, however the corrosion rate increases with time [7].

It sometimes happens that the limiting diffusion flux can be estimated nu-merically and then compared with the rate of metal dissolution measured. An example of this are the results obtained during dissolution of 99.99% purity zinc disks in nitrogen-saturated 0.1 M NaCl solutions in a pH ranging from 12.0 to 13.3, obtained by addition of NaOH. In this pH range zinc dissolution proceeds under diffusion control. This can be seen from Fig. 5.3 which shows

Fig. 5.3. Zinc corrosion rate vs. square root of the disk speed, rps, in a nitrogen-saturated 0.1 M NaCl solution at pH = 12.9, temperature 298 ±0 .2 K; the upper line is the limiting flux of HZnOr ions estimated from the expression: l o g f l z n o â = —15.37 + pH; the bottom line is the flux estimated from the expression: l o g a H z n o J = —16.68 + pH; the assumption was made that DHZno~ = 10"

6 cm

2/s (after [8, 21]).

the results obtained for a solution of pH 12.9 [8]. It follows from the potential vs. pH diagram [9] for the Z n - H 20 system that at this pH value the HZnOj ion is stable. Its activity in a saturated solution with respect ot the ZnO phase is: tog^Hzno; = -15.374-pH, or l ogaH Z no; = -16 .68+pH, depending on the thermodynamic data used for the calculations. Results of the rough evaluation using equation (5.12) (aH Z n C.2 ~ ^Hznoi) of the limiting diffusion flux for H Z n 0 2


ions are shown in Fig. 5.3. It was assumed that diffusion occurs from the satu-rated solution at the zinc surface into the bulk of solution, where the concentra-tion of HZnOi" ions is nearly zero. On a further assumption that the metal dis-solution rate is controlled by diffusion of OH" ions to the electrode surface, the rough evaluation of the limiting diffusion flux for OH" ions yields values that are some three orders of magnitude larger than the zinc dissolution rate measured [10].

In the latter case, i.e. the corrosion of zinc in strongly alkaline NaCl solu-tions, the diffusion suppression of the corrosion process was observed. It has been suggested that diffusion from the reaction surface into the bulk of solu-tion is responsible for this. An analysis of this idea leads to a number of con-ceptual complications. These are related to the indispensable assumption, with-out which the limiting flux would not occur, that the solution at the metal surface is saturated with its salts. Salt layers should therefore precipitate from saturated solution at the metal surface.

If the metal surface is coated with a salt layer, there are two alternatives. One of these represents the case of a porous layer that seals off a substantial portion of the electrode surface. In this case an agreement between the value of the diffusion flux of zincate ions calculated from equation (5.8) (on the assump-tion that the diffusion surface equals the geometrical surface area of the disk), and the value of the zinc corrosion rate measured is fortuitous. The other alter-native is to assume that the salt layer is non-porous. Here, however, the corro-sion rate is determined by the phenomena occurring in the solid, rather than by the transfer processes in the liquid.

In any event, the concept of corrosion accompanied by a suppression of the process by a slow diffusion of metal complexes or salts from the surface into the bulk of solution are uncertain enough to require a carefully planned and com-prehensive experimental confirmation.

During a metal corrosion process there are a number of diffusion fluxes, as many soluble reactants in the solution participate in the reaction. In the case of diffusion suppression of the reaction, its rate is controlled by that flux whose limiting value, y l i m, is the least under the specific conditions.

An example of this situation is the dissolution of a zinc disk in air-saturated, acidic 0.1 M NaCl solutions. In solutions of this nature zinc corrodes with sim-ultaneous oxygen and hydrogen depolarization. This latter parallel corrosion reaction somewhat obscures the picture of the whole process. By taking suitable measurements, however, which will be dealt with later, the corrosion with hydro-


gen depolarization can be measured separately and thus it can be singled out from the overall corrosion process.

The acid corrosion of zinc is kinetically controlled [10]; i.e. it is independent of the angular velocity of the disk. Conversely, zinc corrosion with oxygen de-polarization proceeds in the diffusion region, although the hydrogen ion is in-volved in both reactions. With increasing pH of solution the limiting flux of hydro-gen ions, initially a few orders of magnitude higher than the corresponding flux of the dissolved oxygen, rapidly decreases, while the latter remains virtually unaltered [11, 12] In a 0.1 M air-saturated NaCl solution both fluxes become roughly equal at a pH of 3.5 (Fig. 5.4); at higher pH values the limiting flux

20

Fig. 5.4. Zinc corrosion rate in 0.1 M NaCl solutions of different pH: circles are for air-satu-rated and triangles are for nitrogen-saturated solutions; crosses denote dissolution rate with oxygen depolarization; solid lines denoted by symbols H

+ and 0 2 refer to the limiting current

density for the cathodic reduction of H

+ ion and of the four-electron reduction of the oxygen

dissolved in solution. Results are reported for a disk speed of 16 rps (dissolution rate in the nitrogen-saturated solutions is independent of stirring rate). Temperature 298 ± 0.2 Κ (after [8]).

of hydrogen ions is already smaller than that of the oxygen dissolved in the solution. Accordingly, in solutions of a pH > 3.5 the zinc dissolution rate should be controlled solely by the diffusion of hydrogen ions to the metal surface. At the same time, all of the hydrogen ions are removed by the corrosion with oxy-


gen depolarization, so that the acidic corrosion of the metal virtually disappears.* This indicates that the metal corrodes exclusively under oxygen depolarization; the rate of the process is, however, controlled by hydrogen ion diffusion. At pH < 3.5 the hydrogen ion flux is high enough for both corrosion reactions to proceed, i.e. both the corrosion with oxygen depolarization controlled by the limiting flux of dissolved oxygen, and the corrosion with hydrogen depolariza-tion that operates in the kinetic control region.

Experimental results confirming this argument are shown in Fig. 5.4. Circles in this figure denote the measured values of the zinc corrosion rate in air-satu-rated 0.1 M NaCl solutions of different pH. Triangles correspond to the values measured for zinc corrosion rate in these same solutions but saturated with nitrogen. The rate of zinc dissolution with oxygen depolarization alone (crosses) was obtained by subtracting the value of the corrosion rate with hydrogen depolarization (triangles) from the value of corrosion rate with simultaneous hydrogen and oxygen depolarization (circles). The solid lines with symbols H

+ and 0 2 correspond to the measured values [10-12] of the limiting current

density for the cathodic reduction of hydrogen ions and for the four-electron reduction of the oxygen dissolved in the solutions studied. This approach has been discussed in the study by Zembura and Fulinski [13].

It will be seen from Fig. 5.4 that the rate of zinc corrosion with oxygen de-polarization (crosses) equals the value of the limiting current of the cathodic oxygen reduction, and at pH > 3.5, it equals the value of this current for the cathodic reduction of hydrogen ions. The plot of the corrosion rate with oxygen depolarization controlled by the hydrogen ion diffusion versus the square root of the speed of rotation of the disk would, of course, be a straight line passing through the origin. The effect described here is analogous to those treated by Kemula and Grabowski [14] and related to the latent limiting current of the hydrogen ion.

The results so far described are related to the diffusion control of a corrosion process for the case where the entire disk surface is diffusion-active, i.e. the surface through which diffusion occurs is the same as the geometrical surface of the disk. The situation becomes more involved however, if the corrosion

* In the system considered the following two overall cathodic reactions proceed simulta-neously: 2 H

+ + 2e = H 2 (hydrogen depolarization) and 0 2 + 4 H

+ + 4e = 2 H 20 (oxygen

depolarization) and one anodic reaction: Zn = Z n

++ + 2e. The sum of the rates of the former

two reactions equals the rate of the latter. At the same time the first reaction is slow compared with the second and third one, as shown in Fig. 5.4. Hydrogen ion therefore is used up pri-marily in the second reaction and only excess H

+ participates in the first reaction.


process can no longer proceed on part of the disk surface as a result of, say, precipitates deposited, adsorption, etc. In this situation two extreme cases can be distinguished:

1. the rotating disk active fragments and the spacings between these fragments are much larger than the diffusion layer thickness; or

2. the active fragments and their spacing are much smaller than the diffusion layer thickness.

Case 1 is qualitatively similar to that described by Levich [1] for the disk with a partly covered surface. In the event of a diffusion control of the process occurring on the active fragments of the disk surface, the rate of the process satisfies equation (5.11) or (5.12), although the slope of the experimental line is smaller than the value found for the geometrical disk surface, and depends on which fraction of the surface is covered.

In case 2 the geometrical disk surface may be regarded as equal to the diffusion surface: accordingly, the whole disk surface is diffusion-active.

For intermediate cases, when the control is of a diffusion character, and the radius r" of the active surface fragments and the mean distance between their centres 2r" are comparable with the diffusion layer thickness, Landsberg et al [15-17] derived an approximate equation:

where A„ and xn are constants calculated by Scheller et al [15]. If the disk ro-tates at a moderate speed, r" <̂ δ, then

tanh(x„/r") « 1

and from equation (5 .17) we obtain

ι 1 . 6 1 V

1/

6 ι Σλ„

= ^ + (5.18)

him nFD2l3c° \/ω nFDc

0

Thus, a relation analogous to that found for the case of the mixed kinetics of the process is obtained (a further discussion will be continued in subsequent sec-tions).

For high disk speeds, r" > δ and tanh(xn<?/r") # x„ô/r" hence:

+ (5.17) /lim nFD

2l3c° \/ω nFDc

0


and a straight line is obtained in plotting 1 / /Um vs. /(l/j/co) which, as distinct from the line of equation (5.18), passes through the origin.

As a result, whenever the electrode surface becomes partially covered as described above, together with a diffusion control of the process, we should obtain two intersecting lines, one of which will pass through the origin, if the results are plotted as l / / l im vs./(l/j/co). However, an alternative hypothesis to blocking of the disc surface can be put forward to explain such an outcome; this involves the assumption of mixed kinetics (considered later), with a conclu-sion that this interpretation of the results is the correct one. The two mechanisms (surface blocking and mixed kinetics) can be distinguished by determining the temperature dependence of thé corrosion rate. In the case of mixed kinetics the corrosion rate depends strongly on temperature (the activation energy of the kinetic component is at least 35-40 kJ/mol, whereas with surface blocking com-bined with diffusion control at the uncovered areas, the temperature dependence of corrosion rate is much less pronounced.

This is probably how Armco iron corrosion with oxygen depolarization

Fig. 5.5. The corrosion rate of an Armco iron disk vs. oxygen concentration in solution and square root of the disk speed, rps in 0.1 M solutions of NaCl, NaBr, N a 2S 0 4, N a H P 0 4 and NaCIO* of a pH = 2.2. Measurement temperature 298 ± 0.2 K. The solid line refers to the limit-ing current of the four-electron cathodic reduction of the oxygen dissolved in a 0.1 M Na2S04 solution. The expression Cg2 ]/ω/2π was derived on the assumption that the oxygen concentra-tion in air-saturated solutions equals unity, and hence in oxygen-saturated solutions at an at-mospheric pressure it is 4.8 (after [20]).


occurs in acidic solutions* of various salts, such as: N a 2S 0 4, NaCl, NaBr, N a H 2P 0 4, etc. [18]. The metal corrosion rate in these salts is given by the rela-tion shown in Fig. 5.5.

The plot resembles the mixed kinetics curve, i.e. the case when the rate of the chemical transformation at the interface is comparable with the transport rate, being of the same order of magnitude. However, discussion of the curve of Fig. 5.5 on the assumption of the reaction proceesing in the mixed kinetics region, proceeding in the mixed kinetics region leads to physically meaningless results [19]. On measuring the temperature dependence of the corrosion rate and evalu-ation of the activation energy of the alleged kinetic component** of the iron corrosion with oxygen depolarization, this was found to be approx. 8.5 kJ/mol. The result demonstrates that the concept of the reaction proceeding in the mixed kinetics region is fallacious, and is thus unacceptable. Elaboration of the re-sults of Fig. 5.5 using equations (5.18) and (5.19), i.e. on the assumption of a covering of part of the metal surface, which is illustrated in Fig. 5.6, solves the problem.

<

8 0.2

8

& 0.5 1

Fig. 5.6. The results of Fig. 5.5 in the coordinates: /<Toî =/( l / j /û>). According to equations (5.18) and (5.19) two intersecting lines are derived of which one, upon extrapolation, passes through the origin (after [20]).

* The acidic dissolution component for iron was accounted for in the manner described previously in the case of zinc corrosion with simultaneous oxygen and hydrogen depolariza-tion.

** Determination of the kinetic component of the reaction proceeding in the mixed kin-etics region is dealt with in Section 5.5.


The question of identification of the substances blocking the iron surface remains open. They are perhaps impurities present in the metal, but this re-quires further investigation.

At any rate, replacement of Armco iron in the measurements by zone-re-fined iron shifts the reaction to the diffusion control region on a uniform, non-covered electrode surface. The corrosion rate then coincides with the line shown in Fig. 5.5 for the limiting current of the four-electron reduction of the oxygen dissolved in solution. The same effect is exerted by an addition to the Armco iron system of such ions as Ag

+, C u

2 +, etc. in quantities not lower than 10~

8

to 10~7 mole/1. This effect of such ions precipitating on the iron can be con-

sidered as causing the number of cathodic sites to increase. At a specific surface area of the electrode this results in a gradual decrease in the average spacings between the active sites, and thus in a diffusion reactivation of the whole sur-face area of the disk.

In the same way, during copper or α-brass corrosion in oxygen- or air-satu-rated ammonia solutions of concentrations lower by approx. 0.5 M and 0.15 M, respectively, the diffusion-active electrode surface certainly becomes smaller. This manifests itself in that in similar solutions, during the corrosion process the rate of specimen weight loss gradually decreases. This obviously means a drop in corrosion rate* with time, possibly as a result of precipitation of salts on the metal surface.

5.4 Corrosion in the activation controlled region

It has already been mentioned in previous sections, that the corrosion rate in the activation (kinetic) controlled region is independent of the disk speed. The concentration of the reactants at the reaction surface is virtually equal to their concentration in the bulk of solution. The majority of commercially important metals and alloys corrode with hydrogen depolarization in this manner. The investigation of such systems, using the rotating disk electrode, provides data for the discussion of corrosion processes involving mixed kinetics.

In the case of the metals under discussion, the situation is as follows. Copper is known not to undergo corrosion in non-complexing, oxygen-free acid sol-utions. In these same solutions, but containing oxygen, the corrosion occurs in the activation control region, provided the temperature is low enough. As the

* It seems that copper and α-brasses exhibit a fairly strange behaviour in oxygen-containing ammonia solutions. In more concentrated solutions the corrosion rate increases with time, whereas in more dilute ones it decreases.

CORROSION IN THE ACTIVATION CONTROLLED REGION 167

temperature is raised (cf. subsequent pages), a change through mixed control to diffusion control regions is observed [20]. As stated previously, copper cor-rodes under diffusion control in oxygen-containing ammonia solutions [5].

In acidic electrolytic solutions not containing oxygen or other oxidants, the acid corrosion of iron proceeds in the activation control region. If oxygen is dissolved in these solutions, an additional oxygen-depolarization corrosion appears, which runs in the diffusion control region with a partial coverage of the surface of the corroding metal [8]. This process proceeds in solutions of a pH from 2.5 to 3.5 depending on measurement conditions. At higher pH values, the picture of iron corrosion with oxygen depolarization changes completely, as a result of salt precipitation on the metal surface.

The widest variety of effects (Fig. 5.7) is observed for zinc [21]. In oxygen-free

1.0

CM

1

ε

a 0.0

-1.0

-2.0,

Λ b Τ

Ό 4 8 12 PH

Fig. 5.7. The logarithm of the zinc corrosion rate (A/m

2) vs. pH in 0.1 M NaCl solution not

containing dissolved oxygen. Temperature 298 ± 0.2 K. The line marked by H is the limiting current of the cathodic reduction of hydrogen ions for one rps of the disk, a, (/—kinetic con-trol including: b—the mixed-kinetics region, c—probably two simultaneous corrosion reactions of one which is kinetically controlled and the other is diffusion controlled, e—diffusion control region (after [8]).

non-complexing salt solutions, depending on the degree of purity, the zinc disk electrode corrodes in the kinetic region up to a pH of 3.5. From this to a pH of approx. 4.5, there is a mixed-kinetics region, then from a pH of 5.0 to approx. 11 two different regions of kinetic control appear; and finally, beyond a pH approx. 12, the process is diffusion controlled. If oxygen is dissolved in the sol-


ution, this process becomes overlapped by the oxygen-depolaritazion corrosion. These latter effects, treated in some detail in subsequent sections, are insuffi-ciently understood.

5.5 Corrosion in the mixed-kinetics region

It has already been mentioned in preceding sections that the diffusion control of metal corrosion can be experimentally observed (with a few per cent error), provided that the rate of the activation controlled stage of the process is several times larger than the value of the limiting diffusion flux of the reactants diffusing under specific conditions [13]. Conversely, kinetic control of the corrosion rate will be involved, if under specific experimental conditions the lowest of the limiting diffusion fluxes of reactants is still several times larger than the rate of the kinetic stage of the corrosion process. In the intermediate cases a mixed kinetics of the corrosion process can be observed.

Considering equations (5.8) and (5.10) for the corrosion rate with hydrogen depolarization in the mixed kinetics region for a bivalent metal for example, we derive the relation:

/ M E Î+ . / H+ = ^>H+0&-4») = F k ( c ^ ) p ( 5 2 0 )

The following relations are simultaneously satisfied:

f /V(cH+-4+)

J » \ = 1 - 4 - (5 ·2 1

)

δΗ+ = l . o l D ^ V '

6^

1'

2

/H+ _ Fk(c^y _ / 4 Λ ρ ( 5 · 2 2 )

where c£+ and are the hydrogen ion concentrations in the bulk of solution and at the electrode surface, respectively, k is the reaction rate constant of spon-taneous hydrogen ion reduction on the metal surface, ρ = d(logirtnm)/d(logc^+) is the order of the corrosion reaction with respect to hydrogen ions, and the quantity i r t l im = Fk(c^+)

p is the kinetic component of the corrosion process,

i.e. the corrosion rate if c«+ # This value is achieved in the pH range or stirring rate within which the corrosion rate is independent of stirring intensity.

CORROSION IN THE MIXED-KINETICS REGION 169

Form formulas (5.21) and (5.22):

Ch+ *H

+, lim \

CH

+ I

(5.23)

If ^H

+,iim > /r, «m > ̂ Η+ΛΉ+ ~ 1>

a nd the corrosion process passes into the kinetic

control region; the corrosion rate is no longer dependent on stirring. Converse-ly, if J*HMim ^ U,nm> ^

e corrosion runs in the diffusion control region; its

rate is a linear function of the square root of the disk angular velocity. This latter statement is derived from equation (5.23) on dividing both sides by z M im and on the assumption that the reaction rate constant k oo. Then CH+/CH+ ~ 0, which corresponds to c«+ # 0, and thus to diffusion control of the rate of the process. The way in which equations (5.21) to (5.23) can be used depends on the extent of the measuring data available. If ir,llm9 /H

+, i i m >

as w eU

as the reaction

order ρ have been found from separate measurements, we can solve equation (5.23) in terms of <?Η+/^Η

+· Substituting the result into, say, equation (5.21)

we obtain /H+ = / Μ β2 + .

1.0

ο 05

0.0

ο ο

-0.5

•hi V \ I Χ \

V Τ V

u

Γ 1\

ρΗ

Fig. 5.8. Dissolution rate of a zinc disk at 298 ±0 .2 Κ in nitrogen-saturated 0.1 M NaCl

solutions of various PH. Circles—dissolution rate in the kinetic control region, triangles—dis-

solution rate in the mixed-kinetics region, solid line—limiting current density for the cathodic

reduction of hydrogen ions in the solutions studied. The data for the mixed kinetics and re-

duction of the hydrogen ion are reported for unit disk speed (1 rps). The figure is an enlarged

section of Fig. 5.7 (after [8, 21]).

170 ROTATING DISK STUDÎES IN THE TRANSPORT OF REACTANTS

An illustration of this approach to the use of such measurements are the data shown in Figs. 5.8 and 5.9. Figure 5.8 shows the relation between pH and the logarithm of the corrosion rate of a zinc disk in nitrogen-saturated 0.1 M NaCl solutions made acidic to the required pH with HC1. At pH < 3.5 the disk dis-solves under the kinetic control conditions (corrosion rate is independent of stirring and equals / r, i l m) . At a pH > 3.5 the reaction proceeds in the mixed-kinetics region with respect to hydrogen ion diffusion. Figure 5.8 also shows the logarithm of the limiting current density for the cathodic reduction of hydro-gen ions in the solutions studied, so that / HM i m was also measured.

Figure 5.9 gives the corrosion rate of a zinc disk at pH of 3.9 vs. the square root of disk speed. The curve has the typical shape of a mixed-type kinetics situation. The same figure shows also the curves calculated from equations (5.21) to (5.23).

The procedure used in the calculations is as follows. The / M im line of Fig. 5.8, when extrapolated to a pH value of 3.9, gives log / r, l lm = —3.77. The slope of this line is ρ = d(logirtllm)/d(pH) = 0.33+0.03. The / HM i m value at a pH of 3.9 was also read from Fig. 5.8.

In cases where the quantities: fr,nm, faMim

a nd Ρ

a re difficult to measure,

they can be evaluated from the experimental curve of the mixed kinetics by the methods described in the monograph by Levich [1] or in the book by Pleckov

classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which included slow growing pits (pit diameter


and Filinovskii [2]. Of the various methods mentioned, the most convenient in use is the graphic method.

Let us assume for this purpose that the corrosion process occurs in the mixed kinetics region, for example with respect to the oxygen dissolved in the corro-sion solution. We can evaluate c^2 from equation (5.10) ançl we introduce the expressions thus derived into equation (5.8). Following a few transformations we obtain:

iUF = %f = nFkW - 177 4 ^ (5.24)

where m is the disk speed of rotation, r.p.s. At a constant temperature and a con-stant oxygen concentration in the bulk of solution we have:

tXi = W? = A - B ^ (5.25) y m

where A and Β are the constants defined by equation (5.24). In the coordinates

iUf and iC0Tl\/

m a straight line should therefore be obtained.

Equation (5.24) is simplified if ρ = 1. From equation (5.25) we then obtain:

— = L+M—L- (5.26) hot \m

which gives a straight line using the coordinates l / /c or and l/ | /m. If m -» oo, we obtain from equation (5.24):

H!Sm = (nFkyiPc

0 = A (5.27)

which signifies that c1 « c°, thus the reaction occurs in the kinetic control re-

gion. The ( i r, n m)

1 ,p value is cut off on the ordinate axis by the experimental

curve plotted in the coordinate system: ill?

a nd horlV

m- ^

n dividing equation

(5.24) by nFk

ll

p we obtain :

&p 1 i

*cor λ A *cor /β 28) nFk

llP ~ 0.62i/nnFD^2

3v^

6 j/m

If k oo, i.e. the reaction goes into the diffusion region, then:

*'cor/j/w = io2, nj^m = 0.62]/ïiznFDy

3v^

6c

0 (5.29)

which is identical with, for example, equation (5.12). Equation (5.29) describes the limiting diffusion current of the cathodic reduction of the oxygen dissolved in solution. The experimental curve on the mixed kinetics plot in coordinates ill?


and i c o r/ j /m cuts off, on the abscissa axis, the value of expression (5.29). Ob-viously an additional direct determination of / o 2, i i m / | / ^ is also desirable, which is generally possible. An agreement between the measured and extrapolated values is an additional proof in support of the validity of this interpretation of the results.

Determination by the above approach of the reaction order ρ (which is done by trial and error) and of the reaction rate constant k, furnishes values that are not accurate enough, in view of the generally considerable experimental error involved, if the corrosion process proceeds in the mixed kinetics region.

A good example of metal corrosion [22] in the mixed kinetics region is copper dissolution at a temperature of 318 Κ in a 1.7 M air-saturated sulphuric acid solution. The experimental curves are given in Fig. 5.10, and the same for the

318 Κ line results, are re-drawn in the coordinates: l/jCu (ordinate) and l/j/m (abscissa) in Fig. 5.11.

In this example, copper corrosion is readily assigned to the mixed kinetics region in view of the results presented in Figs. 5.10 and 5.11. Moreover, direct measurements of the corrosion rate in 1.7 M air-saturated sulphuric acid sol-

4 j j j 1 1 ; 1 1

1 2 3 1 4 5

Fig. 5.10. Corrosion rate of a copper disk in air-saturated 1 . 7 M H 2S 0 4 solutions vs. square root of disk speed, rps, at various temperatures (after [22]).


utions under various oxygen perssures show that the process is of the first order. Naturally, the activation energy of the kinetic component of copper corrosion with oxygen depolarization is high and amounts to approx. 60 kJ/mol.

It should be emphasized that the reason why the results of measurements of copper corrosion rate with oxygen depolarization cannot be evaluated in the same manner as for the zinc corrosion with hydrogen depolarization (cf. earlier in this chapter), is the fact that /o2,r,!»m cannot be measured directly at 323 K. This value can however be estimated, provided the activation energy and the *o2,r,iim values are known at temperatures below 323 Κ (the results obtained at 308 Κ are reported in Fig. 5.10). This estimation is, however, too inaccurate as the error in determining the activation energy is large, because of the small tem-perature range over which activation control occurs. In the case of zinc corrosion with hydrogen depolarization, a direct measurement of / H

+r , i i m is possible,

as the CH+ value can be varied over a few orders of magnitude.

It is worth recalling here that the coordinate system in which a straight line is obtained for the mixed kinetics with participation of the first-order reaction is identical with the system involved in a study of the diffusion control on a partly covered electrode. An unambiguous differentiation of these two cases is possible only by measurement of the activation energy for the reaction con-sidered to be a kinetic component. By determining / M im at several tempera-tures we can estimate the activation energy. For the kinetic component in the systems we are concerned with, this is a few times greater than 8-16 kJ/mol— the value which is characteristic of the diffusion component of a corrosion pro-

classed into two groups: one, which included slow growing pits (pit diameter classed into two


cess. In the case of iron corrosion with oxygen depolarization (described earlier in this chapter), we initially considered the reaction to proceed in the mixed kin-etics region and we had evaluated the results accordingly [18, 19]. Only when the activation energy of the apparent kinetic component was measured, did it because possible to arrive at a better interpretation of the results.

5.6 Effect of temperature on the type of corrosion rate control

The activation energy of the chemical step of a heterogeneous reaction con-siderably exceeds the diffusion activation energy. If therefore at a moderate temperature, for instance room temperature, a reaction proceeds under kinetic control, it may happen that with increasing temperature the slow step of the reaction will successively change from activation control, through the mixed-type to pure diffusion control.

An example of this is the corrosion of an electrolytic-copper disk in an air-saturated 1.7 M H 2S 0 4 solution. The relevant results are reported in Fig. 5.10. As can be seen from the data of this figure, the type of control of the copper corrosion with oxygen depolarization varies with temperature. At a temperature of 308 Κ the reaction is activation controlled; at 318 Κ the metal corrodes in the mixed-kinetic control region; and at 338 Κ the rate-determining step in the reaction is the diffusion of the oxygen dissolved in solution to the copper sur-face [22].

5.7 Two simultaneous corrosion reactions

If the rotating disk of the metal under examination corrodes as a result of two parallel reactions, they occasionally can be distinguished. This depends on the type of the rate-determining step involved in these reactions.

If the rate-determining step of both the corrosion processes is a chemical reaction, the rotating disk is clearly unsuitable for the purpose of solving this problem. If, on the other hand, the rate-determining step of both reactions is diffusion, the rotating disk can be used to measure or to evaluate the limiting diffusion fluxes of the reactants belieyed to participate in the corrosion reaction, followed by a comparison of the sum of the fluxes with the measured value of corrosion of the metal.

In the case where one reaction is diffusion controlled, and the other is acti-vation controlled or proceeds in the mixed-control region, the discussion is more complex.

TWO SIMULTANEOUS CORROSION REACTIONS 1 7 5

A good example of this is the corrosion of a metal, for example zinc, in an acidic electrolyte solution with simultaneous hydrogen and oxygen depolar-ization. The former reaction in solutions of a sufficiently low pH values is acti-vation controlled, the latter is diffusion controlled. Thus the corrosion rate of the metal can be written as follows:

*zn*+ = iH+ = Fk(ck+y + io2,lim (5.30)

Since the hydrogen ion is used up in both reactions, its concentration at the disk surface, Ch+, will differ from that in the bulk of solution, c£+. The solution of equation (5.30) therefore resolves itself into determination of the 4+ value. Making use of relations (5.21) and (5.22) we obtain:

CH

+ ι ?Q2,Hm *r, lim /

CH +

C &

+ *H

+,lim * H

+, l i m \

cH

+ (5.31)

The value of câ+ obtained from this is then used in equations (5.30) or (5.21) and (5.22).

An illustration of this argument is the results obtained for zinc disks dissolved in air-saturated 0.1 M NaCl solution at pH = 2.3, shown in Fig. 5.12.

10

c ο ω ο

S 5

ο

ο

ο

~ ο

V 2 π '

. 1 S 2

Fig. 5.12. Dissolution rate of a zinc disk in an air-saturated 0.1 M NaCl solution, pH 2.3, at a temperature of 298 ±0 .2 K. The solid line was plotted on the basis of equations (5.30) and (5.31) and the data of Figs. 5.8 and 5.4. Circles represent the experimental results (after [8]).

In this figure the zinc dissolution rate as calculated from equations (5.30) and (5.31), or (5.21) and (5.22) is shown by the solid line. The data for /r, i im> P* and k, taken from Fig. 5.8, were obtained from a series of measurements in a nitrogen-saturated 0.1 M NaCl solution at different pH values. Similarly, the

1 7 6 ROTATING DISK STUDIES IN THE TRANSPORT OF REACTANTS

limiting current density was measured separately for the hydrogen ions (Fig. 5 . 8 ) and oxygen dissolved in solutions studied (Fig. 5.4).

It should be noted that the straight section of curve of Fig. 5 . 1 2 , on linear extrapolation, intersects the ordinate axis, which is an indication* that in this system, apart from the diffusion-controlled reaction, another parallel reaction occurs which is kinetically controlled.

The solution of equation ( 5 . 3 1 ) is particularly straightforward [ 1 3 ] for first-order (p = 1) or second-order (p = 2 ) reactions. In other cases, equation ( 5 . 3 1 )

is most conveniently solved by the iterative method. If the corrosive medium contains no dissolved oxygen, io l im = 0 and

equation ( 5 . 3 1 ) converts to equation ( 5 . 2 3 ) . If iO2,iim > *H

+,iim> equation ( 5 . 3 1 )

no longer holds and the metal corrosion rate is then controlled by hydrogen ions diffusion. This case has already been treated in a previous section and was illus-trated in Fig. 5.4.

As already pointed out, during corrosion with simultaneous oxygen and hydrogen depolarization, with a diffusion control of the former process and a kinetic control of the latter, the hydrogen ion is used up primarily in the oxy-gen depolarization corrosion. In consequence, the course of the latter process causes the acidic dissolution of metal to slow down. A comparison of equation ( 5 . 3 1 ) with equation ( 5 . 2 3 ) will show that the drop in rate of the corrosion of metal with hydrogen depolarization depends in the first place on the value of the expression / o 2, i i m / * H M . m.

In the borderline case, if ï'o2.iim ~ *H

+.iim» equation ( 5 . 3 1 ) no longer holds,

but at the same time the acidic corrosion of metal vanishes, as all the hydrogen ions are used up in the process involving corrosion with oxygen depolarization. This situation exists in air-saturated solutions at a pH of approx. 4.0. And at a pH ^ 3 . 0 the value of expression /o2,i im /*V , i im

1S negligible compared with

the measuring error. A decline in the rate of acid corrosion of metal in air-saturated solutions can therefore be observed experimentally if, simultaneously, corrosion with oxygen depolarization occurs in solutions of pH ^ 3 . 0 . In the same way, at pH < 3 . 0 for the majority of metals undergoing acidic corrosion, the value of expression irtumlln

+,iim ~ 0 · Since at the same time /o2, u m / 4 i

+, um ~ 0 ,

therefore under these conditions, in equation ( 5 . 3 1 ) CH+/CH+ ~ 1 , and hence in equation ( 5 . 3 0 ) Fk(c^+)

p « Fk(c%+)

p = / r, i l m. Thus it is sufficient to sum up

the separately measured values of the limiting currents of the cathodic reduc-

* This is a necessary but an insufficient condition. A similar result could be obtained, for instance, from a partial coverage of the electrode surface. The question would be resolved by identification of the reaction and determination of the activation energy of the other reaction.

TWO SIMULTANEOUS CORROSION REACTIONS 177

tion of oxygen and the corrosion rate of metal in oxygen-free solutions to obtain the corrosion rate of metal with simultaneous hydrogen and oxygen depolariz-ation.* As already mentioned, this is restricted to the case when the first reaction is kinetically controlled and the second one is diffusion controlled.

This situation occurs in the corrosion of copper containing oxygen for example, at a concentration of 3700 ppm in air-saturated sulphuric acid at a temperature of 338 Κ [22]. As can be seen in Fig. 5.13, the corrosion rate of this metal is

Fig. 5.13. The corrosion rate of copper containing 3700 ppm of oxygen vs. square root of disk speed; air-saturated 1.7 M H 2S 0 4 solutions; 338±0.2 Κ (after [22]).

a linear function of the disk speed; however, on extrapolation the line fails to pass through the origin. This indicates that in this system, apart from a diffu-sion-controlled reaction, another reaction proceeds which is activation controlled.

The overall reaction of copper dissolution with oxygen depolarization can be written thus:

2Cu + 0 2 + 4 H+ = 2 C u

2+ + 2 H 20 (5.32)

and the overall reaction of C u 20 dissolution in the acid is :

C u 20 + 2 H+ = Cu + C u

2+ + H 20 (5.33)

The former reaction occurring at 338 Κ is diffusion controlled, whereas the latter runs in the activation control region. A superposition of these two reactions is

* This also applies to the case of two parallel corrosion processes without a common reac-tant.


represented by the line in Fig. 5.13. The equation of the straight line of Fig. 5.13 has the form:

y C ux l 06 = 12+(6.0+0.8)]/m '

whereas the equation of the line for 338 Κ of Fig. 5.10 is

7 c ux l 06 = (7.2+0.5)j/w

The measurements reported in Fig. 5.10 were made on oxygen-free copper. Apparently in either case the diffusion component of metal dissolution is the same within experimental error.

5.8 Summary

The results described in this chapter fail to cover the whole body of problems related to the use of the rotating disk electrode in investigations of metallic corrosion. Further information on this subject can be found in the monograph by Pleskov and Filinovskii cited previously [2], devoted to the use of the rotating disk electrode in studies on the kinetics of heterogeneous reactions in general and electrode processes in particular.

The experimental data treated in this chapter provide no direct indications as to the techniques of corrosion control. Such techniques, however, are preceded by a profound understanding of the mechanisms of metallic corrosion.

5.9 References

[1] Levich, V. G., Physical Hydrodynamics (in Russian), Nauka, Moskva 1959. [2] Pleskov, Yu. V., Filinovskii, V. Yu., Rotating Disk Electrode (in Russian), Nauka, Moskva

1972. [3] Bierowski, M., Pawelkowa, M., Zembura, Z., Wiad. Chem., 16, 437 (1962). [4] Zembura, Z., Bierowski, M., Wiad. Chem., 18, 215 (1964). [5] Zembura, Z., Opyrchal, M., Corr. Sci., 22, 1097 (1982). [6] Zembura, Z., Maraszewska, Α., Polish J. Chem., 59, 907 (1985). [7] Zembura, Z., Maraszewska, Α., Polish J. Chem., 61, 891 (1987). [8] Zembura, Z., / . Electroanal. Chem., 46, 243 (1973). [9] Pourbaix, M., Atlas of Electrochemical Equilibria, Pergamon Press, Oxford 1966.

[10] Zembura, Z., Corr. Sci., 8, 703 (1968). [11] Kolny, H., Zembura; Z., Roczniki Chem., 46, 301 (1972). [12] Zembura, Z., Kolny, H., Roczniki Chem., 41, 1629 (1967). [13] Zembura, Z., Fulinski, Α., Electrochim. Acta, 10, 859 (1965). [14] Kemula, W., Grabowski, Z., Roczniki Chem., 26, 266 (1952). [15] Scheller, F., Muller, S., Landsberg, R., Spitzer, H. J., / . Electronal. Chem., 19, 187 (1968).

REFERENCES

[16] Landsberg, R., Muller, S., Thiele, R., Acta Chim. Acas. ScL Hung., 51, 85 (1967). [17] Landsberg, R., Thiele, R., Electrochim. Acta, 11, 1243 (1966). [18] Glodzinska, W., Zembura, Z., Ochrona przed korozjq, 16, 297 (1973). [19] Zembura, Z., Ziolkowska, W„ Roczniki Chem., 45, 1053 (1971). [20] Zembura, Z., Glodzinska, W., Roczniki Chem., 42, 1525 (1968). [21] Zembura, Z., Burzynska, L., Corr. Sci., 17, 871 (1977). [22] Zembura, Z., Bugajski, J., Corr. Sci., 21, 69 (1981).

181

6. Displacement as a Particular Case of Corrosion

6.1 Introduction

The precipitation of metals from aqueous solutions of their salts by other metals is called displacement or cementation. It is, in fact, a special case of electrochemi-cal corrosion [1, 2]. In the most commonly encountered corrosion cases the cathodic process involves the reduction of oxygen or hydrogen ions (oxygen or hydrogen depolarization corrosion). In the displacement process, the corre-sponding process is the reduction of metal ions leading to the precipitation of metal. The analogy allows the use of a number of research methods and lines of reasoning related to the mechanism of general electrochemical corrosion for a better understanding and interpretation of the patterns of cementation processes.

Such processes have for long been recognized [1-13]. The covering of iron objects with copper during their contact with mine waters in ore mines was orig-inally believed by alchemists to be one of the proofs of the feasibility of trans-mutation of metals. As early as 1587, however, Ercker wrote that "iron displaces metallic copper from copper solutions, the latter does the same to silver and gold, gold is displaced by silver" [5]. A Treatise on the Philosophers* Stone, published by Michaelis Sendigovius in 1604, specifies in more detail the sequence of precipitating power of metals (Fe, Pb, Sn, Cu, Hg, Ag, Au) [5, 14, 15]. In quali-tative terms therefore this represents the incipience of the electrochemical series of metals, which was not firmly established until about 200 years later, by Volta.

In 1865 Beketov published what is considered to be the first work concerned exclusively with investigations on the precipitation of metals by metals [10, 11].

The studies made by Nernst [16] and Boguski [17, 18] have become a starting point for the majority of semiempirical equations used in displacement kinetics. The concept of an immobile layer (diffusion layer) that persists at the liquid-solid interface even under conditions of mixing of the liquid, was put forward by Nernst (1889). An early quantitative approach to the effect of the surface area on the reaction rate in such a system was made by Boguski (1876).

INTRODUCTION 181

Recently, substantial progress in the studies and description of the kinetics of reactions at the liquid-solid interface has been possible, due to the use of the rotating disk technique. The method is helpful in making interpretation possible in terms of the Levich theory [19]; the factor controlling the reaction rate can hence be unambiguously determined.

An analysis of the semiempirical equations that describe displacement kin-etics, as derived by various authors, is given in the monograph by Rotinyan and Kheifets [2].

Studies on the kinetics and mechanism of displacement process in short-cir-cuit model cells have been originated by Sçdzimir and Pawelkowa [20] and by Sçdzimir [21]. The system used permits a direct measurement of the process rate. The results obtained were the starting point for a discussion on displace-ment kinetics in terms of the theoretical shape of the polarization curves of the component reactions (Evans's diagrams). Evans's diagrams are now commonly used in discussing the effect of the relevant parameters on the rate of the displace-ment process. This is treated at length by Power and Ritchie [22, 23]. Applica-tion of the hypothetical runs of the polarization lines for the discussion of the displacement kinetics is also presented by Miller [24].

Attempts were made by Rickard and Fuerstenau [25] to determine the displace-ment kinetics from polarization curves of the component reactions (cathodic and anodic) determined by independent measurements.

The practical significance of the precipitation of metals by metals dates as far back as Roman times. Pliny [3] describes a procedure for the separation of gold from silver which consists in roasting granulated alloy with sodium chlor-ide, followed by displacement of silver by mercury. A similar principle reported in 1557 by Bartolomeo de Medina [4], underlies the process for silver obtained from Mexican ores. Agricola [6] recommends the separation of gold from silver by prolonged boiling of a granulated alloy in concentrated sulphuric acid, fol-lowed by precipitation of silver by copper after dilution of the solution. He also mentions the mine waters in Hungary which "turn" iron into copper. The re-covery of copper from the waters of the Spanish Riotinto mine by precipitation with iron was a commercial process as early as 1670 [12].

About the mid-19th century, a commercial process employed the precipita-tion of gold and silver by zinc from solutions obtained by cyanide leaching of ores. Today, displacement processes are mainly used for recovery of copper, tin, lead, etc., from spent etching liquors, for the recovery of silver from photo-graphic materials, and for the removal of copper, cadmium and thallium from zinc sulphate solutions to be used in electrolysis. The displacement of copper

182 DISPLACEMENT AS A PARTICULAR CASE OF CORROSION

by less noble metals is still applied in the hydrometallurgy to recover copper from the dilute solutions [24, 26, 27] or from some complex solutions [28],

The process of metal displacement by hydrogen, a commercial process which has been going for about twenty years, may also be classed as displacement

Mention should also be made of the possible use of displacement for ana-lytical purposes [32-34], as well as for metal coating [35, 36] without applied current.

6.2 Equilibria of displacement processes

6.2Λ Thermodynamic displacement equilibria

It is often stated that a metal situated higher in the electrochemical series pre-cipitates another metal that is below in that series (a nobler one) from solu-tions of its salts. This is true, provided that the system involved meets the condi-tions appropriate to the values of the standard potentials £Ό°Μ (the activities of the metals and their ions equal to unity, temperature 298 K). A more general statement is therefore that a more electronegative metal, if brought into contact with a particular solution, will precipitate a more electropositive metal from the latter.

The dependence of the equilibrium potential on the activity of a metal and of its ions in solution is expressed by the Nernst equation:

During the displacement process, the concentration of metal ions being precipitated decreases, whereas the concentration of the precipitating metal ions rises. This results in a gradual drop in potential of the former metal, and a rise of potential of the latter. As one would expect, equilibrium will be attained as the potentials of both metals become equal.

From the known values of standard potentials, the equilibrium ratio of the activities of ions of the precipitating and the precipitated metals can be evaluated (provided the activities of the metals equal unity, i.e. the metals do not occur as solid solutions).

Consider, for example, the precipitation of copper by iron. It is assumed that the sole cathodic reaction is the reduction of copper ions, viz. :

[29-31].

(6.1)

C u2+

+ Fe = Cu + F e2+

(6.2)

EQUILIBRIA OF DISPLACEMENT PROCESSES 183

In conformity with the above reasoning, equilibrium will be attained when the concentrations of ions of both metals reach the values that result in equal equilibrium potentials of both the anodic (Fe -» F e

2+ + 2e) and the catho-

dic (Cu2+ + 2e -> Cu) reactions.

Considering the values-of the standard potentials, the following relation is obtained, corresponding to the cementation equilibrium:

Substituting the Ε ο values at a temperature of 298 Κ one arrives at:

-£*L ~ <£*L = 6.9 x . lO-27 (6.3a)

The value obtained indicates that for practical calculations concentrations can be used instead of activities. The error resulting from disregard of the activ-ity coefficients may affect the value of the logarithm of the quotient, at the most, a few per cent, which justifies this simplification. In practice this error is probably much lower, as both the ionic species are in the same solution, and hence their activity coefficients, being a function of ionic strength, should be similar in value.

It follows from equation (6.2) that if the iron ions appear in solution solely as a result of the displacement process, their concentration can be expressed by equation (6.4), where C 0 C u

2+ denotes the initial concentration of copper ions

Cpe

2+ = Q ) c u

2 +~ C C u

2+ (

6·

4)

Equation (6.3) suggests that the C C u

2+ values is negligible relative to CF e

2+.

Hence CFe*+ = C 0 C u

2 +.

Accordingly, the dependence of the equilibrium concentration can be ex-pressed as follows:

logCcV+ = -26.16 + logCF e« = ~26.16+logCO Cu2+ (6.5)

- ^ ^ = 6 . 9 x l 0 - 27 (6.5a) C 0C u 2+

Relations of this kind, calculated under the conditions specified above, are illustrated in Fig. 6.1 for a number of the commonest displacement reactions. As can be seen from the graph, the concentrations of the zinc-precipitated copper, cobalt and lead ions should be lower than the initial concentrations of these ions by several orders of magnitude. We would therefore have a condition of virtually

184

-2

DISPLACEMENT AS A PARTICULAR CASE OF CORROSION

-1

+10

5 - t I-20

mx +y mx + y

y

pu =0.1 MPa ay = 1 "2

298 Κ

-30

J - A O

Fig. 6.1. Logarithmic dependence of the concentration of reduced ions (which remain at equi-librium with the reducing metal), on their initial concentrations (concentrations in mol/1).

1 1 1 2 1 3 4 5 6 1 7 1 8 j 9

X C u

2+ P b

2+ C o

2+ C u

2+ P b

2+ C o

2+ X

Y Fe Zn Fe Zn Co

m 1 2

η 2 1

complete precipitation. The effect of iron would be less pronounced. In the case where the precipitating metal reacts with a solution which con-

tains ions of two, or more, metals they participate simultaneously in the cathodic reaction. Equilibrium is attained, if the concentrations of the ions produced

log cQX


in the anodic process or reduced in the cathodic processes reach the values cor-responding to the same Nernst potentials of individual electrodic reactions.

Displacement is usually carried out in aqueous solutions, in which, it should be remembered, the reduction of hydrogen ions may be one of the cathodic processes. By way of example, we shall consider the precipitation of copper and lead by iron, from acidic solutions of their salts. Here the anodic reaction is the oxidation of metallic iron into its ionic form, whereas the cathodic reactions are the reduction of copper, lead and hydrogen ions to their elemental form. The initial iron ions concentration is assumed to be zero ( C p e

2+ = 0).

As a result of the proceeding reactions, the concentration of the reduced ions diminishes, whereas that of ferrous ions rises. Taking into account the stoichio-metry of the reactions involved, the ferrous ions concentration which corresponds to the displacement equilibrium can be expressed by :

C F e- = C&a* - C C a 2+ + Cft," - C PV + i ( C S + - C H +) (6.6)

It is known from the previous calculations and from experimental data that the equilibrium concentrations of the reduced species involved in displacement processes are very low (C° > C). Thus equation (6.6) can be simplified to

C F e- = C & « + C ? b+ i C g + = ]Tc° (6.6a)

By expressing equilibrium potentials corresponding to individual electrode reactions in terms of the Nernst equation, and considering the condition for the displacement equilibrium, which involves the equality of these potentials, equa-tions (6.7) are obtained :

Figure 6.2, constructed on the basis of these equations, can be used to read off the equilibrium concentrations of the displacement reaction concerned, ob-tained in solutions of various overall initial concentrations.

It is emphasized that the relations derived here hold good only for cementa-tion equilibria. They do not take into account other possible reactions that may occur with participation of the ions concerned.


Fig. 6.2. Solid lines corresponds to the dependence of the cementation equilibrium concentra-tions of copper, lead and iron ions and of pH on the sum of the initial concentrations of the reduced ions. The broken line represents the iron ions concentration as calculated from the solubility product of ferrous hydroxide, /F E(OH)2 = 1.65 χ 10"

15 (after [37]) as a function

ofpH.

6.2.2 Apparent equilibrium. Inhibition of the displacement process as a result of secondary reactions

Experiments show that on many occassions the equilibria expected from thermo-dynamical evaluations are not attained. Specifically, zinc fails to precipitate iron,* while the precipitation of copper by nickel or iron, or the precipitation of cobalt

* A contrary belief of Shakhov [45] that a displacement process does occur in the zinc/iron system at 80°C raises doubts. Perhaps the decreasing iron concentration in the solution was due to precipitation of the hydroxide as a result of gradual drift to higher pH values. The fact that zinc is used for reduction of iron (III) to iron (II) in quantitative analysis is further evi-dence that zinc does not reduce iron II to iron.


by zinc, can stop at concentrations above those corresponding to equilib-rium [ 8 , 11, 38-41].

Other observations suggest that anions may substantially affect the displace-ment process. As an example, displacement processes involving aluminium, zinc or iron in nitrate solutions occasionally become suppressed as a result of passivation, or due to the probable reduction of nitrate ions in these systems [11, 42]. The precipitation of copper by iron from chloride solutions follows a different pattern than for sulphate solutions [13, 43, 44].

The discrepancies observed are often ascribed to the inaccuracies inherent in the calculations, which are due to the use of the concentrations of ions in the Nernst equation, rather than their activities. As already mentioned, such an approximation changes the equilibrium constant by (no more than) one order of magnitude. This fails, therefore, to account for the lack of displacement in the iron-zinc system (see Fig. 6.1).

It thus appears that elucidation of the reasons for the discrepancies observed calls for a closer look at the mechanism of cementation processes. And in addi-tion, the presence of hydrogen ions in solution and their effect on the establish-ment of equilibrium must be taken into account.

6.2.2.1 Cessation of the displacement process as a result of isolation of the precipitating metal from solution

Interruption of the displacement process at a concentration higher than the equilib-rium one may result from isolation of the precipitating metal as a consequence of a compact layer of the reaction products deposited on its surface. Usually, the metals precipitated as a result of displacement are in the form of loose, pow-dery or spongy, deposits (dendritic grains exhibiting low tenacity). It seems, however, that compact coatings can be obtained by displacement from sol-utions of complex salts or containing some complex-forming compounds. This effect can also be achieved by conducting the displacement with concomitant mechanical action (e.g. rubbing the cementing metal with a paste containing salts of the metal being precipitated or with a cloth soaked in a solution of these salts [35]). Even in these cases, however, because of the simultaneous occurrence of the anodic and cathodic reactions on the same surface, the coatings produced exhibit considerable microporosity. In other words, the displacement products layer may slow down the process but it cannot stop it before equilibrium is attained.

The displacement process is also sometimes interrupted as a result of deposi-tion of an insoluble product of secondary reactions (hydroxides, salts) on the


precipitating metal surface, so that the metal becomes isolated from the solution. This may be the case in the above-mentioned copper precipitation by iron (or copper and lead coprecipitation) from acidic solutions. If the sum of initial con-centrations X]C

0 (cf. Fig. 6.2) is, say, about one mole per litre as a result of the

cementation equilibrium condition established, the expected pH value of the solution should be about 7.3. The iron concentration corresponding to the ce-mentation equilibrium should be equal to unity, whereas the copper and lead concentrations should be 1 0 "

25 and 10~

10 times lower, respectively.

The iron concentration vs. pH relation calculated from the solubility prod-uct of the hydroxide is shown as a broken line in Fig. 6.2. The iron concentra-tion C Fe corresponding to the cementation equilibrium is nearly thirty-fold higher than the C Fe value resulting from the solubility product. It follows from equation (6.8), where / w is the ionic product for water, and / F e( o H > 2 is the sol-ubility product for ferrous hydroxide

28.2

«^Fe(OH)2 e X p j ^ o F e (~/^pJ J ^.8)

It would therefore be expected that precipitation of the hydroxide would start before the cementation equilibrium is set up. The process will proceed on the iron surface i.e. at the place where the metal is converted into its ionic form. The copper and lead concentrations corresponding to the cementation equi-librium were calculated to be ΙΟ

20 and 10

9 times lower, respectively, than

the values derived from the solubility products of their hydroxides (/C U(OH)2 = 5.6χ Ι Ο "

2 0; / P b ( O H) 2 = 1.0χ ΙΟ"

1 5) [37, 46].

In this system, therefore, it follows that the process may be stopped as a re-sult of precipitation of ferrous hydroxide. Copper and lead hydroxides, on the other hand, will not precipitate under these conditions.

Passivation of the precipitating metal is also likely, on account of the deposi-tion of insoluble salts on the metal surface. This would be the case of copper precipitation by lead from sulphate solutions.

6.2.2.2 Interruption of the displacement process due to the establishment of mixed potentials

Zinc is sometimes used in chemical analysis for reduction of ferric to ferrous ions. In this process no displacement of iron is observed, contrary to the results of thermodynamical calculations shown in Fig. 6.1. This is even more unex-pected in view of the fact that zinc precipitates cobalt whose electrochemical properties are very close to those of iron. These facts can be explained in terms


of the reaction involving acidic dissolution of metals, which occurs in parallel to the cementation process. An explanation following this line of reasoning is illustrated in Fig. 6.3. This shows the polarization curves for the dissolution*

Fig. 6.3. Polarization curves corresponding to the processes likely to occur in the systems: (a) zinc—solution containing zinc, ferrous and hydrogen ions; (b) zinc—solution containing zinc, cobaltic, and hydrogen ions. Broken line concerns the component reactions, solid lines are the resultant of polarization curves, Ει and Ji—mixed potential and the current flowing in the system at the first moments upon immersion of zinc in solution, Eu and Ju—mixed potential and the current that establish as a result of the appearance of trace iron (a) or trace cobalt (b) on the zinc surface. The dashed arrowed line indicates the direction in which the anodic polarization curve moves in the case of gradual blocking of zinc surface with a layer of ce-mentation prodUCtS: (a) / I K = / F e + / H 2/ Z n , «ΛΑ = /zn, ΛίΚ = / H 2/ Z n + / H 2/ F e / l I A = Jzn + JFe',

(b) Λ κ = / C O + / H2/ Z i i , JllK — «/CO + / H2/ Z i i + /H2/CO, / l , I I A = Jzn-

of zinc, cathodic deposition and anodic dissolution of iron, as well as the curves corresponding to hydrogen evolution on zinc and iron. The relative positions of the equilibrium potentials correspond to a solution approx. 0.1 M with respect to the acid and approx. 1 M with respect to the metal ions.

If pure zinc is immersed in a solution free of ferrous ions, the anodic reac-tion would be dissolution of zinc, and the cathodic reaction would be evolu-tion of hydrogen. A corrosion potential EC0T and a small (because of a very large overpotential of hydrogen) corrosion current 7c o r, can be seen from the intersection of the polarization curves.

If ferrous ions are present in solution, an alternative cathodic reaction in-volving iron deposition may come into play. The resultant cathodic polariza-


tion curve is obtained by adding the currents corresponding to both the reac-tions. Thus, as seen in Fig. 6.3a, in the first moments of the reaction the system will assume potential El which corresponds to the resultant current Jt. Following hydrogen evolution, iron starts to deposit. The iron that appears on the zinc surface, even in minute quantities, causes the hydrogen overpotential to drop dramatically.* The requirement of equality of the anodic and cathodic currents in systems not connected to an external current source hence causes the po-tential of the system to move towards the Eu value. In this case, as seen from Fig. 6.3a, the anodic current Ju is the sum of the predominant zinc dissolution current and of the dissolution current of the initially deposited iron. The cathodic cur-rent / „ , on the other hand, is concerned exclusively with hydrogen evolution. Thus the reason for the unexpected suppression on thermodynamical grounds, of the process of precipitation of iron by zinc is the fact that a mixed potential, whose value exceeds the equilibrium potential of the cementation process (cal-culated without regard to hydrogen ions present), is set up in the system. In the system considered, iron may only periodically deposit (and subsequently re-dissolve) in minor quantities on the zinc surface.

The above reasoning can also be invoked to explain the experimentally ob-served precipitation of cobalt by zinc, despite the low value of the hydrogen overpotential on cobalt. The standard potential of cobalt is about 0.16 V higher than the standard potential of iron. Accordingly, the polarization curves for cobalt are shifted towards more positive values, compared with the correspond-ing curves for iron. The relevant polarization curves are shown in Fig. 6.3b.

Figure 6.3b suggests that initially, upon immersion of zinc in a cobalt salt solution, the expected cathodic processes are deposition of cobalt and hydrogen evolution, the share of the latter process being negligible. The displacement rate is virtually determined by the current value / j . The appearance of cobalt on the zinc surface brings about a considerable decrease in the hydrogen overpoten-tial. In consequence, the mixed potential Ex moves towards a more positive value of En (which, however, still remains lower than the equilibrium potential of cobalt). Zinc dissolution continues to be the anodic process, and the cathodic current is the sum of the cobalt and hydrogen deposition currents; the former decreases in value, whereas the latter becomes substantially higher. This explana-tion is consistent with the observed vigorous evolution of hydrogen during co-balt precipitation by zinc.

* The effect is much more pronounced than shown in Fig. 6.2a calculations using suitable Tafel's coefficient [47] indicate that at the same overpotential values, the rate of cathodic evol-ution of hydrogen on iron is higher than on zinc by over five orders of magnitude.


A similar line of reasoning leads to the conclusion that even for the displace-ment of copper by iron, the thermodynamically calculated concentrations should not be expected in actual processes. Polarization curves for this system are sche-matically shown in Fig. 6.4. The anodic process is the dissolution of iron, whereas

Fig. 6.4. Hypothetical polarization curves corresponding to the processes that are likely to

occur in various systems: iron—solution containing ferrous, cupric, and hydrogen ions. £Ί...ιν

and /ι.,.ιν are respectively the mixed potentials and the currents flowing in the system at the suc-

cessive moments of the process, /Ι,ΙΙ,ΙΠΚ = / C U + / H 2, /ivtc = Λ<η — Jn2, /I...IV = /pe (descrip

tion of the lines as given in Fig. 6.3).

the cathodic processes are precipitation of copper* and evolution of hydrogen. The hydrogen overpotential on iron is lower than on copper. Thus the appearance of precipitated copper has practically no effect on the polarization curves, and hence on the hydrogen evolution rate. As the copper ion concentration decreases as a result of the cementation process, its equilibrium potential moves towards increasingly lower values (Eicu ^ncu -* ···)· The polarization curve thereby shifts, and the rate-determining limiting current of copper ions diminishes. This is associated with an increase in the percentage share of hydrogen evolved in the cathodic process. When the equilibrium potential of copper ESU attains

Cui rent intensity (/=/S)

0

* The shape of the curves accounts for the experimentally observed fact of copper pre-cipitation in the diffusion control region (i.e. in the limiting currents region).


the value of the isc or potential, the anodic process will be entirely compensated by the cathodic evolution of hydrogen. The copper precipitation therefore will be completely suppressed at a potential higher than that given by thermodynamic calculations.

In the foregoing reasoning concerned with the zinc-iron, zinc-cobalt and iron-copper systems, a number of simplifying assumptions have been made. One of these is that the anodic curve (precipitating metal) remains constan during the cementation process. Actually the curve is displaced, though only slightly, towards more positive potential values. This is due to concentration of ions of the precipitating metal which pass into the solution. The hydrogen ion concentration was assumed to remain unchanged during the displacement process. This condition can be easily met in practice by replenishing the acid, as it is used up, in the acidic dissolution of iron.

Furthermore, it should be borne in mind that the area of both the anodic and cathodic surfaces may vary during the process. The anodic surface may, for instance, become smaller as a result of a partial blocking by displacement products, whereas the cathodic one may increase, since reduction of ions may also occur on the enlarging surface of the metals just precipitated.

6.3 Displacement kinetics

6.3.1 Fundamental kinetics equations

The reaction rate in a solid-liquid system depends upon the ratio of the surface reaction rate to the transport rate of the reactants. In the case where the surface reaction rate is much higher than the rate of transport (the rate constant kt is a few orders of magnitude larger than the diffusion resistance kD = D/δ), then the resultant kinetics are controlled by the transport of ions (the case of diffusion control). In opposite cases the reaction is activation controlled. If kT and kD are of the same order, the situation is referred to as mixed control. The re-spective concentration distributions are shown in Fig. 6.5.

The kinetics of metal-by-metal precipitation have been the subject of numer-ous investigations. The established fact that cementation rate depends on the effectiveness of mixing underlies the concept of the diffusion control of the kin-etics of these processes. There is general agreement that the factor controlling the rate is the transport of reactants (ions of the metal being precipitated) to the reaction surface (Centnerszwer and Heller [48-50]), Plaksin et al. [51-53], Drozdov [38, 54] and a number of other workers [13, 20, 21, 40, 43, 44, 55-59]).

DISPLACEMENT KINETICS 193

Distance from reaction surface

Fig. 6.5. Relation of the concentration of the precipitated metal ions vs. the distance from the reaction surface (precipitating metal surface). C and Cs are concentrations in the bulk of solution and on the precipitating metal surface, respectively, at successive instants t;ô denotes thickness of the diffusion layer; (a) diffusion control—the rate constant for the surface re-action kr is a few orders of magnitude larger than the diffusion resistance ko = D/ô; (b) mixed control—kt and kD are of the same order of magnitude; (c) reaction control—kt is a few orders of magnitude less than kO.

In certain cases the displacement rate is thought in some quarters to be deter-mined by the transport of the reaction products (dissolving metal ions) from the reaction surface into the bulk of solution [60-62].

In the case of diffusion control the concentration of the ions, which are re-duced, is at the reaction surface virtually zero (the ions, which move through the diffusion layer at a fairly moderate velocity, react immediately in the rapid reaction). By assuming a linear drop of concentration across the diffusion layer (a constant value of the concentration gradient at a given time, dC/dx = (C-0/ό)) the displacement rate can be expressed by equation (6.9)

where C denotes the concentration of the precipitated metal ions in the bulk of solution, D is the diffusion coefficient of these ions, S is the reaction surface area (generally identified with the geometrical surface of the precipitating metal), V is the volume of solution, and δ is the thickness of the diffusion layer.

Integration of equation (6.9) from t = 0 to t and from C = C0 to C leads to equation (6.10):

C = C0exp


or logC = logC0-(SD/23VÔ)t (6.10)

In the case of mixed control, the concentration gradient is determined by the equation: dC/dx = (C—Cs)/ô (Cs denotes concentration of the precipitated metal ions at the reaction surface). The kinetic equation therefore assumes the form:

^Mwc~Cs) (6ΛΙ)

Integration of equation (6.11) is not possible as it stands, as the relation between the concentration C in the bulk of solution and the concentration at the reaction surface Cs is unknown, since the latter quantity cannot be measured.

Equation (6.11) becomes integrable on the assumption that the surface reac-tion is of the first order (equation (6.12)) and that both the surface reaction rate and the transport rate are equal. These assumptions lead to equation (6.12):

dm dC KS Q ( 6 1 2) MVdt dt V

kTS „ DS V ( C - O (6.13)

where m is the mass of precipitated metal, and M is its atomic mass. Calculating the value of the surface concentration Cs from equation (6.13)

and substituting it in equation (6.12), one arrives at equation

dC S krD dt V krô+D

integration of which leads to

(6.14)

c - ^ ( - 7 W ' l ( 6 1 5 a )

l o s c = l o ^ - W W * < ( 6 ' I 5 b )

Comparison of equation (6.10) with equation (6.15) indicates that in the case of both diffusion and mixed control, the logarithm of concentration of the metal ions being precipitated should be a linear function of time. The absolute value of the slope of this line in either case should rise with mixing intensity, as the


Time, min 20 40 60

Fig. 6.6. The effect of stirring on the kinetics of copper precipitation from sulphate solutions using low-carbon steel. The stirrer is placed next to the reacting steel plate; temperature 306 K; nitrogen atmosphere (after [13]).

Time, min

I ι 1 •

200 400

Stirrer speed, rpm

Fig. 6.7. The effect of stirring on the kinetics of lead precipitation from a PbCI2 solution in 20% NaCl with Armco iron. The stirrer is placed next to the iron plate; temperature 368 Κ (after [57]). Broken line corresponds to the relation of rate constant (slope of the log C vs. time relation) vs. stirrer speed.

1 % DISPLACEMENT AS A PARTICULAR CASE OF CORROSION

diffusion layer thickness then decreases. In the case of activation control, the logarithm of concentration should also be a linear function of time (provided the reaction is of the first order). However, the slope value in this case corre-sponds to the reaction rate constant and is independent of the mixing rate.

Consider, for example, Figs. 6.6, 6.7 and 6.8. The results, which are con cerned with copper precipitation from sulphate solutions by iron [13] and lead precipitation from chloride solutions by Armco iron [57], support the reason-ing of the kinetics of the process corresponding to equation (6.10) or (6.15).

The linearity of the relation between the logarithm of the precipitated metal concentration and time has also been found in a number of other studies [40, 52-54]. It should, however, be emphasized that some non-linear cases of this dependence are also known [43, 44, 57, 63]. These are treated in subsequent sections of this chapter.

As can be inferred from equations (6.10) and (6.15), from the linearity of the logarithm of concentration vs. time relationship and from the dependence of the slope on mixing rate, it is only possible to establish that either diffusion control or mixed control of the process kinetics takes place. A number of at-tempts have been made to decide which of these mechanisms is predominant in specific cases. One approach has been via a study of the dependence of the rate of the process on temperature or rate of mixing. An unambiguous solution of this problem has been achieved only in certain cases from the rotating disk studies and elaboration of the results in terms of the Levich theory [19].

classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which included slow growing pits (pit diameter classed into two groups: one, which inc


6.3.1.1 Attempted determination of the type of kinetics control from the temperature dependence of the precipitation rate

In the case of diffusion control, the activation energy of the displacement pro-cess is expected to be less than 15 kJ/mol. This is because the activation energy of diffusion of ions in aqueous solutions is of this order. As mixed control takes over, this value should gradually increase to the values typical of chemical reac-tions.

In the majority of cases, the values of the activation energy of displacement reactions are quite small. Specifically, Episkoposyan and Kakovskii, in precipi-tating copper and silver, determined the activation energy to be about 12 kJ/mol [56, 58]. Nadkarni et al, studying the precipitation of copper by iron, found the activation energy to be in the range 12 to 25 kJ/mol [13, 44].

Cases of higher values of activation energy are also known. Hahn and Ingra-ham [60] found the activation energy of palladium precipitation by copper, over a 298 to 318 Κ temperature range, to be about 40 kJ/mol; the process rate here is independent of mixing rate. Raising the temperature to 338 Κ the activation energy decreases to about 8 kJ/mol (the rate becames affected by mixing). This indicates that at low temperatures, the kinetics of the process are determined by the surface reaction rate, whereas with rising temperature the kinetic control is assumed by the transport of the rectants. Drozdov [64] similarly claims that in the process of copper precipitation by nickel over the temperature range below 323 K, the kinetic controlling factor is the rate of the surface reaction, whereas above 348 Κ it is diffusion (these results, however, have raised some objections on methodological grounds). Hamdorf [65], in his study on lead pre-cipitation by iron from chloride solutions, found the activation energy of the process to be about 50 kJ/mol. Data from other authors [57] point to a still higher value of approx. 75 kJ/mol. These rather high values of activation energy would seem to suggest activation control of the process kinetics. This conclusion, however, is contradictory to the simultaneously observed relation of the process rate vs. mixing rate (cf. Fig. 6.7).

Evaluation of the activation energy is essentially sufficient to decide which factor controls the displacement kinetics. At the same time, however, a number of points remain unclear. For example the activation energy value for precipi-tation of copper by iron from chloride solutions is nearly twice that found for sulphate solutions; and the high activation energy of displacement of lead from chloride solutions by iron accompanied by a dependence of the process rate on stirring intensity.

The reasons for these inconsistencies may be due to the neglect of the effect


of the precipitation products and variation of their form, depending on the experimental conditions, on the kinetics. Many investigators [20, 21, 63, 66, 67] have called attention to this. Moreover, it should be realized that the experi-mentally determined slope values of the logarithm of concentration vs. time relations in the case of diffusion control only permit the evaluation of D/δ. In this expression both the numerator and the denominator are temperature dependent (D increases and δ decreases with temperature). Thus the activation energy estimated from the logarithm of the slope of equation (6.9) vs. T~

l is

not, strictly speaking, the activation energy for diffusion. The diffusion activation energy can be evaluated only in cases where the

diffusion coefficient D vs. thickness of diffusion layer relation is known (i.e. when the rotating disk method is used) [56, 58].

6.3.1.2 Attempted determination of the type of displacement kinetics control from the precipitation rate vs. solution mixing rate relation

It is necessary to start with a discussion of the term "mixing rate", commonly used by a number of investigators. In the author's opinion the "mixing rate" is too ill-defined to be of value, as no quantitative criterion of this parameter is available, thus making a quantitative comparison of results of different in-vestigators not feasible. Even for the same stirrer speeds or the same flow rates of the liquid through the measuring vessel, entirely different degrees of mixing may be obtained. This is because the mixing also depends on the geometry of the system, that is on such factors as: the shape, design and position of the stirrer and of the samples under reaction, the vessel size and shape, the volume of solution, and the positioning of the liquid inlet and outlet in the vessel, etc.

Nor is there any reason to believe that even in the same measuring system the reactant transport rate to the reaction surface is proportional to the stirrer speed or to the flow rate of the liquid through the measuring vessel.

Attempts to determine the factor controlling cementation kinetics from the effect of mixing rate, encountered in the literature, are therefore usually of a qualitative nature. They can be used only to distinguish the case of activation control from that of diffusion or mixed control, leaving the latter two undis-tinguishable.

As one would expect, in the diffusion control range the rate constant, evalu-ated from the slope of the line showing logarithm of concentration vs. time, should increase (following a straight or curved line) with the mixing rate; this is because the thickness of the diffusion layer δ in equation (6.10) decreases with mixing rate. On passing from diffusion control to mixed control, the rise in rate


constant should be increasingly slower and the curve should asymptotically approach the value corresponding to activation control. This is illustrated by the results reported by Nadkarni et al [13, 44] in Fig. 6.8.

The rotating disk theory developed by Levich [19], which relates the effect of mixing conditions to the thickness of the diffusion layer δ has made possible a quantitative discussion of the mixing effect on displacement kinetics.

The relevant studies have been made by Episkoposyan and Kakovskii [56, 58], who dealt with precipitation of copper and silver on an iron rotating disk. They found that over the range of variables examined, the displacement rate constants are a linear function of the square root of the disk rotation speed, and the lines plotted pass through the origin of the coordinate system, as shown in Fig. 6.9.

0 1 2 3 A ~yfn , rev 2 s"2

Fig. 6.9. Rate constant k vs. square root of speed of the rotating disk. Precipitation of copper from sulphate (after [56]) or chloride solutions (after [58]) using an iron disk; temperature 298 K; k = kD.

In conformity with the Levich theory, this is indicative of an exclusively diffu-sive nature of the control of the process rate.

An explanation of the faster rate of copper precipitation by iron from chloride


solutions, compared with the analogous process in sulphate solutions, has been reported by Kamecki and Sçdzimir [43]. Precipitation of copper from chloride solutions occurs not only as a result of cementation but it is also due to the for-mation of insoluble cuprous chloride in the reaction of the deposited copper with the solution. This may be inferred from Fig. 6.10 which shows the varia-

E

"Ï0 20 30 ία Time, min

Fig. 6.10. Precipitation of copper by iron from chloride solutions (after [43]). Relation of the

sum of copper and iron concentrations, in mol/1, to the precipitation .time. Initial concentra-

tions Cocu = 0.5; Cote = 0.

tion of total concentration of copper and iron with the time of the process. If cementation were the only process, the total concentration should maintain a constant value. The initial drop of the curve is explained by two reactions, the displacement reaction (equation (6.2)) and the reaction Cu + CuCl2 -> 2CuCl, running in parallel. At a subsequent stage of the process, the precipitated cuprous chloride undergoes reduction to metallic copper: 2CuCl + Fe -> 2Cu + FeCl2. As a result, the curve, having passed the minimum, regains its initial level. In the case of precipitation from chloride solutions the rate constant determined by Episkoposyan and Kakovskii is therefore the resultant value that includes the two parallel reactions.

Using the same method, Strickland and Lawson [63] also found diffusion control of the displacement process but only in its earliest stage. The rate "constant" increases with the quantity of precipitated metal. This effect is explained by the authors by turbulent flow and increased effective surface area of the disk as a result of deposition of the displacement products on its surface. An empirical equation has been advanced to account for the effect of this factor:

DISPLACEMENT KINETICS 2 0 1

v= k0C(l+fim) (6.16)

where ν is the displacement rate, C is the concentration of the metal being pre-cipitated, m is the weight of precipitated metal, and k0 and β are empirical con-stants. The k0 value corresponds to the rate constants at the earliest moments of the process when the reaction surface is in practice still equal to the disk sur-face. The k0 values found from kinetic measurements were compared with the values calculated from the Levich equation on the assumption that the activa-tion energy of the process was 14 kJ/mol. The results, shown in Fig. 6.11, pro-

0.8

£ ^ 2 i5

Cu-Zn Cu-Fe Cu-Cd Cd-Zn P b - Z n Ag-Zn Ag-Cd Ag-Cu Ag - Fe

SOf c r ο · Δ V • 0 ο c 9

Χ

0.6 cm/min Rate constant as evaluated from the

Levich equation

Fig. 6.11. Rate constants k* of the process (equation (6.14)) found from cementation kinetics

measurements (the rotating disk, initial stages of the process) as a function of the values of

the constant k as evaluated from the Levich equation (activation energy A of the process as-

sumed to be 14 kJ/mol); k = 0.622>|

/

9

3

8v-2

1

9

/l

w l / 2 e x p[ ~ ~ ~ ^ ) ]

( a f t er [ 6 3 ] )*

vide strong quantitative evidence in favour of the diffusion nature of the control of cementation kinetics.

The rotating disk method and discussion of the results in terms of the Levich theory is a step forward in studies of displacement kinetics. The following dis-


cussion, however, deals with some problems that have arisen in investigations using the rotating disc and similar methods.

One of the fundamental assumptions of the Levich theory is the laminar flow of solution with respect to the disk surface. This condition can only be accom-plished if the surface roughness is much less than the diffusion layer thickness. This condition is manifestly not fulfilled, if displacement products are precipi-tated on the disk surface. Considerable disturbances in the flow would be ex-pected, if the deposit formed is of a dendritic structure (a powder, or spongy de-posit). If the metal being precipitated, on the other hand, deposits on the disk as a uniform layer, the surface area on which the anodic process proceeds gradually decreases. This may, in extreme case, result in inhibition of the process as a con-sequence of the kinetic control being taken over by the anodic reaction. That this may occur has been pointed out by Power and Ritchie [22].

The condition of continuing smoothness of the disk surface in displacement studies was met only in the case of the investigation of the kinetics of mercury precipitation by copper, carried out by Power and Ritchie [68]. The results obtained under these conditions lend full support to the view that the factor determining the cementation rate is the transport of the metal ions being pre-cipitated. A detailed survey on advantages and also on disadvantages of the ro-tating disk method, for the studies of the displacement kinetics, presents Miller [24].

The form of the deposit (grain shape and size, and their arrangement) de-pends, among others factors, on the rotation speed of the disk. This may be an additional parameter that affects the process kinetics, not accounted for by the theory [67].

The question of the effect of the conditions of displacement on the form of grains of depositing metal has been covered by Barth et al [66]. The effect of the morphology of the deposit on the displacement kinetics of copper by iron has been recently the subject of research by Donchenko and Antropov [62], Annamalai and Murr [67], Miller and Beckstead [69] and others.

Mention should also be made of the inconsistency involved in the opinion as to the diffusion or mixed control of cementation kinetics in the light of the experimental data obtained from prolonged measurements (for ten or so, to several scores of minutes). If the displacement rate were determined solely by transport of the metal ions being precipitated towards the reaction surface (diffusion control), the nature of the precipitating metal would be expected to have no bearing upon the process kinetics. In fact, when precipitating lead, for example, from chloride solutions by zinc, the rate of the process was found to


be over two orders of magnitude higher than for precipitation by iron under similar conditions [40].

The diffusion controlled precipitation of various metals from solutions of nearly equal viscosities is another case in point. Such reactions should proceed at similar rates, because of generally minor differences in the diffusion coeffi-cients of individual ions in aqueous solutions. Episkoposyan and Kakovskii [56, 58], in precipitating copper and silver on a rotating iron disk, obtained the rate constants kD for sulphate solutions of 1.3 χ 10~

6 and 2.9 χ 10~

6 respectively;

and for chloride solutions, the rate constants were 2.06 χ 10"6 and 4.88 χ 10"

6

cm~2s~

1 /2 rev."

1 / 2. The relatively high silver precipitation rate here is striking.

The case of precipitation from chloride solutions (AgCl in concentrated NaCl solutions), where a lower rate would be expected, as silver occurs in the form of large complex ions and the viscosity of the solution is thus higher, is par-ticularly strange.

The diffusion mechanism of the kinetics cannot be invoked to explain the relation between precipitation rate and sulphuric acid concentration foi- copper and silver precipitation by iron (cf. Fig. 6.12). An explanation of these ambi-

-J L 1 ι > ) ι

W 20 30 4 0 * ^ 500 H2SO4 concentration, g/l

Fig. 6.12. Rate constants k for copper and silver precipitation by iron from sulphate solutions vs. H 2S 0 4 concentration (after [56]).

guities is attempted in Section 6.3.1.3. Interestingly, the rate of displacement of copper by iron, evaluated from the independently determined diffusion coeffi-cient of copper ions (Episkoposyan and Kakovskii), is smaller than the experi-mental value.


Mention should also be made of the attempts made by Strickland and Lawson [63] to use a rotating cylinder in displacement kinetics studies. A detailed analy-sis of this approach has been made by Higgs et al. [70]. The empirical equation arrived at agrees well with the relation derived by Gave and Robinson [71]. The equation describes the relation between displacement rate and the speed of rotation, diameter and surface area of the cylinder, solution viscosity, con-centration and diffusion coefficient of the ions undergoing reduction.

Ritchie [70] has compared the displacement kinetics results from the ro-tating cylinder method with those obtained from the rotating disk. He has de-monstrated that the rising rate of displacement process, as the products of the reaction build up on the disk surface, cannot be accounted for solely by a tur-bulent flow of solution [72].

The study made suggests that the displacement rate on a rotating cylinder is independent of the system geometry over a wide range. Comparable results can hence be obtained from experiments carried out in different laboratories.

6.3.1.3 Effect of displacement products on kinetics

Theories concerned with the process of precipitation of metals by metals differ among themselves largely in their views of the effect of displacement products on the subsequent course of these processes; this effect is seen differently by various investigators. The assumptions made form the basis for a number of semi-empirical equations derived to describe the kinetics.

A detailed review of these theories and their ensuing kinetic equations has been made by Rotinyan and Kheifets [2].

The majority of investigators agree that the rate of a displacement process is determined by the rate of transport of the metal ions being precipitated, to-wards the reaction surface. The kinetics can accordingly be described by equation (6.9). The differences in interpretation of this equation are generally due to the character of variations in the diffusion layer thickness or the area of the reaction surface S as the precipitating metal is deposited (and thus indirectly with displacement time).

The first attempt at a quantitative approach to the effect of precipitation prod-ucts on the kinetics of the process was made by Drozdov [54]. He suggested that the more or less porous layer of the precipitated metal would hinder the transport of ions towards the metal surface where the cementation process pro-ceeds. In Drozdov's opinion, the thicker this layer is, the slower is the transport. The layer thickness is proportional to the quantity of precipitated metal. Hence:

δ = fc(Co-C) (6.17)


Having substituted this value into equation (6.9) and integrating it, equation (6.18), similar to that derived by Drozdov, is obtained:

where C0 denotes the initial concentration of the precipitated metal ions, C is the concentration after time t, D is diffusion coefficient, k is a proportionality constant, and V and S are the volume of the solution and the reaction surface area (Drozdov omits the latter two quantities; he also uses somewhat different symbols).

The equation was constructed by Drozdov on the basis of results from in-vestigations of the precipitation of copper by nickel powder. However, as a pow-dery material was used, the area of precipitating metal surface could not be determined. Moreover, this surface area was continuously changing during the experiment, with decreasing diameter of the reacting nickel grains.

The reasoning itself also raises some doubts. The precipitating metal and the one depositing on its surface form a short-circuited cell. Solutions in which such processes are conducted, generally exhibit a good electrical conductivity. Thus it would be expected that both the metals have virtually the same short-circuit potential. Considering that precipitation usually proceeds in the limiting current range, it can be inferred that this potential is not much different from the equilib-rium potential of the anodic reaction. Hence, contrary to Drozdov's view, it follows that the ions transported from the bulk of solution need not pass through the layer of cementation products as far as the precipitating metal surface. The ions would be reduced on the surface of the previously precipitated metal. This has been demonstrated by observations which show that the layer of cementation products builds up on the solution side of the system.

A form of the equation analogous to that obtained by Drozdov can be de-duced, according to Rotinyan and Kheifets [2], on the assumption that the pre-cipitated metal gradually reduces the area of the reaction surface.

Berman [61] claims that reduction of the ions being precipitated occurs on the inner side of the precipitated metal layer. He based his reasoning on the questionable assumption that the factor controlling displacement rate is the transport of the precipitating metal ions from its surface to the bulk of the sol-ution. The concentration of these ions varies, across the layer thickness, from Cs (concentration of a saturated solution of the salt of the precipitating metal) to the concentration C which persist in the bulk of solution. Additionally, by assuming a large volume of solution, he considers the value of C to be constant.

DS kV

r(C0-C) (6.18)


This being the case, the concentration gradient varies solely as a result of the layer of the reaction products that builds up on the precipitating metal surface. On these assumptions, integration of equation (6.9) results in equation (6.19)

C = C0-B\/T (6.19)

An analogous equation was derived by Morozenko et al. [73], who started from similar assumptions but followed somewhat different reasoning. Accord-ing to Rotinyan and Kheifets [2] the equation may be derived by considering the change in the resistance in the short-circuited cell produced by displacement to be the limiting factor of the rate of displacement. This resistance is deter-mined by the resistance of the electrolyte contained in the pores of the layer of displaced products.

It should be mentioned that equation (6.19) is an analogue of the well known Tamman's equation, which describes metal oxidation kinetics at high tempera-tures, and is based on assumptions that are similar to those made by Berman.

A common feature of the displacement mechanisms discussed above, and of their ensuing kinetic equation (6.18) and (6.19), is the concept that the prod-ucts are the retarding factor in the process. It is generally believed that this is the result either of an extended route of ion transport to or from the reaction surface, or of a decreased area of the reaction surface.

These views, however, raise doubts. As pointed out above, both the analysis of the short-circuited cell formed during the displacement process and the experimental results indicate that the reduction of ions takes place on the external surface of cementation products. Therefore, the fact that this layer builds up does not elongate the route of ion transport. Acting as a cathode, the displacement products surface builds up, as an increasing quantity of such products is deposited. If the reaction proceeds in the range of diffusion control, an increased surface area, according to equation (6.9), results in a higher reac-tion rate. Should this reasoning be sound, an analysis of equations (6.9) and (6.10) leads to the conclusion that the following three cases are possible (cf. Fig. 6.13). The results illustrated in Fig. 6.14 which were obtained in the experiments related to the processes in which the displacement products exhibited an es-pecially highly developed surface (loose, very fine grain spongy deposits) seem to confirm the above assumptions. The inflection points on the curves indicate that the rate of the process initially increases, passes through a maximum, and then declines.

An experimental series of studies of lead precipitation by iron was also carried out in a model cell [20]. The current determined in this system is a direct measure


Time Time

Fig. 6.13. An analysis of a possible effect of the cementation product (increased size of the reaction surface) on the rate of the process controlled by diffusion of the precipitated metal ions (a) and on the variation in concentration of these ions in solution (b): / —the cathodic surface remains constant U î ; S = const; C |CSi; the reaction kinetics is described by equations (6.9) and (6.10); 2—the cathodic surface increases with precipitation time t\\ S\\ Q ; CS[ (slower than in the case 1); 5—a large increase in the cathodic surface (spongy, powder, de-posits): / | ; S Î Î ; C S | to a maximum, then drops (Î = increases, f f = increases considerably, I — decreases).

0 10 20 30 40 50 Time, min

Fig. 6.14. The time dependence of concentration of the precipitated metal ions. Solid lines— precipitation of copper from chloride solutions with Armco iron; temperature 298 or 368 Κ (after [27, 43]). Broken lines—precipitation of lead from PbCl2 solution in a 20% NaCl using technical or high-purity (99.99) zinc; temperature 298 Κ (after [57]).


of the displacement rate. The results shown in Fig. 6.15 are a direct proof of the fact that the rate may pass through a maximum. These results thereby lend validity to the concept that displacement products may exert an accelerating effect on the kinetics of the process.

Time, min

Fig. 6.15. Dependence of the current intensity which represents the reaction rate on time and on the Pb initial concentration in the model cementation system: Fe/PbCl2 solution in 20% NaCl; temperature 298 Κ (after [20]).

The derivation of an equation to account for the effect of increased cathodic surface area on the kinetics of the process has been attempted [21]. The best agreement with the experimental results was obtained on the assumption that the relation between the cathodic surface area and time is described by

S = bS0Qxp(at) (6.20)

where S0 denotes the geometrical surface area of the precipitating metal, and a and b are constants. Substitution of equation (6.20) into equation (6.9) fol-lowed by integration provides the concentration vs. time relation expressed by

C = C0 jexp -*5^2- [l -exp(a0]l (6.21)


Verification of this equation is possible on the basis of the results obtained in the model cell. Considering the equation that relates the displacement rate to current intensity in the model cell

^ J J L \ J (6.22) dt \VFj

V

where M is equivalent weight, V is the volume of solution, / is the current in-tensity and F is Faraday's constant and combining equations (6.9), (6.21) and (6.22), one obtains

j = * * D M O Ç , β χρ ^ _ E X P ( A I ) ]J ( 6. 2 3)

By differentiating equation (6.23) and equating the derivative to zero, the de-pendence of Jmx and its corresponding time tmx on the process parameters can be evaluated. It follows from these calculations that Jmx should be a linear func-tion of C0, the logarithm of Jmx should rise linearly with S0, whereas tmx should be a linear diminishing function of the logarithm of S0. The results of experi-ments have been in close agreement with these assumptions [20, 21].

The mechanism that assumes increasing cathodic surface area as a result of deposition of products resolves a number of the inconsistencies mentioned in Section 6.3.1.1. These are treated in turn:

a. Episkoposyan and Kakovskii found the copper precipitation rate on the iron rotating disk to be higher than calculated on the basis of an independently estimated diffusion constant of copper ions. One of the probable reasons for this discrepancy is the fact that in the calculations the value of geometrical sur-face area of the disk was used. In fact, however, the surface on which the catho-dic reaction took place was larger (an additional factor may be the turbulence of flow caused by the appearance of displacement products on the disk surface).

b. In the case of precipitating lead from chloride solutions by zinc a much finer precipitate is produced (larger specific surface) than for precipitation by iron. As a result, in spite of the fact that the diffusion rate of the metal ions being precipitated is the same in either case, the rate (proportional to the product SC) is considerably higher for zinc precipitation [57]. A similar explanation can prob-ably be given also for the nearly two-fold higher rate of precipitation of silver by iron, in comparison with precipitation of copper by iron [43, 56, 58].

c. Iron precipitates copper and silver from chloride solutions at a higher rate than from sulphate [56, 57]. This is due to the fact that in the former systems, much finer deposits are formed. For copper this results from the precipitation


of cuprous chloride as an intermediate product [43, 74], whereas for silver, the reason is that complex ions are discharged.

d. The precipitation rate of copper and silver by iron as a function of acid concentration in solution, shown in Fig. 6.12, is most probably connected with the fact that the simultaneous evolution of hydrogen is a factor that favours the formation of fine cathodic deposits. In highly acidic solutions the metal pre-cipitation process may be suppressed by a strong shift of the hydrogen polariz-ation curve towards higher potential values. This can result in a rise in the share of the hydrogen current in the overall cathodic current. And in addition, blocking of the surface by the hydrogen evolved may take place under these circumstan-ces.

The knowledge of the fact that the cathodic surface on which the displaced metal ions become discharged increases with time has constituted a step forward in the understanding of the mechanism and in a quantitative description of the kinetics of displacement reactions. It should be emphasized, however, that the time dependence of the cathode surface size, assumed in equation (6.20), is not satisfying. Independent studies on the cathodic deposition of powders under a constant potential, which corresponds to the range of the limiting currents of the ions reduced, indicate that equation (6.20) describes the growth of the cathodic surface solely in the first minutes of the process. In subsequent stages, the rate of growth of the area diminishes, since the system tends to reach a constant value of effective surface area [55, 75].

The accelerating effect of the displacement products on the process has re-cently been emphasized by a number of investigators [23, 63, 66].

Rotinyan and Kheifets [2] made an interesting attempt at deriving the equa-tion for cementation kinetics on the assumption that both the cathodic and anodic processes proceed in the activation control range. Their reasoning re-sulted in equation (6.24), where χ is the actual order of the cathodic reaction, and α and β the respective transfer coefficients for the cathodic and the anodic reactions

If χ equals unity, the reaction order w, found from variations in solution con-centration with time, should be less than, or at the most equal to, unity.

Although the majority of displacement processes occur in the diffusion control range, the transition to activation control, postulated by Rotinyan and Kheifets, can also take place. A detailed analysis of the conditions for such a transition has


been made by Power and Ritchie [22]. An effect of this kind has been also found experimentally by Sedzimir and Pawelkowa [20], Sedzimir [21], and by Annamalai and Murr [67],

Ivanov et al [76] deduced an equation that describes the kinetics of displace-ment by metal powder. These authors assumed the rate determining factor to be the formation of nuclei of a new phase (precipitated metal) which gradually falls off the precipitating metal surface. They also assumed a spherical shape of the precipitating metal grains. Considering the decreasing grain diameter, as a result of dissolution, they derived a kinetic equation whose experimental verification, however, is difficult in view of its complexity.

Mention should also be made of a typical dependence provided by Tseft and Klukova [77], By assuming a topochemical character of the reaction, they arrived at an equation which suggests that the logarithm of the logarithm of the C to C0 ratio should be a linear function of time. However, almost any function should achieve linearity in such a coordinate system.

The effect of the presence of displaced copper on the kinetics of its precipi-tation by iron from sulphate solutions was dealt with by Miller and Beckstead [69]. Annamalai and Murr [67], Donchenko and Antropov [62, 78]. These last workers suggested the qualitative concept that the precipitated copper may accel-erate (increased reaction surface) or slow down (hindrance of iron transport into the bulk of solution) the process.

From the foregoing review it follows that, depending upon the assumptions made, the kinetics equations derived and subsequently verified experimentally differ substantially from one another. The following two reasons may account for this paradox:

(a) one of the factors that may well be critical to the form of the kinetics equation is the nature of the precipitated metal [67]. This depends on the phys-icochemical nature of both metals, that which is precipitated and that which precipitates, as well as on the experimental conditions (solution composition, stirring rate, temperature, etc.). It is hence understandable that different workers conducting investigations in different systems arrive at different relation-ships;

(b) in a number of instances the combination of the scatter of results and the flexibility of the equations makes it difficult to arrive at firm conclusions.

Once again it needs to be remembered that the agreement of experimental results with an equation, derived on the basis of an assumed model of a phenom-enon, fails to represent an unambiguous proof of the model. A further veri-fication of the hypothesis can only be made, if there is a method available for


the evaluation of relevant data gathered from measurements by independent methods. One example of this is the comparison of the Levich equation coeffi-cient, determined in the investigation of the reaction kinetics by the rotating disk method, with the value of the same coefficient calculated on the basis of independently determined diffusion coefficient of the reacting substances and of the viscosity of the solution.

Additional verification may sometimes be obtained by analysing to what an extent the effects accompanying the phenomena under study are consistent with the working hypothesis. In the case of studies of displacement kinetics, some additional information may be provided by measurements of the potential of the system during the reaction.

6.4 Displacement kinetics and the time dependence of the potential of the system

Some additional information for the discussion of the kinetics can be obtained from measurements of mixed potential, carried out during precipitation. Un-fortunately, only a few investigators have made such measurements. Miller pre-sents in his monography [24] a short survey of the principles of the application of the hypothetical run of the polarization lines for the discussion of the metals displacement kinetics.

The theoretical polarization curves shown in Fig. 6.3b and 6.4, are useful in forecasting how the mixed potential would change, if the cathodic surface remains unchanged during the reaction. In the case illustrated in Fig. 6.4 the precipitation kinetics is controlled by transport of ions towards the cathode surface. The mixed potential, determined in practice by the point of intersection of the anodic and cathodic polarization curves, should move towards more negative values as the reaction proceeds, to reach the Ely value. This behaviour is the result of the decreasing concentration of copper ions, which lowers the cathodic limiting current and shifts the equilibrium potential (and hence also the polarization curves) towards more negative values.

The reverse change in the mixed potential is also possible. In the case where the precipitating metal is covered with a strongly adherent layer of cementation products, the anodic surface decreases in size. This brings about enhanced ano-dic polarization and may be the reason for the shifting of the mixed potential towards positive values (cf. Fig. 6.3b). A relation of this kind has been experi-mentally observed in the precipitation of a compact cobalt deposit by zinc [40].

In the majority of cases, however, more complex curves of mixed potential during displacement reactions are obtained than those treated above. Figure

DISPLACEMENT KINETICS AND THE TIME 213

6.16 shows examples of the curves obtained during displacement of copper by iron and of lead by zinc or iron. These relations may be explained in terms of the concept of the effect of the cementation products on the precipitation kinetics

Time, min 50 100 150 ^

ι : 1 1 —

Fig. 6.16. The time dependence of mixed potential (potential of the precipitating metal that

remains in contact with the precipitated metal). Precipitation from chloride solutions at 298 Κ

[43, 57],

discussed previously. The polarization curves corresponding to this concept in the subsequent process stages are depicted in Fig. 6.17a and b. For the sake of simplicity, the possible concomitant evolution of hydrogen ions was deliberate-ly disregarded, and the anodic curve was assumed to remain unshifted during the process.

At the outset (t0 = 0) the cementation rate is J0 and is proportional to the product of the initial concentration of the ions being precipitated and to the area of the precipitating metal surface (cf. equation (6.23)). The intersection point of the polarization curves determines the mixed potential value, E0.

As displacement proceeds, the concentration of precipitated ions decreases, and the cathodic reaction surface increases. It determines the resultant rate of the process. If the increase in surface size exceeds the drop in the concentration,


Fig. 6.17. Hypothetical polarization curves at successive instants /. As a result of precipita-tion of cementation products the surface available for the cathodic reaction increases: E° are the equilibrium potentials; Ε are the mixed potentials. The current intensities corresponding to the point of intersection of the anodic and cathodic polarization curves determine the dis-placement rate, at successive instants t0, h ... In (b) the initial concentration of the precipi-tating metal ions is much higher than in (a) (after [21]).

the product CS increases. This entails a rise in the rate of displacement and an increase in the value of the mixed potential (Fig. 6.17a, increase of time from t0 to ti). Considering that, in practice, the concentration of the metal being precipitated tends to zero, it is to be expected that, beginning from a certain instant, the value of CS product will start decreasing. This will lead to a decline in the rate (down to zero) and of the mixed potential down to a constant value corresponding to the equilibrium potential (E£ = E£ in the case of Fig. 6.17).

In the case of high initial concentrations of the metal ions being precipitated (cf. Fig. 6.17b), the rise in the cathodic surface size may result in such high rates (high values of current intensity / ) that the process may become no longer diffu-sion controlled. Thus an initial increase in the process rate and in the potential is to be expected until constant values are reached, which stay unchanged for some time (intersection of the curves beyond the limiting current range—the process rate is then determined by the rate of the electrode reactions). A con-tinuous drop in concentration of the metal ions being precipitated must result, after some time, in a decline in value of the product CS to such an extent that the polarization curves will come to re-intersect in the cathodic limiting current

SUMMARY 215

region (diffusion control). Accordingly, the process rate and the potential wil gradually start decreasing just as in the previous case (cf. Fig. 6.17a). Behaviour of this kind has been clearly seen for certain cases when conducting the precipi-tation in a model system (Fig. 6.15).

Recently a mathematical model of the kinetics of the displacement pro-cesses, which takes into account also migration of ions, has been evolved by Khudenko [79]. This author discusses also possibility of the metal hydride formation in the considered process.

An interesting attempt to determine the displacement rate from independent electrochemical measurements was made by Rickard and Fuerstenau [25], The points of intersection of the cathodic polarization curves (deposition of copper from solutions of various C u

2+ concentrations) and of the anodic polarization

curves (dissolution of iron) led to a relation giving the dependence of the ex-pected rate of displacement of copper by iron, on copper ion concentration. The semi-empirical equation obtained shows a satisfactory agreement with the result of direct determination of the rate (for the reaction surface area remaining virtually constant throughout the experiment). The displacement reaction was also studied by means of electrochemical measurements (Schalch et al [80]).

6.5 Summary

Investigations carried out in the last fifty years have provided a considerable amount of data, which has led to a better understanding of the mechanism of displacement processes and a more precise description of the kinetics.

a. Adoption of the reasoning used in the discussion on electrochemical corro-sion can be used to explain the complete suppression of a number of displace-ment processes at concentrations of the metal ions being precipitated that sub-stantially exceed those expected from thermodynamical considerations. This is usually due to secondary reactions which produce deposits, which separate the precipitating metal from the solution. It may sometimes be the establishment of a mixed potential as a result of acidic dissolution of the metal, simultaneously proceeding with the cementation process, which is responsible for suppression of this process.

b. Theories of the kinetics of the displacement process have undergone a re-markable evolution. The consensus has for long been that the factor that deter-mines the rate of these processes is the transport towards the reaction surface of the metal ions being precipitated. This is supported by both the reaction rate vs. solution stirring rate relationship and the low values of the activation energy.


Nevertheless, some pieces of experimental evidence appear to contradict this hypothesis, particularly the following: a substantially different rate depending on the precipitated and the precipitating metals; a distinct effect of the nature of the anionic species in some systems; and deviation of the logarithm of concentration vs. time observed in extended experiments. These effects are usually explained by the influence of the products on the kinetics of the process. It was originally believed that the precipitated metal, by blocking a part of the reaction surface and by hampering the transport of metal ions to and from the surface, slows down the process. The current opinion is now that the products accelerate the reaction, especially if they deposit in a powdery or spongy form, due to the increased area of the cathodic surface. Such acceleration of the rate occurs where the reaction is diffusion controlled by the ions of the precipitating metal.

c. Attempts to modify the classical equation that describes the kinetics as governed by diffusion control resolve themselves into studies of the effect of the diffusion layer thickness, or of the size of reaction surface, on the mass of the precipitated metal. A number of semi-empirical equations have been derived; these differ substantially among themselves.

d. Major steps forward in the field of investigations on displacement rates are the use of the rotating disk method and the experiments in the model cell. The rotating disk, and the rotating cylinder, are very helpful in obtaining com-parable results at various research laboratories. However, attempts to discuss the results in terms of the Levich theory should be viewed with caution, on ac-count of the difficulty with maintaining the conditions for one of the major assumptions of this theory in experiments; this is that roughness of the disk surface should be much less than the diffusion layer thickness.

The model displacement cell has been useful for a straight measurement of the process rate. Experiments have given direct information concerning the accelerating effect of the products on the process. The results obtained have been helpful in explaining the variations in the potential in the system during the displacement process. Thus some interferences can be made as to the form of the kinetics on the basis of potential measurements.

In recent years attempts have also been made to obtain displacement rates from the results of independent electrochemical measurements (polarization curves). These attempts have so far failed to account for the effect of the reaction products on the subsequent stages of the process. An interesting survey of the techniques of the metal displacement kinetics investigation has been presented recently by Xiang and Ritchie [81].

REFERENCES 217

6.6 References [1] The Hydrogen-Induced Deterioration and Corrosion of Metals (in Polish), collective work

Ed. J. Flis, PWN, Warszawa 1979, pp. 277-302. [2] Rotinyan, Α., Kheifets, V., Theoretical Fundamentals of the Contact Displacement of Me-

tals (in Russian), Leningrad 1979, p. 47.

[3] Mellor, J., A Comprehensive Treaties on Inorganic and Theoretical Chemistry, Vol. Ill, New York 1960.

[4] Binz, Α., Edelmetalle, Berlin 1941. [5] Bugaj, R., Michael Sendigovius, His Life and Writings (in Polish), Wroclaw 1968. [6] Georgii Agricolae de Re Metallica Libri XII, Basilea 1556; after Georg Agricola Zwofl

Bûcher vom Berg- u. Hûttenwesen, Dusseldorf 1953. [7] Luckenbacher, F., Die Chemische Behandlung der Rohstoffe, Leipzig 1872. [8] Kamecki, J., Sedzimir, J., Hutnik, XIX, 383 (1952). [9] Habashi, F., Principles of Extractive Metallurgy, Vol. 2, Hydrometallurgy, New York 1973.

[10] Beketov, N., Issledovaniya nad yavleniyami vytesneniya odnikh elementov drugimi, Kharkov 1865, after [11].

[11] Bykhovskii, Yu., Veller, R., The Fundamentals of Metallurgy (in Russian), Vol. 1, Moskva 1961.

[12] Argall, P., Mining Sci. Press, 93, 111 (1906); after [13]. [13] Nadkarni, R., Jelden, C , Bowles, K., Flanders, H., Wadsworth, M., Trans. AIME, 239,

581 (1967). [14] Hubicki, W., "Michael Sendigovius Theory, its Origin and Significance in the History

of Chemistry", Proc. of the X International Congress of Hist, of Science, Ithaca 1962, Paris 1964, Vol. II, p. 829.

[15] Hubicki, W., Sendigovius (Sedzimir or Sçdziwoj) Michael (1566-1636) Dictionary of Scien-tific Biography, New York 1975, Vol. XI, p. 306.

[16] Nernst, W. H., Zeitschrift f. Physikalische Chemie, 4, 129 (1889). [17] Boguski, J., Kosmos, 1, 528, 576 (1876). [18] Boguski, J., Ber. Wiss. Chem. Ges., 9, 1646 (1876); after [9]. [19] Levich, V., Physicochemical Hydrodynamics (in Russian), Moskva 1969. [20] Sedzimir, J., Pawelkowa, M., Bull. Acad. Polon. Sci., CI. ΠΙ, IV, 717 (1956). [21] Sedzimir, J., Bull, Acad. Polon. Sci., CI. ΙΠ, IV, 723 (1956). [22] Power, G., Ritchie, I., Austr. J. Chem., 29, 699 (1976). [23] Power, G., Ritchie, I., Mod. Aspects. Electrochem., 11, 199 (1975).

[24] Miller, J., "Cementation", in: Rate Processes of Extractive Metallurgy, Eds. H. Sohn., M. Wadsworth, New York 1979.

[25] Rickard, R., Fuerstenau, M., Trans. AIME, 242, 1487 (1968). [26] Fischer, W., Hydrometallurgy, 16, 55 (1986). [27] Biswas, Α., Davenport, W., Extractive Hydrometallurgy of Copper, Oxford 1980. [28] Sahoo, P., Rao, K., Hydrometallurgy, 8, 223 (1982). [29] Forward, F., Mackiw, V., / . Metals, 7, 457 (1955). [30] Peters, E., Hahn, E., Hydrometallurgy (in Russian), Moskva 1971, p. 55; Proceedings of the

International Conference on Metallurgy, Dallas, Texas 1963.


[31] Meddings, B., Mackiw, V., Hydro-metallurgy (in Russian), Moskva, p. 131; Proceedings of the International Conference on Metallurgy, Dallas, Texas 1963.

[32] Schleicher, Α., Todoroff, T., Zeitschrift f. Elektrochemie, 50, 2 (1944). [33] Davies, W., Key, C , Ind. Chem., 20, 544 (1944). [34] Fu, B„ Ure, Α., West, T., Anal. Sci. Chim. Acta, 152, 95 (1983); C. A. 99, 205 191d (1983) [35] Kutzelning, Α., Zeitschrift f. Elektrochemie, 38, 154 (1932). [36] Lowenheim, F., Modern Electroplating, New York 1974. [37] Zhuk, I., Corrosion and Protection of Metals (in Russian), Moskva 1957. [38] Drozdov, B., Tsvetnye metally, 1, 1949; after [11]. [39] Chizhikov, V., Tratsevitskaya, I., Izv. Akad. Nauk SSSR, Otd. Tekhn. Nauk, 1653, 1947. [40] Kamecki, J., Sçdzimir, J., Arch. Gorn. i Hutn., m, 379 (1955). [41] Kamecki, J., Sçdzimir, J., Roczniki Chem., 26, 288 (1952). [42] Glicksman, R., King, C , / . Electrochem. Soc, 100, 580 (1953). [43] Kamecki, J., Sçdzimir, J., Bull. Acad. Polon. ScL, Ser. A, 379, 1951. [44] Nadkarni, R., Wadworth, M., Trans. AIME, 239, 1066 (1967). [45] Shakhov, Α., Zhurn. Fiz. Khim., 8, 525 (1936); after [2]. [46] Physicochemical Handbook (in Polish), PWN, Warszawa 1974. [47] Antropov, L., Theoretical Electrochemistry (in Russian), Moksva 1969. [48] Centnerszwer, M., Heller, W., Zeitschrift f. physikalische Chemie, A161, 113 (1932). [49] Centnerszwer, M„ Heller, W., Roczniki Chem., 15, 525 (1935). [50] Centnerszwer, M., Heller, W., / . Chim. Phys., 34, 217 (1937). [51] Plaksin, I., Suvorovskaya, N., Budnikova, O., Izv. AN SSSR, Otd. Tekhn. Nauk, 1, 131

(1948). [52] Plaksin, I., Suvorovskaya, N., Shikhova, V., Nauchnye Doklady Vysshei Shkoly, Metallur-

giya, 2, 69 (1959). [53] Plaksin, I., Suvorovskaya, N., Izv. AN SSSR, Otd. Tekhn. Nauk, 3, 407 (1949). [54] Drozdov, B., Zhurn. Prikl. Khim., ΧΧΠ, 483 (1949). [55] Bogacz, Z., Sçdzimir. J., Arch. Hutn. XIX, 289 (1974).; [56] Episkoposyan, M., Kakovskii, I., Izv. Vysshihk Uchebnykh Zav., Tsvetnaya Metallurgiya,

9, 34 (1966). [57] Kamecki, J., Sçdzimir, J., Gmytryk, M., Arch. Hutn., I, 95 (1956). [58] Episkoposyan, M., Kakovskii, L, Tsvetnye Metally, 10, 15 (1965). [59] Episkoposyan, M., Izv. AN Arm. SSR, Otd. Khim. Nauk, 17, 447 (1961); after [9]. [60] Hahn, E., Ingraham, T., Trans. AIME, 236, 1089 (1966). [61] Berman, I., Zhurn. Fiz. Khim., 32, 1971 (1958). [62] Donchenko, M., Antropov, L., Vestnik Kievskogo Polit. Inst., 15, 85 (1978). [63] Strickland, P., Lawson, F., International Symposium on Hydrometallurgy, AIME, Chicago,

Illinois, 293, 1972. [64] Drozdov, B., Zhurn. Prikl. Khim., 33, 633 (1960). [65] Hamdorf, C , Proc. Austral. Inst. Min. Met., 199, 1950 (1961); after [9]. [66] Barth, H., Gans, W., Knacke, O., International Symposium on Hydrometallurgy, AIME,

Chicago, Illinois, 1080, 1972. [67] Annamalai, V., Murr, L., Hydrometallurgy, 4, 57 (1979). [68] Power, G., Ritchie, I., Electrochim. Acta, 22, 365 (1977). [69] Miller, J., Beckstead, L., Trans. AIME, 4, 1967 (1973); after [67].

REFERENCES 219

[70] Higgs, J., Lawson, F., Power, G., Ritchie, L, Austr. J. Chem., 30, 395 (1977).

[71] Gave, D. , Robinson, D., Electrochim. Acta, 17 (6), 1121, 1129 (1972); Chemical Abstracts, 77, 36985, 36986 (1972).

[72] Cornet, I., Lewis, W., Kapesser, R., Trans. Inst. Chem. Eng., 47, T222 (1969); after [70]. [73] Morozenko, E., Antonov, S., Gorodskii, Α., Ukr. Khim. Zhurn., 41, 1127 (1975); after [2]. [74] Basinski, Α., Roczniki Chem., 14, 31 (1934). [75] Sedzimir, J., Bogacz, Z., Szymanska, H., Arch. Hutn., XX, 193 (1975). [76] Ivanov, V., Lisovskii, D. , Tekev, V., Izv. Vuzov, Ser. Metallurgiya, 2, 159 (1965). [77] Tseft, Α., Klukova, V., Vestnik AN Kazan. SSR, 183, 1968; after [2]. [78] Donchenko, M., Antropov, L., Zhurn. Prikl. Khim., 45, 291 (1972); after [62]. [79] Khudenko, B., Water Sci. Technol, 17, 719 (1985); C. A. 103, 57, 145a (1985). [80] Schalch, E., Nicol, M., Balestra, P., Stapelton, W., National Inst, for Metallurgy, Johan-

nesburg, S. Africa, Report, 1799, 1976; after (70). [81] Xiang, J., Ritchie, M., Hydrometallurgy, 16, 301 (1986).

7. Effects of the Interaction of Hydrogen from the Gas Phase with the Surface of Transition Metals

7.1 Adsorption of hydrogen on the surface of transition metals

Transition metals, especially those of the Group VIII of the Periodic Table of the Elements, exhibit a considerable surface reactivity towards hydrogen. The hydrogen molecule undergoes dissociation and chemisorption on their surfaces as H adatoms, which may then take part in a number of catalytic reactions (or electrocatalytic reactions in the particular case of electrode processes) with oxygen, nitrogen, carbon monoxide, or hydrocarbons, etc. Moreover, hydrogen can penetrate into the metal and can dissolve in it to yield either disordered solutions or ordered systems known as hydrides. The adsorption of hydrogen may, of course, affect physical and chemical properties of the parent metal, both in its bulk and surface phases.

A knowledge of the forms of the adatoms and admolecules on the surface of a given metal under specific conditions of gas pressure, temperature and, especially, the structure and composition of the adsorbent surface, is of paramount significance for the interpretation of the results of studies on the subsequent history of the adsorbate and of its chemical reactivity or its ability to penetrate into the structure of the metal-adsorbent.

A tremendous advance in the science of the solid-gas interface in recent years, which has been made possible by the development of novel research methods and of commercially available apparatus, has furnished a wealth of experimental data in this field. This advance is paralleled by the progress in theoretical research observed in the pertinent world literature (e.g. [1]).

The problem of the variety of the effects of interaction of hydrogen from the gas phase with the transition metal surface has already grown to a proportion that makes its comprehensive treatment impracticable in this brief treatise. Accordingly, our intention has been to touch on some selected aspects of the

224 EFFECTS OF THE INTERACTION OF HYDROGEN

problem, largely on the basis of the results of our own research in the Institute of Physical Chemistry of the Polish Academy of Sciences.

A gas molecule, in approaching a metal surface, at a close distance from it (below 0.1 nm) undergoes an interaction with this surface that is strong enough for the adsorption effect to appear. Dispersion forces that are operative at a dis-tance as short as approx. 0.5 nm, though weak in nature, cause physical adsorp-tion, which results in the formation on a solid surface of non-localized, mobile molecules that form a kind of a two-dimensional liquid. A molecule that gets closer to the metal surface (0.15-0.3 nm), overcoming the activation energy barrier, may become bound to this surface by forces that are much stronger and whose nature is characteristic of covalent or ionic chemical bonds. In fact, this is a chemical reaction, referred to in this case as chemical adsorption or simply chemisorption.

Physically adsorbed hydrogen maintains its molecular form. The H 2 ad-molecules, however, exhibit a minor electric polarization, which is positive with respect to the metal surface.

The chemisorption on the transition-metal surface, on the other hand, which results in the development of chemical forces between the adsorbate and the adsorbent, is preceded by a complete or partial dissociation of molecules such as H 2, 0 2, H2S, N 2 or hydrocarbons. As a result, their atoms or radicals are chemisorbed on the metal surface. Cases of dissociation-free molecular chemi-sorption are rare, the notable exception being the adsorption of CO on such metals as Group IB metals, as well as on cobalt, nickel, or the platinum-group metals at room temperature.

An exceptionally strong interaction of the metal surface with the gas mole-cules, which results in their dissociation and the formation of a generally strong chemisorptive bond, is typical of transition metals owing to the presence of lo-calized and specifically arranged d orbitals, which are available to form a chemisorptive bond with an adsorbate.

Hydrogen dissociated and chemisorbed on the surface of a transition metal forms a layer of adatoms with a characteristic covalent chemical bond with the adsorbent. This bond is negatively polarized with respect to the metal over a wide range of surface concentrations of hydrogen. This effect is characteristic not only of hydrogen but also of other atoms whose ionization potential is high, compared with the electron work function of the metal involved.

In spite of the fact that the energy of a chemisorptive hydrogen-transition metal bond is considerable (it amounts to approx. 60-200 kJ/mol [2, 3]), the hydrogen adatoms may exhibit a surface mobility, diffusing from one specific

ADSORPTION OF HYDROGEN ON SURFACE OF METALS 225

adsorption site to another. The activation energy of surface diffusion is generally much lower than the adsorption energy, which is the reason for this behaviour.

If a chemisorption process occurs on a strictly specified surface of a solid, i.e. on a specified plane of its single crystal, here again several forms of the ad-sorbate are generally observed. These forms differ in the way they are bound to the surface and by the nature of their surface structures. Both the structure of the adsorbent surface and the adsorbate pressure and temperature influence the diversity of the adsorbate forms and the population of the adatoms (or admolecules) of which they are composed.

The research methods currently recognized and generally used make it possible in most cases to describe the geometrically ordered structure of the surface adsorption layer formed, the dipole moment of its component molecules, their electron structure and the energy of chemisorptive bond involved, together with the surface concentration of the adsorbate.

In the case of chemisorption of hydrogen, however, the experiment usually fails to provide a description of the adsorption layer, even on the single-crystal surface. The method of low energy electron diffraction (LEED) is useful in demonstrating the structure of the surface layer and its interpretation, but usually only for elements of higher atomic number (or for molecules). Nor is Auger-electron spectroscopy (AES) of any use in detecting the presence of hydrogen on the surface, a method that is recognized in studies of composition of surface layers. However, thanks to the fact that hydrogen readily adsorbs on the surface of transition metals with a virtually zero activation energy, the adsorption energy, characteristic of the metal-hydrogen chemical adsorption bond, can be directly determined from the value of the desorption activation energy.

For this purpose, the thermodesorption method is widely used. In this method an adsorbate-covered solid is heated in a controlled manner under reduced press-ure. An analysis of the quantitative desorption process recorded as pressure vs. time, allows one to distinguish the individual forms or conditions of the adsorb-ate, to derive the desorption kinetics equation and, in consequence, to determine the value of the desorption activation energy. The order of the reaction provides data on the molecular (for the first order) or atomic (for the second order) form of adsorbate [4].

An illustration of the use of this method in a study of the forms of the hydro-gen adsorbed on nickel is the spectrum of hydrogen desorption from a thin nickel layer, as recorded by Nowacki and Dus [5] and shown in Fig. 7.1. The spectrum was obtained during the heating of a previously sintered, thin (50 nm) nickel film with hydrogen adsorbed on it at 78 K, at a heating rate of 10 deg/min.


0.81

06

120 170 22ΪΓ 270 ~ 320 370 420 470* Temperature. Κ

Fig. 7.1. Thermodesorption curve of hydrogen-nickel system (Ni film 50-60 nm thick); hydro gen was previously adsorbed at 78 K; the desorbing hydrogen pressure, pn2, is given in ar-bitrary units (after [4]): A—weakly adsorbed molecular hydrogen, D and £"-—adatoms of hydro-gen adsorbed on nickel.

The least strongly adsorbed molecular hydrogen form A is the first to be de-sorbed; at higher temperatures the more strongly bound hydrogen adsorbate forms, D and E, successively desorb. The latter forms have a desorption acti-vation energy (corresponding to desorption energy) of ED = 41 kJ/mol H 2 and EE = 75 kJ/mol H 2. Both these regions, D and E, correspond to a desorption of hydrogen atoms, which are, however, linked differently to the nickel-ad-sorbent surface.

Of a number of the transition metals investigated which do not absorb hydro-gen to yield a hydride phase, only in the case of platinum did Dus find that hydro-gen adatoms occur on the surface. These were positively polarized relative to the metal surface. In this form hydrogen adatoms occur on the platinum surface above a certain value of their surface concentration [6],

In addition to the physisorbed molecular form of hydrogen on the surface of transition metals, a molecular, but chemisorbed, form of hydrogen may also occur. This form is polarized positively (on nickel, platinum, or niobium) or negatively (on iron) and is invariably weakly bound to the surface. This form appears in the case where the hydrogen ad- or absorption process is substantially developed, at a higher hydrogen pressure [6-8].

The electron structure of the adsorption bond can be studied by infrared spectroscopy in transmission or reflection [4]. Unfortunately, the hydrogen-metal bond can rarely be recorded. An exception is the case of hydrogen adsorp-tion on platinum that is highly dispersed on a silica surface [9]. The result ob-tained by Ratajczykowa by reflection adsorption infrared spectroscopy [4, 10] is interesting. It indicates the presence of two hydrogen sorbate forms on (or perhaps in) the palladium film surface. A more promising method for this pur-pose may be electron energy loss spectroscopy (EELS) [11].

FORMATION OF HYDRIDE PHASES 227

7.2 Adsorption of hydrogen and formation of hydride phases

7.2.1 Intraduction

A particular effect of the interaction of hydrogen with a number of metals is the absorption of hydrogen in their bulk, which results in the formation of a solid binary solution of a definite concentration under given ρ and Τ conditions. However in the case of certain transition metals, as well as lanthanides and acti-nides, above a specific limiting hydrogen concentration an individual, particular phase of a given metal-hydrogen system is formed, known as the hydride phase. Certain metals (e.g. Zr, Y) are capable of forming more than one hydride phase.

The disordered solid solution of hydrogen in a metal is called the alpha phase, whereas the hydride phase is referred to as the beta phase. Any further hydride phases are denoted as gamma, delta, etc. In the hydride phase, the hydrogen occupies interstitial positions in the lattice of the parent metal, thereby forming a kind of an ordered alloy of two metals, one of which is hydrogen. It occurs in the hydride lattice as the proton, even though it is screened, to some extent, by electrons.

Hydrides of the following transition metals have been found to exist: in Group III of the Periodic Table—Se, Y, lanthanides and actinides; in Group IV—Ti, Zr, Hf; in Group V—V, Nb, Ta; in Group VI—Cr, Mo; in Group VII—Mn; and in Group VIII—Fe, Co, Ni, Rh, Pd.

Hydrogen in metals has been the subject of numerous monographs and pro-ceedings of conferences [12].

The problem of the formation and properties of hydride phases has been the continuing objective of a vast number of studies. The consequences of the changes in properties of the parent metal following absorption of hydrogen and its possible conversion to a hydride are of considerable interest. These may be critically important, among others in chemical technology, where the metal catalyst may change its catalytic properties owing to the "hydride" hydrogen; also in materials engineering, in connection with the altered mechanical prop-erties; and, at some future date, in power engineering, owing to the supercon-ductivity of hydrides [13] or an ability to force hydrogen.

7.2.2 Study of the metal hydride formation by measurement of the metal surface potential and by infrared spectroscopy

In a study of the formation and detection of this particular form of hydrogen (the "hydride" hydrogen), in the surface layer of the analogous methods have been used as for investigation of the forms of adsorbed hydrogen. By measuring


the variation of the metal surface potential, Dus [7, 14] recorded the formation of the metal surface and incorporation into its lattice of a considerable number of hydrogen atoms polarized positively with respect to metal. He ascribed the observed effect to the formation of a hydride phase of the metal. The study was made on polycrystalline films of nickel or palladium deposited at 78 K. A simi-lar behaviour of hydrogen was noted for niobium and cobalt [8, 15].

Figure 7.2 shows results obtained by Dus [7, 14] which illustrate variations

200x10

17 ' 400x1ο

17

The number of adsorbed H2 molecules

Fig. 7.2. Variation in the surface potential, ASP, of a thin palladium film deposited at 78 Κ during increasing hydrogen adsorption at 78 K. The figures provided next to and above the curve denote the gaseous hydorgen pressure over the palladium-hydrogen system studied (after [11]).

in the surface potential of a palladium film as induced by successive amounts of hydrogen introduced onto its surface. The initial hydrogen is seen to be po-larized negatively (i.e. the chemisorbed adatoms are negatively polarized), the subsequent portions cause the surface potential to shift towards positive values (on account of a positive polarization of the hydrogen sorbate), and finally the system attains the relative value of the surface potential that is characteristic of palladium hydride. The measurements were carried out at 78 K, which made it possible to show clearly the complexity of the effect of interaction of hydrogen with palladium.

The "hydride" hydrogen can also be observed by means of the thermodesorp-tion spectrum, which shows that this hydrogen forms chemisorbed as adatoms.

Ratajczykowa has recorded the "hydride" lattice hydrogen in the surface layer of palladium in an infrared spectrum using its reflection-adsorption vari-ant (RAIR) [10]. This was the first achievement of this kind in the spectroscopy of the adsorption systems to be published.

FORMATION OF HYDRIDE PHASES 229

7.2.3 Factors affecting the kinetics of hydrogen sorption by metals

In the research series discussed above some factors affecting the kinetics of hydro-gen sorption by metals were of interest.

The recognized effect of accelerating the rate of hydrogen transport through the hydrogen-metal interface occurs in the case of adsorption of the hydride compounds of S, Se, Te, As, Sb, and Bi, known as promoters of hydrogen pen-etration into metals or transfer catalysts [10, 16], These compounds dissociate completely or partially on the surface of transition metals; thus the actual pro-moters are the corresponding adatoms or radicals produced from such com-pounds.

From a study of the variation in the electron work function for thin iron or nickel films as induced by adsorption of hydrogen sulphide in minor quantities, followed by co-adsorption of hydrogen it was found [16, 17] that the work function decreases, thus the surface potential of the metal rises (cf. Fig. 7.3),

25 75 125 175 225 275

"H2S<

P Q , ML

Fig. 7.3. Variation in the electron work function, Αφ, for a thin nickel film during dissociative adsorption of hydrogen sulphide (after [1<3). The relation between the work function and the surface potential is Αφ = — ASP; /iH 2s—the quantity of adsorbed H 2S is given as a product of pn2sx Vh2s (at 293 K) of the H2S introduced onto the film surface Q>H2S is given in Pa, Vn2s in ml).

which is an indication of the presence of positively polarized hydrogen. As men-tioned previously, the formation of the metal hydride phase is invariably linked with the presence of hydrogen adatoms thus polarized [7, 8, 14].*

* The much earlier results of studies of the kinetics of hydrogen permeation from the gas phase through iron membranes showed a much higher rate of permeation of the metal by the protons (compared with atoms) generated by high-frequency electric discharges in rarified hydrogen [18].


Studies on the effect of carbon monoxide on the hydrogen adsorbed on the (111) surface of a Pd single crystal demonstrated that carbon monoxide causes hydro-gen adatoms to enter the film of palladium atoms or to penetrate beneath this film and to remain there until the CO admolecules are removed [19].

The effect of co-adsorbed foreign molecules on the adsorption and hence on the surface reactivity of hydrogen (or other gases) on transition metals has been the subject of considerable experimental and theoretical research effort. The interaction of promoters of hydrogen penetration is just one example.

A quite different factor that affects the hydrogen sorption rate is the degree of dispersion of the metal sorbent. The mechanism of acceleration of the hydro-gen transport through the gas-metal interface with increasing dispersion of the metal is entirely different from the mechanism of action of the promoters. In-tuitively one would expect the hydrogen absorption rate to be higher, when dispersion of the metal is considerable.

These expectation are confirmed by experiment. In the case of polycrystalline nickel specimens, Bovier et al [20] found "mosaic" hydrogen saturation and conversion of the nickel specimens to the hydride phase. Fine crystal grains had been transformed into the hydride (beta phase), while the larger ones still remained as a hydrogen solution in nickel (alpha phase).

By varying palladium dispersion over a wide range (by precipitating its crys-tals on y -Al203 from PdCl2 solutions of various concentrations), Frçk [21] studied the hydrogen sorption isotherms. The mean size of palladium crystal grains in individual specimens examined was 1.3-3.0 nm as determined by a num-ber of methods such as: chemisorption, electron microscopy, small-angle X-ray scattering (SAX), and X-ray diffraction line broadening. For metal of such a high dispersion a very fast attainment of sorption equilibrium was observed together with complete reproducibility of the run of sorption isotherms studied at a temperature of 300 Κ by a volumetric method. No disintegration of the crystal grains, as induced by hydrogen charging, which is so characteristic of large nickel or palladium grains, was observed [22, 23]. On the other hand, from consideration of the form of the desorption isotherms it was found that for the systems of such a high dispersion the hydrogen pressure over the hydride phase of palladium is higher than over macrocrystallites of this metal at the same temperature.

Investigation of a highly disperse palladium subjected to repeated treatment with hydrogen revealed a new and interesting effect. This palladium films, from 12 to 50 nm in thickness, characterized by a completely random orientation of crystal grains became strongly textured with {111} orientation with respect

CATALYTIC EFFECT OF HYDRIDE PHASE HYDROGEN 231

to substrate upon repeated conversion to the hydride phase, followed by its decomposition at room temperature [24]. This effect is additionally associated with a point of considerable practical significance, which is a several times higher catalytic activity in hydrogénation reactions of the material thus obtained. A more favourable energy state of the palladium crystal grains can be invoked to explain this effect. Such crystal grain reconstruction is usually effected at high temperatures, which facilitate the diffusion of metal atoms and their rear-rangement. Here the repeated dilatation and shrinking of the palladium crystal lattice by the alternate formation and decomposition of the hydride plays an analogous role at as low as room temperature.

In the same way, in the case of an initially homogeneous Ni-Cu alloy sample of a high Ni concentration, a repeated hydride formation, followed by decompo-sition, resulted in the separation of two phases at room temperature: one, rich in nickel, the other, rich in copper, of a composition thermodynamically expected for the Ni-Cu system. This was hitherto unattained experimentally in view of the fact that a homogeneous alloy was obtained of a composition correspond-ing to the supercooled condition [25].

7.3 Catalytic effect of the hydride phase hydrogen

7.3.1 Introduction

As mentioned previously, a number of transition metals dissolve hydrogen in appreciable amounts to form a solid solution known as the alpha phase. Under specific temperature, pressure and composition conditions for certain metals (Pd, Ni, Ti, V, Ta etc.) a phase transition may take place to produce a new phase, referred to as the beta phase9 which is often of a complex crystallographic struc-ture and of variable H/Me stoichiometry [12]. The hydride beta phase differs in a number of physical properties from the parent metal in consequence of a changed electron an geometrical structure. Also, the catalytic properties engen-dered by these factors may likewise undergo a change. For a reaction with hydro-gen participation (hydrogénation, dehydrogenation, hydrogenolysis), a special role may be played by the lattice (or "hydride") hydrogen located at interstitial positions, which are octahedral or tetrahedral, depending on the particular. Mobility of this hydrogen in the lattice is often considerable; thus it may participate in the reactions mentioned in parallel the hydrogen adsorbed on the surface, as is the classic case of common metal catalysts.

If, therefore, the hydride beta phase is used as a catalyst of a reaction in-volving hydrogen, a change in catalytic activity must be taken into account.


This change is due, on the one hand, to an altered electron structure relative to the parent metal and, on the other hand, to the consequences of participation of the active lattice hydrogen in the reaction, which affects its mechanism and kinetics ([19] and the literature cited herein).

7.3.2 Catalyst poisoning

As early as 1938, Farkas [26], in his study of the kinetics of hydrogen diffusion through palladium, noted that for high concentrations of hydrogen dissolved in palladium the activation energy of conversion of para hydrogen on this metal is higher than in the case of low hydrogen content. This was probably due to the formation of palladium hydride at high hydrogen concentrations, but the author gave an alternative explanation of this effect.

The work by Couper and Eley reported in 1950 [27] has been of prime significance in the elucidation of the role of the hydrogen dissolved in palladium in the para hydrogen conversion. These authors investigated the catalytic activ-ity of palladium-gold alloys, finding an increased activation energy with the gold content in the alloy. The authors accounted for this effect in terms of the filling of the d band of palladium by the gold s electrons. Although, from the present-day standpoint, this interpretation seems unsatisfactory, because the electronic structure of alloys cannot be oversimplified as in the above model, and because the surface composition must have differed from the mean alloy compo-sition, it is remarkable to note that the activation energy for palladium increased with the quantity of hydrogen dissolved. Figure 7.4 illustrates the results obtained for a palladium wire over a temperature range of 170-350 K. The activation energy under a pressure of 0.16 kPa is 14.6 kJ/mol, and under a pressure of 0.81 kPa it is 28 kJ/mol. When the palladium wire was exposed to atomic hydro-gen, thus increasing the proportion of hydrogen in the metal, the activation energy increased to 46 kJ/mol (pHi = 0.16 kPa). The authors assumed that in the thin surface layer the concentration of the dissolved hydrogen is high enough for the formation of surface palladium hydride, which is responsible for the rise in activation energy. Couper and Eley accounted for the mechanism of the poisoning effect of the "hydride hydrogen as follows: the palladium d band may be filled by the s electrons of hydrogen atoms, thus increasing the activation energy in much the same way as the gold s electrons do in Pd-Au alloy.

Another very important paper was published by Scholten and Konvalinka [22], who investigated the kinetics and mechanism of para hydrogen and ortho deuterium conversion and of hydrogen-deuterium equilibration on palladium


1000/Γ.Κ"

1

Fig. 7.4. The Arrhenius plot for the para hydrogen conversion on palladium wire (after [25]): O—Ph2 = 0.16 kPa; A—pH2 = 0.81 kPa, • — u p o n interaction with atomic hydrogen.

and on palladium hydride sponge or wire. In the case of para hydrogen, the temperature dependence of the reaction rate was measured for several pressures (0.8-35 kPa), and a ten- to hundred-fold drop in the rate was observed on the Arrhenius line near the point where the temperature and pressure suggest a possible alpha to beta phase transition. In the alpha phase region the activa-tion energy varied from 26 kJ/mol for pHi = 0.8 kPa to 40 kJ/mol for Ph2 = 35 kPa, in agreement with the results reported by Couper and Eley. For beta palladium hydride the activation energy was high, amounting to 52 kJ/mol. It should be noted that beta palladium hydride was obtained at a tem-perature higher than the critical temperature for the Pd-H2 system to avoid decomposition of the crystal grains, which accompanies transition through the coexistence region of both phases. This may increase the sample surface thus making it impossible to compare the results.

The poisoning effect of the hydrogen absorbed in palladium or gold-defi-cient Pd-Au alloys was also observed by Dickens, Linnett and Palczewska in a reaction involving recombination of atomic hydrogen [28]. The study was carried out on metal foils at room temperature and the recombination coefficient was measured; is the ratio of the number of collisions of hydrogen atoms with the surface that result in recombination to the total number of collisions. The method for determination used in the measurements, known as the Smith-Linnett method (detailed, for example, in reference [28]) requires a low total pressure of hydrogen of the order of 10""

2 kPa, and the degree of dissociation

of the molecules in the radio frequency electromagnetic field is approx. 10%.

234 EFFECTS GF THE INTERACTION OF HYDROGEN

The authors observed a lowered catalytic activity of palladium and its alloys with gold, induced by atomic hydrogen. The recombination coefficient for palla-dium is reduced tenfold upon saturation of the foil with atomic hydrogen, which may result in the formation of beta phase at least in the surface layer and, in consequence, in a fall in catalytic activity.

Recombination of atomic hydrogen was also used as a test reaction in de-termining the catalytic activity of nickel and nickel hydride under comparable conditions [29]. In view of the equilibrium position in the nickel-hydrogen system, nickel hydride is thermodynamically unstable outside the high pressure region. Nickel hydride can be produced by electrolytic saturation of foils with cathodic hydrogen in the presence of certain promoters, for example thiourea [30]. The nickel hydride thus obtained can be identified by X-rays and stored at low temperatures. The β nickel hydride prepared by this procedure was used as a catalyst of the 2 Η ^ H 2 reaction at a temperature of 195 K. From the determi-nation of γ and decomposition of nickel hydride in situ by allowing it to warm up, the recombination coefficient on nickel was measured. At the same tempera-ture, the value of was found to be an order of magnitude lower on the hydride than on nickel alone. Even partly decomposed hydride had an activity close to that of the original /?-Ni-H specimen.

Similar results were obtained for this reaction also for copper nickel alloys with a low copper content [31]. Alloys underwent poisoning more readily upon their conversion to the hydride and also they recovered their activity more easily upon decomposition of the hydride. This is probably due to an easier entry of

1 2 3 4 5 1000/Γ,Κ"

1

Fig. 7.5. The Arrhenius plot for recombination of atomic hydrogen on Ni-Cu alloy containing 23% by wt. copper (after [32]).


hydrogen into the lattice, as the change (a rise) in lattice parameters connected with hydride formation is less than in the case of nickel.

Apart from the nickel-copper alloys in the form of foils, the recombination of atomic hydrogen was also studied on films of these alloys [32]. The films were prepared by vacuum deposition of Ni and Cu under a pressure of 10"

7 kPa

onto the inner walls of a glass cylinder at 720 K. The films were made homoge-neous by heating for some time at constant temperature, and then transferred to the Smith-Linnett apparatus, where the recombination coefficient was meas-ured. A typical Arrhenius plot for an alloy containing 23% Cu is illustrated in Fig. 7.5. As seen from the plot, below room temperature the catalytic activity drops sharply as a result of a higher energy of activation. This drop could be expected to be due to the formation of the hydride phase under these conditions. This supposition is rendered more likely by the finding by Palczewska and Janko [33] that nickel or Ni-Cu hydride can be prepared by the treatment with atomic hydrogen of fine metal films approx. 100 nm thick on glass at low temperatures. Formation of the /3-hydride phase was identified by X-ray methods or, as in [34], by means of recording variations in the electrical resistance of the film when treated with atomic hydrogen.

The kinetics of para hydrogen conversion on thin nickel and nickel hydride films produced in situ was investigated in [34]. A nickel film 20-40 nm thick was deposited at 77 K, and the conversion rate was measured over the 165-195 Κ temperature range and under pressures varied between 0.13 and 3.9 kPa. The nickel film was subsequently exposed to atomic hydrogen produced from electro-deless discharges to convert it to nickel hydride, as determined from variations of electrical resistance of the film. The conversion rate was measured again in the same temperature and pressure range. The conversion rate on nickel hydride was found to be approx. 300-fold less than on the initial nickel film. Here the activation energy remained unaltered, and the fall in the activity was entirely due to a decrease in the value of the preexponential coefficient in the Arrhenius equation.

The poisoning effect of hydrogen in the form of hydride has also been ob-served in the case of ethylene hydrogénation. This reaction was studied at 233 Κ in a stationary apparatus by Leszczynski et al [35], who followed the pressure drop associated with the progress of the reaction. A thin nickel film deposited on the reaction vessel wall, or this film previously treated with atomic hydrogen, was used as the reaction catalyst. Figure 7.6 shows the drop in pressure with time. The reaction rate constant per unit mass was found to have been decreased by an order of magnitude, as the nickel film was exposed to atomic hydrogen.


4.00

0 2 4 6 8 10 12 14 Time, min

Fig. 7.6. Kinetics of ethylene hydrogénation on thin nickel (open symbols) and nickel hydride (solid symbols) films; m denotes weight of film (after [35]).

From the results of the studies reviewed here it follows that the hydrogen incorporated into the nickel or Ni-Cu alloy structure in the form of /Miydride acts as a catalytic poison, reducing the activity of the catalyst in various reac-tion such as parahydrogen conversion, recombination of atomic hydrogen, and ethylene hydrogénation. A similar effect occasionally occurs for palladium hydride. It should, however, be pointed out that nickel hydride, as distinct from PdH, can be used only under conditions, far from equilibrium in the Ni-H system, and rather strenuous conditions are required for its preparation. Never-theless, the effect of hydrogen treatment of a nickel catalyst and the subsequent weak catalytic activity often encountered in practice (see, for example, Emmett's study [36]), could be due to the formation of a surface hydride that acts as a catalytic poison.

733 Catalyst activation effect

As mentioned in Section 7.3.1, it may also happen that the lattice hydrogen, also known as the hydride hydrogen, directly participates in the catalytic reaction

CATALYTIC EFFFCT OF HYDRIDE PHASE HYDROGEN 237

by affecting its mechanism and its rate. If the reactivity of this hydrogen (pri-marily dependent on the Me-H bond energy in the lattice) is high, the conver-sion of metal to its hydride, and the accompanying process of hydride decompo-sition which is associated with the appearance of active hydrogen, may cause the catalyst activity in a given reaction to increase.

An experimental proof of this possibility along with its interpretation was presented for the first time by Borodzinski et al. [37]. In a study of acetylene hydrogénation on a P d / A l 20 3 catalyst, the catalytic activity of the /?-PdH phase is much higher than that of palladium, whereas its selectivity towards ethylene, a reaction intermediate, is substantially reduced.

A direct quantitative relationship between the phase composition and the rate of the ethylene hydrogénation reaction on thin palladium or palladium hydride films was observed by Fr^ckiewicz et al. [38]. The rate of this reaction was measured in a differential flow reactor operating at the same time as an X-ray camera [39], with simultaneous determination of the catalyst phase compo-sition, as found from the height of the corresponding peaks on the X-ray diffrac-tion patterns. The catalyst used was in the form of a thin palladium film, approx. 50 nm thick, and deposited on glass. The phase composition of the film could be changed by varying the hydrogen partial pressure in the reaction mixture at a constant temperature. In Fig. 7.7 the reaction rate is plotted against the rela-tive content of the hydride phase at 348 K. As can be seen, the reaction rate on

Fig. 7.7. The ethylene hydrogénation rate versus phase composition of the catalyst (after [38]): * P d H —mole fraction of the hydride phase, O—increasing hydrogen pressure, Δ —decreasing hydrogen pressure.


palladium hydride is about ten times higher than on palladium. The hysteresis occurring here results from the fact that the phase composition, as determined by X-rays, does not coincide with the surface composition, which may be differ-ent for the same xPdH, depending on the direction of the phase transition im-posed by the rising or falling hydrogen pressure (marked with arrows in Fig. 7.7). It has been supposed that hydrogénation occurs by means of the active "hydride" hydrogen following the overall scheme;

2Pd + H 2 -> 2Pd-H (beta phase) (7.1)

2Pd-H + C 2H 4 -> C 2H 6 + 2Pd (7.2)

An analogous scheme may also be postulated for acetylene hydrogénation [37]. Under steady-state conditions the phase composition depends on the rate con-stants of both partial processes. If the rate constant of process (7.1) is much higher than that of process (7.2), as is the case of ethylene, the phase composition is not affected by the chemical reaction. The reverse is true in the case of acety-lene hydrogénation, and the phase composition shifts towards pure palladium, as observed experimentally [37, 38].

Catalytic properties of metal-hydrogen-metal hydride systems were exten-sively reviewed in [19, 40].

7.4 References

[1] Roberts, M. W., McKee, C. S., Chemistry of the Metal-Gas Interface, Clarendon Press, Oxford 1978; Samorjai, G. Α., Chemistry in Two Dimensions: Surfaces, Cornell Univ. Press, 1981; Lâzniéka, M. (Ed.), Physics of Solid Surfaces, Elsevier Scientific Publ. Co., Amsterdam 1980; Bénard, J. (Ed.), Adsorption on Metal Surfaces, Elsevier Scientific Publ. Co., Amsterdam 1983; Renouprez, A. J., "Chemisorption on Metals", in: Catalyse par les Métaux, CNRS, Paris 1984; Ertl, G., Ber. Bunsengesel Phys. Chem., 90, 284 (1986).

[2] Bond, G. C , Catalysis on Metals, Academic Press, New York 1962. [3] Wedler, G., Brocker, F. J., Surface Sci., 26, 454 (1971). [4] Anderson, R. B., Dawson, P. T. (Eds.), Experimental Methods in Catalytic Research,

Vol. Ill, Academic Press, New York 1976. [5] Nowacki, P., Dus, R., Ann. Chem., 51, 103 (1977). [6] Dus\ R., Tompkins, F. C , / . Chem. Soc. Faraday Trans. I, 71, 930 (1975). [7] Due, R., / . Chem. Soc. Faraday Trans. I, 70, 877 (1974). [8] Due, R., Surface Sci., 52, 440 (1975). [9] Eley, D . D . , Morgan, D . M., Rochester, C H . , Trans. Faraday Soc, 64, 2168 (1968);

Primet, M., Basset, J. M., Mathieu, M. V., Prettre, M., / . Catal, 28, 368 (1973). [10] Ratajczykowa, I., Surface Sci., 48, 549 (1975). [11] Brundle, C. R., "The Application of Electron Spectroscopy to Surface Studies", in: Surface

and Defect Properties of Solids, Vol. 1, Chem. Soc , London 1972. [12] Smith, D. P., Hydrogen in Metals, Univ. Chicago Press, Chicago 1948; Smialowski, M.,

REFERENCES 239

Hydrogen in Steel, Pergamon, Oxford 1961; Mackay, Κ. M., Hydrogen Compounds of the Metallic Elements, Spon, London 1966; Libowitz, G. G., Binary Metal Hydrides, New York 1965; Lewis, F. Α., The Palladium I Hydrogen System, Academic Press, New York 1967; Andersen, A. F., Maeland, A. J. (Eds.), Proceedings of the International Symposium, Geilo, Norway 1977, Pergamon Press, Oxford 1978, pp. 19, 33; Alefeld, G., Vôlkl, J. (Eds.), Hydrogen in Metals, Springer-Verlag, Berlin 1978, Vol. 1, pp. 101, 131, 201.

[13] Skoskiewicz, T., Phys. Status Solidi (B), 59, 329 (1973). [14] Due, R., Surface Set, 42, 324 (1973). [15] Dus, R., Lisowski, W., Surface Sci., 61, 635 (1976). [16] Smialowski, M., Jarmotowicz, H., Catal., 1, 165 (1962); Kussner, Α., Wicke, Ε., Z. Phys.

Chem. (Neue Folge), 24, 152 (1960); Wicke, F., Meyer, Κ., Z. Elektrochem., 66, 675 (1962); Kûssner, Α., Ζ. Phys. Chem. (Neue Folge), 64, 225 (1969).

[17] Karpinski, W., Palczewska, W., Szymerska, I., Bull. Acad. Polon. Sel, sér. sci. chim., 19, 27 (1971).

[18] Palczewska W., Ratajczykowa, L, Bull. Acad. Polon. Sci., sér. sci. chim., 24, 673 (1966). [19] Palczewska, W., "Catalytic Properties of Metal Hydrides", in: Hydrogen Effects in Ca-

talysis, Eds. Z. Paâl, P. G. Menon, Marcel Dekker Inc., New York, in press. [20] Bovier, C , Janko, Α., Michel, P., Proc. 4th. Eur. Reg. Conf. Electron Microscopy, 1968. [21] Frak, R., Ph.D. Thesis, Institute of Physical Chemistry of the Polish Academy of Sciences,

Warsaw 1979. [22] Scholten, J. J. S., Konvalinka, J. Α., / . Catal, 5, 1 (1966). [23] Janko, Α., Szummer, Α., Bull Acad. Polon. Sci. sér. sci. chim., 14, 885 (1966; Janko, Α.,

ibid., 10, 612 (1962). [24] Janko, Α., Szymerska, I., Palczewska, W., / . Catalysis, 61, 264 (1980). [25] Palczewska, W., Majchrzak, S., Bull. Acad. Polon. Sci. sér. sci. chim., 17, 681 (1969). [26] Farkas, Α., Trans. Faraday Soc, 32, 1967 (1938). [27] Couper, Α., Eley, D. D., Disc Faraday Soc, 8, 172 (1950). [28] Dickens, P. G., Linnett, J. W., Palczewska, W., / . Catal, 4, 140 (1965). [29] Palczewska, W., Frackiewicz, Α., Bull. Acad. Polon. Sel, sér. sci. chim., 14, 67 (1966). [30] Baranowski, B., Smialowski, M., / . Phys. Chem. Solids, 12, 206 (1959). [31] Palczewska, W., Frackiewicz, Α., Karpinski, Z., Bull. Acad. Polon. Sci., sér. sci. chim.,

17, 678 (1969).

[32] Karpinski, Z., Palczewska, W., Frackiewicz, Α., Bull Acad. Polon. Sel, sér. sci chim., 21,

203 (1973). [33] Palczewska, W., Janko, Α., Thin Solid Films, 18, 105 (1973). [34] Frackiewicz, Α., Poleski, M., Palczewska, W., / . Catal, 40, 184 (1975). [35] Leszczynski, Α., Frackiewicz, Α., Palczewska, W., Proc Symp. Mechanisms Hydrocarbon

Reactions, Siofok, Budapest 1973, p. 187. [36] Emmett, P. H., Hall, W. K., / . Phys. Chem., 63, 1102 (1959). [37] Borodzinski, Α., Due, R., Frak, R., Janko, Α., Palczewska, W., Proc. Sixth Intern. Congress

on Catalysis, London 1976, Eds. G. C. Bond, P. B. Wells, F. C. Tompkins, The Chem. Soc., London 1977.

[38] Frackiewicz, Α., Fra_k, R., Janko, Α., Roczniki Chem., 51, 2395 (1977). 139] Borodzinski, Α., Janko, Α., React. Kinet. Catal Lett., 7, 163 (1977); Zielinski, J., Boro-

dzinski, Α., Appl Catal, 13, 305 (1985). [40] Palczewska, W., Adv. Catalysis, 24, 235 (1975).

8. Electrochemical Aspects of Hydrogen Entry into Iron and Steel from Aqueous Solutions

8.1 Introduction

There are many processes in which metals may absorb considerable amounts of hydrogen derived from aqueous environments. For instance, acid pickling, electroplating and other various surface treatments may introduce hydrogen into a metal. In turn, hydrogen may enter into a metal in service by circumstances that lead to hydrogen evolution, for instance cathodic protection or any type of corrosion with hydrogen or mixed depolarization.

The entry of hydrogen into the metal phase is the necessary step in the pro-cess resulting in degradation of metals (hydrogen embrittlement). Much of interest is attributed to the deleterious effect of hydrogen on iron and its alloys. High-strength steels are particularly susceptible to this type of corrosion.

Experiments show that the entry of hydrogen into iron or steel depends on many environmental factors such as the potential of the electrolyte/metal in-terface, the current density of hydrogen evolution, the pH of the surrounding electrolyte, as well as the electrolyte composition, including impurities. Certain substances may strongly accelerate this process and hence they are referred to as promoters of hydrogen entry.

The mechanism of the entry of hydrogen and the effect of promoters on this process has not been completely elucidated. This is undoubtedly due to the com-plexity of the metal/aqueous electrolyte boundary. Particularly complex condi-tions occur on the corrosive and heterogeneous surface of iron and its alloys, on which localized corrosion can take place. The electrode potential and the pH existing inside pits or cracks can differ so much from those at the outer surface that even anodic polarization does not completely remove hydrogen evolution localized at these sites [1]. Consequently, pitting, stress corrosion cracking, and hydrogen embrittlement can overlap each other in some cases [2, 3],

The first topic covered in this chapter concerns techniques used for the esti-

DETERMINATION OF THE RATE OF HYDROGEN ENTRY 241

mation of the rate of hydrogen entry into metals. Subsequently, views on the mechanism of hydrogen entry are discussed, including the correlation of this process with the hydrogen evolution reaction. Considerable attention has been paid to the hydrogen entry promoters, especially to the most potent ones, be-longing to Groups VA and VIA of the Periodic Table. Finally, the action of inhibitors of hydrogen entry is treated briefly.

8.2 Determination of the rate of hydrogen entry

An unambiguous definition of the process of hydrogen entry into a metal offers some difficulty. This is because there has been no conclusive opinion as to what is the exact form of hydrogen passing the electrolyte/metal interface. Generally, the entry of hydrogen constitutes one stage of a complex process which involves conversion of the hydrogen, in the form of a hydrated proton or of a water molecule in electrolyte, to the hydrogen dissolved in a metal.

The rate of hydrogen entry can be defined in terms of the hydrogen concentra-tion just beneath the surface of the metal, which remains in contact with the electrolyte. Because this concentration cannot be measured directly, the rate of hydrogen entry must be assessed from measurements of other quantities, namely :

(a) the quantity of the hydrogen absorbed by the specimen in a given time; (b) physical and mechanical properties of the metal which are affected by

hydrogen; (c) the permeation rate of hydrogen through a membrane. Results obtained in such measurements are not equivalent and, moreover,

they are not always a quantitative characteristic of the hydrogen entry itself. This question requires a brief discussion.

Thus the amount of hydrogen absorbed may depend not just on the entry rate, but equally on a shape of the specimen and the time of hydrogen charging. This time may be insufficient to attain the steady state which corresponds to hydrogen saturation of the specimen under given circumstances. Other problems relate to the fact that the hydrogen absorbed by iron or steel may occur in at least three different forms, specifically;

I—a component of the interstitial solid solution; II—hydrogen connected with structural defects, especially γ/ith dislocations;

III—molecular hydrogen accumulated in crevices and blisters. Practically, the concentration of only the first of these forms may depend

directly on the charging conditions. Concentrations of forms II and III primarily depend on the properties of the metal itself; its structural defects and its sus-

242 ELECTROCHEMICAL ASPECTS OF HYDROGEN ENTRY INTO IRON

ceptibility to the formation of cracks and blisters. If the proportion of forms II and III in the total amount of absorbed hydrogen is considerable, the experi-mentally found hydrogen content in the metal provides information on the ability of the metal to trap hydrogen, rather than on the rate of hydrogen entry.

The forms of hydrogen mentioned above differ among themselves in their mobility. This fact may in turn be the reason for errors in the determination of hydrogen content. The quantity of hydrogen absorbed is very often measured by means of vacuum extraction at elevated temperatures. The mobility of form I may be so high that during the transfer of the hydrogen-charged specimen to a suitable apparatus, part of the hydrogen may escape. Conversely, on account of the relatively low mobility of form III, even a high temperature and extended extraction time may be insufficient to release the whole of the hydrogen from the specimen.

As regards effects of hydrogen on the physical and mechanical properties of a metal, such as electrical resistance and ductility, these effects may depend primarily on the character and the extent of the irreversible changes in structure (cracks and blisters). Measurements of variations in the mechanical properties of a metal as induced by hydrogen (tensile, and bending tests) have hitherto been commonly used for the quantitative estimation of hydrogen embrittlement. However, the results of these measurements may fail to correlate with the rate of hydrogen entry.

The corroding or cath-odically polarized side of membrane

Aqueous solution

(HjO)

H30+

Fig. 8.1. Diagram showing the electrochemical method for studying the permeation of hydrogen through a membrane.

DETERMINATION OF THE RATE OF HYDROGEN ENTRY 243

The most useful method for studying entry of hydrogen is the measurement of hydrogen permeation through a membrane, in particular by the electrochemical technique developed by Devanathan and Stachurski [4, 5]. In this technique one side of the membrane is brought into contact with a hydrogen source, for example by being cathodically polarized in the solution under study, whereas the other side is anodically polarized at constant potential in an alkaline solution to ionize the hydrogen emerging from the membrane (Fig. 8.1). The anodic current density is a highly sensitive measure of the hydrogen flux. The accu-racy of this technique depends on the value of the background current, i.e. the current measured when the membrane is not subjected to hydrogen charging. Usually the density of this current does not exceed 1 χ 10~

3 A/m

2, which indi-

cates that the hydrogen flux can be determined with an accuracy of l x l O-8

mole H/m2s = 6 χ 10

15 atom H/m

2s.

One must consider however to what extent the measured hydrogen flux through the membrane depends exclusively on the hydrogen entry process, because hydrogen permeation may be treated as being composed of three main stages:

(1) entry of hydrogen into the membrane; (2) diffusion of hydrogen in the membrane; (3) exit of hydrogen from the membrane. The above three stages may be characterized formally by the respective re-

sistances: hydrogen entry Rl9 diffusion R2, and exit R$ resistances. Under steady-state conditions the hydrogen permeation rate Ρ should be proportional to the reciprocal of the sum of these three resistances

1 (8.1)

(Rt+R2 + R3)

In the case of iron or steel membranes, in order to accelerate the oxidation of hydrogen and to reduce the corrosion current of the exit side, this side is initially coated with a thin palladium layer. If a suitable anodic potential is maintained, the whole of the hydrogen emerging from the membrane immediately undergoes ionization. Therefore, the concentration of hydrogen just below the exit side <?L is zero and the hydrogen exit resistance R3 can thus be disregarded.

The diffusion resistance is defined by the ratio L/D, where L is the thickness of the membrane and D is the diffusion coefficient of hydrogen in the metal. For very thin membranes the diffusion resistance may be small in comparison to the hydrogen entry resistance (R2 <^ Rt). In others words, the permeation rate would be independent of the membrane thickness and is determined by the entry rate of hydrogen. This is a convenient situation in the study of the hydro-


gen entry stage. Usually, however, the membranes employed are of such thickness that R2 > Ri, which means that the permeation rate is determined by the diffu-sion rate in the membrane. In this case the quantity of hydrogen m passing through the membrane of thickness L and surface area S in time /, when equi-librium is established, equals

m = SD ^ Z ^ t (8.2)

where c0 and cL are the concentration of hydrogen in the metal at the entry and exit side, respectively, and the hydrogen diffusivity D is constant (Fig. 8.1). The permeation rate is expressed as the quantity of hydrogen passing through a unit area in unit time

m = D(c0-cL) St L

P = - = \ (8-3)

In the electrochemical technique the pereamtion rate ip is more conveniently expressed in units of electric current density

i = zFP (8.4)

where ζ is the ionic charge and F is Faraday's constant. Since ζ = 1 and cL = 0,

ip = DFco/L . (8.5)

In the above equations only the concentration c0 is directly dependent on the rate of hydrogen entry. Thus, when studying the effect of the electrochemical conditions at the electrolyte/metal interface on hydrogen entry the remaining parameters, i.e. S, D and L should be kept constant. The problem resolves itself into the use of a constant-thickness membrane made of a material that is re-sistant to the formation of blisters or cracks. Both blisters and cracks may con-siderably diminish the effective cross-section of the membrane S and the effective diffusivity of hydrogen D.

Non-metallic inclusions may play an essential role in nucleation of hydrogen blisters. As demonstrated experimentally, the susceptibility of steel to develop blisters results primarily from the orientation of the inclusion layers relative to the hydrogen flux [6]. If the membrane was machined out from a steel rod perpendicular to its axis (i.e. normal to the direction of plastic working) then the hydrogen flux ran along the non-metallic inclusion layers, so that the blister-ing was strongly suppressed and the diffusion of hydrogen was not disturbed. The reverse was true for membranes which were cut from rolled sheets or strips, as was the case in the majority of investigations.

HYDROGEN ENTRY AND HYDROGEN EVOLUTION 245

Thus none of the available quantitative techniques provide direct information on the stage of hydrogen entry into the metal phase. The most convenient and commonly used electrochemical permeation technique also requires certain conditions to give reliable results.

8.3 Hydrogen entry and hydrogen evolution

Absorption of hydrogen by a metal during its cathodic polarization or occur-ring in corrosion processes is directly linked to the hydrogen evolution reaction. The quantity of hydrogen absorbed usually makes a small part of the total quan-tity of hydrogen evolved.

8.3.1 Hydrogen evolution reaction

It is generally, recognized that the overall reaction of hydrogen evolution from acid solutions

2 H 30+ + 2e -> H 2 + 2 H 20 (8.6)

and from alkaline solutions

2 H 20 + 2e -> H 2 + 20H~ (8.7)

proceeds in two successive stages. The first stage, which occurs on all metals, consists in discharge D of hydronium ions in acid solution

H 30+ + M + e -> M H a ds + H 20 (8.8)

or water molecules in alkaline solutions

H 20 + M + e -* M H a ds + OH- (8.9)

and leads to the formation of adsorbed hydrogen atoms M H a d s. The second stage of the hydrogen evolution reaction may involve chemical desorption (CD)

M H a ds + M H a d8 -> H 2 + 2M (8.10)

or electrochemical desorption (ED) (a) in acid solutions

M H a ds + H 30+ + e -> H 2 + H 20 + M (8.11)

(b) in alkaline solutions

M H a ds + H 20 + e -» H 2 + OH"" + M (8.12)

The individual stages may have different rates. The rate of the slowest stage controls the rate of the overall reaction and determines the overpotential. In


order to determine the mechanism of the hydrogen evolution reaction, it should first be established whether the discharge of hydronium ions or water molecules (D) is followed by the chemical desorption (CD) or the electrochemical de-sorption (ED). Next, the slowest stage determining the rate of the overall reac-tion should be identified.

Bockris and Reddy [7] and Bockris and Subramanyan [8] have distinguished the following six possible mechanisms of the hydrogen evolution reaction (from acid solutions) :

1. Slow discharge—fast chemical desorption

H 30 — M H a ds = H 2 (8.13)

2. Slow discharge—fast electrochemical desorption

H 30+

- ^ M H a d s^ H 2 (8.14)

3. Fast discharge—slow chemical desorption

H 30+ = M H a ds ™ H 2 (8.15)

4. Fast discharge—slow electrochemical desorption

H 30+ = M H a ds ™ H 2 (8.16)

5. Coupled discharge—chemical desorption

H 30 + M H a d s 4 ^ % ^ H 2 (8.17) v1 -> 0 v2 -> 0

where ν is the rate of the overall reaction, v± and vx are the rates of the discharge stage to the right and to the left, respectively, and v2 and v2 the corresponding rates of the chemical desorption.

6. Coupled discharge—electrochemical desorption

B>

0 +i ^

M H> * j ^ r

u>

( 8·

1 8)

where ν3 and v3 are the respective rates of the electrochemical desorption. In the case of mechanisms 1 to 4, it is assumed that the fast stage is at equi-

librium. This means that the rate of the forward reaction equals the rate of the reverse one ( ^ ) . In mechanisms 5 and 6 the rates of both the reverse reactions are negligible, whereas the rates of both the forward reactions are equal.


On the basis of electrochemical parameters such as the Tafel slope, the cover-age of the electrode surface by hydrogen atoms, the exchange current density, the transfer coefficient, or the stoichiometric number, it can be established which of the above mechanisms is probable. The results of electrochemical measure-ments, however, are affected by additional factor (current density, electrolyte composition) and often are inconsistent. Therefore, it is not surprising that dis-crepancies occur in the literature. Only in some cases, notably for Hg, Pb, Ni, W, Au, and Rh is there agreement with regard to the mechanism of hydrogen evolution [7]. The hydrogen evolution reaction on iron is not sufficiently under-stood. Because in the case of iron and its alloys both hydrogen evolution and absorption occur simultaneously, studies on hydrogen entry can provide valuable information on the mechanism of the hydrogen evolution reaction.

8.3.2 Hydrogen entry

One of the parameters which critically affects the rate of hydrogen entry into metals is the cathodic current density. As demonstrated by Bodenstein as long ago as 1922 [9], the permeation rate of hydrogen through metallic membranes is directly proportional to the square root of the cathodic current density. This relation has been confirmed experimentally in later studies [10-13]. Hence, some investigators believe that it is highly characteristic of hydrogen permeation for electrolytically produced hydrogen. An analogous rule holds for the permeation rate of gaseous hydrogen as a function of its pressure [14-17], Because the last process is accompanied by dissociation of hydrogen molecules into atoms, some authors are therefore of the opinion that chemical desorption is the rate-de-termining step of the hydrogen evolution reaction (mechanism 3). However, for iron the experimentally determined values for electrochemical parameters contradict this idea. For instance, the Tafel slope in mechanism 3 should be RT/2F ~ 0.030 V/decade, whereas the slope observed is RT/2F ~ 0.120 V/dec-ade (at 298 K).

Bockris et al. [11] have supposed that the hydrogen evolution reaction on iron proceeds along the lines of mechanism 5 (coupled discharge-chemical de-sorption), and that hydrogen enters the metal lattice from the adsorbed state

M H a d8 ^ M H a bs (8.19)

The authors [11] offered the following reaction sequence during the cathodic polarization of Armco iron membranes :

2 4 8 ELECTROCHEMICAL ASPECTS OF HYDROGEN ENTRY INTO IRON

H 30 + + Fe + é > - Î ! > F e H a ( J S- ^ i H 2 + Fe

4 -FeHabs (X = 0)

j diffusion (8.20) F e Ha bs (X = L)

where kl9 k2, and k3 are the rate constants for the respective steps and X is the distance through the membrane. The Tafel slope for Armco iron in 0.05 M H 2S 0 4 was found to be 0.115 V/decade (at 298 K), which coincides with the value expec-ted for mechanism 5 on the assumption of a low hydrogen coverage (0H < 0.1). The permeation rate-determining step was the diffusion of hydrogen in the membrane, and the entry process, being faster, was at equilibrium. Assuming in addition a constant diffusivity and a low concentration of hydrogen in the metal, the authors [9] arrived at the following relationships:

DF k* 0 non ip = 0 H (8.21)

DF ^ 3 _ f c j l / 2 / c- l / 2 (8 2 2)

L k_

DF L k_

k^kâU2 - exp [ - - ^ 2 ^ . ] (8.23)

where, besides the previously defined symbols, ic is the cathodic current density, aH the activity of hydrogen ions in the electrolyte, α the transfer coefficient of the discharge reaction, and η is the overpotential.

Thus the permeation rate should be a linear function of the square root of the cathodic current density (equation (8.22)). This relationship has been confirmed [11] for acid electrolytes with and without a KI addition (Fig. 8.2) and also for an alkaline electrolyte (Fig. 8.3). Assuming that α = 0.5, the following rela-tionship results from equation (8.23):

ên _ — 2.3034RT ^ —0.240 V/decade/„ (8.24) rflog/j

which shows that the permeation rate should increase by a factor of ten for every increment of 0.240 V in the hydrogen overvoltage. This result was ob-


V£,(A/m

2)2

Fig. 8.2. Hydrogen permeation rate vs. square root of the cathodic current density relationships; Armco iron, L = 0.77 mm, 0.05 M H 2S 0 4 + KI, 298 Κ (after [11]).

10 ±2 0

\/y7,(A/m

2)2

Fig. 8.3. Hydrogen permeation rate as a function of the square root of the cathodic current density; Armco iron, L = 0.77 mm, 0.1 MNaOH, 298 Κ (after [11]).

tained in 0.05 M H 2S 0 4 with additions of the following compounds: potassium iodide, benzonitrile, valeronitrile, naphthonitrile, and naphthalene [11]. In alka-line solutions the latter relation holds good only as far as a potential of about — 1000 m V N H E, whereas at a higher polarization the permeation rate increases slowly. This could be explained by a change in the mechanism of the hydrogen evolution reaction—from the coupled discharge and chemical desorption into thé slow discharge-fast electrochemical.


An alternative hypothesis put forward by the others workers is that hydrogen enters the metal as protons, i.e. without the intermediate chemisorption stage of the hydrogen atoms on the electrode surface [18, 19]

H+ + M + e -> MHahs (8.25)

This concept seems useful in accounting for generally greater permeation rates when hydrogen is generated in acid rather than in alkaline solutions. For it can be considered that in the case of acid solutions hydrogen enters the metal as protons, whereas in alkaline solutions hydrogen enters mainly through the ad-sorbed hydrogen atoms.

The model of a direct entry of protons into the metal lattice was considered by Bockris et al. [11]. Their analysis leads to the following relationships:

άη RT dlogip

(8.26)

and / 2 (8.27)

The last criterion, showing that the steady state permeation rate should be pro-portional to the square of the cathodic current density, has not been confirmed in permeation experiments. Therefore, Bockris et ai. [11] concluded that the direct entry mechanism is inadmissible.

Others argue, however, that relationship (8.27) is not a necessary condition in validating the concept of the entry of protons. In particular, Evseev [20] ob-served a linear relation between permeation rate and current density over a wide

400 800 Cathodic current density, A /m

2 1200

Fig. 8.4. Hydrogen diffusion rate as a function of the cathodic current density; Armco iron, 0.05 M H 2S 0 4, L = 0.17 mm, 298 Κ (after [20]).

HYDROGEN ENTRY A N D HYDROGEN EVOLUTION 251

range of its values (Fig. 8.4). This fact, in the author's opinion, may be explained by assumption that the discharge of hydrogen ions and the entry of hydrogen occur at the same sites on the metal surface. These sites are unsaturated and hydrogen enters the metal in accordance with equation (8.25).

In the light of other studies, the linear relationship of ip vs. l^2 appears to

V^(A/m

2)2

Fig. 8.5. Effect of cathodic current density on hydrogen permeation rate through steel mem-branes rotated in 0.05 M H 2S 0 4 + 5 χ 10"

5 M A s 20 3, 298 Κ. Membranes (L = 0.6 mm,

0 = 11.3 mm) were machined out from a steel rod perpendicular to its axis. A constant rota-tion speed of 9 rev. per second (after [13]).

151 -τ 1 1 1 1 r~

Fig. 8.6. Effect of cathodic current density on hydrogen permeation rate through steel mem-branes rotated in 0.1 MNaOH + 5 x 10~

5 M A s 20 3. Other conditions as in Fig. 8.5 (after [13]).


be true only for the medium range of current density. As demonstrated by Ra-czynski [10], a complete graph of the permeation rate vs. current density rela-tion encompasses four ranges. For very small ic values the permeation rate is proportional to /c. For higher /c values zp is proportional to i

lJ

2\ with ic increasing

further, the permeation rate rises increasingly slowly, and finally to become independent of current density. This complex character of the permeation rate vs. cathodic current density relation was supported by studies in acidic and alka-line solutions with or without promoters of hydrogen entry [13, 21]. By way of illustration, the effect of the cathodic current density on the permeation rate of the hydrogen evolved from electrolytes containing A s 20 3 is given in Figs. 8.5 and 8.6. In these measurements, the membranes were machined out from the steel rod in such a way that their active surfaces were perpendicular to the rod axis. This has been found to eliminate irreversibility of the behaviour of steel membranes when charged electrolytically with hydrogen [6].

Further information on the mechanism of hydrogen entry was obtained from study of the effect of pH of the electrolyte on this process [22]. The rotating disk was used, which enabled the study to be made over a wide pH range without needing to change the composition of the electrolyte, and for the pH < 7 it allowed the exact determination of the pH at the electrode surface (pHs). The hydrogen permeation rate was found to increase with the concentration of hydro-

10

1

Ε <

10

L

10"

—ι—ι—I t ι ι ι ι ν 1 ι I ι I · ι j

/ 0 . 1 1 3

- * \ o . Q 5 M H 2S 0 4 + N Q O H + 5 * 1 0 ~

5M A S2O :

0.05MH2S04 + No0H

ι ι ι ι 111 I I I I I I 1

Ι Ο " 10

-4 10"' 1CT

The H30+ ions concentration at the electrode surface, mol/l

Fig. 8.7. Effect of H 30

+ concentration at metal surface on hydrogen permeation rate through

steel membranes in various solutions with and without A s 20 3; h = 22.5 A/m

2, other condi-

tions as in Fig. 8.5 (after [22]).


nium ions at the surface (cs), both with and without the addition of a promoter (As203) (Fig. 8.7). For both cases in the pH range examined (from 1 to 3 pH units) the log/p vs. logcs relationship was linear. It could thus be expressed by the fol-lowing equation:

ip = KcsT (8.28)

where k and m are constants. It would seem that this equation remains valid over the whole pH range, as the values of the permeation rate found by extrapo-lation of the lines in Fig. 8.7 to high pHs values are close to the permeation rate observed experimentally for high pH.

At this point attention should be called to one more significant fact. In acid solutions (or for ic < iL, where iL is the limiting current density of the discharge of H 30

+) H 30

+ ions are only species being discharged and these ions exclus-

ively produce the hydrogen that enters the metal. Here both the proton mechan-ism (8.25) and the adsorbed hydrogen atom mechanism (8.20) could be possible. For an alkaline electrolyte (/c > /L) the cathodic current can be divided into three parts :

1. current corresponding to the discharge of H 30+ ions transferred from

the bulk of the electrolyte to the cathode surface; 2. current corresponding to the reduction of H 30

+ produced by the disso-

ciation of water molecules; 3. current corresponding to the direct reduction of water molecules. Therefore, the hydrogen entering into the metal could originate both from

H 30+ ions and water molecules. However, according to Delahay [23], the maxi-

mum current density corresponding to the first two components, in solutions of a high pH, is very small. For instance, in 0.1 M NaOH it is about 2.6 χ x l 0 ~

5 A/m

2. Since, at a high pH, the observed permeation rate was of the

order of 1 x 1 0 "2 A/m

2 in the promoter-free electrolyte, and of the order of

1 χ 10~1 A/m

2 in the promoter-containing electrolyte, it may be supposed that

the entering hydrogen comes exclusively from the discharging water molecules. If, therefore, the hydrogen entry is to proceed via the same mechanism, inde-pendent of pH, then the above fact require rejection of the mechanism of proton penetration through the electrolyte/metal interface, and acceptance of the mech-anism of hydrogen atom entry.

The question remains as to whether this model is sound in its original form, that is, whether hydrogen goes through an adsorbed intermediate state before the metal lattice. The linear relation of permeation rate vs. square root of the cathodic current density holds true for electrolytes with or without promoters


of hydrogen entry [11, 13, 21, 24]. Another point worth noting can be seen in Fig. 8.7. The slopes of both of the log/p = /(logcc) lines are virtually identical, and the promoter effect manifest itself only in the value of the proportionality constant k in equation (8.28). These facts suggest that the presence of a promoter does not qualitatively affect the character of the entry process, it just increases its rate. Before this observation is used in the explanation of the mechanism of hydrogen entry some fundamental aspects related to promoters will be briefly treated below.

8.4 Role of promoters of hydrogen entry

As mentioned earlier, certain substances, referred to as promoters of hydrogen entry, when present in the electrolyte even in small quantities, may significantly accelerate the process of the cathodic hydrogen entry into a metal. A similar phenomenon is observed, when the entry process occurs under silent discharge conditions in rarefied hydrogen [25].

The most effective promoters are compounds of certain elements of the Group VA (Ρ, As, Sb) and VIA (S, Se, Te) of the Periodic Table. There is also a similar, yet much weaker effect with Hg, Pb, Sn, CN~, halide anions, naphthalene, and some other organic compounds. In the discussion that follows, attention will be focused on the promoters of the VA and VIA Groups of the Periodic Table.

8.4.1 Active chemical form of promoters

Determination of the actual chemical form of a promoter in which it manifests its activity was one of the major tasks in the study of hydrogen entry into metals. An essential contribution to the knowledge of this problem was the work of Smialowski and his associates. Studying the expansion of iron wires and coils cathodically charged in sulphuric acid solutions containing additions of various compounds of Group VA and VIA elements, Smialowski and Szklarska-Smia-lowska [26, 27] have found a distinct correlation between the magnitude of the expansion due to the absorption of hydrogen, and the bond strength of the hydri-de of the given promoting element. The expansion increases linearly with the bond strength of hydride, within one periodic group. It has, therefore, been con-cluded that the hydrides of the promoters are decisive for their catalytic action in the overall process of hydrogen entry into metals.

Angerstein-Kozlowska [28] reported that addition of A s 20 3 to 0.2 M H 2S 0 4 accelerates hydrogen entry only at potentials that are more negative than the

ROLE OF PROMOTERS OF HYDROGEN ENTRY 255

potential at which arsenic begins to reduce to arsine. At more positive poten-tials, the surface of an iron cathode becomes covered by a film of arsenic de-posits which protects the metal against corrosion, as well as making hydrogen entry more difficult.

The above observations have been supported by measurements of the per-meation rate of hydrogen through steel membranes cathodically polarized in electrolytes containing suitable compounds of Group VA and VIA elements [21]. Figure 8.8 shows an example of the effect of the cathode potential on the perme-

E < A

cu ο

c 3 ο

ο Οι ε

I I I I — Τ -

ο0.5Μ H2S04 + 5x10"

5MAs203

ρΗ = 0.35

I ρ ι I

. Δ 0.05Μ H2S04 + 5 χ 10"

5Μ As203

ρΗ = 1.2 Λ / /

• 0.05MNa2S04+H2S04+ /

" ρΗ=2.6

5«

1 0"

5 M A SA /

-0.3 -0Λ -0.5 Potential, VN HE

-0.6

Fig. 8.8. Effect of potential on hydrogen permeation rate through steel membranes in solutions of various pH with A s 20 3; other conditions as in Fig. 8.5.

ation rate of hydrogen evolved from electrolytes of various pH and containing A s 20 3. Initially, the permeation rate is low and increases slowly with increasing cathodic polarization. A prompt acceleration of hydrogen permeation occurs when the critical potentials (denoted as E'cy E'c' and E'0") are reached. A similar change of the permeation rate with potential was observed in electrolytes with additions of N a H 2P 0 2, S b 20 3 and T e 0 2.

In the presence of sodium sulphide, which in acid solutions hydrolyzes with evolution of hydrogen sulphide, the permeation rate starts increasing at poten-tials as low as the corrosion potential. A similar situation was noted in the pres-ence of Se02 despite the fact that substantially more negative values of the


equilibrium potential EQ result from thermodynamical data [29] for the Se/H2Se system than for the S/H2S system (Table 8.1).

On the other hand, no measurable effect of a B i 20 3 addition on the permeation rate was found. The reason for this is most probably the very negative equilib-rium potential for the Bi/BiH3 system (Table 8.1), which may not have been reached under the applied conditions.

TABLE 8.1

COMPARISON OF THE EQUILIBRIUM POTENTIALS FOR HYDRIDE FORMATION E0 (after [29, 30]) WITH THE VALUES OF CRITICAL POTENTIAL Ec FOUND EXPERIMENTALLY (after [22]); pH = 2.6

Element Equilibrium potential, Reference Critical potential,

• Eo, V N H E

Reference Ec, V N H E

Ρ -0 .217-0 .0197 log /7pH3 [29] - 0 . 4 2 As -0 .762-0 .0197 l o g i ?A SH3 [29] -0 .51 As -0 .379-0 .0197 log/?A s„3 [30] -0 .51 Sb -0 .664-0 .0197 log />S BH3 [29] -0 .835 Bi -0 .954-0 .0197 log />BiH3 [29] — S + 0.017-0.0295 log/>H2s [29] > - 0 . 3 Se -0 .523-0 .0295 log /?H2se [29] > - 0 . 2 5 Te -0 .871-0 .0295 log /?l l 2Te [29] - 0 . 3 9

As seen in Table 8.1, values of the critical potential Ec at which a clear increase in the permeation rate was observed, due to the initiation of the formation of the hydride of the given element, disagree with the values of the corresponding equilibrium potentials E0. For Ρ and Sb more negative, whereas for Se and Sb more positive Ec values were observed, compared with the E0 values. For As either condition is possible, depending on which value of the equilibrium poten-tial [29, 30] is taken into consideration. Even though the value of the partial pressure of the given hydride at which the latter acts as a promoter is unknown, it would, however, be expected that at least a qualitative relation exists between the calculated and the experimentally found values of the respective potentials. A question arises here as to why, for instance, in the electrolyte containing A s 20 3, at potentials just slightly more positive than the critical potential, generally higher permeation rates are observed than in an electrolyte without added pro-moter. To clarify the situation, experiments were performed to determine the potentials at which the volatile hydride of the element used as promoter begins to appear in the gaseous electrolysis products. The results of these experiments


have shown [22] that the hydrides begin to evolve vigorously at potentials of the steel cathode coinciding with the EQ values, yet hydride traces are formed at even more positive potentials. Arsine traces were identified already at the corro-sion potential of steel in a sulphuric acid solution. These observations lend validity to the view that the hydrides of Group VA and VIA elements are effec-tive promoters.

A contrary view was expressed by McCright and Staehle [31] with regard to arsenic, who claim elemental As, rather than its hydride to be the active form. These workers used as a basis the potential calculated from thermodynamical data and the fact that the relative permeation efficiency, defined as the ratio of permeation rate to cathodic current density (/p//c), was relatively larger in the range of potentials of the thermodynamic stability of elemental arsenic. This explanation, however, fails to account for the effect of corrosion processes at low polarization and also the formation of arsine in minor quantities in the potential range where the elemental form prevails.

The complex relationship between the permeation rate and the cathodic current density (Figs. 8.5 and 8.6) holds true, if hydrogen evolves from acid solutions (ic < iL) or from alkaline solutions (ic > /L). The situation becomes even more involved, if the membrane is polarized under conditions which make it possible to change from the acid region to the alkaline region through the limiting current. This take place when the transport of H 30

+ ions to the cathode

becomes too slow to permit their discharge at a rate corresponding to the applied current density. At the limiting current density the activity of H 30

+ at the cath-

ode surface falls to a value equal to that characteristic of pure water (pH = 7). At higher current densities than iL the electrolyte at the cathode becomes strong alkaline and practically only electrolysis of water molecules occurs [32]. As seen in Fig. 8.9a, a marked peak appears on the curve of permeation rate vs. poten-tial. The position of the peak and the extent to which the permeation rate de-creases with continued polarization are closely related to whether the promoter present in the electrolyte is from Group VA or VIA. For example, for antimony the permeation rate does not decrease until the potential corresponding to the limiting current density is reached (Fig. 8.9b), after which it still remains at a considerable level. In the case of selenium, on the other hand, the maximum permeation rate occurs before the limiting current is achieved. Then, the per-meation rate diminishes dramatically to attain very low permeation rates that are close to those observed in the promoter-free electrolyte.

The above differences result from the fact that in the case of Group VA ele-ments only the effect of pH on the entry of hydrogen is manifested, whereas


10

3

Potential, V N HE Fig. 8.9. Permeation rate of hydrogen (a) and cathodic current density (b) vs. potential;

0.05 M N a 2S 0 4 + H 2S 0 4 (pH = 2.6) with S b 20 3 or S e 0 2; other conditions as in Fig. 8.5.

for Group VIA elements and additional effect observed is a decreased concentra-tion (and hence efficiency) of the hydride resulting from its dissociation in aqueous solution. It is known that the hydrides of Group VIA elements, as distinguished from those of Group VA, are stable only in acid solutions. At higher pH the dissociated anionic forms increasingly prevail over the molecular form

H 2X -> HX- + H+ (8.29)

HX- -> Χ2" + H

+ (8.30)

where X = S, Se, Te. Since the acidity of the hydrides of Group VIA elements changes with their atomic number, the pH values below which the molecular form prevails (and conversely above which HX" or X

2~ predominate) are differ-

ent for the various elements (Table 8.2). Using a rotating disc electrode, the pH value at the electrode surface (pHs,e x pt . )5 corresponding to the maximum value of the permeation rate, can be determined. For Group VA elements p H s, e x p t. = 7; for Group VIA elements the relevant data are provided in Table 8.2. A satis-


TABLE 8.2

COMPARISON OF THE LIMITING pH VALUES CORRESPOND-ING TO THE REGIONS OF RELATIVE PREVALENCE OF H2X, HX- OR Χ

2" (after [29]) WITH THE p H s, e x p l. VALUES

FOUND EXPERIMENTALLY WHICH CORRESPOND TO MAXIMUM PERMEATION RATES (after [22])

Element P H H 2x / H x - ρΗΗχ-/χ2- ρΗ8, expt.

S 7.00 13.90 7.0 Se 3.74 14.01 3.3 Te 2.64 11.00 2.9

factory agreement can be seen for the values of p Hs > exP t. and ρ Η Η 2χ / Ηχ - . This confirm the essential role of the hydrides in accelerating the entry of hydrogen. On the other hand, in contradiction to the viewpoint of some authors [33, 34], who attributed promoting effect of H2S to the action of HS~ or S

2~, these ions

(and similarly HX~ and X2~) are inactive as promoters of hydrogen entry into

metals.

Figure 8.10 shows schematically the typical curves of the permeation rate vs.

Ε

pHs < 7 pHs = 7 . Ί

pHs>7

/ I | S e , T e \ \ \

/ / ' \ V _

] P,As,Sb

S , Se /

ι ι

/ P , A s , S b ! ! / Te j j

! ι I ι t ! ι

^ ^ ^ S e j e

^ c o r ^ H 2 X / H X "

Potential —

Fig. 8.10. A schematic representation of the shape of hydrogen permeation rate vs. potential curves for various promoters (after [22]).


cathode potential for the promoters studied. The shape of the curves result from variations in permeation rate which occur as a result of changes in cathodic current density, pHs, or promoter activity. The potentials at which a given pro-moter has a clear accelerating effect on the entry of hydrogen correspond to the following ranges:

Ε < EC VA Group (8.31)

£H2X/HX- < Ε <EG VIA Group (8.32)

where EC denotes the potentials at which the hydrides of Group VA and VIA begin forming at a considerable rate, whereas £H2X/HX- are the potentials at which dissociation of Group VIA elements hydrides is pronounced.

Occasionally, the complex processes occurring at a polarized iron (steel) cathode in solutions containing additions of As, Sb, Se, and Te compounds, become even more involved as a result of the formation of deposits of these elements on the cathode surface [13, 21, 22, 28, 35, 36]. Layers of these deposits, make hydrogen entry more difficult. This effect depends on the coverage of the electrode surface by the deposit, its thickness, porosity and adhesion to the metal substratum. Moreover, the hindering effect of the deposit depends on the region in which the hydrogen evolution reaction takes place [21]. If it occurs in the kinetic region, that is, if it is controlled by the discharge rate of H 30

+ ions (or

H 20 molecules), then the action of the deposit resolves itself into blocking a fraction of the metal surface. Part of the H 30

+ ion (or H 20 molecule) flux

becomes discharged on the deposit surface, and the hydrogen atoms produced are more likely to recombine and then escape, rather than to enter into the metal. In the diffusion region, i.e. where the rate of hydrogen evolution is controlled by the rate of transport of H 30

+ ions from the bulk of electrolyte to the electrode

surface, hindrance of this transport occurs in addition to the blocking of the surface. Under galvanostatic conditions (Ic = const), the gradient of H 30

+ ion

concentration in the layer encompassing the porous deposit and the diffusion layer on the electrolyte side increases. In consequence, the pH at the metal surface increases too, which also causes a lower hydrogen permeation.

To sum up the above discussion of the active chemical form of promoters of hydrogen entry, it can be stated that P, As, Sb, S, Se and Te accelerate the entry of hydrogen into steel under conditions that ensure the presence of the hydrides of these elements in the system. Their anionic or elemental forms do not exhibit the promoting effect. If the elemental form produces a deposit on the metal surface, then the rate of hydrogen entry decreases.


8.4.2 Mechanism of the promoter action

There are many concepts related to the mechanism of promoter activity. Here only a few will be considered.

Some authors attributed the accelerating effect of promoters to "poisoning" of the recombination reaction of the adsorbed hydrogen atoms into molecules, which was thought to increase the concentration of atomic hydrogen on the cath-ode surface [24, 37 , 38] .

Others thought that promoter action is connected with a lowering of the M - H a ds bond energy, owing to which the activation energy is decreased for the process of passage of hydrogen atoms from the surface into the metal phase [ 1 1 , 3 9 , 4 0 ] .

A concept which has also been put forward is that promoter action is due to the interaction of the promoter molecules with the electrons of the surface metal atoms [41]. Chemisorption of the promoter molecules on the cathode surface causes the deformation of the orbitals in which the electrons move in the surface atoms of metal. If, under no-interaction conditions, the electric fields of atoms in the crystal lattice of the metal are of very nearly spherical symmetry, strong chemisorption of a promoter may cause these fields to deform towards an elliptical symmetry. As a result, the forces acting between the neigh-bouring atoms of the metal may become weaker which, in consequence, lowers the energy barrier hindering hydrogen transfer into the bulk of the metal phase.

The hydrides of elements belonging to the VA and VIA Groups are strongly chemisorbed by clean metal surfaces [42]. This suggests that the catalytic action of these hydrides is connected with the so-called dissociative chemisorption of their molecules [43]

where M is a surface metal atom and X is the promoting element of valence n. It is assumed [42] that one radical of the Η ^ . ^ Χ type covers two sites in chemisorption, while the detached hydrogen atom occupies another single site on the metal surface. When such a film is attacked by more hydride molecules, these attempt to dislodge the hydrogen atoms from their positions, and this might facilitate their entry into the metal phase.

The force constant of water molecules is greater than that of both H 2 and H„X molecules [44], but H 2 molecules are bound more strongly than those of HnX. It is thought, therefore, that in comparison with water, the HnX molecules

H 20 + 2 M = M O H + M H

H „ X + 2 M = Μ Η Ο , - D X + M H

(8 .33)

(8 .34)


chemisorbed dissociatively have a more pronounced ability to transmit hydrogen atoms by producing larger amount of sites for that transmission. The possibility that the H 20 molecules themselves can promote hydrogen entry was suggested by Newman and Shreir [40].

It can be assumed that the rate of hydrogen entry, expressed by the permeation rate through the membrane ip is proportional to the fugacity of the hydrogen at the cathode / and to the number of sites Ν through which hydrogen atoms are transmitted into the metal phase

ip~(f,N) (8.35)

It may subsequently be supposed that / depends on the cathodic current density, but is independent of pH and the presence of a promoter. It follows that for a given metal, the number of active sites JV depends upon both pH and the promoter. Therefore, the relationship between permeation rate and activity is complex. This might be one of the reasons for the observed change from the linear dependence of ip on ic at low cathodic current densities, into a rela-tionship between ip and iG

12 at higher ic values and, eventually, into the inde-

pendence of ip on ic at high degrees of cathodic polarization. A decreasing permeation rate with increasing pH is most probably due to the

adsorption of hydroxyl ions. As a result, the number of the sites on the elec-trode surface on which the dissociative chemisorption (of both the hydride and H 20 ) can take place is decreased. The inhibitive effect of hydroxyl ions can, however, be suppressed by raising the promoter (hydride) concentration. This finding can be used to explain the relatively high permeation rates through steel membranes which have been polarized in alkaline solutions containing pro-moter in large quantities [13, 45]. This is not, however, true of the group VIA hydrides which are inactive in alkaline solutions.

Although the above model of hydrogen entry into a metal, together with the effect of the hydrides of Group VA and VIA elements on this process, ac-counts for a number of experimental facts, it requires verification and further development. In the first place, the assumption of dissociative chemisorption on the metal surface that is in the contact with the aqueous electrolyte needs to be verified. Previous studies on the dissociative chemisorption of various gas-eous molecules were carried out under conditions of high vacuum and by physi-cal methods which provide no possibility of studying the process of hydrogen entry into metals. Furthermore, it should be emphasized that the model ad-vanced here is applicable only to pure electrolytes or those containing the hydri-des of the Group VA and VIA elements. As would be expected, in the experi-

INHIBITORS OF HYDROGEN ENTRY 263

ments whose results were used in deriving this model, the effect of hydrides was predominant. Generally, however, the role of other factors cannot be en-tirely disregarded. For it can be considered that every change occurring at the electrolyte/metal interface as a result of the presence of molecules of any species may affect the energy barrier (by lowering or raising it) that hinders the passage of hydrogen into the bulk of the metal.

8.5 Inhibitors of hydrogen entry

As distinct from promoters, some aubstances added to the solution hinder the entry of hydrogen into metals. These are primarily polar organic compounds, whose molecules contain sulphur, nitrogen or oxygen atoms.

For instance, Rozenfeld et al [46], in their study of hydrogen charging of steel polarized in sulphuric acid, noted an inhibitive effect of sulphoxides, hyd-roxy amino acids, and organosilicon compounds containing chlorine atoms. One of the best known and effective inhibitors of hydrogen entry is dibenzyl sulphoxide or, more precisely, the product of its cathodic reduction—dibenzyl sulphide [47-51]. Fairly extensive investigations into the effect of a wide var-iety of organic compounds on the hydrogen charging of steel were carried out by Beloglazov [52]. The studies included, among other substances, aliphatic alcohols, aldehydes, ketones, acids and amines; aromatic alcohols (phenols), aldehydes and amines; amino acids and proteins. The inhibition activity of the compounds examined was found to depend not only on the kind of functional group but also on the length of the aliphatic hydrocarbons chain.

A number of the inhibitors of acid corrosion of metals (used in pickling solutions and in electroplating bath for deposition of protective coatings), as a rule, also suppress the entry of hydrogen. This would be expected, since the overall amount of hydrogen evolved during corrosion is thereby decreased. Exceptions to this rule are, however, known. This is particularly true of some sulphur-containing inhibitors which, apart from decreasing the rate of corrosion (dissolution) of metal, can, at the same time, increase the amount of hydrogen that enters the metal. This is connected with the inhibitor decomposition pro-ducing hydrogen sulphide, which is known to strongly promote hydrogen entry from acid solutions. Among these dangerous inhibitors are, for instance, thiourea [47, 53], tiophene and thiazole [53]. The higher derivatives of thiourea and other sulphur-containing inhibitors, which do not undergo reduction to hydrogen sulphide on the cathode, are free from this shortcoming.

The mechanism of the inhibitor action is not sufficiently understood. This


is beacause the behaviour of the various organic compounds, when brought in contact with the metal, may be very different, depending on the structure and polarity of their molecules, the potential of the metal surface, and other factors. It is generally believed, however, that inhibition of the entry of hydrogen is the result of a strong chemisorption of the inhibitor molecules on the metal surface. The inhibitor adsorbed there may affect hydrogen entry into the metal phase through one of the following ways :

a. The inhibitor forms a compact layer on the metal surface, which makes the transport of hydronium ions or water molecules to the metal surface more diffi-cult. As a result, the discharge process does not occur directly at the bare metal surface, but mainly on the outer surface of the inhibitor layer. The hydrogen atoms produced, which have no direct contact with the solid metal, are thus incapable of going into its bulk.

b. The inhibitor molecules partially block the metal surface. The higher the degree of coverage of the metal surface by the adsorbed inhibitor molecules, the lower is the hydrogen charging of the metal.

c. The inhibitor affects the mechanism of the hydrogen evolution reaction, for example, the recombination velocity of adsorbed hydrogen atoms. The coverage of the metal surface by hydrogen atoms may decrease and, in conse-quence, the hydrogen absorption reaction might decrease too.

d. The inhibitor may affect the strength of the bonding of the adsorbed hydro-gen atoms to the metal surface atoms.

A simultaneous occurrence of all the above mechanisms of inhibitor action is likely.

8.6 Prevention of hydrogen charging

In order to prevent the hydrogen charging of steel, in the first place, a nearly com-plete elimination of those substances, which are capable of stimulating the hydrogen entry process from the environment surrounding this metal is highly desirable. Of these substances particular mention should be made of the hydrides of Group VA and VIA elements. The presence of substances which during the cathodic process may undergo reduction to any of these hydrides is obviously also un-desirable. Among the number of such substances, mention can here be made of sulphites, selenites, arsenites, hypophosphites, as well as certain organic compounds of sulphur, such as thiourea. For the same reason, efforts should be made to use grades of steel which are the least contaminated with sulphur and phosphorus. On the other hand, compounds of Group VA and VIA elements

REFERENCES 265

which do not undergo reduction on a steel cathode, such as sulphates or ortho-phosphates, are not harmful.

Another method of preventing the hydrogen charging of steel is to raise the pH of the medium. This is because hydrogen enters the metal more easily at higher hydrogen ions concentration in solution. Furthermore, if compounds of Group VIA elements are present in the solution, an increased pH value also eliminates their effect as entry promoters.

In many situations favourable results can be achieved by using inhibitors of the hydrogen entry. Such inhibitors are widely accepted in commercial prac-tice, in acid pickling or the electrolytic coating of steel. Another branch of tech-nology where serious problems of hydrogen embrittlement can be reduced by the application of inhibitors is the petroleum and gas industry.

8.7 References

[I] Ateya, B. G., Pickering, H. W., / . Electrochem. Soc, 122, 1018 (1975). [2] Seys, Α. Α., Brabers, M. J., van Haute, Α. Α., Corrosion, 30, 47 (1974). [3] Shively, J. H , Heheman, R. F M Troiano, Α., Corrosion, 23, 215 (1967). [4] Devanathan, Μ. Α. V., Stachurski, Z., Proc. Roy. Soc, A270, 91 (1962). [5] Devanathan, Μ. Α. V., Stachurski, Z., / . Electrochem. Soc, 111, 619 (1964). [6] Smialowski, M., "Hydrogen Blistering and Surface Microcracks", in: Proc. Intern. Confer.

on Stress Corrosion Cracking and Hydrogen Embrittlement of Iron-Base Alloys, Unieux-Firminy 1973.

[7] Bockris, J. O' M., Reddy, Α. Κ. N., Modern Electrochemistry, Plenum Press, New York 1970.

[8] Bockris, J. O'M., Subramanyan, P. K., Electrochem. Acta, 16, 2169 (1971). [9] Bodenstein, M., Z. Electrochem., 28, 517 (1922).

[10] Raczynski, W., Arch. Hutnictwa, 3, 59 (1958). [II] Bockris, J. O'M., McBreen, J., Nanis, L., J. Electrochem. Soc, 112, 1025 (1965). [12] Raczynski, W., Symposium on Hydrogen Embrittlement and Corrosion of Metals, Jablonna

1977, p. 67. [13] Zakroczymski, T., Szklarska-Smialowska, Z., £mialowski, M., Werkstoffe u. Korrosion,

26, 617 (1975). [14] Richardson, O. W., Phil. Mag., 7, 266 (1904). [15] Kumnick, A. J., Johnson, H. H , Met. Trans., 6A, 1087 (1975). [16] Raczynski, W., Phys. Status Solidi A, 48, K27 (1978). [17] Raczynski, W., / . Less-Comm. Met., 101, 383 (1984). [18] Bagotskaya, I. Α., Zhurn. Fiz. Khim., 32, 2667 (1962).

[19] Frumkin, A. N., Advances in Electrochemistry and Electrochemical Engineering, Vol. 3, Ed. P. Delahay, Interscience Publishers Inc., New York 1963.

[20] Evseev, A. F., Elektrokhimiya, 6, 841 (1970). [21] Zakroczymski, T., Ph.D. Thesis, Institute of Physical Chemistry, Polish Academy of

Sciences, Warszawa 1975.


[22] Zakroczymski, T., Szklarska-Smialowska, Ζ., Smialowski, M., Werkstoffe u. Korrosion,

27, 625 (1976). [23] Delahay, P., / . Amer. Chem. Soc, 74, 3497 (1952). [24] Radhakrishnan, T. P., Shreir, L. L., Electrochim. Acta, 11, 1007 (1966). [25] Palczewska, W., Ratajczykowa, I., Bull. Acad. Polon. Sci., Sér. Sci. Chim., 14, 673 (1966). [26] Smialowski, M., Szklarska-Smialowska, Z., Bull. Acad. Polon. Sci. CI. Ill, 1, 327 (1953);

2, 73 (1954). [27] Smialowski, M., Szklarska-Smialowska, Z., Rocz. Chem., 29, 85 (1955). [28] Angerstein-Kozlowska, H., Bull. Acad. Polon. Sci. Cl. Ill, 6, 739 (1958); 7, 881 (1959). [29] Pourbaix, M., Atlas of Electrochemical Equilibria in Aqueous Solutions, Pergamon Press

and Cebelcor, Brussels 1966. [30] Davis, J. H., J. Electrochem. Soc, 124, 722 (1977). [31] McCright, R. D. , Staehle, R. W., / . Electrochem. Soc, 121, 609 (1973). [32] Kortum, G., Treatise on Electrochemistry, 2nd ed., Elsevier, 1965, p. 492. [33] Le Boucher, B., Rev. de VInstitut Français de Pétrole, 18, No. 4 (1963). [34] Shparber, I. S., Shreider, Α. V., Zashchita Met., 6, 343 (1970). [35] Beloglazov, S. M., Uchen. Zap. Perm. Univ., 19, 23 (1975).

[36] Lunarska, E., Szklarska-Smialowska, Z., Smialowski, M., Werkstoffe u. Korrosion, 26,

624 (1975). [37] Aten, A. H. M., Zieren, F. M., Rec Trav. Chim., 49, 641 (1930). [38] Beck, W., Glass, A. L„ Taylor, E., / . Electrochem. Soc, 112, 53 (1965). [39] Kovba, L. D. , Bagotskaya, I. Α., Zhurn. Fiz. Khim., 37, 161 (1963). [40] Newman, J. F., Shreir, L. L., Corr. Sci., 9, 631 (1969). [41] Smialowski, M., Szklarska-Smialowska, Z., The 17th Meeting of CITCE, Tokyo 1966. [42] Hayward, D . O., Trapnell, B. M. W., Chemisorption, Butherworths, 1964, p. 235. [43] Inai, H., Kemball, C , Proc. Roy. Soc, A302, 399 (1968). [44] Siebert, Η., Z. anorg. allg. chem., 24, 274 (1953). [45] Angerstein-Kozlowska, H., Bull. Acad. Polon. Sci., Sér. Sci. Chim., 8, 49 (1960). [46] Rozenfeld, I. L., Kramarenko, D . M., Lantseva, Ε. N., Zashchita Met., 3, 172 (1967). [47] Smialowski, M., 2nd European Symposium on Corrosion Inhibitors, Ferrara 1965, Univer-

sity of Ferrara, 1966, p. 203. [48] Ertel, H., Horner, L., 2nd European Symposium on Corrosion Inhibitors, Ferrara 1965,

University of Ferrara, 1966, p. 71. [49] Lorentz, J. W., Fischer, H., 2nd European Symposium on Corrosion Inhibitors, Ferrara 1965,

University of Ferrara, 1966, p. 81. [50] Schwabe, K., Leonard, W., Chem. Ing. Techn., 38, 59 (1966). [51] Trabanelli, G., Zucchi, F., Gullini, G., Carassiti, V., Brit. Corr. J., 4, 212, 267 (1969). [52] Beloglazov, S. M., The Charging of Steel with Hydrogen during Electrochemical Processes

(in Russian), Izd. Leningradskogo Universiteta, Leningrad 1975. [53] Hudson, R. N., Warning, G. J., Metal Finishing, 64, 58 (1966).

9. Dissolution and Diffusion of Hydrogen in Iron and Steel

9.1 Introduction

The interaction of hydrogen with metals has for long aroused considerable interest in both academics and industrialists. This is evident from the abundance of literature on the subject and the number of international symposia. In view of the importance of iron and its alloys in modern technology the iron-hydrogen system is given particular attention. The major reason for this is the detrimental effect of hydrogen on the useful properties of steel, which manifests itself in the lowering of the plastic and strength properties of steel at room temperature. The extent of this adverse effect of hydrogen on iron alloys generally depends on the content of this element in the metal and on its distribution, as well as its form, in the metallic phase. In order to assess the degree of harmfulness of the hydrogen, information is required on the concentration and quantity of the hydrogen which is dissolved, segregated on lattice defects, or held in a molecular form in micro-cracks and blisters. The results so far obtained are fairly modest. One reason for this is that reliable measurements of hydrogen solubility in iron can only be made at temperatures above approx. 573 K. This is due to a low solubility of this gas in iron at lower temperatures, its possible interaction with a wide variety of defects in the crystal structure of the metal, as well as to the effect of surface processes on the quantities measured. Some quantitative data on the various forms of hydrogen in the metal can be obtained by indirect measurements which, however, do not always provide unambiguous results.

The fact that hydrogen embrittlement of iron alloys is not observed at low temperatures and at high deformation rates suggests that the transport of hydro-gen to particular regions of the metal is of prime importance.

The diffusion of hydrogen in iron and its alloys has been studied over a wide temperature range, starting from 30 K. It is noteworthy, however, that while the experimental values of the diffusivity for hydrogen reported by numerous authors at high temperatures differ only slightly, at room temperatures these

268 DISSOLUTION AND DIFFUSION OF HYDROGEN IN IRON A N D STEEL

differences reach 2 to 3 orders of magnitude. The reasons for this are diverse, and it is not always easy to determine which is the major factor. Among the poss-ible reasons are: surface processes, interaction of hydrogen with linear defects, and the development of micro-cracks and blisters in metal as induced by hydro-gen.

The lack of agreement in the relevant literature as to the behaviour of hydro-gen in iron and its alloys is largely due to the fact that the majority of investiga-tions in this field have been carried out using insufficiently pure materials whose complex structure depended on the composition of the materials and the history of their mechanical and thermal treatment.

This chapter is concerned with some of the properties of the iron-hydrogen system at room temperatures. The topics covered are: the concentration of the dissolved hydrogen; the critical concentrations of hydrogen which bring about irreversible changes in the properties of metal; the interaction of hydrogen with lattice defects; and the diffusion of hydrogen in iron of various degrees of purity, or subjected to various thermal treatments or plastic deformation.

9.2 Experimental methods

The principle underlying the measurement of hydrogen solubility S in a metal consists of the determination of the quantity of the absorbed gas by a specimen of known weight, at a constant temperature and at a constant partial pressure of gaseous hydrogen pHi after establishing the equilibrium in the system [1]. The solubilities found at various temperatures and at constant pH2 can be used to evaluate the heat (enthalpy) of dissolution of hydrogen from the relation:

5 - S0exp(-Q/RT) (9.1)

where S0 is an approximately constant quantity. The effect of the pressure of molecular hydrogen on the solubility of this gas

in the metal is expressed by the square-root formula of Sieverts law:

S = Sx]/pu2 (9.2)

where St is the solubility of hydrogen at pHz = 0.1 MPa. The solubility of hydrogen in α-iron at room temperature at pHi = 0.1 MPa is of an order of 10~

3 cm

3 H2/100 g Fe, and its determination by volumetric method thus is not

feasible. Under conditions of electrolytic charging of iron with hydrogen, the concentra-

tion c of the dissolved (also called diffusible or lattice) hydrogen can be found

EXPERIMENTAL METHODS 269

from steady-state measurements of the rate of permeation Poo of hydrogen through membranes of a known thickness /, from the equation

where D is the diffusivity of hydrogen in the material examined. In the determi-nation of Poo if the source of the permeating hydrogen is molecular hydrogen from the gaseous phase, the concentration c in equation ( 9 . 3 ) can be replaced by the solubility S.

The total content of hydrogen Vc in the metal can be determined by various methods [ 1 ] . High-temperature vacuum extraction is frequently preferred.

For the determination of diffusivity D9 the following five methods are gen-erally used.

1. In the steady-state flow method, equation ( 9 . 3 ) is used, which allows the calculation of D from the experimental values of the steady-state rates of hydro-gen permeation Poo through a metal layer, provided that the solubility S or con-centration c of the gas in the material studied is known under the conditions of the experiment.

2 . Measurements of hydrogen absorption or desorption rates are used for the determination of D, particularly at high temperatures. Solutions of the diffusion equations under suitable initial and boundary conditions for cylindrical- [2 ] or spherical-shaped [ 3 ] specimens yield simple formulae for the determination of D values.

3 . The classical time lag method [4 ] consists of measuring the quantity of gas V which diffuses through a metal layer of a known thickness in time t and determining from this the time lag tL. The diffusivity of hydrogen is estimated from the formula

If the hydrogen permeation rate is measured as a function of time, i.e. Ρ = / ( / ) , the time lag i L then equals the time after which Ρ = 0 . 6 3 P o o [ 5 ] .

4 . Application of the electrochemical method for the measurement of Ρ [5] has allowed the development of a number of new techniques for the determination of diffusivity of hydrogen [ 6 ] . One of these is: the method of standardized curves of the rise (absorption) P a/ P o o = f(i) and decay (desorption) P d/ P o o = ftt) of the hydrogen permeation rate with time, which occur after the cathodic po-larization current is switched on and off respectively [ 7 ] . In this method, D is found from relation

Poo = Dc/l ( 9 . 3 )

D = l2/6th ( 9 . 4 )

τ = Dt/l2

( 9 . 5 )

2 7 0 DISSOLUTION A N D DIFFUSION OF HYDROGEN IN IRON A N D STEEL

using the r parameter values evaluated for a range of Pt/Poo = f(t) values. An advantage of the method is that D can be evaluated for various concentrations of the diffusing gas from a single experimental curve of the rise or decay of the hydrogen permeation rate with time [8, 9].

5. The internal friction method, which is based on the view that the diffu-sivity value for hydrogen at low temperatures can be found from a study of the internal friction of hydrogen-charged iron specimens as a function of tempera-ture [10]. More recent studies, however, have cast doubts on this approach and have shown that the Snoeck peak at 40 to 50 Κ is due to the interaction of hydro-gen with impurities, rather than to simple diffusion [11]. The dependence of diffusivity D and permeation rate Ρ of hydrogen on temperature Τ can be de-scribed in terms of equations [4] :

D = D0exp(-AO/RT) (9.6)

D = P0Gxp(-AJRT) (9.7)

where AD and Ap are the respective activation energies for the diffusion and per-

meation processes, respectively, and D0 and P0 are constants. From relations (9.3), (9.6), (9.1) and (9.7)

Ap = AD + Q (9.8)

Equation (9.8) is used for the determination of the activation energy of diffusion, AD9 if the values of Ap and Q are known.

9.3 Dissolution and occlusion of hydrogen in iron

α-Iron has a body-centred-cubic lattice with a space filling of 68%. Octahedral and tetrahedral interstices occur in the crystal lattice of iron. By analogy with palladium [12] and nickel [13] hydrides it is assumed that, in α-iron, hydrogen also occupies octahedral sites. Some authors claim, however, that hydrogen occupies mainly tetrahedral sites at moderate temperatures [14],

The electronic state of hydrogen dissolved in iron has not as yet been estab-lished. It is supposed that hydrogen in iron occurs in the form of electron-screened protons [15],

The solubility of hydrogen in iron and its alloys has been investigated over a wide range of temperatures and gaseous hydrogen pressures. The theory of the process of hydrogen dissolution in metals and the experimental results of studies are extensively treated in [1, 16-18]. The overall reaction of dissolution of hydrogen from a gaseous phase in iron and its alloys is endothermal. The values of the heat of dissolution Q determined at temperatures from 573 to 1173 Κ are

DISSOLUTION OF HYDROGEN IN IRON 271

in the range 22.0-30.8 kJ/mol [19-29], and the most probable value proposed by Oriani [30] is Q = 27.2 kJ/mol. More recent studies [14] on the dissolution of hydrogen in mono- and polycrystalline zone-refined (ZR) iron of variable purity have demonstrated that the solubility Sis slightly dependent on the chemi-cal composition of the metal at temperatures exceeding 573 K, whereas the effect of grain boundaries on the solubility above 673 Κ is imperceptible. The authors cited above observed deviations from linearity of the log S = f(l jT) relation and explained them in terms of the temperature-dependent degree of occupation by hydrogen of the tetra- and octahedral interstices in the metal.

In agreement with formula (9.2), the solubility of hydrogen in the metal is proportional to the square root of the pressure of the gas. At elevated tempera-tures formula (9.2) was found to describe well the experimental data for iron and steel over the pressure range 3 χ 10"

3 to 50 MPa [31, 32]. It follows from the

study of the hydrogen permeation rate through ZR iron membranes coated with a thin palladium layer that Sieverfs law is obeyed over the temperature range 273 to 333 Κ under gaseous hydrogen pressures of 0.001 to 0.1 MPa [33]. In

Temperature, Κ

(1 /Γ )χ10

5 , Κ"

1

Fig. 9.1. Logarithm of solubility S and concentration c of diffusible hydrogen vs. reciprocal absolute temperature: i—from formula (9.9) a t / = 0.1 MPa (after [30]), 1'—from formula (9.3) and from the rate of hydrogen permeation from gaseous phase (after [36]), 2, 3, ^/—dur-ing the cathodic polarization of ZR iron at / = 160 A / m

2 in solutions: 2—1 M KOH, 3—

1 M KOH + 2 χ 1 0 -

5 MHgO, 4—0.1 M H 2S Q 4 + 3 χ 10"

5 M A s * 03 (after [42]).


calculations of the solubility of hydrogen in iron at room temperature and at high pressures, the fugacity / of the gas should be used. According to Oriani [30] the eifect of temperature and fugacity on the solubility is described by the equation:

S = 47 / '1/2exp( -27200 / i?r) (9.9)

where S is expressed in cm3H2/100 g Fe. It will be seen that equation (9.9) com-

bines formulae (9.1) and (9.2). Based on this equation, the relation S = f(\/T) for / = 0.1 MPa is shown in Fig. 9.1, curve 7. The temperature range for which the S measurements were made is marked by a solid line. From extrapolation to lower temperatures (dashed portion of curve 1) it follows that at 293 Κ S = 6 .8x l0~

4 cm

3 H2/100 g Fe. This indicates that one H atom occurs per

approx. 3 χ 107 Fe atoms.

Studies on the solubility of hydrogen at temperatures below approx. 573 Κ provide unreliable results. The reasons for this are: the low solubility of hydro-gen in iron, the difficulties of eliminating the effect of surface processes, and the possible interaction of hydrogen with the defects of the host lattice of the metal. Hill and Johnson [34] carried out such a study over the temperature range 416 to 858 Κ at various hydrogen pressures pHi. At pHj = 10 MPa an abnormally high solubility of hydrogen was noted at a temperature below 663 K, which was accounted for by the formation of hydrogen traps in the metal. Plastic deforma-tion of a carbon-iron alloy (0.16% C) not only increases the equilibrium quan-tity of the absorbed gas but it also causes it to rise with decreasing temperature [35]. The main reason for this behaviour of the system is the presence of micro-pores and microcracks developed in the metal during cold plastic deformation.

In high-purity iron the concentration of diffusible hydrogen (curve Γ) cal-culated from the measured steady-state permeation rates of the gas through membranes from a gas phase of fugacity / = 0.1 MPa agrees with the solu-bility determined by extrapolation of the high-temperature data from equation (9.9) [36].

Under conditions of electrolytic charging, the total amount Vc of hydrogen taken up by commercial-grade iron or steel may attain tens or even hundreds of cm

3 H2/100 g Fe. An analysis of the data from a number of investigations

[1, 37-39] shows that the saturation of a metal with hydrogen increases with the lowering of the pH of the electrolyte, addition of hydrogen entry promoters, higher densities of the cathodic polarization current, higher degrees of cold plastic deformation of the metal, and lower thicknesses of specimens. It should be emphasized that only part of the hydrogen taken up by commercial-grade

PARTIAL MOLAR VOLUME OF HYDROGEN IN IRON 273

iron is in solid solution. A quantity of hydrogen equal to the difference V0-c is segregated on dislocations and is precipitated in the molecular form in micro-cracks and voids [40]. The concentration of the diffusible (dissolved) hydrogen can be estimated by using formula (9.3) from measurements of the steady state permeation rate P^ of the gas through membranes. In a thoroughly annealed high-purity iron, in which no micro-cracks or voids are formed on vigorous hydrogen charging, there is virtually only the dissolved form of hydrogen [41]. The temperature dependence of hydrogen concentration in ZR iron during cathodic polarization at a current density / = 160 A/m

2 in a pure alkaline

electrolyte, and in the same electrolyte with a promoter (HgO) added are shown by curves 2 and 5 in Fig. 9.1 [42]. From calculations using equation (9.9) in which the hydrogen fugacities were taken into account [43], it turns out that at 298 Κ the concentration of hydrogen in iron under the conditions relevant to curves 2 and 3 corresponds to equilibrium pressures of gaseous hydrogen /?H2 of approx. 4 and 110 MPa, respectively. In an acid electrolyte with added As2Os and at / = 160 A/m

2, a high concentration of diffusible hydrogen in ZR

iron is obtained, which corresponds to pHz > 103 MPa at 298 Κ (Fig. 9.1,

curve 4). The charging of commercial-grade iron and steel with hydrogen under such

conditions invariably results in the formation of micro-cracks and voids in the metal, accompanied by plastic deformation of the material. Increased tempera-ture raises the concentration of diffusible hydrogen, but it decreases the total quantity of the gas absorbed [40].

9.4 Partial molar volume of hydrogen in iron

In 1966 two papers appeared which reported equations for the determination, in metals, of the partial molar volume of hydrogen, VH [44, 45]. In one of these [44] a suggestion was made as to the determination of the VH value from measure-ments of the rise or fall of the rate of hydrogen permeation through cathodi-cally polarized membranes, at a constant overpotential, which were subjected to tensile or compressive stresses. For Armco iron and an A.I.S.I. 4340 steel the VH values at room temperature were 2.66 and 1.96 cm

3/mol Η [46]. The

method and the results of the study, however, have been questioned by other investigators [47, 48].

In the other paper mentioned [45], attention was called to the fact that the dissolution of hydrogen in a metal should cause a change in the lattice par-ameter and in the specimen size. Wagenblast and Wriedt [49] found iron wires to

274 DISSOLUTION A N D DIFFUSION OF HYDROGEN IN IRON A N D STEEL

elongate during hydrogen charging from the phase in a temperature range of 723 to 1073 K. Above 870 Κ the Vn value estimated on this basis was 2 cm

3,

while below that temperature it was 1.3 cm3/mol H. The author of ref. [50] ob-

served some reversible dilatation effects of ZR iron during electrolytic hydrogen charging and desorption of hydrogen from the metal. From relation [49] :

VH = 3VFea/ca (9.10)

where VFe is the volume of a mol of iron, and ε is the fractional elongation of specimens caused by hydrogen (of a concentration ca9 expressed in mol H/mol Fe) introduced into the metal. The present author has determined the V values at various hydrogen concentrations in iron. Results of the studies are presented in Table 9.1. As the data show, for higher current densities (i.e. for

TABLE 9.1 VALUES OF ca ε, AND KH FOR VARIOUS CATHODIC CURRENT DENSITIES Î; 0.5 M H 2S O 4 + 3 x l 0 ~

5 M A s 20 3, Γ = 293 Κ

ι, A /m

2 10 40 90 160 250

ca χ 10

5, mol H/mol Fe 2.27 4.80 7.35 9.9 12.6

ε χ Ι Ο

6 1.5 3.90 6.90 9.40 12.0

V», cm

3/molH 1.4 1.7 2.0 2.0 2.0

higher hydrogen concentrations) the partial molar volume of hydrogen in iron is constant and at 293 Κ it is 2 cm

3/mol H, whereas for lower values of ca, as

in [49] for lower temperatures, the experimental VH values are lower. The reasons for this have not yet been clarified.

It will be noted that the value of the hydrogen diffusivity D found and used in the evaluation of the hydrogen concentration in iron in [50] was 10""

4 cm

2/s,

whereas in later studies [36] a lower value of D was obtained. Accordingly, the VH value corrected for D is less than 2 cm

3/mol Η [50] and equals

1.6 cm3/mol H.

The results of the investigations discussed suggest that the inclusion of hydro-gen atoms (ions) in α-iron bring about a considerable dilatation of the crystal lattice of the metal. The lack of measurable changes in the lattice parameter

INTERACTION OF HYDROGEN WITH THE STRESS FIELD 275

of iron during hydrogen entry, as observed by X-ray analysis [51], is the con-sequence of an exceedingly low hydrogen solubility in this metal. Taking the Vn value into consideration, it follows from simple calculation that the dilatation effect observed at a hydrogen concentration ca = 12.6 χ 10"

5 mol H/mol Fe

corresponds to a mean change in the lattice parameter of approx. 3 χ 10" 6 nm.

9.5 Interaction of hydrogen with the stress field

It follows from thermodynamic considerations that in the elastic region, tensile stress decreases the chemical potential of hydrogen in a metal, whereas the re-verse is true for compressive stress [45]. In other words, tensile stress should increase the hydrogen solubility and vice versa for compressive stress. At equilib-rium this is expressed by the relation [52]:

ca = c0cxp(aVH/3RT) (9.11)

where ca is the hydrogen concentration in the region where the hydrostatic com-ponent of tensile stress a is operative, cQ is the hydrogen concentration in the region where a = 0, VH is the partial molar volume of hydrogen in the metal studied. A rise in hydrogen solubility induced by tensile stress was experimentally observed in palladium-silver alloys [53]. For the Fe-H system the steady-state permeation rate of hydrogen through iron membranes under galvanostatic and potentiostatic conditions was shown to increase under tensile stress and to de-crease under compressive stresses [44, 54].

Assuming the diffusivity of hydrogen in iron to be independent of stress [44], variations in the permeation rates reflect the variations in the diffusible hydro-gen concentration in the metal as induced by stress. The results have indirectly lent validity to the assumptions of the hydrogen embrittlement theory as to the segregation of hydrogen in the region of triaxial tensile stresses in the vicinity of crack tips [55].

Relation (9.11) also holds for internal stresses set up in the metal, e.g. around dislocations, inclusions, etc. At the dislocations cores, however, Fermi-Dirac instead of Boltzmann statistics should be used [56]. So for dislocation cores the relation (9.11) must be replaced by

l - c d l - c 0

cd CQ - exp (9.12)

276 DISSOLUTION AND DIFFUSION OF HYDROGEN IN IRON AND STEEL

where cd and c0 are equilibrium atom fractions of solute at dislocations and in free lattice sites, respectively, and W is the bonding energy of hydrogen with dislocations.

Plastic deformation of palladium results in an increased solubility of hydro-gen in this metal. The main reason for this is the stress field set up around dislo-cations, rather than the dislocation lines themselves [57]. Plastic deformation of high-purity iron increases the hydrogen concentration in this metal also under conditions of electrolytic charging with hydrogen [42]. According to the data reported in [58], dislocation cores play an important part in the segrega-tion of the interstitial atoms dissolved in iron.

In conformity with equation (9.12), the distribution of hydrogen between the normal interstices and dislocations depends on the value of the bonding energy, W, between hydrogen and the dislocations, which is not precisely known [11, 56]. In calculations of the number of hydrogen traps it is assumed that W = -26.8 kJ/mol H [30]. Investigations on plastically deformed ZR iron have provided information on the dependence of the sum of concentrations of the hydrogen bonded to dislocation cores and hydrogen in the field of stress of

2 0

I ι ι ι ι 1 0 0.2 0.4 0.6 0.8 1.0

c,cm

3H2/100gFe

Fig. 9.2. Concentration of hydrogen connected with dislocations, c Hi vs. diffusible hydrogen concentration, c, in cold-worked (27, 40%) ZR iron: 1—27%, 2 -40%; Τ = 296 Κ (after [59]).

dislocations, cH±9 on the diffusible hydrogen concentration, c (Fig. 9.2). As ex-pected, c Hl increases with the concentration of diffusible hydrogen and with the amount of cold work Z.

CRITICAL CONCENTRATION OF HYDROGEN 277

9.6 Critical concentration of hydrogen

The dissolved hydrogen present in iron in the form of atoms (ions) causes the metal lattice to dilate [50], interacts with dislocations to bring about a peak of internal friction [10, 11], raises the residual electrical resistivity [60] and is re-sponsible for the eifect of reversible hydrogen embrittlement of polycrystalline iron [61]. If the hydrogen concentration attains a critical value cc, micro-cracks and voids filled with molecular hydrogen develop in the metal. The presence of this form of hydrogen in the iron results in plastic deformation of the metal [62,63], decreased density [40], irreversible increase in electrical resistance [64, 65] and the appearance of partly irreversible hydrogen embrittlement [66].

According to the data of [44], the micro-cracks and voids are formed in poly-crystalline Armco iron, ZR (zone refined) iron and in single crystals virtually at the same critical concentration of diffusible hydrogen. The nucleation sites of these defects are dislocations. Another paper [46] reports that the major reason for iron blistering is the expansion of the metal lattice by the dissolved hydrogen. X-ray evidence, however, demonstrates no structural changes in iron whiskers during electrolytic hydrogen charging [67]. The development of micro-cracks and voids due to hydrogen was observed neither in an annealed high-purity iron [50, 68] nor in iron subjected to a plastic deformation to 71% [69]. Results of more recent studies [70] on the behaviour of ZR iron during cathodic polarization are shown in Fig. 9.3. For annealed membranes (curve 7)

Fig. 9.3. Concentration of diffusible hydrogen c vs. square root of current density i in cold-worked (0,75 and 89%) ZR iron membranes: 1—0%, 2—75%, 3—89%; Τ = 295 Κ (after [59]).


and membranes cold worked to Ζ = 75% (curve 2), the concentration of diffu-sible hydrogen is a linear function of the square root of the current density. In the concentration range up to 2.5 cm

3 H2/100 g Fe no blistering was noticed

in any of the methods used. Nor were blisters formed in specimens of this iron cold worked to 89% at c < L5 cm

3 H2/100 g Fe. As soon as the critical con-

centration is reached, cc » 1.5 cm3 H2/100 g Fe, blisters form vigorously and

develop, and the average concentration of diffusible hydrogen decreases with current density (curve 5). The critical concentration found in the study corre-sponds to an equilibrium pressure of molecular hydrogen pHi = 1.1 x 10

3 MPa.

At the critical concentration of diffusible hydrogen in iron cold worked to 89% the total concentration of hydrogen associated with dislocations was found to be 3 cm

3 H2/100 g Fe. Thus a considerable segregation of hydrogen on the

accumulations of dislocations in highly deformed pure iron results in a local loss of cohesion of the metal.

In Armco iron the formation of micro-cracks and voids, as judged from a significant hydrogen uptake and an irreversible dilation of the hydrogen-charged specimens, occurs at the critical concentration of diffusible hydrogen, cc = 0.03-0.04 cm

3 H2/10 g Fe. The equilibrium pressure of gaseous hydrogen

corresponding to this concentration was calculated to range from 120 to 160 MPa. The origin of the sites where microcracks and voids develop in Armco iron

and steel have not so far been determined. Most probably, however, it is asso-ciated with the micropores at the boundaries of non-metallic inclusions, most probably oxides [71, 72]. The results discussed above indicate that for high-purity iron the value of the critical hydrogen concentration is related to the nucleation of micropores. For Armco iron or steel, it is most probably related to the onset of the further development of micropores already present in the metal into micro-cracks and blisters.

Finally, mention should be made of the result of calculations based on the quantum theory of the solid state, which showed that the introduction of pro-tons or hydrogen atoms into approximately 10% of the unit cells of the lattice of a light metal results in a possible loss of cohesion of the metal [73], Unfortu-nately, no such calculations have been made in the case of iron.

9.7 Diffusion of hydrogen

Theoretical approaches to the quantitative problem of the diffusion of hydrogen in metals in terms of the classical [74] or quantum mechanical [75] models have failed to afford completely satisfactory results. They have been useful, however,

DIFFUSION OF HYDROGEN 279

in the more detailed interpretation of experimental data [18]. Results of the measurements of hydrogen diffusivity D dit various temperatures in general agree with equation (9.6). It follows from the theory that for a body-centred cubic lattice:

D0 = (l/6)a2fdexp(ZlS/*) (9.13)

where a is the lattice parameter, fd is the oscillation frequency of an impurity atom about its equilibrium position and AS is the activation entropy. If we take AS χ 0, the D0 value calculated for hydrogen in iron is 16 χ 10~

4 cm

2/s [76].

On the other hand experimental values from high-temperature diffusion studies [77-89], fall within the range 3 χ 10~

4 to 22 χ 10"

4 cm

2/s and the average value

is 9 . 85x l0 -4cm

2/ s .

Theoretical attempts to evaluate the activation energy of hydrogen diffusion in α-iron, AD9 give differing results, viz. 5.4 [90] and 14.0 kJ/mol [91]. The ex-perimental AD values are within the range 4.6 to 13.8 kJ/mol, the average being 10 kJ/mol [77-89]. Taking the average values of D0 and AD values it follows that:

D = 9.85 χ 10~4exp(- 10.000/ΛΓ) (9.14)

A plot of log D = / ( l /T) is shown in Fig. 9.4 (curve 1), where the solid line refers to the range of temperatures measured. Extrapolation to lower tempera-tures shows that at 293 Κ D20 = 1.6 χ 10"

5 cm

2/s.

Temperature, Κ 1173 773 573 473 373 273

I 1 ! 1 1 1 —

Ï0Ô Ï5Ô 200 25Ô 30Ô 350 (1/Γ)χ10

5,Κ"

1

Fig. 9.4. Logarithm of hydrogen diffusivity D in iron vs. reciprocal absolute temperature on the basis of various data: 1—average value from [77] to [89], 2—from formula (9.15), 3—[83], 4— [95], 5— [40], 6— [44], 7—formula (9.16) [42].

280 DISSOLUTION A N D DIFFUSION OF HYDROGEN IN IRON AND STEEL

From a critical analysis of the results of measurements of the hydrogen per-meation rate for iron, Gonzalez [92] gave the value of the activation energy for this process to be AP = 35.2 kJ/mol. The activation energy of diffusion was evaluated using this value (Oriani [30]) along with the heat of hydrogen dissol-ution value (27.2 kJ/mol); the author also gives a formula for the temperature dependence of hydrogen diffusivity:

D = 7.8 χ 10~4 exp(-7950/i*r) (9.15)

This formula shows that D20 ~ 3 x l 0 ~5 cm

2/s (Fig. 9.4, curve 2). Formulae

(9.14) and (9.15) describe fairly well the D vs.f(T) relation at high temperatures, but fail to fit to experimental data obtained at near ambient temperatures.

Johnson and Hill [83] obtained interesting results in the range 298 to 1053 Κ (Fig. 9.4, curve 5). The data suggest a high activation energy (32.7 kJ/mol). It should be noted, however, that in the measurements the iron specimens were charged with hydrogen at high temperatures, followed by a rapid quenching in water and storage in liquid nitrogen. During these operations in a metal con-taining dissolved hydrogen in considerable amounts, the formation of micro-defects (hydrogen traps) is unavoidable; the presence of these traps lowers the experimental D values and raises the AD values [93, 94].

The temperature dependence of diffusivity over a 283 to 363 Κ temperature range has been studied using a number of materials and various conditions for the electrolytic hydrogen charging process. Several examples are shown in Fig. 9.4. Curve 4 relates to En8 steel [95]. The low D values found from the absorp-tion curves may be explained by the large number of hydrogen traps present in steel of considerable carbon content (0.45%) and by a low concentration of the hydrogen introduced at a low current density in an alkaline electrolyte. Under vigorous hydrogen charging in an acid electrolyte with a promoter added, the diffusivity values for Armco iron are high (curve 5) [40], but even so they are lower than the values obtained at low concentrations of diffusible hydrogen for electrolytes without any additions (curve 6) [44]. The differences are due to surface phenomena as well as to irreversible changes in iron at hydrogen con-centrations above the critical value. These questions are considered in more detail later in this chapter. The highest values of hydrogen diffusivity were ob-tained for ZR iron (curve 7) [42]. If the effects of surface phenomena are elim-inated, the diffusivity values in annealed ZR iron are independent of the con-centration of the diffusing gas.

The diffusion activation energies corresponding to slopes of curves 4 to 7 are 24.7, 12.5, 5.9, 3.1 kJ/mol, respectively. The last value refers to ZR iron.

SURFACE EFFECTS IN DIFFUSION STUDIES 281

Results of measurements of hydrogen diffusivity in this iron obey the following equation [42]:

D = 3.57 x 10-4exp(-3100/#r) (9.16)

Extrapolation to 973 Κ yields a D100 value coincident with the hydrogen diffu-sivity value for ZR iron measured at this temperature by Heumann and Domke [96].

A study of hydrogen diffusion in high-purity iron, however, using gaseous hydrogen as a charging medium gave other values, in particular D0 = 1.23 χ χ ΙΟ"

3 cm

2/s and AD = 7.05 kJ/mol [36].

The low activation energy of hydrogen diffusion in α-iron suggests that, at low temperatures, the quantum mechanism of hydrogen diffusion in the metal must be taken into account [87, 97].

9.8 Surface effects in diffusion studies

One of the reasons for the differences observed in the experimental values of diffusivity of hydrogen for iron and its alloys [1, 18] is the effect of surface processes. These processes may affect the kinetics of hydrogen absorption or desorption, thereby introducing errors into the measurements. The desorption rate of the gas from steel specimens immersed in water were, for instance, found to be higher than the rate for those placed in a nitrogen atmosphere [98]. Fur-thermore it was shown that the D values depend on the thickness of the mem-branes used [99, 100]. The D = f(l) dependence cannot be explained simply on the basis of the larger number of hydrogen traps in thin membranes [10], as this relation is observed in both steel [99] and ZR iron [100].

Significant factors in the measurement of D are: the manner in which the membrane surface is prepared, the pH of the electrolyte, the cathodic current density during polarization and the presence of hydrogen entry promoters in the electrolyte [101]. This is illustrated by the curves of rise and fall of the hydro-gen permeation rate Pt/Poo = / ( / ) shown in Fig. 9.5. For ZR iron membranes with mechanically polished surfaces, an increased current density in a sulphuric acid solution only slightly affects the position of the Λ/Λο = fit) curve, which shifts towards shorter times (Fig. 9.5a). In an alkaline solution the character of the curves and their position depend strongly on current density (Fig. 9.5b). The reason for this is a thin layer of deformed metal present on the membrane surface. Measurements of hydrogen diffusivity under these conditions will thus be seriously in error. It has been shown [7, 8] that on switching off the cathodic polarization Vd = 0.33 V0, i.e. the volume of hydrogen evolved through the

Fig. 9.5. Curves P%IP& = / ( / ) f o r ZR iron membranes with mechanically polished surfaces: (a) 0.1 MH2S04, i , 1'—5 A/m

2, 2, 2'—250 A/m

2, (b) 1 M KOH, i , i'—1 A/m

2, 2, 2'—

10 A/m

2, 3,3'—250 A/m

2, Γ = 295 K.

VT, (A/m

2

Fig. 9.6. Plots of Vi/Vo vs. }/7 for ZR iron membranes polarized in various electrolytes: i - 0 . 1 M H 2S 0 4, / ' - 0 . 1 M H 2S 0 4 + A s 20 3, 2—1 M KOH, 2'—1 M KOH + A s 20 3, Γ = 295 Κ.

DIFFUSION OF HYDROGEN 283

diffusion surface should be equal to 0.33 of the total quantity of diffusible hydro-gen in the membrane (for a unit surface of V0 = cl/2). If no hydrogen traps are present in the metal, the numerical value of the coefficient m = Vd/V0 > 0.33 suggests that some surface processes occurring on the polarization side of mem-brane suppress the desorption of hydrogen.

The values of the coefficient m in acid and alkaline electrolytes are plotted in Fig. 9.6. It is clear that addition of A s 20 3 results in a considerable increase in the value of m. This is evidence that the poisoning of iron membranes with arsenic strongly suppresses hydrogen desorption. Determination of diffusivity values under these circumstances is hardly feasible.

9.9 Diffusion of hydrogen at varying concentrations

At high temperatures no effect of the hydrogen concentration c on the diffu-sivity D of the gas in iron was observed [80]. Studies of electrolytic hydrogen charging showed the experimental D values in Armco iron [93] and in ZR iron [100] to increase with the current density L Namboodhiri and Nanis [9] reported the diffusivity in Armco iron to increase with hydrogen concentration, i.e. with current density. These investigators emphasized that the concentration depend-ence D = f{c) observed by them is not due to the presence of hydrogen traps in the metal. Hydrogen diffusion studies, for Armco and ZR irons, however, performed under conditions of polarization at various current densities and in a number of electrolytes, have shown that the D = /(/) relation should not be identified with the D = /(c) relation [101], as can be seen from the results of D

12

2 h

I ι ι ι 1 0 10 2 0 χ 3 0 AO

λ/Γ, (A/m

2 y

Fig. 9.7. Effect of current density \/Ton hydrogen diffusivity D in ZR iron: / , i—0.1 M H 2S 0 4, 2, 2'—\ M KOH, 5,3'—0.1 M H 2S 0 4 + A s 20 3, Τ = 295 Κ.

284 DISSOLUTION A N D DIFFUSION OF HYDROGEN IN IRON A N D STEEL

Time, s

Fig. 9.8. Curves Λ / Λ ο = / ( / ) for Armco iron membranes in 1 M KOH (7, Γ) and in

0.1 M H 2S 0 4 + A s 20 3 (2, 2' to 5, 5') at a hydrogen concentration c < cK (7, / ') and c > cK (2,2' to 5 ,50, Γ = 295 K.

solution at a diffusible hydrogen concentration of 3.5 χ 10"3 cm

3 H2/100 g Fe.

The diffusivity values estimated from the curves lie between 7.5 χ 10"5 and

9x 10"5 cm

2/s, and are close to the D value in ZR iron. Curves 2 to 5, and 2'

measurements in ZR iron presented in Fig. 9.7. It should be noted that the D values determined from the rise (7, 2, 3) or fall (Γ, 2', 3') curves of the rate of hydrogen permeation through membranes for Pt/P^ = 0.5 coincide only in the case where measurements are carried out in a solution of pure acid (curve 1, V). The differences are substantial, however, if the acid contains added promoter (curves 5, 3'). The main significance of the results shown in Fig. 9.7 is that the experimental values of diffusivity found in the diffusible hydrogen range 0.004 to 0.3 cm

3 H2/100 g Fe (curve 1, Γ) are the same as those obtained at high con-

centrations c ^ 2.7 cm3 H2/100 g Fe (curve 5). It can therefore be stated that

the diffusivity of hydrogen in iron that is virtually free of hydrogen traps, and in which such traps are not hydrogen-induced, is independent of the concentra-tion of this gas in the metal. The high D value of (10±0.5) χ 10~

5cm

2/s at 295 Κ

should be noted. The diffusion of hydrogen in Armco iron proceeds in either a simple or a

complex manner, depending on the concentration of the diffusing gas. This is illustrated by the rise and fall curves for the rate of hydrogen permeation through the membranes shown in Fig. 9.8. Curves 7, Γ were determined in an alkaline

DIFFUSION OF HYDROGEN IN A METAL 285

to J', were obtained for the same membrane with an increasing concentration of diffusible hydrogen, greater than the critical concentration cc. The character of the curves clearly suggests that the absorption or desorption processes are not determined by the hydrogen diffusion rate. The D values calculated on the basis of these curves do not at all compare with the real value of hydrogen diffu-sivity in this material. This is due to the formation of micro-cracks and voids in Armco iron.

9.10 Diffusion of hydrogen in a metal in the presence of traps

The effect of traps in a metal on hydrogen diffusion was first considered by Darken and Smith [102]. The general diffusion equation for hydrogen under these conditions was given by McNabb and Foster [103]. Hydrogen traps may be caused by various defects of the crystal structure of the metal, such as: dislo-cations, inclusion boundaries, voids, and micro-cracks, etc.

Allen-Booth and Hewit [104] advanced a model for the diffusion of hydrogen in iron and steel containing hydrogen traps in the form of voids. They assumed firstly that hydrogen in iron alloys occurs in two forms, either dissolved in the crystal lattice of the metal, or molecular hydrogen present in voids, and sec-ondly that an equilibrium exists between these two forms. For a quantitative description of the dependence of dissolved hydrogen concentration on the equilib-rium pressure pHi of molecular hydrogen in voids, Phragmen's equation [105] was used, which accounts for the real properties of gaseous hydrogen. On the basis of the assumptions made, the authors derived an equation which relates the value of the apparent hydrogen diffusivity Dz, to the total hydrogen con-centration in a metal, the trap (void) concentration, the coefficient of trap expan-sion due to the pressure of the gas in the traps, and temperature. The solution of this equation for various cases showed that at room temperature, increased volume concentration of voids from zero to 10"

4 lowers the value of the apparent

diffusivity by about two orders of magnitude [106]. The calculations also show that the Dz value should depend to a considerable extent on the trap expansion coefficient. Authors of reports [104, 106] claim that their diffusion model accu-rately fits the kinetics of hydrogen desorption from steel at room temperatures.

McNabb and Foster [103] demonstrated theoretically that in the case of diffu-sion of hydrogen through metal membranes containing hydrogen traps, the time lag tz is a function of not only the membrane thickness and the real diffu-sivity D, but that it also depends on the number of hydrogen traps Nx, concentra-tion of diffusible hydrogen c, and the kinetic parameters of the processes of the


hydrogen trapping and releasing. Oriani [30] assumed dislocations to be hydro-gen traps and claimed that there is an equilibrium between the trapped hydro-gen and the diffusible hydrogen in the metal lattice. This assumption simplifies the expression for the time lag given in [103] and leads to the following formula for the value of the apparent diffusivity of hydrogen ;

Dz==~== DiU-HNx)-

1 (9.17)

where

1 ί 6iVL 6NI Γ Kc , L Kc\ c [ Kc K

2c2 [ # L \ Nh j

where Κ = cd/ch = exp(— WIRT) is the equilibrium constant. According to [30], W = -26.8 kJ/mol, and thus Κ = 5.5 χ 10

4. Using this,

the value of Η can be estimated for a given concentration c and, in consequence the number of traps Nx can be determined, provided D and Dz are known from experiment.

Results of the measurement of the apparent diffusivity of hydrogen in plas-tically deformed steels [30], Armco iron [107] and ZR iron [108] are in qualita-tive agreement with the dislocation model of traps. The determination of quanti-tative relations, however, requires more accurate experimental data. Results obtained for Armco iron with various amounts of cold work Ζ are shown in

(9.18)

Percentage of cold work, %

Fig. 9.9. Effect of the percentage of cold work of Armco iron on the values of apparent diffu-sivity ( A , A» and Dd) and on hydrogen concentrations (c, Vdi and Vc), 1 M KOH, i = 90 A /m

2,

Τ = 295 Κ.

DIFFUSION OF HYDROGEN IN A METAL 287

Fig. 9.9 [109]. The values of the apparent diffusivity of hydrogen, Z>a, Dd and Dt found from the curves of rise or fall of hydrogen permeation rate, for Λ/Λο = 0.5, decrease strongly with the amount of cold work as do diffusivity values found from the breakthrough-time data. The concentration of diffusible hydrogen c in the membranes on the charging side is independent of Z, whereas the hydrogen concentration Vd increases by over two orders of magnitude as Ζ rises from zero to 25%. The total hydrogen concentration VG in the membranes studied is much higher than Vd. The number of hydrogen traps per cm

3, NX9 estimated using

formula (9.17) and the experimental values of Dd and c in iron cold worked to 1.3, .5, 15, and 25% is 1.8 χ 10

1 9, 7.9 χ 10

1 9, 1.35 χ 10

20 and 1.97 χ 10

2 1, respect-

ively. It follows, from a comparison of the Vc and Vd curves in Fig. 9.9, that in Armco iron during plastic deformation, voids are formed which accumulate hydrogen in a quantity that increases with Z. Studies of the aging-temperature effect of plastically deformed Armco iron on the values of the apparent diffu-sivity [107, 109] and absorption of hydrogen [109] showed that voids formed in the metal were not completely removed during aging at high temperatures. The total hydrogen concentration in membranes cold worked to various degrees and aged at 1113 and 1273 Κ is shown by curves 1 and 2 of Fig. 9.10, respectively.

Percentage of cold work, %

Fig. 9.10. Hydrogen concentration Ve in Armco iron of different percentage of cold work after ageing at various temperatures: i—1113 K, 2—1273 K, 1 M KOH, / = 90 A/m

2, Τ = 295 Κ.

As can be seen, the α-y phase transformation in iron does not lead to the destruc-tion of voids. The presence of voids may lower the hydrogen diffusivity values by several orders of magnitude [110].

Some information on the dislocation trap model has been derived from studies on hydrogen diffusion in plastically deformed ZR iron in which disloca-


tions are already present, but in which voids and micro-cracks have not yet formed. Curves of the apparent diffusivity DZ vs. degree of cold work for this type of iron at various hydrogen concentrations are presented in Fig. 9.11. An

TO 20 30 AO 50 60 70 80~ Percentage of cold work,%

Fig. 9.11. Effect of the percentage of cold work on the values of the apparent diffusivity Dz of hydrogen in ZR iron at various hydrogen concentrations: 1—0.013, 2—0.127, 3—0.254,

4—0.635, 5—1.27 cm

3 H2/100 g Fe, Τ = 295 Κ.

increased degree of cold work enhances the density of dislocations in the metal, thereby decreasing DZ. The DZ values depend strongly on hydrogen concentra-tion, in particular at low concentrations and high degrees of cold work.

According to Oriani [30], for low trap occupancy by hydrogen, the number of traps can be evaluated from the relation:

DZ = d{\ + ^ Y (9.19)

It can be easily verified that, for low hydrogen concentrations, formulae (9.17) and (9.19) yield virtually the same results. For high concentrations formula (9.17) or its simplified form

(9.20)

should be used.

The Nx values found from formula (9.17) raise doubts, since for high hydro-gen concentrations Nx values are much lower than for the low concentrations.

REFERENCES 289

It should be noted that the calculated Nx value depends on the values of the apparent diffusivity and equilibrium constant; the value of the latter in turn depends on the bonding energy of hydrogen, W, with traps. The W value is not precisely known [48] and there is no information as to whether or not it is de-pendent on the degree of occupancy of the traps by hydrogen. The discrepan-cies in Nx values in materials of comparable degrees of cold work reach five to six orders of magnitude [30, 107]. Thus, despite the considerable progress in the elucidation of the role of traps in the hydrogen dissolution and diffusion pro-cesses in metals [111-115] witnessed in recent years, experimental difficulties have greatly restricted the explanation of this complex phenomenon.

The results of hydrogen diffusion studies discussed above indicate that, at room temperature, the diffusivity of hydrogen in iron with a small number of traps is high (~ 10~

4 cm

2/s) and independent of hydrogen concentration. In

iron containing dislocations in the form of traps, experiments lead to values of the apparent diffusivity. This value decreases with the number of traps and increases with hydrogen concentration. In diffusion studies with technological materials hydrogen concentration is important. At concentrations exceeding the critical one, micro-cracks and blisters are formed, and when describing hydrogen diffusion in such cases the dislocation-trap model should be replaced by the void-trap model.

A large number of research work concerning the thermodynamic and kinetic behaviour of the hydrogen-iron system recently appeared [116-121]. For example Kiuchi and McLellan [122] have analysed the complications arising in the measurements of the solubility and diffusivity of hydrogen in annealed and deformed iron due to the both non-regular mixing statistics and the efects of defects in the bulk and on the surface of the samples. The problem of diffusion and trapping of hydrogen in iron has been extensively discussed by Kçdzie-rzawski [123].

Despite of the progress in this research field many problems are not yet re-solved and more precise data are necessary to explain a number of fundamental questions.

9.11 References

[1] Smialowski, M., Hydrogen in Steel, Pergamon Press, Oxford 1962. [2] Johnson, E. W., Hill, M. L., Acta Met., 3, 99 (1955). [3] Dus, R., Smialowski, M., Acta Met., 15, 1611 (1967). [4] Barrer, R. M., Diffusion in and through Solids, Cambridge University Press, Cambridge

1941.


[5] Devanathan, Μ. Α. V., Stachurski, L., Proc. Roy. Soc. Α., 270, 90 (1962). [6] Boes, N., Zuchner, H., Hydrogen in Metals, Intern. Meeting, Birmingham 1976. [7] McBreen, J., Nanis, L., Beck, W., / . Electrochem. Soc, 113, 1218 (1966). [8] Nanis, L., Namboodhiri, T. K. G., / . Electrochem. Soc, 119, 691 (1972). [9] Namboodhiri, T. K. G., Nanis, L., Acta Met., 21, 663 (1973).

[10] Gibala, R., Trans. Met. Soc AIME, 239, 1574 (1967). [11] Hirth, J. P., Metallurg. Trans., TLA, 861 (1980). [12] Worsham, J. E., Wilkinson, M. G., Shull, C. G., / . Phys. Chem. Solids, 3, 303 (1957). [13] Wollam, E. O., Cable, J. W., Kochler, W. C , / . Phys. Chem. Solids, 24, 1141 (1963). [14] da Silva, J. R. G., Stafford, S. W., McLellan, R. B., Hydrogen in Metals, Intern. Meeting,

Birmingham 1976. [15] Oriani, R. Α., Proc. Symp. on Stress Corrosion Cracking, Columbus, Ohio 1967. [16] Fast, J. D. , Interaction of Metals and Gases, Vol. 1, Philips Techn. Library, Eindhoven

1965. [17] Fast, J. D. , Interaction of Metals and Gases, Vol. 2, Philips Techn. Library, Eindhoven

1971. [18] Geld, Ρ. V., Ryabov, R. Α., Hydrogen in Metals (in Russian), Metallurgiya, Moskva 1974 [19] Armbruster, M. H , / . Amer. Chem. Soc, 65, 1043 (1943). [20] Dyakonov, A. G., Samarin, A. M., Izv. AN SSSR, OTN, 9, 813 (1945). [21] Martin, E., Metals and Alloys, 1, 831 (1930). [22] Luckemeyer-Hasse, L., Schenk, H., Arch. Eisenhiittenw., 6, 209 (1932-1933). [23] Baukloch, W., Muller, R., Arch. Eisenhiittenw., 11, 509 (1937). [24] Sieverts, Α., Zapf, G., Morits, Η , Z. phys. Chem., A183, 19 (1938-1939). [25] Morozov, A. N., Hydrogen in Steel (in Russian), Metallurgizdat, Moskva 1950. [26] Eichenauer, N., Kunzig, H., Pebler, Α., Z. Metalik., 49, 220 (1958). [27] Blake, P. D. , Jordan, M. F., Pumprey, W. I., Spec. Rep. Iron Steel Inst., 73, 76 (1961). [28] Deshman, S., Scientific Basis for Vacuum Technology (in Russian), Mir, Moskva 1964. [29] Wagner, R., Sizman, R., Z. angew. Phys., 18, 193 (1964). [30] Oriani, R. Α., Acta Met., 18, 147 (1970). [31] Smith, D. P., Hydrogen in Metals, University Press, Chicago 1950. [32] Archakov, Yu. L, Teodorovich, V. P., Zhur. Prik. Khim., 32, 2667 (1959). [33] Kumnick, A. J., Johnson, Η. H., Met. Trans., 6A, 1087 (1975). [34] Hill, M. L., Johnson, E. W., Trans. Met. Soc. AIME, 221, 222 (1961). [35] Hill, M. L., Johnson, E. W., Trans. Met. Soc AIME, 215, 717 (1959). [36] Raczynski, W., Phys. Stat. Sol., (a) 48, K27 (1978). [37] Beloglazov, G. M., Hydrogen Charging of Steel in Electrochemical Processes (in Russian),

Izd. Leningrad. Univ., Leningrad 1975. [38] Karpenko, G. W., Kripyakevich, R. S., The Effect of Hydrogen on Steel Properties (in

Russian), Metallurgizdat, Moskva 1962. [39] Raczynski, W., Smialowski, M., Bull. Acad. Polon. Sci., sér. sci. chim., 8, 209 (1960). [40] Raczynski, W., Bull. Acad. Polon. Sci. sér. sci. chim., 15, 25 (1967). [41] Raczynski, W., Talbot-Besnard, G., C. R. Acad. Sc. Paris, série C, 269, 1498 (1969). [42] Raczynski, W., not published, 1975. [43] Baranowski, B., Ber. Bunsengesel. Phys. Chem., 76, 714 (1972). [44] Beck, W., Bockris, J. Ο. M., McBreen, J., Nanis, L., Proc. Roy. Soc, A 290, 220 (1966). [45] Li, J. C. M., Oriani, R. Α., Darken, L. S., Z. physik. Chem., N. F., 49, 271 (1966).

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[46] Bockris, J. Ο. M., Beck, W., Genshaw, M. Α., Subramanyan, Williams, F. S., Acta Met., 19, 1209 (1971).

[47] Townsend, H. E., Jr., Corrosion, 26, 361 (1970). [48] Hirth, J. P., Stress Corrosion Cracking and Hydrogen Embrittlement of Iron Base Alloys,

Eds. R. W. Staehle et al., NACE, Houston, TX, 1977, p. 1. [49] Wagenblast, H., Wriedt, H. Α., Met. Trans., 2, 1393 (1971). [50] Raczynski, W., Stress Corrosion and Hydrogen Embrittlement, Intern. Symposium, Dres-

den 1975. [51] Straumanis, M. E., Kim, D. Κ., Z. Metallkunde, 60, 272 (1969). [52] Oriani, R. Α., Trans. Met. Soc. AIME, 236, 1368 (1966). [53] Wriedt, H. Α., Oriani, R. Α., Acta Met., 18, 1753 (1970). [54] de Kazinczy, F., Jernkont Ann., 139, 885 (1955). [55] Steigerwald, Ε. Α., Schuller, F. W., Troiano, A. R., Trans. Met. Soc. AIME, 218, 832

(1960). [56] Hirth, J. P., Carnaham, B., Acta Met., 26, 1795 (1978). [57] Flanagen, T. B., Lynch, J. F., Clewley, J. D., Turkowicz, B., Hydrogen in Metals, Intern.

Meeting, Birmingham 1976. [58] Podgurski, H. H., Oriani, R. Α., Met. Trans., 3, 2056 (1972). [59] Raczynski, W., Symposium Hydrogen Embrittlement and Corrosion of Metals, Jablonna

1977, p. 67. [60] Raczynski, W., Elefant, D., Phys. Stat. Sol., 14, 139, 1972. [61] Cornet, M., Raczynski, W., Talbot-Besnard, S., Mém. Sci. Rey. Métal, 69, 27 (1972). [62] Krupkowski, Α., Smialowski, M., Proceedings of the Electrochemical Conference, PWN,

Warszawa 1955, p. 271. [63] Tetelman, A. S., Wagner, C. N. J., Robertson, W. D. , Acta Met., 9, 205 (1961). [64] Van Ooijen, D. J., Fast, J. D., Acta Met., 11, 211 (1963). [65] Flis, J., Bull. Acad. Polon. Sci., sér. sci. chim., 12, 809 (1964). [66] Raczynski, W., Smialowski, M., Zashchita Met., 5, 482 (1969). [67] Flis, J., Janko, Α., Bull. Acad. Polon. ScL, sér. sci. chim., 12, 51 (1964). [68] Kufudakis, Α., Raczynski, W., Czech. J. Phys., Β 26, 1330 (1976). [69] Raczynski, W., Talbot-Besnard, S., C. R. Acad. Sci. Paris série C, 270, 602 (1970). [70] Raczynski, W., unpublished results, 1976. [71] Bernstein, I. M., Mater. Sci. Eng., 6, 1 (1970). [72] Smialowski, M., ref. [48], p. 405. [73] Olszewski, S., Bull. Acad. Polon. Sci., sér. sci. chim., 20, 443 (1972). [74] Wert, C , Zener, C , Phys., Rev., 76, 1169 (1949). [75] Flynn, C. P., Phys. Rev., Β 1, 3966 (1970). [76] Hill, M. L., Hydrogen Embrittlement in Metal Finishing, Reinhold Publ. Corp., New York

1961.

[77] Sykes, C , Burton, H. H., Gegg, C. C , / . Iron Steel Inst., 156, 155 (1947). [78] Geller, W., HoSun-Tak, Arch. Eisenhuttenw., 21, 423 (1950). [79] Ryabov, R. Α., Geld, P. V., Izd. Vuzov. Chernoi Metal, 2, 83 (1959). [80] Bryan, N. L., Dodge, B. F., AIChE Journal, 9, 223 (1963). [81] Salii, V. I., Geld, P. V., Ryabov, R. Α., Fiz.-Khim. Mekh. Mat., 5, 96 (1970). [82] Zvezdin, Yu. I., Belyakov, Yu. N., Fiz.-Khim. Mekh. Mat., 3, 349 (1967). [83] Johnson, E. W., Hill, M. L., Tfans. Met. Soc. AIME, 219, 1104 (1960).


[84] Stross, T. M., Tompkins, F. S., / . Chem. Soc, 2, 230 (1956). [85] Eichenauer, W., Kunzing, H , Pebler, Α., Z. Metallkunde, 49, 220 (1958). [86] Carmichael, D. C , Hornuday, J. R., Morris, A. E., Parlée, Ν. Α., Trans. Met. Soc.

AIME, 218, 826 (1960). [87] Volkl, J., Alefeld, G., in: Diffusion in Solids, Recent Developments, Eds. A. S. Nowick,

J. J. Burton, Academic Press, New York 1975, p. 232. [88] Wagner, R., Sizmann, R., Z. Angew. Phys., 18, 193 (1964). [89] Heumann, Th., Primas, D. , Z. Naturforsch., A 21, 260 (1966). [90] Ferro, Α., / . Appl. Phys., 28, 895 (1957). [91] Vykhodets, V. B., Goltsov, V. Α., Geld, Ρ. V., Fiz. Met. Metallov, 25, 1090 (1968). [92] Gonzalez, O. D. , Trans. Met. AIME, 245, 667 (1969). [93] Raczynski, W., Stelmach, S., Bull. Acad. Polon. ScL, sér. sci. chim., 9, 633 (1961). [94] Frank, R. C , Swets, D . E., Fry, D. L., / . Appl. Phys., 29, 1262 (1958). [95] Newman, J. F., Shreir, L. L., Corr. Sci., 11, 25 (1971). [96] Heumann, Th., Domke, E., Jul-Conf—6, 2, 492 (1972). [97] Eremeev, V. S., Fiz. Met. Metallov, 25, 1095 (1968). [98] Hudson, R. M., Stragand, G. L., Corrosion, 16, 253t (1960). [99] Raczynski, W., Arch. Hutn., 3, 59 (1958).

[100] Raczynski, W., Talbot-Besnard, S., C. R. Acad. Se. Paris., série C, 269, 1253 (1969). [101] Raczynski, W., Symposium Hydrogen Embrittlement and Corrosion of Metals, Jablonna,

1977, p. 63. [102] Darken, L. S., Smith, R. P., Corrosion, 5, 1 (1949). [103] McNabb, Α., Foster, P. K., Trans. Met. Soc, AIME, 227, 618 (1963). [104] Allen-Booth, D. M., Hewit, J., Acta Met., 22, 171 (1974). [105] Phragmen, G., Jernkont Ann., 128, 537 (1944). [106] Allen-Booth, D. M., Atkinson, G., Bilby, Β. Α., Acta Met., 23, 371, 1975. [107] Kumnick, A. J., Johnson, Η. H., Met. Trans., 15, 1199 (1974). [108] Cornet, M., Thèse, Université Paris VI, 1973. [109] Raczynski, W., unpublished results, 1977. [110] Evans, G. M., Rollason, E. C , / . Iron Steel Inst., 207, 1484 (1969). [ I l l ] Riecke, E., Arch., Eisenhiittenw., 47, 247 (1976). [112] Kumnick, A. J., Johnson, H. H., Acta Met., 28, 33 (1980). [113] Pressouyre, G. M., Bernstein, I. M., Acta Met., 21, 89 (1979). [114] de Avillez, R. R., Lauf, R. J., Altstetter, C. J., Scripta Met., 15, 909 (1981). [115] Wert, Ch. Α., in: Hydrogen in Metals, Topics in Applied Physics, Eds. G. Alefeld, J. Volkl,

Springer-Verlag, Berlin, 29, 305 (1978). [116] McLellan, R. B., Acta Met., 27, 1655 (1979). [117] Hirth, J. P., Met. Trans., HA, 861 (1980). [118] Birbaum, Η. K.: "Hydrogen in Metals" in: Proc. JIMIS 2, Minakami 1979, Trans. Jap.

Inst. Met., 21, 1 (1980). [119] Azou, P. Y., Third Intern. Conf. on Hydrogen and Materials, Paris, 1, 61 (1982). [120] Raczynski, W., Less Common Met., 101, 383 (1984). [121] Puis, M. P., in: Hydrogen Degradation of Ferrous Alloys, Eds. R. A. Oriani, J. P. Hirth,

M. Smialowski, Noyes Publications, Park Ridge, USA, 1985, p. 114. [122] Kiuchi, K., McLellan, R. B., Acta Met., 31, 961 (1983). [123] Kçdzierzawski, P., ref. [121], pp. 251-288.

10. Effect of Hydrogen on Mechanical Properties and Dislocation Structure of Iron and Steel

10.1 Introduction

The hydrogen charging of metals results in a serious deterioration of their mech-anical properties. This effect is known as hydrogen embrittlement.

Hydrogen embrittlement is a common problem in a considerable number of industries such as chemical and petrochemical industries, aircraft and construc-tion engineering, and in the transport of liquid and gaseous fuels. Among the multitude of metals and alloys that undergo hydrogen embrittlement are auste-nitic, martensitic and ferritic steels, nickel, titanium and aluminium alloys, be-ryllium bronzes, tantalum, niobium, vanadium and zirconium. The most serious consequence of hydrogen embrittlement, however, is in the case of steels, es-pecially the high strength ones, because of their wide use in construction and apparatus and possibility to be charged with hydrogen due to exposing to the various aggresive environments [1-3].

It is difficult to make a simple classification of the types of deterioration of a material induced by hydrogen, due to the diversity of variations in mechanical properties, damage symptoms and the different hydrogen sources encountered in practice.

Hirth and Johnson [4] distinguish the following kinds of hydrogen embrittle-ment:

(1) delayed failure, (2) reduced plasticity, (3) brittleness in contact with gaseous hydrogen. Delayed failure (also referred to as static fatigué) consists of the formation

of brittle cracks in a hydrogen-charged material subjected to stresses below the yield stress. This kind of damage is characterized by an induction time and a critical stress below which cracks do not occur. The crack develops gradually: the failure is accelerated by a higher amount of hydrogen, increased load, and the formation of localized stress concentration. Carbon and low-alloy steels are

294 EFFECT OF HYDROGEN ON MECHANICAL PROPERTIES OF IRON

specially subjected to this kind of embrittlement; delayed failure in a high strength steel may occur at a hydrogen concentration as low as 0.1 cm

3/100 g Fe and

at a strain value corresponding to a 10% yield stress. Reduced plasticity of a material manifests itself in a decreased elongation and

a reduction of area in the tensile test of the hydrogen-charged specimens as compared with the hydrogen free material damage or cracking of material ap-pears. The larger the initial strength of steel, the greater is the reduced plastic-ity.

Brittleness occurs in plastically strained steels in contact with gaseous hydro-gen. As distinct from delayed failure, no induction time or critical stress is ob-served, although increased load accelerates the failure.

Hydrogen embrittlement may be reversible in case where the material re-covers its original properties, once the hydrogen source is removed and the gas is desorbed. If, however, the hydrogen absorbed causes micro-cracking in the ma-terial, the recovery of properties is incomplete and the embrittlement effect becomes irreversible.

A thorough review of the various problems concerning hydrogen entry into the iron and steel and its effect on material properties has been recently published [5]. The present summary is focused on some aspects of hydrogen effects on material structure which in the authors' opinion are specially important to understand the hydrogen induced failure of steels.

10.2 Hydrogen induced cracking of iron and steel

During the long history of investigation of hydrogen embrittlement phenomena of iron and steel, a number of concepts such as the formation of blisters [6, 7], reduction of the surface energy of the crack formed [8], interaction of hydrogen with dislocations and hydrogen effect on dislocation motion [9-14], decohesion of the iron crystal lattice by hydrogen atoms [15-17], have been proposed. The mechanism of hydrogen embrittlement is believed to depend strongly on the structure and composition of the material; it is also considered that one par-ticular concept cannot sufficiently explain all various forms of this effect [3, 16].

In high strength steels, the crack formation is accounted for by decohesion of the matrix [15-17]. Hydrogen diffuses into those regions in the material where local triaxial tensile stresses operate. Lattice decohesion and the formation of cracks take place as a result of interaction between the accumulated hydrogen atoms and iron atoms.

In case of mild steel and iron of different purity, hydrogen ingress could also

HYDROGEN INDUCED CRACKING OF IRON A N D STEEL 295

cause the formation of brittle cracks [10, 18, 19] which is thought to be chiefly due to hydrogen interaction with dislocations. One of the proposed mechanism is schematically shown in Fig. 10.1 [10]. In plastically deformed material, disloca-

(α) ο ο

ΤΓδ ο ο ο ο ο oooooôôôo^P

Fig. 10.1. Successive stages of a hydrogen-induced crack formation in iron (after [10]): O — hydrogen atoms, ·—non-metallic inclusions, J_—dislocations.

tions along with hydrogen segregated around them move in crystal. An accumu-lation of dislocations on non-metallic inclusions results in the accumulation of hydrogen in the region between the inclusion and the matrix and may result in a crack nucleation. The formation of micro-cracks is also observed in inclu-sion-free iron [18-20]. Here, at regions of plastic incompatibility (grain boundary, especially points where a number of grains is in contact), a local rise in stress occurs which, together with the local rise in hydrogen concentration, results in crack formation, cf. Fig. 10.2.

In hydrogen charged material, a further development of already formed in-ternal cracks, as well as external fissures produced by corrosion, fatigue or local-ized overload occurs at applied stresses much lower than those required for the crack propagation in the hydrogen-free material. The unit processes of the crack propagation in real materials are not as yet thoroughly understood. The real state of material in the plastic zone formed at the crack tip is not exactly known,


(a) (b)

Fig. 10.2. Model of the crack formation in iron due to hindrance: Ο—dislocation cross-slip by hydrogen; #—hydrogen atoms, J_—dislocations.

nor is the strain and stress distribution in the adjacent material. The results ob-tained in the fracture mechanics study give only scarce and undirect informa-tion concerning the state of metal in the crack tip and its change during the crack propagation, especially in low and medium strength, polycrystalline and multi-phase material. On the basis of assumed or qualitatively demonstrated specific conditions of the material at the crack tip a number of concepts have been pro-posed to account for the tendency of metals affected by hydrogen to undergo brittle fracture.

Similar as in case of crack nucleation, some models based on hydrogen effect on iron-iron interaction [15, 16], surface energy [21], cleavage energy [21] shear and tensile elastic moduli [22] have been proposed to explain the enhanced pro-pagation of cracks in hydrogen charged material.

Figure 10.3 shows a model of atom distribution in the crack tip zone in a me-

Fig. 10.3. Model of hydrogen-induced crack propagation (after [16]).

HYDROGEN INDUCED CRACKING OF IRON A N D STEEL 297

tal [16]. The local stress concentration is responsible for a weakening of the bonding strength between the atoms in the crack tip. Crack propagation con-sists in breaking successive weakened bonds. Hydrogen getting into the region of the weakened bonds causes further reduction of the binding strength which results in a lower value of stress required for the crack to develop [16].

Hydrogen can get to the critical area from the exterior of the metal, if a sui-table aggressive environment is present in the crevice. A reduced cohesion of the metal crack tip may also be a result of the hydrogen already present in the metal, cf. Fig. 10.3. The crack tip zone is believed to be specially preferred for hydrogen absorption and thus for lattice decohesion, either as a result of increased elec-tronegativity of the metal in the stressed state [23, 24] or due to high value of plastic strain build up near crack tip which favours a local increase in hydrogen concentration [25].

Another concept is that the adsorbed hydrogen reduces the surface energy of the metals, which should result in a decrease in energy required to cause the material to separate, and should, hence, favour brittle propagation of the crack [26-29]. For example, the preferential reduction of surface energy of the {110} slip planes and simultaneously a decrease in the elastic moduli of the lattice normal to these planes is assumed to cause the microstep to form and to more readily fracture the metal along {110} planes [21].

In real materials, a drop in stress due to formation of the plastic zone in the crack tip hinders further crack propagation. Thus all the factors that make that stress relaxation more difficult work in the direction of facilitating the brittle fracture of metals. Hydrogen is postulated to prevent slip by blocking the move-ment of dislocation as expected in hydride forming metals [22] or by altering the energy of dislocation formation and metal cleavage energy which renders the slip more difficult [21].

The relaxation of the local stresses arisen around a dislocation pile-ups in iron at room temperature occurs as a result of dislocation escaping from the pile-ups by the cross-slip which is very easy in iron. Consequently, no brittle cracks are observed in iron at room temperature. However, the hindrance of cross-slip by hydrogen results in a brittle development of crack [19]. The for-mation of long pile-ups of dislocations due to the suppression of cross-slip by hydrogen, leading to a rising stress concentration and brittle cracking is shown schematically in Fig. 10.2.

It seems reasonable to assume that, depending on the material, the hydrogen source, and the initial rate of crack propagation, the embrittlement effect of hydrogen may be due to one or several of the mechanisms discussed. At a slow


rate of crack propagation during hydrogen charging, hydrogen enters the metal and causes its decohesion; at a fast cracking rate, on the other hand, the absorp-tion effect is to be expected.

It should be noted, however, that despite the numerous studies no direct experimental proof has been found that hydrogen reduces the surface energy of iron at room temperature. The relevance of absorption theories is also called into question by the fact that no qualitative differences are found in the effect of external hydrogen, compared with the effect of hydrogen already present in the metal.

The concepts reviewed indicate that the critical point is the understanding of the effect of hydrogen on the cohesion of the crystal lattice of iron, and of its effect on the movement of dislocations.

The effect of hydrogen on the lattice cohesion can be estimated from a study of the change in the elasticity of a metal on charging with hydrogen. Informa-tion on the effect of hydrogen on dislocation mobility is obtained by the study of change in

(1) mechanical properties (yield stress, flow stress, strain at yield stress and work hardening in tensile and torsion tests);

(2) dislocation structure during straining; (3) internal friction in comparison with the hydrogen free metal.

10.3 Effect of hydrogen on the elastic properties of iron

The formation of a solid solution brings about changes in the elastic constants of the solvent matrix, both for the substitutional [30-32] and interstitial [33] solutions. For example, carbon dissolved in α-iron lattice lowers the shear mo-dulus by about 0.8% per one atom percent [33].

In studies on the effect of hydrogen on the elastic modulus carried out chiefly on the hydride-forming metals, both lowering of the elastic modulus (for Ti [34] and for Ti-Mo [35]) and rising of modulus (for Nb [36, 37]) were observed. In case of Pd and Pd-Ag alloys, the complex effect of hydrogen and deuterium on elastic constants was established [38]. It should be noted that the forma-tion of hydride in the above materials may affect the results.

Studies on the effect of hydrogen on the elastic properties of α-iron are scarce. This is due to experimental difficulties arising from the negligible solubility of hydrogen in the α-Fe lattice [1]. The introduction of hydrogen by the cathodic charging causes the supersaturation of hydrogen which may result in a plastic strain of the material, and the effects measured will be largely a consequence of

EFFECT OF HYDROGEN ON PROPERTIES OF IRON 299

this. On the other hand, variations in the elastic constants caused by hydrogen dissolved in iron are supposedly quite small requiring very sensitive technique for the measurements.

The effect of hydrogen on the elastic modulus and the elastic limit of high-purity iron single crystals has been investigated in a tensile test [10]. The obtained values of the modulus and the elastic limit for hydrogen-free and cathodically charged single crystals do not differ. The method used is insufficiently accurate, however, because comparison is made between the results obtained for different specimens, and the values of elastic constants are determined from the load-extension graphs, which are not precise enough.

To avoid the above inconveniences, the comparative measurements of the elastic modulus of hydrogen-free and hydrogen-charged iron specimens were carried out by determining the free oscillation frequency of a torsion system con-taining the iron sample in the wire form [39]. The shear modulus is proportional to the second power of the free oscillation frequency, and the change in the fre-quency of system after hydrogen charging of specimen is due to change in the shear moduli. The hydrogen charging and desorption processes were performed directly in the measuring chamber, shown in Fig. 10.4, so that specimen did not have to be removed from the apparatus, and the sequence of measurement was carried out for each specimen: in the initial state, after charging with hydrogen

Fig. 10.4. The apparatus for measurement of the free oscillations frequency and internal fric-tion with simultaneous hydrogen charging of specimen by silent discharge (after [39]): 1— electrolyser, 2—washers, 3—hydrogen gas ionizer, 4—specimen.


at 100 Κ and after desorption at room temperature. Hydrogen ionized by the silent discharge technique was introduced into the apparatus. Simultaneous measurement of internal friction allowed to check whether the hydrogen charging caused any change in the iron structure. It was shown that the used method yielded a suitable amount of dissolved hydrogen and no irreversible change in the metal. This was very important, since the damage formed at hydrogen charging causes itself the change in elastic moduli [40]. The temperature relations of the oscillation frequency and of internal friction are shown in Fig. 10.5 for hydrogen-free and

ο

62 4'—ι 1 ι 1 1 » '39

100 120 140 160 180 200

Temperature, Κ

Fig. 10.5. The temperature dependence of the square of the free oscillation frequency (circles) and of internal friction (triangles) of hydrogen-free ( Ο , Δ ) and hydrogen-charged ( · , A ) iron specimens (after [39]).

hydrogen-charged state of the same specimens. The decrease in both, the oscilla-tion frequency and the internal friction upon hydrogen charging was reversible, and upon desorption of hydrogen their initial values were recovered. The pre-cise value of the hydrogen concentration could not be determined. However, on the basis of comparative measurement of hydrogen content performed for the similar specimens it was estimated that 1% of hydrogen atoms causes the elastic modulus of iron to decrease much more than 1% of carbon [41].

The decrease in elastic moduli due to hydrogen has been confirmed in the

EFFECT OF HYDROGEN ON PROPERTIES OF IRON 301

in situ measurements of surface wave (SW) velocity during hydrogen charging [42], The change of the SW velocity was estimated by measuring an ultra-sonic critical reflective angle yc. The method consists of transmitting an ultra-sonic beam of various angles against the surface of specimen immersed in elec-trolyte, and detection of the amplitude of reflected wave. At a critical angle yc a sharp minimum in the reflectivity associated with generation of the surface wave on the sample is observed. According to Shell's law, the SW velocity VR is equal to Vi/sinyc, where Vx is the velocity of ultrasonic beam in electrolyte. In performed experiments, specimen fitted in revolving goniometer and the transducer for generation and detection of ultrasonic beam were used, as shown in Fig. 10.6a [42]. In Fig. 10.6b, the typical spectra of the reflected wave inten-sity vs. revolving angle for iron specimen are shown. Despite the higher noise detected during cathodic charging, the minima corresponding to the critical

1 1 1 . . . .

/

J

Revolving angle

Fig. 10.6. Measurement of hydrogen effect on surface wave velocity (after [42]): (a) schematic diagram of ultrasonic wave refractiometer: 1—revolving sample, 2—electrolyte, 3—trans-ducer, 4—recorder; (b) reflected wave intensity vs. revolving angle for iron; 1—0.5% H2SO4 no polarization, 2—0.5% H 2S 0 4 polarization 60 A/m

2, 5—0.5% H 2S 0 4 + 4 mg/1 A s 20 3,

polarization 60 A / m

2; yc—critical angle.


angle yc could be distinguished. As compared with the value without polariz-ation, the critical angle increases upon polarization to about 0.5° in pure acid and to about 1.13° in acid with the recombination poison. The effect corresponds to the decrease in the SW velocity of about 1.5% and 3.4%, respectively. As known, the velocity of an ultrasonic wave depends on elastic moduli of material. The decrease in the SW velocity for iron charged with hydrogen indicates, therefore, that hydrogen increases the elastic constants.

The lattice cohesion energy is roughly proportional to the value of the mo-dulus of elasticity. The decreased value found for the elastic modulus of hydro-gen charged iron, therefore, is an experimental support of the decohesion theory. A particularly large reduction in the cohesion energy is to be expected at sites of locally higher hydrogen concentrations. On the other hand, the results of the described above experiments give the support for the assumption that hydro-gen decreases the Peierls stresses of dislocation, since those stresses determine the resistance of a crystal lattice to the motion of dislocations, and are propor-tional to the shear modulus value. Thus it may be expected that the interstitial hydrogen would increase the mobility of dislocations.

10.4 Effect of hydrogen on the plasticity of iron

The hydrogen charging of a material results in considerable change in plas-ticity and strength as measured in the mechanical tests. That change reflects the hydrogen effect of dislocation mobility. Unfortunately, the main conclusion that might be drawn from numerous experiments is that depending on material composition, purity and structure, amount of hydrogen and it partition, method of hydrogen charging as well as mechanical test conditions (mode of loading, strain rate, temperature, etc.) the hydrogen manifests itself in different ways. It was found, for example, that the dimensions of specimens used influence the effect of hydrogen on work hardening observed in steel [43, 44]. No simple re-lationship between the load mode and the obtained results has been found. In [45] the decrease in number of twists to the point of fracture was observed for both the reversible and irreversible brittleness caused by previous hydrogen charging (cf. Fig. 10.7; the reversible brittleness region corresponds to a re-covery of the value to the initial value upon complete desorption of hydrogen). No differences were observed in the tensile tests of the same specimens, as long as no cracks and crevices resulting from hydrogen charging appeared.

In [17] tensile straining of hydrogen charging specimens caused their cracking, while specimens tested in torsion exhibited no brittle cracking.

EFFECT OF HYDROGEN ON PLASTICITY OF IRON 303

Reversible brittleness region

Irreversible brittleness region

8 12 16 20 24 Hydrogen content, cm

3H2/100g Fe

28

Fig. 10.7. The number of twists until fracture for 18% Cr ferritic steel vs. hydrogen content (after [45, 53]): tf—immediately after hydrogen charging, b—after complete desorption.

0 20 40 60 80 100 μ g Original content of hydrogen in cracks

Fig. 10.8. The yield stress vs. hydrogen content in the cracks after the first hydrogen charging, SR (after [47]): O—hydrogen-free, #—1 ppm hydrogen dissolved, approx. 0.75R in cracks, •—approx. 0.1SR in cracks.

In a torsion test of iron with simultaneous electrolytic hydrogen charging, a drop in the number of twists to the point of fracture and in the value of tor-sional strength were observed; work hardening, however, was increased [46]. The reason for such discrepances is the complexity of processes taking place


during the straining of the metal as well as the possibility of the formation of macro-defects arising from the introduction of hydrogen, especially with materials of inadequate purity.

This is supported by the results of a study [47] of the yield stress and strain for Armco and high-purity (70 ppm C + N and 450 ppm 0) iron S. By means of electrolytic charging, cracks were produced in the specimens, then hydrogen was desorbed, and again the cathodic polarization. Figure 10.8 shows the rela-tion of yield stress and strain vs. degree of damage in the material, using as a measure of the damage the amount of hydrogen in the cracks after the first hydrogen charging. As can be seen, elongation decreases with the degree of damage. Moreover, the hydrogen content in the specimen upon its recharging with hydrogen has no effect on the relationship. The author of the work cited draws the conclusion that a major part in the variations of the stress-strain curves observed in other studies should be ascribed to irreversible hydrogen-charging effects, that is, to the formation of cracks and plastically strained re-gions.

Attention has recently been called to studies on high-purity iron, often as single crystals, and charged with hydrogen under mild conditions [10, 14, 48-51]. In this way the secondary effects mentioned above can be avoided.

The variation in flow stress and work hardening (da/de) vs. strain observed

Plastic strain

Fig. 10.9. Stress and work hardening da/σε vs. strain for hydrogen-free high-purity iron (solid line) and for similar specimens hydrogen-charged during straining (broken lines) (after [10]).


during tensile tests of high-purity iron without hydrogen and hydrogen charged is shown in Fig. 10.9 [10].

In the initial stages of the deformation process, the hydrogen content facili-tates plasticity. As the critical strain value is passed, work hardening increases much faster in the hydrogen-charged material. On hydrogen charging, tensile strength and elongation also increase. All the above effects are more pronounced for single crystal specimens. In later works, the hydrogen-induced decrease in flow stress and increase in strain in stage I of single crystal deformation was observed [51]. In stage III, hydrogen increased work hardening and flow stress; the strain until fracture was smaller in hydrogen-charged than in hydrogen-free samples [51]. However, the considerable rise in elongation to fracture upon hydrogen charging was observed for the iron whiskers [50],

In studies on high-purity iron (90 ppm C + N + O ) [5] the variation in flow stress in tensile tested material charged with hydrogen of high fugacity was found to depend strongly on grain size [52]. It can be seen from Fig. 10.10 that the hydrogen charging of iron of a grain size above 30 μιη results in a decrease in flow stress σΑ, internal stress Ci (i.e. the stress due to long range dislocation interaction) and the effective stress of motion of dislocations at. In finer-grain

Grain size d, 500 300 50 20 10

0 100 300 400

Fig. 10.10. Yield stress <rA, internal stress au and effective stress ae vs. grain size of hydrogen-charged and hydrogen-free iron (after [52]).


c ο

φ

0.9

1.1

1.3

0.8

0.9

1.2

1.1

0 8±1 15±1 crrr »

3H2/100g FeCr

Fig. 10.11. Variation of the relative half-width of the (110) reflection with hydrogen con-tent in 18% Cr ferritic steel (after [53]): β0—the half-width value for the initial condition of material.

iron the quantities studied were found to increase in value upon hydrogen-charg-ing.

A hydrogen-induced drop in internal stress was observed in high purity 18% Cr. ferritic steel [53], A measure of this stress was the half-width of the (110) X-ray line, β. Figure 10.11 shows that hydrogen charging of a metal in a range outside the irreversible brittleness region (i.e. up to approx. 8 cm

3 H2/100 g Me) re-

sults in reduced internal stress. A rise in internal stress was apparent, where the hydrogen introduced brought about permanent material damage (hydrogen content about 15 cm

3/100 g Me).

Studies have also been made on the effect of hydrogen on stress relaxation by plastic strain over a wide range of strain [49]. During relaxation, as distin-guished from active straining, the dislocation structure of the metal remains unchanged. Thus the dislocation motion can be studied without being compli-cated by the processes of dislocation multiplication, submicro-crack formation, etc. Figure 10.12 shows the drop in load with time, recorded for a strained single-crystal specimen of iron. The hydrogen charging accelerates relaxation. An analysis of the relaxation curves obtained demonstrated that hydrogen reduces


0

11 ι ι 11 ii I . ι ι ι ι n u l 1—ι ι ι n u i 1—ι—ι ι ι ml 1

0.1 1 10 100

Time, min Fig. 10.12. The effect of hydrogen on the constant relaxation at a plastic strain of single crystal iron (after [49]).

the internal stress of the dislocation network at [19]. This brief review of the results of works of different authors seems to indicate that hydrogen facilitates dislocation slip along preferred directions, under conditions when a complex dislocation structure has not yet been formed. ;

— Hydrogen enhances the motion of screw dislocations [11, 13, 14] due to facilitation of kink formation [13] or by altering Peierl's potential [11],

8 0

Strain, %

Fig. 10.13. The effect of hydrogen on the stress-strain curve of high-purity iron (after [48]): cathodic current density 80 A/m

2, strain rate 8.3 χ 10" V"

1; (a) comparison of the curves for

hydrogen-free and hydrogen-charged specimens; (b) the effect of turning the polarization on (Î) and off (J,).


— Hydrogen decreases the annihilation distance of dislocation the elastic field around the dislocation and thus their interaction in parallel slip planes [51].

For large strain and complex slip pattern, the hydrogen effect is highly in-volved. However, it may be assumed that hindering of cross-slip by hydrogen [50] leads to higher work hardening than in the case of iron without hydrogen.

Recently, the effect of hydrogen charging on the flow stress of iron, iron differently pure polycrystalline, single crystal and whiskers tested in tension [14, 48, 50, 54, 55], have been carried out by different groups of scientists. The example of such a result is shown in Fig. 10.13 [48]. It can be seen that the cessation of hydrogen charging results in increasing flow stress corresponding to the hydrogen-free specimen (broken line). On the reapplying of the polariz-ation, a momentary rise in stress is followed by its decrease. Detailed investiga-tion has demonstrated that interstitial impurities are responsible for the instanta-neous rise in stress on the application of polarization.

According to [14, 48, 54], application of hydrogen charging causes the decrease in flow stress only in the case of high purity iron tested at temperature below 200 K. However, the hydrogen induced softening has been observed also for steel [56], as well as very distinct softening due to hydrogen charging is exhibited by iron whiskers tested at room temperature [50]. The softening of iron at applica-tion of cathodic polarization is accounted for by acceleration of screw dislocations [13, 14, 54]; creation of new dislocation on metal surface due to hydrogen charg-ing [55]; facilitating of dislocation escape from the metal [57].

10.5 Effect of hydrogen on the dislocation structure of strained iron

Variations in dislocation structure on straining, caused by hydrogen, have been investigated primarily with the aid of a transmission electron microscope. In that case previously charged and strained specimen were studied. The results, however, are scarce because of the following factors:

— the complexity of the dislocation structure that arises in the straining of hydrogen-free iron (compared with the structure of metals of a f.c.c. lattice), even for single crystals;

— the difficulty in distinguishing between the effects induced by hydrogen and the secondary effects caused by the presence of cracks formed by cathodic polarization;

— the preparation of specimens involves numerous operations, during which hydrogen desorption, recovery processes etc. may occur.

The evidence from relevant research suggests that hydrogen may cause an

EFFECT OF HYDROGEN ON DISLOCATION STRUCTURE OF IRON 309

increase in dislocation density. This effect was observed in the tensile testing of a hydrogen-charged foil specimen [48]. The increase observed, however, does not exceed the scatter of this quantity as determined in individual grains of the hydrogen-free metal. In later work [14], the dislocation density in a spe-cimen deformed during hydrogen-charging was found to be larger than that in the specimens deformed without hydrogen.

In torsion tests of iron wires during charging, a much faster rise in disloca-tion density was noted in hydrogen atmosphere than in air [58].

In studies of the variation of electrical resistivity and of the thermoelectric effect during straining of iron in tension accompanied by hydrogen charging, it was found that the charging brings about an increase in the number of dislo-cations and a fall in their mobility [59].

Hydrogen-induced changes in dislocation structure can be also inferred from the change on the slip pattern on the surface of deformed material. However, diverse results have been also obtained. At temperature below 200 K, the wavy slip lines were observed on the surface of the single crystals deformed in ten-sion during hydrogen charging [60]. At room temperature, the hydrogen charg-ing of twisted polycrystalline iron caused the slip line to become finer and close-ly spaced that in hydrogen-free iron [19]. In hydrogen charged whiskers the straight slip line homogeneously distributed along the whole specimens were observed in contrary to wavy slip band concentrated near the fracture surface in hydrogen-free crystals [50].

A thorough study of the slip pattern of massive iron single crystals and its change due to the hydrogen charging was done [51]. It was found that in hydro-gen charged crystals deformed in stage I, straight and more closely situated slip lines were formed than those in hydrogen-free crystals. At higher deforma-tion the specific "layer" structure of slip pattern was observed on the surface of hydrogen-charged crystals. It was confirmed by the TEM study [61] that the finer slip steps in stage I manifested the formation of dislocation sheets inside the metal during stage I, thinner than those in pure iron.

Results of studies of the variation of dislocation structure due to hydrogen charging of strained metal are difficult to interpret, as hydrogen affects the mo-tion of dislocations, the number of dislocations, and their arrangement, and these factors in turn determine the hydrogen absorption by strained metal. During torsion at a low strain rate, an enhanced hydrogen absorption was observed to occur in specimens. At higher strain rate, on the other hand, hydrogen uptake declined as a result of a considerable increase in dislocation density [62].

During simultaneous hydrogen-charging and plastic straining of a metal,

3 1 0 EFFECT OF HYDROGEN ON MECHANICAL PROPERTIES OF IRON

not only is the dislocation motion affected by the interaction of hydrogen, but also hydrogen is transported by the dislocations. This process may lead to the development or growth of blisters containing hydrogen at a pressure higher than the equilibrium value [63]. This causes the re-emission of hydrogen to the metal and it may affect the motion of successive dislocations. A thorough theor-etical consideration of this process, referred to as kinetic supersaturation, for the hydrogen-ferritic or austenitic steel system demonstrated, however, that super-saturation in these metals is negligible [64].

The progress in High Voltage Electron Microscopy has recently allowed to study the effect of hydrogen on dislocation motion and structure during in situ tests. It was shown [65] that screw dislocation velocity as well as dislocation multiplication rate increased with increasing hydrogen gas pressure for foil iron specimens tensile strained in microscope chamber.

10.6 Determination of the interaction of hydrogen with dislocations in iron and steel by measurement of internal friction

10.6.1 The internal friction method

The internal friction method consists in measuring the attenuation of the oscilla-tions produced in a solid by a periodically variable stress. The phase shift be-tween strain and stress occurring in the elastic region is caused by dissipation of energy in the system as a result of a number of processes, such as dislocation oscillations, diffusion of point defects, and the interaction of dislocations with point defects.

The temperature effect of internal friction is studied at a constant (moderate) amplitude of oscillations, or alternatively the internal friction effect is investi-gated as a function of the amplitude at a constant temperature. In the former case the relaxation spectrum is obtained, consisting of a series of peaks result-ing from the relaxation processes mentioned [66]. In the latter case, the in-ternal friction background is measured. This is independent of the temperature up to a certain value, beyond which it starts to increase. The occurrence of an internal friction background is due to vibrations of dislocation loops (in the amplitude-independent region) and to the escape of dislocations from their pinning atmospheres of foreign atoms (in the amplitude-dependent region).

The internal friction method is widely used in studies on the behaviour of atoms of foreign elements in metals and their alloys, as it is one of the few that permit determination of the binding energy of foreign atoms with dislocations

INTERACTION OF HYDROGEN WITH DISLOCATIONS IN IRON 311

and the determination of the partition of these atoms between the lattice and dislocations, without imposing irreversible changes in the metal.

The segregation of interstitial atoms around dislocations (the formation of impurity atmospheres) shows up in internal friction investigations in the form of the

(1) appearance of the Snoek-Koster peak; and (2) variation in the internal friction background. The effect of interstitial atoms on the motion of dislocations results in a

changed background (in the amplitude-independent region). Unfortunately, the iron(steel)-hydrogen system is once again too complex. Consequently, the results of studies of hydrogen-charged iron by this method often show large discrepancies. However, a valuable characteristic of the method is its selectiv-ity, that is, by a suitable choice of the frequency and the temperature an indi-vidual relaxation process can be studied with no interference from other factors. For example, the interaction of hydrogen with dislocations can be studied in iron containing hydrogen-induced blisters and cracks.

10.6.2 The hydrogen Snoek-Koster peak

In plastically strained hydrogen-charged pure iron, at a frequency of 1 Hz, a peak is observed at about 130 K, on the curve showing the temperature dependence of internal friction. The height and position of the peak, referred to as the Snoek-Koster peak, depend on the quantity of hydrogen absorbed, the degree of strain, the kind of material, applied frequency, etc. Some results are shown in Fig. 10.14.

The Snoek-Koster peak has its origin in the interaction of hydrogen atoms with non-screw dislocations (most probably with 71-deg edge dislocations). Studies on the Snoek-Koster peak are aimed, on the one hand, at establishing the effect of the condition of a material on the behaviour of the peak, and on the other hand, at setting up theoretical models describing the relaxation processes corresponding to this peak. The latter are considered to be helpful in estimating the binding energy of hydrogen with dislocations, and the hydrogen concentra-tion at dislocations.

Most research has been devoted to the determination of the activation energy of the peak, usually by studying the variation of oscillation frequency with the position of the peak. The value of the activation energy of the peak is usually assumed to be 0.30 to 0.44 eV [67, 69-76].

Despite a number of studies, considerable discrepancies still exist between


Temperature, Κ

Fig. 10.14. Variation in the height and position of the hydrogen Snoek-Kôster peak in iron vs. hydrogen concentration and strain: (a) 90% strain (after [67]); (b) 16% strain (after [68]); (c) 45% strain (after [69]); (d) constant hydrogen concentration, variable strain (after [57]); Δ—hydrogen-free specimens, Ο—maximum hydrogen content χ, φ—lower hydrogen content.

the results obtained by individual workers (cf. Fig. 10.14). This is particularly true of the height and shape of the peak, as well as its position. The discrepancies cannot be explained solely by the different frequencies used, or by differing hydrogen concentrations or strain values.

Fairly coherent results were obtained from studies on the effects of hydrogen concentration and strain on the peak position and height, carried out on rela-tively pure materials. A drop in hydrogen concentration was shown to reduce the peak height and to shift its position towards lower temperatures [67, 68]. Increased plastic strain, on the other hand, raises the height of the peak, as well as lows the temperature at which it appears [67].

It has been suggested that discrepancies in the results from studies on ma-


terials of lesser purity may be due to the effect of other interstitial atoms on the interaction of hydrogen with dislocations [68, 69].

In a ferritic chromium steel, the peak appears when free dislocations are present. It is absent, if dislocations are blocked by nitrogen atoms; thus the formation of saturated Cottrell atmospheres prevents binding between hydrogen and the dislocations [45, 53].

Similar results which show an analogous effect of carbon were obtained else-where [70]. Occasionally, however, the peak did appear, despite the use of ma-terials containing impurities in larger quantities [68, 77].

The presence of carbon or nitrogen atoms may affect not only the height or position of the peak but also its shape [70].

It was also noted that in underaged specimens, the position of the peak is variable and depends on the kind of material, the ageing time, and the degree of plastic strain [68, 72, 78]. In Fig. 10.15, the shift in the position of hydrogen

Snoek-Koster peak for material strained at; 15% and 50% and aged for 10 and 30 hrs is shown [72]. The peak shift was caused by the introduction of hydrogen in the differently aged metal. A study of the amplitude-dependent internal fric-tion of differently aged samples demonstrated that the variations in peak charac-


teristics are related to the variation in the dislocation loop length which in turn is related to the concentration of C and Ν atoms around dislocations [72] :

6 ppmC + Ν

In the case of high purity iron, no change in the loop length during ageing after deformation and also no change in internal friction peak temperature in differ-ently aged samples were found, as seen in Fig. 10.16 [72],

ΊθΟ 120 140 160 180 200 Temperature, Κ

Fig. 10.16. The temperature dependence of internal friction of high-purity iron with 15% strain

and various ageing times for hydrogen-charged specimens (after [72]).

It was found also that the peak exhibits a fine structure. Saturation with hydrogen from the gaseous phase gave rise to two peaks, whereas saturation by the electrolytic method resulted in a single broad peak [73].

Two peaks were also observed in studies made by the magnetic disaccomoda-tion method [74, 75]; the peaks appeared in specimens charged with small quan-tities of hydrogen. Increased hydrogen concentration caused the peaks to come closer until they coalesced into a single broad peak.

The mechanism of the Snoek-Koster relaxation has for many years been described in terms of the Schoeck model [79]. This model views a dislocation as a vibrating string whose vibrations are hindered by the interstitial atoms form-

er

ο


ing the Cottrell atmospheres. The relaxation time of the process responsible for the peak formation can be described by the following equation:

9nkTlèCd (10.1) 16j/3Gé

3D*

where Τ is the absolute temperature, l0 is the mean length of the dislocation loop, G is the shear modulus, b is the Burgers vector, k is the Boltzmann constant, cd and Z>* are concentration of the dislocation-bound interstitial atoms around dislocations and their diffusivity in the vicinity of dislocations, respectively.

In conformity with the model, the concentration of the interstitial atoms at dislocations and their diffusivity are related to temperature by an exponential dependence :

cd = ciexp(- | fr) (10.2)

D* = A ) e x p ( — ( 1 0 . 3 )

using Fermi-Dirac statistics to describe the distribution of interstitial atoms around dislocations and in the lattice, cd and ci9 respectively, these concentra-tions may be defined as follows:

where Nd, Ni are the numbers of the interstitial atoms around dislocations and in the lattice, respectively, and Gd, G{ are the number of interstices available for these atoms near dislocations and in the lattice, respectively.

As is shown in [80, 81], a fairly complex expression is obtained from the above assumptions, where the activation energy of the Snoek-Kôster peak depends also on the degree of saturation of dislocations (and hence on the binding energy, temperature and frequency, etc.). For concentrations close to saturation, the activation energy Q was claimed to be near the HD value, whereas for dilute solutions it reaches a value HD + HB [80]. If it is additionally assumed that HD equals the sum of the activation energy of diffusion of interstitial atoms in the perfect lattice and of binding energy, the activation energy of the Snoek-Koster peak should vary from a value of H0 + HB to H0+2HB [80]. This would indicate that the binding energy of hydrogen with dislocations should lie between 0.11 and 0.22 eV.


A detailed analysis [82], however, reveals that the Schoeck model based on the above assumptions disagrees with the experimental evidence. The value of the pre-exponential term r 0, which also depends on the degree of dislocation saturation, evaluated on the basis of this model, agrees with experiment only with the assumption of a dilute solution of hydrogen atoms around dislocations, whereas in actual fact it is considered that this solution approaches saturation. In the latter case, however, the τ 0 factor estimated from the Schoeck model, is higher by some orders of magnitude than the experimental value [81].

Currently the Snoek-Koster peak is increasingly widely interpreted on the basis of the models developed by Seeger [83] for interstitial impurities and by Hirth [13], especially for the hydrogen Snoek-Koster peak. The models are based on the assumption that the impurity atoms next to dislocations form double kinks on the latter. In accordance with [13] the activation energy of the Snoek-Koster peak is then expressed by the formula

Q = 2Fk + H0/2-(G(fb6/2nyi

2 (10.5)

where Fk is the free energy of formation of an isolated kink, a is the shear stress driving the bowout, and G is the shear modulus. In this equation the bind-ing energy of hydrogen to dislocations does not appear at all. It should be empha-sized that the model affords a good agreement between the value of r 0 evaluated with its use, and the experimental value.

As has been mentioned, some recent papers report that the Snoek-Koster peak may be composed of two separate peaks, which may account for its con-siderable breadth. The authors tend to agree that the peak observed at lower temperatures is due to the interaction of hydrogen atoms with non-screw dislo-cations, so that it corresponds to the classical Snoek-Koster peak. In the case of the high-temperature peak, there no consensus as to its origin and its behaviour with varying hydrogen concentration. Among the various expla-nations are: hydrogen interacting with screw dislocations [75], and clusters of hydrogen interacting with edge dislocations [73]. A more involved mechanism is also suggested in [74],

The above considerations show why, at present, the internal friction method is of no help in an unambiguous determination of either the binding energy of hydrogen with dislocations or the value of the absolute hydrogen concentration in the region of dislocations. There is no way of deciding whether the shifting of the peak position observed in impure materials is due to variations in the binding energy of hydrogen with dislocations, to variations in the hydrogen diffusion processes in the dislocation stress field, or whether it originates from


a much more involved phenomenon. Similarly, the fine structure of the Snoek-Koster peak raises some doubts, and an entirely satisfactory mathematical description of the relaxation process is lacking. Despite these problems, however, investigations of the peak undoubtedly indicate that there is a strong interaction between hydrogen and dislocations which most likely results in the strong satu-ration of dislocations by hydrogen atoms, and that there are differences be-tween the interaction of hydrogen with screw and non-screw dislocations.

10.6.3 The internal friction background

Relatively few studies have been concerned with measurements of the effect of hydrogen on the internal friction background [78, 84-87] and on the behaviour of dislocation peaks occurring in hydrogen-free high-purity iron [76, 88-90].

Investigations carried out with high-carbon iron and steels on the internal friction background at room temperature and above demonstrated a reduced background, induced by hydrogen. The background reduction is assigned to locking of the dislocations by hydrogen atoms. This reasoning, however, cannot

Strain amplitude ε0 Fig. 10.17. The effect of hydrogen on the amplitude dependence of internal friction of high-pu-rity iron at 100 Κ : Ο—-on hydrogen-charging, · — o n desorption.


be applied to high temperature studies where hydrogen is known to move with the dislocations. The rise in internal friction on hydrogen charging occasionally observed is most probably the result of plastic strain which, in turn, is due to a strong charging with hydrogen of the contaminated material. Concentrations as low as 0.03 g/cm

3 Fe are known to bring about some visible plastic strain

in iron [1]. Studies have been made by authors of present paper with a weak hydrogen-

charging source, i.e. gaseous hydrogen excited by electric discharge under low pressure, on high-purity unstrained iron, to avoid some secondary, hydrogen-unrelated effects. As seen from Fig. 10.17, internal friction under these condi-tions decreases, in both the amplitude-dependent and amplitude-independent regions. The effect is not entirely reversible. However, the internal friction back-ground undergoes no measurable change, when steel is hydrogen-charged under the same conditions. This may be explained by a strong locking of dislocations by atoms of other impurities.

Hydrogen also brings about background variations in the temperature range where α (30 Κ) and γ (300-350 Κ) dislocation peaks occur in pure iron. These peaks are known to be due to double kink generation on the screw and non-screw dislocations, respectively [89, 90], so the altering of the position or the height of those peaks may indicate the interaction of hydrogen atoms with screw or edge dislocations. In fact, the effects of hydrogen charging on the position and height of above peaks have been found. However, because of diversity of ob-tained results no exact conclusions on hydrogen effect on movement of disloca-tion, especially of the non-screw ones could be made. α-Peak was found to de-crease [76, 88] and shift to higher temperature [76] upon hydrogen charging. In the case of γ peak, its decrease due to hydrogen charging [76] as well as increase [73] have been observed.

10.7 Concluding remarks

From the variety of experimental results summarized here and elsewhere [5], the following conclusions concerning hydrogen effect on iron lattice and dislo-cation motion leading to the metal fracture can be drawn:

1, Hydrogen decreases the elastic constants of iron, so the hydrogen-induced decohesion of iron-iron atoms may be expected, as well as reduction of Peierls stresses of dislocations.

2 . Hydrogen increases the iron plasticity. This is restricted to the conditions under which the easy glide of dislocation takes place. The hydrogen-induced

REFERENCES 319

softening is more pronounced in whiskers and massive single crystals than in polycrystals where the dislocation free paths are quite small. The facilitation of dislocation motion leads at further stages of deformation to the occurrence of dislocation structure much more dense and complex than in hydrogen-free metal.

3 . Hydrogen increases work hardening of iron and causes brittle fracture. Hydrogen renders the complex movement of dislocations (cross-slip, climbing) more difficult. This prevents the relaxation of complex dislocation structure already formed at previous stage of deformation of hydrogen-charged metal. As a result, the high stress concentration is build locally causing the brittle frac-ture of iron.

10.8 References

[1] Smialowski M., Hydrogen in Steel, Pergamon Press, Oxford 1960. [2] Smialowski M., Zashchita Met., 3, 3 (1967). [3] Smialowski M., Preprints Conf. Stress Corrosion Cracking and Hydrogen Embrittlement

of Iron Base Alloy, Firminy 1973, B-l. [4] Hirth, J. P., Johnson, H. H , Corrosion, 32, 3 (1976). [5] Hydrogen Degradation of Ferrous Alloys, Eds. R. A. Oriani, J. P. Hirth, M. Smialowski,

Noyes Publ., New York 1985. [6] Zapfe, C , Sims, C , Trans. AIME, 145, 225 (1941). [7] de Kazinczy, F., Iron Steel Inst., 177, 85 (1954). [8] Petch, N. J., Phil., Mag. 1, 331 (1956). [9] Louthan, M. R., Caskey, G. R., Donovan, J. Α., Rawl., D . E., Mat. Sci. Engng., 10, 357

(1972).

[10] Gourmelon, Α., Mem. Sci. Rev. Met., 72, 475 (1975). [11] McMahon, C. J., Jr., in: Hydrogen Effects in Metals, Eds. I. M. Bernstein, A. Thompson,

Met. Soc. AIME, 1981, p. 219. [12] Shin, K. S., Park, C. C , Meshii, M., in ref. [11], p. 209. [13] Hirth, J. P., Met. Trans., HA, 861 (1980). [14] Kimura, H., Trans. Japan bits. Metals, 26, 527 (1985). [15] Troiano, A. R., Trans. ASM, 52, 54 (1960); Hydrogen in Steel, British Iron and Steel

Inst., Harrogate, England, 1961, p. 1. [16] Oriani, R. Α., Josephic, P. H , Acta Met., 72, 1065 (1974). [17] John, C-St., Gerberich, W. W., Met. Trans., 4, 589 (1973). [18] Rath, Β. B., Bernstein, I. M., Met. Trans., 2, 2972 (1971). [19] Lunarska, E., Scripta Met., 15, 1149 (1981). [20] Vehoff, H., Neuman, P., in ref. [5], p. 686. [21] Bernstein, I. M., Mat. Sci. Engng., 6, 1 (1970). [22] Gilman, J. J., in ref. [3], C-l. [23] Tiller, W. Α., Schrieffer, R. Scripta Met., 4, 57 (1970).


[24] McLellan, R. B., Harkins, C. C , Mat. Sci. Engng., 5, 18 (1975). [25] Nanis, L., Namboodhiri, T. K. G., Acta Met., 21, 663 (1975). [26] Karpenko, G. V., Proc. 1st Intern. Congr. Hydrogen and Materials, Paris 1972, p. 355. [27] Litvin, A. K., Tkachev, V. I., Fiz.-khim. mekh. mat. 2, 27 (1976). [28] Marichev, V. Α., Fiz.-khim. mekh. mat., 3, 3 (1974). [29] Shchukin, Y. D., Fiz.-khim. mekh. met., 8, 3 (1976). [30] Kôster, W., Rauscher, W. Z., Metallkunde, 39, 111 (1948). [31] Zener, C, Acta Cryst., 2, 163 (1949). [32] Speich, G. R., Schwoeble, A. J., Leslie, W. C , Met. Trans., 3, 2031 (1972). [33] Speich, G. R., Schwoeble, A. J., Leslie, W. C , Met. Trans., 4, 318 (1973). [34] Kolachev, B., Konstantinov, K., Bukhanova, Α., Fiz. Met. Metallov., 26, 572 (1968). [35] Paton, Ν. E., Buck, O., Williams, J. C, Scripta Met., 9, 687 (1975). [36] Buck, O., Thompson, D. O., Wert, C. Α., / . Phys. Chem. Soc, 32, 2331 (1971). [37] Wriedt, Η. Α., Oriani, R. Α., Scripta Met., 8, 203 (1974). [38] Mazzolai, F. M., Papago, G., Franco, R., / . de Physique-Colloque, 44, C9-411 (1983). [39] Lunarska, Ε., Zielinski, Α., Smialowski, M., Acta Met., 25, 305 (1977). [40] Carpenter, S. H., Fawks, J. E., Scripta Met., 15, 699 (1981). [41] Smialowski, M., Lunarska, E., Zielinski, Α., Proc. 2nd Intern. Congr. Hydrogen ana Ma-

terials, Paris 1977, 2C-3. [42] Zielinski, Α., Lunarska, E., Starzynski, G., / . de Physique-Colloque, 44 C9-481 (1983). [43] Asano, S., Otsuka, R., Scripta Met., 12, 287 (1978). [44] Smialowski, M., Scripta Met., 13, 393 (1979). [45] Lunarska, E., Smialowski, M., Proc. 5th Intern. Congr. Metallic Corrosion, Tokyo 1972,

p. 177. [46] Yaremchenko, N., Shved, M., Karpenko, G., Fiz.-khim. mekh. mat., 7, 96 (1971). [47] Kusch, H.-G., Proc Intern. Symposium Stress Corrosion Cracking and Hydrogen Embrit-

tlement, Dresden 1975, p. 71. [48] Matsui, H., Morija, S., Kimura, H., Proc. 4th Intern. Congr. Strength of Metals and Alloys,

Nancy 1976, p. 3.

[49] Lunarska, E., Scripta Met., 11, 283 (1977).

[50] Lunarska, E., Wokulski, Z., Acta Met., 30, 2173 (1982).

[51] Lunarska, E., Zarubova, N., Novak, V., Kadeckova, S., Scripta Met., 17, 705 (1983).

[52] Bernstein, I. M., Scripta Met., 8, 343 (1974).

[53] Lunarska, E., Smialowski, M., Fiz.-khim. mekh. mat., 9, 3 (1973).

[54] Matsui, H. Kimura, H., Morija, S., Mat. Sci. Engn., 40, 207 (1979).

[55] Wada, H., Sakamoto, K., Proc, 3rd Intern. Congr. Hydrogen and Materials, Paris 1982,

p. 527.

[56] Lunarska, E., Scripta Met., 15, 1077 (1981).

[57] Lunarska, E., in ref. [5], p. 321.

[58] Shved, M., Yaremchenko, N., Fiz-khim. mekh. mat. 3, 113 (1973).

[59] Shved, M., Fiz.-khim. mekh. mat., 1, 107 (1975).

[60] Kimuta, Α., Matsui, H., Kimura, H., Mat. Sci. Engng., 58, 211 (1983).

[61] Lunarska, E., Zarubova, N., Libovicky, S., unpunblished results, 1983.

[62] Erdman-Jesnitzer, F., in ref. [26], p. 215.

REFERENCES 321

[63] Tien, J. Κ., Thompson, A. W., Bernstein, I. M., Richards, R. J., Met. Trans., 7A, 821 (1976).

[64] Johnson, H. H., Hirth, J. P., Met. Trans., 7A, 1543 (1976). [65] Tabata, T., Birnbaum, H. K., Scripta Met., 17, 947 (1983). [66] de Batist, R., Internal Friction of Structure Defects in Crystalline Solids, North-Holland,

Amsterdam 1972. [67] Gibala, R., Trans. AIME, 239, 1574 (1974). [68] Sturges, Ο. M., Miodownik, A. P., Acta Met., 17, 1197 (1969). [69] Miodownik, A. P., Achar, B. S., in ref. [26], p. 84. [70] Conophagos, E., Plusquellec, J., Azou, P., Bastien, P., in ref. [26] p. 97. [71] Kikuta, J., Sugimoto, K., Ochiai, S., Iwata, K., in ref. [26], p. 144. [72] Zielinski, A. Lunarska, E., Smialowski, M., Acta Met., 25, 551 (1977). [73] Sakamato, K., Shimada, M., / . de Physique-Colloque, 42, C5-109 (1981). [74] Ke, T. S., Kronmùller, H., Seeger, Α., Sun, Z. Q., in ref. [11], p. 281. [75] Vetter, K., Steeb, H., Kronmùller, H., in ref. [41], 205.

[76] Moser, P., Dufresne, J. P., Ritchie, L G., Internai Friction and Ultrasonic Attenuation in Solids, University of Tokyo Press, Tokyo 1977, p. 473.

[77] Weiner, L., Gensamer, M., Acta Met., 5, 692 (1957). [78] Heller, W. H., Acta Met., 9, 600 (1961). [79] Schoeck, G., Acta Met., 11, 617 (1963). [80] Miner, R. E., Gibala, R., Hultgren, F. Α., Acta Met., 24, 233 (1976). [81] Lunarska, E., Zielinski, Α., Scripta Met., 12, 721 (1978). [82] Lunarska, E., Zielinski, A. in ref. [5], p. 289. [83] Seeger, Α., Phys. Stat. Sol, 55a, 457 (1979). [84] Ke, T. S., Young, P. T., Scientia Sinica, 3, 261 (1954). [85] Mima, G., Mizuta, M., / . Iron Steel Inst., Jap., 43, 132 (1957). [86] Syrykh, L., Ryabov, R., Geld, P., Fiz.-khim. mekh. mat., 10, 51 (1974). [87] Shumilov, M., Vachev, Z., Izv. VUZ, Chernaya Met., 1, 136 (1970). [88] Takita, K., Sakamoto, K., Scripta Met., 10, 399 (1976). [89] Dufresne, J. F., Seeger, Α., Groh, P., Moser, P., Phys. Stat. Sol, 36a, 579 (1976). [90] Hivert, V., Groh, P., Frank, W., Ritchie, I., Moser, P., Phys. Stat. Sol, 46a, 89 (1978).

11. Change in the Real Structure of Nickel due to the Ni ^ βΝϊΗ Phase Transformation

11.1 Introduction

Changes in the structure of metals, induced by hydrogen, have a significant effect on their physical and chemical properties. A particularly strong effect is observed in those metals that form hydride phases.

The concept of the real structure of a solid has been developed especially in X-ray topography, to embrace a description of the imperfections in the crystal lattice. In the present context the "real" structure will be taken to include deviations from the perfect crystal structure such as low- and high-angle grain boundaries, dislocations, and stresses.

The solubility of hydrogen in nickel at room temperature under the standard pressure of 0.1 MPa is very low [1-8]. It is about 5 χ 10"

3 at %, rising to about

5 x l 0 ~2 at % around 1273 Κ [9]. Charging with hydrogen from the gaseous

phase at 1273 K, followed by a rapid cooling of the specimen freezes the hydro-gen in the nickel lattice and allows atomic concentrations of the order of hun-dredths per cent to be reached. The hydrogen present at such concentrations in nickel forms a solid solution, which is referred to as the α phase of the Ni-H system. This solid solution hydrogen has an important influence on the plastic and mechanical properties of nickel. This question has been discussed by numerous authors [1, 2, 6, 7], and has been the subject of a specific review [8].

Many studies lead to the interesting conclusion that hydrogen can interact with dislocations [5-7, 10] and that the hydride phase can locally form in re-gions near dislocations. The data for the solubility of hydrogen in the α phase indicate that the concentration of the hydrogen dissolved in a perfect lattice differs only slightly from that of the hydrogen present in the defect areas of the lattice.

NICKEL HYDRIDE 323

11.2 Nickel hydride

Hydrogen concentrations exceeding by several orders of magnitude those achieved by hydrogen charging from the gas phase can be obtained by charging nickel electrolytically in a sulphuric acid solution. Under a suitable current density and in the presence of substances that act as penetration promoters [11, 12], a new hydrogen-rich phase detectable by X-rays, is formed [13]. This phase, a nickel hydride referred to as the β phase, is produced in the surface layer with a thickness that depends on the electrolytic charging conditions [14] and on the type of specimen used (single or polycrystal). The hydride phase can be identified by X-ray techniques, if the H/Ni atomic ratio exceeds 0.05 [15]. Considering the limits at which crystal phases may be detected by X-ray techniques, the β phase may in fact be present at lower H/Ni ratios.

It should also be mentioned that the hydride phase can be obtained by bom-barding nickel with protons [16] or by saturating it with hydrogen from the gaseous phase under high pressures [17]. When nickel is charged with hydrogen from the gas phases [17], it is difficult to determine the influence of pressure on the distribution and density of the lattice defects (e.g. dislocations), though effect is not negligible [18]. The particular difficulty in the non-electrolytic formation of the hydride phase is that specimens cannot be made large enough to permit structural and mechanical investigations. For this reason we shall confine our-selves to studies related to nickel hydride obtained by electrolytic charging in the following discussion.

Nickel hydride has the same type of Bravais lattice as nickel, the lattice par-ameter being about 6% larger than that of the pure metal [13]. A neutron diffrac-tion study has shown hydrogen in nickel hydride to occupy octahedral sites [19]. Nickel hydride therefore has an NaCl type structure.

The atomic ratios of hydrogen to nickel in nickel hydride, as determined by volumetric measurements, range from 0.6 to 0.8 [12]. This suggests the forma-tion of a non-stoichiometric nickel hydride. Experimental observations show however that in polycrystalline nickel some grains fail to form any hydride phase at all [20]. It is therefore more likely that there are some regions-where a stoichiometric or close-to-stoichiometric nickel hydride is formed, along with regions that are hardly or not at all charged with hydrogen. Under standard conditions (room temperature, atmospheric pressure) nickel hydride is an unstable compound which starts to decompose as soon as the charging process is stopped [21]. Decomposition is accompanied by desorption of hydrogen from the specimen. The lifetime of nickel hydride, i.e. the time for which the hydride

3 2 4 CHANGE IN THE REAL STRUCTURE OF NICKEL

phase can still be detected by X-ray techniques, was found to depend on the type of specimen used. It is longest for single crystals, and shortest for nickel layers deposited electrolytically [22].

In addition to nickel hydride nickel deuteride is also known. It may be ob-tained by electorlytic charging of nickel with deuterium in a heavy water sol-ution of D 2S 0 4 [23]. The lattice parameter of β-NÏD is somewhat smaller than the corresponding value for NiH. The values of lattice parameters for a- and /?-nickel hydrides and deuterides compared with the lattice parameter of the initial nickel specimen are [24] :

Ni 0.35213+0.00005 nm a-NiH 0.35218 ± 0.00005 nm a-NiD 0.35219+0.00004 nm /J-HiH 0.37251+0.00008,nm β-NiD 0.37191+0.00008 nm

In view of the insignificant differences in the lattice parameters of nickel hydride and nickel deuteride it can be assumed that the results obtained for the changes in the real structure of nickel hydride also hold good for nickel deuteride.

113 Micro-cracking in the surface layer

X-ray studies of a nickel single crystal by the Bragg method [25] show the broad-ening of individual lines of the diffraction pattern, when nickel is transformed to nickel hydride and subsequently decomposed. This broadening became more pronounced with increasing numbers of transitions to the hydride phase culminating in an X-ray diffraction pattern characteristic of a polycrystal. This fact suggests that decomposition of the hydride phase is accompanied by a désintégration of the specimens surface layer. The disintegration becomes more pronounced as the number of succesive phase transformations increases. This process is reflected in the formation of micro-cracks on the specimen surface, which can be seen under a metallographic microscope. The cracks make their appearance up to ten minutes or so after the charging process is stopped. For polycrystalline specimens the cracks are both inter- and transgranular in char-acter. The intergranular micro-cracks are not essentially different from those sometimes observed in specimens charged with hydrogen to average concentra-tions corresponding to the presence of the α-phase. On oriented surfaces of single crystals the cracks run in definite crystallographic directions, and the crack lines are parallel to the intersection of {100} planes with the sample surface [26].

Fig. 11.1. Micro-cracks on the nickel surface produced during decomposition of the hydride phase (after [27]); (a) polycrystalline specimen: inter- and transgranular cracks; (b) micro-cracks formed in the region of a non-metallic inclusion.

P.N.

L U J ι ι

8 10 12 20 30

Indentation diagonal,

Fig. 11.2. Microhardness, μΗΥ, of a nickel single crystal prior to and after charging with hydrogen vs. indentation diagonal length (after [26]); solid line—before Η charging, broken line—after the charging, Δ— (HO) orientation, O—(100) orientation.

326 CHANGE IN THE REAL STRUCTURE OF NICKEL

Two kinds of cracks may be distinguished: fine cracks appearing about ten minutes after the cessation of charging, and larger ones that show up much later. As shown by the scanning electron microscope [27] (Fig. 11.1), the large cracks appear at points of high stress concentration and at the sites of non-metallic inclusions. Since {100} planes are not the slip planes in f.c.c. metals, the cracking observed along these planes was unexpected [26].

Scanning electron microscope studies of specimens charged with hydrogen and strained to the point of failure [27] confirmed the earlier observations [26] (Fig. 11.2) on the hardness and brittleness of the hydride phase, and they showed the surface hydride layer to be of limited thickness. As seen from Fig. 11.3, a

Fig. 11.3. A scanning electron micrograph of a nickel single crystal of surface orientation (110) charged with hydrogen for 48 h and strained to the point of failure (after [27]).

plastic nickel layer underlies the cracked brittle hydride surface layer, whose thickness was estimated here to be approx. 17 μπι. An intermediate layer is also visible in the photograph. Repeated conversion of a thin nickel specimen to its hydride may result in its complete destruction. Thus the formation and decomposition of nickel hydride causes permanent changes in the specimens structure.

11.4 Desorption of hydrogen and decomposition of nickel hydride

As mentioned before the hydride phase decomposition observed by X-ray techniques is accompanied by evolution of hydrogen from the specimen. Volumet-ric measurements have shown the hydrogen desorption from polycrystalline

DESORPTION OF HYDROGEN 327

specimens to follow the first order reaction kinetics only in samples that were repeatedly charged with hydrogen ("cycled") [28]. For these the rate constant for hydrogen desorption was equal to the rate constant for decomposition of the hydride phase, as determined by X-ray analysis [29] (Fig. 11.4).

1

. !

2 j CP

ο

0 10 20 30 AO 50 60 Time, min

Fig. 11.4. The kinetics of decomposition of cycled nickel hydride (after [29]); 1—results of vol-umetric study, 2—result of X-ray study; Km ax—total volume of hydrogen desorbed, Ft—vol-ume of hydrogen desorbed after time t, measured from the start of the process, /—diffraction line intensity, in arbitrary units.

Later studies proved however that, independent of the kind of specimen and of the number of chargings, hydride phase decomposition as determined by X-ray analysis, has first-order reaction kinetics [30, 31].

The kinetics of hydrogen desorption as determined by the volumetric method vary with the number of the nickel-hydrogen-nickel transitions, that is with cycling of the specimen. This means that, with incerasing cycling, the factor controlling the hydrogen desorption kinetics must vary. From the discrepancy observed, during the first chargings, between the kinetics of hydride phase de-composition as determined by X-ray analysis and the kinetics of hydrogen de-sorption as determined volumetrically, it follows that this factor is not the rate of actual hydride decomposition but the rate of removal to the atmosphere of the hydrogen released during /?-phase decomposition. This question will be more fully considered in the next section on the role of dislocations in the process of hydrogen charging and hydride decomposition.


11.5 The role of dislocations in the formation and decomposition of the hydride phase

Each successive conversion of nickel to the hydride phase, followed by its de-composition, causes an increase in the surface area of the specimen, as a con-sequence of the increased number and size of cracks.

To determine the effect of dislocations initially present in the nickel specimen on the hydride decomposition process and on hydrogen desorption kinetics, the copper decoration method was used [32], This method of revealing sites of dislocation emergence on a metal surface consists of the electrolytic deposition of copper on nickel from 10~

5 to 10"

6 M CuS04 solution in

0.5 M H 2S 0 4. At a current density of 2 χ 10~6A/m

2, copper is selectively de-

posited and decorates the sites of emergence of dislocations or their accumula-tions on the nickel surface. At these sites deposited copper alters the conditions for hydrogen diffusion from the bulk of specimen. The diffusivity of hydrogen in copper is comparable with the corresponding value for nickel, but is much lower than that along dislocations. As a result, the copper selectively deposited at the sites of dislocation emergence on the metal surface slows the removal of hydrogen from the dislocation regions. Blocking these dislocation emergence

Fig. 11.5. The kinetics of hydrogen desorption from polycrystalline specimens at 343 Κ (after [33]); 7, 3— specimens with blocked sites of dislocation emergence; 2, 4—with unblocked sites; curves3and 4 correspond to specimens in which the hydride phase was not observed (Δρ—pressure variations in the apparatus, in arbitrary units).

σ><|

_ j l ι ι ι ι 1 1 L-i

20 40 60 80 100 120 140 160 180 Time, min

THE ROLE OF DISLOCATIONS IN FORMATION OF HYDRIDE PHASE 329

sites reduces the contribution of dislocations to the removal of hydrogen from the sample during decomposition of the hydride phase.

Nickel hydride obtained in specimens with blocked dislocations has prop-erties quite different from those discussed previously. In the first place, its lifetime is much longer. For single crystal specimens the hydride may last for one month [31] or even longer, whereas in the specimens with unblocked dislo-cations the life is only a few days. The kinetics of decomposition of nickel hydride in dislocation-blocked specimens, as determined by X-ray analysis, still follows the first-order kinetics; but its rate constant is considerably lower. The kinetics of hydrogen desorption is different, however, as compared with that for the dislocation-blocked specimens. The relation expressed by curve 1 in Fig. 11.5 is linear with the chosen system of coordinates, which signifies that the hydrogen desorption process follows first-order kinetics. Curve 2 corresponds to hydrogen desorption from specimens with unblocked dislocations. Curve 4 in Fig. 11.5 represent the kinetics of hydrogen desorption from a specimen in which β phase was not observed by X-ray examination, so that hydrogen may be assumed to have been present as the α phase. Curve 3 in this figure corresponds to curve / for the dislocation-blocked specimen. Comparison of curves 5, 4 and 1 shows that they are of the same character. The blocking of dislocations in the nickel must therefore result in lowering of the rate of hydride decomposition, while hydrogen desorption follows the same mechanism as desorption from the α-phase. Metallografic studies of specimens with blocked sites of dislocation emergence showed that hydride phase decomposition was predominantly ac-companied by the apperance of large cracks.

The conclusion resulting from these experiments, combined with other supple-mentary ones, made it possible to advance the following dislocation model to explain the character of the effects observed [33]. During the first electrolytic charging with hydrogen, nickel hydride is originally formed at some sites in the surface layer, most probably in the defect regions, producing nuclei of a new phase. As a result of the formation of nuclei with a lattice parameter consider-ably different from that of the matrix, a large stress gradient appears at the in-terface. The stresses are responsible for displacement and multiplication of dis-locations in the stress interaction regions, thereby facilitating a further forma-tion of the hydride phase in the deformed areas of specimen.

Dislocations may assist the penetration of hydrogen into the bulk of the specimen not merely by diffusion along dislocation lines. An additional kinetic model for hydrogen transport has been proposed, invoking Cottrell atmospheres;


and this makes plausible a much faster movement of hydrogen than that of regular diffusion through the undisturbed nickel lattice [34].

The processes that occur in nickel durnng electrolytic charging with hydrogen may be visualized as follows. Hydride phase nuclei form and cause strong stresses within the specimen, with a resultant change in dislocation structure. Further charging raises the hydrogen concentration in the metal and causes the hydride phase to spread within the specimen. This change in the dislocation structure cannot however proceed indefinitely, and its last stage is the formation of a dislocation cell structure.

Such a structure is a characteristic feature of deformed metals. It represents the most favourable dislocation arrangement with a minimum energy per unit volume [35]. In a cell structure most of the dislocations accumulate along bound-aries which separate areas of a very low dislocation density. The appearance of a cell structure with dislocation walls parallel to {100} planes was observed un-der the electron microscope in nickel specimens deformed in tension [36], the cell structure forming at as low a deformation as 3% [37].

In the present case the process of dislocation interaction and multiplication starts in the surface layer in the regions neighbouring the hydride phase nuclei and spreads into the bulk of specimen during hydrogen charging, that is with the quantity of the hydride phase. Thus the formation of the dislocation cell structure will be possible only in regions of finite thickness, which depend pri-marily on the magnitude of the stress field at the interface. Much higher stresses than those present in the material will be required to change the dislocation pat-tern of the cell structure. The formation of a dislocations cell structure therefore, will limit the possibility of a further propagation of the hydride phase, as the process of further multiplication of dislocations and their displacement into the bulk of specimen will be more difficult. During nickel hydride decomposi-tion the hydrogen liberated diffuses both through the dislocation-free regions and through dislocations. As a result of a difference in diffusion rate through the dislocation-free metal and along dislocations, a situation may arise where all the hydrogen from the dislocation stress field is removed; this will result in the formation of regions of pure nickel, inside the hydride phase. Consequently, tensile stress will be set up, which will favour the development of the micro-cracks observed at the sites of emergence of dislocations, or their systems on the surface. These stresses also account for the fact that the cracks are observed some time after the cessation of hydrogen charging process. The blocking of the sites of emergence of dislocations is equivalent to arresting the dislocation mechanism of hydrogen removal from the specimen, so that the hydrogen de-

THE ROLE OF DISLOCATIONS IN FORMATION OF HYDRIDE PHASE 331

sorption would be diffusion controlled through undisturbed regions of the crys-tal lattice. The character of the kinetics of hydrogen desorption from specimens with blocked dislocations, agrees with this explanation.

The model put forward here also makes sense of the previously discussed fact (cf. Fig. 11.4) that reproducible result from the volumetric hydrogen desorp-tion kinetics, which agree with the hydride phase decomposition kinetics found by the X-ray method, are only obtained in cycled specimens [29]. In such speci-mens, the number of defects is so large than in practice the whole of the hydro-gen may be desorbed from them. The role of diffusion through the undisturbed crystal lattice of nickel then becomes negligible.

The nickel-hydride-nickel phase transition affects the real structure of a single crystal (or crystallites in polycrystalline material). The ease of formation of the hydride phase however depends also on the degree of strain of the nickel. In annealed specimens the dislocation density is about 10

10 to 1 0

u dislocations/m

2.

This value drastically increases with hydrogen charging of specimens up to concentrations corresponding to the β phase: in the cases studied [38] the den-sity was estimated at approx. 10

14 dislocations/m

2.

The cycling of a thin polycrystalline nickel specimen (nickel foil), in which the hydride is liable to form throughout the entire bulk, leads to disintegration of the specimen, as a result of grain disintegration and intergranualar cracking.

X-ray studies [38] showed that the initial hydrogen charging produces the largest change in the real structure of nickel resulting in a dramatic change in the size of the areas of coherent scattering (of the size of crystallites).

An interesting finding is that the size of crystallites evaluated in the <001> and < 111> directions decreases during the first and second charging from approx. 1000 nm to 30-40 nm, hence remaining nearly constant during the subsequent cycles, reaching a limiting value of approx. 20 to 30 nm. X-ray measurements [38] have also shown the lack of stress in the <001> direction and the presence of stress in the <111> direction. Although the X-ray diffraction measurements were carried out in a much shorter time than the lifetime of nickel hydride in the specimens examined, the stress relaxation corresponded to the time when the first micro-cracks on the specimen surface were observable under a metallographic mic-roscope. Since it is unlikely that such an anisotropy of stresses is involved at the moment of cessation of the charging, it is concluded that as soon as the hydride phase starts to decompose, the stresses along the {100} planes relax. This re-laxation may be due to submicro-cracks formation. It is the development of these submicro-cracks that makes the micro-cracks metallographically observable. It should be noted that the terms used here such as submicro-cracks, micro-


cracks, fine cracks and large cracks describe definite stages in the process of crack formation, that is, nickel disintegration process as resulting from the re-laxation of stresses of various magnitudes. The results are compatible with the previous statement that the micro-cracks form along the {100} planes [26]. This in turn gives support to the inference of the formation of a cellular structure of dislocations along these planes.

11.6 Structural properties of nickel compared with palladium hydride

Nickel hydride is structurally an analogue of palladium hydride, so that the prop-erties of the two systems are often compared. The lattice parameter of palla-dium hydride is about 3.5% larger than that of pure palladium [39], whereas for nickel hydride the corresponding value is about 6% larger [13, 24]. The hy-drogen in palladium hydride, and in its nickel analogue, occupies octahedral sites [40]. It has been pointed out that the much higher lattice parameter differ-ence for nickel hydride may be explained by the difference in size of the octa-hedral sites in nickel and palladium. This geometrical factor has been held responsible for the more difficult preparation of nickel hydride, compared with palladium hydride [13].

The difference in the extent of lattice dilatation has a substantial effect on the change in the real structure of nickel and palladium subjected to hydrogen charg-ing.

The increased lattice parameter of 6% in the case of nickel hydride formation results in stresses being formed in the boundary regions, which substantially exceed the yield stress, and which correspond to stresses of stage II or III on the stress-strain curve. During the charging of nickel with hydrogen, therefore, no slip lines are observed, unlike the analogous process for palladium [41].

In the case of palladium hydride the effect of an increased lattice parameter is much less; hence, the stresses and strain at the palladium hydride-palladium interface are much lower. Also, the change of dislocation density is slower. Under-standably the formation of a cellular structure of dislocations will be more diffi-cult in this case and such a structure will be formed only in some regions of speci-men. Although the cracks in palladium also grow in the {100 } planes during hydride decomposition [42], they are are less observable and only appear at a few points. In the case of a palladium single crystal, therefore, the accumulation of dislo-cations along the slip lines caused by the initial stage of hydrogen charging can be observed [43] using the Berg-Barrett technique of X-ray reflection topo-

EFFECTS OF HYDRIDE FORMATION ON STRUCTURE OF NICKEL 333

graphy. For nickel, as mentioned previously, electrolytic charging with hydro-gen under similar conditions results in a vastly greater change in the arrange-ment of dislocations [44].

The formation of the cellular structure in nickel is a factor that controls the depth to which the hydride phase is formed. In the case of palladium this factor is insignificant, and consequently a limiting depth of the hydride phase formation is not involved. If the metal-hydride phase transition were accompanied by a negli-gible change in the value of lattice parameter, and hence by a minor lattice dis-tortion, the change in dislocation density in this process should also be insigni-ficant. Indeed, when charging with hydrogen a Pd-20% Ag alloy, for which the difference in lattice parameters of the hydride and of the alloy is only 1.6%, no change in dislocation density was found by the positron annihilation method [45]. An electron microscopic study also showed that, in Pd-Ag alloys charged with hydrogen, the dislocation density is lower than in pure Pd charged with the gas [46].

11.7 Comparison of the effects of hydride formation on the structure of nickel and Cr-Ni austenitic steels

The chromium-nickel steels, with an austenitic (f.c.c.) structure also undergo hydrogen-induced phase transition processes which result in the formation of hydride phase and in some changes in the real structure of the alloy. In austenitic steels however the processes are naturally more complex as compared with nickel due to their alloy nature.

X-ray studies showed the formation of three alloy hydride phases in these steels as a result of a cathodic charging with hydrogen [47-52]:

1. The H y phase [47-50], also known as y* [51] or Yt [52], with a face-centred cubic (f.c.c.) lattice, i.e. with the same Bravais lattice as in austenite but with a 5% larger lattice parameter. The phase is entirely analogous to nickel hydride.

2. The Hy phase [47-50], also known as Y2 [52], with a face-centred cubic (f.c.c.) lattice like the H y phase but with a lattice parameter about 1.25% larger than in austenite.

3. The He phase [47-50], also known as ε' [51] or ε* [52], with a hexagonal close-packed (h.c.p.) lattice, with the same Bravais lattice as martensite ε, but with a lattice parameter about 5% larger.


The phases, as with β nickel hydride, are formed with a thickness of the order of a few μιη in the surface layer on charging the alloy with hydrogen. They are unstable at room temperature and atmospheric pressure, and they start to de-compose as soon as the charging process is stopped. At low temperatures (be-low 273 K) the phases are much more stable [47, 49, 51].

The formation and decomposition of the hydride phases result in some ir-reversible effects and some changes in the structure of the austenitic Cr-Ni steels. On electrolytic charging with hydrogen, followed by desorption, the X-ray diffraction lines of austenite are strongly broadened and somewhat shifted with respect to their initial 20 Bragg angles. The broadening is due to the formation of finer crystallites of austenite, and/or lattice distortions. The linear shifting, on the other hand, is connected with a change in the lattice parameter resulting from the solid solution of the hydrogen in the austenite [50]. Hence there is a marked analogy with the effect of hydrogen on nickel.

In type 310 austenitic steels (i.e. containing about 25% Cr and 20% Ni) which feature "stable austenite", ε martensite is formed as a result of the austenite -> alloy hydrides phase process. After some ageing at room temperature the ε martensite almost completely disappears [47,49, 53]. In turn, in type 304 steels (i.e. those containing approx. 18% Cr and 8% Ni), in which austenite is me-tastable, both α martensite and ε martensite are produced [49-54] and remain stable at room temperature.

The alloy hydrides in austenitic steels, as with nickel hydride, have a lattice parameter that differs significantly from the corresponding value in the original material with increases up to 16% in the volume of the unit cell. As a result, considerable stress gradients and plastic deformation occur. The stresses bring about the generation and movement of dislocations which significantly increases their density.

The alloy hydrides produced during cathodic charging yield a hard and very brittle surface layer in austenitic steels [47, 49, 55, 56]. This results in the devel-opment in the surface layer of multiple micro-cracks, which are formed during hydride decomposition and hydrogen desorption processes, and their appearance is connected with the high tensile stresses brought about by the decreasing value of the lattice parameter. In contrast with nickel the cracks run along the {111} and {110} planes. A possible reason for these differences lies in the difference in structual changes caused by hydrogen in austenitic steels in which, unlike nickel, three hydride phases are formed, and additionally martensitic transi-tions are involved [57-62].

REFERENCES 335

11.8 Summary

1. The hydride phase nuclei are primarily formed in the local, distorted regions of the crystal lattice of nickel.

2. The formation of hydride phase nuclei results in strong stresses in the surface layer which affect the dislocation structure of the specimen. The final stage is the formation of a cellular structure with walls parallel to the {100} planes. This is equivalent to a rise in dislocation density in specimens during hydrogen charging. The cellular structure formed limits propagation of the hydride phase into the bulk of metal.

3. The electrolytic charging of nickel with hydrogen result in inhomogeneous formation of the hydride in the surface layer, primarily in the regions of disloca-tion stress field.

4. The nickel hydride produced is hard and brittle. 5. Micro-cracks are formed along the {100} planes during the hydride phase

decomposition. These result from tensile stresses caused by a faster removal of hydrogen from the dislocation regions.

6. In polycrystalline specimens, as a result of multiple transformations of nickel to the hydride, the crystallites are reduced in size (the coherent scattering zones become smaller) down to a limiting value of approx. 20-30 nm.

7. The effect of the hydride formation on the structure of Cr-Ni austenitic steels is different to those in nickel. This is caused by the formation of three hydride phases and a martensite transformation involved.

One of the authors (A. J.) acknowledges gratefully the contribution of Dr A. Szummer in preparation of Section 11.7.

11.9 References

[1] Blakemore, J. S., Met. Trans., 1, 145 (1970). [2] Blakemore, J. S., Met. Trans., 1, 151 (1970). [3] Boniszewski, T., Smith, G. C , Chem. Phys. Sol, 21, 115 (1961). [4] Boniszewski, T., Smith, G. C , Acta Met., 11, 165 (1963). [5] Lauthan, M. R., Jr., Donovan, J. Α., Caskey, G. R., Jr., Acta Met., 23, 745 (1975). [6] Windle, A. H., Smith, G. C , Metal Science Journal, 2, 187 (1968). [7] Windle, A. H., Smith, G. C , Metal Science Journal, 4, 136 (1970). [8] Smith, G. C , Hydrogen in Metals, Am. Soc. for Metals, 1973. [9] Smith, D. P., Hydrogen in Metals, Chicago Press, 1948.

[10] Combette, P., Grilhe, J., Communication présenté aux Journée d'automne, Paris 1969. [11] Baranowski, B., Szklarska-Smialowska, Z., Smialowski, M., Bull Acad. Polon. Sci., sér.

sci. chim. géol. et géogr., 6, 179 (1958).


[12] Baranowski, B., Smialowski, M., Bull. Acad. Polon. Sci., sér. sci. chim. géol. et géogr., 7, 663 (1959); J. Phys. Chem. Solids, 12, 206 (1959).

[13] Janko, Α., Bull. Acad. Polon. Sci., sér. sci. chim., 8, 131 (1960); Naturwiss., 47, 225 (1960). [14] Jarmolowicz, H., Smialowski, M., / . Catalysis, 1, 165 (1962). [15] Majchrzak, S., Bull. Acad. Polon. Sci., sér. sci. chim., IS, 485 (1967). [16] Palczewska, W., Bull. Acad. Polon. Sci., sér. sci. chim., 12, 817 (1964). [17] Baranowski, B., Wisniewski, R., Bull. Acad. Polon. Sci. sér. sci. chim., 14, 273 (1966). [18] Mashushchenko, L. Α., Shmatov, V. T., Beresev, Β. I., Fizika Metallov i Metallovedenye,

41, 1290 (1976). [19] Wollan, E. O., Cable, J.W., Koehler, W. C , J. Phys. Chem. Solids, 24, 1141 (1963). [20] Bovier, C , Janko, Α., Michel, P., Electron Microscopy 1968, Proc. Eur. Reg. Conf. 4th,

Roma 1968. [21] Baranowski, B., Bull. Acad. Polon. ScL, sér. sci. chim. géol. et géogr., 7, 891 (1959). [22] Janko, Α., Thesis IChF PAN, 1965. [23] Stroka, Α., Baranowski, B., Bull. Acad. Polon. ScL, sér. sci. chim., 10, 147 (1962). [24] Janko, Α., Pielaszek, J., Bull. Acad. Polon. ScL, sér. sci. chim., 15, 569 (1967). [25] Janko, Α., Bull. Acad. Polon. ScL, sér. sci. chim., 10, 613 (1962). [26] Janko, Α., Szummer, Α., Bull. Acad. Polon. ScL, sér. sci. chim., 14, 885, 1966. [27] Szummer, Α., Janko, Α., Pielaszek, J., Bull. Acad. Polon. Sci sér. sci. chim, 19, 33 (1971). [28] Baranowski, B., Bull. Acad. Polon., Sci., sér. sci. chim. géol. et géogr., 7, 891 (1959). [29] Janko, Α., Bull. Acad. Polon. ScL, sér. sci. chim., 10, 617 (1962). [30] Pielaszek, J., Bull. Acad. Polon. ScL sér. sci. chim., 20, 487 (1972). [31] Pielaszek, J. Ph.D. Thesis, Inst. Chem., Polish Academy of| Sciences, Warszawa 1972. [32] Pielaszek, J., Phys. Stat. Sol., 10, Κ 145 (1972). [33] Pielaszek, J., Bull. Acad. Polon. ScL, sér. sci. chim., 20, 611 (1972). [34] Tieu, J. K., Thompson, A. W., Bernstein, J. M., Richards, R. J., Met. Trans., A7, 821

(1976). [35] Bassim, M. N., Kuhlmann-Wilsdorf, D., Cryst. Lattice Defects, 4, 9 (1973). [36] Hirsch, P. B., Internal Stresses and Fatigue in Metals, Elsevier, New York 1958. [37] Spruiel, C. R., du Bose, K., Mat. Res. Bull., 3, 381 (1968). [38] Janko, Α., Pielaszek,'J., Polukarov, Yu. M., Semenova, Ζ. V., Elektrochimiya, 13, 878

(1977). [39] Ubbelohde, A. R., Proc. Roy. Soc, A159, 174 (1937). [40] Worsham, J. E., Willinson, M. K., Shull, C. G., Phys. Chem. Sol., 3, 303 (1957). [41] Tidema, T. J., Kocoy, C , Burgers, W. G., Proc. Kon. Ned. Akad. Weiensch, Ser. B, 62,

34 (1959). [42] Janko, Α., Pielaszek, J., Szummer, Α., Bull. Acad. Polon. Sci. sér. sci. chim., 22, 959 (1974).

[43] Pielaszek, J., unpublished work.

[44] Pielaszek, J., Phys. Stat. Sol., 31, Κ 95 (1975). [45] Dlubek, G., Brumer, O., private communication. [46] Wise, M. L. H., Farr, J. P. G., Harris, I. R., Hirst, J. R., Congr. Intern. Vhydrogène dans

les métaux, Paris 1972. [47] Szummer, A. Janko, Α., Corrosion, 35, 461 (1979).

[48] Szummer, Α., Janko, Α., XI International Congress of Crystallography, Collected Abstracts,

06.3-7, 178, Warszawa 1978.

REFERENCES 337

[49] Szummer, Α., Janko, Α., Proc. IX Confer, on Applied Crystallography, Kozubnik, Poland,

August 13-16 1978, Katowice 1979, Vol. II, p. 691. [50] Szummer, Α., Janko, Α., Acta Crystallographica, A 34, Part 54, 178 (1978). [51] Kamachi, K., Transaction ISIJ, 18, 485 (1978). [52] Narita, N., Altsletter, C. J., Birnbaum, Η. K., Metallurgical Trans. Α., 13A, 1355 (1982). [53] Okada, K., Hosoi, Y., Abe, S., Corrosion, 26, 183 (1970). [54] Holzworth, M. L., Louthan, M. R., Jr., Corrosion, 24, 110 (1968). [55] Szummer, Α., Bull. Acad. Polon. Sci., sér. sci. chim., 12, 651 (1964). [56] Smialowski, M., Szummer, Α., Metallurgia Italiana, 47, 144 (1965).

[57] Szummer, Α., Janko, Α., Proc. X Confer, on Applied Crystallography, Kozubnik, Eds.

Z. Bojarski, T. Bold, 1980, p. 452. [58] Szummer, Α., Janko, Α., Acta Crystallographica, A 37, Supplement, C-101 (1981). [59] Szummer, Α., Acta Crystallographica, A 40, Supplement, C-123 (1984). [60] Szummer, Α., Janko, Α., Proc. XI. Confer, on Applied Crystallography, Kozubnik, Uni-

wersytet Sl^ski, Katowice, 1, 92 (1984). [61] Szummer, Α., in: Hydrogen Degradation of Ferrous Alloys, Eds. R. A. Oriani, J. P. Hirth,

M. Smialowski, Noyes Publication Inc. Park Ridge, New York 1985, p. 512. [62] Szummer, Α., Slawinski, Α., Stybel, T., Janko, Α., Proc. XII Confer, on Applied Crystallo-

graphy, Cieszyn, Eds. Z. Bojarski, T. Bold, 1986, p. 50.

Added in proof 1. Flanagan, T. B., Bowerman, B. S., Biehl, G. E., Scripta Met., 14,443 (1980). 2. Borbe, P. Ch., Erdmann-Jenitzer, F., Schoebel, W., Z. Fur Metallkunde, 71, 227 (1980). 3. Mclnteer, W. Α., Thompson, A. W., Bernstein, I. M., Acta Metallurgica, 28, 887 (1980). 4. Ashok, S., Duquette, D . J., Stolloff, N. S., Verpoort, C , Scripta Metallurgica, 15,1329 (1981).

5. Brossard, L., Thome, L., Traverse, Α., Bernas, H., Chaumont, J., Lalu, F.,Phys. Stat. Sol. (a),

68, 619 (1981). 6. Flanagan, T. B., Clewley, J. D., / . Less-Common Metals, 83, 127 (1982). * 7. Flanagan, T. B., Oates, W. Α., Kishimoto, S., Scripta Met., 16, 293 (1982). 8. Tsuru, T., Latanision, R. M., Scripta Met., 16, 575 (1982). 9. McLelîan, R. B., Sutter, P. L., Acta Metallurgica, 32,2233 (1984). 10. Lengeler, B., Zeller, R., Solid State Communication, 51, 889 (1984). 11. Birnbaum, H. K., / . Less-Common Metals, 104, 31 (1984). 12. Pielaszek, J., „Nickel hydride", in : Hydrogen Embrittlement of Ferrous Alloys, Eds. R. A.

Oriani, J. P. Hirth and M. Smialowski, Noyes Publ. N. Y., USA, Ch. 8. 13. Frankel, G, S,, Latanision R. M., Metallurgical Transactions A, 17A, 861 (1986). 14. Frankel, G. S., Latanision, R. M., ibid., 17A, 869 (1986).

12. Ionization of Hydrogen on the Surface of Metals in Aqueous Solutions of Electrolytes

12.1 Introduction

The problems related to hydrogen entry into a metal from the gas phase and from aqueous solutions of electrolytes have been dealt with in Chapters 7 and 8. In the first case the initial form of hydrogen was either molecular hydrogen H 2 or atomic hydrogen H present in the gas phase, in the second case it was the hydrogen ion H

+ (in acid solutions) or else, in alkaline solutions, hydrogen

bound in the water molecule H 20 .

In the ionization of hydrogen on the surface of metals, which is the subject of this chapter, the initial hydrogen form is molecular hydrogen dissolved in the electrolyte, that is, hydrogen that has passed into the electrolyte from the gas

Fig. 12.1. A schematic diagram showing various possible ways in which hydrogen passes through the interfaces in the metal-gas-electrolyte system: i—hydrogen entering from the gas phase to the metal, 2—from the electrolyte to the metal (cathodic saturation or charging), 3— from the gas phase to the electrolyte (hydrogen ionization).

INTRODUCTION 339

phase: this then yields the hydrogen ion in the electrolyte; in alkaline electro-lytes the water produced as a result of a reaction of this ion with OH"" ions may be considered as the product of this process. All three processes, gas phase entry, electrolyte entry, and ionization, are shown in Fig. 12.L

The hydrogen ionization process is the reverse of the process of evolution of hydrogen from an electrolyte:

The symbol Me denotes that the process occurs on the surface of an electrode metal, which acts as a catalyst. Hence this is a heterogeneous reaction. The left to the right reaction represents the anodic oxidation (ionization) of hydrogen and is sometimes referred to as electrochemical combustion. The reverse reaction represents the cathodic reduction (evolution) of hydrogen. Investigation of this reaction has greatly contributed to the development of the fundamentals of the modern electrochemistry; moreover, it is of considerable significance in the majority of corrosion processes, so that the reaction continues to be the sub-ject of an extensive research effort, despite the abundance of the literature already devoted to this field.

The hydrogen ionization reaction has not received the same attention in electrochemistry and related branches of physical chemistry as the reverse pro-cess of cathodic evolution of hydrogen. This is because the ionization reaction is much more difficult to study, largely because of the low solubility of hydrogen in aqueous solutions; another reason is that at potentials, at which it can be studied, the secondary processes which proceed on the majority of metals disturb the stationary conditions (e.g. oxide film formation or corrosion of metal). Exceptions to this are the platinum group metals, to some extent the iron group metals, and also the recently used non-stoichiometric compounds of the tran-sition metals, among which tungsten carbide is the major representative. In the platinum group metals the majority of experimental data concern platinum itself, which is the most effective catalyst of hydrogen electrode reactions known. Based on the results obtained for platinum, detailed electrochemical theories have been developed for this reaction, in particular the thermodynamic double layer theory for the reversible hydrogen electrode due to Frumkin and his co-workers [1].

The anodic reaction of hydrogen on platinum is particularly useful for the investigation of the thermodynamics and kinetics of adsorption from electrolyte solutions, the results of which can be also compared with the results of studies

340 IONIZATION OF HYDROGEN ON THE SURFACE OF METALS

on hydrogen adsorption from the gas phase. One important example has been the development of the Temkin isotherm [2], which accounts for the change in adsorption energy with coverage.

12.2 Practical significance of the hydrogen ionization reaction

The theoretical importance of the hydrogen ionization reaction has been treated briefly above. Early in the sixties, extensive research effort began on the devel-opment of fuel cells. It was found, however, that reliable data on the rate and mechanism of the hydrogen ionization process were not available even for the "neatest" case of platinum in acid or alkaline solutions, let alone nickel in alka-line solutions where the process of interest is invariably accompanied by the competitive reaction of the formation of the oxide film.

The continuing technological interest in the hydrogen ionization reaction stems from the fact that hydrogen constitutes an almost ideal "fuel" for the fuel cell. The fuel cell is a specific kind of a primary cell with a continuous supply of reactants. The fuel, i.e. hydrogen, is supplied to the anode, while an oxidan is fed to the cathode, the latter being oxygen from the air in the simplest case. When the terminals of such a cell are shortened, the reaction proceeding on the anode, in an acid solution, can be written as

The normal potential of reaction (12.2) is about 1.23 V R HE (RHE—reversible hydrogen electrode scale of potentials), and this is also the theoretical value of the voltage of an open-circuit hydrogen-oxygen cell under standard condi-tions [3].

The oxidation of one mole of hydrogen should give, under reversible condi-tions, electric energy in a quantity

W = - AG = 2FÂE - 2 χ 1.23 V χ 96500 C/mol = 237 kJ/mol (12.3)

which is equivalent to approx. 119 kJ per gramme of fuel. This very high energy density makes hydrogen an attractive fuel, especially in fuel cells.

Additional advantages of hydrogen oxidation as the electrode process are: the fast rate of the reaction, which greatly exceeds the electrode oxidation rate for other fuels, the lack of environmental contaminants among the reaction prod-ucts, and the fact that the water formation in the reaction is a significant factor

H 2 ( B) -* 2 H+ + 2e

and on the cathode the oxygen reduction is given by the equation

2 H+ + i 0 2 + 2e-+ H 20

(12.1)

(12.2)

SIGNIFICANCE OF HYDROGEN IONIZATION REACTION 341

in space exploration involving manned vehicles. Hydrogen is one of the most abundant elements on the Earth; so there is no danger of it ever being in short supply. Further practical advantages of the hydrogen-oxygen (air) fuel cell are: its noiseless operation, suitability for work at roughly ambient temperatures and the lack of moving mechanical parts subject to wear.

One major advantage of fuel cells in general is their high energy efficiency, compared with traditional electric power generators which operate on the heat cycle principle, i.e. with fuel combustion in thermal engines with all the ensuing thermodynamic limitations. For classical thermal energy sources an energy W can theoretically be obtained from one mole of fuel, thus in Carnot's cycle:

W = AH T l

~Tl

(12.4)

where T2, Tx are the temperatures of the heat source and of the sink, respectively. This value is normally very much smaller than the energy obtained theoretically from fuel cells and defined by formula (12.3). So far this general advantage of fuel cells is of little practical significance, since the electrode polarization pro-cesses in fuel cells, apart from those containing platinum-group catalysts, res-trict their energy efficiency to a considerable extent.

All these advantages of hydrogen as a fuel in fuel cells are responsible for the extensive and renewed interest in the hydrogen oxidation reaction. A research and development effort, involving considerable outlay, was conducted in many countries, aimed at constructing a cost-effective fuel cell suitable for daily use. A great number of prototypes and production cells have been made, but chiefly for special purposes that have little to do with the large-scale production of electric power.

There are two reasons why the successful use of the anodic hydrogen ioniza-tion for power generation has not yet been achieved: firstly, the unavailability of a suitable (inexpensive, efficient and stable) electrocatalyst, i.e. an electrode material that would catalyze the hydrogen ionization reaction, and secondly the continuing high price of hydrogen.

The first problem very probably will be solved by the development of new catalysts with the above-mentioned properties or by minimization of the plati-num amount used in fuel cell anode. This is the subject of basic research carried out worldwide in the field of electrochemistry. A promising research direction is the use of typical and little known transition metals compounds, for example tungsten carbide, which is discussed at the end of this chapter.

The other problem which interferes with a more widespread use of fuel cells


of the hydrogen-oxygen type, i.e. the price and availability of hydrogen, will also hopefully be solved in view of the increasing role to be played by hydrogen as a potential energy carrier in the future [3-6]. The attractiveness of hydrogen on the one hand and the mounting energy hunger of the world on the other, suggest that hydrogen will eventually become one of the major energy carriers. The energy would be recovered from the gas either by anodic oxidation or even in combustion engines [1, 2]. Consequently, the price of hydrogen should de-crease as its availability will increase. This, then, is the justification for the con-tinuing vigorous interest in hydrogen displayed by researchers all over the world.

12.3 Ionization of hydrogen on surfaces of various metals

12.3.1 General

The overall reaction for the anodic oxidation of hydrogen in an acid solution proceeds as in equation (12.1), whereas in an alkaline solution, the following takes place:

In these overall processes a distinction may be made between the stage of reac-tant transport to and from the electrode (diffusion, convection, migration in the electric field) and the stage of the electrode process itself. The latter may further be subdivided into the steps involved in the electrode process. The elucidation of the electrode reaction kinetics consists of determining these steps and their subsequent combination into a sequence that describes the overall process. Taking the hydrogen ionization reaction in an acid solution, the following are possible individual steps in the overall process :

The respective subscripts denote: g—gas phase, b—bulk of solution, s—the charge transfer site in the double layer, ads—the form (physically or chemically) adsorbed.

H 2 ( g) + 2 0 H - -» 2 H 20 + 2e (12.1a)

+ e

(12.5)

(12.6)

(12.7)

(12.8)

(12.9)

(12.10)

(12.11)

IONIZATION OF HYDROGEN ON SURFACES OF METALS 343

Reactions (12.8), (12.9) and (12.10) are the electrode reactions proper; among these reaction (12.8) is a catalytic reaction of dissociative chemisorption, whereas reactions (12.9) and (12.10) are charge transfer reactions. Processes (12.6) and (12.11) are the processes of transport of reactants to and from the electrode, and reaction (12.7), seldom considered in the discussion of hydrogen ioniza-tion kinetics, is the physical adsorption of the hydrogen molecule proceeding dissociation. Little attention is also generally given to the process of transfer of gaseous hydrogen into the solution as in equation (12.5). The principal steps in the kinetics of the electrode process are the reactions (12.8), (12.9) and (12.10), but from the point of view of technical applications all the steps are of equal importance.

The major electrochemical problem is the elucidation of the mechanism of the overall reaction proceeding on various electrode and to determine the kinetics of individual reaction stages. As the kinetics of reactions are essentially affected by the structure of the double layer, investigation of the latter should supplement studies on the kinetics of the hydrogen ionization reaction.

There is a large imbalance in the literature between investigations on the cathodic evolution of hydrogen and on its anodic oxidation. An abundance of experimental and theoretical studies, together with reviews, has been devoted to the cathodic process [7-13]. The anodic process discussed here is covered by relatively few investigations, a large proportion of which are technologically, rather than pure-science oriented. A major reason for this is the fact that a di-rect determination of the kinetic parameters of the hydrogen ionization reac-tion can not be easily deduced from simple current-overvoltage relations ob-tained from steady-state polarization curves for various electrode materials in the anodic potential regions. This is due, on the one hand, to the low solubility of hydrogen in aqueous solutions of the electrolytes, and on the other hand, to the deactivating passivation of a great many electrode materials in the region of potentials of interest.

In practice the only electrode materials featuring a relatively low overvoltage for the reaction of anodic oxidation of hydrogen, thus enabling this reaction to be studied over a fairly wide potential range, are the platinum-group me-tals. Their use in research work helped in the elucidation of the kinetics of the process in question, and in the development of the double-layer theory for the reversible hydrogen electrode. This theory, which is largely the achievement of Frumkin and his co-workers [1, 14, 15], has been a major step forward in relation to the classical double-layer theory for ideally polarizable electrodes (in practice on mercury) derived by Graham [16] and Parsons [17, 18] in the


forties and the fifties. The theory of the double layer for the reversible hydrogen electrode mentioned led to the development of a generalized theory of the double layer on a reversible electrode with respect to any redox reaction [15].

The major assumption of the thermodynamic double-layer theory for the reversible hydrogen electrode is that the condition of an electrode-solution sys-tem, at a given temperature, pressure and concentration of substances other than hydrogen, is determined by the sum of the excess charge at the metal-solution interface and the quantity of the hydrogen adsorbed (in the Gibbs adsorp-tion sense) at the interface. Classical theory takes into account only the first of these two factors. In other words, in order to increase the interface surface area of a reversible hydrogen electrode with constant intensive thermodynamic parameters, a definite number of hydrogen atoms, hydrogen ions and electrons should be supplied to the system from the outside. For such an electrode the concepts of total charge and free charge are introduced as given by formula:

q = FrH = FAH-e (12.12)

where q is the total charge, ε is free charge of an electrode, ΓΗ is the excess con-centration of hydrogen on the interface surface in the sense of the Gibbs's de-finition, An is the concentration of hydrogen atoms at the interface.

From the difference between the total and free charges the total capacitance and free capacitance of the double layer can be derived. Also, the electrocap-illary curves of the first and second kind connected with the constancy of AH or ΓΗ can be derived. These concepts are foreign to the classical double-layer theory for ideally polarizable electrodes. The thermodynamic double-layer theory for a reversible electrode can be resolved into the classical theory for the ideally polarizable electrode by removing from the electrode-solution system one degree of freedom related to the possible passage of the charge through the interface during the electrode reaction.

12.3.2 Ionization of hydrogen on platinum

12.3.2.1 Introductory remarks

Mention has already been made of the problems involved in the study of the hydrogen ionization reaction caused by the low solubility of hydrogen. For example, in 0.5 M H 2S 0 4, the equilibrium concentration of hydrogen at a pressure of 0.1 MPa and temperature of 278 Κ is only 0.71 χ 10"

3 mol/1 and

in 1 M KOH, 0.57 χ 10~3 mol/1. The diffusion coefficient of hydrogen in both


cases is the same and equals approx. 5 χ 10~9 m

2/s. The limiting diffusion cur-

rent evaluated from these data, with a moderate rate of stirring at which the thickness of the laminar layer δ = 50 μηι, is

/d = 2FD-ψ- - 1 5 A/m2 (12 .13)

ο

In the case of an electrode reaction whose rate is transport- and activation-con-trolled by the chemical reaction proceeding on the surface, we have (ignoring the cathodic reaction) :

H'-ibMw) <|2I4)

and

J - = - 1 - 4 - (12.15) ^lim h *d

where iT is the limitting current of the activation step, zd is the limiting diffusion current, illm is the resultant limiting current.

The exchange current density i0 and the limiting currents of the reaction, determined under cathodic polarization, exceed considerably the limiting diffu-sion current evaluated from equation (12.13). Thus on the basis of equations (12.14) and (12.15) it would be expected that the hydrogen ionization reaction is strongly diffusion controlled. Indeed, it has long been known that a limiting current appears on a platinum electrode at fairly moderate anodic polarizations in both acid and alkaline solutions saturated with hydrogen. The nature of this limiting current, however, is obscure and investigators have differed in their interpretations (see e.g. [19] and [20]).

The first reliable results were obtained fro the mrotating disc electrode method [21, 22] which allows control of the limitations related to transport of the reactants in the solution. In particular it was found that with increasing speed of rotation of the disc, and hence with an increase of limiting diffusion current, on platinum in acid or alkaline solutions, reactions (12.1) and (12.1a) become activation controlled upon elimination of the diffusion limitations. It was also found that adsorbed anions have a profound effect on the activation controlled current. The anions S O I " , Cl", Br"" and I" are responsible for the passivation of a platinum electrode for the hydrogen oxidation reaction even at low anodic polarizations, and the process ceases to be diffusion-controlled.

Additionally, an interesting dependence of the anodic current on the square root of the disc rotation speed was observed under conditions of no passivation,


Square root of disk speed

Fig. 12.2. Current density vs. square root of the rotation speed curves for hydrogen ionization on a rotating disc electrode (after [22]): 1—reaction of zero order with respect to H 2, 2—diffu-sion control over the entire range of disc speed, 3—reaction of the first order with respect to H 2.

as illustrated by curve 1 of Fig. 12.2 [22]. On this curve, the sharp bend corre-sponds to a change from pure diffusion control to a condition that is independent of the disc rotation speed. In the flat portion of the curve, therefore, the current is independent of the rate of transport of molecular hydrogen to the electrode surface, and thus independent of the hydrogen concentration at the surface. In other words, the electrode process in the activation region is of zero order with respect to molecular hydrogen. The most probable reason for this is hin-drance of the process in the activation region by the reaction that follows the step of fast dissociative adsorption given by formula (12.8). The charge transfer reaction (12.9) or a slow diffusion of H a ds atoms over the metal surface from chemisorption sites to ionization reaction sites, have been put forward as the hindering step. For the process in the activation region every other mechanism with a reaction order larger than zero should, according to [22], lead to a rela-tion i = f(m) illustrated by curve 3 in Fig. 12.2; in other words, what should be observed is a smooth passage from diffusion control leading to mixed (activa-tion-diffusion) control and thence to the limiting current of the reaction.

These ideas have been widely accepted for many years and can be found in standard text-books and monographs [23]. More recent studies have, however, raised doubts, and the present opinion, following [24], is that in acid solution free from adsorbing ions and containing an active platinum electrode the ob-served effect depends simply on the hindrance of step (12.6) over the entire po-tential range from zero to approx. 700 mVR H E. The activation controlled process itself, being many times faster than the process of hydrogen diffusion to the elec-trode surface, if experimentally observable at all, is of first order with respect to molecular hydrogen. Thus, the current-disc speed relation recorded for these


conditions is a straight line 2 in Fig. 12.2. It follows that practically no conclu-sion can be drawn either from steady-state measurements of current vs. poten-tial or from current vs. disc rotation speed relations as to the kinetics of the activation steps. At more anodic potentials, both in acids and alkalis, the hydro-gen ionization current drops, due to passivation of the electrode (deactivation of the electrocatalyst) by the adsorbing anions and the oxide film formed on the metal surface.

The current vs. overpotential plot over a wide potential range, typical of the platinum electrode in various hydrogen-saturated solutions, is shown in Fig. 12.3.

I 1 — , 1 — : 1—. 1 1 I ι L_

0 0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6 Potential, V N HE

Fig. 12.3. The current vs. potential dependence for platinum electrode in a hydrogen-saturated 0 . 5 M H 2S O 4 solution at 298 K; (origin-^)—the diffusion control range for the electrode in the active condition; (a-b)—the transition between the active and passive conditions; (b-c)— residual current in the passivity range; (c-d)—increasing current due to oxygen evolution.

The curve has four characteristic regions: — (origin-fl), a region following the formula for diffusion overvoltage:

— (a-b)9 passivation of the ionization reaction caused by the mixed effect of anion adsorption and oxide formation,

— (b-c), residual current in the passive region, — (c-d)9 increasing current caused by oxygen evolution. The hydrogen ionization current in the passive region is about one order of


magnitude less than the limitting diffusion current, and is more or less independent of the conditions of reaction transport (disc rotation speed).

Alternatives to the rotating disc technique have been tried, in attempts to reduce the diffusion limitations in reactions (12.1) and (12.1a). Examples are the semi-immersed electrode [25-27] and the method where a thin film of elec-trolyte separates a microelectrode from the gas phase [28]. Higher values of the limiting diffusion current density were achieved using such techniques, but even so, it was not possible to reach the activation controlled region of the reac-tion.

In the presence of adsorbing anions or in alkalis, the kinetic currents can be determined at moderate speeds of the rotating disc electrode [22, 24]. It is also known that, for expample in 1 M NaOH the exchange current of the hydrogen ionization reaction is about 2 A/m

2 under standard conditions (pHi = 0.1 MPa,

Τ = 297 Κ) [19], whereas in acids, depending on the anionic species and their concentration, the exchange current may range from 50 to approx. 120 A/m

2 [22].

It should be emphasized that one reason for the frequent discrepancies be-tween results of studies of these reactions, obtained by various workers, is their considerable sensitivity to the presence or even trace quantities of contamination in the reaction medium, as is true for all catalytic processes. Even with the highly sophisticated and often laborious methods for the purification of the reac-tion medium available today, it is often impossible to remove contaminants completely. One recommended procedure is to thoroughly standardize the ex-perimental procedure, including the preparation of the electrode, by prelimi-nary electrode polarization under successive potentiostatic conditions at several selected polarization values, maintaining these polarization values for definite, strictly reproducible periods. This technique of electrode preparation should be followed prior to every single measurement [29].

There is reason to believe, however, that the time-dependent changes in the catalytic properties of the platinum electrode are related not only to the adsorp-tion of contaminants but also to a process of slow reconstruction of the surface layer, leading to a stronger bonding of the adsorbed hydrogen to the metal, which in turn may result in a gradual increase in the overvoltage of the hydrogen ioni-zation reaction. Another factor is that the slowly adsorbing ions may bring about gradual changes in the energy state of the surface hydrogen. The investi-gation of effects of this kind, however, requires a departure from the steady-state measuring technique. The alternative is to use one of the various pulse methods, which are particulary useful in studying the intermediate stages in the overall reaction.


12.3.2.2 Theory of hydrogen ionization on platinum

By analogy with the cathodic reaction [12], the following possible mechanisms may be put forward for the hydrogen ionization reaction:

(a) H 2 ^ 2 H a ds dissociative chemisorption (Tafel) reaction H a ds ^ H

+ + e ionization (Volmer) reaction

(b) H 2 ^ H a ds -f H+ + e electrochemical adsorption (Heyrovsky)

reaction H a ds H

+ + e

(c) H 2 HJ + e ionization of the H 2 molecule Ht ^ 2 H

+ + e further ionization followed by dissociation of the

molecular ion. The kinetic step preceding each of the reaction pairs written is sometimes in-

cluded, viz.;

H2(S) ^ H 2 ( a d s) For the reaction in an alkaline medium, the respective mechanism can be modi-fied by adding OH~ to both sides of the above reactions and by considering the water dissociation equilibrium, H

+ 4- OH~ ^ H 20 . In an alkaline solution

mechanism (c) has not so far been considered. The ionized molecule HJ was theoretically supposed to be involved in the

hydrogen ionization mechanism as early 1936, by Horiuti et al. [30, 31], but it was not identified experimentally as an intermediate product of the overall reaction. The participation of this reaction was suggested on the basis of experi-mental evidence, from the activation current vs. pH dependence [28]. In fact there are no conclusive proofs in favour of any of the proposed mechanisms, and the debate continues. The prevailing opinion, however, is that in a neutral or alkaline solution mechanism (a) holds for the hydrogen ionization process on platinum. For acid solution one is faced with a difficult choice between the mechanism (c) suggested by Bagotsky [28], and that traditionally assumed for this system, mechanism (a), which has been substantiated in recent studies by Ross and Stonehart [32]. A possible change in the mechanism of the process as a function of overvoltage, which is often assumed in order to explain the experi-mentally observed variations in the Tafel slopes in the cathodic reactions [33], or the likelihood of a mixed mechanism, can not be disregarded.

12.3.2.3 Determination of the kinetics of individual steps in the hydrogen ionization reaction

The majority of experimental data are related to Volmer's reaction, both as regards its kinetics and the related adsorption equilibria. The platinum electrode,


both in acid and alkaline solution, in the presence of dissolved molecular hydro-

gen, attains a steady potential equal, on the standard hydrogen electrode scale, to

RT Λ a& ENHE = I n — ^ (12.17)

where and aH2 are the activities of hydrogen ions and of molecular hydrogen, respectively. Hence, under standard conditions

i?NHE = -0.059 pH-0.0295 logaHl (12.18)

From this relation, the potential with respect to the reversible hydrogen electrode

in a solution of a given pH is :

^RHE = ^NUE

+ 0.059 pH (12.19)

£RHE = -0.0295 loga„a S -0.0295 log/?H2 (12.20)

If an adsorption equilibrium is assumed, μ Ηΐ = 2μΗ ( 3 ( ί 8) and if to take μη2 = 2 μ Η( 3( ΐ 8 ) then

£RHE = -0.059 logaH ( a d S) (12.21) At equilibrium with gaseous hydrogen at a pressure activity of 0.1 MPa, i.e. at zero potential on the RHE scale, platinum is assumed to be covered with chemisorbed hydrogen in an atomic ratio of H:Pt = 1:1, i.e. hydrogen forms a monoatomic layer on the platinum surface. This has been demonstrated from studies of adsorption from the gas phase [34]. If such an electrode is polarized by a small anodic current, variation of the potential with time will correspond to variation in the activity, following formula (12.21), and hence also to varia-tion in the concentration of atomic hydrogen on the surface as a result of Vol-mer's reaction. It follows from equation (12.21) that

dFRHE = d(-\naH) = -RTdaH

dt F dt FaHdt

1

'

;

where aH stands for alHads). For this process the faradaic current passing through the electrode-solution interface will be the ionization current of the adsorbed hydrogen :

dqQ άθ / n o-' - - i r

= ^

( 1 2·

2 3)

where qQ is the charge required to produce a complete hydrogen monolayer as a result of Volmer's reaction, and qQ is the charge required to form an adsorption layer of coverage θ by the same reaction. Considering therefore the relation of potential vs. the charge passing through the electrode-solution


interface under polarization but near equilibrium conditions, the coverage 0 ( = qelqo) vs. potential relation is obtained. Here the double-layer charging current is disregarded, as it is at least one order of magnitude smaller than the faradaic charge for measurements carried out with sufficiently small variations in the potential. Thus, under a slow potential variation, the hydrogen adsorp-tion isotherm on platinum in an electrolyte solution is obtained. This isotherm may be presented in various equivalent forms :

Θ=/Ρ(η2) (12.24a)

6=f(ERHE) (12.24b)

qo = / ( £ R H E ) (12.24C)

The forms (12.24b) and (12.24c) are the most convenient in use, as they are directly related to electrochemical measurements. Curve b in Fig. 12.4 shows the relation θ = f(ERHE) for a platinum electrode in a 0.5 M H 2S 0 4 at 298 K.

If the electrode is polarized under potentiodynamic conditions, i.e. if linear variation in the potential with time is imposed, occurring at a small and constant rate dEjdt = s, the current flowing through the electrode will be expressed by relation

. _ dqe dqQ dE _ dqQ _ dd ~dË~W~ ~

S~dË~

Sq°~dË

(12.23a)

The current vs. potential curve in this case (curve a in Fig. 12.4) is the deriva-tive of the curve b in the same figure. The peaks on the curve which are not

0.2 0.3 Potential, VN HE

Fig. 12.4. The ionization current of adsorbed hydrogen and coverage with hydrogen atoms vs. potential on a platinum electrode in 0.5 M H 2S 0 4 (s = 100 mV/s). Curve a—i vs. £"RHE; curve b—0 vs. JERHB.


seen on curve b, are due to the anodic current caused by hydrogen oxidation. If the process is carried out in reverse, that is, if the potential is varied linearly towards zero starting from 500 mV, the current vs. potential curve is roughly as shown in Fig. 12.5. The curve is more or less a mirror image of curve a from Fig. 12.4. Two valleys (negative peaks) can be seen on this curve which corre-spond to the adsorption of two forms of hydrogen.

There are no direct indications that the hydrogen adsorption and desorption peaks of Fig. 12.4 and 12.5 can be related to the various forms of chemisorbed hydrogen observed in the adsorption from gas phase studies [35-37], despite the suggestion of some investigators that such a relation exists. The properties of the platinum-electrolyte interface are clearly different from those of the P t -H2 (gaseous) interface, if for no other reason than for the presence of water dipoles strongly bonded to the metal surface. To allow identification of the two (or three) hydrogen forms observed in the electrochemical experiments, much more refined methods, still awaiting development, will be required for the study of the ad-sorbed molecules in situ.

The recent results obtained with Pt monocrystals suggest that the different hydrogen adsorption peaks on polycrystalline electrode should be related to different crystallographic structures on the metal surface [38, 39].

From the results of potentiodynamic measurements, using a suitable kinetic model for Volmer's reaction, numerical values of the equilibrium constants and of the rates of this reaction can be obtained on the basis of the theoretical analysis by Srinivasan and Gileadi [40] with further refinements by Stonehart [41] and

N HE Fig. 12.4. The ionization current of adsorbed hydrogen and coverage with hydrogen atoms vs. potential on a platinum electrode in 0.5 M H 2S 0 4 (s = 100 mV/s).


Angerstein-Kozlowska et al [42-44]. In the case of Volmer's reaction under conditions close to equilibrium, simple formulae can be derived after these authors to describe the current-potential relation. The formula for the faradaic current in Volmer's reaction for the case of adsorption that follows the Langmuir isotherm is :

h = /,-ic = M e x p - ^ - f c c O -0)exp (12.25)

Under conditions close to equilibrium ( i F <̂ / a, / c, o r /a £ /c) and from the balance condition (equation (12.23)) it follows that:

FE Ί0 RT

F RT

\ „ FE FE 1 ~ e x p ——

RT F

RT

(12.26)

where Κ = ka/kc, Ε = E0 + st, s = dEjdt. Formula (12.26) shows that the current is proportional to the rate of vari-

ation of the potential, thus it is purely capacitive in character. We refer to this case under conditions close to an equilibrium as "pseudocapacitive behaviour" of the hydrogen electrode. This is why some abnormally high capacitance values are measured on the platinum electrode in the hydrogen region. The equilibrium constant of Volmer's reaction, K, can be easily found from formula (12.26) or, better perhaps, after differentiation from the formula for the peak current (see equation (12.29)).

If the rate of potential variation is increased, the current, starting from a cer-tain value is no longer proportional to this rate, and one gradually passes into a logarithmic relation. This is connected with departure from the equilibrium condition towards the range of full irreversibility of Volmer's reaction. For the anodic sweep of the potential and for /a > zc, i.e. ip = za, formula (12.25) re-duces to

/F = * a0 e x p ^ p - (12.27)

Clearly the balance condition, i.e. formula (12.23), must be fulfilled and, using this and formula (12.27), a new formula is obtained for the current vs. potential relation and the current vs. potential variation rate relation in irreversible case:

. . . 7 KRT (xFE0l (xFst \ ocFE .


Differentiation of this gives the value of the peak current and, on simplifying, this assumes the form:

(12.29)

(12.30)

in; , „ q0«sF , KRT «E0F Imp.™, = l n ^ + - ^ - e x p ^ - - l

and the value of the potential that corresponds to this peak current is

RT , q0otF RT , ^ = ^ I n u f +

~ a 7

I nS

From the dependence of the peak current potential on the rate of the potential variation, the anodic reaction rate constant can be found using equation (12.30). In the same way, by conducting a rapid potentiodynamic measurement with potential variation towards the cathodic side, the cathodic reaction rate constant can be determined, provided that the complete irreversibility condition is sat-isfied. In this case, however, the onset of the cathodic hydrogen evolution makes the analysis of the cathodic peaks much more difficult. Thus a complete set of data may be, in principle, obtained for use in the kinetic equation for the rate of Volmer's reaction. In the case where the adsorption cannot be assumed to follow the Langmuir isotherm, but heterogeneity of the surface must be accounted for in terms of Temkin's or Frumkin's isotherm, the derivation of the above for-mulae becomes much more involved and the calculations are best made numeri-

0.1 Potential, V N HE

Fig. 12.6. Theoretical dependence of the oxidation current of the adsorbed hydrogen upon the electrode potential, as determined from formula (12.26) (q0 = 2.1 C/m

2, ka = 1 x 10"

2 A/m

2,

ke = 1 A/m

2, s = 100 mV/s).


cally. In this case also, however, the equilibrium constant and the rate constants for Volmer's reaction can be evaluated. Examples of such calculations can be found in papers by Angerstein-Kozlowska et al [42-44].

As illustration, the theoretical curve for the current vs. potential relation given by formula (12.26) is plotted in Fig. 12.6. The curve corresponds to the values of constants α = 0.5, Jfc. = 1 χ 10"

2 A/m

2, k0 = 6.5 χ 10"

2 A/m

2, s = 0.1 V/s.

A comparison of this curve with the experimental graphs of Figs. 12.4 and 12.5 shows that, at least formally, hydrogen adsorption on polycrystalline platinum can be described in terms of two overlapping Langmuir isotherms, that is by two pairs of rate constants in equation (12.25). This idea was advanced by Eucken and Weblus [45] and developed by Breiter [46, 47]. These authors clearly distinguish two forms of adsorbed hydrogen, referring to the weakly bonded hydrogen (a peak in the lower potential region) and the strongly bonded hydrogen (a peak at higher potentials).

The potentiodynamic method is classified, with many others which require considerable potential or current variation, as a non-linear method. This is due to the non-linear form of equation (12.25) in common with the majority of electrochemical kinetic equations where the relation between the "driving force" of the process, i.e. the electrode potential, and the "flow", i.e. current density, is non-linear. There is another group of techniques in which the poten-tial variations imposed are so small that in this range a linearization of the kin-etic equation is possible by expansion into a Taylor series, followed by rejec-tion of the non-linear terms. The most commonly used of these methods is the alternating-current or impedance method. In contrast with non-linear methods, such techniques often do not require an exact form of the kinetic equation such as equation (12.25), merely the knowledge of the general form of the relation of individual parameters. This again can be illustrated using Volmêr's reaction as an example. Considering that

(12.31)

the formula for a total differential of /F is

(12.32)

The balance condition is (equation (12.23))

dd dt

i L 9o


If low amplitude sinusoidal variations of potential are imposed, described in complex form by formula (12.33), the current will also vary following the sine function, as described by equation (12.34):

dE = \dE\cxp(ja)t) (12.33)

diF = \diF\exp(jœt) (12.34)

Hence, upon rearrangement, using formulae (12.32), (12.23), (12.33), and (12.34):

q0j<u \3eJQ

\M JE diF \diF\ _ 1 dË \dE\ ~~ ZF q0jo>

(12.35)

tdiA \ se JE

where./ = ( - 1 )0'5, ω = 2π/.

Formula (12.35) may be re-written as

Z F = Rs+1̂ ~ (12.36)

where Z F stands for faradaic impedance, and Rs and Cs are series-connected resistive and capacitive equivalents of this impedance. From the above deriva-tion it follows that in Volmer's reaction, using very small potential and/or cur-rent variations, a series equivalent RC circuit can represent the impedance of this reaction. Moreover,

χ* = τ τ γ γ (

1 2·

3 ?)

and \SEJe

\39 ) E

(12.38)

Thus, if the conservative assumptions made and contained in equations (12.23) and (12.31) are valid, the values of Rs and Cs determined from the impedance measurements should be independent of the frequency of oscillation of the elec-trode potential. And hence, using the complete kinetic equation (12.25) and


deriving from it the derivatives of expressions (12.37) and (12.38), formulae for Rs and Cs can be written as functions of potentials and rate constants:

Γ, * *F I *F Λ / /ι m ( l - a ) F ( α - Ι ) ^ Γ

1

/ 1 0 < 5m

L Λ aF aF£ t y , f m ( l - a ) F ( a - l ) F £ _ ^ P ^ e x P^ +/ : c ( l ^ e ) ^ ^ e x p - - ^ r - - |

. " , " aFE . (oc-l)FE A C XP " Λ Γ "

+^

E XP —Rf—

Since equations (12.39) and (12.40) are independent, values of the rate constants ka and kc of Volmer's reaction can be evaluated from the experimentally found Rs and Cs values.

Unfortunately, measurements show that Rs and Cs are frequency dependent; however this dependence can be accounted for by introducing into the formula for the electrode impedance an expression of the type B\/f, where is a constant factor, and / i s the frequency [48]. The physical meaning of this term is not clear, but it may be noted that a square-root dependence on frequency is character-istic of diffusion processes, inclusive of the surface diffusion processes [49]. These therefore should perhaps be taken into account in deriving the kinetic relations in the anodic reactions of hydrogen on the platinum surface, as has been attempted by Novoselskii [50]. But, apart from these complicating cir-cumstances, it is an important finding that the values of rate constants and ex-change currents of the Volmer's reaction can be, at least, estimated from impe-dance measurements. Lukyantseva et al [51] showed that the exchange currents are of the order of 500 A/m

2. This indicates that Volmer's reaction is very much

faster that the overall reaction of hydrogen ionization, thus contradicting the earlier view that this reaction is the limiting step in the overall process of hydro-gen ionization on active platinum in both acid and alkali.

Another elementary reaction connected with the anodic oxidation of mole-cular hydrogen on platinum and platinum-group metals, is the dissociative chemisorption reaction (Tafel's reaction). This reaction is not related to a net flow of electric charge through the interface and thus, it cannot be directly studied from current-potential relations, as in the case of Volmer's reaction. Inferences as to the participation and kinetics of this reaction are usually drawn indirectly from an analysis of the steady-state polarization curves, and in parti-cular on the basis of a study of the order of the reaction with respect to hydrogen ions, as well as by determination of Tafel slopes. These methods, however, are


confined exclusively to the hydrogen evolution reaction and cannot be used to study the ionization of molecular hydrogen. An alternative approach is as fol-lows. The decay of potential with time is recorded after the polarization current had been turned off, a solution containing hydrogen in a definite concentration, and under suitable conditions of hydrogen transport towards the surface (e.g. by using the rotating disc electrode). The rate of hydrogen accumulation on the surface, i.e. the rate of chemisorption, can thence be determined on the as-sumption of a unique dependence of the electrode potential on the coverage with atomic hydrogen (the isotherm in formula (12.24b)), i.e. by assuming a priori the equilibrium of Volmer's reaction. Such studies were conducted by Gromyko et al [52] and Rosen and Schuldiner [53], who showed the reaction to be first order with respect to molecular hydrogen and to follow kinetic equation which assumes the validity of the Temkin isotherm. It is interesting, incidentally, that Bagotsky found that hydrogen adsorption on palladium follows the Frumkin isotherm. It is not known, however, to what an extent the results of measurements for palladium were rendered inaccurate by the permeation of hydrogen into the bulk of metal, which may be inferred from the graphs given in the paper. Ba-gotsky proposed the following kinetic equation for platinum in both 0.5 M H 2S 0 4 and 1 M KOH:

*>ads = kpÎÎ2exp(-βgθH) (12.41)

where β is a symmetry factor, and g is the surface inhomogeneity factor in the Temkin isotherm. The rate constant evaluated from this formula, approx. 45 A/m

2, was almost the same for both of the solutions studied. A similar

value for Pt in H 2S 0 4 was obtained by Rosen and Schuldiner [53]. If follows from these data that the dissociative adsorption reaction on a platinum electrode is much slower than Volmer's reaction. Therefore, if mechanism (a) above (i.e. the Tafel-Volmer mechanism) is assumed to be operative, Tafel's reaction has to be the rate determining step.

It needs to be said, however, that other measurements of adsorption rate, carried out by recording the decay of potential on the rotating disc electrode after cutting the current off [54], fail to confirm relation (12.41). In particular, these findings shows that, as with Volmer's reaction, hydrogen adsorption on platinum proceeds in two consecutive stages, and that the adsorption rate does not decline monotonously with coverage. Nevertheless, from the lack of con-tradictory data, it can be assumed that it is Tafel's reaction that controls the hydrogen ionization rate on platinum and platinum-group metals in alkaline solution, disregarding diffusion limitations. The assumption of mechanism (a)


for an acid solution, on the other hand, leads to an ionization current vs. pH relation (diffusion limitations disregarded) which disagrees with the experimental data. This relation instead can be satisfactorily explained by mechanism (c), which suggests a reaction order of 0.5 with respect to H

+ ions. Many re-

searchers, however, have doubts about mechanism (c), as it fails to fit into the whole overvoltage theory for the hydrogen evolution reaction, which is by now well developed. Only Horiuti et al [31] have given attention to this mechanism.

A number of aspects of the kinetics of hydrogen ionization on platinum or platinum-group metals are yet to be clarified [23, 55-58].

It is important to realize that, for a thorough analysis of the kinetics of both ionization and evolution of hydrogen, an essential item of information is a de-scription of the adsorption of atomic hydrogen on the electrode.

Current experimental techniques do not permit the development of theor-etical models to describe the equilibrium and kinetics of the hydrogen adsorption reaction on metals, under electrochemical conditions. These conditions pose serious problems in attempts to maintain a well defined surface throughout the experiment, which would ideally have a simple structure, i.e. monocrystalline planes of low index. Only such a surface allows theoretical models, based on the theory of the metal-hydrogen, metal-water and metal-electrolyte ions chem-isorption bonds, to be applied at least in the initial stage of the theoretical analysis. Electrochemical adsorption, for both the experimenter and theore-tician alike, is a much more complex phenomenon that the "pure" adsorp-tion of atoms or molecules on a well defined, recrystallized metal surface, equilibrated under high vacuum. Nevertheless, electrochemical studies must follow the current trend of the physical chemistry of surface and heterogeneous catalysis, where the specimens used must be clean and they must have a surface with a known and simple structure. This requirement has long been recognized; in 1965 Will [59] carried out some of the earliest measurements of hydrogen adsorption on platinum single crystals of orientation {100}, {110} and {111} in 8 Ν H 2S 0 4. He found clear differences in the shape of the voltammograms and, in particular, different ratios of concentrations of strongly and weakly adsorbed hydrogen, for the various monocrystalline planes. Following this work, similar studies have been made by other researchers [38, 39, 60-68]. But despite the use of standard procedures, combining electromechanical techniques~and surface chemistry methods, there has been little agreement between the results obtained. The results of some studies [60, 61] have called into question the observations made by Will as to the unequal hydrogen adsorption on different platinum crystallographic planes. In these studies, however, as in some recently published


ones [64, 65], the standard, electrochemical redox cleaning of Pt surface was used which, in the light of more recent findings [39], brings about recrystalliza-tion of the surface and a corresponding loss of its homogeneous structure.

In the latest studies [38, 62-65] the platinum surface structure and chemical composition were checked before, during and after the electrochemical experi-ment using LEED, RHEED or AES methods. Even these studies, however, leave some doubts as to the equilibrium hydrogen concentration on the plati-num monocrystalline planes, especially as to the specific différencies in the energy of adsorption (potentials of peaks on voltammograms) for hydrogen on various platinum crystal planes. Agreement is found only in the case of the {100} plane [38, 62-65], associated with a peak indicating strongly bonded hy-drogen at a potential of approx. 250 mV on the RHE scale. The authors of all these papers claim that a prolonged cycling of monocrystalline electrodes be-tween the hydrogen and oxygen adsorption potentials results in a reconstruction of the surface and a loss of its homogeneity. An interesting observation in the three references [38, 62, 63] is that the hydrogen adsorbs much more weakly and at lower equilibrium concentrations on the {111} plane that on the other planes. This is in contradiction to the results of studies of adsorption from the gas phase. It should, however, be borne in mind that in an electrolyte environ-ment, hydrogen adsorption occurs in competition with the adsorption of elec-trolyte ions, water and oxygen.

The importance of the character of the ionic environment on the hydrogen adsorption voltammograms was pointed out by Ross [66]. His findings demon-strated that the difficulties in comparing the results obtained by various workers arise from the fact that they used solutions of different ionic compositions and concentrations.

The emerging outline of a consensus at the end of the seventies has been re-cently disturbed by a series of works by Clavilier et al [39, 67-69] which show some major conflict with all the preceding works on this subject. This conflict con-sists in a more extensive potential range for hydrogen adsorption and in a divi-sion of the voltammogram into two regions in the absence of strongly adsorbing anions. Moreover, Clavilier's results are impressive and convincing, since they are in line with the present understanding of the adsorption of anions on platinum. The technique of preparation of Pt monocrystals and of contacting them with an electrolyte differed considerably in the Clavilier's work. Clavilier claims that due to specifities of his experiments the purity conditions are reached which previously were unatainable. There is a need for these results to be confirmed by inde-


pendent research team; until then the problem of electrochemical adsorption of hydrogen on platinum single crystals still has to be regarded as open.

It should be noted that no data are as yet available on the kinetics of hydro-gen adsorption arid hydrogen evolution on platinum (or other metals) single crystals. And there are in the literature no reliable data concerning the electro-chemical adsorption of hydrogen on monocrystalline surfaces of metals other than platinum.

12.3.3 Ionization of hydrogen on nickel

12.3.3.1 Electrochemical properties of nickel

Nickel is a metal that is unstable in solution of pH < 7. The corrosion rate in aqueous acid solutions depends on the anion, and the corrosion products are soluble N i

2+ compounds. The corrosion potential of nickel in a 0.5 M H 2S 0 4

is about - 5 0 m V R H E. If nickel specimen is polarized by an anodic current above the critical passivation potential (110 to 150 mVR H E) , the metal becomes covered with a protective oxide film.

In solutions of pH > 7 nickel is usually resistant to corrosion. The rest po-tential of nickel in a deoxygenated 1 M KOH solution is 600 to 800 mVR H E. The corrosion resistance and the value of the rest potential given above are a proof of the presence of an oxide film on the nickel surface [70].

In view of the widespread use of nickel in technology, the electrochemical properties of the metal have been the subject of a continuous reasearch effort. Some of this has been concerned with the hydrogen evolution or ionization reac-tion. For example Lukovtsev et al [71] showed that the hydrogen evolution process follows Tafel's relation, the slope value b in alkaline solutions at room temperature being 0.115 V/decade, and the value of the exchange current rang-ing from 10~

3 to 10""

1 A/m

2. Other authors have obtained values between 0.100

to 0.120 for the slope b [72, 73]. Nickel is therefore classified as a metal with a low overvoltage for hydrogen evolution.

The considerable corrosion resistance of nickel in concentrated aqueous solutions of high pH, combined with the low hydrogen evolution overvoltage on nickel electrodes, indicate that nickel may be a promising electrocatalyst for the hydrogen ionization reaction. In the last two decades a number of papers have appeared on the subject. These are chiefly studies of an engineering char-acter, in which the electrode material was predominantly nickel powder or Raney nickel, a widely recognized catalyst in a vast number of hydrogénation


processes of organic compounds, rather than solid nickel in the form of sheet or wire.

Raney or skeleton nickel, with other skeleton metals (referred to as Raney metals), is a material of a huge surface area. Unfortunately, electrochemistry has as yet no methods for exact determination of the fundamental electrochemi-cal quantities for the processes occurring on porous electrodes. Accordingly, so far it is not clear whether the characteristic properties of skeleton nickel are the consequence of the highly developed nickel surface, or whether they are specific result from the manner in which it is prepared.

Skeleton nickel usually contains a large proportion of adsorbed and absorbed hydrogen; the atomic ratio of hydrogen to nickel in skeleton is 0.2 to 0.3 [74, 75]. In spite of this, the value of the exchange current for the oxidation-reduction reaction of hydrogen on the skeleton nickel surface resembles the corresponding values for this reaction on the surface of solid metal [76].

The rest potential of skeleton nickel in a concentrated alkaline solution saturated with hydrogen of a pressure of 0.1 MPa is 0 m V R HE [77], a value which is only attainable on a smooth nickel surface by a careful surface treatment [78, 79]. One way to achieve a stable rest potential of 0 m V R HE for solid nickel is a prolonged cathodic polarization of a nickel electrode, which possibly brings about both the reduction of the surface oxides and some penetration of hydro-gen into the metal phase [76]. The process of hydrogen penetration can be sub-stantially accelerated by adding "catalyst poisons" to the electrolyte. The treat-ment however results in a rest potential that is negative [80].

An important feature of skeleton nickel electrodes is that they allow the continuation of the hydrogen ionization reaction at a fairly high current density for hundreds of hours without the current-potential characteristics of the elec-trode being affected. On a smooth-nickel electrodes, on the other hand, steady-state conditions for this reaction at room temperature have not yet been achie-ved.

It is currently believed that in the potential range from —150 to +150 mV on the nickel electrode, the oxygen chemisorption process occurs as follows:

OH~ ^ O H a ds + e (12.42)

O H a ds + OH" ^ O a ds + H 20 + e (12.43)

and at higher potentials (above + 150 mV) a film of hydrated NiO, Ni (OH)2 is formed:


Ni + 2(OH)~ ^ Ni(OH)2 + 2e (12.44)

which, at much higher potentials (above +600 mVR H E) , undergoes reconstruc-tion into a form that is difficult to reduce electrochemically [81, 82].

It is therefore considered that in the region of small positive (anodic) poten-tials, which are of interest from the viewpoint of the hydrogen ionization pro-cess, two competitive processes proceed, namely the hydrogen ionization process and a strongly irreversible oxygen adsorption process or the formation of a pass-ive film. This is true of both the smooth and skeleton nickels [71, 75]. The si-multaneous occurrence of both these processes hinders substantially the deter-mination of the hydrogen adsorption isotherm on a nickel electrode in aqueous solutions. This is the reason for the major disagreement in the estimation of the coverage of a nickel electrode with hydrogen, observed at potentials near zero VRHE- The coverage under conditions of an equilibrium with gaseous hydrogen of an activity of 0.1 MPa is about 0.005 according to Makrides [78] and about unity according to Burshtein and co-workers [83, 84]. Some investigators claim that hydrogen is adsorbed on nickel in two forms: the atomic one (predominant form at positive potentials) and the molecular one (predominant form at nega-tive potentials) [85].

The skeleton nickel electrodes lose their activity with respect to the hydrogen ionization reaction, if the anodic polarization is too high. This is due to the forma-tion of an almost irreducible nickel oxide on the nickel surface [86].

This sensitivity to higher anodic potentials, which results in a decline or even a complete loss of electrocatalytic activity, seriously restricts the usefulness of this metal as a hydrogen ionization catalyst. Furthermore, in view of its highly developed surface and defect structure, skeleton nickel has a tendency to recrystallization, which also results in its decreasing activity with time. Some doping of the skeleton nickel with minor quantities of promoters such as Ti, V , Mo, or W has been found advisable [87], A titanium addition probably in-creases the stability of the catalyst structure [88], whereas vanadium helps in enhancing the catalyst activity in the hydrogen ionization reaction [87]. The presence of promoters in the catalyst also affects the overall energy of activation of hydrogen reactions on skeleton nickel [87]. According to Fasman and Sokolsky [87], doping NiAl alloys (from which skeleton nickel is produced by etching he aluminium away) by the .y-, d-, and /-group metals affects the phase composi-tion of the alloy. With nickel or aluminium the alloying additions form two- or three-component crystallites, whose presence essentially affects the physico-chemical properties of the catalyst surface.


12.3.3.2 Mechanism of hydrogen ionization on nickel

The mechanism of the electrode reactions of hydrogen on nickel in alkaline solutions is not yet elucidated. A mixed-type mechanism is belived to be likely, where the controlling step varies with the potential and concentration of the solution. The prevailing opinion, however, is that advanced by Lukovtsev et al. [71] who claim that the controlling step for not too dilute solutions is Volmer's reaction

N iHa ds + OH- ^ Ni + H 20 + e (12.45)

Ignoring the effects of diffusion, a kinetic formula has been given by this author which accounts for the effects of current density and OH"" concentration on the overvoltage:

2RT RT η = — = - In ι - - = - l n [ O H " ] + const (12.46)

The validity of this accounting for the effect of OH" concentration on overvoltage has been confirmed by O'Brien and Rao [79].

Much controversy has arisen from the fact that the limiting current for hydro-gen ionization was observed on both smooth and porous nickel electrodes. In Lukovtsev's opinion the limiting current is due to the stoppage of hydrogen transport to the electrode surface. This, however, disagrees with the fact that the limiting current is generally a few orders of magnitude lower than the limit-ing diffusion current of hydrogen in alkaline solutions (approx. 10 to 50 χ χ 10~~

2 A/m

2 as against approx. 10 A/m

2) [88]. Some authors (e.g. Horiuti and

Sugavara [89]) believe that the electrochemical reactions of hydrogen on nickel are hindered by recombination, i.e. by Tafel's reaction and it is this process that they consider responsible for the limiting anodic current. This view is partly shared by Makrides [78], who claims the limiting current to be of a mixed diffu-sion-activation character and the activation part to be due to Tafel's reaction. It should be noted, however, that not all authors take into account the Tafel reaction to be a step in the hydrogen reactions on nickel [85]. The present con-sensus is (cf. e.g. [81, 90]) that the appearance of a limiting current is not con-nected with any particular mechanism, so that it cannot be used in support of, for example, the hypothesis of hindrance by Tafel's reaction.

According to Breiter [90], the apparent limiting current originates from a de-activation of the electrode due to the irreversible formation of an oxide film at low potentials approaching zero in the RHE scale. This accounts for the often encountered ^reproducibility of the current-potential curves for the anodic


oxidation of hydrogen, as well as the hysteresis that appears once potentials close to 100 mVR HE are reached.

Figure 12.7 shows potentiostatic curves for the current vs. potential relation

0 0.04 0.08 "0.12 0.16 Potential, VN HE

Fig. 12.7. Potentiostatic anodic polarization curves for a smooth nickel electrode in hydrogen-saturated NaOH solutions (after [79]): 1—0.1 M, 2—0.5 M, 3—1 M.

for a nickel electrode in hydrogen-stirred NaOH solutions [79]. Three character-istic potential ranges may be distinguished on these curves :

zero to 20 mV: a fast rise of the hydrogen ionization current with potential; 20 to 70 mV: a gradual suppression of the rise of ionization current with po-

tential (apparent limiting current); 70 to 200 mV: declining current with potential. For some nickel samples these ranges are moved to more positive potentials

[85]. The fact that in the 20 to 70 mV region the limiting current is strongly de-

pendent on the OH" concentration is an additional confirmation that the limit-ing current is not the hydrogen diffusion current. As mentioned previously, for the potential range 0 to 200 mV in a hydrogen-saturated solution, two com-petitive processes occur: the ionization and the build-up of a passivating oxide film. In the potential range from zero to 20 mV the degree of passivation of the nickel electrode surface, which has previously attained an active condition char-


acterized by a reversible hydrogen potential under no-current conditions, is insignificant; for this reason the hydrogen ionization current is not measurably aifected by passivation. In the 20 to 70 mV potential range the degree of passi-vation increases (though the passivation is still reversible in nature), and this clearly hinders the rise in current value with overvoltage. For potentials above 70 mV the current falls; with increasing irreversibility of the oxidation of the metal surface, this fall is more and more strongly marked. The passivating effect of the oxide film has not as yet been thoroughly studied. For instance it is not known whether this a strictly mechanical blocking of part of the surface with no effect on the electrochemical properties of the remainder, or else whether the adsorbed oxygen acts in a manner analogous to some catalytic poisons by deactivating even the oxygen-free nickel surface regions.

The oxide films on nickel in the potential range where hydrogen ionization occurs are the most probable reason for the present lack of an entirely reliable overvoltage theory of hydrogen ionization reaction on nickel. It is also the reason why experimental results are irreproducible and often show discrepancies. It seems therefore that the development of theoretical models in this field has run well ahead of experimental verification. The elucidation of the mechanism of the ionization reaction of hydrogen on nickel, as on the other iron-group metals, may be possible only when the composition and structure of the oxide films are known in detail, together with their effect on the electrocatalytic prop-erties of the surface of a nickel electrode [91].

The reason for the exceptional resistance to oxide passivation of skeleton nickel, particularly with the addition of various promoters added, also remains unclear.

12.3.4 Ionization of hydrogen on tungsten carbide

Tungsten carbide, WC, has long been recognized and widely used because of its hardness and refractory properties. The bulk properties of this solid material are fairly well known. The tungsten atoms make up a close-packed hexagonal lattice, with the carbon atoms occupying interstitial voids. As with a number of other transition-metal carbides and nitrides, tungsten carbide exhibits a large and typically metallic electrical conductivity [92, 93].

The surface properties of tungsten carbide aroused interest only after the discovery by Boehm and Pohl in 1968 [94] that it can be an efficient electro-catalyst of hydrogen ionization. The discovery was soon confirmed by many others [95-101],


Investigation of the hydrogen ionization reaction on the tungsten carbide surface and the surface properties of this material were carried out both on its powder form [94-106] and on a fairly thin, compact layer obtained by carbi-dizing solid tungsten [107-111]. The tungsten carbide used was prepared by reacting either tungsten, or the products of the reduction of hydrated or anhyd-rous tungsten oxides, with carbon black, carbon monoxide or methane, often in the presence of gaseous hydrogen, at temperatures between 970 and 1770 K. Under these conditions the material obtained is of considerably variable quality, and hence the results provided by various investigators are difficult to compare.

The rest potential of a tungsten carbide electrode (specially of a smooth electrode, in the solid form) in low pH electrolytes ( H 2S 0 4, H 3P 0 4, HCi) satu-rated with gaseous hydrogen under atmospheric pressure departs slightly towards the anodic side (up to approx. 60 mV) from the reversible hydrogen electrode po-tential in the same electrolyte [102, 103, 107, 109, 111, 112]. This deviation seems to decrease when the temperature is increased [103]. For smooth (solid) elec-trodes, the rest potential is highly sensitive to cleanliness and, for smooth elec-trodes, to the tightness of their mounting [112]. Three potential ranges can be distinguished on the anodic polarization curve in a hydrogen-saturated electro-lyte. In the range up to a potential of approx. 300mVR HE the hydrogen ionization rate rapidly increases. In the range from approx. 300 to 500-700 m V R HE the hydrogen ionization rate declines visibly, flattening out at a fairly low level. Simultaneously, above a potential of 300 to 600 m V R HE an additional parallel anodic process is observed, which is referred to as the tungsten carbide corrosion process. The rate of this latter decreases with time [103, 109, 111, 113]. Daring the corrosion process, evolution of a gas on the electrode was observed [103,113]; this gas was identified as C 0 2 [113]. The corrosion process was interpreted as the oxidation of the carbide carbon to carbon dioxide [103]. As a result of this corro-sion process, the rate of the hydrogen ionization reaction in the potential range from zero to 300 m V R HE is henceforth reduced to a permanently low level.

Under certain conditions, however, maintaining a tungsten carbide elec-trode at high anodic potentials (approx. 600 to 700 mVR H E) for a long time (approx. 1 to 2 hours) results in a considerable rise in the hydrogen ionization rate, up to two times or more, in a potential range from zero to 303 m V R HE [95, 103, 110-112]. This process is referred to as the activation process of tung-sten carbide.

For powder specimens, the activation process was found to bring about a roughly 100% increase in the electrical capacitance of the electrode in the region up to 300 mVR H E. In turn, etching the surface with a hot solution of con-


centrated potassium hydroxide for several minutes causes this capacitance to fall substantially; this, however, has no effect on the activity of the electrode for the hydrogen ionization process in acid electrolytes [103]. For smooth elec-trodes, the activation process causes an increase of the capacitance of the elec-trode by 10 to 50 percents [111].

In the potential range up to 300 m V R HE the hydrogen ionization rates remain constant for literally thousands of hours. Carbon monoxide, even if present in considerable quantities in the hydrogen supplying the electrode, has no effect on the activity of the tungsten carbide electrode [102, 103].

The exchange current for the hydrogen ionization reaction on the surface of a tungsten carbide specimen, in sulphuric acid of concentrations from 0.5 to 1 M, is reported as being in the range from 0.5 χ 10"

2 A/m

2 at 303 Κ to approx.

5x 10~2 A/m

2 at 353 Κ (on the assumption of an ideally smooth surface) [107,

109, 112]. In H 2S 0 4, H3PO4, HC1, or HBr the reaction is of first order with respect

to molecular hydrogen. The apparent activation energies of the reaction at a potential of 300 m V R HE in H 2S 0 4, H 3P 0 4 and HC1 were found to be 34, 31, and 54 kJ/mol, respectively [109]. The apparent activation energy of the reaction at zero overpotential in H 2S 0 4 was found to be about 40 kJ/mol [111].

It seems certain that in the potential range below 300 m V R HE the tungsten carbide surface which exhibits a catalytic activity for the hydrogen ionization reaction is to some extent covered with the adsorbed hydrogen. This can be in-ferred from the results of studies carried out by both the electrochemical method of the so-called charging curves [104, 107] and the thermodesorption method [104, 105]. The coverage of the surface by hydrogen in the electrolyte is hard to determine both because of the lack of a "double-layer region" on the poten-tiodynamic curves of tungsten carbide, and the lack of an adequate method of measurement of the actual surface area of the material under these conditions.

A correlation was found between the catalytic activity of specimens in the hydrogen ionization reaction and the extent of hydrogen adsorption in the gas phase [105]. The coverage with hydrogen in the gas phase for the most active samples of tungsten carbide (from the viewpoint of its catalysis of the hydrogen ionization reaction in the electrolyte) was determined to be approx. 20% [114, 115]. The hydrogen adsorption process on tungsten carbide is possibly much slower than on platinum [107]. Many authors believe the tungsten carbide sur-face to be highly inhomogeneous [101, 103, 104, 114]. It was for instance sug-gested that, as a result of the above-mentioned corrosion process, the tungsten carbide surface becomes coated with tungsten oxides which are inactive in the


hydrogen ionization reaction. In the opinion of Benda et al [103], as a result of the so-called activation process, the WC surface becomes covered with tung-sten-hydrogen bronze, H xW 0 3 (bronzic acid). It is experimentally proved that the activation process does not develop considerably the active surface area of the electrode, but it changes the chemical composition of the surface [101,

The hydrogen ionization process on a tungsten carbide surface was initially supposed to follow the Tafel-Volmer mechanism, i.e. to proceed in two following elementary steps :

and the rate determining step was thought to be the step of dissociative chemi-sorption of molecular hydrogen (equation (12.47)) [107, 109]. But more recent works [110,112] suggest that even if this mechanism holds, the overall process rate is also affected by the rate of the charge transfer reaction (equation (12.48)).

It is worth noting that tungsten carbide also catalyses the electro-oxidation of carbon monoxide, hydrazine, formaldehyde, acetaldehyde and formic acid [103, 116]. The tungsten carbide activation process mentioned previously for hydrogen oxidation also increases the rate of hydrazine oxidation [103].

Horanyi et al [117-119] have found that tungsten carbide catalyzes an elec-trode reduction of some inorganic anions. The same authors have observed a catalytic activity of tungsten carbide in liquid phase heterogeneous hydro-génation [120].

Levy and Boudart [115] point out that tungsten carbide exhibits some sur-face properties hitherto observed only in the case of platinum and some plati-num-group metals, such as the ability to catalyze the reaction of oxygen and hydrogen at room temperature, the reduction of tungsten trioxide by hydrogen in the presence of water wapour, and the isomerization of 2,2-dimethylpropane to 2-methylbutene. On the basis of these observations they have suggested that the introduction of carbon into the tungsten lattice results in a modification of the surface electron structure of tungsten such that it becomes similar to the electron structure of a platinum surface. However, a study made by Houston et al [121] has shown this idea to be fallacious. They found the degree of fill-ing of the 5d band in tungsten carbide surface layer to be less than for tungsten, whereas for platinum the 5d band is nearly complete.

The results of the most recent works [101, 105, 110-112, 114] suggest that the catalytic properties towards the hydrogen ionization reaction are displayed

111, 114].

H 2 -> 2 H a ds H a ds -* H

+ + e

(12.47)

(12.48)


solely by those tungsten carbide specimens which exhibit a capacity for chemi-sorption of large quantities of hydrogen, and that both these characteristics are related to the peculiar stoichiometry of the surface and to the energy of the surface bonds of tungsten atoms. Spectroscopic measurements indicate that on the surface of a catalytically active WC specimen, the ratio of the number of W atoms to C atoms is analogous to the corresponding ratio for the surface of bitungsten carbide, W2C. The spectra of bond energy of the surface tungsten atoms for these two materials are clearly different. It should be mentioned, how-ever, that W2C itself exhibits some ability for hydrogen chemisorption [114].

A correlation between the catalytic activity of tungsten carbide, WC, in the hydrogen ionization reaction and the deficiency of carbon on the surface rela-tive to the stoichiometric WC appears to explain to some extent the process of activation of tungsten carbide.

12.4 References

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Chem. Soc , Washington 1969, p. 70.

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[89] Horiuti, J., Sugavara, S., / . Res. Ins. Catalysis Hokaido Univ., 4, 1 (1956). [90] Breiter, M. W., Electrochemical Processes in Fuel Cells, Springer-Verlag, Berlin 1969. [91] Wolf, J. F., Yen, L.S.R., Damianovic, Α., Electrochim. Acta, 26, 409 (1981). [92] Storms, Ε. K., The Refractory Carbides, Academic Press, New York 1967. [93] Toth, L. E., Transition Metal Carbides and Nitrides, Academic Press, New York 1971. [94] Boehm, H., Pohl, F. Α., Wiss. Ber. AEG Telefunken, 41, 46 (1968). [95] Binder, H., Koehling, Α., Kuhn, W., Linder, W., Sandstede, G„ Nature, 224,1299 (1969). [96] Mund, K., Richter, G., von Sturm, F., CITCE Meeting Prague 1970, Ext. Abstr., p. 347. [97] Schultz-Ekloff, G., Baresel, D. , Heidemeyer, J., ibid., p. 350. [98] Armstrong, R. D., Douglas, A. F., Keene, D . E., J. Electrochem. Soc, 118, 568 (1971). [99] Sokolsky, D. V., Palanker, V. Sh, Baybatyrov, Ε. N., Bogdanova, A. E., Elektrokhimiya,

8,1754 (1972). [100] Svata, M., Zabransky, Z., Coll. Czechosl. Chem. Comm., 39, 1015 (1974). [101] Ross, P. N., Jr., Macdonald, J., Stonehart, P., / . Electroanal. Chem., 63, 450 (1975). [102] Pohl, F. Α., Boehm, H., / / / Journées Internat. d'Etude des Piles a Combustible, Presses

Acad. Européennes, Bruxelles 1969, p. 180, 183. [103] von Benda, K., Binder, H., Koehling, Α., Sandstede, G., in: From Electrocatalysis to

Fuel Cells, Ed. G. Sandstede, Univ. Washington Press, Seattle 1972, p. 87. [104] Palanker, V. Sh., Baybatyrov, E .N. , Sokolsky, D .V. , Electrochim. Acta, 20, 51(1975). [105] Bashtavelova, E., Nikolov, I., Vitanov, T., J. Power Sources, 7, 257 (1982). [106] Nikolova, V. Nikolov, L, Vitanov, T., Yotova, L., ibid., 12, 1 (1984). [107] Boehm, H., Electrochim. Acta, 15, 1273 (1970). [108] Baybatyrov, Ε. N., Palanker, V. Sh., Sokolsky, D. V., Elektrokhimiya, 10, 158 (1974). [109] Sokolsky, D. V., Palanker, V. Sh., Baybatyrov, Ε. N., Electrochim. Acta, 20, 71 (1975). [110] Zoltowski P., 31st ISE Meet., Ext. Abstr., Venice 1980, p. 326. [111] Zoltowski, P., Electrochim. Acta, 31, 103 (1986). [112] Zoltowski, P., ibid., 25, 1547 (1980). [113] Voorhies, J. D., / . Electrochem. Soc, 119, 219 (1972). [114] Ross, P. N., Stonehart, P., / . Catalysis, 39, 298 (1975). [115] Levy, R. B., Boudart, M., Science, 181, 547 (1973). [116] Schultz-Ekloff, G., Baresel, D. , / . Electroanal. Chem., 35, 73 (1972). [117] Horanyi, G., Vertes, G., / . Electroanal. Chem., 63, 359 (1975). [118] Horanyi, G., Rizmayer, E. M., ibid., 70, 377 (1976). [119] Horanyi, G., Rizmayer, E. M., ibid., 132, 119 (1982). [120] Horanyi, G., Rizmayer, E. M., React. Kinet. Catal. Lett., 13, 21 (1980). [121] Houston, J. E., Laramore, G. E., Park, R. L., Science, 185, 258 (1974).

Added in proof

1. Appleby, A. J., Chemla, M., Kita, H., Bronoël, G., "Hydrogen", in: Encyclopedia ofElectro» chemistry of the Elements, Ed. A. J. Bard, Vol. 9, Part A. Marcel Dekker, New York, 1982.

2. Bai, L., Harrington, D. Α., Conway, Β. E., Electrochim. Acta, 32, 1713 (1987). 3. Clerc, C , Landolt, D. , Electrochim. Acta, 33, 859 (1988). 4. Zoltowski, P., / . Electroanal. Chem., 260, 269 (1989); ibid., 260, 287 (1989).

13. Effect of the Oxide and Sulphide Inclusions on the Absorption and Permeation of Hydrogen in Steel

13.1 Introduction

The destructive effect of hydrogen on very many types of steel is generally re-cognized. It is generally agreed that the fewer impurities are present in steel, the less hydrogen is taken up by this metal. This is particularly significant in such processes as electroplating and the chemical treatment of steel (e.g. pickling in mineral acid solutions), and also in the formation of hydrogen impurity as the result of secondary reactions (e.g. thermal decomposition of hydrated salts and reactions involving gases that contain hydrogen compounds). The conse-quences of hydrogen in steel are many and varied; e.g. impaired strength, snow-flake formation, hydrogen embrittlement, delayed fracture, or damaged pro-tective coatings and finishes. Hydrogen absorption, occlusion, and diffusion occur in three main categories: macroscopically (e.g. at discontinuities in the structure and inclusions which are visible at low magnification), microscopically (fine cracks, crevices, micro-cracks and inclusions approx. 15 μηι in diameter), and in the submicroscopic domain (lattice defects, dislocation accumulations, or voids in the crystal lattice).

In this chapter we are concerned primarily with the effect of inclusions in the macroscopic domain. From the physicochemical standpoint the effects involved are extremely complex, so that often the effects of macro- and microstructural factors are opposite.

Hydrogen is present in liquid steel in varying proportions, because of the different reactions occurring during steel melting in individual manufacturing processes (open-hearth, Bessemer and the electric-arc furnace process) and related factors such as slag composition, refining rate, furnace gas composition, and the moisture content of the furnace lining. The concentration of hydrogen in liquid iron exceeds the hydrogen solubility in finished steel products as the result of the temperature decreasing and of the phase transition of iron (y-iron is formed

FORMATION OF OXIDE INCLUSIONS IN STEEL 375

upon solidification, followed by conversion to α-iron, which is accompanied by a decrease in H 2 solubility). The associated effects are treated at a greater length in the preceding chapters and in a comprehensive monograph by Smia-lowski [1]. In this chapter a closer look is taken at the effect of oxide and sul-phide inclusions present in carbon steel, and on the absorption of hydrogen and its permeability through steel containing such inclusions in varying quan-tities. The effects are usually studied by using specimens (sheet metal, thin films, rods or wires) which are electrolytically (cathodically) charged with hydrogen in acid, neutral or alkaline solutions with or without the addition of hydrogen entry promoters (usually arsenic or sulphur compounds). Under these conditions much important basic research on the behaviour of hydrogen in steel has been carried out (e.g. studies of diffusion, the formation of hydrides as a separate phase, and the effect of alloy additions on hydrogen absorption, permeation and solubility). Apart from this approach, which is experimentally convenient, a considerable amount of work in the field of hydrogen in steel has been concerned with the determination of the effect of gaseous hydrogen, especially at elevated temperatures and pressures, on the properties of steel, such as the finding by Raczynski and Smialowski [2] that the hydrogen absorption vs. degree of cold work relation is strongly affected by distribution of the impurities. This chapter will deal with the origin of impurities in the form of inclusions in steel, and the effect of such inclusions on the absorption and permeation of hydrogen in steel.

13.2 Formation of oxide and sulphide inclusions in steel

The notion of "pure steel", i.e. one that is free from non-metallic impurities, is an abstract concept. According to Kiessling [3], in a 3-tonne commercial steel ingot there are 3 χ 10

13 oxide inclusions of varying size; and one gramme

of steel with a total sulphur concentration of 30 ppm contains 6 χ 1017 sulphur

atoms [4]. Such quantities are closely related to the steel melting process used. It is a well known fact that one of the major metallurgical treatments is the re-moval from the liquid melt of dissolved oxygen and excess sulphur, phosphorus and sometimes nitrogen. This treatment is generally accomplished by adding appropriate elements with a greater oxygen affinity than iron, to the liquid metal. Silicon (ferro-silicon), aluminium, and titanium are widely used as deoxi-dants in melts which at the same time contain other elements (alloy additions), both metallic and non-metallic, such as Mn, Cr, S, and P. The method is known as deoxidation by precipitation, as the solid or semi-solid reaction products

376 EFFECT OF THE OXIDE AND SULPHIDE INCLUSIONS

resulting from the interaction of the dissolved oxygen (as FeO [5]) and the impur-ities, which come up to a slag on the surface of the bath.

The deoxidation process itself is complex. Its success depends not only on the oxygen affinity of the deoxidant elements, but also on the reaction progress and its kinetics. Because various phases are involved (liquid metal, the deoxidant additions that occur in the solid phase, and the deoxidant oxides produced) the reactions involved are heterogeneous in character, and the reac-tion rate is controlled by the slowest step. Consequently, a treatment of the deoxidation process ought to involve, at the outset, a consideration of its indi-vidual stages, so that tentative conclusions could be drawn as regards the compo-sition of inclusions and their distribution in a particular steel product. Such an approach, however, is hardly practicable in view of the lack of basic data related to the partial processes.

By assuming a number of simplifying steps and very fast reactions between the deoxidants and liquid steel, the following four major physicochemical pro-cesses are to be considered :

(a) distribution or separation of phases, i.e. production of oxides; (b) coalescence of small particles into larger ones; (c) adsorption of non-metallic impurities (e.g. sulphides) on the surface of

the oxide "nuclei" formed; (d) flotation of non-metallic particles onto the surface of the melt for in-

clusion in the slag that covers the steel surface. Whereas basic research on the deoxidant-dissolved oxygen equilibrium has

already provided much valuable material, the phenomena related to phase separ-ation (metal-insoluble non-metallic particles) and to the conditions for coaeles-cence and flotation of the particles formed, are much less understood. Their importance however is considerable. The slag particles that come from the fur-nace, or the particles of refractories entrained by steel on its way from ladle to the mould, usually represent only a minor part of the total impurities in the finished steel. The steel purity therefore depends primarily on those products of the deoxidizing process which fail to float and remain instead in suspension. Research on the production of high purity steel is considerably hampered by the elevated temperature (approx. 1870 K) at which measurements must be made. As a result, a number of the data for the physicochemical properties of inclu-sions such as surface tension, viscosity, surface tension at the metal interface, etc are associated with considerable error.

Furthermore, the rate of flotatoin of inclusions and thus the degree of contamination of the product depends to a large extent on their chemical compo-

REVIEW OF TYPE OF INCLUSIONS IN STEEL 377

sition. Since inclusions which differ in their chemical composition may be pro-duced during addition and solution of oxidants, the flotation rate must also be different. The inclusions remaining in steel will therefore differ substantially in chemical composition from the average composition of the oxides produced as a result of reaction with the deoxidants.

Whereas the iron-oxygen-deoxidant system is generally fairly well under-stood, the iron-sulphur-deoxidant one is much more complex and the picture is less clear. Although sulphur dissolves in liquid steel, its solubility in solid steel is negligible; hence, on the solidification of steel, solid or liquid sulphides separate. The presence of sulphides in steel, with the exception of free-machining steels, is generally undesirable in view of a detrimental effect of these inclusions on the mechanical properties of steel. The presence of iron sulphide is particu-larly undesirable (the hot brittleness effect), because of the formation of a low-melting point eutectic (1261 K) which surrounds the metal grains at hot rolling temperatures (approx. 1270 K) or during ingot annealing, and results in the cracking of rolled sheet steel. Manganese, which forms manganese sulphide of a much higher melting point (1883 K), is commonly added to overcome this problem. The occurrence of pure MnS in commercial steel is however fairly infrequent. This is because in liquid iron, containing sulphur and oxygen, complex compounds are produced such as MnS-FeS-FeO-Si02. These com-pounds tend to accumulate on the oxide nuclei (i.e. so-called primary oxide inclusions). As a result, heterogeneous structures occur in solidified steel inclu-sions, whose external layers are usually sulphide rich.

The complexity of the processes occurring during steel melting and solidi-fying, together with the range of subsequent treatments (rolling, forging, etc.), leads to varying degrees of dispersion of the oxide and sulphide inclusions in the steel. The distribution of such inclusions in any given steel is non-uniform, being generally higher in the bulk of the metal than on its surface. However, even on the surface the impurities may reach high concentrations. Consequently, the presence, quantity and composition of oxide and sulphide inclusions has an unquestionable influence on the susceptibility of steel to absorb hydrogen, and also on corrosion processes (general, pitting, or structural corrosion etc.).

13.3 Review of the methods for determining the quantity and type of inclusions in steel

Due to the widely recognized detrimental effect of inclusions on steel properties, the methods for qualitative and quantitative determination of inclusions have been the subject of numerous studies over the last fifty years. Chemical methods


were originally used, commencing with the solution of steel chips in inorganic acids (HC1, H 2S 0 4, H N 0 3, H 3P 0 4) . Thus the major part of the material ma-trix goes into solution, whereas the remaining residue, a mixture of insol-uble carbides and inclusions, primarily oxides and nitrides, was analysed fol-lowing the removal of carbides. These elementary chemical methods provided very general, and often erroneous, information, of little value for the purposes of physical metallurgy. A major step forward in this field has been the electrolytic extraction procedure, the now standard method involving the selective disso-lution of the matrix in mainly neutral electrolytes containing iron-complexing agents.

Cells of special design are commonly used, in which the anode space is separ-ated from the cathode space by a diaphragm (e.g. cellophane or collodion)

Fig. 13.1. A schematic diagram of the electrolytic cell for isolation of the oxide and sulphide inclusions: 7—cover, 2—terminals for current supply, 5—an 18/8 grade steel cathode, 4—spe-cimen, 5—diaphragm, 6—funnel, 7—anode space, top section, $—•cathode space, 9—anode space, bottom section, 10—stopcock.

REVIEW OF TYPE OF INCLUSIONS IN STEEL 379

and the anolyte is typically a solution of sodium citrate with added salts (e.g. KC1, NaCl, KCNS, KBr). The chemical composition of the anolyte and its pH are matched to the type of steel being examined. For example, for separation of sulphides a 5% sodium citrate solution is used with a 2% addition of NaCl at pH 7, whereas for separation of the oxide inclusions in bearing steel, a 5% solution of sodium citrate with 1% additions of KBr, KSCN, and NaCl. Figure 13.1 shows schematically a typical electrolysis cell for the isolation of carbides and inclusions [6].

A 5% sodium citrate solution containing some HC1 to make the pH about 5 is typically used as catholyte. After removal of the insoluble residue (isolate) from the specimen (either mechanically or in an ultrasonic field), the specimen surface should be free from pits, and the weight loss of the specimen should be uniform all over its surface. Metal dissolution should proceed under active conditions; consequently, for passivation-prone steels the electrolysis should be carried out in deoxygenated anolytes, under an inert atmosphere (pure nitrogen or argon), and with the use of a potentiostat which automatically controls the current conditions of the electrolysis process. Other electrolytes used include FeS04 solutions with 0.2% Seignette salt added, and a 5% KBr solution with an addition of 1% KI, etc. [7-9]. The citrate electrolytes, however, are the most widely used.

The electrolysis process should be carried out under conditions that ensure the dissolution of the anode in a fresh solution. For this reason former methods involved recirculation of the anolyte, while today use is made of a differ-ence in specific density between (a) the solution saturated with iron salts and salts of other metals produced from the dissolved matrix and (b) fresh ano-lyte, by placing a funnel in the cell directly under the anode. A self-acting pump is thus formed and the used anolyte (i.e. saturated with salts of the matrix me-tals) falls to the cell bottom, being replaced by fresh solution around the anode. More detailed information is to be found in the monographs by Klinger and Koch [7], and by Koch [8] and in numerous original papers, e.g. Fitterer [9], Treje and Benedicks [10], Heiskanen [11], Foryst et al [6], or Koch and Sun-derman [12]. Following the anodic dissolution the isolate is analysed to de-termine the content of carbides, sulphides, and oxides. From this fairly complex procedure the concentration in the steel and the average chemical composition of the oxide and sulphide inclusions are successfully determined by microchemical methods. The method is laborious, and it requires considerable care and accu-racy in working up the isolate (removal from the anode surface, washing off traces of electrolyte, removal of carbides, microchemical analyses, etc.).


A further substantial improvement in the determination of composition of the inclusions has been the application of the X-ray microprobe [13-15] as the standard technique for the precise determination of the content of the major-ity of the elements in a particular inclusion. Studies made with the help of this device have clearly demonstrated the heterogeneity of the composition of the oxide inclusions, and inclusions of sulphides, phosphides and carbides, informa-tion which is unattainable by other methods.

Nowadays the X-ray microprobe (ΕΡΜΑ, SEM) is a commonplace instru-ment in the laboratories of major metallurgical plants and research institutes, and the results obtained have furnished much valuable information on inclu-sions. As a result of the introduction of such in situ procedures, the electrochemi-cal methods for the isolation of inclusions and the microchemical procedure for composition determination have been out of favour for some time. In recent years however there has been a renewed interest in the elctrolytic isolation pro-cedure, primarily for the following two reasons:

1. The cost of a high performance X-ray microprobe capable of covering a wide range of chemical elements including oxygen, carbon and nitrogen is fairly high;

2. The X-ray microprobe is suitable in the main for the examination of indi-vidual inclusions and it requires special arrangements for the analysis of a spe-cimen. Routine studies of inclusions (e.g. typically with the use of cross sections of five microsections of a specimen of steel) are too laborious to be generally used.

Another device that is helpful in quantitative metallurgical analysis aimed at inclusion studies is the Quantimet TV microscope. This gives rapid and auto-matic determination of the average surface area occupied by inclusions over the microsection together with their size distribution. A very important step forward in the study of inclusions was the improvement made in the method for determination of total (i.e. dissolved and bound) oxygen. The basic method consists in reducing the oxides of a specimen molten in a graphite crucible under vacuum. The improvement in the apparatus was mainly concerned with a suitable choice of the reduction process variables, followed by oxidation of the initially produced CO to C 0 2 and a precise micro-scale determination of the carbon dioxide (e.g. the devices by Strôhlein in 1960 and Balzers in 1965, or, more re-cently, by Leco).

Such apparatus is widely used today for the estimation of oxygen in steel during the melting process and also in rolled steel, in the laboratories of major ironworks and metallurgical research institutes. An equally important tech-

EFFECT OF IMPURITY CONTENT OF IRON 381

nique, used primarily in steel mills for on-line oxygen analysis, and thus also for determining the amount of inclusions, is the oxygen actimeter [16, 17] which operates on the principle of continuous measurement of the EMF of the con-centration cell using solid electrolytes of variable composition (MgO, stabilized Z r 0 2 + CaO, T h 0 2 + Y 20 3 + CaO).

Adequate information can often be obtained by quantitative metallography, especially using the stereological method. This is particularly useful for deter-mination of carbide content.

All the methods discussed above are in general use today and, as pointed out previously, they require expensive equipment backed up by suitably trained personnel. The major drawback to the instrumental methods is that a zonal determination of the inclusions content and composition cannot be made over the entire cross section of the specimen examined, nor can their average chemical composition be established. Only recently has this information become available by means of the SIMS technique. As pointed out in Section 13.1, the distribution of inclusions is generally uneven, the bulk of the specimen contains more inclu-sions than the surface. Both because of this and in view of the considerable cost of sophisticated apparatus, there has recently been a renewed interest in the method of electrolytic isolation, followed by microanalysis, where the main objective is to find the average content of metal ions bounded with oxygen or sulphur. This is a very valuable tool that supplements instrumental methods. It is often highly useful in studying the effect of the inclusions content on the properties of the end steel products (strength or corrosion properties, linked to the level of rejects, etc.).

In conclusion, it should be noted that the problem of the effect of inclusions on the specific properties of steel products needs to be tackled from a statisti-cal standpoint. This is due to the fact that the inclusions, which differ substan-tially in chemical composition, are present in considerable quantity.

13.4 Effect of impurity content of iron on the absorption of hydrogen

The effect of the degree of contamination of iron or steel on hydrogen absorp-tion has been the subject of few investigations. Much more attention has been given to the analogous effect of sulphide inclusions on the pitting corrosion of steel [18].

Studies on the solubility and kinetics of hydrogen absorption in metals have been relatively abundant [19-21] but they have been concerned mainly with the effect of the content of major alloying components of steel. More comprehensive


studies on the effect of inclusions on hydrogen absorption and permeation were carried out by Ciszewski et al. [22], Hewitt and Murray [23] and by Brooksband and Andrews [24]. The studies show that hydrogen-included cracks develop most readily at inclusions of the type 3A1203 * 2Si02 or S i 0 2, due to the much lower expansion coefficient than martensitic steel. In the case of MnS-rich inclusions with a much higher dilatability, internal stresses do not appear but narrow crevices develop at the interface of the inclusions, from which gaseous hydrogen can evolve.

As mentioned earlier, the majority of the studies on the effect of hydrogen on steel properties have been made by taking steel samples of varying compo-sition and shape saturating them with hydrogen in electrolytes with addition of hydrogen entry promoters, and recording the relation between the effect meas-ured and the specimen charging time. Specifically, Foryst and Smialowski [25] investigated the effect of chemical composition of steel (in particular the overall content of oxide inclusions) on its sensitivity to hydrogen charging.

The property studied was the elongation of Armco iron rods as the function of charging time. It was shown that a more contaminated iron, when charged with hydrogen at approx. 293 K, elongates over three times more than killed iron (Fig. 13.2). The authors ascribed the difference in elongation to a larger number of hydrogen traps in the unkilled steel which are responsible for the for-mation of hydrogen blisters and crevices in the surface layers of rods. Prior to this, Foryst and Harazin [26] dealt with the determination of the hydrogen ab-

2

ο 20 40 60 i

Polarization time, min 80 100

Fig. 13.2. Elongation curves for the Armco iron specimens for unkilled (curve i ) or killed

(curve 2) sorts during cathodic polarization.

EFFECT OF IMPURITY CONTENT OF IRON 383

sorbed in steel and iron during cathodic hydrogen charging in 0.5 M H 2S 0 4 with A s 20 3 added as a promoter. In this study specimens of two kinds were used, namely rods made of pure electrolytic iron containing other metals in trace amounts and very little sulphur (0.001% S), yet with a substantial amount of oxygen bound nearly exclusively as FeO (0.0825% 0 2) , and Armco iron rods with typical contamination by other elements (Mn, Si, C, S) but containing about half as much oxygen bound in the form of various oxides (0.040% 0 2) and more sulphides (0.026% S). The determination of hydrogen was performed during the process of cathodic hydrogen charging, and the arsenic correction was accounted for, i.e. the reduction to AsH3 of the arsenic that deposits simulta-neously with hydrogen on the cathode (evolution of AsH3). It was demonstra-ted that in pure iron containing more oxygen bound as FeO, the quantity of hydrogen absorbed is over twice as high as in the Armco iron, despite the fact that the latter contained much more sulphur (Fig. 13.3). It therefore seems that

60

0 AO 80 120 160 200 240 Polarization time, min

Fig. 13.3. The volume of the absorbed hydrogen in pure iron and Armco iron cathodes (after [26]).

it is the total oxygen bound in iron in the form of iron oxide inclusions that is largely responsible for the quantity of the hydrogen absorbed. In the paper men-tioned no attention was paid to the relation between the quantity of hydrogen absorbed and the resulting defects. For comparative reasons, however, the quantity of hydrogen was measured in specimens of a high-carbon steel (0.83% C, 0.50% Mn, 0.22% Si) with a high tendency to delayed fracture, which contained approx. 0.01% bound oxygen. This steel, compared with pure iron, absorbed

384 EFFECT OF THE OXIDE A N D SULPHIDE INCLUSIONS

relatively little hydrogen (7.7 cm3/0.1 kg Fe, as against 56.1 cm

3/0.1 kg Fe for

pure iron). It follows clearly from the foregoing that the amount of the hydrogen ab-

sorbed depends on the degree of contamination of the iron (steel) with inclusions; even though the consequences of the presence of inclusions in iron in a quantity above a specified level seem to be unrelated to the quantity of hydrogen ab-sorbed.

13.5 Permeation of hydrogen through a low-carbon steel with variable impurity content

The first observations on permeation of hydrogen through steel were reported by Cailletet as early as 1868 [27], More detailed studies of the phenomenon came much later [28, 29], using membranes of various metals including iron and steel. Results of these studies are comprehensively covered in the monograph by Smia-lowski [1], and in the review by Riecke [30].

In recent years the problem has attracted increasing attention for practical reasons connected with the manufacturing flaws which develop in protective coatings, in particular in ceramic coatings (enamels). The enamel baking process is known to entail the decomposition of the hydrated enamelling compounds. At the working temperatures above 1073 K, atomic hydrogen is formed in the decomposition of the hydrated compounds and the water that comes from the enamelling mass, and it enters the steel substrate. Under the specific conditions of enamel finishing and also in hot-dip galvanizing and tinning, this hydrogen is later released from the steel substrate at elevated temperatures, and results in substantial stresses in the coating which gradually becomes detached from the finished piece giving rise to the fish scale effect. For this reason a consider-able number of studies have appeared in recent years, aimed at determining the effect of steel composition (usually for cold-rolled deep-drawn steels) on hydrogen absorption and permeation [31-33]. On the assumption that the re-sidence time of hydrogen in coat-protected steel is highly critical, the authors of these papers determined the so-called breakthrough time for hydrogen passing through steel membranes, using electrolytic cells of relatively simple design. The conclusion is that the shorter the time lag, the more suitable is the specific steel for hot surface coating. Generally, the hydrogen permeation rate is roughly inversely proportional to the quantity of hydrogen traps, of which oxide and sulphide inclusions are a significant part, since clearly, the more traps are there in steel, the more permanently they bind molecular hydrogen.

PERMEATION OF HYDROGEN 385

- 6 2 5 w A

^ 110

Fig. 13.4. A schematic view of the cell for the determination of time lag (after [34]): 1—body, 2—cover, 3—fastening nut, 4—gasket, 5—cooler, 6—anode, 7—supply pipe, 8—capillary, 9-—measuring capillary, 10, 11—electrical connections, 12—epoxy glue, 13—connecting tubing, 14—specimen, 75—the ultrathermostated jacket, 16—graduated scale.

Polarization time min

Fig. 13.5. The volume of hydrogen diffused through steel membranes vs. cathodic polarization time (after [36]): specimen /—suitable for enamel finishing (time lag approx. 20 min), specimen 2—unsuitable for enamel finishing (virtually no hydrogen flow).


The high temperatures of the enamel baking process (the steel is first sprayed with a primer coat, also baked at 1073 K) naturally cause considerable stresses to be set up in the steel matrix. This results in the coating gradually scaling away from the enamel-finished material during storage or normal use. Considering the above, Foryst and Przybylo [34] carried out measurements on the time lag [35] for hydrogen permeation through cold-rolled unkilled steel membranes cut along the direction of rolling using a number of steels of variable composition and impurity content. The authors assumed that the shorter the time lag, the less traps are present in the steel. An electrolytic cell of special shape (Fig. 13.4) and commercial steel of a carbon content of 0.04 to 0.08% were used in the study. On the basis of the results obtained, the steel sheets sampled from individual melts were classified as [36] : totally suitable for enamel finishing (short time lag), of medium suitability (a time lag of 20 min) and unsuitable (virtually no hydro-gen getting through membranes) (Fig. 13.5). In the measurements the mem-branes were charged with hydrogen in 12% H 2S 0 4 at 313 Κ without addition of hydrogen entry promoters. Commercial-scale enamel-finish testing was per-formed on the steels studied, paying particular attention to the occurrence of typical enamel-finish defects. An agreement was found between the classification of the laboratory-tested steel sheets and the results of the commercial-scale testing. Other findings in the study are that the total sulphur in steel should not exceed 0.025%, and that the combined oxygen and sulphur should stay below 0.07% [36].

It should be noted that in order to obtain information on the effect of hydro-gen on more purely scientific properties (e.g. determination of hydrogen diffu-sivity in steel or iron of a variable and known impurity content) the electrolytic cells are of a much more complex design [30]. This is largely to achieve a more uniform hydrogen charging and to determine the hydrogen effusion rate with a high precision [37-39].

As well as the overall sulphur content of unkilled steel, another significant factor in the permeation of hydrogen appears to be [34, 36] the distribution of sulphide inclusions in a given specimen. To investigate this effect, a method was developed to determine the zonal distribution of total combined sulphur, fol-lowed by a topogram taken from the specimen cross section. It was found that in sheets where the concentration of combined sulphur varies considerably (if, for instance, the sulphide content in the middle of the cross-section is ten times as high as near the surface), the steady-state flow time becomes substantially re-duced. The number of the measurements required to achieve an unambiguous estimation for a given temperature should be sufficiently large (at least ten).

REFERENCES 387

In conclusion, it must be said that the effect of inclusion content on the mechanism and kinetics of hydrogen permeation and absorption is not yet suf-ficiently understood. Future studies aimed at clasification of this topic will need to take into account the quantity, chemical and phase composition of inclusions in steel, as well as their distribution and their effect on mechanical properties and related effects (e.g. evolution rate of the absorbed hydrogen, total fixed hydrogen, internal stresses etc.).

13.6 References

[1] Smialowski, M., Hydrogen in Steel (in Polish), PWNT, Warszawa 1961. [2] Raczynski, W., Smialowski, M., Bull Acad. Polon. Sci. sér. sci. chim. 8, 209 (1960). [3] Kiessling, R., Non-metallic Inclusions in Steel, Sw. Inst. Met. Res. JSJ, London 1968, publ.,

115, p. 6. [4] Kiessling, R., Sulphide Inclusions in Steel Proc. Intern. Symp. No. 4,1974, ASM No. 6,

p. 104. [5] Krupkowski, Α., Ptak, W., Arch. Hutn., 6, 265 (1961). [6] Foryst, J., Mchedlishvili, V., Orzechowska, J., Prace IH, 17, 35 (1965). [7] Klinger, P., Koch, W., Beitrage zur Metallkundliche Analyse, Verl. Stahleisen, Dusseldorf

1949. [8] Koch, W., Metallkundliche Analyse, Verl. Stahleisen, Dusseldorf 1965. [9] Fitterer, G. R., Trans. AIME, Iron Steel Div., 95, 196 (1931).

[10] Treje, R., Benedicks, C , Jernkont Ann., 116, 165 (1932). [11] Heiskanen, S., Jernkont Ann., 139, 78 (1955). [12] Koch, W., Sundermann, H., Radex Rund., 679, (1957). [13] Casting, P., Z. Anal. Chem., 192, 51 (1963). [14] Cosselett, V. E., Duncop, P., Nature, 177, 1172 (1956). [15] Kiessling, R., Lange, Ν., Non-metallic Inclusions in Steel, Sw. Inst. Met. Res. JSJ, London

1964, publ., 90, p. 5. [16] Pargeter, J. K., Faurschou, D. K., / . Met., 21, 46 (1969). [17] Gatellier, C. et al, JRSJD, com. 303, May 1975, p. 1. [18] Janik-Czachor, M., Szummer, Α., Szklarska-Smiaiowska, Ζ., Brit. Corr., J., 2, 90 (1972). [19] Karnaukhov, M. N., Morozov, A. N., Izv. Akad. Nauk SSSR, Otd. Tekh. Nauk, 12, 1845

(1948). [20] Baukloh, W., Muller, R., Arch. Eisenhiittenw., 11, 509 (1938). [21] Luckenmayer-Hasse, L., Schenck, H., Arch. Eisenhiittenw., 6, 1861 (1932). [22] Ciszewski, Α., Radomski, Α., Smialowski, M., Stress Corrosion Cracking and Embrittle-

ment, Unieux-Firminy 1973. [23] Hewitt, J., Murray, J. D., ibid., 1965. [24] Brooksband, D., Andrews, K. W., JJS Inst., 206, 206 (1968). [25] Foryst, J., Smialowski, M., Proc. Electrochem. Conf., Warszawa 1955, p. 319. [26] Foryst, J., Harazin, J., Bull. Acad. Polon. Sel, sér. sci. chim., 23, 689 (1975). [27] Cailletet, L., Comp. Rend., 58, 965 (1868).


[28] Bodenstein, Μ. Ζ., Z. Elektrochem., 28, 269 (1922). [29] Edwards, C. Α., JJS Inst., 110, 9 (1924). [30] Riecke, E., Arch. Eisenhuttenw., 49, (11), 509 (1978). [31] Petzold, W., Bergakademie, 7, 343 (1955). [32] Buckel, E., Leonaritis, L., Arch. Eisenhuttenw., 39, 343 (1968). [33] Koning, P., Schmitz, K. H., Thieman, L., Arch. Eisenhuttenw., 40, 991 (1969). [34] Foryst, J., Przybylo, W., Proc. Conference Stress Corrosion and Hydrogen Embrittlement,

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SUBJECT INDEX

Absorption of hydrogen, 227-231, 247-265, 269-273 diagnostic criteria for mechanism, 248, 250 effect of impurities, 381-384 effect of metal dispersion, 230, 231 mechanism of hydrogen entry, 247-253

Acetylen hydrogénation, 237 Activation energy of

hydrogen adsorption, 225 hydrogen diffusion, 280, 281, 315 pit nucleation, 26 precipitation, 197, 201 Snoek-Kôster internal friction peak, 311, 315-316

Active-passive region of potentials determination method, 65, 66 pitting in, 44 stress-corrosion in, 65,66, 75

Adsorption isotherms, 351 Frumkin isotherm, 354, 358 Langmuir isotherm, 353 Temkin isotherm, 340

Adsorption of hydrogen, 223-226 electrochemical adsorption, 352, 360 Gibbs adsorption, 344

Adsorption of hydrogen sulphide, 229 Aggresive ions, 25, 28, 29 Apparent activation energy for stress-corro-

sion-cracking, 70-73 Apparent diffusion coefficient of hydrogen

in iron, 244, 286, 288 Atomic hydrogen, 232-236

Boundary layer, 154 hydrodynamic, 154 diffusion, 154, 155, 163, 216, 261

Brittleness, 293, 294

Brittleness of hydride phase, 326

Carburization of steel, 116-117 Catalytic effect of hydrides, 231-238

catalyst poisoning, 232-236 catalyst activation, 236-238

Cementation, see displacement Charge-transfer reaction, 343 Chloride ions, 10-21, 23, 26, 30, 33, 36, 38,

39, 43,44 concentration in pits, 34-37

Chloride salt layer, 21, 25, 31, 42 Competitive ion adsorption, 24,25 Corrosion cracks

formation, 61-63 growth rate, 63-65

Corrosion equilibria in molten carbonates, 140-146 molten halides, 133-135

effect of oxide ions on, 135 molten sulphates, 146-148

Corrosion rate, 152 control of, 174

activation control, 41,174 diffusion control, 41, 42, 152, 153, 156, 159, 162, 163, 169, 176, 177, 193 kinetics control, 152, 169, 170, 176 mixed kinetics control, 164, 165, 168, 169, 170, 174

Cotrell atmosphere, 313-315, 329 of hydrogen, 315, 329

Decarburization of steel, 116, 117 Delayed failure, 293, 294 Dislocation structure, 276, 298, 306, 329, 330,

332,333 effect of hydrogen on, 308-310, 334

390 SUBJECT INDEX

Dislocations

as traps for hydrogen, 241, 275, 276, 287-289 binding energy with hydrogen, 275, 276, 311, 315,316 cross-slip, effect of hydrogen, 296, 308, 319 double kink formation on, 316

Displacement (cementation), 180-215 equilibria, 182-192 kinetics, 192-212

diffusion control, 198,200,203,206, 207 mixed control, 198

rate, 200,201 Dissociative chemisorption, 261,262, 343

of promoting hydrides, 261,262 Dissolution of metal in liquid metals, 106-111,

118 rate of, 107, 110

Double-layer theory, 343, 344

Electrocatalytic activity of Ni, 363 of WC, 368-370

Electron work function, 224, 229 Ellipsometric study, 15, 17 Entry of hydrogen from aqueous solutions

environmental factors, 240 sources of hydrogen, 240 surface effects, 252

Ethylene hydrogénation, 235-237

Flow stress, effect of hydrogen on, 304, 305, 307, 308

Forms of hydrogen absorbed, 227, 228, 241, 242, 268, 270

Grain boundary

as pit nucleation sites, 12,13 effect on hydrogen solubility, 271 role in corrosion in liquid metals, 114,115, 117, 120

role in stress corrosion cracking, 61-63, 80, 82, 85-87

Hardening of iron and steel by hydrogen, 304-308

Hardening-softening phenomena, 302-308 Hematite, 16, 76 "Hot corrosion", 146, 148 "Hydride" hydrogen, 227, 228, 232 Hydride phase

decomposition of, 324-332 formation of, 227-231, 326, 328, 330 in austenitic steels, 333, 334 nuclei of, 330, 335

Hydrides

Ni hydrides, 230, 234-236, 323-332 Pd hydrides, 228, 231-233, 236-238, 332, 333 promoting hydrides, 256

equilibrium potentials, 256 Hydrogen adatoms, 223,224, 229 Hydrogen adsorption, 223-226, 229

chemical adsorption, 224,225 activation energy of, 225

physical adsorption, 224, 226

Hydrogen concentration in iron, 242, 267,

274-278, 283-285, 287 critical, 277, 278

Hydrogen depolarization, 152, 161, 168,

173, 176,177

Hydrogen desorption, 225,226, 230, 282,283, 285, 326-328 activation energy, 225, 226 isotherms of, 230 kinetics of, 328

Hydrogen diffusion coefficient in electrolytes, 344

Hydrogen diffusion in Fe, 243, 244, 269, 274, 278-289 activation energy, 280, 281 diffusion coefficient 242-244, 269,270, 274, 278-289

around dislocations, 315 determination methods, 269,270 effect of concentration on, 283-285 effect of surface, 281-283 effect of traps, 285-289

Hydrogen embrittlement, 240, 267, 293, 294

SUBJECT INDEX 391

Hydrogen evaluation reaction, 245-247, 249, 251,253, 339 diagnostic criteria for mechanism, 247 possible mechanism, 246 247, 253 stages, 245,246

Hydrogen fugacity, 262, 272 Hydrogen induced cracks, 273, 294-298, 308

319 initiation, 294-296 propagation, 296-298

Hydrogen ionization reaction, 341 steps of, 342, 350 mechanism, 349-369

Heyrovsky mechanism, 349 Tafel mechanism, 349, 357, 358, 364, 369 Volmer mechanism, 349, 357, 358, 369

Hydrogen recombination, 233-236 Hydrogen Snoek-Kôster internal friction peak,

311-314, 317 effect of hydrogen concentration, 314 effect of other interstitials, 312-314

Hydrogen solubility in iron, 268, 271,272 in nickel, 322

Hydrogénation reactions, 231, 235-238

Inhibition of hydrogen entry, 263-265 Inhibition of corrosion in liquid metals, 120,

121 Impedance

method, 355 faradaic impedance, 356

Internal friction background, effect of hydrogen, 300, 311, 317-319 method, 310,311 method to measure hydrogen diffusivity, 270

method to study hydrogen-dislocation interaction, 311-318

Ionic fraction, 130 Ionic liquids, 125-129

properties of, 127

Kinetics components, 165

Latent limiting current, 162

Levich theory (Levich equations), 155, 163, 171, 199, 201,216

Limiting current, 257, 345-347, 364, 365 Limiting diffusion flux, 159, 160, 168 Limiting oxygen flux, 157, 161 Liquid metals, 96-98, 105-118, 120, 121

properties, 97, 98 Localized corrosion, 9, 63, 114-116

Magnetite, 16, 59, 76, 77, 80, 81 Mass transfer, 107, 110, 115-117 Microcracking due to hydride decomposition,

324-326, 331, 334 Mixed potential, 188, 190, 212, 214

effect on displacement, 188 Model displacement cell, 209, 216 Molten alkali halides, 125 Molten Carbonate Fuel Cells (MCFC), 140,

143, 144

Nernst equation, 155, 182, 185, 187 Nernst potential, 185 Nuclear reactor fuel as a corrosive agent,

117, 118

Occluded cell, 36 Oxide film, 10, 14, 18, 21, 23, 25, 39, 70, 75-

77, 89-91, 111, 112, 116 effect in stress corrosion cracking, 70, 75-77, 89-91

Oxide inclusions, 12, 374-378 as pit nucleation sites, 12 determination method, 378-381 effect on hydrogen absorption, 382, 383 overall content of, 382

Oxides, 12, 17, 36, 59, 60, 111, 115, 118, 120 Oxygen depolarization, 152, 158, 162, 168,

173, 176,177

Para-hydrogen conversion, 232, 235, 236 Partial molar volume of hydrogen in iron,

273-275 Passivation, 41, 42, 345, 347, 365

current of, 41

392 SUBJECT INDEX

Passive film, 10, 14,15, 23, 29, 30, 33, 39, 45 breakdown of, 15, 19,20, 26-30, 44 defect sites in, 23, 27, 29, 30 growth kinetics, 15, 17, 19

logarithmic, 15, 17, 19 parabolic, 15, 17, 19

thickness of, 14-19 Permeation of hydrogen through membrane

effect of cathodic current, 248-251 effect of electrode potential, 248, 255, 256, 258, 259

effect of hydrogen gas exitation by electric discharge, 229 effect of inclusions, 374, 384 effect of promoters, 249, 252, 255, 258 electrochemical method, 242-245

PH at electrode surface, 252-254, 257 changes with cathodic current density, 253 effect on corrosion rate, 161, 167, 169 effect on H2X , HX~ and S

2~ distribution,

258,259 effect on pit formation, 28 effect on promoter efficiency, 252, 257

Pit growth, 31-35, 41 Pit morphology, 32, 37-41 Pit nucleation

critical potential, 20, 22, 24-26 induction time, 18, 19, 24-26 sites of

carbide, 12,13 grain boundary, 12, 13 Fe sulphide, 10, 11 Mn sulphide, 10-14, 20,21, 30, 44 mixed sulphide-oxides, 31 mixed sulphides, 10, 11, 31 surface inhomogeneity, 9,10, 13 Ti sulphide, 12

Pitting in borates, 44 chloride containing solutions, 9, 11, 18-23, 25, 28-32, 42-45, liquid metals, 113 methanol, 44 silicates, 44

sulphate solutions, 38, 44 water, 44

Precipitated metal, 184 effect on displacement process, 187 effect on displacement kinetics, 204-212 morphology, 187, 202, 209, 211, 216

Promoters of hydrogen entry, 229, 230, 234, 240, 250, 251, 254-262 active chemical form, 254-260 effect of potential on efficiency, 252-260 mechanism of promoter action, 261-263

Radiation effect on metal mechanical prop-erties, 119,120

Rate determining steps for hydrogen evol-ution, 245, 246

Redox potential, 133, 134, 136, 139, 147 Redox system, 17 Reduced ions, 184-186 Repassivation, 27, 28, 30, 31

rate of, 28, 30 Rotating disc, 152, 155, 156, 159, 163, 165,

170, 172, 175, 177, 199-202, 216, 251

Secondary reaction, 186, 187, 215 Semi-immersed electrode, 348 Shear moduli of iron, 298

effect of hydrogen on, 298-300 Snoek-Koster relaxation, 314-317 Softening of iron and steel by hydrogen,

304-308 Solubility of iron in liquid metals, 106-109 Solubility of oxygen in liquid metals, 120,121 Standard potential, 130, 131, 134, 135, 183

of metallic electrodes in molten salts, 144 Stress intensity factor, 56 Stress corrosion cracking, 50, 52-92

effect of cold work, 87, 88 heat treatment, 84-87 mechanical stress, 88,89 potential, 74,75 steel composition, 77-84 temperature, 69-73

environments, 36, 52, 55, 66-69

SUBJECT INDEX 393

in practice, 52 mechanism, 90,91 prevention, 91,92

Stress field, hydrogen interaction with, 275 Stress relaxation, effect of hydrogen on, 307 Sulphide inclusions in steel, 10-14, 20, 30, 31,

44, 374, 375, 377, 382 as pit nucleation sites, 10-14, 20, 30,31, 44 determination method, 378, 379 distribution of, 386 effect on hydrogen absorption, 382 effect on hydrogen permeation, 386

Surface inhomogeneity as pit nucleation sites, 9,10, 13

Surface potential, 227, 228 Surface wave velocity, effect of hydrogen on,

301,302

Susceptibility to pitting, 9, 15, 17, 18, 45

Tafel slope for hydrogen evolution, 247, 248

X-ray study of hydrogen effects in metals, 237, 275, 306, 323, 324, 326, 329,331-334

X-ray study of hydrides, 324, 329, 331-334

Yield stress for iron, effect of hydrogen on,

303-305

Zero charge potential, 28

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