Group Title: Differential capacity of stainless steel in potassium chloride solutions during potentiostatic and galvanostatic polarization /
Title: Differential capacity of stainless steel in potassium chloride solutions during potentiostatic and galvanostatic polarization
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Title: Differential capacity of stainless steel in potassium chloride solutions during potentiostatic and galvanostatic polarization
Physical Description: xiii, 187 leaves : ill. ; 28cm.
Language: English
Creator: Fiorino, M. Elaine Curley-, 1944-
Publication Date: 1975
Copyright Date: 1975
 Subjects
Subject: Steel, Stainless -- Testing   ( lcsh )
Potassium chloride   ( lcsh )
Polarization (Electricity)   ( lcsh )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
 Notes
Thesis: Thesis--University of Florida.
Bibliography: Bibliography: leaves 179-186.
General Note: Typescript.
General Note: Vita.
Statement of Responsibility: by M. Elaine Curley-Fiorino.
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Bibliographic ID: UF00098874
Volume ID: VID00001
Source Institution: University of Florida
Holding Location: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: alephbibnum - 000157592
oclc - 02573890
notis - AAS3898

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DIFFERENTIAL CAPACITY OF STAINLESS STEEL
IN POTASSIUM CHLORIDE SOLUTIONS DURING
POTENTICSTATIC AND GALVANOSTATIC POLARIZATION











By

M. ELAINE CURLEY-FIORINO


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











UNIVERSITY OF FLORIDA


1975

































The author dedicates this dissertation to her

husband, John A. Fiorino.














ACKNOWLEDGEMENTS


The author would like to express her gratitude

to the members of her committee; Dr. E. D. Verink, Jr.,

Dr. J. D. Winefordner, Dr. W. S. Brey, Dr. R. Bates, and

especially to her research director, Dr. G. M. Schmid for

the interest and assistance given to her in the course

of this investigation and in the preparation of this

manuscript.

She would also like to thank Mr. R. Dugan and

Mr. A. Grant and their associates for their help in the

technical aspects of this work.

Steel samples were provided by the United States

Steel Corporation.

The author is also grateful for the financial assist-

ance received from the University of Florida in the form

of Graduate School Fellowships.

Finally, the author would like to thank

Dr. and I:rs. James L. Fortuna and Mr. and Yrs. Willis Bodine

for their friendship and encouragement during her

graduate career.
















TABLE OF CONTENTS


ACKNOWLEDGEENNTS. . . . .

LIST OF TABLES . . . .

LIST OF FIGURES . . . .

ABSTRACT. . . . . . .

Chapter

I. INTRODUCTION. . . .

II. EXPERIMENTAL. . . .

Experimental Design .


Experimental Technique. . . . .

Fotentiostatic polarization . .

Galvanostatic polarization . .

III. RESULTS . . . . . . . .

Current-Potential Behavior During
Fotentiostatic Polarization . . .

Primary solutions . . . . .

Secondary solutions . . . .

Capacity-Potential Behavior During
Potentiostatic Polarization .. ..

Primary solutions . . . . .


Secondary solutions . . . . .


Fage

. . . . . . iii

. . . . . . vi

. . . . . . viii

. . . . . . x



. . . . . . 1

. . . . . . 13

. . . . . . 13


. . 18

S . 21

. . 25

. . 34


. 92

. 92

. 102








Chapter

Galvanostatic Polarization . . . . .

Potential-time behavior . . . . .

Capacity-potential behavior. . . . .

IV DISCUSSION . . . . . . . . . .

Potentiostatic Polarization . . . . .

Active dissolution and passivation. . .

The hydrogen evolution reaction . . .

Rest Potential. . . . . . . .

Passive region. . . . . . . .

Pitting . . . . . . . .

Transpassive dissolution. . . . . .

Galvanostatic Polarization. . . . . .

V SUMMARY . . . . . . . . . .

LITERATURE CITED . . . . . . . . .

BIOGRAPHY . . . . . . . . . . .


Page

103

103

122

131

131

131

135

142

144

150

156

160

173

179

187








LIST OF TABLES


Table Page

1. Fotentiostatic current-potential behavior in
primary solutions (average curves) ..... 38

2. Fotentiostatic current-potential behavior in
primary solutions (individual experiments) . 39

3. Cathodic loop in primary solutions . . .. .40
4. Transpassive behavior in primary solutions . 41

5. Solution composition parameters investigated
in primary solutions . . . . . ... .42

6. Interrelationships in potentiostatic
experiments in primary solutions . . . 43

7. Potentiostatic current-potential behavior in
secondary solutions (average curves) ..... 69

8. Potentiostatic current-potential behavior in
secondary solution (individual experiments). 70

9. Interrelationships in potentiostatic
experiments in secondary solutions . . . 83

10. Cathodic loop in secondary solutions .... .95

11. Potentiostatic capacity-potential behavior in
primary solutions (average curves). ... 96

12. Potentiostatic capacity-potential behavior in
primary solution (individual experiments). . 97

13. Potentiostatic capacity-potential behavior in
secondary solutions (average curves) .... .98

14. Potentiostatic capacity-potential behavior in
secondary solutions (individual experiments) 99

15. Capacity-potential behavior in the potential
range of linearity of 1/C versus E . . .. 101

16. Galvanostatic potential-time behavior (0.0 M
potassium chloride). . . . . . .. 106

17. Galvanostatic potential-time behavior (0.100
[M potassium chloride). . . . . . ... 107










Table


Fage


1J. Galvanostatic potential-time behavior (0.303
M potassium chloride). . . . . . ... 108

19. Galvanostatic potential-time behavior (0.518
M potassium chloride). . . . . . ... 109

20. Tafel behavior during galvanostatic
polarization . . . . . . . . . 111

21. Average final breakdown potentials during
galvanostatic polarization . . . . .. 116

22. Open circuit decay behavior following
galvanostatic polarization . . . . ... .118










LIST OF FIGURES


Figure Page

1. Stainless steel electrode . . . . . 14

2. Electrochemical cell. ... . . . . 17

3. Kel-F electrode holder. ..... . . . . 20
4. Block diagram of the potentiostatic
polarization circuit . . . . . .. 23

5. Block diagram of the galvanostatic
polarization circuit . . . . . .. 27

6. Block diagram of the circuit employed in
capacity-potential data correlation . .. 31

7. Control panel ...... .. . . . 32
8. Schematic potentiostatic current-potential
behavior ... .............. * 36

9. Variation of the logarithm of the critical current
density with the concentration function,
pH + log ( [S042] + [C1-] ) . . . 48

10. Potentiostatic polarization curve showing
cathodic loop obtained in 0.334 M sodium
sulfate at pH 2.40 . . . . . . 51

11. Variation of the potential of total passivity
with chloride ion concentration . . . 54

12. Variation of the potential of total passivity
with the concentration function,
[C1-]/(2 [SC4 -] + [01- ) . . . . . 56

13. Variation of the potential of total passivity
with the concentration function,
[C1-]/(2 [S042-j + [C1-] ) ... . . . . 58

14. Variation of the potential of total passivity
with the concentration function, pH +
log ( [S42-] + [C1-] ) . . . . 60


viii









Figure Page

15. Variation of the potential of total passivity
with the cQncentration function,
log ( [S04 -] / [Cl-] ) . . . . . 62
16. Potentiostatic polarization curve showing the
influence of pitting obtained in 0.301 M
potassium chloride at pH 2.35 . . . . 65

17. Variation of the critical current density with
chloride ion concentration . . . . 72

18. Variation of the logarithm of the critical
current density with rest potential . .. 74

19. Variation of the logarithm of the critical
current density with the concentration function,
pH + log ( [S04-] + [C1-] ) ... . .. 76

20. Variation of the potential of total passivity
with rest potential . . . . . . 78

21. Variation of the logarithm of the critical
current density with the potential of total
passivity . .... . . . . . . 80

22. Variation of the primary passivation
potential with pH . . . . . * 85

23. Variation of the rest potential with pH . 87

24. Variation of the logarithm of the critical
current density with pH . . . * * 89

25. Variation of the logarithm of the critical
current density with the concentration function,
pH + log ( [S04j-] + [Cl-] ) . . . ... 91

26. Potentiostatic capacity-potential behavior
in the absence of hydrogen interference . 94

27. Schematic representation of the galvanostatic
potential-time curve in the presence (1)
and absence (2) of pitting breakdown. ... . 105

28. Schematic representation of the capacity-time
behavior of a system not subject to pitting
breakdown during galvanostatic polarization 124













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



DIFFERENTIAL CAPACITY OF STAINLESS STEEL
IN POTASSIUM CHLORIDE SOLUTIONS DURING
POTENTICSTATIC AND GALVANCSTATIC POLARIZATION

By

M. ELAINE CURLEY-FIORINO

June, 1975

Chairman Gerhard M. Schmid
Major Department: Chemistry

Potentiostatic and galvanostatic polarization of

AISI 304 stainless steel was performed in deaerated

solutions of 0.0 through 1.0 M potassium chloride at

pH 2.4. Sodium sulfate was added to an ionic strength

of one. Cylindrical electrodes were mechanically

polished and prepolarized at -0.700 V for twenty minutes.

The differential capacity of the metal-solution interface

was determined as a function of potential using the

single pulse technique.

All potentials are given versus the saturated calomel

electrode. Current densities and differential capacities

are referred to the electrode geometric area.

During potentiostatic polarization, an active-passive

transition is observed for all solutions. The rest

potential (-0.48 V) and primary passivation potential








(-0.42 V) are independent of solution composition. The

critical current density increases with the total anion

concentration. The capacity peak occurring slightly

negative to the primary passivation potential is attributed

to specific adsorption of the anions involved in the

dissolution-passivation mechanism. The potential at which

the electrode becomes totally passive shifts positive

with increasing chloride ion concentration and is dependent

on kinetic rather than thermodynamic factors.

Transpassive dissolution initiates between 0.57 and

0.64 V and exhibits Tafel behavior. Both the current

density maximum (10-4 A cm-2) and the corresponding

potential (0.95 V) associated with the onset of secondary

passivity are independent of solution composition. The

capacity peak observed in this region is attributed to

adsorption of the passivating species.

In solutions of ) 0.3 M potassium chloride, an

increase in current density caused by pitting starts

at potentials normally in the passive region. The pitting

potential depends on the relative amounts of chloride and

sulfate ions present. No capacity peak is observed

prior to pitting breakdown.

The capacity, between the potential of total passivity

and the capacity minimum occurring at 0.60 V, is approxi-

mately independent of the stability of the passive state.

This independence results from the domination of the

interfacial capacitance by that of the passive film, obscuring
xi









any changes in the film-solution capacity which may occur.

The sudden decrease in capacity observed at 0.30 V is

attributed to a change in the dielectric constant of the

film. Between 0.3 and 0.6 V, film growth follows an

inverse logarithmic law.

The potential-time behavior of samples subjected

to galvanostatic polarization depends on both current

density and solution composition. Systems not susceptible

to pitting reach and maintain a positive steady state

potential. In systems subject to pitting, the maximum

potential attained is unstable and a shift in potential

to more active values occurs.

In 0.3 M potassium chloride intermediate arrests

produced by a given current density as well as the

maximum potential achieved prior to breakdown correspond

to potential arrests in the non-pitting systems. In

0.5 M potassium chloride pitting breakdown occurs from

a potential maximum which is considerably below that

observed in 0.3 M but the behavior at more negative

potentials is similar.

It is assumed, therefore, that the initial effect

of the anodic current on the metal surface at a given

potential is the same in all cases and that pitting

succeeds through the perturbation of these initial

surface conditions by chloride ion.









Tafel behavior is associated with the arrests at

-0.4, 0.8, 0.85, and 1.12 V. Active dissolution at

-0.4 V exhibits a slope of 0.060 V decade-1 and is

thought to proceed by the Heusler mechanism. Arrests

at 0.8, 0.85 and 1.12 V correspond to transpassive

dissolution, secondary passivity, and oxygen evolution,

respectively. Capacity peaks are associated with the

latter two effects and with the arrest at -0.4 V but

not with pitting breakdown.


xiii













CHAPTER I

INTRODUCTION



A metal is in the passive state when it is inert

in an environment in which, on the basis of thermodynamics,

it should corrode readily. An ennoblement of the

potential of the metal-environment interface accompanies

the onset of passivation. Passivity has been recognized

since 1836, when Faradayl observed the stability of

iron metal immersed in concentrated nitric acid.

However, the nature of the surface species leading

to the onset and maintenance of the passive state is

still not well understood. The two major theories

advanced to explain the phenomenon are the bulk oxide

theory and the adsorption theory.

The oxide film theory proposes that a bulk oxide

is formed directly on the metal surface from the
2,3,4
products of the metal dissolution reaction.3 This

oxide then acts as a physical barrier between the metal

surface and the aggressive environment.

The adsorption theory of passivity attributes

passivation of the metal surface to the adsorption of

an "oxygen" species in less than or equal to monolayer

quantities.5, Two variations of the adsorption theory







2

have been described. In the chemical variation proposed

by Uhlig,8 adsorbed oxygen atoms are believed to

satisfy the surface valences of all atoms on the metal

surface. The correlation of the theoretical and

experimental concentrations of components of transition

metal alloys needed to bring about a sharp increase in

the ease with which the alloys are made passive lends

support to this hypothesis. The electrochemical

variation ascribes the retardation of the metal dis-

solution reaction to a change in the double layer

structure caused by the dipolar character of chemisorbed

oxygen.9 Orientation of the dipole with its positive

end towards the solution increases the activation

energy necessary for the metal dissolution reaction.

Hackerman10 has proposed a theory intermediate to

the above two. Here, the adsorption of oxygen atoms on

the metal results in a metastable state lending

temporary protection to the surface. Following electron

transfer from the metal to adsorbed oxygen atoms, an

amorphous bulk oxide is formed by cation migration through

the adsorbed array. It is this bulk oxide which provides

long-term protection. This theory, originally postulated

for metals in oxygen-containing solutions, has also been

proposed to explain the passivation mechanism of iron-

chromium alloys in deaerated acidic sodium sulfate

solutions.1 A similar mechanism has been proposed by

Frankenthall2 for the passivation of an iron-24 chromium








alloy. Electron diffraction studies of the surface of

stainless steels exposed to oxidizing acids or the atmosphere

for short times at intermediate temperatures (25 to 600C)

have shown that the passive film formed is non-crystalline.13

Transmission microscopy studies of films formed on an iron-24

chromium alloy during potentiostatic polarization in 0.5 M

sulfuric acid show that they are also amorphous.14

The corrosion resistance of austenitic stainless steels

in diverse environments is a result of the high degree of

passive state stability conferred by the presence of chromium

in amounts ) 12 percent.1 16 However, in specific media,

especially in halide solutions, this corrosion resistance is

lost, as evidenced by the onset of intense local attack,

i. e., pitting. The mechanism by which pits nucleate and

propagate in the presence of chloride ion has been investigated

in great detail. Excellent reviews are given by Kolotyrkin17

and Szklarska-Smialowska.18

The pitting phenomenon has been characterized by

three parameters the pitting potential, negative to

which no pits can nucleate; the critical chloride ion

concentration, the minimum concentration needed to

initiate pitting in a given system; and the induction

time, the time, at a given potential, which passes

prior to breakdown. Thus, studies have been carried out








19, 20, 21, 22, 23
to examine the effect of metal composition, 20, 21, 22, 23

defect density of a metal surface,24, 25 grain size,26

solution composition,27, 28, 29, 22 sulfide inclusions,30' 31, 32

and temperature28 on the pitting potential and/or the location

and number of pit sites. The critical chloride ion

concentration has been shown to depend strongly on the

concentration of inhibiting anions (e.g., sulfate, hydroxyl

and nitrate ions) in the solution33 35, 28 as well as

on alloy composition.36 The induction time has been found

to decrease with increasing chloride ion concentration,37
38, 39, 40 potential of passivation,39' 40 and

temperature.38

The influence of an induction time is seen both in

potentiostatic and galvanostatic polarization measurements.

Previous works on austenitic stainless steels in chloride-

containing solutions have shown that the application of a

constant anodic current to an initially active electrode

shifts the potential positive to a maximum value.41

Despite the continued application of the current, the

potential then decreases rapidly to a value normally in

the passive region. The length of time to potential

breakdown is a function of the current density and the

chloride ion concentration. Examination of the metal

surface shows that pitting has occurred.








5
The steady state potential achieved after breakdown

has been called the protection potential of the system
4?2
under study.42 At potentials between the protection and

pitting potentials previously nucleated pits can continue

to grow but no new pits may form. At potentials negative

to the protection potential, all existing pits become

inactive.35

Anodic potentiostatic polarization to values more

positive than the pitting potential also causes pit

formation. A change in potential to values just positive

to the pitting potential results in a decrease in current

density with time from the initially high value associated

with double layer charging.40 This current decrease

represents the readjustment of the electrode-solution

interface to maintain the passive condition. After a

time interval which depends on solution composition

and potential, the current decrease is replaced by current

oscillations signifying the pitting-induced breakdown of

passivity. As the potential is made still more positive,

the time for which passivity is maintained decreases

until an induction period is no longer apparent. 3

Mechanisms proposed to explain the dependence of

pitting parameters on experimental variables are a

function of the assumed nature of the passivating film.

On the basis of a bulk oxide, Hoar3 has described a

mechanical breakdown process in which the mutual repulsion









of anions adsorbed on the oxide surface leads to the

formation of cracks. A relationship between critical

breakdown stress and surface tension as influenced by

anion specific adsorption has been derived by Sato.44

Hoar and Jacob45 have also suggested that a metal cation

is dissolved from the oxide through the formation of a

metal-chloride complex containing 2.5 to 4.5 chloride

ions. Cation migration through the film then allows

continuation of the process. If the passive film is

assumed to be an adsorbed "oxygen" species, then its

replacement by chloride ions will lead to activation

of the metal when a critical surface concentration of
28, 17
chloride ion is attained.

The presence of an induction time associated with

pitting breakdown suggests that a time consuming change in

surface structure is occurring. If the specific adsorption

of chloride ion is involved in this change, then its

incorporation into the electrical double layer should

be reflected in the differential capacity-potential

behavior of the interface.

The electrical double layer, in its ability to store

charge, acts as a capacitor. The magnitude of the

capacity associated with it is a complex function of

potential and is therefore defined as a differential

capacity, q/6E. In the absence of specific adsorption,

only water molecules populate the compact double layer,










the region between the metal and the Outer Helmholtz

Plane. The capacity of the interface can then be

represented by two capacitors in series; that of the

region between the metal and the Outer Helmholtz

Plane (OHP); and that of the region between the OHP

and the bulk of the solution (the diffuse double layer).

In dilute solutions ( (0.001 M), the capacity of the

latter is small and dominates the interfacial capacity.

In concentrated solutions, all of the diffuse double

layer charge is located close to the OHP. The interface

then behaves like a single parallel plate condenser, i.e.,

its capacity is approximately independent of potential.

When specifically adsorbed ions populate the IHP

in a concentrated solution (> 0.001 M), the interface

again functions as two capacitors in series. Bockris and
46
Reddy, in their treatment of the effect of contact

adsorption on the total capacity of the metal-solution

interface, derive the equation

1/C = 1/KM-OHP (I- (qCA ) (1)
KM_-0OHP KM-IHP qao

where C is the total measured differential capacity and

KMOHp and KM-IHp are the integral capacities associated
with the metal to Outer Helmholtz Plane and metal to Inner

Helmholtz Plane regions, respectively. The function,

ZOqCA/'qOM, represents the change in the amount of

specifically adsorbed ion, qCA, with the charge on the

metal, qp. Thus, as the charge on the metal becomes more










positive with increasing potential, an increase in the

degree of specific adsorption with metal charge

(-2qCA/ qc2 >0) should occur, producing an increase

in the measured differential capacity of the interface.

As growth continues, however, the buildup of lateral

repulsion forces tends to decrease the degree to which

specific adsorption occurs at a given metal charge.

This inflection in the qCA_-q relationship results in a

peak in the differential capacity-potential curve

(-2qCA/ qRM2 = 0). Differential capacity measurements

have been successfully applied to the determination of

the specific adsorption of sulfate, perchlorate and

chloride ions on iron.47' 48, 49. 50

The presence of a positive charge on the metal

surface implies polarization at potentials positive to

its zero point of charge. Studies of binary alloys

suggest that the zero point of charge of an alloy

should approach that of its component with the most

negative zero point of charge if it is present in

sufficient concentration.51 The zero points of charge

for nickel, chromium and iron have been determined
52
in acidic sulfate solutions.52 The corresponding values

are -0.57, -0.69 and -0.62 V, respectively. It is

expected then, that the zero point of charge of active

stainless steel should be close to that of chromium,

i.e. -0.69 V.








9
In order to explain the potential dependence of the

activating effect of chloride ion in terms of specific

adsorption to a critical surface concentration on the

passive electrode, the zero point of charge of the

passive surface must lie in the passive region. For

passive iron, the zero point of charge occurs at -0.125 V

in 0.01 M sodium hydroxide.52 This positive shift in

value from that observed on active iron can be attributed

to an increase in work function resulting from the

presence of an adsorbed "oxygen" species or an oxide

film. The extension of this phenomenon to stainless

steel seems logical.

The primary purpose of the experiments conducted in

this study was the determination of the effect of

chloride ion on the capacity-potential behavior of

stainless steel observed during potentiostatic and

galvanostatic polarization in solutions initiating pitting.

In order to explain the potential arrests observed during

constant current polarization, the potentiostatic studies

were extended to cover the potential range from active

dissolution to oxygen evolution. Corresponding capacity

values were determined and the current-potential and

capacity-potential data correlated.

The relationship between the rate of an electrochemical

reaction and the potential difference across the interface

at which it occurs is discussed in detail by Bockris and










Reddy. When electron transfer is rate-determining,

i.e., the system is under activation control, the current-

potential relation is given by the general form of the

Butler-Volmer equation,

i = io exp 2 LF i exp ( F )i (2)
RT RT
where io is the exchange current density, the oC's are the

transfer coefficients, and n is the overvoltage. All other

terms have their usual meaning. The first term on the

right hand side of equation 2 is the current density resulting

from the oxidation anodicc) reaction. The second term

pertains to the reduction cathodicc) reaction.

At the equilibrium potential of the rate-determining

reaction, the overvoltage, which represents the potential

difference across the interface in excess of the equilibrium

potential difference ( E Eo), is zero. The net current

density observed (i) will therefore also be zero since at

equilibrium the rates of the anodic and cathodic reactions

will be equal. The transfer coefficients determine what

fraction of the potential difference across the interface

is operative in changing the energy barrier for the oxidation

and reduction reactions.

As the potential difference across the interface is

made more positive (v> 0), the contribution of the anodic

current density to the total current density will

increase. At a sufficiently positive overvoltage

(~0.120 V for a one-electron transfer reaction), the






11

influence of the cathodic current density becomes negligible.

The current-potential relationship is then given by

i =i exp aF_ (3)
RT (3)

which can be put in logarithmic form and rearranged to give

= 2.3 RT log i 2.3 RT log i (4)
SF ocF
A plot of overvoltage versus log i is therefore linear.

Such plots are known as Tafel lines. The slope of the

line contains the transfer coefficient which is a complex

function of the total number of electrons transferred

during the reaction as well as the mechanism by which the

reaction proceeds. From the slope of the line and its

intercept, the exchange current density for the reaction

can be calculated.

In systems in which a faradaic current can flow, i.e.,

charge can cross the metal-solution interface, the

electrical behavior of the interface can be represented by

a resistor in series with a capacitor and resistor in

parallel. The series resistor represents the resistance

of the solution to current flow. The capacity is the

differential double layer capacity. The parallel resistor

represents the polarization resistance of the faradaic

reaction, decreasing as the rate constant of the reaction

increases. In cases of low polarization resistance, the

determination of the differential capacity is difficult

since the polarization resistance can act as a leakage









path for the signal measuring the capacity. Faradaic

current also interferes indirectly by causing a change

in the true electrode surface area as well as in the solution

composition.

The use of the classical alternating current technique

in which the interface forms one arm of an impedance

bridge to determine the double layer capacity on a solid

electrode is precluded because of the dependence of the

capacity value measured on signal frequency.5 Most of

the direct interference can be eliminated, however, by

the use of the single pulse method of differential

capacity measurement developed by Riney, Schmid and
54
Hackerman. Analysis of the linear segment of the

potential transient resulting from a single current

pulse allows calculation of the capacity at the point

from which the pulse initiates since

C = i(dt/dE)t= (5)

where i is the pulse magnitude and dt/dE is the slope of

the Dotential-time transient evaluated at t = 0.













CHAPTER II

EXPERIMENTAL

ExDerimental Design

The material investigated was stainless steel, AISI

304, provided by the United States Steel Corporation. Its

composition was given as 0.03 C, 0.027 P, 1.10 Mn, 0.022 S,

0.43 Si, 9.26 Ni, 18.6 Cr, 0.39 Mo, and 0.04 N (weight

percent). Bar stock was machined to cylinders with a

diameter of 6 mm and a height of 9 mm (Figure 1). The

cylinders were tapped, threaded,and mechanically polished

at 2400 ram with 400 followed by 600 grit emery paper.

They were then degreased with spectral grade benzene in

an ultrasonic cleaner, rinsed with triply distilled water,

and stored in a closed polyethylene container until needed.

Primary studies were carried out in solutions

containing 0.0, 1.17 x 10-2, 9.97 x 10-2, 0.301, 0.508,

and 1.0 M potassium chloride. Solution pH was measured

with a Beckman pH meter and adjusted to 2.4 with

concentrated sulfuric acid. Sodium sulfate was added as

required to maintain an ionic strength of one (0.318, 0.313,

0.284, 0.232, 0.147, and 0.00 M, respectively). Secondary

experiments involved solutions of pH 2.4 containing 0.102,

0.123 and 1.0 M potassium chloride with no sodium sulfate'

additions as well as 0.3 M potassium chloride at pH 1.52

















































Figure 1. Stainless steel electrode.








15

and 6.22. All chemicals used in solution preparation were

reagent grade. Recrystallization of potassium chloride

from triply distilled water had no effect on experimental

results. The water employed was distilled from alkaline

potassium permanganate and then from a two-stage Heraeus

quartz still and collected in a two-liter Pyrex volumetric

flask. Its maximum conductivity, determined with a General

Radio Impedance Bridge, was 2 x 10-6 1 -lcm-.

Platinum electrodes, generally used in pre-electrolysis

to remove electroactive impurities from the solution, have

been shown to dissolve when polarized anodically in both

sulfate and chloride containing solutions.55 Because of

the possibility of contaminating both the solution and the

stainless steel surface with platinum, a pre-electrolysis

step was therefore omitted.

The electrochemical cell was made of Pyrex and was of

conventional design (Figure 2). A Luggin capillary

connected the saturated calomel reference electtode (SCE)

to the cell via two solution-lubricated mercury-seal
2
stopcocks and a potassium chloride salt bridge. A 1 cm

platinum flag auxiliary electrode was mounted on the cell

with a standard taper joint. For use in constant current

polarization and capacitance measurements, a platinum

gauze basket, approximately 100 cm2 in area (Engelhard

Industries), was mounted concentric to the test electrode.

The cell cap incorporated a 24/40 standard taper joint for

mounting the test electrode.



























Figure 2. Electrochemical cell.















SOLUTION INLET--
















GAS INLET


Pt FLAG -
ELECTRODE
COMPARTMENT


TEST ELECTRODE RECEIVER

-GAS OUTLET













REFERENCE ELECTRODE-
COMPARTMENT


Pt GAUZE ELECTRODE

Hg SEAL STOPCOCKS


LUGGIN CAPILLARY







18

The electrode holder (Figure 3) was made from a Kel-F

rod machined to approximately 1 cm in diameter in which a

3 mm center hole was drilled. The rod was then heated and

a threaded stainless steel rod inserted so that thread

protruded on both ends. The Kel-F rod was shrunk in an

ice bath to facilitate its insertion into a 1 cm ID glass

stirrer bearing sleeve which had a 24/40 standard taper

joint attached. The sample was then affixed to the

protruding steel rod. The electrode area exposed to
2
solution was approximately 2 cm2. The sample-sample

holder seal could be tightened by turning a finger nut

at the top of the assembly. Teflon washers were placed

between both the sample and holder and the nut and holder

to reinforce the seal. Spot checks of the shielded top

surface of the test electrode showed no evidence of

corrosion indicating the absence of leakage.

Experimental Technique

Solutions were deaerated with helium (99.99 percent)

for a minimum of eight hours prior to use. Prepurification

and water saturation of the gas was accomplished by passing
0
it through a 12 cm column of Linde 5 A molecular sieve

pellets into a gas wash bottle containing triply distilled

water. The gas then flowed into a two-liter reservoir

containing the solution and continued through a dispersion


































Figure 3. Kel-F electrode holder.










































24/40 STANDARD
TAPER JOINT
















TEFLON WASHER-----


FINGER NUT




TEFLON WASHER









KEL-F ROD




STAINLESS STEEL ROD







21

tube into the electrochemical cell. The cell gas outlet

terminated in a gas wash bottle to prevent contamination

of the cell contents with ambient air.

Immediately before each experiment, the cell was

washed with hot chromo-sulfuric acid cleaning solution

and rinsed with triply distilled water. Gas pressure was

then diverted to fill the cell with deaerated solution

from the reservoir. The stainless steel sample was

secured on the Kel-F holder, rinsed with triply distilled

water and the solution to be studied, and immersed in

solution. All samples were pretreated at -0.700 V for

twenty minutes to reduce air-formed surface films.

Solution stirring was accomplished with a magnetic

stirring bar and the helium flow continued throughout

the experiment. All solutions were at room temperature

and all potentials are reported relative to the saturated

calomel electrode. Current density and differential

capacity were calculated using electrode geometric

areas. The test electrode was grounded during all

experiments.

Potentiostatic polarization

The block diagram of the system employed in potentio-

static experiments is shown in Figure 4. Polarization

was accomplished with a slightly modified Harrar56

potentiostat. (Two each of obsolete transistors 2N333A

and HA7534 were replaced with the equivalent circuit elements

2N3568, 2N5869, and 2N3644,2N5867.)





















Figure 4. Block diagram of the potentiostatic polarization circuit.

A. Test electrode
B. Reference electrode
C. Platinum flag electrode
D. Platinum gauze electrode

















PULSE GENERATOR









After prepolarization at -0.700 V the potential of the

test electrode was shifted anodic in step-wise increments.

The magnitude of the imposed step depended on the potential

range under investigation and the electrochemical reactions)

associated with it. In regions where changes in potential

caused significant changes in current density, steps of 20

to 30 mV were generally employed. In regions of approx-

imately constant current, 50 mV steps were used. A

Keithley 610B electrometer was used to monitor the applied

potential. After an arbitrary time interval of 10 minutes,

the current flowing in the auxiliary-test electrode circuit

was determined from the potential drop across a 1.18 kGl

or a 10 k1 wire-wound precision resistor (1 percent)

shunting the input of a Keithley 660 electrometer (1014 _

input impedance).

The differential capacity of the stainless steel-solution

interface was determined as a function of potential using

the single pulse method.5 A voltage pulse from the gate

of a Tektronix 549 storage oscilloscope was used to trigger

a Tektronix 114 Pulse Generator. The 100 ,sec square wave

current pulse produced was applied to the platinum

basket-test electrode circuit. The resulting potential-time

transient was recorded on the oscilloscope operated in the
-l
storage mode at a sensitivity of 2 or 5 mV cm- and 2 or

5 ysec cm-1. Differential capacity values were calculated

from the linear segment of the transient slope during the

initial 10 usec of the pulse. Since the transient is







25

linear between 4 and 10 l sec, the slope determined between

these two points is equivalent to that of the tangent drawn

to the curve at t = 4 sec. Measuring times of 4,ysec

correspond to an alternating current frequency of

approximately 1.2 x 105 Hz. Since a frequency of 105 Hz
57
is required to eliminate faradaic interference completely,

very little interference is expected.

Prior to each experiment, the pulse magnitude was

calibrated using standard capacitors and resistors. The

current density of the pulse could be calculated from the

known value of capacity in the calibration circuit, the

measured slope of the transient, and the electrode geometric
-2
area. Pulse magnitude was approximately 5 mA cm and was

independent of capacity between 1 and 0ldfarads.

Galvanostatic polarization

After potentiostatic prepolarization the system was

switched to galvanostatic control using a two-position

toggle switch. A block diagram of the circuit used

during galvanostatic polarization experiments is given

in Figure 5. A constant current was supplied to the

platinum basket-test electrode circuit by a Hewlett Packard

881A power supply operated in the constant voltage mode

in series with a bank of resistors (4.7 to 47.8 kl ). The

positive terminal of the power supply was grounded to

establish the test electrode as the anode. The current

magnitude was determined from the voltage drop developed























Figure 5. Block diagram of the galvanostatic polarization circuit.

A. Test electrode
B. Reference electrode
C. Platinum flag electrode
D. Platinum gauze electrode



















ELECTROMETER


OSCILLOSCOPE 2


POTENTIOSTAT


ELECTROSCAN 30


POWER SUPPLY








across a lO lprecision resistor (1 percent) using a

Keithley 6-0 electrometer and was adjusted by varying

the series resistance and the voltage output of the

rower supply. Applied current densities ranged from

0.99 x 10 to 2.0 x In-3 A cm-2. The resulting

potential-time curve was displayed on the recorder of a

Beckman Electroscan 30. Potential values were monitored

with a Keithley 61(B electrometer.

A study of differential capacity as a function of

potential was also carried out during constant current

polarization. As described above, a square wave current

pulse was delivered from the pulse generator to the

platinum basket-test electrode circuit. The resulting

potential-time trace was displayed on the storage

oscilloscope and the capacity calculated from the initial

slope of the transient and the predetermined pulse

magnitude.

The development of the theory of the single pulse

method assumes that a perturbation current is applied

to a system under constant current polarization at a steady-

state potential so that the current before and after the

pulse remains the same. Since, during potentiostatic

polarization of austenitic stainless steel, either a steady

state current or a very slow decrease in current with time

is observed, the method should be directly applicable. For

galvanostatic polarization the constant current condition








29

is satisfied but a continuous change in potential with time

occurs.' However, for small measuring times ( 10 yOsec),

the change in potential will be negligible.

Because of the rapid change in potential with time

during constant current polarization, difficulties were

encountered in the correlation of capacity values and the

potentials at which they were determined. A system was

therefore developed which facilitated the sequential

storage of pulse-produced transients and supplied a marker

signal to the Electroscan recorder whenever the pulse

generator was triggered (Figure 6).

Two oscilloscopes were used sequentially. Depression

of PG 1 or PG 2 on the control panel (Figure 7) first

changed the DC level of the input of the oscilloscope in

use by 10 mV. This provided automatic downward displacement

of the successively produced potential-time transients.

A two-step generator was associated with PG 1, an eight-

step generator with PG 2. In most cases, the oscilloscope

sensitivity needed for precise determination of the

transient slope (2 or 5 mV cm-) precluded the use of

automatic stepping and vertical displacement was accomplished

manually instead.


























Figure 6. Block diagram of the circuit employed in capacity-potential data
correlation.
























3 CHART
RECORDER













O VERT.
INPUT
(OSCILLISCOPE)













00000
SET CLIP TP PG2 PG, RESET
REC CELL VERT TI T2 PG
000000


Figure 7. Control panel.









Then, after a 50osec delay, the horizontal sweep

of the oscilloscope was triggered. Simultaneously, 15 V,

100 msec pulse was supplied to the external trigger

of the pulse generator to produce the current pulse

needed to determine the capacity. At the same time a 50 mV,

100 msec pulse was superimposed on the galvanostatic

potential-time response signal from the cell to the recorder.

Since the potential values recorded during constant

current polarization reached 1.4 V, while the marker pulse

superimposed on this signal was only 50 mV, the cell signal

was passed through a buffered amplifier and could be

clipped at a preset value. The recorder sensitivity could

therefore be adjusted to display the pulse marker while the

entire potential-time curve remained on chart.













CHAPTER III

RESULTS

Current-Potential Behavior During
Potentiostatic Polarization

The response of the stainless steel samples to poten-

tiostatic polarization is given schematically in Figure 8.

The region ABC represents the transition from the active

to the passive state characteristic of this material in

the solutions employed. An adjoining region of approximately

constant current is then usually observed (CC'D). The

increase in current on further polarization is caused by

loss of passive state stability as the result of pitting (DE)

or of transrassivity (FG). The potential and current

density notations on the schematic are defined below. They

are presented in the text in approximately the same order

in which they arise during polarization.

Possible interrelationships between experimentally

determined current densities, potentials, and/or solution

composition variables were examined. The various data

were plotted. If a plot appeared to be linear, then a

linear regression analysis was performed. Only curves

with correlation coefficients (r2) of ) 0.90 were

reported as linear.

























Figure 8. Schematic potentiostatic current-potential behavior,

















[crtl,2
B
cri t, /I



ifp


i A




V-i


Etr Epp,2
I I


POTENTIAL


r pp,
Er E,,
I


I 1 I '









Primary solutions

Characteristic potentials and current densities

determined from the potentiostatic polarization curves

obtained in primary solutions are presented in Tables 1 4.

The solution composition parameters investigated are listed

in Table 5. The logarithms of these functions have also

been considered. Relationships with correlation coefficients

) 0.90 are reported in Table 6.

After the twenty-minute prepolarization period at

-0.700 V, net currents are cathodic with values between

2.1 x 10 3 and 5.3 x 10-3 A cm2. No correlation between

the hydrogen evolution current and solution composition at

constant pH and constant potential is apparent (Table 1).

As the potential is shifted anodic, the measured current

decreases and becomes zero at the rest potential, Er.

The rest potential is defined as that potential at which

the absolute values of the internal anodic and cathodic

currents become equal, resulting in a net external current

flow of zero. Values of the rest potential determined here

are -0.450 to -0.510 V. The maximum average deviation for

data obtained for a given solution composition is 16 mV for

0.0 M potassium chloride (Table 2). For all other

compositions it is substantially less ( < 7 mV).

Polarization at potentials positive to Er results in

a net anodic current which increases with potential to a

maximum, signifying the onset of passivation. This maximum

current density and its corresponding potential are the










Potentiostatic current-potential
i-70Q
[C-"] [S042-] (A cm ) Er
x 101 x 101 pH x 103 (V,SCE)


Table 1

behavior in primary solutions (average curves).
icrit i
Epp,l (A cm (A pm-2) ETP Epit(V,SCEI"
(V,SCE) x 105 x 107 (V,SCE) inc gr


12.90 -0.285

5.25 -0.255

6.40 -0.242

7.85 -0.230
-0.205

-0.160


0

0.117

0.997

3.01

5.08

10.0


3.34

3.29

3.00

2.48

1.63

0.16


2.40

2.40

2.40

2.35

2.42

2.40


-2.20

-3.10

-3.50

-2.23

-2.06

-5.31


-0.453

-0.475

-0.480

-0.470

-0.510

-0.460


-0.420

-0.425

-0.432

-0.420

-0.400

-0.400


2.26

2.73

3.13

3.43

8.49

6.73


+0.395

0.00

-0.090
-o. 090o


+0.360

-0.012

-0.155











Potentiostatic
experiments).


Table 2

current-potential behavior in primary solutions (individual


i-700
(A cm-2)
pH x 103


0 3.34 2.40 -2.3


-2.1

0.117 3.29 2.40 -2.7

-3.5

0.997 3.00 2.40 -3.8
-3.2

3.01 2.48 2.35 -1.8
-2.7

5.08 1.63 2.42 -2.4
-1.9

-1.9

10.0 0.160 2.40 -4.9

-5.8


ocritsl
Er Epr,l (A cm-lt
(V,SCE) (V,SCE) x 105


-0.420

-0.460

-0.460

-0.473

-0.473

-0.480

-0.475

-0.475

-0.480

-0.508

-0.520

-0.520

-0.460


-0.400

-0.430

-0.410

-0.410

-0.420

-0.430

-0.430

-0.420

-0.415

-0.400

-0.400

-0.400

-0.400


1.31

2.65

1.90

3.4

2.3

3.3

2.9

3.55

3.39

8.3

7.6

9.5

7.3


i
(A cm-2)
x 107


10.7

21.5

65.9

5.32

5.18

7.19

5.01

7.23

8.48


x[ 1 J [S041l
x d1 x i1


Epit(V,SCE)
inc gr


ET1p
(V,SCE)

-0.278

-0.287

-0.287

-0.255

-0.265

-0.237

-0.250

-0.230

-0.226

-0.200

-0.200

-0.210

-0.197


0.365

0.360

-0.020

-0.047

-0.020

-0.145 o


- -0.197 -0.05 -0.175


-0.460 -0.400 6.2


0.420

0.370

0.00

0.00

0.00

-0.10










Cathodic loop

70Q2
(A cm ) C-7002
pH x 103 (if cm2)


Table 3

in primary solutions.

Negative
Cl loop
(pf cm-2) (V,SCE)


0 3.34 2.40 -2.31



-2.10

0.117 3.29 2.40 -2.67

-3.54

0.997 3.00 2.40 -3.83

-3.18

3.01 2.48 2.35 -1.8

-2.67

5.08 1.63 2.42 -2.45

-1.89

-1.P5

10.0 0.160 2.40 -4.86

-5.77


-0.30 to 0.00 -13.6



-0.250 -2.27



-0.27 to -0.20 -3.89










-0.20 to -0.15 -4.25


] 01 SO 21
x x 0


ic
(A cm-2)
x 107


1-200-2
(A cm
x 10


13.4

27.3

-9.13

4.54

3.10

10.4

-0.09

9.66

11.5

5.97

8.15

10.6

-4.25

3.42








Table 4

Transpassive behavior in primary solutions.

ba E Ep icrit,
pH V decade (V, E) (V,bCE) (A cm-)
x 104

2.40 0.114 0.671 0.950 1.12

2.40 0.114 0.641 0.950 1.02

2.40 0.150 0.575 0.950 1.03


rso
x


0.0

0.117

0.997


3.34

3.29

3.00


A i,
A cm-2)
x 106

21

14

17








Table 5

Solution composition parameters investigated in primary solutions.

Cla S04/C1 C1/(C1 + 04) C1/(C1 + 2 S04) C1/(C1 + S04) pH + log(Cl + S04)


10-2

10-2


28.1

3.00

0.82

0.32

0.016


0

0.034

0.249

0.548

0.755

0.984


0.017

0.142

0.377

0.609

0.968


0.066

0.400

0.707

0.861

0.982


1.92

1.93

2.00

2.09

2.25

2.40


a
Cl and S04 represent the analytical concentration of chloride and sulfate, respectively.
The brackets and charge designations have been omitted for clarity.


0.0

1.17 x

9.97 x

0.301

0.508

1.00








Table 6

Interrelationships in potentiostatic experiments in primary solutions.


Independent
vr Qhl1 0


---;aba


pH + log(S04 + Cl)a

Cl

C1/(C1 + SO4)
Cl/(Cl + 2 SO4)

pH + log (SO4 + Cl)

log(S04/Cl)





Cl

C1/(C1 + SO4)
C1/(C1 + 2 SO4)

C1/(C1 + S04)
log (SO/Cl)

pH + log (S04 + Cl)


Dependent
variable


log icrit

ETP

ETp

ETp
ETp

ETp






ip

ip
ip
i
P

log ip

log ip


Correlation
Slope coefficient


1.10

0.108 V mole-1

0.103 V

0.110 V

0.218 V decade-1

a. 2 segments -
-0.016 V decade
-0.039 V decade-1

b. 1 segment -1
-0.030 V decade

8.68 x 10-7 A cm2 mole-1

5.03 x l0-7 A cm"2

7.09 x 10" A cm2

4.05 x l0-7 A cm-2

-0.11

1.05


(0.83)

0.93
0.92

0.94

0.94


0.98
0.99

0.94

0.97
1.00

0.99

0.99

0.99

0.99


Data
points


variabl










Table 6 continued


Independent
variable


Dependent
variable


Correlation
Slope coefficient


log Cl

SO0/Cl

C1/(C1 + S04)
log{Cl/(Cl + SO4))

log[Cl/(Cl + 2 SO4)

C1/(C1 + j SO4)

log Cl/(Cl + i S04)1

pH + logS04 + Cl))

SO /Cl

log1Cl/(C1 + S04)}

logfCl/(Cl + 1 SC 4)


Data
points


Epit(gr)

Epit(gr)

Epit(gr)
Epit(gr)

Epit(gr)

Epit(gr)

E .
pit(gc)
Epit(inc)

Epit(inc)
Epitine)


-0.96
0.65

-1.17
-2.09

-1.26

-1.90

-3.73
-1.67

0.62

-1.98

-3.54


decade1


V decade-1

V decade-1

V
decade-1
V decade-1

V

V decade-1

V decade-1


0.91

0.99
0.92

0.97
0.94

0.97

0.97

0.95
0.96

0.92

0.93










Table 6 continued


Independent
variable


Dependent
variable


Correlation
Slope coefficient


C1/(C1 + 2 SO4)

pH + log(SO4 + Cl)


Epit(inc)

Epit(inc)


-1.80 V


-1.57 V decade-1


a
Cl and S04 represent the analytical concentration of chloride and sulfate, respectively.
The brackets and charge designations have been omitted for clarity.


Data
points


0.92

0.95









critical current density (icrit) and the primary passivation

potential (E p ), respectively. All values of ic

measured here are between 1.3 x 10-5 and 9.50 x 10-5 A cm-2.

Although, in some cases, the lowest value of critical

current density observed at a given chloride ion concentration

approximates the highest seen in the next lower chloride ion

concentration, there is a definite trend in the average

values toward a reproducible maximum at 0.508 M potassium

chloride (Table 2).

The only relationship for which linearity is suggested

is that between the logarithm of critical current density

and the solution composition function, pH + log ( [SO42- +

[Cl1] ), where [S 42-] and [C1-] are the analytical

concentrations of sulfate and chloride, respectively. A

slope of 1.10 is measured (Figure 9). No linearity is

found between the logarithm of icrit and E contrary to
crit r
the work of Wilde.5 However, the most negative rest

potential and highest critical current density values

occur in the same system.

From its value at the maximum the current density

then decreases to a low value ( 10 A cm2 ) characteristic

of the passive region, i and is approximately independent

of potential. In some cases, prior to the attainment of

the passive current value, a negative current loop is

observed (CicC'). Experimentally determined characteristics

of the loop are presented in Table 3 and an experimental





























Figure 9. Variation of the logarithm of the critical
current density with the concentration
function, pH + log ( [S042-] + [ClI- ).


























-4.0-

0


0







0
-4.5-



0 0



0




-5.0-







2.0 2.5



pH + Log ([SOP + [ci])







49

curve containing a cathodic loop is presented in Figure 10.

The potential range in which it occurs, -0.30 to 0.0 V, is

independent of solution composition. The maximum

cathodic current density associated with the loop, ic

(1.36 x 106 A cm-2), is largest in solutions with no
chloride ion (0.334 M sulfate). Solutions containing

chloride ion show significantly lower current values

(2.3 x 10-7 to 4.2 x 10-7 A cm-2) which increase slightly

with increasing chloride ion concentration.

To minimize the effect of the negative current loop,

the magnitude of i is evaluated at 0.3 V except in 0.508

and 1.0 M potassium chloride where pitting occurs below

this potential. The dependence of the passive current

density on solution composition is qualitatively similar

to that of the cathodic loop current. The addition of

1.17 x 10-2 M chloride ion causes a substantial lowering

of i from the value observed in 0.334 M sulfate (5.25 x 10-7
P
versus 1.29 x 10-6 A cm-2, respectively) (Table 1). A

further increase in chloride content is accompanied by

a slight increase in passive current density. In this

region of increasing current, several linear relationships

between i and solution composition are observed (Table 6).

Because of the narrow range covered by the three values of

the passive current density available (5.25 x 10-7 through

7.85 x 10-7 A cm-2), the validity of these relationships

is questionable.

























Figure 10. Potentiostatic polarization curve showing cathodic loop obtained
in 0.334 M sodium sulfate at pH 2.40.










































-..
- -7.0





-5.0-





-3.0- /


-0.6 -0.4 -0.2 0.0 0.2 0.4 0.6 0.8 1.0 1.2


POTENTIAL (V, SCE)








The classical evaluation of the potential of total
passivity, ETp, involves the determination of the point
at which the current increases from its value in the
passive region into the anodic loop of the active-passive
59, 60
transition during a cathodic potential scan. In
the present work the potential of total passivity is
defined as the intersection of the line defined by the
decrease in anodic current from the maximum anodicc
potential scan) with the value of the passive current at
0.30 V. This method gives more reproducible results because
of the random cathodic loop observed here.
Values of ETp observed undergo a positive shift (from
-0.285 to -0.160 V) as the chloride ion concentration of
the solution is increased (Table 1). Examination of the
interrelation of E and solution parameters suggests
TP
several possibilities (Figures 11-15):


Relationship Slope r2

1. -ETp// [ClI] 0.108 V mole 1 0.93
2. -TEp/ [Cl-] /( [S042- + [C1-] ) 0.103 V 0.92
3. -Elp/ [Cl /(2[S0 2-+ [C1-] ) 0.110 V 0.94
4. -Ep/ pH + log ([042-] +[C1] ) 0.218 V 1 0.94
S+decade1

5. ETp/ I log ( [042-] / [Cl )o

a. two segments with slopes -0.016 and -0.039 V

decade-1 and correlation coefficients of 0.98

and 0.99, respectively or





























Figure 11. Variation of the potential of total passivity with chloride ion
concentration.






















0.15










c.O
LL 0.25-





I I I I I I -
0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0























Figure 12. Variation of the potential of total passivity with the concentration
function, [Cl] /( Is42] + [Cl )
( S0 + [Cl7 ]











015











0.20- -


FC0




0. 0
w

0

0.25--










01 02 03 0.4 05 0.6 0.7 0.8 0.9 1.0
[c ] /([so^] + [CL-])
[C L71 1 4J






















Figure 13. Variation of the potential of total passivity with the concentration
function, [C1-] /(2 [s0421 + [Cl] )*









0.15- -











0.20--
0

0

00

I-



0.25-
0








0.1 0.2 0.3 04 0.5 0.6 0o7 0.8 0o9 10 00
[c E-] / (2[so 0l + Ec ])





















Figure 14. Variation of the potential of total passivity with the
concentration function, pH + log ( [S04 -] + [Cl-] ).

































0.20+


pH + Log ([S0o:]+ [CL- )


I

























Figure 15. Variation of the potential of total passivity with the
concentration function, log ( [S042-] / [C1- )'
























0 15-






O 0.20--




.I-
I 0.25-







-1.0 0.0 1.0

Log 0LS4 r]/[Cl-









b. one segment with slope -0.030 V decade-I

and correlation coefficient 0.94.

The latter plot lends itself to two interpretations. It

may be considered to have two linear segments whose point

of intersection occurs at ETp = -0.230 V and 0.301 M

potassium chloride which is the smallest chloride ion

concentration investigated in which pitting occurs.

However, evaluation of data in terms of one single line is

also feasible. Differentiation between the two possibilities

is difficult because only five data points are available.

In solutions whose sulfate to chloride ion concentration

ratio is greater than one ( 0.3 M chloride), a continuing

shift to more anodic potentials results in the breakdown

of the passive state with the onset of pitting. An

experimentally determined polarization curve showing the

influence of pitting is given in Figure 16. Two values

are recorded for the potential at which pitting initiates,

Epit Epit(inc) is the most negative potential at which
any increase in current density with time is observed within

the ten minute waiting period. At potentials slightly

positive to Epit(inc) ( 0 to 150 mV), an initial decrease

in current followed by current spikes is usually observed.

At more positive potentials an immediate increase in

current density occurs. Epit(gr) is the potential obtained

by extrapolation of the increasing pitting current density-

potential curve back to the point at which it intersects

the value of the passive current at 0.30 V.

























Figure 16. Potentiostatic polarization curve showing the influence of
pitting obtained in 0.301 M potassium chloride at pH 2.35.























-3.0-






-5.0-






-7.0--
c'J
E
o




-7.0


0

-1

-5.0-







-3.0-


-0.6 -0.4 -0.2 0.0 0.2 04 0.6


POTENTIAL (V, SCE)










Both values of E pit are arbitrary because of the
pit
imposed time limitation for current measurement and the

heavy fluctuations in current of up to an order of

magnitude observed in this potential range, respectively.

Thus, the induction period for pitting may be greater than

ten minutes at potentials more negative than Epit(inc) and

the current magnitude, after ten minutes at a given potential,

is a random function of time as a result of the fluctuations

present. The graphically determined value of the pitting

potential, Epit(gr), is always the more negative. A

three-fold increase in chloride ion concentration shifts

Epit(inc) from 0.395 to -0.090 V and Epit(gr) from 0.360
to -0.155 V.

Several linear relationships are observed between

Epi and the solution composition parameters examined.
pit
Their corresponding slopes, intercepts, and correlation

coefficients are reported in Table 6. Because only three

data points are available, it is difficult to judge which

of these relationships, if any, are valid. It does appear,

however, that the potential at which pitting initiates

shifts toward more active values as the chloride ion

concentration of the solution is increased.

In systems not subject to pitting attack (< 0.3 M

potassium chloride), loss of passive state stability occurs

with the onset of transpassive dissolution (Table 4). The

disolution reaction in this region involves oxidation of










chromium(III) in the passive film to chromium(VI) and

exhibits Tafel behavior over 1 to 1.5 decades of current.

An anodic Tafel slope of 0.114 V decade-1 is observed in

0.0 M and 1.17 x 10-2 M potassium chloride, while a value

of 0.150 V decade-1 is found for 9.97 x 10-2 M potassium

chloride. The potential at which transpassivity initiates,

Etr, is taken as the intersection of the computed Tafel
line with the passive current density at 0.30 V. Etr shifts

negative from 0.671 to 0.575 V with increasing chloride

ion concentration.

Continued anodic polarization produces a slight current

maximum as a result of secondary passivity. Neither the

current magnitude at the maximum, irit,2 (1 x 10 A cm),
crit, 2
nor its associated potential, E pp,2 (0.95 V), is affected

by solution composition at constant pH. The current then

decreases to a minimum prior to oxygen evolution. The

degree of stability of secondary passivity is represented

by the difference in the magnitude of the maximum and

minimum current values, a i. Present results indicate that

the stability decreases slightly on addition of chloride

ion to solutions originally 0.334 M in sulfate ion at

pH 2.4. Current differences,A i, of 14 to 17 pA cm-2 and

21 A cm-2 are observed for 1.17 x 10-2 to 9.97 x 10-2 WM
potassium chloride and 0.0 M potassium chloride (0.334 M

sulfate), respectively.










Secondary solutions

A brief survey of the effect of chloride ion concentra-

tion in the absence of sulfate ion on the polarization

curve was carried out at pH 2.4 (hydrochloric acid) in

solutions containing 0.102, 0.123 and 1.0 M potassium chloride.

Characteristic potentials and current densities obtained

from the polarization curves are presented in Tables 7 and

8. Linear relationships with correlation coefficients

) 0.90 are reported in Table 9 with the corresponding

slopes.

Cathodic current magnitudes at -0.700 V are signifi-

cantly lower than those in sulfate containing solutions,

varying from 1.25 x 10-3 to 7.06 x 10- A cm-2 in the

chloride range 0.102 to 1.0 M. An increase in concentration

from 0.102 to 0.123 M causes only a slight shift in rest

potential (-0.470 to -0.480 V) while the rest potential

for 1.0 M potassium chloride occurs at -0.505 V, a negative

shift of 45 mV. The primary passivation potential is not

a function of chloride concentration. In the absence of

sulfate ion icrit changes linearly with chloride ion

concentration (Figure 17) and with the rest potential

(Figure 18). The corresponding slopes are 3.11 x 10-5 A

cm-2 mole-1 1 (r2 = 0.98) and -15.7 decades V-1 (r2= 1.0),

respectively. A linear relationship is also observed

between the logarithm of the critical current density and

the concentration function, pH + log ( [S04 2- + [Cl )

(Figure 19). The corresponding line has a slope of 0.48














[Cl-] [s4 2-]
x 101 x 101


Table 7

Potentiostatic current-potential behavior in secondary solutions
(average curves).

i-700 icrit 1 1
(A cm-2) Er Eppl (A cm- ) (A cm-2) Ep Ept(V,SCE)
pH x 104 (V,SCE) (V,SCE) x lo5 x 107 (V,SCE) inc gr


2.40

2.35

2.42


Polarization

-12.5 -0.470

-12.5 -0.480

-7.06 -0.505


in the

-0.420

-0.400

-0.420


Polarization

1.52 -203.0 -0.420 -0.370

2.35 -22.30 -0.470 -0.430

6.22 -0.0013 -0.760 -0.500


absence

1.18

1.69

4.17


of sulfate ion

-0.287

-0.273

-0.250


at varied pH


4.65

3.43

0.147


-0.681

7.85


-0.205

-0.230


1.02

1.23

10.0


3.01

3.01

3.01


0.233

0.248

0


0.140

-0.025

-0.100


0.070

-0.120

-0.132




0.397

0.360


0.425

0.395








Table 8

Fotentiostatic current-potential behavior in secondary solutions
(individual experiments).


[C 1 [S04x
x 101 x 101


i-700
(A cm 2)
pH x 104


icrit,l p
Er Epp,l (A cm ) (A cm-2)
(V,SCE) (V,SCE) x 105 x 107


ETp Epit(V,SCE)
(V,SCE) inc gr


Polarization in the absence of sulfate ion


1.02 0 2.40 -12.3

-12.6

1.23 0 2.35 -15.3

-9.61

10.0 0 2.42 -5.71

-8.42



3.01 2.33 1.52 -189.0

-196.0

-224.0

3.01 2.48 2.35 -18.0

-27.0

3.06 0 6.22 -0.0013


-0.460 -0.420


-0.480

-0.480

-0.480

-0.500

-0.500


-0.420

-0.400

-0.400

-0.420

-0.420


Polarization at

-0.400 -0.350

-0.440 -0.370

-0.400 -0.370

-0.475 -0.420

-0.480 -0.415

-0.760 -0.500


0.881

1.49

1.7

1.68

4.69

3.70


varied pH

4.1

5.8

5.25

3.55

3.39

0.147


-0.287

-0.282

-0.285

-0.270

-0.250

-0.235



-0.240

-0.230

-0.155

-0.230

-0.226


7.61

7.49

2.80

7.23

8.48


- 0.075

- 0.120

-0.05 -0.107

0.00 -0.120

-0.10 -0.155

-0.10 -0.120


0.450

0.400

0.400

0.420

0.370


0.335

0.400

0.460

0.365

0.360


































Figure 17. Variation of the critical current
density with chloride ion concentration.




































4.0--


2.0-4-


I II I I


[crt-


6.0--









74



















-4.8--







-4.6- -







-4.4--



0.50 0.49 0.48 0.47



-4.2








0.50 0.49 0.48 0.47


-Er (V,SCE)






























Variation of the logarithm of the critical
current density with the Boncentration
function, pH + log ( [S04 9 + [Cl-] ).


Figure 19.


















































2:0 2:

pH + Log ( [S04o]+ [I-] )

































Figure 20. Variation of the potential of total passivity
with rest potential.































0.24-t-


0.26--


0.28-
c 0.28-


0.30--


I I I I


-Er (V,SCE)



































Figure 21. Variation of the logarithm of the
critical current density with the
potential of total passivity.























































0.25


0.28 0.27 0.26


-ETP (V,SCE)








81
and a correlation coefficient of 0.93. No cathodic current

loop is observed in these solutions.

Little change in the potential of total passivity is

observed in going from 0.102 to 0.123 M potassium chloride

when data from individual runs are examined. Values are

-0.287 V and -0.282 V for the former and -0.285 V and

-0.270 V for the latter. However, average values show a

definite positive trend with increasing chloride ion

concentration in agreement with the value of -0.250 V

found in 1.0 M potassium chloride.

A one-to-one correlation between ET and Ep is

suggested (Figure 20). The corresponding slope is -1.03.

The log of icrit increases as Ep becomes more positive

at the rate of 15.1 decades V-1 (Figure 21).

In the absence of sulfate ion, pitting occurs at all

chloride ion concentrations investigated. A negative shift

in the pitting potential occurs with increasing chloride

ion concentration but no statistically valid variation in

this relationship is apparent.

The effect of pH on polarization curve parameters was

also examined in 0.3 M potassium chloride at pH 1.52, 2.35

and 6.22 (Table 7,8). In the former two cases, sodium

sulfate additions of 0.233 M and 0.248 M, respectively,

maintained the ionic strength at one. No sulfate was

added to the latter and polarization in this system was










terminated after passivation occurred. Possible linear

relationships are presented in Table 9.

Net currents at -0.700 V are cathodic and decrease

with decreasing hydronium ion activity. Values of

2.03 x 10-2, 2.23 x 10-3 and 1.34 x 10-7 A cm-2 are recorded

for pH 1.52, 2.35, and 6.22, respectively. A linear

relationship is observed between both the rest potential

and pH (-0.073 V pH-1) (Figure 22) and the primary passiva-

tion potential and pH (-0.025 V pH-1) (Figure 23). The

slope of the straight line resulting from a plot of log

icrit s pH is -0.328 (Figure 24). A linear relationship
is also observed between the logarithm of the critical

current density and the concentration function, pH +

log ( [S0 2-] + [Cl-] ) (Figure 25). A slope of -0.35

and a correlation coefficient of 0.99 are associated with

the line.

No cathodic loop is found in the pH 2.35 solutions

but one does occur between -0.15 and 0.0 V in two of three

experiments at pH 1.52 (Table 10). The maximum associated
-7 -2
cathodic current density is 4.5 x 10 A cm-

The increase in pH from 1.52 to 2.35 shifts the

potential of total passivity to more active values (-0.205

to -0.230 V) and changes the current density at 0.300 V from

cathodic to anodic values. An active shift in pitting

potential also occurs.








Table 9

Interrelationships in potentiostatic experiments in
secondary solutions.

Independent Dependent
Variable Variable Slope

[C1- i crit 3.11 x 10-5 A cm-2 moles
Er log i crit 15.7 decades V-1

Er ETP -1.03

ETp log i crit 15.1 decades V-1
pH + log ([O042-] + [Cl] ) log i crit 0.48

pH Er -0.073 V pH-1

pH Epp -0.025 V pH-1

pH log i crit -0.328

pH + loo ([SO42-j + [Cl] ) log i crit -0.35


Correlation
Coefficient

-11 0.98

1.0

0.99

0.99

0.93
1.0

0.91

0.99

0.99


Data
Points

3

3

3

3

3

3

3

3

3


~


























-Hd 1q;TmM T'ET9uaGod uoTOATlsspd ReuTjid eq. jo uoTeBaltjA "'3 ajnrTJ
















0.30--


0.40--


0.50--


I 1























*H-d M1qTM -[-reTueaod qseJ GqO Jo UOTLBTJeA *CZ aJnSTj



























0.5 0--


0,704-




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