Group Title: Influence of applied potential, fluid velocity, pH and temperature on formation of calcareous deposits under impressed current cathodic protection /
Title: Influence of applied potential, fluid velocity, pH and temperature on formation of calcareous deposits under impressed current cathodic protection
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Title: Influence of applied potential, fluid velocity, pH and temperature on formation of calcareous deposits under impressed current cathodic protection
Physical Description: x, 144 leaves : ill. ; 28 cm.
Language: English
Creator: Lee, Rupert Utak, 1951-
Publication Date: 1984
Copyright Date: 1984
 Subjects
Subject: Cathodic protection   ( lcsh )
Seawater corrosion   ( lcsh )
Corrosion and anti-corrosives   ( lcsh )
Materials Science and Engineering thesis Ph. D
Dissertations, Academic -- Materials Science and Engineering -- UF
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
 Notes
Thesis: Thesis (Ph. D.)--University of Florida, 1984.
Bibliography: Bibliography: leaves 138-143.
Additional Physical Form: Also available on World Wide Web
General Note: Typescript.
General Note: Vita.
Statement of Responsibility: by Rupert Utak Lee.
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Bibliographic ID: UF00097416
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 - 000491170
oclc - 11959391
notis - ACQ9672

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INFLUENCE OF APPLIED POTENTIAL, FLUID VELOCITY,
pH AND TEMPERATURE ON FORMATION OF CALCAREOUS DEPOSITS
UNDER IMPRESSED CURRENT CATHODIC PROTECTION






BY






RUPERT UTAK LEE


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


UNIVERSITY OF FLORIDA


1984

























































UNIVERSITY OF FLORIDA

I3 1262 08552 3404l
3 1262 08552 3404





















To God,

who has guided me

through my life











ACKNOWLEDGMENTS


I am grateful to

Dr. John R. Ambrose for his guidance through my

graduate work as well as this research;

to Drs. Ellis D. Verink, Jr., Paul H. Holloway, E. Dow

Whitney, and Gerhard M. Schmid for serving as members of my

supervisory committee;

to Mr. Suribabu Jayanti for his suggestions and

discussion;

and to Messrs. E.J. Jenkins and Guy P. La Torre and Dr.

Michael A. Kosinski for their assistance in chemical

analysis.

This work was supported by the National Program on

Marine Corrosion Grant No. NA 81AA-DOOO11.

Special thanks are extended to my wife, Clara, and my

parents for their support and encouragement.


iii


___












TABLE OF CONTENTS

Page

ACKNOWLEDGMENTS ........................................iii

LIST OF TABLES............................................vi

LIST OF FIGURES......................................... vii

ABSTRACT...................................................... ix

CHAPTER

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

2 REVIEW OF LITERATURE .............................. 5

2.1 Mechanism of Deposit Formation...............5
2.2 Properties of Calcareous Deposits ...........19

3 EXPERIMENTAL.....................................36

3.1 Materials........................ ..... . ... 36
3.2 Surface Preparation..........................38
3.3 Equipment ............................ .........58
3.4 Current Measurements ........................ 39
5.5 Morphology and Chemical Analysis............41

4 RESULTS........................................... 42

5 DISCUSSION.......................................66

5.1 Discussion of Background....................66
5.1.1 Rotating Disc Electrode .............. 66
5.1.2 Effects of Parameters................68
5.1.3 Ratio of Currents at Two Different
Rotation Speeds...................... 69
5.2 Discussion of Influence of Parameters
on Deposit Formation.........................74
5.2.1 pH...................................74
5.2.2 Temperature ..........................81
5.2.3 Potential.............................89
5.2.4 Fluid Velocity........................ 99
5.2.5 Deposit Morphology.................. 103
5.2.6 Chemical Analysis ...................105
5.2.7 Interaction Effects.................111






5.3 Surface pH Model........................... 113

6 CONCLUSIONS.....................................124

7 RECOMMENDATIONS FOR FURTHER RESEARCH............126

APPENDIX

1 CATHODIC PROTECTION ...........................128

2 24 FACTORIAL DESIGN ................................132

3 COMPOSITION OF IRON USED FOR ELECTRODES......... 137

BIBLIOGRAPHY.. ............. ............................... 138

BIOGRAPHICAL SKETCH....................................... 144












LIST OF TABLES


Table Page

1 Solubility product of various compounds
in seawater. ..................................... 7

2 Degree of supersaturation at various pH
values.............................................. 12

3 Ionic composition of ASTM D-1141-75
artificial seawater...............................13

4 Variation of solubility with temperature.........35

5 Parameters and their levels...................... 40

6 Current measurements under various conditions....43

7 Main and interaction effects.....................44

8 Effective surface coverage........................46

9 Combination of different experimental
conditions........... ............................... 47

10 Relative viscosity of seawater....................85

11 Values of k...................................... 108

12 Limiting equivalent ionic conductivities in
aqueous solutions at 250C and calculated
diffusion coefficients...........................120

13 Arrangement of 24 factorial design experiment,
duplicated......................................... 154

14 Signs for calculating effects of 24 factorial
design............................................. 136











LIST OF FIGURES


Figure Page

1 Decrease in the Ratio of i1000/i500 With Time....48

2 Variation of Hydrogen Current (i_1.0 i_0.8)
and Its Percentage With Time..................... 51

3 Variation of i-1.0 i_0.8 With Time.............52

4 SEM Micrograph of Deposit Formed at (0100);
-0.8V, 1000 rpm, pH=8.3, 23 C .....................53

5 SEM Micrograph of Deposit Formed at (0000);
-0.8V, 500 rpm, pH=8.5, 230C ..................... 53

6 SEM Micrograph of Deposit Formed at (1000);
-1.OV, 500 rpm, pH=8.3, 230C .....................54

7 SEM Micrograph of Deposit Formed at (0001);
-0.8V, 500 rpm, pH=8.3, 16C .....................54

8 SEM Micrograph of Deposit Formed at (0010);
-0.8V, 500 rpm, pH=7.9, 23C .....................55

9 Electrode Surface After Deposit Was Removed
With a Kimwipe Paper. Specimen Was Stored in a
Desiccator for Five Days .........................55

10 Energy Dispersive X-ray Spectra of Globular
Particle .......... .............................. 57

11 Energy Dispersive X-ray Spectra of Globular
Particles .......... .............................. 58

12 Energy Dispersive X-ray Spectra of Background
Layer............................................... 59

13 ESCA Survey Scan of Deposit Formed at (0000).....60

14 ESCA Mg Scan of Deposit Formed at (0000).........61

15 ESCA Fe Scan of Deposit Formed at (0000)..........62

16 ESCA Ca Scan (Low Resolution) of Deposit
Formed at (0000)..................................63


vii







17 ESCA Ca Scan (Medium Resolution) of Deposit
Formed at (0000).................................. 64

18 FTIR Spectra of Deposit Formed at (0000) for
2 (right scale) and 6 (left scale) hours.........65

19 Three-Dimensional Flow of Liquid Near
Rotating Disc Electrode...........................67

20 Cathodic Current vs. (Electrode Rotation
Speed) '12 ...................................... ... 71

21 Cathodic Polarization Curves. The Top Curve
Was Obtained After Deposit Was Formed for 4 1/2
Hours and the Bottom Two Curves Were Obtained
in Deposit-Free Solution. Electrode Rotation
Speed Was 500 rpm ................................73

22 Schematic Presentation of the pH Effects......... 75

23 Cathodic Polarization Curve ......................78

24 Schematic Presentation of the T Effects..........82

25 Schematic Presentation of the E Effects..........90

26 Schematic Presentation of the RPM Effects....... 100

27 Schematic Layouts of Cathodic Protection
System: Sacrificial Anode Method and
Impressed Current Method................. .. ..... 129

28 Diagram of Relative Energy Levels of Fe and
Fe++ (or Fe++ ) Before and After Cathodic
Protection....... ................. ................. 130


viii














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

INFLUENCE OF APPLIED POTENTIAL, FLUID VELOCITY,
pH, AND TEMPERATURE ON FORMATION OF CALCAREOUS DEPOSITS
UNDER IMPRESSED CURRENT CATHODIC PROTECTION

BY

RUPERT UTAK LEE

December, 1984

Chairman: Dr. John R. Ambrose
Major Department: Materials Science and Engineering

Formation of calcareous deposits is one of the

characteristic features associated with cathodic protection

in seawater environments. The deposits have two beneficial

aspects in cathodic protection: they decrease cathodic

current requirements and mitigate corrosion reactions.

The objective of this research was to characterize the

effects of cathodic protection parameters on formation of

calcareous deposits. Four parameters were selected:

cathodic potential, fluid velocity, pH, and temperature. A

factorial design technique was employed to investigate the

influence of each parameter and also to assess the order of

importance among the parameters. Within the experimental

conditions studied, the largest influence on deposit

formation was found to be associated with pH, with







temperature, potential, and flow velocity, decreasing in

this order.

Morphological studies showed that the size of deposit

particles increases with increase in deposition rate, which

is a function of surface pH and degree of supersaturation.

These findings could be explained by a heterogeneous

nucleation and growth mechanism.

Chemical analysis did not show any trend in the

variation of chemical composition with the variation of

parameters.

Qualitatively, the influence of the four parameters

could be explained in terms of surface pH and degree of

supersaturation. For quantification of the influence, a

surface pH model was proposed. It was found that the

quantification cannot be fully completed until the kinetics

of Mg(OH)2 formation is fully described.













CHAPTER 1
INTRODUCTION


Corrosion is the destructive attack on a metal by its

electrochemical reactions with its environment. Recent

technical innovations have enabled man to carry out more

activities in seawater environments, thereby increasing the

demand for performance from materials. Metals are parti-

cularly susceptible to high chloride concentration, high

electrical conductivity, and biological activities of

seawater.

One of the successful ways to cope with seawater

corrosion is to apply cathodic protection (Appendix 1) by

impressed current or by sacrificial anode methods. One of

the characteristic features associated with cathodic

protection in marine environments is formation of a white

scale mainly comprised of calcium and magnesium compounds on

metal surfaces. These deposits, called calcareous deposits,

can not only reduce cathodic current requirements by

physically forming a resistive barrier between metal and

seawater but also mitigate corrosion reactions when cathodic

protection current is removed by accident or for main-

tenance. Although these aspects of calcareous deposits have

been well known for decades, little effort has been expended

to take better advantage of the deposits. Consequently, the








formation of these deposits has not been introduced into the

engineering of cathodic protection systems as a design

parameter. In this respect, it is important to know the

nature of deposit formation. This research has been

directed toward providing practical information applicable

to cathodic protection design, especially with respect to

the following questions:

What is the most important parameter in determining
the rate of calcareous deposit formation?

What is the relative degree of influence of
parameters, e.g., temperature and flow velocity, on
formation of deposits?

What is the optimum condition for deposit
formation? Can a better utilization of the
protective nature of calcareous deposits be
achieved when cathodic potential (or current
density), practically speaking, the only
controllable parameter, is properly adjusted?

Four parameters were selected: cathodic potential,

solution flow velocity, pH, and temperature. Instead of

circulating solution at different flow rates, the specimen

electrode was rotated. In addition to the capability to

simulate various flow velocities, the rotating electrode has

the advantage of controlling mass transport. By means of a

24 factorial design method (Appendix 2), influence of the

four parameters on cathodic current requirements and on

formation of calcareous deposits were investigated. The

same experiments were also performed in solutions which did

not have deposit forming elements such as Ca++, Mg++ and

Sr++ in order to find influence of the four parameters on








rate of cathodic reactions, uninfluenced by occurrence of

surface deposits.

Degree of coverage of the metal surface with deposit

was studied from two different points of view. First,

effective surface coverage which was defined as cathodicc

current requirement without deposit that with deposit

formed on surface)/(cathodic current requirement without

deposit) was studied. Because the numerator is the decrease

in cathodic current caused by formation of deposit and the

denominator is the cathodic current in the absence of

deposit, the ratio multiplied by 100 is actually the

percentage decrease in cathodic current contributed by

calcareous deposit. Therefore, variation of effective

surface coverage can be related to the influence of any

parameter on formation of calcareous deposits. Secondly,

electrode rotation speed was varied. At high rotation

speed, i.e., at high flow velocity, reactants and products

of cathodic reactions are supplied and removed at an

increased rate, thus increasing cathodic current. Because

the calcareous deposit functions as a barrier to the

movement of these species, different types of deposit will

have different barrier efficiencies, thus different cathodic

currents. Two electrode rotation speeds were used, 500 and

1000 rpm. After a deposit was formed at one of the two

rotation speeds, rotation speed was changed to the other.

Current was measured before and after the rotation speed

change. (Current at 1000. rpm)/(current at 500 rpm) was








plotted against time of cathodic polarization. This was

repeated until the ratio dropped to an arbitrary value of

1.050.1 at which the deposit assumes almost full control of

mass transport. By comparing decay rate of the ratio at one

set of conditions with that at the other, the influence of

the four parameters on deposit formation rate could be

studied.

Chemical compositions of calcareous deposits were

investigated with various techniques such as energy dis-

persive X-ray analysis, X-ray photoelectron spectroscopy and

Fourier Transform infrared analysis.

The pH near cathodically protected metal surfaces has

been proposed by several researchers. This surface pH is a

very important parameter in controlling deposit formation

rate. A mathematical model, including reactions affecting

surface pH, was established.













CHAPTER 2
REVIEW OF LITERATURE

2.1 Mechanism of Deposit Formation

Cathodic protection systems supply negative charges to

a metal to be protected from corrosion attack (See Appendix

1). At the negative potentials of cathodic protection, two

cathodic reactions can take place on the metal surface in

seawater environments.


02 + 2H20 + 4e- + 40H~ [1]


2H+ + 2e- + H2 [2]


Depending upon the potential, oxide film on the surface and

other reducible species, if present, can also be reduced.

The result of reaction [1] is exactly the same as that of

reaction [2] in increasing OH- concentration at the metal-

solution interface area. Increase in OH- concentration

dissociates HCO.
3.

HCO + H+ + CO [3]


As OH- and CO concentrations increase beyond the

solubility product limit of Mg(OH)2 and CaCO3, these two

compounds precipitate on the metal surface.


Mg++ + 20H- + Mg(OH)2 [4]








Ca++ + C0 + CaCO3 [5]


The pH above which Mg(OH)2 is supersaturated can be

calculated from


]2 Ks,Mg(OH)2
[OH-] =
[Mg+t]

where Ks,Mg(OH)2 = solubility product of Mg(OH)2.

From the solubility product of Mg(OH)2 in Table 1 (1-4) and

from the Mg++ concentration in seawater of 5.46x10-2M, the

[OH-]=2.1x10-5M or pH=9.3 can be calculated above which

Mg(OH)2 is supersaturated. Pytkowicz et al. (5) observed

that Mg(OH)2 precipitation started in seawater between

pH=9.8 and 9.9. If the precipitation process is aided by

the presence of a metal surface in its nucleation and

growth, it may require a lower degree of supersaturation,

i.e., a lower surface pH than the above calculation for

initiation.

A very high surface pH, greater than 11, under cathodic

protection conditions in seawater has been suggested by

several authors. On the assumption that OH- is transported

away from the metal surface by diffusion, Engell and

Forchhammer (6) and Wolfson and Hartt (7) equated


flux 1 flux
02 4 OH-

and calculated such a high surface pH. The factor of 1/4 is

to account for the fact that one mole of 02 generates four












Table 1 Solubility product of various compounds in seawater


Compounds

CaCO3 (calcite)

CaCO3 aragonitee)

SrCO3

MgCO3

MgCO3.H20

Mg(OH)2

CaSO4.2H20


log Ks

-6.2

-6.05

-4.46

-3.52

-3.15

-10.16

-4.37


Temp (oC)

25

25

25

0



25

25


Reference

1

1

Note

2

2

3

4


Note: Based upon reported solubility product (1), corrected
for ionic strength at seawater (2).








moles of OH- as in equation [1]. At a more negative

cathodic potential, the hydrogen evolution reaction also

takes place supplying more OH- to the surface region, and

surface pH could go even higher. When a [HCO ] of 2x10-3M

in seawater is considered, such a high pH is unrealistic.

The actual pH will be substantially lowered by the buffer

capacity of HCO The method to calculate surface pH used

by Guillen and Feliu (8) is probably more accurate. Using

the following mass balance equation for OH-


Nf = Nd + Np

where Nf = amount generated by cathodic reactions)

Nd = amount diffused away from surface

Np = amount consumed by precipitation,

they suggested a surface pH of greater than 10 which varied

with current density. From the weight and the chemical com-

position of deposit, the amounts of Mg(OH)2 and CaCO3 were

calculated. Subsequently, they calculated the amount of OH-

associated with Mg(OH)2 and necessary to produce the amount

of CO-- associated with CaCO This procedure for surface

pH estimation is valid only under the assumptions that all

the precipitate particles are deposited on the cathode

surface and that the precipitation of CaCO3 does not need

any supersaturation which is not true. However, as will be

discussed later in this section, precipitation of CaCO3 is

retarded by the presence of other species in seawater.

Therefore, all the generated C0-_ are not coordinated with







Ca++ to form CaCO3 and a certain amount of CO- will be

transported away from the surface. When this effect is

taken into account, a lower surface pH will be calculated.

A brucite structure was found from the X-ray analyses

of calcareous deposits (5,9,10) indicating that Mg(OH)2 in

calcareous deposits is present in the crystalline form.

On the other hand, CaCO3 is supersaturated in

seawater. From [Ca++]=1.04x10-2M and [C03-]=1.4x10-4M at

pH=8, which are the [Ca++] and [COy-] concentrations of the

solution used for this study and also represent the concen-

trations of the two species in seawater, the product of
[Ca++]'[CO~-]=1.5x10-5 can be obtained which is substan-

tially greater than its solubility product of 6.3x10-7

(calcite). Once a solution is supersaturated with CaCO ,

the rate of its deposition depends on the kinetics of the

deposition reaction. There are several applicable experi-

mental equations (11-15). However, those suggested rate

equations can be generalized into

d[CaCO ]
rate = -- -
dt

L K [Ca++]a[CO-]a 1/a
= K }
s,CaCO3 K a
s,CaCO3

where L = constant including rate constant, activity
coefficient and/or surface area of seed
crystals
Ks,CaCO3 = solubility product of CaCO3

[Ca++] = concentration of Ca++








[CO0-] = concentration of CO~

a = experimental constant.

From the equation, it can be concluded that the rate of

CaCO3 formation depends upon the degree of supersaturation,

which is defined as


[Ca+ ][CO3 ]
KS


Also, complex ion formation has to be considered. Studies

(16-18) show that about 90% of COT- are coordinated with

Na+, Mg++, and Ca +. Therefore, the degree of supersatura-

tion could be less than the above calculation indicates.

Also, as will be discussed later, the precipitation of CaCO3

from naturally occurring supersaturated seawater does not

occur in reasonable experimental periods (19-21) because of

the effect of Mg++ and other species on its deposition

kinetics.

At a cathodic potential less negative than the poten-

tial for the initiation of hydrogen evolution reaction, only

reaction [1] takes place on a steel surface whereas at a

potential more negative than that, both reactions [1] and

[2] take place leading to a higher pH near the metal sur-

face. From the pH consideration, it is conceivable that at

a low surface pH (still higher than bulk pH), CaCO3 is the

predominant species of the deposits whereas at a high sur-

face pH both CaCO3 and Mg(OH)2 precipitate. Guillen and








Feliu (8) found that Ca-to-Mg ratio decreased with increase

in surface pH. He calculated degree of supersaturation at

various pHs for CaCO3 and Mg('OH)2. As can be seen in Table

2, the degree of supersaturation of Mg(OH)2 increases far

more rapidly than that of CaCO3. From the solubility data

of Table 1 and concentration of ions in seawater as shown in

Table 3 (22), it is very unlikely that other compounds would

form as precipitate. However, MgCO3 and SrCO3 can form a

solid solution with CaCO3 due to the increase in activity

coefficient of minor constituents in solid solution.

The incorporation of Mg++ in CaCO3 can be explained

from the study of the effect of Mg++ on CaCO3 nucleation and

growth. It has been known for decades that CaCO3 formation

is retarded in the presence of Mg++. There are two

hypotheses about the inhibitory effect of Mg++. The first

hypothesis is that Mg++ is a poison to CaCO3 formation.

Pytkowicz (19,23) suggested that Mg++ acts as a surface

poison by being adsorbed as a hydrated ion, and its presence

increases the nuclei size and the number of collisions

between Ca++ and CO -. Also, Mg++ adsorbed on active growth

sites such as kinks inhibits the spread of monomolecular

steps on the crystal surface. This hypothesis was also

provided by Lippman (24) in a slightly different way. He

suggested that Mg++ ions are adsorbed on active growth sites

and congest the site due to their large sheath of

hydration. Pytkowicz (19) studied the time of CaCO3












Table 2 Degree of Supersaturation at various pH values


++ --
[Ca ][CO3 ]

s, CaCO3


18.1

28.8

53.7

79.0


174

281

349


2
[Mg++][OH-]

K s, Mg(OH)2

2.17x10-3.2

2.17x10-3

2.17x10-2.4

2.17x10-2

2.17x10-1

2.17

2.17x101


Data used for these


calculations:


pK1 = 6.0, pK2 = 9.1 (2)

Ks,CaCO3 = 6.2 (1)

Ks,Mg(OH)2 = 10.16 (3)
[Ca++] = 1.04x10-2, [Mg++] = 5.46x10-2


7.9

8.0

8.3

8.5

9.0

.9.5

10.0











Table 3 Ionic composition of ASTM D-1141-75 artificial
seawater


Ions Concentration (mole/liter)

Na+ 6.52x10-1

Mg++ 5.46x10-2

Ca++ 1.04x10-2

K+ 1.02x10-2

Sr++ 9.03x10-5

Cl- 7.24x10-1

SO 3.30x10-2

HCO 2.39x10-2*

Br- 8.49x10-4

F- 7.14x10-5

H3B03 4.37x10-4


Note: Chlorinity of this artificial seawater is 19.38 o/oo
(22). Salinity is 35 0/oo, calculated from salinity
= 0.03+1.805 chlorinity.

* as added







nucleation in natural seawater, Mg++ free artificial sea-

water and Mg++ enriched natural seawater after adding

CO03. The induction period for nucleation was determined by

monitoring the time lapse before the increase in pH. From

these observations, it was concluded that Mg++, at its

nominal seawater concentration, inhibits the formation of

CaCO3 and is the predominant factor in determining the time

for the initiation of nucleation. Reddy and Wang (25)

studied the growth rate of CaCO3 on a seed CaCO3 crystal in

CaCO3 supersaturated solutions with and without Mg++. The

rate of deposition was determined by monitoring the

concentration of ions as a function of time. Mg++, at

10-5M, had almost no inhibitory effect on crystallization.

In contrast 10-1M nearly stopped CaCO3 growth. However, the

introduction of induction period was not observed in their

seed crystal experiments, apparently due to the presence of

the seed crystal. A Langmuir adsorption isotherm fitted

their data well, confirming that an adsorption mechanism can

explain the dependence of crystal growth rate on the

adsorption of Mg++ on crystal growth sites. Similar

observations as to the inhibitory effect of Mg++ on CaCO3

nucleation and growth have been made by several authors

(20,24,26).

The second hypothesis was conceived more recently.

Berner (27) suggested that Mg++ may serve as a surface

poison, but is also incorporated into the growing crystals

to such an extent that the solubility of the








magnesium-containing calcite is markedly increased. This

increase in solubility slows down the formation of CaCO3.

As more studies have been carried out in natural

environments, the second hypothesis is becoming the more

acceptable mechanism to explain the effect of Mg++ on CaCO3

deposition in seawater environments.

It has been known for decades that natural CaCO3

sediments have some magnesium concentration. Goldsmith et

al. (28) found from X-ray analysis that natural carbonate

sediments contain up to 9% of MgCO3. From the study of

natural carbonate skeletal materials of marine organisms

containing up to 30% MgCO3, Chave et al. (29) reported that

the order of solubility is calcite with large MgCO3,

aragonite and low MgCO3 calcite, decreasing in this order.

In a study of the stability of magnesium containing calcite,

i.e., magnesian calcite, Plummer and MacKenzie (30)

monitored the variation of pH with magnesian calcite

dissolution and, from the extrapolation of time to infinity,

they found that 2 and 24 mole % of magnesian calcites are

stable or at least metastable. Thorstenson and Plummer (31)

emphasized the thermodynamic stability of magnesian calcite

system. Based upon the thermodynamic data provided by

Plummer and MacKenzie (30), they found that in two magnesian

calcite ranges of 1-5 and 20-25 mole %, magnesian calcite is

stable and metastable, respectively. Thorstenson and

Plummer also stated that the solubility of magnesian calcite

is controlled by [Ca++]'[Mg++] in seawater. This concept







that [Ca"+]'[Mg"+] controls magnesian calcite stability or

solubility was disputed by Lafon (32), who questioned the

validity of the extrapolation method used in Plummer and

MacKenzie's experiments (30)(also disputed by Stumm and

Morgan (1)). The relative stability of magnesian calcite

and pure calcite can be studied from simple thermodynamic
equilibrium equations as provided by Stumm and Morgan (1).

The equation for the dissolution of magnesian calcite

containing x mole fraction of MgCO3 can be expressed as
reaction [6].


Ca(lx)MgxCO3 + (1-x)Ca++ + xMg++ + COT- [6].

Also, dissolution of pure calcite, expressed in the reverse

direction, is


Ca++ + CO + CaCO3 [7].

The equilibrium constants are K(x) and (Ks,CaC3)1 for

reaction [6] and [7], respectively. To compare the
stability of magnesian calcite to that of pure calcite,

reaction [6] and [7] can be added to give reaction [8].

Ca(1x)MgxCO3 + xCa++ + CaCO + xMg+ [8]

The equilibrium constant is

([Mg++])x
[Ca++]

which is also equal to the product of K(x) and (Ks,CaCO3)-1







Therefore,


([Mg++])x K (K )-1
[Ca++] (x) s,CaCO3

Rearranged,
([Mg++])x K
K(x) = [Ca*+] s,CaCO5


This equation states that the relative solubility of mag-

nesian calcite containing x mole fraction of MgCO5 is

governed not only by the mole fraction (x) of MgCO5 in

calcite but also by the ionic concentration ratio between

Mg++ and Ca++ in seawater. Because [Mg++]/[Ca++] is always

larger than 1 in seawater, K(x) becomes greater than
Ks,CaCO3 as mole fraction (x) increases. Although effect of

solid solution formation and kinetic aspects are not

considered in this simple thermodynamic equilibrium, it can

be said that naturally observed magnesian calcite with a

high MgCO3 mole fraction is not a stable form. This

assertion is supported by the observation made by Winland

(35) that high magnesian calcite underwent transformation

into aragonite. In seeded crystal experiments, Berner (24)

found that magnesian calcite containing 7-10 mole% of MgCO5

grew on the pure calcite seed crystals. Recently, Morse et

al. (34) observed magnesian calcite formed on calcite seed

crystals in surface seawater. Auger depth profiling showed

no significant variation of magnesium signal. Quantitative

analysis showed about 4 mole % of MgCO3 in calcite.








Whether the stability of magnesian calcite is con-

trolled by kinetics or thermodynamic equilibrium is still in

dispute. However, observations of low magnesian (low-

magnesium) calcite formation (24,30,31,34), even though

magnesian calcite is more soluble than calcite, are

generally believed to be due to the inhibiting effect of

Mg++. Using the solubility of calcite and dolomite reported

by Hsu (35), Stumm and Morgan (21) stated that dolomite,

CaMg(C03)2, is the more stable form. However, the

comparison of the solubilities obtained in Florida under-

ground water to seawater environments is not appropriate.

Crystalline CaCO3 has another structure, metastable

aragonite. As shown in Table 1, and also measured by Morse

et al. (36), aragonite is more soluble than pure calcite.

It is generally believed that calcite nucleation and growth

are retarded in the presence of magnesium ions and, as a

result, aragonite precipitation is kinetically favored.

Formation of aragonite is reported by several authors in

various environments (20,24,27,37). However, it is hard to

believe that naturally occurring CaCO3 is the pure form of

aragonite or calcite (1,31). The formation and stability of

aragonite in seawater environment can be characterized in

the same way as was calcite; even though metastable

aragonite formation is kinetically favoured over calcite,

incorporation of Mg++ into calcite is the dominant factor in

CaCO3 deposition (10,24,27,30,31,38).








There are other species which are known to have a

similar inhibitory function to that of Mg++ for CaCO3

precipitation: organic materials (39,40), sulfate (25) and

phosphate (41). These species are generally believed to

behave like Mg++: occupation of growth site, charge

neutralization, and growth site congestion. Even though a

large amount of these species retard or prevent CaCO3

precipitation, the influence of these species at seawater

level concentration is not well known.

2.2 Properties of Calcareous Deposits

Beneficial aspects of calcareous deposits can be viewed

as bifold:

i) reduction of surface area directly exposed to
seawater, and consequent decrease in cathodic
current requirements.

ii) mitigation of corrosion attack during power
interruption, either accidental or for maintenance.

These advantageous aspects of calcareous deposits have

been known for decades. In the early 1940's, Cox realized

the beneficial effects of the deposits and suggested

applying high current density of several hundred mA/ft2 to

obtain a coating mainly comprised of Mg(OH)2 and Ca(OH)2,

the latter eventually transforming to CaCO3 in the presence

dissolved CO2 (42,43). He viewed the deposit as a mixture

of calcium salts dispersed in a matrix of Mg(OH)2 which is

serving as a bonding agent as cement bonds gravel in con-

crete. The application of this deposit to offshore drilling

platforms with initial high current densities functioned








successfully in lowering total cathodic current requirements

and also in mitigating corrosion problems after cathodic

current was stopped for several months (44).

The first empirical relationship between properties of

calcareous deposits and current densities was reported by

Humble (45). He analyzed the deposits formed in seawater

and found the variation of Ca-to-Mg ratio with current

density. He concluded that low current densities provide

the conditions favorable for CaCO3 deposition as expected

from solubility limits (see also Table 2). Because the

deposits are formed in almost any situation where current

densities are large enough to maintain protective potentials

of steel, he claimed that self-healing ability of the

deposits is one of the inherent properties that can be

effectively employed to protect almost any surface

regardless of shape and also seal any breakdown in the

deposit which may occur from time to time.

The protective property of the deposits was also

observed on galvanized steel. Denison and Romanoff (46)

attributed the high corrosion resistance, exhibited by most

of their galvanized specimens after outer Zn coating was

removed by corrosion, to a thin coating deposited

cathodically by the galvanic action between the outer Zn

coating and the underlying steel.

Cox's view on calcareous deposits as an analogy of

cement-gravel structure was well described in La Que's

discussion (47). Cox noted that calcium-rich deposits were







more porous and magnesium-rich deposits were easily

spalled. By using Humble's results (45), he suggested a

medium range of current density for a proper mixture of the

two properties. The range of 80-300 mA/ft2 was recommended

for "best" coatings and 50-400 mA/ft2 for "useful"

coatings. Preiser and Silverstein (48) agreed with Cox's

view of calcareous deposits as mixtures of calcium and

magnesium compounds. They recommended 100-200 mA/ft2 for

the best coating in which microscopic calcium (compound)

particles prevent shrinkage cracks in the magnesium

(compound) matrix. This was expected to result in a

relatively tough and nonporous coating. They also reported

several cases in which calcareous deposits were actually

applied to Navy warships. According to this report, as far

as the effectiveness of the deposits is concerned, the

applications were very successful in mitigating corrosion

problems at the sites where painting was damaged. However,

because ship hulls were repainted during the dry docking

period, the presence of the deposits caused a problem in

cleaning the ship hulls. Another difficulty was the control

of biofoulings. It was found that fouling problems could

not be handled with calcareous deposits as easily as with

anti-fouling paints.

Application of calcareous deposits to a ship hull was

also tried by Harvey and Streever (49). The paint on the

hull was removed in three places and the surface cleaned.

Initially 100 mA/ft2 was applied for 80 hours and then







current was maintained over 110 days with 12 hours on and 12

hours off. They reported no perceptible corrosion on the

bare steel surface at any time during the experiment. From

the context, it could be found that they tried to obtain a

uniform coating which would replace paint or be overcoated

with paint. Such uniform coatings could not be observed.

Apparently, the object of the Cox process (42,43) was to

develop a deposit which would obviate the need for

painting. From the corrosion standpoint, such a deposit is

satisfactory. However, its application to ship hulls could

not remove the problems of nonuniformity and biofouling

(48,49).

Calcareous deposits were also found on aluminum under

cathodic protection conditions (50,51). Kole (51) studied

the relationship between the deposit weight and the total

amount of charge on steel, aluminum and galvanized iron.

Potential varied between -0.992 and -1.473V vs. Ag/AgC1.

The weight to charge plot showed a linear relationship

between them without any noticeable difference among the

three different materials. This indicates that the deposit

gained weight from the cathodic charge throughout the 804

hour period and that the difference in substrate material

did not affect the deposit formation.

Using the IR internal reflection technique, Smith and

Mattson (52) studied calcareous deposits formed on a

germanium plate in seawater at -1.15V vs. SCE. The peaks of

C0O- and SO were observed. Deposits had two different








morphological shapes, one a forming plate-like background

layer and the other acicular particles on top of the

background layer. In X-ray analysis, the plate-like

background showed strong signals of magnesium and

chlorine. The acicular crystals showed magnesium and

calcium along with chlorine, potassium, sulfur, and

germanium. From these analyses, they concluded that the

deposits had carbonates, hydroxides, and sulfates of calcium

and magnesium.

Pirogov et al. (53) studied calcareous scales formed in

Black Sea. Three potentials of -0.85, -1.0, and -1.2V vs.

Ag/AgCl electrode were employed for their 300 hour

experiments. At -0.85V, a fine-grained, dense, and thin

film with good adhesion to the metal was observed. A

continuous film was completed after 210 hours. At -1.OV,

the deposit became more porous and coarse grained. A

continuous film was observed after 140 hours. At -1.2V, a

70-hour period was necessary to form a completely continuous

film. The thick film was coarse-grained and friable and was

easily separated from metal surface. They stated that

hydrogen evolution reaction caused the easily-spalling film

and provided frequent fluctuation of current as a supporting

evidence. In an analogy to Ohm's law, R = V/i, they

calculated the resistance of the deposit after the 300 hour

polarization was completed.









R = applied
deposit i


The deposit formed at -0.85V showed Rdeposit=5.0 (rnm2), at

-1.0V, Rdeposit=2.5 and at -1.2V, Rdeposit=1.2. It was

concluded that -0.85V deposit had the highest resistance

and, therefore, was the best deposit. This method of

comparing one resistance to another needs careful appli-

cation for two different potentials because the resistance

is a function of potential. The electrochemical reactions

at -0.85V are different from that at -1.OV and -1.2V in

their rates. Also, the hydrogen evolution reaction is

activation controlled and current increases exponentially

with potential. As a result, if two imaginary deposits with

exactly same properties are tested with this method, a

higher cathodic potential always yields a lower resist-

ance. To validate such resistance measurements, the same

potential has to be used.

Observation of deposit formation in city water

environments was made by Schwerdtfeger and Manuelle (54).

Spectrochemical analysis of the coatings on specimens

cathodically protected for 5 years showed strong signals of

Ca, Mg and Si. The coatings were called carbonaceous or

silicious deposits, apparently because of the richness in

silicon.

Using artificial seawater, Grigorev and Popov (55)

studied the deposits formed at different current densities








of 50, 100, 150, 200 and 250 pA/cm2 up to 280 hours.

Deposit weight was found to be proportional to deposition

time and current density. Degree of protectiveness of the

deposits was studied by comparing deposit weights to a

"minimum protective" current density which was defined as

the cathodic current density to achieve -0.58V vs. SHE

("minimum protective" potential). The 280 hour polarization

time was equally divided into four divisions. At the end of

each division, the cathodic current density to achieve the

"minimum protective" potential was measured and plotted

against deposit weight. The curves had the shape of a y=1/x

curve. The inflection point was at about 2 mg/dm2, which

corresponded to approximately 70-210 hours depending upon

applied current density. After the inflection point, the

current leveled off and further increase in deposit weight

had little effect on the "minimum protective" current.

Their observation indicates that after the film develops to

a certain extent, the film may break or develop cracks so

that further increase in weight provides no additional

resistance to the film in retarding the flow of current

across the film. Also the presence of the "best" current

density (150 WA/cm2) in terms of the "minimum protective"

current density indicates that properties, physical or

chemical, vary with current density applied.

Since calcareous deposits started to attract interest

from scientists and engineers, Ca-to-Mg ratio has been

continuously discussed. The belief is, as shown in Humble's







analysis (45), that as current density increases, surface pH

increases and the ratio decreases. Because CaCO3 is already

supersaturated in seawater while Mg(OH)2 is not, richness in

calcium content has been thought to indicate an enduring

deposit. The variation of Ca-to-Mg ratio with current

density and temperature was also reported by Rodrigo (56).

The ratio increased up to 5, which was observed at 50 pA/cm2

and 60CC. Guillen and Feliu (8) also reported similar

behavior of the ratio with current density and tempera-

ture. Deposits were formed in Barcelona seawater at 0, 20,

40 and 600C by supplying cathodic currents of 10, 50, 100

and 150 vA/cm2 with each experiment duplicated in the static

and in the aerated (intended for agitation) condition. For

example, at 200C and aerated, the ratio increased from 0.1

at 150 pA/cm2 to 2.4 at 50 PA/cm2. In Table 2, variation of

degree of supersaturation of Mg(OH)2 and CaCO3 with pH is

listed. From the rapid increase in Mg(OH)2 degree of

supersaturation with surface pH increase, the variation of

the ratio was explained. Deposit weights were measured at

different temperatures. They stated that, at high

temperature, the diffusion coefficient increases whereas

diffusion layer thickness decreases, according to Glasstone

(57). Increase in diffusion coefficient and decrease in

diffusion layer thickness accelerated the arrival rate of

Ca++ and Mg++ from the solution. An accompanying phenomenon

of accelerated OH- diffusion rate from metal surface to bulk

solution was not considered. Noting that the concentration








of OH- is the primary driving force for deposit formation,

it has to be considered for a complete explanation. From

the context, however, it can be read that the authors

believed that the positive effect (accelerated diffusion of

Ca++ and Mg++) overcomes the negative effect of accelerated

diffusion of OH-. Hartt et al. (58) suggested, as an

explanation to the deposit weight-temperature relationship,

that precipitation kinetics is enhanced at higher

temperatures.

Recently a few papers have been published in this

area. Wolfson and Hartt (7) studied the effect of potential

and seawater velocity on the protective properties of the

deposits. Potentials were varied among -0.78, -0.93 and

-1.03 vs. SCE and velocities were varied among 8, 31 and 107

cm/sec. Current density dropped gradually and leveled off

after 100 hours. Film thickness increased with increase in

time and cathodic potential and decreased with increase in

solution flow rate. Film thickness was larger at -1.03 than

at -0.93V and final current density was lower at -1.03.

However, the same relationship did not always hold. Their

observations are very similar to those of Grigorev and Popov

(55) in two aspects:

1) Wolfson and Hartt noted the presence of "better"
current density or potential which yielded "better"
deposit in terms of potential or current
measurement

2) Wolfson and Hartt noted that as the deposit grew in
thickness with time, current density or "minimum
protective" current density leveled off and further








increase in thickness had little effect on the
protective property of the deposit measured as
current.

A critical review paper was published recently (58).

Discussing chemical compositions of deposits, the authors

stated that several weight percentages of strontium were

reported in one case (59). Though the coprecipitation of

SrCO3 with CaCO3 forming a solid solution is not

unanticipated, such a concentration is too high. When two

solid materials form a solid solution in equilibrium with

solution, the solubility of the minor constituent is

appreciably reduced by the effect of the increase in its

activity in the solid solution. In the case of SrCO3

forming a solid solution with CaCO as illustrated by Stumm

and Morgan (1):



K [Ca++] X f
s,CaCO3 SrCO3 SrCO3
Ks,SrCO3 [Sr++] X fCaCO
5 CaCO^ 3



fSrCO3
= D .
obs fCaCO



where Ks: solubility product

D: distribution constant calculated from
solubility data

Dobs: observed distribution constant

X: mole fraction in solid solution

f: activity coefficient in solid solution.








It can be seen from this equation that the increase in

activity coefficient of the minor constituent (SrCO3) of the

solid solution increases its mole fraction in the solid

solution over the ratio that is calculated from solubility

data. Using the reported values of D=10 and Dobs=0.14 (33),
XSrCO 0.2 mole% can be obtained. This value is in good

agreement with Humble's analysis (45) of 0.1-0.5 mole%. The

cited value is far off from the expectation of equilibrium

thermodynamics, probably controlled by precipitation

kinetics. Wolfson and Haart reported that Hilbertz (60)

found aragonite structure in calcareous deposits from X-ray

analysis. On Guillen and Feliu's data about deposit weight-

temperature relation (8), Wolfson and Haart stated that

"consistent with this, it is generally recognized by the

cathodic protection industry that calcareous deposits form

more readily upon metal surface in warm water than in cold

water" (4, p. 8), supporting their suggestion of thermally

activated deposit formation mechanism.

Recently another study has been done by Ambrose et al.

(61) concerning chemical analysis of calcareous deposits.

They used rotating disc electrode to simulate the variation

of seawater flow rate by rotating an electrode screwed into

a spindle. In their "galvanodynamic" experiments in which

current was controlled to maintain the electrode potential

within the range between -0.85 and -0.90V vs. Ag/AgC1

electrode in every two minute period, initial current

density of -13 mA/cm2 was applied for a few seconds. As








they pointed out, in galvanostatic experiments in which

constant current was applied throughout experiments (most of

the studies previously discussed have been done in this

mode), severe corrosion occurred with iron compounds formed

around small pits. Under these conditions, the contribution

by deposits to decrease in cathodic current is obscured by

the effect of rust because rust itself can function as a

surface film.

Artificial seawater according to ASTM D-1141 (ASTM),

natural seawater taken from St. Augustine Beach (SAB) and

natural seawater taken from Crescent Beach (CB) were used.

The "protective" current density to maintain the

"protective" potential between -0.85 and -0.90V varied among

the different solutions. At the electrode rotation speeds

of 100, 250, 500, 1000 and 2500 rpm, the "protective"

current densities after one hour, at which current density

vs. time curve leveled off, were 50, 300, 480, 550 and 700

uA/cm2 whereas 40, 120, 160, 250 and 360 PA/cm2 were

necessary to maintain the protective potential in SAB

seawater. Interestingly, in CB seawater, the current

densities were similar to those observed in ASTM seawater.

Though the authors did not give explanation to this

phenomenon, it is apparently due to the variation of

seawater from place to place. Studies have found that

seawater varies in temperature, dissolved matter, pH, etc.

with geographical location (2,62-63). Also it is

conceivable that small variation of micro-organisms can








influence the formation of calcareous deposits from their

photosynthesis and respiration, changing the distribution of

carbonate system among its elements (19,33).

Morphology studies showed that globular particles have

grown, effectively covering electrode surface. The

difference in morphology between zero and 250 rpm group and

1000 and 2500 rpm group was that the globular particles in

the low rpm group were smaller (diameter 0.5 um) than

those in high rpm group (diameter = 1 um) under SEM.

Electron microprobe (EMP) analysis showed even

distribution of the elements of Ca, Mg, Si and S on

electrode surface. Energy dispersive X-ray analysis results

showed a difference between the deposits formed under static

condition and those formed in flowing solution condition.

In the former, distinct Mg and Ca peaks were observed

whereas in the latter Ca peak was not shown, indicating low

Ca deposits formed under flowing solution conditions.

However, in the energy dispersion X-ray spectra obtained

from the deposits formed under flowing solution condition,

very large Fe peaks (Ka and K ) could be seen. This

indicates that the thicknesses of the deposits were so small

that strong signals were obtained from the iron substrate.

As a result, a small peak from Ca, if there was any, could

be obscured by the continuous spectra of Fe. Information

about the difference in chemical composition between flowing

condition and static condition could not be drawn. The

attempt by Ambrose et al. to relate Ca-to-Mg'ratio to the








conditions under which the deposits were formed did not show

any consistent trend. Electron microprobe analysis was

employed for the deposits and atomic absorption (AA) was

used for the solution into which the deposits were

dissolved. In EMP analysis, the ratio varied from 0.50 to

1.97 (natural seawater) and in AA analyses, the ratio varied

from 0.16 to 23.0. No consistent trend could be found from

these scattered data.

Their Auger depth profiling showed chemical variation

across the deposit thickness. By using a 2KeV argon beam

ion sputtering technique, Ca-to-Mg ratio (actually Mg-to-Ca

ratio in their paper) was plotted against the time of

sputter etching. The ratio increased up to 24 minute period

and decreased again. Assuming the sputtering rate of

several hundred angstroms per minute, it is hard to think

that the depth profile represents the variation of Ca-to-Mg

ratio across the total thickness of the deposit. Instead,

it is very likely that the depth profile represents the

chemical composition across a few globular particles.

Back in 1958, Klas (9) studied calcareous deposits

formed in artificial seawater which had 0.1 g/l of CO -.

Natural seawater has about 0.15 g/l of HCO3 and about 0.1

mg/l of CO0-. In the high [CO -] solution, they obtained

deposits containing more than 80 weight % of CaCO3. They

also found that both aragonite and calcite were formed as

well as brucite (Mg(OH)2). Small well-defined particles

were deposited at -0.780V vs. SCE and, as the cathodic








potential increased (more negative) up to -1.080V, the

particles became larger and more blurry. This behaviour was

attributed to the increase in Mg(OH)2 concentration in the

deposit as potential became more negative.

Culberson (64) studied the effect of several parameters

on cathodic current requirement and deposit formation. By

comparing the cathodic current density in seawater and that

in Mg++ free artificial seawater, he found that the presence

of Mg++ markedly increased current requirements. This was

attributed to the inhibitory effect of Mg++ on CaCO3

precipitation. By controlling the partial pressure of N2,

02 and CO2 gas mixture bubbled into solution, he found that

an increase in dissolved 02 amount up to the air saturation

level increased the cathodic current requirement and then

further increase in 02 amount did not increase current. pH

was varied between 8.2 and 8.65 by controlling partial

pressure of CO2 of the gas mixture. Current was found to

decrease as pH increases. At high pH, carbonate system has

more CO- concentration increasing the degree of CaCO3

supersaturation. Investigation of current requirements at

various CO concentrations showed a linear relationship and

verified that the rate of CaCO3 precipitation is controlled

by its degree of supersaturation.

Research efforts have also been done in the desalina-

tion area (10,65-75). However, most of the work was done at

high temperatures at which the desalination process is

carried out. The deposits were very rich in CaCO3 and







differently hydrated CaSO4 compounds because these species

have reduced solubilities at high temperatures as shown in

Table 4 (76,77). Nair and Misra (10,65) studied deposits

formed at 300C in seawater with 3 mA/cm2 current density.

After 6 hours, deposits were analyzed. By assigning

magnesium to Mg(OH)2 and Ca to CaCO3 and CaSO4, 76.7% of

Mg(OH)2, 19.6% of CaCO3 and 2.6% of CaSO4 could be

obtained. X-ray diffraction pattern showed mostly calcite

and brucite structures. Scale weight increased with time

with diminishing rate because the concentrations of the

deposit forming elements were gradually depleted from

solution.


























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CHAPTER 3
EXPERIMENTAL

3.1 Materials

These studies were performed using pure iron made by a

vacuum remelting process as electrode material. Its

composition is shown in Appendix 3. Three specimens with

different geometry were used. For electrochemical studies,

specimen A and specimen B were used. Specimen A has a

diameter of 4.3 mm. The stock iron rod was machined to size

and press-jointed to a brass base which has a thread. The

assembly was mounted into a nylon tubing. The gap of about

2 mm between the center piece and the nylon tubing was

sealed with epoxy resin in vacuum. Specimen A was seriously

damaged during the course of this research because of the

accidental disconnection of a salt bridge. The

discontinuity of the circuitry of the electrochemical cell

system caused a very large cathodic current to flow.

Blistering was observed along the edge, apparently due to

the hydrogen penetrated into the iron piece. Another

specimen was made for further electrochemical measure-

ments. Specimen B has the same geometry as specimen A

except that the diameter of the center piece was slightly

increased to 5.0 mm. This was done to reduce the effect of

gap opening between the center piece and the surrounding








epoxy layer. Except for the factorial design experiments,

specimen B was used throughout. Specimen C with different

geometry was used for X-ray photoelectron spectroscopy

analysis (XPS or ESCA) and infrared analysis. Because of

the high vacuum requirement of XPS system, specimen B, which

was permanently mounted in such outgassing polymeric

materials as nylon and epoxy, could not be directly put into

XPS vacuum. Therefore, a removable type of specimen became

necessary. An iron disc was made from the same pure iron

stock into 12 mm in diameter and 5 mm in thickness. This

disc was attached to a brass adaptor by applying a small

amount of polyester resin on the side wall of both the

specimen and the adaptor. After experiment, the thin

polyester layer was peeled off and the disc was separated

from the adaptor. This specimen C not only met the high

vacuum requirement but also was larger than the size of the

ESCA X-ray beam. Also, its large surface area was suitable

for infrared analysis.

The solutions used in this study were prepared

according to ASTM D-1141-75 Substitute Ocean Water using ACS

reagent grade chemicals. Ionic composition is shown in

Table 3. Distilled and deionized water which had a conduc-

tivity of 4x10-7 ohm-lcm-1 was used. After adding

chemicals, 4.8x10-2 ohm-lcm-1 was measured. Before mixing,

water was stirred with a magnetic stirrer for ten minutes to

ensure air saturation. The procedure of mixing was modified

from the ASTM standard. Instead of adding NaHCO3 into the







No. 2 stock solution, it was weighed and kept in a plastic

container with the appropriate amount of NaCI and Na2SO4.

0.1 N NaOH was slowly added with stirring to avoid possible

precipitation. For deposit-free experiments, MgC12.6H20,

CaC12 and SrCl2.6H20, which are the source chemicals for

Mg++, Ca++ and Sr+, were substituted with an additional 6

grams of NaC1 and the same conductivity of 4.8x10-2

ohm1 cm- could be maintained.

3.2 Surface Preparation

Specimens were ground with SiC paper down to 600 grit

and polished with 6 um diamond paste. Because alcohol and

acetone damage the epoxy layer of the specimen assembly,

electrodes were ultrasonically cleaned in soap solution

between each polishing step. Though such fine surface

preparation is not employed in industry, it was necessary

for surface area reproducibility and surface sensitive

analytical work. Specimens for electrochemical measurements

were frequently ground with 320 grit SiC paper to eliminate

any possible gap opening between the epoxy layer and the

iron piece.

3.3 Equipment

Electrochemical measurements were made using a rotating

disc electrode system. The electrode was attached to a

spindle which was designed to fit a Pine Instruments ASR-2

rotator unit. A Princeton Applied Research (PAR) potentio-

stat and a PAR programmer were used for electrochemical

measurements.







The electrochemical cell has a 1 liter capacity with an

outer water jacket connected to a temperature controller.

Solutions were made with distilled deionized water, the

temperature of which was adjusted to experiment temperature

prior to mixing.

Scanning electron microscopy was carried out with JEOL

JSM 35C and JSM 35CF microscopes. X-ray analysis was done

with an Ortec EEDS II attached to JSM 35C microscope. A

Kratos XSAM 800 with 1253.6 eV mg Ka source was used for X-

ray photoelectron spectroscopy. Infrared analysis was

carried out using a Nicolet Fourier transform infrared

(FTIR) MX-1 with a 1200S data station in diffuse reflection

mode in nitrogen gas environment. The stage was manu-

factured by Barnes Analytical Co.

3.4 Current Measurements

A 24 factorial design technique was employed (Appendix

2) to study the influence of parameters on cathodic current

and deposit formation. The four parameters selected for

these studies are cathodic potential (E), electrode rotation

speed (RPM), solution pH (pH) and solution temperature

(T). Each parameter was given two levels which are shown in

Table 5. Therefore, a total of 16 different conditions were

investigated. Potential had two levels of -0.8 and -1.OV vs

Ag/AgCl. At -0.8V, oxygen reduction reaction takes place

and, at -1.OV, both oxygen reduction and hydrogen evolution

occur. For the two levels of rotation speed, 500 and 1000

rpm were studied. Solution pH had two values of 8.3 and 7.9








which represent high and low seawater pH values, respec-

tively. The two temperatures used were 23 and 160C, high

and average temperature of natural seawater, respectively.

Each of the two levels will be abbreviated to level 0 (zero)

and level 1 as defined in Table 5. Each experiment was

duplicated. Experimental period was 2 1/2 hours. For other

experiments, the period was appropriately selected. Current

was measured with Keithley 602 Electrometer. The output

terminal which provides 1V full-scale deflection voltage

output was connected to an Easterline Angus stripchart

recorder.


Table 5 Parameters and their levels


Level E (V) RPM (rpm) pH T (oC)

0 -0.8 500 8.3 25

1 -1.0 1000 7.9 16


After cathodic polarization is over, specimens were

removed from the cell and dried under N2 stream and kept in

a desiccator for further studies. An Ag/AgCl reference

electrode was used throughout these experiments. To measure

current at a different potential or at a different rotation

speed, thumbwheels were slowly turned to the desired setting

and a few minutes were allowed before reading until the

current became stable.








3.5 Morphology and Chemical Analysis

For SEM studies, deposits were sputter coated with an

Au-Pd alloy. This coating was employed to provide an

electron flow path and to enhance secondary ion yield

(78). To avoid strong Fe signals (characteristic and

continuous) from the substrate iron, deposits were peeled

off using acetate replicating tape and acetone and coated

with carbon for X-ray analysis. Because of incomplete

removal of deposits, quantitative analysis was not

attempted. The deposits which were to be subsequently

analyzed with EDX were not coated. Instead, 5 KeV electron

beam was used to minimize charge build-up at the expense of

imaging.

For ESCA work, specimen C was polarized under the same

conditions as electrochemical experiments. Specimen was

usually stored in desiccator for a few days before

analysis. Resolution was controlled by properly adjusting

the opening of the hemispherical spectrum analyzer.

Inductively coupled plasma (ICP) was employed for

solution analysis to detect any change in concentration

after experiment. However, this study was discontinued once

its incapability to detect very small change was realized.













CHAPTER 4
RESULTS


Cathodic currents, measured after the deposit was

formed on the electrode surface for 21/2 hours, are listed in

Table 6 under the heading of "with deposit." From the

experiments performed in the solution which does not contain

deposit forming elements (Ca++, Mg++ and Sr++), cathodic

currents were measured and also listed in Table 6 under the

heading of "without deposit." From these measurements, four

main effects and eleven interaction effects were calculated

for both "with deposit" experiments and "without deposit"

experiments (Table 7). The methods to calculate the effects

and the standard errors are shown in Appendix 2. Zeroes (0)

and ones (1) in Table 6 represent the values preassigned to

each parameter (Table 5).

Taking the measured current "without deposit" as the

reference, the effectiveness of surface coverage by

calcareous deposit (effective surface coverage) was

calculated by subtracting the current with deposit from the

reference current and dividing the result by the reference

current. For example, in case of (1110), reference current

was 146 PA and the current "with deposit" was 125 vA.

Therefore, the effective surface coverage for (1110)

experimental condition is (146-125)/146 = 14%. Sixteen

















LC

MC- (M T


CL- 0 l l -
CO *oo
cOo o-


IC'


C00
T-


00 o-C CO M C cON O-Cj
T 1- ^- -- ^ ^


00O


M Ln -- CO 10 Cm 00 1O -0 -0 0 0 M LO-u ,--
N[n- 000 --1- O00N-- L-"O 0--O
N-Nr -O^- -OON -O


n *n n
. . LC
LC r 7-
,o M(3)


Ln
t'-O -


Ln LnLn

o G \n
L>- 0 C


mOl~n >- o0 cocm
00


D3 Ln T- MD
KO-) 0 %,0


O K'-0 n I OL Ot0
Lt- 0 M L- 0O 0 -









10- O 0Ol- OLnc Nn
Ln\0 lCO M O mr 0 N
L

0000 0000 -r-r- r-T


0000 r-r- 0000 -V-r


00 000 00 000


0T-0 OtOT 00O 0 T- 0:4


0 *- *


Wh LC nP

n 110nc-


L n ";I m
I'D (7) m CMj










Table 7 Main and interaction effects


Effect

E

RPM

pH

T

E*RPM

E*pH

E*T

RPM*pH

RPM*T

pH*T

E*RPM*pH

E*RPM*T

E*pH*T

RPM*pH*T,

E*RPM*pH*T


With deposit (vA) Without density (PA)

21.2 1.3 21.9 1.4

24.3 32.9

30.1 -0.1

9.6 -5.4

-1.9 -2.3

5.1 -0.1

-5.2 -3.4

4.7 0.2

2.2 -0.4

-7.6 0.6

1.4 2.2

0.7 -0.4

-1.3 1.8

0.1 -1.4

-1.9 -2.7


Note: 1.3 and 1.4 are standard errors. Calculation
procedure is given in Appendix 2.








effective surface coverages obtained from 16 different

experimental conditions were divided into two groups: Group

Oxy (OXXX) and Group Hydr (1XXX). The letter "X" in the

brackets indicates that the level of the corresponding

parameter was either 0 or 1. Each group was divided into

two subgroups in three different ways: RPM-O-Subgroup

(XOXX) and RPM-1-Subgroup (X1XX), pH-0-Subgroup (XXOX) and

pH-1-Subgroup (XX1X), and T-O-Subgroup (XXXO) and T-1-

Subgroup (XXX1). Therefore, the subgroup of RPM-O-Subgroup

in Group Oxy covered four different experimental conditions

of (0000), (0010), (0001), and (0011). Likewise, each of

the other subgroups represented four different experimental

conditions. The grouping was done in a similar way. The

result is shown in Table 8.

In Table 9, the currents were measured after one

experimental condition was followed by another. The

schedule for Run 3 will be explained as an example. Ini-

tially, deposit was formed at 1000 rpm (0100). After two

hours, rotation speed was reduced to 500 rpm (0000) without

interruption and another two hour cathodic protection was

maintained. Current was measured at the end of the experi-

ment. It should be noted that from this set of experiments

on, Specimen B was used because Specimen A was damaged

during this study.

In Figure 1, the deposit was formed under one set of

conditions. Intermittently, the rotation speed was changed

either from 500 to 1000 rpm or from 1000 to 500 rpm,











Table 8 Effective surface coverage


Group Oxy Group Hydr
Uroup (Oxxx) (1XXX)

RPM-O-Subgroup (XOXX) 31 % 26 %
RPM-1-Subgroup (X1XX) 30 26

pH-0-Subgroup (XXOX) 42 40
pH-1-Subgroup (XX1X) 19 12

T-0-Subgroup (XXXO) 38 30
T-1-Subgroup (XXX1) 23 22

Average 31 26


Note: E.S.C. = without deposit with deposit
without deposit












Table 9 Combination of different experimental conditions


Initial (2 hrs)

0100
0000
0100
0000

1000
0000
1000
0000


Followed by (2 hrs)


0100
0100
0000
0000

1000
1000
0000
0000


For the convenience of comparison, each notation can be
interpreted in the following way omitting common parameters:


Run 1 Run 4



Run 5 ~ Run 8


0000 : 500 rpm
0100 : 1000 rpm


0000 : -0.8 V
1000 : -1.0 V


No.

Run 1
Run 2
Run 3
Run 4

Run 5
Run 6
Run 7
Run 8


Current

52 uA
46
35
35


























z

O
(Li
0



O
I-




r- I




w
0












0-

0


- 0


oosl / ooo000


r, cN


I I







depending upon the initial rotation speed. The ratio of the

current measured at 1000 rpm (i11000) to the current measured

at 500 rpm (i500) was plotted against time. Each of the

five curves of Figure 1 represents (0000), (1000), (0100),

(0010) and (0001). Except for (1000) and (0100), rotation

speed was increased from 500 to 1000 rpm. For (0100),

rotation speed was decreased from 1000 to 500 rpm. For

(1000), after the cathodic potential was decreased from -1.0

to -0.8V, current was measured for i500. Then rotation

speed was increased to 1000 rpm for i1000 measurement. For

all of the experiments, rotation speed and/or cathodic

potential were set back to original values after i1000

measurement. In case of rotation speed change, a stable

current was reached almost instantaneously and current went

back to the previous value when the rotation was set back to

the original speed. In the case of potential change, the

current did not stabilize even after several minutes and the

probable error caused by using the unstable current could

obscure the exact evaluation of the ratio and the exact

nature of the deposit formation at (1000). It has to be

noted that i1000 and i500 were measured at -0.8V for all of

these experiments. Experiments were discontinued when

i100/i500 reached 1.050.01. Throughout the (0010)
experiment, 0.1N H2SO4 was added at the rate of 1 drop for

every hour to maintain pH 7.9. Because no automatic pH

adjusting equipment was available, the (0010) experiment








could not be extended beyond thirteen hours so that the

ratio could reach the same value as the others.

In Figure 2, the current vs. time relationship is

shown. Deposit was formed at (1000) or -1.0V. Inter-

mittently, the potential was changed to -0.8V. The current

measured after the potential change (i_0.8) was subtracted

from the current measured before the current change (i_1.0)

and this difference was plotted against time on the right

vertical scale. It has to be noted that positive values are

always used for cathodic currents throughout this thesis.

Therefore, i_1.0 is larger than i-0.8 and the difference is

always positive. The difference was divided by i_1.0 and

multiplied by 100 and is shown on the left vertical scale.

Figure 3 is similar to Figure 2. But the deposition

condition was (0001), or 16C. Intermittently, the

potential was changed from -0.8V of (0001) to -1.OV. The

definitions for the currents, i-1.0 and i_0.8, are the same

as in Figure 2. To avoid a large transient current, the

potential was changed by slowing turning the potential

control thumbwheel at the rate of 0.1 V/min, approximately.

Figures 4 through 8 are the scanning electron

microscope (SEM) pictures, each of which was taken after the

experiment for Figure 1 was finished. Therefore, the SEM

pictures were taken after different periods of time.

Figure 9 is the electrode surface after deposit was

removed with a Kimwipe paper. The specimen was stored in a

dessicator for five days.


















Vri

0o D
reo o.


8'0-


I -


00


0 <
0 []


0 0 io 0
1- r) r) c\j J

o'.- I
cooco





O'-












% 80-
8 at 16 C (
o
S60


40-

020
20 I -

0 2 4 6

Hour


0001)


8 10


Variation of i_1.0 i_0.8 With Time


Figure 3






























Figure 4


Figure 5


SEM Micrograph of Deposit Formed at (0100);
-0.8V, 1000 rpm, pH=8.3, 230C


SEM Micrograph of Deposit Formed at (0000);
-0.8V, 500 rpm, pH=8.3, 230C






























Figure 6


Figure 7


SEM Micrograph of Deposit Formed at (1000);
-1.OV, 500 rpm, pH=8.3, 230C


SEM Micrograph of Deposit Formed at (0001);
-0.8V, 500 rpm, pH=8.3, 16C












































Figure 8


SEM Micrograph of Deposit Formed at (0010);
-0.8V, 500 rpm, pH=7.9, 230C


C w ? t
^ ^ " ----- ~s ,
<* * % S


'p
V


*a1
Q "4;.


0 ~


0

C


u *E *4
*4
0g
g



4,;
- *

(t


-A 1 ".'46


Figure 9


Electrode Surface After Deposit Was Removed With
a Kimwipe Paper. -Specimen Was Stored in a
Desiccator for Five Days


t. I


5, .O


.








Energy dispersive X-ray (EDX) spectra are shown in

Figures 10 through 12. From the globular particles as can

be seen in the SEM pictures (Figures 8 through 12), Fig'ures

10 and 11 were obtained and, from the elongated-particle

background, Figure 12 was obtained.

Figures 13 through 17 are the ESCA spectra obtained

from the deposit formed at (0000). Figure 13 is the survey

scan and the others are single element scans for Mg, Fe, and

Ca.

Fourier transform infrared (FTIR) spectra were studied

from the deposits formed at (0000) for 2 hours and 6 hours

in order to obtain the chemical information for the anionic

species in the deposits (Figure 18).














































n>
m Q)


o 0
o 0
O0
o OD
















01 O
C\i
C\J -4

(0
d






-
E-,



0



0
00





(0
C-i






a
to
tO






a



*rI


oj a,
CL]


N c00
























































































0 0
0 0
0 0
0oo
,i,-















RUN:-DVT6 / 20-APR-83


8000

I




6000


REG: 1.001 ST:0.30 #C:4000 +SW: 5 DWELL: 0.100
BACKGROUND : 108


0(A)


K2s
"I'g2p


B.E.


400 200 0


Figure 13 ESCA Survey Scan of Deposit Formed at (0000)


1000





61








RUN:-CS21 / 16-NOV-83 REG: 2.001 ST:0.18 IC: 480 #SW: 50 DWELL: 0.200
BACKGROUND : 13857





20000
I


















8 | I _\


B.E.


Figure 14 ESCA Mg Scan of Deposit Formed at (0000)















16-NOV-83 REG: 3.001


ST:0.10 tC: 400 #SW: 50 DWELL: 0.200
BACKGROUND : 27566


0
730


720 710 700
B.E.


ESCA Fe Scan of Deposit Formed at (0000)


Figure 15













RUN:-CS21 / 16-NOV-83 REG: 1.001 ST:0.10 IC: 400 tSW: 50 DWELL: 8.200
BACKGROUND : 21344
3000

:I 1 A


360 350


B.E.


Figure 16


ESCA Ca Scan (Low Resolution) of Deposit Formed
at (0000)


2000














RUN:-BH20 / 25-JUN-84


1500

I


REG: 2.001 ST:0.10 #C: 400 +SW:100 DWELL: 0.200
BACKGROUND : 13448


360 350


Figure 17


ESCA Ca Scan (Medium Resolution) of Deposit
Formed at (0000)


B.E.





65







.-I






r--i
aoNvaosav I o
oZL- 960' CLO' 6eo' seo'



CC\
Q)

o









0 C
O








0 0
0
4- o



q c
-0 o
C\J







O)

C0- 0









---- -o
0





Sl- 0
0 C
L.0 O 0 "





3 Do3
z C u



9lV2 8 'O2 0L5
3DN r8WiO98lt













CHAPTER 5
DISCUSSION

5.1 Discussion of Background

5.1.1 Rotating Disc Electrode

When an electrode of a flat disk shape (Figure 19)

rotates, the solution near the electrode follows the motion

of the electrode because of the viscous nature of the

liquid. As the distance from the electrode increases, the

movement of the solution starts to deviate from that of the

electrode. The three dimensional movement of the solution

is shown in Figure 19. Because the perpendicular motion

decreases as the distance from electrode surface decreases,

at certain point, the perpendicular motion becomes almost

zero and convection does not contribute to mass transport

any more. Therefore, mass transport is solely driven by

concentration gradient, i.e., diffusion. The term "diffu-

sion layer" is defined as the solution layer from zero to

this distance from electrode surface. The presence of a

diffusion layer is not limited to a rotating disc electrode

system. Every electrochemical system inherently containing

a solid-solution interface can have this phenomenon

depending upon reaction kinetics. Therefore, by a employing

rotating disc electrode system, mass transport across the

diffusion layer can be successfully controlled.















































Figure 19 Three-Dimensional Flow of Liquid Near Rotating
Disc Electrode








5.1.2 Effects of Parameters

The factorial design technique is well established and

detailed information can be found in experimental statistics

books (79,80). The magnitude of an effect can be

interpreted in this way: when a parameter is changed from

level 0 to level 1, average current increases by that

amount. Therefore, a positive effect indicates increase in

current and a negative effect indicates decrease in

current. Detailed information on the procedure of

calculation can be found in Appendix 2.

In the case of "with deposit" (Table 7), temperature

has the smallest effect on current whereas pH has the

largest effect. But these effects have to be carefully

interpreted. A parameter has two different natures of in-

fluence; one on the rate of electrochemical cathodic reac-

tions taking place on electrode surface and the other on the

formation of deposit which effectively retards the passage

of current through the deposit. The variation of a para-

meter will inevitably have these two aspects. Therefore,

the effects reflect both, cathodic reaction and deposit

formation. In order to study the influence on deposit

formation only, the influence on the cathodic reaction has

to be subtracted from the total influence. For this

purpose, a solution which did not contain deposit forming

elements (deposit-free solution) such as Ca++, Mg++ and Sr++

was used. Because deposits did not form on metal surfaces,

the variation of a parameter had an influence on the







cathodic reaction only. Therefore, the figures under the

heading of "without deposit" in Table 7 represent the

influence (or effect) of the four parameters on the rate of

electrochemical reactions only. Unfortunately, the right

column ("without deposit") can not be directly subtracted

from the left column ("with deposit"), even though the

former represents the influence on reaction rate and the

latter represents the influence on both reaction rate and

deposit formation. Because the original electrode surface

area did not change in the "without deposit" experiments

whereas the area was reduced in the "with deposit"

experiments, direct subtraction should not be used unless

the "without deposit" effects (or currents) are compensated

for the change in the surface area. However, a "without

deposit" effect can be compared to the corresponding "with

deposit" effect in order to evaluate the influence of the

parameter on deposit formation.

5.1.3 Ratio of Currents at Two Different Rotation Speeds

When an electrochemical reaction is diffusion

controlled, the diffusion limited current i can be expressed

as (81,82)


DC
i = m C [1]

where m= constant

D= diffusion coefficient

C= bulk concentration of the rate limiting species

6= diffusion layer thickness.










Also, diffusion layer thickness 5 can be expressed as

(81,82)



1/3 1/6 -1/2 [
6 = s D v 0 [2]



where s= constant

v= kinematic viscosity

,= rotation speed.

Substituting [2] into [1]



2/3 1/2
i m D C [3]
s 1/6
v



It can be seen from [3] that a diffusion limited current is

proportional to the square root of the electrode rotation

speed (w). The linear relationship between cathodic current

and w/2 is shown in Figure 20, based upon which it can be

said that the electrochemical reaction at -0.8V is diffusion

controlled. The ASTM standard artificial seawater does not

have any other reducible species at the potential used but

dissolved oxygen. The presence of the plateau in the

polarization curves both in the artificial seawater and in

the deposit-free solution confirms that oxygen reduction




71
















()

0 0 0
o r-

o 0
H .-I-


(.0 ,
0
o
0 a
c VO E O
o V .. ILt a.
4-- -I
C3 0 E





Q) OD


<- O
ro
04- Co

U- \ 0




0 O O O oU
0 O O
\ 0

4-)


0 0 0 0 0

o o0 0 0 cj

(Vri) uaiijjn o!pol Wo








reaction is diffusion controlled at -0.8V and the more

negative potentials (Figure 21), even with deposits formed

on the electrode surface.

The dependence of oxygen reduction current on electrode

rotation speed is, as shown in [2], due to the fact that

increase of rotation speed results in the decrease in

diffusion layer thickness which, in turn, increases the

arrival rate of dissolved oxygen. According to [3], the

current should increase by the factor of /2 when the

rotation speed is raised from 500 rpm to 1000 rpm.

Because it is inconceivable that electrons leave

metallic electrode and become hydrated in solution and

because it is equally inconceivable that calcareous

deposits, which can be characterized as ionic compounds,

conduct electrons, all electrochemical reactions must take

place on electrode surface. Though the deposits can not

conduct electrons, they can conduct electricity through the

water which the deposits contain; trapped or adsorbed.

Insofar as the deposits contain water, oxygen molecules can

travel through the water in the deposit leading to any

oxygen current; hydrogen current also becomes a factor at

more negative applied potentials. As the deposit increases

its physical coverage of electrode surface, the current path

from the bare metal to the bulk solution loses its role as

the major conduction path and the current path from the

metal surface through the solution contained in the pores

and adsorbed deposit particles becomes the more important




73








*H
0



-P *H

.0 0
(1) 0
(CO
iuej) jua no o!poquC) (1 o
4-) > Co
O O O O i 8
0 0 0 -o u
(0 0 c co

o > o


E 3 E >0
O c co
^^ 3 E-1 0
\0 E
> ,,, > J4-)
SQ.O0
o ~ E- 4-P 0
0 0o c-- O
CO Oa Ocat
S\\ -C
0' \ .., 4 U-
0 4-) -P





.C '
.a +- Cr O N -I





Q DI \ N -i
-o % I
O rOCo
I (I I) 0



I 0 0

0 E

Pt S, 4-)





cZI
w 0
4 *-O C 0
CU4 cu
0 O



*-0


4u0
I








conduction path. Also, as deposits grow in thickness, mass

transport across the deposits becomes more difficult and

mass transport across diffusion layer, whose thickness can

be calculated from the equation [2], gradually loses its

effect on total mass transport. Therefore, as deposit

grows, increase in rotation speed, even though it decreases

the original diffusion layer thickness proportional to the

inverse one-half power of rotation speed, does not

accelerate oxygen diffusion through the deposit as much as

can be predicted from [2] and the amount of current increase

will become smaller and smaller. As a result, the V2

relationship between the current at 1000 rpm (i1000) and the

current at 500 rpm (i500) predicted from [3] will fail and

the ratio, ii1000/500 will decrease to approach unity at

which mass transport across the deposit becomes rate

controlling step. Therefore, observation of decrease in

i1000/i500 against time will be able to show the development
process of deposit on the electrode surface (Figure 1).

5.2 Discussion of Influence of
Parameters on Deposit Formation

5.2.1 pH

The effects of pH of Table 7 are shown schematically in

Figure 22, "without deposit" and "with deposit." "Without

deposit," the current at level 1 was almost the same as that

at level 2. In other words, variation of bulk pH had little

influence on the rate of cathodic reaction which was

measured in terms of current, though [OH-] is in the











pH effect :


Level








0


-0.1 ->30.1 pA


Without deposit


With deposit


Figure 22 Schematic Presentation of the pH Effects







reaction path for both oxygen reduction reaction and

hydrogen evolution reaction.


02 + 2H20 + 4e- + 40H- [4]


2H20 + 2e- + H2t + 20H- [5]


The fact that the cathodic current is mainly comprised of

hydrogen evolution current and oxygen reduction current

allows us to interpret the measured pH effect ("without

deposit") as the sum of the two corresponding parts.

As was discussed in 5.1.3, the oxygen reduction

reaction is diffusion controlled at the potentials of this

study. Therefore, variation in pH should not have any

influence on oxygen current, though [OH~] is in the reaction

path of oxygen reduction reaction [4].

In contrast, hydrogen evolution is kinetically

controlled. Therefore, decrease in pH from 8.3 (level 0) to

7.9 (level 1) is expected to yield a slight increase in

hydrogen evolution current. The amount of hydrogen current

increase can be calculated from the relationship between

current and overpotential. For hydrogen evolution reaction,

equilibrium potential is related to pH in the following

equation:


E = EO 2.303RT pH [6]
F


where E is the equilibrium potential and the other symbols

are conventional. As pH decreases from 8.3 to 7.9, at








T = 230C,



AE = El EO = 0.023V


where the subscripts are the two levels of pH. From the

following equation between current and overpotential, for

cathodic reaction:


n = B 2.303 RT log i [7]
aF


where n is the overpotential and the other symbols are

conventional. Therefore,



2.303 RT lg [8]
1 0 aF iO


where the subscripts are the two levels of pH. From the

slope of the cathodic polarization curve (Figure 23),


2.303 RT = 0.18V [9
aF


at T = 23C. From above calculation, n1 nO = -0.023V.

Therefore, from [8], iI = 1.34 iO. Because the average

current increase as cathodic potential increases can be

regarded as an average hydrogen current,


Ai = ii i0 = 0.34 i0 7 uA


can be obtained. The effect of 22 uA, the average hydrogen

current, was used for iO. Effect is calculated from 16








Cathodic Potential (V)
-1.3 -1.2 -1.I -1.0 -0.9 -08 -0.7 -0.6 -0.5














10
0)
SWEEP RATE 0.3 V/ MIN -
DEPOSIT-FREE SOLUTION








Io_


Figure 23 Cathodic Polarization Curve








different experimental conditions among which 8 have oxygen

current only and the other 8 have both oxygen and hydrogen

currents. Therefore, pH effect, which is the average

current increase with pH variation from level 0 to level 1,

is half of the sum of the influences on oxygen current (zero

expected, theoretica-lly) and on hydrogen current (7 uA

expected, theoretically). So the measured pH effect of -0.1

,A "without deposit" can be theoretically explained by the

preceding discussion.

This almost zero pH effect changed to +30.1 uA when

deposit was introduced into the system. And ii became

larger than i0 ("with deposit" in Figure 22). If the two

deposits formed at the two different pH values are identical

in their current blocking capability, il has to be almost

the same as i0 ("with deposit" in Figure 22), which is

suggested by the negligible difference in currents "without

deposit." The fact that il became larger than iO ("with

deposit" in Figure 22) allows us to draw the conclusion that

the deposit formed at level 1 is poorer than the deposit

formed at level 0 in terms of current blocking capability.

The effect of pH decrease from 8.3 (level 0) to 7.9

(level 1) on the formation of deposit is obvious. The pH at

the immediate vicinity of electrode surface (surface pH) is

much higher than bulk pH and such a small increase or de-

crease in bulk pH may not be substantial in increasing or

decreasing surface pH. However, if surface pH does not

change with bulk pH, the OH- concentration gradient across








the diffusion layer will be greater at lower bulk pH and OH-

diffusion rate will be accelerated. To maintain same

surface pH at a lower bulk pH as at a higher bulk pH, OH-

generation rate at a lower bulk pH must be greater to keep

up the increased diffusion rate. However, the nil pH effect

on the rate of cathodic reactions) indicates that OH-

generation is not accelerated at a lower bulk pH. As a

result, surface pH decreases as bulk pH decreases. As was

discussed in Section 2.1, the deposit forming reactions are

driven by high surface pH near metal surface. The decrease

in surface pH will decrease the rates of the deposit forming

reactions. Therefore, at a lower bulk pH, rates of the

deposit forming reactions decrease and, in the same period

of time, the quantity of the deposit will be smaller leading

to a lower degree of protection, in terms of cathodic

current decrease.

The lower degree of protection of the electrode surface

by the deposit at the low bulk pH, in terms of the reduction

of cathodic current, is also indicated by the large differ-

ence between the effective surface coverages (ESC) of the

two pH subgroups (Table 8). Because ESC is the percentage

of current decrease from the base metal condition, the small

percentage at the low bulk pH (level 1) suggests that the

deposit formed at the low bulk pH is poorer than the deposit

formed at the high bulk pH. This can be explained by the

low deposition rate, as was discussed previously. It can

also be noted that the largest difference in ESC was between







the two pH subgroups, indicating that the variation of

solution pH has the most significant influence on deposit

formation under the experimental conditions employed for

this study.

This conclusion is also supported by Figure 1. Because

the decay of the ratio ii000/i500 was due to the fact tha

the deposit on the metal surface controls the transportation

of oxygen, the slow decay of the ratio at (0010), (low pH),

compared to (0000), (high pH), indicates that the deposit

formed at the low bulk pH is not as effective as the deposit

formed at the high bulk pH in its capability to control the

transportation of oxygen. This can be explained by the low

deposition rate at the low bulk pH.

5.2.2 Temperature

The T effects of Table 7 are shown schematically in

Figure 24, "without deposit" and "with deposit." "Without

deposit," the current at level 1 (160C) was smaller than the

current at level 0 (23C). At low temperatures, more oxygen

is dissolved and oxygen current should increase, as a re-

sult. However, there are compensatory effects which occur

as temperature is lowered: increase in diffusion layer

thickness and decrease in diffusion coefficient. Semi-

quantitative predictions can be made using well established

equations that apply to diffusion limited reactions in

rotating disc systems. The fact that oxygen reduction

reaction is diffusion limited was previously discussed.










T effect : -5.4 -- 9.6


Without deposit


With deposit


I~L>


Figure 24 Schematic Presentation of the T Effects


pA


Level




83


Based upon the diffusion limited oxygen current, the

negative effect of temperature will be discussed.

The dependence of diffusion layer thickness on

temperature can be expressed by (83)


6= C [10]


where p is a constant. Equation [3] is

m2/3 1/C

s 1/6

Substituting [3] to [10],

1/6
Sps T v [11]
D /3w /2


To compare the thickness at level 0 (60) to that at level 1

(61) where w is constant,


61 T 1 1/6 DO 2/3 [12]
6 To DVO

Kinematic viscosity v is related to dynamic viscosity p by


P

where P = density. Neglecting the small difference in p

between 160C and 23C,


V1i .
VO 0




84


Using the interpolated data from Table 10, V1 = 0.660 and WO

= 0.559. Therefore,


V 0.660
V0 0.559

Glasstone (57) reported that diffusion coefficient of

most substances in aqueous solutions is experimentally found

to increase about 2.5% per degree (C), because of the

increase in viscosity. In fact, the increase in viscosity

is of the same magnitude. From the relationship,


DO = 1.175 D1

can be calculated. Therefore, from [12],


1 1/6 1.175 Di 2/3 .
1 289 (0.660) 1.175 D1) 1.12
0 296 0.559 D1

From equation [1],

il_- D1 60 C1
i0 D0 61 CO

where superscripts are used same as elsewhere. Using the

oxygen solubility of 8.0 and 7.0 mg/l at 16 and 230C,

respectively, in water containing 20 gm/l of chloride (84),


ii DI 60 8.0 0.87
io 1.175 D1 1.12 60 7.0












Table 10 Relative viscosity of seawater


Temp (oC)

0

5

10

15

20

25

30


Relative viscosity*

1.056

0.891

0.772

0.675

0.599

0.533

0.481


Note: From Handbook of Marine Science (2)

* Relative viscosity = viscosity of seawater

viscosity of pure water at OOC (1.787
centipose)








Therefore, oxygen current is expected to decrease by 13%

with temperature decrease from 23 to 160C.

The direction of the change of hydrogen current will be

a negative increase because hydrogen evolution reaction is

activation controlled. Cathodic current is exponentially

related to temperature in the following equation:



i = A i exp [- F n] [13]
ex RT


where A = surface area

iex = exchange current density

a = symmetry coefficient

F = Faraday constant

R = gas constant

T = temperature

n = overpotential.

Lack of information about the dependence of iex and a on

temperature variation does not allow good approximation.

Theoretically expected temperature effect is one half of the

sum of temperature influences on oxygen current and on

hydrogen current. Because the temperature influence on

oxygen current is about 13% decrease or about -12 pA, its

influence on hydrogen current should be relatively small to

explain the measured T effect of -5.4 pA.

This negative T effect was changed to a positive T

effect when deposit was introduced into the system. And i1

*became larger than i0 ("with deposit" of Figure 24). If the








two deposits formed at the two different temperatures are

identical in their current reducing capability, ii has to be

smaller than i0 ("with deposit" of Figure 24), which is

suggested by the relationship between currents "without

deposit." The fact that il became larger than i0 ("with

deposit" of Figure 24) allows us to draw the conclusion that

the deposit formed at level 1 is poorer than the deposit

formed at level 0 in terms of current blocking capability.

In other words, the lower temperature (160C) caused the

deposit to become less efficient in obstructing current flow

through the deposit than at the higher temperature (23C).

The effect of temperature on formation of calcareous

deposits is believed to be bifold. As temperature de-

creases, the solubility products of both Mg(OH)2 and CaCO3

increase (Table 4) and the deposition rate of these com-

pounds is expected to decrease. Also, a decrease in

temperature may cause a decrease in the rate of the

deposition reaction. In general, a chemical reaction can be

expressed in an Arrhenius type relationship between reaction

rate and temperature like the following:


Rate = A exp (- ) [13]
RT


where A = rate constant

Q = activation energy

R = gas constant

T = temperature.







Therefore, for the deposition of calcareous scales, an

exponential decrease in the rate with temperature decrease

may be substantial in the process of calcareous deposit

formation.

The low deposition rate at low temperatures caused by

the increase in the solubility products of both Mg(OH)2 and

CaC3, and also possibly caused by the thermal activation

mechanism yields a less protective deposit. This can also

be shown in Table 8. Because effective surface coverage

(ESC) is the percentage of current decrease from the bare

metal condition, the smaller percentage at the low

temperature (level 1) suggests that the deposit formed at

the low temperature is poorer than the deposit formed at the

high temperature in its current reducing capability. This

can be explained by the low deposition rate at the low

temperature, as was discussed previously. It can also be

noted that the difference in ESC between the two temperature

subgroups is the second largest, indicating that the

variaton of temperature has the second most significant

influence on deposit formation under the experimental

conditions employed for this study. This can also be drawn

from Figure 1. Because the decay of the ratio, i1000/i500

was caused by the development of the calcareous deposit

which functions as a barrier to the transportation of

oxygen, the slow decay shown by the low temperature curve,

(0001), indicates that the deposit is not as effective as

the deposit at the high temperature (0000), with respect to








capability to slow down the oxygen diffusion rate through

the deposit. The reduced deposition rate at the low

temperature results in the smaller amount of the deposit

formed on the metal surface and thus the lower degree of

protection in terms of cathodic current decrease.

5.2.3 Potential

The potential effects (E effects) of Table 7 are

schematically presented in Figure 25. Potential (E) effect

did not vary significantly from "with deposit" experiments

to "without deposit" experiments. Potential effect can be

viewed as the average increase in cathodic current through

the 16 different experimental conditions as potential is

raised from -0.8V to -1.OV. Because oxygen current is

limited by the diffusion of oxygen molecules onto electrode

surface, increase in cathodic potential has little influence

on it. In Figure 4 is shown a typical polarization curve

with deposit formed on electrode surface. Oxygen current

does not increase in the plateau region. Therefore, the

increase in current (E effect) is believed to be due to

hydrogen current. E effect of 21.9 pA "without deposit"

does not significantly decrease as deposit is formed on

electrode surface as shown by 21.2 pA "with deposit." From

this, the following can be suggested:

Though formation of deposit effectively decreases

oxygen current by its nature of transportation

barrier, hydrogen evolution reaction does not need

transport of reacting species and as far as the




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