part one. a study of the dissolution of tin tubes in

University of Windsor University of Windsor

Scholarship at UWindsor Scholarship at UWindsor

Electronic Theses and Dissertations Theses, Dissertations, and Major Papers

1-1-1971

Part One. A study of the dissolution of tin tubes in hydrochloric Part One. A study of the dissolution of tin tubes in hydrochloric

acid solutions. Part Two. A potentiostatic investigation of the acid solutions. Part Two. A potentiostatic investigation of the

polarization and cathodic protection of tin tubes in hydrochloric polarization and cathodic protection of tin tubes in hydrochloric

acid solutions. acid solutions.

David S. P. Poa University of Windsor

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Recommended Citation Recommended Citation Poa, David S. P., "Part One. A study of the dissolution of tin tubes in hydrochloric acid solutions. Part Two. A potentiostatic investigation of the polarization and cathodic protection of tin tubes in hydrochloric acid solutions." (1971). Electronic Theses and Dissertations. 6104. https://scholar.uwindsor.ca/etd/6104

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PABT ONE. A STUDY OF THE DISSOLUTION OF TIN TUBES IN HYDROCHLORIC ACID SOLUTIONS

PART TWO. A POTENTIOSTATIC INVESTIGATION OF THEPOLARIZATION AND CATHODIC PROTECTION OF TIN TUBES IN HYDROCHLORIC ACID SOLUTIONS

A DissertationSubmitted to the Faculty of Graduate Studies Through the Department of Chemical Engineering in Partial Fulfilment

of the Requirements for the Degree of Doctor Of Philosophy at the

University of Windsor

byDavid S.P. Poa

Windsor, Ontario 1971

Reproduced with permission of the copyright owner. Further reproduction prohibited without permission.

UMI Number: DC52684

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ACKNOWLEDGMENTS

The author wishes to express his sincere gratitude to his advisor, Dr. A.W. Gnyp, for his suggestions, guidance, constant encouragement and understanding throughout the course of this work.

Special thanks are due Mr. Otto Brudy, Central Research Shop, Mr. Wolfgang Sberhart, glassblower, and Mr. George Ryan, departmental technician, for their assistance with the construction of the apparatus. The author is also indebted to Mr. William Bear, former graduate student of the Department of Engineering Materials, for his help in remoulding the tin cylinders.

The financial assistance offered by the National Research Council of Canada has been appreciated.

Ill


TABLE OF CONTENTSPage

ACKNOWLEDGMENTS i üTABLE OF CONTENTS ivLIST OF TABLES ixLIST OF FIGURES X

PART ONE. A STUDY OF TBE DISSOLUTION OF TIN TUBES INHYDROCHLORIC ACID SOLUTIONS 1

ABSTRACT 2CHAPTER

I. INTRODUCTION 4II. LITERATURE REVIEW 7

A. Kinetic Studies of Metal Dissolution ?B. General Review of Studies on Tin Dissolution 8C. Hydrodynamic and Diffusional Mass Transfer

Effects 10III. THEORY 16

A. Reaction Mechanism 16'B. Mass Transfer in Electrochemical Dissolution

Processes 231. Fundamental Equations 232. Nernst Diffusion Layer Theory 263. Application of Dimensional Analysis 274. Diffusional Mass Transfer in a

Circular Tube • 29IV. EXPERIMENTAL 34

A. Equipment 341. Design of Liquid Recycle System and

Accessories 34iv


CHAPTER Page2. Details of Test Section Construction 383. Temperature Control and Measurement 38

B, Chemicals and Materials 40C. Procedure 40

V. EXPERIMENTAL RESULTS AND DISCUSSION 43A. Method of Analysis of Experimental Data; A

Typical Example 43B. Reproducibility 46C. Effects of Different Factors on the

Dissolution Rate of Tin 461. Rate Dependence on Tin Ion Concentration 462. Rate Dependence on Fluid Velocity 463. Rate Dependence on Temperature 494. Effect of Hydrochloric Acid

Concentration 493. Effect of Oxygen Partial Pressure in

Gas Phase 376. Effect of Tube Inside Diameter 377. Effect of Tube Length 60

D. Proposed Mechanism for the Dissolution ofTin in Hydrochloric Acid Solutions 64

E. Empirical Rate Equation for Tin Dissolution 71

VI. SUMMARY AND CONCLUSIONS 74A. Comparison of Present Results with Previous

Work 74B. Comparison of Present Results with Mass

Transfer Correlations 77C. Conclusions 78


PART TWO. A POTENTIOSTATIC INVESTIGATION OF THEPOLARIZATION AND CATHODIC PROTECTION OFTIN TUBES IN HYDROCHLORIC ACID SOLUTIONS 81

ABSTRACT 82CHAPTER

I. INTRODUCTION 84II. LITERATURE REVIEW 8?

A. Electrochemical Theory of Corrosion 8?B. Polarization Studies of Tin 96C. Cathodic Protection of Metals 97D. Current Density Requirements for Cathodic

Protection 100E. Overprotection 101F. Studies on Cathodic Protection 102G. Polarization Studies Employing Potentiostatic

Techniques 103III. EXPERIMENTAL 103

A. Details of the Design of the StationarySystem 103

B. Details of the Design of the Flow Systemand the Tubular Electrolysis Cell 107

C. Materials and Chemicals 112D. Procedures for Polarization Measurements and

Cathodic Protection Tests Using a' Stationary System 112

E. Procedures for Polarization Measurements and Cathodic Protection Tests Using aFlow System 116

IV. EXPERIMENTAL RESULTS AND DISCUSSION II9A. Results of Stationary System Tests 1i 9

Vi


CHAPTER1. Polarization Tests2, Cathodic Protection Tests

B, Results Obtained from Flow System1. Polarization Tests2. Cathodic Protection Tests3» Protection Current Dependence on

Temperature4. Protection Current Dependence on

Fluid Velocity3. Protection Current Dependence on

Oxygen Concentration6. Protection Current Dependence on

HCl ConcentrationV. SUMMARY AND CONCLUSIONS

APPENDIX 1. Calibration of RotametersAPPENDIX 2. Analytical Methods and CalibrationsAPPENDIX 3. Calculations of Rate Constants for

Tin DissolutionAPPENDIX 4» An Example of the Application of the

Empirical Rate EquationAPPENDIX 3» Diffusional Mass Transfer CorrelationsAPPENDIX 6. Preparation of Saturated Calomel

ElectrodeSpecifications of the Model 4700M Research Potentiostat

8. Data of Tin Dissolution9. Data of Polarization and Cathodic

Protection Tests

APPENDIX 7.

APPENDIXAPPENDIX

REFERENCES

Page

119124126126131

141

144

144

147148

131

134

137

161164

167

168

169

206

223

Vll


PageNOMENCLATURE 228VITA AUCTORIS 232

viix


LIST OF TABLE

TABLE Page1-3-1• Effect of HCl Concentration on Tin

Dissolution 36A-3-1. Evaluation of Velocity Constant

(k° for hydrogen evolution) 138A-3-2. Evaluation of Velocity Constant

(kg for oxygen depolarization) 139A-3-3* Evaluation of Velocity Constant

(k^ for autocatalytic reaction) 160

IX


LIST OF FIGURES

FIGURE Page1-3-1• Mass Transfer in a Tube with a Soluble

Metal Wall 291-4-1. Schematic Diagram of Fluid Recycle

System 331-4-2. Detailed Drawing of Test Section 391-3-1. A Proposed Graphical Analysis for Tin

Dissolution in HCl Solutions 451-3-2. Reproducibility of Experimental Data 471-3-3» Dissolution as Function of Tin Ion

Concentration in the Bulk Solution (for autocatalytic reaction) 48

I-3-4 . Dissolution as Function of Fluid Velocity(for Hg evolution reaction) 30

I-5-5. Dissolution as Function of Fluid Velocity(for oxygen depolarization reaction) 31

1-3-6. Dissolution as Function of Fluid Velocity(for autocatalytic reaction) 32

1-3-7. Dissolution as Function of Temperature(for hydrogen evolution reaction) 33

1-3-8. Dissolution as Function of Temperature(for oxygen depolarization reaction) 34

1-5-9. Dissolution as Function of Temperature(for autocatalytic reaction) 35

1-3-10. Dissolution as Function of OxygenConcentration(for Oxygen depolarization) 58

1-3-11. Dissolution as Function of OxygenConcentration(for autocatalytic reaction) 39

1-3-12. Dissolution as Function of Tube InsideDiameter(for Hg evolution) 61

1-3-13. Dissolution as Function of Tube InsideDiameter(for oxygen depolarization) 6

X

0


FIGURE PageI-3-I4. Dissolution as Function of Tube Inside

Diameter(for autocatalytic reaction) 63I-3-I3* Dissolution as Function of Tube Length

(for Hg evolution) 651-5-16. Dissolution as Function of Tube Length

(for oxygen depolarization) 66I-5-1?. Dissolution as Function of Tube Length

(for autocatalytic reaction) 6?II-2-1. Activation Polarization Curve for a

Reversible Electrode $2II-2-2. Concentration Polarization Curve

(reduction process) 94

II-2-3. Combined Polarization Curve - Activation and Concentration Polarization(reduction process) 95

II-2-4. Cathodic Protection by Superposition ofImpressed Current on Local Action Current 99

II-2-5. Electric Circuit for PolarizationMeasurements 103

II-3-1. Schematic Diagram of Stationary TestCell Assembly 106

11-3-2. Schematic Diagram of Test Electrode 10811-3-3, Block Diagram of Model 4700M Magna

Potentiostat 109II-3-4* Schematic Diagram of Tubular Electrolysis

Cell 11111-3-5* Schematic Diagram of Counter Electrode II311-3-6. Schematic Diagram of Electrical Circuitry

of Flow System ' II4II-4-I. Cathodic and Anodic Polarization of

Stationary Tin Samples in Nitrogen Saturated HCl Solutions 120

xi


FIGURE Page11-4-2. Cathodic and Anodic Polarization of

Stationary Tin Samples in Air SaturatedHCl Solutions 121

II-4-3. Cathodic Protection Effects Related to Protection Current Requirements for Stationary Tin Samples in Air Saturated HCl Solutions 123

II-4-4. Effect of Oxygen Concentration on Cathodic and Anodic Polarization of Tin in a Flow System 12?

II-4-3* Effect of Fluid Velocity on Cathodic andAnodic Polarization of Tin in a Flow System 128

11-4-6. Effect of Oxygen Concentration on CorrosionCurrent for Tin in IN HCl Solutions 129

II-4-7. Effect of Fluid Velocity on Corrosion Current for Tin in Air Saturated IN HCl Solutions I30

II-4-8A. Cathodic Protection of Tin in High Velocity,Air Saturated HCl Solutions 132

II-4-8B. Cathodic Protection Effects Related toProtection Current Requirements for Tin inHigh Velocity, Air Saturated HCl Solutions 133

II-4-9A. Cathodic Protection of Tin in High Velocity,Oxygen Saturated HCl Solutions 134

II-4-9B. Cathodic Protection Effects Related toProtection Current Requirements for Tin inHigh Velocity, Oxygen Saturated HClSolutions 133

II-4-10A. Cathodic Protection of Tin in High Velocity,Nitrogen Saturated HCl Solutions 136

II-4-10B. Cathodic Protection Effects Related toProtection Current Requirements for Tin inHigh Velocity, Ng Saturated HCl Solutions 137

II-4-IIA. Cathodic Protection of Tin in Lov/ Velocity,Air Saturated HCl Solutions 138


FIGURE PageII-4-11B. Cathodic Protection Effects Related to

Protection Current Requirements for Tin inLow Velocity, Air Saturated HCl Solutions 139

II-4-12. Protection Current Dependence onTemperature 142

II-4-13. Protection Current Dependence onFluid Velocity 143

II-4-I4. Protection Current Dependence onOxygen Concentration 145

11-4-15. Protection Current Dependence onHCl Concentration I46

A-1-1. Calibration Curves for Rotameter No.10-1000(Tube:R-10M-23-1) 152

A-1-2. Calibration Curves for Rotameter No.12-1000(Tübe:R-12M-23-3) 153

A-2-1. A Typical Calibration Curve for AtomicAbsorption Analysis I56

A-4-1. An Example of the Application of theEmpirical Rate Equation 163

A-6-1. Saturated Calomel Electrode 16?

Xlll


PART ONE

A STUDY OF THE DISSOLUTION OF TIN TUBES IN HYDROCHLORIC ACID SOLUTIONS


ABSTRACT

Horizontally positioned tin tubes, with hydrodynamic entry and exit lengths, were incorporated into a closed loop flow circuit, which was designed for the kinetic study of the dissolution of tin in aerated and deaerated hydrochloric acid solutions circulating through the system.

The effects of total tin ion concentration in the bulk solution; acid concentration, temperature, oxygen concentration, flow velocity of the corroding solution, and the variation of the inside diameter and length of the tin tubes on the dissolution rate of tin in HCl solutions were investigated.

Because the dissolution of tin in aerated hydrochloric acid solutions appears to occur through three simultaneous processes ( hydrogen evolution, oxygen depolarization, and an autocatalytic reaction ), the experimental results obtained in this investigation have been analyzed in terms of three parallel effects. Over the range of conditions studied the dissolution rate of tin has been correlated by an empirical rate equation showing the contributions of the three possible processes.

Low temperature coefficients and significant effects of fluid velocity indicate that diffusional mass transfer


3

effects play a significant role in the overall dissolution process.


CHAPTER I INTRODUCTION

The systematic study of corrosion and corrosion prevention is quickly becoming an important engineering science. People have learned that it is more economical to prevent corrosion initially by proper material selection and protective systems rather than pay for replacements and repairs made necessary by the damages of corrosion. Consequently, the corrosion engineer has established himself as a necessary member of the industrial engineering team.

The challenge in Chemical Engineering lies in predicting the behaviour of a given material in a given environment.For many years, corrosion technology was an art rather than a science, with experience and empirical data the only tools available. Stimulated and facilitated by parallel advances in other fields of technology and science, the present trend is toward deepening the understanding of the mechanisms of corrosion and the physical laws which govern corrosion processes in the hope of improving predictions by a coherent theoretical model of physical and chemical behaviour.

In recent years, many investigators have studied the mechanism and kinetics of the dissolution of metals with rotating cylindrical systems. Empirical correlations have been established and various criteria have been postulated for different situations for prediction of dissolution rates. A

4


survey of the literature indicates that little data exist for the interpretation of the mechanism and kinetics of dissolution of metals in recycling flow systems.

In studies on the dissolution and corrosion rates ofS3rotating metal cylinders, Lui^^ found the dissolution rate of

tin to be directly proportional to the square root of the12surface area of the specimen, while Bodner observed that the

dissolution rate of titanium increased with 3/4 power of the surface area of the metal sample. There are examples of other correlations between metal dissolution rate and sample surface area.^^'^^ Interpretations of these correlations are open to question. The effects of cylinder diameter and length have never been incorporated into these correlations, although the significance of these factors should not be Ignored.

The present work was undertaken with the following aims in mind: (a) to experimentally establish correlations between physical properties of a system, geometrical and hydrodynamic conditions, and rates at which metallic tin is dissolved from the inner surface of a tube into the hydrochloric acid flowing through a recycling system; (b) to compare the rate correlations from this flow system with those obtained from rotating cylindrical systems, and examine whether they share any important characteristics.

The particular system studied, tin-hydrochloric acid was chosen for reasons of convenience and simplicity. Previously, using rotating cylindrical systems, very extensive and


6

systematic studies on the dissolution rates of tin in HCl solutions were done by Luif^ and the present a u t h o r . T h e s e , works provide readily available rate correlations v/hich facilitate the second objective. Also, high purity tin can be easily obtained and cast into any desired geometry. Hydrochloric acid is one of the simpler acids and there is precise information available concerning many of its physical properties.


CHAPTER II LITERATURE REVIEW

Over the last three decades, a large number of articles has appeared in the literature on the general problem of dissolution and corrosion rate of metals. Since the purpose of this study is to obtain a better understanding of the dissolution kinetics of tin in HCl solutions flowing through a recycling system, the following survey will review only those papers pertinent to this subject.A. Kinetic Studies of Metal Dissolution

Lu and Graydon^^studied the mechanism and kinetics of copper dissolution in aqueous ammonium hydroxide and aqueous

O Qsulfuric acid solutions. Weeks and Hills ^ investigated the initial corrosion rate of copper in HCl solutions. Their work was extended by Gnyp. Other systematic researches on the dissolution kinetics of metals including brass, by Kagetsu and Graydon^^ and Bumbulis and Graydon,^^ titanium, by .Bodner, and iron, by Taneja have been carried out. Using rotating metal cylinders, the above investigators studied the rate of metal dissolution as a function of temperature, oxygen partial pressure, rotational speed, sample surface area, corroding solution volume, and acid concentration. Over a wide range of conditions, all of them found an autocatalytic effect, with dissolution rate increasing with increasing metal ion concentration in solution.

7


8

B. General Review of Studies on Tin DissolutionAs early as 1813, the dissolution of tin in acid solutions

was examined by Berzelius?’^ According to his report, tin dissolves in deaerated hydrochloric and sulfuric acid solutions in the stannous form with the evolution of hydrogen gas.

The effect of oxygen on tin dissolution in various acidic media was studied by Whitman and R u s s e l . T h e i r results showed that, in most cases, oxygen acts as a depolarizer, and if added to acids, will increase the attack on tin. In the absence of air or other oxidizing agents, tin was very resistant to dilute acids. This is because tin has a high hydrogen overpotential, and becomes quickly polarized by hydrogen which prevents the flow of current that accompanies corrosion.

Kohman and Sanborn^^ reported that, in air-free solutions, tin is depolarized by increasing the temperature of the solution, and hydrogen evolution occurs. The effect of temperature on the corrosion of tin in acid solutions was investigated by Khitrov and Shotalova.^^ The rate of tin corrosion in 1-?M solutions of hydrochloric acid and sulfuric acid was observed to increase with increasing temperature.The increase of corrosion rate with increasing temperature is associated with a fall in hydrogen overpotential, decrease in polarization, decrease in viscosity of the solution, and destruction of protective films.

The phenomenon of localized corrosion of tin was examinedZQelectrochemically by Hoar. He studied the attack on tin by


nearly neutral solutions containing numerous different anionsand cations. Brennert^^ investigated the formation of "blackspots" on tin in sodium chloride, sulphate, and nitrate

1 5solutions. Britton and Michael studied the local corrosion of tin in chloride solutions. Their results showed that the formation of black spots was an actual building up of the oxide film originally present on the tin surface. Anodic attack at a weak spot on the oxide film covering the metal would not cause tin cations to pass into liquid, but rather cause an increase in the film thickness by deposition of oxide and hydroxide. After a certain time the accumulated acidity at these points apparently became sufficient for the formation of soluble stannous ions, which ruined the film so that breakdown occurred with the formation of black spots. The blackness of these spots, according to Britton, was probably due to the absence of reflection from the locally roughened surface.

Hagymas and Quintin^^ studied the corrosion of tin in sulfuric acid. In 0.1 to 1.0 M solutions the hydrogen overvoltage on the tin electrode did not vary with acid concentration. Kohman and Sanborn^^ also reported that there was no apparent influence of acid concentration on the corrosion of tin in air-free solutions.

Probably the most extensive and systematic investigation53on the dissolution of tin was done by Lui.^ He reported

that over a wide range of conditions the dissolution of tin in.hydrochloric acid solutions proceeds in two autocatalytic


10

stages with the rate during each stage being dependent on the square root of the tin ion concentration in the corroding solution. His work was extended by this anther.

C. Hydrodynamic and Diffusional Mass Transfer EffectsA review of the literature concerning the relative

movement between the corroding solution and metal surface in corrosion processes is given on the following pages.

In 1921 Friend^^ studied the effect of velocity on iron dissolution using a stirrer method and a weight-loss technique, He found that in natural water the corrosion rate decreased with increasing velocity and at high velocity no weight losses were observed. However, in acid solutions the corrosion rate increased with increasing velocity. The rate of dissolution of a rotating disc in dilute acid with no depolarizer wasfound to be linear with rotational speeds up to 1 ,000 r.p.m.

75Speller and Kendall'^ investigated the flow of natural water in steel pipes in 1923 and published results contradictory to those of Friend. They observed that the corrosion rate was low in the laminar flow region, increased rapidly in the transition region, and continued to increase but at a lower rate in the turbulent region.

Romeo et al.^^, in 1958, reported that for turbulent flow, in small steel pipes, the corrosion rate was found tobe proportional to the velocity raised to the power of 0.88.

7?Roethelli and Brown'" rotated steel specimens in oxygenated water. Their results showed an increase in the


11

corrosion rate at low velocities followed by a decrease at intermediate velocities. At higher velocities the corrosion rate once more increased with increase in velocity. This rate enhancement they attributed principally to erosion. The minimum at intermediate velocities was due to a change in the corrosion product, ferrous ions being oxidized to ferric.

Both V/hitman^^ and Wilson^^, in 1913, stated that velocity was important in determing the thickness of the film through which oxygen must diffuse.

Makrides and Hackerman^^, using a stirrer technique to investigate the behaviour of iron in hydrochloric acid, observed that the dissolution rate was small and independent of the stirrer speed in air-free solutions, but showed amarked dependence on stirrer speed when oxygen was present.

37Hatch and Rice^', in 1943, stressed the importance of velocity as a means of regulating the rate of supply of oxygen or inhibitor to a corroding surface. They investigated the problem by using a weight loss technique on a mild steel pipe in aerated tap water. At high values of Reynolds number the corrosion rate attained a steady value. The oxygen was being supplied to the metal surface in sufficient quantities for the actual rate of cathodic discharge to be the controlling step.

71Ross and Jones' , in 1962, studied the effect of flow on mild steel in sulfuric acid with and without added inhibitors. They observed that;


12

(1) dissolution rate (Re) - for laminar flow conditions.(2) dissolution rate oL (Re) for turbulent flow conditions.At Re values of 100,000, the dissolution rate still

increased rapidly, the systems apparently still being diffusion controlled. It was not stated whether the system was oxygenated, and if so, to what degree.

The dissolution of rotating steel specimens in 1.0M HgSO, was studied by Uhlig^^ under both aerated and deaerated conditions. In deaerated acid low carbon iron corroded at the same rate at all flow velocities. This was attributed to the activation energy of the cathodic reaction being the controlling factor. In aerated acid, velocity very much accelerated corrosion but at a decreasing rate as the velocity increased.

Levich^^ was one of the first workers to present his findings in terms of dimensionless groups. He studied rotating disc electrodes and derived the following expression for the thickness of the Nernst diffusion layer;

6 = 1.62 (= 1.62 (Sc)-'33(%e)-.50

where: h = diffusion layer thickness;w = angular velocity of the electrode;

Riddiford and Bircumshaw^^ reported on the kinetics of dissolution of zinc in aqueous iodine solutions. The rate of dissolution was found to be a function of the stirrer speed.The reactions being diffusion controlled, they established the following relationships between the dissolution constant


13k and the stirrer speed;

k = 0.0103 (r.p.m.)0'50or in dimensionless groupings

Nu = 0.558 (Re)°'56(2c)0.27

Wilke et al^^ evaluated mass transfer coefficients for free convection from limiting currents and correlated their data by;

N*av = 0.66 (Sc)0'25(Gr)0'Z5The results were in good agreement with those predicted

from the boundary layer theory for mass transfer by free convection.

25In 1934, Eisenberg et al investigated the dissolution of metal cylinders by varying, systematically, the rotational speed, the limiting current, and the cylinder diameter. It was possible to correlate their results by;

Nu = 0.079 (Re)0'70(Sc)0'356

Makrides^^ used the rotating electrode technique in his study of the anodic dissolution of iron in sulfuric acid and ferric sulphate solutions. He expressed the dependence on hydrodynamic flow conditions at a constant oxidant concentration as;

E = g log v ” ) const,where v = linear velocity

<f> = const,Tafel slope


14

Lin et al^^, in 1931,and Von Shaw et al^^, in 1963,studied several diffusion controlled processes under forced flow conditions using circular tube flow systems. In the laminar region their results were in excellent agreement with the Leveque^^heat transfer equation in which the Prandtl group was replaced by the Schmidt group;

Nu = 1.614 (sc)1/3( %

In the region of turbulent flow, Von Shaw, Reiss, and 87Hanratty ' expect the average Nusselt number in circular tubes

to be given by;Nu^^ = 0.276 (Re)0'58(Sc)0'33( d ^0.33«.V Ju

20Cornet et al studied the effect of flow rate on a copper tube in 2.IN HgSO^. They found that at high Reynolds numbers there was little increase in corrosion with increase in flow rate. This behaviour was attributed to control by a slow cathodic discharge process. In laminar flow the system was diffusion controlled. Cornet’s group correlated their data in the turbulent region by;

Nu = 0.033 (R e )0 '6 8 (g c )0 .33

Wranglen and Nilsson^^, in 1962, studied cathodic deposition on horizontal plates under forced convection.Under limiting current density conditions they expressed their results as;

Nu = 0.34 (Re)^'^^(Cc)^'^^ for laminar flow, andNu = 0.17 (Re)°'G0(Sc)°'33 for turbulent flow


13xqTwo papers published recently, one by Ibl in 1939 and

26the other by Elder and Wranglen in 1964, present theoretical studies of the application of dimensional analysis to electrochemistry. The latter paper gives a detailed theoretical treatment of mass transfer at plane plate electrodes under conditions of both forced and natural convection.

The agreement between the results of all investigations is quite good considering the difficulty of the problem. In general the power of the Reynolds number varies from O.3O to 0.70, and the Schmidt number from 0.2,3 to 0.336 under turbulent flow conditions. Most of the later workers have compared their correlations with standard heat transfer correlations and obtained excellent agreement.


CHAPTER III THEORY

In the first section of this chapter, the mechanism of dissolution of a metal in electrically conductive solutions (electrolytes) is briefly presented by examining a typical case of electrochemical corrosion - dissolution of metal in hydrochloric acid solutions. In the second section, effects of diffusional mass transfer on the dissolution process are discussed. The reason for this special emphasis is that previous workers^^' found the reaction of tin with hydrochloric acid to be mass transport controlled.

A. Reaction MechanismIt was established many years ago that metals corrode in

aqueous environments by an electrochemical mechanism. On a piece of metal that is corroding there are both anodic and cathodic sites. These may be permanently separated' from each- other, but in many instances the whole of the metal surface consists of anodic and cathodic sites which are continually shifting. At an anodic site an oxidation process occurs through a loss of electrons and the metal goes into solution by a reaction that can be depicted as(for example):

Sn - 2e ----- Sn'"'*(metal) (ionized in solution)

At a cathodic site a reduction process occurs (a gain of electrons). This will result in the reduction of dissolved

16


17

oxygen or the liberation of hydrogen gas (particularly from acid solutions) or other less common reactions. The two most common reactions can be depicted as:

Op + 4e + 4OH"

2H + 2e '— ;---^ Hp

Very detailed descriptions of the mechanism of dissolution of metals were given elsewhere by Taneja^^ and this author^^.

The fundamental mechanism of such a process can be represented in the following manner:

1. The dissolution of a metal, capable of existing in two oxidation states, by a hydrogen evolution process can be described in terms of the following elementary stages:

a. First, when an acid such as hydrochloric acid dissolves in water, we have

H c i + + C l"

The bulk solution, therefore, consists of water molecules, hydronium ions, and chloride ions. In general, contact between two immiscible phases is accompanied by an increased concentration in the fluid phase close to the interface. This tendency leads to the formation of a "diffusion layer". Therefore, when a metal is placed in the deaerated acid solution, the hydrogen ions will diffuse through this layer, from the bulk solution to the metal-solution interface,

Tt“ TT'“1


18

where subscript b denotes bulk solution, and subscript i denotes metal-solution interface. (Hereafter we use

instead of for convenience.)b. The adsorption of hydrogen ions on the metal surface

h ! + M

where M-H^ denotes a hydrogen ion adsorbed on an active * ®site, M,on the metal surface:

*c. Discharge of adsorbed hydrogen ions by virtue of

electron transfer through the metal:(cl) Electrons transferred from a neighbouring vacant

site on the metal surface

H H! I I I^ ta Mc -■» j Ml.

# * * * * *(c2) Electrons transferred directly from the surface

atoms on which adsorption occurred

M-H+ ---* 8 * &

d. Formation of molecular hydrogen from atomic hydrogen, by one of the following possible modes:

(d1)

H H-A - M - M"*"- M - M - -M - M - M"*- M - M -* * * * * *

(d2) 2 ( M+-H_)-----* 8 *s *


19

+ HÎ — ------------ - M*-Hp .* 8 1 *

e. The removal of molecular hydrogen from the metal surface by one or more of the following possibilities:

(el) Dissolution of molecular hydrogen into the solution at the metal-solution interface

M-Hp ---^ M + Hp* s * i,dissolved

(e2) Absorption of molecular hydrogen from the metalsurface into the lattice structure of the metal

-M-Hp -M-— M — M — M — M — — M — M « M — M —

I I I I I I ^21 I- M - M - M - M - - M - M - M - M -(e5) -Formation of hydrogen gas bubbles on the metal

surface and their subsequent detachment from the metal surface

n( M-Hp) — - + (n - 1)%* ^ li *

M-(Hp)^ ---- M + (Hp)^^ ^ XI ^ ^ XI

where M-(Hp) represents a hydrogen gas bubble attached on* c. nthe metal surface:

f. Desorption of metal ions from the metal surface

g. Diffusion of metal ions(probably hydrated) from the metal-solution interface to the bulk solution

MÎ+ + nH^O M+lnH^O1 2 2


20

2. In aerated solutions, two other processes can occur together with hydrogen evolution, namely, oxygen depolarization and autocatalysis.

Investigations^-^have shown that the following successive stages are in good agreement with experimental data for the reaction of oxygen depolarization:

a. Diffusion of molecular oxygen from the bulk solution to the metal-solution interface

Where b denotes bulk solution, and i denotes metal- solution interface.

b. Adsorption of oxygen on the metal surface; there aretwo types of adsorption possible for oxygen on metal surface^^*^^:

(b1) Molecular adsorptionM + Op ===== M-Op * i * s

(b2) Dissociative adsorption2M + Op 2( M-0^)* “

c. Formation of a mono-valent oxygen ion(cl) _

M-Op M:Op-M - - M “ M - -M - A - M - M -(c2)

M-0 M-0 M^-O" M^-0"* * * *I—M — M — M — M — ---—M — M — M — M —

d. Formation of perhydroxyl radical or hydroxyl radical


21

( d 1 )

M*-0" M-He ^ M*-HOp +* 2g * 8 * *8(d2)m"-0" + — M*-HOe H-

e. Formation of perhydroxyl ion by virtue of electron transfer

M -HOp -^ M -HO:* s *

f. Formation of hydrogen peroxide by one of the following possibilities:

(f1)

M+-HO: + M-HÎ — M-HpOp +* * 8 * 2 2g *8(f2)

2( M' -HO ) ^ M-HpOp -1-* * s *

g. Reduction of hydrogen peroxide with the formation of hydroxyl ions. This step is itself a complex process and can be further broken down into the following sequence:

(g1) Reduction of hydrogen peroxide to hydroxyl ion and hydroxyl radical

M-HpOp -- ^ M+-HO= + OH:* 2 2g * 6 ].(g2) Reduction of hydroxyl radical to hydroxyl ion

M+-HO= ---^ + OH:* s *S 1

h . Desorption of the metal ions from the metal surface


22

i. Diffusion of the metal ions from the metal-solution interface to the bulk solution

m T- - nHgO --- -

Many of these elementary stages are experimentally proven facts. For instance, the existance of mono-valent oxygen ions (0 ) in the cathodic reduction of oxygen in aqueous solutions has been verified by Krasilshchikov^^.

3. For metals that can exist in two or more oxidation states in acidic media, an autocatalytic reaction should be considered as a possibility. The following reaction scheme is suggested as a possible mechanism for this process when the ion of higher oxidation state is quadrivalent of the1 y - 4 4- 4- 4-type M :

a. An oxidation process^ caused by dissolved oxygen with the formation of peroxide intermediate according to

M++ + Og + 2H+--- + HgOg

This process can occur in the bulk of the solution, at the metal-solution interface, or at the metal surface

' M y . HgOg(a1 ) «r ^ °2b(a2) °2.1(a3) M-0*

b

1 ■2H' — — Mp''" + H,0

2, ' T " :


23

(a4) + M-0= + M-lC ^*8 * s * s

Ml*** + M-OH" + ML*8 * *S

b. An oxidation process caused by the peroxide intermediates^^(produced in stage a and in the oxygen depolarization reaction) according to

M-HÎ + MÎ* + + M-OH" + HpO

orM-H* + MÎ+ + M-HpOp ------* 8 *8 * ^

M++++ + M-OH% + M-HpO=* S ^ CL S

c. An equilibrium^ between metal ions of higher and lower oxidation states by virtue of electron transfer, either at the metal surface or at the metal-solution interface:

(cl )-M-* -M*- M- _#!*** -M - M-* * * * * *I I I I I I—M “ M — M — M — M— —n.—.—- —M — M — M — M — M—

(c2) + Mg 2M2+

Stages a, b, and c constitute an autocatalytic reactionsequence.B. Mass Transfer in Electrochemical Dissolution Processes

1. Fundamental EquationsThe laws of transport in dilute electrolytic

solutions have been known for many years and have been discussed in detail e l s e w h e r e . The flux of a species is


24

due to migration in an electric field, diffusion in a concentration gradient, and convection with the fluid velocity.

N. = -z.u F C V ^ - D.VC. + vC. (1-3-1)V Ü V V V «J Ü

2where = flux of species j (mole/cm.-sec)

zj = charge number of species ju . = mobility of species j(sq.cm.-mole/joule-sec)F = Faraday's constant(coulomb/equiv.)Cj = concentration of species j(mole/c.c.)0 = electrostatic potential(volt)Dj = diffusion coefficient of species j(sq.cm./sec)V = fluid velocity(cm./sec.)

A material balance for a small volume element leads to the differential conservation law:

+ B (1-3-2)O V J J

where R . = homogeneous rate of production of species j (mole/c.c.-sec)

Since reactions are frequently restricted to the surfaces of electrodes, the bulk reaction term R_. is often zero in electrochemical systems. To a very good approximation the solution is electrically neutral hence

I]z,C. = 0 (1-3-3)j J J

The current density in an electrolytic solution is due to the motion of charged species:

i = F]Cz,N. (1-3-4)j J J


25

where i = current density(amp/sq.cm.)These laws provide the basis for the analysis of electro

chemical systems. The flux relation, Equation(1-3-1), defines transport coefficients - the mobility u^ and the diffusion coefficient Dj of an ion in a dilute solution. Many electrochemical systems involve flow of the electrolytic solution. The fluid velocity is to be determined from the Navier-Stokes equation

2C( + v'Vv ) = - V p - (1-3-5)where Ç = fluid density(gm./c.c.)

p = fluid pressure(gm./cm.-sec.)= fluid viscosity(gra./cm.-sec.)

pg = gravitational acceleration(cm./sec.) and the continuity equation

= 0 (1 -3 -6 )

For the reaction of minor species in a solution containing excess supporting electrolyte, it should be permissible to neglect the contribution of ionic migration to the flux of the reacting ions, so that Equation(I-3~1) becomes

N. = -D . \7 c . + vC. (1 -3 -7 )J J J Jand substitution into Equation(1-3-2) yields

^ 6 . 2^^4 + v'^/Cj = Dj\7Cj (1-3-8)

This is the general equation of convective diffusion.For a full solution of the equations of hydrodynamics and


26

convective transfer, we must also know the boundary and initial conditions that satisfy the different systems under consideration. It would, of course, be more satisfactory to obtain a complete solution of these equations, but, this is, in most cases,impossible. Therefore, in practice, most of the work in this field has been done by dimensional analysis and empirical approach.

2. Nernst Diffusion Layer TheoryEssential to the understanding of convective transfer

problems is the concept of the diffusion layer. Frequently,due to the small value of the diffusion coefficient, theconcentrations differ significantly from their bulk valuesonly in a thin region near the surface of an electrode.According to Nernst^^, a practically stationary liquid layermust be assumed in contact with the electrode surface,regardless of whether the electrolytic solution is stirred ornot. Within this layer molecular diffusion and ionic transportare of primary importance to the transport process, whileoutside this layer the convective transport dominates and theconcentration is maintained at a constant value by convection.The concentration gradient was assumed to be linear andrelated to the solute mass transfer flux N . by the equation

D,Nj = ) Cl-3-9)

where G. = concentration of species j at the solution- electrode interface(mole/c.c.)


27

= concentration of species j in bulk solution (mole/c.c.)

S = diffusion layer thickness(cm.)In general, the flux N^ of the ions through a cross

section parallel to the surface and away from the surface is additively composed of the diffusion terra and the migration term. The flux is therefore, given by

( c - C ) i.t= “3 " - Ï

where tj = transport number of species jNernst's theory postulates that the liquid is stationary

within the diffusion layer, which is contrary to experimental data obtained for fluid flow near solid surfaces^^. In reality the transition from diffusion'to convection takes place continuously. Furthermore, the theory does not allowquantitative predictions to be made as to the dependence of ô

on the flow conditions.

3« Application of Dimensional AnalysisThe diffusional flux is now usually written in the

formez = kj( ) (1-3-11)

where k . = mass transfer coefficient of species j.The value of k . can be determined by dimensional analysis

Jand the correlations have the general form

Nu = const. (Re/* (Sc) (1-3-12)for forced convection


28

Nu = const.(Gr)^(8c)*for natural convection

The exponents a , f , 7 , and e are determined by experiment where the dimensionless groups are

Grashof number, Gr = ^

Reynolds number,Re =k.L

Nusselt number, Nu = —^MSchmidt number. Sc =

where L = characteristic length(cm.)The physical significance of the dimensionless groups is

of interest. The Reynolds number is defined as the ratio of the non-viscous or inertial forces to the viscous forces acting on the element of fluid. At low values of Reynolds number the inertial forces are small compared with the viscous forces and the flow is streamlined. At high values of Reynolds number the inertial forces predominate and the flow becomes turbulent.

The Schmidt number is the ratio of the fluid property governing the transfer of momentum by viscous forces, the kinematic viscosity, to the fluid property governing mass transfer by diffusion, the diffusivity.

It is possible to obtain the dimensionless correlations (1-3-12) and (1-3-13) by;

(a) direct dimensional analysis,


29

(b) writing the fundamental equations of motion and diffusion in dimensionless terms,

and (c) employing the Prandtl-Taylor or Chilton-Colburn extensions of the Reynolds Analogy as applied to mass transfer.

By analogy with their jy factor for heat transfer Chilton and Colburn deduced a factor for mass transfer and showed that the two were approximately equal. Thus it is possibleto calculate mass transfer coefficients from data obtainedfrom heat transfer experiments

4. Diffusional Mass Transfer in a Circular TubeIn this section we consider a special simple case

as an example of the application of the dimensional analysis method by writing the fundamental equations of motion and diffusion in dimensionless terms. Consider the steady isothermal flow of an acid solution inside a tube of circular cross section as shown in Figure 1-3-1.

11 II 112"

4 ^ 'C V ^ 4 4/ ^ y 1Fluid r Fluid leaves wiNfchenters witld^ z _ bulk composition—uniform composition-----------------7U-&-4 / / ; // ? / /v?i ' ^

Plastic Fluid composition next Velocity of dissolvedtube to metal wall surface metal ions away from

maintains constant value metal wall is assumed to beC» very small1

FIGURE 1-3-1. Mass Transfer in a Tube With a Soluble Metal Wall.


30

Assume that the flow pattern is fully established and the velocity distribution is known at plane"1" and that the fluid concentration is uniform at in the region z < 0. From z = 0 to z = L, the tube is made of a metal or other soluble solid, which dissolves slowly and maintains the liquid composition along the dissolving surface at constant value C» , We further assume that the physical properties p, p, and are constant.

The diffusional mass transfer at the solution-metal interface can be described by starting from Equation(1-3-7)•

®M = 'fM (1-3-14)in this case j = M, and at the metal surface, within the diffusion layer, the fluid velocity is very small and the term vC^ can be neglected. Therefore, the flux of M away from the metal surface is

Nj, = -D,^VC„ = -D , (1-3-15)

The total molar rate of addition of species M by diffusion over the whole metal tube surface between plane "1" and "2" in Figure 1-3-1 is given by the following expression, valid for either laminar or turbulent flow:

Total molar rate of addition of species M'*L ran dC

( D, »O J 0 “ àr

where R is the radius of the tube. But from Equation(1-3-11), we have

)

)Rdg dz (1-3-16)

Repro(juc8(j with permission of the copyright owner. Further reproctuction prohibitect without permission.

31and therefore;

The total molar rate of addition of species M= A ( Cm^ M

where d denotes the inside diameter of the metal tube, and A denotes the inner surface area of the metal tube.

Substitution of Equation(1-3-17) into Equation(1-3-16)gives

1n id ( cM.

LOJ

2D0 (D,

doM p^g)Rd8 dz (1-3-18)M dr1 "b

This equation can be reduced to a dimensionless form by*introducing the dimensionless quantities; r = r/d, z = z/d,

and CM Using these quantitiesin the equation and rearranging it gives

Nu = 1DM -

'L/df2n dC0 JOJc (“ M

dr* r )d0 dz (1-3-19)

The general equation of convective diffusion for anyspecies is

dt v.T^C = D\fC (1-3-20)This equation can be simplified if radial symmetry is

present to the convective diffusion equation in cylindrical coordinatesacat V.- ( l? ) "

(1-3-21)

ReprocJucecJ with permission of the copyright owner. Further reproctuction prohibitect without permission.

32

This equation may now be reduced to dimensionless form by introducing a characteristic flow velocity , which is the uniform velocity of the solution in the tube.

Multiplying Equation ( 1-3-21 ) by , where is theuniform concentration in the bulk of the solution, we have

V(1-3-22)

Introducing the dimensionless quantities; v = v / U ,

V* = v / U , t" = tUg/d , r = r/d, z = z/d, and C = C/C

we obtainIf •

D / 22c*\ D /3Zc*\ D /3C*\= U , d r * U r V

(1-3-23)0 dThe dimensionless ratio, Pe = - y , is known as the

Peclet number. Introducing the Peclet number into Equation(1-3-23), we have


33

at'

Pe

But Pe becomes

3G*

VD

3 r

loi J:y ' D

*23z

1 /ac^ \? r *

(1-3-24)

Re*Sc. Therefore, Equation(1-3-24)

1 ?2c*?r Re • Sci

aZç**2

~dz

1 3C*

* * r 9r- V* ?C

* 3c*

Equation(1-3-23) shows clearly

3c '3r

2z 2 t

(1-3-23)

(1-3-26)= f ( Re, Be)

and consequently, Equation(1-3-19) is o f the form

Nu = * ( R e , 8c, d /L )

or Nu = con st.(R e)M (3c)B ( ^ )P (i_3_2? )

where the exponents, m, n, and p for the dissolution of metal tubes of circular cross section can be experimentally determined. Previous workers^^'^7,93 gj owed that m varies between 0.4 and 0.8, n between 0.3 and O.4, and p between0.3 and 0.4.


CHAPTER IV

EXPERIMENTAL

A. Equipment1. Design of Liquid Recycle System and Accessories

The schematic diagram of the experimental equipment is shown in Figure 1-4-1. The flow system is a closed loop which was constructed from plexiglas tubing. A three- fifteenth horsepower Dynalab Model 4-ML magnetic drive pump, capable of delivering 660 gal./hr. water against 3 feet head, was used to circulate the corroding solution.

Adjustment o f the electrolyte flow rate was achieved by means of globe valve B in the pump by-pass. Two rotameters equipped w ith stainless steel floats provided an in d ic a tio n

of actual flow rates.( One was calibrated to 4 USGPM of 1.0N

HCI a t 23°C and the o th er to 13 USGPM.)

The 2 0 - l i t r e QVF glass supply tank was equipped with an immersion Vycor brand glass h e a te r, a thermo-regulator thermometer, a Teflon coated stirrer, and a gas saturation system. The 3 5 - l i t r e polypropylene make-up tank was equipped with a Teflon coated electric stirrer.

The test section, mounted in a horizontal position, was located on the downstream side of .the pump. The only part of the system on the suction side was about 6 feet of PVG pipe leading from the top supply tank to the pump. This meant, that if any leaks developed, escaping electrolyte

34


CD

cyM

o o O

1 r

- Q =

I 1 Ï

I I d

;0302

IQMgpAo

MA0 M1MO02

IIH

Is


36

LEGEND - For Figure 1-4-1

A, F, G, H - PVC ball valvesD, E, B - PVG globe valvesC - QVF glass globe valveI - Sealless, magnetic drive pumpJ - Make-up tank, 33 litre, polypropyleneK - Teflon coated electric stirrerL - Compressed gas cylinderM - Test sectionN - Rotameter0 - Supply tank, QVF glass, 20 litreP " Constant temperature water bathQ - RelayR - Electric heating coilS - StirrerT - Thermo-regulator thermometerU - Glass washing towerV - Gas dispersion tubeW - Vycor brand glass heaterX - Thermometer


37

would be easier to see than air being drawn into the system.Eight feet of clear plexiglas piping immediately

preceded the test section to allow development of the hydrodynamic boundary layer. Three feet of clear plexiglas piping immediately followed the test section to ensure steady flow in the test section.

Gas saturation of the acid solution was achieved by passing the appropriate gases from compressed gas cylinders through a series of wash bottles containing HGl solution of the same concentration as in the supply tank. All wash bottles were kept at the same temperature as the corroding solution by immersing them in a constant temperature bath.

The return inlet to the supply tank was located below tbe solution surface and directed away from the gas dispersion tube to prevent large gas bubbles from being drawn into the system.

A thermometer, incorporated near the return inlet by means of a polypropylene T-piece, provided a check on the temperature variation of the system.

Except for the calming and test section, the circulation loop consisted only of standard half inch plexiglas and PVC pipes, valves, and joints. As a result of the choice of materials and equipments the only metal components contacting the corroding solution other than the surface of the test section are two Type 304 stainless steel floats in the rotameters. The dissolution of Type 304 stainless steel in


38

dilute HCl solution (up to 6.ON) is negligible.2, Details of Test Section Construction

A detailed drawing, illustrating the design and materials used in the construction of the test section, is provided as Figure 1-4-2.

Tin bars, originally about 1.02 cm. in diameter, were melted and remoulded into 1 inch and 1,3 inch diameter rods under argon gas in a vacuum furnace. Tin tubes of 0.23 to 1.0 inch inside diameter, 2.3 to 3.3 inches in length were produced from these remoulded tin rods.

Pairs of spigot were machined from 1.3 and 2 inch diameter polypropylene rods to different dimensions to accommodate the tin tubes.

Both ends of the tin specimen were recessed into the spigots and secured firmly with plastic cement. In addition, functioning as watertight seals, two Teflon washers were placed between the end surfaces of tin specimen and the stopping edges of the spigots.

Before any test section assembly was incorporated into the system, it was held together, under pressure, so that any Teflon washers projecting beyond the inner surface could be machined flush with the tube walls.

3. Temperature Control and MeasurementThe temperature control of the corroding solution

was achieved by setting the desired temperature on the thermoregulator-thermometer which sensed the temperature


:)9

P4 EH

§HEHO

EhcqgPh0C5M

§

<t|1

•V■4-I

g§ I—I P4


40

variation of the corroding solution and maintained it at the set-point by activating a 5OO watt immersion Vycor brand glass heater through a relay. The temperature control used in this system was supplied by the Chemical Rubber Company who claimed that it was capable of controlling the temperature within + 0.01°C.

The temperature of the solution at the return inlet was also checked and recorded. If the difference between this temperature reading and the temperature reading of the bulk solution exceeded 0.1°C the average value would then be used for data analysis.

B. ' Chemicals and MaterialsThe Analar grade tin bars used in this study were supplied

by British Drug House LTD. The purity of metal was 99*92 per cent tin. The exact analysis according to the manufacturer showed 0.01 per cent lead, 0.0025 per cent copper, 0.002 per cent bismuth, 0.002 per cent iron, O.OO4 per cent total foreign metals, 0,0001 per cent arsenic, and 0.025 per cent antimony.

Hydrochloric acid and all other reagents used were of analytical grade. Dilute HCl solutions were made with distilled water.C. Procedure

1. Experimental PreparationThe two rotameters were calibrated by maintaining

a particular reading and measuring the volume of effluent


41

collected during a known time interval. Details of rotameter calibration are given in Appendix 1.

The make-up tanli and the QVF glass supply tank were also calibrated. The liquid capacities at different liquid levels were measured and marked.

Before the tin tube was integrated into the flow system it was polished manually to 3/0 emery paper smoothness(average surface roughness less than 20 microinches) and washed initially with distilled water and dried with filter paper, then washed with acetone for degreasing, and finally rewashed with distilled water. After each run, the tin tube was washed and dried. A check on the material balance was maintained by weighing the clean dry tin tube before and after each run.

When not in use, the tin tube was filled with distilled water. At the start of a run, 0.3N HCl was flushed through for about 3 minutes. This removed any oxide built up during storage in distilled water.

2. Procedure for Dissolution RunThe HCl solutions were made up and mixed in the

make-up tank with valves A and H closed(refer to Figure 1-4-1). Valves A, D, E, F, and G were then opened, valve C closed, and the solution was pumped from the make-up tank into the supply tank. After the volume of solution in the system was adjusted by letting the excess solution drain through valves D, E, F, and G into the make-up tanli, valve A was closed and C opened.At this stage the system was operational. The appropriate gas


42

was then introduced into the solution through the dispersion tube.

After the solution was flushed overnight, the temperature control system and the stirrer in the supply tank were switched on. During the time required for the acid solution to reach the desired temperature, the test section was prepared and connected to the flow system as previously described.

After the solution reached the desired temperature, the pump was switched on and the flow rate was regulated by adjusting valves B, D, and E.

A 5 ml. sample of the corroding solution was withdrawn for routine analysis at convenient intervals of time from the top supply tank. For every sample of corroding solution withdrawn, an equal volume of HCl solution of the same concentration was added to the supply tank to minimize solution volume change during the dissolution run.

Tin concentrations were determined by means of a Jarrell- Ash Model 82-526 Maximum Versatility Atomic Absorption Spectrophotometer. Details of the analytical procedure are given in Appendix 2.

Upon termination of a dissolution run, the system was drained of corroding solution, flushed with city water three or four times and then flushed with distilled water once.The test section was then disconnected from the system and stored in distilled water until further use.


CHAPTER V EXPERIMENTAL RESULTS AND DISCUSSION

A. Method of Analysis of Experimental Data; A Typical ExampleIn Chapter III details of the mechanism of dissolution of

a metal, capable of existing in two oxidation states in acid solutions, were presented. On the basis of that discussion the following characteristics of dissolution can be visualized:

The corroding solution concentration-time plot for the dissolution of such a metal in aerated solutions will show an autocatalytic effect, with dissolution rate increasing with increasing metal ion concentration in solution. The over-all dissolution rate is a combination of the following three parallel corrosion effects: hydrogen evolution, oxygen depolarization, and autocatalysis. The rates of hydrogen evolution and oxygen depolarization are independent of the metal ion concentration in the solution(zero order dependency on metal ion concentration). Therefore, the corroding solution concentration-time plots for these two reactions should be straight lines. The oxygen depolarization and autocatalytic reactions can occur only when oxygen gas is introduced into the solution. The autocatalytic effect is negligible at the initial stage because no metal ions are present in the solution at the start of a dissolution process.

The analysis of the dissolution data obtained in this

43


44

study is based on the above discussion with experimental data recalculated to a metal dissolved per unit surface area basis.A typical model of this analysis is illustrated in Figure 1-5-1, where OA shows the plot of the rate data for the dissolution of tin in deaerated hydrochloric acid solution, a condition under which only hydrogen evolution can occur. Curve 00 shows the rate data obtained under the same conditions as those of OA except that the acid solution was air saturated.

The shape of OA, a straight line, conforms to a zero order plot of metal ion concentration vs time. The exponential concentration-time plot of curve 00 is a confirmation of an autocatalytic process. The straight line OB has been drawn asymptotically to curve 00 at its initial stage where autocatalysis is negligible. This linearity essentially represents an extrapolation of the initial corrosion rate under aerobic conditions in the absence of autocatalysis.

The total area under curve 00 is divided by straight lines OA and OB into three regions, each of which corresponds to a specific effect on the over-all dissolution reaction as shown in Figure 1-5-1, where slope of line OA = r^, slope of line OB = r + rg, and slope of curve 00 = r.j + r^ + r^. ( r.j, and r? are the reaction rates in moles of metal dissolved per square centimeter per minute for hydrogen evolution, oxygen depolarization, and autocatalysis respectively.)

This analysis provides a pattern for the interpretation of the tin dissolution data obtained during this investigation.


45

12,000 ml. 1M HCl, 5,730 cm./min 30^0, d = 1,27 cm., L = 10.16 cm A Bn]A Un]

ALSn]

2.8due to Eg evolutiondue to oxygen depolarizationdue to autocatalytic reaction

2.4

2.0

o r, +rw<Dr 4

0

X

800600 10000 200TIME - minutes

FIGURE 1-5-1. A PROPOSED GRAPHICAL ANALYSIS FOR TIN DISSOLUTION IN HCl SOLUTIONS


46

B. ReproducibilityIn order to check the reproducibility of experimental

determinations duplicate runs were made under various conditions. A typical plot of a duplicate run is shown in Figure 1-5-2.C. Effects of Different Factors on the Dissolution Rate

of Tin1. Rate Dependence on Tin Ion Concentration

As shown in Figure 1-5-5, the autocatalytic dissolution effect, which is represented by the difference in tin ion concentration between curve OC and the linear zero- order initial rate plot OB as shown in Figure 1-5-1, can be correlated by a half-order plot. The reaction rate of autocatalysis, r,, is proportional to the total tin ion concentration in the bulk solution raised to the 0.5 power. Experimental evidence shows that the autocatalytic reaction becomes important only after an elapsed time of 40 to 120 minutes. The beginning of noticeable autocatalysis, as determined by the deviation from the asymptotic line OB in Figure 1-5-1, depends on the dissolution conditions, but in general the bulk tin ion concentration must be greater than0.35x10”^ moles/litre.

2. Rate Dependence on Fluid VelocityThe effect of the velocity of the acid solution

flowing through the tube was studied for velocities ranging from 1,175 to 16,600 cm./rain, in tin tubes of different


47

0)u-p'H

g 4I—IOs

-cJ-OX

0

12,000 ml. IN HCl, T = 30°C V = 4,150 cm./min. d = 1.27 cm. L = 10.16 cm. AIRA - RUN A O - RUN B

O/▲6

.0

./

. /

10 200 800400 600

TIME - minutes FIGURE 1-5-2. REPRODUCIBILITY OF EXPERIMENTAL DATA

1,000


48

12,000 ml. Air Saturated 1M HCl = 1.905 cm., L = 7.62 c: 16,600 cm./min.8,500 cm./min.4,700 cm./min.

20

#

i<D

r HOa

A

501 10 202[in] X 10^, (moles/liter)

FIGURE 1-5-5. DISSOLUTION AS FUNCTION OF TIN IONCONCENTRATION IN THE BULK SOLUTION

( FOR AUTOCATALYTIC REACTION )


49

diameters and lengths, at three different temperatures.The results shown in Figures I-5-4, 1-5-5, and 1-5-6

indicate that r^, the reaction rate for hydrogen evolution, rg , the reaction rate for oxygen depolarization, and r^, the reaction rate for autocatalysis are proportional to fluid velocity raised to the 0.67, 0.59, and 0.61 powers respectively.

This behaviour indicates that resistance to diffusional mass transfer is operative for the range of velocities (turbulent flow in all cases) investigated.

3. Rate Dependence on TemperatureThe rate dependence on temperature was studied over

the range 25 to 50°C under nitrogen, air, and oxygen saturation at four different flow velocities. The results are illustrated by the Arrhenius' activation energy plots shown in Figures 1-5-7, 1-5-8, and 1-5-9.

The temperature variations correspond to apparent activation energies of 4.25 kcal per mole for the hydrogen evolution reaction, 5*45 kcal per mole for the oxygen depolarization reaction, and 5*8 kcal per mole for the autocatalytic reaction. These low values of activation energy suggest that the temperature coefficients for these reactions are low and the controlling step is more likely physical than chemical.

4. Effect of Hydrochloric Acid ConcentrationThe effect of hydrochloric acid concentration was

studied in aerated and deaerated solutions over the


50

100

8I8Ü>

CDI— !O8

ooOM

50

20

10

5

.5

12,000 m l. Ng SATURATED IN HCl □ - d=1.27 cm. 1=10.16 cm. T=40°C O - d=1.27 cm. L=10.16 cm. T=50°C A - d = 1 .905cm. L=7.62 cm. T=25°C

J L1 2

FIGURE 1-5-4.

10 20VELOCITY X ICT^, (cm ./m in .) DISSOLUTION AS FUNCTION OF FLUID VELOCITY( FOR H EVOLUTION REACTION)

50


51

1 0 -12,000 ml. AIR SATURATED IN HCl o - d=1.27 cm. L=10.16 cm. T=40°C O - d=1.27 cm. 1=10.16 cm. T=30°C A - d= 1..905cm. 1=7.62 cm. T=25°C

8Ü(DHO

OoX

.2

. 1 J L-J 1 LI

FIGURE 1-5-5.

5 10 20VELOCITY X icr3\ (cm./min.)DISSOLUTION AS FUNCTION OF FLUID VEL0CITY( FOR OXYGEN DEPOLARIZATION REACTION )

50


52

•douaü>0)iHOaIN-oX

20

AT

10

.5

12,000 ml. AIR SATURATED IN HCl d = 1.27 cm. L = 10.16 cm. T = 30°C" - (Snj = 7x10”^moles/litre^ - [Sn] = 3*5x10'"' raoles/litre• - |Sn[ = 1.75xlO“^moles/litre

I I I

FIGURE 1-3-6,

3 10 20 VELOCITY X 10"3, (cm./min.) DISSOLUTION AS FUNCTION OF FLUID VELOCITY(FOR AUTOCATALYTIC REACTION)

50


53

12,000 ml. N_ SATURATED IN HCl

16,600 cm./min. 10,280 cm./min. 4,150 cm./min. 1,175 cm./min.

20 -

(\t*aÜ

■■'A-

00o

5.1

FIGURE 1-5-7.(1/T) X 10^, (1/°K)

DISSOLUTION AS FUNCTION OF TEMPERATURE ( FOR HYDROGEN EVOLUTION REACTION)


54

12,000 ml. AIR SATURATED IN HCl d = 1.27 cm. L = 10.16 cm.□ ~ V = 16,600 cm./min.A _ V = 10,280 cm./min.- V = 4,150 cm./min.

A - V = 1,175 cm./min.

FIGURE 1-5-8.

5.2 5.5 5.4(1/T) X 10^, (1/°K)

DISSOLUTION AS FUNCTION OF TEMPERATURE ( FOR OXYGEN DEPOLARIZATION REACTION )


55

ü

I( Maü(Q(DiHOa

o-oX

50

20

10

5.0

12,000 ml. AIR SATURATED IN HCl d = 1.27 cm. L = 10.16 cm.V = 16,600 cm./min.° - [Sii] = 2.0 X 10”^ moles/litreA - {Siïj = 1.0 X 10“^ moles/litre• - Csn] = 0.5 % 10 ^ moles/litre

5.1 5.4 5.5

FIGURE 1-5-9.

5.2 5.5(1/T) X 10^, (1/°K)

DISSOLUTION AS FUNCTION OF TEMPERATURE ( FOR AUTOCATALYTIC REACTION )


%

concentration range 0.10 to 4*0 M HCl. Analysis of the data obtained in these investigations shows that the rate of dissolution of tin tubes is essentially independent of HCl concentration. Table 1-5-1 indicates that there is no apparent trend for r^, rg, and r^ over the range of concentration studied.

TAELS 1-5-1Effect of HCl Concentration on Tin Dissolution T = 30°C, V = 16,600 cm./min., d = 1.27 cm.L = 10.16 cm., [bn] = 2.0 x 10”^moles/litre

HCl

(N)*HC1(N)

1/ moles \

^2f moles \

^5/ moles \'■2cm.-min. V 2 / cm.-min. I 2 ' cm.-min.

0.10 0.079 7.95 X 10-G 2.60 % 10"? 4.05 X 10“'?0.25 0.194 7.52 - 2.55 - 4.18 -U.35 0.253 6.95 - 2.50 - 5.96 -0.50 0.379 7.15 - 1.95 - 5.85 -0.75 0.586 7.65 - 2.05 - 4.22 -T.OO 0.809 6.82 — 2.86 - 5.75 -2.00 2.018 7.95 - 2.25 - 4.15 -3.00 4.080 7.20 - 1.85 - 4.20 -4.00 7.280 6.75 - 2.55 - 3.80 -

Dissolution rates independent of acid concentration were also reported by other i n v e s t i g a t o r s ^ T h i s behaviour is an indication that the diffusion of hydrogen ions from the bulk solution to the solution-metal interface and the adsorption of hydrogen ions on the active sites on the tin surface are fast steps. Therefore, the variation of hydrogen ion concentration in the bulk solution has little influence on the dissolution rates of tin in hydrochloric acid solutions.


57

5. Effect of Oxygen Partial Pressure in Gas PhaseThe variation of the concentration of oxygen in the

gas mixtures used to saturate the acid solutions and purge the flow system was achieved by adjusting the ratio of volumetric flow rates of oxygen and nitrogen from gas cylinders.

The influence of oxygen concentration on the dissolution rate was determined by saturating the acid solutions with air, and gas mixtures of oxygen and nitrogen. The proportion of oxygen in these gas mixtures was varied over the range 0 to 1.0. This investigation was performed at three different velocities; 16,600cm./min., 10,280 cm./min., and 4,150 cm./min.

The results shown in Figures 1-5-10 and 1-5-11 indicate that the rate of dissolution can be well correlated in terms of the square root of the oxygen partial pressure in the gas phase with which the solution is equilibrated. The reaction rates r^ and r^ are directly proportional to the square root of the oxygen partial pressure. This behaviour suggests that the adsorption of oxygen on the metal surface is possibly of the dissociative type.

6. Effect of Tube Inside DiameterThe effect of the inside diameter of the tube was

studied over the range 0.955 to 2.54 cm., under nitrogen and air saturation, and at the following three flow velocities;1,175 cm./min., 5,750 cm./min., and 16,600 cm./min. It was

2found that the dissolution rate in moles/cm.-rain, decreased


^8

<\jaooiHa

i>-oX

6

12,000 ml. IN HCl 30°C

16,600 cm./min 10,280 cm./min 4,150 cm./min• -

4

3

0/

2

1

0 0 2 6 8 1.0( Pp. in Gas Phase - (atm.)^0^

FIGURE 1-5-10. DISSOLUTION AS FUNCTION OF OXYGENCONCENTRATION( FOR OXYGEN DEPOLARIZATION )


59

1

12,000 ml. IN HCl 30°C

V = 15,800 cm./min.□ - |Cn) = 2.5 X 10”^ moles/litre^ - [Sr = 1.5 X 10”^ moles/litreO - (jsi = 1.0 X 10”- moles/litre

oj •

8 1 . 060 .2 4 1 4-

( Pp. in Gas Phase ) - (atm.)^^2

FIGURE 1-5-11. DISSOLUTION AS FUNCTION OF OXYGENCONCENTRATION(FOR AUTOCATALYTIC REACTION)


60

slightly with increasing tube diameter.The results given in Figures 1-5-12, 1-5-13, and 1-5-1A

show that r , r^, and r^ are proportional to and d"" * ^respectively,

This behaviour can be explained by the effect of tube diameter on the thickness of the diffusion layer s .

Levich^^, who examined the case of the inside surface of a tube serving as the reaction surface under conditions of both laminar and turbulent flow through the tube, derived the following relationships for the thickness of the diffusion

for laminar flow and ge6(d)T/G

for turbulent flow The range of Reynolds number encountered in this study is

between 3,940 - 74,300. As a result all experimental data fall within the region of turbulent flow. Therefore, according to Levich, the relationship, 8 o6(d)^^^, should hold true forthis investigation. In other words, the diffusional mass

— 1 /Pitransfer flux should follow the relationship; N oC (d)“ .

A comparison of the results obtained in this investigation with the above relationship shows that they are of the right order of magnitude.

7. Effect of Tube LengthThe effect of tube length on the dissolution rate


61

a

i o(I)HOS

°o

10

. 6

12,000 ml. IN HCl, NITROGEN SATURATED 30°CL = 10.16 cm.^ - V = 16,600 cm./min.® - V = 5,730 cm./min.■ - V = 1,175 cm./min.

_L5.01.0 2 .0

DIAMETER - cm.FIGURE 1-5-12. DISSOLUTION AS FUNCTION OF TUBE

INSIDE DIAMETER( FOR EVOLUTION)


62

OJsü© r—I Oa

!>-oX(V

.5

.3.0

12,000 ml. AIR SATURATED IN HCl 20°C, L = 10.16 cm.A V = 16,600 cm./min.• - V = 3,730 cm./min.o - V = 1,173 cm./min.

JL_J L ±

FIGURE 1-3-13.

2 3 DIAMETER - cm.DISSOLUTION AS FUNCTION OF TUBE INSIDE DIAMETER(FOR OXYGEN DEPOLARIZATION)


63

CMaolo©Ha

M

10

. 6

12,000 ml. AIR SATURATED IN HCl 30°C, L = 10.16 cm.V = 16,600 cm./min.A _ [sn] = 1.2 X 10‘"^moles/litre• - [sn] = 0.7 X lOT^moles/litre□ - [srQ = 0.33% lOT^moles/litre

-p- — □-

I1 2 3DIAMETER - cm.

FIGURE 1-3-14. DISSOLUTION AS FUNCTION OF TUBEINSIDE DIAMETER(FOR AUTOCATALYTIC REACTION)


64

was studied over the range 3*073 - 13*930 cm., under nitrogenand air saturation, at 30 and 40°C, and at the followingthree flow velocities: 2,330 cm./min., 8,300 cm./min., and16,600 cm./min.

Figures 1-3-13, 1-3-16, and 1-3-17 show that the ratesof dissolution are influenced by the length of the tube. Thereaction rates, r^, rg, and r^ are proportional to p,"0'30^•j -0.33 and respectively.

This phenomenon can be interpreted in terms of aprogressive decrease of concentration gradient of reactingspecies in the immediate vicinity of the metal surface asthe solution moves from one edge of the tube to the other edge.

According to Levich^®, the diffusional mass transfer flux,against the Inside surface of a tube, is a function of thedistance along the tube in the inlet section of the masstransfer entry region. It extends for a distance of about100 diameters for laminar flow. The mass transfer region ismuch shorter in turbulent than in laminar flow. Von Shaw,

8?Reiss, and Hanratty , indicate that this length ranges from 2 diameters to 0.3 diameters as the Reynolds number varies from 3,000 to 73,000.

D. Proposed Mechanism for the Dissolution of Tin in Hydrochloric Acid Solutions On the basis of the established experimental results, tin

dissolution may now be interpreted in terms of the following mechanism:


65

OJBo

©r 4a

COoX

fT

20

10

5

1

12,000 ml. Ng SATURATED IN HGl 30°C, d = 1.27 cm.A - V = 16,600 cm./min,• - V = 8,300 cm./min.o - V = 2,330 cm./min.

'D*

I . I I X3 3 10 20

LENGTH - cm.FIGURE 1-3-13. DISSOLUTION AS FUNCTION OF TUBE

LENGTH(FOR H? EVOLUTION)

30


10

CMSüCO ©

r—) OS

X

.3

3

12,000 ml. AIR SATURATED IN HCl d = 1.27 cm.

A A _ V = 16,600 cm./min.• - V = 8,300 cm./min.D - V = 2,330 cm./min.

"A.

'O,

10LENGTH - cm.

40°C

30°C

30°C

30"C

20

66

30

FIGURE 1-3-16. DISSOLUTION AS FUNCTION OF TUBELENGTH(FOR OXYGEN DEPOLARIZATION)


67

10

PI

CM •i©iHO

IN-O

12,000 ml. AIR SATURATED IN HCl 30°C, d = 1.27 cm.V = 16,600 cm./min.A - [Sn] = 1.0 X 10“^ moles/litre• - [Sn] = .33 X 10"’ moles/litre□ - [Sn] = .23 X 10”^ moles/litre

14 3

J I L_L_120 3010

LENGTH - cm.FIGURE 1-3-17. DISSOLUTION AS FUNCTION OF TUBE

LENGTH(FOR AUTOCATALYTIC REACTION)


681. The reaction scheme proposed for the hydrogen

evolution reaction is:a. Diffusion of hydrogen ions from the bulk solution to

the metal-solution interface

K — " H ib. Adsorption of hydrogen ions on the metal surface

+ Sn 3n-Ht

c. Discharge of adsorbed hydrogen ions by virtue of electron transfer through the metal:

(cl) Electrons transferred from a neighbouring vacant site on the metal surface

H+ H HI I, I + + I—Sn—Sn—Sn—Sn—Sn— ---—Sn—Sn—Sn —Sn—Sn—* # * * * *(c2) Electrons transferred directly from the surface

atoms on which adsorption occurred

Sn-HÎ a - 3n/-H_* S * 8d. Formation of molecular hydrogen from atomic hydrogen,

by one o f the following modes:

(d1)

H HI - M - !

-Sn-Sn-Sn -Sn-Sn- —* * -Xr

(d2) 2 (Sn+-H= ) — =* °(d3) Sn+-H_ 4 S&+-H^* ^

&Sn-Sn-8n -Sn-Sn-* * *

H- Sn-Hp*


69

e. Removal of hydrogen molecules from the metal surface, either by dissolution into the solution or by the formation of gas bubbles and their subsequent detachement from the metal surface.

(el) Sn-Hp --- ^ Sn + Hp* s * i,dissolved

(e2) n ( Sn-Hp ) ---8n-(H_)^ + (n - 1 )Sn* ^8 * ^ "s *

Sn-(Hp)^ --- ^ Sn + (H_)_* s *

where Sn-(Hp) represents a hydrogen gas bubble attached

on the metal surface.f. Desorption of stannous ions from the metal surface to

the metal-solution interface-, -i- „ 4- 4-Sn^

g. Diffusion of hydrated stannous ions from the metal- solution interface to the bulk solution

S n - . nHgO —

2. In aerated hydrochloric acid solution, the dissolution of tin is accelerated by two other simultaneous reactions.These are oxygen depolarization and autocatalysis.

The mechanism proposed for the oxygen depolarization can be described by the following elementary stages:

a. Diffusion of oxygen from the bulk solution to the metal-solution interface

^


70

b. A dissociative adsorption of oxygen on the metal surface

0_ + 2 Sn 'V 2 ( Sn-0^ )®c. Formation of mono-valent oxygen ions by virtue of

electron transferSn-0 --- =— 8n*-0%* s * s

d. Formation of hydroxyl radical

Sn^-0% + Sn-H+ ---=- Sn*-HO= + Sn=* S * G * 8 * 8e. Formation of hydrogen peroxide

2 ( 8n+-0H ) — 8n-HpOp + 8n%* * A

f. Reduction of hydrogen peroxide with formation of hydroxyl ions

Sn-HpOp a- Sn*-HO^ + 0H~* ^

8n*-H0= ------- 8n++ + OH"s * ^g. Desorption of stannous ions from the metal surface

to the metal-solution interface

* S 2.

h. Diffusion of hydrated stannous ions from the metal- solution interface to the bulk solution

S „ - . n H j O —

3. The following reaction scheme is proposed for the autocatalytic process:

a. An oxidation process^ caused by dissolved oxygen


71

with the formation of peroxide intermediate according to

Sn++ + Op + Sn+++ + H_0_* ^ ^ ^ ^

b. An oxidation process caused by the atomic oxygen adsorbed on the metal surface with the formation of hydroxyl

5nl + Sn-0= + Sn-H+ -=-8nl** + 8n-0H% + Sni* G * G * S * S * Sc. An oxidation process caused by the peroxide

intermediates^"^ (produced in stage a and in the oxygen depolarization reaction) according to

5n-H+ + + HpOp =- + + + 5n-0H- + H,0* G 1 2 1 * 6 2or

8n-Hl + 3nl + Sn-H-Op =- Snî^* + Sn-OH" + Sn-H_0_* S * S * 2 Ag * S * S * 2 87d. An equilibrium' between the stannous and stannic ions

is established at all times at the metal surface by virtueof electron transfer:

,, «+ + + + _ „+ +bUg + Sn^ 25 ■5 56As reported by other investigators^^'^ , the specimen

surface became coated with a grayish film during tin dissolution in acidic media. Chemical analysis^® and X-ray diffraction analysis^^ of the surface film showed that it consisted of redeposited pure tin.

E. Empirical Rate Equation for Tin DissolutionThe experimental data for the dissolution of tin in

hydrochloric acid solutions in the presence or absence of


7 2

oxygen may be summarized by the following equation:

1 dSn _ V d[Sn]A dt " 1 dt

= k°[HCg° ( v ) ' G ? ( d ) " ' ° 5 ( L ) - ' 3 0 g -

4 k2 BUi]°(v)'59(d)-'11(L)"'35(pQ )'50g-

+ k2[HC^°(v)'Gl(d)"'10(L)"'32(pQ )'50[8n]'50e"

where the dissolution rate, , is expressed in moles

of tin dissolved per square centimeter of metal surface per minute. The tin ion concentration, [Sn], is expressed in moles per litre. The values of k°, kg, and kS have been found to be 3*37x10"^, 3.05x10” , and 6.43x10"^ respectively. The detailed calculations are given in Appendix 3.

Therefore, the empirical rate equation for the dissolution of tin in HCl solutions can be established as:

1 M il _ 1 d[sn]A dt - A dt= 3.37x10"?[HUg °(v)'G7(d)"'°5(L)-'30g-

= 3.03x10"5[Hcg °(v) "59(d)"'T1(L)"'33(pQ )'30e-

= 6.43x10"4[HCg °(v)'GT(d)"'T0(L)-'32(p )'50[ga]'30Q-^§gr

in which the third term(autocatalysis) applies only after an elapsed time of 40 to 120 minutes. The beginning of noticeable autocatalysis depends on the dissolution conditions, but in general the bulk tin ion concentration must be greater than


73

0.33x10“^ moles/litre.

Range of experimental conditions:= 298 .13 '- 323.13°K

V = 1,173 - 16,600 cm./min.d = 0 .933 - 2.34 cm.L = 3 .073 - 13.930 cm.

V = 8 ,000 - 14,000 ml.

^0. = 0 — 1.0 atm.[hcï| =0.1 - 4.0 M

An example of the application of the above empirical rate equation to predict the changes in tin ion concentration in the corroding solution is given in Appendix 4.


CHAPTER VI SUMMARY AMD CONCLUSIONS

A. Comparison of Present Results with Previous WorkIn earlier studies with rotating cylindrical systems^^^

experimental data were correlated by the following empirical rate equation designed to express the changes in tin ion concentration in the corroding solution:

= 3.31x10"? ^^HCg°(v)'9Ge"3060

dt " V+ 2.88x10"5 ^^HCg°(v)'9G(Po^)'50e" ^§0^

where: n = 0.30 for 2,840<Cv<C21,230 cm./min.n = 0.90 for 21,230<Cv<:34,100 cm./min.

Rearranging the above equation gives

% 33^ = 3.31x10"? [HC^°(v)'9Ge"+ 2.88x10"5[Ëci]°(v)'98(Po )'50e" ^§5^

+ 1.23x10"li9ÇiLÎ_(v)G(p )'50(&[8^ )'50e" ^RT~(A).pu ug j

The rate equation obtained in the present investigationis very similar to the above correlation. Both rate equations show that, (a) the dissolution rate of tin is independent of the hydrochloric acid concentration, (b) in aerated acid solutions, the dissolution reactions show an autocatalytic

74


73

effect, with dissolution rate increasing with increasing tin ion concentration in corroding solution, and the autocatalytic process shows a half order dependence on tin ion concentration, (c) the rates of oxygen depolarization and autocatalytic processes are proportional to the square root of the oxygen partial pressure, and (d) there is no direct proportionality between the dissolution rate and the apparent surface area of the dissolving metal for the autocatalytic component.

The respective values of apparent activation energies appearing in these two equations fall within the same order of magnitude. It is also interesting to note that the values of activation energies in both correlations show the same important characteristic; that is, the hydrogen evolution process has a smaller value of activation energy than those of the oxygen depolarization and autocatalytic processes.

The low activation energy, along with independence on hydrogen ion concentration and other chemical factors, but significant dependence on flow velocity suggest that the hydrogen evolution component is essentially under diffusional mass transfer control.(Probably the diffusion of hydrated tin ions from the metal-solution interface to the bulk solution is a very slow step.)

Because the rates of the oxygen depolarization and autocatalytic processes are proportional to oxygen partial pressure raised to the 0.3 power, it would appear that the adsorption of oxygen on the metal surface is a dissociative


76

type, and also a slow step. Therefore, the oxygen depolarization and autocatalytic processes have a combination of physical(the diffusion of hydrated tin ions from the metal- solution interface to the bulk solution) and chemical(the dissociation of oxygen molecules on adsorption) slow steps involved in their elementary stages. In other words, the oxygen depolarization and autocatalytic processes are under mixed control.

Significant differences are shown between the effect ofvelocity on dissolution rates from cylindrical and tubularspecimens. For the rotating cylindrical system, the reactionrates are proportional to peripheral velocity raised to the0.90 - 0.98 powers.(The rate of the autocatalytic process isproportional to peripheral velocity raised to the O.3O powerfor a narrow range of lower velocities.) For the tubularflow system the reaction rates are proportional to linearvelocity of the corroding solution raised to the 0.59 - 0.67powers. This difference in magnitude of the flow velocityeffect is probably due to the significant differences in flowpatterns in these two systems. The flow pattern in therecycling flow system was well established, while that in therotating cylindrical system was made far more complex by the

65baffles incorporated into the solution container ' . The violent turbulence of the corroding solution caused by these baffles surely enhanced the effect of velocity on diffusional mass transfer. Therefore, the same apparent increase in


77

relative movement between the metal surface and corroding solution would cause a much higher turbulence (consequently a much larger effect on dissolution rate of tin) in the case of the rotating cylindrical system than in the tubular flow system.B. Comparison of Present Results with Mass Transfer

CorrelationsDiffusional mass transfer correlations for the dissolution

of solids into fluids moving through a tube of circular cross section, as discussed in Chapter III, have the following general form:

Nu = const.(Re)*(8c)B( ^ )P (1-6-1)

where the exponents m, n, and p are usually determined experimentally.

To compare some of the characteristics of the rate equation obtained in this investigation with those of the above general correlation, consider the case of any general dissolution run of tin in HCl solution. At any given instant, the following properties of the corroding solution; temperature, oxygen partial pressure, and tin ion concentration are fixed. Under these conditions, the empirical rate equation reduces to

• ilfsl = o,.(v)'67(d)-°5(j-)-.30+ C2.(v)'59(d)-'11(L)-'53+ C2.(v)'6T(d)"'10(L)"'52 (1-6-2)

where c^, c^, and c^ are constants.


78

With some further assumptions and rearrangements, the above equation can be converted to the following form:

l^ = Ku = o;.(Ee)-67(So)-67(f )-29 + C2.(Re)'59(Sc)'59( ^ )'31

+ C2.(Re)'Gl(Sc)'Gl( d ).30 (1-6-3)t I fwhere c.j, Cg, and c% are constants. (A detailed derivation

along with the necessary assumptions is given in Appendix 5.)Equation(1-6-3) is essentially similar to Equation(1-6-1).

This general agreement is an indication that some very slow diffusional mass transfer step (or steps) is involved in the dissolution processes of tin in hydrochloric acid solutions.

C. Conclusions1. The dissolution of tin in hydrochloric acid solutions

appears to occur through three parallel reaction paths. In deaerated acid solutions there is a single slow hydrogen evolution process. In aerated acid solutions there are three • simultaneous reactions; hydrogen evolution, oxygen depolarization, and an autocatalytic process. The last reaction becomes significant only after sufficient tin has been dissolved to create a tin ion concentration of about0.35x10”^ moles/litre.

The hydrogen evolution process is essentially under diffusion control, while the oxygen depolarization and autocatalytic processes are under mixed control, because they


79

have a combination of slow physical and chemical steps involved in their elementary stages.

2. On the basis of the experimental results it is reasonable to believe that, (a) the diffusion of hydrogen ions from the bulk solution to the solution-metal interface and the adsorption of hydrogen ions on the metal surface are fast steps;

fastK ------ «1

fast ,H"' + Sn Sn-H%1 s

(b) the adsorption of oxygen on the tin surface is a dissociative type of adsorption, and is probably a slow step,

tolow0? + 2 Sn . 2 Sn-0c.. * * ®

and (c) the diffusion of hydrated tin ions from the metal- solution interface to the bulk solution is a slow step.

slowSn^ + n HjO -----— Sn

3. The dissolution rate of tin in HCl solutions is not directly proportional to the inner surface area of the tin tube. It is, instead, a function of the inside diameter and length of the tin tube.

4. Formation of gray film on the tin surface was observed. Other investigators5^'^3;67^ayg reported this film to be composed of pure tin.

5. The stannous species, stannic species equilibrium:


80

8n_ + 2 Sn** o 1 1

is assumed to be established at the metal surface very quickly.


PART TWO

A P0TENTI08TATIC INVESTIGATION OF THE POLARIZATION AND CATHODIC PROTECTION OF TIN TUBES IN HYDROCHLORIC ACID SOLUTIONS

81


ABSTRACT

A horizontally mounted tubular electrolysis cell, incorporated into a closed loop flow circuit, was used for the potentiostatic investigation of the polarization and cathodic protection of tin tubes in aerated and deaerated hydrochloric acid solutions.

The counter electrode consisted of a coaxial length of 17 s.w.g. platinum wire.

Some preliminary studies on stationary tin cylinders immersed in HCl solutions were carried out to determine the magnitude of the required protection current and the appropriate potential range likely to be encountered.

For both the stationary and flow systems plots of potential“log(current density) showed well defined linear portions (Tafel lines). The values of the Tafel constants,

and 0^ , given by the slopes of these Tafel lines decreased with increasing hydrochloric acid and oxygen concentrations and flow velocities of the corroding solutions.

The corrosion current (I ) of tin in HCl solutions, as determined by the intersections of the anodic and cathodic Tafel lines, was found to be proportional to the fluid velocity raised to the 0.65 power, and also directly proportional to the square root of the oxygen concentration, but essentially independent of the hydrochloric acid concentration.

82


83The relationships between imposed cathodic potential on

the tin tubes, protection current densities, and dissolution rates of tin were presented in the form of |snj~electrode potential-protecticn current density plots.

The experimental results showed that cathodic protection of tin can be achieved in both aerated and deaerated HCl solutions. The protection current requirements increased with increasing oxygen concentrations, flow velocities, acid concentrations, and temperatures of the corroding solutions.


CHAPTER I INTRODUCTION

The application of cathodic protection is encountered in many industrial and civic facilities such as petroleum, gas, and water pipelines, hulls of ships, chemical storage tanks, boilers, heat exchangers, cooling towers, drying cylinders in paper-making machines, and other equipment.

Cathodic protection is perhaps the most effective of all approaches to corrosion abatement. By means of an externally applied electric current, the corrosion rate is reduced virtually to a negligible value, thus a buried or immersed metallic structure can be maintained in a corrosive environment for an indefinite time without deterioration.

This technique is usually concerned with the protection of ferrous materials, because these constitute the bulk of the objects that are used in industrial situations. As a result of this general dependence on iron containing metals, most of the earlier works, on the applicability of cathodic protection to control corrosion, have been done with ferrous materials in a variety of systems. A survey of the literature indicates that little attention has been paid to the study of cathodic protection of tin in electrolytic solutions.

In the application of cathodic protection, either in industrial or research projects, the coupled anode is usually connected externally to the system which is under protection.

84


85

There are two important disadvantages to this type of connection. The first results from the uneven protection- current distribution on the metal surface of the protected structure. As a result the most accessible parts of the structure are made more cathodic than necessary, while the current density on the metal surface of the least accessible parts is insufficient to offer protection. The second disadvantage is due to the inability of the protection current to enter electrically screened areas such as interiors of water condenser tubes and inner pipe surfaces of double pipe heat exchangers.

This study was conducted to investigate the applicability of cathodic protection to control corrosion of tin in HCl solutions under controlled flow patterns. This was accomplished through the following program:

1. Experimental examination of the behaviour of tin samples during anodic and cathodic polarization.

2. Experimental measurements to determine the dissolution rate of tin as a function of imposed cathodic potential.

3. Experimental measurements to determine the dependence of protection-current density on temperature, corroding solution velocity, acid concentration, and oxygen partial pressure,

An additional purpose of this study was to seek a way of obviating the difficulties presented by uneven protection- current distribution and electrically screened areas. To achieve this objective, a coaxial length of platinum wire


86

was used as the counter electrode(I.e. an internally coupled anode was employed in the specially constructed tubular electrolysis cell).


CHAPTER II LITERATURE REVIEW

A. Electrochemical Theory of CorrosionIn aqueous media, metallic corrosion is the result of

basic electrochemical reactions which can be separated into two or more partial reactions. These partial reactions are divided into two classes: anodic reactions and cathodic reactions.

The electrode at which oxidation reactions occur is called the anode An example of anodic reaction is:

Sn ------» Sn+ + 2e“ (II-2-1)This reaction is the basis of the corrosion of tin.The electrode at which chemical reduction occurs is

85called the cathode . One of the commonest cathodic processes during corrosion is that of hydrogen evolution:

2 H* + 2e" ------ » Eg (II-2-2)To view the over-all reaction, one combines the oxidation

and reduction reactions.

Sn + 2 H+------ » Sn/ * + Hg (II-2-3)1. Thermodynamics of Electrochemical Corrosion Processes

In any electrochemical corrosion reaction, thedriving force is the potential difference which exists betweenthe anode and cathode.

For these electrode potential calculations, the concept87


88

of standard half-cell potentials has been developed®^.Arbitrarily, the hydrogen electrode has been universallychosen as the standard reference electrode, and assigned astandard potential of zero volts at all temperatures. Thestandard potential of an electrode, E°, is that potentialwhich exists between the electrode and a standard hydrogenelectrode when all reactants and products are at unit activity.If the activity of any of the products or reactants is notunity, a new electrode potential can be calculated from theNernst equation:

E = (II-2r4)#red.where :

E° = standard electrode potential, voltsE = electrode potential, voltsR = gas constant, cal,/g.mole-°KT = absolute temperature,F = Faraday,coulomb/g.equiv,n = number of electrons transferred in the

reactionCLq- - activity of oxidized species^red “ activity of reduced species

The potential of a reaction is related to its free energy bv28,73." AG = -nFE (II-2-5)

where A G is the change in free energy for the reaction.A negative value for the change in free energy corresponds to a spontaneous r e a c t i o n ' " ^ I t can be seen also from


89Equation(II-2“5) that a reaction which has a positive potential will proceed as written.

The application of thermodynamic principles to corrosion studies has been further generalized by means of potential-pH plots. Pourbaix and his co-workers-^^ have calculated thephases at equilibrium for metal-water,and for other systems at 25°C, from the chemical potentials of the species concerned in the equilibrium, and have expressed the data in the form of equilibrium diagrams having pH and potential, E, as ordinates.

The main uses of the potential-pH diagrams are (1)predicting the spontaneous direction of reactions, (2)estimating the composition of corrosion products, and (3)predicting environmental changes which will prevent or reduce

75corrosion attack .The potential-pH diagrams for tin have been constructed

by Deltombe, Zoubov, Vanleugenhaghe, and Pourbaix^^. These diagrams predict a domain of passivation in moderately acid, neutral and slightly alkaline solutions, in the pH range between 3*3 to 9,0, This passivation would correspond to the metal being covered with a protective film of stannic oxide in the absence of substances capable of forming soluble complexes with tin or insoluble compounds. On the other hand, acid solutions and moderately alkaline soD.utions( pl-l value between about 0 and 12) are passivated if they contain oxidizing agents capable of raising the potential to about +0.2v in acid media to -0.7v in alkaline media. From these


90

diagrams it also appears to be possible to bring about thecathodic protection of tin by lowering its potential to valuesincluded between two well defined limits, which are about-0.3 and 1.0 v in very acid solutions, and -1.1 and -1.8 v invery alkaline solutions.

Potential-pH diagrams are subject to the same limitationsas any thermodynamic calculations. They represent equilibriumconditions and hence can not be used to predict the velocityof a possible reaction.

2. Kinetics of Electrochemical Corrosion ProcessesOne of the keystones of electrochemistry is Faraday's

27law which relates chemical change and electrical energy .For every equivalent of chemical reaction 96,300 coulombs must pass through the cell. This equivalence between chemical reaction and electrical charge makes it possible to write the rates for electrochemical reactions in terms of electrical currents.

The discussion of the kinetics of electrochemicalreactions is based largely on the mixed potential theory of

88electrode kinetics as stated by Wagner and Traud . The basic assumptions of this theory are (a) the kinetics of the various partial reactions can be treated separately and (b) no net current flows from an electrode which is in equilibrium or at steady state. The condition of no net current flow means the total rate of reduction must equal the total rate of oxidation

on the electrode surface if the electrode is at steady state


91or equilibrium.

The relationship between exchange-reaction rate and current density can be directly derived from Faraday's law;

oxd. - /Cred. ~ nF (II-2-6)•^oxd. /t

where and are the equilibrium oxidation andreduction rates and i^ is the exchange-current density.

When a reaction is forced away from equilibrium, the potential at which the reaction is securing changes. The amount by which the potential changes is the overvoltage,

defined as??: _ % _ g, (II-2-?)GCJ • 1

where :■n - overvoltage, volts ®eq ~ equilibrium potential, volts E. = polarized( current flowing ) potential, volts

The current applied to cause the departure from equilibrium is the net rate of reaction. Thus:

whereiapp.= ^ 2

2i = anodic current density, amp./cm.p

ig = cathodic current density, amp,/cm.2i_ = applied current density, amp./cm.^PP •

The overvoltage, exchange-current density, and the ratesof the various partial processes can be related in the form

73of a chemical rate equation2.3= i exp( — ) (II-2-9a)

o 4a


92

ioGxpC/ c

) (II-2-9b)

where and are constants (Tafel constants) of the system and may be equal to one another. For corrosion andelectrochemical studies, Equation(II-2-9) is usually written

7Q 8lin logarithmic form and called the Tafel equation .

U “ i / i o V ^ o ( n - 2- 10)

where; = activation overvoltage, volts In Figure II-2-1, polarization curves have been constructed

for a reaction whose rate is described by Equation(II-2-9), which may be applied whenever corrosion is solely electrochemical and governed by activation overvoltage28

■prHO1>

bOoJ+>rHO

>or!o•H

-PcCl>•H-P

E = Corrosion potential = Corrosion current

2

X

— ♦ 1

-.2

Log (current density)Figure II-2-1. Activation Polarization Curve for a

Reversible Electrode


9In Figure II-2-1, the zero-net-current point for a

reversible system is the equilibrium potential. In a corroding system, the no-current point is the corrosion potential

Besides activation polarization, it is possible to have concentration polarization at an electrode.

Concentration polarization occurs when one or more of the reactants are consumed at an electrode faster than they can be supplied from the bulk of the solution or when products accumulated at the electrode surface. The rate of the reaction is then limited by diffusion of ions or dissolved species toward or from the metal-solution interface. This limiting rate can be expressed as a current density( the limiting current density ij ) by the equation " ;

il = ^ AC (II-2-11)

where:?ij = limiting current density, amp./cm.

D = diffusion coefficient of reacting species, CBu/sec.

AC = concentration difference between ions in the bulk solution and that at the metal-solutioninterface.

The amount of concentration polarization overvoltage,Vg, is related to the limiting current density by the

77equation :


94

log( 1 - 1 )where: concentration polarization overvoltage, voltsA schematic plot of this equation is shown in Figure

II-2-2. It is seen that concentration polarization does notbecome effective until the net reduction current density

28approaches the limiting current density .

+

0 “N

Log (current density)Figure II-2-2. Concentration Polarization Curve

(Reduction Process)

It was observed^^ that, in practice, both activation and concentration polarization usually occur at an electrode. At low reaction rates, activation polarization usually controls, while at higher reaction rates( higher current densities ) concentration polarization becomes controlling.

For corrosion reactions it is usually sufficient to consider the measured overvoltage to be the sum of the


95

concentration and activation overvoltage*^^:

= fa + -Pc + fa (II-2-15)

The last term in Equation(II-2-13), ^ , is the resistance polarization which is merely an error produced by the potential measuring circuit. It is important only at high current densities or in high resistance solutions.

During reduction processes such as hydrogen evolution or oxygen reduction, both types of polarization are present, and the over-all equation for the overvoltage of the reduction process is formed by combining Equations (II-2-10) and (II-2-12)

foIA

= -/^%log ic/io + 2.3 log(1 - i ) (II-2-14)II

Equation(II-2-l/f) is graphically illustrated in Figure

II-2-3.

+

0Activationpolarization

Concentrationpolarization

!

log(current density)d Polarization ' centration Pola:(Reduction Process)

Figure 11-2-3» Combined Polarization Curve - Activation and Concentration Polarization


96

B. Polarization Studies of TinIn 1938 Kerr^^ Investigated the formation of oxide

films on tin during anodic polarization in sodium hydroxide solutions. He observed that the nature (such as colour, composition, etc.) and the rates of growth of the films depended on various factors, including alkali concentration, temperature, and current density. He also found that, in the presence of oxidizing agents such as chromâtes, chlorates, etc., the anodic polarization curves showed that tin became passive for any current density.

%c;In 1968 Hampson and Spencer ^ examined the anodic polarization of tin in potassium hydroxide solutions. They reported that the orientation of the anode played an important part in the observed behaviour; permanent passivity occurred when the potential rose to that required for oxygen evolution from a film of stannic oxide. In potassium hydroxide solutions of concentration higher than ?M, anodic films formed with less and less dependence on mass transfer in the electrolyte.

70In 1963 Ross and Firoiu/' studied the cathodic polarization behaviour of tin in potassium hydroxide solutions at four concentrations and three temperatures.They reported that the hydrogen overvoltage values, obtained in solutions of 0.01, 0.03, 0.10, and 1.0-N KOH, were in full agreement with Tafel's equation.. The slope of the


97

potential-log(current density) p l o t s , , was 0.12v at 23°C for all concentrations. The effect of temperature was investigated at 23°, 40°) and 35°C in 0.1N KOH solution. The hydrogen overvoltage decreased by approximately 2mv/°C, whilst the constant /5 increased approximately in direct ratio to the absolute temperature.C. Cathodic Protection of Metals

The rate of electrochemical corrosion of buried or immersed metallic structures can be reduced by cathodic polarization or by contact with an additional electrode serving as an anode relative to the corroding system. This

ppphenomenon has long been known( first noted by Davy in 1824)and serves as a basis for a wide spectrum of protectivemethods, generally known as cathodic protection. It isaccomplished (a) by cathodic polarization through theapplication of an impressed potential from an external sourcesuch as a d-c generator or a rectifier, or (b) by connectingthe metal to be protected with another metal which has amore negative( active ) electrochemical potential^

Many different theories have been advanced to explain?7the mechanism of cathodic protection. Evans" assumed that

the basic reason for cathodic protection was the formation of alkali which would produce protective hydroxide films oncathodically polarized surfaces.

36Barker et al"" have expressed opinions that to obtain complete protection it is sufficient to provide a cathodic


98

current density equal to the total local current density onthe metal surface. Cathodic protection, according to theseauthors, is due to a reverse electrolysis of the dissolvingmetal. Later, similar points of view on the mechanisms of

57 58cathodic protection were expressed by Hears and BrownStender, Artamonov, and Bogoyavlenskii^^ suggested an

entirely different approach* They proposed that cathodic protection could be explained by the fact that the atomic hydrogen evolved on the protected surface during cathodic polarization completely tied up the oxygen diffusing to the corroding surface. This retardation process would limit the access of the corroding surface to the oxygen essential for depolarization. However, this does not completely explain the mechanism of cathodic protection. For instance, it is well known that cathodic protection can be attained in the absence of oxygen and also in acid media where the oxygen supply is not the controlling factor of corrosion.

In general, the principles of cathodic protection maybe explained by considering the corrosion of a typical metalM in an acid environment. The electrochemical reactionssecuring are the dissolution of the metal and the evolutionof hydrogen gas:

M ---- — - + ne" (II-2-13)

2 h '’ + 2e"--------------------- (II-2-16)

Cathodic protection is achieved by supplying electrons to the metal structure to be protected. Examination of


99

Equations(II-2-15) and (II-2-16) indicates that the addition of electrons to the metal will tend to suppress metal dissolution and increase the rate of hydrogen evolution. If current is considered to flow from (+) to (-), as in conventional electrical theory, then a structure is protected if current enters it from the electrolyte. Figure II-2-4 illustrates cathodic protection by supplying an external current to the corroding metal on the surface of which local action cells operate. Current leaves the auxiliary anode and enters both the cathodic and anodic areas of the corroding surface, returning to the source of d-c current B. Thus the corroding metal will be made more negative by electrons flowing to it. These electrons will attract the positive ions and thus reduce the tendency for these ions to go into solution (i.e. corrosion rate will be reduced).

Electrolyte

+

AuxiliaryAnode

CorrodingMetalFigure II-2-4. Cathodic Protection by Superposition ofImpressed Current on Local Action Current


100

There are three mechanisms which cause reduction ofoxcorrosion when cathodic protection is applied'^:

1. The potential of the metal is lowered so that cathodic processes actually occur on all surface areas of the metal, i.e. M — > + ne” is prevented.

2. The electrolyte adjacent to the metal surface becomes more alkaline owing to the reduction of oxygen or hydrogen ions, and for ferrous metals this increase in pH will cause inhibition of corrosion.

3. The increase in pH will cause the precipitationof insoluble salts, i.e. GaCO^ and Mg(OH)2 > which may deposit on the metal surface and produce a protective film.

D. Current Density Requirements for Cathodic ProtectionThe current density required to protect a metal surface

is usually determined empirically, because it depends on a large number of factors. According to Ctern^^, Shreir^^, Morgan^^, and Tomashov^^, the current requirement depends to a large extent on the following environmental factors;

1. Oxygen accessibility. This factor is associated with oxygen concentration in the electrolyte and degree of turbulence. The latter reduces the thickness of the diffusion layer through which the oxygen must diffuse. In highly turbulent oxygenated liquids substantially higher current densities are required to protect an uncoated metal surface.

2. The nature of the electrolyte. In the presence of high concentrations of cathode depolarizers very large current


101

densities may be required to produce protection. The most important instance is that of acid solutions in which hydrogen ions can be discharged.

3. The temperature of electrolyte. Because an increase in temperature of the electrolyte is associated with a fall in hydrogen overvoltage, increase in depolarization, decrease in viscosity of the solution, and destruction of protective film, the current requirement is usually increased with increase in temperature.

E. OverprotectionCurrent used in excess of that required does no good, and

may do harm to metals. In 1931 Akimov^ observed that in acase of extraordinarily high protection current in the cathodicprotection of Duralumin in a 3% solution of NaCl, instead ofthe expected increase in the protective action there occurredjust the reverse. This overprotection phenomenon was lateralso observed by Slomyanskaya^^ with stainless steel in sea

AOwater. Muller pointed out that the rate of dissolution of chromium in acid solutions can be increased by cathodic polarization. Kabanov and Kokoulina^\in 1938, observed that the cathodic protection of iron and stainless steel by high current densities in nitric acid solutions can considerably increase their rates of dissolution.

It is generally accepted that the negative protective effect( overprotection ) is caused by the destruction of the protective film and reactivation of the metal by one of the


102

following destructive actions:1. Chemical dissolution of the protective oxide films by

the alkali formed at the cathode during cathodic polarization( as may occur with amphoteric metals such as aluminum).

2. Cathodic reduction of the protective film( for metalswith not too high negative potential - Cu, Ni, Fe).

3. Purely mechanical destruction of the protective film by the liberation of hydrogen on the cathode.

The critical cathodic current density at which thenegative protective effect appears depends on the operatingconditions of the protected metal surface.

F. Studies on Cathodic ProtectionMost of the earlier work on the applicability of cathodic

protection to control corrosion was done with ferrous materials. In addition, the emphasis was generally focussed on special applications under very limited experimental conditions.

Recently, a systematic study of cathodic protection ofa rotating vertical Monel cylinder in a Zf% aqueous solution

21of sodium chloride was carried out by Cornet and Kappesser They observed that at -0.9?v vs S.C.E. the Monel cylinder was cathodically protected from chemical attack in aerated sea water. Constant voltage runs at this potential were made to determine the correlation of Sherwood number with Reynolds number for the diffusion of oxygen to the surface of the cylinder under limiting current conditions.


1 0 3

The anodic behaviour of tin in alkaline solutions plays an important part in the process of tin-plating and has been the subject of various studies, because over half of the tin produced goes into coating other metals, primarily steel to make the tin can. The cathodic behaviour of tin in acid media has not been paid much attention.

G, Polarization Studies Employing Potentiostatic Techniques A potentiostat, as the name implies, maintains an

electrode at a pre-set potential with respect to a reference electrode. If the potential drifts from this value, an error signal is generated, amplified, and fed to an output stage where the error is corrected by changing the current to the auxiliary electrode^^.

A schematic diagram of a potentiostatic circuit for use in the study of metallic polarization behaviour is shown in Figure II-2-5.

kV

PefAux.

mV

Figure II-2-5. Electric Circuit for Polarization Measurements


104

The metal sample is termed the working electrode(W.E.) and current is supplied to it by means of the auxiliary electrode composed of some inert material such as platinum. Current is measured by means of an ammeter A, and the potential of the working electrode is measured with respect to a reference electrode. In practice, current is increased by reducing the value of the variable resistance R.

To obtain polarization curves, the metal sample is made the cathode or anode by adjusting the potential of the sample to the desired potential which is then maintained between the working electrode and the reference electrode. This is accomplished by automatic regulation of the current which flows between the auxiliary electrode and the working electrode. After the desired potential is adjusted, the current is allowed to stabilize and noted^^. A new potential is set, and the process is repeated until the desired portion of the polarization curve is obtained.

The chief use of a potentiostat is in elucidating the kinetics and mechanism of corrosion processes. This normally involves the study of potential-current characteristics as a function of the corroding system and of time^^.


CHAPTER III EXPERIMENTAL

A. Details of the Design of the Stationary SystemThe following is a description of the equipment used

for conducting preliminary polarization measurements and cathodic protection tests of tin in stationary HCl solutions. These preliminary experiments were carried out to determine the proper potential range,and the magnitude of the required protection current likely to be encountered in order to limit the surface area of the specimens to be used in the flow system,so that the current output would fall well within the range of the potentiostat.

1. The Assembly of the Test CellThe test cell was assembled as shown in Figure

II-3-1• The cell consisted of a 1,500 ml. reaction vessel fabricated from plexiglas pipe and plate. The four sockets on the top cover were used to hold the test electrode, gas bubbler, counter electrode,, and reference electrode probe.The reference electrode probe was connected by a salt bridge of ECl to a saturated calomel reference electrode contained in a separate constant temperature water bath. A description of the preparation of the saturated calomel electrode used in this study is given in Appendix 6.

2. Test Electrode DesignThe design of the test electrode is shown in

105


106

m

AHC£5MA

ctS<DrH® (D"d %3 O oA+)ü(D

r - i(Dfn®+)dOO

4-)ü©iH©ü,©4->Ozsoü

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X)(0©H©TlO+>ü©

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• H- P

a©xsoCO

I?HHM©&gj

I I I I I Io A W Pth o td


1 0 7

Figure I1-3-2.3 . The Counter Electrode

The counter electrode was an eight inch length of 17 s.w.g. platinum wire.

4. The PotentiostatThe potentiostat used in this study was a Model

47OOM Research potentiostat produced by Magna Corporation,Santa Fe Springs, Calif, The essential features of the circuit of this potentiostat are shown in Figure II-3-3.

The principle of operation of th is potentiostat is that a desired potential is pre-set into the potential co n tro l

circuit. The potential d iffe re n c e between the reference and test electrode is then read and balanced against this p re -s e t

potential. The difference or error s ig n a l is fed into the control A m p lif ie r . The resultant control signal varies the output of the Current Supply as necessary to maintain the pre-set potential, A chopper c ir c u i t is incorporated into the Control A m p lifie r to counteract any drift and to prevent accumulation of error signal. Because of the extremely high

input impedance, essentially no current is drawn through the reference e le c tro d e , and its potential remains stable. Specifications of the Model 4700M potentiostat are given in Appendix 7.B. Details of the Design of the Flow System and the Tubular

Electrolysis Cell1. Design o f the L iq u id Recycle System and i t s Accessories


108

Copper wire Copper nut—

Teflon washer

Teflon tube

0.125

Teflon washer-

Tin bar-

0.3

Figure II-3-2. Schematic Diagram of Test Electrode


109

Potential recorder output

AAAA--Potentialamplifier Potential

meterPotential measuring circuit Potential control circuit

"PoweramplifieControl

amplifierCurrentmeter

Potential control setpointadjustment

Current recorder output__-Volt —

meterPower

ProbeReference electrode Aux cLW.E.

Figure II-3-3* Block Diagram of Model 47OOM MagnaPotentiostat


110

Details of the design of the liquid recycle system and its accessories have been given in Part One, Chapter IV.

2. Tubular Electrolysis Cell DesignThe construction of the tubular electrolysis cell

is shown in Figure II-3-4. A tin tube of half inch inside diameter and four inches in length was produced from a remoulded one inch diameter tin rod.

A pair of spigots were machined from 1,3 inch diameter polypropylene rod to different dimensions to accommodate the tin tube as shown in Figure II-3”4*

Both ends of the tin tube were recessed into the spigotsand secured firmly with plastic cement. In addition two Teflon washers, functioning as watertight seals, were placed between the end surfaces of tin specimen and the stopping edges of the spigots.

A capillary tube was machined from 0.23 inch diameterTeflon rod to the dimensions shown in Figure II-3-4* One endof the capillary was inserted tightly into a hole drilled through the wall of the tin tube, and fixed firmly with plastic cement. This system constituted the probe for the reference electrode. A salt bridge of saturated KOI solution contained in a length of Tygon tubing connected the short capillary tube to the saturated calomel reference electrode.

A portion of the outside surface of the tin tube between the two connecting polypropylene spigots was exposed for electrical contact with the control lead of the potentiostat.


111

“inb— - *

O7/

/ X

/ X

% X

X X

X

X X

X X

X X

aO 0) I—I«H CO© n5 EH &

r4r-l©Oto>>HOU

• PÜ©r-lHP(0r-4

(HOgW(0

Ü•rl-Pa©OCO

-4-I?MM©

&El


112

3. Design of the Counter ElectrodeThe construction of the counter electrode is shown

in Figure II-3-3. This electrode was a ten inch length of 17 s.w.g. platinum wire. The main body of this platinum wire projected through the centre of the tin tube when the electrolysis cell was assembled. This central location of the counter electrode ensured a uniform electrical field.

One end of the platinum wire was inserted tightly into and led through the wall of a five inch length polypropylene tube. This polypropylene tube was used as a supporting base for the platinum wire.

4. Electrical CircuitryThe electrical circuitry of the assembled

electrolytic cell is shown in Figure 11-3-6.C. Materials and Chemicals

A description of the materials and chemical reagents used in this study has been given in Part One, Chapter IV.D. Procedures for Polarization Measurements and Cathodic

Protection Tests Using a Stationary SystemThe corrosion behaviour of the stationary tin bar was

determined by potentiostatic polarization (the variation of current was examined as a function of imposed cathodic or anodic potential).

1. PreparationPrior to testing, the test cell was filled with

1,000 milliliters of test solution, placed in the constant


113

bO« -H O 'H H “H A4® 4-) E4 CO

E>-iH

■sp-1

l O

CD'CJO?44Ü®Hf4®4goA4O

kc3

Ü• H+>a®oCO

LT\I?H1-4®g,


ao•H ^ +> O ü H <D «HA•H *H O O

114

-P

H

-P

+>Q> 0) 'Zi -P O

-P

-PU -p pl ü -p o (ü H (0 ®

* s■p

-p TJ r 4 - Hcd !-i tn XI «M

-p

■H -p -p p d o

I—i o rHP k o (D


115

temperature water bath, and saturated with appropriate gas for three hours.

The test electrode(tin bar),polished manually to 3/0 emery paper smoothness, was washed initially with distilled water, then washed with acetone for degreasing, and finally rewashed with distilled water.

The power supply of the potentiostat was then turned on and allowed to warm up in the "standby" position for approximately 30 minutes prior to testing.

The test electrode and counter electrode were then placed in the test cell. The probe of the salt bridge was inserted into the test cell and the probe tip positioned approximately half inch from the test electrode surface. Precautions were taken to ensure that the probe faced the test electrode without blocking the current path from the counter to test electrode.

All electrical leads were then connected as shown in Figure II-3-3. The zeroing of the potential measuring circuit was adjusted prior to the test and approximately every 10 minutes through the test period.

2. Polarization TestingAfter the preparatory steps were completed, the

potentiostat was first adjusted to the rest potential with respect to the saturated calomel reference electrode and then switched to the "on" position. At the rest potential no current flowed between the test electrode and the counter


1 1 6

electrode,The imposed potential of the test electrode was then

changed by approximately 50 millivolt increments, allowing the current to stabilize for 5 minutes, and recording the potential and current. This process was repeated until the desired potential range was covered.

3. Cathodic Protection StudyThe cathodic protection study was conducted by

making the tin sample(test electrode) the cathode as shown in Figure II-3-3. The potential of the test electrode was adjusted to the desired potential which was then maintained between the test electrode and the reference electrode. After the desired potential was reached, a record of current versus time was made.

When the termination point (after 3 hours) of the test was reached, an appropriately sized sample of the corroding solution was withdrawn for routine Atomic Absorption Analysis.

After the test was completed, the potentiostat was switched off, all electrical leads were disconnected and the purge gas was turned off. The test cell was disassembled, cleaned, and the acid solution discarded.

E. Procedures for Polarization Measurements and Cathodic Protection Tests Using a Flow System1. Preparation

Prior to testing, hydrochloric acid solutions were made up and mixed in the make-up tank as described in Part One,


11?

Chapter IV. A proper volume of the solution was then pumped into the top supply tank. After the solution v/as saturated with the appropriate gas overnight, the temperature control system and the stirrer in the supply tank were switched on. During the time required for the acid solution to reach the desired temperature, the electrolytic cell was prepared and connected to the flow system. The power supply of the potentiostat was turned on and allowed to warm up in the "standby" position for approximately 30 minutes.

After the acid solution in the supply tank reached the desired temperature, the Dynalab Model Zf-MD pump used to deliver the corroding solutions was switched on and the flow rate v/as regulated. All electrical leads of the electrolytic cell were then connected to the potentiostat as shown in Figure II-3-6.

2. Polarization MeasurementsAfter the preparatory steps as described in the

above section were completed, the potentiostat was adjusted to the rest potential with respect to the reference electrode. The imposed potential of the test electrode was then changed by approximately 30 millivolt increments, allowing the current to stabilize for 3 minutes, and recording the potential and current. This process was repeated until the desired potential range was covered.

3. Cathodic Protection StudyAfter the preparatory steps as described in Section 1


118

were completed, the test electrode(tin tube) was made cathodic by adjusting its potential to the desired value which was then maintained between the test electrode and the reference electrode. After the desired potential was reached, a record of current versus time was made. Besides the current-time record, an appropriately sized sample of the corroding solution was withdrawn for routine Atomic Absorption Analysis at convenient intervals of time,

4. Shut-down ProcedureWhen the test was terminated, the potentiostat was

switched to the "standby" position, and all power to the circuit was turned off. All electrical leads were then disconnected. The acid solution was drained from the flow system. After the flow system was flushed with water twice, the electrolytic cell was disconnected, cleaned with distilled water, and stored.


CHAPTER IV EXPERIMENTAL RESULTS AND DISCUSSION

The results obtained from the polarization and cathodic protection studies are presented in this chapter.

For each polarization test, a separate plot of current density versus electrode potential with respect to a saturated calomel electrode was constructed. Current density values at each potential were calculated from the measured current and the exposed surface area of the electrode.

For each cathodic protection test, in addition to the current density-electrode potential plot, a separate plot of metal ion concentration in the corroding solution versus time was made for the purpose of examining the protective effect of the impressed current on the test electrode at each potential, A, Results of Stationary System Tests

1. Polarization TestsPreliminary polarization tests were conducted at

23°C by immersing stationary tin samples in aerated and deaerated solutions over the concentration range 0.0623 to2.0 M HCl. The results are illustrated by the electrode potential-log(current density) plots shown in Figures II-4-1 and II-4-2. The polarization curves shown in these Figures are of similar form. This similarity in form of the polarization curves confirmed that the electrical circuitry of the electrolytic cell, including the potentiostat,

119


120

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121

o •

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122functioned satisfactorily.

The dashed lines in Figures II-Zf-1 and 11-4-2 are the estimated Tafel lines for the anodic and cathodic processes. These linear portions of the polarization curves are well defined. The intersection of the anodic and cathodic Tafel lines gives both the corrosion current (I ) and corrosion potential (E ) as specified by theory^^.

Figure II-4-I shows that the anodic slopes of the potential-log(current density) plots ( in Tafel equation) for tin in 1.OM and 2.0M HCl solutions are O.I7 volt and 0,14 volt, respectively. The cathodic Tafel slopes, , are 0.19volt and 0.17 volt, respectively. Figure II-4-2 shows that at 25* 0, in aerated 1 .OM HGl solution, for tin is 0.12 volt,

^ clis 0.14 volt.A literature s u r v e y ^ 64»89 showed that the value of yg

for electrochemical reactions ranges between O.O9 and 0.19 volt per tenfold increase in current. However, in the presence of readily adsorbed impurities, the values ofyS may be twice the values quoted above and may also vary irregularly with current density^^*^^. Conway^^ has reported a y6 value of 0.19 volt for the hydrogen overvoltage of tin in deaerated 1.OM HCl solution at 20°C, and Potter^^ has given a y3 value of0.13 volt for hydrogen overvoltage of tin in deaerated 2.ON HgSO. solution at 20°C.2 4

When compared with the data in the literature, the presently determined values of /5 are of the right order of


123

magnitude.Figures II-4-1 and II-4-2 indicate that at lower reaction

ratesdower current densities), activation polarization appears to he controlling for both anodic and cathodic reactions, because the linear portions of the potential- log (current density) plots are well defined. At higher reaction rates(higher current densities) the effect of concentration polarization becomes significant. It is also noted that the value of current density at which the concentration polarization effect starts to show up is always lower for the cathodic reaction than that for the anodic reaction. For instance, the polarization curves obtained at 29^C, in nitrogen saturated 1.OM HCl solution(as shown in Figure II-4-1) indicate that for the cathodic reaction the effect of concentration polarization starts to show up at a

pcurrent density of 9 ma/cm. , for the anodic reaction theconcentration polarization effect starts to show up only when

2the current density reaches a value of about 20 ma/cm.

Both Figures II-4-1 and II-4-2 show that changes in the HCl concentration of the corroding solutions produce no apparent variation in the corrosion current. This behaviour is consistent with the results obtained in Part One.

Theoretically the corrosion current should be directly proportional to the dissolution rate of a metal. In Part One it has been shown experimentally that the dissolution rates of tin are essentially independent of acid concentration over the range 0.10 to 4*0 M HCl.


124

Figures II-4-I and II-4-2 also show that the values of the cathodic and anodic Tafel slopes, and , decrease withincreasing HCl concentrations in the corroding solutions. This behaviour is consistent with theory which predicts that an increase in HCl concentration generally results in a decrease in the overvoltage of metals.

2. Cathodic Protection TestsA series of preliminary cathodic protection tests

were carried out at 29°C by immersing stationary tin samples in air saturated 1.OM HCl solutions. At each test potential, the tin ion concentration in the corroding solution was determined after the test electrode had been immersed in the solution for three hours. The current corresponding to the applied potential was recorded after it became stabilized.

Figure II-4-3 shows these results in the form of a [CiJ- potential-protection current density plot. This Figure indicates clearly that the corrosion rate of tin in HCl solutions can be reduced sharply by means of an externally applied electric current.

In Figure II-4-3 the point labeled "from unprotected sample" represents the tin ion concentration obtained by immersing a stationary, unprotected( no external protection current was impressed on the specimen) sample in the corroding solution for three hours.

The main advantage of representing the experimental data according to the plot of ilgure 11-4-3 is that it reveals the


125

ao'1u®

«HOtoëw

kfl)+>«H<lj

h•p•H

$®Hi-d-oMa'co.

81000 ml. 1N HCl 25°C

From unprotectet sample

.6

4

OJ •

.2

.0

•p.8

6

4

_ 0.2

or— I— -1.5 -1.5 -0.51 . 1 -0.9 -0.7 -0.5

Potential,(volts) vs S.C.E.Figure II-if-5. Cathodic Protection Effects Related to

Protection Current Requirements for Stationary Tin Sample in Air Saturated HCl Solutions


126

optimum region of applied cathodic potential. From this Figure, it becomes easy to select the most desirable potential which, if imposed on the test electrode, will offer the most effective protection and at the same time require the least protection current. For instance, Figure II-4-5 shows that the most desirable range of applied cathodic potential is -0.8 to -O.95 volt for the conditions specified.

B. Results Obtained from Flow System1. Polarization Tests

The cathodic and anodic polarization tests were conducted at 50^0, in nitrogen, air, and oxygen saturated HCl solutions, at the following three flow velocities; 2,650 cm./min., 8,550 cm./min., and 14,500 cm./min. The results are presented by the potential-log(current density) plots shown in Figures Il-if-Zj. and II-lf-5* The general form and characteristics of these polarization curves are internally consistent and are consistent with those for stationary specimens.

It can be observed that an increase in the fluid velocity, or an increase in the oxygen concentration of the corroding solutions lowers the values of the cathodic and anodic Tafel slopes. It can also be observed that the values of corrosion currents given by the intersections of these Tafel slopes increase with an increase in the fluid velocity, or an increase in the oxygen concentration in the corroding solutions.

Figures II-Zf-6 and II-4-7 indicate that the corrosion current (I ) obtained from the intersections of the anodic


127

VD+

CIS a

/ CIS m

® ch

H PM

Ooo0

H t jS O - d CD 'd

m (D -P■ p CIS -P

s % CIS osÜ U d U

d - p do « - p CIS + >o iH CIS CO CISt A a c/a CO

F ,o CM •H CMo % < OoII 1 1 1>

f v m e o

-d"+

nj+

' a ' 0 ' 2 SA (s%%OA) 'te T ^ n a ^ o j


128

o s p;

'd «H

® «H

s a sü ü ü CD CO

\ o\ / W rH (H O H p4o o oIXN lA O VO K\ \ 'A '°T<\j co ^

Il I I II

f> > t»I I I

o CM -d-1 1 l

'a'O'S SA (sq-IOA)


129

12,000 ml. IN HCl 30°CV = 1 if,300 cm./min

a

Ü

-pa

Î0Üao

oKÜ

. 6(atm. )

.80 2 4 1.0

Figure 11-4-6. Effect of Oxygen Concentration onCorrosion Current for Tin in IN HCl Solutions


130

12,000 ml. Air Saturated IN HCl

Velocity - cm./min.Figure II-4-7. Effect of Fluid Velocity on Corrosion

Current for Tin in Air Saturated IN HCl Solutions


131

and cathodic Tafel lines plotted in Figures II-4-4 and II-Zf-3» is proportional to the fluid velocity raised to the O.64 power, and also directly proportional to the square root of the oxygen concentration in the corroding solutions. This behaviour is completely consistent with the results obtained in Part One, which showed similar proportionalities between the spontaneous dissolution rate of tin, fluid velocity, and oxygen concentration in the corroding solution.

2. Cathodic Protection TestsThe cathodic protection tests were carried out at

30°C, in nitrogen, air, and oxygen saturated HCl solutions, and at flow velocities of 2,630 and 14,300 cm./min. The results of these tests are given in Figures 11-4-8 through II-4-I1 which are presented in pairs. An explanation of Figures II-4-8A and II-4-8B is given here as an example to show how these Figures were constructed.

Figure II-4-8A shows nine plots, of tin ion concentration versus time, which represent the results of a series of nine testing runs conducted at 30^C, in air saturated 1.OM HCl solutions flowing at 14,300 cm./min. Each curve is the result of a separate run conducted by maintaining the test electrode (tin tube) at a pre-assigned potential. These plots clearly indicate that the tin tube received cathodic protection of different effectiveness at different potentials.

In the construction of Figure II-4-8B, the (Siij -electrode potential-protection current density plot, tin ion


132

9

8

0)f-i+>

m0 H1

-d-oX

0

12,000 ml. 1N HCl, 30°C V = 14,300 cm./min.AO□▲■AOQ

From unprotected sample -O.3 volts vs S.C.E. -1.0 volts - -1.2 volts - -1.1 volts - -0.8 volts - -O.9 volts - -0.7 volts - -0.6 volts -

EEEEEEEE

■ A '

0_ * •'

0 a J q Aq A O O G Q^nO.^E! o_L _L 1 _L JL _L

0 100 200 300 600 700300 400TIME - min.

Figure II-4-8A. Cathodic Protection of Tin in High Velocity,Air Saturated HCl Solutions


flo•H+>rsfHOCQto

(MOCO

§wOO

o•p«Hcîj

+>

0)rH

a

i)X

to.

912,000 m l. IN HCl, 30"C V = 14,300 cm ./m in.

8 30

Fromunprotectedsample7

N '6 20

■p4

3

2

0— .8 — .61 .2 1 . 0 .2

Potential, (volts) vs S.C.E.Figure II-4-8B. Cathodic Protection Effects Related

to Protection Current Requirements for Tin in High Velocity, Air Saturated HCl Solutions


1342.0

% 1.2-p

0.8

12,000 ml. 1N HCl, 30°G V = 14,300 cm./min.

□ - E = -Ô.3 volts vs S.C.E.m - E = -O.9 volts -A - E = -0.6 volts -

• E = -0.7 volts -

200 300 400 300TIME - min.

Figure II-4-9A. Cathodic Protection of Tin in HighVelocity, Oxygen Saturated HCl Solutions

700


135

oo•H+>73tHO(ûCQa«M01W00Q)•P«H4

0)ÎH+»

O I—fo8

12,000 ml. IN HCl V = 14,300 cm/min 30°C

8Fromunprotected“ sample

. 6

4

2

0

20.8

6

.2

2001--- 1 L.1.4 -1.2 -.8 —. 6 -.21 .0 - . 4

fvj

Potential, (volts) vs S.C.E Figure II-4-9B. Cathodic Protection Effects Related

to Protection Current Requirements for Tin in High Velocity, Oxygen Saturated HCl Solutions

80

1+>«0)«aiodo•H+>o(D•PO


136

1.8

1.6

1.4

1 .20)

Ha 1 .0Q)Ha

X

(8 .6

.4

0

12,000 ml. IN HCl, 30"cV = 14,300 cm./min.A - From unprotected sample O - S = -O.3 volts vs S.C.E.■ - E = -0.6 volts vs S.C.E,• - E = -O.7 volts vs S.C.E.A - E = -0.9 volts vs S.C.E.□ - E = -1.0 volts vs S.C.E.

L

/

'6'h^a Hn®^* ®D^ o Aq',®B A ■ ■ A ® o ■1 _L JL

0 100 200 300 400 300 600Time - min.

Figure II-4-10A. Cathodic Protection of Tin in HighVelocity, Nitrogen Saturated HCl Solutions

700


137

12,000 ml. IN HCl V = 1 4 ,3 0 0 c m ./m in

30°C

Fromunprotectedsample

20

+>

0)

—, 8-1 .2 - 1 .0 - . 6

fvj80

1>5+>

«A+>d(D(4odo•H+>o0)+>o&

Potential, (volts) vs S.G.E.Figure II-4-IOB. Cathodic Protection Effects Related to

Protection Current Requirements for Tin in High Velocity, Ng Saturated HCl Solutions


1 8

12,000 ml. IN HCl, 30"C V = 2,650 cm./min.^ - From unprotected sample O - E = “0.5 volts vs S.G.E.0 “ E = -1.1 volts vs -A ». E = -0.6 volts vs -# - E = -1.0 volts vs -

200 300 400TIME - min.

Figure II-4-11A. Cathodic Protection of Tin in LowVelocity, Air Saturated HGl Solutions


U 9

0ors'ÉCQtoAo

3.0

12,000 ml. IN HCl V = 2,650 cm./min. 30°G

Fromunprotectedsample

25

- 1.4 - 1.2 -.2

ovj

—1.0 —0.8 —.6 —.4Potential, (volts) vs S.C.E.

Figure II-4-11B. Cathodic Protection Effects Relatedto Protection Current Requirements for Tin in Low Velocity, Air Saturated HCl Solutions

a

II

dAdiodo•H

-PO(1)

•PO


140

concentrations corresponding to 48O minutes of exposure to the corroding solution at the nine different electrode potentials were adopted from Figure II-4-8A. The time of 48O minutes at which the values of tin ion concentrations were taken is purely arbitrary. Figure II-4-8B will show the same general characteristics if tin ion concentrations are taken at any other exposure time. The protection current densities were calculated from the stabilized currents recorded at each potential.

The general form and shape of Figure II-4-8B is consistent with Figure II-4-3. The point labeled "from unprotected sample" in Figure II-4-8B represents the tin ion concentration obtained from a spontaneous dissolution run which was carried out under a condition of no external protection current being impressed on the specimen.

An examination of Figures II-4-8 to II-4-I1 reveals the following characteristics:

(a) Cathodic protection of tin can be achieved in both aerated and deaerated hydrochloric acid solutions.

(b) The effect of overprotection appears in certain situations. For instance, Figure II-4-8B shows that, at 30°C, in air saturated 1.ON HCl solution, at a velocity of 14,300 cm./min., when the impressed protection current density

pexceeds 20 ma/cm., instead of the expected decrease in the dissolution rate of tin there may be just the reverse, an increase in the dissolution rate.


141

This damaging overprotection effect was probably caused by (1) excess alkalies generated at the metal surface. Since tin is one of the amphoteric metals, excess alkalies can damage tin by causing increased attack rather than reduction of corrosion, and (2) mechanical destruction of the protective film and disturbance of diffusional mass transfer layer by the liberation of hydrogen on the electrode. Because the overprotection effect was more significant in the aerated systems, the first reason appears to be the more damaging factor.

(c) The range of electrode potential within which the metal can be satisfactorily protected becomes narrower when the fluid velocity or the oxygen concentration of the corroding solution increases.

3. Protection Current Dependence on TemperatureThe protection current dependence on temperature was

studied over the range 23 to 33°C, in air saturated 1,0M HCl solutions, at three different velocities and two different potentials. The results presented in Figure II-4-12 show that the protection current density increases with increasing temperature of the corroding solutions at a rate of approximately .12 ma/cm?-°K.

An increase in temperature of electrolyte is usually associated with a fall in hydrogen overvoltage, decrease in viscosity of the solution, and destruction of protective films. These are probably the factors which contribute to the increase


142

32

28

nj.ao

aQ+>asoAo•H+>o®

■Po&

24

20

16

12

8

12,000 ml. Air Saturated IN HCl ■ - E = -0.8 volts vs S.C.E., V□ - E = -0.6 volts vs S.C.E., VO - E = -0.6 volts vs S.C.E., V

14.300 cm./min.14.300 cm./min. 8,330 cm./min.

I0293 300 303 310 313 320 323

TEMPERATURE - Figure II-4-I2. Protection Current Dependence on

Temperature

330


143

HOWk:ot—Td0+J M H H

o o O• < •+> 03 03 03(003 ra CQ tol> i> >P t> t> >CO 1>- v£>

« o O OH 1 1 1a II II IIoo M w wo o•V O 1 1 1(VI o

01— < □

<

oo ■cf(V i

orvj KO (V i

X Pi

1

o(V I

oo

('mo/Tsm) *jÇq.Tsu8G q.U8JjnQ noT^oe^oj^


144

in the protection current,4, Protection Current Dependence on Fluid Velocity

The investigation of the protection current dependence on fluid velocity was carried out at 30^0, in air saturated 1,0M HGl solutions, and at three different electrode potentials.The fluid velocity was varied over the range 1,130 to 14,300 cm./min.

The results, as shown in Figure II~4“13> indicate that the protection current increases with increasing flow velocity of corroding solution. This behaviour can be interpreted by the fact that a higher solution velocity can cause a higher degree of turbulence which reduces the thickness of the diffusional mass transfer layer through which the reactants must diffuse. Most previous investigators observed that in highlyturbulent fluids substantially higher currents were required to protect metal surfaces.

3. Protection Current Dependence on Oxygen concentration Tlie protection current dependence on oxygen

concentration in the corroding solution was studied at 30°C, in 1 .CM HCl solutions, and at tv/o different fluid velocities.The results are presented in Figure 11-4-14"

It was found that the protection current increased continuously with increasing oxygen concentration in theacid solutions. The same effect was observed by most other

59 85 •corrosion researchers^^'Because oxygen is a strong depolarizer, its accessibility


rg •ao

-p'Éfl«+>Bh5BO•rl•PO<D

•P

6

12,000 ml. IN HCl 30°CE = -0,6 volts vs S.C.E o - V = 14,200 cm./min. • - V = 2,650 cm./min.

atm.Figure II-4-I4. Protection Current Dependence

on Oxygen Concentration


146

au

-p

a+>«

ono•H-PO0+>&

12,000 ml. Air Saturated 30°C, E = -0.6volts 78.S.C

V = 14,300 cm./min.□ - V = 9,100 cm./min.A - V = 8,350 cm./min.- V = 2,630 cm./min.

.2 .4 .6 .8[hcÎ] - molar

Figure II-4-I3 . Protection Current Dependence on HGlConcentration


147

to the metal surface is a very important factor in determining the density of protection current.

6. Protection Current Dependence on HCl Concentration The protection current dependence on the HCl

concentration in corroding solution was investigated at 30°C, in air saturated solutions, and at four different fluid velocities. The concentration of the HCl solutions was varied over the range 0.0623 to 1.OM, The results, as plotted in Figure II-4-I3, illustrate that the protection current initially increases sharply with increasing HCl concentrations. The rate of increase of the protection current decreases when the acid concentrations become higher than about 0.3M.

This phenomenon has been observed by many other investigators^^'7^*^^. They pointed out that in acid solutions in which hydrogen ions can be discharged, the presence of high concentrations of such acids generally creates a much larger current requirement for protection of metal surfaces.


CHAPTER V SUMMARY AND CONCLUSIONS

An examination of the experimental results obtained from the present polarization and cathodic protection studies revealed the following characteristics:

1. The polarization curves (potential-log(current density) plots) are of similar form. They are internally consistent.At lower reaction rates (lower current densities), activation polarization appears to be controlling for both anodic and cathodic reactions, because the linear portions (Tafel lines) of the potential-log(current density) plots are well defined.At higher reaction rates (higher current densities) the effect of concentration polarization becomes significant. It is also noted that the value of current density at which the concentration polarization effects starts to show up is always lower for the cathodic reaction than that for the anodic reaction. For instance, the polarization curves obtained at 30°C, in air saturated 1.OM HCl solution flowing at 14,300 cm./min.(as shown in Figure II-4-3) indicate that for the cathodic reactions the effect of concentration polarization

pstarts to show up at a current density of about 9 ma/cm.; for the anodic reaction the concentration polarization effect starts to show up only when the current density reaches a value of 20 ma/cmf

148


149

2. The values of Tafel constants, /^'s and /5’'s ,(0.12 - 0.32 volts as given by the slopes of the Tafel lines) are of the right order of magnitude when compared with the data reported by other i n v e s t i g a t o r s ^ T h e present results show that increasing HCl concentrations, flow velocities, and oxygen concentrations of the corroding solutions lower the values of Tafel constants. This behaviour is an indication that the activation overvoltages for both anodic and cathodic reactions occurring at the tin electrode decrease with increasing HCl concentrations, flow velocities, and oxygen concentrations of the corroding solutions.

3. The corrosion current 1^, as indicated by the intersections of the anodic and cathodic Tafel lines, is found to be proportional to the fluid velocity raised to the 0.64 power, and also directly proportional to the square root of the oxygen concentration; but essentially independent of the HCl concentration. This behaviour is consistent with the results obtained in Part One, which show similar proportionalities between the spontaneous dissolution rates of tin, HCl concentrations, flow velocities, and oxygen concentrations of the corroding solutions.

4. The cathodic protection of tin can be accomplished in both aerated and deaerated hydrochloric acid solutions. To achieve the most effective cathodic protection (maximum protective effect with lowest protection current), optimum regions of applied cathodic potentials under different


130

conditions have been determined; for the stationary system, the most effective range of applied potential is -0.8 to -0.93 volt, for the flow system the effective range is generally -0.6 to -0.9 volt. The range of applied cathodic potentials within which tin can be satisfactorily protected becomes narrower when the flow velocity or oxygen concentration of the corroding solution increases.

3. Negative effects of overprotection were observed in the flow system with air and oxygen saturated solutions. This damaging overprotection effect was probably caused by (1) excess alkalies generated at the metal surface, and (2) mechanical destruction of the protective film and disturbance of the diffusion layer by the liberation of hydrogen on the cathode. Since tin is one of the amphoteric metals, excess alkalies can damage tin by causing increased attack rather than reduction of corrosion. Because the overprotection effect was more significant in the aerated systems, the first reason appears to be the more damaging factor.

6. The protection current density required to protect the tin tube increases with increasing temperature at a rate of approximately 0.12 ma/cm?-°K, and it also increases with increasing hydrochloric acid concentration, flow velocity, and oxygen concentration of the corroding solution.


APPENDIX 1 CALIBRATION OF ROTAMETERS

Two rotameters were used to measure the full range of flow rates of corroding solutions.

1. 12-1000 Brooks Glass Tube Rotameter(Tube;R-12M-23-3)Flow capacity: 1.3 - 13 USGPMAccuracy: + 2% of maximum flowTube: Borosilicate glassPacking: TeflonFloat: Hasteloy B

2, 10-1000 Brooks Glass Tube Rotameter(Tube:R-1OM-23-1)Flow capacity: .4 - 4 USGPMAccuracy: _+ 2% of maximum flowTube: Borosilicate glassPacking: TeflonFloat Hasteloy B

The rotameters were calibrated by maintaining a particular reading and measuring the volume of effluent collected in a known time interval. The calibration curves are shown in Figures A-1-1 and A-1-2.

131


o O Oo O oo lA o.CM .lA

< 4

HOM

1^2

oCM

OO

M

I I IOIT\CM

oOCM

J 1 L OlAJ I I L OO

I I I L J I I L

“ X>

“M3

CM

OlA

O

IlACM

ÈIKO,Q

'— /OOOToo uCD+iCD

§ -P O

CD W -P•H A H O(H

(D•P

:Is

tûCD

tK O

gO0

■H

1

4îa'I


w

\, <\J lf\

1 I I I I I I I J I I I I I L J I I I I I L

PlAI«<D

êo8I

O OJlA —

Oko

Q)-P

RoaA0)13

<D -p -P O•HiH

O M fA Eh

OOJ

OAOJ

oA Oo oA

O o

(una) ‘SuxpeaH Ja^ame^os

O

§I 4H)%g

IICjo.r

a)


APPENDIX 2 ANALYTICAL METHODS AND CALIBRATION

A. InstrumentationA Jarrell-Ash Model 82-526 Maximum Versatility Atomic

Absorption Spectrophotometer was used for the analysis of tin ion concentrations in the corroding solutions. This instrument features a ^00 mm focal length Ebert Mount Monochrometer with two interchangeable gratings with 1,180 grooves/mm. A choice of three burners is available; a total consumption Hetco burner, a 5 cm.and a 10 cm. slot laminar flow burner. The instrument can simultaneously accommodate two fuels and two oxidants.

B. Analytical MethodsThe analysis of tin by the Atomic Absorption Spectro-

5 2photometric method has been developed by Allan^, Agazzi ,17Amas and Willis , Capacho-Delgado and Manning , Gatehouse

and Willis^^, Gibson and Grossman^^, and Vollmer^^.

C. Preparation of Standard SolutionA stock of 2x10 " M tin ion solution was made by dissolving

Analar Grade tin bars in 1M HCl. Standard tin ion solutions were prepared by diluting aliquot samples of the stock solution with 1M HCl.

A new calibration curve was drawn each time a series of samples was to be analyzed. The calibration curves were made

154


155

by plotting tin ion concentration against absorbance on linear scales, A typical calibration curve is shown in Figure A-2-1.

D. Analytical Conditions and Instrument SettingsThe 10 cm. laminar flow burner was used in this analytical

work. The fuel and oxidant used were hydrogen and air, respectively.

Following are the instrument settings which were used in the analysis of tin:

Wave length:Hollow cathode current:High voltage:Hg tank regulator pressure:Air tank regulator pressure:H^ flow:Air flow: Burner height: Damping:

2246 212 ma0 .6 5 kv

20 psi 30 psi 50 SCFH 17.5 5CFE 10 mm minimum


1 %

o\

oo

tf\

-d*

oVOO C OOJ

fa

I«o•H+>PiPOCQ

g

ISu o«M(DP oao

urQ •H rH

< Odo

ICM

<kÜ)&id

H01

(aJ4TT/s@T0m) <^oi % [*2]


APPENDIX 3CALCULATIONS OF RATE CONSTANTS FOR TIN DISSOLUTION

The following equations;

kg = ( *5 (L)*^^(PQ^)"*^e ' RT

have been used for the determination of k^, kg, and k^ to give the average values of these constants, with an average deviation of _+ 4 per cent as shown in Tables A-3-1, A-3-2, and

157


158

TABLE A-3-1 Evaluation of Velocity Constant

(k® for hydrogen evolution)/ 1 dSn \ A ïït" M V d L T k?

/ moles \ / cm. \( cm. ) ( cm. ) ( °K )cm.-nun min.

3.31x10"G 4,700 1.905 7.620 298.15 3.27x10-73.97 - 5,730 1.270 10.160 303.15 3.30 -5.81 - 10,280 — — — — — — 3.55 -2.11 - 2,350 1.902 7.620 298.15 3.42 -1.35 - 1,175 " — — — — — 3.31 -7.20 - 14,300 1.270 10.160 303.15 3.57 -7.93 - 16,600 — — - - — — 3.37 -8.89 - — — — - - 308.15 3.36 -9.93 - — — — — 313.15 3.29 -1.23x10-7 323.15 3.35 -7.21x10-8 — "" — — 13.970 303.15 3.37 -7.42 - — — — — 12.700 •••— 3.39 -8.25 - — — 8.900 — — 3.45 -9.76 - " — 5.075 3.38 -4.03 - 5,730 0.953 10.160 — — 3.17 -3.93 - — — 1.588 — — 3.43 -3.84 - — 2.540 *- — — — 3.32 -3.22 - 4,150 1.270 — — — 3.19 -4.87 - 8,500 1.905 7.620 298.15 3.42 -5.51 - 10,280 — — — — — — 3.47 -

Average = 3 •37x10**'


159

TABLE A-3-2 Evaluation of Velocity Constant ( kg for Og depolarization )

f moles \( Y 7 ^cm.-min

V

cm.min. )( cm.) ( cm.)

°2 (atm. )

T

(°K)

-7

-8

1.10x101.37 - 1.95 - 7.29x10 4.80 - 2.38x10"? 2.61 -

3.01 -3.47 - 4.36 - 2.33 - 2.42 -

2.72 - 3.25 -1 .42 -

1.35 -1.28 — 1.75 -2.07 -2.47 -

4,7003,73010,2802,3301,17314,30016,600

3,730

4,130

1.9051.270

1.902

1.270

7.6210.16

7.62

10.16

0.9331.3882.5401.270

13.9712.708.905.0810.16

0.21 298.15303.15

298.15

303.13

308.13313.13323.13303.13

0.500.701.00

2.93x10"53.20 - 3.04 -3.11 -3.13 - 3.10 -2 . 9 7 -

3 .1 4 -2.98 - 3.01 - 3.09 -3.12 -2.96 -3.13 - 3.00 -2.95 -2.97 -3.06 - 2.90 -3.15 -

Average = 3*05x10-3


160

TABLE A-3-3Evaluation of Velocity Constant (k^ for autocatalytic reaction)

/ 1 dSn \( %/ moles ^

V

/ cm. \d

( cm. )

L

( cm. ) ( atm.)

Bn]/moles\

T

(°K)> 3

2 . cm.-min min. \litre/

2.15x10“? 16,600 1.905 7.62 0.21 3 .20x10-4 298.13 6.40x10-42.40 - — — "" — - - 4.00 - — — 6.33 -2.72 - mm mm «- — ». ». 3.13 - — — 6.51 -5 .18 - — •» — -------- — — 7.10 - — — 6.33 -5.61 - •M — > - - — " — — 9.20 - 6.30 -1.18 - 8,300 -------- 2.00 - 6.22 —1.70 - mm mm — — — - — — * 4.43 - 6.14 -

2.20 - mm mm #" mm mm mm — »• 7.20 - ». ». 6.65 -

7 .80x10-8 4,700 mm mm mm mm - - 1.33 - “ — 6.37 -

1.15x10"? — ».mm mm mm - - 2.95 -

- - 6.70 -

1.41 -— •-

— mm mm 4.73 - “ — 6.20 -

1.33 -— — — -• — — — — — — 303.13 6.63 -

3 » 90 — 1,173 1.270 10.16 ». ». 1.70 - — — 6.42 -

6.02 - 2,330 mm mm ». ». ». »“ ». ». mm mm 6 .46 -

1.06x10"8 3,700 mm mm ». ». — ». — — mm mm 6.61 -■1.35 - 10,280 mm mm mm mm — ». ». ». mm mm 6.43 -4 .10x10"? 13,800 — — ». ». ». 1.20x 10"^ mm mm 6.37 -6.20 - — —

». ». "" 0.50 — — ». — 6.27 -8.51 - — — •» ». — ». 1.00 ». ». — ». 6.39 -4.33 - 16,600 0.933 — W 0.21 — — — — 6.30 -4.05 - ». — 1.388 - « 6.33 -3 .83 - — 2.540 — — — — •* — - - 6.63 -3.25 - — — 1.270 5.08 •* — 6.80x10"4 — — 6.71 -2.68 - — — •* »i 8.90 •" — - .» — 6.14 ~2.41 - «. ». 12.70 — — —' — — — 6.65 -2.32 - ». ». 13.97 — — ». ». — — 6.33 -

Average = 6.45x10“^


APPENDIX 4An Example of the Application of the Empirical Rate Equation

The application of the empirical rate equation to theprediction of changes in tin ion concentration in corrodingsolutions may be illustrated by considering the followingdissolution conditions:

Solution volume = 12 litersFluid velocity = 14,150 cm./min.Temperature = 50°CP^ =0.21 atm.^2&Cl| = 1.0MInside diameter of tin tube = 1.27 cm.Length of tin tube = 10.16cm.

Substitution of these values into the empirical rate equation gives

1 dSn V d [Sn]A dt - A dt

= 7.148 X 10"8 + 2.365 X 10"? + 2.802 x 10"G[Sn]'5Rearranging gives (A-4-1)

Ÿ (7.148xlO"8 + 2.365x10"? + 2.802x10"6 Sn "5)= 1.041 X 10"G + 9.45 X 10"^[Sn]' (A-4-2)

henced[Sn] = (1.041x10"G + 9.43x10"G[3n]'5 )dt (A-4-3)

161


162

andAgiQ = d.OlflxlO”^ + 9.45x10"^[Sn]*^ )At (A-4-4)

To calculate the tin ion concentration for the initial period of dissolution, only the first term on the right hand side of Equation(A-4-4) is applied. When the tin ion concentration thus calculated exceeds the value of .35x10”^ moles/liter, the second term on the right hand side of Equation(A-if-2f) is also applied. By employing appropriate numerical techniques a dissolution curve can be predicted as shown in Figure A-4-I, where an experimental dissolution curve is also drawn for comparison. Figure A-4-1 shows that the rate equation works well even for quite large time increments (about 100 minutes). The accuracy of the prediction can be improved by decreasing the time increments used in the numerical calculations.


163

1.4

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12,000 ml. Air Saturated IN HCl 30°C, d = 1.27 cm.L = 10.16 cm.V = 14,130 cm./min.# - experimental data A - predicted curve

\/ /A

' /A

/

/â

/Z___ L 1.0 200 HÔô"400 600TIME - minutes

Figure A-4-1. An Example of the Application of the Empirical Rate Equation

1000


APPENDIX 5 Diffusional Mass Transfer Correlations

To appreciate the role of diffusional mass transfer in the dissolution of tin in HCl solutions, it is possible to write , [cJl

= k^AC (A-3-1 )

where k, is the mass transfer coefficient for the species concerned.

For the purpose of illustration, let us consider the casein which the diffusion of stannous ions from the metal-solutioninterface to the bulk solution, Sn^* ^Sn^^, is the importantphysical step, then

AC = Cg++ - (A-3-2)

Because the concentration of stannous ion in the bulk of the solution is negligible in comparison to the concentration at the metal-solution interface, Cg^+t <([<cCg+ + . Also ,which is the saturation concentration of stannous ion corresponding to the specific conditions, can be regarded as a constant. Therefore, from Equations(A-3-1) and (A-3-2), we

= kr-const (A-3-3)

%

obtain''tr- = V

By combining Equations (1-6-2) and (A-3-3), we obtain = const]'(v)*^^(d)"'^^(L)~'^^+ constg* (v)‘ ^(d)”* (L)”*- ^+ c o n s t y ( v ) ( d ) " ‘^^(L)“*^^ (A-3-4)

164


165

We further assume that the physical properties of the corroding solution, ^ and are constants, then.Equation(A-5-4) can be changed into the following form by multiplying both sides by d, dividing both sides by Dg^+ , and rearranging slightly:

S6+' M I V l - 3 0 / ^ ) - 3 3

+ const

S&'

' /'(ivf \ / d ,

z \ : ^ l \ J ^ J \L-33/(5g^)-41

Since the diffusion coefficients of metal ions at the metal-solution interface are generally not measurable, the correlations between dissolution rate constants and Dg++ were not directly investigated. However, for a discussion which requires only characteristic comparison it is quite reasonable to assume that the value of Dg++ is constant. Therefore, the three factors of Dg++ at the right-hand side of Equation(A-5-5) can be combined vriLth their respective constant. Thus, we have

kj_d“Sn

= Ku = cj (Re)'G7(Sc)'&?( j )'Z9

. 4 (Iîe)-59(Sc)-59( f ) •3>

Cj (Re)-G'(5c)'G'( I )-3° (1-6-3)


166

The above equation is consistent in form with Equation (1-6-1); which is the generally accepted form of dimensionless correlations for diffusional mass transfer for a fluid moving through a tube of circular cross-section.


APPENDIX 6Preparation of Saturated Calomel Electrode

A sketch of the saturated calomel electrode used in this investigation is shown in Figure A-6-1.

First, pure mercury was added to the reference electrode compartment to produce a layer of 1 to 2 cm. depth. This was covered with an equally thick layer of a paste made by stirring equal weights of mercurous and potassium chlorides with a saturated potassium chloride solution containing a large excess of solid salt. After the cell was filled with

Connection to KCl salt bridge

Hg

Saturated KCl solution

HggClg + KCl (paste)Mercury pool

Platinum wireFigure A-6-1. Saturated Calomel Electrode

saturated potassium chloride solution, two days were allowed for the S.C.E. to reach its equilibrium potential.

167


APPENDIX 7Specifications of the Model 4700M Research Potentiostat

Following are the specifications of the Model 4700M Research Potentiostat used in this investigation:

Potential control range: 1> 3 and 6 volts.Potential control stability: _+ 0.3 mv/24 hours.Potential meter scales: ^ 1, 3 and 6 volts.Potential recorder output: 0 ™ 10 mv.Anodic and cathodic current capacity: 10 amps at _+ 10

volts and 3 amps at _+ 20 volts.Current meter scales: _+ lOOy^a to + 10 amps in 6 decades. Current recorder output: 0 - 10 mv.Voltage meter scales; _+ 10 and 20 volts.Rise time: less than 2 microseconds at rated currents

into a resistance load over the full _+ 6 volt potential range.

Noise: less than 130 microvolts RMSInput impedance: 300,000 ohms at maximum imbalance,

resistance component greater than 100 megohms at balance, constant offset current over full potential control range typically Aa.,

168


APPENDIX 8

DATA OF TIN DISSOLUTION

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APPENDIX 9DATA OF POLARIZATION AND CATHOLIC POLARIZATION TESTS

206


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Metals," MacMillan, New York,(1966).84. Uhlig, H.H., J. Electrochem. Soc., Ill, 13(1964).83. Uhlig, H.H., "Corrosion and Corrosion Control," John Wiley

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NOMENCLATURE

2A apparent surface area of specimen, cm. dL activity^ x d activity of oxidized species ^red activity of reduced species C concentration, moles/liter C^ concentration of metal ion, moles/liter D diffusion coefficient, sq.cm./sec.

diffusion coefficient of metal ion, sq.cm./sec. d diameter of specimen, cm.E electrode potential, volts

standard electrode potential, voltsE^ corrosion potential, volts^eq equilibrium potential, voltsE^ polarized potential, voltsF Faraday, 23,063 nal./volt-mole(96,300 coulomb/equiv.)G Gibbs free energy, cal.

pg gravitational acceleration, 980.663 cm./sec.

?Iç corrosion current density, amp./cm.2i current density, amp./cm.

i „ _ applied current density, amp./cm.app. plimiting current density, amp./cm.

2ig exchange current density, amp./cm.k® reaction rate constant for hydrogen evolution

228


229

kg reaction rate constant for oxygen depolarizationk^ reaction rate constant for autocatalytic reactionkj mass transfer coefficient of species j, cm./sec.kj mass transfer coefficient, cm./sec.L length of specimen, cm.M an atom of metal

pN mass transfer flux, moles/cm.-sec.

pN. mass transfer flux of species j, moles/cm.-sec.J p

mass transfer flux of metal ion, mole s/cm.-sec.n number of electrons transferred in a reactionPy partial pressure of oxygen, atm.^2

R gas constant, 1.98? cal./deg-moler reaction rate of hydrogen evolution, moles/cm?-min.

prg reaction rate of oxygen depolarization, moles/cm.-min.

pr^ reaction rate of autocatalytic reaction, moles/cm.-min.Sn tin ions in solution, moles[Sn] tin ion concentration, moles/literA[Sn].tin ion concentration resulted from Hp evolution,

moles/literA[8n]ptin ion concentration resulted from oxygen depolarization,

moles/literA[sn]vtin ion concentration resulted from autocatalytic

•^reaction, moles/literSn active sites on a tin surface*T absolute temperature, °Kt time

yu . mobility of species j, cm.-mole/joule-sec.


230

V solution volume, literV fluid velocity, cm./min.z. charge number of species j

V

GREEK SYMBOLS

« constant7 constanti constant« constantV» constant

Tafel constant anodic Tafel constant cathodic Tafel constant

J diffusion layer thickness, cm.f density, gm./cm^/U. viscosity, gm./cm.-sec.

? kinematic viscosity, cm./sec.name of function

^ name of functionf overvoltage, volts^ activation overvoltage, volts

concentration overvoltage, volts resistance overvoltage, volts

^ total overvoltage, voltsÜ) angular velocity, r.p.m.e angle, radians


231<î> electrostatic potential

SUBSCRIPTS

1 activationa anodicb bulk solutionC concentrationc cathodic or corrosioni interfacej jth speciesM metal atom or ionR resistancer radial direction of tubes surfaceT totalt timez axial direction of tube

DI^'IENSIONLESS GROUPS

Gr Grashof number,

Nu Nusselt number for mass transfer, ^

Re Reynolds number,

Sc Schmidt number,

Sh Sherwood number, ^


232

VITA AUCTORIS

Birthplace; Date of Birth: Education:

Experience :

Shantung, China February 18, 1938Provincial Yunlin High School, Taiwan, China, 1933-1938, diploma.National Taiwan University, Taiwan, China, 1938-1962, B.S. degree.University of Windsor, Ontario, Canada,1963-1967, M.A.Sc. degree.University of Windsor, Ontario, Canada, 1967-present, Ph.D. candidate.Taiwan Fertilizer Co., Factory No. 6,Ammonia Plant, Oct. 1962 - Oct. 1963,Shift Supervisor.


part one. a study of the dissolution of tin tubes in

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