Anodic Dissolution of Silver in Aqueous Solutions

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Masters Theses Student Theses and Dissertations

1966

Anodic dissolution of silver in aqueous solutions

B. W. Jong

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Recommended Citation Jong, B. W., "Anodic dissolution of silver in aqueous solutions" (1966). Masters Theses. 5780. https://scholarsmine.mst.edu/masters_theses/5780

This thesis is brought to you by Scholars' Mine, a service of the Missouri S&T Library and Learning Resources. This work is protected by U. S. Copyright Law. Unauthorized use including reproduction for redistribution requires the permission of the copyright holder. For more information, please contact [email protected]. ANODIC DISSOLUTION OF SILVER IN AQUEOUS SOLUTIONS

BING-WEN JONG - I 'f! ~ .

A THESIS submitted to the faculty of the UNIVERSITY OF MISSOURI AT ROLLA in partial fulfillment of the requirement for the Degree of MASTER OF SCIENCE IN CHEMICAL ENGINEERING Rolla, Missouri 1966

Aproved by

121407 ii

ABSTRACT

The purpose of this investigation was to study the anodic dissolution of silver in various electrolytes to determine if there existed any deviations from Faraday•s law in nitric acid-silver nitrate solutions. The effect 4 of small additions (lo- and 10-5 N) of Cl-~ so4-~ and cro4- on the apparent valence and dissolution potential was also studied. Current densities were varied from 0.00 to 0.07 amp•cm-2 . Based on the observed apparent valences and a grey

film containing metallic particles on the anode surface~ it was concluded that deviations from Faraday•s law existed under the experimental conditions. In nitric acid solutions, the deviation was due to corrosion or self-dissolution. In

other solutions~ it seemed to be primarily due to disintegration. It was also concluded that the dissolution in these solutions was diffusion controlled. iii

ACKNOWLEDGEMENTS

The author is deeply indebted to Dr. J. w. Johnson~ Associate Professor of Chemical Engineering, who served as research advisor, and Dr. w. J. James, Professor of Chemistry and Director of the Materials Research Center of the University of Missouri at Rolla. Their help, guidance~ and encouragement during the course of this investigation are sincerely appreciated. Thanks are extended to the Department of Chemical Engineering for the use of its equipment. The author gratefully acknowledges a student assist antship from the Department of Chemical Engineering. iv

TABLE OF CONTENTS Page

L·IST OF FIGURES ...... • ...... vi LIST OF TABLES...... vii I . INTRODUCTION. . . . • . . . . • . . . . . • ...... • ...... l II. LITERATURE REVIEW. . • ...... 2 A. Anodic Dissolution of Silver in Acidic and Aqueous Salt Solutions and Anion Effects 2 B. The Relationship between Current Density and Potential of the Silver Anode in Acidic and Aqueous Salt Solutions...... 6 III. EXPERIMENTAL ...... •...... •...... ••...•••..... 12

A. Materials . . . . • . • • . • . . . . . • ...... • . . • • ...... 12 B . Equipment ...... 12 C. The Apparent Valence of Silver Undergoing Anodic Dissolutiona in Various Electro 1 yte s at 50 C . • ...... • • . • . • • • . . . • • ...... 1 2 1 . Apparatus ...... 12 2. Procedure...... 13 3. Data and results ...... 15 a. Nitric acid-silver nitrate solutions...... 15 b. 0.30 N HNO -0.70 N AgNO solution3with KCl~ K 2 ~o4 ~ and K?CrO 11 additions ...••....•...... 16 4. Sample ealculations ...... •.... 16 a. Calculation of the apparent weight of silver dissolved from coulombic data...... 16 b. Calculation of the apparent valence • ...... • ...... 17 D. The Dissolution Potential of Silver Undergoing Anodic gissolution in Various Electrolytes at 50 c .•..•...... 17 1. Apparatus ...... 17 2. Procedure. . • ...... • . • . • • . . . . . 18 3. Data and results ...... •...... 18 a. Nitric acid-silver nitrate solutions...... 18 b. 0.30 N HNO -0.70 N AgNO solution3with KCl~ K~so ~ and K cro additions ...... •.....4 18 2 11 c. 1.0 N HNO~ with KCl~ K 2 so 4 ~ and K?Cro11 additions ..•••.....•.•.... 26 4. Sample ealculations ...•...... ••..... 26 v

IV. DISCUSSION...... 28 A. Anodic Dissolution of Silver in Various Electrolytes...... 28 B. The Relationship between the Overpotential of the Silver Anode and the Current Density in Various Solutions...... 30 V. LIMITATIONS ...... 36 VI . RECOMMENDATIONS ...... 37 VII. APPENDICES...... 38 A. Materials...... 39 B. Eqtlipmen t ...... 40 1. Surface preparation of silver specimen... 40 2. Apparatus for the electrolysis and potential measurements...... 40 C. Surface Preparation of Silver Specimens...... 41 D. Experimental Data (Tables III-XXXIV)...... 42 E. Derivation of an Equation for both Activation and Diffusion Control and its Application to the Silver Dissolution Reaction...... 70 BIBLIOGRAPIIY...... 77 VITA...... 80 vi

LIST OF FIGURES

Figure Page 1. Diagram of apparatus for measurement of apparent valence of silver undergoing anodic dissolution ...... 14 2. Diagram of apparatus for measurement of the anodic dissolution potential of silver...... 19 3. The Tafel curve for the anodic dissolution of silver in 1.0 N HNO~, 0.70 N HN0~-0.30 N AgN03, and 0.30 N HNo -0.70 N AgNo at 50°C.. 20 3 3 4. The Tafel curve for the anodic dissolution of silver in 0.10 N HNo 3 -o.go N AgNo and 0.07 N HN0 -0.93 N AgNo at 50 C ...... 3 21 3 3 5. The Tafel curve for the anodic dissolution of silver in 0.03 N HN03-0.97 N AgNO~, 0.01 ~ HN0 -0.99 N AgNo , and 1.00 N AgN0 at 50 C.. 22 3 3 3 6. The Tafel curve for the anodic dissolution of silver in 0:~0 N HN03-0:7? N AgNO~ ~nd with lo-5 and 10 N KCl add1t1ons at 50 c ...... 23 7. The Tafel curve for the anodic dissolution of sil5er in 0.~0 N HN03-0.70 N AgNO~ and with 10- and 10- N K so additions a£ 50°C ...... 24 2 4 8. The Tafel curve for the anodic dissolution of silver in 0.~0 N HN0~-0.70 N AgNO~ and with lo-5 and 10- N K cr04 additions ~t 50°C ..... 25 2 g. The Tafel curve for the anodic diss~lution of 4 silver in 1.0 N ~o 3 and with 10- N KCl, 10- N K so , and 10- N K cro additions at 50oc. 27 2 4 2 4 10. Current density-overpotential relationship of the silver anode in aqueous solutions at 50°C 32 11. Current density-overpotential relationship of the silver anode in aqueous solutions at 50°C 33 12. Current density-overpotential relationship of the silver anode in aqueous solutions at 50°C 34 vii

LIST OF TABLES

TABLE PAGE I. Summary of the Apparent Valences of Silver Anode in Various Electrolytes at 50oc ...... 29 II. The Exchange Current of Silver Anode in AgNo -HNo Solutions at 50°C ....•...... •. 35 3 3 III. The Weiqht Loss of Silver Electrode in HNO at 50°C with Zero External Current Appli~d .. 43 IV. The Apparent Valence of Silver in 1.00 N HNo at 50°C...... 3 44 v. The Apparent Valence 8f Silver in 0.70 N HNo -0.30 N AgN0 at 50 C ..••...... 3 45 3 VI. The Apparent Valence 8f Silver in 0.30 N HNo -0.70 N AgNo at 50 C...... 3 46 3 VII. The Apparent Valence 8f Silver in 0.10 N HNo -0.90 N AgNo at 50 C ...•...... •..•.•.....•.3 46 3 VIII. The Apparent Valence gf Silver in 0.07 N HNo -0.93 N AgNo at 50 C ..•...... 3 47 3 IX. The Apparent Valence 8f Silver in 0.03 N HN0 -0.97 N AgNo at 50 C ...•...... ••...... 3 47 3 x. The Apparent Valence gf Silver in 0.01 N HNo -0.99 N AgN0 at 50 C ...... 3 48 3 XI. The Apparent zalence of Silver in 1.00 N AgNo at 50 C. . • ...... • ...... 49 3 XII. The Apparent Valence of Silver ig 0.30 N HNo -0.70 N AgNo -lo-5 N KCl at 50 C ...... 3 50 3 XIII. The Apparent Valenc~ of Silver ig 0.30 N HN0 -0.70 N AgN0 -10- N KCl at 50 C ...... 3 50 3 XIV. The Apparent Valence of Silver in 9.30 N HNo -0.70 N AgNo -lo-5 N K so4 at 50 C ...... 3 51 3 2 XV. The Apparent Valenc~ of Silver in 9.30 N HNo -0.70 N AgN0 -lo- N K so at 50 C .•...... 3 51 3 2 4 viii

TABLE PAGE XVI. The Apparent Valencg of Silver in 0 30 N HNo -0.70 N AgNo -10- N K cro at 506 C .....••..3 52 3 2 4 XVII. The Apparent Valenc~ of Silver in 0 30 N HNo -0.70 N AgNo -lo- N K cro at 506 C ...... •..3 52 3 2 4 XVIII. Dissolution Potential of the Silver Electrode in 1.0 N HNo at 50°C •••...... ••...•...... 53 3 XIX. Dissolution Potential of the Silver Electrode in 0.70 N HN0 -0.30 N AgN0 at 50°C ...... 54 3 3 XX. Dissolution Potential of the Silver Electrode 0 in 0.30 N HN0 -0.70 N AgNo at 50 C ...•..•.. 55 3 3 XXI. Dissolution Potential of the Silver Electrode in 0.10 N HNo -o.go N AgNo at 50°c ...... 56 3 3 XXII. Dissolution Potential of the Silver Electrode in 0.07 N HN0 -0.93 N AgNo at 500 C ...•...•. 57 3 3 XXIII. Dissolution Potential of the Silver Electrode in 0.03 N HN0 -0.97 N AgN0 at 500 C ...... 58 3 3 XXIV. Dissolution Potential of the Silver Electrode 0 in 0.01 N HN0 -0.99 N AgNo at 50 C ..•...... 59 3 3 XXV. Dissolution Potential gf the Silver Electrode in 1.0 N AgNo at 50 C .••.•...... •...... 60 3 XXVI. Dissolution Potential of the Silv~r Electrode in 0.30 N HN0 -0.70 N AgNo -lo-~ N KCl at 500 c ...... 3...... 3...... 61

XXVII. Dissolution Potential of the Silv~r Electrode in 0.30 N HN0 -0.70 N AgNo -lo- N KCl at 50 0 c ...... 3...... 3...... 6 2 XXVIII. Dissolution Potential of the Silver Electrode in 0.30 N HN0 -0.70 N AgNo -lo-5 N K so at 500 c ...... 3...... 3...... 2. . . 4. . . . . 6 3

XXIX. Dissolution Potential of the Silv~r Electrode in 0.30 N HN0 -0.70 N AgNo -lo- N K so at 500 c ...... 3...... 3...... 2. . . 4. . . . . 6 4

XXX. Dissolution Potential of the Silver Electrode in 0.30 N HN0 -0.70 N AgN0 -lo-5 N K cro at 500 c ...... 3...... 3...... 2. . . . 4. . . . 6 5 ix

TABLE PAGE XXXI. Dissolution Potential of the Silver El~4trode in 0.30 N HN0 -0.70 N AgN0 - 10 N K cro at 500 C ...... •...... 3 3 66 2 4 XXXII. Dissolution Potential of the ~ilver Electrode in 1.0 N HNo -lo- N KCl at 50°c 67 3 XXXIII. Dissolution Potential of the ~ilver El5ctrode in 1.0 N HNo -lo- N K so at 50 c ...... 3...... 2. . . 4...... 6 8

XXXIV. Dissolution Potential of the ~ilver Elect5ode in 1.0 N HNo -1o- N K cro at 50 c ...... 3...... 2. . . . 4...... 6 9

XXXV. Current .Density-Overpotential Relationships of the Silv5r Anode in 0.70 N HNo -0.30 N AgNo at 50 C. • • • • • • • • • • • • • • • • • • • 3• • • • • • • • • 73 3 'XXXVI. Current Density-Overpotential Relationships of the Silvgr Anode in 0.30 N HNo -0.70 N AgNo at 50 C...... 3 73 3 XXXVII. Current Density-Overpotential Relationships of the Silv5r Anode in 0.10 N HNo -o.go N AgNo at 50 C. • . . • . . . • ...... • . . 3 ...... 74 3 XXXVIII. Current Density-Overpotential Relationships of the Silv5r Anode in 0.07 N HNo -0.93 N 3 AgNO 3 at 50 C . • ...... • . . . • . . . . 7 4 XXXIX. Current Density-Overpotential Relationships of the Silvsr Anode in 0.03 N HNo -0.97 N AgNo at 50 C. • • • • • • • • • • • • • • • • • • • 3• • • • • • • • • 75 3 XL. Current Density -overpotential Relationships of the Silvsr Anode in 0.01 N HNo -0.99 N AgNo at 50 C. . • . . . . . • • ...... • • . 3 . . . • . . . . . 75 3 XLI. Current Density-Overpotential Relationshipg of the Silver Anode in 1.0 N AgNo at 50 C 76 3 1

I. INTRODUCTION

The use of silver metal in industry is extensive. For

example~ it is used in coatings~ in electroplatings~ in cir

cuit panels~ as catalysts for vapor-phase oxidations~ in the

sterling silver and photographic industries~ and in jewelry.

Thus~ the need of an understanding of its properties in solution is important. Reports have been made in the literature that discrep ancies often arise between coulometric data and the weight

loss of metal electrodes when active metals such as magne

sium~ zinc~ aluminum~ or cadmium dissolve anodically in aqueous salt solutions. This deviation from Faraday's law has been ascribed to disintegration of the metal, local

corrosion~ and uncommon valence ion formation. The present work is an investigation of the anodic dissolution of silver metal in both acidic and aqueous salt solutions. The intent is to determine if a similar devia-

tion from Faraday's law exists for silver. If so~ the un-

common valence theory would have to be rejected~ as silver~ being a monovalent ion~ would have no intermediate~ partially oxidized state. The effect of various anions on the anodic

dissolution was also observed. 2

II. LITERATURE REVIEW

The review of literature is divided into two parts: (1) anodic dissolution of silver in acidic and aqueous salt solutions~ including anion effects~ and (2) the re lationship between the current density and potential of the silver anode in acidic and aqueous salt solutions.

A. Anodic Dissolution of Silver in Acidic and Aqueous Salt Solutions and Anion Effects Johnson and Sanghvi (1) investigated the anodic dissolution of silver in 1.0 N solutions of potassium nitrate and ammonium nitrate at temperatures of 25 and 50°C and in 1.0 N potassium sulfate at 25°C. The current density was varied from 10 -3 to 10 -1 amp•cm -2 . The apparent valence of silver varied between 0.95 and 1.00. Small silver particles were found which originated from the anode. However~ the amount was very small.

When a noble metal is dissolved electrolytically~ the surface often becomes covered with a dark-colored material known as "anode slime". This substance either continuously or periodically parts from the surface and accumulates under the anode. Vermilyea (2) investigated the nature of these slimes. When silver (99.99 percent) was anodically dissolved in one normal silver nitrate -2 (pH = 3) at 1 amp•cm ~ he found that the slime consisted of metallic silver. No elements other than silver were detected by x-ray fluorescence. He suggested that~ 3

.. Ordinary solutions contain many impurities~ some of which adsorb on metal surfaces and cause passivation~ some sec tions of the metal surface remain passive and are cut off by dissolution from the sides, forming the silver particles which constitute the anode slime ... Craig, Hoffman, Law, and Hamer (3) investigated the silver anode dissolution in aqueous solutions of perchloric acid (20 wt. percent) containing a small amount (0.5 wt. percent) of silver perchorate. They found that as the sil ver goes into the solution, some particles fall to the bot tom of the anode compartment. They did not explain the origin of the particles. Active metals, such as magnesium, zinc, aluminum, and cadmium dissolving anodically in aqueous salt solutions, often give discrepancies between coulometric and weight loss data. This deviation from Faraday's law has been ex plained by corrosion and disintegration of the metal (4,5, 6,7,8,9,10,11,12), complex ion formation (13), and uncommon valence ion formation (14). The disintegration is most pronounced under conditions which cause high corrosion rates. The negative difference effect of metals can be explained in terms of corrosion which removes blocks or "chunks" of metal containing per haps only a few atoms. Marsh and Schaschl (15) investigat ing steel dissolving in dilute hydrochloric acid (pH= 2), pointed out that about 50 percent of the observed corrosion

rate could be attributed to the chunk effect. They proposed 4

that if Vermilyea (2)~ who dissolved silver anodically and found an uanode slime"~ had also taken corrosion rate and applied current data~ he would have observed the negative difference effect.

The fact that silver particles can be obtained in solution by silver electrode disintegration at high current densities is not a new concept. In 1906~ Burton (16) made colloidal solutions of metals such as platinum, gold~ silver~ bismuth~ lead, and iron by using current densities of 6.5 to 7.5 amp.cm-2 and potentials of 30 to 60 volts. He found that clouds of finely divided metal would scatter from the cathode during sparking and would remain suspended in the water. He explained the mechanism as vaporization of the metal at-the electrode with dispersion and condensa- tion of the vapor in the s0lution. However, large coarse particles were found in the solution. Some metals for which disintegration has been recognized as such are: Mg dissolving in dilute aqueous solutions of HCl and HClo4(4); Mg dissolving in NaCl solutions (5)7 Mg dissolving in HCl: HClo , and H so solutions (6)1 Be dissolving in HCl solu 4 2 4 tions (7,8)~ Al dissolving in KCl-HCl solutions (9)1 Cd dissolving in three percent KN0 solutions (10)7 Zn dis 3 solving in three percent KNo3 solutions (11)~ and Zn dis solving in 0.05 to 2.0 N KN03 solutions (12). Del Boca (13) investigated the dissolution of various metals in liquid ammonia and observed similar phenomena. He suggested a mechanism in which the dissolving metal entered 5 into solution by the formation of complex ions such as zn.zn+ 2 ~ cd.cd+ 2 ~ and A1 ·Al+3. 2 Sorensen, Davidson, and Kleinberg (14) investigated the dissolution of zinc and cadmium in NaNo ~ KClo , and 3 3 11 NaCl-NaNo3 solutions. They suggested that the uncommon valence ion 11 concept explained the deviation from Faraday's law. The first step involved the oxidation of the metal to a unipositive ion at the anode.

M M+ + e (at the anode) ( 1)

This unipositive ion, being very reactive, would readily for a bipositive ion and could occur in two ways, either by further oxidation at the anode or by oxidation by an oxidizing agent in the solution, i.e.,

M+ 2 + e (at the anode) ( 2) or M+ + oxidant M+ 2 + reductant (in solution) (3)

With nonreducible electrolytes, reaction (3) could not occur and a valence of two would be observed. With re- ducible electrolytes, both reactions (2) and (3) could occur~ thus valencies between 1 and 2 would be observed.

The effect of the anions, Cl-, Br , and I-~ on the anodic dissolution of metals was studied by Kolotyrkin (17). He found that they increased the anodic dissolution rate of cadmium and indium amalgams in acid solutions. He pointed out that this is accounted for by a direct participation of 6 these ions in elementary processes of ionizing metal atoms.

The extent of their filling the surface shifts the potential to more positive values. These concepts explain both the dissolution rate acceleration by ions that specifically adsorb and passivation of a metal surface by adsorbed oxygen from water. He suggested that the mechanism of the dissolution of cadmium could, in this case, be expressed by two con- secutive reactions. The first corresponded to the specific adsorption of anions with the formation of a surface complex, and the second corresponds to the dissolution of this complex, i.e.,

Cdi -n (4) Cd + ni n 2 Cdi -n Cdi -(n- ) + 2e ( 5) n n

B. The Relationship Between Current Density and Potential of the Silver Anode in Acidic and Aqueous Salt Solutions

Most studies of polarization phenomena are done in three ways: (1) measurement of electrode potentials at various current densities, (2) polarization growth and decay curves (electrode potentials as a function of time), and (3) depolarization effects at anodes and cathodes upon current reversal. Piontelli, Poli, and Serravalle (26) studied the electrode behavior of silver single crystals in 1 M AgClo4-l M with the surface oriented parallel to the HCl04 mixtures (111), (110), and (100) planes. They found that overvolt- ages were small, though somewhat higher than for Pb, Sn, and 7

Cd~ and tended to be symmetric on the anodic and cathodic sides. The overvoltage was generally independent of orien- tation~ but under some conditions~ an influence existed in complex salt solutions. Reedy (18) investigated silver reaction potentials in various electrolytes. The reaction potential was defined as that at which the electrode began to dissolve in the electrolyte. The following is a partial list of his results:

Electrolyte Concentration Reaction Potential (NHS)

(gmol/liter) {volt)

5 H s·o 0.521 o. 2 4

0.25 K so 0.521 2 4 1.0 HNo 0.520 3 0.521 1.0 KN03 0.222 1.0 KCl - 0.5 H2so4 0.077 1.0 KBr - 0.5 H2so4 -0.152 1.0 KI - 0.5 H2so4

Ferguson and Turner (19) investigated the depolarization effects after current reversal on silver anodes and cathodes in 2 N H so at a constant current density of 2 4 1.00 ma•cm-2 . The electrode reactions were suggested as follows: 8

Anodic + H • (ads) H (sol) + e ( 6) Ag(s) Ag+ (sol) + e (7) 2Ag+(sol) + so =(sol) 4 ( 8) Cathodic

Ag so (s) 2Ag+(sol) + so =(sol) 2 4 4 (9) 2Ag+(sol) + 2e 2Ag(s) (10) H+(sol) + e H. (ads) (11)

Jones and Thirsk (20) also investigated the behavior

of anodically polarized silver in 2 N H . The system 2so4 was studied by following potential variations during con 2 stant current (40 ma·cm- ) polarization at 25°C. They suggested that the stepwise anodic reaction occurred as follows:

2Ag 2Ag+ + 2e (12)

+ - Ag so 2Ag + so4 2 4 (13) Ag so + 60H- 2Ag0 + o + 2H+ + 804- 2 4 2 + 2H + 6e 2o (14) Redey (21) investigated the behavior of silver

upon anodic polarization in 1.0 N silver nitrate~ 1.0 N potassium nitrate~ and 1.0 N potassium chloride solutions. He found that stirring decreased the overpotential. The variation of polarization was greater in potassium nitrate 9 than in silver nitrate. This was attributed to the more constant silver ion concentration in silver nitrate solu tions. The silver anode in 1.0 N potassium chloride solution acted as a Ag/AgCl electrode. This is commonly used as a reference electrode and should be only slightly polarizable~ but it was found to be significantly polarized at low current densities. Hickling and Taylor (22) used polarization growth and decay curves to investigate the anodic behavior of silver in 1.0 N NaOH at 18°c. They found three main stages in the polarization: (1) charging of the double layer~ (2) formation of Ag o as a film~ and (3) formation of an oxide of 2 silver with an oxygen content higher than Ag2o 2 which decom poses to give this latter substance. They pointed out that at pH•s lower than 9~ the silver anode does not become pas sive and the anodic process is merely the dissolution of silver. van Norman (23) used a chronopotentiometer technique to investigate the dissolution of silver in molten silver chlo ride and silver bromide. He found the reaction to be a re versible one-electron oxidation process that could be ac counted for by several mechanisms. Mehl and Bockris (24) investigated the mechanism of the electrolytic deposition and dissolution of silver in 0.2 N AgClo - 1.0 N HClo solutions. They suggested the 4 4 surface diffusion of adions as the principle rate-determin ing step at low overpotentials. At high overpotentials~ the transfer of ions from the solutions to the electrode was 10 suggested as rate-determining.

In another study (25)~ they gave the relationship between overpotential and current density of fl ( 60 mv as:

e (l-d)zF~/RT -~zF~/RT i = i - e 0 . -~zFtl/RT (15) l... e 1 + ~o______v FeFrt/RT 0

2 i = current density~ amp·cm- -2 i = exchange current~ amp•cm 0

z = number of charges transferred = 1

F = Faraday•s constant = 96~500 coulombs/gmequivalent

R = gas constant= 8.314 joules/gmole·°K

~ = fraction of the overpotential assisting the dissoltuion of the electrode

~ = overpotential~ volts

-2 -1 v = exchange velocity~ em ·sec .grnequivalent 0

They also pointed out that similar relations could be obtained for the dissolution process. Johnson and Sanghvi (1) investigated the anodic dissolution of silver in potassium nitrate-silver nitrate mixtures. They suggested a dissolution sequence as:

(a) 'l < 0.05 volts

Ag(s) Ag+(sol) + e ( 16) ll

(b) '1 > 0.05 volts

fast + Ag(s) Ag (s) + e (17)

Ag+(s) + AgO(s) slow AgO(s) + Ag+(sol) (18) 12

III. EXPERIMENTAL

The purpose of this investigation was to study the anodic dissolution of silver in acidic salt solutions and the effect of various anions on the anodic dissolution. The experimental plan consisted of: (1) the effect of concentration~ various anions~ and current density on the apparent valence of silver undergoing anodic dissolution in nitric acid-silver nitrate solutions and (2) the effect of concentration~ various anions~ and current density on the dissolution potential of silver in nitric acid-silver nitrate solutions. The studies were carried out at 50°c.

A. Materials The list of materials used in the study is given in Appendix A.

B. EguiEment The list of equipment used in the study is given in Appendix B.

C. The AEparent Valence of Silver Undergoing Anodic Dissolution in Various Electrolytes at 50°C 1. Apparatus. The apparatus consisted of an electro lytic cell with separated compartments of 300 milliliter capacity~ a silver anode~ and a platinized-platinum cathode in series with a sensitive milliammeter~ a power source~ and a decade type power resistor. A diagram of the arrangement 13

is shown in Figure 1. A timer with one-second divisions was used for measuring the time. The cell was immersed in a constant-temperature water bath controlled at 5Q±O.l0 c. 2. Procedure. A cylindrical silver electrode was prepared by melting silver shot in a graphite mold in a vacuum furnace. The silver was of 99.999 percent purity. 0 It was melted at about 1200 C. Before each run~ the electrode was polished according to the procedure listed in

Appendix C~ etched in dilute nitric acid~ and rinsed in distilled water. It was then again rinsed with acetone and dried to constant weight in a dessicator. Approximately 140 milliliters of electrolyte were transferred into the cell. The cell was placed in the constant temperature bath in such a position as to ensure complete submergence of the solution. The solution remained in the water bath for about 30 minutes to bring the system to constant temperature before starting a run. The silver anode and platinized-plat- inurn cathode were put into the cell~ and connected into the external circuit. The cylindrical anode was suspended in the anolyte so that a known cross sectional area contacted the solution. Nitrogen gas was bubbled through both compartments. The timer and milliammeter were used to measure the number of coulombs passed. At the end of each run~ the silver anode was removed from the cell~ rinsed with distilled water and acetone~ and then dried to constant weight in a dessicator. The experimental weight loss of silver during the anodic dissolution was determined in this manner. G E F ~ 000 LTI @

A - Silver anode n n I II I I I I 1\ D I B - Platinum cathode C - Electrolytic cell D - Nitrogen inlet E - Power resistor F - Milliammeter c I I l~o I II I G - Power supply H - Knife blade switch

Constant Temperature Water Bath

Figure 1. Diagram of apparatus used for measurement of apparent valence of silver undergoing anodic dissolution.

~ -+==- 15

During this part of the investigation~ the chemical dissolution of silver in nitric acid solution was also checked to see if an appreciable weight loss occurred. The apparatus was essentially the same as described above except that no external circuit was used and only the silver electrode was placed in the cell. It was immersed in the nitric acid solution for 24 hours. It's weight loss for this period is shown in Table III~ Appendix D. Below

0.07 N HNo 3 ~ no detectable weight loss occurred. Above this concentration~ it can be seen that the weight loss increased with nitric acid concentration. However~ even at the higher concentration~ the chemical dissolution rate was very small compared to the anodic dissolution rate. The dissolution of the electrode was also checked by adding a small amount of KCl to the electrolyte to detect the presence of silver ions. 3. Data and results a. Nitric acid-silver nitrate solutions. The silver electrode was anodically dissolved in nitric acid- silver nitrate mixed electrolyte. The concentration ranged from 1.0 N HNo to 1.0 N AgNo . The ionic strength 3 3 was held constant at one. The current densities ranged -2 from 0.001 to 1.0 arnp•cm for 1.00 N HNo ~ 0.70 N HNo - 3 3 0.30 N AgNo ~ 0.30 N HN0 -0.70 N AgN0 ~ and 1.0 N AgNo . 3 3 3 3 The current densities ranged from 0.001 to 0.1 amp•cm-2 for 0.1 N HN0 -0.9 N AgN0 ~ 0.07 N HN0 -0.93 N AgNo ~ 3 3 3 3 0.03 N HN0 -0.97 N AgN0 , and 0.01 N HN0 -0.99 N AgNo . 3 3 3 3 16

The data are shown in Tables IV to XI~ Appendix D. It can be seen that the apparent valence of silver varied between 0.87 and 1.01. A gray colored film was observed on the surface of the silver electrode at all current densities and in all solutions except 1.0 N HNo • 3 b. 0.30 N HNo -0.70 N AqNo solutions with KCl, 3 3 K cro : and K cro additions. The silver electrode was 2 4 2 4 anodically dissolved in 0.30 N HNo -0.70 N AgNo with small 3 3 additions (10-5 and 10-4 N) of KCl~ K so ~ and K cro . 2 4 2 4 Current densities of 0.01 and 0.1 amp•cm-2 were used. The data are shown in Tables XII to XVII~ Appendix D. It can be seen that the apparent valence of silver varied between 0.94 and 0.96. A gray film was again observed on the electrode surface. 4. Sample calculations. The data from the experiment in 1.0 N HNo (Table IV~ Appendix D) have been used to 3 illustrate the calculation of the apparent valence. a. Calculation of the apparent weight of silver dissolved from coulombic data. The apparent weight of silver anodically dissolved according to Faraday's law assuming a normal valence of one was calculated from the expression:

(19)

where~

wa = apparent weight of silver dissolved~ gm

I = current = 0.760 amp (see Table IV) 17

t = time of run = 300 sec

A = atomic weight of silver = 107.88 gm

V normal valence of silver 1.0 n = =

F =Faraday's constant = 96,500 coulombs/gmequivalent

Therefore,

0.2549 gm

b. Calculation of the apparent valence. The apparent valence was calculated by the equation:

(20) where,

V apparent valence a = W weight of silver dissolved experimentally e = = 0.2545 gm

therefore,

- (0.2549~ = V a- (0.25~~ 1.00

D. The Dissolution Potential of Silver Undergoing Anodic Dissolution in Various Electrolytes at 50°C 1. Apparatus. The apparatus was the same as described previously, except that the anode was incorporated in an additional circuit so that the potential difference between it and a reference electrode could be measured. A calomel electrode (1 N KCl) was used as the reference. A salt 18 bridge (saturated ammonium nitrate) was also used to prevent chloride ions from diffusing into the electrolysis cell. The potential difference was measured with a high impedance electrometer. A diagram of the arrangement is shown in Figure 2. 2. Procedure. The silver electrode and cell were prepared as previously described except that when the silver anode and platinized-platinum cathode were placed in their respective compartments~ the silver anode was also connected to the calomel reference electrode through the electrometer. The potential of the silver electrode was measured at inter- vals of 15 minutes until it reached a constant value at cur rent densities ranging from 0 to 0.070 arnp·cm -2 .

3. Data and results. a. Nitric acid-silver nitrate solutions. Polariza- tion curves for silver undergoing anodic dissolutions in these solutions are shown in Figure 3 to 5. The concentrations ranged from 1.0 N HNo to 1.0 N AgNo . The ionic 3 3 strength was held constant at one. The data are shown in

Tables XVIII to xxv~ Appendix D. The previously mentioned gram film began to spall off the electrode at 0.03 amp·cm-2 •

b. 0.30 N HN0 -0.70 N AgNo solution with KC1 3 3 2 K so and K cro additions. Polarization curves for silver 2 42 2 4 undergoing anodic dissolution in these solutions are shown in Figures 6 to 8. The concentrations of the additions of 4 KCl~ K so _, and K cro were 10-5 and 10- N. The data are 2 4 2 4 shown in Tables XXVI to XXXI_, Appendix D. A gray film was G E F J [f 000 ITJ

A - Silver anode I I B - Platinum cathode C - Electrolytic cell D - Salt bridge E - Power resistor F - Milliammeter G - Power supply H - Knife-blade switch I - Nitrogen inlet J - Electrometer K - Calomel electrode th I

Figure 2. Diagram of apparatus used for measurement of the anodic dissolution potential of silver.

j-1 \0 0.08

1.00 N HN03 0.70 N HN0 -0.30 N AgN0 o.o6 3 3 0.30 N HN0 -0.70 N AgN0 3 3 Ul +l r-1 0 :> 0. 04 "" s::-'

0.02

0 . 00 I I I I I. I I I I I I I ! I - ! I I I I I I I - I I I ' I I i I I

i, amp·cm-2

Figure 3. The Tafel curve for the anodic dissolution of silver in 1.00 N HN03, 0.70 N HN0 -0.30 N AgN0 , and 0.30 N HN0 -0.70 N AgNo at 50°C. 3 3 3 3 "'0 o.o8r------~

0 0.10 N HN0 -0.90 N AgN0 0.06 3 3

8 0.07 N HN03-0.93 N AgN03 fJl ..j.J M 0 >0.04 ... s::'

0.02

0.00 10-l . -2 1, amp·cm

Figure 4. The Tafel curve for the anodic dissolution of silver in 0.10 N HNo3-

0.90 N AgN0 and 0.07 N HN0 -0.93 N AgNo at 50°C. 1--' 3 3 3 "' o.o8r------~

O.o6 1- 0 0.03 N HN03-0.97 N AgNo3

8 0.01 N HN0 -0.99 N AgNo 3 3 Ul .&J ...-1 [!] 1.00 N AgNo 0 3 :> 0. 04 c:-l

0.02

I I 1 1 l 1 r -,-1 ! I I I I IT 1 1 I I I I I I ' o.oo ~ [3 ~ 1'\ 10-l . -2 1, amp·cm

Figure 5. The Tafel curve for the anodic dissolution of silver in 0.03 N HNo - 3 o 0.97 N AgNo . 0.01 N HN0 -0.99 N AgN0 , and 1.00 NAg N0 at 50 C. 3 3 3 3 N N 0.08.------~

0 0.30 N HN0 -0.70 N AgN0 0.06 3 3 A 0.30 N HN0 -0.70 N AgNo -lo-5 N KCl 3 3 ell 4-J 4 r-f 8 0.30 N HN0 -0.70 N AgNo -1o- N KC1 go.o4 3 3

0.02

I :r I I I I I I I I I I I I I I I I I I I I I I j I -I o.ool111§. 10 1 . -2 1, amp·cm Figure 6. The Tafel curve for the anodic dissolution of silver in 0.30 N HN0 - 3 4 0.70 N AgNo and with 10-5 and 10- N KC1 additions at 50°C. 3 N lJJ o.osr------~

O.o61- 0 0.30 N HN03-0.70 N AgN03 8 0.30 N HN0 -0.70 N AgN0 -lo-5 N K so 3 3 2 4 [I) +J -4 r-1 0 0.30 N HN0 -0.70 N AgN0 -10 N K so 0 3 3 2 4 :> 0. 04 c::>

0.02

I I I. I I I I I I I I I I I I I I I I ! ! ! 0 . 00 I A ctr- i ; I I i - I X I

i, amp·cm-2

Figure 7. The Tafel curve for the anodic dissolution of silver in 0.30 N HNo3- 0.70 N AgNo and with 10-5 and 10-4 N K so additions at 50°C. 3 2 4 "'+:- 0.08~------~

0 . 06 1- 0 0 . 30 N HN0 -0 . 70 N AgNo 3 3 A 0.30 N HN0 -0.70 N AgNo -10 -5 N K cro 3 3 2 4 Cll ..j...l r-i G 0.30 N HN0 -0.70 N AgNo -10 -4 N K Cro 0 2 4 ~ 0.04 3 3 c:-'

0.02

0.00~~~--~--~~~~~~~--~--~~~~~~~--~--~~._~~

1.,. arnp·crn -2 Figure 8. The Tafel curve for the anodic dissolution of silver in 0.30 N HN0 - 3 0.70 N AgNo and w1.th. 10 -5 and 10 -4 N K cro additions at 50 0 C. 2 4 3 N \Jl 26 also observed on the surface of the silver electrode in all of these solutions and began to spall at the same current density as before. c. 1.0 N HN0 with KC1 K so , and K cro 3 3 2 4 2 4 additions. Polarization curves for these studies are shown in Figure g. The contrations of the additions of KCl, K so , and 2 4 K cro were 10-4 N. The data are shown in Table XXXIV, 2 4 Appendix D. No gray films were observed. The polarization curves were shifted a little in the positive direction from the one in 1.0 N HNo without additions. 3 4. Sample calculations. The method used to calculate the overpotentials from these studies are illustrated using data from the experiment in 1.0 N HN0 (Table XVIII, Appen 3 dix D). The overpotential was calculated from the equation: (21) = v - v.l.= 0 where,

= overpotential at a given current density V = steady dissolution potential at a given current = density = 0.581 volts vi=O = steady dissolution potential at a current density of zero

= 0.546 volts

Therefore,

= 0.581 - 0.546 0.035 volts 0.40.------~

0.30 1- 0 1.0 N HN0 3 A 1.0 N HN0 -Kcl til 3 +J r-1 8 1.0 N HN0 -K so 0 I 3 2 4 !> 0. 20 "' 0 1.0 N HN0 -K cro c"' I 3 2 4

0.10

1 1 1 I I I I I I I I 1 1 1 I I I I 1 0 • 00 I I I I ' I I I I I I I I I ~ ') 10- . -2 1, amp•cm

Figure 9. The Tafel curve for the anodic dissolution of silver in 1.0 N HNo3 and with 10 -4 N KCl, 10 -4 N K and 10 -4 N K additions at 50 0 C. 2so4, 2cro4 N -..J 28

IV. DISCUSSION

The discussion of results is presented in two parts:

(1) anodic dissolution of silver in various electrolytes, and (2) the relationship between the dissolution overpotential of the silver anode and current density in various electrolytes.

A. Anodic Dissolution of Silver in Various Electrolytes The anodic dissolution of silver was studied in the various electrolytes to determine if any deviation from

Faraday's law existed. The apparent valences of silver from thes~ studies have been summarized in Table I. It can be seen that the largest deviation from the normal valence occurred at low current densities in 1.0 N HNo . 3 It also can be seen from Table IIIJ Appendix DJ that in this solutionJ the silver electrode is chemically dissolved to some extent. With external current applied, the chemical dissolution of the silver electrode may also be increasedJ if some of its resistance to self-dissolution is due to passivating oxide films. ThusJ the chemical dissolution of silver probably accounts for these lower valences. For other HNo -AgNo solutions~ the apparent 3 3 valence varied between o.g4 and 1.01. A gray-colored film was observed on the surface of the silver electrode at all current densities. This film was removed and exam- ined. It was found to contain small metallic particles as TABLE I

SUMMARY OF APPARENT VALENCES OF THE SILVER ANODE IN VARIOUS ELECTROLYTES AT 50°C

Electrolyte Current Density, ma•cm -2 1 3______lO 30 _ __70 100 300 1000 1.0 N HNo o.87 0.92 o.94 0.94 0.95 0.96 1.oo 1.oo 3 0.93 0.94 0.95 0.95 1.00 1.00 0.91 0.95 0.70 N HN0 -0.30 N AgNo 0.95 1.00 1.00 3 3 0.95 0.95 0.95 0.94 0.96 0.30 N HN0 -0.70 N AgN0 0.97 1.00 1.00 3 3 0.97 0.99 0.97 0.96 0.97 0.10 N HN0 -0.90 N AgNo 0.97 1.00 0.97 0.97 0.99 3 3 0.07 N HN0 -0.93 N AgNo 3 3 0.97 0.98 0.99 0.96 0.97 0.97 0.02 N HN03-0.97 N AgNo3 0.96 0.96 0.98 0.97 0.98 0.01 N HN0 -0.99 N AgNo 0.99 0.97 o.g8 0.98 0.97 1.oo 3 3 1.0 N AgN0 1.00 o.g6 0.98 0.95 0.99 0.95 1.00 3 o.g6 0.99 o.g6 1.01 0.96 0.96 1.00 1.01 0.30 N HN0 -0.70 N AgNo -lo-5 N KCl 0.95 0.95 3 3 0.95 0.95 -4 0.30 N HN0 -0.70 N AgNo -10 N KCl 0.96 o.g6 3 3 0.30 N HN0 -0.70 N AgNo -lo-5 N K2so o.g4 0.96 3 3 4 0.95 0.95 -4 0.30 N HN0 -0.70 N AgN0 -10 N K so 0.94 0.95 3 3 2 4 0.30 N HN0 -0.70 N AgNo -lo-5 N K Cro 0.94 0.95 3 3 -4 2 4 0.30 N HN0 -0.70 N AgN0 -10 N K cro 0.94 o.g6 3 3 2 4 0.94 o.g6 N \..0 30

reported previously by several investigators (1~2,3). In 0.30 N HNo -0.70 N AgNo with small additions (10-5 3 3 4 and 10- N) of KCl~ K 2 so 4 ~ and K 2 cro 4 ~ the apparent valence of silver varied between 0.94 and 0.96. In the same solution without the addition~ the apparent valence was 0.97.

Precipitation was observed after the additions~ but a small amount of the added ions would have still remained in solution. The gray film was still observed on the silver anode surface in these solutions. Thus, the anion additions seemed to have very little, if any, effect on the anodic dissolution. Also, since Johnson and Sanghvi (1) observed similar gray films and valences ranging from 0.95 to 1.00 in 1.0 N potassium and ammonium nitrate solutions, it must be concluded that the presence of Ag+ does not effect the anodic dissolution.

B. The Relationship between the Overpotential of the Silver Anode and the Current Density in Various Solutions Overpotential measurements as a function of current density were made in the various electrolytes. The purpose of this part of the investigation was to determine the effect of Ag+ and small additions of Cl-~ so4=, and cro4- on the dissolution potential. The amount of polarization of the anode with current density decreased as the concentration of Ag+ increased. This is as expected, since the silver ion concentration at the electrode interface is much more stable in the more 31 concentrated AgNo solutions. Redey (21) has also reported 3 similar observations in 1.0 N KNo and 1.0 N AgNo solu- 3 3 tions. Silver~ in aqueous salt solutions~ is noble enough that changing the dissolution potential by changing the concentration of Ag+ is not sufficient to cause corrosion. Estimation of the exchange currents from Figure 10

through 12 for the anodic dissolution of silver in these studies gave the values shown in Table II. These values are sufficiently high to indicate a diffusion controlled

reaction. This is consistent with the low overpotentials at low current densities. No appreciable effect was noted on the dissolution potential by the small additions of Cl-~

so 4=~ and cro4=. 32

50 1.0 N AgNo 0 3

A 0.70 N HN0 - 3 40 0.30 N AgNo LC\. 3 0 -+01

()< ttSH 30 H

0 0 + ...... - 20

3 4 -Ft1/2RT) - e

Figure 10. Current density-overpotential relationship of the silver anode in aqueous solutions at 50°C. 33

50 0 0.01 N HN0 3 -0.99 N AgNo3 40 A 0.30 N HN0 3 -0.70 N AgNo3

0 0 1 2 3 4

Figure 11. Current density-overpotential relationship of the silver anode in aqueous solutions

at 50°C. 34

50 0 0.03 N HN0 3 -0.97 N AgNo 3 0.07 N HNo 40 3 -0.93 N AgNo 3 +O'l 0.10 N HNo .::X: 3 -0.90 N AgNo 30 3

0 0 1 2 3 4 ( eF'l/2RT e -F'V2RT)

Figure 12. Current density-overpotential relationship of silver anode in aqueous solutions at 50°C. 35

TABLE II THE EXCHANGE CURRENT OF THE SILVER ANODE IN AgNo -HNo 3 3 SOLUTIONS AT 50°C

Electrolyte Exchange Current (i ) 0 (amp·cm-2 ) 1.00 N AgNo 0.041 3 0.01 N HNo - 0-99 N AgNo 0.028 3 3 0.03 N HN0 - 0.97 N AgNo 0.027 3 3 0.07 N HN0 - 0.93 N AgNo 0.022 3 3 N N AgNo 0.017 0.10 HN03 - 0.90 3 0.30 N HN0 - 0.70 N AgNo 0.016 3 3 N HN0 - 0.30 N Ag:No 0.009 0.70 3 3 V. LIMITATIONS

Because the surface area (and hence the current density) of the silver anode varied during dissolution~ the relationship between overpotential and current den sity was not definite. This was also complicated by the geometric configuration of the electrode. 37

VI. RECOMMENDATIONS

It is recommended that polarization and valence studies be made in KClo -AgClo solutions, as no faradaic 3 3 efficiencies have been reported. These studies might be helpful in determining the anodic dissolution mechanism. It is also recommended that a different type of electrode of non-variant surface area be used. 38

VIII. APPENDICES

Five appendices are included in this thesis. Appendix A contains a list of materials. Appendix B contains a list of equipment. Appendix C contains a procedure for the surface preparation of silver specimens. Appendix D contains the experimental

data~ and Appendix E contains the derivation of an equation for both activation and diffusion control and its application to the silver dissolution reaction. 39

A. Materials

The following is a list of the major materials used in this investigation:

1. Silver nitrate 2 nitric acid 2 potassium chloride 2 potassium sulfate, potassium chromate, and ammonium nitrate. Reagent grade, meets ACS specifications, Fisher

Scientific Company, Fair Lawn, N. J. 2. Silver. 99.999 percent purity, Asarco, Central Research Lab., Americal Smelting and Refining Company,

South Plainfield, N. J. 3. Mercurous chloride. Reagent grade, meets ACS specifications, Merck and Company, Inc., Rahway, N. J. 40

B. Equipment

1. Surfacer preparation of silver specimens.

a. Belt surfacer. Buehler No. 1250~ Buehler Ltd., Evanston, Illinois.

b. Hand grinder. Handimet~ 4 stage~ Buehler No.

1470~ Buehler Ltd.~ Evanston, Illinois. 2. Apparatus for the electrolysis and potential measurements. a. Power supply. Model EUW-15, 0-500 volts,

Heath Company~ Benton Harbor, Michigan.

b. Power resistor. Model No. 240-C, 1 to 999~999 ohms in one ohm divisions, Clarostat Mfg. Co.~ Inc., Dover,

New Hampshire. c. Electrometer. Model 610 B, Keithley Instruments,

Inc.~ Cleveland, Ohio. d. Ammeter. Model 931, Weston Electric Instrument

Corporation~ Newark~ N. J. 41 c. Surface Preparation of Silver Specimens

The following procedure was used for the silver metal surface preparation: 1. All pits were removed from the metal surface with a wet belt surfacer equipped with a No. 150 grit abrasive cloth belt. 2. The sample surface was finished on a water-flushed four-stage hand grinder equipped with No. 's 240~ 320~ 400 and 600 abrasive strips, proceeding from the coarsest to the finest. 3. The sample was rinsed with distilled water, etched in dilute nitric acid~ and again rinsed with distilled water. 42

D. Experimental Data (Tables III-XXXIV) 43

TABLE III

THE WEIGHT LOSS OF SILVER ELECTRODE IN HN0 AT 50°C WITH 3 ZERO EXTERNAL CURRENT APPLIED

Solution Weight of Silver Electrode Loss* 2 (gm/cm ·24 hours2

1 N HN0 0.00146 3 0.7 N HN0 0.00146 3 0.3 N HN0 0.00106 3 0.1 N HN0 0.00027 3 0.07 N HN0 0.00000 3 0.03 N HN0 0.00000 3 0.01 N HN0 0.00000 3 2 *exposed area of the electrode = 0.754 em 44

TABLE IV

THE APPARENT VALENCE OF SILVER IN 1.00 N HN0 AT 50°C 3

Time Current Density* Weight of Silver Apparent Calc. Expt. (sec) (amE·Cm-2 ) (gm) (gm) Valence

9~000 0.00305 0.0231 0.0267 o. 87 3~000 0.00995 0.0252 0.0273 0.92

3~000 0.00995 0.0252 0.0270 0.93

3~000 0.0100** 0.0335 0.0367 0 .. 91

3~000 0.0300** 0.1006 0.1073 0.94

3~000 0.0300** 0.1006 0.1068 0 .. 94 500 0.0703 0.0296 0.0314 0 .. 94 500 0.0703 0.0296 0.0313 0.95

3~000 0.0700** 0.2348 0. 2480 0.95

3~000 0.0700** 0.2348 0. 2466 0.95

3~000 0.0700** 0.2348 0. 2463 0.95 300 0.0995 0.0252 0.0262 0.96 1,000 0.300 0. 2527 0.2528 1.00 1,000 0.300 0.2527 0.2532 1.00 300 1.008 0. 2549 0. 2568 1.00 300 1.008 0. 2549 0.2545 1.00 2 *area of the electrode = 0.754 em 2 **area of the electrode = l.OOcm 45

TABLE V

THE APPARENT VALENCE OF SILVER IN 0.70 N HNo -0.30 N AgNo 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp·cm-2 ) (gm) (grn) 30~000 0.00099 0.0252 0.0266 0-95 9~000 0.00305 0.0231 0.0243 0.95 3~000 0.00995 0.0252 0.0266 0.95 900 0.0305 0.0231 0.0245 0.94 500 0.0703 0.0296 0.0308 0.96 300 0.0995 0.0252 0.0264 0.95 1~000 0.300 0.2527 0. 2536 1.00 300 1.008 0 .. 2549 0.2558 1.00 2 *area of the electrode = 0.754 em 46

TABLE VI

THE APPARENT VALENCE OF SILVER IN 0.30 N HN0 -0.70 N AgNo 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp·crn-2 ) (grn) (grn) 30,000 0.00099 0.0252 0.0260 0.97 9,ooo 0.00305 0.0231 0.0233 0.99 3~000 0.00995 0.0252 0. 0260 0.97 900 0.0305 0.0231 0.0243 0.95 500 0.0703 0.0296 0.0306 0.97 300 0.0995 0.0252 0.0260 0.97 1,000 0.300 0.2527 0. 2537 1.00 300 1.008 0. 2549 0. 2551 1.00 2 *area of the electrode = 0.754 em

TABLE VII

THE APPARENT VALENCE OF SILVER IN 0.10 N HN0 -0.90 N AgNo 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp·cJ1"! -2) (grn) (qrn) 9,000 0.00305 0.0231 0.0239 0.97 3,000 0.00995 0.0252 0.0252 1.00 900 0.0305 0.0231 0.0238 0.97 500 0.0703 0.0296 0.0307 0.97 300 0.0995 0.0252 0.0254 0.99 2 *area of the electrode = 0.754 em 47

TABLE VIII

THE APPARENT VLAENCE OF SILVER IN 0.07 N HN0 -o.g3 N AgNo 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (arnp-crn-2 ) (gm) (gm) 30!'000 O.OOOg9 0.0252 0.0259 o.g7 9!'000 0.00305 0.0231 0.0236 o.g8 3!'000 o.oogg5 0.0252 0.0254 o.gg goo 0.0305 0.0231 0.0240 o.g6 500 0.0703 0.0296 0.0305 0.97 300 0.0995 0.0252 0.0260 0.97 2 *area of the electrode = 0.754 cm

TABLE IX

THE APPARENT VALENCE OF SILVER IN 0.03 N HN0 -0.97 N AgNo 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp·cm -2) (gm) (gm)

g~ooo 0.00305 0.0231 0.0242 0.96 3.,000 0.00995 0.0252 0.0263 0.96 goo 0.0305 0.0231 0.0237 0.98 500 0.0703 o.o2g6 0.0304 o.g7 300 0.0995 0.0252 0.0258 0.98 2 *area of the electrode = 0.754 em 48

TABLE X

THE APPARENT VALENCE OF SILVER IN 0.01 N HN0 -0.99 N AgN0 3 3 AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp-cm-2 ) (g:m) (gm)

30~000 0.00099 0.0252 0.0255 0.99 9~000 0.00305 0.0231 0.0238 0.97 3~000 0.00995 0.0252 0.0256 0.98 500 0.0305 0.0129 0.0132 0.98 500 0.0703 0.0296 0.0305 0.97 300 0.0995 0.0252 0.0252 1.00 2 *area of the electrode = 0.754 em 49

TABLE XI

THE APPARENT VALENCE OF SILVER IN 1.00 N AgNo AT 50°C 3

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence 2 (sec) (amE·cm- ) (gm} (gm} 20~000 0.00102 0.0224 0.0234 0.96 5!)000 0.00306 0.0168 0.0169 0.99 7!)000 0.00306 0.0235 0.0236 1.00 2!)000 0.0102 0.0224 0.0233 0.96 2!)000 0.0102 0.0224 0.0232 0.96 1!)000 0.0306 0.0335 0.0343 0.98 1!)000 0.0306 0.0335 0.0331 1.01 1!)000 0.0700** 0.0593 0.0627 0.95 500 0.0713 0.0391 0.0409 0.96 500 0.0713 0.0391 0.0396 0.99

1~000 0.100 0.0850 0.0891 0.95 1,000 0.100*** 0.1207 0.1260 0.96 1!)000 0.300** 0.2527 0.2518 1.00 l!JOOO 0.300 0.2527 0.2531 1.00 300 1.008 0. 2549 0.2231 1.01 2 *area of the electrode = 0.982cm 2 **area of the electrode = 0.754 em 2 ***area of the electrode = 0.54 em 50

TABLE XII

THE APPARENT VALENCE OF SILVER IN 0.3 N HN0 -0.7 N AgNo 3 3 -10-5 N KCl AT 50°C

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence 2 (sec) (amp·cm- ) (g:m) (g:m) 3~000 0.0100 0.0335 0.0353 0.95 3~000 0.0100 0.0335 0.0354 0.95 3~000 0.100 0.3354 0.3513 0.95 3~000 0.100 0.3354 0.3521 0.95 2 *area of the electrode = 1.00 em

TABLE XIII

THE APPARENT VALENCE OF SILVER IN 0.3 N HN0 -0.7 N AgN0 3 3 -10-4 N KCl AT 50°C

Time Current Density * Weight of Silver Apparent Calc. Expt. Valence (sec) (amp·cm-2) (g:m) ( gm)

3~000 0.00995 0.0252 0.0263 0.96 300 0.0995 0.0252 0.0263 0.96 2 *area of the electrode = 0.754 em 51

TABLE XIV THE APPARENT VALENCE OF SILVER IN 0.3 N HN0 -0.7 N AgN0 3 3 -10-5 N K so AT 50°C 2 4

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence 2 (sec) (arnp·crn- ) (g:m) (g:m) 3~000 0.0100 0.0335 0.0355 0.94 3~000 0.0100 0.0335 0.0354 0.95 3~000 0.100 0.3354 0.3512 0.96 3,000 0.100 0.3354 0.3524 0.95 2 *area of the electrode = 1.00 em

TABLE XV

THE APPARENT VALENCE OF SILVER IN 0.3 N HN0 -0.7 N AgNo 3 3 -10-4 N K so AT 50°C 2 4

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec} (arnp·cm -2).. (g:m} (g:m}

3~000 0.0100 0.0335 0.0357 0.94 3,000 0.100 0.3354 0.3540 0.95 2 *area of the electrode = 1.00 ern 52

TABLE XVI

THE APPARENT VALENCE OF SILVER IN 0.30 N HN0 -0.70 N AgNo 3 3 -10-5 N K cro AT 50°C 2 4

Time Current Density* Weight of Silver Apparent Calc. Expt. Valence (sec) (amp.cm-~) (gm) (gm) 3:JOOO 0.0100 0.0335 0.0358 0.94 3,000 0.100 0.3354 0. 0562 0.95 *area of the electrode 1.00 em 2

TABLE XVII

THE APPARENT VALENCE OF SILVER IN 0. 30 N HN0 70 N AgN0 3 -o. 3 -10-4 N K cro AT 50°C 2 4

Time Current Density* Weight o£ Silver Apparent Calc. Expt. Valence (sec) (amp· em -2) (gm) (gm) 3:JOOO 0.0100 0.0335 0.0358 0.94 3:JOOO 0.0100 0.0335 0.0357 0.94 3:JOOO 0.100 0.3354 0.3501 0.96 3:000 0.100 0.3~54 0.3503 0.96 *area of the electrode = 1.00 em 53

TABLE XVIII

DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 1.0 N HN0 3 AT 50°C

i X 103 t v i X 103 t v 2 (arnE·Cm-2 )* (min} (volt)** (amE·cm- }* (min} (volt}** 0.0 0 0.541 3.0 0 0.641 15 o. 546 15 0. 6L11 30 o. 546 30 0.641 45 0.546 45 o.6lJl 0.1 0 0.576 10.0 0 0.676 15 0.581 15 0.676 30 0.581 30 0.681 45 0.681 0.3 0 o. 581 15 o. 586 30.0 0 0.716 30 0.586 15 0.720 30 0.720 1.0 0 0.606 45 0.720 15 0.611 30 0.611 70.0 0 0.740 45 0.611 15 0.740 30 0.7ll0 45 O.I40 2 *area of the electrode = l.ll5 crn **normal hydgrogen scale 54

TABLE XIX DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.70 N HN0 3 -0.30 N AgN0 AT 50°C 3

i X 103 t v i X 103 t v -2 2 (amE·cm }* (min} (volt}** (am:e-cm- }* (min} (volt}** 0.0 0 0.736 3.0 0 0.756 15 0.741 15 0.761 30 0.741 30 0. 761 0.1 0 o. 746 10.0 0 0.766 15 o. 746 15 0.771 30 o. 746 30 0.771 0.3 0 0.747 30.0 0 0.781 15 0.747 15 0.781 30 0.747 30 0.781 1.0 0 0-751 70.0 0 0.801 15 0-751 15 0.801 30 O-I21 30 0.801 2 *area of the electrode = 1.45 cm **normal hydrogen scale 55

TABLE XX DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.30 N HN0 3 -0.70 N AgN0 AT 50°C 3

i X 103 t v i X 103 t v 2 (am:e·cm-2 } (min} (volt)** (am:e-cm- } (min) (volt}** 0.0 0 0.767 3.0 0 0.780 15 0.768 15 0. 781 30 o. 768 30 0.781 0.1 0 0.773 10.0 0 0-791 15 0.774 15 0.791 30 0.774 30 0.791 0.3 0 0.774 30.0 0 0.801 15 0.774 15 0. 801 30 0.774 30 0.801 1.0 0 0.774 70.0 0 0.806 15 0.777 15 0.806 30 o.zzz 30 0.806 2 *area o£ the electrode = 1.45 em **normal hydrogen scale TABLE XXI DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.10 H HN0 3 -0.90 N AgN0 AT 50°C 3

i X 103 t v i X 103 t v ( arnE · em-2 }* (min} (volt}** {amE·Cm-2}* (min} (volt}** o.o 0 0.772 3.0 0 0.782 15 0.772 15 0.787 30 0.772 30 0.787 0.1 0 0.772 10.0 0 0.794 15 0.772 15 0.794 30 0.772 30 0.794 0.3 0 0.772 30.0 0 0.797 15 0-776 15 0. 806 30 0-776 30 0. 806 1.0 0 0.776 70.0 0 0. 821 1.5. 0..78l 15 0. 821 30 o.:z:81 30 0. 821 2 *area of the electrode = 1.45 cm **normal hydrogen scale 57

TABLE XXII

DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.07 N HN0 3 -0.93 N AgN0 AT 50°C 3

i X 103 t v i X 103 t v (amE·cm-2 }* (min} (volt}** (amE·cm-2 }* (min} (volt)** o.o 0 0.773 3.0 0 0.781 15 0.773 15 0.785 30 0.773 30 0.785 0.1 0 0.773 10.0 0 0.791 15 0.773 15 0.791 30 0.773 30 0.791 0.3 0 0.773 30.0 0 0.801 15 0.773 15 0.801 30 0.773 30 o. 801 1.0 0 0.776 70.0 0 o. 806 15 0.781 15 o. 806 30 o.z81 30 o. 806 2 *area of the electrode = 1.45 em **normal hydrogen scale TABLE XXIII DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.03 N HN0 3 -0.97 N AgNo AT 50°C 3

i X 103 t v i X 103 t v (amE·cm-2 )* (min} (volt}** (amE·Cm-2 }* (min} (volt}** 0.0 0 0.776 30.0 0 0.798 15 0.776 15 0.798 30 0.776 30 0.798 0.1 0 0.780 40.0 0 0.803 15 0.781 15 0. 803 30 0.781 30 0.803 0.3 0 0.783 50.0 0 0.804 15 0.783 15 0.804 30 0.783 30 0. 804 1.0 0 0.784 60.0 0 0.809 15 0.784 15 0.809 30 0.784 30 0.809 3.0 0 0.788 70.0 0 0.811 15 0.788 15 0.813 30 0.788 30 0.813 10.0 0 0.791 15 0.791 30 O.I21 2 *area of the electrode = 1.45 em **normal hydrogen scale 59

TABLE XXIV

DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.01 N HN0 3 -0.99 N AgNo AT 50°C 3

i X 103 t v i X 103 t v 2 (... am:e·cm-2 )* (min} (volt}** (amE·cm- }* (min} (volt}** o.o 0 o. 776 3.0 0 0.786 15 0.776 15 0.786 30 o. 776 30 0. 786 0.1 0 0.780 10.0 0 0.791 15 0.781 15 0.790 30 0.781 30 0.791 0.3 0 0.781 30.0 0 0.798 15 0.781 15 0.798 30 0.781 30 0.798 1.0 0 0.782 70.0 0 0.807 15 0.783 15 0.807 30 o.:z:83 30 o.so:r 2 *area of the electrode = 1.45 crn **normal hydrogen scale 60

TABLE XXV DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 1.00 N AgNo 3 AT 50°C

i X 103 t v i X 103 t v 2 2 (amE·cm- )* (min} (volt)** (amE·cm- }* (min} (volt)** 0.0 0 0.781 3-0 0 0.782 15 0.781 15 0.782 30 0.781 10.0 0 0.783 0.1 0 0.781 15 0.783 15 0.781 30 0.781 30.0 0 0.796 15 0.796 0.3 0 0.781 15 0.781 70.0 0 0.811 15 0.811 1.0 0 0.781 30 0.811 12 o.z81 2 *area of the electrode = 1.45 em **normal hydrogen scale 61

TABLE XXVI

DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.30 N HN0 3 -0.70 N AgNo -lo-5 N KC1 AT 50°C 3

i X 103 t v i X 103 t v 2 (am:e·cm -2).. * (min} (volt}** (amE·cm- }* (min} (volt}** 0.0 0 0.764 3.0 0 0.773 15 0.765 15 0.773 30 0. 765 30 0.773 45 0. 765 60 0. 765 10.0 0 o. 776 15 0.776 0.1 0 0.765 30 0.776 15 0.765 30 0. 765 30.0 0 0.781 15 0.781 0.3 0 0.768 15 0.768 70.0 0 0.798 30 0. 768 15 0.798 1.0 0 0.771 15 0.771 30 0.771 2 *area of the e1ectr0de = 1.00 em **normal hydrogen scale 62

TABLE XXVII DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.3 N HNo 3 -0.7 N AgNo -1o-4 N KC1 AT 50°C 3

i X 103 t v i X 103 t v -2 2 (arnE· em }* (min~ (volt~** (am:e·cm- } (min} (volt}** 0.0 0 o. 761 3.0 0 0.766 15 0.761 15 0.766 30 o. 761 30 0.766 0.1 0 0.761 10.0 0 0.771 15 0.761 15 0.771 0.3 0 0.766 30.0 0 0.781 15 0.766 15 0.781 1.0 0 0.766 70.0 0 0.796 15 0.766 15 0.796 30 o.zg_6 2 *area of the electrode = 1.45 em **normal hydrogen scale TABLE XXVIII DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.30 N HN0 3 -7.70 N AgN0 -lo-5 N K so AT 50°C 3 2 4

i X 103 t v i X 103 t v 2 (amE·cm-2 )* (min) (volt)** (amE·cm- )* (min) (volt)** 0.0 0 0.766 3.0 0 0.771 15 0.766 15 0.771 30 0.766 30 0.771 0.1 0 0.766 10.0 0 0.766 15 0.766 15 0.776 30 0.766 0.3 0 0.771 30.0 0 0.785 15 0.771 15 0.785 30 0.771 1.0 0 0.771 70.0 0 0.805 15 0.7_7_1 15 0.805 2 *area of the electrode = 1.45 cm **normal hydrogen scale 64

TABLE XXIX DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.3 N HNo 3 -0.70 N AgN0 -10 -4 K so AT 500 C 3 2 4

3 i X 10 t v i X 103 t v (am;e·cm-2 )* (min} (volt)** (am:e·cm-2 )* (min} (volt}** 0.0 0 o. 761 1.0 0 0.766 15 0.761 15 0.766 30 0.761 30 0.766 45 0.761 3.0 0 0.771 0.1 0 0.761 15 0.771 15 o. 761 10.0 0 0.776 0.3 0 0.761 15 0.776 15 0.761 30.0 0 0. 781 0.7 0 0.761 15 0.781 15 o. 761 70.0 0 0.801 12 0.801 2 *area of the electrode = 1.45 cm **normal hydrogen scale TABLE XXX

DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.30 N HN0 3 00.70 N AgNo -lo-5 N K cro AT 50°C 3 2 4

i X 103 t v i X 103 t v (amE·cm-2 }* (min} {volt)** (amE·cm-2 )* {min) {volt)** 0.0 0 o. 766 3.0 0 0.771 15 0.766 15 0.773 30 0. 766 30 0.773 0.1 0 0.771 10.0 0 0.776 15 0.771 15 0.776 30 0.771 0.3 0 0.771 30.0 0 0.781 15 0.771 15 0.781 1.0 0 0.771 70.0 0 0.801 15 0.771 15 0. 801 2 *area of the electrode = 1.00 em **normal nhdrogen scale 66

TABLE XXXI DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 0.3 N HN0 3 -0.7 N AgN0 -lo-4 N K cro AT 50°C 3 2 4

i X 103 t v i X 103 t v 2 (arnE·cm-2 }* (min} (volt}** (amE·cm- }* (min} (volt}** 0.0 0 0.761 3.0 0 0.771 15 0.761 15 0-771 30 0. 766 45 0. 766 10.0 0 0.776 15 0-776 0.1 0 0.771 15 0.771 30.0 0 0.781 30 0.771 15 0.781 0.3 0 0.771 70.0 0 0.796 15 0.771 15 0.801 30 0.801 1.0 0 0.771 12 o.zzl 2 *area of the electrode = 1.45 cm ** normal hydrogen scale TABLE XXXII DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 1 N HN0 3 -10-4 N KCl AT 50°C

i X 103 t v i X 103 t v 2 (amE·cm-2) ~* (min} ( vo1 t}** (amE·cm- } (min} (volt}** o.o 0 0. 565 3.0 0 0.671 15 0. 565 15 0.676 30 0. 565 30 0.676 0.1 0 0.611 10.0 0 0.706 15 0.611 15 0.711 30 0.611 30 0.711 0.3 0 0.651 30.0 0 0.736 15 0.651 15 0.736 30 0.651 30 0. 736

1.0 0 0.661 70.0 0 0.761 15 0.661 15 0.761 30 0.661 30 o.z61 2 *area of the electrode = 1.45 em **normal hydrogen scale 68

TABLE XXXIII DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN 1.0 N HNo 3 -4 K so AT 0 C -10 2 4 50

i X 103 t v i X 103 t v 2 2 (arn:e·crn- )* (min) (volt)** (arn:e.crn- }* (min} (volt}** 0.0 0 0.541 3.0 0 0.651 15 0.541 15 0.651 30 0.541 30 0.651 0.1 0 0.571 10.0 0 0.686 15 0. 586 15 0.686 30 0.586 30 0.686 45 0. 586 30.0 0 0.711 0.3 0 0.611 15 0.711 15 0.611 30 0.711 1.0 0 0.631 70.0 0 0.741 15 0.636 15 0.746 30 0.636 30 0.746 42 0.636 2 *area of the electrode = 1.45 em **normal hydrogen scale TABLE XXXIV DISSOLUTION POTENTIAL OF THE SILVER ELECTRODE IN l N HN0 3 -10-4 N K 2cro4 AT 50°C

i X 103 t v i X 103 t v 2 (amE·cm-2) .. * (min} (volt}** (amE·cm- 2* (min} (volt}** 0.0 0 o. 531 3.0 0 0.681 15 0.541 15 0.681 30 0.541 30 0.681 45 0.541 10.0 0 0.721 0.1 0 0.581 15 0.721 15 0.581 30 0.721 30 0.581 30.0 0 0.756 0.3 0 0.616 15 0.756 15 0.616 30 0.756 30 0.621 70.0 0 0.791 1.0 0 0.651 15 0.791 15 0.651 30 0.791 30 0.621 2 *area of the electrode = 1.45 cm **normal hydrogen scale 70

E. The Derivation of an Equation for Both Activation and

Diffusion Control and Its Application to the Silver Dis- solution Reaction For the reaction,Ag ----- Ag+ + e~ the following equations apply: i FK a.FV/RT -(1-a)FVlRT a = n aaAge nFk c aAg + e

V=ll·+V 'L rev

aFV 1 RT aF~IRT -(1-a)FV rev i k 'a e e -k 'a + e rev a = a Ag c Ag

X e-(1-a)Fo/R.T

v rev = v o

aFV /RT -alna la + a.Frt.IRT 0 Ag· Ag i = k 'a e e e a a Ag

a A - ( 1-a ) F '1 IRT -(1-a.)FVo.IRT e(1-a)lnaA~ e -k c 'aAg + e '=' ·

aAg+)a aFV lRT a,F(\/RT = k 'a ( e 0 e a Ag aAg

a l - ( 1-a ) FV IRT _ ( l-a) F (\ IRT - k ' a + (aAg.) -ae o e c Ag Ag .

aFfl 1 RT e 71

-(1-a)FV IRT k '~a )1-a(a ~)a e o = c Ag Ag (23)

ZF~ i a = b (aAg- aAg+)

ZF.O (a +) iL = ~ Ag

aAg aAg+

(24)

From equations (22), (23)~ and (24)

Let a = 0.5,

i a i ( eFfli2RT _ e -Ffl /2RT) ( 25) 0.01 i 0 5 = 0 (1 + 0.396 ~Ag+) . where, 2 i = anodic current density, rna·crn- a i = exchange current, rna.crn -2 0 72

CAg+ = silver ion concentration in bulk solution, gmols/1

F =Faraday's constant = 96~500 coulombslgmequiva1ent

~ = anodic overpotential~ volts R = gas constant = a.314 joules/gmol·°K T = temperature, °K a = fraction of the overpotential assisting the dissolution of the silver anode = diffusivity - 10-5 em 2 ·sec -1 z = number of charge transferred = 1

diffusion layer's thickness ~ 0.01 em rAg+ = activity coefficient of Ag+ in solution of unit ionic strength = 0.396 = limiting current density, ma·cm-2

* Reference (27) 73

TABLE XXXV CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF THE SILVER

ANODE IN 0.70 N HN0 -0.30 N AgN0 AT 50°C 3 3

i a L1._ 0.01 i 0 5 -~ i 'l (e2RT 2RT) a ( l + 0 • 396 ~ Ag+ ) . - e (ma·cm-2 ) (volts) 0.0 0.000 0.0000 o.ooo 0.1 0.005 0.0996 0.180 0.3 0.006 0.296 0.216 1.0 0.010 0.960 0.361 3.0 0.020 2.68 0.734 10.0 0.030 7-37 1.13 30.0 0.040 16.0 1.56 70.0 0.060 26.7 2.60

TABLE XXXVI CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF TIE SILVER ANODE IN 0.30 N HN0 -0.70 N AgN0 AT 50°C 3 3

i a .E.1._ 0.01 i 0.5 - L1_ (e2RT 2RT) i a 1 - e ( + 0 . 396 ~ Ag+) (ma·cm-2) (volts) o.e 0.000 0.0000 0.000 0.006 0.0998 o. 216 0.1 0.216 0.3 0.006 0. 298 1.0 0.009 0.982 0.325 3.0 0.013 2.85 0.471 10.0 0.023 8.57 0.850 0.033 20.8 1.26 30.0 1.47 70.0 0.038 37-3 74

TABLE XXXVII

CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF THE SILVER ANODE IN 0.10 N HN0 -0.90 N AgNo AT 50°C 3 3

i a 0.01 i .E!L F'1 i '1 a )0.5 a (1 (e2RT - e - 2RT) 2 +a. 396 CAg+ (ma·cm- ) (volts) o.o 0.000 0.0000 0.000 0.1 o.ooo 0.0999 0.000 0.3 0.004 0.299 0.144 1.0 0.009 0.986 0.325 3.0 0.015 2.88 0.546 10.0 0.022 8.84 0.811 30.0 0.034 22.1 1.30 70.0 0.049 4o.z 2.00

TABLE XXXVIII

CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIP OF THE SILVER ANODE IN 0.07 N HN0 -0.93 N AgNo AT 50°C 3 3 r------~~-a Fll. Ffl i ~ 0.01 ia 0.5 - 2RT) a (1 + o:3§0_C_+) (e2RT - e 2 ) . Ag --~~-a.·cm-~l~(~v~o~1~t~s~~----~------·------0.0 o.ooo 0.0000 0.000 0.1 0.000 0.0999 0.000 0.3 0.000 0.298 O.OQO 1.0 0.008 0.987 0.289 3.0 0.012 0.288 0.435 10.0 0.018 8.87 0.658 30.0 0.028 22.3 1.05 __Y;.._o ...... __.o __ Q,:.Q3.3._ 41 . 1 1 . 26 75

TABLE XXXIX

CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF THE SILVER ANODE IN 0.03 N HN0 -0.97 N AgNo AT 50°C 3 3

i a K.!L - .E!L i 0.01 i 0 5 a '1 (e2RT 2RT) ( l_ + 0 . 396 ~ Ag+) . - e (rna·cm-2 ) (volts) 0.0 0.000 0.0000 0.000 0.1 0.005 0.0999 0.180 0.3 0.007 0.299 0.252 1.0 0.008 0.987 0. 289 3.0 0.012 2.89 0.435 10.0 0.015 8.91 0. 546 30.0 0.022 22.5 0.811 70.0 0.037 41.7 1.43

TABLE XL

CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF THE SILVER

ANODE IN 0.01 N HN0 -0.99 N AgN0 AT 50°C 3 3

i a f!L - E..1_ 0.01 i 0 5 (e2RT 2RT) i '1 (1 + a ) . - e a 0.396 CAg+ (ma·cm-2 ) (volts) 0.0 0.000 0.0000 0.000 0.1 0.005 0.0999 0.180 0.180 0.3 0.005 0.299 1.0 0.007 0.988 0.252 3.0 0.010 2.89 0.361 10.0 0.015 8.93 0.546 0.022 22.6 0.811 30.0 1.17 70.0 0.031 41.9 76

TABLE XLI

CURRENT DENSITY-OVERPOTENTIAL RELATIONSHIPS OF THE SILVER ANODE IN 1.0 AgN0 AT 50°C 3

i a i 0.01 1 0 5 a ( 1 + 0 396 +) . . ~ Ag (ma·cm-2) (volts) 0.0 0.000 0.0000 0.000 0.1 0.000 0.0999 0.000 0.3 0.000 0. 299 0.000 1.0 0.000 0.988 0.000 3.0 0.001 2.89 0.0359 10.0 0.002 8.93 0.0719 30.0 0.015 22.6 o. 546 70.0 0.030 42.1 1.13 77

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VITA

The author of this thesis~ Bing-wen Jong~ was born on

May 1~ 1935 in Kaohsiung~ Taiwan~ China. He received his elementary and high school education in Kaohsiung~ Taiwan. He entered the National Taiwan .university in September,

1954~ and graduated in June~ 1958 with a B.S. degree in Chemical Engineering. He attended the Chinese Army Chemical

School in October~ 1958~ and then served in the Chinese Army for a period of 18 months. Afterwards he was employed by the Taiwan Fertilizer Co. for one year and by the Chinese Petroleum Corporation for three years in Kaohsiung, Taiwan, China. He entered the University of Missouri at Rolla in

September~ 1964 as a graduate student in the Chemical Engineering Department.