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Kinetics of the dissolution of copper metal in some chelating systems under oxygen pressure Milants, Henri Yves Jean 1958

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KINETICS OF THE DISSOLUTION OF COPPER METAL IN SOME CHELATING SYSTEMS UNDER OXYGEN.PRESSURE by HENRI YVES MILANTS A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE. OF MASTER OF SCIENCE i n the Department of MINING AND METALLURGY We accept t h i s thesis as conforming to the standard required from candidates f o r the degree of MASTER OF SCIENCE. Members of the Department of Mining and Metallurgy THE UNIVERSITY OF BRITISH COLUMBIA October 195S. ABSTRACT An investigation was conducted on the d i s s o l u t i o n of copper metal i n aqueous solutions of ethylenediamine, glycine, a-alanine and p-alanine, under oxygen pressure. The kinetics of these reactions were investigated over a wide concentration range of the corresponding ionized species. The rate of d i s s o l u t i o n of copper, i n a l l solutions, has been found to be independent of the i n i t i a l copper concentration, the volume of the solu t i o n and the area of the copper sample. No intermediate products, i . e . , cuprous ions, were observed. Two regions were observed, having d i f f e r e n t dependence on oxygen pressure. In one, the rate depends on the f i r s t power of the oxygen pressure, and i s independent of the concentration of the chelating agent. In the other region, the reaction i s f i r s t order i n chelating agent and independent of oxygen pressure. The rate of the reaction i n t h i s second region appears to be chemically controlled at the copper surface. The neutral and charged species of the chelating agent were found to have independent rates. These two d i s s o l u t i o n reactions were found to be f i r s t order with respect to the concentration of the respective complexing species. The mechanism proposed by Halpern previously for the ammonia system was found to be applicable to the systems studied i n the present work. The rate constants f o r each chelating agent have been computed and appear to be related to the s t a b i l i t y constants. No s t e r i c effect was observed. I n p r e s e n t i n g t h i s t h e s i s i n p a r t i a l f u l f i l m e n t o f t h e r e q u i r e m e n t s f o r an advanced degree at the U n i v e r s i t y o f B r i t i s h C olumbia, I agree t h a t the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e f o r r e f e r e n c e and s t u d y . I f u r t h e r agree t h a t p e r m i s s i o n f o r e x t e n s i v e c o p y i n g o f t h i s t h e s i s f o r s c h o l a r l y purposes may be g r a n t e d by t h e Head o f my Department o r by h i s r e p r e s e n t a t i v e . I t i s u n d e r s t o o d t h a t c o p y i n g o r p u b l i c a t i o n o f t h i s t h e s i s f o r f i n a n c i a l g a i n s h a l l not be a l l o w e d w i t h o u t my w r i t t e n p e r m i s s i o n . : Ti The U n i v e r s i t y o f B r i t i s h Columbia, Vancouver S, Canada. Department o f TABLE OF CONTENTS Page INTRODUCTION » 1 Scope of the Present Investigation 3 EXPERIMENTAL . . . . • • • • . • • • « . Q O « . . « O . . . . . . . . 4 A. Preparation of the Copper Samples . . 4 B„ „ Apparatus <, . 5 C. Temperature Control 7 Do Geometry 7 E. Chemical Reagents and Solutions ..* 7 F. A n a l y t i c a l Procedures . . . . . 9 G. Measurements of the Rates 10 RESULTS AND DISCUSSION 14 A. Ethylenediamine System 14 1. Introduction 14 2. Effect of s t i r r i n g v e l o c i t y 14 3« Effect of surface area and solution volume 16 4. Effect of cupric ion on the d i s s o l u t i o n rate 17 5. Effect of oxygen pressure . . 17 6. Effect of ethylenediamine concentration 25 7. Effect of hydrogen and ethylenediaminium ion . . . . . . 25 8. Effect of NaOH on the rate 31 9. Summary of results 32 B, Ammonia System 33 TABLE OF CONTENTS (continued) page C. Glycine System • ... 33 1. Introduction <>. 33 2. Effect of oxygen pressure and glycinate concentration. . 35 3. Effect of oxygen pressure and concentration of glycine i n water 40 4. Effect of hydrogen ion 40 5. Summary • 43 D. Alpha-Alanine System j • 44 1. Introduction 44 2. Effect of oxygen pressure and alpha alaninate concentration • • 47 3. Effect of oxygen pressure and concentration of alpha alanine i n water . . . . . 47 4. Effect of hydrogen ion 52 5. Summary 52 E. Beta-Alanine System 54 1. Introduction 54 2. Effect of oxygen pressure and beta alaninate concentration « 54 3. Effect of oxygen pressure and concentration of beta alanine i n water 57 4. Effect of hydrogen ion 57 5. Summary 57 CONCLUSIONS 60 BIBLIOGRAPHY , . . . . 65 FIGURES Figure Page No. 1„ Diagram of Pressure Vessel 6 2. Calibration Curve f o r Copper Carbamate . . 11 3. Typical Rate Curve f o r the Dissolution of Copper 13 4. Plot showing the effect of Cupric Ion on the Dissolution Rate . . 18 5. Rate Curves for Copper at Various 0 2 pressures Ethylenediamine 0.475 Molar 19 6. Rate Curves f o r Copper at Various 0 2 pressures Ethylenediamine 0.1425 Molar o . 20 7. Rate Curves f o r Copper at Various 0 2 pressures Ethylenediamine 0.095 Molar 21 8. Rate Curves f o r Copper at Various 0 2 pressures Ethylenediamine 0.0712 Molar . . . . . . . . . . . 22 9. • Rate Curves f o r Copper at Various 0 2 pressures Ethylenediamine 0.0475 Molar 23 10. Effect of Oxygen Pressure on the Rate of Dissolution of Copper . . 24 11. Rate Curves f o r Copper with 0 2 Pressure Constant at Various Concentrations of Ethylenediamine . . . . 26 12. Effect of Ethylenediamine Concentration on the Rate of Solution,, . 27 13. Effect of H + Ion on the Dissolution Rate of Copper i n Ethylenediamine . . . . . 29 14. Variation of the Rate with the Concentration of Ethylenediaminium. 30 15. Effect of Concentration of Aqueous Ammonia on the Rate of 16. T i t r a t i o n Curve of Glycine with Sodium Hydroxide . 36 17o Rate Curves f o r Dissolution of Copper i n Glycinate Solution . . . 37 Figures (cont'd.) Figure ' Page  No. 18. Effect of Oxygen Pressure on the Rate of Dissolution of Copper i n Glycinate Solution 38 19. Effect of Glycinate Concentration on the Rate of Solution 39 20. Rate Curves f o r Dissolution of Copper i n Aqueous Glycine 41 21. Effect of Concentration of Aqueous Glycine on the Rate 42 22. T i t r a t i o n Curve of a-Alanine by Sodium Hydroxide 46 23. Rate Curves f o r the Dissolution of Copper i n a-Alaninate Solution . 48 24. Effect of Total a-Alaninate Concentration on the Rate of Solution . 49 25. Dissolution Rate of Copper i n Aqueous a-Alanine at Various Concentrations • « . • « . . a • o • o . o . o . . . 50 26. Effect of Concentration of Aqueous a-Alanine on the Rate of Solution 51 27. Rate Curves f o r the Dissolution of Copper i n p-Alaninate . . . . . 55 28. Effect of Total p-Alaninate Concentration on the Rate of Solution . 56 29. Effect of Concentration of Aqueous p-Alanine on the Rate of Solution o o e « * o o o o « « o » « o * o « o « 6 o « o o o * 5^ TABLES No. Page '.I , Determination of Reproducibility 12 I I . Effect of S t i r r i n g Velocity on the Dissolution Rate 15 I I I . Effect of Surface Area on Copper Dissolution 16 IV. Effect of Solution Volume on the Rate 16 V. Effect of Ethylenediaminium Ion on the Rate 28 VI. Effect of Sodium Hydroxide on the Rate (Ethylenediamine System) . 31 VII. Effect of Hydrogen Ion i n the Glycine-Water System 43 V I I I . Relative Concentrations of Glycinate, Glycinium and Zwitterion SpGCISS o « e « o c « o « o « o o • o e o o o o o * « o o • A-3 IX. Effect of Hydrogen Ion i n the Alpha Alanine-Water System . . . . 52 X. Relative Concentrationsof Alpha Alaninate, Alpha Alaninium and Zwitterion Species 53 XI. Relative Concentrations of Beta .Alaninate, Beta Alaninium and Zwitterion Species 59 XII. Summary Table of Rate Constants 62 XI I I . Correlation of Rate Constants and S t a b i l i t y Constants 63 ACKNOWLEDGEMENT The author i s indebted to the National Research Council of Canada f o r the f i n a n c i a l aid which enabled t h i s project to be carried out, and to the Union Miniere du Haut Katanga f o r permission to take a year's leave of absence. The author i s grat e f u l to the members of the Department of Mining and Metallurgy for t h e i r assistance throughout t h i s work, and is especially grateful to Dr. D.R. Wiles, who ably directed t h i s investigation. Thanks are extended to Dr. J. Halpern for many helpful discussions. KINETICS OF THE DISSOLUTION OF COPPER METAL IN SOME CHELATING SYSTEMS UNDER OXYGEN PRESSURE INTRODUCTION In the past two or three decades, much attention has been focused on corrosion of metals i n oxidizing aqueous media (Evans yUhlig ) but very l i t t l e has been directed towards the theo r e t i c a l aspects of the processes involved. The mechanisms of reactions of the d i s s o l u t i o n of a metal i n aqueous media, which involve simultaneous oxidation and di s s o l u t i o n 'of the metal, were not understood u n t i l recent years^*-' Only a few systems have been studied i n d e t a i l . E a r l i e r investigations of the kin e t i c s of the diss o l u t i o n of copper i n aqueous ammonia have provided c o n f l i c t i n g information about the order of the reaction and the effect of the different variables. Yamasaki"*" investigated the rate of dis s o l u t i o n of metallic copper i n ammonium hydroxide by rotating copper specimens at constant v e l o c i t y and supplying a steady current of p u r i f i e d a i r . His work indicated that auto-catalysis by cuprammonium was involved i n the d i s s o l u t i o n of copper, but showed the rate of di s s o l u t i o n to be independent of the concentrations of ammonia or of s a l t s i n solution. 2 Z a r e t s k i i and Akimov investigated the mechanism of corrosion of copper i n aqueous solutions of ammonium compounds ahd attributed the corrosion to three processes, the f i r s t being the electrochemical disso l u t i o n of copper with formation of cuprous ammonia complex, the second, oxidation of the complex to form cupric ammonia complex, and the t h i r d , electrochemical reduction of the cupric ammonia complex to cuprous ammonia complex. In a more recent study, Lane and McDonald, f a i l e d to observe any evidence of auto-c a t a l y s i s , and found instead a strong dependence of the rate on ammonia concentration. A l l the e a r l i e r investigators found that the rate increased with the s t i r r i n g v e l o c i t y , indicating dependence on transport of oxygen to the copper surface. As was pointed out more recently by Halpern,^ when d i f f u s i o n of a reactant controls the rate of a chemical reaction, the effects of other variables are masked and k i n e t i c results provide l i t t l e information about the mechanism of the reaction. Therefore i t appeared desirable to investigate the region where the transport of oxygen to the copper surface was s u f f i c i e n t l y f ast that i t did not affect the rate. This condition was achieved i n a study of the k i n e t i c s of the d i s s o l u t i o n of copper i n ammonia solutions recently conducted i n t h i s laboratory by Halpern^" and Fisher.'' In the course of these l a t t e r investiga-t i o n s , i t was shown that the observed rate of d i s s o l u t i o n of copper i n ammonia solutions was the sum of the rates of two independent reactions, f i r s t order with respect to the ammonia and ammonium ion concentrations, respectively. The rates were found to be independent of the oxygen concentration, provided the oxygen was present i n excess (so that the transport of oxygen to the copper surface did not l i m i t the rate of the reaction). On the basis of these r e s u l t s , the following mechanism f o r the d i s s o l u t i o n of copper was proposed: (1) Adsorption of dissolved oxygen on the copper surface: f a s t Cu + 1/2 0 2 — Cu . . . . 0 - 3 -(2) Reaction of an ammonia molecule or ammonium ion with the copper-oxygen complex on the surface: slow-. /NH3 fast Cu...O + NH3 — |^Cu+^/ J+ HOH — Cu(NH 3) + + + 2 OH" slow, /F* 3 » fast or Cu . , . 0 + NH^+ | C u + ^ H | + 0H~ Cu(NH 3) + + + 2 OH" * v ' - o * ; Following t h i s work on copper i n ammonia solution, i t seemed of interest to compare the dissolution rates i n ethylenediamine solution with those determined previously i n ammonia solutions and to extend the study to other complexing agents, p a r t i c u l a r l y chelating agents. Scope of Present Investigation Although any of several physical processes can control the rate of a heterogeneous reaction,(e.g., absorption of gaseous reactant by the solution, transport of dissolved reactant from the bulk of the solution to the s o l i d -solution interface, reaction on the surface or transport of the soluble desorbed products into the solution), only the chemical reaction rate aspects are of interest i n the present work. Therefore preliminary experiments were carried out to determine the conditions under which the reaction i s controlled by the rate of chemical action at the surface. This study was undertaken with the object of establishing the effect of several chelating agents of the amine type. I f the substance which combines with the metal contains two or more donor groups so that one or more rings are formed, the r e s u l t i n g structure i s said to be a chelate compound. Structural differences might show up r e a d i l y and s t i l l be comparable to the general pattern of mechanism proposed by 0 Halpern. In order to determine the effect of the -C group, several amino 'OH acids, e.g., glycine a-and B- alanine, were chosen and compared with the ethylenediamine and ammonia systems. The variables examined included copper c r y s t a l size differences, effects of oxygen pressure, s t i r r i n g rate, area, and chelating "agent concentration. The "temperature used was . 25.0°G. ' Other ' tempera-Were not studied*. ; L ,. Lv.t j.. c .. EXPERIMENTAL A. Preparation of the Copper Sample The work reported i n this . t h e s i s was performed on conductivity copper (about 99.97% Cu) but some check measurements were made on high-purity copper (99.999%). Preliminary measurements indicated that there was no measurable effect due to grain size (see below) or to the difference i n purity. High-purity copper dissolution specimens were prepared as follows: Ingots about 3/4*' diameter and 1'' length were melted by a high-frequency induction unit (H.F. furnace of 300 k i l o c y c l e s and power of the order of 1 kw) i n spectrographically pure graphite crucibles. The melting operation was performed i n vacuo; a Welch Duo-Seal vacuum pump was used. A vacuum of better than 50 microns was maintained i n the melting unit. The copper was allowed to s o l i d i f y i n the crucible. The s o l i d i f i e d ingots were cold worked by squeezing i n a press and annealed at either 500°C for f i v e minutes or 700°C for three minutes. The diameter of the r e c r y s t a l l i z e d grains was about 0.120 mm (A.S.T.M. non-ferrous grain size standard^) a f t e r the 500°C anneal and about 0.090 mm. after the 700°C anneal. The conductivity copper had an average grain diameter of 0.065 mm. Dissolution samples were prepared by mounting the copper specimens i n bakelite i n a metallographic specimen mounting press. This preparation was done i n order to insulate the copper samples e l e c t r i c a l l y and to leave - 5 -exposed only a single plane face of convenient size and shape. After mounting, the specimens were polished down to a No. 00 emergy paper. P r i o r to each experiment, the dimensions of the polished sections were caref u l l y measured by means of a microscope eyepiece calibrated to 0.01 mm. The area was computed from the mean of several measurements. The exposed surface area of di f f e r e n t copper samples studied ranged from 1.386 to 2.248 cm. After measurement, the samples were stored i n a desiccator u n t i l needed. Alternative methods of preparing the polished surface, such as etching with ammonium persulphate-ammonium hydroxide or with HN03, were found to give i d e n t i c a l d i s s o l u t i o n rates to those experiments i n which the copper sample was not pre-etched. In the actual experiments, the ammonium persulphate-ammonium hydroxide etch was used. B.„ Apparatus The investigation reported i n t h i s thesis was conducted i n an auto-clave (Figure 1) designed f o r working at pressures up to 8 atmospheres. The autoclave was fabricated from 316 stainless s t e e l but a titanium l i n e r was used, which was shown to be inert to the solutions used. The mounted copper specimen was held by means of a stainless s t e e l rod, the l a t t e r being screwed into the l i d of the autoclave. Agitation was provided by two turbine-type impellers operating on a single shaft. The ro t a t i o n a l speed of the impeller could be selected by choosing the diameter of the drive pulley on the motor and could be varied between 500 and 820 R.P.M. A solution volume of two l i t e r s was normally used. Oxygen was supplied from a standard cylinder and the desired pressure held constant by means of a standard oxygen pressure regulator f i x e d on the - 6 -Figure 1. Schematic diagram of the autoclave and heating control system. A. Shaft B. Thermometer w e l l and.thermoregulator C. Cooling c o i l D. Sampling tube E. Copper sample i n bakelite mount F. Impeller G. Pump H. Relay I. Heater J. Thermostat - 7 -cylinder. An additional pressure check was maintained by a low-pressure gauge mounted on the oxygen supply line. C, Temperature Control. In order to keep the temperature constant, a cooling coil which was fitted inside of the autoclave was connected with a thermostatically controlled water bath. Water at the desired temperature was forced through the cooling line by means of a centrifugal pump (see Figure l ) . A mercury contact thermo-regulator set in the stainless steel temperature well of the autoclave was used to control the temperature of the water bath. The temperature in a l l the experiments was maintained at 25*0,1°C. D, Geometry One of the aims of this work was to compare the dissolution rates in ethylenediamine solution with those determined previously in ammonia solutions.-* Thus, i t was of primary importance to determine whether the geometry (i.e. effective contact between the copper surface and the solution) of this system was similar to that used by Fisher. The geometry w i l l be a function of the efficiency of agitation, both at the copper surface and in the bulk of the solution, also of the effectiveness of the contact between the solution and the oxygen atmosphere. Preliminary measurements indicated that there was no measurable effect due to difference in geometry. In fact, the results obtained in the ammonia system were found to give good agreement with those of Fisher.^ E, Chemical Reagents and Solutions. The glycine and a-and 3-alanine- used in this work were reagent grade, ammonium-free,and supplied by Nutritional Biochemicals Corp, ' The ammonia and sodium hydroxide were chemically pure, supplied by Nichols Chemical Company - 8 -(Baker and Adamson Reagent Grade). The ethylenediamine, supplied by Carbide and Carbon Chemicals Company was 98.98$ pure, the remaining 1 . 0 2 $ presumably being water. In order to detect any disturbing impurities i n the l a t t e r reagent, the following tests were performed? ( 1 ) Detection of heavy metal cations. The ethylenediamine was examined for heavy metals by X-ray f l u o r -escence. The technique used was to evaporate a 2 0 0 ml sample of ethylene-diamine on a small amount of ' C e l i t e ' , which had previously been c a r e f u l l y washed with perchloric acid and water. No heavy metals were detectable. ( 2 ) Fractional d i s t i l l a t i o n The ethylenediamine was f r a c t i o n a l l y d i s t i l l e d according to the method of Clarke and Blant i n order to obtain 1 0 0 % purity. This d i s t i l l a t i o n was conducted i n a Todd S c i e n t i f i c Co. column. The f i n a l product obtained was examined using the r e f r a c t i v e index and b o i l i n g points as c r i t e r i a . The figures obtained were, n 1 . 4 5 4 0 , B.P. 116°C, i n exact agreement with published data f o r pure anhydrous ethylenediamine. (3) Comparison of dissolution rates. Dissolution rates were compared i n solution using a. the 98.98$ ethylenediamine b. the 1 0 0 $ d i s t i l l e d ethylenediamine c. the 1 0 0 $ reagent grade (Eastman Kodak) ethylenediamine. Differences i n the rates observed were about 0 . 5 $ , less than the expected "'V, experimental deviation. From the results of these tests i t was concluded that no disturbing impurities were present. - 9 -Solutions were prepared by d i l u t i n g a measured volume of reagent to 2,0 l i t e r s with d i s t i l l e d water. In the case of the ethylenediamine and ammonia systems, the normality of these solutions was checked by t i t r a t i o n with standard s u l f u r i c acid to a methyl red end point. In order to avoid a s a l t effect as observed by Halpern, sodium perchlorate was added to the solution i n the autoclave to give a t o t a l s a l t concentration of 0.1 moles per l i t e r . F. A n a l y t i c a l Procedures. Analysis of copper i n samples was done c o l o r i m e t r l c a l l y using the sodium diethyldithiocarbamate method. ' Golorimetric determination of copper by means of the yellow chelate compound formed by sodium diethyldithiocarbamate with cupric ions, has been described by many authors. In most cases, measure-ments are made aft e r extracting the coloured complex into carbon tetrachloride 13 or amyl alcohol. Since effective extraction may be very time consuming, several methods of avoiding i t have been devised f o r use i n cases when i n t e r -f e r i n g elements, e.g. i r o n , are absent. Since the s t a b i l i t y of the copper carbamate complex i s affected by daylight, a small amount of gum arabic was added and a clear golden-brown colour w e l l s t a b i l i s e d by the protective c o l l o i d , ^ . "1.4,15, 16 was obtained. The carbamate-gum arabic solution was prepared as follows; 0.2 grams of carbamate dissolved i n 200 ml. of d i s t i l l e d water was mixed with 5 grams of gum arabic of 1 ml. toluene i n 1000 ml. d i s t i l l e d water. This solution (henceforth abbreviated CGA) was f i l t e r e d and stored i n the dark, The a n a l y t i c a l solutions were prepared i n an a r t i f i c i a l l y lighted room to avoid the effect of daylight. The s t a b i l i t y of the chelate compound i s very dependent on the pH. However, i t was found that the o p t i c a l density was not affected between pH 7.5 and 9.2. Samples for analysis were prepared as follows: To an aliquot of the - 1 0 -sample taken from the autoclave, 5 ml. of ammonium c i t r a t e was added to complex accidental traces of i r o n , and the pH adjusted to 9 . 0 with 1 : 1 ammonium hydroxide. F i n a l l y , 1 0 ml of CGA solution was added and the mixture d i l u t e d with d i s t i l l e d water to 5 0 ml i n a volumetric f l a s k . Depending on the copper concentration, the aliquot varied from 2 to 1 0 ml. The concentration of the coloured complex may be re a d i l y determined by spectrophotometric comparison with a series of s i m i l a r solutions of known concentration. For t h i s purpose, a Beckman model DK-2 Ratio-Recording Spectrophotometer was used to determine the o p t i c a l density of the solutions at a wavelength of 4 3 7 millimicrons. The o p t i c a l density was found to be proportional to the copper concentration over a s u f f i c i e n t range to allow convenient analysis. The c a l i b r a t i o n curve i s given i n Figure 2, The res u l t s obtained by t h i s method could be duplicated w i t h i n less than ± 0 . 5 % . The presence of ethylenediamine, glycine, or aV and (3r alanine, ,was found'not to affectiAthe c a l i b r a t i o n curve; : '\'\ G. Measurement of the Rates of Dissolution. The method used f o r measuring the dissol u t i o n rates of copper i n the various solutions was as "follows:: The autoclave, charged with solution and sample, was flushed with nitrogen u n t i l the desired temperature was reached. The nitrogen was then flushed out three times with oxygen and the desired oxygen pressure applied to the solution. Zero time was taken when the.desired oxygen pressure and temperature were obtained. To follow the course of the reaction, 4 0 ml samples were withdrawn by means of the sampling tube at 1 5 minute in t e r v a l s and analysed f o r copper. From the known surface area of the sample and the known volume of the solution at the time of each sampling, the t o t a l amount of copper dissolved per unit area was computed. The experiments lasted between 1 . 5 and 2 hours. During that time, approximately 0 . 5 0 to 2 . 0 0 x 1 0 ~ 3 - 11 -I I I I 0.000 0.001 0.002 0.003 0.004 0.005 Copper concentration, (grams liter -"*") Figure 2. Calibration curve f o r analysis of copper by the carbamate method. - 12 -moles of copper was dissolved. This amount was i n s u f f i c i e n t to cause any appreciable change i n the concentration of the chelating agent, through the formation of a chelate or amine: complex. A t y p i c a l rate curve f o r the dissolution of copper i n ethylenediamine solutions under oxygen pressure i s shown i n Figure 3. The amount of copper dissolved was d i r e c t l y proportional to the reaction time. The l i n e a r i t y of the rate curves indicates a zero-order reaction at constant oxygen and ethylenediamine concentration. Rates.of solution were calculated from the slopes of the rate plo t s . A series of experiments was carried out at constant ethylenediamine concentration, oxygen pressure, agitation and temperature, i n order to deter-mine the r e p r o d u c i b i l i t y of the experimental technique. The results of these tests are summarized below: TABLE 1 . . .. Determination of Reproducibility Run No. . Rate of Solutibn (MG/cm.2/hr.) Deviation f0 1 22.5 -0.44 2 22.3 -1.33 3 23.0 +2.20 Average 22.6 1.3 Conditions: ethylenediamine, 0.1 M; Temp. 25.0°Cj S t i r r i n g v e l o c i t y 775 RPM; Oxygen pressure 5.0 atm. The rates were generally found to be reproducible to within 1 or 2% i n duplicate experiments. In the following sections are described the results of experiments designed to determine the influence of the dif f e r e n t variables on the reaction rate. - 13 -15 30 45 50 75 Time (minutes) Conditions: 0.1 M ethylenediamine - 775 R.P.M - 25°C 0.1 M NaClO^ - Pressure oxygen 6.5 atm. Figure 3. Typical curve f o r the d i s s o l u t i o n of pure copper. RESULTS AND DISCUSSION - 14 -A. Ethylenediamine System 1. Introduction Ethylenediamine i s a strong chelating agent, and with cupric ions forms the complex: CH2 - NH2 y NH2 - CH2 1 I ^ C U I CH2 - NH2 ' NH2 - CH2 It i s also a strong base, and r e a d i l y takes on either one or two hydrogen ions. 17 18 ± The equilibrium constants ' f o r the formation of the various species are: [Cu e n + + ] - 10 * (kj) (!) (1:1 chelates) [ C u ^ F e n ] C enH 2 + + J - 1 0 " 7 ' 4 9 (2) tenH^[H +3 [enH +] - 1 0 " 1 0 , 1 7 ' (3) 4 For convenience, ethylenediamine w i l l be abbreviated 'en* i n formulae. S i m i l a r i l y one has for the acid species enH + and enH 2 +t. Most of the present work was done at about pH 11.5 at which the ethylenediamine i s present almost e n t i r e l y as the free base. The corrosion action of the ethylenediamine solutions may be formulated according to the equation: Cu + 2 en + 1/2 0 2 + H 20 — Cu e n 2 + + + 2 OH" (4) No intermediate products such as cuprous ions were detected during the course of the experiments. 2. Effect of s t i r r i n g v e l o c i t y I t i s w e l l known that when d i f f u s i o n controls the rate of the - 15a -A series of tests was performed with the object of determining whether the reaction rate i s agitation dependent i n the low-pressure region, as was found by Halpern.^ The res u l t s are given i n Table Ha. TABLE I l a . Effect of S t i r r i n g Rate on Dissolution  Rate i n the Low Pressure Region. S t i r r i n g rate (R.P.M.) Reaction rate (Mg/cm.2/hr.) Deviation % 550 18.6 +0.8 675 18.3 -0.8 775 18.3 -0.8 810 18.6 +0.8 Average reaction rate 18.45 0.8% Conditions; Temp. 25°C., ethylenediamine 0.475 M/l. Na perchlorate 0.1 M, oxygen pressure 3,7 a t m . The results show c l e a r l y that the rate i s not dependent on the s t i r r i n g rate. This i s at variance with Halpern*s results f o r the same pressure region. No reason has been found f o r t h i s discrepancy. One possible explanation i s that the sample i n the present investigation was not located i n the same place as was Halpern*s, but was closer to the t i p of the impeller blades, where agitation i s expected to be more v i o l e n t . It i s possible, then, that the agitation i n t h i s p o s i t i o n reaches a maximum at f a i r l y low s t i r r i n g rates, and that further increase i n s t i r r i n g rate causes no increase i n turbulence near the sample surface. - 15 -chemical reaction, the k i n e t i c results give l i t t l e information about the mechanism of the reaction. When such a physical process i s predominant, the effect of other variables on the rate i s d i f f i c u l t to be determined. It i s , rather, important to f i n d a range of conditions such that the rate i s independ-ent of s t i r r i n g and therefore controlled by factors other than the transport of reactants or products. I t has already been shown by Halpern that, f o r the ammonia system at low ox/gen pressures the rate increases with s t i r r i n g v e l o c i t y , whereas at higher oxygen pressures such a dependence i s not observed. Thus, i t was of primary importance to determine at what s t i r r i n g v e l o c i t y the present work should be done. The results of tests at di f f e r e n t s t i r r i n g rates are summarized i n Table I I . TABLE I I Effect of S t i r r i n g V elocity on Dissolution Rate S t i r r i n g Rate (R.F.M.) Reaction Rate (Mg/cm.2/hr) Deviation 550 18.6 -0.16% 675 18.3 -1.77% 775 19.0 +2.00% 810 18.6 -0.16% Average Reaction Rate: 18.63 1.04% Conditions: Temp. 25°C, ethylenediamine 0.075 M, Na perchlorate 0.1 M, oxygen pressure 6.5 atm. The data given i n Table I I show that the reaction rate i s not limited by transport of reactants w i t h i n the solution. As w i l l be discussed l a t e r , simpler c r i t e r i a were made available as the systems studied became better understood. 1. Independence of oxygen pressure indicates absence of control by oxygen transport. 2. Linear dependence of the rate on the concentration of complexing agent: I t was shown by Halpern and corroborated by the present work that i n - 16 -the absence of transport control, the reaction i s f i r s t order i n the complexing agent. This w i l l be shown more c l e a r l y i n a l a t e r section. In order to eliminate any extraneous e f f e c t , a s t i r r i n g rate of 775 RPM was used, 3. Effect of surface area and s o l u t i o n volume. Preliminary measurements indicated that effects due to the change of volume and to changes i n the surface area upon the rate of solution were not s i g n i f i c a n t . Results of these experiments are given i n Table I I I and IV. TABLE I I I Effect of Surface Area on Copper Dissolution Surface area of copper (cm.2) Rate of solution (Mg/cm.2/hr)' Deviation (*) 1.386 19.00 +2.86 1.606 18.00 -2.57 2.066 18.45 -0.135 2.243 18.50 +0.135 Average rate 18.475 1.42 Conditions: Temperature 25°C, ethylenediamine concentration 0.5M, oxygen pressure 3.8 atm, s t i r r i n g v e l o c i t y 775 RPM. TABLE IV. Effect of Solution on the Rate Volume of Solution (ml) Concentration (Mg/l/hr.).. Rate (Mg/hr) Rate of Solution: (Mg/cm.Vhr.) Deviation (%) 1,500 2,000 2,456 33.6 24.1 19.2 40.1 40.5 40.2 19.20 19.00 19.30 +0.26 -0.75 +0.75 Average rate of solution 19.15 0.585 These results show that the r a t e of dissolu t i o n i s independent of the volume - 17 -and of the apparent surface area of the copper. The fact that the rate of reaction was always invariant with time indicates that the e f f e c t i v e area of the copper remains constant during the experiment. This fact also gives evidence that the rate c o n t r o l l i n g step occurs at the surface i t s e l f and not i n the solution, since the volume of the solution necessarily diminished during the course of each experiment, 4. Effect of cupric ion on the d i s s o l u t i o n rate The zero-order reaction indicated i n a l l cases shows that the cupric ions dissolved during the corrosion process i n the form of a chelate compound such as C\i(en^'+', do not affect the rate i t s e l f . As a further t e s t , an experi-ment was done i n which 0.015 g/1 of copper perchlorate was added at the s t a r t . The r e s u l t s , plotted i n Figure 4, show that the rate of d i s s o l u t i o n i s the same (within .2%)) as when no cupric ions were added i n i t i a l l y . 5. Effect of oxygen pressure Some t y p i c a l rate curves showing the effects of varying the oxygen pressure from 2.3 to 7.8 atmospheres are shown i n Figures 5 to 9 f o r various concentrations of ethylenediamine. The complete system i s summarized i n Figure 10, where the rate of solution f o r d i f f e r e n t ethylenediamine concentrations i s plotted veraus the oxygen, pressure. At low.pressures, the rate of solution i s proportional to the pressure and independent of the reagent, i n t h i s case, the ethylenediamine. Thus i n t h i s region, the rate may be controlled by the transport of oxygen to the surface of the copper. At higher pressures, however, a point i s reached, where the rate of sol u t i o n levels o f f . In t h i s region, the rate i s independent of the - 18 -15 30 45 60 Time (minutes) Conditions: ethylenediamine 0.1 M - 775 R.P.M., temperature 25°C., pressure 6.45 atm. Figure 4. Plot showing the effect of cupric ions on the dis s o l u t i o n rate. - 19 -T j j ! r 15 30 45 60 • 75 Time (minutes) Conditions: Temperature 25°C, s t i r r e r v e l o c i t y 775 R.P.M., NaClO^ 0.1 M. Figure 5. Rate curves f o r the d i s s o l u t i o n of pure copper i n ethylenediamine solution at various 0 2 pressures. ~ 20 -Time (minutes) Conditions: Temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., NaC104 0.1 M. Figure 6 . Rate curves f o r the d i s s o l u t i o n of pure copper i n ethylenediamine solution at various 0 2 pressures. - 21 -Time (minutes) Conditions: Temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., NaC10 4 0.1 M. Figure 7. Rate curves f o r the d i s s o l u t i o n of pure copper i n ethylenediamine solution at various 0 2 pressures. - 22 -35 30 • 25 20 -15 10 Oxygen pressure O 7.8 atm. O 3.7 " O 2.4 [enj: 0.0712 Molar 30 45 60 75 Time (minutes^ Conditions: Temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., NaCIO/, 0.1 M. Figure 8. Rate curves f o r the di s s o l u t i o n of pure copper i n ethylenediamine solution at various 0 2 pressures. - 23 -Time (minutes) Conditions: Temperature 25°C, s t i r r e r v e l o c i t y 775 R.P.M., NaClO^. 0.1 M. Figure 9 . Rate curves f o r the dissol u t i o n of pure copper i n ethylenediamine solution at various oxygen pressures. - 2k -0 2 pressure ( atmospheres) Conditions: Temperature 2 5 ° C , stirrer velocity 775 R . P . M . , NaC104 0.1 M . Figure 10. Effect of oxygen pressure on the rate of dissolution of copper. - 25 -concentration of oxygen and appears to be controlled by the chemical reaction at the copper surface and w i l l be shown i n the subsequent section to be zero order with reference to oxygen pressure. The data presented i n Figure 10 demonstrate that the dual control found by Halpern f o r the copper ammonia system i s v a l i d i n the present system. 6. Effect of t o t a l ethylenediamine concentration The rate of di s s o l u t i o n of copper was investigated using a series of solutions containing ethylenediamine i n concentrations ranging from 0,05 to 0.5 M/l(see Figures 5 to 10). Since the region of chemical control was the more intere s t i n g from the standpoint of the present work, a high p a r t i a l pressure of oxygen was maintained above the sol u t i o n . The results are shown i n Figures 11 and 12 and the data indicate quite c l e a r l y that at low concentration of the chelating agent, the reaction i s f i r s t order with respect to the concentration of ethylenediamine. Zero-order dependence i n oxygen was demonstrated i n Figure 10. These factors can be expressed as follows: Rate - k [ e n J 1 [o 2 ]° (5) The experimental rate constant K®^ 3, was calculated from Figure 12 and found t 0 b e : exp _2 -1 _1, k n = 245 Mg Cu cm. A Hr M 71. The deviation from l i n e a r i t y at high ethylenediamine concentration could r e f l e c t the t r a n s i t i o n to transport control (Figure 12). 7. Effect of hydrogen ion and ethylenediaminium ion. The effect of ethylenediaminium ion (enH +) was investigated, keeping - 26 -I I I Time (minutes) Conditions: Oxygen pressure 7.8 atm., s t i r r e r v e l o c i t y 775 R.P.M., temperature 25°C. Figure 11. Plot showing the rate of solution of copper i n aqueous ethylenediamine at various concentrations of en. - 27 -T i 1 : 1 } 0.025 0.05 0.075 0.100 0.125 en concentration (molar) Conditions! Oxygen pressure 7 .8 atm., temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., pH 11.5. Figure 12. Effect of t o t a l en concentration on the rate of solution. - 28 the concentration of ethylenediamine constant. This was done by adding excess ethylenediamine and an equivalent amount of HCIOz,. to the reference system of ethylenediamine. The series of solutions i s described i n Table V. TABLE V. Effect of Ethylenediaminium Ion on the Rate en HCIO4. en enH + NaClO^ Rat©2 (moles/1) (moles/l) (moles/l) (moles/l) (moles/l) (Mg/cm. /hr) 0.0475 0.00 0.0475 0.00 0.1 11.7 0.0575 0.01 0.0475 0.01 0.09 17.0 0.0675 0.02 0.0475 0.02 0.08 22.0 0.0775 0.03 0.0475 0.03 0.07 27.0 The equation representing the equilibrium f o r the ionization of ethylene-diamine can be rewritten i n the form; [enH+J - 10^°- 1 7 - I O 3 ' 9 3 l^W) *w or [enH + D = c f e n l "FT I t follows that as long as an appreciable concentration of the free base ethylenediamine remains i n solution, any added H + ions are almost qua n t i t a t i v e l y transformed into enH + ions. The increase i n enH + concentration i s of course associated with a proportional (although very small) increase i n H + concentra-t i o n . The results of t h i s investigation are shown i n Figures 13 and 14. It i s seen that the rate of solution increases l i n e a r l y with increasing enH +. I t appears, thus, that the t o t a l rate of dissolution i s made up of the contributions of two separate reactions. These two reactions are f i r s t order with respect to the concentration of en and enH +',respectively. The t o t a l d i s s o l u t i o n rate can be expressed as follows: R T = R e n + R e n H+ (6) or _ R T - k e n(en) - k ^ (enH+) (7) - 29 -T Ethylenediaminium concentration Time (minutes) Conditions: S t i r r e r v e l o c i t y 775 R.P.M., oxygen pressure 7 . 8 atm., ethylenediamine concentration 0.0475 molar. Figure 13. Effect of copper i n hydrogen ion on the dis s o l u t i o n rate of ethylenediamine. T 30-20 0.01 0.02 0.03 0.04 Ethylenediaminium (moles/1.) Conditions: Oxygen pressure 7.8 atm., s t i r r e r v e l o c i t y 775,R.P.M., ethylenediamine concentration. 0.0475 molar. Figure 14. Variation of the rate with concentration of ethylenediaminium (enH+)„ 31 -From the slope of the straight l i n e i n Figure 14, the f i r s t order rate constant f o r the d i s s o l u t i o n of copper i n ethylenediaminium solutions i s found to be 515 1 Ig/ cm.2/:hr/:rtdle's/l. For a t o t a l e l e c t r o l y t e concentration of 0.1 mole/liter, equation 7 can now be written: R T = 245 [en] + 515 [enH+] (8) This equation i s v a l i d i n the region of chemical rate control, where oxygen i s present i n excess. 8. Effect of NaOH on the rate. A few experiments were also performed to investigate the effect of sodium hydroxide on the rate of solution of copper i n ethylenediamine solution. A decrease of the rate was observed. The data are summarized i n Table VI. TABLE VI Effect of Sodium Hydroxide on the Rate (Ethylenediamine System) No. en NaOH NaGlO^ moles/l moles/l moles/l Mg/cm. /hr. 1 0.095 0.00 0.1 23.00 2 0.095 0.01 0.09 20.00 3 0.095 0,02 0.08 19.40 4 0.095 0.04 0.06 18.50 A s l i g h t excess of hydroxyl ion w i l l transform any enH + into en species and therefore decrease the r a t e . A further decrease (Experiments Nos. 3 and 4) i s probably due to a passivating effect of the copper suface by 0H~ ion. 9 . Summary of results f o r the ethylenediamine system. - 32 -The investigation performed and the results obtained i n the ethylene-diamine system may be summarized as follows; (a) The rate of dissolution of copper has been found to be independ-ent of i n i t i a l copper concentration, of the volume of the solution and of the area of the sample. (b) No intermediate products, e.g., cuprous ions, were observed. (c) Two regions of rate control have been found: the low oxygen pressure region i n which oxygen transport may l i m i t the rate, and the high oxygen pressure region i n which the rate i s chemically controlled at the copper surface. (d) In the region of chemical control (at low concentrations of the chelating agent), the reaction i s f i r s t order i n ethylenediamine and zero order i n oxygen. (e) pH effect. Any added H + i s almost quantitatively transformed into enH + and increases the r a t e . The rate f o r enH + was found to be f i r s t order with respect to the concentration of enH +, and independent of oxygen concentration provided that i t s transport to the surface of copper does not l i m i t the reaction rate. A decrease of the rate was observed by increasing the pH with sodium hydroxide. At t h i s point, i t seems evident that the general pattern of the experimental re s u l t s , which were obtained i n the course of the present work with the ethylenediamine system, i s very similar to that of the ammonia system studied by Halpern^ and Fisher.^ So as to be better able to discuss the interpretation and the mech-anism of the reaction i n the region i n which the rate i s chemically controlled at the copper surface, several other systems were compared i n that same chemically controlled region. - 33 -B. Ammonia System A few experiments were performed i n the ammonia system so as to be able to determine and to compare i t s rate constant k j j H ^ with the one found i n the other systems. The results obtained are summarized i n Figure 15. Rate measurements made at different ammonia concentrations i n the region of chemical control give a f i r s t order reaction kinetics i n ammonia and zero-order reaction k i n e t i c s i n oxygen. The experimental rate constant, kjjjj calculated from the li n e a r slope of Figure 15 was found to be: 5 N H 3 T The results obtained were found to be i n good agreement with those of Fisher,' C. Glycine System. 1. Introduction Glycine - NH2.CH2„COOH i s an amino acid. Since the publication of 17 Bjerrum's paper,, i t has been generally agreed that the ami.no acid, i n water solutions, e x i s t l a r g e l y as amphions and not as undissociated molecules. In 9 10 11 aqueous solution the 'Zwitterion' * * NH3+CH2C00~, can i t s e l f act as an acid and form a glycinate ion. Glycinate ion i s a bidentate chelating agent which forms one covalent 12 and one ionic bond with the cupric ion r e s u l t i n g i n five-membered r i n g . The metal chelate formation may be represented as follows: „ GO— 0 ^  XNH ? — CH? C u + + + 2 NH2 CH2-C00- 5 ^ , ^ Cu' 2 i 2 CH2-NH2 N 0 — CO It has been shown^7>1^A9 that, i n general, any additional hydrophilic groups such as hydroxyl ion or ionic group would increase the, s o l u b i l i t y of amino acids •in water. ' ' • • 0.5 1.0 Ammonia (Moles/l.) Conditions: Oxygen pressure 6.45 atm., temperature 25°C, s t i r r e r v e l o c i t y 775 R.p.M. Figure 15. Effect of concentration of aqueous ammonia on the rate of so l u t i o n . - 35 -For convenience, the different species w i l l be abbreviated i n formulae as follows: the cation species NH3+CH2C00H = G l + the zwitterion species NH3+CH2C00~ = Gl± the anion species NH2CH2C00~ *» Gl~ 18 The chelate s tab i l i ty constant with glycinate and cupric ions would be expressed as follows: [Cu Gl +J - 1 0 8 * 6 2 (ki) chelates 1:1 [Cu+U par] The ionization cohstants^>18 f o r glycine are given by: [ G 1±]H - I O " 2 ' 3 0 8 and [G1-] [H+] = IO " 9 - 7 8 In a strongly acidic solution the substance is present as the posi t ively charged ion (a) , . V - H + . - H + H3N+CH2C00H ^ ± H3N+CH2C00- ^ H2NCH2C0CT (a) (b) ' (c) The neutral dipolar ion (b) i s formed i n aqueous solution. In order to work with a well defined and completely dissociated species (c), sodium hydroxide was added to the glycine solution up to the second equivalent mid-point (pH 11.3)• (It was found that any excess of hydroxyl ion, beyond the equivalent mid-point, decreases the dissolution rate, possibly by the formation of an insoluble layer on the copper surface^-). A typ ica l t i t r a t i o n curve of glycine i n aqueous solution with sodium hydroxide i s shown i n Figure 16. 2. Effects of oxygen pressure and glycinate concentration. The effects of varying the oxygen pressure of the glycinate concentra-t ion were investigated and are summarized i n Figures 17, 18, and 19. At low oxygen pressure, the rate was found to be proportional to the pressure and independent - 36 -i i i i / / ml. NaOH (1.2 M.) Figure 16 T i t r a t i o n curve of glycine by sodium hydroxide, showing the equivalent points for the carboxylic and amino groups, respectively. - 37 -Time (minutes) Conditions: Oxygen pressure 6.45 atm., temperature 25 ° C , stirrer velocity 775 R.P.M. Glycine plus equivalent amount of NaOH, pH 11.3. Figure 17. Rate curves for the dissolution of copper in glycinate solutions. - 38 -0 2 Pressure (atmosphere) Conditions; Temperature 25 °C, s t i r r e r v e l o c i t y 775 R.P.M., pH adjustment to 11.3 with NaOH. Figure 18. Effect of oxygen pressure on the rate of di s s o l u t i o n of copper. - 39 -0.1 0.2 0.3 Glycinate (moles/1.) Conditions? Oxygen pressure 6.45 atm., temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., pH adjustment to 11.3 by NaOH. Figure 19. Effect of glycinate concentration on the rate of solution. - 10 -of the chelating agent concentration„ At high oxygen pressure the dissolu t i o n rate was found to be independent of oxygen pressure, thus chemically controlled at the copper surface. In t h i s l a t t e r region, the rate i s f i r s t order i n glycinate and zero order i n oxygen pressure as shown i n Figure 19. As i n the ethylenediamine system, the t o t a l rate f o r copper d i s s o l u t i o n i n glycinate solution can be expressed as follows? Rate = k [ G l y c i n a t e ] 1 [o2]° The experimental constant k^ F^ , was calculated from the l i n e a r portion of Figure 19 and was found to be? exp o kG1- = 50 Mg/Gu/cm. /hr/Mrrj-A. 3 . Effect of oxygen pressure and concentration of glycine i n water. The rate of dissolution of copper was investigated using a series of solutions containing glycine i n water i n concentration ranging from 0.1 to 0.3M. These experiments were performed with no sodium hydroxide (Figures 20, 21). The pH of glycine i n aqueous solution i s 6.0. From it-hisllatter, value and pK-j_ constant i t can be seen, that the species present i s almost completely i n the form of the zwitterion Gl± « l O 3 ' 6 2 G l * In the high pressure region, where the rate i s chemically controlled at the copper surface, the reaction i s f i r s t order i n glycine (Gl~) and the exp experimental rate constant calculated from Figure 21 i s found to bet e xP o T -1 , kGl± = 3 2 ' 5 M g Gvi°'cm° H M / l e 4. Effect of H + ion The effect of H + ion was investigated, keeping the concentration of » V I * Time (minutes) Conditions! Oxygen pressure 6.45 atm., s t i r r e r v e l o c i t y 775 R.P.M., temperature 25°C, pH 6.0. (No sodium hydroxide has been added). Figure 20. Plot showing the d i s s o l u t i o n rate of copper i n aqueous copper glycine at various concentrations. - 42 -Glycine concentration (moles/l.) Conditions: Oxygen pressure 6.45 atm., temperature 25 °C , , stirrer velocity 775 R.P.M., pH 6.0. Figure 21. Effect of concentration of glycine on the rate. - 43 -molecular glycine in water constant. This was performed by adding excess of glycine and an equivalent amount of HCIO^ to the reference system. In the present case the concentration of glycinium ion (Gl+) in solution has increased one hundred to one thousand fold over the glycine water system, the concentration of free amine being decreased proportionately. The results obtained are summarized in Table VII, TABLE VII Effect of Hydrogen Ion in Glycihe-Water System Glycine (moles/l) 0.10 0,11 0.12 HCIO^ (moles/l) 0,00 0,01 0,02 NaC104 (moles/l) 0.10 0.09 0.08 PH 6,0 4.0 2.6 Rate (Mg/cm//hr) 3.4 0,2 0 The data show that the rate decreases very rapidly to almost zero with increased H + ion concentration, 5. Summary The investigation performed in the glycine system is summarized in Table VIII. TABLE VIII Relative Concentration of Glycinate. Glycinium and Zwitterion Species. Total Glycine (Moles/l) PH Gl-(moles/1) Gl± (moles/l) G1+ (moles/l) Plot Rate (Mg/cm.2/hr/M) 0,100 0.100 0.110 0.120 11.3 6.0 4.0 2.6 0.0972 0(10-4) 0,000 0,000 0.0028 0,100 0.108 0.073 0,000 0(10-4) 0.002 0.047 50.0 32.5 0.2 0 From data shown in Table VIII i t is clear that in the presence of significant concentration of glycinium (Gl+) the copper surface is passivated. In the - %k-absence of such passivating effect i t appears that the rate of di s s o l u t i o n can be expressed by the equation: R - k G 1 ± [Gl±] + k Q 1- [ G l - J The values computed f o r the rate constants are: For Gl±(ph 6) R = 3.25 = kQ 1 ±[o.l] + k G 1-["0j and k G 1 ± = 32.5 Mg Cu cm. - 2 H-1M£J-+/1. For Gl"(pH 11.3) R = 5.0 = 32.5 [0.0028] + kG 1_fo.0972j k G 1 _ = 50.5 Mg Cu cm. - 2 H ' V ^ . / l . D. Alpha-Alanine System  Introduction The anionic form of a-alanine forms two bonds with cupric ion re s u l t i n g i n five-membered r i n g . The ion i z a t i o n of the l a t t e r amino acid i s seen by considering the ampholyte cation as a dibasic acid: -H -H CH3 —.CE-—COCH CH3 — CH^COO" CH3 -~ CH — COO" (1) NH 3 + PKX=2.34 N i 3 + ' pK2-9.87 NH2 The neutral form written as an amphion indicates that the io n i z a t i o n of the carboxyl group probably occurs when an acid solution containing the cation i s pa r t l y neutralized.by hydroxyl ion, while the removal of hydrogen ion from the substituted amino group probably occurs when a l k a l i i s added to the neutral 9,10,11,17 form.' 17 18 The i o n i z a t i o n constant^ * for a-alanine can, thus, be expressed - 45 -by: 4 rAla±]LH-tl IO" 2' 3 4 (2) and [Alg-JfH+J - 1 0 7 9 , 8 7 (3) [Ala + J L A l a ±J ± For convenience, the zwitterion, the a-alaninium and the a-alaninate species w i l l be abbreviated respectively as follows: Ala* ; A l a + and Ala". The bidentate chelate formation with cupric ion may be represented as follows: Cu + 2 + 2 CH3 - CH - C00~=± CO - 0 / NH2 - CH - CH3 k ^ \ \c\ / M • CH3 - CH - - NH2 .O.r CO 18 The s t a b i l i t y constant for 1': 1. chelates [ Cu Ala] = l O 8 , 4 0 ^ ) [Cu + +] [Ala"] In the present system, as with glycine, sodium hydroxide was added to alpha-alanine solution up to the second equivalent mid-point (pH 11.3), as shown i n Figure 22,, The re fo r e t h e only species present was the alaninate ion* and the bidentate chelate_,fprmation may be expressed as follows: Cu + 2 + 2 CH3 - CH - COO" ^ CO - 0 ^ NH2 - CH - CH3 / | ) Cu ^ | NH2 CH3 - CH - NH2 0 — CO •k From pH and pK-^  values, the relative concentration of Gl~ and Gl± are 96.3 and 3.7% respectively. As with glycine^excess of hydroxyl ion was found to decrease the dissolution rate. The apparent passivation effect could occur by the forma-tion of an insoluble layer on the copper surface.^" Indeed, Ley and • '~?^y -Ephraim 19» 2 0i2L,a2h ave shown that a slight excess of sodium hydroxide would hydrolyse the chelate to copper hydroxide.- , ^ - 46 » Figure 22 Titration curve of a-alanine by sodium hydroxide, showing the equivalent points for the carboxylic and amino groups, respectively. - 47 -2. Effects.of oxygen pressure and a-alaninate concentration. The effect of varying oxygen pressure and anion (CH 3 - CH - C00~) NH2 concentration was studied. The results are summarized i n Figures 23 and 24. •At high oxygen pressure, the rate of solution was found to be independent of the oxygen p a r t i a l pressure. In t h i s region, i n which the rate i s chemically controlled at the copper surface, the rate of solution was found to be f i r s t order in a-alaninate and zero order i n oxygen. The t o t a l rate f o r copper--a-alaninate i s expressed as follows: R - k [ a - a l a n i n a t e ] 1 [o a]° exp _2 -3,-1/ \ The experimental rate constant k = 56 Mg Cu cm. H M (Al JL a - A l - a 3. Effects of oxygen pressure and concentration of d-alanine i n water. Making use of equation ( l ) , pK, andlthe•Ibhizationbconstant i t can be shown that in the absence, of: additionali'sod'ium- hydroxide, the species; present f o r a-alanine i n aqueous solution (pH 6.0) i s almost completely i n the' form of the zwitterion: [Al±] = 10 3.55 It has been found, as i n the previous case, that the rate of solution of copper i n aqueous a-alanine, at different oxygen pressures, i s made up of the contributions of two separate regions, BBut a further investigation was carried out, only i n the region of high oxygen pressure, i n which.it appears that the concentration of oxygen and consequently, i t s rate of transport to the surface of the copper are s u f f i c i e n t l y high so that the rate of disso l u t i o n becomes controlled by the chemical reaction at the surface (Figures .25, 26). exp The experimental rate' constant k^^± , calculated from Figure 26, where the rate of solution i s a lin e a r function, of a-alanine (zwitterion) a-Alaninate concentration O 0.1 Molar O 0.2 .» O . 0.3 15 30 45 60 75 Time (minutes) Conditions? Oxygen pressure 6.45 atm., temperature 2 5 ° C . , s t i r r i n g v e l o c i t y 775 R.P.M., pH adjustment to 11.3 by NaOH. Figure 23. Rate curves for. the d i s s o l u t i o n of copper i n a-alaninate. 0.1 0.2 0.3 Alpha-Alaninate (moles/l.) Conditionss Oxygen pressure 6 . 4 5 atm., temperature 25°C., s t i r r e r v e l o c i t y 7 7 5 R.P.M., pH adjustment to 11.3 by NaOH. Figure 24. Effect of t o t a l a-alaninate concentration on the rate of solution. - 50 Time (minutes) Conditions; Oxygen pressure 6„A.5 atm„, temperature 25°C 0 S s t i r r e r v e l o c i t y 775 RoP .M , , pH 6,00,, No sodium hydroxide has been added. Figure 2 5 , Dissolution rate of copper i n aqueous a-alanine at various concentrations. 0.1 0.2 0.3 Alpha-Alanine (mole) Conditions: Oxygen pressure 6.45 atm., temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., pH 6.00. Figure 26. Effect of concentration of aqueous a-alanine on the rate of s o l u t i o n . - 52 -concentration and independent of oxygen pressure, was found to be: exp 9 _1 _ i k A 1 + » 37.0 Mg Cu cm. H M /l. a" 4. Effect of H + ion A series of experiments was performed, i n which an excess of a-alanine and an equivalent amount of HCIO^ were added. The molecular a-alanine was kept constant. The results reported i n the following Table show that the rate f e l l r a p i d l y to near zero with increasing of H + ions. TABLE IX Effect of Hydrogen Ion i n the Alpha*. Alanine-Water System. a-Alanine (moles/1) 0.10 0.11 0.12 H G I O * (moles/1) 0.00 0.01 0.02 NaC104 (moles/1) 0.10 0.09 0.08 PH 6.30 4.20 2.7 Rate (Mg/cm.^/hr) 3.70 0.22 0 By addition of H + ion, the concentration of a-alaninium ion (A1+) i n solution has increased more than one hundred to one thousand f o l d over the a-alanine water system, the concentration of the free amine (Al~) being decreased proportionately. 5. Summary The results obtained f o r the different species i n Alpha Alamine system are summarized i n Table X. - 53 -TABLE X Relative Concentrations of Alpha-Alanlnate  Alpha-'Alaninium and Zwitterion Species. Total a-Alanine (moles/l) PH Ala* (moles/l) A l * (moles/l) Ala" (moles/l) Plot Rate (Mg/cm.2/hr/M) . 0.100 11.3 0.0963 0.0037 0.000 56.0 0.100 6.3 vlO-4 0,100 37.0 0.110 4.2 0.000 0.108 0.002 A / 0 0.120 2.7 0.000 0.089 0.031 0 The rate of dissolution can be related as follows: R - kA 1, [A3|] + kA1. | Al*] The computed rate constants are: Zwitterion species (Al±)-pH 6.3 E - 3.7 - k [o.l] • k [o] . a <x and kA]^ - 37.0 Mg Cu cm."2 H " 1 Mml£i±/1, Anion species (Al~)- pH 11.3 R = 5.6 = 37.0 [0.OO37] + kA1-[o.0963 0 -1 -1 and kA1- = 56.5 Mg Cu cm.""* H M Alj A» As in the previous glycine system, we can see that in the presence of significant increase of alpha-alaninium concentration (and corresponding decrease of alpha alanlnate concentration) the copper surface is passivated. - 54 -D, Beta-Alanine System 1. Introduction The bidentate chelate formed between cupric ions and B-alaninate ions may be written as follows: CO - 0 NH2 - CH2 Cu+/d + 2 NH2 CH2- CH2 - C00-^±CH 2 N Cu | CH2 NCH 2 - NH2 X 0 — CO The s t a b i l i t y constant f o r the complex and the hydrogen i o n i z a t i o n 17 IS A constants '* respectively are given by: feu A1+ ] = 1 0 7 , 1 5 ( k 1 ) 1:1 chelates [Alt] [H* ] . 10-3-6 A N D [ A 1 - ] [H+] = i o - 1 0 - 3 6 Sodium hydroxide was added up to the second equivalent mid-point (pH 11.6), so as to have d e f i n i t e species i n solution, e.g. 94»6$ of Al£ and 5.4$ of Al±. 2. Effects of oxygen pressure and B-alaninate concentration. In as much as the region of chemical control was the more i n t e r e s t i n g from the standpoint of the present work, a high p a r t i a l pressure of oxygen was maintained above the solution. Rate curves showing the effect of varying the B-alaninate concentration, on the di s s o l u t i o n of copper, are shown i n Figures 27 and 28. As i n the other systems, i n the region of chemical control, the reaction was found to be f i r s t order i n B-alaninate of zero order i n oxygen. A The B-alaninate, the zwitterion and the p-alaninium species w i l l be abbreviated, respectively, as follows: Aig ; A l | and A l ^ B-Alaninate concentration O o.l Molar .2 15 30 45 Time (minutes) Conditions: Oxygen pressure 6.45 atmw, temperature 25°C, s t i r r e r v e l o c i t y , 775 R.P.M. B-alamine plus equivalent amount of NaOH, pH 11.6. Figure 27. Rate curves f o r the d i s s o l u t i o n of copper i n B-alaninate. - 56 -B-Alaninat e (mole s/1.) Conditions: Oxygen pressure 6.45 atm., temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., pH 11.6. Figure 28. Effect of t o t a l ; B-alaninate concentration on the rate of so l u t i o n . - 57 -For A l " , the slope was calculated to be: exp 11.4 Mg Cu cm, 3 . Effect of oxygen pressure and concentration of 3-alani.ne i n water. No sodium hydroxide was added i n i t i a l l y to the media. The pH of these solutions was 6 . 6 . At t h i s pH i t can be computed that a l l but 0 . 1$ (Al± = 10 ) of the 3-alanine. i s in.the zwitterion form. The experimental rate constant f o r the dissolution of copper i n solution of 3-alanine i n aqueous solution, determined from the plot i n Figure 2 9 , i s equal to: 4 . Effect of H + ion The effect of H + ion was investigated, keeping the concentration of 3-alanine constant. I t was found that the rate decreased v i r t u a l l y to zero. By addition of H + ion, the concentration of 3-alaninium ion (Al*) has increased more than ten f o l d over the 3-alanine water system. 5 . Summary exp k A 1± = 3 .05 Mg Cu cm, 3 i . - 2 H"1 M"1/!. The concentration of the d i f f e r e n t species of respective constants obtained i n Beta- Alanine system are summarized i n Table XI, - 58 -3-Alanine (moles/l.) Conditions: Oxygen pressure 6.45 atm., temperature 25°C., s t i r r e r v e l o c i t y 775 R.P.M., pH 6.6 Figure 29. Effect of concentration of 6-alanine on the rate of s o l u t i o n . - 59 -TABLE X I Relative Concentrations of Beta-Alaninate s  Beta-Alaninium and Zwitterion Species. Total 3-Alanine (moles/l) PH A 1p' mole's/1 Al|,.. moles/l moles/1 Rate constant Mg/cm.2/hr/M/ w. 0.100 0.100 0.110 11.6 6.6 4.65 0.0946 0(10-3) 0.000 0.0054 0.100 0.101-0.000„ 0(10~ 3) 0.009 11.40 3.05 0.00 The rate of so l u t i o n can be related as follows? R "B - ^ B The computed rate constants are l i s t e d belows Zwitterion species (Al±) pH 6.6 R = 0.305 = k A1e + k A 1 i therefore k A 1 + = 3.05 Mg Cu cm. - 2 H"1 M _ ^ / l 4 0 Anion species (Al~) pH 11.6 R » 1.14 = 3.05 0.0054 + kAl4 - B -0.0946 3 and k - 11.8 Mg Cu cm. - 2 H"1 M _ 1 / i # A ± B As i n the previous amino acid systems studied, i n the presence of si g n i f i c a n t Beta-Alaninium (Alp) concentration, a passivating effect on the copper surface i s observed. - 60 -CONCLUSIONS The observations and correlations are s u f f i c i e n t l y consistent that i t i s possible to account f o r some of the observed effects and to draw some conclusions concerning the k i n e t i c s and mechanisms of the reactions involved i n the d i s s o l u t i o n i n the presence of the several chelating agents on copper surface. In a l l the systems studied.in the present investigation, two oxygen pressure regions were found. In the low-pressure region, the transport of oxygen to. the copper surface l i k e l y controls the rate of the chemical reaction and the rate i s d i r e c t l y proportional to the pressure and independent of the concentration of the chelating agent. In so f a r as the l a t t e r i s concerned the results provide only information on the transport mechanism. Therefore, i n - a l l the systems studied, i t was desirable to investigate the k i n e t i c s under conditions where the chemical process i t s e l f was rate- c o n t r o l l i n g . In t h i s l a t t e r region (high pressure), i t appears that the rate of transport of oxygen to the surface of copper i s s u f f i c i e n t l y high that the rate of dissolu t i o n becomes controlled by the chemical reaction and independent of the concentration of oxygen. For lim i t e d conditions, with low concentration of the chelating agent, the reaction f i t s an empirical equation of the fbrms Rate = k [chelating agent] [Wj° (l) The measured rate of d i s s o l u t i o n of copper metal i n each of the dif f e r e n t systems studied, i s made up of the contribution of two separate reactions. The species giving r i s e to these reactions appear i n a l l cases to be the ligand and i t s acid form. For example, i n ethylenediamine system - ethylene-diamine and ethylenediaminium; and i n glycine system - glycinate (Gl~) and the zwitterion species ( G l * ) . 6 1 -I n e a c h s y s t e m t h e r e a c t i o n i s i n d e p e n d e n t l y f i r s t o r d e r w i t h r e s p e c t t o t h e c o n c e n t r a t i o n o f e a c h o f t h e s e two s p e c i e s w h i c h f i t t h e f o l l o w i n g e q u a t i o n r a t e : R = k X [ L ]+ k 2 [LH*] where L i s t h e c h e l a t e f o r m i n g s p e c i e s . The mechan i sm demised b y H a l p e r n t o e x p l a i n t h e r e s u l t s i n t h e ammoniax s y s t e m c a n be a p p l i e d e q u a l l y w e l l t o t h e s y s t e m s s t u d i e d i n t h e p r e s e n t w o r k , as i l l u s t r a t e d f o r am ines and amino a c i d s as f o l l o w s , u s i n g e t h y l e n e d i a m i n e and g l y c i n e as e x a m p l e s , r e s p e c t i v e l y , I. Oxygen i s c h e m i s o r b e d r a p i d l y on t h e c o p p e r s u r f a c e : f a s t Cu + 1/2 0 2 -+ Cu.,.,0 m e t a l d i s s o l v e d I I , R e a c t i o n o f e t h y l e n e d i a m i n e m o l e c u l e o r e t h y l e n e d i a m i n i u m i o n w i t h t h e c o p p e r o x y g e n c o m p l e x : s l o w Cu.,.,0 + e n y N H 2 - C H 2 - C H 2 - N H 2 Cu + t o ~ + HOH^CufNHp - C H 2 - C H 2 - N H 2 ) + + + 2 OH" f a s t s l o w o r Cu.,.,0 + e n H + y N H 2 - C H 2 - C H 2 - N H 2 Cu++" \ ^ 6-+ O H - y C u ( N H 2 - C H 2 - C H 2 - N H 2 ) + + + 2 OH" f a s t I l a , R e a c t i o n o f g l y c i n a t e a n i o n o r z w i t t e r i o n s p e c i e s w i t h t h e c o p p e r o x y g e n comp lex ( t he f i r s t s t e p i s t h e same i n b o t h s y s t e m s ) . s l o w C u . . . . O + G l ~ N H 2 - C H 2 - C00-C u + + + HOH f a s t ' C u ( N H 2 C H 2 C 0 0 ) + + 2 OH" slow or Cu....O + Gl± -* Cu(NH2 - CH2 - C00) + + 2 OH" The zero-order dependence on complexed cupric ion indicates that the cupric ions dissolved during.the corrosion process do not affect the rate i t s e l f and shows also that any mechanism involving oxidation of copper to cuprous by cupric ion can be ruled out. The fa c t that the rate of dissolution f o r the ethylenediamine, glycine, a-and B-alanine systems :is "independent of the concentration of oxygen gives evidence that the f i r s t step i s f a s t , also that the surface i s covered with a f i l m of oxygen. The formation of the proposed intermediate activated complex would involve the coordination of the NH2 group to the copper surface, which would be accompanied by a transfer of electrons from the copper to the adsorbed oxygen atom. Table XII below gives the computed rate constants obtained f o r the different species studied. TABLE XII  Summary Table of Rate Constants System Species Computed rate constant (Mg/cm,2/hr/M/l.at 25°C.) ; Ammonia Ammonium NH3 NH 4 + 61 (experimental) 1550 (Halpern) Ethylenediamine Ethylenediamine en enH + 245 515 Glycine Glycine Gl-Gl± 50.5 32.5 a-Alanine A 1 5 56,5 a-Alanine A l t a 37.0 3-Alanine A l f i i ; e , B-Alanine A l * 3 3I05 X N H 2 - CH2 - C00-Cu + + " H x -6-+ OH-fast - 63 -With respect to the preceding data, there are certain s i g n i f i c a n t trends which warrant further additional remarks. It w i l l be observed from Table XII that the addition of a proton to NH3 or 1 en'increases the rate of i t s attack on copper. On the other hand, in the case of the amino aKids studied, the opposite effect i s obtained. I t i s further observed that In a l l cases the neutral species has a lower rate than the corresponding charged species, regardless of the sign of the charge. In Table XIII a comparison has been made between the s t a b i l i t y constants of the species involved with copper and t h e i r respective computed rate constants. I t appears that a s i g n i f i c a n t relationship exists between these two constants. For instance, the neutral species 1, 2, 3, and 4 show the s t a b i l i t y constant (log k-|) to be a function of the rate constant. The corresponding protonated species show a displaced similar functional r e l a t i o n s h i p . I f ad d i t i o n a l points corroborated these li m i t e d curves, the results would be s i g n i f i c a n t . One could u t i l i s e t h i s to predict the rate constant from the s t a b i l i t y constant value. TABLE XIII Correlation of Rate Constants  and S t a b i l i t y Constants. Species S t a b i l i t y constant log k]_ (chelates 1:1) Rate constant Mg/cm,2/hr/M/l.; (1) en (2) G l -(3) A l -(4) Alp 10.7 8.2 8.4 7.15 245 50.5 56.5 11.8 Surprisingly, comparison of the rate constants obtained for glycine and a-alanine, respectively, indicates that the expected s t e r i c e f f e c t s , due to the branched amino group of a-alanine, does not occur. In fa c t , the rate constants obtained f o r the d-alanine species are even s l i g h t l y greater than - 64 -the ones for glycine. Recommendat ions There remain several questions pertaining to the corrosion of copper i n amines and amino acid solutions which have not been yet answered. These include (a) the a c t i v i t y of the diprotonated ethylenediamine cation which must be obtained i n strong acid.solutions and (b) the effect of the carboxylate end of the amino acid may possible form purely carboxylate complexes with cupric ions.., . Experimentally, t h i s part of the ethylenediamine system can be completed by a series of experiments i n acid solutions. The amino acid systems can probably be c l a r i f i e d by a c a r e f u l l y chosen sequence of experiments i n which both a carboxylate, i . e . , acetate and an amine such as ammonia or ethylenediamine are used to simulate the carboxylate and the amine ends of the amino acid. This would c l a r i f y the nature of the corrosion i n h i b i t i o n observed i n the amino acid systems where H + ions have been added, although i t i s believed that i f the concentration of free amine goes below a c r i t i c a l value, passivation results from hydrolysis of the corrosion product. I t would also be of great interest to study other amines and amino acids, and relate more close l y the s t a b i l i t y constant, the number of carbon of the ligand, the possible s t e r i c e f f e c t , etc. on the rate constant. - 65 -BIBLIOGRAPHY 1. E. Yamasaki, Science Repts., Tohoku Imp. Univ., 2» 169 (1920). 2. E. Zaretskii and G. Akimov, J. Appl. Chem., U.S.S.R., 11, 1161 (1936). 3. R.W. Lane and H.J. McDonald, J. Am. Chem. Soc , 68, 1699 (1946). 4. J. Halpern, J. of Electrochem. Soc , vol. 100, 10, 421 (1953). 5. T.I . Fisher, ''Reaction of Copper Gold Alloy in Aqueous NH3 Under 02 P r e s s u r e 1 » , M.A.Sc. Thesis. Univ. of British Columbia, (1953). 6. A.S.T.M., Grain Size Standards for Non Ferrous Metals (Introduction to X-ray Metallography, A. Taylor). 7. J.T. Clarke and E. Blant, Journal of Folymer Sc., 1, 427 (1946). 8. Handbook of Chemistry, Standard Table, Rubber Publ. Co, (1947). 9. C. Schmidt, The Chemistry of the Amino Acids and Proteins, C.L.A, Schmidt Springfield, Springfield, 111. (1945) 10. L. Ebert, Z. Physik Chem., 121, 385 (1926). 11. H.H. Weber, Biochem. Z, 218, 1 (1930). 12. D.N, Sen, San-Ichiro Mizushima, ,Columba Curran, and J.V. Quagliano, J, Am. Chem. Soc , 22, 211 (1955). 13. C. Haddock and P. Evans, Analyst, $_2, 495 (1922). 14. T.P. Hoar, Analyst, 62, 657 (1937). 15. T.H. High, Analyst, 22, 60 (1947). 16. E.B. Sandell, » 'Colorimetric Determination of Traces of Metals, Interscience (1950). 17. J. Bjerum, Metal Ammine Formation in Aqueous Solution, P. Haas© and Son, Copenhagen (1941). 18. A. Martell and M. Calvin, Chemistry of Metal Chelate Compounds, Prentice Hall , New York (1952). 19. H. Ley, Z. Elektrochem., 10, 954 (1904). 20. 0. Pleiffer, Angew Chem. «£, 93 (1940). 21. L. Liebhafsky, J. Chem. Education 23., 341 (1946). 22. Jr. Ephraim, Inorganic Chemistry, Norderman Publ, Co. Inc., New.York (1943). - 66 -Bibliography (cont'd.) 2 3 . E. Peters, The Homogeneous Cat a l y t i c Activation of Molecular Hydrogen by Cupric Salts i n Aqueous Solution, Ph.D. Thesis, Univ. of B r i t i s h Columbia, A p r i l 1956. 24 . U. Evans, M e t a l l i c Corrosion, P a s s i v i t y and Protection, E. Arnold and Co., London (1951) . H.H. Uhlig, The Corrosion Handbook, J.'Wiley, New York (1948) . 

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