Kinetics of the Dissolution of Copper Metal in Some
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
in the Department
of
MINING AND METALLURGY
We accept this thesis as conforming to the
standard required from candidates for 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 dissolution of copper metal in 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 dissolution of copper, in all solutions, has been found to be
independent of the initial copper concentration, the volume of the solution and the area of the copper sample. No intermediate products, i.e., cuprous
ions, were observed.
Two regions were observed, having different dependence on oxygen pressure. In one, the rate depends on the first power of the oxygen pressure,
and is independent of the concentration of the chelating agent. In the other
region, the reaction is first order in chelating agent and independent of oxygen pressure. The rate of the reaction in this 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 dissolution reactions were found to be first 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 in the present work. The rate constants for each chelating
agent have been computed and appear to be related to the stability constants.
No steric effect was observed. In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my
Department or by his representative. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.
Department of : Ti The University of British Columbia, Vancouver S, Canada. TABLE OF CONTENTS
Page
INTRODUCTION » 1
Scope of the Present Investigation 3
EXPERIMENTAL ....••••.•••«.QO«..«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. Analytical Procedures ..... 9
G. Measurements of the Rates 10
RESULTS AND DISCUSSION 14
A. Ethylenediamine System 14
1. Introduction 14
2. Effect of stirring velocity 14
3« Effect of surface area and solution volume 16
4. Effect of cupric ion on the dissolution 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 in 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 in 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 in 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 for Copper Carbamate . . 11
3. Typical Rate Curve for 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 02 pressures
Ethylenediamine 0.475 Molar 19
6. Rate Curves for Copper at Various 02 pressures
Ethylenediamine 0.1425 Molar o. 20
7. Rate Curves for Copper at Various 02 pressures
Ethylenediamine 0.095 Molar 21
8. Rate Curves for Copper at Various 02 pressures
Ethylenediamine 0.0712 Molar ...... 22
9. • Rate Curves for Copper at Various 02 pressures
Ethylenediamine 0.0475 Molar 23
10. Effect of Oxygen Pressure on the Rate of Dissolution of Copper . . 24
11. Rate Curves for Copper with 02 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 in
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. Titration Curve of Glycine with Sodium Hydroxide . 36
17o Rate Curves for Dissolution of Copper in Glycinate Solution ... 37 Figures (cont'd.)
Figure ' Page No.
18. Effect of Oxygen Pressure on the Rate of Dissolution of Copper in
Glycinate Solution 38
19. Effect of Glycinate Concentration on the Rate of Solution 39
20. Rate Curves for Dissolution of Copper in Aqueous Glycine 41
21. Effect of Concentration of Aqueous Glycine on the Rate 42
22. Titration Curve of a-Alanine by Sodium Hydroxide 46
23. Rate Curves for the Dissolution of Copper in a-Alaninate Solution . 48
24. Effect of Total a-Alaninate Concentration on the Rate of Solution . 49
25. Dissolution Rate of Copper in 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 for the Dissolution of Copper in 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 ooe«*oooo««o»«o*o«o«6o«ooo* 5^ TABLES
No. Page
'.I, Determination of Reproducibility 12
II. Effect of Stirring Velocity on the Dissolution Rate 15
III. 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 in the Glycine-Water System 43
VIII. Relative Concentrations of Glycinate, Glycinium and Zwitterion
SpGCISS o«e«oc«o« o « o o • oeooooo* « o o • A-3
IX. Effect of Hydrogen Ion in 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
XIII. Correlation of Rate Constants and Stability Constants 63 ACKNOWLEDGEMENT
The author is indebted to the National Research Council of
Canada for the financial aid which enabled this project to be carried out, and to the Union Miniere du Haut Katanga for permission to take a year's leave of absence.
The author is grateful to the members of the Department of
Mining and Metallurgy for their assistance throughout this work, and is especially grateful to Dr. D.R. Wiles, who ably directed this 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 in oxidizing aqueous media (EvansyUhlig ) but very little has been directed towards the theoretical aspects of the processes involved.
The mechanisms of reactions of the dissolution of a metal in aqueous media, which involve simultaneous oxidation and dissolution 'of the metal, were not understood until recent years^*-' Only a few systems have been studied in detail. Earlier investigations of the kinetics of the dissolution of copper in aqueous ammonia have provided conflicting information about the order of the reaction and the effect of the different variables.
Yamasaki"*" investigated the rate of dissolution of metallic copper in ammonium hydroxide by rotating copper specimens at constant velocity and supplying a steady current of purified air. His work indicated that auto- catalysis by cuprammonium was involved in the dissolution of copper, but showed the rate of dissolution to be independent of the concentrations of ammonia or of salts in solution.
2
Zaretskii and Akimov investigated the mechanism of corrosion of copper in aqueous solutions of ammonium compounds ahd attributed the corrosion to three processes, the first being the electrochemical dissolution of copper with formation of cuprous ammonia complex, the second, oxidation of the complex to form cupric ammonia complex, and the third, electrochemical reduction of the cupric ammonia complex to cuprous ammonia complex. In a more recent study, Lane and McDonald, failed to observe any evidence of auto- catalysis, and found instead a strong dependence of the rate on ammonia concentration.
All the earlier investigators found that the rate increased with the stirring velocity, indicating dependence on transport of oxygen to the copper surface. As was pointed out more recently by Halpern,^ when diffusion of a reactant controls the rate of a chemical reaction, the effects of other variables are masked and kinetic results provide little information about the mechanism of the reaction. Therefore it appeared desirable to investigate the region where the transport of oxygen to the copper surface was sufficiently fast that it did not affect the rate.
This condition was achieved in a study of the kinetics of the dissolution of copper in ammonia solutions recently conducted in this laboratory by Halpern^" and Fisher.'' In the course of these latter investiga• tions, it was shown that the observed rate of dissolution of copper in ammonia solutions was the sum of the rates of two independent reactions, first 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 in excess (so that the transport of oxygen to the copper surface did not limit the rate of the reaction).
On the basis of these results, the following mechanism for the dissolution of copper was proposed:
(1) Adsorption of dissolved oxygen on the copper surface:
fast
Cu + 1/2 02 — 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(NH3) + 2 OH"
slow, /F*3» fast + + ++ or Cu.,.0 + NH^ |Cu ^H | + 0H~ Cu(NH3) + 2 OH" * v'-o*;
Following this work on copper in ammonia solution, it seemed of interest to compare the dissolution rates in ethylenediamine solution with those determined previously in ammonia solutions and to extend the study to other complexing agents, particularly 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 solid- 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 in the present work. Therefore preliminary experiments were carried out to determine the conditions under which the reaction is 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. If the substance which combines with the metal contains two or more donor groups so that one or more rings are formed, the resulting structure is said to be a chelate compound. Structural differences might show up readily and still 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 crystal size differences, effects of oxygen pressure, stirring 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 in this.thesis 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 in 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 kilocycles and power of the order of
1 kw) in spectrographically pure graphite crucibles. The melting operation was performed in vacuo; a Welch Duo-Seal vacuum pump was used. A vacuum of better than 50 microns was maintained in the melting unit. The copper was allowed to solidify in the crucible. The solidified ingots were cold worked by squeezing in a press and annealed at either 500°C for five minutes or
700°C for three minutes. The diameter of the recrystallized grains was about
0.120 mm (A.S.T.M. non-ferrous grain size standard^) after 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 in bakelite in a metallographic specimen mounting press. This preparation was done in order to insulate the copper samples electrically 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. Prior to each experiment, the dimensions of the polished sections were carefully 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 different copper samples studied ranged from 1.386 to 2.248 cm. After measurement, the samples were stored in a desiccator until needed.
Alternative methods of preparing the polished surface, such as etching
with ammonium persulphate-ammonium hydroxide or with HN03, were found to give identical dissolution rates to those experiments in 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 in this thesis was conducted in an auto• clave (Figure 1) designed for working at pressures up to 8 atmospheres. The autoclave was fabricated from 316 stainless steel but a titanium liner was used, which was shown to be inert to the solutions used. The mounted copper specimen was held by means of a stainless steel rod, the latter being screwed into the lid of the autoclave.
Agitation was provided by two turbine-type impellers operating on a single shaft. The rotational 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 liters was normally used.
Oxygen was supplied from a standard cylinder and the desired pressure held constant by means of a standard oxygen pressure regulator fixed on the - 6 -
Figure 1. Schematic diagram of the autoclave and heating control system.
A. Shaft F. Impeller B. Thermometer well and.thermoregulator G. Pump C. Cooling coil H. Relay D. Sampling tube I. Heater E. Copper sample in bakelite mount 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 all 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, it 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 will 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 1.02$ presumably being water. In order to detect any disturbing impurities in the latter reagent, the following tests were performed?
(1) Detection of heavy metal cations.
The ethylenediamine was examined for heavy metals by X-ray fluor• escence. The technique used was to evaporate a 200 ml sample of ethylene- diamine on a small amount of 'Celite', which had previously been carefully washed with perchloric acid and water. No heavy metals were detectable.
(2) Fractional distillation
The ethylenediamine was fractionally distilled according to the method of Clarke and Blant in order to obtain 100% purity. This distillation was conducted in a Todd Scientific Co. column. The final product obtained was examined using the refractive index and boiling points as criteria. The figures obtained were, n 1.4540, B.P. 116°C, in exact agreement with published data for pure anhydrous ethylenediamine.
(3) Comparison of dissolution rates.
Dissolution rates were compared in solution using
a. the 98.98$ ethylenediamine
b. the 100$ distilled ethylenediamine
c. the 100$ reagent grade (Eastman Kodak) ethylenediamine.
Differences in the rates observed were about 0.5$, less than the expected "'V, experimental deviation. From the results of these tests it was concluded that no disturbing impurities were present. - 9 -
Solutions were prepared by diluting a measured volume of reagent to
2,0 liters with distilled water. In the case of the ethylenediamine and ammonia systems, the normality of these solutions was checked by titration with standard sulfuric acid to a methyl red end point. In order to avoid a salt effect as observed by Halpern, sodium perchlorate was added to the solution in the autoclave to give a total salt concentration of 0.1 moles per liter.
F. Analytical Procedures.
Analysis of copper in samples was done colorimetrlcally 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 after extracting the coloured complex into carbon tetrachloride
13 or amyl alcohol. Since effective extraction may be very time consuming, several methods of avoiding it have been devised for use in cases when inter• fering elements, e.g. iron, are absent. Since the stability of the copper carbamate complex is affected by daylight, a small amount of gum arabic was added and a clear golden-brown colour well stabilised by the protective colloid, ^ . "1.4,15, 16 was obtained. The carbamate-gum arabic solution was prepared as follows; 0.2 grams of carbamate dissolved in 200 ml. of distilled water was mixed with 5 grams of gum arabic of 1 ml. toluene in 1000 ml. distilled water.
This solution (henceforth abbreviated CGA) was filtered and stored in the dark,
The analytical solutions were prepared in an artificially lighted room to avoid the effect of daylight. The stability of the chelate compound is very dependent on the pH. However, it was found that the optical density was not affected between pH 7.5 and 9.2. Samples for analysis were prepared as follows: To an aliquot of the - 10 - sample taken from the autoclave, 5 ml. of ammonium citrate was added to complex accidental traces of iron, and the pH adjusted to 9.0 with 1:1 ammonium hydroxide. Finally, 10 ml of CGA solution was added and the mixture diluted with distilled water to 50 ml in a volumetric flask. Depending on the copper concentration, the aliquot varied from 2 to 10 ml. The concentration of the coloured complex may be readily determined by spectrophotometric comparison with a series of similar solutions of known concentration. For this purpose, a Beckman model DK-2 Ratio-Recording Spectrophotometer was used to determine the optical density of the solutions at a wavelength of 437 millimicrons. The optical density was found to be proportional to the copper concentration over a sufficient range to allow convenient analysis. The calibration curve is given in Figure 2,
The results obtained by this method could be duplicated within less than ±0.5%. The presence of ethylenediamine, glycine, or aV and (3r alanine,
,was found'not to affectiAthe calibration curve; : '\'\
G. Measurement of the Rates of Dissolution.
The method used for measuring the dissolution rates of copper in the various solutions was as "follows:: The autoclave, charged with solution and sample, was flushed with nitrogen until 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,
40 ml samples were withdrawn by means of the sampling tube at 15 minute intervals and analysed for copper. From the known surface area of the sample and the known volume of the solution at the time of each sampling, the total amount of copper dissolved per unit area was computed. The experiments lasted between 1.5 and 2 hours. During that time, approximately 0.50 to 2.00 x 10~3 - 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 for analysis of copper by the carbamate method. - 12 - moles of copper was dissolved. This amount was insufficient to cause any
appreciable change in the concentration of the chelating agent, through the formation of a chelate or amine: complex. A typical rate curve for the
dissolution of copper in ethylenediamine solutions under oxygen pressure is
shown in Figure 3. The amount of copper dissolved was directly proportional to the reaction time. The linearity 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 plots.
A series of experiments was carried out at constant ethylenediamine
concentration, oxygen pressure, agitation and temperature, in order to deter• mine the reproducibility of the experimental technique. The results of these
tests are summarized below:
TABLE 1..
... Determination of Reproducibility
Run No. . Rate of Solutibn Deviation 2 (MG/cm. /hr.) 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 Stirring velocity 775 RPM; Oxygen pressure 5.0 atm.
The rates were generally found to be reproducible to within 1 or 2% in
duplicate experiments. In the following sections are described the results
of experiments designed to determine the influence of the different 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 for the dissolution of pure copper. - 14 - RESULTS AND DISCUSSION
A. Ethylenediamine System
1. Introduction
Ethylenediamine is a strong chelating agent, and with cupric ions forms the complex:
1 CH2 - NH2 y NH2 - CH2
I ^ CU I
CH2 - NH2 ' NH2 - CH2
It is also a strong base, and readily takes on either one or two hydrogen ions. 17 18 ± The equilibrium constants ' for the formation of the various species are:
[Cu en++] - 10 * (kj) (!) (1:1 chelates) [Cu^Fen]
7 49 ++ C enH2 J - 10" ' (2) tenH^[H+3
[enH+] - 10"10,17' (3)
4 For convenience, ethylenediamine will be abbreviated 'en* in formulae. + + Similarily one has for the acid species enH and enH2 t.
Most of the present work was done at about pH 11.5 at which the ethylenediamine is present almost entirely as the free base. The corrosion action of the ethylenediamine solutions may be formulated according to the equation:
++ Cu + 2 en + 1/2 02 + H20 — Cu en2 + 2 OH" (4)
No intermediate products such as cuprous ions were detected during the course of the experiments.
2. Effect of stirring velocity
It is well known that when diffusion controls the rate of the - 15a -
A series of tests was performed with the object of determining whether the reaction rate is agitation dependent in the low-pressure region, as was found by Halpern.^ The results are given in Table Ha.
TABLE Ila.
Effect of Stirring Rate on Dissolution Rate in the Low Pressure Region.
Stirring rate Reaction rate Deviation (R.P.M.) (Mg/cm.2/hr.) % 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 atm.
The results show clearly that the rate is not dependent on the
stirring rate. This is at variance with Halpern*s results for the same
pressure region. No reason has been found for this discrepancy. One possible
explanation is that the sample in the present investigation was not located in
the same place as was Halpern*s, but was closer to the tip of the impeller
blades, where agitation is expected to be more violent. It is possible, then,
that the agitation in this position reaches a maximum at fairly low stirring
rates, and that further increase in stirring rate causes no increase in
turbulence near the sample surface. - 15 - chemical reaction, the kinetic results give little information about the mechanism of the reaction. When such a physical process is predominant, the effect of other variables on the rate is difficult to be determined. It is, rather, important to find a range of conditions such that the rate is independ• ent of stirring and therefore controlled by factors other than the transport of reactants or products. It has already been shown by Halpern that, for the ammonia system at low ox/gen pressures the rate increases with stirring velocity, whereas at higher oxygen pressures such a dependence is not observed. Thus, it was of primary importance to determine at what stirring velocity the present work should be done. The results of tests at different stirring rates are summarized in Table II.
TABLE II
Effect of Stirring Velocity on Dissolution Rate
Stirring Rate Reaction Rate Deviation (R.F.M.) (Mg/cm.2/hr)
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 in Table II show that the reaction rate is not limited by transport of reactants within the solution. As will be discussed later, simpler criteria 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: It was shown by Halpern and corroborated by the present work that in - 16 - the absence of transport control, the reaction is first order in the complexing agent. This will be shown more clearly in a later section. In order to eliminate any extraneous effect, a stirring rate of 775 RPM was used,
3. Effect of surface area and solution volume.
Preliminary measurements indicated that effects due to the change of volume and to changes in the surface area upon the rate of solution were not significant. Results of these experiments are given in Table III and IV.
TABLE III
Effect of Surface Area on Copper Dissolution
Surface area Rate of solution Deviation of copper (cm.2) (Mg/cm.2/hr)' (*) 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, stirring velocity 775 RPM.
TABLE IV.
Effect of Solution on the Rate
Volume of Solution Concentration Rate Rate of Solution: Deviation (ml) (Mg/l/hr.).. (Mg/hr) (Mg/cm.Vhr.) (%)
1,500 33.6 40.1 19.20 +0.26 2,000 24.1 40.5 19.00 -0.75 2,456 19.2 40.2 19.30 +0.75
Average rate of solution 19.15 0.585
These results show that the rate of dissolution is 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 effective area of the copper remains constant during the experiment. This fact also gives
evidence that the rate controlling step occurs at the surface itself and not in the solution, since the volume of the solution necessarily diminished during the
course of each experiment,
4. Effect of cupric ion on the dissolution rate
The zero-order reaction indicated in all cases shows that the cupric ions dissolved during the corrosion process in the form of a chelate compound such as C\i(en^'+', do not affect the rate itself. As a further test, an experi• ment was done in which 0.015 g/1 of copper perchlorate was added at the start.
The results, plotted in Figure 4, show that the rate of dissolution is the same
(within .2%)) as when no cupric ions were added initially.
5. Effect of oxygen pressure
Some typical rate curves showing the effects of varying the oxygen pressure from 2.3 to 7.8 atmospheres are shown in Figures 5 to 9 for various concentrations of ethylenediamine. The complete system is summarized in Figure
10, where the rate of solution for different ethylenediamine concentrations is plotted veraus the oxygen, pressure.
At low.pressures, the rate of solution is proportional to the pressure and independent of the reagent, in this case, the ethylenediamine.
Thus in this region, the rate may be controlled by the transport of oxygen to the surface of the copper.
At higher pressures, however, a point is reached, where the rate of solution levels off. In this region, the rate is 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 dissolution rate. - 19 -
T j j ! r
15 30 45 60 • 75 Time (minutes) Conditions: Temperature 25°C, stirrer velocity 775 R.P.M., NaClO^ 0.1 M.
Figure 5. Rate curves for the dissolution of pure copper in
ethylenediamine solution at various 02 pressures. ~ 20 -
Time (minutes) Conditions: Temperature 25°C., stirrer velocity 775 R.P.M.,
NaC104 0.1 M.
Figure 6. Rate curves for the dissolution of pure copper in
ethylenediamine solution at various 02 pressures. - 21 -
Time (minutes) Conditions: Temperature 25°C., stirrer velocity 775 R.P.M.,
NaC104 0.1 M. Figure 7. Rate curves for the dissolution of pure copper in
ethylenediamine solution at various 02 pressures. - 22 -
35 Oxygen pressure O 7.8 atm.
O 3.7 "
30 • 2 O .4
25
20 -
15
10
[enj: 0.0712 Molar
30 45 60 75 Time (minutes^ Conditions: Temperature 25°C., stirrer velocity 775 R.P.M., NaCIO/, 0.1 M. Figure 8. Rate curves for the dissolution of pure copper in
ethylenediamine solution at various 02 pressures. - 23 -
Time (minutes)
Conditions: Temperature 25°C, stirrer velocity 775 R.P.M., NaClO^. 0.1 M.
Figure 9. Rate curves for the dissolution of pure copper in ethylenediamine solution at various oxygen pressures. - 2k -
02 pressure ( atmospheres) Conditions: Temperature 25°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 will be shown in the subsequent section to be zero order with reference to oxygen pressure.
The data presented in Figure 10 demonstrate that the dual control found by Halpern for the copper ammonia system is valid in the present system.
6. Effect of total ethylenediamine concentration
The rate of dissolution of copper was investigated using a series of
solutions containing ethylenediamine in concentrations ranging from 0,05 to
0.5 M/l(see Figures 5 to 10). Since the region of chemical control was the more
interesting from the standpoint of the present work, a high partial pressure of oxygen was maintained above the solution.
The results are shown in Figures 11 and 12 and the data indicate quite clearly that at low concentration of the chelating agent, the reaction is first order with respect to the concentration of ethylenediamine.
Zero-order dependence in oxygen was demonstrated in Figure 10. These
factors can be expressed as follows:
1 Rate - k[en J [o2 ]° (5)
The experimental rate constant K®^3, was calculated from Figure 12 and found
t0 be: exp _2 -1 _1, A kn = 245 Mg Cu cm. Hr M 71.
The deviation from linearity at high ethylenediamine concentration could reflect the transition 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., stirrer velocity 775 R.P.M., temperature 25°C.
Figure 11. Plot showing the rate of solution of copper in 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., stirrer velocity 775 R.P.M., pH 11.5. Figure 12. Effect of total 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 is described in 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 for the ionization of ethylene-
diamine can be rewritten in the form;
[enH+J - 10^°-17 - IO3'93 l^W) *w
or [enH+D = c fenl "FT It follows that as long as an appreciable concentration of the free base
ethylenediamine remains in solution, any added H+ ions are almost quantitatively
transformed into enH+ ions. The increase in enH+ concentration is of course
associated with a proportional (although very small) increase in H+ concentra•
tion. The results of this investigation are shown in Figures 13 and 14. It
is seen that the rate of solution increases linearly with increasing enH+. It
appears, thus, that the total rate of dissolution is made up of the contributions
of two separate reactions. These two reactions are first order with respect to
the concentration of en and enH+',respectively. The total dissolution rate
can be expressed as follows:
RT = Ren + RenH+ (6)
or _ RT - ken(en) - k^ (enH+) (7) - 29 -
T
Ethylenediaminium concentration
Time (minutes) Conditions: Stirrer velocity 775 R.P.M., oxygen pressure 7.8 atm., ethylenediamine concentration 0.0475 molar. Figure 13. Effect of hydrogen ion on the dissolution rate of copper in ethylenediamine. T
30-
20
0.01 0.02 0.03 0.04 Ethylenediaminium (moles/1.) Conditions: Oxygen pressure 7.8 atm., stirrer velocity 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 line in Figure 14, the first order rate constant for the dissolution of copper in ethylenediaminium solutions is
2 found to be 515 1 Ig/ cm. /:hr/:rtdle's/l. For a total electrolyte concentration of
0.1 mole/liter, equation 7 can now be written:
+ RT = 245 [en] + 515 [enH ] (8)
This equation is valid in the region of chemical rate control, where oxygen is present in 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 in ethylenediamine solution. A decrease of the rate was observed. The data are summarized in
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 slight excess of hydroxyl ion will transform any enH+ into en species and therefore decrease the rate. A further decrease (Experiments Nos. 3 and
4) is probably due to a passivating effect of the copper suface by 0H~ ion. - 32 -
9. Summary of results for the ethylenediamine system.
The investigation performed and the results obtained in the ethylene- diamine system may be summarized as follows;
(a) The rate of dissolution of copper has been found to be independ• ent of initial 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 in which oxygen transport may limit the rate, and the high oxygen pressure region in which the rate is chemically controlled at the copper surface.
(d) In the region of chemical control (at low concentrations of the chelating agent), the reaction is first order in ethylenediamine and zero order in oxygen.
(e) pH effect. Any added H+ is almost quantitatively transformed into enH+ and increases the rate. The rate for enH+ was found to be first order with respect to the concentration of enH+, and independent of oxygen concentration provided that its transport to the surface of copper does not limit the reaction rate. A decrease of the rate was observed by increasing the pH with sodium hydroxide.
At this point, it seems evident that the general pattern of the experimental results, which were obtained in the course of the present work with the ethylenediamine system, is 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 in the region in which the rate is chemically controlled at the copper surface, several other systems were compared in that same chemically controlled region. - 33 -
B. Ammonia System
A few experiments were performed in the ammonia system so as to be able to determine and to compare its rate constant kjjH^ with the one found in the other systems. The results obtained are summarized in Figure 15. Rate measurements made at different ammonia concentrations in the region of chemical control give a first order reaction kinetics in ammonia and zero-order reaction kinetics in oxygen. The experimental rate constant, kjjjj calculated from the linear slope of Figure 15 was found to be:
NH3T 5 The results obtained were found to be in good agreement with those of Fisher,'
C. Glycine System.
1. Introduction
Glycine - NH2.CH2„COOH is an amino acid. Since the publication of 17 Bjerrum's paper,, it has been generally agreed that the ami.no acid, in water solutions, exist largely as amphions and not as undissociated molecules. In 9 10 11
+ aqueous solution the 'Zwitterion' * * NH3 CH2C00~, can itself act as an acid and form a glycinate ion. Glycinate ion is a bidentate chelating agent which forms one covalent 12 and one ionic bond with the cupric ion resulting in five-membered ring. The metal chelate formation may be represented as follows:
++ „ GO— 0 ^ XNH 2? — CH ?2 Cu + 2 NH2 CH2-C00- 5^ , ^ Cu' i N CH2-NH2 0 — CO
It has been shown^7>1^A9 that, in general, any additional hydrophilic groups such as hydroxyl ion or ionic group would increase the, solubility of amino acids
•in water. ' ' • • 0.5 1.0 Ammonia (Moles/l.)
Conditions: Oxygen pressure 6.45 atm., temperature 25°C, stirrer velocity 775 R.p.M.
Figure 15. Effect of concentration of aqueous ammonia on the rate of solution. - 35 -
For convenience, the different species will be abbreviated in
formulae as follows:
+ + the cation species NH3 CH2C00H = Gl
+ the zwitterion species NH3 CH2C00~ = Gl±
the anion species NH2CH2C00~ *» Gl~ 18
The chelate stability constant with glycinate and cupric ions would be expressed as follows:
[Cu Gl+J - 108*62 (ki) chelates 1:1 [Cu+U par]
The ionization cohstants^>18 for glycine are given by:
[G1±]H - IO"2'308 and [G1-] [H+] = IO"9-78
In a strongly acidic solution the substance is present as the positively
charged ion (a) , . V -H+ . -H+
+ + H3N CH2C00H ^± H3N CH2C00- ^ H2NCH2C0CT (a) (b) ' (c)
The neutral dipolar ion (b) is formed in 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 typical titration curve of glycine in aqueous
solution with sodium hydroxide is shown in Figure 16.
2. Effects of oxygen pressure and glycinate concentration.
The effects of varying the oxygen pressure of the glycinate concentra• tion were investigated and are summarized in 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 Titration 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 -
02 Pressure (atmosphere) Conditions; Temperature 25 °C, stirrer velocity 775 R.P.M., pH adjustment to 11.3 with NaOH.
Figure 18. Effect of oxygen pressure on the rate of dissolution of copper. - 39 -
0.1 0.2 0.3 Glycinate (moles/1.) Conditions? Oxygen pressure 6.45 atm., temperature 25°C., stirrer velocity 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 dissolution rate was found to be independent of oxygen pressure, thus chemically controlled at the copper surface. In this latter region, the rate is first order in glycinate and zero order in oxygen pressure as shown in Figure 19.
As in the ethylenediamine system, the total rate for copper dissolution in glycinate solution can be expressed as follows?
1 Rate = k [Glycinate] [o2]°
The experimental constant k^^F , was calculated from the linear 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 in water.
The rate of dissolution of copper was investigated using a series of solutions containing glycine in water in concentration ranging from 0.1 to 0.3M.
These experiments were performed with no sodium hydroxide (Figures 20, 21).
The pH of glycine in aqueous solution is 6.0. From it-hisllatter, value and pK-j_ constant it can be seen, that the species present is almost completely in the form of the zwitterion
Gl± « lO3'62 Gl*
In the high pressure region, where the rate is chemically controlled at the copper surface, the reaction is first order in glycine (Gl~) and the exp experimental rate constant calculated from Figure 21 is found to bet
ex P o T -1 , kGl± = 32'5 Mg Gvi°'cm° H M /le
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., stirrer velocity 775 R.P.M., temperature 25°C, pH 6.0. (No sodium hydroxide has been added).
Figure 20. Plot showing the dissolution rate of copper in 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 PH Gl- Gl± G1+ Plot Rate (Moles/l) (moles/1) (moles/l) (moles/l) (Mg/cm.2/hr/M)
0,100 11.3 0.0972 0.0028 0,000 50.0 0.100 6.0 0(10-4) 0,100 0(10-4) 32.5 0.110 4.0 0,000 0.108 0.002 0.2 0.120 2.6 0,000 0.073 0.047 0
From data shown in Table VIII it 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 it appears that the rate of dissolution can be expressed by the equation:
R - kG1± [Gl±]+ kQ1- [Gl-J
The values computed for the rate constants are:
For Gl±(ph 6)
R = 3.25 = kQ1±[o.l] + kG1-["0j
-2 -1 and kG1± = 32.5 Mg Cu cm. H M£J-+/1.
For Gl"(pH 11.3)
R = 5.0 = 32.5 [0.0028] + kG1_fo.0972j
-2 kG1_= 50.5 Mg Cu cm. H'V^./l.
D. Alpha-Alanine System
Introduction
The anionic form of a-alanine forms two bonds with cupric ion resulting in five-membered ring. The ionization of the latter amino acid is seen by considering the ampholyte cation as a dibasic acid: -H -H
CH3 —.CE-—COCH CH3 — CH^COO" CH3 -~ CH — COO" (1)
+ NH3 PKX=2.34 Ni3+ ' pK2-9.87 NH2
The neutral form written as an amphion indicates that the ionization of the carboxyl group probably occurs when an acid solution containing the cation is partly neutralized.by hydroxyl ion, while the removal of hydrogen ion from the substituted amino group probably occurs when alkali is added to the neutral 9,10,11,17 form.' 17 18 The ionization constant^ * for a-alanine can, thus, be expressed - 45 - by:4
±]LH-tl IO"2'34 (2) and [Alg-JfH+J 1079,87 (3) rAla Ala± - [Ala+J L J
± For convenience, the zwitterion, the a-alaninium and the a-alaninate species
will be abbreviated respectively as follows: Ala* ; Ala+ and Ala".
The bidentate chelate formation with cupric ion may be represented as follows:
+2 Cu + 2 CH3 - CH - C00~=± CO - 0 / NH2 - CH - CH3 k ^ \ \c\ / M •
CH3 - CH - - NH2 .O.r CO
18 The stability constant for 1': 1. chelates
[ Cu Ala] = lO8,40^) [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 in Figure 22,, The re fo rethe only species present was the alaninate ion* and the bidentate chelate_,fprmation may be expressed as follows:
+2 Cu + 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 -
20 1 Ephraim 9» i2L,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 (CH3 - CH - C00~)
NH2 concentration was studied. The results are summarized in Figures 23 and 24.
•At high oxygen pressure, the rate of solution was found to be independent of the oxygen partial pressure. In this region, in which the rate is chemically controlled at the copper surface, the rate of solution was found to be first order in a-alaninate and zero order in oxygen. The total rate for copper--a- alaninate is expressed as follows:
1 R - k [a-alaninate] [oa]° exp _2 -3,-1/ \ The experimental rate constant k = 56 Mg Cu cm. H M (Al JL a-Al- a
3. Effects of oxygen pressure and concentration of d-alanine in water.
Making use of equation (l), pK, andlthe•Ibhizationbconstant it can be shown that in the absence, of: additionali'sod'ium- hydroxide, the species; present for a-alanine in aqueous solution (pH 6.0) is almost completely in the' form of the zwitterion:
[Al±] = 103.5 5
It has been found, as in the previous case, that the rate of solution of copper in aqueous a-alanine, at different oxygen pressures, is made up of the contributions of two separate regions, BBut a further investigation was carried out, only in the region of high oxygen pressure, in which.it appears that the concentration of oxygen and consequently, its rate of transport to the surface of the copper are sufficiently high so that the rate of dissolution 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 is a linear 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 25°C., stirring velocity 775 R.P.M., pH adjustment to 11.3 by NaOH.
Figure 23. Rate curves for. the dissolution of copper in a-alaninate. 0.1 0.2 0.3 Alpha-Alaninate (moles/l.)
Conditionss Oxygen pressure 6.45 atm., temperature 25°C., stirrer velocity 775 R.P.M., pH adjustment to 11.3 by NaOH.
Figure 24. Effect of total a-alaninate concentration on the rate of solution. - 50
Time (minutes)
Conditions; Oxygen pressure 6„A.5 atm„, temperature 25°C0S stirrer velocity 775 RoP.M,, pH 6,00,, No sodium hydroxide has been added.
Figure 25, Dissolution rate of copper in aqueous a-alanine at various concentrations. 0.1 0.2 0.3 Alpha-Alanine (mole)
Conditions: Oxygen pressure 6.45 atm., temperature 25°C., stirrer velocity 775 R.P.M., pH 6.00.
Figure 26. Effect of concentration of aqueous a-alanine on the rate of solution. - 52 - concentration and independent of oxygen pressure, was found to be:
exp 9 _1 _i
kA1 + » 37.0 Mg Cu cm. H M /l. a"
4. Effect of H+ ion
A series of experiments was performed, in which an excess of a-alanine and an equivalent amount of HCIO^ were added. The molecular a-alanine was kept constant. The results reported in the following Table show that the rate fell rapidly to near zero with increasing of H+ ions.
TABLE IX
Effect of Hydrogen Ion in the Alpha*. Alanine-Water System.
a-Alanine (moles/1) 0.10 0.11 0.12 HGIO* (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+) in solution has increased more than one hundred to one thousand fold over the a-alanine water system, the concentration of the free amine (Al~) being decreased proportionately.
5. Summary
The results obtained for the different species in Alpha Alamine system are summarized in Table X. - 53 -
TABLE X
Relative Concentrations of Alpha-Alanlnate Alpha-'Alaninium and Zwitterion Species.
Total a-Alanine PH Ala* Al* Ala" Plot Rate (moles/l) 2 (moles/l) (moles/l) (moles/l) (Mg/cm. /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 - kA1, [A3|] + kA1. | Al*]
The computed rate constants are:
Zwitterion species (Al±)-pH 6.3
E - 3.7 - k [o.l] • k [o] . a 2 1 ml and kA]^ - 37.0 Mg Cu cm." H" M £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 +/d N Cu + 2 NH2 CH2- CH2 - C00-^±CH2 Cu | CH2 N X CH2 - NH2 0 — CO The stability constant for the complex and the hydrogen ionization 17 IS A constants '* respectively are given by: 7,15 feu A1+ ] = 10 (k1) 1:1 chelates 10 6 [Alt] [H*] . 10-3-6 AND [A1-] [H+] = io- -3 Sodium hydroxide was added up to the second equivalent mid-point (pH 11.6), so as to have definite species in 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 interesting from the standpoint of the present work, a high partial pressure of oxygen was maintained above the solution. Rate curves showing the effect of varying the B-alaninate concentration, on the dissolution of copper, are shown in Figures 27 and 28. As in the other systems, in the region of chemical control, the reaction was found to be first order in B-alaninate of zero order in oxygen. A The B-alaninate, the zwitterion and the p-alaninium species will be abbreviated, respectively, as follows: Aig ; Al| and Al^ B-Alaninate concentration O o.l Molar .2 15 30 45 Time (minutes) Conditions: Oxygen pressure 6.45 atmw, temperature 25°C, stirrer velocity, 775 R.P.M. B-alamine plus equivalent amount of NaOH, pH 11.6. Figure 27. Rate curves for the dissolution of copper in B-alaninate. - 56 - B-Alaninat e (mole s/1.) Conditions: Oxygen pressure 6.45 atm., temperature 25°C., stirrer velocity 775 R.P.M., pH 11.6. Figure 28. Effect of total; B-alaninate concentration on the rate of solution. - 57 - For Al", the slope was calculated to be: exp 11.4 Mg Cu cm, 3. Effect of oxygen pressure and concentration of 3-alani.ne in water. No sodium hydroxide was added initially to the media. The pH of these solutions was 6.6. At this pH it can be computed that all but 0.1$ (Al± = 10 ) of the 3-alanine. is in.the zwitterion form. The experimental rate constant for the dissolution of copper in solution of 3-alanine in aqueous solution, determined from the plot in Figure 29, is equal to: exp i.-2 H"1 M"1/!. kA1± = 3.05 Mg Cu cm, 3 4. Effect of H+ ion The effect of H+ ion was investigated, keeping the concentration of 3-alanine constant. It was found that the rate decreased virtually to zero. By addition of H+ ion, the concentration of 3-alaninium ion (Al*) has increased more than ten fold over the 3-alanine water system. 5. Summary The concentration of the different species of respective constants obtained in Beta- Alanine system are summarized in Table XI, - 58 - 3-Alanine (moles/l.) Conditions: Oxygen pressure 6.45 atm., temperature 25°C., stirrer velocity 775 R.P.M., pH 6.6 Figure 29. Effect of concentration of 6-alanine on the rate of solution. - 59 - TABLE XI Relative Concentrations of Beta-Alaninates Beta-Alaninium and Zwitterion Species. A1 Total 3-Alanine p' Al|,.. PH /1 (moles/l) mole's moles/l moles/1 Rate constant Mg/cm.2/hr/M/ w. 0.100 11.6 0.0946 0.0054 0.000„ 11.40 0(10-3) 3 0.100 6.6 0.100 0(10~ ) 3.05 0.00 0.110 4.65 0.000 0.101- 0.009 The rate of solution can be related as follows? R "B - ^ B The computed rate constants are listed belows Zwitterion species (Al±) pH 6.6 R = 0.305 = k + k A1 A1e i -2 1 _ therefore kA1+ = 3.05 Mg Cu cm. H" M ^/l4 0 Anion species (Al~) pH 11.6 R » 1.14 = 3.05 0.0054 + kAl4 0.0946 - B - 3 -2 1 _1 and k - 11.8 Mg Cu cm. H" M /i# A±B As in the previous amino acid systems studied, in the presence of significant Beta-Alaninium (Alp) concentration, a passivating effect on the copper surface is observed. - 60 - CONCLUSIONS The observations and correlations are sufficiently consistent that it is possible to account for some of the observed effects and to draw some conclusions concerning the kinetics and mechanisms of the reactions involved in the dissolution in the presence of the several chelating agents on copper surface. In all 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 likely controls the rate of the chemical reaction and the rate is directly proportional to the pressure and independent of the concentration of the chelating agent. In so far as the latter is concerned the results provide only information on the transport mechanism. Therefore, in-all the systems studied, it was desirable to investigate the kinetics under conditions where the chemical process itself was rate-controlling. In this latter region (high pressure), it appears that the rate of transport of oxygen to the surface of copper is sufficiently high that the rate of dissolution becomes controlled by the chemical reaction and independent of the concentration of oxygen. For limited conditions, with low concentration of the chelating agent, the reaction fits an empirical equation of the fbrms Rate = k [chelating agent] [Wj° (l) The measured rate of dissolution of copper metal in each of the different systems studied, is made up of the contribution of two separate reactions. The species giving rise to these reactions appear in all cases to be the ligand and its acid form. For example, in ethylenediamine system - ethylene- diamine and ethylenediaminium; and in glycine system - glycinate (Gl~) and the zwitterion species (Gl*). 61 - In each system the reaction is independently first order with respect to the concentration of each of these two species which fit the following equation rate: R = kX[L ]+ k2 [LH*] where L is the chelate forming species. The mechanism demised by Halpern to explain the results in the ammoniax system can be applied equally well to the systems studied in the present work, as illustrated for amines and amino acids as follows, using ethylenediamine and glycine as examples, respectively, I. Oxygen is chemisorbed rapidly on the copper surface: fast Cu + 1/2 02 -+ Cu.,.,0 metal dissolved II, Reaction of ethylenediamine molecule or ethylenediaminium ion with the copper oxygen complex: NH - CH - CH - NH slow y 2 2 2 2 Cu.,.,0 + en Cu+t o~ ++ + HOH^CufNHp - CH2 - CH2 - NH2) + 2 OH" fast slow y NH2 - CH2 - CH2 - NH2 or Cu.,.,0 + enH+ Cu++" \ ^ 6- ++ + OH- y Cu(NH2 - CH2 - CH2 - NH2) +2 OH" fast Ila, Reaction of glycinate anion or zwitterion species with the copper oxygen complex (the first step is the same in both systems). slow NH2 - CH2 - C00- Cu....O + Gl~ Cu++ + HOH fast' + Cu(NH2 CH2 C00) + 2 OH" slow XNH2 - CH2 - C00- or Cu....O + Gl± -* Cu++" H + OH- x -6- fast + 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 itself and shows also that any mechanism involving oxidation of copper to cuprous by cupric ion can be ruled out. The fact that the rate of dissolution for the ethylenediamine, glycine, a-and B-alanine systems :is "independent of the concentration of oxygen gives evidence that the first step is fast, also that the surface is covered with a film 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 for 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 NH3 61 (experimental) + Ammonium NH4 1550 (Halpern) Ethylenediamine en 245 Ethylenediamine enH+ 515 Glycine Gl- 50.5 Glycine Gl± 32.5 a-Alanine A15 56,5 a-Alanine Alt a 37.0 3-Alanine Alf ii;e, B-Alanine Al* 3I05 3 - 63 - With respect to the preceding data, there are certain significant trends which warrant further additional remarks. It will be observed from Table XII that the addition of a proton to 1 NH3 or en'increases the rate of its attack on copper. On the other hand, in the case of the amino aKids studied, the opposite effect is obtained. It is further observed that In all 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 stability constants of the species involved with copper and their respective computed rate constants. It appears that a significant relationship exists between these two constants. For instance, the neutral species 1, 2, 3, and 4 show the stability constant (log k-|) to be a function of the rate constant. The corresponding protonated species show a displaced similar functional relationship. If additional points corroborated these limited curves, the results would be significant. One could utilise this to predict the rate constant from the stability constant value. TABLE XIII Correlation of Rate Constants and Stability Constants. Species Stability constant Rate constant log k]_ (chelates 1:1) Mg/cm,2/hr/M/l.; (1) en 10.7 245 (2) Gl- 8.2 50.5 (3) Al- 8.4 56.5 (4) Alp 7.15 11.8 Surprisingly, comparison of the rate constants obtained for glycine and a-alanine, respectively, indicates that the expected steric effects, due to the branched amino group of a-alanine, does not occur. In fact, the rate constants obtained for the d-alanine species are even slightly greater than - 64 - the ones for glycine. Recommendat ions There remain several questions pertaining to the corrosion of copper in amines and amino acid solutions which have not been yet answered. These include (a) the activity of the diprotonated ethylenediamine cation which must be obtained in 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, this part of the ethylenediamine system can be completed by a series of experiments in acid solutions. The amino acid systems can probably be clarified by a carefully chosen sequence of experiments in 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 clarify the nature of the corrosion inhibition observed in the amino acid systems where H+ ions have been added, although it is believed that if the concentration of free amine goes below a critical value, passivation results from hydrolysis of the corrosion product. It would also be of great interest to study other amines and amino acids, and relate more closely the stability constant, the number of carbon of the ligand, the possible steric effect, 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 Pressure1», M.A.Sc. Thesis. Univ. of British Columbia, (1953). 6. 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