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The dissolution of cupric hydroxide films from copper surfaces
THE DISSOLUTION FROM
OF CUPRIC HYDROXIDE COPPER SURFACES
FILMS
D. W. SHOESMI’M and W. LEE Research
Chemistry Branch, Whiteshelf Nuclear Research Establishment, Pinawa, Manitoba ROE ILO, Canada (Received 25 May 1976)
Ah&act-The dissolution of cupric hydroxide films grown under well defined potentiostatic conditions has beau studied in Lithiumhydroxide as a function of hydroxide ion concentration, electrode rotation speed (at a r&r), and formation potential The dissolution reaction was followed by monitoring the amount of cquic hydroxide on the copper surface. eleotrochemically at any given time. The open-circuit potential was used as an indication of when the dissolution reaction was complete. The dissolution reaction was found to bc jointly controlled by surface kinetics and diffusion of the dissolution uroduct awav from the surface. The effect of incomplete dissociation of LiOH was also considered’
6.
NOMENCLATURE A
surface
area
of cupric
hydroxide
t
crystals (ml) [Cu(OH):-“],=,
CCU’I R E0 EOC F h
Kd
ki k-d k. L
k, KS0 n
QX= Qcuoo) R,
R*
concentration of the nth cupric species at the electrode surface (mol dm-s) concentration of the nth cupric species in the bulk of the solution (mol dnm3) concentration of cuprous ion (mol dm-3) diffusion coefficient of the nth cupric species (m’s_‘) standard electrode potential (mV) open circuit potential (mV) the Faraday (=9649OC mol-‘) flux of the nth cupric species (mol s-‘md2) equilibrium constant for the disproprtionation reaction (mol m- 1 disproportionation forward rate constant (s- ‘) disproportionation reverse rate constant Im%no-‘s-l) rate constant for Cu(O& dissotution (m4mol-is-‘: for n = 4) rate ‘constant for deposition of cupric species (ms-t) rate constant for Cu10 precipitation (m4mol-‘s-‘)
Y
w
layer thickness cm) transition time(s) sohltion viscosity (In%- I) electroderotation speed (Hz) steady state difision
INTRODUCIION
The dissolution of oxide/hydroxide films is important in the power industry in a variety of contexts For example a knowledge of the kinetics of the diss&ttion of nuclear fuel deposits and boiler circuit corrosion products is essential to an un&rstanding of the transport of radioactive material around a nuclear power system, and in developing methods of removing such activity during a plant decontamination. In the present paper the kinetics of dissolution of cupric hydroxide films grown on copper electrodes in lithium hydroxide are studied Lithium hydroxide was chosen, because this is the electrolyte used for pH control in the CANDU (Canada Deuterium Uranium) reactor. When a copper electrode is potentiostatically oxidised in alkaline solution a film of cupric hyclroxide is formed, consisting of two distinct layers[l], a porous base layer of cupric hydroxide grown by a solid state mechanism and an upper layer of crystals formed by nucleation and growth from aoh~tion. If solubility product for Cu,O the electrode is then allowed to stand on open-circuit (moPIn- ‘2) in alkaline solution, the electrode potential eventually number of complexed OHions undergoes a sharp transition which has been shown total anoclic charae 1C) to be due to a phase transition from Cu(OH), to charge due to a&e ‘cupric ion Cu,O[2]. A detailed investigation of this transition dissolution during film formation showed that it involved the steps: (1) CUE disso(C) charge due to CL@ (C) lution; (2) diaproportionation of the cupric ion to ~,(~~~~ dissoluuon rate (mol cuprous ion; (3) precipitation of CuzO on the electrode. If the solution is stirred, the majority (>90”/,) p?O_ 2’precipitation rate (mot of the dissolved cupric species are transported into m ) solution and steps (2) and (3) are relatively unimportime(s) tant. Under these conditions the open-circuit potensol&i& volume (m’) tial transition can be used as an indication of when number of moles of Cu,O distance from electrode surface (m) the Cu(OH)? dissolution is complete. 1411
D. w.
1412
~HOESMIT’H
AND
w.
LEE
This paper uses this transition as the basis of a method for following the kinetics of the dissolution reaction. EXPERIMEmAL
Equipment, cell design, electrode preparation, and solution specifications have previously been described[l, 23. All potentials are referred to the saturated calomel electrode. The procedure used in measuring dissolution rates was as follows: Cupric hydroxide films were grown poteutiostatically on cylindrical (A = 5.2~4 or rotating disc (A = 2.8 cm’) copper electrodes. The majority of the films were grown at a potential of -25OmV, but a series of measurments were performed at potentials in the range -180 mV to - 250 mV. For LiOH solutions of concentraton 1.0 mol drnm3 up to 5.3 mol dm- 3 dissolution experiments were performed in the same solution in which the films were grown. For dissolution measurements in solutions of hydroxide concentrations less than 1.0 mol dmm3 the surface Urns were grown in 1.0 mol drnm3 solution. The electrode was then removed from this solution, rinsed quickly with triply distilled water, and transferred to another cell containing a previously deoxygenated solution of the desired hydroxide concentration. The open-circuit electrode potential was then fot lowed for a known period of time during which dissolution of the film occurred. Then the potential was pulsed to E = - 1300 mV and the reduction current transient recorded. If dissolution had been followed in solutions of [OH-] =z 0.5 mol dmm3, the electrode was quickly transferred back to the i.Omol dmm3 solution before the reduction was performed. Figure 1 shows the potential profile used in the experiinents. RESULTS AND DISCUSSION
0
nwo
4wJ
eew
Time(*)
Fig. 2. Open-circuit potential-time transient obtained in 1 mol dm-” LiOH after oxidation of a copper electrode at -250 mV on a cylindrical electrode (A = 5.2 cm’). Points (a) to tj) indicate the times at which the potential was pulsed to - 1300 mV. transients obtained for each of ‘these times, t. Integration of these plots yields Q(t), the amount of film remaining on the electrode surface at the time, t, the dissolution was terminated. The amount of film originally present on the electrode (Qr) is obtained by integrating the anodic transient, The anodic charge, QA, is the sum of two contributions,
QA = QAD + QF
(1)
where QP is the actual charge consumed in forming the film and QAo is the charge due to active dissolution during film formation. QAr, was estimated by reducing the film immediately after the anodic growth. Under these conditions the cathodic charge is given by Qo=QF=
QA -QAD
(2)
Q usually accounts for 1545% of QA depending or%e potential used to grow the film. The reproducibility of QAD is poor, and therefore the correction
Figure 2 shows the behavionr of the open-circuit potential in 1.0 mol dnm3 LiOH. The letters (a) to (j) on the potential curve indicate where individual runs were terminated and the potential pulsed back to - 1350 mV. Figure 3 shows the series of reduction
Fig. 1. Potential-time profile used to measure the dissolution rate.
Fig. 3. Potentiostatic reduction transients were recorded at -1300mV in 1 mol dm-” LiOH after varying times on open circuit.
The dissolution of cupric hydroxide
films from copper surfaces
Fig. 4. Plot of QA (o, 0) and Q. + QAD(e, I) as functions of the
at -250
1413
Data recorded for oxidation Dissolution was also ~~3.3Hz.
mV on a cyliidrical electrode (A = 5.2 cm*) in 1 mol dnm3 LiOH. in 1 mol dme3 LiOH; +-no stirring; D-magnetically stirred at
is only approximate. The potential dependence of QAD of -365 mV to -390 mV. Similarly in 0.5 mol dm-” has been discussed pretiously[l]. Under the conclithe full transition does not occur for w > 50 Hz. tions of the present experiment, Nevertheless the partial transition which does occur can still be used as an indication of the completion Q(t) = QF - Qo = QA - QAD - QD (3) of cupric hydroxide dissolution. The final potential where Q. is the charge lost by dissolution on openof - 365 mV to - 390mV would indicate that once circuit. the hydroxide has dissolved the electrode returns to In Fig. 4, QA and Q, + QADare plotted as a funca more active state. tion of time for two different conditions of stirring. Not only was this method more convenient for The effect of stirring on the transition time was noted measuring the dissolution rate it was the only viable pretiously[2]. These results were recorded on the method to use for films grown on a rotating disc eleccylindrical electrodes since the reproducibility in QA trode. This is because the values of QA for films grown was much better on this electrode. than on the rotaton disc electrodes is much less reproducible than at ing disc electrode. QADis shown as an intercept on the cylindrical electrode. This leads to large variations the charge axis. Thi reproducibility in the transition in T and hence excludes the application of the method time z is indicated in this plot and is in the range suggested for the data of Fig. 4. This k-reproducibility is demonstrated by the data of Table 1 which show 5-6%. The slope of such plots yields a value for the dissothree runs under the same conditions at a disc eleclution rate of cupric hydroxide. The clissolution is trode. The data show that, despite the fluctuations complete by the time the transition is reached, and in Q,, and T, the dissolution rates, calculated using occurs linearly with time except for times very close (4) are fairly reproducible. to the transition, Hence to a close approximation, This variation in QA at the disc electrode may be the dissolution rate can be calculated from the equadue to the variation in convective diffusion conditions tion, R =
(4)
UQA - QAD)- Q,l/r
where r is the transition time, and Q, is measured immediately after the transition has occured. Equation (4) was used to calculate the dissolution rate as a function of electrode rotation speed and hydroxide ion concentration. In 2.0 mol dm- 3 a transition is always observed even at the highest rotation speeds. This is not always the case for lower [OH-]. For 1.0 mol dm-3 the transition is much less pronounced for w > 66 Hz the final potential being in the vicinity Table
1. Data recorded on a rde in 1.0 mol dm-’ R
from equation (4) (CKl)
0.963 0.366 0.666
10 10 10
0.089 0.035 0.048
3270 1200 2040
LiOH
x 104 2.67 2.80 2.98
Fig. 5. Dilution rate recorded on a rde as a function of electrode rotation speed; (1) 0.5 mol drnm3 UOH; (2) 1.0 mol dmL3; (3) 2.0 mol dmm3.
D. W. Srroawrt~ ANDW. LEE
1414
at the electrode surface during the nucleation and growth of the films. The precipitation of material iu the solution close to the electrode leads to the development of convective pathways in the solution. These effects are less reproducible at the horizontal disc surface than at the vertical walls of the cylindrical elec-
Under these conditions, bined to yield
trOdJ3.
where S,, has been taken for the case of the rotating disc. Hence for dissolution as n different species, we can write,
dissolution rate data are plotted as a function of electrode rotation speed in Fig. 5 for 2.0, LO and 0.5 mol dme3 hydroxide concentrations. The data is typical of a process controlled jointly by diffision and surface kinetics At high electrode rotation speeds, the rate is tending to become independent of rotation speed, indicating that control is via a surface kinetic step. Dissolution can be represented by the equation, The
Cu(OH),(S) + (n - 2)OH-
+
.
CU(OH),~-~
(5)
For the nth species, the rate of dissolution can be written, R, = IcJOH-~“-~
- k_,[Cu(OH),2-“I,=, - k$c,[Cu(OHx
-“I,= 0
(6)
where the last term in (6) represents the rate of formation of CuxO as described in the previous paper[2]. Equation (6) also assumes that the deposition reaction is first order with respect to Cu(OH)k2-“. Since [OH-J,=e %=[Cu(OH),2-q,,, in all the experiments, diffusion control is due to the flux of cupric species away from the electrode surface and not to the flux of OH- to the electrode and the flux is given, for the nth species, by
j. = D,, $[Cu(OH),-‘1)
r=*
= +J(OH):-q,=. n
% k$i,g
(7)
(8)
then the 6naI term in (6) is negligible. This is true iu the present cast where >900/, of the dissolved cupric hydroxide is carried away to the bulk of solution, and less than 10% appears on the electrode as CtlxO.
At equilibrium, R, + 0 and, k,,[OH-I(“-”
= k_,[Cu(OH~-“l,
Rm = I$,[OH-]“-~
+ 0.62v- r’6~“2D.2’“[CU(OH~-~,
R,’
= R,l
(9)
where [Cu(OH~-“I,, denotes the equilibrium saturation concentration of CIJ(OH~~~. This condition is only applicable providing that the last term in (6) is negligible compared to the deposition rate, ie
(II)
+ Ro-1’2
(12)
where R, = f k,[OH-I”-’
(13)
0
and B = 0.62~+“~ $ (Df’“[Cu(OH),Z-“I,,,)-’
(14)
lf it is assumed that the diffusion coe&ients, 0,. for the various cupric species are approximately the same, then a single diffusion coefficient, D, can te written and (14) becomes B =
0.62v- 1’6D”3 $ [Cu(OH):-“,,.,)
-1
(15)
where the concentration term in (15) represents the solubility of cupric hydroxide at a given hydroxide ion concentration. This concentration term can be calculated using the known constants for the following equilibria Cu(OH)&) + CL?+ + 20H-
Equation (7) assumes that a steady diffusion layer thickness, S, is attained at the rotating disc electrode. The bulk concentration of soluble cupric species is assumed to be close to zero for all times during the experiment. This assumption will apply under the present conditions where the solution volume is very large compared to the volume of the diffusion layer. If, D.&l
(6), (7) and (9) can bc corn-
pZ&= 18.8 (16)
Cu’+ +OH-
2 CuOH+
pK, = -6
Cu*+ -t-20H-
a Cu(OH),
pK, = - 13.33 (18)
Cu(OH),(s) + OH - e Cu(OH); Cu(OH)&)+20H-
P Cu(OH):-
p&
(17)
= 3.6 (19)
p&., = 2.7 (20)
where (s) denotes solid cupric hydroxide. The constants I& Kr, & and K,.+ were taken from the values listed by Feitknecht and Schindler[3]; K, was c&xl&cd from the free energy data of Spivakovskii and Makovskaya[4]. For completeness, polynuclear species should also be eonsidered[5l but this omission, under the present circumstances leads to no sign&ant error, because they are relatively unimportant at alkaline pH values. A plot of the rate data according to (12) yields a means of separating the kinetic contribution to dissolution from the diffusion contribution. This is shown in Fig. 6. The intercepts in Fig 6 yield values of R, from which the reaction order with respect to hydroxide ion concentration can be determined. A logarithmic plot of R, against hydroxide ion concmtration is shown in Fig. 7(a). The slope of this,plot is 2 f 0.2 yielding a value of 4 for n, the number of complexed [OH-] ions in the product The distribution diagram for the hydrolysis of cupric ion, Fig. 8, shows that, in the range 0.5 -=z[OH-] < 2.0 mol dm-‘, the pro
0.Sd
r
I 0.1 w-s
( rada.
9-1
dd3
1415
e
The dissolution of cupric hydroxide films from copper surfaces
_.-’
I
I
0.2
a3
PH
1
Fig 8. Fraction of individual hydrolyzed species present as a function of pH.
) =!!
Fig. 6. Data of Fig. 5 plotted according to equation (12). The right hand scale refers to the data represented by open circlea
then the slopes in Fig. 7(b) and (c) yield n-2
= 2.3 + 0.2 for I&
= 1.2[4,9]
and cjominant hydrolysed species is Cu(OH)f-. Hence the dissolution mechanism could be represented by Cu(OH)&)
+ 2OH- e Cu(OH):-
(21)
There is a substantial amount of evidence for the incomplete dissociation of lithium hydroxide[&lO], and this should be taken into account in determining the dissolution kinetics, There is, however, doubt over the value of the dissociation constant depending on the method used to measure it. If the actual hydroxide ion concentration is determined using alternative values of K,&& for LiOH
a
Li+ + OH-
(22)
Fig. 7. Logarithmic pl,ot of dimolution rates obtained by extrapolation of the plots of Fig. 6 to OJ-“* = 0, (a) K ai” = co; (b) s&j‘, = 1.2: (c) Kd,” = 0.66.
n - 2 = 2.8 + 0.2 for K&_ = 0.66[7,8,
IO]
A reaction order greater than 2 with respect to OHcould be taken to indicate that dissolution involves not only the formation of the species Cu(OH& but also a contribution from a species CI.I(OH~~’ with n > 4 (possibly 6). No such species has so far been postulated, and if further hydrolysis did occur it is unlikely to be important in the range 0.5-2.0 mol dm-J. An alternative, and much m&e likely, explanation for the high reaction order is that the hydroxide ion attacks the film locally leading to the development of cracks and fractures Film break-up in this manner would lead to au increase in the surface area of cupric hydroxide and hence an increase in dissoh~tion rate. Break-up of the cupric hydroxide base layer as dissolution proceeds was shown to occur in the previous paper[2]. Such an effect would be more predominant at high [OH-]. Figure 9 shows a series of dissolution rate values calculated using (4) from data recorded on a cylindrical electrode. The data are plotted .for Kdil = 0.66. In this case the only stirring present was due to a magnetic stirrer revolving at 3.3 Hz in the bottom of the cell For [OH-] > 0.3 mol dmm3 there is nochange in reaction order up to a lithium hydroxide concentration of 5.3 mol dmm3 (saturation). The actual value of 2.1 cannot be taken as an accurate measure of the reaction order because the rate contains a contribution from diffusion under these conditions. For [OH-] < 0.3 mol dn-‘, the reaction order changes and tends towards 1.0. Under these condtious the surface rate equation is composed of two terms;
R = k,[OH-]
+ kJOH-]
(23)
ANDW.hE D. W. SKO~SWTH
1416
. .
.
/
1Icqm= 2.1
f,/
,tis1opa=
1.0
i
Fig 9. Logarithmic plot of dissolution rate as a function of hydroxide ion concentration. Data recorded on cylindrical electrodes (A = 5.2 cm*) in solutions stirred magnetically (3.3 Hz).
Fig. 10. Piot of QA (0) and Q, + QAn (0) as a function of time for dissolution in 0.1 mol dm-’ LiOH. Data recorded on a cylindrical electrode (A = 5.2cm’). Solution stirred magnetically (3.3 Hz).
with the first term becoming predominant for [OH-] c 0.3 mol dn-3. As shown in Fig. 8 this corresponds to the pH range where CutOH); is .becoming the predominant hydrolysed @es. For 0.1 mol dmm3 [OH-] no potential transition was observed and the dissolution rate was measured using the alternative method described above. The data are plotted in Fig. 10 and is less accurate than for higher [OH-]. Table 2. Diffusion coefficients from the data of Fig. 6 D in cm%-’
x 105
The slopes of Fig. 6 should yield a value of the diffusion coelikient for the species Cu{OH)i-. The values are given in Table 2 and depend significantly on the chosen value of Kdiss. The fact that the values are in the range 10-5-10-6cmzs-’ confirms that diffusion is by a solution species. The variation of D with [OH-] suggests that the diffusion co&icients of the two species involved, Cu(OH)z- and Cu(OH);, are not equal. This last condition was assumed in using (15). Equation (14) should be used and for two species can be written B- ’ = 0.62~ - I” (D;‘3 [Cu(OH);]
LiOH mol dmm3
K dk, = co
1.2
0.66
2.0 1.0 0.5
0.17 0.16 0.2
0.9 0.57 0.4
1.7 0.9 0.6
+ D:‘3 [Cu(OH):-1)
(24)
An analysis of the data using this equation (and a vahle of Kd& = 0.66) yields D4 = (1.3 f 0.5) x iOe5 and D3 = (3 + 2) x 10-5cmzs-1. The value for D4 is around the value expeeted[il].
The dissolution
of cupric hydroxide films from copper surfaces
1417
potential region - 180 mV r E L -230 mV, the dissolution rate is independent of potential. At potentials < -230 mV the crystals grow larger and there are less of them. Hence the surface area to volume ratio is smaller. This trend to fewer but larger crystals continues as the potential is decrcased[l], and is reflected in the decrease in dissolution rate at the more negative potentials. Acknowledgement~The author would like to thank Dr. T. E. Rummery and Dr. P. Tremaine for helpful discussion.
-
REFERENCES
-
I -200
I
I
-220 Pol*ntial
I -240
I
I
-2bO
I
(rn.V.)
Fig. 11. E&t of oxidation potential on the dissolution rate in 1.0 mol dm-’ UOH. Solution stirred magnetically (3.3 Hz). The effect of film formation
potential
on the clisso-
lution rate is shown in Fig. 11. For potentials 2 -230 mV, the surface layer of Cu@H), is composed of a large number of small crystals whose size and morphology are almost independent of poteutial[ 1J. consequently the surface area to volume ratio of the film will be large. Consequently in the
I!.*. 22/12--F
1. D. W. Shoesmith, T. E. Rummery, D. Owen and W. Lee, J. electrochem. Sot. 123, 790 (1976). 2. D. W. Shoesmi th, T. E. Rummery, D. Owen and W. Lee, submitted to Electrochin Actu. 3. W. Feitknecht and P. Schindler, Pure appl. Chem. 6, 130 (1963) 4. V. BI Spivakovskii and G. V. Makovskaya, Russ. J. inorg. Chem. if 815 (1968). 5. D. D. Perrin, J. them. Sot. 3189 (1960). 6. L. S. Darken and H. F. Meier, J. Am them. Sot. 64, 621 11942). 7. F. d. R. &mblett and C. B Monk, Trans. Faraday. Sot. So, 965 (1954). &r&y Sot. 55, 1746 (1959). 8. M. Spiro, 7&s. 9. J. M. Wright, W. T. Lindsay and T. R. Druga Jr., AEC Research and Development Report, WAPD-TM-204 (1961). 10. T. Ohtaki. Actu them. scand. I%, 521 (1964). at Solid Electrodes, pp. 11. R. N. Adams, Electrochemistry 220-l. Marcel Dekker, New York (1969).
OF CUPRIC HYDROXIDE COPPER SURFACES
FILMS
D. W. SHOESMI’M and W. LEE Research
Chemistry Branch, Whiteshelf Nuclear Research Establishment, Pinawa, Manitoba ROE ILO, Canada (Received 25 May 1976)
Ah&act-The dissolution of cupric hydroxide films grown under well defined potentiostatic conditions has beau studied in Lithiumhydroxide as a function of hydroxide ion concentration, electrode rotation speed (at a r&r), and formation potential The dissolution reaction was followed by monitoring the amount of cquic hydroxide on the copper surface. eleotrochemically at any given time. The open-circuit potential was used as an indication of when the dissolution reaction was complete. The dissolution reaction was found to bc jointly controlled by surface kinetics and diffusion of the dissolution uroduct awav from the surface. The effect of incomplete dissociation of LiOH was also considered’
6.
NOMENCLATURE A
surface
area
of cupric
hydroxide
t
crystals (ml) [Cu(OH):-“],=,
CCU’I R E0 EOC F h
Kd
ki k-d k. L
k, KS0 n
QX= Qcuoo) R,
R*
concentration of the nth cupric species at the electrode surface (mol dm-s) concentration of the nth cupric species in the bulk of the solution (mol dnm3) concentration of cuprous ion (mol dm-3) diffusion coefficient of the nth cupric species (m’s_‘) standard electrode potential (mV) open circuit potential (mV) the Faraday (=9649OC mol-‘) flux of the nth cupric species (mol s-‘md2) equilibrium constant for the disproprtionation reaction (mol m- 1 disproportionation forward rate constant (s- ‘) disproportionation reverse rate constant Im%no-‘s-l) rate constant for Cu(O& dissotution (m4mol-is-‘: for n = 4) rate ‘constant for deposition of cupric species (ms-t) rate constant for Cu10 precipitation (m4mol-‘s-‘)
Y
w
layer thickness cm) transition time(s) sohltion viscosity (In%- I) electroderotation speed (Hz) steady state difision
INTRODUCIION
The dissolution of oxide/hydroxide films is important in the power industry in a variety of contexts For example a knowledge of the kinetics of the diss&ttion of nuclear fuel deposits and boiler circuit corrosion products is essential to an un&rstanding of the transport of radioactive material around a nuclear power system, and in developing methods of removing such activity during a plant decontamination. In the present paper the kinetics of dissolution of cupric hydroxide films grown on copper electrodes in lithium hydroxide are studied Lithium hydroxide was chosen, because this is the electrolyte used for pH control in the CANDU (Canada Deuterium Uranium) reactor. When a copper electrode is potentiostatically oxidised in alkaline solution a film of cupric hyclroxide is formed, consisting of two distinct layers[l], a porous base layer of cupric hydroxide grown by a solid state mechanism and an upper layer of crystals formed by nucleation and growth from aoh~tion. If solubility product for Cu,O the electrode is then allowed to stand on open-circuit (moPIn- ‘2) in alkaline solution, the electrode potential eventually number of complexed OHions undergoes a sharp transition which has been shown total anoclic charae 1C) to be due to a phase transition from Cu(OH), to charge due to a&e ‘cupric ion Cu,O[2]. A detailed investigation of this transition dissolution during film formation showed that it involved the steps: (1) CUE disso(C) charge due to CL@ (C) lution; (2) diaproportionation of the cupric ion to ~,(~~~~ dissoluuon rate (mol cuprous ion; (3) precipitation of CuzO on the electrode. If the solution is stirred, the majority (>90”/,) p?O_ 2’precipitation rate (mot of the dissolved cupric species are transported into m ) solution and steps (2) and (3) are relatively unimportime(s) tant. Under these conditions the open-circuit potensol&i& volume (m’) tial transition can be used as an indication of when number of moles of Cu,O distance from electrode surface (m) the Cu(OH)? dissolution is complete. 1411
D. w.
1412
~HOESMIT’H
AND
w.
LEE
This paper uses this transition as the basis of a method for following the kinetics of the dissolution reaction. EXPERIMEmAL
Equipment, cell design, electrode preparation, and solution specifications have previously been described[l, 23. All potentials are referred to the saturated calomel electrode. The procedure used in measuring dissolution rates was as follows: Cupric hydroxide films were grown poteutiostatically on cylindrical (A = 5.2~4 or rotating disc (A = 2.8 cm’) copper electrodes. The majority of the films were grown at a potential of -25OmV, but a series of measurments were performed at potentials in the range -180 mV to - 250 mV. For LiOH solutions of concentraton 1.0 mol drnm3 up to 5.3 mol dm- 3 dissolution experiments were performed in the same solution in which the films were grown. For dissolution measurements in solutions of hydroxide concentrations less than 1.0 mol dmm3 the surface Urns were grown in 1.0 mol drnm3 solution. The electrode was then removed from this solution, rinsed quickly with triply distilled water, and transferred to another cell containing a previously deoxygenated solution of the desired hydroxide concentration. The open-circuit electrode potential was then fot lowed for a known period of time during which dissolution of the film occurred. Then the potential was pulsed to E = - 1300 mV and the reduction current transient recorded. If dissolution had been followed in solutions of [OH-] =z 0.5 mol dmm3, the electrode was quickly transferred back to the i.Omol dmm3 solution before the reduction was performed. Figure 1 shows the potential profile used in the experiinents. RESULTS AND DISCUSSION
0
nwo
4wJ
eew
Time(*)
Fig. 2. Open-circuit potential-time transient obtained in 1 mol dm-” LiOH after oxidation of a copper electrode at -250 mV on a cylindrical electrode (A = 5.2 cm’). Points (a) to tj) indicate the times at which the potential was pulsed to - 1300 mV. transients obtained for each of ‘these times, t. Integration of these plots yields Q(t), the amount of film remaining on the electrode surface at the time, t, the dissolution was terminated. The amount of film originally present on the electrode (Qr) is obtained by integrating the anodic transient, The anodic charge, QA, is the sum of two contributions,
QA = QAD + QF
(1)
where QP is the actual charge consumed in forming the film and QAo is the charge due to active dissolution during film formation. QAr, was estimated by reducing the film immediately after the anodic growth. Under these conditions the cathodic charge is given by Qo=QF=
QA -QAD
(2)
Q usually accounts for 1545% of QA depending or%e potential used to grow the film. The reproducibility of QAD is poor, and therefore the correction
Figure 2 shows the behavionr of the open-circuit potential in 1.0 mol dnm3 LiOH. The letters (a) to (j) on the potential curve indicate where individual runs were terminated and the potential pulsed back to - 1350 mV. Figure 3 shows the series of reduction
Fig. 1. Potential-time profile used to measure the dissolution rate.
Fig. 3. Potentiostatic reduction transients were recorded at -1300mV in 1 mol dm-” LiOH after varying times on open circuit.
The dissolution of cupric hydroxide
films from copper surfaces
Fig. 4. Plot of QA (o, 0) and Q. + QAD(e, I) as functions of the
at -250
1413
Data recorded for oxidation Dissolution was also ~~3.3Hz.
mV on a cyliidrical electrode (A = 5.2 cm*) in 1 mol dnm3 LiOH. in 1 mol dme3 LiOH; +-no stirring; D-magnetically stirred at
is only approximate. The potential dependence of QAD of -365 mV to -390 mV. Similarly in 0.5 mol dm-” has been discussed pretiously[l]. Under the conclithe full transition does not occur for w > 50 Hz. tions of the present experiment, Nevertheless the partial transition which does occur can still be used as an indication of the completion Q(t) = QF - Qo = QA - QAD - QD (3) of cupric hydroxide dissolution. The final potential where Q. is the charge lost by dissolution on openof - 365 mV to - 390mV would indicate that once circuit. the hydroxide has dissolved the electrode returns to In Fig. 4, QA and Q, + QADare plotted as a funca more active state. tion of time for two different conditions of stirring. Not only was this method more convenient for The effect of stirring on the transition time was noted measuring the dissolution rate it was the only viable pretiously[2]. These results were recorded on the method to use for films grown on a rotating disc eleccylindrical electrodes since the reproducibility in QA trode. This is because the values of QA for films grown was much better on this electrode. than on the rotaton disc electrodes is much less reproducible than at ing disc electrode. QADis shown as an intercept on the cylindrical electrode. This leads to large variations the charge axis. Thi reproducibility in the transition in T and hence excludes the application of the method time z is indicated in this plot and is in the range suggested for the data of Fig. 4. This k-reproducibility is demonstrated by the data of Table 1 which show 5-6%. The slope of such plots yields a value for the dissothree runs under the same conditions at a disc eleclution rate of cupric hydroxide. The clissolution is trode. The data show that, despite the fluctuations complete by the time the transition is reached, and in Q,, and T, the dissolution rates, calculated using occurs linearly with time except for times very close (4) are fairly reproducible. to the transition, Hence to a close approximation, This variation in QA at the disc electrode may be the dissolution rate can be calculated from the equadue to the variation in convective diffusion conditions tion, R =
(4)
UQA - QAD)- Q,l/r
where r is the transition time, and Q, is measured immediately after the transition has occured. Equation (4) was used to calculate the dissolution rate as a function of electrode rotation speed and hydroxide ion concentration. In 2.0 mol dm- 3 a transition is always observed even at the highest rotation speeds. This is not always the case for lower [OH-]. For 1.0 mol dm-3 the transition is much less pronounced for w > 66 Hz the final potential being in the vicinity Table
1. Data recorded on a rde in 1.0 mol dm-’ R
from equation (4) (CKl)
0.963 0.366 0.666
10 10 10
0.089 0.035 0.048
3270 1200 2040
LiOH
x 104 2.67 2.80 2.98
Fig. 5. Dilution rate recorded on a rde as a function of electrode rotation speed; (1) 0.5 mol drnm3 UOH; (2) 1.0 mol dmL3; (3) 2.0 mol dmm3.
D. W. Srroawrt~ ANDW. LEE
1414
at the electrode surface during the nucleation and growth of the films. The precipitation of material iu the solution close to the electrode leads to the development of convective pathways in the solution. These effects are less reproducible at the horizontal disc surface than at the vertical walls of the cylindrical elec-
Under these conditions, bined to yield
trOdJ3.
where S,, has been taken for the case of the rotating disc. Hence for dissolution as n different species, we can write,
dissolution rate data are plotted as a function of electrode rotation speed in Fig. 5 for 2.0, LO and 0.5 mol dme3 hydroxide concentrations. The data is typical of a process controlled jointly by diffision and surface kinetics At high electrode rotation speeds, the rate is tending to become independent of rotation speed, indicating that control is via a surface kinetic step. Dissolution can be represented by the equation, The
Cu(OH),(S) + (n - 2)OH-
+
.
CU(OH),~-~
(5)
For the nth species, the rate of dissolution can be written, R, = IcJOH-~“-~
- k_,[Cu(OH),2-“I,=, - k$c,[Cu(OHx
-“I,= 0
(6)
where the last term in (6) represents the rate of formation of CuxO as described in the previous paper[2]. Equation (6) also assumes that the deposition reaction is first order with respect to Cu(OH)k2-“. Since [OH-J,=e %=[Cu(OH),2-q,,, in all the experiments, diffusion control is due to the flux of cupric species away from the electrode surface and not to the flux of OH- to the electrode and the flux is given, for the nth species, by
j. = D,, $[Cu(OH),-‘1)
r=*
= +J(OH):-q,=. n
% k$i,g
(7)
(8)
then the 6naI term in (6) is negligible. This is true iu the present cast where >900/, of the dissolved cupric hydroxide is carried away to the bulk of solution, and less than 10% appears on the electrode as CtlxO.
At equilibrium, R, + 0 and, k,,[OH-I(“-”
= k_,[Cu(OH~-“l,
Rm = I$,[OH-]“-~
+ 0.62v- r’6~“2D.2’“[CU(OH~-~,
R,’
= R,l
(9)
where [Cu(OH~-“I,, denotes the equilibrium saturation concentration of CIJ(OH~~~. This condition is only applicable providing that the last term in (6) is negligible compared to the deposition rate, ie
(II)
+ Ro-1’2
(12)
where R, = f k,[OH-I”-’
(13)
0
and B = 0.62~+“~ $ (Df’“[Cu(OH),Z-“I,,,)-’
(14)
lf it is assumed that the diffusion coe&ients, 0,. for the various cupric species are approximately the same, then a single diffusion coefficient, D, can te written and (14) becomes B =
0.62v- 1’6D”3 $ [Cu(OH):-“,,.,)
-1
(15)
where the concentration term in (15) represents the solubility of cupric hydroxide at a given hydroxide ion concentration. This concentration term can be calculated using the known constants for the following equilibria Cu(OH)&) + CL?+ + 20H-
Equation (7) assumes that a steady diffusion layer thickness, S, is attained at the rotating disc electrode. The bulk concentration of soluble cupric species is assumed to be close to zero for all times during the experiment. This assumption will apply under the present conditions where the solution volume is very large compared to the volume of the diffusion layer. If, D.&l
(6), (7) and (9) can bc corn-
pZ&= 18.8 (16)
Cu’+ +OH-
2 CuOH+
pK, = -6
Cu*+ -t-20H-
a Cu(OH),
pK, = - 13.33 (18)
Cu(OH),(s) + OH - e Cu(OH); Cu(OH)&)+20H-
P Cu(OH):-
p&
(17)
= 3.6 (19)
p&., = 2.7 (20)
where (s) denotes solid cupric hydroxide. The constants I& Kr, & and K,.+ were taken from the values listed by Feitknecht and Schindler[3]; K, was c&xl&cd from the free energy data of Spivakovskii and Makovskaya[4]. For completeness, polynuclear species should also be eonsidered[5l but this omission, under the present circumstances leads to no sign&ant error, because they are relatively unimportant at alkaline pH values. A plot of the rate data according to (12) yields a means of separating the kinetic contribution to dissolution from the diffusion contribution. This is shown in Fig. 6. The intercepts in Fig 6 yield values of R, from which the reaction order with respect to hydroxide ion concentration can be determined. A logarithmic plot of R, against hydroxide ion concmtration is shown in Fig. 7(a). The slope of this,plot is 2 f 0.2 yielding a value of 4 for n, the number of complexed [OH-] ions in the product The distribution diagram for the hydrolysis of cupric ion, Fig. 8, shows that, in the range 0.5 -=z[OH-] < 2.0 mol dm-‘, the pro
0.Sd
r
I 0.1 w-s
( rada.
9-1
dd3
1415
e
The dissolution of cupric hydroxide films from copper surfaces
_.-’
I
I
0.2
a3
PH
1
Fig 8. Fraction of individual hydrolyzed species present as a function of pH.
) =!!
Fig. 6. Data of Fig. 5 plotted according to equation (12). The right hand scale refers to the data represented by open circlea
then the slopes in Fig. 7(b) and (c) yield n-2
= 2.3 + 0.2 for I&
= 1.2[4,9]
and cjominant hydrolysed species is Cu(OH)f-. Hence the dissolution mechanism could be represented by Cu(OH)&)
+ 2OH- e Cu(OH):-
(21)
There is a substantial amount of evidence for the incomplete dissociation of lithium hydroxide[&lO], and this should be taken into account in determining the dissolution kinetics, There is, however, doubt over the value of the dissociation constant depending on the method used to measure it. If the actual hydroxide ion concentration is determined using alternative values of K,&& for LiOH
a
Li+ + OH-
(22)
Fig. 7. Logarithmic pl,ot of dimolution rates obtained by extrapolation of the plots of Fig. 6 to OJ-“* = 0, (a) K ai” = co; (b) s&j‘, = 1.2: (c) Kd,” = 0.66.
n - 2 = 2.8 + 0.2 for K&_ = 0.66[7,8,
IO]
A reaction order greater than 2 with respect to OHcould be taken to indicate that dissolution involves not only the formation of the species Cu(OH& but also a contribution from a species CI.I(OH~~’ with n > 4 (possibly 6). No such species has so far been postulated, and if further hydrolysis did occur it is unlikely to be important in the range 0.5-2.0 mol dm-J. An alternative, and much m&e likely, explanation for the high reaction order is that the hydroxide ion attacks the film locally leading to the development of cracks and fractures Film break-up in this manner would lead to au increase in the surface area of cupric hydroxide and hence an increase in dissoh~tion rate. Break-up of the cupric hydroxide base layer as dissolution proceeds was shown to occur in the previous paper[2]. Such an effect would be more predominant at high [OH-]. Figure 9 shows a series of dissolution rate values calculated using (4) from data recorded on a cylindrical electrode. The data are plotted .for Kdil = 0.66. In this case the only stirring present was due to a magnetic stirrer revolving at 3.3 Hz in the bottom of the cell For [OH-] > 0.3 mol dmm3 there is nochange in reaction order up to a lithium hydroxide concentration of 5.3 mol dmm3 (saturation). The actual value of 2.1 cannot be taken as an accurate measure of the reaction order because the rate contains a contribution from diffusion under these conditions. For [OH-] < 0.3 mol dn-‘, the reaction order changes and tends towards 1.0. Under these condtious the surface rate equation is composed of two terms;
R = k,[OH-]
+ kJOH-]
(23)
ANDW.hE D. W. SKO~SWTH
1416
. .
.
/
1Icqm= 2.1
f,/
,tis1opa=
1.0
i
Fig 9. Logarithmic plot of dissolution rate as a function of hydroxide ion concentration. Data recorded on cylindrical electrodes (A = 5.2 cm*) in solutions stirred magnetically (3.3 Hz).
Fig. 10. Piot of QA (0) and Q, + QAn (0) as a function of time for dissolution in 0.1 mol dm-’ LiOH. Data recorded on a cylindrical electrode (A = 5.2cm’). Solution stirred magnetically (3.3 Hz).
with the first term becoming predominant for [OH-] c 0.3 mol dn-3. As shown in Fig. 8 this corresponds to the pH range where CutOH); is .becoming the predominant hydrolysed @es. For 0.1 mol dmm3 [OH-] no potential transition was observed and the dissolution rate was measured using the alternative method described above. The data are plotted in Fig. 10 and is less accurate than for higher [OH-]. Table 2. Diffusion coefficients from the data of Fig. 6 D in cm%-’
x 105
The slopes of Fig. 6 should yield a value of the diffusion coelikient for the species Cu{OH)i-. The values are given in Table 2 and depend significantly on the chosen value of Kdiss. The fact that the values are in the range 10-5-10-6cmzs-’ confirms that diffusion is by a solution species. The variation of D with [OH-] suggests that the diffusion co&icients of the two species involved, Cu(OH)z- and Cu(OH);, are not equal. This last condition was assumed in using (15). Equation (14) should be used and for two species can be written B- ’ = 0.62~ - I” (D;‘3 [Cu(OH);]
LiOH mol dmm3
K dk, = co
1.2
0.66
2.0 1.0 0.5
0.17 0.16 0.2
0.9 0.57 0.4
1.7 0.9 0.6
+ D:‘3 [Cu(OH):-1)
(24)
An analysis of the data using this equation (and a vahle of Kd& = 0.66) yields D4 = (1.3 f 0.5) x iOe5 and D3 = (3 + 2) x 10-5cmzs-1. The value for D4 is around the value expeeted[il].
The dissolution
of cupric hydroxide films from copper surfaces
1417
potential region - 180 mV r E L -230 mV, the dissolution rate is independent of potential. At potentials < -230 mV the crystals grow larger and there are less of them. Hence the surface area to volume ratio is smaller. This trend to fewer but larger crystals continues as the potential is decrcased[l], and is reflected in the decrease in dissolution rate at the more negative potentials. Acknowledgement~The author would like to thank Dr. T. E. Rummery and Dr. P. Tremaine for helpful discussion.
-
REFERENCES
-
I -200
I
I
-220 Pol*ntial
I -240
I
I
-2bO
I
(rn.V.)
Fig. 11. E&t of oxidation potential on the dissolution rate in 1.0 mol dm-’ UOH. Solution stirred magnetically (3.3 Hz). The effect of film formation
potential
on the clisso-
lution rate is shown in Fig. 11. For potentials 2 -230 mV, the surface layer of Cu@H), is composed of a large number of small crystals whose size and morphology are almost independent of poteutial[ 1J. consequently the surface area to volume ratio of the film will be large. Consequently in the
I!.*. 22/12--F
1. D. W. Shoesmith, T. E. Rummery, D. Owen and W. Lee, J. electrochem. Sot. 123, 790 (1976). 2. D. W. Shoesmi th, T. E. Rummery, D. Owen and W. Lee, submitted to Electrochin Actu. 3. W. Feitknecht and P. Schindler, Pure appl. Chem. 6, 130 (1963) 4. V. BI Spivakovskii and G. V. Makovskaya, Russ. J. inorg. Chem. if 815 (1968). 5. D. D. Perrin, J. them. Sot. 3189 (1960). 6. L. S. Darken and H. F. Meier, J. Am them. Sot. 64, 621 11942). 7. F. d. R. &mblett and C. B Monk, Trans. Faraday. Sot. So, 965 (1954). &r&y Sot. 55, 1746 (1959). 8. M. Spiro, 7&s. 9. J. M. Wright, W. T. Lindsay and T. R. Druga Jr., AEC Research and Development Report, WAPD-TM-204 (1961). 10. T. Ohtaki. Actu them. scand. I%, 521 (1964). at Solid Electrodes, pp. 11. R. N. Adams, Electrochemistry 220-l. Marcel Dekker, New York (1969).