Active/Passive Behavior of Copper in Strong Sulfuric Acid
Desmond Tromans* and Tawfk Ahmed
Department of Metals and Materials Engineering, University of British Columbia,
Vancouver, British Columbia V6T 1Z4, Canada
ABSTRACT
A combination of thermodynamic analyses and potentiodynamic polarization tests have been used to study the anodic behavior of Cu in strong H2S04 solutions in the concentration range 1—10 M. The studies were supplemented by chemical analyses of surface films. It was found that concentration-dependent changes in the activity of water played a major
role in determining the anodic behavior and relative stability of corrosion product films. The anodic Tafel slope decreased
from —41 to —31 mV with increasing acid concentration. The onset of limiting current and active—passive behaviors at
higher potentials was determined by the formation of films of hydrated copper sulfate, CuSO4xH2O, and not by formation of oxides. Limiting current behavior was observed in 1 M solutions, where the degree of hydration was x = 5. Welldeveloped passivity occurred in 10 M solutions where x = 1. The results are relevant to industrial electrorefining operations for Cu and indicate that chloride contamination, if present in sufficient amounts, could cause the premature onset
of limiting current behavior (anode "passivity") during refining.
Introduction
occurrence of passivation-type phenomena on Cu
anodes is a troublesome problem during electrorefining of
copper in acidic copper sulfate solutions, because it leads
to increased power consumption.'4 The refining process is
essentially a galvanostatic operation, using acid concenThe
trations of 1.4—1.66 M H2504,2'3 and the term "passivation"
is applied somewhat loosely to describe the situation
where there is a sudden increase in applied voltage across
the cell. The potential increase may be due to the onset of
relatively constant current behavior, where the anodic current density, i, attains a limiting value, L, with increasing
electrode potential, E, or a true active/passive transition
caused by the onset of protective film formation at a critical current density, i, after which a decreases to a lower
passive value, ii,, with increasing E. Often, the formation
of CuSO45H2O and/or Cu20 has been considered to be a
major cause of passivation phenomena. For example, Abe
et al.5 analyzed electrorefining passivation in terms of
mass-transport-limited anodic dissolution leading to precipitation of "CuSO4" on the anode surface. Chen and
Dutrizac'' detected copper sulfate and Cu20 crystals on
industrial Cu anodes. Jin and Ghali2 conducted laboratory
tests on pure and impure Cu anodes and concluded that
passivation was associated with CuSO45H2O formation
under some circumstances and jointly associated with
CuSO4•51320, Cu2O, and CuO under other conditions. Increasing amounts of CuO were observed in stirred solutions relative to static conditions.
The industrial situation is complicated by slime formation on the impure Cu anodes. Slimes are an integral part
of the refining operation and contain noble metal impurities, arsenates, and selenates.3'4 However, their presence
may cause differences in polarization behavior between
pure and impure Cu anodes.3 For example, Cheng and
Hiskey4 suggested the major cause of anode passivation to
be the formation of Cu20 adjacent to the anode surface
and that slimes influence the stability of the oxide film.
Abe et al. believed the slime layer exerted its influence by
increasing the effective diffusion layer thickness, thereby
promoting mass-transport-controlled precipitation of
"CuSO4." A fundamental understanding of passivation phenomena in industrial refining solutions is helped considerably by studying the polarization behavior of pure copper in
simple sulfuric acid (H2S04) solutions and relating the
observed behavior to the likely formation of corrosion product films based on aqueous thermodynamics and E-pH equilibria in the Cu-H2O-SO system. In this manner, Tromans
and Silva9 showed that pure Cu anodes exhibited limiting
current behavior in a solution of 0.1 M H2S04 + 1 M Na2SO4
that was due to the formation of a film of CuSO45H2O.
The current work extends the previous study9 and is
concerned with the anodic behavior of Cu in more concen*
Electrochemical Society Active Member.
trated H2S04 solutions ranging from 1 to 10 M. Polari-
zation behavior in this concentration range has been
reported earlier by Leckie,1° who observed limiting current
and true passivation behavior which he suggested were
associated with formation of "CuSO4" and Cu20, respec-
tively. However, Leckie did not consider fully the thermodynamics of the formation of the corrosion products, particularly the different hydrated forms of CuSO4xH2O, and
he neglected the effects of acid concentration on activities
of H20 and the soluble sulfate-containing species. Activities
in strong HZSO4 solutions are not known reliably, although
the studies of Awakura et al.'113 provide some guidance in
1 M solutions. Thus, a re-examination of the anodic behavior of Cu in concentrated H2SO, solutions is merited.
Experimental
Materials and solutions—Annealed Cu electrodes, with
a mechanically polished surface (1 p.m diamond paste) and
a working area of 100 mm2, were prepared from commer-
cially pure metal (99.96% by weight), using the same
material, heat-treatment, and electrode preparation procedures as those reported previously.9'14 Distilled water
and reagent-grade H2S04 were used to produce four test
solutions with acid concentrations of 1 M (1.04 m), 5 M
(6.28 m), 7 M (9.96 m), and 10 M (17.91 m). where the values
in parenthesis refer to molal concentrations based on the
concentrative properties of sulfuric acid.'1 All solutions
were purged with nitrogen for 1200 s prior to placement in
the polarization cell.
Polarization tests—Linear scan potentiodynamic tests
were conducted in a single-compartment glass cell containing a Pt counter electrode and 800 mL of solution at
24
1 °C. The working electrode was mounted vertically
and potentials were measured with respect to an external
saturated calomel electrode (SCE) interfaced to the test
solution via a salt bridge. The bridge was filled with test
solution to minimize contamination of the cell by chloride
ions from the SCE and was connected to a polyethylene,
Luggin-type capillary that terminated —-2 mm from the
working electrode. The capillary was end-sealed with a
flush-fitting, porous wooden plug. A microprocessor-controlled potentiostat was used to control E and measure
currents, which were transferred to a personal computer
for final analysis.
All tests commenced with a cathodic pretreatment at
— 0.900
VSCE for 1200 s, to reduce any initial surface oxide
films, followed by anodic polarization at a scan rate (E) of
0.1 mV/s, unless stated otherwise. All solutions were
purged with nitrogen throughout each test and stirred
with a polytetrafluoroethylene_coated magnetic stirrer.
Experimental E, reported as V, were referenced to the
standard hydrogen electrode (SHE) via the conversion
V=
VSCE + 0.242 V, which does not include corrections
for liquid junction potentials (Eu,,) between the test solu-
J. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
601
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602
J. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, tnc.
Table I. Estimated activities of species in H2S04 solutions at 298 K (25 °C).
Concn. of H2S04
m
M
0.96
1
1
1.04
7
6.28
9.96
10
17.92
5
[H]
[H20]
0.00207
2.02'
(ref 9)
0.01
0.01
0.01
0.01
(ref 9)
1.04
5.93
9.13
15.5
0.962'
(ref 9 and 10)
[HSOfl
[SO]
>0.00633
o.411a
(ref 11)
(ref 9)
[H,S04]
0.01
0.36
0.84
2.43
1.02
5.91
9.11
15.48
0.965
0.606
0.388
0.147
Experimental values obtained by Awakura et al.11-'3
tion and saturated KC1 (—4.66 M)'5 in the SCE. Corrections
for
were estimated from the Henderson formula,16
using molar concentrations of ionic species and the limiting (infinite dilution) ionic diffusion coefficients of H,
Na, K, CY, SO, and HSO ions.'6 The concentrations of
ionic species in the 1 M solution were calculated from the
second stoichiometric association constant given by
Baes,'7 and concentrations in the other test solutions were
estimated from the ion activities listed in Table I. The
corrections were +17, +31, +34, and
resulting
+39 mV in the 1, 5, 7, and 10 M H2S04 solutions, respecshould be considered approximate and
tively. The
indicative of their order of magnitude. Corrections for the
ohmic potential drop (i,R0) due to the effective ohmic solu-
tion resistance (R0) between the Luggin capillary and
working electrode are discussed in the text.
It was determined from observations of liquid levels in
the reference electrode and bridge assembly that chloride
contamination of the test solution due to leakage of KC1
from the SCE was 3 X iO' M C1. Consequently, addi-
tional polarization tests were undertaken with a mercurous sulfate reference electrode (MSE), consisting of
HgIHg2SO4/saturated K2S04, in order to avoid chloride
contamination. Tests were conducted in 1 and 10 M H2S04
solutions at 0.1 mV/s using the same salt bridge arrangement and experimental procedures previously described.
The concentration of saturated K2S04 at 25 °C is 0.694 m
(0.672 M)," with a mean activity coefficient of 0.233."
Hence, based on the standard reduction potential (E°) of
+0.615 VSHE for the MSE at 25 °C and unit activities of
species,20 the calculated reference potential of the saturated MSE was +0.638 VSHE and all potentials measured as
VMsE were referenced to the SHE via the conversion VS}OE =
+ 0.638 V. The estimated EL,? corrections between the
saturated K2S04 and the 1 and 10 M H2S04 solutions were
+26 and +52 mV, respectively.
VMsE
Corrosion product analysis.—Following polarization,
specimens were washed in water and ethanol and stored in
a desiccator until required for surface examination.
Surfaces were analyzed by X-ray photoelectron spec-
troscopy (XPS), using Alt Ka radiation, or subjected to
chemical analysis in a scanning electron microscope
(SEM) interfaced to an energy-dispersive X-ray (EDX)
system. The EDX detector was fitted with a 0.3 p.m window of low atomic number to allow detection of Ka X-rays
of oxygen. In some cases, the corrosion products were ana-
lyzed by EDX after being scraped from the Cu electrode
with a wooden pick and placing on a graphite substrate.
Results and Discussion
Activities of species and aqueous equilibria—Activities
of aqueous species were required for thermodynamic cal-
culations relating to E-pH equilibria in order to compare
the polarization behavior of Cu with the potentials at
which different corrosion products were likely to form.
Activity data were not available for the component species
in strong H2S04 solutions, although Akawura et al."'3
measured activities in solutions near 1 M which are listed
in the first row of data in Table I. There are some inherent
errors in their results, which they recognize, due to the
unavoidable necessity to apply estimated EL,0 corrections
to their experimental measurements. Consequently, for
very strong solutions 1 M HZSO4 it was necessary to esti-
mate activities. In all subsequent text, activity is distinguished by placing the species within square brackets, e.g.,
water [H20].
The activity of the solvent, [H20], was estimated with
reasonable confidence from the ratio of the equilibrium
vapor pressure of H20 in the acid solution to that of pure
water (Raoult's law). Ratios based on published vapor
pressure data for the aqueous H2S04 system at 1 atm pressure'5 gave [H20] values listed in Table I.
Estimations of the activities of [H], [HSOfl, [SO], and
[H25041 were based on the first and second dissociation
equilibria in eq 1 and 2, respectively
H2S04 = H + HSO, k, = {[HJ[HSOfl}/[H2SO4] = 98.789
[1]
HSO = H + SO, k2 = {[H][SOI}/[HSOfl
= 1.0123 x
10-2 [2]
where the values of the equilibrium constants, k, and k2,
were calculated from published values of standard chemical potentials, p.°, of the relevant species listed in Table iT2'
= AG° = —RT ln k, where v
(i.e.,
is the stoichiometric coefficient of the corresponding reac-
tant and product species in the reaction, G° is the change
in standard chemical free energy of the reaction, and R
and T have their usual meaning). All soluble species were
considered to be in equilibrium and their activity coefficients were assumed to be unity so that dissociation in eq 1
and 2 could be treated stoichiometrically. Thus, in a solution containing a total molal concentration of N H2S04,
stoichiometric dissociation obeys eq 3 and 4
N=
+ [HSOfl + [SOt]
[3]
2N = 2[H2504] + [HSOJ + [Hi
[4]
[H2504]
Hence, knowing N, the four simultaneous equations, eqs 1-4,
may be solved to determine the individual activities.
The estimated activities are listed in Table I for N values
of 1.04 m (1 M), 6.28 m (5 M), 9.96 m (7 M), and 17.92 m
(10 M) solutions. Activity errors are expected because of the
assumption of unit activity coefficients. However, compari-
son of estimates for the 1.04 m solution with the experimental activities for the 1.0 m solutio&''3 in Table II indicated that the estimates are of the right order of magnitude.
Furthermore, since electrode potentials are dependent upon
the logarithm of the activity, any errors are likely to be tolerable for our purposes.
The effects of [H20] on equilibria between the different
hydrated forms of CuSO4xH2O are now examined. Equilibrium between CuSO43H2O and CuSO45H2O is given by
CuSO43H2O + 2H20 = CuSO45H2O, k = 1/[H20]2 = 14.584
[5]
where the equilibrium constant, k, was calculated from data
in Table II. Equation 5 shows that the trihydrate and pen-
tahydrate are in equilibrium when [H20] is 0.262, corresponding closely to a 8.4 M H2S04 solution.'5 Similarly, equi-
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603
i. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
Table II. Standard chemical potentials, jc, of species at
298 K (25 °C).2'
State'
Species
H20
H2504
JI' (kJ/mol)
S0
H
aq
aq
aq
aq
aq
s
s
s
s
s
s
s
•
Cu2
Cu
Cu20
CuO
CuSO45H20
CuSO43H20
CuSO4H20
Potential, V5
112504 solution
(M)
—237.18
—744.63
—756.01
—744.63
1
1150;
Table Ill. Calculated electrode potentials for oxidation of
copper at 25 °C.
1
5
7
10
0
+65.7
F020
+0.463
+0.514
+0.530
+0.557
E0504 Ecuso 1120 ECUSOIH2O F0250 5110
+0.482
+0.482
+0.482
+0.482
+0.399
+0.404
+0.410
+0.422
+0.357
+0.374
+0.392
+0.429
+0.323
+0.353
+0.382
+0.444
0
—148.1
—134.0
—1880
—1399
—916.4
—663.0
They show that as the potential of Cu is raised anodically,
the first oxidation product film to form should be a
hydrated sulfate, not Cu20, and the film should be
libria between CuS04H20 and CuS043H20, and CuSO4 and
CuSO4'1120 are given by eq 6 and 7, respectively
CuS04H20 + 21120 = CuS043H20, k = 1/[H20]2 = 27.82
CuS045H20 in the 1, 5, and 7 M solutions and CuS04H20
in the 10 M solution, consistent with conclusions arising
from eq 5—7. Note that the data in Table II show that oxidation of Cu20 to CuO should not occur until the potential
is 146 mV above F0050 in all test solutions.
During anodic polarization, ion migration occurs in the
electrical field so that the concentrations (and activities) of
ions near the electmde surface will not be exactly the same
as the bulk solution. The effect will be more pronounced for
[6]
migration of it to the Cu anode, because it has a much
CuSO4 + 1120 = CuS04H20, k = 1/[H20] = 696.45 [7]
showing that the trihydrate and monohydrate forms are in
equilibrium when [1120] is 0.19, corresponding closely to
9.2 M H2504.'5 The monohydrate and unhydrated forms
are in equilibrium when [1120] is 1.435 x lO, correspond-
higher mobility than other ionic species.16 Inspection of
CuSO4
's = solid, 1 = liquid, aq = aqueous.
ing to solution concentrations far in excess of 14 M
112504.15 hence, with regard to Table I, the monohydrate,
CuSO4+120, should be the most stable cupric sulfate in
10 M 112504 and the pentahydrate, CuS045H20, should be
the most stable form in the other test solutions.
Equilibrium potentials for oxidation of Cu to Cu20,
CuSO4, CuS04H20, CuS043H50, and CuS045H20 were
determined from the electrode reactions in eq 8a—12a and
their corresponding Nernst equations, eq 8b—12b, using the
data in Table II to calculate the standard electrode potentials (E°) at 298 K (25 °C) via the relationship AG° = —nE°F,
where n is the number of electrons transferred in the stoichiometric electrode reaction and F is the Faraday constant
Cu50 + 2W + 2e =
F002C = +0.4615
+ —ln
2F
2Cu + 1120
[8a]
[11+12
[8b]
[1120]
CuSO4 + it + 2e = Cu + HSO;
FT
F1.50
=+a4818+-.—-ln
.0
2F
[it]
[9a]
V
[9b]
[1150;]
eq of 8—12 show that any resulting increases in [111 on the
potential for oxidation of Cu are likely to be small because
of the logarithmic term, and any increase in potential will
be a little more for oxide formation than for hydrated sulfate formation. Consequently, the first oxidation product
film should still be a hydrated sulfate and not Cu20.
Anoclic polarization behavior—All anodic curves exhibited an initial region of active behavior where i, increased
with increasing F. Further increases in F produced either
a limiting current plateau (iL), or an active—passive transition. Limiting L behavior was the dominant characteristic
in 1 M 112504, which is close to the acid concentrations
used in electrorefining, and well-developed passivity with
the lowest i, occurred in the 10 M solution. These effects
are shown in Fig. 1 for the as-measured (uncorrected) data
obtained with the SCE, where the initial E corresponds to
the open-circuit potential (F00). The behaviors were similar
to those reported by Leckie1° with an SCE. For example,
the F00 range (—0.05 to 0 VECE), the potential region of 0ffi
at the active-passive transitions (+0.1 to +0.15 VEcE) and
the magnitude of L in the 1 M solution (_4Ø3 Aim2) were in
reasonable agreement with those of Leckie. The magnitude
of i. in the 10 M solution (—20 Aim2) was comparable to but
higher than that of —2 Aim2 reported by Leckie.
Polarization data obtained with the MSE electrode for
duplicate tests in 1 and 10 M 112S04 are shown in their
uncorrected form in Fig. 2. The general effect of solution
concentration on the transition from ZL to i, behavior was
the same as that observed in Fig. 1.
CuS04H2O + H + 2e = Cu + HSO; + 1120 [ba]
Ecusono = +0.3978
+ FT
—ln
[it]
V5
[Hso;][1120} ,
[lOb]
CuS043H20 + 11 + 2e = Cu + 1150; + 31120 [ha]
Ecuso 3F10 = +0.3551
+ FT
— ln
[it]
Ui
0
V561 [lib]
[HSo;]{1120J2
CuSO45H20 + it + 2e = Cu + HSO; + 5H0 [12a]
F005051111 = +0.3207
+FT-ln
[it]
,
V5[12b]
[HSo;][H2o]
Equilibrium potentials for the test solutions, based on
eq 8—12 and activities in Table I, are listed in Table III.
Log(), (A/rn2)
Fig. 1. Uncorrected anodic polarization behavior of Cu in 1, 5, 7,
and 10 M H2S04. E = 0.1 mV/s, SCE electrode.
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604
J. Electrochem. Soc., Vol. 145, No.2, February 1998 The Electrochernical Society, Inc.
O.5
LU
I
LU
0.4c
m
0.3
Lu
0
1
Log(i), (Aim2)
Log(), (A/rn2)
Fig. 2. Uncorrected anodic polarization behavior of Cu in 1 and
10 M H2504. E = 0.1 mV/s, MSE reference electrode.
More thorough analyses of the polarization data required
and ohmic drop (iaTh) corrections to be applied to the
as-measured E. The effective value of R0 for each test was
obtained from a plot of a against E in the active region prior
to the onset of L behavior or passivity, as described previously.9 At the higher a' each plot approached linearity (i.e.,
a constant gradient with i oc E), indicative of ohmic-con-
trolled behavior Under these conditions, the gradient is
dominated by the ohmic resistance and aE/ai, R,1. In this
manner, R0 values were obtained for each test and most
were in the range 2 X i0 to 4 X 10 0 m2, depending on
the solution concentration and the placement of the Luggin
probe. The resulting iaR corrections were applied to E and
values were added to obtain the fully
the estimated
corrected potentials. The corrected polarization curves cor-
responding to Fig. 1 and 2 are shown in Fig. 3 and 4,
respectively, where the potentials are referenced with
respect to the SHE on the second ordinate.
The corrected curves indicated an overall trend for the
highest E0 to be associated with the most concentrated
solutions. However, Eec values were controlled primarily
by traces of dissolved oxygen in the solution that were not
removed by purging, causing E0 to vary between duplicate tests (e.g., see the 1 M tests in Fig. 4).
Active region—Corrected polarization curves exhibited
well developed Tafel behavior in the active region with E
log a' Figure 5 shows the experimental Tafel slopes, b, =
OE/d log a, for different test conditions. Clearly, ba de-
creased gradually from —41 to —31 mV as H2504 increased
from 1—10 M. Such concentration influences have not been
reported previously.
Under high field conditions, a is given by eq 132223
'
b—
2.O3RT
Fig. 4. E0 + i0R0 corrected anodic polarization behavior of Cu in
1 and 10 M H2S04. E = 0.1 mV/s, MSE reference electrode. (0)
potentials in Table Ill.
Corresponding
where cia is the anodic transfer coefficient.
The theory of multistep electrode reactions developed by
Bockris et al.22'23 shows that cia has a value of 1-3 when the
anodic rate-determining step (rds) is charge transfer of the
first electron (Cu -. Cu +
e) and 2-p when the rds is
charge transfer of the second electron (Cut — Cu2 + e),
where p is the symmetry factor measured from the outer
Helmholtz plane (01fF). When the rds is the transport of
Cu2 from the UHF into the bulk solution (i.e., no charge
transfer), a, is 2. Applying the usual assumption that 3 is
Ø52223 at 25 °C (298 K), these ci, values lead to predicted
b, slopes of 118, 39.4, and 29.5 mV when the rds is the first
charge transfer; the second charge transfer, and the transport of Cu2, respectively. (N.B. When the rds is a chemical
step preceding anodic electron transfer, such as mass
transport of solvating water molecules to the surface, cia is
zero22'23 and b, becomes infinite).
Examination of Fig. 5 suggests that the first charge
transfer was not the rds under any test condition, because
all experimental a values were considerably less than 118
mV. Transfer of the second charge was the likely rds in
solution concentrations near 1 M H2S04, because the
experimental b were reasonably close to the predicted
39.4 my, consistent with the reported behavior of Cu in
acidic sulfate9 and sulfuric acid solutions.'6'24'25
The decrease in b, to —31 mV as the H2504 concentration
increased to 10 M indicated one of two possibilities: (i) the
second charge transfer remained the rds, but there was a
continual decrease in p with increasing concentration due
to the effect of the decreasing 1H301 on the structure of the
electrical double layer and the hydration of the activated
complex (e.g., a p of 0.25 predicts a ba of 33.4 mV when ci,
is 2-p), or (ii) the electrode reaction moved toward an rds
involving mass transport of Cu2 from the OHP (i.e., cia is 2).
[131
ci,F
>
0
Lii
C
E
0
lb
LU
2
-0.1
.3
-2
.1
0
1
8
4 5 6 7
H2S04 Concentration (M)
3
Log(19), (Aim2)
Fig. 3. E + i0R correctqd anodic polarization behavior of Cu in
1, 517, and 10 M H2504. E= 0.1 mV/s1 SCE reference electrode.
(0) Corresponding E,20 potentials in Table Ill.
Fig. 5. Effect of H,S04 concentration on E1, + i0R corrected
anodic Tofel slopes (b0) of Cu. Data include tests with SCE and MSE
reference electrodes and Eequal to 0.1 and 1 mV/s.
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605
J. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
The possible influence of transport effects on ha in 10 M
H2504 was examined by polarizing Cu at 0.1 mV/s under
rapid stirring conditions and switching the stirrer off and
on again while in the Tafel region. Test results, using the
SCE, are shown in Fig. 6 for uncorrected and fully corrected data. Switching the stirrer off caused a small de-
crease in a, followed by a positive displacement of the
Tafel region by —10 mV. Thus, a higher overvoltage (fla)
was required to produce the same a as that which would
have been obtained if the initial stirring conditions had
been maintained. Switching the stirrer on again caused a
displacement toward higher i, and the subsequent Tafel
region moved toward a position coincident with that anticipated if stirring had been maintained throughout the
test. All these displacements were consistent with changes
in the liquid diffusion layer thickness (6) at the OHP,
caused by changes in stirring rate (fluid flow), and the
resulting effects upon transport of Cu2 away from the
electrode surface. Such effects suggested that Tafel behavior was controlled by transport processes to some degree,
if not as the single rds then at least in mixed rate control
with the second charge-transfer step.
The relative viscosity of H2504 increases by a factor of 5
as the concentration increases from ito 10 M,'2 consistent
with the anticipated effects of the decreasing [H20}. Such
effects increase the possibility of mass transport (diffusion) controlled processes becoming important at higher
acid concentrations. It is known that transport effects may
control the b value of other metals during oxidation to the
+2 state, as shown by recent studies on the oxidation of Cd
to CdClt in acidic chloride solutions.26
Limiting current behavior.—The initiation of L behavior in 1 M H2S04 occurred at potentials below E20 and
nearer to ECUSO4.5H20 (+0.323 V5 in Table III), as shown by
positive deviations from Tafel behavior in Fig. 3 and 4.
This indicated that the onset of L behavior was associated
with nucleation and growth of the pentahydrate film.
metal sulfates,'9 causing the saturation concentration to be
>0.254 M. In fact, saturation is likely to be about 1 M,
because industrial electrorefining solutions typically contain approximately 0.65 M dissolved Cu2 in 1.7 M H2S04.27
In the presence of a large concentration of supporting
electrolyte (e.g., H2S04), electrical migration effects may be
neglected and L is given by the well-known relationship9
= zFD(C, — C,j/6,
film/electrolyte interface according to eq 14
Cu504-5H20 + H' = Cu2t + HSO; + 5H20,
k = {[Cu2"][HSO;][H2O]5}/[H1 = 0.2166 [14]
where the equilibrium constant was calculated from thermodynamic data in Table II.
Based on eq 14 and the activities in Table I, the saturation activity of the cupric ion [Cu2t]5 in 1 M H2S04 is 0.264,
corresponding to a molal concentration of O.264/-y
(approximately equivalent to a molar concentration of
0.254/-)') where y is the molal activity coefficient. The value
of -y is uncertain in the acid solution but is known to be
less than unity for concentrated solutions of divalent
i0
Cu2t),29 C5 is the saturation concentration of Cu2 (mol/m3),
CB is the concentration of Cu2 (mol/m3) in the bulk solution, and 6 is the diffusion layer thickness. Substituting for
the estimated C 1 x io mol/m3 (equivalent to 1 M),
C5 0, and 6 5 >< iO m under gently stirred conditions9
gives an estimated L of approximately 1900 A/m2. Comparisons with the limiting current regions in the 1 M solutions in Fig. 3 and 4 show that at potentials just prior to
Ecu2o (+0.463 V5,), where the pentahydrate is the only
solid oxidation product, i, is 2260 A/m2 for the SCE tests
(Fig. 3), and 2300 and 2380 A/m2 for the two MSE tests
(Fig. 4). These values are in good agreement with the estimated L based on eq 15. Consequently, the onset and magnitude of limiting current behavior is determined by formation of the pentahydrate film. Furthermore, it may be
concluded from eq 15 that any factors increasing the effective value of 6, such as formation of anode slime layers
during electrorefining, decrease the value of a at which
industrial electrodes begin to experience limiting current
phenomena, in general agreement with the conclusions of
Abe et al.5
Once the applied E exceeds
it is possible to form
Cu20 at the surface of the electrode, which then undergoes
anodic oxidation to the soluble Cu2t species
UncenectedloM
[16]
Thus, the final oxidation product, Cu2", is the same as that
produced by formation and dissolution of the pentahydrate, with the maximum value of L remaining under the
control of eq 15.
The Cu20 may become the dominant film on the surface
when E> E20, if its film formation kinetics are faster than
the formation kinetics of CuSO45H2O. This appeared to be
the case, because EDX analyses of corrosion products that
were present at the termination of the 1 M H2S04 polarization tests in Fig. 4 failed to detect any sulfur but did detect
oxygen. The corrosion product was in the form of fine nodules, approximately 0.5 p.m diam. The EDX spectra of cor-
rosion products scraped from the Cu surface contained a
large Cu ICe and a small 0 Ka peak that were very similar
to those observed in spectra obtained from Cu2O powder.
The formation of Cu20 at higher E probably accounts for
confusion in the literature as to whether the origin of passivation phenomena during industrial electrorefining is relatformation of Cu oxides and/or sulfates.2'4-8
At electrode potentials where both CuSO45H2O and
Cu20 may occur, the faster kinetics of formation of Cu20
Conected 1DM
0.15
Cu20 + 2H -* 2Cu2t + H2O + 2e'
ed to
0.45
0.2
[15]
where z is the number of electrons transferred (2) in the
stoichiometric oxidation of Cu to Cu2, D is the aqueous
diffusion coefficient of Cu2 (0.5 X
m2 s above 0.1 M
Under these circumstances, the highest z1 is determined by
the rate at which the film dissolves by transport-controlled reactions in the liquid phase when the saturation
condition (solubility limit) for Cu2 is reached at the
A/m2
0.4
relative to the pentahydrate may be attributed to the
0.35
w 0.1
P1
0.3
smaller number of reacting species in the stoichiometric
equation for formation of Cu2O, particularly H20 molecules (cf. eq 8a and 12a) and the smaller molar volume 1TM
of Cu in the oxide relative to the pentahydrate. Table IV
0.25
0.0:
0.2
-0.05
-0.1
-3
-2
- 0.15
0
1
2
3
4
Log(), (Nm2)
lists VM for several compounds, as calculated from eq 17
VM = (VNA)/nCU
[171
where NA is Avogadro's number, v is the volume of the unit
cell, and n, is the number of Cu atoms per unit cell. The v
and n were obtained from crystallographic data.29 The
ratio VM/(VM)CU is the relative molar volume increase in Cu
Fig. 6. Effect of stirring on changes in the uncorrected and EUP +
iR corrected anodic polarization behavior of Cu in 10 M H2504.
E=O.1 mV/s.
accompanying oxidation, being 15.35 for the pentahydrate
and only 1.647 for Cu20. The VM data is used later in calculations relating to thickness of films.
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606
J. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
Table IV. Molar volume (VM} of Cu in different solid phases.°
Phase
Ratio (VM/(VM)cU)
TIM, (m1)
Cu
7.1096 )< 10-6
Cu20
CuO
CuSO4'5H20
CuSO4'3H20
Cu304'H20
10
10.913 )< 10
7.9845 >< 1Q
5.2955 >< 10
4.0871 )< 10
2.4029 >< 10
1.1713 x
1.2216 x
CuSO4
CuC1
1
1.647
1.718
15.35
11.23
7.448
5.749
3.380
1Q
The EDX analyses of surfaces passivated in 10 M H2S04
during the MSE tests in Fig. 4 produced very strong S Ka
and Cu Kcs peaks, together with a well-defined 0 Kci peak,
that were strongly supportive of a CuSO4xH2O film being
responsible for passivation. Other indirect evidence for
sulfate film formation in 10 M 113504 was the observation
that the passivated surfaces were poor catalysts for the
oxygen evolution reaction. Electrochemical oxidation of
water occurs via eq 19a at potentials above E01 in eq 19b
2H20 —>02 + 4W + 4e
a Calculated from JCPDS-ICDD.2>
= +1.229
Passive behavior.—The increased tendency to exhibit
active—passive transitions with increasing concentration
of 112S04 correlated with the declining activity of [1120]
and its effect on the degree of hydration of Cu504. Figures
3 and 4 show that at the active—passive nose (i6,J, E was
always well below Ecu60 and closer to the potentials listed
in Table III for formation of a hydrated sulfate. The i
decreased with increasing acid concentration and was
lowest in 10 M 112504 where CuS04H20 was the most stable hydrate. For example, at potentials just below Ecuao in
the passive region, i, values were 225, 140, and 20 A/rn2 in
the 5, 7, and 10 M solutions, respectively, in Fig. 3 (SCE
tests). Similarly, i was 50—80 A/m2 in the 10 M solution in
Fig. 4 (MSE tests).
Additional passivation characteristics in 10 M 112S04 are
detailed in Table V for different testing conditions. In particular, the potential (E5)6> at which the minimum passive
occurred was always below Ecaao.
current density
These details, and the general positions of Ecu2o on the
curves in Fig. 3 and 4, indicated that active/passive behavior was always associated with hydrated sulfate films
______
+ RT
—in
'
2
4.?
[19a]
[19h]
VSHE
[1120]
where eq 19b is based on thermodynamic data in Table H.
Assuming unit fugacity (1 atm pressure) for [021 and values of [H] and [1120] listed in Table I, E0 is 1.324 V in
io M 113304. Polarization tests at potentials above 1.324 V5>
failed to reveal any evidence of anodic currents (increasing
i>) corresponding to oxygen evolution, as confirmed by the
two fully corrected tests at 0.1 and 1.0 mV/s in Fig. 7. The
i, remained at approximately 20 A/rn2 up to test potentials
of + 1.580 V5 at 0.1 mV/s, and +3.278 V5 at 1 rnV/s, corresponding to oxygen overvoltages of 256 rnV and 1,954 V,
respectively. The absence of increased currents at oxygen
overvoitages close to 2 V was convincing evidence that the
passive surface was an inefficient electrode for chargetransfer reactions associated with oxygen formation. This
was consistent with a filmed surface having poor electrical properties (e.g., CuS0>xH20). The passive film was
unlikely to be a copper oxide because Cu20 and CuO are
close similarity between the pentahyrate and trihydrate
semiconductors32 and oxygen evolution currents are readiiy detected on oxide-passivated surfaces of Cu in other
environments, such as weakly alkaline solutions.31
Overall, it may be concluded with fair conviction that
passivation of copper in strong 112S04 solutions is associated with formation of hydrated sulfate films, leading to
rnonohydrate formation, CuS04H20, in the 10 M solution.
potentials, Ecuso4.iH2o and Ecasoa.2H6O, in the 5 and 7 M solu-
Films in this solution are stable to high potentials and
whose tendency to promote passivation (lower i) in-
creased with decreasing degree of hydration, which in turn
was controlled by the activity of [1120] in the acid. The
tions in Table 111 suggests a gradual transition in the average hydration (x) of the sulfate films, CuS04xH20, from 5
through 3 to 1 as the acid concentration increased over the
range 1—10 TtiI.
If the total charge passed in the passivation peak (QT)
during the transition from active to passive behavior contributed solely to film formation, the thickness (t) of the
film should be
t = (QV6)/(21i)
[18]
resist dissolution. The prime reason for the differences in
dissolution characteristics between the sulfate films as the
acid concentration increases, and the related transition
from limiting current to true active/passive behavior, is
believed to be directly associated with changes in the
[1120] activity. A more detailed explanation is proposed in
the Appendix.
Chloride-contaminated solutions.—The potential for the
oxidation of Cu to CuCi at 298 K (25 >C) is given by eq 2014
Values of QT were obtained from the peak area (integral) of
Ecaci = + 0.117 — .:!2i ln[Cl],
current-time plots of the polarization data and are listed
[20]
in Table V. After substituting for VM and Q.6 from Tables IV
and \/ respectively, eq 18 shows that passive films composed of Cu504H20 in 10 M 112S04 should vary in thick-
ness between 5.7 and 48.6 sm, depending on the testing
conditions. However, the similar magnitudes of (1p)m,n in
3
Table V suggest that the thickness variations are not as
large as calculated and that the higher QTvalues probably
include some oxidation of Cu to soluble Cu2 during the
filming process.
3
0.1 mV/s laM
2.5
lmV/s 1CM
25
I
U)
2
uj-1.5
Table V. Passivation of Cu in 10 M H2S04. Total charge passed (01)
during active-to-passive fransition, minimum passive current (iii)mj.,
0.5
0.5
0
0.1
0.1
0.1
0.1
1.0
1.0
MSE
MSE
5CR
5CR
SCE
5CR
1.77 X 10
1.18 >< 10
7.45 )< io
5.69 X l0
3.05 >< 10
2.06 X i04
(ip),j> (A/rn2)
20
18
9
8
23
18.5
I
m
and corresponding potential (Ep)mj.
E (rnV/s) Ref. electrode QT (C/rn1)
IT'
(E>), (V5)
+0.439
+0.515
1-0.434
+0.404
+0.506
+ 0.486
.4
4.
-1
0
1
2
.1. 0
3
4
Log(i), (Aim2)
Fig. 7. E0 iR0 corrected anodic polarization behavior of Cu in
10 M H2504, E = 0.1 and I mV/s, showing no oxygen evolution
currents above 1.324 V1.
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607
J. Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
Assuming the CL concentration in the SCE tests was close
to the estimated maximum value of 3 X 1O M due to leakage from the reference electrode and approximating this
concentration to the [Cl-], then
was close to
+ 0.3 84 VSHE. Therefore, comparison with the corrected
potential ranges of the SCE tests in Fig. 3 suggests that
traces of CuC1 could have formed during the anodic polar-
ization tests. Consistent with this, EDX analyses of the
specimens tested in Fig. 3 showed trace peaks corresponding to Cl Kct.
al solutions. These potentials are comparable to or lower
than any of the ECUSO4.XH2O potentials listed in Table III,
indicating that CuCl could form on the electrode before a
hydrated sulfate film, particularly if inadequate process
control allowed higher CL contamination. Under the galvanostatic operating conditions in electrorefining practice, formation of CuC1 could trigger the onset of limiting
behavior by decreasing the effective electrode area and
causing the applied potential to rise into the regime where
films of CuS045H2O form. These proposed effects provide
The presence of chloride was also confirmed by XPS
studies on Cu that had been potentiostatically polarized in
a thermodynamic basis for the observations of Jin and
Ghali,34 who reported that additions of CL to simulated
10 M H2S04 in the passive region at +0.4 VSCE (equivalent
to a corrected E of + 0.681 V5) for 5400s. During this time,
i, remained close to 23 A/rn2. The results of the XPS analy-
electrorefining solutions (1.66 M H2504 at 65 °C) promoted
by Lindberg et al.33
Figure 8 shows that 5, 0, and Cl were present on the surface, consistent with a surface passivated by a sulfate film
and containing traces of CuCl. The vertical line at 198.4 eV
Thermodynamic analyses and anodic polarization stud-
sis are shown in Fig. 8, where the positions of the peaks
were identified with reference to published tables compiled by Wagner32 and 0 and S peaks in sulfates reported
corresponds to the position of the Cl 2p peak in inorganic
chlorides32 and the vertical line at 169.1 eV corresponds
closely to the S 2p peak in inorganic sulfates.32'33 The doublet 0 is peak is attributed primarily to oxygen in a hydrated sulfate film. The three labeled vertical lines passing
through the doublet correspond to reported positions of
the 0 is peak in H2O, inorganic sulfate,33 and Cu20.32
The detection of CuC1 on passivated surfaces where the
CL concentration in the solution was 3 x
M has
i0
important ramifications for industrial electrorefining
operations where contaminant concentrations are usually
in the range 0.03—0.05 g/L [(4.7 to 7.8) >< io M).27 Approx-
imating these concentrations to activities and inserting in
eq 20 shows that CuC1 could form on the electrode surface
at potentials between +0.314 and +0.301 VsF in industri-
the onset of "passivation" of Cu anodes and accelerated
the formation of sulfate and oxide films when the CL con-
centration was 7 X
io- M.
Conclusions
ies on Cu in strong 113504 solutions led to the following
conclusions
1. The decrease in [1130] activity with increasing acid
concentration exerts a major influence on the polarization
behavior of Cu, especially with respect to effects on the
relative stability of anodic corrosion products.
2. Copper exhibits active Tafel behavior with a slope
that changes from '—41 to —31 mV as the acid concentration increases from 1 to 10 M. The higher slope is consistent with charge-transfer controlled anodic dissolution.
The lower slope indicates changes in the symmetry factor
in the electrical double layer and/or diffusion-controlled
transport of Cu2 from the electrode.
3. Transitions from active to limiting current (iL) behavior and active-to-passive current (in) behavior occur with
rising potential and are associated with formation of
hydrated sulfate films, CuS04xH20. The pentahydrate
(x = 5) is favored in acid concentrations near 1 M, leading
to L behavior. The monohydrate (x = 1) is favored in acid
concentrations approaching 10 M, causing passive behavior characterized by i, <<
4. The results are relevant to galvanostatic electroref in-
i.
ing processes for Cu and indicate that the onset of L
behavior in refining solutions (termed "anode passivity"
in industrial practice) is associated with formation of sulfate films and not oxide formation, although oxides may
form later as the electrode potential increases. Furthermore, the results suggest that CuCi formation may occur
prior to sulfate film formation in chloride-contaminated
electrorefining solutions and promote the premature onset
of L behavior
U,
a.
0
U,
C
C,
C
Acknowledgments
Financial support for the work was provided by the
Natural Sciences and Engineering Council of Canada and
TA. received a postgraduate scholarship from the Canadian
Bureau for International Education, Libyan Educational
Program. The experimental assistance of Mary Mager and
Dr. P. Wong during the EDX and XPS analyses, respectively, is gratefully acknowledged.
Manuscript submitted June 30, 1997; revised manuscript
received November 4, 1997.
Binding Energy (eV)
60000
Ols
CUL3MJAQ
Cu
50000
(0
&
40000
U,
30000
C
4,
C
20000
10000
c*20
960
H20
-580
-560
-540
-520
APPENDIX
Dissolution a1 Sulfate Films
-500
Binding Energy (eV)
Fig. 8. XPS spectrum of passivated Cu surface polarized at
+0.4 VE (+0.68 1 VE) in 10 M H2504. Photoelecfron peaks for 0
(1 s}, Cl (2s and 2p), S (2s and 2p), and Cu (3s), together with Cu
LMM Auger peaks.
The significant differences between anodic current densities on filmed surfaces in 1 M H2504 (L behavior) and
those in 10 M H2504 (i behavior) may be explained in
relation to changes in the rate-controlling steps during
chemical dissolution of the film. In the 1 M solution, the
results indicated that film dissolution and L behavior was
determined by the transport of dissolution products from
the electrode surface. At higher acid concentrations, it is
proposed that film dissolution and i, behavior is deter-
mined by the rate (R) at which reactants interact with
the film. The complete chemical reaction for film dissolution is shown below
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608
J Electrochem. Soc., Vol. 145, No. 2, February 1998 The Electrochemical Society, Inc.
CuSO4xH2O + HF + (6-x)H30 -* Cu6H3O2 +
HS0,
<5 AP
—
Equation A-i recognizes that cations are stabilized in
solution by formation of ligands with H20 and that the
number of ligands associated with Cu2 is usually consid-
ered to be 6 (i.e., Cu6H202).35 Hence, it follows from eq A1 that Re,, is expected to be dependent upon the product of
the reactants H and H30 according to eq A-2
[H][H2OV6
REFERENCES
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IA
I -"I
Therefore, changes in the activity of [1130] and the degree
of hydration of the sulfate, x, have significant effects on Re.
that would not be as obvious if hydration of the cation was
neglected. For example, consider 10 M H2S04, where x is
close to 1. Alter substituting the corresponding activities for
the 10 M solution (Table I) into eq A-2, Re. is found to be
proportional to 1.064 >< iO. Similarly, using the relevant
activities for 1 M H2S04, where xis 5, eq A-2 shows that Re.
and Winning of Copper, J. E. Hoffman, R. G. Bautista,
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ba
Cb
C,
LIST OF SYMBOLS
anodic Tafel slope, V
bulk solution concentration, mol/m3
saturation concentration, mol/m3
diffusion coefficient, m2 s
D
E
electrochemical potential, V
E
scan rate, V/s
standard electrode potential at 298 K, V
E°
liquid junction potential, V
ELJP
open circuit potential, V
potential in passive region, V
potential at (ip),,,, V
EDX energy dispersive X-ray system
E
Faraday constant, 9.6485 X i0 equiv'
crit
is
(ip)mjn
k
anodic current, A/m2
active/passive peak current, A/m2
limiting current, A/m2
passive current, A/m2
minimum ii,, A/rn2
equilibrium constant
mercurous sulfate electrode
Avogadro number 6.023 )< 10 moY1
NA
number of Cu atoms/unit cell
ne.
number of electrons in reaction
n
OHP outer Helmholtz plane
total charge in passivation peak, C/m2
QT
gas constant, 8.3144 3 K1 mo11
R
rate at which reactants interact with film
ohmic resistance of the solution, ft m
rds
rate-determining step
SCE saturated calomel electrode
SHE standard hydrogen electrode
SEM scanning electron microscope
T
temperature, K
film thickness, m
t
MSE
x
XPS
V
z
Greek
a
-Y
j0
V
[1
molar volume of Cu in crystal, m3
degree of hydration, moles of H20
X-ray photoelectron spectroscopy
volume of unit cell, m
number of electrons transferred in oxidation
anodic transfer coefficient
symmetry factor
diffusion layer thickness, m
molal activity coefficient
anodic overvoltage, V
standard chemical potential, J/mol
change in standard chemical free energy, 3
stoichiometric coefficient
activity denoted by square brackets
373 (1983)
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21. Standard Potentials in Aqueous Solutions, A. J. Bard,
R. Parsons, and J. Jordan, Editors, IUPAC, Marcel
Dekker Inc., New York (1985).
22. J. O'M. Bockris and A. K. N. Reddy, Modern Electrochemistry, Vol. 2, Chap. 9 and 10, Plenum Press, New
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23. 3. O'M. Bockris, Z. Nagy, and A. Damjanovic, This
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24. E. Mattson and J. O'M. Bockris, Trans. Faraday Soc.,
55, 1586 (1959).
25. W. H. Smyrl, in Comprehensive Treatment of Electrochemistry, Vol. 4, 3. O'M. Bockris, B. E. Conway, E.
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26. D. Tromans, B. Foster, and F. Weinberg, Corros. Sci.,
39, 1291 (1997).
27. J. H. Schloen, in Hydrometallurgy and Electrometallurgy of Copper, W. C. Cooper, D. J. Kemp, G. E.
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28. F. R. McLarnon, R. H. Muller, and C. W. Tobias, Electrochim. Acta, 21, 101 (1976).
29. Powder Diffraction File, JCPDS - International Center
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30. 0. Kubaschewski and B. H. Hopkins, Oxidation of Metals and Alloys, 2nd ed., pp 22-29 and 182-194, Butterworth and Co., Ltd., London (1962).
31. D. Tromans and H-h. Sun, This Journal, 139, 1945 (1992).
32. C. D. Wagner, in Practical Surface Analysis, 2nd ed., D.
Briggs and M. P. Seah, Editors, pp 595-634 and 639642, John Wiley & Sons, Ltd., Chichester.
33. 3. Lindberg, K. Hamrin, G. Johansson, U. Gelius, A.
Fahiman, C. Nordling, and K. Siegbahn, Phys. Scr.,
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34. 5. Jin and E. Ghali, Can. Metall. Q., 31, 259 (1992).
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