Electrochemical quartz crystal microgravimetry study of metal deposition from EDTA complexes

Article abstract of DOI:10.1149/1.1393318

The electrochemical deposition of metals from Me^II (ethylenediaminetetraacetic acid, EDTA) complexes was studied by electrochemical quartz crystal microgravimetry. These investigations are important for the development of environmental clean-up processes, such as remediation of metal-contaminated soils by EDTA and subsequent electrochemical recovery of metal, in situ electrokinetic soil remediation, electrochemical gas purification, etc. Deposition of metals frequently encountered in hazardous waste sites, such as Pb, Cu, Cd, Zn, and Ni, was investigated. The potential regions of metal deposition were detected from a distinct increase in electrode mass. We observed reduction of Pb, Cu, Cd, and Zn from both protonated and nonprotonated EDTA complexes. No indications were found for Ni deposition. While diffusion-limited currents of Cu, Cd, and Pb deposition could be achieved, the Zn deposition current was much lower. This can be explained in terms of electrode passivation during Zn deposition. Side reactions were identified, namely, proton reduction, EDTA reduction, and water decomposition.

Full text of DOI:10.1149/1.1393318

Electrochemical Quartz Crystal Microgravimetry Study of
Metal Deposition from EDTA Complexes
Eimutis Juzeliunas, Howard W. Pickering and Konrad G. Weil
J. Electrochem. Soc. 2000, Volume 147, Issue 3, Pages 1088-1095.
doi: 10.1149/1.1393318
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1088
Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
S0013-4651(98)12-051-7 CCC: $7.00 © The Electrochemical Society, Inc.
Electrochemical Quartz Crystal Microgravimetry Study of Metal
Deposition from EDTA Complexes
Eimutis Juzeliunas,a,z Howard W. Pickering,* and Konrad G. Weil**
Department of Materials Science and Engineering, The Pennsylvania State University, University Park,
Pennsylvania 16802, USA
The electrochemical deposition of metals from MeII (ethylenediaminetetraacetic acid, EDTA) complexes was studied by electro-
chemical quartz crystal microgravimetry. These investigations are important for the development of environmental clean-up
processes, such as remediation of metal-contaminated soils by EDTA and subsequent electrochemical recovery of metal, in situ
electrokinetic soil remediation, electrochemical gas purification, etc. Deposition of metals frequently encountered in hazardous
waste sites, such as Pb, Cu, Cd, Zn, and Ni, was investigated. The potential regions of metal deposition were detected from a dis-
tinct increase in electrode mass. We observed reduction of Pb, Cu, Cd, and Zn from both protonated and nonprotonated EDTA com-
plexes. No indications were found for Ni deposition. While diffusion-limited currents of Cu, Cd, and Pb deposition could be
achieved, the Zn deposition current was much lower. This can be explained in terms of electrode passivation during Zn deposition.
Side reactions were identified, namely, proton reduction, EDTA reduction, and water decomposition.
© 2000 The Electrochemical Society. S0013-4651(98)12-051-7. All rights reserved.
Manuscript submitted December 12, 1998; revised manuscript received November 22, 1999.
Ethylenediaminetetraacetic acid (EDTA), a nontoxic chelating
where Me(EDTA) complexes are stable. In contrast, for processes of
electrochemical soil remediation the optimal conditions for metal-
ion reduction from the Me(EDTA) complexes are of great interest.
It is noteworthy that the study of metal-ion reduction from
Me(EDTA) complexes is complicated if conventional dc measure-
ments are used. The electrochemical polarization behavior of
FeII(EDTA) and FeIII(EDTA) was studied over a wide potential
range at different pH values in our previous work.15,16 Three differ-
ent electrode reactions proceed simultaneously with the metal depo-
sition, namely, proton reduction, reduction of EDTA, and water de-
composition. The partial current of iron deposition is small com-
pared to that of hydrogen evolution. Therefore, voltammetric meas-
urements supply only very limited information about the iron ion
reduction. The problem was overcome by the use of electrochemical
quartz crystal microgravimetry (EQCM). It could be shown that
EQCM gives reliable data for mass changes even in the presence of
a rather large hydrogen evolution current.38
In the present study we investigated EDTA complexes of metal
ions which are frequently present in hazardous waste sites such as
Cu, Cd, Pb, Zn, and Ni. The attention was focused on the determi-
nation of the partial current density and the potential region of metal
ion deposition from the MeII(EDTA) complexes. Metal deposition
was detected by an electrochemical quartz microbalance (EQCM).
The balance supplies data on electrode mass changes with high sen-
sitivity and thereby provides for the separation of the partial deposi-
tion current from contributions of side reactions.
agent, forms very stable complexes with most metal ions (Me) and is
widely used in a variety of industrial processes. The chelating prop-
erties of EDTA are also very useful in processes of environmental
protection. Recently, new electrochemical concepts were suggested
for various cleanup applications employing EDTA as complexing
agent for different pollutant species. The advantage of electrochemi-
cal processes is the employment of electrodes for the supply of elec-
trons instead of chemicals, which usually have to be employed in
large amounts in conventional techniques.
FeII(EDTA) is commonly known as a useful agent for the removal
of H2S, 1,2 NO, 3-10 and NO ϩ SO2 11-13 from waste gases. NO con-
version to nontoxic compounds was achieved by use of dithionite as
redox mediator with subsequent electrochemical dithionite regenera-
tion.13,14 It has also been shown that direct electrochemical NO con-
version on electrodes from FeII(EDTA)NO is possible.15,16
Numerous investigations were conducted over the last years to
study the extraction of toxic metals from polluted soils with chelat-
ing agents.17-31 Different metals were removed from contaminated
17-24
19,20,24-31
20,29
20,22,27,29
soils using EDTA: Cd,
Pb,
Co,
Ni,
Cu, 20,24,27,29,31 Fe, 20,29 Zn, 19-22,29,31,32 Mn, 20,29 As, 31 Hg, 31 and
Cr. 25 Subsequent electrochemical reduction of the Me-EDTA com-
plexes seems to be an effective extension of the remediation
process.33,34 The conditions need to be found under which metal
ions can be electrochemically reduced to the environmentally friend-
ly metallic state and EDTA can be recovered.
A promising alternative in the electrochemical soil cleanup is the
in situ electrokinetic soil remediation.35-37 Electrodes are embedded
in the contaminated soil and metals are driven to the electrode by an
electric field. Transport of the contaminant occurs by electromigra-
tion, electro-osmosis, and electrophoresis.35 However, electrochem-
ical water decomposition can cause a pH increase near the collecting
electrodes. Therefore, the solubility of most heavy metals decreases
and their mobility is reduced. Introducing EDTA into the reaction
zone36 can help to overcome these effects. Hence, the possibility to
reduce metal ions from Me(EDTA) complexes at electrodes should
be explored.
Experimental
Experiments were conducted using an EQCN-700 electrochemi-
cal quartz crystal nanobalance coupled with a PS-205 potentiostat
(both Elchema, USA).
Quartz disks, 14 mm diam and 160 ␮m thick, having a funda-
mental frequency of f0 ϭ 10 MHz were used. According to Sauer-
brey’s equation,39 the frequency-to-mass conversion factor is
0.22 Hz cm2/ng when f0 ϭ 10 MHz. The quartz disks were coated
with a 10 nm thick chromium underlayer and a 90 nm thick gold
overlayer. The geometric area of the working electrode was 0.3 cm2.
The specimens were used as received from Elchema. Quartz disks
were glued to the special window of the electrochemical cell.
The experiments were conducted in 0.5 M Na2SO4 supporting
electrolyte. The Me(EDTA) complexes were produced by dissolving
Na2EDTA (Fisher Chemical, 100.9%) and metal sulfate (p.a. grade)
in the supporting electrolyte at an equimolar ratio. The measurements
were taken in solutions deoxygenated with nitrogen gas. The pH cor-
rections were made by using 0.1 N NaOH or 0.1 N H2SO4.
This overview demonstrates that further development of an elec-
trochemical cleanup process requires basic knowledge regarding the
electrochemical behavior of MeII(EDTA) complexes. For electro-
chemical gas purification one wants to know the potential region
** Electrochemical Society Fellow.
** Electrochemical Society Active Member.
a
*
*
Permanent address: Institute of Chemistry, 2600 Vilnius, Lithuania.
E-mail: ejuzel@ktl.mii.lt
z
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Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
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Bruckenstein and Shay44 relate this mass increase to anion adsorp-
tion. Gordon and Johnson,45 however, attribute this mass increase to
increased surface hydration as a consequence of the formation of
Au(OH)a (a << 1). The second mass-gain region in Fig. 1 and the
corresponding limiting current (above ca. 0.8 V) are due to the gold
oxide formation on the surface.44,45 Recently, this conclusion was
confirmed also by simultaneous measurements by EQCM and opti-
cal reflectivity.46 The sharp cathodic peak, which appears at ca.
0.7 V during the reverse potential sweep, is due to the reduction of
the oxide layer developed during the positive potential sweep.
Figure 2 shows a voltammogram and the corresponding mass-
change curve obtained from a solution containing Na2EDTA. Anod-
ic EDTA decomposition is observed above ca. 0.5 V. Numerous pro-
ducts of the anodic EDTA decomposition in Na2SO4 ϩ H2SO4 solu-
tion (pH 0.35-4.0) were identified by Johnson et al.,47 e.g., CO2,
formaldehyde, iminodiacetic acid, ethylenedinitrilotriacetic acid,
ethylenediglycine, glycine, 2-oxo-1-piperazineacetic acid, amino-
ethylglycine, 2-oxopiperazine, and ethylenediamine. Thus, the reac-
tion pathways seem to be rather complicated.
Two distinct cathodic currents are observed during the negative
potential sweep in acid EDTA solution (A and B in Fig. 2). Peak A
appears at the same potential as in the EDTA-free solution (Fig. 1).
Furthermore, peak A is absent in solutions with pH values of 4.6 and
6.9 (Fig. 3). These data imply that current A is due to H3Oϩ reduc-
tion. Current B should be attributed to the cathodic decomposition of
EDTA, which was discussed in detail.16 The distinct current in
pH 4.6 solution indicates easy EDTA decomposition, whereas the
process is retarded in neutral media (pH 6.9, Fig. 3). These data con-
firm the well-known fact that the protonated forms of weak organic
acids can be reduced more easily than the deprotonated forms.48
The mass-change curve in the solution containing EDTA (Fig. 2)
is similar to that in the blank supporting electrolyte (Fig. 1). Thus,
An Ag/AgCl reference electrode was used; the counter electrode
was a platinum wire. The Rotacell electrochemical cell system RTC-
100 and the Faraday cage 702 (both designed by Elchema, USA)
were used.
Results and Discussion
Na2SO4 and EDTA solutions.—Prior to the study of the
Me(EDTA) complexes one has to elucidate the EQCM behavior in the
blank supporting solution as well as in the metal-free EDTA solution.
Figure 1 shows a cyclic voltammogram and the simultaneous mass
change for a gold electrode in the supporting electrolyte adjusted to
pH 3 (this value is typical for the Me(EDTA) solutions studied).
While no remarkable faradaic process is observed during the potential
sweep from E ϭ 0.0 V to E Ϸ Ϫ0.4 V, an increase in electrode mass
is observed in this region. A similar phenomenon was also reported
for platinum electrodes in acid media.38,40,41 One may assume that
hydrogen atoms are adsorbed on the electrode during the negative
potential sweep and water molecules are bonded with the surface via
these hydrogen atoms.40,41 As a consequence, the surface becomes
more hydrophilic, i.e., develops a thicker adhering solution layer. This
explanation is supported by infrared spectroscopy, which shows that
hydrogen atoms are adsorbed at potentials below 0 VAg/AgCl in acid
solutions.42,43
At potentials below E Ϸ Ϫ0.4 V a cathodic current with a peak
at E Ϸ Ϫ0.8 V appears. This current is not observed in neutral sup-
porting solution and therefore may be ascribed to proton reduction.
The mass gain stops when the proton reduction starts. The slight
decrease in mass below ca. Ϫ0.5 V may be due to replacement of
liquid by gas bubbles during hydrogen evolution on the electrode.
Two mass-gain zones are observed above ca. 0.3 V in Fig. 1.
There is no remarkable faradaic process in the first mass-gain zone
(ca. 0.3-0.8 V), which is the so-called double-layer region. A similar
phenomenon can be observed in HClO4 and HNO3 solutions.44-46
Figure 1. EQCM data and cyclic voltammogram recorded simultaneously in
Figure 2. EQCM data and cyclic voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 10 mM Na2EDTA solution adjusted to pH 3. Cathodic
potential sweep 20 mV sϪ1 was started at 0.0 V.
0.5 M Na2SO4 solution adjusted to pH 3. Cathodic potential sweep 20 mV
s
Ϫ1 was started at 0.0 V.
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Figure 4. EQCM data and cyclic voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 10 mM CuII(EDTA) solution (pH 2.9). Cathodic potential
Figure 3. Cyclic voltammograms recorded in 0.5 M Na2SO4 ϩ 10 mM
Na2EDTA solution adjusted to pH 4.6 (upper) and pH 6.9 (lower). Cathodic
potential sweep 20 mV sϪ1 was started at 0.0 V.
sweep 5 mV sϪ1 was started at 0.0 V.
the mass vs. potential change in the EDTA solution may be ex-
plained in the same way as was done for the supporting electrolyte.
In the region E > 1.0 V EDTA oxidation is accompanied by oxide-
layer formation. It may be assumed that the oxide film has an inhibit-
ing effect on the EDTA oxidation, which leads to the hysteresis in
the reverse potential sweep. An analogous conclusion was drawn
studying EDTA oxidation on Pt electrodes.16
CuII(EDTA) solutions.—Figure 4 shows the polarization behav-
ior and the mass-change curve for a solution containing 10 mM
CuII(EDTA) at pH 2.9. When the potential is swept from E ϭ 0.0 V
in negative direction, a cathodic current, i1, appears below ca.
Ϫ0.4 V. A simultaneous increase in the electrode mass is observed,
which indicates copper deposition
CuII(EDTA) ϩ 2eϪ } Cu ϩ EDTA
[1]
The deposition rate reaches a limiting value below Ϫ0.6 V. This can
be seen from the appearance of a limiting current i ϭ Ϫ0.4 mA
cmϪ2 and a constant mass gain dm/dt ϭ 0.11 ␮g sϪ1. According to
Faraday’s law, this mass gain rate corresponds to a current i ϭ
Ϫ0.33 mA cmϪ2, which is in satisfactory agreement with the ob-
served value.
From earlier voltammetric measurements a limiting diffusion
current for iron deposition from 5 mM FeII(EDTA) solution was
found to be i ϭ 0.17 mA cm2, while EQCM data lead to i ϭ 0.2 mA
cm .
Ϫ2 38 These rates are about half those measured for copper reduc-
tion from 10 mM CuII(EDTA) solutions. This indicates that both
ions are reduced under diffusion control and that the diffusivities of
CuII(EDTA) and FeII(EDTA) do not differ very much.
A further current increase is observed in Fig. 4 below ca. Ϫ0.75 V,
which does not affect the mass gain rate. In order to make current
contributions more visible, a higher potential sweep rate was applied
(Fig. 5). Indeed, two current peaks A and B appear in Fig. 5 after the
Figure 5. EQCM data and voltammetric characteristic recorded simultane-
ously in 0.5 M Na2SO4 ϩ 10 mM CuII(EDTA) of pH 2.9. Potential sweep
rate 20 mV sϪ1
.
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Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
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Figure 6. Voltammogram of a copper electrode in 0.5 M Na2SO4 adjusted to
pH 2.9. Potential sweep rate 20 mV sϪ1. The gold electrode was preplated
with copper from a 0.5 M Na2SO4 ϩ 10 mM CuII(EDTA) of pH 2.9 at E ϭ
Ϫ0.7 V for 4 min. Then the solution was replaced by the deaerated solution
under investigation.
copper deposition current (i1). i1 is higher in the first deposition stage.
This may reflect the development of the diffusion layer during the
polarization. It should be mentioned that a limiting current and a
steady-state diffusion profile can only be observed at low scan rates,
at a scan rate of 20 mV/s the voltammogram indicates a t1/2 behavior,
in contrast to what is seen in Fig. 4 at a scan rate of 5 mV/s.
The current A can be attributed to H3Oϩ reduction on deposited
copper. A similar current peak can be seen in Fig. 6 at a copper elec-
trode in EDTA-free solution. Consequently, current B can be attrib-
uted to EDTA reduction, as was already reported by Müller et al.49
They found the following reduction products from CuII(EDTA): for-
maldehyde, formiate, glycoacid, acetate, ammonium, monomethyl-
amine, and ethylenediamine.
Figure 7. EQCM data and cyclic voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 10 mM CuII(EDTA) solution (pH 7.0). Cathodic potential
sweep 20 mV sϪ1 was started at 0.0 V.
Cadmium deposition in the pH 6.6 solution starts at E Ϸ Ϫ1.0 V
(Fig. 9), i.e. at more negative potentials than in the pH 3.28 solution
(ca. Ϫ0.6 V, Fig. 8). Consequently, the process is more favored in
acid medium.
The EQCM curves yield the limiting deposition current i ϭ
Ϫ0.23 mA cmϪ2 in acid solution (Fig. 8) and i ϭ Ϫ0.19 mA cmϪ2
in the neutral one (Fig. 9). These values are very close to the values
established for iron and copper deposition which were concluded to
be diffusion limited currents. Therefore, it may be supposed that
electrochemical deposition of cadmium from CdII(EDTA) solutions,
takes place as a diffusion-limited process as well.
During the reverse potential sweep, anodic cadmium dissolution
starts actually at the same potentials at which the deposition started,
i.e., at E Ϸ Ϫ0.6 V in the acid solution (Fig. 8) and E Ϸ Ϫ1.0 V in
the neutral one (Fig. 9). The mass decrease vanishes at about 0.0 V,
and the mass remains constant over a wide potential region (until
E Ϸ 1.2 V). Hence, some cadmium is still present on the electrode
in form of scarcely soluble compounds. A further stripping process
takes place above 1.2 V, i.e., in the region where anodic EDTA and
water decomposition takes place. During a subsequent reverse po-
tential sweep, a cathodic peak at E Ϸ 0.7 V and the simultaneous
mass decrease indicate the reduction of a gold oxide layer, as dis-
cussed previously (Fig. 1). Even after the complete cycle at E ϭ 0.0
V, a residual mass gain is observed indicating strong bonding be-
tween the nonsoluble compound and the electrode surface (Fig. 8).
PbII(EDTA) solutions.—According to Fig. 10, lead deposition
from a solution of pH 3.12
Figure 7 depicts results obtained in CuII(EDTA) solution at pH 7.
Copper deposition starts below ca. Ϫ0.7 V, as indicated by an increase
in electrode mass and a corresponding cathodic current. Thus, copper
deposition in neutral solution occurs at more negative potentials as
compared to acid solutions (Fig. 4 and 5). However, the diffusion-lim-
ited reduction rate, derived from the steady-state slope dm/dt (Fig. 7),
is of the same magnitude as in more acid solutions (Fig. 4).
The anodic copper dissolution takes place above ca. 0.0 V in both
acid and neutral solutions, as can be seen from Fig. 4 and 7. The dis-
solution is complete since the electrode mass returns to its initial
value at E ϭ 0.0 V.
CdII(EDTA) solutions.—Figures 8 and 9 show simultaneous
EQCM and voltammetric curves in acid and almost neutral
CdII(EDTA) solutions, respectively. The cadmium deposition reaction
CdII(EDTA) ϩ 2eϪ } Cd ϩ EDTA
[2]
is indicated in the acid solution below ca. Ϫ0.6 V by the distinctive
increase in the electrode mass. The current peak at E Ϸ Ϫ0.8 V is
typical for proton reduction (Fig. 1 and 2). The partial current of
Reaction 2 is very low compared to the peak current. So, the mass
gain rate at E Ϸ Ϫ0.8 V yields a partial ion reduction current i ϭ
Ϫ0.072 mA cmϪ2, whereas the total current is much higher. The
small peak at E Ϸ Ϫ1.1 V should be attributed to EDTA reduction,
as it follows from current B in Fig. 2, established in a metal-free
EDTA solution.
PbII(EDTA) ϩ 2eϪ } Pb ϩ EDTA
[3]
takes place below ca. Ϫ0.7 V (Fig. 10). The current peak at E Ϸ
Ϫ0.8 V is due to proton reduction, as discussed earlier. The slope
dm/dt around E ϭ Ϫ0.9 V leads to a current density of Reaction 3 of
i ϭ Ϫ0.29 mA cmϪ2. The total current at this potential is i ϭ
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Ϫ0.45 mA cmϪ2. The difference of these currents equals the rate of
proton reduction. At E > Ϫ1.2 V, a further current contribution
appears in Fig. 10 which can be assigned to EDTA reduction. The lim-
iting deposition current in Fig. 10 is i ϭ Ϫ0.17 mA cmϪ2, which cor-
responds well to the diffusion-limited currents discussed previously.
One can see from Fig. 10 that anodic lead oxidation starts at
about Ϫ0.6 V, indicated by the anodic current A and the simultane-
ous decrease in mass. Lead ions produced by the anodic current may
be stabilized by reaction with EDTA, which is present in the vicini-
ty of the electrode after the cathodic Reaction 3. Two sharp current
peaks B and C are accompanied by distinct increases in electrode
mass. Obviously, scarcely soluble compounds are formed on the
electrode surface associated with these currents. Current B can be
attributed to the reaction
2Pb ϩ SO42Ϫ ϩ H2O } PbOиPbSO4 ϩ 4eϪ ϩ 2Hϩ
[4]
The equilibrium potential for this reaction can be calculated from50
2Ϫ
E ϭ Ϫ0.113 Ϫ 0.0295 pH Ϫ 0.0148 log aSO
[5]
4
Under the conditions of Fig. 10 one obtains E ϭ Ϫ0.426 VAg/AgCl
,
compared to E Ϸ Ϫ0.44 V for the onset of current B. Current C may
be ascribed to the reaction
4Pb ϩ SO42Ϫ ϩ 4H2O } 3PbOиPbSO4иH2O ϩ 8eϪ ϩ 6Hϩ [6]
Its equilibrium potential is50
2Ϫ
E ϭ 0.03 Ϫ 0.044 pH Ϫ 0.0074 log aSO ϭ Ϫ0.330 VAg/AgCl [7]
4
under the conditions of Fig. 10. The onset potential of current C is
E ϭ Ϫ0.35 V
Figure 11 shows EQCM and voltammetric data obtained in
PbII(EDTA) solution at pH 6.34. According to the mass curve, lead
deposition starts below ca. Ϫ1.1 V. Comparing this potential with
Figure 8. EQCM data and cyclic voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 5 mM CdII(EDTA) solution (pH 3.28), potential sweep rate
20 mV sϪ1
.
Figure 9. EQCM data and cyclic voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 5 mM CdII(EDTA) solution (pH 6.6), potential sweep rate
Figure 10. EQCM data and voltammetric characteristic recorded simultane-
ously in 0.5 M Na2SO4 ϩ 5 mM PbII(EDTA) solution (pH 3.12), potential
20 mV sϪ1
.
sweep rate 20 mV sϪ1
.
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Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
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5 mM CuII(EDTA)], cadmium i ϭ 0.23 mA cmϪ2 (Fig. 8), lead i ϭ
38
0.17 mA cmϪ2 (Fig. 10), iron i ϭ 0.2 mA cmϪ2
.
The deposition
reaction is inhibited in the region of water decomposition. Indeed, no
mass change is observed between Ϫ1.6 and Ϫ1.4 V in the reverse
potential sweep. Zinc dissolution only occurs above ca. Ϫ1.1 V, as can
be seen from the pronounced mass decrease and the corresponding
anodic current. Zinc stripping is completed at E Ϸ Ϫ0.4 V.
The data in Fig. 12 show that the partial current of zinc deposition
is only a small fraction of the total current. Therefore, the dominant
cathodic process is H3Oϩ reduction. This is a good example to demon-
strate the usefulness of the EQCM in cases where an ion deposition
current is masked by a much higher hydrogen evolution current.
Figure 13 gives data obtained for a ZnII(EDTA) solution of pH
7.2, i.e., when H3Oϩ ion concentration is negligible. Zinc deposition
in this solution takes place below Ϫ0.8 V, as is evident from the mass
increase and the corresponding cathodic current. The current peak at
ca. Ϫ1.05 V is most likely due to the reduction of EDTA (Fig. 2),
which is present in the vicinity of the electrode after the cathodic
ZnII(EDTA) deposition. The dissolution reaction starts at E Ϸ
Ϫ1.0 V. The accumulation of scarcely soluble compounds on the
surface is evident above Ϫ0.75 V.
The steady-state mass gain in Fig. 13, dm/dt ϭ 0.02 ␮g sϪ1 , cor-
responds to a current density i ϭ 0.06 mA cmϪ2. The limiting cur-
rent obtained from the voltammogram is i ϭ 0.08 mA cmϪ2. Much
higher currents were determined in neutral solutions for other met-
als: i ϭ 0.17 mA cmϪ2 for copper [Fig. 7, when corrected to 5 mM
CuII(EDTA)], i ϭ 0.19 mA cmϪ2 for cadmium (Fig. 9), and i ϭ
0.2 mA cmϪ2 for lead (Fig. 11). This difference implies that under
the conditions studied the zinc electrode is quite passive. A possible
reason for this is considered.
A gold electrode was immersed in 5 mM ZnII(EDTA) of pH 7.4
and polarized to E ϭ Ϫ0.7 V by a potential step (Fig. 14). The
steady-state cathodic current is negligible at this potential. While a
Figure 11. EQCM data and voltammetric characteristic recorded simultane-
ously in 0.5 M Na2SO4 ϩ 5 mM PbII(EDTA) solution (pH 6.34), potential
sweep rate 20 mV sϪ1
.
the analogous one in Fig. 10, we can conclude that lead deposition
is more favored in acid media. The current maximum i Ϸ Ϫ0.4 mA
cmϪ2 at E Ϸ Ϫ1.3 V coincides well with the current derived from
the dm/dt slope. Upon sweeping the potential in positive direction, a
limiting current i ϭ Ϫ0.15 mA cmϪ2 also coincides well with the
value i ϭ 0.13 mA cmϪ2 calculated from dm/dt slope. Thus, there
were no side reactions found along with lead deposition prior to the
beginning of water decomposition.
The anodic lead dissolution in almost neutral solution (Fig. 11)
starts at about Ϫ0.85 V. The anodic current A and the simultaneous
decrease in mass suggest the anodic reversal of Reaction 3. A sub-
sequent mass gain above ca. Ϫ0.34 V occurs when the current
decreases after peak B. This behavior may be explained by lead
oxide formation
Pb ϩ H2O } PbO ϩ 2eϪ ϩ 2Hϩ
the potential of which is50
[8]
E ϭ Ϫ0.356 Ϫ 0.0295 log aSO2Ϫ ϭ Ϫ0.349 V
[9]
4
The nature of currents C and D needs to be identified in further
investigations.
The EQCM data in both acid and neutral solutions indicate that
in order to dissolve the deposited lead the formation of hardly solu-
ble species must be avoided.
ZnII(EDTA) solutions.—Results obtained in acid ZnII(EDTA) solu-
tion of pH 3.0 are shown in Fig. 12. A zinc deposition reaction starts
below ca. Ϫ0.7 V. A constant mass gain rate dm/dt ϭ 0.038 ␮g sϪ1 is
observed in the region between Ϫ1.1 and Ϫ1.5 V, which corresponds
to an ion reduction current i ϭ 0.11 mA cmϪ2. This current is some-
what lower than ion reduction currents observed in acid solutions of
other metals: copper i ϭ 0.17 mA cmϪ2 [Fig. 4, when corrected to
Figure 12. EQCM data and cyclic voltammogram recorded simultaneously
in 0.5 M Na2SO4 ϩ 5 mM ZnII(EDTA) solution (pH 3.0), potential sweep
rate 20 mV sϪ1
.
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1094
Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
S0013-4651(98)12-051-7 CCC: $7.00 © The Electrochemical Society, Inc.
faradaic process is absent, an electrode mass gain can be seen from
Fig. 14. This mass gain occurs very slowly at the beginning (dm/dt
ϭ 3.77 ϫ 10Ϫ4 ␮g sϪ1), however, with time it reaches an approxi-
mately ten times higher value. The total mass gain is ⌬m Ϸ 0.38 ␮g
cmϪ2, which remains for a long time unchanged when the potential
is returned to the open-circuit value. (Let us remember that the mass
of a monolayer of a metal in the underpotential region is 0.16 ␮g
cmϪ2 for Cd and 0.14 ␮g cmϪ2 for Cu. 51) Hence, the mass gain is
caused by a few monolayers only. It is likely that a scarcely soluble
compound precipitates onto the surface, which gives rise to the pas-
sivity of the electrode in the region of zinc deposition. Presumably,
Zn(OH)2 could precipitate on the electrode.
52
It follows from the stability constant log K ϭ 16.5 that about
5.6 ϫ 10Ϫ9 mol/L of noncomplexed Zn(II) ions are present in a solu-
tion containing 5 mM ZnII(EDTA). At negative potentials, Zn(II)
ions can adsorb on the Au surface and ZnII(EDTA) can act as a
Zn(II) ion donor. In spite of the very low concentration, the transfer
rate of Zn(II) to the electrode could be high enough due to a very
thin diffusion layer caused by a relatively high donor concentration.
With time, the surface concentration of Zn(II) increases sufficiently
to cause precipitation of Zn(OH)2. Certainly, this proposition needs
further verification by experiments.
NiII(EDTA) solutions.—There is no indication from either cur-
rent measurement or from mass determination that Ni deposition
occurs from acid NiII(EDTA) solution (Fig. 15). The mass change
curve in this case is similar to that in metal-free EDTA solution
(Fig. 2). The mass increase between 0.0 and Ϫ0.8 V is due to proton
adsorption and the increase in surface hydrophilicity as discussed
previously. There is no distinctive mass increase indicating metal
deposition below ca. Ϫ0.8 V. The electrode was polarized down to
E ϭ Ϫ1.8 V and the gold film on the quartz was destroyed by hydro-
gen gas; however, no indications of Ni deposition were detected.
Figure 13. EQCM data and cyclic voltammogram recorded simultaneously
in 0.5 M Na2SO4 ϩ 5 mM ZnII(EDTA) solution (pH 7.2), potential sweep
rate 5 mV sϪ1
.
Figure 15. EQCM data and voltammogram recorded simultaneously in
0.5 M Na2SO4 ϩ 5 mM NiII(EDTA) solution (pH 3), potential sweep rate
Figure 14. Current and electrode mass vs. time for a gold electrode after a
potential step from 0.18 VAg/AgCl to Ϫ0.7 VAg/AgCl in 0.5 M Na2SO4 ϩ 5 mM
ZnII(EDTA) solution (pH 7.4).
5 mV sϪ1
.
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Journal of The Electrochemical Society, 147 (3) 1088-1095 (2000)
1095
S0013-4651(98)12-051-7 CCC: $7.00 © The Electrochemical Society, Inc.
Therefore, future efforts should be focused on searching for a cata-
lyst, which facilitates Ni reduction.
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Acknowledgments
This work was supported by the National Science Foundation,
grant no. 9612303. E.J. thanks the American Chemical Society for a
stipend.
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