Chemical Behavior of Electrochemically Generated Nanostructured

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Chemical Behavior of Electrochemically Generated Nanostructured Silver Surfaces Farkhondeh Fathi,† Meghan Schlitt,† David B. Pedersen,‡ and Heinz-Bernhard Kraatz*,† †

Department of Physical and Environmental Sciences, University of Toronto at Scarborough, 1265 Military Trail, Toronto, Ontario, M1C 1A4 Canada ‡ Defence Research and Development Canada, Suffield, Suffield Medicine Hat, AB T1A 8K6 Canada

bS Supporting Information ABSTRACT: A simple electrochemical method has been proposed for the preparation of silver nanoparticles that can be used for the detection of cyanide. Both the electrochemical behavior and morphology of the Ag nanoparticles have been characterized in the presence of KCN or diethyl cyanophosphonate (DECP) as well as in alkaline media. These were investigated by cyclic voltammetry and scanning electron microscopy (SEM). DECP is a simulant of the chemical warfare agent tabun.

’ INTRODUCTION In recent years, the synthesis of metal nanostructures has attracted much interest in different fields of chemistry because of the nanostructures’ physicochemical properties, which differ significantly from those of macroscopic metal phases.1,2 Nanostructured Ag, like other nanometals, is very different from its bulk counterpart in physicochemical properties.3,4 Nano-Ag is also extremely stable and has an electrical conductivity that is significantly higher than that of other metals. These properties make Ag attractive as an electrode material in electrochemical cells.4 In previous studies, Ag nanostructures have been used for the electrochemical detection of ammonia,5 thiocyanate,6 hydrogen peroxide,7 and cyanide.8,9 Previously, a number of Ag nanostructures with various sizes and shapes were fabricated in aqueous and nonaqueous media by different methods including the chemical reduction of Ag precursors, photochemical methods, electron bombardment, laser ablation, and electrochemical methods. Among these methods, electrochemical strategies typically provide a more economical, easily adaptable, and higher purity of nanostructured Ag surfaces.10,11 An electrochemical method for silver nanoparticle synthesis was first proposed by Reetz and Helbig in 1994. In their work, silver metal sheets were dissolved anodically and metal salts were formed as an intermediate. These were subsequently reduced and deposited as silver nanoparticles on the cathode.12 On the basis of this work, Sanchez et al. reported the dissolution of a metallic Ag anode in a nonprotic solvent in the presence of a r 2011 American Chemical Society

Scheme 1. Chemical Drawings of Tabun (Left) and Diethylcyanophsphonate (DECP, Right)

stabilizer as a method of producing Ag nanoparticles/structures on Pt surfaces.13 Similarly, Shang et al. reported the application of a high negative voltage to an aerated solution of AgNO3 and KNO3 in the presence of a stabilizer. The result was the formation of 20 to 40 nm Ag nanoparticles on indium tin oxide surfaces. An extended application of the potential led to the formation of aggregated Ag nanostructures.14 Kaydarov et al. reported the formation of Ag NPs at a constant potential between two Ag electrodes under stirring. In this approach, the current polarity between the electrodes was changed every 30 to 300 s.15 It is important to notice that no method is ideal and all of those mentioned have drawbacks, such as the need for organic solvents or chemical stabilizers, heating of the solution, or deaeration of the solution. In this context, we also mention the use of roughened metal surfaces for surface-enhanced Raman spectroscopy Received: April 26, 2011 Revised: July 26, 2011 Published: August 25, 2011 12098

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Langmuir (SERS) studies, which makes it critical to develop simple, cheap, and reproducible methods for their preparation.16
18 Here we present a simpler method for the preparation of nanostructured Ag surfaces based on cyclic voltammetry (CV) using a conventional three-electrode setup with a Ag working electrode. The nanostructured Ag surfaces prepared were then explored for the electrochemical detection of cyanide ions and

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diethylcyanophosphonate (DECP), a tabun mimic (Scheme 1) that can dissociate to release cyanide. Our own interest in this topic stems from a study involving gas-phase-generated Ag NPs deposited on indium tin oxide that underwent significant morphological changes in the presence of DECP.19 In that study, it was hypothesized that the dissolution of Ag NPs is driven by a reaction with cyanide. To substantiate this proposed mechanism of cyanide detection, it was necessary to carry out a more detailed investigation that goes beyond previous studies of Ag surfaces in alkaline20
22 and alkaline cyanide solutions9,23
27 and that evaluates changes in the surface topography, on the nanoscale, caused by cyanide ions reacting with Ag2O to form [Ag(CN)2]
under cyclic voltammetry conditions. Here we present our results of a combined electrochemical and structural study that shows (a) that electrochemical cycling leads to the formation of a nanostructured surface, (b) that the grain size of the nanostructured surface is affected by cyanide and DECP concentration, and (c) that the electrochemical properties of this surface are affected significantly by the cyanide and DECP concentrations.

’ EXPERIMENTAL METHODS

Figure 1. Typical cyclic voltammetry scan of the Ag foil in a solution of 8 M KOH at a scan rate of 0.15 V s
1 in the potential range of
0.5 to 0.9 V vs Ag/AgCl. During the anodic sweep, peaks are observed because of the oxidation of Ag to the Ag(I) and Ag(II) oxides. During the cathodic sweep, these oxides are reduced back to Ag(0).

Reagents. Diethylcyanophosphonate (DECP, Aldrich), KOH (Caledon), and KCN (Fischer) were used as received. Deionized water (18.2 MΩ 3 cm resistivity) from a Millipore Milli-Q system was used throughout this work. K[Ag(CN)2] was purchased from Strem Chemicals and used as received. The Ag foil (thickness = 0.28 mm, 99.9% metal basis) and platinum wire were purchased from Alfa Aesar. Electrochemistry. Cyclic voltammetry experiments were performed using a CHI660B electrochemical workstation (CH Instruments Inc.) with a 10 mL homemade Teflon cell having a three-electrode setup.

Figure 2. Results of a combined SEM and EDX study of the effects of electrochemically cycling a Ag surface at a scan rate of 0.15 V s
1 in the potential range of 0.5 to 0.9 V vs Ag/AgCl for 15 CV scans in 8 M KOH. Please note the roughness of the surface. EDX analysis (right) shows a composition of 98.83% Ag and 1.17% C. 12099

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Langmuir All experiments were carried out in different aqueous KOH solutions (0.001
8.0 M). A coiled Pt wire (0.25 mm diameter, Alfa Aesar, Ward Hill, MA, 99.9% metal basis) was used as an auxiliary electrode, and silver foil was used as the working electrode. Ag/AgCl (3 M KCl, CH Instruments, Inc.) served as the reference electrode, which was connected to the electrochemical cell via a homemade agar salt bridge (1 M KNO3). All electrochemical measurements were carried out in a grounded Faraday cage. Surface Characterization. The silver surface morphology was investigated at the Nanofabrication Facility of the University of Western Ontario using scanning electron microscopes (SEMs, Leo 1540XB FIB/SEM and Leo 1530 SEM) that were equipped with energy dispersive X-ray spectroscopy for composition analysis. X-ray diffraction analysis (XRD, Rigaku Rata Flex 300 RC X-ray diffractometer using a Co source) was employed to determine the crystal structure of the deposits on the surface. X-ray photoelectron spectroscopy (XPS, Kratos Axis Ultra spectrometer) was performed at the Surface Science Center of the University of Western Ontario with Al Kα (15 mA, 14 kV) as the photosource for structural analysis. The size of the Ag NPs was measured with ImageJ 1.43 software.

’ RESULTS AND DISCUSSION Cyclic voltammetry (CV) experiments were carried out using Ag foil immersed in an 8 M KOH solution, and a number of CV

Figure 3. XPS of the white nanostructured material showing signals at binding energies of 368.3 and 374.0 eV characteristic of the Ag 3d3/2 and Ag 3d5/2 signals of Ag(0).

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scans were recorded in the potential range of
0.5 to 0.9 V versus Ag/AgCl. Essentially, there are no noticeable changes in the CV curves after extended cycles. A typical CV curve is shown in Figure 1 and compares well with the CVs reported by Burke.20 The electrochemical behavior of Ag in alkaline solution is complex, and the formation of Ag2O was reported earlier to involve at least three steps.20
22 The first step has been assigned to either the formation of monolayer Ag2O or AgOH or the dissolution of material as [Ag(OH)2]
. The second step was attributed to the electrodissolution of [Ag(OH)2]
, the formation of Ag2O, or the sublayer trapping of O atoms on the Ag surface. The third step was attributed to the nucleation and 3-D growth of Ag2O on the base layer.20
22 This was followed by the formation of AgO on top of the Ag2O layer, which is thought to be formed by advancing nucleation and 3-D growth stages or by direct nucleation, 3-D growth, and Ostwald ripening.20
22,28 Because of some loss of material from the surface, as described above, there is a charge imbalance between the charge of the first anodic and second anodic peaks. In addition, some of the Ag may be dissolved as AgO
and AgO+ in the anodic sweep, and the product may be cathodically reduced and redeposited on the surface as Ag metal at the end of every CV cycle with other silver oxide forms.20
22 We have some indication of a dissolution process and have observed the formation of a black particulate in solution. We did not analyze this material further. Accordingly, signals were assigned to the various oxidation events of Ag to Ag2O and at the higher potential to AgO as labeled in Figure 1. On the cathodic sweep, these oxides are reduced stepwise back to Ag(0). After the completion of the CV cycles, the silver foil was no longer silvery and shiny but had a matte white appearance. This material could be scraped off of the Ag foil and appeared as a loose white powder, opposite to what is expected of oxidation because silver oxides are dark materials in the bulk.29 Scanning electron microscopy (SEM), energy-dispersive X-ray spectroscopy (EDX), X-ray photon spectroscopy (XPS), and X-ray powder diffraction (XRD) studies of this material were implemented to determine the structure and morphology of the white material as well as its chemical composition. Figure 2 shows an SEM image of the Ag foil surface after electrochemical cycling in

Figure 4. XRD pattern spectrum of the deposited Ag NPs on the surface of the Ag foil formed by electrochemical cycling. The XRD pattern of the material exhibits 2θ values of 44.5, 51.9, 76.4, and 93.0° corresponding to the (111), (200), (220), and (311) planes, respectively, of face-centered-cubic silver. Inset: JCPDS card file no. 4-783. 12100

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Figure 5. SEM study of the effects of the electrochemical cycling of the silver foil in the potential range of
0.5 to 0.9 V vs Ag/AgCl at a scan rate of 0.150 mV s
1 at a supporting electrolyte concentration of 8 M KOH. After only one cycle (frames a and e), there are well-defined Ag grains and a clear boundary between the nanostructured Ag and the underlying bulk Ag. After five cycles (frames b and f), the grain size and film thickness increase. For cycles 10 (frames c and g) and 15 (frames d and h), the grains are now strongly fused and there is no longer any clear demarcation between the nanostructured surface material and the underlayer.

the potential range of
0.5 to 0.9 V versus Ag/AgCl in 8 M KOH. In comparison to a fresh Ag surface, the image shows a roughened surface with nanoscopic deposits that are readily removed from the surface. EDX analysis showed that the deposit consists of silver only without any significant contributions from elements other than some small carbon and oxygen contaminants, presumably because of CO2 adsorption on the silver surface or hydrocarbon impurities.30 XPS analysis of the white material showed two peaks at 368.3 and 374.0 eV at core binding energies typical of the Ag 3d5/2 and 3d3/2 signals of elemental silver (Figure 3). The literature values for Ag(0) are Ag 3d5/2 = 368.0 ( 0.2 and Ag 3d3/2 = 374.0 ( 0.2 eV.30
32 In comparison, the binding energies for Ag2O (Ag 3d5/2 = 367.3 (

0.2 eV)30
32 and for AgO (Ag 3d5/2 = 367.7 ( 0.2 eV and Ag 3d 3/2 = 373.2 ( 0.2 eV)30,31 are sufficiently different so that the white nanostructured surface material can be identified as a deposit of elemental silver. Further investigation of the silver nanomaterial by XRD shows (Figure 4) that the silver exhibits reflections that can be indexed to a face-centered-cubic structure. The diffraction patterns around 2θ values of 44.5, 51.9, 76.4, and 93.0° correspond to the (111), (200), (220), and (311) planes of silver, respectively.33 Furthermore the higher intensity ratio of (111) over the others indicates the enrichment of the Ag(111) plane in Ag NPs. This can presumably be caused by defects in the form of steps and kinks in the Ag(111) plane.34 12101

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Figure 6. SEM study of the effects of the supporting electrolyte concentration on the surface morphology of the deposited Ag NP film. Shown are the results after 15 CV cycles over a potential range of
0.5 to 0.9 V vs Ag/AgCl at a scan rate of 0.150 mV s
1 at supporting electrolyte concentrations of (a) 8.0, (b) 1.0, (c) 0.1, (d) 0.01, and (e) 0.001 M KOH. The grain size at 1.0 M KOH is clearly smaller compared to that in the 8.0 M KOH solution.

The electrochemical cycling of a silver electrode in the potential range from
0.5 to 0.9 V versus Ag/AgCl in 8 M KOH as a supporting electrolyte clearly results in the formation of a nanostructured metallic silver surface that is based on work by Conway et al., where the roughness factor for this roughened nanostructure surface might be 1.5.35 Two obvious factors potentially influencing the formation of the nanostructured surface and the thickness of the nanostructured layer are the number of electrochemical cycles and the concentration of supporting electrolyte KOH. Figure 5 shows the SEM images acquired as the number of electrochemical cycles was varied. Both the size of the Ag grains on the surface and the thickness of the nanostructured layer deposited on unperturbed bulk silver are affected by electrochemical cycling. After only one scan, a layer of silver is deposited with a grain size of approximately 31 ( 9 nm (Figure 5a) and a uniform thickness of the deposit of about 300 nm. The population of Ag grains or nanoparticles is not particularly monodisperse, but the number of fused particles is small. After five cycles, the particles are clearly fused and the thickness of the deposit remains uniformly distributed across the surface with a thickness of about 630 nm and a clear boundary between the nanostructured Ag overlying the bulk silver phase below. After 10 and 15 cycles, the particles on the surface continue to fuse and there is no clear demarcation between the nanostrucured surface and the bulk silver. From the cross-sectional images showing a clear boundary between

the nanostructured surface and the bulk silver, it is clear that the Ag nanostructures are the product of continuous deposition/ oxidation at growing Ag grains during the anodic and cathodic sweeps. Figure 6 shows a series of SEMs at KOH concentrations of (a) 8.0, (b) 1.0, (c) 0.1, (d) 0.01, and (e) 0.001 M. As seen, lowering the concentration of KOH resulted in a less compact Ag NP film being formed. Particles are spaced out more, and the silver foil underlayer is clearly visible. At a concentration of 0.001 M KOH, no significant accumulation of Ag NPs was observed. These results combined with the cycling data confirm the disruption of Ag surfaces that occurs with electrochemical cycling under basic conditions.20 In terms of the mechanism, the cross sectional data shown in Figure 5 demonstrate that the nanoparticles formed are the result of a redeposition process, as opposed to a direct roughening of the Ag electrode surface. The observed deposition of layers of nanoparticles up to 650 nm in thickness is also indicative of a significant concentration of dissolved Ag. These aqueous ions are likely in the form of hydroxides, (Ag(OH)n)
n+1 because no nanoparticle formation occurs in the absence of significant hydroxide concentrations (e.g., Figure 6e). Decreasing the electrolyte concentration results in a higher ohmic resistance. Thus, under these conditions the surface is not exposed to the same potential range. 12102

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Figure 7. Electrochemical behavior of Ag exposure to (a) KCN (no KCN added,
; 1 pM KCN, ---; and 100 μM KCN, 3 3 3 ) and (b) DECP (no DECP added,
; 1pM DECP, ---; and 100 μM DECP, 3 3 3 ) in 1 M KOH at a scan rate of 0.100 V s
1 in the potential range of
0.5 to 0.9 V vs Ag/AgCl. Please note that for CN
and DECP addition the peak intensities decrease. This is rationalized by the passivation of the silver surface by CN
, decreasing the anodic peaks. The cathodic peaks decrease in intensity because of the dissolution of the material as [Ag(CN)2]
.

Having successfully formed Ag nanostructures, their response to potassium cyanide and diethylcyanophosphonate (DECP) exposure was of interest. A previous study on gas-phase Ag NPs deposited on ITO substrates has demonstrated that the addition of DECP decreases the number of Ag NPs present on indium tin oxide surfaces.19 That study speculated that cyanide anions released upon the hydrolysis of DECP will interact with Ag surfaces or the various silver oxides formed during cycling, causing the formation of dissolved dicyanoagentate [Ag(CN)2]
according to eq 1. AgO þ 2CN
þ H2 O f ½AgðCNÞ2 
þ 2OH


ð1Þ

Alternatively, Ag can be oxidized in the presence of O2 and

cyanide according to 2Ag þ 4CN
þ

1 O2 þ H2 O f 2½AgðCNÞ2 
þ 2OH
2

ð2Þ

As can be seen in Figure 7, the addition of either KCN or DECP to the supporting electrolyte (1 M KOH) had a significant effect on the electrochemical behavior of the system. In a report by Taheri et al., a decrease in the cathodic current was chosen as an analytical signal of cyanide determination. They assumed that the cathodic current decrease results from the formation of a silver cyanide complex that prevents the oxidation of the Ag surface.9 Previous speciation analysis of silver cyanide complexes in an alkaline cyanide solution has shown that the predominant ion is [Ag(CN)2]
at pH >5 and a CN
concentration of less than 12103

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Figure 8. SEM images of the silver surface as a function of KCN and DECP addition. Images were recorded after a total of 15 electrochemical cycles in the range of
0.5 to 0.9 V vs Ag/AgCl at a KOH concentration of 1.0 M and a scan rate of 0.100 V s
1. (a) SEM image of a Ag surface exposed to KOH only and the formation of a nanostructured Ag film. (b) The addition of KCN (1 pM) causes significantly smaller grains. (c) An increased concentration of KCN (100 μM). The grains are smaller and less evenly distributed across the surface. (d) DECP (1 pM) addition affects the grains. (e) At a DECP concentration of 100 μM, the surface is very heterogeneous.

10 mM. [Ag(CN)3]2
is the predominant species for [CN
] > 20 mM, and [Ag(CN)4]3
dominates only for high cyanide concentrations exceeding 2.5 M. The compositions of species such as [Ag(OH)2]
and [Ag(OH)(CN)]
are small in comparison with those of previous silver cyanide complexes mentioned.36,37 Previous mechanistic studies of Ag dissolution in a cyanide solution have shown that at a low concentration of cyanide (less than 0.1 M), Ag dissolution follows first-order kinetics. The formation and decomposition of surface complex [Ag(CN)2]
is involved in the charge-transfer step, which is typically rate-limiting.38,39 Consistent with previous work, the addition of either KCN or DECP resulted in a decrease in the redox peak intensity, as shown in Figure 7. The similar effects of both chemicals suggest analogous chemistry, indicating that CN
is the active species and the hydrolysis of DECP is occurring. In the context of Taheri et al., the decreased CV intensity observed (Figure 7) is expected to result from the formation of a redox-inactive silver cyanide coating that prevents the oxidation of the Ag electrode and effectively passivates the surface.9 CN
does have a higher affinity for Ag surfaces than does the hydroxide anion.40 Alternatively, the formation of significant quantities of aqueous [Ag(CN)2]
could also explain the loss in redox activity.33,34 In support of the latter, the presence of [Ag(CN)2]
was confirmed by taking mass spectra (MS) of samples of the cyanide/KOH solution after exposure to Ag surfaces subjected to CV

experiments (Supporting Information). Interestingly, [Ag(CN)2]
can serve as a source of elemental silver, and at sufficiently high concentrations, elemental silver is deposited (Supporting Information). Under the experimental conditions, however, it appears to behave as a relatively redox-inactive sink for Ag. Clearly, the addition of CN
either by direct addition of KCN or by addition of DECP followed by hydrolysis alters the electrochemical properties of the nanostructured Ag foil surface. The effect of cyanide on the surface topography is evident from the SEM images. Figure 8 shows a series of such SEM images recorded as a function of added KCN and DECP. As seen, exposing the Ag surface to 1 M KOH results in a significant growth of NPs on the surface, but as KCN and DECP are added to the solution, a decrease in the number of deposited Ag NPs on the surface is observed. At higher concentrations of CN
, the freshly formed Ag nanostructures begin to dissolve. This effect increases as the concentrations of KCN and DECP increase in the solution. This observation is consistent with the dissolution of silver as [Ag(CN)2]
, as suggested above.

’ CONCLUSIONS A nanostructured Ag surface has been prepared by the electrochemical cycling of a Ag foil in the potential range of 12104

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0.5 to 0.9 V versus Ag/AgCl in KOH solution. The surface morphology prepared has been shown to depend on both the KOH concentration and the number of electrochemical cycles. Repeated electrochemical cycling was found to increase the grain size on the surface as well as blur the boundary between the bulk and the nanostructured surface. At low hydroxide concentrations, no nanostructures were observed. Ag nanostructures fabricated by the electrochemical cycling of Ag foil under strongly basic conditions were found to be sensitive to the addition of KCN or DECP. This chemistry is similar to that previously observed to occur on gas-phase Ag nanoparticles deposited onto ITO electrodes. In both systems, a loss of redox activity can be used for the sensitive detection of cyanide.

’ ASSOCIATED CONTENT

bS

Supporting Information. Particle size distribution for prepared Ag NPs. Nanoparticle sizes and their standard deviations for different concentrations of KOH. SEM images of Ag NP-decorated ITOs. EDX analysis of the ITO surface. Mass spectra of K[Ag(CN)2] in water and KCN in KOH. SEM images of the silver surface as a function of KCN and DECP addition. Chronoamperometry measurements and SEM images for two steps of Ag oxidation. Typical cyclic voltammetry scan of a Ag foil in KOH. This material is available free of charge via the Internet at http://pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*E-mail: [email protected].

’ ACKNOWLEDGMENT We thank the University of Western Ontario for financial support. In addition, we thank the following individuals for their invaluable help with SEM and EDX (Dr. Todd Simpson and Dr. Tim Goldhawk of the University of Western Ontario Nanofabrication Center), with XPS (Dr. Mark Biesinger, the Surface Science Center at the University of Western Ontario), with XRD (Dr. Kimberly R. Law in Department of Earth Science, the University of Western Ontario), and with mass spectroscopy (MS, Dr. Doug Hairsine). ’ REFERENCES (1) Shao, F. C.; Jian, P. L.; Kun, Q.; Wei, P. X.; Yang, L.; Wei, X. H.; Shu, H. Y. Nano Res. 2010, 3, 244–255. (2) Tapan, K. S.; Andrey, L. R.; Frank, J.; Thomas, A. K.; Jochen, F. Adv. Mater. 2010, 22, 1805–1825. (3) Luo, X.; Morrin, A.; Killard, A. J.; Smyth, M. R. Electroanalysis 2006, 18, 319–326. (4) Campbell, F. C.; Compton, R. G. Anal. Bioanal. Chem. 2010, 396, 241–259. (5) Guo, H.; Tao, S. Sens. Actuators, B 2007, 123, 578–582. (6) Wang, G. F.; Li, M. G.; Gao, Y. C.; Fang, B. Sensors 2004, 4, 147–155. (7) Safavi, A.; Maleki, N.; Farjami, E. Electroanalysis 2009, 21, 1533–1538. (8) Sun, H.; Zhang, Y. Y.; Si, S. H.; Zhu, D.; Fung, Y. Sens. Actuators, B 2005, 108, 925–932. (9) Taheri, A.; Noroozifar, M.; Khorasani-Motlagh, M. J. Electroanal. Chem. 2009, 628, 48–54. (10) Kumar, C. S. S. R., Ed. Metallic Nanomaterials; Wiley-VCH: Weinheim, Germany, 2009; Vol. 1, Chapter 4, pp 149
171.

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(11) Meng, X. K.; Tang, S. C.; Vongehr, S. J. Mater. Sci. Technol. 2010, 26, 475–522. (12) Reetz, M. T.; Helbig, W. J. J. Am. Chem. Soc. 1994, 116, 7401–7406. (13) Sanchez, L. R.; Blanco, M. C.; Quintela, M. A. L. J. Phys. Chem. B 2000, 104, 9683–9688. (14) Shang, L.; Chen, H.; Dong, S. J. Phys. Chem. C 2007, 111, 10780–10784. (15) Khaydarov, R. A.; Khaydarov, R. R.; Gaprova, O.; Estrin, Y.; Scheper, T. J. Nanopart. Res. 2009, 11, 1193–1200. (16) Compion, A.; Kambhampati, P. Chem. Soc. Rev. 1998, 27, 241–250. (17) Lian, H.; Wang, W.; Wu, Y.; Xu, H. Adv. Mater. 2009, 21, 1–5. (18) Brenja, P. R.; Griffiths, P.|R.; Yang, J. Appl. Spectrosc. 2009, 63, 396–400. (19) Marenco, A. J.; Pederson, D. B.; Wang, S.; Petryk, M. W. P.; Kraatz, H. B. Analyst 2009, 134, 2021–2027. (20) Nagle, L. C.; Ahern, A. J.; Burke, L. D. J. Solid State Electrochem. 2002, 6, 320–330. (21) Ukhur, T.; Chung, W. S. J. Electrochem. Soc. 2005, 152, A996–A1000. (22) Eirehim, S. S. A.; Hassan, H. H.; Ibrahim, M. A. M.; Amin, M. A. Monatsh. Chem. 1998, 129, 1103–1117. (23) Li, J.; Wadsworth, M. E. J. Electrochem. Soc. 1993, 140, 1921–1927. (24) Senanayake, G. Hydrometallurgy 2006, 81, 75–85. (25) Bek, R. Y.; Kov, N.A. R. J. Electroanal. Chem. 1998, 447, 109–115. (26) Baltrunas, G.; Pakalniene, E. Chimija (Vilnius) 2002, 13, 64–69. (27) Baltrunas, G.; Popkirov, G. S.; Schindler, R. N. J. Electroanal. Chem. 1997, 435, 95–101. (28) Lutzenkirchen-Hetch, D.; Strehblow, H. H. Surf. Interface Anal. 2009, 41, 820–829. (29) Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth-Heinemann: Oxford, U.K., 1997; Chapter 28, pp 1180
1181. (30) Weaver, J. F.; Hoflund, G. B. J. Phys. Chem. 1994, 98, 8519–8524. (31) Bielmann, M.; Schwaller, P.; Ruffieux, P.; Groning, O.; Schlapbach, L.; Groning, P. Phys. Rev. B 2002, 65, 235431. (32) Tjeng, L. H.; Meinders, M. B. J.; van Elp, J.; Ghijsen, J.; Sawatzky, G. A. Phys. Rev. B 1990, 41, 3190–3199. (33) Jingquan, C.; Suwei, Y. Rare Met. 2005, 24, 376–380. (34) Shaikhutdinov, S. K.; Savinova, E. R.; Scheybal, A.; Doblhofer, K.; Schl€ogl, R. J. Electroanal. Chem. 2001, 500, 208–212. (35) Tremiliosi-Filho, G.; Jerkiewicz, G.; Conway, B. E. Langmuir 1992, 8, 658–667. (36) Li, J.; Wadsworth, M. E. J. Electrochem. Soc. 1993, 140, 1921–1927. (37) Senanayake, G. Hydrometallurgy 2006, 81, 75–85. (38) Bek, R. Y.; Kov, N.A. R. J. Electroanal. Chem. 1998, 447, 109–115. (39) Baltrunas, G.; Pakalniene, E. Chimija (Vilnius) 2002, 13, 64–69. (40) Baltrunas, G.; Popkirov, G. S.; Schindler, R. N. J. Electroanal. Chem. 1997, 435, 95–101.

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