Iridium Oxide for the Oxygen Evolution Reaction: Correlation between

Article pubs.acs.org/cm

Iridium Oxide for the Oxygen Evolution Reaction: Correlation between Particle Size, Morphology, and the Surface Hydroxo Layer from Operando XAS Daniel F. Abbott,†,# Dmitry Lebedev,‡,# Kay Waltar,†,⊥ Mauro Povia,† Maarten Nachtegaal,§ Emiliana Fabbri,† Christophe Copéret,*,‡ and Thomas J. Schmidt*,†,∥ †

Electrochemistry Laboratory, Paul Scherrer Institut, CH-5232 Villigen PSI, Switzerland Department of Chemistry and Applied Biosciences, ETH Zürich, Vladimir Prelog Weg. 1-5, CH-8093 Zürich, Switzerland § Paul Scherrer Institut, CH-5232 Villigen PSI, Switzerland ∥ Laboratory of Physical Chemistry, ETH Zürich, CH-8093 Zürich, Switzerland ‡

S Supporting Information *

ABSTRACT: A current challenge faced in water electrolysis is the development of structure−activity relationships for understanding and improving IrOx-based catalysts for the oxygen evolution reaction (OER). We report a simple and scalable modified Adams fusion method for preparing highly OER active, chlorine−free iridium oxide nanoparticles of various size and shape. The applied approach allows for the effects of particle size, morphology, and the nature of the surface species on the OER activity of IrO2 to be investigated. Iridium oxide synthesized at 350 °C from Ir(acac)3, consisting of 1.7 ± 0.4 nm particles with a specific surface area of 150 m2 g−1, shows the highest OER activity (E = 1.499 ± 0.003 V at 10 A gox−1). Operando X-ray absorption spectroscopy (XAS) and X-ray photoelectron spectroscopy (XPS) studies indicate the presence of iridium hydroxo (Ir−OH) surface species, which are strongly linked to the OER activity. Preparation of larger IrO2 particles using higher temperatures results in a change of the particle morphology from spherical to rod-shaped particles. A decrease of the intrinsic OER activity was associated with the predominant termination of the rod-shape particles by highly ordered (110) facets in addition to limited diffusion within mesoporous features.

■

INTRODUCTION

then later be introduced back to the energy grid during lulls in energy generation. The production of hydrogen via water electrolysis is expected to play a key role in scalable energy storage to meet this demand.6−8 Germany, for instance, very recently implemented a 6 MW water electrolysis plant (Energiepark Mainz) capable of storing up to 780 kg (26 MWh) of hydrogen produced from surplus electricity generated by wind farms.9 Polymer electrolyte membrane (PEM) water electrolysis operating in the acidic environment was shown to be a promising technology, which makes it possible to produce pure hydrogen at high pressures (over 150 bar) without the need for significant additional compression. Moreover, the cathodic generation of hydrogen from water at low pH is well-known to be extremely facile on Pt-based catalysts.10−15 The coupled oxygen evolution reaction (OER) occurring at the anode, however, is limited by sluggish kinetics and requires a

The increasing demand for energy combined with the decreasing fossil fuel reserves and the safety concerns regarding nuclear energy point to finding a long-term solution focused on the development of clean, renewable power supplies, such as wind farms, solar power stations, and hydropower plants. Many countries have already taken the initiative to progress toward a sustainable future based on renewable energy technology, such as Germany’s decision to cease the production of nuclear energy by 2022 or Denmark’s goal to be completely supported by renewables by 2050.1,2 China being the world’s largest producer of photovoltaic power and the largest investing country for renewables in 2015 ($102.9 billion excluding hydropower) is planning to achieve 5−10% (without hydropower) or 15−20% (with hydropower) of their total energy consumption from renewables by 2020.3,4 Additionally, the U.S. ($44.1 billion investment for renewables 2015)3 is planning to achieve 30% of all consumed energy coming from renewables by 2025.5 A large concern surrounding this movement, however, is the consequential rise in demand to store excess energy produced from intermittent power sources, which can © 2016 American Chemical Society

Received: June 28, 2016 Revised: August 29, 2016 Published: August 29, 2016 6591

DOI: 10.1021/acs.chemmater.6b02625 Chem. Mater. 2016, 28, 6591−6604

Article

Chemistry of Materials Table 1. Synthesis Conditions and BET and BJH Analyses of the N2 Adsorption/Desorption Data

average pore diameter, nm

pore volume, cm3 g−1 sample IrO2-150 IrO2-110 IrO2-90 IrO2-30

synthetic conditions 350 500 600 500

surface area, m2 g−1 (BET)

°C/30 min °C/1 h °C/1 h °C/1 h + calcination, 600 °C/1 h (air)

149 107 86 28

± ± ± ±

significant overpotential to achieve modest current densities.16 In addition, the harsh acidic environment and high anodic operating potentials limit the choice of stable electrocatalyst materials to those of the noble metal oxides.17 The current state-of-the-art anodes employed in commercial PEM electrolyzers are comprised of mixed Ir−Ru metal oxides on a TiOx support.1,18 The high cost and scarcity of both Ir and Ru metals, however, may limit the widespread adoption of PEM electrolysis. To this end, the development of cost-effective PEM electrolyzers is a key factor in the implementation of a hydrogen-based infrastructure. The reduction of the noble metal loading at the anode and enhancing catalyst stability for OER in PEM electrolyzers remains a challenge. The current approaches typically consider reducing the catalyst particle size,19,20 partial substitution of the noble metals with cheap and abundant elements (by synthesis of mixed or complex oxides),21 or the use of high surface area oxide supports.22,23 Reducing the noble metal content through a reduction in particle size is perhaps one of the most widely implemented approaches. Several authors have reported on the synthesis of IrO2 nanoparticles with particle sizes ranging down to approximately 1.6 nm.19,24−27 This approach effectively increases the number of surface atoms with respect to bulk and thus increases the OER activity per gram of catalyst. A major issue that arises from this approach is that it becomes increasingly difficult to establish structure−activity relationships for the OER due to the complex nature of the surface species, which are difficult to probe experimentally. Moreover, the surface of the catalyst changes under the electrocatalytic conditions, that is, applied potential, and thus the in situ techniques (such as in situ XPS or XAS) are extremely important for establishing structure−activity−stability relationships. Previous reports28−32 have typically aimed at understanding the relationship between activity and a restricted number of parameters, for example, particle size or electronic structure. Here, we aim to understand the influence of particle size, morphology, surface structure, and the nature of surface species on the electrochemical OER activity of nanocrystalline IrO2 prepared by a chloride-free modified Adams method. In doing so, we develop a synthesis that yields a wide range of particle sizes, and we use a broad range of physical characterization techniques including N2 adsorption/desorption measurements, high-resolution transmission electron microscopy (HRTEM), powder X-ray diffraction (XRD), X-ray photoelectron spectroscopy (XPS), and operando X-ray absorption spectroscopy (XAS) to relate the structural and morphological changes to the observed electrochemical OER activity and catalyst stability. Furthermore, the use of a chloride-free synthesis allows us to avoid chloride contamination that may affect the catalytic properties and performance of iridium oxide due to the possibility of catalytic chlorine evolution (CER, V = 1.36 V vs NHE),33 which is competitive with OER at low pH values and

2 3 1 1

BET

BJH

BET

0.09 ± 0.02 0.04 ± 0.02 0.18 ± 0.02 0.18 ± 0.02 0.17 ± 0.05 0.17 ± 0.05 nonporous

BJH

IrO2-110 > IrO2-90 ≈ IrO230. As shown by the slightly higher activity of IrO2-30 as 6597


Article

Chemistry of Materials capacitance. IrO2-150 shows the highest TOF; while the increase in the BET surface area is about 5-fold higher for IrO2150 as compared to IrO2-30, the increase in current density at 1.525 V (see Table 4) increases by about a factor of 11. This highlights the importance of the nature of the surface sites as suggested by XAS and XPS, that is, the presence of a surface hydroxide layer.69 −1

TOF [s ] =

J1.525V [A g −1] Q DL(1.0 − 1.4V)[C g −1]

(1)

The second important observation originating from the TOF evaluation is related to the activities of IrO2-90 and IrO2-110. Both samples show a lower intrinsic activity than IrO2-30, although the effect is more pronounced for IrO2-90. A possible explanation for this phenomenon arises from the domination of the {110} surfaces in IrO2-90 as revealed by the HRTEM (Figure 1b) and XRD (Figure 2) studies discussed above. In fact, previous reports72 have shown that the (110) IrO2 surface displays a lower OER activity than other terminations such as the (100). In this case, it is important to recognize that the high synthesis temperature (600 °C) of IrO2-90 results in the formation of highly ordered surfaces. In comparison, IrO2-30 and IrO2-110 were both synthesized at a lower temperature (500 °C; Table 1) and show a higher TOF than IrO2-90, which can be attributed mainly to the exposure of other surfaces. Another parameter that can also play role in decreasing the intrinsic activity of the IrO2-90 and IrO2-110 is associated with the porosity of the samples. Contrary to IrO2-30, both IrO2-90 and IrO2-110 are mesoporous, which can lead to diffusion related problems, for example, gas bubble formation and entrapment within the pores formed within the catalyst layers. The stability of each sample was evaluated after the polarization of the sample. The stability tests are intended to simulate the start/stop behavior of an electrolyzer by stepping between a potential close to the open circuit value, that is, where no OER current is observed (1.00 V), and a potential where an appreciable OER current is observed, that is, J > 10 A g−1 (1.60 V). Shown in Figure 6 are the normalized current densities over the course of 500 potential step cycles for the IrO2 samples. It is clearly shown that the normalized current density at the end of 500 cycles is about the same for samples IrO2-30, IrO2-90, and IrO2-110 when considering the margin of error. This can be attributed to the fact that these three samples all display a very similar morphology. In all cases, the samples display a high degree of bulk crystallinity and a rod-like morphology as evidenced by the X-ray diffractograms and TEM images. This is also further supported by the fact that the surface charge capacitance scales linearly with surface area but the electrochemical OER activity does not. If one compares the trend between stability (normalized current density after 500 cycles, see Figure 6b) and activity (E at J = 10 A g−1, see Figure 5a) as a function of BET surface area, then it becomes clear that the resemblance is strikingly similar. The stability of the measured samples scales in the same manner as activity with surface area, indicating that the stability is directly related to the OER active area. IrO2-150 shows the lowest stability, losing more than 50% of the current density after 500 cycles. While the sample shows a high OER activity, it comes at a cost in terms of catalyst stability. This further supports claims in the literature that activity and stability are correlated to one another.73−75 Generally speaking, higher electrochemical OER

Figure 6. Stability measurements for the IrO2 samples (a): electrodes were stepped between 1.00 and 1.60 V for 500 cycles; the normalized current density at 500 cycles as a function of surface area (b).

activity is associated with lower electrocatalyst stability and vice versa. Operando XAS measurements were conducted at the Ir−LIII edge to investigate the development of the iridium oxide surface as a function of the applied electrode potential. Figure 7a presents the normalized XANES spectra for sample IrO2-150 recorded during a typical OER polarization measurement. As the electrode potential is gradually stepped into the OER regime, there are two effects that manifest themselves in the XANES spectra: one is that the Ir-edge position gradually shifts toward higher energies, and the other is that the peak magnitude gradually decreases. The first effect can be attributed to the fact that IrO2-150 initially exists in a mixed Ir3+/4+ oxidation state as was inferred from the dry catalyst analysis of the XANES (Figure 3a) and XPS (Figure 4a). The observed shift in edge energy upon anodic polarization now corresponds to the transition of the initial fraction of Ir3+ to higher oxidation states. On the basis of the Pourbaix diagrams76 and reported electrochemical redox transitions in acidic media for IrO2,31,77−80 it is most likely that the initial fraction of Ir3+ is being oxidized to Ir4+, although it is possible that some fraction is being oxidized to Ir5+ at higher potentials (ca. 1.50 V) coinciding with the onset of OER.32,57 This is in good agreement with previous XANES studies of hydrous iridium oxide films undergoing transitions from Ir3+ to higher oxidation states upon anodic polarization.32,81 It is interesting to note that this effect is reversible, such that the edge position shifts to lower energies upon returning to 1.00 V. This indicates that the oxidized fraction of iridium can be reduced back to the initial state during the cathodic polarization. Considering that the presence of Ir3+ in IrOx films is accompanied by the presence of 6598


Article

Chemistry of Materials

Figure 7. Normalized XANES spectra for IrO2-150 (a) and IrO2-30 (b) recorded at different electrode potentials. Insets show an expansion of the absorption peak to highlight the change in position. Fourier transformed Ir EXAFS spectra for IrO2-150 and IrO2-30 at given electrode potentials (c) and the EXAFS-determined Ir−O bond distance shown over the range of applied potentials (d). Within the figure legends, A and C indicate anodic and cathodic polarization, respectively.

surface hydroxide31,32,53,55 and taking into account that hydroxide species were detected via the O1 s XPS, we propose that the shift to higher edge energies represents a transition from a hydroxide covered surface to an oxide covered surface. The observed transition is reversible, such that the newly formed oxide surface is reduced back to a hydroxide layer upon cathodic polarization. This phenomenon, however, cannot be observed for IrO2-30. The XANES spectra corresponding to the polarization of this sample are shown in Figure 7b. In contrast to IrO2-150, the lack of any development in the Ir−LIII edge position indicates that there is no measurable change in the iridium oxidation state. It is important to recognize that this does not indicate that there is no change in the oxidation state of surface iridium atoms in the IrO2-30 sample. The larger particle size of IrO2-30 ultimately results in a larger ratio of bulk iridium species as compared to surface iridium species. Because XAS is a bulk technique, it must also be considered that the measurements performed might lack the required sensitivity to detect changes in oxidation state for processes that are typically restricted to the particle surface. An observed decrease in the LIII edge peak magnitude during electrode polarization can also be observed for both IrO2-150 and IrO2-30 in Figure 7a and b, respectively, which is typically representative of changes occurring within the d-band structure of absorbing iridium atoms. The white line (WL) peak at the LIII-edge results from transitions from 2p to empty localized 5d states (localized with respect to the absorbing atom).32 Interactions between surface iridium atoms and the adsorbing species, that is, oxo species, can therefore have a profound effect

on the peak magnitude as electron density is donated or withdrawn from the iridium 5d orbitals. This effect has been previously recognized on Pt,82,83 in which case the binding of oxide species to the Pt surface with increasing electrode potential (ca. 0.5−1.0 V vs RHE) results in an increase of the Pt−LIII edge peak magnitude as the Pt d-band is modified. Similar changes in WL peak intensity that occur during polarization have also been previously recognized for IrO2.49,53,54 However, rather than observing an increase in the WL intensity for IrO2-150 and IrO2-30 as the potential is raised, we instead observed a large initial decrease in WL peak intensity with the minimum occurring at ca. 1.44 V. This loss in intensity cannot be fully recovered upon returning to the starting potential (1.00 V), indicating that the materials undergo some irreversible changes before reaching a more stable oxide state. The changes in WL peak intensity are more pronounced for IrO2-150 than for IrO2-30, which again can be rationalized in terms of particle size and the relative ratio of surface-to-bulk absorbing atoms. Although the exact nature of this process cannot be determined, it may be associated with an initial change in the iridium d-states within the structure (bulk and surface) as the electrode is first polarized. Depicted in Figure 7c are the corresponding iridium EXAFS spectra for IrO2-150 and IrO2-30 in R-space recorded at different electrode potentials. The EXAFS data were refined using the same structural model applied for the dry catalysts IrO2-30 and IrO2-150 in Figure 3c and d, respectively. Further information regarding the fitting parameters and the calculated results of the refinement can be found in Table S1 and Figures 6599


Article


Figure 8. Normalized XANES spectra recorded at 1.00 V for IrO2-150 (a) and IrO2-30 (b) after 0, 100, and 250 stability cycles. Insets show an expansion of the absorption peak to highlight the change in position. Fourier transformed Ir EXAFS spectra for IrO2-150 and IrO2-30 recorded after 0, 100, and 250 cycles (c) and the EXAFS-determined Ir−O bond distance shown over the course of 250 potential step cycles (d).

potential, which again is consistent with the observed XANES signal. This is further supported by the fitted Ir−O bonding distance as shown in Figure 7d, showing that the average bond distance remains relatively close to the initial value at all potentials investigated. It should be stressed again, however, that due to the larger particle size of IrO2-30, a change in the bonding distance of iridium−oxygen pairs located near the surface cannot be ruled out. The development of the iridium oxide surface in terms of catalyst stability was further investigated by in situ XAS. Similar to the stability measurements performed via the RDE methodology, the potential was stepped between 1.00 and 1.60 V for 250 cycles while stopping at 1.00 V to record the XAS after every 50 cycles. Shown in Figure 8 are the XANES spectra for IrO2-150 (Figure 8a) and IrO2-30 (Figure 8b) at the initial state (0 cycles) and again after 100 and 250 potential step cycles. Figure 8a presents the measured XANES spectra of IrO2-150 corresponding to the electrochemical stability. Interestingly, we observe that over the course of 250 potential step cycles, the position of the Ir−LIII edge now shifts to lower energies. Considering that it was shown that IrO2-150 initially exists in a mixed Ir3+/4+ state, the negative edge shift with increasing cycle number suggests that the fraction of Ir3+ increases with respect to that of Ir4+ as the electrode is repeatedly stepped in and out of the OER regime. This is further supported by the analysis of the corresponding IrO2-150 EXAFS spectra (Figure 8c) after 0, 100, and 250 cycles. The best fit parameters of the Ir EXAFS spectra show that the average Ir−O bond length increases

S5 and S6. The EXAFS refinement for IrO2-150 (see Figure 7c) shows a considerable development of the peak located within the range of 1−2 Å, which mainly represents the 6 nearest neighboring oxygen atoms forming the distorted octahedron. The Ir−O bond length obtained from the EXAFS fitting (see Figure 7d) gradually changes from ca. 2.02 Å at 1.00 V to 2.01 Å at 1.50 V. As noted previously, the average bond length for Ir−O in IrOx containing Ir3+ has been reported to be roughly 2.02 Å, whereas the bond length of Ir4+ in the IrO2 rutile oxide is closer to 1.98 Å. Therefore, the shortening of the Ir−O bonds within the octahedron represents the transition of IrO2150 from a lower oxidation state, that is, a mixed Ir3+/4+ state, to a structure more heavily populated by Ir4+. This further supports the XANES data indicating that the Ir undergoes a transition to a higher oxidation state upon anodic polarization. Furthermore, it can be observed that the shortened bond distance again relaxes to the initial value of ca. 2.02 Å upon returning to 1.00 V during the cathodic polarization. In essence, we propose that the observed fluctuations in the Ir−O bonding distance can be interpreted as a transition from a hydroxide covered Ir surface to a surface covered predominantly by oxide when the electrode is cycled into the OER regime. We then observe the subsequent reduction of the surface back to a hydroxide covered state as the electrode is returned to 1.00 V. This is in further agreement with recent in situ XPS studies that observed a decrease in IrO2 surface hydroxide species as the electrode is cycled into the OER regime.57 In contrast to the EXAFS spectra for IrO2-150, however, there is no significant development in the structure of IrO2-30 with the applied 6600


Article

Chemistry of Materials monotonically from approximately 2.02 to 2.03 Å as the number of cycles is increased (Figures 8d and S7 and Table S2), which signifies the transition of Ir to a lower average oxidation state. Referring back to the RDE-measured electrochemical stability presented in Figure 6a, we observe that IrO2150 degrades quite significantly over the course of 250 cycles. Taking this into consideration, the degradation of the material then appears to coincide with the formation of more Ir3+ sites and therefore an increased coverage of surface hydroxide species. Although the presence of an initial hydroxide layer was found to be related to an enhancement in electrochemical OER activity during the RDE measurements, this benefit seems to be outweighed by the rapid degradation of the material and the unavoidable loss of electrochemically active surface area. This observation is in strong agreement with the recent proposition by Cherevko et al.84 that the dissolution of IrO2 is directly related to the formation of unstable Ir3+ complexes, which are susceptible to rapid oxidation/reduction and subsequent dissolution as a result of chemical and/or electrochemical processes. Similar to the results obtained for the operando XAS polarization measurements, sample IrO2-30 shows very little change in XANES as the electrode is cycled (Figure 8b), which implies that the apparent average oxidation state of IrO2-30 remains unchanged. The drop in WL peak amplitude can also be observed in Figure 8b, which again suggests that the material reaches a relatively stable state and does not undergo any further significant developments. This is also in agreement with the stability measurements performed via RDE (Figure 6a), which show that sample IrO2-30 is moderately stable with respect to that of IrO2-150 over the course of 500 potential step cycles. The fitting of the EXAFS spectra shown in Figure 8c further supports this interpretation of the XANES data. Although an increase in the Ir−O bond distance is observed over the course of 250 cycles, the average bond length still remains relatively close to that expected for Ir4+ in the rutile structure.

of an initial hydroxide layer, which is responsible for a significant enhancement in OER activity. While the operando XAS analysis provided little information regarding the nature of IrO2-30 due to the larger size of the particles, valuable insight was gained regarding the surface of IrO2-150. We observe that IrO2-150 initially exists in a mixed Ir3+/4+ oxidation state at 1.00 V vs RHE (ca. OCP) and transitions to a higher oxidation state upon anodic polarization above 1.00 V. This transition was followed by a shift of the edge energy in the XANES spectra and also by monitoring the Ir−O distances in the first coordination sphere obtained from the fitted EXAFS spectra recorded as a function of the applied potential. The presence of Ir3+ coincides with a higher fraction of surface hydroxide species with respect to IrO2-30, as was detected by XPS. These hydroxo species are likely converted to oxide species as the oxidation state is increased. The use of the surface-sensitive XPS (O 1s region) for IrO2-30 also revealed the presence of the hydroxide surface species (O:OH ratios of 0.51 and 0.77 for IrO2-150 and IrO2-30, respectively), but the initial hydroxide layer appears more pronounced for IrO2-150 due to a high surface-to-bulk ratio. In terms of stability, it was shown that the fraction of Ir3+ sites increases on IrO2-150 as the electrode is repeatedly stepped in and out of the OER potential regime. Any enhancement due to the formation of additional surface hydroxide species, however, seems to be outweighed by the loss in electrochemically active surface area that inevitably occurs as the catalyst degrades. It is unclear if the decreased stability is directly related to the presence of Ir3+ sites, although the relationship between BET surface area and stability suggests that this is a possibility. This work thus shows that small IrO2 nanoparticles with a hydroxide surface layer are particularly efficient OER catalysts. Stabilization of the nanoparticle state could lead to significant improvements, allowing further development of PEM water electrolysis technology.

CONCLUSIONS A series of chlorine-free iridium oxide electrocatalysts were prepared by a modified Adams fusion method. Varying the synthesis conditions allowed for control over the particle size and morphology, which range from small, spherical particles (d = 1.7 nm) with high surface area (150 m2 g−1) to larger, rodshaped particles with well-defined crystal facets. The OER mass-based activity generally increases with decreasing particle size, that is, increasing surface area, although it is observed that the surface charge capacitance scales linearly with BET surface area while the OER activity does not. Anomalies in the intrinsic catalyst activity (TOF) can be rationalized in terms of the exposure of different crystal facets. The lower intrinsic activity of the rod-shaped IrO2-90 was explained by the domination of the {110} particle terminations, which is in line with the previous reports on oriented IrO2 samples highlighting the higher activity of other IrO2 facets (particularly the (100) facet). Moreover, the mesoporosity of samples IrO2-90 and IrO2-110 can lead to the entrapment of gas bubbles within the pores of the material during OER in the catalyst layers, thereby blocking some of the electrochemically active surface area. According to the calculated TOF, samples IrO2-150 and IrO2-30 still show large differences in the OER activity that cannot be accounted for simply by differences in the measured surface area. This is correlated with the presence

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemmater.6b02625. Additional EDX, N2 adsorption/desorption, XPS, and EXAFS characterization data, including Figures S1−S7 and Tables S1 and S2 (PDF)

■

■

ASSOCIATED CONTENT

S Supporting Information *

■

AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. Present Address ⊥

Surface Physics Laboratory, Department of Physics, University of Zürich, Winterthurerstrasse 190, CH-8057 Zürich, Switzerland.

Author Contributions #

D.F.A. and D.L. contributed equally.

Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We would like to acknowledge and thank the Competence Center Energy & Mobility (CCEM-CH, Project Renerg2), the 6601


Article


(20) Lee, Y.; Suntivich, J.; May, K. J.; Perry, E. E.; Shao-Horn, Y. Synthesis and Activities of Rutile IrO2 and RuO2 Nanoparticles for Oxygen Evolution in Acid and Alkaline Solutions. J. Phys. Chem. Lett. 2012, 3, 399−404. (21) Nong, H. N.; Gan, L.; Willinger, E.; Teschner, D.; Strasser, P. IrOx core-shell nanocatalysts for cost- and energy-efficient electrochemical water splitting. Chem. Sci. 2014, 5, 2955−2963. (22) Mazúr, P.; Polonský, J.; Paidar, M.; Bouzek, K. Non-conductive TiO2 as the anode catalyst support for PEM water electrolysis. Int. J. Hydrogen Energy 2012, 37, 12081−12088. (23) Nong, H. N.; Oh, H.-S.; Reier, T.; Willinger, E.; Willinger, M.G.; Petkov, V.; Teschner, D.; Strasser, P. Oxide-Supported IrNiOx Core−Shell Particles as Efficient, Cost-Effective, and Stable Catalysts for Electrochemical Water Splitting. Angew. Chem., Int. Ed. 2015, 54, 2975−2979. (24) Felix, C.; Maiyalagan, T.; Pasupathi, S.; Bladergroen, B.; Linkov, V. Synthesis and Optimisation of IrO2 Electrocatalysts by Adams Fusion Method for Solid Polymer Electrolyte Electrolysers. Micro Nanosyst. 2012, 4, 186−191. (25) Zhao, Y.; Hernandez-Pagan, E. A.; Vargas-Barbosa, N. M.; Dysart, J. L.; Mallouk, T. E. A High Yield Synthesis of Ligand-Free Iridium Oxide Nanoparticles with High Electrocatalytic Activity. J. Phys. Chem. Lett. 2011, 2, 402−406. (26) Nakagawa, T.; Bjorge, N. S.; Murray, R. W. Electrogenerated IrOx Nanoparticles as Dissolved Redox Catalysts for Water Oxidation. J. Am. Chem. Soc. 2009, 131, 15578−15579. (27) Nakagawa, T.; Beasley, C. A.; Murray, R. W. Efficient ElectroOxidation of Water near Its Reversible Potential by a Mesoporous IrOx Nanoparticle Film. J. Phys. Chem. C 2009, 113, 12958−12961. (28) Otogawa, R.; Morimitsu, M.; Matsunaga, M. Effects of microstructure of IrO2-based anodes on electrocatalytic properties. Electrochim. Acta 1998, 44, 1509−1513. (29) Lervik, I. A.; Tsypkin, M.; Owe, L.-E.; Sunde, S. Electronic structure vs. electrocatalytic activity of iridium oxide. J. Electroanal. Chem. 2010, 645, 135−142. (30) Kötz, R.; Neff, H.; Stucki, S. Anodic Iridium Oxide Films: XPSStudies of Oxidation State Changes and. J. Electrochem. Soc. 1984, 131, 72−77. (31) Hüppauff, M.; Lengeler, B. Valency and Structure of Iridium in Anodic Iridium Oxide Films. J. Electrochem. Soc. 1993, 140, 598−602. (32) Minguzzi, A.; Locatelli, C.; Lugaresi, O.; Achilli, E.; Cappelletti, G.; Scavini, M.; Coduri, M.; Masala, P.; Sacchi, B.; Vertova, A.; Ghigna, P.; Rondinini, S. Easy Accommodation of Different Oxidation States in Iridium Oxide Nanoparticles with Different Hydration Degree as Water Oxidation Electrocatalysts. ACS Catal. 2015, 5, 5104−5115. (33) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications; Wiley: NJ, 2001. (34) Man, I. C.; Su, H.-Y.; Calle-Vallejo, F.; Hansen, H. A.; Martínez, J. I.; Inoglu, N. G.; Kitchin, J.; Jaramillo, T. F.; Nørskov, J. K.; Rossmeisl, J. Universality in Oxygen Evolution Electrocatalysis on Oxide Surfaces. ChemCatChem 2011, 3, 1159−1165. (35) Hansen, H. A.; Man, I. C.; Studt, F.; Abild-Pedersen, F.; Bligaard, T.; Rossmeisl, J. Electrochemical chlorine evolution at rutile oxide (110) surfaces. Phys. Chem. Chem. Phys. 2010, 12, 283−290. (36) Adams, R.; Shriner, R. L. Platinum Oxide as a Catalyst in the Reduction of Organic Compounds. III. Preparation and Properties of the Oxide of Platinum Obtained by the Fusion of Chloroplatinic Acid with Sodium Nitrate. J. Am. Chem. Soc. 1923, 45, 2171−2179. (37) Oakton, E.; Lebedev, D.; Fedorov, A.; Krumeich, F.; Tillier, J.; Sereda, O.; Schmidt, T. J.; Coperet, C. A simple one-pot Adams method route to conductive high surface area IrO2-TiO2 materials. New J. Chem. 2016, 40, 1834−1838. (38) Shirley, D. A. High-Resolution X-Ray Photoemission Spectrum of the Valence Bands of Gold. Physical Review B 1972, 5, 4709−4714. (39) Scofield, J. H. Hartree-Slater subshell photoionization crosssections at 1254 and 1487 eV. J. Electron Spectrosc. Relat. Phenom. 1976, 8, 129−137. (40) Muller, O.; Nachtegaal, M.; Just, J.; Lutzenkirchen-Hecht, D.; Frahm, R. Quick-EXAFS setup at the SuperXAS beamline for in situ X-

Commission for Technology & Innovation Switzerland, and the Swiss Competence Center for Energy Research (SCCER) Heat & Electricity Storage for their financial support. In addition, we also thank ScopeM (ETH Zürich) for the use of their electron microscopy facilities. PSI is acknowledged for the provision of beam time at the SuperXAS beamline of the Swiss Light Source SLS.

■

REFERENCES

(1) Carmo, M.; Fritz, D. L.; Mergel, J.; Stolten, D. A comprehensive review on PEM water electrolysis. Int. J. Hydrogen Energy 2013, 38, 4901−4934. (2) Lins, C.; Williamson, L. E.; Leitner, S.; Teske, S. The First Decade: 2004−2014 - 10 Years of Renewable Energy Progress; 2014; p 48. (3) Global Trends in Renewable Energy Investment 2016; Frankfurt School of Finance & Management gGmbH: Frankfurt, Germany, 2016; p 84. (4) Renewable Energy and Energy Ef f iciency in China: Current Status and Prospects for 2020; Worldwatch Institute: Washington, DC, 2010; p 48. (5) Executive Order: Planning for Federal Sustainability in the Next Decade. In The White House: Washington, DC, 2015. (6) Bertuccioli, L.; Chan, A.; Hart, D.; Lehner, F.; Madden, B.; Standen, E. Development of Water Electrolysis in the European Union; Fuel Cells and Hydrogen Joint Undertaking: 2014; p 83. (7) Quadrennial Technology Review 2015; U.S. Department of Energy, 2015; p 489. (8) Commercialisation of Energy Storage in Europe; The Fuel Cell and Hydrogen Joint Undertaking (FCH JU), 2015. (9) Green Light for Green Hydrogen at Energiepark Mainz; Theurer, M., Ed.; Stadtwerke Mainz AG: Mainz, Germany, 2015. (10) Sheng, W.; Gasteiger, H. A.; Shao-Horn, Y. Hydrogen Oxidation and Evolution Reaction Kinetics on Platinum: Acid vs Alkaline Electrolytes. J. Electrochem. Soc. 2010, 157, B1529−B1536. (11) Trasatti, S. Work function, electronegativity, and electrochemical behaviour of metals: III. Electrolytic hydrogen evolution in acid solutions. J. Electroanal. Chem. Interfacial Electrochem. 1972, 39, 163−184. (12) Norskov, J. K.; Bligaard, T.; Rossmeisl, J.; Christensen, C. H. Towards the computational design of solid catalysts. Nat. Chem. 2009, 1, 37−46. (13) Parsons, R. The rate of electrolytic hydrogen evolution and the heat of adsorption of hydrogen. Trans. Faraday Soc. 1958, 54, 1053− 1063. (14) Conway, B. E.; Bockris, J. O. apos; M., Electrolytic Hydrogen Evolution Kinetics and Its Relation to the Electronic and Adsorptive Properties of the Metal. J. Chem. Phys. 1957, 26, 532−541. (15) Herranz, J.; Durst, J.; Fabbri, E.; Patru, A.; Cheng, X.; Permyakova, A. A.; Schmidt, T. J. Interfacial effects on the catalysis of the hydrogen evolution, oxygen evolution and CO2-reduction reactions for (co-)electrolyzer development. Nano Energy 2016, in press, corrected proof; doi: 10.1016/j.nanoen.2016.01.027. (16) Fabbri, E.; Habereder, A.; Waltar, K.; Kotz, R.; Schmidt, T. J. Developments and perspectives of oxide-based catalysts for the oxygen evolution reaction. Catal. Sci. Technol. 2014, 4, 3800−3821. (17) McCrory, C. C. L.; Jung, S.; Ferrer, I. M.; Chatman, S. M.; Peters, J. C.; Jaramillo, T. F. Benchmarking Hydrogen Evolving Reaction and Oxygen Evolving Reaction Electrocatalysts for Solar Water Splitting Devices. J. Am. Chem. Soc. 2015, 137, 4347−4357. (18) Millet, P.; Mbemba, N.; Grigoriev, S. A.; Fateev, V. N.; Aukauloo, A.; Etiévant, C. Electrochemical performances of PEM water electrolysis cells and perspectives. Int. J. Hydrogen Energy 2011, 36, 4134−4142. (19) Cruz, J. C.; Baglio, V.; Siracusano, S.; Ornelas, R.; Ortiz-Frade, L.; Arriaga, L. G.; Antonucci, V.; Aricò, A. S. Nanosized IrO2 electrocatalysts for oxygen evolution reaction in an SPE electrolyzer. J. Nanopart. Res. 2011, 13, 1639−1646. 6602


Article

Chemistry of Materials ray absorption spectroscopy with 10 ms time resolution. J. Synchrotron Radiat. 2016, 23, 260−266. (41) Newville, M. IFEFFIT: interactive XAFS analysis and FEFF fitting. J. Synchrotron Radiat. 2001, 8, 322−324. (42) Ravel, B.; Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537−541. (43) Schmidt, T. J.; Gasteiger, H. A.; Stäb, G. D.; Urban, P. M.; Kolb, D. M.; Behm, R. J. Characterization of High-Surface-Area Electrocatalysts Using a Rotating Disk Electrode Configuration. J. Electrochem. Soc. 1998, 145, 2354−2358. (44) Binninger, T.; Fabbri, E.; Patru, A.; Garganourakis, M.; Han, J.; Abbott, D. F.; Sereda, O.; Kötz, R.; Menzel, A.; Nachtegaal, M.; Schmidt, T. J. Electrochemical Flow-Cell Setup for In Situ X-ray Investigations: I. Cell for SAXS and XAS at Synchrotron Facilities. J. Electrochem. Soc. 2016, 163, H906−H912. (45) Mayousse, E.; Maillard, F.; Fouda-Onana, F.; Sicardy, O.; Guillet, N. Synthesis and characterization of electrocatalysts for the oxygen evolution in PEM water electrolysis. Int. J. Hydrogen Energy 2011, 36, 10474−10481. (46) Song, S.; Zhang, H.; Ma, X.; Shao, Z.; Baker, R. T.; Yi, B. Electrochemical investigation of electrocatalysts for the oxygen evolution reaction in PEM water electrolyzers. Int. J. Hydrogen Energy 2008, 33, 4955−4961. (47) Sen, F. G.; Kinaci, A.; Narayanan, B.; Gray, S. K.; Davis, M. J.; Sankaranarayanan, S. K. R. S.; Chan, M. K. Y. Towards accurate prediction of catalytic activity in IrO2 nanoclusters via first principlesbased variable charge force field. J. Mater. Chem. A 2015, 3, 18970− 18982. (48) Standard x-ray Diffraction Powder Patterns; U.S. Dept. of Commerce, National Bureau of Standards: Washington, DC, 1965; Vol. 25, p 19. (49) Hillman, A. R.; Skopek, M. A.; Gurman, S. J. X-Ray spectroscopy of electrochemically deposited iridium oxide films: detection of multiple sites through structural disorder. Phys. Chem. Chem. Phys. 2011, 13, 5252−5263. (50) Choy, J.-H.; Kim, D.-K.; Demazeau, G.; Jung, D.-Y. LIII-Edge XANES Study on Unusually High Valent Iridium in a Perovskite Lattice. J. Phys. Chem. 1994, 98, 6258−6262. (51) Sardar, K.; Petrucco, E.; Hiley, C. I.; Sharman, J. D. B.; Wells, P. P.; Russell, A. E.; Kashtiban, R. J.; Sloan, J.; Walton, R. I. WaterSplitting Electrocatalysis in Acid Conditions Using Ruthenate-Iridate Pyrochlores. Angew. Chem., Int. Ed. 2014, 53, 10960−10964. (52) Cruz, A. M.; Abad, L.; Carretero, N. M.; Moral-Vico, J.; Fraxedas, J.; Lozano, P.; Subías, G.; Padial, V.; Carballo, M.; CollazosCastro, J. E.; Casañ-Pastor, N. Iridium Oxohydroxide, a Significant Member in the Family of Iridium Oxides. Stoichiometry, Characterization, and Implications in Bioelectrodes. J. Phys. Chem. C 2012, 116, 5155−5168. (53) Pauporté, T.; Aberdam, D.; Hazemann, J.-L.; Faure, R.; Durand, R. X-ray absorption in relation to valency of iridium in sputtered iridium oxide films. J. Electroanal. Chem. 1999, 465, 88−95. (54) Mo, Y.; Stefan, I. C.; Cai, W.-B.; Dong, J.; Carey, P.; Scherson, D. A. In Situ Iridium LIII-Edge X-ray Absorption and Surface Enhanced Raman Spectroscopy of Electrodeposited Iridium Oxide Films in Aqueous Electrolytes. J. Phys. Chem. B 2002, 106, 3681−3686. (55) Pfeifer, V.; Jones, T. E.; Velasco Vélez, J. J.; Massué, C.; Arrigo, R.; Teschner, D.; Girgsdies, F.; Scherzer, M.; Greiner, M. T.; Allan, J.; Hashagen, M.; Weinberg, G.; Piccinin, S.; Hävecker, M.; KnopGericke, A.; Schlögl, R. The electronic structure of iridium and its oxides. Surf. Interface Anal. 2016, 48, 261−273. (56) Bozack, M. J. Sputter-Induced Modifications of IrO2 During XPS Measurements. Surf. Sci. Spectra 1993, 2, 123−127. (57) Sanchez Casalongue, H. G.; Ng, M. L.; Kaya, S.; Friebel, D.; Ogasawara, H.; Nilsson, A. In Situ Observation of Surface Species on Iridium Oxide Nanoparticles during the Oxygen Evolution Reaction. Angew. Chem., Int. Ed. 2014, 53, 7169−7172. (58) Pfeifer, V.; Jones, T. E.; Velasco Velez, J. J.; Massue, C.; Greiner, M. T.; Arrigo, R.; Teschner, D.; Girgsdies, F.; Scherzer, M.; Allan, J.;

Hashagen, M.; Weinberg, G.; Piccinin, S.; Havecker, M.; KnopGericke, A.; Schlogl, R. The electronic structure of iridium oxide electrodes active in water splitting. Phys. Chem. Chem. Phys. 2016, 18, 2292−2296. (59) Smith, R. D. L.; Sporinova, B.; Fagan, R. D.; Trudel, S.; Berlinguette, C. P. Facile Photochemical Preparation of Amorphous Iridium Oxide Films for Water Oxidation Catalysis. Chem. Mater. 2014, 26, 1654−1659. (60) Chen, H.; Trasatti, S. Oxygen evolution on aged IrOx /Ti electrodes. I. Acidic solutions. J. Indian Chem. Soc. 1993, 70, 323−330. (61) Marshall, A. T.; Haverkamp, R. G. Nanoparticles of IrO2 or Sb− SnO2 increase the performance of iridium oxide DSA electrodes. J. Mater. Sci. 2012, 47, 1135−1141. (62) Reier, T.; Oezaslan, M.; Strasser, P. Electrocatalytic Oxygen Evolution Reaction (OER) on Ru, Ir, and Pt Catalysts: A Comparative Study of Nanoparticles and Bulk Materials. ACS Catal. 2012, 2, 1765− 1772. (63) Yeo, R. S.; Orehotsky, J.; Visscher, W.; Srinivasan, S. Ruthenium-Based Mixed Oxides as Electrocatalysts for Oxygen Evolution in Acid Electrolytes. J. Electrochem. Soc. 1981, 128, 1900− 1904. (64) Vuković, M. Oxygen evolution reaction on thermally treated iridium oxide films. J. Appl. Electrochem. 1987, 17, 737−745. (65) Angelinetta, C.; Trasatti, S.; Atanasoska, L. D.; Minevski, Z. S.; Atanasoski, R. T. Effect of preparation on the surface and electrocatalytic properties of RuO2 + IrO2 mixed oxide electrodes. Mater. Chem. Phys. 1989, 22, 231−247. (66) Frazer, E. J.; Woods, R. The oxygen evolution reaction on cycled iridium electrodes. J. Electroanal. Chem. Interfacial Electrochem. 1979, 102, 127−130. (67) Burke, L. D.; O’Sullivan, E. J. M. Oxygen gas evolution on hydrous oxides  An example of three-dimensional electrocatalysis? J. Electroanal. Chem. Interfacial Electrochem. 1981, 117, 155−160. (68) Ouattara, L.; Fierro, S.; Frey, O.; Koudelka, M.; Comninellis, C. Electrochemical comparison of IrO2 prepared by anodic oxidation of pure iridium and IrO2 prepared by thermal decomposition of H2IrCl6 precursor solution. J. Appl. Electrochem. 2009, 39, 1361−1367. (69) Chandra, D.; Takama, D.; Masaki, T.; Sato, T.; Abe, N.; Togashi, T.; Kurihara, M.; Saito, K.; Yui, T.; Yagi, M. Highly Efficient Electrocatalysis and Mechanistic Investigation of Intermediate IrOx(OH)y Nanoparticle Films for Water Oxidation. ACS Catal. 2016, 6, 3946−3954. (70) Rasten, E.; Hagen, G.; Tunold, R. Electrocatalysis in water electrolysis with solid polymer electrolyte. Electrochim. Acta 2003, 48, 3945−3952. (71) Marshall, A.; Børresen, B.; Hagen, G.; Tsypkin, M.; Tunold, R. Hydrogen production by advanced proton exchange membrane (PEM) water electrolysersReduced energy consumption by improved electrocatalysis. Energy 2007, 32, 431−436. (72) Stoerzinger, K. A.; Qiao, L.; Biegalski, M. D.; Shao-Horn, Y. Orientation-Dependent Oxygen Evolution Activities of Rutile IrO2 and RuO2. J. Phys. Chem. Lett. 2014, 5, 1636−1641. (73) Danilovic, N.; Subbaraman, R.; Chang, K.-C.; Chang, S. H.; Kang, Y. J.; Snyder, J.; Paulikas, A. P.; Strmcnik, D.; Kim, Y.-T.; Myers, D.; Stamenkovic, V. R.; Markovic, N. M. Activity−Stability Trends for the Oxygen Evolution Reaction on Monometallic Oxides in Acidic Environments. J. Phys. Chem. Lett. 2014, 5, 2474−2478. (74) Chang, S. H.; Connell, J. G.; Danilovic, N.; Subbaraman, R.; Chang, K.-C.; Stamenkovic, V. R.; Markovic, N. M. Activity-stability relationship in the surface electrochemistry of the oxygen evolution reaction. Faraday Discuss. 2014, 176, 125−133. (75) Binninger, T.; Mohamed, R.; Waltar, K.; Fabbri, E.; Levecque, P.; Kötz, R.; Schmidt, T. J. Thermodynamic explanation of the universal correlation between oxygen evolution activity and corrosion of oxide catalysts. Sci. Rep. 2015, 5, 12167. (76) Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions; National Association of Corrosion Engineers: Houston, TX, 1974. 6603


Article

Chemistry of Materials (77) Andreas, H.; Elzanowska, H.; Serebrennikova, I.; Birss, V. Hydrous Ir Oxide Film Properties at Sol-Gel Derived Ir Nanoparticles. J. Electrochem. Soc. 2000, 147, 4598−4604. (78) Mozota, J.; Conway, B. E. Surface and bulk processes at oxidized iridium electrodesI. Monolayer stage and transition to reversible multilayer oxide film behaviour. Electrochim. Acta 1983, 28, 1−8. (79) Juodkazytė, J.; Šebeka, B.; Valsiunas, I.; Juodkazis, K. Iridium Anodic Oxidation to Ir(III) and Ir(IV) Hydrous Oxides. Electroanalysis 2005, 17, 947−952. (80) Birss, V. I.; Bock, C.; Elzanowska, H. Hydrous Ir oxide films: the mechanism of the anodic prepeak reaction. Can. J. Chem. 1997, 75, 1687−1693. (81) Minguzzi, A.; Lugaresi, O.; Locatelli, C.; Rondinini, S.; D’Acapito, F.; Achilli, E.; Ghigna, P. Fixed Energy X-ray Absorption Voltammetry. Anal. Chem. 2013, 85, 7009−7013. (82) Kongkanand, A.; Ziegelbauer, J. M. Surface Platinum Electrooxidation in the Presence of Oxygen. J. Phys. Chem. C 2012, 116, 3684−3693. (83) Teliska, M.; O’Grady, W. E.; Ramaker, D. E. Determination of O and OH Adsorption Sites and Coverage in Situ on Pt Electrodes from Pt L23 X-ray Absorption Spectroscopy. J. Phys. Chem. B 2005, 109, 8076−8084. (84) Cherevko, S.; Geiger, S.; Kasian, O.; Mingers, A.; Mayrhofer, K. J. J. Oxygen evolution activity and stability of iridium in acidic media. Part 2. − Electrochemically grown hydrous iridium oxide. J. Electroanal. Chem. 2016, 774, 102−110.

6604


Iridium Oxide for the Oxygen Evolution Reaction: Correlation between

Recommend Documents