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Nickel–silver alloy electrocatalysts for hydrogen evolution and oxidation in an alkaline electrolyte† Maureen H. Tang,a Christopher Hahn,ab Aidan J. Klobuchar,ab Jia Wei Desmond Ng,a Jess Wellendorff,ab Thomas Bligaardab and Thomas F. Jaramillo*a The development of improved catalysts for the hydrogen evolution reaction (HER) and hydrogen oxidation reaction (HOR) in basic electrolytes remains a major technical obstacle to improved fuel cells, water electrolyzers, and other devices for electrochemical energy storage and conversion. Based on the free energy of adsorbed hydrogen intermediates, theory predicts that alloys of nickel and silver are active for these reactions. In this work, we synthesize binary nickel–silver bulk alloys across a range of compositions and show that nickel–silver alloys are indeed more active than pure nickel for hydrogen evolution and, possibly, hydrogen oxidation. To overcome the mutual insolubility of silver and nickel, we employ electron-beam physical vapor codeposition, a low-temperature synthetic route to metastable alloys. This method also produces flat and uniform films that facilitate the measurement of intrinsic catalytic activity with minimal variations in the surface area, ohmic contact, and pore transport.
Received 30th March 2014, Accepted 30th June 2014
Rotating-disk-electrode measurements demonstrate that the hydrogen evolution activity per geometric
DOI: 10.1039/c4cp01385a
has comparable stability and hydrogen oxidation activity. Our experimental results are supported by
area of the most active catalyst in this study, Ni0.75Ag0.25, is approximately twice that of pure nickel and density functional theory calculations, which show that bulk alloying of Ni and Ag creates a variety of
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adsorption sites, some of which have near-optimal hydrogen binding energy.
1 Introduction The electrochemical evolution and oxidation of hydrogen are reactions of both practical importance and fundamental interest. The hydrogen oxidation reaction (HOR) is employed in polymerelectrolyte-membrane fuel cells that convert chemical energy to electricity, while the hydrogen evolution reaction (HER) is relevant to water and chlor-alkali electrolysis, electrodeposition, and corrosion. As early as 1957, the hydrogen evolution reaction has also contributed to the fundamental understanding of electrochemistry and electrocatalysis.1 Recently, the HER has been used in exploring the fundamental differences between electrocatalysis in acid and base.2 It has long been recognized that the reaction rates of the HOR and the HER are often slower in basic electrolytes than acidic electrolytes, even though the surface intermediate of adsorbed a
Department of Chemical Engineering, Stanford University, Stanford, CA 94305, USA. E-mail: [email protected] b SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, CA 94025, USA † Electronic supplementary information (ESI) available: Detailed information on the synthesis procedure and experimental characterizations, including stability testing, high-resolution XPS spectra, XRD fitting results, and OH adsorption experiments. See DOI: 10.1039/c4cp01385a
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hydrogen (Had) is independent of solution pH.3 Early fundamental studies suggested that, at high pH, OH competes with Had for surface sites, effectively poisoning the electrode and reducing overall rates.4,5 More recent work has focused on the recombination step between Had and OHad during hydrogen oxidation, suggesting that optimal adsorption of OH as well as H is necessary to facilitate recombination of Had and OHad to form H2O.2 In the reverse direction, protons must be removed from water before adsorbing onto the electrode and recombining as hydrogen; this water-splitting step may also be influenced by OHad species on the electrode.6 While the role of adsorbed OH in the alkaline HOR and HER is still unclear, the effect of hydrogen adsorption strength is much less ambiguous. Fig. 1 plots previously published values of measured hydrogen evolution activity under basic conditions versus the calculated free energy of hydrogen surface binding (DGHad). The activity follows a volcano-type relation in which the best catalysts exhibit hydrogen binding free energy closest to the optimum value of zero.7 As in acid, platinum and palladium are the most active elemental catalysts, and nickel is the most active non-precious elemental electrocatalyst. Thus, from a catalystdevelopment point of view, appropriately tuning the hydrogen binding energy is viewed as a necessary, but not sufficient, requirement for an active catalyst in alkaline environments.
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Fig. 1 Plot of hydrogen evolution activity under basic conditions (measured) versus hydrogen adsorption energy (calculated). As in acid,7 the data show a volcano trend. Platinum, nickel, and palladium bind hydrogen with the energy closest to optimum, but still offer room for improvement. All hydrogen binding energies are from ref. 10, except for Ag and Ni values, which were calculated in this work. Alkaline measurements are from ref. 11 (red squares, 30% KOH, 80 1C) and ref. 12 (blue circles, 0.1 M KOH, 25 1C). The dashed lines are to guide the eye. Arrows mark the hydrogen binding energy to surface alloys of (a) Ni on the Ag substrate and (b) Ag on the Ni substrate.9
Recent work on non-noble electrocatalysts for alkaline hydrogen oxidation has also attributed extremely high activity to optimal hydrogen binding.8 Computational screening is a useful tool to make predictions for new materials, which helps in guiding the development of active and stable catalysts. Previous theoretical calculations of DGHad predicted that surface alloys of silver and nickel bind hydrogen with an energy very close to optimum, as indicated by the arrows in Fig. 1.9 The optimal binding energy of NiAg surface alloys suggests that bulk NiAg alloys may be similarly active and motivates our further experimental and theoretical investigations into these catalysts. Experimentally, bimetallic NiAg systems have shown promise for hydrogen generation both electrocatalytically13 and in homogenous catalysis.14 However, synthesis of binary NiAg alloys is challenging because thermodynamic instability causes Ag and Ni to phasesegregate into pure elemental phases at all temperatures.15 This immiscibility, combined with a large lattice mismatch (0.352 vs. 0.408 nm) and the lower surface energy of Ag, renders the Ag–Ni core–shell the most thermodynamically favorable arrangement for nanoparticles.16,17 Circumventing the phase diagram and forming a metastable alloy require kinetically controlled materials synthesis, such as rapid reduction or condensation at low temperature, where Ag and Ni atoms do not have enough kinetic energy to diffuse and segregate. Previously, various thin film and nanoparticle synthesis techniques have been applied for this purpose, including electrodeposition, pulse laser deposition, e-beam PVD, DC sputtering, and gamma-irradiation.18–22 In this work, e-beam PVD was chosen for its ability to access a wide range of alloy compositions. PVD has the additional advantage of producing flat, conformal films on a variety of substrates. Although alkaline electrolysis is an established technology for which
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many materials have been tested as catalysts, many previous studies are developed around high surface area structures, such as RANEYs nickel, and tested as composite electrodes in a fuel cell or electrolyzer.23–26 Although such conditions most accurately represent those during device operation, distinguishing between catalysis, reactant transport, and electron transport effects in a complicated device geometry is extremely difficult. For example, nickel–molybdenum and nickel–molybdenum–cobalt catalysts show the best HER activity of all non-precious metals, but the reason for the improvement over pure nickel has been alternatively attributed to increased microscopic porosity and increased intrinsic catalytic activity.8,27–30 In other intermetallic systems, improved activity in the presence of dopant metals has been attributed to a variety of reasons, including better ohmic contact between particles, better mass transport through catalyst pores, better mass transport through electrode pores, and better catalysis at the particle surface.26,31–33 Physical vapor deposition produces flat, conformal films of uniform composition, avoiding such complications and enabling fundamental studies to provide a greater insight. Additionally, synthesizing thin films of catalysts on a glassy carbon substrate permits alloys to be tested in a rotating-disk configuration that separates catalytic activity from the surface area, ohmic contact, and pore transport.
2 Methods 2.1
Computational details
All theoretical results are based on density functional theory calculations. These are performed using Quantum ESPRESSO,34,35 a plane-wave36 pseudopotential37 code, and the Atomic Simulation Environment.38 We used the RPBE39 exchange–correlation functional and ultrasoft40 pseudopotentials from the PS Library project.41 Kinetic energy cutoffs were 700 eV for wave functions and 6500 eV for charge densities and potentials, and the surface Brillouin zone was sampled using a 6 6 1 Monkhorst–Pack42 k-point mesh. A Gaussian smearing of 0.2 eV was applied to the electron occupation numbers near the Fermi level. Periodically repeated 5-layer slabs were used to model hydrogen adsorption onto three close-packed fcc(111) NiAg alloy surfaces (Ni3Ag, NiAg, and NiAg3) as well as onto the pure Ni and Ag(111) facets. A 2 2 surface unit cell with one H atom per cell was applied. This resulted in a 1/4 monolayer coverage of atomic H on the transition-metal substrates and several structurally non-equivalent adsorption sites, particularly, on the Ni3Ag and NiAg3 alloys. The three alloyed surfaces were alloyed throughout the bulk, and the top two layers of all surfaces and the H adsorbate were relaxed until all residual forces acting on them were less than 0.05 eV Å 1. The bottom three metal layers were fixed in the bulk structure, with lattice constants computed using an equation of state approach. 2.2
Synthesis of NiAg alloys
Catalysts were deposited on 5 mm diameter glassy carbon disks (Hochtemperatur Werkstoffe) and 1.6 cm quartz slides (GM Associates) using a load-locked dual e-beam evaporator equipped
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with separate quartz crystal microbalances for each source (Temescal VES2550). Additional Ni films were separately deposited using an Innotec ES26C e-beam evaporator. The total deposition rate was maintained at approximately 2 Å s 1 for all depositions. To improve substrate–catalyst adhesion, glassy carbon disks were anodized in air at 400 1C for 2 hours before sonication in acetone, isopropanol, and water, followed by mounting on wafers for deposition. An 80 nm Ti adhesion layer was deposited on the quartz substrates, but was found to increase delamination from the carbon substrates during catalyst testing, possibly due to hydrogen embrittlement of the Ti. Therefore, no Ti layer was used on glassy carbon substrates. Ni and Ag were deposited in different ratios by separately adjusting the electron beam power for each source. The quartz-crystal monitors were calibrated for cross-talk between the metal sources at the beginning of each deposition; depending on the speed of equilibration, this calibration period resulted in a 5–30 nm layer of pure Ag underneath the 70–85 nm NiAg alloy catalyst layer. 2.3
Physical and chemical characterizations
After deposition, catalysts were characterized by scanning electron microscopy (SEM) using a FEI XL30 Sirion with a beam voltage of 5.0 kV. X-ray diffraction (XRD) was performed on a PANanalytical X’Pert Pro using Cu Ka radiation (l = 1.541 Å). X-ray photoelectron spectroscopy (XPS) was performed using a PHI Versaprobe Scanning XPS using Al Ka radiation (hn = 1486.6 eV, spot size 200 mm, 451 collection angle). The roughness factor of samples was also measured by atomic force microscopy (AFM), using a Park Systems XE-70 in non-contact mode. 2.4
Electrochemical testing
Electrochemical tests were conducted in 0.1 M KOH (Sigma) using a rotating disk electrode (RDE) set-up (Pine Instruments) and a three-electrode potentiostat (Biologic Instruments) with ohmic resistance compensation. A cell made of corrosion-resistant quartz (Pine) was used to reduce contamination from KOH etching in conventional borosilicate glass.43 All potentials were measured with respect to a mercury/mercuric oxide reference electrode (1.0 M KOH, Koslow Scientific), which was separated from the main cell by a potassium nitrate salt bridge. The reference electrode was calibrated to 0.869 V versus the reversible hydrogen electrode (RHE); all potentials in this work are reported vs. RHE. The counterelectrode was Ni wire (Alfa, 99.99%). Although heat treatment of the glassy carbon disks improved catalyst–substrate adhesion, catalyst delamination still occurred during extended periods of hydrogen bubbling. Longer stability tests were conducted on quartz substrates with better adhesion properties. Further details of sample preparation are given in the ESI.† The competition between the hydrogen reactions and the formation and removal of Ni hydrides and oxides can lead to erroneous activity measurements.44 The testing protocol was developed to isolate the catalytic activity for HER and HOR from these competing reactions. After insertion into the electrolyte, the working electrode was activated with a sweep from open circuit potential to 300 mV vs. RHE at 10 mV s 1. This step reduced surface oxides while permitting minimal time for
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hydrogen absorption. The electrode was then held at 20 mV for seven minutes while purging hydrogen gas (Praxair, 99.999%) through the electrolyte. Linear sweep voltammograms of the hydrogen evolution activity were then recorded between 300 and 20 mV. Next, HOR activity was measured by chronoamperometry steps from +50 to +200 mV. Chronoamperometry instead of cyclic voltammetry was used to separate steady-state hydrogen oxidation from transients due to double-layer charging, oxidative removal of Ni hydrides, and oxidation of Ni catalysts. To demonstrate that the observed anodic current was due to hydrogen oxidation, separate samples were tested using an identical procedure under a nitrogen atmosphere as control experiments.
3 Results and discussion 3.1
Uniform deposition of metastable NiAg alloys
SEM was used to determine the morphology of the catalysts synthesized on glassy carbon substrates (Fig. 2). The images show that all of the films are flat and conformal. The grain size, which is approximately 20 nm for Ni and the Ni-rich alloys, increases to about 40 nm for pure Ag. The XPS spectra, shown in Fig. 3, demonstrate the composition range of the depositions. The integrated areas of the Ni2p region from 850–890 eV and the Ag3d region from 364–378 eV were used to calculate the relative surface compositions, which are plotted versus nominal composition in Fig. 4. Depositions on both quartz and glassy carbon show good agreement with the nominal composition. In contrast to studies in which NiAg systems were synthesized at high temperature or by a wet chemical method, no surface enrichment of Ag was observed, suggesting that atoms were deposited with insufficient thermal energy to separate into the thermodynamically preferred arrangement.14,22 The oxygen peak at 532 eV indicates that the catalysts form a native oxide in air, and that Ni is more oxidized than Ag, as expected; the presence of Ni surface oxides was confirmed by high-resolution XPS.† XRD measurements showed that the e-beam deposition method was able to achieve partial metastable alloying of Ni and Ag. All samples, except for the pure metals, were found to have both Ni-rich and Ag-rich alloy phases. Diffractograms of the catalysts deposited on quartz are shown in Fig. 5. The counts are
Fig. 2 SEM images of catalysts as synthesized on glassy carbon substrates. Films are flat and conformal, with Ag depositing on larger grains than Ni or the NiAg alloys.
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Fig. 3 XPS survey spectra of the vapor-deposited NiAg samples (glassy carbon substrate). The trends in Ni2p and Ag3d peaks show the compositional variation of the depositions.
Fig. 5 XRD diffractograms recorded on quartz substrates. The curves are normalized to the maximum intensity and corrected for the substrate diffraction. Although the metals segregate into two phases, shifts in the Ni(111) and Ag(111) peaks demonstrate that each phase is a metastable alloy. Also shown are reference peaks for Ni (JCPDS 4-0850), Ag (JCPDS 4-1783), and Ti (JCPDS 44-1294).
Fig. 4 Nominal versus measured Ni concentration. Bulk concentrations of the Ni-rich and Ag-rich alloy phases are calculated from Vegard’s Law. Average surface concentrations are measured using XPS. Closed symbols are for quartz substrates, and open symbols are for glassy carbon. The maximum concentration of Ag in the Ni-rich alloy phase is about 13 at%.
corrected for the quartz substrate and shown to be normalized to the maximum intensity. For pure Ni, the Ni(111) reflection is observed at a = 0.353 nm, which is consistent with the literature value. As the Ag deposition rate increases, incorporation of larger Ag atoms into the Ni lattice shifts the peak to smaller values of 2y. Alloying Ag also decreases the crystallinity and crystallite size, as evidenced by the smaller peak size and an increase in FWHM. Similarly, the peak position of the Ag(111) reflection increases from that of pure Ag (a = 0.409 nm), showing incorporation of the smaller Ni atoms in Ag-rich films. The shifting of the Ag(111) peak can only be observed visually for the Ag-rich films in Fig. 5 because the (002) reflection of the Ti adhesion layer overlaps for Ni-rich films with lower intensity Ag(111) reflection.
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To quantitatively determine the extent of alloying between Ni and Ag, Gaussian curves were fit to the Ni(111) and Ag(111) peaks. The peak location and FWHM were used to calculate the alloy lattice constant and grain size for each of the two phases according to Bragg’s law and the Scherrer equation, respectively; calculations and results are shown in the ESI.† Grain sizes were calculated between 7 and 14 nm for alloys, slightly smaller than indicated by the SEM images. Inhomogeneities in composition may explain the slight disparity between the two methods. The relative fractions of Ag and Ni as computed by Vegard’s law are shown in Fig. 4. Fig. 4 indicates that the maximum Ag content introduced into the Ni-rich phase by this method is about 13%, at a bulk ratio of Ni0.25Ag0.75. The Ag-rich phase reaches a maximum solubility of approximately 6% Ni, at a bulk ratio of Ni0.125Ag0.875. At a nominal composition of Ni0.125Ag0.875, the solubility of Ag in the Ni-rich phase decreases. This may be because, at higher concentration, Ag atoms need to diffuse shorter distances in order to segregate into a separate phase. Similar solubility trends in metastable NiAg thin-films have been observed previously.19,20 3.2
Electrochemical performance of alloys
The combined materials characterization results of Fig. 2 through 5 show that the films synthesized by e-beam codeposition are flat, uniform, and composed of a Ni-rich phase and an Ag-rich phase.
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Fig. 6 Electrochemical activity measurements of Ni, Ag, and selected NiAg alloys on glassy carbon substrates in 0.1 M KOH/hydrogen gas at 900 rpm. (a) Cathodic HER sweep at 10 mV s 1. (b) HOR at potential steps from +50 to +200 mV. No oxidation is seen during a control experiment in nitrogen.
Fig. 7 Current density at 300 mV (left axis, blue diamonds) and +150 mV (right axis, red squares) vs. nominal Ni concentration. The best-performing catalysts have nominal compositions between 50 and 85% Ni. Error bars represent two standard deviations, based on measurements of at least three samples per composition.
To relate the chemical composition to hydrogen evolution activity, linear sweep voltammetry was performed in an RDE configuration. Fig. 6 and 7 show the HER and HOR activities of Ni, Ag, and selected NiAg alloys. Fig. 6 shows that, as expected, the hydrogen evolution activity of pure Ag is negligible. Pure Ni displays an activity consistent with the literature,12 while the Ni0.75Ag0.25 alloy shows the highest HER activity for these samples (Fig. 6(a)). The current density at 300 mV overpotential is shown in Fig. 7, demonstrating a broad maximum in HER activity at an alloy composition of Ni0.75Ag0.25. The HOR activity at constant potentials is shown in Fig. 6(b) for Ni, Ni0.75Ag0.25, and Ni0.5Ag0.5. The Ni0.75Ag0.25 alloy shows higher current than
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the pure Ni sample at overpotentials of 150 mV or less, but the difference in current disappears at 200 mV overpotential. This change may be caused by a different HOR mechanism at higher overpotentials, or because oxidation of the catalyst at these potentials changes the catalytic properties of the surface. Similar behavior in other Ni-based systems was ascribed to catalyst oxidation.8,24,27 Noise in the data is caused in part by bubbles of hydrogen gas contacting the electrode. Also shown in Fig. 6(b) are oxidation steps for a separate Ni sample tested in a nitrogensaturated atmosphere. The difference in current shows that the anodic current in Fig. 6(b) is due to HOR and not simply because of catalyst oxidation. While the uncertainty associated with measuring small currents in the presence of vigorous bubbling makes the HOR trend less obvious, the Ni-rich alloys have an activity comparable to, if not better than, the pure Ni films, as shown in Fig. 7. Stability testing also shows a similar behavior between the NiAg alloys and pure Ni.† The plot of the current density at 300 mV vs. nominal Ni concentration in Fig. 7 shows that, for HER, the optimum catalyst concentration is in between 50 and 85% Ni. The peak alloy activity is about 9.4 mA cm 2, approximately twice that of pure Ni on a geometric-area basis. Examination of Fig. 6(a) shows that obtaining a current density of 2 mA cm 2 on elemental Ni requires about 260 mV of overpotential, slightly higher than the overpotential for Ni as plotted in Fig. 1. Consideration of surface roughness may account for slight differences in the measured activity from previously published values. Fig. 7 also shows that the HER activity of Ni0.5Ag0.5 exceeds that of pure Ni by about 60%, but that the HOR activity of this alloy is negligible. This difference suggests that the hydrogen binding energy for the NiAg alloy is stronger than optimum, so that adsorbate–adsorbate repulsions between hydrogen atoms can improve the HER activity but not the HOR.45 It is challenging to separate the intrinsic catalytic activity from the effects of the microscopic electrode surface area for non-noble electrocatalysts. Because of the low oxidation potential of Ni, commonly used methods for measuring the surface area, including underpotential deposition of H, Cu, and Pb, are not applicable to Ni-based systems. The low overpotential for the HER also eliminates the possibility of finding a non-faradaic region of the voltage window in which the double-layer capacitance is measured. In this work, however, the vacuum deposition synthetic approach creates flat, uniform films with negligible porosity. AFM images, shown in Fig. 8, show negligible differences in the surface area among the different films. Therefore, improvements in HER and HOR activity were caused by intrinsic effects on catalysis and not by the increased surface area. Fig. 8 shows that for both samples, grains are only a few nm tall, yielding calculated surface roughness factors of 1.04 and 1.03, respectively, for the Ni0.825Ag0.175 and Ni samples. Measurements of OH adsorption also suggested that the NiAg catalysts have similar or smaller areas of electrochemically active Ni (ECSANi) than films of pure Ni.† 3.3
Comparison with theoretical values
Although the AFM image in Fig. 8 shows negligible surface roughness, determining the intrinsic catalytic activity for the
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Fig. 8 AFM images and line scans of Ni and Ni0.825Ag0.375 on glassy carbon. Differences in line scans appear larger than actual because of the axes aspect ratio.
NiAg alloys is still not straightforward because the active site of the alloys is not known. To develop a qualitative insight into the active site for HER and HOR catalysis, DFT calculations were used to determine the hydrogen binding energy at different sites of Ni, NiAg3, NiAg, Ni3Ag, and Ag. The results of all of the calculated sites are shown in Table 1, and selected binding energies and geometries are shown in Fig. 9. The theoretical hydrogen evolution current density was calculated using the `lason, et al.46 and normalized to the experimentally model of Sku measured current density for pure Ni. Because the model was developed for acidic electrolyte, not basic electrolyte, it can be used only for qualitative interpretation in studying trends. As previously discussed, the effects of hydroxide ions on alkaline HER and HOR are not entirely understood, and a quantitative microkinetic model that incorporates their effects is beyond the scope of this work. Table 1 and Fig. 9 indicate that simple alloying of Ni and Ag can obtain many energetically different adsorption sites, and that several NiAg structures have sites where hydrogen is predicted to bind with energy close to the optimum value of 0 eV. The full 4d-shell of the Ag atoms is expected to shift the projected d-band downwards from the pure-Ni case, leading to
Table 1 Calculated hydrogen binding free energies (eV) on uniformly alloyed NiAg(111) surfaces. Selected images of the calculated surface structures are shown in Fig. 9
Composition
Binding site
Ni
Hollow On-top Hollow (Ni) Hollow (Ag) On-top (Ni) On-top (Ag) Bridge Hollow (Ag) On-top (Ni) Hollow (Ni) On-top (Ni) Hollow (Ag) Hollow On-top
Ni3Ag
NiAg NiAg3 Ag
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DGHad (eV) 0.19 0.38 0.25 0.07 0.34 1.04 0.15 0.16 0.32 0.09 0.26 0.47 0.46 0.99
Fig. 9 Above, combined theoretical-experimental volcano plot for mostactive sites. Markers are simulations, and the horizontal lines are measured currents for: pure Ni, geometric area (grey); Ni0.75Ag0.25, geometric area (green); Ni0.75Ag0.25, area of Ag in the Ni phase (blue). The dotted red line represents the predicted current.46 Below, selected images of the calculated surface structures are shown.
weaker hydrogen binding and more active catalysts. This effect is observed for binding sites which include both Ni and Ag atoms, such as the Ni hollow site on NiAg3 and the Ag hollow site on Ni3Ag. Such binding sites generally have the DGHad values closest to the optimum of 0 eV. Another competing effect is seen at the Ni hollow site of Ni3Ag, which binds more strongly than the pure Ni site. Similarly, the Ag hollow site of NiAg3 binds more weakly than the pure Ag site. This is due to the effect of the alloyed atoms not in the binding site. For example, in the Ni3Ag case, there are Ni atoms in the binding site which interact with an Ag atom. This Ni–Ag bond has lower saturation compared to a Ni–Ni bond, leading to a higher saturation of the Ni–H bond, and thus to stronger hydrogen binding. The interplay of these two effects results in a variety of surface sites with differences in DGHad, some of which are favorable and some of which are unfavorable for the HER. However, due to the exponential dependence of activity on DGHad, we expect the most active sites to dominate the overall alloy activity. Additionally, the results of Table 1 assume a perfectly ordered lattice, and a random distribution of Ag and Ni may change both the geometric frequency and electronic structure of binding sites. Fig. 9 also shows experimental data as extracted from Fig. 7 for Ni0.75Ag0.25. Because the exact surface structure is not known, the experimentally measured current density is represented as a horizontal green line instead of a discrete point. All experimental and theoretical values are normalized to the experimentally measured activity on pure Ni (grey line). The alloys contain both Ag-rich and Ni-rich phases (Fig. 5), and it is not immediately clear in which phase the active site is located. Inspection of Table 1 shows that the most optimal sites for
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hydrogen binding are Ag atoms in a Ni-rich environment, or Ni atoms in an Ag-rich environment. Based on this theoretical insight, the measured current density for Ni0.75Ag0.25 was scaled to the area fraction of the Ag component in the Ni-rich phase. This value is represented by a blue dashed line in Fig. 9. Scaling in this manner would suggest an enhancement by 100 for each of these sites relative to pure Ni. Because this calculation neglects the activity of all other sites, the estimated current is therefore an upper bound of the possible current per active site. Although the Ni atoms in the Ag-rich phase are expected to be similarly active, they make up a much smaller fraction of the surface, and scaling to this area would yield a current density that exceeds the theoretical limit. Other active sites may occur at the Ni-rich and Ag-rich grain boundaries; hydrogen adsorption at the bridge site between Ni and Ag is predicted to be very close to optimum. Tuning the grain size by controlling synthesis parameters such as deposition rate and temperature might elucidate the roles of different sites and suggest strategies for increasing the density of the most active sites. Finally, the effects of hydroxide adsorption cannot be ignored. Because hydrogen evolution requires an additional step of watersplitting, a bifunctional catalyst is required.2 The calculated values for OH adsorption energy differ by 1.8 eV between Ni and Ag.47 This difference may have synergistic effects on the HER independent of hydrogen adsorption. More work is necessary to understand fully the effect of hydroxide adsorption on alkaline HER and HOR catalysis, and the control afforded by e-beam codeposition motivates further use of the NiAg system to probe these effects.
4 Conclusions Guided by theoretical calculations, we identified NiAg alloys as promising electrocatalysts for the HOR and HER in base. To mitigate the thermodynamic phase segregation of Ag and Ni, e-beam codeposition was used to deposit thin films of metal catalysts on glassy carbon and quartz substrates. This technique permits finely-tuned catalyst compositions in a well-defined morphology. Incorporation of up to 13% Ag into the Ni lattice and 6% Ni into the Ag lattice was demonstrated via XRD. Rotating disk voltammetry showed that Ni-rich NiAg samples were more active than pure Ni for the HER, with the most active catalyst, Ni0.75Ag0.25, outperforming pure Ni by approximately a factor of two. Because of the planar electrode geometry (AFM-measured roughness factors of approximately unity), the improvement may be attributed to intrinsic catalytic activity, not to the higher surface area of Ni. The HOR activities of the Ni-rich NiAg catalysts were comparable to pure Ni. The improved HER activity was corroborated by DFT calculations, which demonstrate that alloying Ni and Ag creates a variety of adsorption sites, several with near-optimal hydrogen binding energies. Future work to elucidate the active site(s) in the NiAg system for alkaline HER and HOR catalysis would help in enabling rational catalyst design, which will in turn facilitate the incorporation of advanced catalysts into functional devices.
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Acknowledgements This work was supported by a TomKat Grant for Energy Research at Stanford University. We also thank the Department of Energy for SUNCAT funding. A.J.K. was supported by an Abbott Laboratories Stanford Graduate Fellowship. Part of this work was performed at the Stanford Nano Center (SNC)/ Stanford Nanocharacterization Laboratory (SNL), part of the Stanford Nano Shared Facilities. This work was also supported by the facilities of the UCSB Nanotech, a member of the National Nanofabrication Infrastructure Network partially supported by the National Science Foundation.
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Nickel–silver alloy electrocatalysts for hydrogen evolution and oxidation in an alkaline electrolyte† Maureen H. Tang,a Christopher Hahn,ab Aidan J. Klobuchar,ab Jia Wei Desmond Ng,a Jess Wellendorff,ab Thomas Bligaardab and Thomas F. Jaramillo*a The development of improved catalysts for the hydrogen evolution reaction (HER) and hydrogen oxidation reaction (HOR) in basic electrolytes remains a major technical obstacle to improved fuel cells, water electrolyzers, and other devices for electrochemical energy storage and conversion. Based on the free energy of adsorbed hydrogen intermediates, theory predicts that alloys of nickel and silver are active for these reactions. In this work, we synthesize binary nickel–silver bulk alloys across a range of compositions and show that nickel–silver alloys are indeed more active than pure nickel for hydrogen evolution and, possibly, hydrogen oxidation. To overcome the mutual insolubility of silver and nickel, we employ electron-beam physical vapor codeposition, a low-temperature synthetic route to metastable alloys. This method also produces flat and uniform films that facilitate the measurement of intrinsic catalytic activity with minimal variations in the surface area, ohmic contact, and pore transport.
Received 30th March 2014, Accepted 30th June 2014
Rotating-disk-electrode measurements demonstrate that the hydrogen evolution activity per geometric
DOI: 10.1039/c4cp01385a
has comparable stability and hydrogen oxidation activity. Our experimental results are supported by
area of the most active catalyst in this study, Ni0.75Ag0.25, is approximately twice that of pure nickel and density functional theory calculations, which show that bulk alloying of Ni and Ag creates a variety of
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adsorption sites, some of which have near-optimal hydrogen binding energy.
1 Introduction The electrochemical evolution and oxidation of hydrogen are reactions of both practical importance and fundamental interest. The hydrogen oxidation reaction (HOR) is employed in polymerelectrolyte-membrane fuel cells that convert chemical energy to electricity, while the hydrogen evolution reaction (HER) is relevant to water and chlor-alkali electrolysis, electrodeposition, and corrosion. As early as 1957, the hydrogen evolution reaction has also contributed to the fundamental understanding of electrochemistry and electrocatalysis.1 Recently, the HER has been used in exploring the fundamental differences between electrocatalysis in acid and base.2 It has long been recognized that the reaction rates of the HOR and the HER are often slower in basic electrolytes than acidic electrolytes, even though the surface intermediate of adsorbed a
Department of Chemical Engineering, Stanford University, Stanford, CA 94305, USA. E-mail: [email protected] b SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, CA 94025, USA † Electronic supplementary information (ESI) available: Detailed information on the synthesis procedure and experimental characterizations, including stability testing, high-resolution XPS spectra, XRD fitting results, and OH adsorption experiments. See DOI: 10.1039/c4cp01385a
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hydrogen (Had) is independent of solution pH.3 Early fundamental studies suggested that, at high pH, OH competes with Had for surface sites, effectively poisoning the electrode and reducing overall rates.4,5 More recent work has focused on the recombination step between Had and OHad during hydrogen oxidation, suggesting that optimal adsorption of OH as well as H is necessary to facilitate recombination of Had and OHad to form H2O.2 In the reverse direction, protons must be removed from water before adsorbing onto the electrode and recombining as hydrogen; this water-splitting step may also be influenced by OHad species on the electrode.6 While the role of adsorbed OH in the alkaline HOR and HER is still unclear, the effect of hydrogen adsorption strength is much less ambiguous. Fig. 1 plots previously published values of measured hydrogen evolution activity under basic conditions versus the calculated free energy of hydrogen surface binding (DGHad). The activity follows a volcano-type relation in which the best catalysts exhibit hydrogen binding free energy closest to the optimum value of zero.7 As in acid, platinum and palladium are the most active elemental catalysts, and nickel is the most active non-precious elemental electrocatalyst. Thus, from a catalystdevelopment point of view, appropriately tuning the hydrogen binding energy is viewed as a necessary, but not sufficient, requirement for an active catalyst in alkaline environments.
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Fig. 1 Plot of hydrogen evolution activity under basic conditions (measured) versus hydrogen adsorption energy (calculated). As in acid,7 the data show a volcano trend. Platinum, nickel, and palladium bind hydrogen with the energy closest to optimum, but still offer room for improvement. All hydrogen binding energies are from ref. 10, except for Ag and Ni values, which were calculated in this work. Alkaline measurements are from ref. 11 (red squares, 30% KOH, 80 1C) and ref. 12 (blue circles, 0.1 M KOH, 25 1C). The dashed lines are to guide the eye. Arrows mark the hydrogen binding energy to surface alloys of (a) Ni on the Ag substrate and (b) Ag on the Ni substrate.9
Recent work on non-noble electrocatalysts for alkaline hydrogen oxidation has also attributed extremely high activity to optimal hydrogen binding.8 Computational screening is a useful tool to make predictions for new materials, which helps in guiding the development of active and stable catalysts. Previous theoretical calculations of DGHad predicted that surface alloys of silver and nickel bind hydrogen with an energy very close to optimum, as indicated by the arrows in Fig. 1.9 The optimal binding energy of NiAg surface alloys suggests that bulk NiAg alloys may be similarly active and motivates our further experimental and theoretical investigations into these catalysts. Experimentally, bimetallic NiAg systems have shown promise for hydrogen generation both electrocatalytically13 and in homogenous catalysis.14 However, synthesis of binary NiAg alloys is challenging because thermodynamic instability causes Ag and Ni to phasesegregate into pure elemental phases at all temperatures.15 This immiscibility, combined with a large lattice mismatch (0.352 vs. 0.408 nm) and the lower surface energy of Ag, renders the Ag–Ni core–shell the most thermodynamically favorable arrangement for nanoparticles.16,17 Circumventing the phase diagram and forming a metastable alloy require kinetically controlled materials synthesis, such as rapid reduction or condensation at low temperature, where Ag and Ni atoms do not have enough kinetic energy to diffuse and segregate. Previously, various thin film and nanoparticle synthesis techniques have been applied for this purpose, including electrodeposition, pulse laser deposition, e-beam PVD, DC sputtering, and gamma-irradiation.18–22 In this work, e-beam PVD was chosen for its ability to access a wide range of alloy compositions. PVD has the additional advantage of producing flat, conformal films on a variety of substrates. Although alkaline electrolysis is an established technology for which
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many materials have been tested as catalysts, many previous studies are developed around high surface area structures, such as RANEYs nickel, and tested as composite electrodes in a fuel cell or electrolyzer.23–26 Although such conditions most accurately represent those during device operation, distinguishing between catalysis, reactant transport, and electron transport effects in a complicated device geometry is extremely difficult. For example, nickel–molybdenum and nickel–molybdenum–cobalt catalysts show the best HER activity of all non-precious metals, but the reason for the improvement over pure nickel has been alternatively attributed to increased microscopic porosity and increased intrinsic catalytic activity.8,27–30 In other intermetallic systems, improved activity in the presence of dopant metals has been attributed to a variety of reasons, including better ohmic contact between particles, better mass transport through catalyst pores, better mass transport through electrode pores, and better catalysis at the particle surface.26,31–33 Physical vapor deposition produces flat, conformal films of uniform composition, avoiding such complications and enabling fundamental studies to provide a greater insight. Additionally, synthesizing thin films of catalysts on a glassy carbon substrate permits alloys to be tested in a rotating-disk configuration that separates catalytic activity from the surface area, ohmic contact, and pore transport.
2 Methods 2.1
Computational details
All theoretical results are based on density functional theory calculations. These are performed using Quantum ESPRESSO,34,35 a plane-wave36 pseudopotential37 code, and the Atomic Simulation Environment.38 We used the RPBE39 exchange–correlation functional and ultrasoft40 pseudopotentials from the PS Library project.41 Kinetic energy cutoffs were 700 eV for wave functions and 6500 eV for charge densities and potentials, and the surface Brillouin zone was sampled using a 6 6 1 Monkhorst–Pack42 k-point mesh. A Gaussian smearing of 0.2 eV was applied to the electron occupation numbers near the Fermi level. Periodically repeated 5-layer slabs were used to model hydrogen adsorption onto three close-packed fcc(111) NiAg alloy surfaces (Ni3Ag, NiAg, and NiAg3) as well as onto the pure Ni and Ag(111) facets. A 2 2 surface unit cell with one H atom per cell was applied. This resulted in a 1/4 monolayer coverage of atomic H on the transition-metal substrates and several structurally non-equivalent adsorption sites, particularly, on the Ni3Ag and NiAg3 alloys. The three alloyed surfaces were alloyed throughout the bulk, and the top two layers of all surfaces and the H adsorbate were relaxed until all residual forces acting on them were less than 0.05 eV Å 1. The bottom three metal layers were fixed in the bulk structure, with lattice constants computed using an equation of state approach. 2.2
Synthesis of NiAg alloys
Catalysts were deposited on 5 mm diameter glassy carbon disks (Hochtemperatur Werkstoffe) and 1.6 cm quartz slides (GM Associates) using a load-locked dual e-beam evaporator equipped
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with separate quartz crystal microbalances for each source (Temescal VES2550). Additional Ni films were separately deposited using an Innotec ES26C e-beam evaporator. The total deposition rate was maintained at approximately 2 Å s 1 for all depositions. To improve substrate–catalyst adhesion, glassy carbon disks were anodized in air at 400 1C for 2 hours before sonication in acetone, isopropanol, and water, followed by mounting on wafers for deposition. An 80 nm Ti adhesion layer was deposited on the quartz substrates, but was found to increase delamination from the carbon substrates during catalyst testing, possibly due to hydrogen embrittlement of the Ti. Therefore, no Ti layer was used on glassy carbon substrates. Ni and Ag were deposited in different ratios by separately adjusting the electron beam power for each source. The quartz-crystal monitors were calibrated for cross-talk between the metal sources at the beginning of each deposition; depending on the speed of equilibration, this calibration period resulted in a 5–30 nm layer of pure Ag underneath the 70–85 nm NiAg alloy catalyst layer. 2.3
Physical and chemical characterizations
After deposition, catalysts were characterized by scanning electron microscopy (SEM) using a FEI XL30 Sirion with a beam voltage of 5.0 kV. X-ray diffraction (XRD) was performed on a PANanalytical X’Pert Pro using Cu Ka radiation (l = 1.541 Å). X-ray photoelectron spectroscopy (XPS) was performed using a PHI Versaprobe Scanning XPS using Al Ka radiation (hn = 1486.6 eV, spot size 200 mm, 451 collection angle). The roughness factor of samples was also measured by atomic force microscopy (AFM), using a Park Systems XE-70 in non-contact mode. 2.4
Electrochemical testing
Electrochemical tests were conducted in 0.1 M KOH (Sigma) using a rotating disk electrode (RDE) set-up (Pine Instruments) and a three-electrode potentiostat (Biologic Instruments) with ohmic resistance compensation. A cell made of corrosion-resistant quartz (Pine) was used to reduce contamination from KOH etching in conventional borosilicate glass.43 All potentials were measured with respect to a mercury/mercuric oxide reference electrode (1.0 M KOH, Koslow Scientific), which was separated from the main cell by a potassium nitrate salt bridge. The reference electrode was calibrated to 0.869 V versus the reversible hydrogen electrode (RHE); all potentials in this work are reported vs. RHE. The counterelectrode was Ni wire (Alfa, 99.99%). Although heat treatment of the glassy carbon disks improved catalyst–substrate adhesion, catalyst delamination still occurred during extended periods of hydrogen bubbling. Longer stability tests were conducted on quartz substrates with better adhesion properties. Further details of sample preparation are given in the ESI.† The competition between the hydrogen reactions and the formation and removal of Ni hydrides and oxides can lead to erroneous activity measurements.44 The testing protocol was developed to isolate the catalytic activity for HER and HOR from these competing reactions. After insertion into the electrolyte, the working electrode was activated with a sweep from open circuit potential to 300 mV vs. RHE at 10 mV s 1. This step reduced surface oxides while permitting minimal time for
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hydrogen absorption. The electrode was then held at 20 mV for seven minutes while purging hydrogen gas (Praxair, 99.999%) through the electrolyte. Linear sweep voltammograms of the hydrogen evolution activity were then recorded between 300 and 20 mV. Next, HOR activity was measured by chronoamperometry steps from +50 to +200 mV. Chronoamperometry instead of cyclic voltammetry was used to separate steady-state hydrogen oxidation from transients due to double-layer charging, oxidative removal of Ni hydrides, and oxidation of Ni catalysts. To demonstrate that the observed anodic current was due to hydrogen oxidation, separate samples were tested using an identical procedure under a nitrogen atmosphere as control experiments.
3 Results and discussion 3.1
Uniform deposition of metastable NiAg alloys
SEM was used to determine the morphology of the catalysts synthesized on glassy carbon substrates (Fig. 2). The images show that all of the films are flat and conformal. The grain size, which is approximately 20 nm for Ni and the Ni-rich alloys, increases to about 40 nm for pure Ag. The XPS spectra, shown in Fig. 3, demonstrate the composition range of the depositions. The integrated areas of the Ni2p region from 850–890 eV and the Ag3d region from 364–378 eV were used to calculate the relative surface compositions, which are plotted versus nominal composition in Fig. 4. Depositions on both quartz and glassy carbon show good agreement with the nominal composition. In contrast to studies in which NiAg systems were synthesized at high temperature or by a wet chemical method, no surface enrichment of Ag was observed, suggesting that atoms were deposited with insufficient thermal energy to separate into the thermodynamically preferred arrangement.14,22 The oxygen peak at 532 eV indicates that the catalysts form a native oxide in air, and that Ni is more oxidized than Ag, as expected; the presence of Ni surface oxides was confirmed by high-resolution XPS.† XRD measurements showed that the e-beam deposition method was able to achieve partial metastable alloying of Ni and Ag. All samples, except for the pure metals, were found to have both Ni-rich and Ag-rich alloy phases. Diffractograms of the catalysts deposited on quartz are shown in Fig. 5. The counts are
Fig. 2 SEM images of catalysts as synthesized on glassy carbon substrates. Films are flat and conformal, with Ag depositing on larger grains than Ni or the NiAg alloys.
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Fig. 3 XPS survey spectra of the vapor-deposited NiAg samples (glassy carbon substrate). The trends in Ni2p and Ag3d peaks show the compositional variation of the depositions.
Fig. 5 XRD diffractograms recorded on quartz substrates. The curves are normalized to the maximum intensity and corrected for the substrate diffraction. Although the metals segregate into two phases, shifts in the Ni(111) and Ag(111) peaks demonstrate that each phase is a metastable alloy. Also shown are reference peaks for Ni (JCPDS 4-0850), Ag (JCPDS 4-1783), and Ti (JCPDS 44-1294).
Fig. 4 Nominal versus measured Ni concentration. Bulk concentrations of the Ni-rich and Ag-rich alloy phases are calculated from Vegard’s Law. Average surface concentrations are measured using XPS. Closed symbols are for quartz substrates, and open symbols are for glassy carbon. The maximum concentration of Ag in the Ni-rich alloy phase is about 13 at%.
corrected for the quartz substrate and shown to be normalized to the maximum intensity. For pure Ni, the Ni(111) reflection is observed at a = 0.353 nm, which is consistent with the literature value. As the Ag deposition rate increases, incorporation of larger Ag atoms into the Ni lattice shifts the peak to smaller values of 2y. Alloying Ag also decreases the crystallinity and crystallite size, as evidenced by the smaller peak size and an increase in FWHM. Similarly, the peak position of the Ag(111) reflection increases from that of pure Ag (a = 0.409 nm), showing incorporation of the smaller Ni atoms in Ag-rich films. The shifting of the Ag(111) peak can only be observed visually for the Ag-rich films in Fig. 5 because the (002) reflection of the Ti adhesion layer overlaps for Ni-rich films with lower intensity Ag(111) reflection.
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To quantitatively determine the extent of alloying between Ni and Ag, Gaussian curves were fit to the Ni(111) and Ag(111) peaks. The peak location and FWHM were used to calculate the alloy lattice constant and grain size for each of the two phases according to Bragg’s law and the Scherrer equation, respectively; calculations and results are shown in the ESI.† Grain sizes were calculated between 7 and 14 nm for alloys, slightly smaller than indicated by the SEM images. Inhomogeneities in composition may explain the slight disparity between the two methods. The relative fractions of Ag and Ni as computed by Vegard’s law are shown in Fig. 4. Fig. 4 indicates that the maximum Ag content introduced into the Ni-rich phase by this method is about 13%, at a bulk ratio of Ni0.25Ag0.75. The Ag-rich phase reaches a maximum solubility of approximately 6% Ni, at a bulk ratio of Ni0.125Ag0.875. At a nominal composition of Ni0.125Ag0.875, the solubility of Ag in the Ni-rich phase decreases. This may be because, at higher concentration, Ag atoms need to diffuse shorter distances in order to segregate into a separate phase. Similar solubility trends in metastable NiAg thin-films have been observed previously.19,20 3.2
Electrochemical performance of alloys
The combined materials characterization results of Fig. 2 through 5 show that the films synthesized by e-beam codeposition are flat, uniform, and composed of a Ni-rich phase and an Ag-rich phase.
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Fig. 6 Electrochemical activity measurements of Ni, Ag, and selected NiAg alloys on glassy carbon substrates in 0.1 M KOH/hydrogen gas at 900 rpm. (a) Cathodic HER sweep at 10 mV s 1. (b) HOR at potential steps from +50 to +200 mV. No oxidation is seen during a control experiment in nitrogen.
Fig. 7 Current density at 300 mV (left axis, blue diamonds) and +150 mV (right axis, red squares) vs. nominal Ni concentration. The best-performing catalysts have nominal compositions between 50 and 85% Ni. Error bars represent two standard deviations, based on measurements of at least three samples per composition.
To relate the chemical composition to hydrogen evolution activity, linear sweep voltammetry was performed in an RDE configuration. Fig. 6 and 7 show the HER and HOR activities of Ni, Ag, and selected NiAg alloys. Fig. 6 shows that, as expected, the hydrogen evolution activity of pure Ag is negligible. Pure Ni displays an activity consistent with the literature,12 while the Ni0.75Ag0.25 alloy shows the highest HER activity for these samples (Fig. 6(a)). The current density at 300 mV overpotential is shown in Fig. 7, demonstrating a broad maximum in HER activity at an alloy composition of Ni0.75Ag0.25. The HOR activity at constant potentials is shown in Fig. 6(b) for Ni, Ni0.75Ag0.25, and Ni0.5Ag0.5. The Ni0.75Ag0.25 alloy shows higher current than
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the pure Ni sample at overpotentials of 150 mV or less, but the difference in current disappears at 200 mV overpotential. This change may be caused by a different HOR mechanism at higher overpotentials, or because oxidation of the catalyst at these potentials changes the catalytic properties of the surface. Similar behavior in other Ni-based systems was ascribed to catalyst oxidation.8,24,27 Noise in the data is caused in part by bubbles of hydrogen gas contacting the electrode. Also shown in Fig. 6(b) are oxidation steps for a separate Ni sample tested in a nitrogensaturated atmosphere. The difference in current shows that the anodic current in Fig. 6(b) is due to HOR and not simply because of catalyst oxidation. While the uncertainty associated with measuring small currents in the presence of vigorous bubbling makes the HOR trend less obvious, the Ni-rich alloys have an activity comparable to, if not better than, the pure Ni films, as shown in Fig. 7. Stability testing also shows a similar behavior between the NiAg alloys and pure Ni.† The plot of the current density at 300 mV vs. nominal Ni concentration in Fig. 7 shows that, for HER, the optimum catalyst concentration is in between 50 and 85% Ni. The peak alloy activity is about 9.4 mA cm 2, approximately twice that of pure Ni on a geometric-area basis. Examination of Fig. 6(a) shows that obtaining a current density of 2 mA cm 2 on elemental Ni requires about 260 mV of overpotential, slightly higher than the overpotential for Ni as plotted in Fig. 1. Consideration of surface roughness may account for slight differences in the measured activity from previously published values. Fig. 7 also shows that the HER activity of Ni0.5Ag0.5 exceeds that of pure Ni by about 60%, but that the HOR activity of this alloy is negligible. This difference suggests that the hydrogen binding energy for the NiAg alloy is stronger than optimum, so that adsorbate–adsorbate repulsions between hydrogen atoms can improve the HER activity but not the HOR.45 It is challenging to separate the intrinsic catalytic activity from the effects of the microscopic electrode surface area for non-noble electrocatalysts. Because of the low oxidation potential of Ni, commonly used methods for measuring the surface area, including underpotential deposition of H, Cu, and Pb, are not applicable to Ni-based systems. The low overpotential for the HER also eliminates the possibility of finding a non-faradaic region of the voltage window in which the double-layer capacitance is measured. In this work, however, the vacuum deposition synthetic approach creates flat, uniform films with negligible porosity. AFM images, shown in Fig. 8, show negligible differences in the surface area among the different films. Therefore, improvements in HER and HOR activity were caused by intrinsic effects on catalysis and not by the increased surface area. Fig. 8 shows that for both samples, grains are only a few nm tall, yielding calculated surface roughness factors of 1.04 and 1.03, respectively, for the Ni0.825Ag0.175 and Ni samples. Measurements of OH adsorption also suggested that the NiAg catalysts have similar or smaller areas of electrochemically active Ni (ECSANi) than films of pure Ni.† 3.3
Comparison with theoretical values
Although the AFM image in Fig. 8 shows negligible surface roughness, determining the intrinsic catalytic activity for the
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Fig. 8 AFM images and line scans of Ni and Ni0.825Ag0.375 on glassy carbon. Differences in line scans appear larger than actual because of the axes aspect ratio.
NiAg alloys is still not straightforward because the active site of the alloys is not known. To develop a qualitative insight into the active site for HER and HOR catalysis, DFT calculations were used to determine the hydrogen binding energy at different sites of Ni, NiAg3, NiAg, Ni3Ag, and Ag. The results of all of the calculated sites are shown in Table 1, and selected binding energies and geometries are shown in Fig. 9. The theoretical hydrogen evolution current density was calculated using the `lason, et al.46 and normalized to the experimentally model of Sku measured current density for pure Ni. Because the model was developed for acidic electrolyte, not basic electrolyte, it can be used only for qualitative interpretation in studying trends. As previously discussed, the effects of hydroxide ions on alkaline HER and HOR are not entirely understood, and a quantitative microkinetic model that incorporates their effects is beyond the scope of this work. Table 1 and Fig. 9 indicate that simple alloying of Ni and Ag can obtain many energetically different adsorption sites, and that several NiAg structures have sites where hydrogen is predicted to bind with energy close to the optimum value of 0 eV. The full 4d-shell of the Ag atoms is expected to shift the projected d-band downwards from the pure-Ni case, leading to
Table 1 Calculated hydrogen binding free energies (eV) on uniformly alloyed NiAg(111) surfaces. Selected images of the calculated surface structures are shown in Fig. 9
Composition
Binding site
Ni
Hollow On-top Hollow (Ni) Hollow (Ag) On-top (Ni) On-top (Ag) Bridge Hollow (Ag) On-top (Ni) Hollow (Ni) On-top (Ni) Hollow (Ag) Hollow On-top
Ni3Ag
NiAg NiAg3 Ag
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DGHad (eV) 0.19 0.38 0.25 0.07 0.34 1.04 0.15 0.16 0.32 0.09 0.26 0.47 0.46 0.99
Fig. 9 Above, combined theoretical-experimental volcano plot for mostactive sites. Markers are simulations, and the horizontal lines are measured currents for: pure Ni, geometric area (grey); Ni0.75Ag0.25, geometric area (green); Ni0.75Ag0.25, area of Ag in the Ni phase (blue). The dotted red line represents the predicted current.46 Below, selected images of the calculated surface structures are shown.
weaker hydrogen binding and more active catalysts. This effect is observed for binding sites which include both Ni and Ag atoms, such as the Ni hollow site on NiAg3 and the Ag hollow site on Ni3Ag. Such binding sites generally have the DGHad values closest to the optimum of 0 eV. Another competing effect is seen at the Ni hollow site of Ni3Ag, which binds more strongly than the pure Ni site. Similarly, the Ag hollow site of NiAg3 binds more weakly than the pure Ag site. This is due to the effect of the alloyed atoms not in the binding site. For example, in the Ni3Ag case, there are Ni atoms in the binding site which interact with an Ag atom. This Ni–Ag bond has lower saturation compared to a Ni–Ni bond, leading to a higher saturation of the Ni–H bond, and thus to stronger hydrogen binding. The interplay of these two effects results in a variety of surface sites with differences in DGHad, some of which are favorable and some of which are unfavorable for the HER. However, due to the exponential dependence of activity on DGHad, we expect the most active sites to dominate the overall alloy activity. Additionally, the results of Table 1 assume a perfectly ordered lattice, and a random distribution of Ag and Ni may change both the geometric frequency and electronic structure of binding sites. Fig. 9 also shows experimental data as extracted from Fig. 7 for Ni0.75Ag0.25. Because the exact surface structure is not known, the experimentally measured current density is represented as a horizontal green line instead of a discrete point. All experimental and theoretical values are normalized to the experimentally measured activity on pure Ni (grey line). The alloys contain both Ag-rich and Ni-rich phases (Fig. 5), and it is not immediately clear in which phase the active site is located. Inspection of Table 1 shows that the most optimal sites for
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hydrogen binding are Ag atoms in a Ni-rich environment, or Ni atoms in an Ag-rich environment. Based on this theoretical insight, the measured current density for Ni0.75Ag0.25 was scaled to the area fraction of the Ag component in the Ni-rich phase. This value is represented by a blue dashed line in Fig. 9. Scaling in this manner would suggest an enhancement by 100 for each of these sites relative to pure Ni. Because this calculation neglects the activity of all other sites, the estimated current is therefore an upper bound of the possible current per active site. Although the Ni atoms in the Ag-rich phase are expected to be similarly active, they make up a much smaller fraction of the surface, and scaling to this area would yield a current density that exceeds the theoretical limit. Other active sites may occur at the Ni-rich and Ag-rich grain boundaries; hydrogen adsorption at the bridge site between Ni and Ag is predicted to be very close to optimum. Tuning the grain size by controlling synthesis parameters such as deposition rate and temperature might elucidate the roles of different sites and suggest strategies for increasing the density of the most active sites. Finally, the effects of hydroxide adsorption cannot be ignored. Because hydrogen evolution requires an additional step of watersplitting, a bifunctional catalyst is required.2 The calculated values for OH adsorption energy differ by 1.8 eV between Ni and Ag.47 This difference may have synergistic effects on the HER independent of hydrogen adsorption. More work is necessary to understand fully the effect of hydroxide adsorption on alkaline HER and HOR catalysis, and the control afforded by e-beam codeposition motivates further use of the NiAg system to probe these effects.
4 Conclusions Guided by theoretical calculations, we identified NiAg alloys as promising electrocatalysts for the HOR and HER in base. To mitigate the thermodynamic phase segregation of Ag and Ni, e-beam codeposition was used to deposit thin films of metal catalysts on glassy carbon and quartz substrates. This technique permits finely-tuned catalyst compositions in a well-defined morphology. Incorporation of up to 13% Ag into the Ni lattice and 6% Ni into the Ag lattice was demonstrated via XRD. Rotating disk voltammetry showed that Ni-rich NiAg samples were more active than pure Ni for the HER, with the most active catalyst, Ni0.75Ag0.25, outperforming pure Ni by approximately a factor of two. Because of the planar electrode geometry (AFM-measured roughness factors of approximately unity), the improvement may be attributed to intrinsic catalytic activity, not to the higher surface area of Ni. The HOR activities of the Ni-rich NiAg catalysts were comparable to pure Ni. The improved HER activity was corroborated by DFT calculations, which demonstrate that alloying Ni and Ag creates a variety of adsorption sites, several with near-optimal hydrogen binding energies. Future work to elucidate the active site(s) in the NiAg system for alkaline HER and HOR catalysis would help in enabling rational catalyst design, which will in turn facilitate the incorporation of advanced catalysts into functional devices.
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Acknowledgements This work was supported by a TomKat Grant for Energy Research at Stanford University. We also thank the Department of Energy for SUNCAT funding. A.J.K. was supported by an Abbott Laboratories Stanford Graduate Fellowship. Part of this work was performed at the Stanford Nano Center (SNC)/ Stanford Nanocharacterization Laboratory (SNL), part of the Stanford Nano Shared Facilities. This work was also supported by the facilities of the UCSB Nanotech, a member of the National Nanofabrication Infrastructure Network partially supported by the National Science Foundation.
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