Trends in Catalysis and Catalyst Cost Effectiveness for N2

Article pubs.acs.org/JPCC

Trends in Catalysis and Catalyst Cost Effectiveness for N2H4 Fuel Cells and Sensors: a Rotating Disk Electrode (RDE) Study David A. Finkelstein,† Régis Imbeault,† Sébastien Garbarino, Lionel Roué, and Daniel Guay* Institut National de la Recherche ScientifiqueÉnergie, Matériaux et Télécommunications (INRS - EMT), 1650, Boul. Lionel-Boulet, Varennes, Québec, Canada J3X 1S2 S Supporting Information *

ABSTRACT: Hydrazine (N2H4) is a promising high-power energy carrier for fuel cells, combining the energy density of methanol (MeOH) with the rapid oxidation kinetics of hydrogen (H2). N2H4 does not require expensive Pt group metals nor Au for low-potential (high voltage) oxidation, offering significantly lower fuel cell materials costs compared to H2, MeOH, ethanol (EtOH), and ammonia (NH3). In our study, we use rotating disk electrode (RDE) voltammetry to explore N2H4 oxidation at a wide variety of catalysts, including first-row transition metals (Co, Ni), coinage metals (Ag, Au) and Pt group metals (Ru, Rh, Pd, Ir, Pt). While several groups have focused on Co, Ni, or CoNi alloys, we find that other metals, including Ag, Ru, and Pd, offer much higher electron recovery and have more stable reactions, and still cost far less than Pt, Au, Rh, or Ir. We analyze our findings in terms of cost vs performance for the metals, developing a guide for the design of N2H4 fuel cell systems and sensors to suit various application spaces. The many metals studied also reveal an important trend for the theoretical understanding of catalysis: the onset and passivation of N2H4 oxidation in nearly every system were directly tied to the appearance or disappearance of specific metal surface states (e.g., hydrides and oxides). Indeed, metals with multiple surface states frequently showed multiple mechanisms for N2H4 oxidation, each with separate values for electron recovery. These observations provide support for the continued development of electrocatalytic theory in which different metal surface states are treated as independent materials with distinct reaction mechanisms.

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conditions,3,6−8 where catalysts less expensive than Pt, such as Pd and Ag, may be used for O2 reduction.9−13 For these reasons, N2H4 fuel cells were of tremendous industrial and governmental interest internationally from the approximate time period of 1962 to 1984. This technology was pursued by the corporations Matsushita14 and Hitachi15 (Japan), Siemens16 (Germany), and the Canadian defense department.17 In the United States, N2H4 fuel cells were extensively developed for the military and space program, and much of the research was contracted to chemical, energy, and military corporations. These included Union Carbide,18,19 AllisChalmers,20 Shell,21 Monsanto,5,22,23 and Lockheed,3,11 among others. The U.S. National Aeronautic and Space Administration (NASA) has explored N2H4 fuel cells to use on-board N2H4 rocket fuel to power satellite electronics both historically24 and as recently as 2010.25 The U.S. Department of Energy (Argonne National Laboratory)6 and the U.S. Naval Research Lab26 both conduced detailed studies. The U.S. Army, however, not only engaged in intensive benchtop research,5,22,23,27 but

INTRODUCTION

N2H4 is an extremely energy dense fuel, making it highly suitable for portable power applications. As a pure liquid, its volumetric energy density is greater than that of MeOH and 2.5 times higher than H2 at 10 000 psi (5.5 vs 4.8 and 2.0 kWh/L, respectively), while its gravimetric energy density is similar to that of MeOH (5.4 vs 6.1 kWh/kg). N2H4 is fully miscible in water, allowing high concentration (32 M) fuel cell electrolyte, and can be alternatively stored as the highly soluble salt dihydrazine sulfate (24.7 M N2H5+, 12.3 M salt). The high solubility and rapid kinetics of N2H4 allow it to obtain current and power densities that are simply unachievable for H2, MeOH,1 EtOH,2 or NH3. Just 0.1 M N2H4 was shown to have 165 times the mass transport limited current of H2.3 Similarly, 5 mM N2H4 at a smooth Pt electrode at 3000 rpm produces 8 times the geometric current density as 5 mM NH3 does at high-surface area, Pt(100) preferentially oriented thin films (RDE data below vs Finkelstein et al.4). The reactivity of N2H4 is so high that its oxidation is catalyzed at inexpensive commodity metals, such as Ni and Co,5,6 a feat unimaginable for more commonly studied fuels. Furthermore, N2H4 produces greater current and is more stable in strongly alkaline © XXXX American Chemical Society

Received: October 16, 2015 Revised: February 8, 2016

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DOI: 10.1021/acs.jpcc.5b10156 J. Phys. Chem. C XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry C

intermediates at the surface.45 That is, N3 and N4 intermediates are not formed during heterogeneous oxidation of N2H4 at Pt, unlike homogeneous oxidations of N2H4.49,50 Decomposition of N2H4 to H2 and N2 (eq 2, X = 0) is desired when N2H4 is used as an H2 storage agent, and the reaction proceeds well at Rh-based catalysts.51 The oxidation mechanisms of N2H4 at the various metals are critically analyzed in the Discussion.

also performed rigorous testing of prototype vehicles and actually deployed N2H4 fuel cells in combat zones. A 3/4 ton M37 military truck powered by N2H4 fuel cells reportedly achieved 100 miles/gallon fuel and outperformed a conventionally fueled vehicle during track tests. Thirty N2H4 fuel cells were deployed in the Vietnam War to provide silent power generation in applications as diverse as fire support bases and drift boats, and military personnel found them easy to use and handle.28 Although N2H4 is attractive as a fuel source, it is acutely toxic in all common forms, including pure N2H4 liquid,29 hydrazine hydrate,30 and hydrazine sulfate.31 This poses a challenge to its widespread use. Daihatsu has developed polymer materials that can reversibly bind N2H4 in the nontoxic forms of hydrazone or hydrazide,32 although this procedure imposes additional materials and energy costs to the system. Despite the health hazards associated with N2H4, fuel cells using it appear to have been safe enough to use on boats and in live-fire situations by soldiers with limited training.28 If N2H4 fuel cells are deemed insufficiently safe for civilian use, they may yet find application in space exploration, where high energy devices are needed to power electronics. Currently, many NASA robotic probes are powered by thermoelectrics utilizing heat from the nuclear decay of plutonium (238Pu). This isotope is produced in nuclear reactors, but the U.S. shut down its production capacity in 1988 and stopped receiving Russian supplies in 2010. NASA is paying $75−90 million over 5 years to restart 238Pu production,33,34 and 93% of the resultant nuclear fuel’s energy will be wasted due to inefficient thermoelectrics.35 With the difficulty and expense of using 238 Pu, N2H4 may offer a cost-effective alternative as a means of powering space probes for both NASA and the European Space Agency, the latter of which does not use nuclear materials at all.34 As mentioned above, NASA has already reopened its studies of N2H4 fuel cells for satellites.25 During the early period of N2H4 fuel cell research and deployment (∼1962−1984), asbestos was the primary material used for alkaline fuel cell membranes,26 and likely posed inherent engineering difficulties. Interest seems to have eventually shifted to proton-exchange membrane (PEM) based H2/O2 fuel cells. However, with the recent development of alkaline membranes with improved conductivity and durability,36−39 industries have returned focus to N2H4 fuel cells, with Daihatsu actively publishing on new electrocatalysts40−42 and fuel cell systems1,32 and Samsung reviewing developments in the field.43 Modern academic interest has also been reviewed as well.44 The low potential oxidation of N2H4 involves up to 4e−, with N2 as the final product eq 1, but side products and catalytic decompositions are known. The complete oxidation to N2 has been observed at Pt-group metals45,46 and Au, while first-row transition metals, such as Ni and Co, decompose significant amounts of N2H4 to H2 eq 2,41 greatly decreasing observed current densities. The heterogeneous oxidation of N2H4 to NH3 has been observed in acidic (Nafion-based) N2H4 fuel cells using Pt catalysts,40 which likely follows a decomposition mechanism that occurs under open-circuit and low pH eq 3.45,47,48 The electrocatalytic oxidation of N2H4 at Pt under alkaline conditions has been observed by electrochemical mass spectrometry to generate N2 as the only gaseous product.45,46 Experiments with isotopically labeled N2H4 at Pt demonstrated that both N atoms in each N2 originate from the same N2H4 molecule, precluding the possibility of reactions between N2H4

N2H4 + 4OH− → N2 + 4H 2O + 4e−

(1)

E0 = −1.21 V vs NHE, − 1.35 V vs Hg/HgO at pH 1452,53 N2H4 + X OH− → N2 + (2 −

1 X )H 2 + X H 2O + X e− 2 (2)

2N2H4 → 2NH3 + N2 + H 2

(3)

With the increased recent interest in N2H4 fuel cells and sensors, and alkaline fuel cells in general, we aim to provide a comparison of the cost and performance of a wide array of N2H4 electrocatalysts, identifying application spaces for the various metals. The oxidation mechanisms at each metal are reported primarily in terms of n, the number of e− recovered. We have also paid close attention to the potentials at which various mechanisms begin and end, noting that nearly every mechanism is associated with the presence of various metal surface groups. Such information is useful to providing a deeper understanding of the catalytic processes occurring, which in turn aids the design of alloys with improved or tailored reactions. The data will underscore an important point: native metal surface chemistry is intimately involved with N2H4’s oxidation mechanisms. This finding is unexpected from theoretical approaches to catalysis that (1) assume a given electrochemical species’ reaction is uniform across multiple catalysts and (2) consider a single metal−adsorbate electronic interaction as the primary predictor of the onset of a multi-e− catalytic reaction.54−56 However, our results follow well from combined experimental-theoretical approaches in which variable surface states, along with their corresponding adsorbates, are given primacy in catalytic understanding.57−62 In light of our findings, we propose that distinct metal surface states receive theoretical treatment as different materials altogether, each with the possibility of a unique reaction mechanism. We further suggest that the characterization of such surface states deserves more focus, as does the theoretical prediction of their appearance.

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MATERIALS AND METHODS Catalysts. All RDEs were 5.0 mm in diameter and mirrorpolished. Commercial Pt and Au RDEs (Pine Instruments, 99.99% purity) were used. The Rh RDE was machined from a commercial rod (Goodfellow, 99.9% purity). The remaining RDE electrodes consisted of high purity vapor-deposited metals (Kurt J. Lesker, various individual purities). Solutions Preparation. All chemicals used were reagent grade and included NaOH (Fisher Scientific), N2H4·H2SO4 (Alpha Aesar) and H2SO4 (OmniTrace, Fisher Scientific). Solutions were made with deionized water (Millipore, > 18 MΩ·cm) and were freshly prepared immediately prior to experimentation to avoid significant hydrolysis of N2H4 to H2. Before experiments involving characterization of metal surface chemistry, electrolyte solutions were purged with high purity argon for 30 min, and a flow of argon over the solution was maintained continuously during electrochemical analyses. B


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The Journal of Physical Chemistry C Table 1. Summary of N2H4 Catalysis Properties Determined for the Metals Studieda 1 M NaOH metal

n

EOC (V vs Hg/HgO)

E1/2 at 500 rpm (V vs Hg/HgO)

0.5 M H2SO4 E cleaning (V vs Hg/HgO)

Pt Pd Rh Ru Ir Au Ag Ni Co

4 4 4 3, 2.5, 1.5 3 4 4 2.5, 1.5, 3 1.7

−0.825 −0.755 −0.940 −0.960 −0.840 −0.400 −0.420 −0.910 −0.982

−0.740 −0.690 −0.870 −0.920 −0.810, −0.640 −0.400 −0.215 −0.770 −0.770

−0.30 −0.10 ≤ E ≤ +0.60 −0.30 −0.20 −0.30 −0.50 −0.60 −0.70 −0.70

a These include n, the number of e− involved in oxidation; EOC, the open-circuit potential of 5 mM N2H4 in 1 M NaOH; and E1/2, the half-wave potential of 5 mM N2H4 in 1 M NaOH at 500 rpm. The cleaning potential(s) used in 0.5 M H2SO4 for each metal are also shown; see text for details on the cleaning procedure.

coefficient of the reduced species, ω is the rotation rate, ν is the solution’s kinematic viscosity, and CR* is the concentration of the reduced species. A value of 1.40 × 10−5 cm2·s−1 was used for DR.7 To eliminate error from background currents and verify transport-limited behavior, values for n were determined from the slopes of linear regions of Levich plots (il vs ω1/2).

Solutions of N2H4 were not Ar-purged to avoid accelerating N2H4 hydrolysis via the removal of H2 in solution. Electrochemical Setup. Electrochemical experiments were performed in a conventional three-electrode cell with a Hg/ HgO reference electrode (+0.140 V vs NHE) and a Pt counter electrode. The reference and counter electrode chambers were separated from the working electrode chamber by a Luggin capillary and a medium-porosity glass frit, respectively. Electrochemical measurements were performed using a BioLogic VMP3 multichannel potentiostat/galvanostat and a Pine analytical rotor (RDE, MSRX Speed Control). All experiments were carried out under ambient conditions (20 °C, 1 atm). Electrode Cleaning and Electrochemical Analyses. The surface chemistry of each metal catalyst was thoroughly investigated using stationary CV in 1 M NaOH. A sweep rate of 100 mV s−1 was generally used in order to observe the various electrochemical features with a good resolution. RDEs were polished to a mirror finish using a graded series of alumina pastes (1, 0.3, and 0.05 μm) on polishing cloth (Buehler) and sonicated for 5 min. RDEs were then electrochemically pretreated in 0.5 M H2SO4 solution via application of a static cathodic potential in the H2 evolution region for 5 min. The specific applied potential for each metal is listed in Table 1. This method was used to remove accumulated surface oxides, especially for Co, Ni, Ru, Rh, and Ir (see Discussion). Electrochemical pretreatment was not used for Pd due to its tendency to absorb H2, and electrochemical cycling was substituted instead. RDE voltammetry was performed in solutions of 5 mM N2H4 in 1 M NaOH, at 20 mV s−1, and EOC was recorded prior to the start of each experiment. An extensive set of rotation rates were used to obtain detailed Levich plots and included 50, 65, 80, 100, 150, 250, 350, 500, 750, 1000, 1500, 2000, and 3000 rpm. Select rotation rates of 50, 100, 250, 500, 750, 1500, and 3000 rpm are shown in figures for clarity. Current densities have been calculated using the geometrical area of the approximately flat RDEs. Electrochemical Equations. The number of electrons recovered for N2H4 oxidation at the various metals was determined using the Levich equation:63 il = 0.62nFADR 2/3ω1/2ν−1/6C R*

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DISCUSSION Many catalysts have been used in the various N2H4 application spaces mentioned above. We have divided our catalysts into several groups for discussion based on similarity of electrochemical surface chemistry, which, as we will demonstrate, often translates into similarity of elecrocatalysis as well. We begin with Pt-group metals (PGMs: Ru, Rh, Pd, Ir, Pt; Os was not considered for this study), which are commonly employed electrocatalysts that support surface hydrides, hydroxides, and oxides (only alkaline conditions are discussed herein). Pd and Pt (group 10) form hydroxides and oxides at high potential with a low number of monolayers. They will be considered separately from Ru, Rh, and Ir (groups 8 and 9), which form hydroxides at low potential and oxides at high potential, the latter of which can develop into thick multilayers upon repeated cycling. The coinage metals (CMs) Ag and Au (group 11) do not support hydrides and form hydroxides and oxides at higher potentials than PGMs. As will be discussed, they also support high-energy oxides at low potential, which have frequently been implicated in catalysis.64,65 The first-row transition metals Co and Ni (period 4) form stable hydroxides at very low potential and generate various oxides at higher potential, many of which result in significant etching (metal dissolution and subsequent roughening). Many of these surface states have high hysteresis of formation and removal, are altogether irreversible, or have transitions between hydrated and dehydrated states. While Co and Ni are generally less stable than the PGMs or CMs, they have significantly lower costs, and so are considered for economical technologies.41,42,66 Before experimental data was taken, each metal was subjected to a low-potential, chronoamperometric pretreatment in 0.5 M H2SO4 for 5 min to remove oxides that quickly form upon exposure to air. This was especially important for metals such as Co, Ni, Ru, Rh, and Ir, each of which may generate thick oxides that are removed slowly at low potentials. The potential applied to each metal is listed in Table 1. Due to Pd’s tendency to generate and absorb H2 at low potential, it was subjected to more conventional cycling prior to experimentation.

(4)

where il is the limiting current at a given potential, n is the number of electrons involved in the reaction, F is Faraday’s constant, A is the electrode surface area, DR is the diffusion C


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The Journal of Physical Chemistry C Table 1 also presents a summary of the analytical findings for each metal, including EOC (open-circuit potential) and n (number of e− recovered during N2H4 oxidation) determined from Levich analysis of RDE linear sweep voltammograms (LSVs). Many metals showed multiple oxidation mechanisms for N2H4 and so several values of n are listed. All metals produced gas during oxidation, but without differential electrochemical mass spectrometry (DEMS), it is not possible to differentiate between the gas-producing mechanisms in eq 1−3. As expected, EOC was close to the lowest onset potential for N2H4 oxidation for every metal (EOC = Eonset). Also, each metal participated in a catalytic decomposition of N2H4 when left at EOC. This was evidenced by the accumulation of bubbles at the electrode surface, presumably consisting of N2 and H2, as NH3 is fully miscible in water. Part I: Group 10 Pt Group Metals. Pt. A. Pt Surface Chemistry. A typical CV of Pt in 1 M NaOH is shown in Figure 1, and the anodic sweep shows four major regions of surface

Figure 2. RDE LSVs of 5 mM N2H4 at Pt in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at 0.20 V in the anodic sweep.

Figure 1. CV of Pt surface chemistry in 1 M NaOH, 100 mV/s. Features in the anodic (A1−A4) and cathodic (C1, C2) sweeps are described in the text.

phenomena: (A1) two Pt hydride oxidations from −0.80 to −0.42 V; (A2) weak physisorbtion of OH− from −0.42 to −0.17 V; (A3) oxidation of Pt to PtOH, with strongly chemisorbed OH−, from −0.17 to −0.03 V;67 and (A4) oxidation of PtOH to PtO2 from −0.03 V to +0.70 V.67,68 At E > + 0.60 V, O2 evolution becomes superimposed on (A4) PtO2 formation, manifesting exponentially increasing current. The cathodic sweep shows high hysteresis for (C1) PtO2 reduction. While some Pt oxides are reduced starting immediately negative of +0.7 V, most Pt oxide is reduced from +0.17 and −0.46 V. The (C2) peaks for Pt hydride formation are symmetric with those for hydride oxidation. Hydride symmetry and oxide hysteresis is a common feature for all of the metals studied. B. N2H4 Electrocatalysis at Pt. The RDE LSVs of 5 mM N2H4 at Pt (Figure 2, parts A and B) show a single oxidation mechanism for N2H4 with n = 4e− using the linear portion of the Levich plot (Figure 2A, inset). The reaction begins at low potential (−0.83 V, at EOC), translating into high voltage in fuel cell applications. There is a clear 0.10 V gap from where H2 evolution effectively ceases at −0.90 V and where N2H4 oxidation begins (Figure 3A). This Eonset seems to coincide with the (A1) oxidation of Pt hydrides beginning at −0.80 V (Figure 3B).

Figure 3. Comparison of N2H4 RDE CVs at Pt, 500 rpm, over various potential ranges (A, 20 mV/s) and Pt surface chemistry (B, 100 mV/ s).

Thus the low-potential Pt hydride appears to block the reaction. Generally, this peculiar behavior shifts EOC and E1/2 (Table 1) more positive for N2H4 at Pt than at most other PGMs, decreasing expected fuel cell voltage. D


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−0.05 V, while (C1) PdII reduction is shifted to −0.10 to −0.46 V (Figure 4B). It has been reported that PdII oxide formed from PdIV oxide is more compact than when formed from Pd directly,73 and its different phase likely accounts for the shift in potential of oxide reduction. The formation of Pd hydrides on bulk Pd samples is indistinct because they become absorbed into interstitial vacancies. The (C2) formation of Pd hydride starts in the cathodic sweep at −0.65 V, manifesting a broad peak superimposed on bulk H2 evolution (Figures 4.A and 4B). In the anodic sweep, the (A1) and (A2) oxidations of Pd hydrides from −0.62 to −0.17 V are broad as current is generated slowly while interstitial hydrides diffuse to the surface for oxidation. However, when bulk Pd is not accessible to Pd hydrides, the hydrides are trapped at the surface and become sharp and distinct in CV. This occurs for (a) thin layers of Pd deposited on other materials,67 (b) Pd nanoparticles,77 and (c) highly oxidized/etched Pd with buried, impenetrable Pd oxide layers.78 B. N2H4 Electrocatalysis at Pd. Similar to Pt, N2H4 oxidation at Pd begins at very low potential (Eonset = −0.77 V) with n = 4, making Pd an attractive fuel cell catalyst with high voltage and complete e− recovery from N2H4 (Figures 5, parts A and B,

It has been proposed that residual Pt oxides, which persist even into the Pt hydride region, are responsible for N2H4’s lowpotential oxidation.69,70 Since both Pt hydride oxidation and the activation of the residual oxide are proposed to occur at the same potential, which of the two is responsible for N2H4’s catalytic properties cannot be readily distinguished. Furthermore, the residual oxide may govern the potential of Pt hydride oxidation, so all three processes may be interrelated. Most of the high-potential surface phenomena at Pt do not appear to inhibit N2H4 oxidation, including the physisorption of OH− or the (A3) and (A4) formations of PtOH and PtO2, respectively. Indeed, if the RDE CV is reversed at +0.2 V, the cathodic sweep shows the same current as the anodic sweep (Figure 3.A, blue line). At potentials >+0.2 V, current is gradually diminished. Such potentials support higher coverages of Pt oxide, multilayer Pt oxides, and possible competitive OH− adsorption.70 Current is gradually restored in the cathodic sweep starting at +0.65 V, as Pt oxides are slowly removed. Though the gradual current restoration in the cathodic sweep begins before the large (C1) peak for PtO2 reduction, N2H4 oxidation is blocked only at high coverages of Pt oxide, so it sensible that removing even small amounts of oxide is adequate to restore N2H4 oxidation. Pd. A. Pd Surface Chemistry. A CV of Pd in 1 M NaOH in a limited potential window is shown in Figure 4.A, and a CV

Figure 4. CVs of Pd surface chemistry in 1 M NaOH, 100 mV/s, over low potential (A) and high potential (B) ranges. Features in the anodic (A1−A4) and cathodic (C1−C3) sweeps are described in the text.

across a broader potential range is shown in Figure 4B to reveal high-potential features. Phenomena in the anodic sweep include (A1) oxidation of β-Pd hydride from −0.62 to −0.38 V, (A2) oxidation of α-Pd hydride from −0.38 to −0.17 V, (A3) PdII oxide formation from −0.17 to +0.15 V, and (A4) PdIV oxide formation from +0.15 to +0.63 V, followed by O2 evolution at higher potentials.67,68,71−76 When only PdII oxides are generated, they are (C1) reduced in the cathodic sweep starting at +0.1 V, with most reduction occurring between +0.05 and −0.38 V (Figure 4A). The (C2) formation of Pd hydrides, which are subsequently absorbed, is evident below −0.65 V. When both (A3) PdII oxide and (A4) PdIV oxide are formed, (C3) PdIV is reduced from +0.60 to

Figure 5. RDE LSVs of 5 mM N2H4 at Pd in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at −0.40 V in the anodic sweep.

with inset Levich plot). The Eonset is more negative than the (C2) lower potential limit of Pd hydride formation (−0.65 V, Figure 6), so current below −0.65 V represents a superposition of N2H4 oxidation and H+ reduction. This results in a sharp increase of current in this region of the anodic sweep that does not appear to follow Butler−Volmer kinetics63 because the early exponential increase is obscured by H+ reduction. E


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Part II: Group 8 and 9 Pt Group Metals. Rh. A. Rh Surface Chemistry. Window-opening CVs of Rh in 1 M NaOH are shown in Figure 7, parts A and B, and reveal dynamic

Figure 6. Comparison of N2H4 RDE CVs at Pd, 500 rpm, over various potential ranges (A, 20 mV/s) and Pd surface chemistry (B, 100 mV/ s).

As significant amounts of (A4) PdIV oxide are formed in the anodic sweep between +0.40 and +0.70 V, N2H4 oxidation current decreases, more severely for higher rotation rates (Figures 5.A, 6.A, and 6.B). When (A4) PdIV oxide formation is avoided by sweeping below 0 V, N2H4 oxidation readily proceeds in the cathodic sweep, which is not the case when forming PtIV oxide in sweeps to +0.70 V (Figure 6). While at first it appears that (A4) PdIV oxide is still capable of decreased yet effective oxidation of N2H4 between +0.40 and +0.70 V (Figure 6), chronoamperometry at +0.70 V shows a rapid decrease in current, indicating that high coverages of PdIV oxide will quickly lead to a complete cessation of the reaction (Figure S1). The persistence of current at high potential in the RDE LSVs suggests that PdIV oxide grows in slowly. This slow growth of (A4) PdIV oxide leads to peculiar cathodic sweeps, in which the current is still decreasing from +0.7 to −0.1 V, where (C3) PdIV oxide is reduced to the compact PdII oxide (Figures 5B and 6A). The decrease slows between +0.7 and +0.4 V, where mixed PdII/PdIV oxides are present, suggesting they may inactivate less rapidly. By−0.05 V, however, it is clear that full coverage of compact PdII oxide completely inactivates the surface for N2H4 oxidation. As significant amounts of (C1) compact PdII oxide are quickly removed in a sharp reduction peak in the cathodic sweep from −0.05 to −0.25 V, exposing Pd metal, current for N2H4 oxidation is rapidly restored (Figures 6, parts A and B). Prior to oxide removal, the N2H4 surface concentration increases to nearly the bulk concentration between 0 to −0.1 V, since little to no N2H4 oxidation occurs in this potential region. Thus when N2H4 oxidation is restored, the current waveform appears similar to a potential step, generating peaks in the cathodic RDE LSVs between −0.15 and −0.25 V (Figure 5B). The small dips in current seen for lower rotation rates between −0.25 and −0.33 V is simply the (C1) compact PdII oxide reduction peak superimposed on the N2H4 LSV.

Figure 7. Window-opening CVs of Rh surface chemistry in 1 M NaOH, 100 mV/s, over low potential (A) and high potential (B) ranges. Features in the anodic (A1−A4) and cathodic (C1−C3) sweeps are described in the text.

surface chemistry events. The anodic sweep of the CV to +0.02 V (Figure 7A) shows (A1) two overlapping Rh hydride oxidations from −0.84 to −0.51 V; (A2) OH− physisorption from −0.51 to −0.35 V; and (A3) oxidation of Rh to RhIII oxide from −0.35 to +0.02 V.67,68,71 The cathodic sweep to +0.02 V shows (C1) significant RhIII oxide reduction from −0.1 to −0.6 V and (C2) Rh hydride formation from −0.6 to −0.85 V. When the potential window is increased to +0.63 V (Figure 7B), the anodic sweep additionally shows (A4) oxidation of RhIII oxide to RhIV oxide67,68,71 from +0.02 to +0.54 V, followed by O2 evolution. This causes not only (C3) reduction of RhIV oxide from +0.58 to −0.17 V, but also shifts the reduction of RhIII oxide (C1) to lower potentials, from −0.17 to −0.7 V, with the RhIII oxide reduction peak shifting by over 0.25 V from −0.39 to −0.65 V. The shift in the RhIII oxide peak to lower potentials occurs gradually as the upper potential limit is increased (Figure 7B, gray lines). F


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The Journal of Physical Chemistry C There is also a subtle shift in the (A3) formation of RhIII oxide when scanning to higher potentials. In Figure 7A, the first scan to −0.2 V (dotted gray line) shows that nearly negligible amounts of RhIII oxide have formed, with minor (if any) oxide reduction visible in the cathodic sweep. However, once the CVs have been swept to 0 V or higher, RhIII oxide formation is clearly evident beginning at −0.35 V. As will be discussed below, these dynamic events have a profound influence on N2H4 oxidation at Rh. Oxidizing Rh at high potential can lead to thick layers of (A4) RhIV oxide. The accumulation of oxide layers during repeated cycling at high potential is characteristic of the Group 8 and 9 Pt-group metals. Thick oxide layers are generally unfavorable for catalysis, since it is often difficult or impossible to reconvert all of the layers back to the original metal. B. Electrocatalysis of N2H4 at Rh. The RDE LSVs of N2H4 at Rh show a single oxidation mechanism with n = 4 (Figure 8,

Figure 9. Comparison of N2H4 RDE CVs at Rh, 500 rpm, over various potential ranges (A, 20 mV/s) and corresponding Rh surface chemistry (B, 100 mV/s).

RhIII oxide does not have activity for N2H4 oxidation, strictly limiting the potential range in which Rh catalysts can be effectively used. In Figure 9A, avoiding high coverages of (A3) RhIII oxide in the anodic sweep by scanning only to −0.20 V (Figure 7A, dashed gray line) allows the oxidative current to continue unhindered in the cathodic sweep. When scanning to +0.20 V, N2H4 oxidation becomes inhibited as higher coverages of RhIII are formed. The current is quickly restored as (C1) RhIII removal begins from +0.20 to −0.30 V (Figure 7B and 9B, dashed gray line). While the reduction peak for RhIII oxide does not occur until −0.46 V, complete RhIII oxide removal is unnecessary to restore current, since only high coverages inhibited it. Higher coverages of (A4) RhIV oxide may have limited activity for N2H4 oxidation. When such high coverages are produced by scans to +0.60 V, oxidative current in excess of the RhIV oxide peak and O2 evolution is evident between +0.36 to +0.60 V (Figure 9A). This current diminishes in the cathodic sweep as (C3) the RhIV oxide is converted back to RhIII oxide from +0.60 to +0.30 V (Figure 9B). Since the formation of RhIV oxide shifts the (C1) RhIII oxide reduction peak to lower potentials (Figure 7B), N2H4 oxidation is not restored until potentials of −0.20 V and lower (Figure 9). Since Rh has a limited potential window of utility, it must be alloyed with other metals with high-potential activity when employed in the demanding, variable-voltage environment of a fuel cell. Ru. A. Ru Surface Chemistry. Window-opening CVs of Ru in 1 M NaOH are shown in Figure 10 and closely resemble those of Rh in Figure 7. The anodic sweep to 0 V (Figure 10A) shows (A1) two overlapping Ru hydride oxidations from −0.88 to −0.56 V; (A2) Ru(OH)2 or RuO formation from −0.56 to −0.30 V;67,68 and (A3) RuIII oxide formation beginning at −0.30 V. The corresponding cathodic sweep shows (C1) RuIII oxide reduction from −0.15 to −0.55 V; (C2) Ru(OH)2/RuO reduction from −0.55 to −0.70 V; and (C3) Ru hydride formation from −0.70 to −0.90 V.67,68,71,80

Figure 8. RDE LSVs of 5 mM N2H4 at Rh in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at −0.20 V in the anodic sweep.

parts A and B, with inset Levich plot) with the oxidation beginning at low potential (−0.94 V). Rh has the lowest EOC (highest voltage) of any metal with complete e− recovery (Table 1), making it the highest power catalyst in our study. These attractive characteristics likely led to the utilization of RhPtPd catalysts in industrially produced N2H4 fuel cells.79 The onset of N2H4 oxidation at Rh is very sharp and does not appear to follow Butler−Volmer kinetics. As was the case for Pd, this is because the oxidation of N2H4 at Rh begins at −0.94 V and overlaps with H+ reduction, which begins at −0.86 V (Figure 9). This creates a mixed region of current control at low potentials that obscures the region of kinetically controlled current. G


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formation reduction peak at −0.86 V, which is still visible. While (C5) was not evident in all of our experimental runs, we always observed extra current in the region of (C3), indicating that another, overlapping process occurs after exposure to high potential. B. Electrocatalysis of N2H4 at Ru. Since high potentials change the behavior of Ru’s low potential oxides, electrocatalytic oxidation of N2H4 at Ru is highly dependent on the electrode’s history of potential exposure. When RuIII oxide formation is avoided by sweeping to −0.30 V, N2H4 oxidation involves n = 3 (Figure 11), and Yamada40 detected the presence

Figure 11. RDE CVs of 5 mM N2H4 at Ru in 1 M NaOH, 20 mV/s, over a low potential range.

of NH3. This suggests that N2H4 oxidation may follow a mix of eq 1 and eq 3. This has major repercussions for N2H4 fuel cells: since few metals besides Pt oxidize NH3 effectively,81 Ru would require co-utilization with Pt to serve as an efficient catalyst. This would in turn add to the complexity of a N2H4 fuel cell, since the oxidation of NH3 at Pt results in rapid poison accumulation and requires potentiostatic treatment for removal.4 Both RuIII and RuIV oxides have significant impacts on N2H4 oxidation, as seen in the RDE plots in Figure 12, parts A and B. As discussed above, sweeps to +0.40 V form (A4) RuIV oxide and permanently alter the surface. In the anodic sweeps to −0.30 V, current drops only slightly (n = 2.5) as (A1) Ru hydrides are oxidized from −0.88 to −0.56 V in the anodic sweep (Figure 11), whereas sweeps to +0.40 V show a significant current decrease (n = 2) upon hydride oxidation (Figure 12A). Fully avoiding (A3) RuIII oxide by sweeping to −0.30 V allows N2H4 oxidation current to proceed unhindered in the cathodic sweep (Figure 13A, purple line). Anodic sweeps to +0.40 V show that (A3) RuIII oxide formation from −0.65 to −0.05 V is correlated with a significant drop in current, especially at higher oxide coverages (Figure 13A, green line, vs Figure 13B), indicating that RuIII oxide is inactive for N2H4 electrocatalysis. Current from N2H4 oxidation continues to decrease from −0.05 to +0.10 V (Figure 13.A, green line), likely at very high (A3) RuIII oxide coverages that overlap with the beginning of (A4) RuIV oxide formation. However, oxidation current in

Figure 10. Window-opening CVs of Ru surface chemistry in 1 M NaOH, 100 mV/s, over low potential (A) and high potential (B) ranges. Features in the anodic (A1−A4) and cathodic (C1−C5) sweeps are described in the text.

As was observed for Rh, generating high-potential oxides at Ru changes the nature of oxides formed at lower potentials. The (A4) oxidation of RuIII oxide to RuIV oxide occurs from −0.05 to +0.25 V in anodic sweeps to +0.40 V (Figure 10.B), overlapping with O2 evolution at high potentials, while (C4) reduction of RuIV oxide occurs from +0.24 to −0.22 V in the cathodic sweep. The higher potential range greatly increases the current, and likely coverage, associated with (A3) RuIII oxide formation, which now appears to occur between −0.65 and −0.05 V. The (A2) Ru(OH)2/RuO becomes obscured on the larger (A3) feature, but (C1) and (C2) are still clearly distinct, with (C1) RuIII oxide reduction now occurring from −0.40 to −0.55 V. The(C1) and (C2) reductions in sweeps to +0.40 V have a larger background current than in the CV to 0 V (Figure 10A), possibly reflecting higher capacitive current from the formation of irreversible oxides that occurs during the oxidation from RuIII to RuIV oxide. The high-potential sweeps to +0.40 V also seem to generate a (C5) species that is sharply reduced from −0.66 to −0.83 V (Figure 10B). This is likely an oxide generated at high potential with significant hysteresis, persisting until very low potential, as its peak at −0.81 V is clearly distinct from the (C3) Ru hydride H


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data shows that RuIV oxide performs N2H4 oxidation to some extent, although the value of n ≤ 1.5 from Levich analysis of Figure 12, parts A and B, is very low for consideration as either a fuel cell or sensor catalyst. Current in the cathodic sweep in Figure 13A is finally restored from −0.50 to −0.83 V. In this region, (C1) RuIII oxide and (C2) Ru(OH)2/RuO are removed, although N2H4 current increases most significantly during the (C5) reduction of the high-potential oxide from −0.66 to −0.83 V. Thus both Rh and Ru show that oxides formed at high potentials require very low potentials for removal. This underscores the importance of avoiding such high potentials to preserve lowpotential (high voltage) current with high e− recovery when using these catalysts in fuel cells or sensors. Ir. A. Ir Surface Chemistry. Window-opening CVs of Ir in 1 M NaOH are shown in Figure 14 and are similar to CVs of Rh

Figure 12. RDE LSVs of 5 mM N2H4 at Ru in 1 M NaOH, 20 mV/s, over a high potential range. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at −0.835 V in the anodic sweep.

Figure 13. Comparison of N2H4 RDE CVs at Ru, 500 rpm, over various potential ranges (A, 20 mV/s) and corresponding Ru surface chemistry (B, 100 mV/s).

Figure 13A rises as RuIV oxide coverage increases from +0.10 to +0.30 V. The higher oxidation current at RuIV oxide also appears in the cathodic sweep, slowly decreasing as the majority of RuIV oxide is removed from +0.25 to −0.10 V (C4). This

Figure 14. Window-opening CVs of Ir surface chemistry in 1 M NaOH, 100 mV/s, over low potential (A), medium potential (B), and high potential (C) ranges. Features in the anodic (A1−A5) and cathodic (C1−C5) sweeps are described in the text. I


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The Journal of Physical Chemistry C and Ru in Figures 7 and 10. The anodic sweep of the CV to −0.2 V (Figure 14A) shows (A1) the oxidation of three Ir hydrides from −0.86 to −0.44 V and (A2) the oxidation of Ir to IrII hydroxide, Ir(OH)2, or IrII oxide, IrO, from −0.35 to −0.21 V.82,83 The corresponding cathodic sweep shows (C1) the reduction of Ir(OH)2/IrO from −0.44 to −0.56 and (C2) the formation of Ir hydrides from −0.56 to −0.84 V. Expanding the potential window to +0.13 V (Figure 14B) dramatically increases the size of (A2) the peak for Ir(OH)2/ IrO formation, now from −0.35 to −0.12 V. It also introduces (A3) the oxidation of Ir(OH)2/IrO to IrIII oxide, Ir2O3, from −0.12 to +0.13 V in the anodic sweep, with (C3) its corresponding reduction back to Ir(OH)2/IrO from −0.05 V to −0.44 V in the cathodic sweep.68,71,82,83 The formation of Ir2O3 also results in the appearance of several peculiar features: (A4) an increase in current between −0.52 and −0.35 V, which reflects the formation of a subtle oxide; and (C4) an oxide reduction in the hydride formation region and beyond, from −0.56 to −0.87 V (Figure 14B). The two likely correspond to oxidation and reduction of the same species. The (C4) reduction of the oxide greatly increases cathodic current throughout the hydride formation region, resulting in apparent asymmetric Ir hydride peaks, while the oxide reduction peak itself is evident from −0.75 to −0.87 V. Since sweeps to +0.13 V showed a large increase in (A2) Ir(OH)2/IrO formation without a corresponding increase in (C1) (part B vs part A of Figure 14), the (C4) reduction may also reflect reduction of Ir(OH)2/IrO. The Ir(OH)2/IrO species may dehydrate at high potential and require a larger overpotential to reduce, as was seen previously for Rh and Ru oxides. Repeated cycling of Ir up to +0.53 V (Figure 14C) results in an anodic sweep that shows a greatly increased and expanded (A2) Ir(OH)2/IrO formation peak from −0.38 to −0.10 V; reveals that (A3) Ir2O3 formation occurs over the broad range of −0.12 to +0.22 V; and introduces (A5) the oxidation of Ir2O3 to IrIV oxide, IrO2, from +0.22 to +0.40 V, with its corresponding (C5) reduction from +0.40 to +0.03 V in the cathodic sweep.82 The evolution of O2 is evident at potentials greater than +0.40 V. B. Electrocatalysis of N2H4 at Ir. Similarly to Ru, Ir undergoes irreversible changes upon exposure to high potentials, and correspondingly, N2H4 electrocatalytic oxidation at Ir is dependent upon the electrode’s history of potential exposure. RDE of N2H4 at Ir to −0.20 V, which avoids the formation of (A3) Ir2O3 but not (A2) Ir(OH)2/IrO, shows n = 3, with current fully maintained in cathodic sweeps (Figure 15, parts A and B). However, the current in the anodic sweeps is clearly delayed, beginning only at −0.86 V, where (A1) Ir hydride oxidation begins. This suggests that, as in the case of Pt, Ir binds more strongly to hydrides than to N2H4, and the hydrides must be oxidized off the surface to allow room for N2H4 to bind. The RDE voltammetry to −0.20 V (Figure 15) clearly evidences the accumulation of poisons. As the rotation rates are increased, corresponding to longer periods of experimentation and greater passage of current, the values for E1/2 shift to higher potentials (lower voltage). The limiting currents at higher rotation rates are also smaller than expected from the Levich equation (Figure 15A, inset). Thus, while (A2) Ir(OH)2/IrO does not cause hysteresis in the voltammetry, n is still low and poisons accumulate, severely limiting the effectiveness of Ir as a fuel cell catalyst.

Figure 15. RDE LSVs of 5 mM N2H4 at Ir in 1 M NaOH, 20 mV/s, over a low potential range. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at −0.20 V in the anodic sweep.

RDE of N2H4 at Ir to +0.53 V (Figure 16, parts A and B) shows a number of profound changes in oxidative activity in the anodic sweep. First, N2H4 oxidation in the anodic sweeps begins at −0.90 V, immediately after H2 evolution ceases (Figure 17, parts A and B). The catalysis at these very low potentials may occur at irreversible Ir oxides that accumulated at high potential. It is also possible that the formation of oxides restructured the surface, providing high-energy Ir crystal faces that are more active for catalysis. The next clear change is that oxidative current in the anodic sweeps decreases from −0.60 to −0.40 V (Figure 16A), corresponding closely to the range in which the (A4) oxide is formed (Figure 17), indicating that this subtle oxide poisons the reaction. As (A2) Ir(OH)2/IrO is formed from −0.40 to −0.28 V, current increases, though not to its original level, suggesting that Ir(OH)2/IrO has some degree of activity for N2H4 oxidation. At potentials greater than −0.28 V, past the (A2) Ir(OH)2/ IrO peak, current declines again (Figures 16A and 17). It is possible that N2H4 shows activity at a partially Ir(OH)2/IrO covered surface, but is inhibited by a full monolayer. Thus maximum N2H4 oxidative current should occur at 50% Ir(OH)2/IrO coverage, corresponding to the potential of Ir(OH)2/IrO’s peak current. Current continues to decrease at higher potentials as (A3) Ir2O3 is formed. The (A5) formation of IrO2 oxide may result in further decline above +0.24 V. As described above, the current increase at +0.40 V is due to O2 evolution. J


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positive of −0.20 V, the anode must periodically be brought to low potential to restore current. The cathodic sweeps from +0.53 V in Figure 16B show a very small increase in current as (C5) IrO2 and (C3) Ir2O3 are gradually reduced from +0.40 to −0.30 V (Figure 17). However, significant current restoration begins only when (C3) Ir2O3 becomes fully reduced between −0.30 and −0.44 V and lower. The (C1) reduction of Ir(OH)2/IrO may also help restore current. However, it is clear that (C4) reduction of the subtle oxide and/or Ir(OH)2/IrO below −0.56 V is mandatory for full current restoration. This observation is consistent with (A4) subtle oxide formation greatly decreasing current in the anodic sweep. Part III: Group 11 Coinage Metals. Au. A. Au Surface Chemistry. The surface states associated with electrocatalysis at Au have been described as defects in the lattice induced by high surface energy, and generally appear only after etching Au surfaces to high potential or by subjecting Au to thermal treatment. This description is part of the Incipient Hydrous Oxide/Adatom Mediator (IHOAM) model for Au.64 The model describes the defects, IHOAMs, as Au atoms that oxidize before the rest of the bulk material because of their distinct electronic properties that emerge after lattice displacement. The oxidized Au atoms may then directly take part in unique catalyses and mediate e− between the reactant and the bulk Au electrode. Many of these IHOAMs have been described for Au in 1 M NaOH84 and have been implicated in a number of reactions. These include the oxidations of N2H4, borohydride (BH4−), and formaldehyde (H2CO), and the reductions of dichromate (Cr2O72−), persulfate (S2O82−), and iodate (IO3−).65 A CV of Au in NaOH to +0.35 V (Figure 18A) shows two IHOAMs in the anodic sweep, (A1) Aua from −0.90 to −0.72 V and (A2) Aub from −0.57 to −0.23, with peaks at −0.82 and −0.40 V, respectively. While the IHOAMs are visible when avoiding H2 production (dashed gray line in Figure 18A and inset), they are much more pronounced when the lower limit of the CV is negative of H2 evolution (−1.24 V, solid black line in Figure 18A). They are also easier to observe at lower sweep rates, which have smaller peak separation and distortion, and so are shown at 25 mV/s. In the cathodic sweep, the IHOAMs (A1) Aua and (A2) Aub may have partial reductions (C3) from −0.34 to −0.58 V and (C4) from −0.64 to −0.98 V, respectively. As described above, however, some IHOAM reduction clearly occurs at much lower potentials. At higher potentials, (A3) formation of Au(OH)ads occurs from +0.15 to +0.35 V in Figure 18A, followed by Au2O3 formation at different Au crystal faces from +0.35 to +0.67 V in the broader sweep at 100 mV/s in Figure 18B (Sargent et al.85 and references therein86−90). Finally, O2 evolution manifests at E > + 0.67 V. The reduction of Au(OH)ads at different crystal faces occurs in the cathodic sweep at (C1) from +0.30 to −0.07 V and (C2) from −0.07 to −0.34 V (Figure 18A). When sweeping to higher potentials at faster scan rates, all Au hydroxides and oxides are reduced in a larger (C1) feature (Figure 18B). B. Electrocatalysis of N2H4 at Au. The RDE LSVs of N2H4 at Au show a single oxidation mechanism with n = 4 (Figures 19, parts A and B) with Eonset = −0.60 V.91,92 For the first few sweeps, Eonset = EOC = −0.48 V, after which Eonset shifts negative to −0.60 V. The oxidation of N2H4 at Au provides complete e− recovery across a broad 1 V range, making Au an ideal material for N2H4 sensors. However, Au delivers current only at 0.26 to

Figure 16. RDE LSVs of 5 mM N2H4 at Ir in 1 M NaOH, 20 mV/s, over a high potential range. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at −0.64 V in the anodic sweep.

Figure 17. Comparison of N2H4 RDE CVs at Ir, 500 rpm, over various potential ranges (A, 20 mV/s) and corresponding Ir surface chemistry (B, 100 mV/s).

The final change in the RDE of N2H4 at Ir to +0.53 V is that current in the cathodic sweeps is compromised, and is not restored until very low potentials (Figure 16B). This implies that if N2H4 fuel cells and sensors using Ir push the anode K


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Figure 19. RDE LSVs of 5 mM N2H4 at Au in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at 0 V in the anodic sweep. Figure 18. CVs of Au surface chemistry in 1 M NaOH, over ranges of low potential at 25 mV/s (A) and high potential at 100 mV/s (B). Features in the anodic (A1−A3) and cathodic (C1−C4) sweeps are described in the text.

0.46 V more positive potentials (lower voltages) than the PGMs, representing a significant compromise of power in N2H4 fuel cells. The initial Eonset of −0.48 V correlates well with (A2) the formation of IHOAM Aub, suggesting that this surface species is responsible for catalyzing N2H4 oxidation.65 The interaction of Aub with N2H4 may alter Aub’s structure, resulting in a slight shift in its potential of formation. This could in turn explain the shift in Eonset observed on consecutive RDE LSVs. In RDE LSVs to +0.80 V, N2H4 oxidation current is unaffected by (A3) Au(OH)ads formation but decreases at high coverages of Au2O3 above +0.46 V, indicating that Au2O3 is inactive for N2H4 oxidation (Figure 20, parts A and B). The decay in oxidation current is very slow, continuing in the cathodic sweep, suggesting that Au2O3 formation is a slow, time-dependent reaction once sufficiently positive potentials are reached. Chronoamperometry at +0.77 V confirms this interpretation (Figure S2). Once even small amounts of Au2O3 are reduced below +0.45 V in the cathodic sweep (Figure 20A, green line), current is rapidly restored, consistent with the observation that high coverages of Au2O3 are necessary to block the reaction. Both avoiding Au2O3 entirely by limiting LSV sweeps to +0.05 V (data not shown) and avoiding high coverages of Au2O3 by limiting sweeps to +0.40 V (Figure 20A, blue line) result in sustained oxidation current in the cathodic sweep.

Figure 20. Comparison of N2H4 RDE CVs at Au, 500 rpm, over various potential ranges (A, 20 mV/s) and corresponding Au surface chemistry (B, 100 mV/s).

Ag. A. Ag Surface Chemistry. Like Au, Ag shows subtle IHOAMs in what would otherwise appear to be a featureless double-layer region. The anodic sweep of Ag in 1 M NaOH to +0.03 V (Figure 21A) shows the formations of IHOAMs (A1) Aga from −1.00 to −0.70 V, (A2) Agb from −0.70 to −0.55 V, L


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Figure 22. RDE LSVs of 5 mM N2H4 at Ag in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at 0.10 V in the anodic sweep. Figure 21. CVs of Ag surface chemistry in 1 M NaOH, 100 mV/s, over low potential (A) and high potential (B) ranges. Features in the anodic (A1−A4) and cathodic (C1−C6) sweeps are described in the text.

The Eonset of −0.45 V is closely associated with high coverages of (A3) Agc, indicating that this IHOAM is primarily responsible for N2H4 oxidation at Ag. The Eonset is more than 0.40 V negative of the earliest possible appearance of (A5) AgI oxide (Figure 21A, dashed gray line), precluding this oxide from playing a role in Ag’s low-potential catalysis. To avoid possible etching during RDE, it was necessary to avoid (A5) forming significant AgI oxide, so potentials were kept below +0.20 V (Figure 22). When this was done, the oxidation of N2H4 continued uninterrupted in the cathodic sweep. The small amounts of AgI oxide formed from 0 to +0.20 V in Figure 21A (dashed gray line) do not inhibit the reaction. During certain experimental trials, Eonset shifted to −0.76 V, or 0.31 V lower potential (higher voltage), albeit with only n = 1 e− until −0.45 V was reached (Figure S3). This Eonset is associated with high coverages of (A1) Aga, suggesting that this IHOAM may become catalytically active after the electrode is cycled in certain potential ranges. While this observation was unfortunately not regularly reproducible, it offers an open research area of increasing the maximum voltage obtained while using inexpensive Ag catalysts. Part IV: Period 4/First-Row Transition Metals. Ni. A. Ni Surface Chemistry. CVs of Ni in 1 M NaOH in various potential ranges are shown in Figure 23, parts A and B, and reflect dynamic surface chemistry. The anodic sweep from −1.00 to −0.45 V (Figure 23A) reflects (A1) the reversible formation of α-Ni(OH)2 from−0.83 to −0.45 V, with (C1) its reduction back to Ni0 in the cathodic sweep from −0.55 to −0.92 V. Notably, H2 is generated at E < − 0.92 V, so there is a range where H2 evolution overlaps with (C1) α-Ni(OH)2 reduction.68,98−102 When CVs are expanded positive of −0.45

(A3) Agc from −0.55 to −0.33 V, and (A4) Agd from −0.33 to −0.02 V.93−96 The cathodic sweep shows reversible reductions (C1−C3) for most of these features, with the (C4) area for the reverse reaction for (A1) showing three smaller peaks for unknown reasons. Whether these features reflect reversible formation of Ag oxides or hydroxides is not known. The anodic sweep of Ag in NaOH to +0.33 V (Figure 21B) shows the more classical (A5) formation of AgI oxides from +0.29 to +0.67 V, with its corresponding (C5) reduction in the cathodic sweep from +0.28 to −0.24 V and lower potentials. The (A5) formation of AgI oxides likely starts as low as −0.02 V, immediately following (A4) Agd formation (Figure 21A, dashed gray line). An anodic sweep to +0.80 V (Figure 21B) additionally shows the (A6) formation of AgII oxide from +0.67 to +0.78 V, with its corresponding (C6) reduction in the cathodic sweep from +0.55 to +0.30 V.68 The formation of these oxides can lead to stripping Ag from the surface.68,97 B. Electrocatalysis of N2H4 at Ag. The RDE LSVs of N2H4 at Ag show a single oxidation mechanism with n = 4 with Eonset = −0.45 V (Figure 22, parts A and B). Thus Ag has the highest overpotential for N2H4 oxidation of all the metals studied, serving as a critical drawback to its use in fuel cells. However, its full e− recovery across a 0.4 V range of high current makes Ag an excellent sensor material. While Ag is prone to stripping at lower potentials than Au, Ag nonetheless provides complete e− recovery at a cost 55 times lower. M


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Figure 24. RDE LSVs of 5 mM N2H4 at Ni in 1 M NaOH, 20 mV/s, over low potential and high potential regions. Anodic (A) and cathodic (B) sweeps are shown. The insets show Levich plots taken at −0.64 (left) and +0.57 V (right) in the anodic sweep. Figure 23. CVs of Ni surface chemistry in 1 M NaOH, 100 mV/s, over isolated potential ranges (A) and a wide potential range (B) with shorter ranges from (A) superimposed in gray dashed and dotted lines. Features in the anodic (A1−A4) and cathodic (C1−C2) sweeps are described in the text.

When avoiding β-Ni(OH)2 by sweeping from −1.10 to −0.60 V, N2H4 oxidation in the anodic sweep begins at −0.96 V, shortly after H2 evolution ceases, and increases markedly at −0.90 V, where small amounts of (A1) α-Ni(OH)2 have likely formed (Figure 25A, blue line). Oxidation current begins tapering at −0.74 V, where coverage of α-Ni(OH)2 is significant (Figure 25B, blue line). This suggests that Ni0 and mixed Ni/α-Ni(OH)2 surfaces have the best low-potential catalytic activity for N2H4 oxidation. Higher rotation rates in RDE, which are conducted later in our experimentation, show less than their expected increase in limiting current, likely because the low potential limit was insufficient to prevent αNi(OH)2 from accumulating on subsequent scans. Raising the high potential limit to −0.20 V results in gradual accumulation of β-Ni(OH)2, causing passivation on subsequent scans (Figure S4). Sweeping RDE LSVs of N2H4 from −1.1 to +0.6 V (Figure 25A, dotted purple line) results in the poorest possible performance for N2H4 oxidation, which begins as a slow process at −0.17 V, precisely where (A2) appears. As discussed above, cycling in this potential range greatly decreases surface coverages of (A1) α-Ni(OH)2 and shows muted features for (A2) and (A3). Limiting sweeps to −0.10 to +0.60 V enhances the coverages of (A2) and (A3), as described above, and leads to profoundly higher oxidation current (Figure 25A, green line) precisely in the region of (A2) and (A3) formation (Figure 25B, green line). Current increases further as (A4) NiOOH is formed above +0.34 V, decreasing at high NiOOH coverage above +0.52 V. The cathodic sweep shows some hysteresis, with high

V, as in the CV to +0.63 V (Figure 23B), the α-Ni(OH)2 is dehydrated to compact β-Ni(OH)2.98 In Figure 23B, this process appears to be electrochemically silent. The β-Ni(OH)2 is reduced only at E < − 0.92 V, concomitantly with H2 evolution. The size of current for (A1) α-Ni(OH)2 formation on the return sweep was observed to vary with time spent at E < − 0.92 V. The anodic sweep of the CV to +0.63 V (Figure 23B) shows subtle process (A2) from −0.17 to +0.34 V. Examining this feature via CV from −0.16 to +0.40 (Figure 23A) reveals two distinct processes in this region, (A2) from −0.17 to +0.20 V and (A3) from +0.20 to +0.34 V, which are unknown. Both features are irreversible in this potential range and likely require lower potentials for reduction. Figure 23B shows additionally shows (A4) the formation of NiOOH from +0.34 to +0.57 V, 98−103 followed by O2 evolution. The cathodic sweep shows (C2) the reduction of NiOOH from +0.55 to at least +0.13 V. B. Electrocatalysis of N2H4 at Ni. The RDE LSVs of N2H4 at Ni show that at least three separate oxidation mechanisms are possible, with Eonset = −0.96, − 0.17, and +0.34 V, and n = 2.5, 1.5, and 3, respectively (Figure 24, parts A and B). The potential ranges in Figure 24 were carefully chosen based on Ni’s dynamic surface features in Figure 23. N


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surface species are thus irreversible, causing permanent changes to N2H4 catalysis upon cycling. The first CV sweep of Co in 1 M NaOH from −1.20 to +0.63 V is shown in Figure 27A, and a stabilized CV of Co after

Figure 25. Comparison of N2H4 RDE CVs at Ni, 500 rpm, over the two catalytically active potential regions (A, 20 mV/s) and Ni surface chemistry (B, 100 mV/s).

activity at NiOOH maintained to lower potentials until its (C2) reduction. Recently, a number of studies have explored Ni-based catalysts for N2H4 oxidation,41,42,66 and for good reason: Ni can provide reasonable e− recovery (n = 2.5) at very low potentials (high voltages), providing decent power from N2H4 at exceptionally low cost. Yet 2.5 of 4 e− reflects a 38% loss of current, and Ni tends to quickly form passivating hydroxides at minute overpotentials. These factors may limit its practical employment in durable, high-efficiency N2H4 fuel cells. Co. A. Co Surface Chemistry. The CV of Co, like that of Ni, changes profoundly upon repeated cycling (Figure 26), with a rapid decrease in current for a major low-potential species and sharpening of a high-potential feature. In the presented potential range of −1.20 to +0.63 V, the formations of many

Figure 27. First CV (A) and stable CV (B) of Co surface chemistry in 1 M NaOH, 20 mV/s. Features in the anodic (A1−A5) and cathodic (C1−C6) sweeps are described in the text.

multiple cycles is shown in Figure 27B. The anodic sweep shows (A1) and (A2) the formations of Co(OH)2 and CoO, hydrated and dehydrated, from −0.85 to −0.75 V and −0.75 to −0.05 V, respectively;68,71,104−108 (A3) the formation of an unknown species Cox from −0.05 to +0.11 V; (A4) the formation of CoIII and mixed CoII/CoIII species, including CoOOH, Co 2 O 3 , and Co 3 O 4 , from +0.11 to +0.49 V;68,71,104−107 and (A5) formation of CoO2 from ≤ + 0.49 V to +0.63 V,68,71,104,105 with O2 evolution superimposed. The beginning of (A5) likely overlaps with the end of (A4). The cathodic sweep shows reductions (C1) and (C2) of CoO2 from +0.63 to +0.43 V and +0.43 to +0.10 V, respectively; (C3) reduction of CoOOH, Co2O3, and Co3O4 from +0.10 to −0.52 V, with residual reduction occurring as low as −0.77 V; (C4) reduction of Cox from −0.51 to −0.77; and (C5) reductions of Co(OH)2 and CoO from −0.77 to −1.12 V. The stabilized CV of Co in Figure 27B shows a near-total elimination of Cox (A3) formation and (C4) reduction, the appearance of (C6) another low potential reduction, and sharpening of the other features. Various processes have shifted their range of appearance: (A2) now appears from −0.75 to

Figure 26. Consecutive CVs of Co surface chemistry in 1 M NaOH, 20 mV/s, over a high potential range. O


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The Journal of Physical Chemistry C −0.12 V, (A4) from +0.09 to +0.40 V, and (A5) from ≤ + 0.40 V to +0.63 V. Feature (C6) arises from −0.77 to −0.96 V, likely a reduction of either Co(OH)2 or CoO, while (C5) occurs from −0.96 to −1.12 V. B. Electrocatalysis of N2H4 at Co. The RDE LSVs of N2H4 at Co (Figure 28, parts A and B) show four areas of catalytic

Figure 29. Comparison of 5 mM N2H4 RDE CV at Co at 500 rpm (A) and Co surface chemistry (B), both at 20 mV/s.

N2H4 oxidation catalysts based on CoNi catalysts have been studied recently, with an emphasis on the low-potential (highvoltage) oxidation from −1.00 to −0.30 V (Figure 28A).41 While oxidative current clearly manifests in this region, the RDE LSVs indicate that e− recovery here is quite small. To better understand the low-potential processes occurring at Co, we superimposed the first scan of stationary CV of N2H4 at Co with the first scan of Co in NaOH alone (Figure 30). This data

Figure 28. RDE LSVs of 5 mM N2H4 at Co in 1 M NaOH, 20 mV/s. Anodic (A) and cathodic (B) sweeps are shown. The inset shows a Levich plot taken at 0.60 V in the anodic sweep.

activity in the anodic sweeps with Eonset = −1.00, − 0.85, + 0.10, and +0.35 V. Only Co’s high-potential oxidation at +0.35 V exhibits a strong dependency on rotation rate and shows a maximum n = 1.7; an exact determination of n was prohibited by curvilinear Levich and Koutecky−Levich plots, likely due to passivation at higher rotation rates. While Co exhibits the lowest Eonset of all the metals considered in our study, the lack of significant e− recovery until potentials 1.35 V higher (lower voltage) severely restricts Co’s ability to deliver high performance in N2H4 fuel cells. The first catalytic region of the RDE LSVs in Figure 28A at −1.00 V likely corresponds to a region of Co metal, while the subsequent three catalytic areas (−0.85, + 0.10, and +0.35 V) correspond nearly perfectly to the major surface processes (A1)/(A2), (A4), and (A5) (Figure 29, parts A and B). Cathodic sweeps in Figure 28B show little hysteresis from +0.68 to +0.32 V and minor hysteresis from +0.32 to 0 V, reflecting the hyseresis of their corresponding Co oxides. However, a complete absence of current from −0.30 to −1.00 V is evident, reflecting that catalytically active (A1) and (A2), Co(OH)2 and CoO, remain absent from the surface until the subsequent anodic sweeps.

Figure 30. Comparison of first scans of CVs of N2H4 in 1 M NaOH at Co and Co in 1 M NaOH. The current attributable to N2H4 and not simply Co surface processes is shown in red.

indicates that the vast majority of the charge passed at lowpotential at 5 mM N2H4 may be attributed to the oxidation of Co itself rather than N2H4. Current for N2H4 also rapidly decreases on subsequent scans, just as it does for Co surface species (Figure S5 vs Figure 26), reflecting significant passivation of the Co surface likely responsible for the curvilinear Levich and Koutecky−Levich plots mentioned P


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with Rh and Ir, two of the most expensive catalysts, showing the highest Power Factor. Several catalysts (Ni, Co, Ag) have multiple data points in Figure 31, indicating multiple N2H4 oxidation mechanisms of interest occurring at separate potentials. Negative Power Factors reflect negative fuel cell voltage with O2, and more positive voltages would be obtained when using higherpotential oxidants, such as Ce4+, ClO−, or MnO4−. Several catalysts show performances of high interest and are worthy of detailed discussion. In Figure 31, it is clear that Ni shows the most unexpectedly high performance for its cost, with the Ni outlier at 2.1 mW/cm2 corresponding to its lowpotential oxidation with Eonset = −0.9 V in Figure 24. For this reason, Ni has served as the basis for Daihatsu’s development of low-cost N2H4 fuel cell catalysts.41,42 Yet there exist serious drawbacks for the use of Ni in an N2H4 fuel cell. As discussed earlier, the onset of N2H4 oxidation is directly tied to the appearance of α-Ni(OH)2, but similar potentials also generate passivating β-Ni(OH)2, decreasing current even for CVs limited to −0.6 V in Figure 24. Higher potentials caused irreversible surface passivation (Figure 25). Any fuel cell using Ni catalysts would require strict operating controls to avoid higher potentials and periodically reduce any reversible β-Ni(OH)2 that forms. This would significantly decrease the robustness of a commercial system. Furthermore, the low-potential current at Ni corresponds to a 2.5 e− process, and previous work has shown that the remaining e− are lost as H2.41 Conversion of N2H4 to H2 dramatically decreases power density, since H2 has solubility

Trends in Catalysis and Catalyst Cost Effectiveness for N2

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