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Rate and stability of photocatalytic water oxidation using [Ru(bpy)3]2+ as photosensitizer Bart Limburg, Elisabeth Bouwman, and Sylvestre Bonnet ACS Catal., Just Accepted Manuscript • DOI: 10.1021/acscatal.6b00107 • Publication Date (Web): 29 Jun 2016 Downloaded from http://pubs.acs.org on June 30, 2016
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Rate and stability of photocatalytic water oxidation using [Ru(bpy)3]2+ as photosensitizer B. Limburg, E. Bouwman, S. Bonnet* Leiden University, Leiden Institute of Chemistry, Gorlaeus Laboratories, P.O. Box 9502, 2300 RA Leiden. *Email corresponding author:
[email protected] Abstract The kinetics of homogeneous photocatalytic water oxidation is reported using [Ru(bpy)3]Cl2 as photosensitizer, Na2S2O8 as sacrificial electron acceptor, and three different water-oxidation catalysts: the ruthenium catalyst [Ru(bda)(isoq)2] ([1], H2bda = 2,2’-bipyridine-6,6’-dicarboxylic acid, isoq = isoquinoline), Co(NO3)2 ([2]), or [Ir(Cp*)(dmiz)(OH)2] ([3], Cp* = pentamethylcyclopentadienyl, dmiz = 1,3-dimethylimidazol-2-ylidene). At pH=7.0, in a phosphate buffer, and under blue light irradiation, the production of O2 at the catalyst is rate determining when [2] or [3] are used as water-oxidation catalysts. However, when [1] is used as catalyst in identical conditions the turn over at the wateroxidation catalyst is not the rate-limiting step of the photocatalysis. Instead, the step limiting dioxygen production is the transfer of electrons from the catalyst to the photooxidized photosensitizer [Ru(bpy)3]3+. Due to the instability of [Ru(bpy)3]3+ in neutral aqueous solutions, slow electron transfer results in significant photosensitizer decomposition, which limits the overall stability of the photocatalytic system. When the catalyst [1] is used decomposition of both the photosensitizer and the catalyst [1] occurs in parallel. However, the photosensitizer and the catalyst also stabilize each other, i.e., the TON increases when more photosensitizer is added, while the photocatalytic turnover number PTON increases when more catalyst [1] is added. These data demonstrate that not only new and more stable water oxidation catalysts should be developed in the future, but that new and more stable photosensitizers are needed as well. Keywords: Photocatalysis, water oxidation, electron transfer, kinetics, ruthenium tris bipyridine
Introduction In recent years the number of reported water-oxidation catalysts has increased enormously.1, 2 While in attempts to mimic nature the first catalysts were based on manganese, more recently catalytically active compounds for water oxidation have been reported based on cobalt,3-7 nickel,8-10 iron,11-13 copper,14-20 ruthenium21-34 and iridium.35-39 Generally, the catalytic activity of these compounds is investigated using the oxidant (NH4)2[Ce(NO3)6] (CAN).21-23 In three decades the turnover frequencies (TOF) of water-oxidation catalysts have increased from 2 × 10–3 s–1 to more than 300 s–1 under optimal conditions. Meanwhile, the apparent stability of the catalysts, as reflected by turnover numbers (TONs), has also increased over 3 orders of magnitude.40 The standard activity-screening method for water oxidation driven by CAN provided the scientific community with a rough estimate of which catalysts might be viable candidates for solar fuel production. However, due to the absence of light, the high cerium concentrations used in these experiments, the high oxidation potential of the CeIV / CeIII redox couple (E0 = +1.71 V vs. SHE), and the low pH (less than 1), these testing conditions are far from the operational conditions under which water-oxidation catalysts would operate in a solar fuel-producing photocatalytic system.40 Several 1
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water-oxidation catalysts have also been tested under photocatalytic conditions, usually using the sacrificial electron acceptor S2O82– and a molecular photosensitizer (PS) such as [Ru(bpy)3]2+ or a derivative thereof.6, 7, 23, 28, 41-63 Under such conditions, the TOF is generally much lower (< 1 s–1) than that obtained by CAN-driven chemical oxidation (> 300 s–1). 7, 28, 41-63 More importantly, the stability of photocatalytic systems is also lower (generally TON < 300, compared to >> 1000 for CAN-based systems),7, 28, 40-63 which, independently of the nature of the catalyst, can be attributed to the decomposition of the photosensitizer.64-66 The shift from chemical (CAN) to photocatalytic conditions brings more complexity to the wateroxidation reaction. Whereas in chemical water oxidation the concentration of the catalyst and oxidant, the acid used, and the temperature are sufficient to define the catalytic system, under photocatalytic conditions many more parameters must be optimized, and all of them significantly influence the reaction. For example, photochemical quenching by dioxygen, the light intensity, or the photosensitizer concentration,41 each influence the rate and stability of the system. Other parameters such as the pH of the solution,43 the type of buffer,66 the buffer concentration,67 or the catalyst concentration,6, 7 can also play a critical role. Photocatalytic water-oxidation systems are intricate systems to study; for example, the mechanism of the initial step of the photochemical reaction, i.e., of the oxidative quenching of the excited state of the photosensitizer (PS*) by the electron acceptor S2O82– in absence of any catalyst, is in itself a complex reaction.68 Furthermore, the oxidized photosensitizer [Ru(bpy)3]3+ (PS+), which is produced by oxidative quenching of PS* by S2O82–, is notoriously unstable in neutral and alkaline solution,64, 65 which further complicates the study of photocatalytic water oxidation. In order to further advance towards a device that can be used for photocatalytic splitting of water it is essential to understand what is the bottleneck of photocatalytic water oxidation. Scheme 1 depicts a global kinetic overview of the photocatalytic reaction in which S2O82– is used as sacrificial electron acceptor. In this model reaction, three steps can limit the overall rate of dioxygen production: the oxidative quenching of PS* by the electron acceptor (step r1), the electron transfer from the catalyst to PS+ (step r2), or the catalytic oxidation of water (step r3). As in photocatalytic conditions the wateroxidation catalyst is likely to exist in different oxidation states in the catalytic cycle, we name r2 the slowest step of oxidation of these catalytic intermediates by PS+. For similar reasons step r3 is taken as the rate-limiting step during the dioxygen evolution reaction, generally reported as the liberation of O2 or formation of the O-O bond.11, 22, 23, 33, 69, 70 The sulfate radical SO4•–, which is formed upon oxidative quenching of PS* by S2O82– to produce PS+, is also a strongly oxidizing species that can accept a second electron either from the catalyst, or from PS. Finally, two different decomposition pathways may occur that lead to either the loss of PS (rd1), or of the catalyst (rd2). Herein, we investigate how variations of the main parameters of the photocatalytic system, i.e., light intensity and the concentration of all components of the system, influence the evolution of the overall rate and stability of photocatalytic water oxidation using [Ru(bpy)3]2+ as photosensitizer, and three different water oxidation catalysts based on Ru, Co and Ir. We also discuss which of the three main catalytic steps r1, r2, or r3, may, depending on the conditions, be the bottleneck of the photocatalytic reaction.
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Scheme 1. Simplified scheme for photocatalytic water oxidation. The total oxygen production rate (rtot) is determined by the slowest step of the intermediate reactions r1, r2, or r3. r1: formation rate of PS+ via light absorption by PS and oxidative quenching of PS* by S2O82–. r2: electron transfer rate from the catalyst to PS+. This step needs to be performed four times to generate one molecule of O2. r3: O2 production at the catalyst. Depending on the slowest step (r1 vs. r2 vs. r3), one of the decomposition process rd1 or rd2 will become more prominent. The initial decomposition product of rd1 can be further oxidized by rox1 or rox2, by reacting with PS+ or S2O82–, respectively. Initial reduction of S2O82– leads to the sulfate radical anion, SO4•–, which can directly oxidize either the photosensitizer (rpo), or the catalyst (rco). For some water-oxidation catalysts, the active species must first be formed from the precatalyst (indicated by the grey arrow).
Results Description of the experimental setup. The kinetics of photocatalytic water-oxidation was studied in homogeneous system employing three different water-oxidation catalysts based on ruthenium ([Ru(bda)(isoq)2], [1]), cobalt ([Co(H2O)6](NO3)2, [2], or iridium ([Ir(Cp*)(dmiz)(OH)2], [3]) (Figure 1). In all cases, the photosensitizer was [Ru(bpy)3]Cl2 (PS), the sacrificial electron acceptor was Na2S2O8, and the solvent was a phosphate buffer at pH=7.0. Under such conditions, [1] is reported to be a homogeneous catalyst,22 [2] is reported to form nanoparticles,3, 4, 71 and the nature of catalyst [3] seems to depend strongly on the conditions.35, 39, 72, 73 A lag phase in the photocatalytic water-oxidation reaction might be present as the active catalyst must first be formed by oxidation of the precatalyst (see the grey arrow in Scheme 1). The concentration of the water-oxidation catalyst was kept constant (5 µM) and the samples, additionally containing PS (50 µM) and sacrificial electron acceptor Na2S2O8 (2.5 mM), were irradiated with blue light (λirr = 450 nm). The dioxygen levels were measured simultaneously in solution by an Ocean Optics FOXY-AR fluorescence probe dipped into the reaction mixture, and in the gas phase by sampling the headspace with gas chromatography (GC).
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Figure 1. Chemical structures of the water-oxidation catalysts used in this study, based on ruthenium ([1]),22 cobalt ([2]), and iridium ([3]).39
The details of the experimental setup are reported in the Supplementary Information. As reported elsewhere35 the results of the GC measurements were initially not in agreement with the dissolved dioxygen quantification, which is due to the slow transfer of O2 between the solution and the gas phase. For this reason, the headspace gas was pumped through the solution for 30 seconds prior to each injection into the GC (see ESI). Such pumping ensured quick equilibration of O2 between the solution and the headspace, as well as homogenization of the headspace. Typical traces of O2 evolution as a function of irradiation time are shown in Figure 2. In our setup the calculated TOF (initial rate), based on the dioxygen concentration as measured in solution with the fluorescence probe (solid curves), and calculated vs. the amount of catalyst present in the system, are in close agreement with those based on the GC data (data points). As oxygen is produced in solution the initial TOF was measured at the beginning of the curve obtained from the Ocean Optics fluorescence probe, at the point where the slope of the curve shown in Figure 2 is at its maximum. Initials TOFs of 4.5 × 10–2 s–1, 8.3 × 10–2 s–1 and 1.3 × 10–1 s–1 were found for [3], [2], and [1], respectively, at a light intensity of 8.4 mW·cm–2 . It should be noted that a lag phase due to active catalyst formation was observed for [2]. This lag phase lasted less than five seconds, indicating fast formation of the active catalyst. We therefore assume that our measurements were done at a point where the catalyst is in the active state. Measuring the total dioxygen production was more complicated. As the irradiation progresses, the amount of dissolved dioxygen goes through a maximum and then decreases to 0 ppm. As the pumping procedure removes almost all O2 from the solution and brings it into the headspace, the TON can in principle be determined from the GC data. The quantification by GC is however affected by three factors. First, inevitable air leakage increased the error on the quantification of O2 with increasing conversion (see ESI). Second, inevitable exchange of helium from the purged housing of the pump with the headspace gas decreased the measured amount O2 over time. Third, for every data point measured by GC, 0.5 mL of the headspace was consumed and replaced by the same volume of helium, which led to a decrease in the measured amount of O2. Therefore, the total amount of O2 measured at the end of the reaction (30 minutes) was lower than the amount of O2 actually produced. When we attempted to correct for these problems, the error present on the last data point (t = 30 min) was transferred to the start of the reaction, thus leading to a much larger standard deviation between identical experiments. Therefore, the amount of O2 removed from the system by GC analysis was neglected, and all TONs reported below were calculated from the maximum amount of O2 measured in the headspace by GC (asterisk in Figure 2). 4
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For all catalysts low stabilities were found, characterized by TONs (amount of O2 produced divided by the amount of water-oxidation catalyst) smaller than 50. Both the stabilities and activities observed in these photocatalytic conditions were found to be low compared to chemical oxidation using CAN, for which much higher TOFs and TONs were previously reported. For example, the reported TOF and TON for catalyst [1] in chemical oxidation conditions are 300 s–1 and 8360, respectively.22 These results suggested that the productivity of photocatalytic systems is not limited by the activity and stability of the catalysts, and that further mechanistic studies should be undertaken to understand what is the bottleneck in photocatalytic water oxidation.22, 39
Figure 2. Photocatalytic dioxygen evolution as measured in solution by a fluorescence probe (solid curves) and in the headspace by GC (data points). The red (squares), blue (diamonds) and green (triangles) curves represent [1], [2], and [3] as catalyst, respectively. Conditions: catalyst (5 µM), [Ru(bpy)3]Cl2 (50 µM), Na2S2O8 (2.5 mM) in 5 mL phosphate buffer (10 mM, I = 50 mM, pH = 7.0), λirr = 450 nm, θ0 = 30 nmol·s–1, T = 298 K. The dips in the curves of the probe, and the decrease after a maximum, are caused by the solution-headspace equilibration procedure (see ESI). The asterisks indicate the datapoints that were used for the calculation of the TON.
Rate of the system Varying the light intensity. In order to find out which step in the photocatalytic process (r1, r2, or r3; Scheme 1) determines the rate of dioxygen evolution, first the light intensity was varied under otherwise equal conditions. Low intensity blue light was obtained from an LED, and high intensity white light was obtained from a 1 kW Xenon lamp fitted with a UV and an IR filter (for the lamp spectrum, see ESI). The rate of O2 evolution with both light sources can only be compared if the number of excited states of the photosensitizer PS* formed per unit time, hereafter noted θ (mol.s-1), is known. The number of photons absorbed by the PS per unit time was calculated from the irradiance spectrum of the incoming light, convoluted with the absorption spectrum of the PS (see ESI). Kasha’s rule was assumed to be valid over the entire metal-to-ligand charge transfer (MLCT) absorption band of PS, and the quantum yield of triplet formation was taken to be equal to 1, i.e., the number of 3MLCT excited states produced per second (θ) was taken to be equal to the number of photons absorbed by the PS per second. The overlap between the irradiation spectrum and the absorption spectrum of PS has a great impact on the number of excited states produced per second, and thus it is possible to obtain the same θ for two light sources with a large difference in light intensity. For instance, a θ value of 30 nmol·s–1 was obtained with light intensities of 8.4 mW·cm–2 for the blue LED or of 26 mW·cm–2 for the white Xenon lamp. However, due to the high maximum power of the Xenon lamp much higher values of θ can be obtained compared to the maximum power of the blue LED (380 nmol·s–1 vs. 30 nmol·s–1, respectively).
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Figure 3a shows the dependence of the initial TOF on the initial amount of photons absorbed by the PS per second at t = 0, noted θ0. From these data, it is clear that catalyst [1] behaves differently from catalyst [2] and [3]. The TOF for [2] and [3] was more or less stable upon increasing θ0, which indicates that in such conditions and over the entire range of studied θ0 the O2 evolution rate was not limited by the amount of excited states PS* produced per second. Due to the extremely low dioxygengeneration rate at low light intensities for these two water-oxidation catalysts it was not possible to find the regime where the light intensity would limit the reaction rate. At θ0 19 nmol·s–1) this was not the case, therefore step r3, i.e., dioxygen production at the catalyst, was rate-limiting for systems containing [2] or [3]. For catalyst [1] on the contrary, a plot of the initial TOF vs. θ0 showed two different regimes. For values of θ0 greater than 0.14 µmol·s–1, the TOF decreased with increasing θ0. However, for values of θ0 lower than 0.14 µmol·s–1 the TOF was found to be almost proportional to θ0, showing that in this range of light intensities the photocatalysis is limited by the number of photons absorbed per second by the PS. Coincidentally, the light intensity of the Xenon lamp at the maximum TOF was found to be equal to the “solar constant” of 136 mW·cm–2.74 Thus, in the current system under “working conditions” employing the sun as irradiation source, the rate-limiting step is not the release of dioxygen from catalyst [1] (r3). In other terms, up to intensities of θ0 = 0.14 µmol·s–1 catalyst [1] is capable of producing dioxygen at a rate higher than the other steps in the photocatalytic process (r1 and r2). From here on we will only consider the applicable range where θ0 is smaller than 0.14 µmol·s–1.
Figure 3. a) Plot of the initial TOF for catalyst [1] (squares), [2] (diamonds) and [3] (triangles) as a function of the number of excited states PS* produced per second at t=0, θ0. Irradiation system was either a blue LED (λirr = 450 nm, filled data points) or a white 1 kW Xenon lamp fitted with a UV and IR cutoff filter (open data points). Concentration of catalyst: 5 µM. b) Plot of the rate of dioxygen production (vO2, diamonds, right axis) and TOF (triangles, left axis) as a function of the concentration of catalyst [1] under constant irradiation by a blue LED (λirr = 450 nm, θ0 = 30 nmol·s–1). The error bars denote the standard deviation of at least a duplicate experiment. Common conditions for a and b: [Na2S2O8] (2.5 mM), [Ru(bpy)3]Cl2 (50 µM) in 5 mL phosphate buffer (10 mM, I = 50 mM, pH = 7.0) , T = 298 K.
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Varying the catalyst concentration. In a second step, the light intensity was fixed at a constant value (θ0 = 30 nmol·s–1, supplied by the blue LED) and the concentration of catalyst [1] was varied between 1 and 20 μM. In such conditions the rate of dioxygen production, vO2, was found to be more or less constant; as a consequence the highest TOF was obtained at the lowest concentration tested for [1] (Figure 3b). This counter-intuitive result highlights a critical difference between photocatalysis and classical (dark) catalysis: many experimental parameters that are completely independent from the catalyst may critically influence the overall reaction rate. As an example, a TOF of ~0.5 s-1 was observed both at low concentration of [1] (1 µM) and low θ0 (30 nmol·s–1), or at higher concentration of [1] (5 µM) and higher θ0 (140 nmol · s ). In these experiments, as the rate at which catalyst [1] releases O2 is not limiting the reaction, adding more catalyst leads to the formation of more O2, but does not necessarily lead to faster reactions, so that characterizing the performances of the photocatalytic system only by reporting a TOF calculated vs. the amount of catalyst may be misleading. Furthermore, comparing TOF measured in photocatalytic conditions to that determined in chemical oxidation conditions (e.g. using CAN as an oxidant) is tricky. On the one hand, several conditions, such as the pH of the solution, are different in both type of experiments (7.0 in photocatalysis vs. 1.0 under CAN oxidation). On the other hand, when CAN is used the oxidant concentration is much higher (typically 0.1-1 M) than in photocatalytic conditions, where the oxidant PS+ is produced by light at much lower concentrations (
