Article pubs.acs.org/JPCC
Poly-tetrakis-5,10,15,20-(4-aminophenyl)porphyrin Films as TwoElectron Oxygen Reduction Photoelectrocatalysts for the Production of H2O2 Nicholas U. Day and Carl C. Wamser* Department of Chemistry, Portland State University, Portland, Oregon 97207-0751, United States S Supporting Information *
ABSTRACT: Thin films of electropolymerized porphyrin polymers, poly-tetrakis-5,10,15,20-(4-aminophenyl)porphyrin (pTAPP) and its cobalt derivative (pCoTAPP), were found to be electrocatalysts and photoelectrocatalysts for the photosynthesis of hydrogen peroxide via a two-electron reduction of oxygen. On glassy carbon (GC) electrodes, oxygen reduction potentials were measured at −0.58 (GC), −0.40 (pTAPP), and −0.05 V (pCoTAPP), compared to Pt at +0.34 V (all vs Ag/AgCl). Thin electrode films were tested as photosynthetic electrocatalysts using small positive bias potentials (0.0 to +0.3 V) and specifically measuring H2O2 production. pTAPP achieved turnover numbers (TON) of 5−6, while pCoTAPP showed TON of 14−23. Faradaic efficiency in both cases was initially high, about 50%, but decreased over 1 h.
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capable of selective two-electron reduction of O2.9 Recently Mase et al. reported on the sunlight-driven production of H2O2 from O2 and seawater at an efficiency of 0.55%; the concentration of H2O2 reached 48 mM, suitable for driving a fuel cell at 50% efficiency, giving an overall solar-to-electrical efficiency of 0.28%.10 Recently, Osterloh has reviewed the distinctions between processes that are photocatalytic and those that are photosynthetic.11 When the energy of light is used to drive a reaction that is uphill (ΔG > 0), the process is photosynthetic, as opposed to a photocatalytic process, in which light is used to increase the kinetics of a reaction that is already favorable (ΔG < 0). In this work, we are attempting to develop a successful photosynthetic process, focusing on the reduction of oxygen to hydrogen peroxide as a means of chemical energy storage (ΔG > 0). We will use the corresponding terminology for photosynthetic processes as much as possible, even though most of the prior literature in this area refers to photocatalysts. We have recently synthesized and characterized poly-tetrakis5,10,15,20-(4-aminophenyl)porphyrin (pTAPP) and investigated its novel conductive and electrochromic properties.12,13 Bettelheim et al. also characterized the electropolymerization of variously substituted porphyrins as well as many of their metalated derivatives.14−17 Catalysts immobilized by polymer supports onto electrodes are promising systems for solar fuel production.18,19 They have been shown to catalyze CO2 reduction, hydrogen production, and the reduction of carbon tetrachloride.19−22 Recently, an electropolymerized cobalt
INTRODUCTION The effects of global climate change have been clearly linked to the ongoing output of CO2 from fossil fuel use.1 In order to diminish the current dependence on fossil fuels, cheap and efficient ways to generate fuels from renewable energy have become increasingly important objectives.2 Hydrogen is theoretically one of the best molecules for energy storage, yet storage of hydrogen safely and efficiently at high energy densities is still a significant challenge.3 Nevertheless, the first hydrogen fuel cell vehicles have recently become commercially available.4 In comparison there has been much less research on hydrogen peroxide fuel cells, an alternative fuel with significant potential advantages. Hydrogen peroxide fuel cells operate with H2O2 as both the oxidant and reductant with H2O and O2 as the only byproducts.5 This allows the fuel cell to operate without a semipermeable membrane, which substantially simplifies the architecture; these systems have been recently reviewed by Fukuzumi et al.6 The oxidation of H2O2 to O2 occurs at a standard potential of +0.68 V vs NHE, and the reduction of H2O2 to H2O occurs at a standard potential of +1.77 V vs NHE.7 Therefore, the hydrogen peroxide fuel cell can potentially operate at an open circuit voltage up to 1.09 V, close to that of the hydrogen fuel cell at 1.23 V.8 In addition, H2O2 can be easily stored at room temperature in high concentrations in plastic containers.5 The reverse process, generation of H2O2 from O2, is a key objective for developing this process into a sustainable cyclic system. The two-electron reduction of O2 to H2O2 is to be distinguished from the more well-studied four-electron oxygen reduction reaction (ORR) of O2 to H2O. There has been increasing research on developing catalysts and photocatalysts © XXXX American Chemical Society
Received: February 2, 2017 Revised: May 10, 2017 Published: May 10, 2017 A
DOI: 10.1021/acs.jpcc.7b01071 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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light entered the cuvette and went through the FTO glass before striking the sample, which faced the bulk of the solution. Oxygen was continuously bubbled through the solution for the entire experiment, which lasted from 30 to 120 min. Samples for H2O2 analysis had aliquots removed at 0, 5, and 10 min and at 10 min intervals thereafter. Analysis for H2O2. Samples were tested for H2O2 using double recrystallized titanium(IV)-oxo-tetrapyridylporphyrin (TiOTPyP), which was synthesized according to previously published procedures.24 Testing of each aliquot was performed as follows using methods for the determination of H2O2 optimized from previously published procedures.25 To a scintillation vial 5 mL of a 15 μM TiOTPyP solution in aqueous 50 mM HCl was added, followed by 5 mL of 6.0 M perchloric acid, and then 5 mL of deionized H2O. The solution was then mixed by vortex and let stand for 5−20 min. To a 990 μL aliquot of this solution, a 10 μL aliquot of the sample to be tested was added, thoroughly mixed, and allowed to stand for 10 min. After the 10 min incubation time, a spectrum was taken from 350 to 650 nm. The quantity of H2O2 was determined using the decrease in absorbance at 433 nm relative to a control sample prepared with only 10 μL of the acetate buffer solution.25 Calibration of TiOTPyP as a colorimetric reagent for H2O2 was performed using measured quantities of hydrogen peroxide urea. TiOTPyP has been reported to accurately determine the presence of H2O2 between 0.1 and 10 μM.25 Aliquots (10 μL) of hydrogen peroxide urea solutions at concentrations of 5.0, 4.0, 3.0, 2.0, 1.0, and 0 mM were combined with 990 μL aliquots of TiOTPyP solution as described above. Absorbance changes at 433 nm were determined and plotted in Figure S1 of the Supporting Information. Four different series were prepared: 0−5 nmol, 0−3 nmol linear range, and at each end of the linear range, at concentrations of 0.1, 0.2, 0.3, 0.4, and 0.5 nmol, and at 2.8, 2.9, 3.0, 3.1, and 3.2 nmol. The four curves shown in Figure S1 had an average slope of 0.11 ± 0.01 absorbance unit/nmol of hydrogen peroxide urea; the linear fit equation was used to determine the concentration of H2O2 in all samples tested.
porphyrin was found to have electrocatalytic activity for oxygen reduction.23 Herein we report the photosynthetic activity of pTAPP and pCoTAPP for the two-electron reduction of O2 for H2O2 production. To our knowledge this is the first time a nonmetalated porphyrin polymer has been shown to produce H2O2.
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EXPERIMENTAL METHODS Synthesis of pTAPP and pCoTAPP Films. Fluorinedoped tin oxide (FTO) working electrodes (1 × 6 cm) were cleaned for 5 min by sonication in Sparkleen detergent solution followed by 10 min of sonication in deionized H2O. All electrochemical preparations used a Pt flag counter electrode (2 cm2) and a Ag/AgNO3 reference electrode. pTAPP films were prepared from 0.15 mM TAPP in dichloromethane with 20 mM tetrabutylammonium perchlorate, using either multiple cycles between 0.0 and +0.6 V with a 10 s delay at +0.6 V vs Ag/AgNO3, or using only a single pass from 0.0 to +0.6 V with a 10−200 s hold at +0.6 V, and then cycling back to 0 V. Both procedures ended at +0.4 V vs Ag/AgNO3, found to be the oxidation state at which pTAPP has the lowest charge transfer resistance.13 pCoTAPP films were prepared by immersing pTAPP films in 10 mM cobalt acetate solution in methanol and heating to reflux. Complete metalation was detected spectrophotometrically. Instrumentation. UV−visible spectra were taken on a Shimadzu UV 3600 spectrophotometer. Electrochemical and conductivity measurements were performed with a Gamry Reference 600 potentiostat. Electrochemical Studies. Glassy carbon electrodes were encased in a 1 cm glass tube with Loctite 1C epoxy and exposed by grinding with 300 grit sandpaper. The electrode was then polished using sandpaper of 600 grit followed by 1200 grit, and then Buehler alumina paste of 1.0 μm, 0.3 μm, and finally 0.05 μm. The final electrode surface was determined to be 0.04 cm2 using ImageJ software. Ohmic contact was established using silver paint and tinned copper wire. Cyclic voltammograms (CVs) of pTAPP and pCoTAPP films on glassy carbon electrodes were taken in 0.1 M KCl (electrochemical grade from Sigma-Aldrich) in Milli-Q 18 Mohm water at various pH values, adjusted with concentrated HCl, 4 M HCl, or 10 M NaOH. A Pt coil counter electrode and a Ag/AgCl reference electrode were also in solution. All solutions were bubbled with nitrogen or argon for 20 min before adjusting for each pH and again before each cycle was run. Nitrogen or argon was blanketed over the solution during CV scans. Typical samples were run at multiple scan rates including 100 mV/s. In catalytic tests, oxygen or air was bubbled through solutions for 10 min with stirring before testing. Photosynthetic Studies. pTAPP and pCoTAPP samples for photosynthetic testing were synthesized on FTO glass electrodes using fixed potentiometry, holding for 200 s at +0.6 V vs Ag/AgNO3. The samples were all approximately 2 cm2 in size. Samples were submerged in a 3 mL cuvette along with a Ag/AgCl reference electrode and an electrochemical grade graphite counter electrode with an approximate surface area of 3 cm2. The cuvette was filled with 0.1 M acetate buffer at pH 3.0. All samples were held at 0 V bias vs Ag/AgCl and illuminated with a 5000 K 500 lm LED light; radiant intensity at the sample face was measured with a silicon radiometer at approximately 2 suns of irradiation during photosynthesis of H2O2 studies and 0.8 sun of irradiation for photoresponse measurements. The samples were placed in the cuvette so that
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RESULTS AND DISCUSSION Metalation of pTAPP films occurs rapidly and completely by immersion in methanolic cobalt acetate solution and heating at reflux briefly. Completeness of metalation was monitored by disappearance of the band at 791 nm in pTAPP films. The
Figure 1. UV−vis absorbance spectra of pTAPP and pCoTAPP films. The same film was used for both spectra, before metalation in red and after metalation in blue. The inset shows the same spectra but with the x-axis as energy in eV. B
DOI: 10.1021/acs.jpcc.7b01071 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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Figure 2. Oxygen reduction at glassy carbon (GC) electrodes coated with (a) pTAPP and (b) pCoTAPP taken in aqueous solutions at pH 4 with a scan rate of 100 mV/s. The Co-treated GC control electrode in (b) was treated identically to the metalation procedure for the preparation of pCoTAPP. Figure 4. Photoresponse of 1 cm2 pTAPP films at different bias potentials vs Ag/AgCl. All films were tested at pH 3.0 in 0.1 M acetate buffer saturated with oxygen and under 0.8 sun illumination. The system was allowed to equilibrate for 30 s before the light was turned on for 10 s intervals starting at 20, 40, and 60 s.
Figure 5. Band gap diagram showing photoresponse behaviors in the photosynthetic reduction of oxygen by pTAPP.
absorbance spectra of the free base polymer and the cobalt metalated polymer are shown in Figure 1. The absorption seen at 1300 nm has not been reported before, but is likely due to the presence of polarons, which allow charge hopping and thus conductivity within the polymer film. Similar long wavelength absorptions are seen in other conductive polymers.26−28 Electrocatalytic Studies. Metalated aminophenylporphyrin derivatives have been tested as electrocatalysts for the oxygen reduction reaction and have shown small overpotentials.29 As metal-free catalysis is often more desirable because metals are sometimes susceptible to catalyst poisoning,9 we decided to first elucidate the catalytic behavior of free-base pTAPP films. Nitrogen-doped carbon nanotubes have shown catalytic
Figure 3. CVs of (a) pTAPP and (b) pCoTAPP in saturated O2 solutions of 0.1 M acetate buffer at pH 3.0 with and without illumination at 0.8 sun intensity.
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DOI: 10.1021/acs.jpcc.7b01071 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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Figure 6. Photoresponse of 1 cm2 pCoTAPP films at different biases vs Ag/AgCl. All films were tested at pH 3.0 in 0.1 M acetate buffer saturated with oxygen and under 0.8 sun illumination. The system was allowed to equilibrate for 30 s before the light was turned on for 10 s intervals starting at 20, 40, and 60 s.
Figure 7. TON vs time for samples of (a) pTAPP and (b) pCoTAPP with 2 suns illumination at bias potentials of +0.3 (purple), +0.2 (blue), +0.1 (red), and 0 V (green) vs Ag/AgCl. All samples tested were approximately 2 cm2 in a cuvette with aqueous 0.1 M acetate buffer at pH 3 with a graphite counter electrode and constant O2 bubbling.
behavior for oxygen reduction even better than that of Pt, attributed to the electron-withdrawing effect of nitrogen on neighboring carbon atoms.30 pTAPP has multiple types of aromatic nitrogen heterocyclic structures, including the porphyrin core and the phenazine linkages,12,13 suggesting that it may be a promising material for metal-free oxygen reduction. Figure 2 shows cyclic voltammograms of pTAPP and pCoTAPP films on glassy carbon (GC) electrodes in the presence of oxygen as well as control experiments without O2. The half-wave potential determined by a cyclic voltammogram is typically used to specify the potential at which a redox reaction takes place, although detectable catalysis begins sooner than this at the onset potential. Overpotentials are estimated by comparing measured half-wave potentials using pTAPP or pCoTAPP versus the half-wave potential measured using Pt (+340 mV vs Ag/AgCl). In order to standardize the magnitude of the overpotential, samples were compared to the catalysis peak with a platinum electrode. Current commercial polymer electrolyte membrane fuel cells use Pt-based electrodes for the oxygen reduction electrode.31 Figure S2 in the Supporting Information shows O2 reduction at four different electrodes including glassy carbon, pCoTAPP on glassy carbon, pTAPP on glassy carbon, and a platinum electrode, all at pH 3 and under identical conditions. Platinum shows a half-wave reduction potential of +0.34 V vs Ag/AgCl, while glassy carbon has a half-wave reduction
potential of −0.58 V vs Ag/AgCl. The difference of 0.92 V is due to glassy carbon’s poor catalytic activity and illustrates why GC electrodes are commonly used to study catalytic behavior. A GC electrode coated with pCoTAPP shows a half-wave reduction potential of −0.05 V vs Ag/AgCl at pH 3, indicating that there is still an increase of 0.39 V in potential required as compared to the Pt electrode. pTAPP shows a half-wave reduction potential of −0.40 V indicating an overpotential of 0.74 V compared to the Pt electrode. Photosynthetic Studies. Both pTAPP and pCoTAPP films show an increase in current under illumination, i.e., a photoresponse. This was first determined by cyclic voltammograms under dark and illuminated conditions as shown in Figure 3. Photoresponses can be measured in different ways including a decrease in the resistivity of the material or increased activity of redox reactions in solution, both of which were observed with pTAPP and pCoTAPP samples. Because of the difficulty in accurately measuring the decrease in resistivity of films and the interest in using pTAPP and pCoTAPP as heterogeneous catalysts, the photoresponse as a function of redox reactions in solution was measured. Since this is a catalytic reaction which is being partially or completely driven by light absorption, it has been called a photoassisted catalysis D
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(although it may be better termed a photoassisted synthesis).10,32 Running identical CVs back to back, one in light and one in the dark, higher photocurrents were observed at a range of potentials under illuminated conditions. It should be noted that there is no significant change in current between light and dark under identical conditions without oxygen present. Depending on the applied bias potential, the photoresponse can range from almost no increase in current to around 50 μA/ cm2 at −0.1 V vs Ag/AgCl (Figure 4). The photoresponse maximizes at a potential lower than that at which oxygen reduction is catalyzed electrochemically by pTAPP (−0.37 V vs Ag/AgCl), and even shows significant photocurrent at +0.4 V vs Ag/AgCl, i.e., an underpotential with respect to Pt under identical conditions. The photoresponse correlates with the band gap of pTAPP. Figure 4 shows that the photoresponse is essentially absent when the film is at potentials above +0.5 V or below −0.5 V; these values correspond closely to the oxidation and reduction potentials for pTAPP determined by cyclic voltammetry from our previous findings.13 This suggests that the electrons photoexcited into the lowest occupied molecular orbital (LUMO) of the film have enough energy to reduce oxygen in solution. If the potential bias applied is not within the band gap, then the electrons are either not excited into the LUMO of the polymer or the LUMO of the polymer is already occupied; in the latter case electroreduction takes place, indicated by the large background currents present at negative potentials. The explanation of the photoresponse in relation to band gap theory is shown in Figure 5. In this photosynthetic system, electrons are excited by light from the highest occupied molecular orbital (HOMO) to the LUMO and then leave the LUMO to perform the solution reduction of O2. Holes in the HOMO are filled by electrons from the electrode. Comparable measurements of the photoresponse of pCoTAPP films show their activity to be lower than that of pTAPP films, with a maximum of approximately 10 μA/cm2, as shown in Figure 6. Furthermore, when the pCoTAPP film was in buffer saturated with argon, a significant photoresponse was still detected, approximately 5 μA/cm2, indicating that pCoTAPP is a less selective catalyst than pTAPP. Nevertheless, measurements of photosynthetic activity, discussed in the section Photosynthetic Efficiency, indicate that pCoTAPP films are more efficient than pTAPP films in the production of H2O2. Photosynthetic Efficiency. In order to test the synthetic efficiency (often referred to as catalytic efficiency), the number of electrons transferred into solution and the number of catalytic sites available were compared to the quantity of H2O2 produced. Using pTAPP and pCoTAPP samples on FTO glass irradiated in a cuvette, H2O2 production was measured at different bias potentials as a function of time. It is important to note that FTO can also show catalytic activity for oxygen reduction, with a onset potential of −0.1 V vs Ag/AgCl. Production of H2O2 from pTAPP and pCoTAPP was confirmed on glassy carbon electrodes, but because of the small amounts produced, all quantitative measurements are from polymer films on larger FTO electrodes. In order to remove any reduction of O2 due to FTO from reduction due to the polymers, all measurements were taken at bias potentials of 0 V and more positive vs Ag/AgCl. pTAPP was tested for photosynthetic efficiency at a range of potentials between 0 and +0.4 V vs Ag/AgCl. The photosynthetic activity between 0 and +0.2 V vs Ag/AgCl was found
Figure 8. Average Faradaic efficiencies over time for multiple samples tested for (a) pTAPP and (b) pCoTAPP under 2 suns irradiation and bias potentials of 0.0 or 0.1 V vs Ag/AgCl.
Figure 9. Production of H2O2 by (a) pTAPP and (b) pCoTAPP samples in the dark (blue) and under 2 suns irradiation (red) at +0.1 V vs Ag/AgCl in saturated oxygen at pH 3 in 0.1 M aqueous sodium acetate buffer.
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two identical samples of each were tested for production of hydrogen peroxide under illumination and in the dark. Figure 9a shows that pTAPP samples produced no H2O2 in the dark while a modest amount of H2O2 was produced in the light. Figure 9b shows that pCoTAPP in the light produces approximately twice as much H2O2 as pCoTAPP in the dark and approximately 6 times the amount of H2O2 that pTAPP produces in the light.
to be essentially independent of the potential as was suggested by the photoresponse experiments. The photosynthetic activity at +0.3 V was lower and variable depending on the sample, and at +0.4 V no production of H2O2 was detectable. pTAPP samples tested without irradiation at these potentials show no activity in the production of H2O2 over a 1 h time period. The number of monomeric units of TAPP in a representative pTAPP film was calculated by quantitating the amount of TAPP removed from solution during polymerization. Using an identical polymerization method, seven samples of pTAPP were electropolymerized; an average of 60 ± 6 nmol of TAPP was removed from solution, representing the upper limit of the number of TAPP equivalents deposited on the electrode. Using this quantity of monomeric units and assuming each monomeric unit to be photosynthetically active, turnover numbers (TON) can be determined as the number of products produced divided by the number of photosynthetic sites in the system. Figure 7 shows TON values over time for both pTAPP and pCoTAPP films at various bias potentials. Although showing greater variability, pCoTAPP clearly shows higher TON values than pTAPP. Both polymer films, however, eventually lose activity, as indicated by the leveling off of the TON values with time. From seven of the pTAPP samples tested the average TON was 5.6 ± 1 (Figure 7a). pCoTAPP proved to be a more active photoelectrocatalyst for the reduction of O2 with average TON values of 17.5 ± 4 (Figure 7b). pCoTAPP at 0 V vs Ag/ AgCl produces much of the H2O2 in the first 5 min, while at +0.3 V the TON grows steadily, still producing H2O2 over more than an hour. Since +0.3 V is much higher than the −0.05 V half-wave potential determined for electrosynthesis, it is clear that pCoTAPP is driven at this potential primarily photosynthetically. Since pCoTAPP can produce H2O2 at +0.3 V vs Ag/AgCl, there is almost no overpotential compared to platinum or the thermodynamic potential, making pCoTAPP a promising material for further study. Faradaic efficiency was calculated by integrating the current passed during a photoelectrosynthetic experiment using chronoamperometry. By calculating the number of moles of electrons passed and using a molar ratio of two electrons per mole of H2O2, the theoretical amount of H2O2 that could have been created by that current was determined. Faradaic efficiency is the actual number of moles of H2O2 produced divided by the theoretical number of moles. Faradaic efficiency showed no correlation with bias potentials between 0 and +0.2 V vs Ag/AgCl for either pTAPP or pCoTAPP (Figure 8). The highest Faradaic efficiency observed, for two of the four samples tested, was just over 70% for the pCoTAPP samples held at either 0 or +0.1 V bias vs Ag/AgCl, suggesting this system can be a good photoelectrocatalyst for production of H2O2. Samples held at +0.2 and +0.3 V vs Ag/AgCl maintained production of H2O2 for longer periods of time but showed lower initial Faradaic efficiencies, close to 30%. No sample of pTAPP held at any potential achieved a Faradaic efficiency above 50%. Both pTAPP and pCoTAPP pass more current than is necessary for the quantity of H2O2 produced, indicative of nonselective electrochemical processes. Some other reactions likely to be occurring are the reduction of H2O2 to H2O and the reduction of protons to H2. Testing for other products was not done; no gas was seen to evolve from the electrode system during testing. Considering both pTAPP and pCoTAPP to be photoelectrocatalysts as well as electrocatalysts at positive potentials,
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CONCLUSIONS Both forms of the electrochemically polymerized porphyrin films showed activity for the reduction of oxygen to hydrogen peroxide. Electrochemically, pCoTAPP showed a lower overpotential for oxygen reduction than pTAPP under identical conditions. Photoelectrochemically, pCoTAPP produces H2O2 at close to the thermodynamic potential for the reduction of oxygen to hydrogen peroxide. Faradaic efficiencies were initially good and TON values for both polymers were modest; both measures indicate that the polymers gradually lose effectiveness over about 1 h. Both pTAPP and pCoTAPP remain on the electrode and no treatments have been found to restore activity.
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ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.7b01071. Titrations of TiOTPyP with standard solutions of hydrogen peroxide urea; cyclic voltammograms of oxygen reduction at four electrodes under identical conditions (Pt, glassy carbon, pTAPP, pCoTAPP) (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. ORCID
Carl C. Wamser: 0000-0001-5969-8376 Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS Support from the Oregon Nanoscience and Microtechnologies Institute (ONAMI) and the National Science Foundation (Grant CHE-0911186) is gratefully acknowledged.
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DOI: 10.1021/acs.jpcc.7b01071 J. Phys. Chem. C XXXX, XXX, XXX−XXX
