Laboratory Experiment pubs.acs.org/jchemeduc
Facile Method To Study Catalytic Oxygen Evolution Using a Dissolved Oxygen Optical Probe: An Undergraduate Chemistry Laboratory To Appreciate Artificial Photosynthesis Genesis Renderos, Tawanda Aquino, Kristian Gutierrez, and Yosra M. Badiei* Department of Chemistry, Saint Peter’s University, Jersey City, New Jersey 07306, United States S Supporting Information *
ABSTRACT: Artificial photosynthesis (AP) is a synthetic chemical process that replicates natural photosynthesis to mass produce hydrogen as a clean fuel from sunlight-driven water splitting (2H2O → O2 + H2). In both natural and artificial photosynthesis, an oxygenevolving catalyst (OEC) is needed to catalyze oxygen production from water (2H2O → O2 + 4H+ + 4e−). Herein, we present a simple undergraduate laboratory experiment for monitoring the catalytic oxygen-evolution reaction using a dissolved oxygen optical probe, a luminescent probe that can continuously measure the concentration of dissolved oxygen in aqueous solutions. The laboratory experiment consists of two parts. First, students identify and evaluate the necessary criteria for an active molecular O2-evolving catalyst. Second, students use a kinetic protocol based on the method of initial rates to derive an empirical rate law expression and calculate the turnover frequency and rate constant for a benchmark OEC, [RuII(tpy)(bpy)(H2O)]2+ (tpy is 2,2′:6′,2″-terpyridine; bpy is 2,2′-bipyridine). The use of simple transition metal complexes allows students to easily grasp how the choice of transition metals and the coordinated organic ligands can greatly impact the overall catalytic performance. Students collect timedependent dissolved oxygen plots, and analyze and interpret oxygen-evolution data using modern equipment. This laboratory protocol is simple, and allows students to relate many fundamental topics such as thermodynamics, kinetics, redox and coordination chemistry, and catalysis to advances in current chemical research to help broaden their understanding of artificial photosynthesis. KEYWORDS: Upper-Division Undergraduate, Physical Chemistry, Inorganic Chemistry, Catalysis, Kinetics, Hands-On Learning/Manipulatives, Coordination Compounds, Aqueous Solution Chemistry, Oxidation/Reduction
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INTRODUCTION The global need for energy is constantly increasing, and as such, it is important to develop available renewable energy sources that can deliver energy on a large scale without an adverse environmental impact. Global climate concerns have risen from the rapid increase in emissions of carbon dioxide, a greenhouse gas, produced from the escalating production and use of nonrenewable fossil fuels.1 For over 2 billion years, green plants, algae, and bacteria have been harnessing and converting sunlight energy, water, and carbon dioxide into carbohydrates and molecular oxygen, a process known as photosynthesis, which powers life on Earth.2 Inspired by nature, the synthetic process of using cheap, inexhaustible, and available materials to make chemical fuels is known as artificial photosynthesis (AP).3 In this context, many scientists are highly interested in the sunlight-driven water splitting reaction that separates water into its constituents, elemental oxygen and hydrogen (Figure 1).4−6 Hydrogen is a clean and environmentally friendly fuel because it is free of carbon, and when combusted, it releases energy and forms water as the only byproduct. For example, fuel cells utilize hydrogen as an energy source in combination with oxygen to generate zero-emission electricity. As hydrogen can be produced in the dark by electrolysis of water, the © 2017 American Chemical Society and Division of Chemical Education, Inc.
Figure 1. Diagram depicting artificial photosynthesis.
development of solar water splitting technologies (i.e., by AP) will provide a more commercially viable and sustainable Received: February 4, 2017 Revised: May 25, 2017 Published: June 16, 2017 922
DOI: 10.1021/acs.jchemed.7b00074 J. Chem. Educ. 2017, 94, 922−927
Journal of Chemical Education
Laboratory Experiment
redox potential which can help mediate the water-oxidation reaction.15 Although the long-term goal is to use sunlight for AP, the activities of O2-evolving catalysts are generally evaluated by using a powerful chemical oxidant such as cerium ammonium nitrate [(NH4)2Ce(NO3)6] (CAN). Ce(IV) is a one-electron oxidant and can oxidize water at low pH (E° Ce(IV)/Ce(III) = +1.7 V versus NHE, pH = 0)16 in the presence of an active catalyst as shown in eq 5. CAN is a popular oxidant in the screening of OEC since it is highly reactive, inexpensive, and readily available in high purity.17
method to produce hydrogen on a larger scale, and can help increase the widespread use of fuel cell devices.4−8 overall: 2H 2O(l) → O2 (g) + 2H 2(g) Ecell° = Ered° + Eox ° = − 1.23 V
(1)
O2 production: 2H 2O(l) → O2 (g) + 4H+(aq) + 4e− Eox ° = −1.23 V (2)
H 2 production: 4H+(aq) + 4e− → 2H 2(g) Ered° = 0 V
2H 2O(l) + 4Ce(IV)(aq) catalyst
⎯⎯⎯⎯⎯⎯→ O2 (g) + 4H+(aq) + 4Ce(III)(aq)
(3)
Water splitting is a redox reaction (eq 1) and consists of two half-reactions: the oxygen evolution reaction (eq 2), also known as water oxidation, and proton reduction to hydrogen (eq 3). The oxygen evolution reaction is thermodynamically demanding (ΔG° = +475 kJ mol−1), requiring a potential of 1.23 V versus NHE at pH 0.0 (ΔG° = −nFE°), and is considered the bottleneck of AP due to the high kinetic barrier associated with the removal of four electrons and four protons from water, and the sluggish formation of the oxygen−oxygen bond (2H 2O → O 2 + 4H + + 4e − ).9 The minimum thermodynamic requirement for the reaction can sustainably be provided by sunlight energy. In addition, an effective oxygen evolution catalyst (OEC) is needed to lower the activation energy by stabilizing intermediates critical for the O−O bond formation step to produce oxygen at fast rates, i.e., high turnover frequency (TOF) as defined by eq 4. Furthermore, in electrochemical reactions, kinetic barriers are commonly present due to slow electron transfer that arises at the interface between a solution and an electrode. Reaction rates, as measured by the current density (current divided by the area of electrode) that passes through the cell, typically increase rapidly with increasing applied potential. The additional potential beyond the thermodynamic requirement (for, e.g., 1.23 V in water oxidation) needed to drive a reaction forward at a certain rate is called the overpotential.10,11 An efficient OEC can accelerate the reaction rate of water oxidation and hence allows a high current density to be achieved at a low overpotential. TOF for OEC =
(moles of O2 /moles of catalyst) time
E°Ce(IV)/Ce(III) = 1.7 V
(5)
Despite extensive research efforts from the scientific community on artificial photosynthesis, there are only a few undergraduate hands-on laboratory experiments18−20 that introduce students to this rapidly growing field. However, for real progress to occur, we need to better engage learners and make them aware of scientific approaches and challenges in the commercial viability of renewable clean technologies such as artificial photosynthesis.20,21 Herein, we present a facile protocol that can be implemented in an undergraduate teaching laboratory for studying the oxygen-evolution reaction with an artificial benchmark OEC, [RuII(tpy)(bpy)(H2O)]2+ (tpy is 2,2′:6′,2″-terpyridine; bpy is 2,2′-bipyridine). This complex has been thoroughly examined in the literature as a water-oxidation catalyst.22−24 The Ru(II) complex contains a coordinated water ligand and a mixed set of polypyridine ligands (Figure 2).
(4)
The use of simple molecular-based OECs allows the rational understanding of the structure−activity relationship and provides a wealth of mechanistic insight into the complexity of the water-oxidation/O2-evolving reaction.12,13 Typically, these OECs contain Ru and Ir metal ions, where the coordinated ligands can be structurally and electronically modified to improve the catalytic performance.9,10 Recent progress has also been shown with earth-abundant transition metal complexes containing Co, Fe, Mn, and Cu.14 Important desirable criteria for these active molecular catalysts are (i) to be robust under highly oxidizing conditions in water, (ii) to undergo multielectron and multiproton steps, and (iii) to minimize overpotential, i.e., the difference between the equilibrium potential and the catalytic potential to drive the oxygen evolution reaction (OER).10 The use of electrondonating ligands (e.g., 6,6′-dicarboxy-bpy) has been shown to favor the formation of high-valent intermediates at a lower
Figure 2. Molecular structure of the benchmark O2-evolving catalyst, [RuII(tpy)(bpy)(H2O)]2+ (tpy is 2,2′:6′,2″-terpyridine; bpy is 2,2′bipyridine).
Students perform oxygen-evolution kinetic experiments using a Vernier dissolved oxygen optical probe (DOOP), a luminescence-based oxygen sensor that can continuously measure the increase in the concentration of dissolved oxygen (mg/L) in aqueous solutions. Although the Vernier DOOP has been used in laboratory experiments in biological systems, it has not been commonly studied in a chemical environment. To our knowledge, this is the first undergraduate laboratory experiment that describes a simple protocol for direct measurement of dissolved oxygen and the kinetics of a chemical reaction using 923
DOI: 10.1021/acs.jchemed.7b00074 J. Chem. Educ. 2017, 94, 922−927
Journal of Chemical Education
Laboratory Experiment
from organic combustible materials. Perchloric acid solutions should be handled appropriately and can be very hazardous in the case of skin contact (corrosive, irritant), eye contact, or ingestion. This experiment requires training students on how to handle a nitrogen gas cylinder. The molecular catalyst [RuII(tpy)(bpy)(H2O)]Cl2 is not considered extremely toxic but should be handled with care.
this luminescent probe. We illustrate the robustness of the DOOP sensor in monitoring the oxygen evolution reaction, especially under a strong oxidizing environment, and the practicality of its use as an instructional tool in an undergraduate laboratory setting to identify and study the kinetics of an OEC. Moreover, this undergraduate laboratory experiment is unique as compared to other experiments involving artificial photosynthesis18,19,21 in focusing on the use of a homogeneous catalyst for water oxidation to help illustrate the importance of redox chemistry and the ability of supporting ligands to tune the catalytic properties. In addition, the kinetics experiments in this laboratory will enable students to understand the wateroxidation reaction mechanism on a molecular level.
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RESULTS AND DISCUSSION
Identification of an O2-Evolving Catalyst (Part A)
The O2-evolution reaction was carried out using [Ru(tpy)(bpy)(H2O)]2+ as the catalyst and CAN as the chemical oxidant; O2(aq) is the dissolved oxygen product of the catalytic reaction monitored in 1.0 M HClO4 solution (pH = 0) using DOOP in a closed reaction flask. The experimental setup is shown in Figure 3. Initial experiments were designed to test the
EXPERIMENTAL SECTION
Equipment and Materials
All dissolved oxygen concentration (mg/L) measurements were performed using the Vernier DOOP. The DOOP sensor can be purchased online.25 The Vernier Logger Pro software was used for data collection and analysis and was interfaced to the probe using the Vernier Lab Quest. The oxygen-evolution measurements were performed in a side arm glass flask. The [RuII(tpy)(bpy)(H2O)](Cl2) catalyst was synthesized by undergraduate research assistants according to the literature. For more details on the synthesis, solution preparation, and detailed procedures for using DOOP, see the Supporting Information. General Protocol for Using a Dissolved Oxygen Optical Probe
The dissolved oxygen optical probe was first connected to the computer or laptop by plugging it in one of the channels in the LabQuest Mini interface. The sensor does not need to be calibrated or stirred and can be used right away without warmup. Nevertheless, a miniature stir bar was used to accelerate the rates of oxygen evolution. The oxygen levels read by the probe were first checked by inserting the probe tip in fresh deionized water. Typical values for dissolved oxygen (DO) concentrations in the aqueous phase are 8−9 mg/L at ambient conditions. Oxygen-evolution measurements in solution were mainly performed in a leak-free side arm flask fitted with the DOOP in a temperature regulated water bath at 21 ± 1 °C. A septum was placed on the side arm joint for catalyst injection. In a typical experiment, the flask equipped with a miniature stir bar was first charged with 25.0 mL of a 22 mM solution of (CAN) in 1.0 M perchloric acid. Note that, as an alternative, nitric acid, sulfuric acid, or triflic acid can also be used for this experiment (see Supporting Information). The use of HCl acid is not recommended since chloride is a noninnocent anion in the reaction and can be oxidized to chlorine gas.17 Dissolved oxygen preliminary readings were initially collected (180 s) in acidic solutions of CAN until stable readings (to the nearest 0.2 mg/L) of dissolved oxygen were obtained before the addition of the catalyst (1.0 mL). For more details on the setup for the catalytic reaction see the Supporting Information.
Figure 3. Simple dissolved oxygen measurements setup with a dissolved oxygen optical probe.
response of the optical probe in the CAN solution before and after the addition of the catalyst. We found that the optical probe immersed in aqueous acidic solutions (e.g., in 1.0 M HClO4) containing 25 mL of 22 mM CAN is quite robust, and no significant changes in the concentration of oxygen levels were detected even for prolonged times (4 h) of continuous measurements (Figure S1). In addition, during 4 h of use in a typical laboratory session, the measured baseline concentration of dissolved oxygen in solution remains constant and stable after repetitive measurements. The addition of 1.0 mL of the benchmark molecular catalyst [RuII(tpy)(bpy)OH2)]2+ (ca. 1.86 mM) to 25 mL of deaerated CAN solution (22 mM) in 1.0 M HClO4 resulted in a continuous increase in the concentration of dissolved oxygen signal (mg/L) for the first 1000 s of the reaction (Figure 4 and Figure S2). Control experiments performed for the same duration displayed no significant oxygen production (
