Applied Science, Faculty of
Chemical and Biological Engineering, Department of
DSpace
UBCV
Hejazi, Seyyed Arman
Master of Applied Science - MASc
University of British Columbia
Organic pollution of water is one of the most persistent types of impurities. In municipal water treatment plants, the degradation of organic compounds is achieved by advanced oxidation processes (AOPs), where the exposure of hydrogen peroxide (H₂O₂) containing influent to ultraviolet (UV) lamps converts H₂O₂ to hydroxyl radicals (·OH), which oxidize organic pollutants. However, the need to introduce a chemical (H₂O₂) and the use of bulky UV lamps prevent this process from being scaled down. In this study, an electrochemical cell was developed for the in situ generation of H₂O₂ from water and atmospheric oxygen using a two-electron oxygen reduction reaction (2e-ORR), thus eliminating the need for the external addition of H₂O₂. Moreover, the electrochemical reactor was equipped with novel UV sources, ultraviolet light-emitting diodes (UV-LEDs) with a wavelength of 277 nm, and microplasma lamps with wavelengths of 172 nm, 222 nm, and a broad spectrum of 220–280 nm. The electrochemical cell exhibited a high current efficiency of nearly 90% toward H₂O₂ production. The degradation of 2,4-dichlorophenoxyacetic acid (2,4-D) and methylene blue (MB) was monitored based on various operational parameters. The performance of the system for electro-Fenton (EF) and photoelectro-Fenton (PEF) processes was studied using the three microplasma lamps. In all cases, the system exhibited high removal of the organic compounds (e.g., >70% of 2,4-D with UV-LEDs at the water flow rate of 50 mL min⁻¹). A gas diffusion electrode (GDE) was developed based on expanded polytetrafluoroethylene (ePTFE) substrate to produce H₂O₂. The GDE characterization indicated its tolerance to a pressure difference of >70 psi. The performance of the GDE for the cathodic generation of H₂O₂ was evaluated based on catalyst loading, electrolyte flow rate, and current density. The results indicated a direct relationship between the current and H₂O₂ production. Higher catalyst loading led to increased H₂O₂ production, while the concentration of H₂O₂ decreased with an elevated flow rate. In sum, the electrochemical photoreactor developed in this study demonstrated key advantages such as small footprint, low power consumption, and high contaminant degradation rate, making it suitable for point-of-use (POU) water treatment to eliminate both microbial and chemical contaminants.
DEVELOPMENT OF A POINT-OF-USE WATER TREATMENT SYSTEM FOR ORGANIC CONTAMINANT DEGRADATION by Seyyed Arman Hejazi B.Sc., Sharif University of Technology, 2018 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE in THE FACULTY OF GRADUATE AND POSTDOCTORAL STUDIES (Chemical and Biological Engineering) THE UNIVERSITY OF BRITISH COLUMBIA (Vancouver) December 2020 © Seyyed Arman Hejazi, 2020 ii The following individuals certify that they have read, and recommend to the Faculty of Graduate and Postdoctoral Studies for acceptance, a thesis entitled: Development of a point-of-use water treatment system for organic contaminant degradation submitted by Seyyed Arman Hejazi in partial fulfillment of the requirements for the degree of Master of Applied Science in Chemical and Biological Engineering Examining Committee: Dr. Fariborz Taghipour, Professor, Chemical and Biological Engineering, UBC Supervisor Dr. Előd Gyenge, Professor, Chemical and Biological Engineering, UBC Supervisory Committee Member Dr. Yankai Cao, Assistant Professor, Chemical and Biological Engineering, UBC Supervisory Committee Member iii Abstract Organic pollution of water is one of the most persistent types of impurities. In municipal water treatment plants, the degradation of organic compounds is achieved by advanced oxidation processes (AOPs), where the exposure of hydrogen peroxide (H2O2) containing influent to ultraviolet (UV) lamps converts H2O2 to hydroxyl radicals (·OH), which oxidize organic pollutants. However, the need to introduce a chemical (H2O2) and the use of bulky UV lamps prevent this process from being scaled down. In this study, an electrochemical cell was developed for the in situ generation of H2O2 from water and atmospheric oxygen using a two-electron oxygen reduction reaction (2e-ORR), thus eliminating the need for the external addition of H2O2. Moreover, the electrochemical reactor was equipped with novel UV sources, ultraviolet light-emitting diodes (UV-LEDs) with a wavelength of 277 nm, and microplasma lamps with wavelengths of 172 nm, 222 nm, and a broad spectrum of 220–280 nm. The electrochemical cell exhibited a high current efficiency of nearly 90% toward H2O2 production. The degradation of 2,4-dichlorophenoxyacetic acid (2,4-D) and methylene blue (MB) was monitored based on various operational parameters. The performance of the system for electro-Fenton (EF) and photoelectro-Fenton (PEF) processes was studied using the three microplasma lamps. In all cases, the system exhibited high removal of the organic compounds (e.g., >70% of 2,4-D with UV-LEDs at the water flow rate of 50 mL min–1). A gas diffusion electrode (GDE) was developed based on expanded polytetrafluoroethylene (ePTFE) substrate to produce H2O2. The GDE characterization indicated its tolerance to a pressure iv difference of >70 psi. The performance of the GDE for the cathodic generation of H2O2 was evaluated based on catalyst loading, electrolyte flow rate, and current density. The results indicated a direct relationship between the current and H2O2 production. Higher catalyst loading led to increased H2O2 production, while the concentration of H2O2 decreased with an elevated flow rate. In sum, the electrochemical photoreactor developed in this study demonstrated key advantages such as small footprint, low power consumption, and high contaminant degradation rate, making it suitable for point-of-use (POU) water treatment to eliminate both microbial and chemical contaminants. v Lay Summary Ultraviolet (UV) radiation combined with hydrogen peroxide (H2O2) is one common method of generating free radicals to degrade organic contaminants present in drinking water. The process is used in drinking water treatment plants worldwide. However, the issues associated with bulky conventional UV lamps and the need for H2O2 addition prevent this process from being scaled down for a point-of-use (POU) system application. In this work, an electrochemical cell was developed to produce H2O2 using oxygen of ambient air and water. The electrochemical cell was then coupled with UV sources to convert the generated H2O2 to free radicals. The effect of various operating and design parameters on the performance of the system was investigated. Moreover, a new generation of air-breathing cathodes was developed for practical H2O2 production. The results suggest that this system can successfully degrade low concentrations of organic contaminants, which indicates its suitability for a POU water treatment system. vi Preface I, Seyyed Arman Hejazi, am the principal author of the present thesis. I held the full responsibility for conducting literature review, developing research questions, designing experiments, executing experiments, analyzing results, and writing reports. The whole process was performed under the direct supervision of Dr. Fariborz Taghipour, as the principal investigator of the project. The research presented in this work was conducted in the Photo-reaction Engineering Laboratory of Chemical and Biological Engineering (CHBE) Department at the University of British Columbia (UBC), Vancouver, Canada. A version of Chapter 2 is ready for submission to an academic journal: • Hejazi, A. and Taghipour, F., Integrated UV-LED hydrogen peroxide electrochemical photoreactor for point-of-use water treatment. A version of Chapter 3 is ready for submission to an academic journal: • Hejazi, A. and Taghipour, F., Microplasma ultraviolet radiation integrated with electrochemical in situ generation of hydrogen peroxide for degradation of organic pollutants in water. A version of Chapter 4 is ready for submission to an academic journal: • Hejazi, A. and Taghipour, F., Expanded polytetrafluoroethylene-based gas diffusion electrode for electrochemical generation of hydrogen peroxide. vii A version of Chapter 2 was presented in a national conference as an oral presentation: • Hejazi, A. and Taghipour, F., Point-of-use device for water treatment with no addition of chemicals, 69th Canadian Chemical Engineering Conference (CCEC 2019), Halifax, Canada, Oct. 2019. viii Table of Contents Abstract ......................................................................................................................................... iii Lay Summary .................................................................................................................................v Preface ........................................................................................................................................... vi Table of Contents ....................................................................................................................... viii List of Tables .............................................................................................................................. xiii List of Figures ............................................................................................................................. xiv List of Symbols .............................................................................................................................xx List of Abbreviations ................................................................................................................ xxii Glossary ......................................................................................................................................xxv Acknowledgements .................................................................................................................. xxvi Dedication ................................................................................................................................ xxvii Chapter 1: Introduction ................................................................................................................1 1.1 Background ..................................................................................................................... 1 1.1.1 Advanced oxidation processes .................................................................................... 2 1.1.1.1 Electrochemical H2O2 generation ....................................................................... 3 1.1.1.1.1 GDEs ............................................................................................................. 5 1.1.1.2 Ultraviolet light-emitting diodes ......................................................................... 6 1.1.1.3 Microplasma lamps ............................................................................................. 7 1.2 Research questions .......................................................................................................... 7 1.3 Research objectives ......................................................................................................... 8 1.4 Thesis outline .................................................................................................................. 9 ix Chapter 2: Integrated UV-LED hydrogen peroxide electrochemical photoreactor for point-of-use water treatment .................................................................................................................11 2.1 Introduction ................................................................................................................... 11 2.2 Materials and methods .................................................................................................. 14 2.2.1 Materials ................................................................................................................... 14 2.2.2 Electrochemical photoreactor ................................................................................... 15 2.2.2.1 UV-LED array preparation ............................................................................... 16 2.2.2.2 Nafion membrane preparation .......................................................................... 17 2.2.3 Experimental procedure ............................................................................................ 17 2.3 Results and discussion .................................................................................................. 19 2.3.1 Performance of the cell for the in situ cathodic generation of H2O2 ........................ 20 2.3.2 Effect of UV, H2O2, and UV/H2O2 on degradation of 2,4-D .................................... 29 2.3.3 Effect of incident UV fluence rate on degradation of 2,4-D ..................................... 33 2.3.4 Effect of initial concentration of 2,4-D on degradation of 2,4-D ............................. 35 2.3.5 Effect of initial pH on degradation of 2,4-D ............................................................. 36 2.3.6 Energy and stability analyses of the system ............................................................. 39 2.4 Conclusions ................................................................................................................... 39 Chapter 3: Microplasma ultraviolet radiation integrated with electrochemical in situ generation of hydrogen peroxide for degradation of organic pollutants in water .................41 3.1 Introduction ................................................................................................................... 41 3.2 Materials and methods .................................................................................................. 45 3.2.1 Materials ................................................................................................................... 45 3.2.2 Electrochemical photoreactor ................................................................................... 46 x 3.2.3 Experimental procedure ............................................................................................ 47 3.2.3.1 Microplasma UV lamp operation...................................................................... 48 3.2.3.2 H2O2 concentration measurement method ........................................................ 50 3.2.3.3 MB concentration measurement method .......................................................... 51 3.3 Results and discussion .................................................................................................. 51 3.3.1 Electrochemical cell performance for the generation of H2O2 ................................. 52 3.3.1.1 Effect of electrolyte type and concentration ..................................................... 58 3.3.1.2 Effect of electrolyte temperature ...................................................................... 60 3.3.2 UV/H2O2 process with microplasma lamps for MB degradation ............................. 62 3.3.2.1 Direct photolysis of MB using microplasma lamps .......................................... 63 3.3.2.2 UV/H2O2 process using in situ generation of H2O2 and microplasma lamps ... 65 3.3.3 EF and PEF processes for MB degradation .............................................................. 67 3.3.3.1 Effect of initial pH and Fe2+ concentration ....................................................... 68 3.4 Conclusions ................................................................................................................... 70 Chapter 4: Expanded polytetrafluoroethylene-based gas diffusion electrode for electrochemical generation of hydrogen peroxide ....................................................................72 4.1 Introduction ................................................................................................................... 72 4.2 Material and methods .................................................................................................... 74 4.2.1 Materials ................................................................................................................... 74 4.2.2 Electrochemical reactor and experimental procedure ............................................... 74 4.2.2.1 Nafion membrane preparation .......................................................................... 76 4.2.3 H2O2 concentration measurement method ................................................................ 76 4.2.4 GDE preparation ....................................................................................................... 77 xi 4.2.5 Characterization techniques ...................................................................................... 78 4.2.5.1 SEM imaging .................................................................................................... 78 4.2.5.2 Contact angle measurement .............................................................................. 78 4.2.5.3 Water entry pressure measurement ................................................................... 79 4.2.5.4 MIP ................................................................................................................... 80 4.3 Results and discussion .................................................................................................. 81 4.3.1 Surface morphology analysis .................................................................................... 81 4.3.2 Hydrophobicity analysis ........................................................................................... 82 4.3.3 Leak resistance analysis ............................................................................................ 84 4.3.4 Pore distribution analysis .......................................................................................... 85 4.3.5 Electrochemical cell performance for the generation of H2O2 ................................. 89 4.4 Conclusions ................................................................................................................. 100 Chapter 5: Conclusions and Recommendations .....................................................................101 5.1 Conclusions ................................................................................................................. 101 5.1.1 Chapter 2 ................................................................................................................. 101 5.1.2 Chapter 3 ................................................................................................................. 103 5.1.3 Chapter 4 ................................................................................................................. 104 5.2 Recommendations ....................................................................................................... 106 References ...................................................................................................................................109 Appendices ..................................................................................................................................132 Appendix A ............................................................................................................................. 132 A.1 UV-LED array ........................................................................................................ 132 A.2 Absorption spectra of various solutions .................................................................. 133 xii A.3 H2O2 concentration measurement method .............................................................. 134 A.4 2,4-D concentration measurement method ............................................................. 135 A.5 Incident UV fluence rate measurement method ...................................................... 136 A.6 Kinetic study of UV/H2O2 process for degradation of 2,4-D ................................. 139 A.7 Cyclic voltammetry experiment .............................................................................. 142 A.8 pH change in the cell............................................................................................... 144 A.9 Electric power consumption analysis...................................................................... 145 A.10 Stability and durability of the device ...................................................................... 148 Appendix B ............................................................................................................................. 151 B.1 Average properties of the dilute mixture ................................................................ 151 B.2 H2O2 molecular structure and UV absorption spectrum ......................................... 151 B.3 MB molecular structure and UV-visible absorption spectrum ............................... 152 B.4 Voltage of the system as a function of applied current density and electrolyte flow rate 153 B.5 EEC and H2O2 generation rate as a function of applied current density and electrolyte flow rate .............................................................................................................................. 154 B.6 UV transmittance as a function of concentration of H2O2 and MB ........................ 155 B.7 Key reactions of EF and PEF processes ................................................................. 157 Appendix C ............................................................................................................................. 158 xiii List of Tables Table 2.1 Average properties of the electrolyte used in this study. .............................................. 19 Table 3.1 Properties of microplasma lamps (reported by the manufacturer). .............................. 50 Table 4.1 Physical properties of the bare and coated GDLs. ........................................................ 86 Table A.1 Nominal properties of a single UV-LED used in the UV-LED array. ....................... 133 Table A.2 Pseudo-first-order kinetic data. The error bars represent 95% CIs from three experimental replicates. .............................................................................................................. 141 Table B.1 Average properties of the dilute mixture used as the base solution for most of the experiments. ................................................................................................................................ 151 xiv List of Figures Figure 2.1 (A) The electrochemical photoreactor: different parts and their respective order. Each chamber had a thickness of 2 cm. (B) The experimental setup demonstrates the flow path of water........................................................................................................................................................ 16 Figure 2.2 The effect of flow rate (bottom x-axis) and current density (top x-axis) on the concentration of H2O2 (vertical axis). The inset figure shows the generation rate of H2O2 as a function of flow rate (bottom x-axis) and current density (top x-axis). The error bars represent 95% CIs from three experimental replicates. ........................................................................................ 21 Figure 2.3 The concentration of H2O2 in the cathodic and anodic chambers as a function of the charge density. The error bars represent 95% CIs from three experimental replicates. ............... 24 Figure 2.4 The current efficiency values as a function of charge density showed an average value of 89%. The error bars represent 95% CIs from three experimental replicates. ........................... 26 Figure 2.5 The concentration of electrogenerated H2O2 in the cathodic chamber as a function of applied current density in two cases of open-to-air GDE (red circles) and sealed GDE (black boxes). The error bars represent 95% CIs from three experimental replicates. ............................ 28 Figure 2.6 Electric potential values obtained at various applied currents. The error bars represent 95% CIs from three experimental replicates. ................................................................................ 29 Figure 2.7 The concentration of 2,4-D in different locations of the cell, for various modes of operation: control, UV only, H2O2 only, and UV/H2O2. The error bars represent 95% CIs from three experimental replicates. ....................................................................................................... 31 Figure 2.8 The concentration of 2,4-D in different locations as a function of incident UV fluence rate. The error bars represent 95% CIs from three experimental replicates. ................................ 34 xv Figure 2.9 The degradation percentages as a function of the initial concentration of 2,4-D at two different UV fluence rates. The error bars represent 95% CIs from three experimental replicates........................................................................................................................................................ 36 Figure 2.10 Degradation percentages of 2,4-D as a function of the initial pH. The error bars represent 95% CIs from three experimental replicates. ................................................................ 38 Figure 3.1 The electrochemical photoreactor used in this study. 1: current distributor, 2: carbon cloth cathode, 3: cathodic chamber, 4: PEM, 5: mixed metal oxide anode, 6: anodic chamber, 7: plexiglass separator, 8: UV chamber, 9: quartz disk, and 10: microplasma lamp. ....................... 47 Figure 3.2 The normalized output profiles of the 222 nm (left y-axis) and 220–280 nm (right y-axis) microplasma UV lamps. ....................................................................................................... 49 Figure 3.3 The plot of concentration of H2O2 (right y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at various distances between anode and cathode, shown in the legend. The error bars represent 95% CIs from three experimental replicates. ................................................................................................................ 54 Figure 3.4 The concentration of H2O2 (solid lines, left y-axis) and Faraday efficiency values (dashed lines, right y-axis) as a function of applied charge density (bottom x-axis) at various distances between the anode and the cathode (legend). The inset graph shows the anodic concentration of H2O2 as a function of the cathodic concentration of H2O2. The error bars represent 95% CIs from three experimental replicates. ................................................................................ 56 Figure 3.5 Cathodic (solid lines, left y-axis) and anodic (dashed lines, right y-axis) concentrations of H2O2 as a function of electrolyte type (legend) and concentration (bottom x-axis). The inset plot shows the voltage of the cell obtained for various concentrations of different electrolytes. The error bars represent 95% CIs from three experimental replicates. ........................................................ 60 xvi Figure 3.6 Cathodic (solid lines, left y-axis) and anodic (dashed lines, right y-axis) concentrations of H2O2 as a function of temperature (bottom x-axis). The inset plot shows the cell voltage as a function of electrolyte temperature. The error bars represent 95% CIs from three experimental replicates. ...................................................................................................................................... 62 Figure 3.7 Degradation percentages of MB using direct photolysis, UV/H2O2 process, and EF and PEF processes in the absence of UV and the presence of the three different microplasma UV lamps. The error bars represent 95% CIs from three experimental replicates. ........................................ 64 Figure 3.8 The performance of the EF and PEF processes as a function of the initial Fe2+ concentration. The inset figure shows the degradation of MB as a function of the initial pH in the PEF process. The error bars represent 95% CIs from three experimental replicates. .................. 70 Figure 4.1 The electrochemical reactor used in the current study. The first and second chambers from the left are the anodic and cathodic chambers, respectively. The MMO anode and Nafion membrane were placed between the two chambers (not shown), while the cathode (shown in black) was placed on the air-facing side of the cathodic chamber. ......................................................... 75 Figure 4.2 The FIB-SEM images of the carbon-coated ePTFE-based electrode. (A) and (B) show the surface morphologies at two different magnifications of 2,000 and 100,000, respectively, while (C) and (D) depict cross-sectional images of the same electrode at magnifications of 5,000 and 20,000, respectively. 1: carbon nanoparticles penetrated into the ePTFE structure; 2: ePTFE porous network. The white border in (C) is a platinum strip and was used to protect the underlying material from FIB damage during the milling stage. .................................................................................. 82 Figure 4.3 Contact angle measurement using a goniometer. (A) shows a droplet of water on the bare substrate, while (B) depicts a droplet of water on carbon-coated ePTFE. ............................ 84 xvii Figure 4.4 The results of MIP. (A) Cumulative volume over the entire pressure range; (B) incremental intrusion volume as a function of MIP; (C) cumulative pore area versus pore diameter; (D) differential volume as a function of pore diameter. ............................................................... 89 Figure 4.5 Concentration of H2O2 (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend). The error bars represent 95% CIs from three experimental replicates. ...................................................................................................................................... 92 Figure 4.6 Concentration of H2O2 (solid lines, left y-axis) and current efficiencies (dashed lines, right y-axis) as a function of applied charge density (bottom x-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend). The inset graph shows the anodic concentration of H2O2 versus the cathodic concentration of H2O2, along with the fitted line. The error bars represent 95% CIs from three experimental replicates. ........................................................................................ 94 Figure 4.7 The resultant electric potentials of the cell (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at the various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend). The error bars represent 95% CIs from three experimental replicates. ....................................................................................................... 96 Figure 4.8 EEC (solid lines, left y-axis) and H2O2 generation rate (dashed lines, right y-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend), as a function of (A) applied current density and (B) electrolyte flow rate (bottom x-axis in both cases). The error bars represent 95% CIs from three experimental replicates. ................................................................................ 99 Figure A.1 (A) The UV-LED array used in this study. (B) The normalized relative optical intensity of the UV-LED as a function of wavelength, showing a monochromatic peak at 277 nm and having an FWHM value of 10.7 nm. ...................................................................................................... 132 xviii Figure A.2 Spectra of various solutions: (a) ultrapure water, (b) 25 mM Na2SO4, (c) 1 mg L–1 2,4-D, and (d) the initial solution, containing 1 mg L–1 2,4-D in 25 mM Na2SO4. ........................... 134 Figure A.3 H2O2. (A) Molecular structure. (B) Absorption spectrum of 0.98 mM H2O2 in ultrapure water. ........................................................................................................................................... 135 Figure A.4 2,4-D. (A) Molecular structure. (B) Absorption spectrum of 1 mg L–1 2,4-D in ultrapure water. ........................................................................................................................................... 136 Figure A.5 Quantum yield values of iodide-iodate actinometer reported in the literature, along with the fitted line. .............................................................................................................................. 139 Figure A.6 Batch mode setup used for kinetic studies. (A) Arrangement of various parts of the setup. (B) The five points taken on the surface of the Petri dish to confirm collimation. .......... 141 Figure A.7 Log removal values as a function of time (bottom horizontal axis) and UV fluence (top horizontal axis) for 1 mg L–1 2,4-D and various concentrations of H2O2 shown in the legend. The error bars represent 95% CIs from three experimental replicates. .............................................. 142 Figure A.8 Cyclic voltammetry results for the blank electrolyte (0.1 M Na2SO4) and electrolyte with 500 mg L–1 2,4-D, suggesting a reversible change in the concentration of 2,4-D. ............. 144 Figure A.9 The pH evolution in the system, as a function of applied current density. The error bars represent 95% CIs from three experimental replicates. .............................................................. 145 Figure B.1 H2O2. (A) Molecular structure. (B) UV absorption spectrum. ................................. 151 Figure B.2 MB (methylthioninium chloride). (A) Molecular structure. (B) UV-visible absorption spectrum, showing the highest peak at 664 nm. ......................................................................... 152 Figure B.3 The voltage of the cell (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis), at various distances shown in the legend. The error bars represent 95% CIs from three experimental replicates. .................... 153 xix Figure B.4 The EEC (solid lines, left y-axis) and H2O2 generation rate (dashed lines, right y-axis) as a function of (A) applied current density (bottom x-axis), and (B) electrolyte flow rate (bottom x-axis), at various distances between the anode and the cathode (legend). The error bars represent 95% CIs from three experimental replicates. .............................................................................. 154 Figure B.5 UV transmittances at two wavelengths of 275 nm and 222 nm of solutions containing various concentrations of (A) H2O2 and (B) MB. ....................................................................... 156 Figure C.1 An image of a carbon-coated ePTFE substrate. ........................................................ 158 Figure C.2 Logarithm of differential volume as a function of diameter of the pore for bare GDL and carbon-coated GDL in both intrusion and extrusion modes. ............................................... 159 xx List of Symbols • Chemistry: ℎ𝜈 Radiation ·O2− Superoxide radical ·OH Hydroxyl radical ·OOH Hydroperoxyl radical e- Electron H+ Proton H2O2 Hydrogen peroxide • English letters: A Area cm2 B Geometrical factor to consider the irregularities of the pores CH2O2 Concentration of electrogenerated hydrogen peroxide mM or M CH2O2,t Concentration of H2O2 at time t mM or M C2,4−D,0 Initial concentration of 2,4-D mg L–1 C2,4−D,t Concentration of 2,4-D at time t mg L–1 CMB,i Initial concentration of MB mg L–1 CMB,f Final concentration of MB mg L–1 c Speed of light in vacuum m s–1 E0 Standard potential V Eλ Fluence rate at a wavelength of λ mW cm–2 E̅ Average fluence rate mW cm–2 F Faraday constant C mol–1 h Planck constant J s xxi I Applied current A i Applied cathode current density A m–2 k Second-order kinetic rate constant M–1 s–1 k′ Pseudo-first-order kinetic rate constant s–1 Mw Molar mass g mol–1 NA Avogadro constant E–1 n Number of electrons required for 2e-ORR P Total electrical power for the cell, UV-LEDs and active cooling system kW p(λ) Normalized radiation output ratio at each wavelength Q Flow rate of the electrolyte L s–1 or m3 h–1 rmax Maximum pore radius of the membrane m t Time s or h U Cell potential V Uλ Photon energy per einstein at a wavelength of λ J E–1 V Volume cm3 or L • Greek letters: αλ0 Absorption coefficient at a wavelength of λ before UV exposure cm–1 αλt Absorption coefficient at a wavelength of λ after UV exposure cm–1 γ Interfacial surface tension N m-1 ελ Molar absorption coefficient at a wavelength of λ M–1 cm–1 θ Contact angle of the liquid with the membrane deg. λ Wavelength nm ρe Charge density C L–1 Φλ Quantum yield at a wavelength of λ xxii List of Abbreviations 2,4-D 2,4-dichlorophenoxyacetic acid 2,4-DCP 2,4-dichlorophenol 2e-ORR Two-electron oxygen reduction reaction 4e-ORR Four-electron oxygen reduction reaction ACS American Chemical Society AO Anthraquinone oxidation AOP Advanced oxidation process BET Brunauer–Emmett–Teller BSE Backscattered electron CCD Cosine corrector diffuser CI Confidence interval COD Chemical oxygen demand CV Coefficient of variation CV Cyclic voltammetry DSA Dimensionally stable anode DUV Deep-UV E. coli Escherichia coli EEC Electric energy consumption EEO Electric energy per order EF Electro-Fenton EPA Environmental Protection Agency ePTFE Expanded polytetrafluoroethylene FIB Focussed ion beam FWHM Full width at half maximum GAC Granular activated carbon GDE Gas diffusion electrode GDL Gas diffusion layer xxiii HER Hydrogen evolution reaction HPLC High-performance liquid chromatography KCl Potassium chloride LED Light-emitting diode LEP Liquid entry pressure MAC Maximum acceptable concentration MB Methylene blue MCPCB Metal core printed circuit board MIP Mercury intrusion porosimetry MMO Mixed metal oxide NPTF National Pipe Tapered for Fuels NUV Near-UV OC Open circuit OER Oxygen evolution reaction OFAT One-factor-at-a-time PCB Printed circuit board PDA Photodiode array PEF Photoelectro-Fenton PEM Proton exchange membrane PF Petri factor PP Polypropylene ppm Part per million POU Point-of-use PTFE Polytetrafluoroethylene SCE Saturated calomel electrode SEM Scanning electron microscopy SHE Standard hydrogen electrode TDS Total dissolved solids TOC Total organic carbon UNICEF United Nations Children’s Fund xxiv UV Ultraviolet VUV Vacuum UV WEP Water entry pressure WHO World Health Organization WPE Wall-plug efficiency xxv Glossary UV-A region UV radiation with a wavelength of 315 to 400 nm UV-B region UV radiation with a wavelength of 280 to 315 nm UV-C region UV radiation with a wavelength of 200 to 280 nm xxvi Acknowledgements First and foremost, I would like to express my gratitude to my supervisor, Dr. Fariborz Taghipour, for his phenomenal guidance through the current work. His invaluable advices and innovative ideas were highly impactful in virtually every aspect of this research. Secondly, I would like to thank my thesis examination committee members, Dr. Előd Gyenge and Dr. Yankai Cao, for their precious time attending the examination. Fabulous support of the former and current members of the Photo-reaction Engineering Laboratory is greatly appreciated. Specifically, precious discussions with Dr. Kai Song, Dr. Ataollah Kheyrandish, Dr. Majid Keshavarzfathy, and Dr. Ehsan Espid remarkably helped me shaping the current piece of work. It goes without saying that the tremendous support I received throughout this journey from my beloved family, my mother, father, and brother, was undoubtedly the most crucial factor contributing to my master’s thesis. Without their unconditional love, this journey would have not been feasible. The financial support from Innovative BC is also acknowledged. xxvii Dedication This work is dedicated to all who suffered and sacrificed.1 Chapter 1: Introduction 1.1 Background According to the United Nations, the priority of poor countries should be providing clean water to the population.1 Even though considerable concerns exist regarding water issues in the developing countries, these issues are also present in the developed countries. Currently, Canada is dealing with its long-established water crisis on First Nations reserves and the United States is in the middle of resolving its water crisis in Flint, Michigan.2,3 Growing concern over the waterborne diseases in many remote communities has demanded new water treatment systems to meet their specific needs. Drinking water can be contaminated with various organic compounds, such as pesticides, pharmaceuticals, and dyes. In order to provide clean drinking water, point-of-use (POU) treatment systems can be used, which can be installed under sinks. Home POU water treatment systems have been discussed as “the ultimate barrier” for consumer protection.4 Currently, the POU devices available in the market use physical methods for water purification to remove harmful organics, such as activated carbon adsorption, distillation, ion exchange, particle filters, and reverse osmosis. However, the efficiency of physical methods for water treatment is generally lower than that of chemical methods. Also, the physical methods are considered as economically nonviable options.5 Chemical methods are not suitable for POU systems as they need the addition of chemicals, which could be dangerous, and also increases operational costs.6 2 1.1.1 Advanced oxidation processes Advanced oxidation processes (AOPs) are a set of chemical procedures for water and wastewater treatment to remove organic contaminants by oxidation with hydroxyl radicals (·OH). These strong radical species have a very short lifetime and must be produced right before oxidation. One of the sources for ·OH production in municipal water treatment plants is the exposure of hydrogen peroxide (H2O2) containing influent to ultraviolet (UV) radiation. H2O2 molecules can be homogenously break down to produce two ·OH, by absorbing a UV radiation with a wavelength of lower than 300 nm, as shown in Eq. 1.1. Then, the produced ·OH attack the organic contaminants of water and mineralize them, and the outlet water is a potable water.7 H2O2 + ℎ𝜈 → 2 ·OH Eq. 1.1 The UV/H2O2 AOP is a well-known process, used in the water treatment industry for many years. However, there are certain issues associated with this process, which makes it unsuitable to be utilized in a POU water treatment system. The most significant one is the environmental concerns about the transportation and storage of H2O2 by the end user since it is a strong oxidant.8 The aforementioned problem can be solved by the in situ electrochemical production of H2O2 at low concentrations. H2O2 can be produced from water and air, and then be exposed to the UV radiation to let the homolytic cleavage of H2O2 happen to produce ·OH. Therefore, microbial and organic pollutants of water can be eliminated without the addition of any chemicals. The concentration of H2O2 used in water treatment plants is typically from 6 to 10 mg L–1, and a similar value can be obtained with a POU H2O2 generation system. The efficiency of the electrochemical generation of H2O2 highly depends on the supplement of oxygen (O2) to the aqueous medium. Gas 3 diffusion electrodes (GDEs) are usually used for this purpose since they provide O2 to the solution at a high rate. The second concern about scaling down the traditional UV/H2O2 process is the utilization of bulky UV lamps, which have mercury content and demand high energy consumption. On the other hand, ultraviolet light-emitting diodes (UV-LEDs) and microcavity plasma (microplasma) UV lamps offer several advantages compared to the traditional UV lamps; perhaps most importantly is the flexibility in the reactor design of water purification systems. Because UV-LEDs are small and adjustable sources of radiation, they offer more freedom in the reactor design and optimization compared with UV lamps. Moreover, they are environmentally friendly (no mercury content) and they can be operated with a much lower electrical energy input compared with the UV lamps. Besides, microplasma lamps have certain wavelengths which makes them suitable for water treatment. 1.1.1.1 Electrochemical H2O2 generation In the last decade, a significantly increased attention was drawn toward the electrochemical advanced oxidation processes (EAOPs) for the degradation of organic pollutants from drinking water. These methods are mostly based on the generation of ·OH from water oxidation reaction 6. However, there are a few concerns regarding the practical application of these systems, such as the stability of anodes and the low rate of reaction. Another interesting approach to produce ·OH is the combination of in situ generation of H2O2 with UV-C radiation. This method has several advantages, including elimination of the need to add any chemicals externally, disinfection of 4 water (due to utilization of UV-C radiation), proven stability of the electrodes, and faster reaction rates (due to volumetric reaction, instead of surface reaction used in EAOPs). The principle of electrochemical H2O2 generation is lying on the electro-reduction of oxygen. Dissolved oxygen in water can react with protons (H+) of water in two ways: two-electron oxygen reduction reaction (2e-ORR), which produces H2O2 (Eq. 1.2), and 4e-ORR, which produces water (Eq. 1.3). O2 + 2 H+ + 2 e− → H2O2 E0 = 0.695 V vs. SHE Eq. 1.2 O2 + 4 H+ + 4 e− → 2 H2O E0 = 1.229 V vs. SHE Eq. 1.3 As can be seen, Eq. 1.2 is favourable since it produces the desired product (i.e., H2O2). On the other hand, Eq. 1.3 is the parasitic reaction because it does not lead to H2O2. Since the standard potential of the second reaction is higher than that of the first reaction, from a thermodynamic point of view, Eq. 1.3 is more favoured than Eq. 1.2. Therefore, an appropriate electrocatalyst should be used as the cathode to promote the first reaction and quench the second reaction. Carbon black is a popular catalyst used in the 2e-ORR since it exhibits high catalytic activity for H2O2 production and is highly conductive.9 Several types of cathodes have been tested and reported in the literature for the in situ production of H2O2: carbon felt, graphite, carbon cloth, and GDEs.9–13 A number of studies compared the performance of submerged graphite plates with GDEs for H2O2 generation, and at the end, concluded that the GDEs have a better overall performance,14–16 which is the focus of this study. Some studies have reported electrochemical H2O2 generation from water oxidation or by photocatalytic methods.17–20 Some research have also been done on both modifying the cathode 5 (e.g., introducing metals such as cobalt) and the reactor design (e.g., a fluidized bed, trickle bed, or supplying hydrogen gas to the GDE).21–23 However, the use of a metal-based cathode or a complex reactor design would increase the manufacturing cost and could prohibit the practical use of the H2O2 generation systems. Moreover, H2O2 production in a solid polymer electrolyte fuel cell suffers from some disadvantages, such as a low current efficiency (e.g., 30%), the need to operate at relatively high temperatures (e.g., 60°C), and a low water flow rate (e.g., 5 to 15 mL min-1).24 1.1.1.1.1 GDEs GDEs are porous electrodes with a three-phase boundary interface. They are composed of a gas diffusion layer (GDL), which acts as a support for the catalyst and a pathway for gaseous molecules to pass through, and an electrocatalyst, which catalyzes the desired electrochemical reaction (which is expected to occur on the surface). Every GDE has two sides: one side is in contact with the liquid, which has immobilized catalyst on it, and the other side is in contact with the air (or the gas which is going to diffuse through the GDL). A suitable GDE should be both conductive, to act as the cathode, and hydrophobic, in order to prevent leakage/flooding of the liquid from the liquid-facing side to the air-facing side.25 The main advantage of GDEs compared to submerged electrodes is that they can supply more oxygen to the liquid; since it is porous and can facilitate the passage of oxygen molecules through, thus compensate the depleted oxygen and cause a nearly constant concentration of oxygen in the layer of liquid adjacent to the electrode.26 However, GDEs are not disadvantage-free. The main 6 issue with the commercially available GDEs is their vulnerability to flooding (leaking) when there is a low excess of pressure on their liquid-side over their gas-side.27–30 1.1.1.2 Ultraviolet light-emitting diodes UV-LEDs could revolutionize water treatment systems by replacing UV lamps since LEDs have enhanced dramatically in performance over the past few years and are significantly cheaper. LEDs are semiconductors that emit radiation at a single wavelength. According to this unique feature, they can be designed to produce UV radiation at desired wavelengths for microorganism inactivation or photochemical reaction induction. Several studies have reported the effectiveness of short-wavelength UV-LEDs for Escherichia coli (E. coli) inactivation.31,32 Research has also been conducted on E. coli inactivation by coupling UVA (wavelength 315 to 400 nm) and UVC (wavelength 200 to 280 nm)-LEDs, which has shown a synergistic effect.33 UV-LED photocatalytic reactors have been used for the degradation of phenol and several dyes, such as Rhodamine B.34 UV-LED photoreactors have also been used for the degradation of agriculturally produced contaminants, antibiotics, and endocrine disruptor hormones, with promising results.35 Despite the enormous potential of UV-LED technology, there are only limited studies to date on the application of UV-LEDs for the design of POU water purification systems in the literature. UV-LEDs will play a central role in the future generation of UV photoreactors for various applications, including advanced oxidation processes for water treatment. Water disinfection using large UV lamps is a well-established process in the water treatment industry, however, since the footprint and electric power requirement of UV-LEDs are low, small UV-LED based photoreactors 7 can be developed with high efficiencies by overcoming the limitations of mass transfer and photon transfer in traditional large-scale photoreactors. 1.1.1.3 Microplasma lamps Microplasma lamps are a new flat source of UV radiation which are able to generate UV rays when a potential difference is applied to both sides of the lamp. Based on the gas occupying the gas chamber of these lamps, they generate UV radiation with monochromatic wavelengths of 172 nm (Xe2), 222 nm (KrCl), 308 nm (XeCl), and a polychromatic pattern of 220–280 nm (phosphor).36 These certain wavelengths, in addition to their flat structure, low power consumption, and being mercury free, make them a suitable candidate to be used in POU water treatment systems. 1.2 Research questions This research project was aimed toward designing a novel UV-based POU water purification system through the development of an efficient and compact electrochemical UV photoreactor, operated at low power and practically maintenance-free. The hypothesis was that an integrated UV-LED electrochemical photoreactor can be developed with features suitable for POU degradation of trace organic contaminants in water. In order to evaluate the hypothesis, the following research questions were identified: 1. To what degree an electrochemical cell integrated with UV-LED is capable of generating ·OH for the removal of organic contaminants? What are the key parameters in H2O2 or ·OH generation in such system and how they impact the generation? 8 Recent progress in UV-LEDs has verified their application in water disinfection, however, a particular set of research studies was required to further study UV-LEDs new concepts and reactor design for other applications, including AOPs. 2. How the performance of UV-LEDs is compared to those of different microplasma lamps? The performance of microplasma lamp-based electrochemical photoreactor was evaluated at different operating conditions in order to address this question. 3. How can expanded polytetrafluoroethylene (ePTFE) be utilized in order to develop a new generation of GDEs for 2e-ORR? The hypothesis was that this substrate for the electrode would prohibit the leakage from the GDE, while have a higher catalytic efficiency compared to submerged electrodes, since it acts as an air-breathing substrate. 1.3 Research objectives The idea of this research project was emerged from a recent opportunity made possible by advances in the radiant power of UV-LEDs and microplasma lamps. This research aimed to develop the next generation of POU water purifiers operating with a new type of GDE to produce H2O2, while coupled with UV-LEDs and microplasma lamps for water disinfection and organic contaminant degradation purposes. This goal became feasible by exploring the concepts of UV-based AOPs and the application of those concepts to the development of novel UV-LED and microplasma lamp photoreactors. The three specific objectives of this research project were: 9 1. Develop a portable electrochemical reactor for the in situ generation of H2O2 and examining the parameters that affect the performance; 2. Test the efficacy of the system for organic contaminant degradation, using two newly developed UV sources: the UV-LEDs and microplasma UV lamps; 3. Fabricate a GDE for operating at a pressure of about 70 psi to make the device applicable for practical POU water treatment. 1.4 Thesis outline The structure of this thesis is as follows: • Chapter 1 provides a general overview of the background and objectives of the research; • Chapter 2 presents the results of studying the electrochemical generation of H2O2 and the application of UV-LEDs combined with the electrochemical reactor for the degradation of 2,4-dichlorophenoxyacetic acid (2,4-D); • Chapter 3 describes the effect of design and operational conditions on cathodic H2O2 production, as well as the performance of microplasma UV lamps when coupled with the electrogenerated H2O2 for organic contaminant degradation; • Chapter 4 introduces ePTFE-based GDEs for the production of H2O2, including their characterization and operation; • Chapter 5 presents the detailed conclusions of the research, as well as some recommendations for the future studies. To the best of author’s knowledge, this is the first research project which combines the in situ generation of H2O2 with the newly developed UV sources, the UV-LEDs and microplasma UV 10 lamps, and investigates the related electrochemical and photochemical processes thoroughly. The developed system could offer new approaches toward treating the trace organic contaminants in drinking water at the POU. 11 Chapter 2: Integrated UV-LED hydrogen peroxide electrochemical photoreactor for point-of-use water treatment 2.1 Introduction Advanced oxidation processes (AOPs) are a set of chemical, photochemical, and electrochemical processes that are mainly based on the in situ generation of highly reactive oxidants, such as hydroxyl radicals (·OH), which can unselectively oxidize organic compounds.37 One of the main advantages of AOPs compared with other water treatment methods (e.g., membrane processes) is the degradation of recalcitrant compounds with no generation of a secondary waste stream.38 AOPs include various processes for the generation of ·OH, such as those that are ozone-based and ultraviolet (UV)-based as well as electrochemical AOPs and photocatalysis. Among these processes, the UV-hydrogen peroxide (H2O2) process is well-established for the removal of refractory organic pollutants and is currently operated in drinking water treatment facilities at full-scale.39 H2O2 is a versatile, environmentally friendly compound with no harmful residues, as its decomposition products are oxygen and water.40 The treatment method is based on the photolyzing the H2O2 molecule by UV radiation with a wavelength of 200 to 280 nm (i.e., the UV-C region), which forms two ·OH due to the homolytic scission of the oxygen-oxygen bond of the H2O2 molecule as shown in Eq. 2.1.41 Moreover, the UV irradiation employed in the UV/H2O2 process can effectively inactivate microorganisms and disinfect water, which makes it superior to other processes, as it can be used for both microbial inactivation and chemical contaminant degradation.42 H2O2 + ℎ𝜈 → 2 ·OH Eq. 2.1 12 The UV/H2O2 process is common in large-scale municipal water treatment plants. However, it is not readily feasible to scale down this process for point-of-use (POU) applications due to the costs and hazards related to the transportation, storage, and handling of H2O2. Therefore, the on-site generation of H2O2 has gained attention in recent years. Among the different methods of production of H2O2, which are direct synthesis from hydrogen and oxygen, photocatalysis, and two-electron oxygen reduction reaction (2e-ORR), the latter method is considered the most promising choice since it can be done at moderate pressure and temperature using earth-abundant carbon-based materials as electrodes, as opposed to the traditional anthraquinone method, which requires high pressure and special operational safety considerations.43,44 Carbonaceous cathodes, such as graphite and activated carbon, have several advantages over other materials, including low cost, non-toxicity, high stability, high electrical conductivity, high surface area, and chemical resistance.45,46 The 2e-ORR process is the result of the electrochemical reaction of oxygen, proton, and electron on the surface of the cathode, as illustrated in Eq. 2.2. Carbon-based gas diffusion electrodes (GDEs) are a class of suitable candidates as electrocatalysts for this process since they provide a three-phase boundary to increase reaction rates by favouring the diffusion of oxygen to the reaction zone so that low solubility of oxygen in water can be easily overcome.45–47 This class of electrodes is composed of two layers: the gas diffusion layer (GDL), which is in contact with the ambient air or pure oxygen and has the role of supplying reactant oxygen to the reaction zone on a diffusion basis, and the active layer, which is in contact with the electrolyte and allows the oxygen passed through the GDL to encounter the other reactants (i.e., proton and electron) and thus produce H2O2 on the surface of the electrocatalyst.40 On the other hand, on the surface of the anode, the oxygen 13 evolution reaction (OER; Eq. 2.3) takes place, which produces protons which are one of the reactants in Eq. 2.2. The overall reaction, shown in Eq. 2.4, clearly demonstrates that H2O2 is generated from oxygen and water using an electrochemical system. O2 + 2H+ + 2e− → H2O2 E0 = 0.67 V vs. SHE Eq. 2.2 H2O →12O2 + 2H+ + 2e− E0 = −1.23 V vs. SHE Eq. 2.3 12O2 + H2O → H2O2 E0 = −0.56 V vs. SHE Eq. 2.4 Ultraviolet light-emitting diodes (UV-LEDs) are a new source of UV radiation based on semiconductors that can emit UV rays at a single, tunable wavelength by adjusting the composition of gallium nitride (GaN) and aluminum nitride (AlN) to obtain aluminum gallium nitride (AlXGa1-XN) alloy.48 In contrast to traditional UV lamps, which suffer from some disadvantages, such as having a cylindrical shape, long warm-up time, and high electrical power input, UV-LEDs have a small footprint and instant on-off ability, and they can be operated with low input power, so they are recommended for POU water treatment photoreactors.49,50 Indeed, UV-LEDs have been studied for both water disinfection and organic pollutant degradation with promising results.51,52 2,4-Dichlorophenoxyacetic acid (2,4-D) is a chlorinated phenoxy compound that is used as an herbicide to control broadleaf weeds. Due to its widespread use (currently third most commonly used herbicide in the world) and high solubility in water, 2,4-D and its residues leach from the root zone and have been found in surface and ground waters.53,54 It is considered a carcinogenic agent that affects the heart, liver, and central nervous system.55 14 To tackle issues associated with the POU UV/H2O2 process mentioned above, this study reports the development of a novel water treatment system enabled by the in situ generation of H2O2 which is coupled with UV-LEDs in a single unit. The electrochemical, on-site generation of H2O2 was performed on the surface of an air-breathing cathode, which is a carbon-based GDE. The electrochemical H2O2 generation was studied in a continuous flow-through, single-pass reactor. The degradation of 2,4-D, as a model organic contaminant, was then performed in the electrochemical UV-LED photoreactor. The effect of several operating parameters, such as UV fluence rate, initial concentration of 2,4-D, and initial pH, was investigated. To the best of our knowledge, this is the first study to combine the in situ electrochemical generation of H2O2 with UV-LEDs as the radiation source. The application of UV-LED as the radiation source in an electrochemical cell offers much flexibility in terms of the design, system integration, and operating conditions. The integrated system developed in our research showed promising results in chemical contaminant degradation, and it could be applied to both chemical and microbial water treatment at the POU. 2.2 Materials and methods 2.2.1 Materials All of the following materials, which were supplied by Sigma-Aldrich, were of American Chemical Society (ACS) reagent grade or higher and used with no further purification: 2,4-D (Cl2C6H3OCH2CO2H), sodium sulfate (Na2SO4), acetonitrile (CH3CN), acetic acid (CH3CO2H), hydrogen peroxide solution (H2O2, 30% w/w in H2O), potassium iodide (KI), potassium iodate (KIO3), ammonium molybdate tetrahydrate ((NH4)6Mo7O24.4H2O), sodium tetraborate decahydrate (Na2B4O7.10H2O, also known as borax), sulfuric acid (H2SO4), sodium hydroxide 15 (NaOH), chlorohydroquinone (ClC6H3-1,4-(OH)2), and 2,4-dichlorophenol (Cl2C6H3OH, shortened as 2,4-DCP). 2.2.2 Electrochemical photoreactor Different sections of the electrochemical photoreactor and the experimental apparatus are illustrated in Figure 2.1(A) and (B), respectively. The electrochemical photoreactor used in this study consisted of three chambers: the cathode, UV, and anode chambers. The cathodic chamber was equipped with a cut of carbon cloth (GDL-CT; FuelCellsEtc, College Station, TX, USA) serving as the cathode with an active area of 17 cm2. A cut of stainless steel wire cloth (85385T39; McMaster-Carr, Elmhurst, IL, USA) was placed on the air-facing side of the cathode to increase the mechanical strength of the GDE, as well as to act as the electric current distributor. The UV chamber consisted of an array of 12 UV-LEDs (DF7ZV; DOWA Electronics Materials Co., Ltd., Tokyo, Japan) as the source of UV radiation (description of the UV-LED array preparation is provided in the next section), where the LEDs were separated from the chamber by a quartz polished disk of 90% UV transmittance at 277 nm with a diameter of 10.16 cm and a thickness of 0.32 cm (G.M. Associates, Inc., Oakland, CA, USA). The inner walls of the UV flow channel were made of quartz plate to minimize photon loss in the system. The anodic chamber utilized a titanium mesh coated with mixed metal oxide (MMO) RuO2/IrO2, which had a length of 50 mm, a width of 50 mm, and a thickness of 1 mm (William Gregor Ltd, London, UK), as the anode. Each chamber was made of plexiglass and had an effective volume of 34 cm3. A cut of proton exchange membrane (PEM; Nafion™ 117; Fuel Cell Store, College Station, TX, USA) was placed next to the anode, between the cathodic chamber and the anodic chamber, to allow the diffusion of electrogenerated protons at the anode toward the cathode. Cuts of gasket were used to seal the 16 system. The cell consisted of various components sandwiched between a series of frames which were bolted together. Figure 2.1 (A) The electrochemical photoreactor: different parts and their respective order. Each chamber had a thickness of 2 cm. (B) The experimental setup demonstrates the flow path of water. 2.2.2.1 UV-LED array preparation Twelve UV-LEDs with a nominal wavelength of 275 nm were used in this study, each with a radiant power output of 75 mW, totaling 0.9 W overall. Further discussion about the UV-LED array is provided in Section A.1. 17 One of the considerations when working with UV-LEDs is appropriate thermal management, as it will affect the radiant power of UV-LEDs.49 UVC-LED output power is a function of the junction temperature. To insure proper heat dissipation, each UV-LED was soldered to a copper base printed circuit board (PCB; Nichia 3535; Shenzhen Deshengxing Electronics Co., Ltd., Shenzen, China) employing thermally stable solder paste (TS391SNL10; Chip Quik Inc., Niagara Falls, NY, USA) and by following the soldering profile recommended by the manufacturer.56 Further, the metal core printed circuit boards (MCPCBs) were connected to a 6 cm × 6 cm copper-based heatsink (342945; Aavid, Thermal Division of Boyd Corporation, Santa Barbara, CA, USA) by applying a thin layer of silicone thermal grease (S606C; T-Global Technology, Lutterworth, UK) and were fixed using thermal adhesive (AATA-5G; Arctic Silver, Visalia, CA, USA) around the MCPCBs. The heatsink was then connected to a DC brushless fan (9GV0612P1G031; Sanyo Denki America Inc., Torrance, CA, USA) to promote convection heat transfer. 2.2.2.2 Nafion membrane preparation The Nafion 117 PEM was pretreated by boiling in 3% v/v H2O2 (diluted with ultrapure water), ultrapure water, and 0.5 M H2SO4 (diluted with ultrapure water), respectively, each step for 30 min and rinsing with ultrapure water between each step, similar to a method described in the literature.57 2.2.3 Experimental procedure In this study, ultrapure water with a dilute electrolyte (25 mM Na2SO4) was used to simulate natural and tap water with their typical conductivity. The selection of this electrolyte was due to the fact that it is a common compound used to study 2e-ORR; however, we expect that other 18 electrolytes and water with relatively high conductivity will give a similar result. Solutions of 25 mM Na2SO4 were prepared by dissolving 3.551 g Na2SO4 powder in ultrapure water (electrical resistivity: 18.2 MΩ∙cm, obtained from a PURELAB Option-Q water purification unit) by stirring at 500 rpm for 10 min followed by 1,000 rpm for 5 min. These solutions were used for the H2O2 generation experiments. A stock solution of 500 mg L–1 (also referred to as 500 part per million (ppm), equivalent to 2.26 mM) of 2,4-D was prepared by dissolving 0.5 g of 2,4-D crystal in 1 L ultrapure water and was kept in a dark place. Solutions of degradation experiments consisted of aliquots of the 2,4-D stock solution diluted in 25 mM Na2SO4 to obtain the required initial concentration of 2,4-D, which was 1 mg L–1, unless otherwise specified. Each solution was passed through the continuous photoreactor using a Masterflex peristaltic pump, equipped with a model 7518-00 pump head (Cole-Parmer Canada Company, Montreal, Canada). The influent was first passed through the cathodic chamber, then to the UV chamber, and finally to the anodic chamber. Samples were taken from each chamber after 10 cycles of the solution had passed through the cell, that is, 20 min in base condition flow rate (50 mL min–1), and were further analyzed for various parameters, including concentration of 2,4-D, concentration of H2O2, and pH. Each experiment was conducted at the constant current mode (galvanostatic condition) by means of a DC power supply (FPS-325DU; Zhao Xin, Shenzen, China), and the resultant voltage was recorded. The pH adjustments were made by adding either NaOH or H2SO4. Tests were performed at room temperature (23.0 ± 1.2 °C) and humidity levels of 46 ± 8%. Table 2.1 shows the average properties of the dilute supporting electrolyte used in this study, measured by a UV-visible spectrophotometer (Cary 100; Agilent Technologies, Inc., Santa Clara, CA, USA) and a multiparameter tester (WD-35634-35 PCTSTestr 50; Oakton, Vernon Hills, IL, USA). Absorption spectra of different 19 solutions, including the dilute electrolyte with and without the organic model, are shown in Section A.2. Table 2.1 Average properties of the electrolyte used in this study. Solution Temperature (°C) pH Conductivity (mS cm–1) TDS (ppt) Salinity (ppt) UV277 transmittance (%) 25 mM Na2SO4 23.7 7.47 4.97 3.52 2.57 97.0 Unless otherwise specified, experiments were performed in the base condition: a flow rate of 50 mL min–1 and an applied current density of 134 A m–2. Each set of experiments was performed at least three times, and the error bars shown in figures and tables represent a confidence interval (CI) of 95% (equivalent to a significance level of α = 0.05), using Student’s t-distribution for the samples. The details of the protocols used for measuring H2O2 and 2,4-D concentration, as well as incident UV, are provided in Sections A.3 to A.5. 2.3 Results and discussion The device performance was evaluated for both cathodic reduction of oxygen to H2O2 and organic contaminant degradation when coupled with UV-LEDs. For each set of experiments, the system was characterized according to operational conditions, such as current density and influent flow rate for H2O2 generation, and the effect of UV fluence, initial concentration of 2,4-D, and pH of 20 the influent for 2,4-D degradation. In the end, an electric power consumption analysis was performed, and the longevity of the system was assessed. 2.3.1 Performance of the cell for the in situ cathodic generation of H2O2 Two operating parameters affecting the generation of H2O2, for an identical liquid composition are the flow rate of the electrolyte passing through the cell, which dictates the residence time, and the applied current density, which determines the concentration of electrons as reactants (Eq. 2.2). These two parameters were studied separately to find the effect of each on the concentration of electrogenerated H2O2. At a given current density of 134 A m–2, the concentration of H2O2 increased with a decrease in flow rate (Figure 2.2). This behaviour was expected since, at lower flow rates, higher residence times can be achieved at a constant volume. Moreover, the response of flow rate can be separated into two parts: high (65 to 95 mL min–1) and low (35 to 65 mL min–1) flow rates. At high flow rates, the reaction seemed to be somewhat mass-transfer limited (diffusion-controlled), as there was less time for oxygen molecules to pass through the GDL and less time for the product (H2O2) to diffuse to the bulk of the fluid, resulting in a higher slope of the operating line. At low flow rates, the reaction appeared to be largely charge-transfer limited (kinetic-controlled), since the relatively high concentration of electrogenerated H2O2 promoted mass transfer, leading to a lower slope of the operating line. The Reynolds numbers were 37.5 and 101.5 at flow rates of 35 and 95 mL min–1, respectively. The normalized result of H2O2 generation rate seemed independent of the electrolyte flow rate, as shown in the inset figure of Figure 2.2. This could be attributed to contradictory effects of better mixing and higher H2O2 decomposition rate on H2O2 generation rate at elevated flow rates. 21 Figure 2.2 The effect of flow rate (bottom x-axis) and current density (top x-axis) on the concentration of H2O2 (vertical axis). The inset figure shows the generation rate of H2O2 as a function of flow rate (bottom x-axis) and current density (top x-axis). The error bars represent 95% CIs from three experimental replicates. In regard to the effect of applied current density, as shown in Figure 2.2, with an increase in current density at a constant volumetric flow rate of 50 mL min–1, a higher generation of H2O2 was observed, suggesting that electrons were the limiting reactant in this case. Increasing current density would accelerate electron transfer on the surface of the GDE, promoting 2e-ORR.45 Moreover, by increasing the applied current, the electric potential can be increased, so that a stronger electric field would be established. This electric field forces more protons, generated on 22 the surface of the anode, to migrate toward the catholyte. Further, at higher currents, the anodic reaction, which is OER (Eq. 2.3) is promoted so that more protons are generated. The generated protons will then migrate to the catholyte, due to the electric field, and supplement the protons consumed in 2e-ORR (Eq. 2.2).58 However, at a high current density of 334 A m–2, a small deviation from linear behaviour was observed, which was likely due to the activation of side reactions (Eq. 2.5 to Eq. 2.8) caused by higher overpotentials.45,58–60 Eq. 2.8 is referred to as 4e-ORR, which converts ambient oxygen to water instead of H2O2. H2O2 + H+ + e− → H2O + ·OH Eq. 2.5 H2O2 + 2H+ + 2e− → 2H2O E0 = 1.78 V vs. SHE Eq. 2.6 2H+ + 2e− → H2 E0 = 0.00 V vs. SHE Eq. 2.7 O2 + 4H+ + 4e− → 2H2O E0 = 1.23 V vs. SHE Eq. 2.8 Overall, in the base condition, the rate of production of H2O2 in the cathodic chamber was 7.94 mg cm–2 h–1. This value is higher than most of the values reported to this point in the literature (see Zhou et al., 201944), likely due to the superior reactor design. The inclusion of the parallel channels in the chambers provided a desirable mixing of the fluid, which promoted the diffusion of H2O2 to the bulk of the solution. It is worth mentioning that, unlike many H2O2 generation systems reported previously,44 the system we developed relied only on the natural diffusion of atmospheric oxygen from ambient air through the GDE and did not utilize any aerator or pump to supply oxygen to the electrolyte. This is beneficial for lowering the total electric energy consumption (EEC). 23 The combination of the two aforementioned parameters (i.e., electrolyte flow rate and applied current density) yielded the definition of charge density, which was the amount of electric charge per unit volume, and can be shown as Eq. 2.9. Plotting the concentration of H2O2 as a function of charge density showed a linear relationship, as illustrated in Figure 2.3, suggesting that there was a linear trade-off between the input electric energy (current) and influence flow rate; that is, at higher flow rates, higher currents needed to be applied in order to generate the same concentration of H2O2. Moreover, the removal of excess H2O2 was observed in the anodic chamber, likely due to Eq. 2.10 to Eq. 2.12.59,60 Eq. 2.10 can occur in the bulk of the fluid anywhere in the system, while Eq. 2.11 (one-step oxidation of H2O2) and Eq. 2.12 (two-step oxidation of H2O2) can only happen on the surface of the anode. H2O2 oxidation on the surface of anode is beneficial for drinking water treatment purposes, as it decreases the concentration of H2O2 in the effluent of the system, which is potable water. ρq =IQ Eq. 2.9 In Eq. 2.9, I (A) is the applied current, Q (L s–1) is the flow rate of the electrolyte inside the cell, and ρq (C L–1) is the charge density. 24 Figure 2.3 The concentration of H2O2 in the cathodic and anodic chambers as a function of the charge density. The error bars represent 95% CIs from three experimental replicates. H2O2 → H2O +12O2 Eq. 2.10 H2O2 → HO2· + H+ + e− Eq. 2.11 H2O2 → O2 + 2H+ + 2e− Eq. 2.12 25 Eq. 2.13 and Eq. 2.14 show the linear relationship between the concentration of H2O2 (mM) and charge density (C L–1) in the cathodic and anodic chambers, respectively, which was previously demonstrated in Figure 2.3. R2 values in both cases were higher than 0.998. CH2O2,cathode = 0.0046 × ρq Eq. 2.13 CH2O2,anode = 0.0039 × ρq Eq. 2.14 The current efficiency (also referred to as Faraday efficiency or coulombic efficiency) indicates the efficacy of electron transfer to promote a desired electrochemical reaction and is defined as the ratio of theoretically required charge for the favoured reaction to the experimentally applied charge, as illustrated in Eq. 2.15. For a wide range of different charge density values from 100 C L–1 to 700 C L–1, an average current efficiency of 89% was obtained (Figure 2.4), which confirmed the capability of the system for efficient in situ cathodic production of H2O2. 26 Figure 2.4 The current efficiency values as a function of charge density showed an average value of 89%. The error bars represent 95% CIs from three experimental replicates. Current efficiency (%) =n F CH2O2 VI t× 100% Eq. 2.15 In Eq. 2.15, n (= 2) is the number of electrons required for 2e-ORR, F (96,487 C mol–1) is the Faraday constant, CH2O2 (mol L–1) is the concentration of electrogenerated H2O2 in the cathodic chamber, V (L) is the volume of catholyte, I (A) is the applied current, and t (s) is the hydraulic residence time of the catholyte in the cathodic chamber. 27 To evaluate the efficiency of the GDE to bring oxygen to the reaction zone, a set of experiments was performed in two conditions: open-to-air GDE and sealed GDE. The results, shown in Figure 2.5, suggest that for a sealed system, oxygen is the limiting species, as a very low generation of H2O2 was achieved. Moreover, at higher current densities, lower generation was observed, due to the higher activation of side reactions (Eq. 2.5 to Eq. 2.7). On the other hand, for the open-to-air GDE, a linear relationship between the concentration and the applied current density was a confirmation that electrons are the limiting species. Higher generation of H2O2 when the cathode is open to air, compared with a sealed cathode, is in agreement with the results for similar types of electrodes reported in the literature (e.g., Pozzo et al., 200514). 28 Figure 2.5 The concentration of electrogenerated H2O2 in the cathodic chamber as a function of applied current density in two cases of open-to-air GDE (red circles) and sealed GDE (black boxes). The error bars represent 95% CIs from three experimental replicates. The resultant electric potential of the system as a function of applied current showed a linear relationship (Figure 2.6). This proves that the system follows Ohm’s law with a constant resistivity, which is proportional to the concentration of protons migrating (due to the potential gradient) and diffusing (due to the concentration gradient) from the surface of the anode through the PEM toward the cathode. 29 Figure 2.6 Electric potential values obtained at various applied currents. The error bars represent 95% CIs from three experimental replicates. 2.3.2 Effect of UV, H2O2, and UV/H2O2 on degradation of 2,4-D To identify the concentration of H2O2 suitable for degradation studies, a batch mode setup, shown in Section A.6, was used to estimate the apparent reaction rate constants. Based on the discussion and values shown in Section A.6, an H2O2 generation concentration of 1.2 mM was selected, as it has a relatively high apparent rate constant, while it does not lead to an excessive residual concentration of H2O2 in the effluent. Moreover, it is preferable to avoid high concentrations of H2O2, because it decreases the UV transmittance of the solution. The strong dependence of reaction 30 rates on the initial concentration of H2O2 has been reported.61 However, it should be noted that the addition of excess H2O2 could lead to less efficient oxidation, as due to the scavenging effect of excess H2O2, it consumes the ·OH available to degrade 2,4-D according to Eq. 2.16 and Eq. 2.17.61,62 The apparent pseudo-first-order kinetic rate constants obtained by UV/H2O2 method were significantly higher than those of the photocatalytic method.63 This can be explained by the fact that H2O2 molecules are soluble in water so that a volumetric reaction can occur, while only surface reaction (with generally less reaction zone) can take place over immobilized photocatalysts. ·OH + H2O2 → HO2· + H2O Eq. 2.16 HO2· + ·OH → H2O + O2 Eq. 2.17 The effect of different modes of operation on the removal of 2,4-D at various sampling locations (i.e., the influent, after the cathodic chamber, after the UV chamber, and after the anodic chamber) is shown in Figure 2.7. In the control experiment, both the UV radiation source and the electrochemical cell were turned off. For UV photolysis only (without the electrogeneration of H2O2), a slight degradation of 5.7% was observed, which was insignificant. UV radiation is known to induce a substitution reaction in which hydrogen atoms will replace chlorine atoms in chlorinated aromatics.61 When the electrochemical cell was on, without any UV radiation employed, a reversible change in the concentration of 2,4-D was observed: a decrease in the cathodic chamber and an increase in the anodic chamber. Finally, when both UV radiation and in situ generation of H2O2 were in effect, a reduction of 46.4% was observed at a fluence rate of 2.78 mW cm–2. The significant increase in the conversion of 2,4-D can be attributed to the production of ·OH, which will increase the apparent fluence-based rate constant greatly compared with direct UV photolysis alone. It has been reported that a 20- to 100-fold increase in degradation rate, 31 compared with that of UV photolysis, could be achieved by the addition of optimized H2O2 concentration.64,65 Figure 2.7 The concentration of 2,4-D in different locations of the cell, for various modes of operation: control, UV only, H2O2 only, and UV/H2O2. The error bars represent 95% CIs from three experimental replicates. The observed reversible change in the concentration of 2,4-D after the cathodic and anodic chambers was likely due to the reversible reduction and oxidation of this organic molecule. This was confirmed by observing no degradation of 1 mg L–1 2,4-D with 1.2 mM H2O2 in the absence of UV radiation. This result is in agreement with a reported study,62 where in dark conditions only 32 1% of 2,4-D was converted after 5 h of reaction. It is suggested, however, that in acidic or alkaline media, H2O2 can oxidize 2,4-D according to Eq. 2.18 and Eq. 2.19.66 HO2− + H+ + R → RO + H2O Eq. 2.18 H2O2 + R → RO + H2O Eq. 2.19 where R denotes the organic compound, 2,4-D. A cyclic voltammetry experiment was performed to assess the oxidation and reduction peaks as described in Section A.7. A slow electron transfer behaviour along with the incapability of H2O2 to oxidize 2,4-D suggests a chemically reversible reaction. In fact, the electroreduction of organic compounds is known to be reversible under some conditions.67,68 In addition, it has been suggested that 2,4-D can undergo reversible reduction and oxidation.69 The analysis of the effluent of the UV-electrochemical cell indicated the generation of a less harmful chemical. The product was 2,4-DCP, which is less harmful than 2,4-D. Canadian drinking water quality guidelines state that the maximum acceptable concentration (MAC) for 2,4-D in drinking water is 0.1 mg L–1, whereas that of 2,4-DCP is 0.9 mg L–1.70 Therefore, it can be claimed that the organic contaminations were converted to their less harmful forms by oxidation with ·OH. The formation of 2,4-DCP as the main by-product of 2,4-D degradation is in agreement with the results reported in the literature.71 Regarding the pH effect, the cathodic reaction (Eq. 2.2) consumes protons, so that the pH value rises in the cathodic chamber. UV radiation converts a portion of H2O2 into ·OH, so no change in the pH value is expected. Finally, at the anode, OER (Eq. 2.3) occurs, leading to the generation of 33 protons as the main cause of a decrease in the pH. Further discussions on the pH evolution in the system are provided in Section A.8. 2.3.3 Effect of incident UV fluence rate on degradation of 2,4-D One of the features of UV-LEDs is the adjustability of their optical output by changing the electric current. In order to find the effect of the UV fluence rate in continuous flow mode, the H2O2 concentration and 2,4-D conversion were studied at different fluence rates. The H2O2 concentration was tracked along the flow of water for the two different fluence rates of 2.78 and 4.78 mW cm–2, where 14.1% and 26.6% of the in situ formed H2O2 molecules were dissociated, respectively, after the UV chamber. The concentration of 2,4-D at different locations as a function of incident UV fluence rate was monitored at the different fluence rates of 0.37, 0.74, 1.48, 2.78, and 4.78 mW cm–2 (Figure 2.8). The average degradation values, based on the UV chamber’s influent and effluent streams, were 2.26%, 12.13%, 20.70%, 33.83%, and 52.77%, respectively, while the total degradation values, according to the initial solution and anode effluent, were 4.08%, 15.69%, 31.59%, 48.57%, and 71.00%, respectively. The trend observed in Figure 2.8 suggests that low UV fluence rates result in inefficient conversion of the initial form of molecule to its by-products, and that reversible reaction on the surface of anode increases the 2,4-D concentration. We found that when the concentration of 2,4-D in the anodic chamber was relatively high (i.e., less ·OH was generated in the UV chamber, or the concentration of 2,4-DCP was low), a more reversible conversion of 2,4-DCP to 2,4-D could occur. However, above a certain value of fluence rate (1.20 mW cm–2), more ·OH generation led to more 2,4-DCP formation and a lower concentration of 2,4-D. An 34 explanation can be that, at higher fluence rates, total mineralization will happen, converting the initial form of molecule to carbon dioxide (CO2), water (H2O), and hydrochloric acid (HCl), so that higher fluence rates lead to higher Cl- generation (due to higher mineralization of 2,4-D), which will further oxidize the organic molecule on the surface of anode. Moreover, higher fluence rates will mineralize 2,4-DCP as well, so that there is less 2,4-DCP available to convert back to 2,4-D. This was observed in the peak area changes shown in chromatographs (the results are not shown). Figure 2.8 The concentration of 2,4-D in different locations as a function of incident UV fluence rate. The error bars represent 95% CIs from three experimental replicates. 35 The results of the continuous mode at different fluence rates (Figure 2.8) agreed well with the batch results (shown in Section A.6), given that the concentration of electrogenerated H2O2 was close to 1.2 mM. 2.3.4 Effect of initial concentration of 2,4-D on degradation of 2,4-D The removal percentages of 2,4-D as a function of the initial concentration for the two fluence rates of 2.78 mW cm–2 and 4.78 mW cm–2 are shown in Figure 2.9. As expected, higher initial concentrations yielded lower percentage degradations, given that there were fewer ·OH for each molecule to react with. The decrease in degradation percentage with an increase in the initial concentration is in agreement with the data available in the literature (e.g., Adak et al., 201965). Based on the Beer-Lambert law and according to the data shown in Sections A.3 and A.4, the molar absorption coefficients for 2,4-D and H2O2 at a wavelength of 277 nm were calculated to be 4,796 and 9.8 M–1 cm–1, respectively. Therefore, higher 2,4-D concentrations will lead to more absorption of the radiation by their molecules, and H2O2 molecules have fewer UV rays to absorb, which will further cause homolytic cleavage of the oxygen-oxygen bond, leading to ·OH formation. Based on the results shown in Figure 2.9, we propose an empirical equation, Eq. 2.20, to predict the degradation percentages in the electrochemical photoreactor at various contaminant initial concentrations and UV fluence rates. Degradation (%) = −7.67 × Ln(C2,4−D,0) + 12.67 × fluence rate + 10.66 Eq. 2.20 where C2,4−D,0 (mg L–1) is the initial concentration of 2,4-D, and fluence rate (mW cm–2) is the UV fluence rate of the UV-LED array at a wavelength of 277 nm. This equation is valid for the 36 range shown in Figure 2.9 under the base condition. The values of degradation percentage predicted by the equation are in agreement with the results for various UV fluence rates reported in Section 2.3.3. Figure 2.9 The degradation percentages as a function of the initial concentration of 2,4-D at two different UV fluence rates. The error bars represent 95% CIs from three experimental replicates. 2.3.5 Effect of initial pH on degradation of 2,4-D The effect of initial pH on the removal of 2,4-D was studied at an initial 2,4-D concentration of 1 mg L–1 and fluence rate of 2.78 mW cm–2 (Figure 2.10). The pH value of 2.59 had the highest removal percentage, possibly due to the higher generation of H2O2. The kinetics for 37 electroreduction of oxygen to H2O2 is faster in acidic conditions compared with alkaline conditions,40 so that, generally at lower pH values, higher H2O2 molecules are available for ·OH generation. As can be seen in Eq. 2.2, lower pH values, which are equivalent to higher concentrations of protons, are more favourable for the electro-generation of H2O2. However, it should be noted that a very high concentration of protons will promote the hydrogen evolution reaction (HER; Eq. 2.7) and reduce the current efficiency,58 leading to less generation of H2O2. This may explain the lower degradation at pH 1.46 compared with pH 2.59. Highly acidic conditions can also lead to a low rate of H2O2 photodecomposition to ·OH.61 In addition, lower degradation at pH 1.46 could be due to the dissociation of H2O2 in highly acidic media. Highly acidic conditions will lead to the reaction of excess protons with H2O2, forming water (Eq. 2.6) and promoting 4e-ORR, which again forms water as the product (Eq. 2.8).45 38 Figure 2.10 Degradation percentages of 2,4-D as a function of the initial pH. The error bars represent 95% CIs from three experimental replicates. At higher pH values, the concentration of ·OH decreases in the UV/H2O2 process.65 Lower degradation at higher pH values could be due to significant H2O2 self-decomposition to water and oxygen at alkaline conditions, which is a barrier for the formation of ·OH.61 During UV irradiation at a high pH, the chloro group on the aromatic rings can be replaced by OH- groups, a process called photohydrolysis. This mechanism is slower than a direct attack on organic molecules by free radicals, which is the dominant mechanism at a low pH.61 Moreover, in alkaline conditions, the HO2− ion is predominant which will decompose H2O2, according to Eq. 2.21.45,46 39 O2 + H2O + 2e− → HO2− + HO− E0 = −0.065 V vs. SHE Eq. 2.21 Furthermore, in basic conditions, Eq. 2.22 to Eq. 2.24 could occur, producing ·OH and superoxide (·O2−) on the surface of the cathode, which can degrade the organic compound.72 O2 + 2H2O + 2e− → H2O2 + 2OH− Eq. 2.22 H2O2 + OH− → H2O + HO2− Eq. 2.23 H2O2 + HO2− → ·OH + ·O2− + H2O Eq. 2.24 Among the pH values studied here, the pH value of 2.59, which is close to the pKa value of the organic molecule (2.81), resulted in the highest degradation of 2,4-D. This is in agreement with the reported values in the literature (for example, see Trivedi and Mandavgane, 201873). 2.3.6 Energy and stability analyses of the system Thorough discussions of the electric power consumption of the system, as well as the durability of the device, are provided in Sections A.9 and A.10, respectively. 2.4 Conclusions In this study, a POU electrochemical device was developed with the aim of in situ generation of H2O2 using ambient oxygen and a dilute supporting electrolyte. This would help to produce H2O2 on demand in a safe, inexpensive, and green way. The electrochemical system was characterized based on the applied current density and fluid flow rate. A cathodic H2O2 generation rate of 7.94 mg cm–2 h–1 was obtained at a current density of 134 A m–2 and flow rate of 50 mL min–1. A high current efficiency of ~90% proved the capability of the system for efficient cathodic production of 40 H2O2. The system was coupled with an array of high-power UV-LEDs in a single unit to degrade 2,4-D. At a flow rate of 50 mL min–1, more than 70% of the initial contaminant was degraded. The efficacy of the system for the degradation of organic pollutants can be enhanced by increasing the UV fluence rate, providing more electrical current, and operating at an optimal pH value. Moreover, an energy analysis showed that the system is energy efficient and suitable for application as a POU water treatment system. High stability of the system could be expected using carbon-based cathodes and MMO anodes. Overall, this system showed promising results that can be used as a chemical-free advanced oxidation and disinfection system for the simultaneous degradation of toxic chemicals and inactivation of pathogenic microorganisms. 41 Chapter 3: Microplasma ultraviolet radiation integrated with electrochemical in situ generation of hydrogen peroxide for degradation of organic pollutants in water 3.1 Introduction Safe drinking water is everyone’s right. However, due to the growing population and industrialization of societies, many water sources are being contaminated with various types of pollution, including physical impurity (e.g., sediments), chemical contamination (e.g., organic compounds, such as pesticides and dyes), and biological pollution (e.g., pathogenic viruses and bacteria). Different water treatment methods have been developed to decontaminate the aforementioned types of impurities. For example, filtration is used to remove sediments, and ultraviolet (UV) radiation is employed in order to inactivate pathogens. Chemical contamination of water is considered one of the most persistent types of water pollution. Chemical impurity can be divided into several categories, in which organic compounds account for a large portion. To mineralize organic contaminants of water, a set of processes, called advanced oxidation processes (AOPs), have been developed and are well established in both research and practice.74,75 These processes are based on the oxidation of organic molecules with strong radicals, commonly hydroxyl radical (·OH) among a few others, which can non-selectively convert the organic molecule to a less harmful, mineralized form and carbon dioxide (CO2). The most common reaction mechanisms are addition reactions, hydrogen abstraction reactions, and electron transfer reactions.76 Some of the most prevalent methods for the in situ generation of ·OH 42 are by exposure of low concentrations of hydrogen peroxide (H2O2) to UV radiation with a wavelength of 200 nm to 300 nm (i.e., UV-C region and part of the UV-B region) (Eq. 3.1), direct exposure of water to UV rays with a wavelength of 100 nm to 200 nm (i.e., vacuum UV (VUV) region) (Eq. 3.2), and the Fenton process (using ferrous ion, Fe2+, as the catalyst) (Eq. 3.3).76–78 The difference between the Fenton process and the electro-Fenton (EF) process is that, in the former, H2O2 is externally added to the system, while, in the latter, H2O2 is generated in situ via electrochemical methods. H2O2 + ℎ𝜈 → 2 ·OH Eq. 3.1 H2O + ℎ𝜈 → ·OH + ·H Eq. 3.2 2 H2O2 + Fe2+ → ·OH + ·OOH + H2O + Fe2+ Eq. 3.3 Traditional UV sources include low-pressure and medium-pressure mercury vapour lamps, which can emit UV radiation with wavelengths of 184.45 nm and 253.7 nm (low-pressure lamps), and a polychromatic pattern (medium-pressure lamps). The wall-plug efficiency (WPE; also known as photoelectric conversion efficiency) of mercury vapour lamps is higher than other UV sources; however, some disadvantages such as bulky configuration, high electric energy requirement, relatively short lifetime, and mercury content, have hindered the utilization of UV lamps in point-of-use (POU) devices. On the other hand, UV light-emitting diodes (UV-LEDs) are currently gaining attention as a new source of UV radiation, which have a small footprint, require much less electrical energy, are environmentally friendly, and have instant on-off ability.79,80 Nevertheless, the highest WPE of commercially available UVC-LEDs is 3.0% at the time, which makes them inefficient to be used in practical systems. Further, UV-LEDs of short wavelengths (e.g., less than 250 nm) have even much less efficiency and power output; the radiant power output for UV-LEDs 43 of 220 nm, as an example, is currently at sub-mW ranges. Given that a large portion of the electrical power applied to these LEDs is dissipated in the form of waste heat, they require special heat management.81,82 Recently, a new source of UV radiation, the microcavity plasma (microplasma) UV lamp was introduced.36 A microplasma lamp consists of microcavities and a gas chamber, which is filled with a rare gas. When a potential difference is applied to both sides of the lamp, the gas becomes ionized and generates UV radiation. This resembles the working principle of fluorescent technology and excimer lamps. Microplasma lamps have several advantages over other UV sources, such as instant-on ability, very short warm-up time, flat geometry, dimmability, environmental friendliness (mercury-free), high WPE, and less waste heat generation. Moreover, these lamps are able to generate specific wavelengths monochromatically (e.g., 172 nm or 222 nm), which makes them suitable for AOPs, as shorter wavelengths are preferred to generate ·OH from H2O2. The ability of a microplasma UV lamp to generate various wavelengths depends on the type of the rare gas used to fill the gas chamber, such as Xe2, KrCl, XeCl, and phosphor gases which are used to produce monochromatic peaks at 172 nm, 222 nm, 308 nm, and a broad spectrum of 220–280 nm, respectively.36 One unique feature of microplasma lamps, compared with UV-LEDs, is that they can emit UV rays from both sides of the lamp, with the radiant power output of the backside being ~80% of that of the front. Therefore, a special reactor design can be considered to utilize the generated photons from both sides. Further, the capability of generating UV radiation at shorter wavelengths is advantageous when integrated with H2O2, given that H2O2 absorption of UV radiation is considerably higher at short microplasma lamp wavelengths at 222 and 172 nm, compared with those of UV lamps and UV-LEDs at about 250–270 nm. Considering these 44 advantages, as well as the planar form factor and the compact footprint, microplasma UV lamps could offer special features for a POU UV/H2O2 water treatment system. The UV/H2O2 process possesses several benefits compared with other chemical or biological treatment methods, such as no sludge formation, the ability to run at ambient conditions, and the potential to completely mineralize and convert the organic contaminations to CO2.83 As mentioned earlier (Eq. 3.1), the external addition of H2O2 is needed in the UV/H2O2 process. This requirement increases the operational cost of the process and is an obstacle in scaling down the process, due to the risks related to the transportation and handling of H2O2.84 A feasible solution to this problem is the in situ electrochemical generation of H2O2 based on the two-electron oxygen reduction reaction (2e-ORR, Eq. 3.4),85 which can be dated back to as early as 1939.86 This process can be promoted using porous, carbon-based gas diffusion electrodes (GDEs) as the cathode, which can convert ambient oxygen to H2O2, with the help of protons (H+) coming from the anodic oxygen evolution reaction (OER).85 The process is highly energy-efficient and is a promising method for the on-site generation of H2O2, which eliminates the need for frequent external H2O2 addition to the photoreactor. The alkaline form of Eq. 3.4 can be written as Eq. 3.5,8 in which HO2− is the conjugate base of H2O2. O2 + 2 H+ + 2 e− → H2O2 E0 = 0.700 V vs. SHE Eq. 3.4 O2 + H2O + 2 e− → HO2− + OH− E0 = −0.065 V vs. SHE Eq. 3.5 Our research focused on the development of an electrochemical cell for the cathodic production of H2O2 integrated with various types of microplasma UV lamps and the evaluation of the electrochemical microplasma lamp reactor performance. First, the efficiency of the 45 electrochemical cell was evaluated by studying the effect of various parameters, such as cell configuration and applied current density, as well as the electrolyte flow rate, temperature, type, and concentration. Then, the cell was combined with three different types of microplasma lamps with UV emissions of 172 nm (UV172), 222 nm (UV222), and a wide spectrum (220 to 280 nm; UVwide), in order to study the direct photolysis of water, organic compound, and the UV/H2O2 process to generate oxidative radicals for organic contaminant degradation. Methylene blue (MB) was selected as a model organic contaminant and its degradation was monitored. MB is representative of organic compounds that are degraded by ·OH, which has been commonly applied for assessing the performance of AOPs and is readily measured via spectrometric methods. Moreover, EF and photoelectro-Fenton (PEF) processes were studied by the addition of Fe2+, and the effect of initial pH and Fe2+ concentration on the performance of the EF and PEF processes was investigated. The promising results suggested that the electrochemical photoreactor we developed can degrade organic compounds using the microplasma UV lamps. To the best of our knowledge, this is the first study that reports an integrated microplasma UV lamp with the in situ generation of H2O2. 3.2 Materials and methods 3.2.1 Materials All of the following materials were purchased from Sigma-Aldrich with an ACS reagent grade and used without any further purification: MB (C16H18ClN3S), sodium sulfate (Na2SO4), sodium nitrate (NaNO3), sodium chloride (NaCl), potassium iodide (KI), potassium iodate (KIO3), sodium tetraborate decahydrate (Na2B4O7.10H2O), ammonium molybdate tetrahydrate 46 ((NH4)6Mo7O24.4H2O), iron(II) sulfate heptahydrate (FeSO4.7H2O), sulfuric acid (H2SO4), and sodium hydroxide (NaOH). 3.2.2 Electrochemical photoreactor The experiments were performed in a divided electrochemical cell, equipped with an additional chamber to expose the fluid to UV radiation. The three chambers were made of plexiglass and sandwiched together using frames that were bolted together with screws and nuts. Each chamber had an effective volume of 34 cm3 and a thickness of 2 cm. To study the effect of cell configuration, the thickness of the cathodic chamber was changed to 1, 2, and 3 cm. A cut of ELAT® LT1400W gas diffusion layer (GDL; FuelCellsEtc, College Station, TX, USA) was used as the air-breathing cathode. A cut of Nafion™ 117 proton exchange membrane (PEM; Fuel Cell Store, College Station, TX, USA) was pretreated as described elsewhere57 and placed between the cathodic and anodic chambers to act as an electronic insulator and ionic conductor in order to allow the passage of protons (and other small positively charged cations) to occur, thus maintaining the electroneutrality of the solution. A mixed metal oxide (MMO) mesh, consisting of a layer of 6 μm IrO2 and RuO2 coated on titanium substrate (William Gregor Ltd, London, UK), was used as the anode. The dimensionally stable anode (DSA) had low overpotential for the OER and was corrosion resistant. A fused quartz disk (G.M. Associates, Inc., Oakland, CA, USA) was used between the UV chamber and the UV source. The overall dimensions of the photo-electrochemical reactor were 10.4 cm × 10.4 cm × 8.8 cm. Figure 3.1 shows the schematic of the electrochemical photoreactor. 47 Figure 3.1 The electrochemical photoreactor used in this study. 1: current distributor, 2: carbon cloth cathode, 3: cathodic chamber, 4: PEM, 5: mixed metal oxide anode, 6: anodic chamber, 7: plexiglass separator, 8: UV chamber, 9: quartz disk, and 10: microplasma lamp. 3.2.3 Experimental procedure A stock solution of 0.1 g L–1 MB was prepared by dissolving 100 mg MB in 1 L ultrapure water (18.2 MΩ∙cm, obtained from a PURELAB® water purification system). Unless otherwise specified, the tests were performed at room temperature (23 °C), a constant current density of ~134 A m–2, and a flow rate of 50 mL min–1, using 25 mM Na2SO4 as the supporting electrolyte to increase the conductivity of ultrapure water. Na2SO4 is a common compound used to study 2e-48 ORR, given its inertness and availability. However, other electrolytes (including water) will most likely demonstrate a similar performance. For the EF and PEF processes, an adequate amount of FeSO4.7H2O was dissolved in the dilute electrolyte to get the desired concentration of Fe2+, which was 0.1 mM unless otherwise specified. For the degradation studies, the initial concentration of MB was adjusted to 2 mg L–1 by diluting it with the electrolyte in all the experiments, except in the one described in Section 3.3.3.1, which was set to 4 mg L–1. The temperature of the influent stream was adjusted by a hot plate stirrer and measured by a general-purpose thermometer, as described in Section 3.3.1.2, and pH adjustments of the initial solutions were done by adding either H2SO4 (toward acidic range) or NaOH (toward basic range) as described in Section 3.3.3.1. Table B.1 shows the average properties of the dilute solution used for most of the experiments. Samples were taken at the outlet of the system after the passage of 10 cycles of the initial solution, each cycle being equivalent to one hydraulic residence time. The samples were then analyzed instantly with a UV-visible spectrophotometer (Cary 100, version 9.00; run by the scan software version 4.20; Agilent Technologies, Inc., Santa Clara, CA, USA). The error bars shown in figures and tables represent a confidence interval (CI) of 95% (equivalent to a significance level of α = 0.05). 3.2.3.1 Microplasma UV lamp operation All three microplasma UV lamps were turned on using a power circuit provided by the manufacturer, Eden Park Illumination, Inc. (Champaign, IL, USA). The power circuit was connected to a commercial DC power supply (EX355R; Aim-TTi, Huntingdon, Cambridgeshire, UK) to run at a constant voltage of 12.00 V. The resultant current of each lamp was ~0.7 A, with 49 slight variations over time. Figure 3.2 shows the normalized radiant power output profiles of the 222 nm lamp and the 220–280 nm lamp, which were obtained using an Ocean Optics USB 2000+ UV-Vis spectrometer connected to a Sony ILX511B cosine corrector diffuser (CCD). Table 3.1 displays the properties of the microplasma lamps used in this study as claimed by the manufacturer of these lamps.87 The changes in the irradiance of the microplasma UV lamps were insignificant over time, as monitored for 20 min using the spectrometer. Figure 3.2 The normalized output profiles of the 222 nm (left y-axis) and 220–280 nm (right y-axis) microplasma UV lamps. 50 Table 3.1 Properties of microplasma lamps (reported by the manufacturer). Wavelength (nm) Gas filling the chamber Front intensity (mW cm–2) Total UV output (mW) Lifetime (h) 172 Xe2 ~15 ~430 >2,000 222 KrCl 1.2 to 4 115 to 500 2,000 to 6,000 220–280 Phosphor <10 ~290 >10,000 Iodide-iodate actinometry was used to find the incident UV fluence rates of the 222 nm and 220–280 nm lamps.88 The actual fluence rate of 172 nm lamp was not investigated, as it would have required methanol actinometry,89–92 which was not feasible given the complexity of the analysis. With the quantum yield values obtained from the literature,88 the fluence rate of 222 nm and wide-spectrum lamps were found to be 0.29 mW cm–2 and 1.07 mW cm–2, respectively. 3.2.3.2 H2O2 concentration measurement method H2O2 concentration in the samples was determined by the triiodide (I3−) method.93 The concentrations were found by reading the absorbance of a mixture of 100 μL sample, 100 μL ammonium molybdate tetrahydrate solution (0.01 M), and 5 mL KI solution (0.1 M) at a wavelength of 352 nm against a calibration curve. The molecular structure and UV absorption spectrum of 4.9 mM H2O2 are depicted in Figure B.1(A) and (B), respectively. H2O2 increased absorbance with shorter wavelengths, starting from 300 nm. Therefore, it is anticipated that the 222 nm lamp would yield high removal rates. 51 The Faraday efficiency percentage was calculated using Eq. 3.6 94: Faraday efficiency (%) =n F Q CH2O2i A× 100% Eq. 3.6 where n is the number of electrons transferred in Eq. 3.4 (= 2), F is the Faraday’s constant (= 96,487 C mol–1), Q is the electrolyte flow rate (L s–1), CH2O2 is the concentration of electrosynthesized H2O2 (mol L–1), i is the applied cathode current density (A m–2), and A is the active cathode geometric area (m2). 3.2.3.3 MB concentration measurement method MB concentrations were quantified by reading the absorbance value of the sample at the peak wavelength of MB absorption spectrum, 664 nm, and by comparing it against the calibration curve, prepared by known concentrations of MB. Figure B.2(A) and (B) present the molecular structure and the absorption spectrum of 4 mg L–1 MB in the UV-visible region, respectively. The degradation percentage was calculated using Eq. 3.7: Degradation percentage (%) =CMB,i − CMB,fCMB,i× 100% Eq. 3.7 where CMB,i and CMB,f are the initial and final concentrations of MB in the sample (mg L–1), respectively. 3.3 Results and discussion The performance of the device for both H2O2 generation and MB degradation (when integrated with microplasma lamps) was evaluated. The effect on the electrochemical H2O2 generation of various operational and configurational parameters, such as current density, electrolyte flow rate, 52 and distance between the anode and the cathode, as well as the electrolyte type, concentration, and temperature was evaluated (Section 3.3.1). Further, the electrochemical cell was coupled with three microplasma UV lamps to investigate direct UV photolysis and UV/H2O2 processes (Section 3.3.2), as well as the EF and PEF processes (Section 3.3.3). 3.3.1 Electrochemical cell performance for the generation of H2O2 As the first set of operational experiments, the performance of the electrochemical cell for the in situ generation of H2O2 was studied. Figure 3.3 shows the concentration of electrogenerated H2O2 as a function of electrolyte flow rate and applied current density, measured at various distances of 1, 2, and 3 cm between the anode and the cathode (i.e., different sizes of the cathode chamber). As illustrated, by increasing the current density, a higher generation of H2O2 was achieved, indicating that the electrons, and not oxygen, were the main limiting species. By increasing the electron density, according to Eq. 3.4, a higher concentration of reactant was introduced into the system, which led to a higher generation of H2O2. Moreover, increasing current density led to a stronger electric field between the anode and the cathode, so that the protons generated on the surface of the anode (due to water oxidation, shown in Eq. 3.8) were more attracted to the surface of the cathode. However, it should be mentioned that increasing the current density to very high values was avoided, as an excess amount of electron can lead to the decomposition of H2O2 on the surface of the cathode by triggering a parasitic cathodic reduction reaction, as shown in Eq. 3.995: 2 H2O → O2 + 4 H+ + 4 e− E0 = 1.23 V vs. SHE Eq. 3.8 H2O2 + 2 H+ + 2 e− → 2 H2O E0 = 1.77 V vs. SHE Eq. 3.9 53 As shown in Figure 3.3, by increasing the flow rate, a lower concentration of H2O2 was obtained. This can be justified by considering that, at higher flow rates, the protons generated on the surface of the anode will be washed away by the incoming flow of water and therefore have a lower chance to fully diffuse to reach the surface of the cathode. Therefore, there is less probability for the three reactants to encounter; thus, the reaction (Eq. 3.4) would proceed slower. Moreover, at faster flow rates, the self-decomposition of H2O2 (Eq. 3.10) might happen faster in the bulk of the electrolyte, due to higher turbulence of the fluid96: 2 H2O2 → O2 + 2 H2O Eq. 3.10 An investigation of the effects of applied current density and electrolyte flow rate on the cathodic H2O2 concentration was performed under three the different distances of 1, 2, and 3 cm (different volumes of the cathodic chamber) between the anode and the cathode. As illustrated in Figure 3.3, the values obtained are very close to each other, with slightly more H2O2 concentration at farther distances. When the distance increases, the volume of the cathodic compartment increases, which at a constant flow rate, leads to a higher residence time. This could be the reason for the slight increase in the concentration of H2O2. Moreover, at higher distances, the velocity of water is slower, resulting in higher diffusion of protons. However, another factor is that at higher distances, it takes longer for the protons generated on the surface of the anode to reach the surface of the cathode. Thus, a trade-off is expected to be observed at distances of more than 3 cm. At very low flow rates (i.e., 40 mL min–1), the generation of H2O2 was slightly higher when the distance between the anode and the cathode was 1 cm, likely due to a very short time required for protons to reach the surface of the cathode. 54 Figure 3.3 The plot of concentration of H2O2 (right y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at various distances between anode and cathode, shown in the legend. The error bars represent 95% CIs from three experimental replicates. By amalgamating the two aforementioned parameters (i.e., the electrolyte flow rate and the applied current density), the definition of charge density could be obtained, which is the ratio of the electrical current (implication of electric charge) to the electrolyte flow rate (implication of volume). As shown in Figure 3.4, a nearly linear relationship between the electrogenerated H2O2 concentration and applied charge density at different distances was observed. A higher charge 55 density led to a higher electrogeneration of H2O2, so that potable water could be obtained at higher flow rates, but more electrical energy needed to be supplied to the system. The Faraday efficiency values as a function of charge density (Figure 3.4) indicated that the efficiencies in general are higher at farther distances between the anode and the cathode. The decrease in the Faraday efficiency values by increasing the charge density could be due to higher consumption of protons at higher charge densities (Eq. 3.4), so that the catholyte is more alkaline, which is less favourable for the electroreduction of oxygen.44 A side reaction accounting for the decrease in Faraday efficiency could be hydrogen evolution reaction (HER; Eq. 3.11), which accelerates at higher pH values.44,97 2 H+ + 2 e− → H2 E0 = 0.00 V vs. SHE Eq. 3.11 The inset graph in Figure 3.4 shows the concentration of H2O2 in the anolyte versus that in the catholyte. A line slope of 0.90 proves that the anode is able to remove 10% of H2O2 by its oxidation to oxygen, according to Eq. 3.12 to Eq. 3.14.44 This would be advantageous for water treatment purposes, since it leads to a lesser concentration of H2O2 in the potable water effluent. H2O2 → HO2· + H+ + e− Eq. 3.12 HO2· → O2 (g) + H+ + e− Eq. 3.13 H2O2 → O2 (g) + 2 H+ + 2 e− Eq. 3.14 56 Figure 3.4 The concentration of H2O2 (solid lines, left y-axis) and Faraday efficiency values (dashed lines, right y-axis) as a function of applied charge density (bottom x-axis) at various distances between the anode and the cathode (legend). The inset graph shows the anodic concentration of H2O2 as a function of the cathodic concentration of H2O2. The error bars represent 95% CIs from three experimental replicates. The plot of voltage as a function of electrolyte flow rate and applied current density at various distances is shown in Figure B.3. Changing the flow rate did not affect the resistivity of the system. One the other hand, the electrostatic potential increased linearly with increasing the current density, following Ohm’s law. For 1, 2, and 3 cm distances, the system presented a resistance of 16.80, 24.47, and 34.17 Ω, respectively, which led to a normalized resistivity of 8.69 Ω cm–1 (all 57 R2 values were more than 0.99). When the distance between the anode and the cathode was increased, it took longer for cations (e.g., H+ and Na+) to reach the surface of the cathode on a diffusion-migration basis, leading to the resistivity rises. Based on the normalized resistivity value, the system can be redesigned to have a desired energy consumption, which depends on an estimation of the influent’s conductivity. Two momentous parameters which determine the performance of the developed system are the electric energy consumption (EEC) and H2O2 generation rate. By increasing the applied current density, the normalized H2O2 electrogeneration rate was increased linearly, within the range studied (Figure B.4(A)). In addition, the electrochemical H2O2 generation rate was nearly independent of the configuration (i.e., the distance between the anode and the cathode). Moreover, the EEC had a linear relationship with the current density, due to the constant resistivity of the system at a fixed configuration and influent. Based on these findings, keeping a distance of 1 cm between the anode and the cathode would be the most suitable option, due to its lower energy consumption compared with that of at larger distances. Increasing the flow rate led to a lower EEC (Figure B.4(B)). On the other hand, the H2O2 generation rate increased slightly. Both observations could be justified by higher flow rates resulting in more turbulence, which promotes mixing. In the case of EEC, increased mixing led to greater involvement of proton ions in the catholyte, and for the case of normalized H2O2 generation rate, promoted mixing yielded more efficient detachment of H2O2 molecules generated on the surface of the cathode. 58 The EEC and H2O2 production rates were calculated using Eq. 3.15 and Eq. 3.16, respectively. EEC (kWh kg−1) =U I3600 C Mw Q Eq. 3.15 H2O2 generation rate (mg cm−2 h−1) =3.6 × 106 C Mw QA Eq. 3.16 where U is the cell voltage (V), I is the applied current (A), C is the concentration of electrogenerated H2O2 (M), Mw is the molecular weight of H2O2 (g mol–1), Q is the electrolyte flow rate (L s–1), and A is the active cathode area (cm2). 3.3.1.1 Effect of electrolyte type and concentration The cathodic and anodic concentrations of H2O2 as a function of the electrolyte type and concentration are depicted in Figure 3.5. At low concentrations of various types of electrolytes, the cathodic generation was affected by the nature of the electrolyte, being highest for NaCl, followed by NaNO3 and then Na2SO4. Higher H2O2 generation of a chloride ion (Cl-)-containing solution is in agreement with the data reported in the literature.58 However, by increasing the concentration, all three electrolytes showed a similar response. More specifically, at a concentration of 50 mM, all three mixtures were able to generate the same concentration of H2O2. A further increase to 100 mM was not effective in increasing H2O2 generation, which is in agreement with the data reported in the literature for 50 mM and 100 mM Na2SO4.98 The reason for an increase in the generation of H2O2 by increasing the electrolyte concentration was the lower voltages at higher concentrations of the electrolyte (shown in the inset plot of Figure 3.5), so that lower overpotentials were obtained, which increased the efficiency of the process and prohibited the activation of the side reactions (i.e., Eq. 3.9 and Eq. 3.11). Going above a certain concentration 59 was less helpful in decreasing the overpotential, as demonstrated in the inset plot, so that a higher generation of H2O2 was not observed. A comparable behaviour was observed for the anodic decomposition of H2O2, except when chloride ions were present in the influent, which aided the removal of excess H2O2 on the anode due to the formation of hypochlorous acid (HOCl) which converted H2O2 to oxygen and protons based on the set of reactions illustrated in Eq. 3.17 to Eq. 3.21.99–101 2 Cl− ↔ Cl2 + 2 e− E0 = −1.396 V vs. SHE Eq. 3.17 Cl2 + H2O ↔ HOCl + H+ + Cl− Eq. 3.18 HOCl ↔ H+ + OCl− Eq. 3.19 OCl− + H2O2 → Cl− + H2O + O2 Eq. 3.20 H2O2 + 2 Cl− → O2 + 2 H+ + 2 Cl− + 2 e− Eq. 3.21 The inset plot in Figure 3.5 shows the voltages obtained in various solutions and concentrations. As evident, Na2SO4 containing solution had the lowest voltage, which was due to the release of 3 ions, 2 Na+ and 1 SO4–, from the dissociation of 1 molecule of Na2SO4 salt. The voltage of the NaCl mixture was slightly lower than that of the NaNO3 solution at a concentration of 12.5 mM, which could have been due to Cl– crossover.102 However, at higher concentrations, the responses of NaCl and NaNO3 were similar. It is worth mentioning that the plot of voltage versus concentration was not linear, which could have been due to approaching the cell’s open circuit (OC) potential at more concentrated solutions. 60 Figure 3.5 Cathodic (solid lines, left y-axis) and anodic (dashed lines, right y-axis) concentrations of H2O2 as a function of electrolyte type (legend) and concentration (bottom x-axis). The inset plot shows the voltage of the cell obtained for various concentrations of different electrolytes. The error bars represent 95% CIs from three experimental replicates. 3.3.1.2 Effect of electrolyte temperature By increasing the temperature of the electrolyte, lower cathodic and anodic concentrations of H2O2 were observed, as depicted in Figure 3.6. Although the O2 electroreduction kinetic is faster at elevated temperatures (positive effect), higher temperatures lead to decreased O2 solubility in the electrolyte and an increased H2O2 decomposition rate (negative effect).44 The lower temperature 61 yielded a higher generation of H2O2, as it is more stable at cold conditions. However, the increase in H2O2 generation was not significant at low temperatures, since there was a trade-off between higher stability of H2O2 and lower mobility of H+ which migrated from the surface of the anode toward the cathode. Moreover, as the mobility of H2O2 was also lower at decreased temperatures, higher molecular residence time could be assumed for the H2O2 molecule, because more anodic decomposition of H2O2 was observed. The inset plot in Figure 3.6 delineates the cell voltages obtained at different temperatures. As is clear from the linear trend, by increasing the electrolyte temperature from 9 °C to 56 °C, a 68.4% decrease in the voltage was achieved, which is due to the higher mobility of protons and sodium ions at elevated temperatures. The decrease in H2O2 electrogeneration (i.e., lower Faraday efficiency) and cell voltage caused by increasing the temperature is in agreement with the data reported in the literature for a similar membrane-divided cell.103 62 Figure 3.6 Cathodic (solid lines, left y-axis) and anodic (dashed lines, right y-axis) concentrations of H2O2 as a function of temperature (bottom x-axis). The inset plot shows the cell voltage as a function of electrolyte temperature. The error bars represent 95% CIs from three experimental replicates. 3.3.2 UV/H2O2 process with microplasma lamps for MB degradation The performance of the electrochemical H2O2 generation system coupled with microplasma UV lamps for the degradation of organic contaminants was investigated. The optimum value of H2O2 concentration in the UV/H2O2 process for the degradation of organic compounds is a function of the initial concentration of organic contaminants.83 The electrochemical cell was tuned to generate H2O2 at a concentration of 1.1 mM, since it can address acceptable degradation of trace organic 63 contaminants typically present in water sources. Higher concentrations of H2O2 could lead to a higher degradation rate of contaminants; however, considering the EEC, residual H2O2 in the treated water, and the scavenging effect of H2O2 for ·OH formation, the H2O2 concentration was kept at 1.1 mM. A H2O2 concentration of 1.2 mM was reported by Raeiszadeh and Taghipour as the optimum value in the UV/H2O2 process to degrade trace concentrations of MB.36 3.3.2.1 Direct photolysis of MB using microplasma lamps The degradation of MB using the three different microplasma lamps alone, a process called direct photolysis, was investigated. As shown in Figure 3.7, UV222 was not able to degrade MB within the 40 second residence time inside the photoreactor. However, UVwide was able to remove 3.0% of MB during the same period. This could be due to the direct photolysis of the organic molecule with UV radiation. MB spectrum, shown in Figure B.2(B), had three main peaks at 664 nm, 292 nm, and 246 nm. The UV222 lamp had only one peak wavelength at 222 nm (Figure 3.2), while the UVwide lamp had two main peak wavelengths at 255 nm and 226 nm. Further, UVwide had a higher radiant power output compared with UV222 (Section 3.2.3.1). Given the higher output, as well as a peak at 255 nm (close to the peaks of 246 nm and 292 nm in MB spectrum), it appeared that UVwide was able to decolorize MB in low amounts. The incapability of UVC radiation for removing a significant concentration of MB is in agreement with the data reported in the literature by applying low-pressure and medium-pressure mercury lamps.83,104 VUV photolysis of water (Eq. 3.2) can generate ·OH, ·OOH, ·O2−, ·H, H2O2, HO2−, and O2.105 The quantum yield of water homolysis at the wavelength of 172 nm is 0.42.76 As displayed in Figure 3.7, UV172 alone was able to remove 5.5% of the initial concentration of MB (i.e., 2 mg L–1) within 64 the hydraulic residence time of the organic molecule inside the system. It should be noted that the penetration depth of UV172 in water was a thin layer of a maximum 70 μm. The reaction of ·H with molecular oxygen in this layer produced the ·OOH radical, which can further degrade organic contaminant.76 Moreover, the VUV radiation was easily absorbed by the ambient oxygen to generate ozone (O3), so that the penetration length in a non-vacuum medium was relatively short. The penetration depth of longer wavelengths (e.g., UV222) was greater. Figure 3.7 Degradation percentages of MB using direct photolysis, UV/H2O2 process, and EF and PEF processes in the absence of UV and the presence of the three different microplasma UV lamps. The error bars represent 95% CIs from three experimental replicates. 65 3.3.2.2 UV/H2O2 process using in situ generation of H2O2 and microplasma lamps The results of coupling the microplasma lamps with the in situ generation of H2O2 are delineated in Figure 3.7. As a control test, when there was not any UV radiation employed and only the electrochemical cell was on, no degradation of MB was observed. The inability of H2O2 alone to degrade MB is in agreement with prior studies.83,104 When the cell was coupled with UV172, an average degradation percentage of 5.7% was obtained, which was similar to the case where UV172 alone was employed (Figure 3.7). Apparently, when VUV was combined with H2O2, the result was the same as when only VUV was utilized. This could have been due to lower transmittance of VUV radiation in solutions containing H2O2; however, the radiation can generate ·OH from both water and H2O2, making it have the same overall efficiency as that of VUV alone. When the UV222 lamp was employed besides the electrochemical generation of H2O2, the system was able to remove 22.5% of MB. This was due to the formation of the strong ·OH radicals and weak ·OOH radicals which were responsible for converting the initial organic molecules to by-products and CO2. When a UVwide lamp was used, an MB removal of 33.3% was observed. The lower removal of the UV222/H2O2 process compared with that of the UVwide/H2O2 process was due to the lower radiant power output of the lamp, as shown in Table 3.1. The lower output cancelled out the advantage of higher absorption of UV-C radiation by H2O2 molecules at a lower wavelength (i.e., 222 nm compared with 280 nm) (see Figure B.1(B)). Direct photolysis of MB molecules using the UVwide lamp, which is shown in Figure 3.7, could also have been responsible for higher removal of the UVwide/H2O2 process compared with the UV222/H2O2 process. 66 It was reported that the MB decolorization in the UV/H2O2 process follows the first-order kinetics in low concentrations of H2O2.104 Going to higher concentrations of H2O2 is not beneficial, as it decreases the transmittance of the solution and will eventually yield to comparable or lower degradation (see Figure B.5(A)). Moreover, the initial concentration of organic compound will also change the UV transmittance, with higher impact on 275 nm compared with 222 nm (see Figure B.5(B)). Therefore, compared with the UV222 lamp, the UVwide lamp was expected to be more affected by an increase in the contaminant concentration. Degradation of MB in the UV/H2O2 process is much faster than that of the UV/photocatalyst process,106,107 as the photocatalytic reaction is a surface reaction with limited reaction zones, but H2O2 molecules are dispersed in the solution, and the UV radiation can effectively generate higher concentrations of potent ·OH radicals, compared with the generally weaker radicals generated in the photocatalytic method. The electric energy per order, EEO, for the base condition (2 mg L–1 MB at a flow rate of 50 mL min–1) was calculated to be 136.6, 28.9, and 20.5 kWh m–3 order–1 for the 172 nm, 222 nm, and wide-spectrum lamps, respectively. These values are lower than UV-LED-based AOPs, mainly due to the higher WPE of the microplasma lamps and higher absorbance of UV radiation at a wavelength of 222 nm compared with 275 nm by H2O2 molecules. The EEO was calculated using Eq. 3.22.108 EEO (kWh m−3 order−1) =PQ log (CiCf) Eq. 3.22 67 where P is the total electric power of the electrochemical cell and microplasma lamp (kW), Q is the electrolyte flow rate (m3 h–1), and Ci and Cf are the initial and final concentrations of the contaminant. 3.3.3 EF and PEF processes for MB degradation The EF and PEF processes were investigated using microplasma UV lamps. A set of experiments was performed to study the effect of the existence of Fe2+ in the solution on the direct photolysis of the organic contaminant. The results confirmed that the presence of Fe2+ ions, without the generation of H2O2, has a negligible effect on the direct photolysis of the organic compound and water. The results of the EF and PEF processes for the degradation of MB is depicted in Figure 3.7. For the EF process (i.e., Fe2+/H2O2), a removal percentage of 18.0% was obtained, which is due to the formation of two radicals by the reaction of Fe2+ with H2O2, demonstrated in Eq. 3.3. The PEF process yielded better results than the EF process, as the UV radiation employed could release the ·OH trapped in the complexes formed between the organic molecule and Fe2+ ion. Interestingly, the results were similar for UVwide and UV172 (40.0%), but slightly higher for UV222 (45.2%). This could be due to closer absorption spectrum peaks of the complexes formed to 222 nm. In fact, It is suggested that the excitation wavelength has an important effect on the ligand-to-metal charge transfer as shown in Eq. 3.23.105 Fe(III)OH(H2O)52++ H2O + ℎ𝜈 → Fe(II)(H2O)62++ ·OH Eq. 3.23 68 The photo-Fenton process is a water treatment process in which oxidative radicals are generated from Fe3+/H2O2 employing a UV or visible light source.105 Discussions on the key reactions of the EF and PEF processes are provided in Section B.7. 3.3.3.1 Effect of initial pH and Fe2+ concentration The effect of the initial pH of the solution on the removal of MB in the PEF process, using the UVwide lamp as the radiation source, is depicted in the inset in Figure 3.8. The trend is in agreement with a study reported by Li et al., 2016.105 As noted, by decreasing the pH, a higher degradation was obtained, and at a pH of 3.02, 100% decolorization occurred. This is because, at alkaline conditions, the Fe2+ ions will precipitate in the form of Fe(OH)3. At a pH of 10.7, only 13.4% of MB was degraded. The slow MB degradation in this case can be attributed to the hydrated Fe-hydroxide species that are present on the surface of the precipitate.105 On the other hand, going to highly acidic conditions is not useful either, as it can trigger side reactions, such as those shown in Eq. 3.9 and Eq. 3.11.44 Also, at pH values of less than 3, the Fe(H2O)63+ complex, which is an active specie that leads to faster MB degradation, gets deprotonated, and the degradation slows down.105 At pH values of greater than 4.8 (i.e., equivalent to the pKa of the hydroperoxyl radical), the ·OOH radical is converted to a superoxide radical anion (·O2−) as illustrated by Eq. 3.24. H2O2 can then react with ·O2−, producing ·OH radicals as shown in Eq. 3.25.105,106 ·OOH ↔ H+ + ·O2− Eq. 3.24 H2O2 + ·O2− → ·OH + OH− + ·O2 Eq. 3.25 It has been suggested that for the UV/H2O2 process, a pH value of 4–5 is optimum, as a highly acidic condition would prohibit ·OH formation from the H2O2 photolysis, and an alkaline condition 69 would accelerate the decomposition of H2O2.44,104 For photocatalytic degradation of MB, a pH value of 2 was found to be the optimum value.106 The performance of the EF and PEF processes as a function of the initial concentration of Fe2+ ions is shown in Figure 3.8. By increasing the initial Fe2+ concentration from 0.1 mM to 0.5 mM, both EF and PEF processes demonstrated a higher degradation of MB. Increasing the initial concentration of Fe2+ yielded better results, as there are more radicals produced in the case of higher presence of Fe2+ in the solution. However, when the initial concentration was further increased to 1 mM, the removal declined, since ·OH is consumed when there is an excess amount of Fe2+. Therefore, a trade-off exists between removal percentage and initial concentration of Fe2+, or, to be more accurate, the ratio of H2O2 concentration to Fe2+ concentration. In this study, a concentration of 0.5 mM Fe2+ was found to be an appropriate concentration, when H2O2 was electrogenerated at a concentration of 1.1 mM. Moreover, we found that when the initial concentration of Fe2+ was low, the PEF process could significantly increase the removal, compared with the EF process alone. However, when the concentration of Fe2+ was increased, the gap between the EF and PEF processes became negligible. This could be explained by the fact that when there is a lower amount of Fe2+ in the solution, a significant regeneration of the complexes formed is required for high removal of any organic contaminant, but when the concentration of Fe2+ is high enough, the regeneration is no longer required, so that the difference between the EF and PEF processes becomes insignificant. 70 Figure 3.8 The performance of the EF and PEF processes as a function of the initial Fe2+ concentration. The inset figure shows the degradation of MB as a function of the initial pH in the PEF process. The error bars represent 95% CIs from three experimental replicates. 3.4 Conclusions Our study focused on the applicability of three different microplasma UV lamps in combination with a POU electrochemical system for in situ generation of ·OH. The system was characterized based on the applied current density and the electrolyte flow rate. The reactor exhibited Faraday efficiency values of greater than 80% over a wide range of charge densities. We found that the generation of H2O2 is independent of the nature of the solution passing through the system. The 71 electrochemical cell showed comparable performance in various sizes of 1 cm to 3 cm, but had the minimum EEC when the distance between the anode and the cathode was 1 cm. After coupling the system with different microplasma lamps with wavelengths of 172 nm, 222 nm, and 220–280 nm, the degradation of MB was monitored. The UVwide lamp exhibited the highest removal percentage of 33.3%, probably due to its higher radiant power output. The EF and PEF processes were also investigated using microplasma lamps, with the UV222 lamp showing an MB degradation of 45.2%. Overall, the electrochemical photoreactor developed in this study exhibits promising results for both in situ generation of H2O2 and degradation of organic contaminants, when coupled with microplasma lamps. 72 Chapter 4: Expanded polytetrafluoroethylene-based gas diffusion electrode for electrochemical generation of hydrogen peroxide 4.1 Introduction Hydrogen peroxide (H2O2) is a versatile chemical compound that has various applications in wastewater treatment, chemical manufacturing, electricity generation, and the pulp and paper industry.109–111 It is a powerful, environmentally friendly oxidant that produces water and oxygen as the by-products.112,113 The global market demand for H2O2 is expected to reach 6.0 million metric tons, equivalent to US$6.4 billion in 2024.112,114 Currently, the main industrial process for commercial-scale H2O2 production is the anthraquinone oxidation (AO) process that yields high concentrations of H2O2; however, it requires high energy input, hydrogen gas reactant, and expensive palladium catalysts, and it leaves harmful organic by-products behind.109–111,114,115 As an alternative approach, the in situ electrochemical production of H2O2 enabled by a two-electron oxygen reduction reaction (2e-ORR) has been suggested, which also eliminates the risks of transport, handling, and storage of concentrated H2O2.110,114,116 The electrochemical reaction can be done on carbonaceous cathodes, such as graphite, graphene, and carbon felt.117 Carbon-based electrodes are promising as they have low costs, high stability, and excellent conductivity.110 In the 2e-ORR, shown in Eq. 4.1, the protons are supplied from the anodic reaction (i.e., oxygen evolution reaction (OER)), Eq. 4.2, using a dimensionally stable anode (DSA).13,111 O2 + 2H+ + 2e− → H2O2 E0 = 0.695 V vs. SHE Eq. 4.1 2H2O → O2 + 4H+ + 4e− E0 = 1.229 V vs. SHE Eq. 4.2 73 Due to the low solubility of oxygen in aqueous media, a type of electrode, the gas diffusion electrode (GDE), is typically used in the H2O2 electrosynthesis. The GDE is in contact with the aqueous and gaseous phases at the same time and is considered the most efficient type of electrode for H2O2 production.13,111,113 However, the current generation of commercial GDEs is highly susceptible to flooding when the pressure of the liquid side is higher than that of the gas side,111,113 making their out-of-laboratory application impractical. Polytetrafluoroethylene (PTFE), also known as Teflon, is well known for its high water-repellency (i.e., hydrophobicity) properties. Expanded polytetrafluoroethylene (ePTFE, commercially known as GORE-TEX®), on the other hand, shares the same chemical formula with PTFE, but possesses a high porosity of about 80%.118 It is reported that the ePTFE fabric has more than 106 pores per cm2.119 These two properties together (high hydrophobicity and porosity), which, in general, allow the passage of gases but not liquids, make the use of ePTFE-based substrates a potential candidate as gas diffusion layers (GDLs) in GDEs. The ePTFE was discovered in 1969, and since then, due to its unique features, it has found numerous applications, such as in medical devices and filtration material.120 The goal of the current study was to investigate the potential of breathable, ePTFE-based GDE as a cathode in an electrochemical cell for the generation of H2O2. To the best of our knowledge, this is the first study in which an ePTFE-based GDE has been applied to the electrochemical generation of H2O2. 74 4.2 Material and methods 4.2.1 Materials All of the following American Chemical Society (ACS) reagent grade chemicals were supplied from Sigma-Aldrich and used without any further purification: hydrogen peroxide solution (H2O2, 30% w/w in water), sulfuric acid (H2SO4), potassium iodide (KI), ammonium molybdate tetrahydrate ((NH4)6Mo7O24.4H2O), polytetrafluoroethylene (PTFE) preparation (60% w/w dispersion in water), Nafion™ perfluorinated resin solution (5% w/w in mixture of lower aliphatic alcohols and water, containing 45% water), 2-propanol ((CH3)2CHOH; also known as isopropyl alcohol (IPA)), and sodium sulfate (Na2SO4). Carbon black (Vulcan XC 72R), with an average particle size of 50 nm, was supplied from Fuel Cell Store (College Station, TX, USA) and used for electrode fabrication. 4.2.2 Electrochemical reactor and experimental procedure A divided electrochemical cell was used to study the performance of the fabricated electrodes for the electrochemical generation of H2O2. The cell was a continuous, flow-through, single-pass reactor composed of two chambers: a cathodic chamber and an anodic chamber. Each chamber was selected to be a flow channel in order to increase the mixing of the fluid. The chambers were made of plexiglass with a thickness of 2 cm, which is equivalent to the distance between the anode and the cathode. The input and output threads to the chambers were one-quarter inch National Pipe Tapered for Fuels (NPTF). A proton exchange membrane (PEM), Nafion 117, supplied by Fuel Cell Store, was pretreated (Section 4.2.2.1) and placed between the anode and the cathodic chamber. The anode was mixed RuO2 and IrO2 (referred to as mixed metal oxide (MMO) anode) 75 coated on a titanium mesh, obtained from William Gregor Ltd (London, UK). The MMO anode is a common electrocatalyst for the OER. A thin cut of carbon cloth was placed on the carbon-coated side of the cathode to act as a current distributor. Figure 4.1 shows the electrochemical reactor used in this study. Figure 4.1 The electrochemical reactor used in the current study. The first and second chambers from the left are the anodic and cathodic chambers, respectively. The MMO anode and Nafion membrane were placed between the two chambers (not shown), while the cathode (shown in black) was placed on the air-facing side of the cathodic chamber. 76 A dilute, inert electrolyte (Na2SO4) was used to raise the electrical conductivity of ultrapure water. For each experiment, 1 L of electrolyte (0.1 M Na2SO4) was passed through the cathodic and anodic chambers, respectively, and a sample was taken from each chamber at the end. The applied electrical current and electrolyte flow rate were changed separately to study the effect of each on the performance of the system. All the solutions were prepared with ultrapure water, obtained from a PURELAB® Option-Q unit (ELGA LabWater, High Wycombe, UK). The error bars illustrated in figures show a confidence interval (CI) of 95% (equivalent to a significance level of α = 0.05). 4.2.2.1 Nafion membrane preparation Following a method described elsewhere,57 the Nafion 117 membrane was prepared by boiling in 3% v/v H2O2 for 0.5 h, rinsing with ultrapure water, boiling in ultrapure water for 0.5 h, rinsing with ultrapure water, boiling in 0.5 M H2SO4 for 0.5 h, and finally rinsing with ultrapure water. The membrane was then placed in the reactor. 4.2.3 H2O2 concentration measurement method The concentration of H2O2 in the samples was measured by reading the absorbance of a mixture of 100 μL sample, 5 mL KI solution (0.1 M), and 100 μL ammonium molybdate tetrahydrate solution (0.01 M) at a wavelength of 352 nm against a calibration curve.93 Ammonium molybdate tetrahydrate acts as the catalyst and accelerates the reaction between H2O2 and I-. A UV-Vis spectrophotometer (Cary 100; Agilent Technologies, Inc., Santa Clara, CA, USA) was used to read the absorbances of the samples. 77 4.2.4 GDE preparation An aspire laminated ePTFE membrane with a polypropylene (PP) backing layer was acquired from Sterlitech Corporation (Kent, WA, USA) and used as the substrate of the GDE. The datasheet of this product specifies that it has a water entry pressure of >60 psi and an air flow rate of 2.13–7.62 cm3 min–1 cm–2 at 125 Pa. After some initial trials, we found that a binder content of ~30% would be sufficient for the strong adhesion of carbon black nanoparticles to the substrate. Therefore, 1 mL of Nafion perfluorinated resin solution and 0.1 g carbon black were added to 10 mL IPA, and the mixture was subjected to an ultrasonic homogenizer (Fisher Scientific Sonic Dismembrator Model 500; Thermo Fisher Scientific, Waltham, MA, USA) for 10 min at 50%, leading to a total delivered energy of 120 kJ. The 20 mL vial containing the unhomogenized ink was placed under the horn tip and covered with a cut of Parafilm in order to minimize the IPA evaporation. The homogenized ink was then uniformly coated onto the ePTFE side of the substrate using a Paasche H-Set Airbrush (Paasche Airbrush Company, Kenosha, WI, USA) at a distance of 20 cm, with an air pressure of 30 psia. The substrate was firmly kept in its place using a 3 mm thick aluminum mask. The catalyst loadings were determined by dividing the difference in the initial and final weights of the samples by the coated area. A photo of a sample carbon-coated GDE is provided in Figure C.1. It should be mentioned that prior to airbrushing, two other coating methods were tested: a spray coating with an ultrasonic nozzle (ExactaCoat OP2; Sono-Tek Corporation, Milton, NY, USA) and a sputter coating with a high-vacuum coating system (Leica EM MED020; Leica, Wetzlar, Germany). However, in the former, the electrodes exhibited poor operational performance, 78 probably due to the full blockage of the air flow pathway, while in the latter, the substrate degraded during the coating procedure because of the high temperature of the target (i.e., carbon). 4.2.5 Characterization techniques In order to have a better understanding of the GDE’s molecular structure, which can be related to its performance, the ePTFE-based electrode was characterized for its morphology (scanning electron microscopy (SEM)), hydrophobicity (contact angle measurement), leakage resistance (liquid entry pressure (LEP)), and pore size distribution (mercury intrusion porosimetry (MIP)). It should be mentioned that, due to some limitations, the characterization was performed on only one catalyst loading of 0.1 mg cm–2; however, the key analyses are applicable to different catalyst loadings as discussed in the respective sections. 4.2.5.1 SEM imaging To study the surface morphology of the airbrushed electrode, the SEM imaging technique was employed. Surface and cross-section imaging were performed using a Helios NanoLab 650 focussed ion beam (FIB)-SEM. The samples were coated with iridium using the sputter coater. 4.2.5.2 Contact angle measurement In order to verify the hydrophobicity of the GDE after the deposition of carbon black with the Nafion binder, the contact angle of the GDE with water was measured. The contact angle quantifies the wettability of a surface by a liquid using the Young equation, which is shown in Eq. 4.3. A contact angle goniometer (OCA 15; DataPhysics Instruments, Filderstadt, Germany) was used to study and compare the contact angle values of the bare substrate and carbon-coated ePTFE. 79 γsg = γsl + γlg cos θ Eq. 4.3 where γ is the interfacial surface tension (N m–2), s, l, and g denote the solid, liquid, and gas phases, respectively, and θ is the contact angle (deg.). The contact angles were measured using the plugin Contact Angle of ImageJ software, and the ellipse angles are reported (as opposed to circle angles). 4.2.5.3 Water entry pressure measurement LEP is a measure of membrane wettability and is defined as the pressure that needs to be applied to a dry membrane in order to force a liquid to penetrate through the membrane.121 Conventionally, the Franken model, which is based on the Young-Laplace equation, is used to find the LEP of a membrane, as in Eq. 4.4 as follows122,123: LEP =−2 B γl cos θrmax Eq. 4.4 where B is a unitless geometrical factor to consider the irregularities of the pores (0.4 < B < 0.6 for the stretched PTFE membranes and B = 1.0 for cylindrical pores),121 γl is the liquid surface tension (N m–1), θ is the contact angle of the liquid with the membrane (deg.), and rmax is the maximum pore radius of the membrane (m). The aforementioned SI units result in a LEP value in the units of Pa. When water is used as the liquid, LEP can be replaced by water entry pressure (WEP). In order to measure the WEP values of the developed GDE, a standard test was conducted in a similar apparatus suggested by W. L. Gore & Associates, Inc. in an application note.124 In summary, the test sample was installed on a support screen, and the pressure of water was increased gradually until the first drops of water were observed from the other side, which happened 80 simultaneously with a decrease in the filter inlet pressure. In the experiment, the carbon-coated ePTFE was cut and placed into a PP filter holder with a diameter of 47 mm (Advantec PP 47; Sterlitech Corporation, Kent, WA, USA), followed by a tube connection to a pressure gauge and a syringe pump (RK-74900-20; Masterflex; Cole-Parmer Canada Company, Montréal, QC, Canada). The filter holder is rated for a maximum differential pressure of 71 psi, equivalent to 4.9 bar. 4.2.5.4 MIP To determine the pore size distribution and porosity of porous media, the MIP technique is commonly used. The test involves injecting mercury into the porous media, which is placed in a glass tube (i.e., penetrometer) with a certain volume. Since mercury is a non-wetting fluid, it will resist entering a pore or capillary until sufficient pressure is applied to force its entry, starting with the largest pores. A lower pore size demands a higher mercury pressure. Washburn’s equation has established a simplified relationship between the pore diameter and MIP and is analogous to Eq. 4.4, bearing in mind that the surface tension, contact angle, and intrusion pressure of mercury should be inserted instead. The MIP test was carried out on both bare ePTFE (i.e., the substrate) and carbon-coated ePTFE (i.e., the electrode) using a mercury porosimeter (Autopore IV 9500; Micrometrics, Norcross, GA, USA). 81 4.3 Results and discussion The mechanical properties of the developed GDE were characterized (Section 4.3.1 to 4.3.4), and its efficacy for in situ production of H2O2 was tested by installing it into the electrochemical cell (Section 4.3.5). 4.3.1 Surface morphology analysis Figure 4.2(A) and (B) illustrate the surface structure of the electrode. Figure 4.2(A) shows a macroscopic-level uniform deposition of the electrocatalyst over the surface of the substrate, while Figure 4.2(B) suggests a variety in the size of the carbon black nanoparticles, which could be due to cracking during the ultrasonic homogenization process. Most of the particles indicate a diameter of ~50 nm, as indicated by the provider. Moreover, it was observed that the particles attach to each other and form a cluster, likely as a result of using the binder. The uneven surface of the electrocatalyst layer can promote the reaction zone for the 2e-ORR. The cross-sectional images (Figure 4.2(C) and (D)) were taken in backscattered electron (BSE) mode. The results proved that the carbon black nanoparticles penetrated into the porous network of the ePTFE substrate and filled a portion of the empty space, which was made possible by applying IPA. IPA has a lower surface tension than water; therefore, it will enter the pore structure of the ePTFE membrane and displace the air in the pores. Once the air is forced out of the membrane pores, carbon particles can follow the alcohol into the pores. Once the membrane is allowed to fully dry (i.e., IPA has evaporated), the air will be back in the pore structure and water will no longer be able to enter the pores. Figure 4.2(D) indicates that the deposition of nanoparticles is nearly uniform, with an average thickness of ~1.5 μm, with slight variations. 82 Figure 4.2 The FIB-SEM images of the carbon-coated ePTFE-based electrode. (A) and (B) show the surface morphologies at two different magnifications of 2,000 and 100,000, respectively, while (C) and (D) depict cross-sectional images of the same electrode at magnifications of 5,000 and 20,000, respectively. 1: carbon nanoparticles penetrated into the ePTFE structure; 2: ePTFE porous network. The white border in (C) is a platinum strip and was used to protect the underlying material from FIB damage during the milling stage. 4.3.2 Hydrophobicity analysis As demonstrated in Figure 4.3(A) and (B), 9 μL of water showed 136.6° and 151.2° contact angles on the bare ePTFE and carbon-coated ePTFE, respectively. As the results indicated, a 14.6° increase in the contact angle was obtained when the carbon black nanoparticles with the Nafion binder were applied to an ePTFE substrate. Indeed, Nafion is a proton conductive polymer with a 83 hydrophobic base and hydrophilic ionic clusters, and the proton transfer to the electrocatalyst occurs within a transition region.125 Apparently, the hydrophobic base of the binder decreases the surface energy, causes a higher contact angle, and converts the hydrophobic GDL into a superhydrophobic GDE. It should be mentioned that, during the initial experiments, the goal was to use a PTFE binder, which is a less expensive option and more commonly used as a binder to attach an electrocatalyst to a substrate and to limit water leakage from the GDE, given its hydrophobic properties. However, in this study, the PTFE binder did not yield a proper attachment of the nanoparticles to the substrate, and the electrocatalyst was easily washed off the surface after deposition. This could have been due to the high electronegativity of the fluoro groups on both the substrate (i.e., ePTFE) and the binder (i.e., PTFE), which results in a strong repellant force between the two, leading to detachment of the electrocatalyst from the surface. Therefore, a Nafion binder was found to be a suitable option to adhere electrocatalysts to an ePTFE-based surface, which could increase hydrophobicity as well. 84 Figure 4.3 Contact angle measurement using a goniometer. (A) shows a droplet of water on the bare substrate, while (B) depicts a droplet of water on carbon-coated ePTFE. 4.3.3 Leak resistance analysis Two experiments were performed: a gradual increase of the water pressure using the syringe pump, and a sudden pressure increase (pressure shock). After more than 10 trials in each case, a leakage from the filter holder was observed at ~70 psi. As expected, the filter holder was not able to tolerate pressures higher than the rated value of 71 psi. However, after visual inspection of the GDE, no deformation in its structure was observed, implying that the membrane (i.e., GDE) can withstand a differential pressure of at least 70 psi. Slightly lower values for the WEP of bare, unmodified PTFE membranes have been reported in the literature (e.g.,122,126,127), partially indicating that the use of Nafion binder to attach a carbon black catalyst to the bare ePTFE can increase the WEP of a PTFE substrate to some extent. 85 4.3.4 Pore distribution analysis The pressure and volume data were recorded from 30 to 33,000 psia in the forward direction (i.e., intrusion) and from 33,000 to 14.7 psia in the backward direction (i.e., extrusion). A summary of the physical properties of the two samples is provided in Table 4.1. Some of the properties, such as total intrusion volume, bulk and apparent densities, tortuosity factor, and fractional dimensions remained almost the same, indicating that the deposition of the carbon black electrocatalyst using the Nafion binder did not modify the macroscopic structure significantly. Slight increases in the total pore area and porosity of the coated electrode suggest that the addition of carbon to the porous ePTFE substrate probably created an extended layer of particles on top of the membrane, which along with the nanoparticles penetrated into the ePTFE structure (as described in Section 4.3.1) was the active catalyst layer of the electrode. Moreover, this proves that the catalyst coating method was viable, as there was no reduction in the porosity of the electrode. The average pore radius of the electrode declined from 338 nm to 290 nm, due to the nanoparticles occupying the pores of the ePTFE network. The threshold pressure, which is the pressure at which mercury finds pore network connectivity and starts penetrating into the sample, was 3.03 psi for the bare GDL and 8.14 psi for the coated GDL, denoting that the resistance of carbon-coated GDL toward pressure (i.e., leakage) is much higher than the bare ePTFE. In addition, the permeability decreased from its initial value of 6,850 millidarcy by a factor of 5, implying that the air flow rate was probably reduced as well. Lastly, the tortuosity triplicated, most likely due to manual coating. 86 Table 4.1 Physical properties of the bare and coated GDLs. Row Properties Bare GDL Coated GDL 1 Total intrusion volume (mL g–1) 1.612 1.710 2 Total pore area (m2 g–1) 9.534 11.811 3 Average pore radius (nm) 338.3 289.6 4 Bulk density at 0.53 psia (g mL–1) 0.3828 0.3685 5 Apparent (skeletal) density (g mL–1) 1.0004 0.9966 6 Porosity (%) 61.7342 63.0218 7 Threshold pressure (psia) 3.03 8.14 8 Permeability (millidarcy) 6850 1346 9 Tortuosity factor 1.532 1.518 10 Tortuosity 3.326 9.463 11 Percolation fractal dimension 2.993 2.988 12 Backbone fractal dimension 2.957 2.959 The mercury porosimeter also provided a set of data pairs regarding mercury intrusion pressure and the corresponding mercury intrusion volume. Figure 4.4(A) depicts the cumulative intrusion volume over the pressure range of 0 to 33,000 psia. The plot provides a graphical representation of how mercury intrudes into the sample in the intrusion (forward) step and how much mercury is trapped in the sample in the extrusion (backward) step. From this figure, the breakthrough pressure (i.e., the pressure at which mercury starts intruding into the sample) was found to be around 0.5 psia for both samples. It can also be seen that significant mercury intrusion happens in the low-pressure region, indicating the presence of large pores. The flat curves, which represent the 87 extrusion step for both samples, indicated that most of the mercury is trapped in the sample, which is another proof of a high pore-to-throat aspect ratio. Figure 4.4(B) is a representation of the incremental intrusion volume as a function of the mercury intrusion pressure. The incremental volume is calculated by subtracting the cumulative volume measured at a certain pressure from the cumulative volume measured at a higher pressure. Again, Figure 4.4(B) indicates the presence of large pores that take up the bulk of the intruded mercury in both samples. Nearly 0.38 mL mercury per 1 gram of the samples was used in the pressure range of 0.5 to 10 psia. Figure 4.4(C), which illustrates the cumulative pore area versus the pore diameter, shows that the coated GDL (i.e., the electrode) contains more pores than the bare GDL, which is due to the coating material contributing to the porous medium. It is also evident that both bare and coated GDLs have pore diameters of 5 to 1,000 nm, with most of the contribution coming from pore diameters of 5 to 100 nm, and that the coated GDL has more pores with diameters in the aforementioned region. The trend of the two samples was similar for pore diameters of more than 100 nm, confirming that introducing carbon black nanoparticles into the ePTFE structure will not produce large pores. Figure 4.4(D) illustrates the differential volume as a function of pore diameter. The differential volume is defined as incremental volume divided by the difference in the diameter of pores over the volume increment. The plot indicates the presence of pores with a diameter of ~5 nm in both bare and coated GDLs. The extrusion curves for both samples show that all the mercury is trapped in pores of 5 nm as the mercury gets disconnected from the pore throats (i.e., smaller pores) at higher pressures. However, the differential volume versus pore diameter plot results in a data 88 compression effect in the large diameter range (amplification effect of the small diameter range), which can be eliminated by calculating the differential volume divided by the difference in the logarithm of diameter of pores, as shown in Figure C.2. In summary, in the MIP test, the coated GDL showed the presence of larger pores and a higher pore volume compared with the bare GDL. As an additional measurement, the Brunauer–Emmett–Teller (BET) surface area analysis was performed on the substrate (ePTFE) using a Micromeritics ASAP 2020 system with nitrogen (N2) gas at 77 K. The volume of N2 adsorbed was measured as a function of the pressure ratio (i.e., the equilibrium pressure divided by the saturation pressure of the adsorbate gas), and the BET surface area was calculated using the BET equation. The results indicated a BET surface area of 988.9015 m2 g–1 for the substrate, proving that it was highly porous and that the diffusion of oxygen from the gaseous phase (i.e., air) could occur without much resistance. 89 Figure 4.4 The results of MIP. (A) Cumulative volume over the entire pressure range; (B) incremental intrusion volume as a function of MIP; (C) cumulative pore area versus pore diameter; (D) differential volume as a function of pore diameter. 4.3.5 Electrochemical cell performance for the generation of H2O2 The performance of the electrochemical cell (Section 4.2.2) was evaluated for the cathodic generation of H2O2. For this purpose, three electrodes were prepared with catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2. Each ePTFE-based GDE was installed into the cell and tested for the electrochemical generation of H2O2 at various electrolyte flow rates and applied current densities using a one-factor-at-a-time (OFAT) method. Unless otherwise specified, the experiments were 90 performed at a current density of 134 A m–2 and an electrolyte flow rate of 40 mL min–1. 0.1 M Na2SO4 was selected as the electrolyte. Two important operating parameters that can affect H2O2 production in a continuous electrochemical cell are the flow rate of the fluid within the system and the current density applied to the electrodes. At higher flow rates, given the constant volume of the reactor, less residence time can be achieved, which leads to a lower concentration of H2O2. Our results, depicted in Figure 4.5, confirmed that there is an inverse relationship between the H2O2 concentration and the electrolyte flow rate. Moreover, electrons are among the reactants of 2e-ORR (Eq. 4.1), so that in theory, increasing the current density would lead to a higher concentration of H2O2, as long as oxygen is not the limiting reactant. The experiments proved that there is a linear, direct relationship between the applied current density and cathodic H2O2 concentration for the range studied (Figure 4.5). This proves that oxygen is not the limiting specie and that coating the ePTFE substrate with carbon black using the method described in this work does not lead to the blockage of the air-passing channels of the substrate. Moreover, a higher electrical current yields a higher electric potential for the system, which itself generates a stronger electric field between the anode (positive pole) and the cathode (negative pole). The stronger electric field forces the protons produced as a result of anodic OER (Eq. 4.2) to migrate toward the cathode more intensely, promoting the 2e-ORR. Moreover, due to a concentration gradient, a higher driving force is formed when the concentration of H+ on the surface of the anode is higher, which also contributes to the elevated H2O2 production at higher current densities. 91 Three different catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 were studied. By increasing the catalyst loading, two contradictory effects can be expected: a higher reactive surface area (positive effect), which increases the reaction zone and therefore the concentration of H2O2, and a lower air flow rate (negative effect), which suppresses the transport of oxygen to the three-phase boundary thus reduces H2O2 concentration. A catalyst loading of 0.22 mg cm–2 outperformed the other examined loadings (Figure 4.5), suggesting that even at a relatively high loading, the electrode is still porous enough to generate comparatively high concentrations of H2O2, so that the positive effect surpasses the negative effect. An exact optimum catalyst loading was not found in the current study; however, the concentration of electrogenerated H2O2 obtained using a carbon black loading of 0.22 mg cm–2 is satisfactory for water treatment applications (0.5 to 0.6 mM), and given the high cost of the Nafion binder, a higher catalyst loading seems unnecessary. 92 Figure 4.5 Concentration of H2O2 (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend). The error bars represent 95% CIs from three experimental replicates. In order to evaluate the combined effect of the applied current density and the electrolyte flow rate, charge density, which is the amount of electric charge per unit volume, was considered. The results, shown in Figure 4.6, indicated that the concentration of H2O2 increases almost linearly by increasing the charge density, in the range of 228 to 513 C L–1. A high charge density leads to the activation of side reactions, as shown in Eq. 4.5 to Eq. 4.7,14,111–113,128–130 thus reducing the amount of experimentally electrogenerated H2O2 compared with the theoretically expected value. This can 93 be quantified by the definition of current efficiency, which is the ratio of theoretical energy required for a given reaction to the experimentally applied energy. As depicted in Figure 4.6, at higher charge densities, much lower current efficiencies were obtained, confirming that the efficiency of the electrodes for 2e-ORR (Eq. 4.1) drops at high charge densities, and these electrodes are most efficient at low charge density values. In addition, a higher catalyst loading exhibited a higher current efficiency, which again is due to less activation of side reactions, given that a more conductive electrode requires less potential to reach a certain electrical current; thus, a lower overpotential (as a measure for the triggering of side reactions) is achieved. 2H+ + 2e− → H2 E0 = 0.000 V vs. SHE Eq. 4.5 H2O2 + 2H+ + 2e− → 2H2O E0 = 1.770 V vs. SHE Eq. 4.6 O2 + 4H+ + 4e− → 2H2O E0 = 1.230 V vs. SHE Eq. 4.7 The anode was able to remove 18.0% of the H2O2 molecules generated on the surface of the cathode (the inset graph in Figure 4.6). The main reaction leading to dissociation of H2O2 to the products is shown in Eq. 4.8.128,130 The decrease in H2O2 concentration in the anodic chamber, which contains the effluent of the system, would be useful for water treatment purposes, because it is desired to have the minimum concentration of H2O2 in drinking water. H2O2 → O2 + 2H+ + 2e− E0 = −0.695 V vs. SHE Eq. 4.8 The charge densities and current efficiencies were calculated using the following equations: ρq (C L−1) =IQ Eq. 4.9 94 Current efficiency (%) =n F C VI t× 100% Eq. 4.10 where ρq is the volumetric charge density (C L–1), I is the applied current to the cell (A), Q is the volumetric flow rate (L s–1), n is the number of electrons transferred in 2e-ORR (= 2), F is the Faraday constant (= 96,486 C mol–1), C is the concentration of H2O2 (M), V is the volume of the cathodic compartment (L), and t is the residence time of the catholyte inside the cathodic compartment (s). Figure 4.6 Concentration of H2O2 (solid lines, left y-axis) and current efficiencies (dashed lines, right y-axis) as a function of applied charge density (bottom x-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–95 2 (legend). The inset graph shows the anodic concentration of H2O2 versus the cathodic concentration of H2O2, along with the fitted line. The error bars represent 95% CIs from three experimental replicates. Since the experiments were performed at a galvanostatic (i.e., constant current) mode, the resultant voltage was expected as an output parameter of each test. For each set of experiments, the output electric potentials were recorded in order to analyze the behaviour of the system. As expected, varying the flow rate did not have any influence on the resultant voltage (Figure 4.7). On the other hand, the voltage had a linear relationship with the applied current density, proving that the carbon-coated side of the GDE follows Ohm’s law. As predicted, a higher catalyst loading (i.e., a more conductive electrode) resulted in a lower voltage. In fact, the transition from 0.05 to 0.10 mg cm–2 and from 0.05 to 0.22 mg cm–2 resulted in a 13.4% and 29.2% reduction in the resistivity of the entire system, respectively, which is a direct consequence of increasing the catalyst loading. 96 Figure 4.7 The resultant electric potentials of the cell (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis) at the various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend). The error bars represent 95% CIs from three experimental replicates. Two parameters that can be utilized to normalize the results to scale-up or scale-down the system are electric energy consumption (EEC) and H2O2 generation rate. The former relates to the amount of electricity, typically in units of kWh, to produce 1 kg of the desired product (i.e., H2O2), while the latter states the production flux of H2O2 per cathode area in the unit of mg cm–2 h–1. As discussed before, a higher current density resulted in a higher H2O2 generation rate (Figure 4.8(A)). In addition, a higher current demands a higher voltage, thus increasing electrical energy 97 consumption. A similar discussion is valid for the effect of catalyst loading: a high catalyst loading requires less electrical energy input and yields more production of H2O2. Studying the effect of electrolyte flow rate on EEC and H2O2 generation rate (Figure 4.8(B)) showed that faster flow rates resulted in slightly higher generation of H2O2. This could be explained by the fact that, at a higher flow rate, more mixing will occur, so that H2O2 molecules produced on the surface of the cathode have a higher chance of moving to the bulk of the fluid. The same justification can be applied for the EEC; that is, better mixing of the fluid inside the cell at faster flow rates promotes detachment of charge carrier ions (e.g., H+) and leads to a slightly lower consumption of electricity. The rate of H2O2 generation using ePTFE-based cathodes was up to 20 times higher than those reported in the literature which applied a similar current density but used cathode types different from the one used in our study (e.g.,11,130–133). This could be due to a well-maintained three-phase boundary that the ePTFE substrate provided, in addition to its high surface area, as confirmed by characterization tests (Sections 4.3.1 to 4.3.4). The EEC and H2O2 generation rates were calculated using the following equations: EEC (kWh kg−1) =U I3600 C MW Q Eq. 4.11 H2O2 generation rate (mg cm−2 h−1) =3.6 × 106 C MW QA Eq. 4.12 98 where U is the electric potential of the cell (V), I is the applied current to the cell (A), C is the concentration of electrogenerated H2O2 (M), MW is the molecular weight of H2O2 (g mol–1), Q is the flow rate of electrolyte (L s–1), and A is the active area of the cathode (cm2). 99 Figure 4.8 EEC (solid lines, left y-axis) and H2O2 generation rate (dashed lines, right y-axis) at various catalyst loadings of 0.05, 0.10, and 0.22 mg cm–2 (legend), as a function of (A) applied current density and (B) electrolyte flow rate (bottom x-axis in both cases). The error bars represent 95% CIs from three experimental replicates. 100 4.4 Conclusions In our study, a new generation of GDEs for the electrochemical generation of H2O2 was developed, employing ePTFE fabric as the substrate, carbon black as the electrocatalyst, and Nafion solution as the binder. The electrode was characterized for surface morphology, hydrophobicity, leak resistance, and pore distribution. The SEM imaging revealed that a portion of the electrocatalyst penetrates the pores of the substrate, indicating the suitability of the ePTFE substrate, Nafion binder, and airbrushing as the coating method. The deposition of the electrocatalyst ink increased the hydrophobicity of the electrode, and the final electrode was able to tolerate a differential pressure of more than 70 psi without degrading. Further, we found that the deposition of the electrocatalyst with the method reported in this study did not lead to a blockage in the air-channel pores of the substrate. The operational evaluation showed promising results about 2e-ORR’s ability to produce H2O2. The efficacy of the electrode for H2O2 generation can be enhanced by increasing the catalyst loading (up to saturation level) and providing more electrical current. 101 Chapter 5: Conclusions and Recommendations 5.1 Conclusions The overall objective of the present research was to develop a practical point-of-use (POU) system for the degradation of organic contaminants in drinking water. For this purpose, three distinct topics were explored. First, an electrochemical system was developed to produce hydrogen peroxide (H2O2) as a green oxidizing agent. The system was coupled with ultraviolet light-emitting diodes (UV-LEDs) for the degradation of 2,4-dichlorophenoxyacetic acid (2,4-D), as a model organic contaminant, in an advanced oxidation process. The generation of H2O2 was studied thoroughly under various configurational and operational parameters. The performance of a new UV source, called microplasma UV lamp, when integrated with the in situ generation of H2O2, was investigated. Besides, the efficiency of microplasma UV lamps for the photoelectro-Fenton (PEF) process was assessed. In addition, a new generation of gas diffusion electrodes (GDEs), based on expanded polytetrafluoroethylene (ePTFE) substrate as a leak-resistant material, was developed. Several characterization techniques and experimental procedures proved the applicability of these electrodes for two-electron oxygen reduction reaction (2e-ORR). Based on the research findings, the following conclusions were obtained: 5.1.1 Chapter 2 • An electrochemical cell was developed for the in situ generation of H2O2 from water and oxygen. This process eliminated the need to add H2O2 to a photoreactor externally, thus reduced the expenses and risks associated with the transportation of H2O2. The effect of water flow rate and current density on the output of the system (i.e., the concentration of 102 H2O2) was investigated and in-depth discussions were presented. In general, both a higher current density and a lower flow rate led to an increased concentration of H2O2. At a flow rate of 50 mL min–1 and the applied current density of 134 A m–2, a cathodic H2O2 generation rate of ~8 mg cm–2 h–1 was obtained. The experiments showed that the Faraday efficiency of 2e-ORR can be up to 90%, which proved that the design of the system was highly efficient. The electric energy consumption (EEC) of the system was calculated to be 14 kWh per kg of H2O2. It was also found that an anodic reaction can decompose the excess H2O2 in the effluent to some degree. • The electrochemical cell was then coupled with a UV-LED array to investigate the UV/H2O2 process to produce hydroxyl radicals (·OH) and degrade organic contaminants. The degradation of 2,4-D was studied as a model organic contaminant. At a flow rate of 50 mL min–1, an applied current density of 134 A m–2, an initial 2,4-D concentration of 1 mg L–1, and a total UV output of 900 mW, a degradation percentage of >70% was obtained. • Utilizing the one-factor-at-a-time (OFAT) method, the effect of various operational parameters, such as the UV fluence rate, the initial concentration of 2,4-D, and the initial pH, on the performance of the system was explored. A higher UV fluence rate, a lower contaminant initial concentration, and a pH closer to the pKa of 2,4-D, all contributed to a higher degradation of the organic compound. An empirical equation was suggested to anticipate the degradation percentages as a function of the initial concentration of 2,4-D (logarithmic relationship) and UV fluence rate (linear relationship). Moreover, we found that the organic molecule can undergo a reversible reduction/oxidation in the cathodic and anodic chambers, respectively. In addition, the energy analysis proved that the system is 103 energy efficient and could be improved using passive cooling for UV-LEDs. Stability analysis showed that the system is durable for extended periods of time. • In summary, the UV-LED based electrochemical photoreactor developed in this project demonstrated promising results, making it suitable for POU water treatment applications. In addition, since UV radiation is used in the photoreactor, pathogenic microorganisms will also be inactivated besides chemical contaminant degradation. 5.1.2 Chapter 3 • Effect of the fluid stream and configuration conditions on the performance of the electrochemical cell developed in Chapter 2 was investigated. The type and concentration of the supporting electrolyte did not exhibit a considerable effect on the cathodic H2O2 generation. When chloride ions were present in the fluid, they raised the removal of excess H2O2 in the anodic chamber by producing hypochlorous acid (HOCl). In addition, a low to room temperature was found to be the most suitable region for the electrochemical H2O2 production. The distance between the anode and the cathode did not affect H2O2 generation rate significantly, however, at the shortest distance (i.e., 1 cm), the minimum EEC was obtained. At an electrolyte flow rate of 50 mL min–1 and an applied current density of 134 A m–2, a H2O2 concentration of 1.1 mM was obtained with a Faraday efficiency of close to 90%, which demonstrated the efficiency of the electrochemical system for H2O2 synthesis. • The device was then integrated with newly emerged microplasma UV lamps to produce ·OH. Three microplasma lamps were studied in this work, two with monochromatic wavelengths of 172 nm and 222 nm, and one with a polychromatic pattern of 220–280 nm. The degradation of methylene blue (MB) as a model organic contaminant was monitored. 104 The three aforementioned lamps yielded a removal of 5.7%, 22.5%, and 33.3%, respectively, with a H2O2 concentration of 1.1 mM. We found that the microplasma lamps can directly photolyze MB without the electrogeneration of H2O2, but to a lesser degree. • Electro-Fenton (EF) and PEF processes were also studied with the electrochemical cell in order to enhance the rate of ·OH generation. Ferrous (Fe2+) ions were added externally, and UV radiation from the microplasma lamps was applied in the PEF process. The results revealed that UV radiation at a wavelength of 222 nm can significantly improve the performance of the PEF process, with a MB degradation of 45.2%. Besides, the effect of the initial pH of the influent, as well as the initial concentration of Fe2+ on the efficacy of EF and PEF processes were investigated. A pH of 3 and Fe2+ concentration of 0.5 mM showed the highest degradation of MB, which was >90%. • To summarize, the electrochemical H2O2 generation system equipped with microplasma UV lamps exhibited successful in situ production of ·OH and degradation of low concentrations of organic contaminants, indicating its suitability for a POU water treatment system. 5.1.3 Chapter 4 • During the experiments performed in Chapters 2 and 3, we found that the GDEs used as the cathode were susceptible to leaking when the pressure of the fluid inside the system was slightly higher than the pressure of ambient air. This issue would prohibit developing a practical system. Therefore, in the Chapter 4, a new type of GDE was developed for the cathodic H2O2 production which was leak resistant. The electrode used ePTFE as the substrate, carbon black nanoparticles as the electrocatalyst, and Nafion perfluorinated resin 105 solution as the binder. After several trials, we found airbrushing to be the most effective coating method. The detailed fabrication protocol was reported. • The developed electrode was fully characterized using several characterization techniques, such as scanning electron microscopy (SEM) imaging, contact angle measurement, water entry pressure (WEP) measurement, and mercury intrusion porosimetry (MIP). SEM images revealed that the electrocatalyst penetrated the pores of the substrate. Moreover, the contact angle measurement proved that the Nafion binder increased the hydrophobicity of the electrode. WEP test indicated that the fabricated electrode was able to tolerate a differential pressure of more than 70 psi without leaking and degrading. Besides, we found that the deposition of the electrocatalyst did not lead to a blockage in the air-channel pores of the substrate. • The operational performance of the ePTFE-based electrodes for 2e-ORR was assessed based on catalyst loading, electrolyte flow rate, and applied current density. We found that the electrochemical cell equipped with ePTFE-based electrodes as the cathode can produce 0.6 mM H2O2 using a catalyst loading of 0.22 mg cm–2, flow rate of 40 mL min–1, and applied current density of ~200 A m–2. The operational evaluation exhibited promising results for 2e-ORR, and the cathodic concentration of H2O2 can be enhanced by increasing the catalyst loading (up to saturation level), applying more electric current, and operating at a lower flow rate. • Overall, the GORE-TEX based GDE developed in this work can provide a means of viable flood-free H2O2 generation via 2e-ORR. 106 5.2 Recommendations The followings are recommended to be done in future research and practical works: • It is suggested that the performance of a GORE-TEX based cathode be assessed using real water samples. In the current study, a dilute electrolyte was used to simulate tap water with a relatively similar conductivity, however, the existence of other chemicals in water, such as carbonate, might affect ·OH production. • The performance of the ePTFE-based electrode for 2e-ORR can be optimized using various novel electrocatalyst types and loadings, other than the values studied in this work. • The radicals formed during the H2O2 generation and UV irradiation can be identified using various radicals’ scavengers. Therefore, an in-depth study of the radicals formed in various parts of the system is recommended. • In a future study, the efficacy of this system to reduce the total organic carbon (TOC) and chemical oxygen demand (COD) of the influent is recommended to be investigated. • The performance of the current system against pathogenic contaminants (i.e., disinfection) can be assessed in various modes of H2O2 production, UV irradiation, and the combination of both processes. Moreover, the microorganisms might be susceptible to the electrochemical reactions, which can be confirmed in a detailed study. • A set of systematic experiments can be performed to investigate the long-term stability of the ePTFE-based electrodes. • A simulation of the whole system, including the UV radiation, fluid hydrodynamics, and electrochemical and photochemical reactions kinetics, can be performed in order to obtain a comprehensive model which can be used for a better understanding of the involved phenomena, design optimization, and scalability of the system. Also, the computer-based 107 model can be used to do virtual prototyping on new reactor designs, such as cylindrical and a single chamber design, which could have a smaller footprint and a better performance. • It is advised to perform an economic analysis of the system, considering the cost of various items. • During all the experiments conducted in this research, the pathway of water was from the cathodic chamber to the anodic chamber. However, reversing this path might be beneficial under certain conditions, as the anodic reaction will lower the pH of the liquid, leading to a higher cathodic generation of H2O2. Although, in the new proposed pathway, the excess H2O2 cannot be removed on the surface of the anode anymore. Further studies can give a definite answer to the suitability of this idea. • The microplasma lamp generates UV radiation from both sides, with the radiant power of the back side being ~80% of the power in front. Therefore, a new reactor may be designed to utilize the UV radiation of these lamps from both sides, thus compensating the relatively low output of these lamps. • The inner faces of the photoreactor which are exposed to UV rays can be coated with a photocatalyst, preferably titanium dioxide (TiO2), so that a higher degradation of organic compounds can be achieved. • In this study, active cooling (i.e., heatsink and DC brushless fan) was applied for the thermal management of the UV-LED array. However, this method can be substituted by passive cooling, that is, to use the flow of water as the coolant. Thanks to its extraordinary heat capacity, water is considered as one of the most prominent cooling materials. Changing the active cooling method by the passive cooling method will decrease the energy consumption of the system significantly. 108 • Granular activated carbon (GAC) is used in water treatment plants to absorb and decompose residual H2O2 remained after the UV/H2O2 process. Therefore, a GAC filter can be inserted at the end of the electrochemical cell to remove excess H2O2. • The electrochemical device can be evaluated for removing heavy metals from drinking water. Since an electrochemical system forms an electric field between the anode and the cathode, the positively charged cations can get absorbed, reduced, and electrodeposited on the surface of the cathode. Given that the specific surface area of the GORE-TEX based electrodes is high, a saturation of the electrode might not be foreseeable. • The effect of pulsed applied current to the cell on H2O2 production can be evaluated. Moreover, the response of organic contaminant degradation to a pulsed UV radiation can be explored. 109 References (1) UNESCO. World Water Development Report Volume 4: Managing Water under Uncertainty and Risk; 2012; Vol. 1. (2) Masten, S. J.; Davies, S. H.; McElmurry, S. P. Flint Water Crisis: What Happened and Why? J. Am. Water Works Assoc. 2016, 108 (12), 22–34. https://doi.org/10.5942/jawwa.2016.108.0195. (3) Human Rights Watch. Make It Safe: Canada’s Obligation to End the First Nations Water Crisis; Human Rights Watch, 2016. (4) Abbaszadegan, M.; Hasan, M. N.; Gerba, C. P.; Roessler, P. F.; Wilson, B. R.; Kuennen, R.; Van Dellen, E. The Disinfection Efficacy of a Point-of-Use Water Treatment System against Bacterial, Viral and Protozoan Waterborne Pathogens. Water Res. 1997, 31 (3), 574–582. https://doi.org/10.1016/S0043-1354(96)00263-1. (5) Assumpão, M. H. M. T.; Moraes, A.; De Souza, R. F. B.; Gaubeur, I.; Oliveira, R. T. S.; Antonin, V. S.; Malpass, G. R. P.; Rocha, R. S.; Calegaro, M. L.; Lanza, M. R. V.; Santos, M. C. Low Content Cerium Oxide Nanoparticles on Carbon for Hydrogen Peroxide Electrosynthesis. Appl. Catal. A Gen. 2012, 411–412, 1–6. https://doi.org/10.1016/j.apcata.2011.09.030. (6) Stirling, R.; Walker, W. S.; Westerhoff, P.; Garcia-Segura, S. Techno-Economic Analysis to Identify Key Innovations Required for Electrochemical Oxidation as Point-of-Use Treatment Systems. Electrochim. Acta 2020. https://doi.org/10.1016/j.electacta.2020.135874. (7) Andreozzi, R. Advanced Oxidation Processes (AOP) for Water Purification and Recovery. 110 Catal. Today 1999, 53 (1), 51–59. https://doi.org/10.1016/S0920-5861(99)00102-9. (8) Campos-Martin, J. M.; Blanco-Brieva, G.; Fierro, J. L. G. Hydrogen Peroxide Synthesis: An Outlook beyond the Anthraquinone Process. Angew. Chemie Int. Ed. 2006, 45 (42), 6962–6984. https://doi.org/10.1002/anie.200503779. (9) Yeager, E. Electrocatalysts for O2 Reduction. Electrochim. Acta 1984, 29 (11), 1527–1537. https://doi.org/10.1016/0013-4686(84)85006-9. (10) Drogui, P.; Elmaleh, S.; Rumeau, M.; Bernard, C.; Rambaud, A. Hydrogen Peroxide Production by Water Electrolysis: Application to Disinfection. J. Appl. Electrochem. 2001, 31 (8), 877–882. https://doi.org/10.1023/A:1017588221369. (11) Özcan, A.; Şahin, Y.; Savaş Koparal, A.; Oturan, M. A. Carbon Sponge as a New Cathode Material for the Electro-Fenton Process: Comparison with Carbon Felt Cathode and Application to Degradation of Synthetic Dye Basic Blue 3 in Aqueous Medium. J. Electroanal. Chem. 2008, 616 (1–2), 71–78. https://doi.org/10.1016/j.jelechem.2008.01.002. (12) Vyhnánková, E.; Kozáková, Z.; Krčma, F.; Hrdlička, A. Influence of Electrode Material on Hydrogen Peroxide Generation by DC Pinhole Discharge. Open Chem. 2015, 13 (1), 218–223. https://doi.org/10.1515/chem-2015-0054. (13) Márquez, A. A.; Sirés, I.; Brillas, E.; Nava, J. L. Mineralization of Methyl Orange Azo Dye by Processes Based on H2O2 Electrogeneration at a 3D-like Air-Diffusion Cathode. Chemosphere 2020, 259, 127466. https://doi.org/10.1016/j.chemosphere.2020.127466. (14) Pozzo, A. Da; Palma, L. Di; Merli, C.; Petrucci, E. An Experimental Comparison of a Graphite Electrode and a Gas Diffusion Electrode for the Cathodic Production of Hydrogen Peroxide. J. Appl. Electrochem. 2005, 35 (4), 413–419. https://doi.org/10.1007/s10800-111 005-0800-2. (15) Sim, J.; An, J.; Elbeshbishy, E.; Ryu, H.; Lee, H. S. Characterization and Optimization of Cathodic Conditions for H2O2 Synthesis in Microbial Electrochemical Cells. Bioresour. Technol. 2015, 195, 31–36. https://doi.org/10.1016/j.biortech.2015.06.076. (16) Shen, Z.; Yang, J.; Hu, X.; Lei, Y.; Ji, X.; Jia, J.; Wang, W. Dual Electrodes Oxidation of Dye Wastewater with Gas Diffusion Cathode. Environ. Sci. Technol. 2005, 39 (6), 1819–1826. https://doi.org/10.1021/es049025e. (17) Hou, H.; Zeng, X.; Zhang, X. Production of Hydrogen Peroxide by Photocatalytic Processes. Angew. Chemie Int. Ed. 2020, 59 (40), 17356–17376. https://doi.org/10.1002/anie.201911609. (18) Kaynan, N.; Berke, B. A.; Hazut, O.; Yerushalmi, R. Sustainable Photocatalytic Production of Hydrogen Peroxide from Water and Molecular Oxygen. J. Mater. Chem. A 2014, 2 (34), 13822–13826. https://doi.org/10.1039/C4TA03004D. (19) Viswanathan, V.; Hansen, H. A.; Nørskov, J. K. Selective Electrochemical Generation of Hydrogen Peroxide from Water Oxidation. J. Phys. Chem. Lett. 2015, 6 (21), 4224–4228. https://doi.org/10.1021/acs.jpclett.5b02178. (20) Wu, X.; Zhang, X.; Zhao, S.; Gong, Y.; Djellabi, R.; Lin, S.; Zhao, X. Highly-Efficient Photocatalytic Hydrogen Peroxide Production over Polyoxometalates Covalently Immobilized onto Titanium Dioxide. Appl. Catal. A Gen. 2020, 591, 117271. https://doi.org/10.1016/j.apcata.2019.117271. (21) Li, W.; Bonakdarpour, A.; Gyenge, E.; Wilkinson, D. P. Production of Hydrogen Peroxide for Drinking Water Treatment in a Proton Exchange Membrane Electrolyzer at Near-Neutral PH. J. Electrochem. Soc. 2020, 167 (4), 044502. https://doi.org/10.1149/1945-112 7111/ab6fee. (22) Li, W.; Bonakdarpour, A.; Gyenge, E.; Wilkinson, D. P. Design of Bifunctional Electrodes for Co-Generation of Electrical Power and Hydrogen Peroxide. J. Appl. Electrochem. 2018, 48 (9), 985–993. https://doi.org/10.1007/s10800-018-1232-0. (23) Xia, C.; Xia, Y.; Zhu, P.; Fan, L.; Wang, H. Direct Electrosynthesis of Pure Aqueous H2O2 Solutions up to 20% by Weight Using a Solid Electrolyte. Science (80-. ). 2019, 366 (6462), 226–231. https://doi.org/10.1126/science.aay1844. (24) Li, W.; Bonakdarpour, A.; Gyenge, E.; Wilkinson, D. P. Drinking Water Purification by Electrosynthesis of Hydrogen Peroxide in a Power-Producing PEM Fuel Cell. ChemSusChem 2013, 6 (11), 2137–2143. https://doi.org/10.1002/cssc.201300225. (25) Nguyen, T. N.; Dinh, C.-T. Gas Diffusion Electrode Design for Electrochemical Carbon Dioxide Reduction. Chem. Soc. Rev. 2020, 49 (21), 7488–7504. https://doi.org/10.1039/D0CS00230E. (26) Zhang, Q.; Zhou, M.; Ren, G.; Li, Y.; Li, Y.; Du, X. Highly Efficient Electrosynthesis of Hydrogen Peroxide on a Superhydrophobic Three-Phase Interface by Natural Air Diffusion. Nat. Commun. 2020, 11 (1), 1731. https://doi.org/10.1038/s41467-020-15597-y. (27) Ihonen, J.; Mikkola, M.; Lindbergh, G. Flooding of Gas Diffusion Backing in PEFCs. J. Electrochem. Soc. 2004, 151 (8), A1152. https://doi.org/10.1149/1.1763138. (28) Leonard, M. E.; Clarke, L. E.; Forner-Cuenca, A.; Brown, S. M.; Brushett, F. R. Investigating Electrode Flooding in a Flowing Electrolyte, Gas-Fed Carbon Dioxide Electrolyzer. ChemSusChem 2020, 13 (2), 400–411. https://doi.org/10.1002/cssc.201902547. (29) Nishida, K.; Murakami, T.; Tsushima, S.; Hirai, S. Measurement of Liquid Water Content 113 in Cathode Gas Diffusion Electrode of Polymer Electrolyte Fuel Cell. J. Power Sources 2010, 195 (11), 3365–3373. https://doi.org/10.1016/j.jpowsour.2009.12.073. (30) Lin, G.; Nguyen, T. Van. Effect of Thickness and Hydrophobic Polymer Content of the Gas Diffusion Layer on Electrode Flooding Level in a PEMFC. J. Electrochem. Soc. 2005, 152 (10), A1942. https://doi.org/10.1149/1.2006487. (31) Chatterley, C.; Linden, K. Demonstration and Evaluation of Germicidal UV-LEDs for Point-of-Use Water Disinfection. J. Water Health 2010, 8 (3), 479–486. https://doi.org/10.2166/wh.2010.124. (32) Vilhunen, S.; Särkkä, H.; Sillanpää, M. Ultraviolet Light-Emitting Diodes in Water Disinfection. Environ. Sci. Pollut. Res. 2009, 16 (4), 439–442. https://doi.org/10.1007/s11356-009-0103-y. (33) Chevremont, A. C.; Farnet, A. M.; Coulomb, B.; Boudenne, J. L. Effect of Coupled UV-A and UV-C LEDs on Both Microbiological and Chemical Pollution of Urban Wastewaters. Sci. Total Environ. 2012, 426, 304–310. https://doi.org/10.1016/j.scitotenv.2012.03.043. (34) Natarajan, T. S.; Thomas, M.; Natarajan, K.; Bajaj, H. C.; Tayade, R. J. Study on UV-LED/TiO2 Process for Degradation of Rhodamine B Dye. Chem. Eng. J. 2011, 169 (1–3), 126–134. https://doi.org/10.1016/j.cej.2011.02.066. (35) Malkhasian, A. Y. S.; Izadifard, M.; Achari, G.; Langford, C. H. Photocatalytic Degradation of Agricultural Antibiotics Using a UV-LED Light Source. J. Environ. Sci. Heal. - Part B Pestic. Food Contam. Agric. Wastes 2014, 49 (1), 35–40. https://doi.org/10.1080/03601234.2013.836871. (36) Raeiszadeh, M.; Taghipour, F. Microplasma UV Lamp as a New Source for UV-Induced Water Treatment: Protocols for Characterization and Kinetic Study. Water Res. 2019, 164, 114 114959. https://doi.org/10.1016/j.watres.2019.114959. (37) Bolton, J. R.; Bircher, K. G.; Tumas, W.; Tolman, C. A. Figures-of-Merit for the Technical Development and Application of Advanced Oxidation Processes. J. Adv. Oxid. Technol. 1996, 1 (1). https://doi.org/10.1515/jaots-1996-0104. (38) Dewil, R.; Mantzavinos, D.; Poulios, I.; Rodrigo, M. A. New Perspectives for Advanced Oxidation Processes. J. Environ. Manage. 2017, 195, 93–99. https://doi.org/10.1016/j.jenvman.2017.04.010. (39) Miklos, D. B.; Remy, C.; Jekel, M.; Linden, K. G.; Drewes, J. E.; Hübner, U. Evaluation of Advanced Oxidation Processes for Water and Wastewater Treatment – A Critical Review. Water Res. 2018, 139, 118–131. https://doi.org/10.1016/j.watres.2018.03.042. (40) Sheng, Y.; Song, S.; Wang, X.; Song, L.; Wang, C.; Sun, H.; Niu, X. Electrogeneration of Hydrogen Peroxide on a Novel Highly Effective Acetylene Black-PTFE Cathode with PTFE Film. Electrochim. Acta 2011, 56 (24), 8651–8656. https://doi.org/10.1016/j.electacta.2011.07.069. (41) Oturan, M. A.; Aaron, J.-J. Advanced Oxidation Processes in Water/Wastewater Treatment: Principles and Applications. A Review. Crit. Rev. Environ. Sci. Technol. 2014, 44 (23), 2577–2641. https://doi.org/10.1080/10643389.2013.829765. (42) von Gunten, U. Oxidation Processes in Water Treatment: Are We on Track? Environ. Sci. Technol. 2018, 52 (9), 5062–5075. https://doi.org/10.1021/acs.est.8b00586. (43) Yi, Y.; Wang, L.; Li, G.; Guo, H. A Review on Research Progress in the Direct Synthesis of Hydrogen Peroxide from Hydrogen and Oxygen: Noble-Metal Catalytic Method, Fuel-Cell Method and Plasma Method. Catal. Sci. Technol. 2016, 6 (6), 1593–1610. https://doi.org/10.1039/C5CY01567G. 115 (44) Zhou, W.; Meng, X.; Gao, J.; Alshawabkeh, A. N. Hydrogen Peroxide Generation from O2 Electroreduction for Environmental Remediation: A State-of-the-Art Review. Chemosphere 2019, 225, 588–607. https://doi.org/10.1016/j.chemosphere.2019.03.042. (45) Yu, X.; Zhou, M.; Ren, G.; Ma, L. A Novel Dual Gas Diffusion Electrodes System for Efficient Hydrogen Peroxide Generation Used in Electro-Fenton. Chem. Eng. J. 2015, 263, 92–100. https://doi.org/10.1016/j.cej.2014.11.053. (46) Barros, W. R. P.; Ereno, T.; Tavares, A. C.; Lanza, M. R. V. In Situ Electrochemical Generation of Hydrogen Peroxide in Alkaline Aqueous Solution by Using an Unmodified Gas Diffusion Electrode. ChemElectroChem 2015, 2 (5), 714–719. https://doi.org/10.1002/celc.201402426. (47) Perry, S. C.; Pangotra, D.; Vieira, L.; Csepei, L.-I.; Sieber, V.; Wang, L.; Ponce de León, C.; Walsh, F. C. Electrochemical Synthesis of Hydrogen Peroxide from Water and Oxygen. Nat. Rev. Chem. 2019, 3 (7), 442–458. https://doi.org/10.1038/s41570-019-0110-6. (48) Kneissl, M.; Seong, T.-Y.; Han, J.; Amano, H. The Emergence and Prospects of Deep-Ultraviolet Light-Emitting Diode Technologies. Nat. Photonics 2019, 13 (4), 233–244. https://doi.org/10.1038/s41566-019-0359-9. (49) Kheyrandish, A.; Mohseni, M.; Taghipour, F. Development of a Method for the Characterization and Operation of UV-LED for Water Treatment. Water Res. 2017, 122, 570–579. https://doi.org/10.1016/j.watres.2017.06.015. (50) Keshavarzfathy, M.; Taghipour, F. Radiation Modeling of Ultraviolet Light-Emitting Diode (UV-LED) for Water Treatment. J. Photochem. Photobiol. A Chem. 2019, 377, 58–66. https://doi.org/10.1016/j.jphotochem.2019.03.030. (51) Song, K.; Mohseni, M.; Taghipour, F. Application of Ultraviolet Light-Emitting Diodes 116 (UV-LEDs) for Water Disinfection: A Review. Water Res. 2016, 94, 341–349. https://doi.org/10.1016/j.watres.2016.03.003. (52) Matafonova, G.; Batoev, V. Recent Advances in Application of UV Light-Emitting Diodes for Degrading Organic Pollutants in Water through Advanced Oxidation Processes: A Review. Water Res. 2018, 132, 177–189. https://doi.org/10.1016/j.watres.2017.12.079. (53) Maleki, N.; Safavi, A.; Shahbaazi, H. R. Electrochemical Determination of 2,4-D at a Mercury Electrode. Anal. Chim. Acta 2005, 530 (1), 69–74. https://doi.org/10.1016/j.aca.2004.08.058. (54) Samir, R.; Essam, T.; Ragab, Y.; Hashem, A. Enhanced Photocatalytic–Biological Degradation of 2,4 Dichlorophenoxyacetic Acid. Bull. Fac. Pharmacy, Cairo Univ. 2015, 53 (2), 77–82. https://doi.org/10.1016/j.bfopcu.2015.03.002. (55) Silva, T. M.; Stets, M. I.; Mazzetto, A. M.; Andrade, F. D.; Pileggi, S. A. V.; Fávero, P. R.; Cantú, M. D.; Carrilho, E.; Carneiro, P. I. B.; Pileggi, M. Degradation of 2,4-D Herbicide by Microorganisms Isolated from Brazilian Contaminated Soil. Brazilian J. Microbiol. 2007, 38 (3), 522–525. https://doi.org/10.1590/S1517-83822007000300026. (56) Developing High-power SMD package MODEL xFxZV-0F001 3.5 x 3.5mm SMD Flat Top Type https://www.ultraviolet-led.com/dcms_media/other/CD-7-2-11-ZV0F001 Rev0.2.pdf (accessed May 28, 2020). (57) Pujiastuti, S.; Onggo, H. Effect of Various Concentration of Sulfuric Acid for Nafion Membrane Activation on the Performance of Fuel Cell. In AIP Conference Proceedings; 2016; Vol. 1711, p 060006. https://doi.org/10.1063/1.4941639. (58) Qiang, Z.; Chang, J.-H.; Huang, C.-P. Electrochemical Generation of Hydrogen Peroxide from Dissolved Oxygen in Acidic Solutions. Water Res. 2002, 36 (1), 85–94. 117 https://doi.org/10.1016/S0043-1354(01)00235-4. (59) Lim, J.; Hoffmann, M. R. Substrate Oxidation Enhances the Electrochemical Production of Hydrogen Peroxide. Chem. Eng. J. 2019, 374, 958–964. https://doi.org/10.1016/j.cej.2019.05.165. (60) Khataee, A. R.; Safarpour, M.; Zarei, M.; Aber, S. Electrochemical Generation of H2O2 Using Immobilized Carbon Nanotubes on Graphite Electrode Fed with Air: Investigation of Operational Parameters. J. Electroanal. Chem. 2011, 659 (1), 63–68. https://doi.org/10.1016/j.jelechem.2011.05.002. (61) Chu, W. Modeling the Quantum Yields of Herbicide 2,4-D Decay in UV/H2O2 Process. Chemosphere 2001, 44 (5), 935–941. https://doi.org/10.1016/S0045-6535(00)00556-7. (62) Cruz González, G.; Julcour, C.; Chaumat, H.; Jáuregui-Haza, U.; Delmas, H. Degradation of 2,4-Dichlorophenoxyacetic Acid by Photolysis and Photo-Fenton Oxidation. J. Environ. Chem. Eng. 2018, 6 (1), 874–882. https://doi.org/10.1016/j.jece.2017.12.049. (63) Rezaei, R.; Mohseni, M. Impact of Natural Organic Matter on the Degradation of 2,4-Dichlorophenoxy Acetic Acid in a Fluidized Bed Photocatalytic Reactor. Chem. Eng. J. 2017, 310, 457–463. https://doi.org/10.1016/j.cej.2016.05.086. (64) Alfano, O. M.; Brandi, R. J.; Cassano, A. E. Degradation Kinetics of 2,4-D in Water Employing Hydrogen Peroxide and UV Radiation. Chem. Eng. J. 2001, 82 (1–3), 209–218. https://doi.org/10.1016/S1385-8947(00)00358-2. (65) Adak, A.; Das, I.; Mondal, B.; Koner, S.; Datta, P.; Blaney, L. Degradation of 2,4-Dichlorophenoxyacetic Acid by UV 253.7 and UV-H2O2: Reaction Kinetics and Effects of Interfering Substances. Emerg. Contam. 2019, 5, 53–60. https://doi.org/10.1016/j.emcon.2019.02.004. 118 (66) Badellino, C.; Rodrigues, C. A.; Bertazzoli, R. Oxidation of Pesticides by in Situ Electrogenerated Hydrogen Peroxide: Study for the Degradation of 2,4-Dichlorophenoxyacetic Acid. J. Hazard. Mater. 2006, 137 (2), 856–864. https://doi.org/10.1016/j.jhazmat.2006.03.035. (67) Popp, F. D.; Schultz, H. P. Electrolytic Reduction of Organic Compounds. Chem. Rev. 1962, 62 (1), 19–40. https://doi.org/10.1021/cr60215a002. (68) Harrison, J. A.; Shoesmith, D. W. The Electrochemical Reduction of Organic Acids. J. Electroanal. Chem. Interfacial Electrochem. 1971, 32 (1), 125–135. https://doi.org/10.1016/S0022-0728(71)80239-5. (69) Fontmorin, J.-M.; Huguet, S.; Fourcade, F.; Geneste, F.; Floner, D.; Amrane, A. Electrochemical Oxidation of 2,4-Dichlorophenoxyacetic Acid: Analysis of by-Products and Improvement of the Biodegradability. Chem. Eng. J. 2012, 195–196, 208–217. https://doi.org/10.1016/j.cej.2012.04.058. (70) Health Canada. Guidelines for Canadian Drinking Water Quality—Summary Table. Water and Air Quality Bureau, Healthy Environments and Consumer Safety Branch, Health Canada, Ottawa, Ontario; 2019. (71) Oh, K.-H.; Tuovinen, O. H. Degradation of 2,4-Dichlorophenoxyacetic Acid by Mixed Cultures of Bacteria. J. Ind. Microbiol. 1990, 6 (4), 275–278. https://doi.org/10.1007/BF01575873. (72) Wang, H.; Wang, J. Comparative Study on Electrochemical Degradation of 2,4-Dichlorophenol by Different Pd/C Gas-Diffusion Cathodes. Appl. Catal. B Environ. 2009, 89 (1–2), 111–117. https://doi.org/10.1016/j.apcatb.2008.12.003. (73) Trivedi, N. S.; Mandavgane, S. A. Fundamentals of 2, 4 Dichlorophenoxyacetic Acid 119 Removal from Aqueous Solutions. Sep. Purif. Rev. 2018, 47 (4), 337–354. https://doi.org/10.1080/15422119.2018.1450765. (74) O’Shea, K. E.; Dionysiou, D. D. Advanced Oxidation Processes for Water Treatment. J. Phys. Chem. Lett. 2012, 3 (15), 2112–2113. https://doi.org/10.1021/jz300929x. (75) Comninellis, C.; Kapalka, A.; Malato, S.; Parsons, S. A.; Poulios, I.; Mantzavinos, D. Advanced Oxidation Processes for Water Treatment: Advances and Trends for R&D. J. Chem. Technol. Biotechnol. 2008, 83 (6), 769–776. https://doi.org/10.1002/jctb.1873. (76) Sievers, M. Advanced Oxidation Processes. In Treatise on Water Science; Elsevier, 2011; Vol. 4, pp 377–408. https://doi.org/10.1016/B978-0-444-53199-5.00093-2. (77) Kanakaraju, D.; Glass, B. D.; Oelgemöller, M. Advanced Oxidation Process-Mediated Removal of Pharmaceuticals from Water: A Review. J. Environ. Manage. 2018, 219, 189–207. https://doi.org/10.1016/j.jenvman.2018.04.103. (78) Drzewicz, P.; Afzal, A.; El-Din, M. G.; Martin, J. W. Degradation of a Model Naphthenic Acid, Cyclohexanoic Acid, by Vacuum UV (172 Nm) and UV (254 Nm)/H2O2. J. Phys. Chem. A 2010, 114 (45), 12067–12074. https://doi.org/10.1021/jp105727s. (79) Song, K.; Taghipour, F.; Mohseni, M. Microorganisms Inactivation by Continuous and Pulsed Irradiation of Ultraviolet Light-Emitting Diodes (UV-LEDs). Chem. Eng. J. 2018, 343, 362–370. https://doi.org/10.1016/j.cej.2018.03.020. (80) Keshavarzfathy, M.; Taghipour, F. Computational Modeling of Ultraviolet Light-Emitting Diode (UV-LED) Reactor for Water Treatment. Water Res. 2019, 166, 115022. https://doi.org/10.1016/j.watres.2019.115022. (81) Chen, J.; Loeb, S.; Kim, J.-H. LED Revolution: Fundamentals and Prospects for UV Disinfection Applications. Environ. Sci. Water Res. Technol. 2017, 3 (2), 188–202. 120 https://doi.org/10.1039/C6EW00241B. (82) Song, K.; Mohseni, M.; Taghipour, F. Mechanisms Investigation on Bacterial Inactivation through Combinations of UV Wavelengths. Water Res. 2019, 163, 114875. https://doi.org/10.1016/j.watres.2019.114875. (83) Banat, F.; Al-Asheh, S.; Al-Rawashdeh, M.; Nusair, M. Photodegradation of Methylene Blue Dye by the UV/H2O2 and UV/Acetone Oxidation Processes. Desalination 2005, 181 (1–3), 225–232. https://doi.org/10.1016/j.desal.2005.04.005. (84) Oturan, N.; Oturan, M. A. Electro-Fenton Process: Background, New Developments, and Applications. In Electrochemical Water and Wastewater Treatment; Elsevier, 2018; pp 193–221. https://doi.org/10.1016/B978-0-12-813160-2.00008-0. (85) Li, G.; Zhang, Y. Highly Selective Two-Electron Oxygen Reduction to Generate Hydrogen Peroxide Using Graphite Felt Modified with N-Doped Graphene in an Electro-Fenton System. New J. Chem. 2019, 43 (32), 12657–12667. https://doi.org/10.1039/C9NJ02601K. (86) Berl, E. A New Cathodic Process for the Production of H2O2. Trans. Electrochem. Soc. 1939, 76 (1), 359–369. https://doi.org/10.1149/1.3500291. (87) Thin 222 nm UV Lamps From Eden Park https://edenpark.com/ (accessed Jul 25, 2020). (88) Goldstein, S.; Rabani, J. The Ferrioxalate and Iodide–Iodate Actinometers in the UV Region. J. Photochem. Photobiol. A Chem. 2008, 193 (1), 50–55. https://doi.org/10.1016/j.jphotochem.2007.06.006. (89) Sosnin, E. A.; Zakharova, E. A.; Moskaleva, M. L.; Batalova, V. N. An Electrochemical Methanol Actinometer and Its Application to Measurements of the Intensity of Vacuum UV Radiation of a Xe*2 Excilamp. Instruments Exp. Tech. 2006, 49 (1), 101–105. https://doi.org/10.1134/S0020441206010131. 121 (90) Goldstein, S.; Aschengrau, D.; Diamant, Y.; Rabani, J. Photolysis of Aqueous H2O2: Quantum Yield and Applications for Polychromatic UV Actinometry in Photoreactors. Environ. Sci. Technol. 2007, 41 (21), 7486–7490. https://doi.org/10.1021/es071379t. (91) Heit, G.; Neuner, A.; Saugy, P.-Y.; Braun, A. M. Vacuum-UV (172 Nm) Actinometry. The Quantum Yield of the Photolysis of Water. J. Phys. Chem. A 1998, 102 (28), 5551–5561. https://doi.org/10.1021/jp980130i. (92) Kuhn, H. J.; Braslavsky, S. E.; Schmidt, R. Chemical Actinometry (IUPAC Technical Report). Pure Appl. Chem. 2004, 76 (12), 2105–2146. https://doi.org/10.1351/pac200476122105. (93) Klassen, N. V.; Marchington, D.; McGowan, H. C. E. H2O2 Determination by the I3- Method and by KMnO4 Titration. Anal. Chem. 1994, 66 (18), 2921–2925. https://doi.org/10.1021/ac00090a020. (94) Gyenge, E. L.; Oloman, C. W. Electrosynthesis of Hydrogen Peroxide in Acidic Solutions by Mediated Oxygen Reduction in a Three-Phase (Aqueous/Organic/Gaseous) System Part II: Experiments in Flow-by Fixed-Bed Electrochemical Cells with Three-Phase Flow. J. Appl. Electrochem. 2003, 33 (8), 665–674. https://doi.org/10.1023/A:1025094826791. (95) Sánchez-Sánchez, C. M.; Bard, A. J. Hydrogen Peroxide Production in the Oxygen Reduction Reaction at Different Electrocatalysts as Quantified by Scanning Electrochemical Microscopy. Anal. Chem. 2009, 81 (19), 8094–8100. https://doi.org/10.1021/ac901291v. (96) Zhou, W.; Gao, J.; Ding, Y.; Zhao, H.; Meng, X.; Wang, Y.; Kou, K.; Xu, Y.; Wu, S.; Qin, Y. Drastic Enhancement of H2O2 Electro-Generation by Pulsed Current for Ibuprofen Degradation: Strategy Based on Decoupling Study on H2O2 Decomposition Pathways. 122 Chem. Eng. J. 2018, 338, 709–718. https://doi.org/10.1016/j.cej.2017.12.152. (97) Reis, R. M.; Beati, A. A. G. F.; Rocha, R. S.; Assumpção, M. H. M. T.; Santos, M. C.; Bertazzoli, R.; Lanza, M. R. V. Use of Gas Diffusion Electrode for the In Situ Generation of Hydrogen Peroxide in an Electrochemical Flow-By Reactor. Ind. Eng. Chem. Res. 2012, 51 (2), 649–654. https://doi.org/10.1021/ie201317u. (98) Khataee, A.; Sajjadi, S.; Pouran, S. R.; Hasanzadeh, A.; Joo, S. W. A Comparative Study on Electrogeneration of Hydrogen Peroxide through Oxygen Reduction over Various Plasma-Treated Graphite Electrodes. Electrochim. Acta 2017, 244, 38–46. https://doi.org/10.1016/j.electacta.2017.05.069. (99) Umimoto, K.; Nagata, S.; Yanagida, J. Development of Device Producing Electrolyzed Water for Home Care. J. Phys. Conf. Ser. 2013, 450, 012035. https://doi.org/10.1088/1742-6596/450/1/012035. (100) Shams El Din, A. M.; Mohammed, R. A. Kinetics of the Reaction between Hydrogen Peroxide and Hypochlorite. Desalination 1998, 115 (2), 145–153. https://doi.org/10.1016/S0011-9164(98)00034-4. (101) Connick, R. E. The Interaction of Hydrogen Peroxide and Hypochlorous Acid in Acidic Solutions Containing Chloride Ion. J. Am. Chem. Soc. 1947, 69 (6), 1509–1514. https://doi.org/10.1021/ja01198a074. (102) Li, H.; Wang, H.; Qian, W.; Zhang, S.; Wessel, S.; Cheng, T. T. H.; Shen, J.; Wu, S. Chloride Contamination Effects on Proton Exchange Membrane Fuel Cell Performance and Durability. J. Power Sources 2011, 196 (15), 6249–6255. https://doi.org/10.1016/j.jpowsour.2011.04.018. (103) Foller, P. C.; Bombard, R. T. Processes for the Production of Mixtures of Caustic Soda and 123 Hydrogen Peroxide via the Reduction of Oxygen. J. Appl. Electrochem. 1995, 25, 613–627. https://doi.org/10.1007/BF00241923. (104) Jian-xiao, L. V.; Ying, C.; Guo-hong, X.; Ling-yun, Z.; Su-fen, W. Decoloration of Methylene Blue Simulated Wastewater Using a UV-H2O2 Combined System. J. Water Reuse Desalin. 2011, 1 (1), 45–51. https://doi.org/10.2166/wrd.2011.019. (105) Li, M.; Qiang, Z.; Pulgarin, C.; Kiwi, J. Accelerated Methylene Blue (MB) Degradation by Fenton Reagent Exposed to UV or VUV/UV Light in an Innovative Micro Photo-Reactor. Appl. Catal. B Environ. 2016, 187, 83–89. https://doi.org/10.1016/j.apcatb.2016.01.014. (106) Yao, J.; Wang, C. Decolorization of Methylene Blue with TiO2 Sol via UV Irradiation Photocatalytic Degradation. Int. J. Photoenergy 2010, 2010, 1–6. https://doi.org/10.1155/2010/643182. (107) Tayade, R. J.; Natarajan, T. S.; Bajaj, H. C. Photocatalytic Degradation of Methylene Blue Dye Using Ultraviolet Light Emitting Diodes. Ind. Eng. Chem. Res. 2009, 48 (23), 10262–10267. https://doi.org/10.1021/ie9012437. (108) Bolton, J. R.; Bircher, K. G.; Tumas, W.; Tolman, C. A. Figures-of-Merit for the Technical Development and Application of Advanced Oxidation Technologies for Both Electric- and Solar-Driven Systems (IUPAC Technical Report). Pure Appl. Chem. 2001, 73 (4), 627–637. https://doi.org/10.1351/pac200173040627. (109) Valim, R. B.; Reis, R. M.; Castro, P. S.; Lima, A. S.; Rocha, R. S.; Bertotti, M.; Lanza, M. R. V. Electrogeneration of Hydrogen Peroxide in Gas Diffusion Electrodes Modified with Tert-Butyl-Anthraquinone on Carbon Black Support. Carbon N. Y. 2013, 61, 236–244. https://doi.org/10.1016/j.carbon.2013.04.100. (110) Zhao, K.; Su, Y.; Quan, X.; Liu, Y.; Chen, S.; Yu, H. Enhanced H2O2 Production by 124 Selective Electrochemical Reduction of O2 on Fluorine-Doped Hierarchically Porous Carbon. J. Catal. 2018, 357, 118–126. https://doi.org/10.1016/j.jcat.2017.11.008. (111) Cornejo, O. M.; Sirés, I.; Nava, J. L. Electrosynthesis of Hydrogen Peroxide Sustained by Anodic Oxygen Evolution in a Flow-through Reactor. J. Electroanal. Chem. 2020, 873, 114419. https://doi.org/10.1016/j.jelechem.2020.114419. (112) Zhu, Q.; Pan, Z.; Hu, S.; Kim, J.-H. Cathodic Hydrogen Peroxide Electrosynthesis Using Anthraquinone Modified Carbon Nitride on Gas Diffusion Electrode. ACS Appl. Energy Mater. 2019, 2 (11), 7972–7979. https://doi.org/10.1021/acsaem.9b01445. (113) Giomo, M.; Buso, A.; Fier, P.; Sandonà, G.; Boye, B.; Farnia, G. A Small-Scale Pilot Plant Using an Oxygen-Reducing Gas-Diffusion Electrode for Hydrogen Peroxide Electrosynthesis. Electrochim. Acta 2008, 54 (2), 808–815. https://doi.org/10.1016/j.electacta.2008.06.038. (114) Kim, H. W.; Ross, M. B.; Kornienko, N.; Zhang, L.; Guo, J.; Yang, P.; McCloskey, B. D. Efficient Hydrogen Peroxide Generation Using Reduced Graphene Oxide-Based Oxygen Reduction Electrocatalysts. Nat. Catal. 2018, 1 (4), 282–290. https://doi.org/10.1038/s41929-018-0044-2. (115) Huang, D.; Wang, G.; Zhao, X.; Pang, W.; Wang, M.; Wang, D. Electrogeneration of Hydrogen Peroxide in Neutral Medium Using PVDF-Based Air Cathode for Removing Ofloxacin in Wastewater. J. Solid State Electrochem. 2019, 23 (5), 1455–1462. https://doi.org/10.1007/s10008-019-04241-5. (116) Hu, R.; Yao, F.; Wu, C.; Jin, C.; Guan, L. Enhancing the Production of Hydrogen Peroxide from Electrocatalytic Oxygen Reduction Reaction by Tailoring the Electronic States of Single-Walled Carbon Nanotubes: A Synergistic Effect from Interior Filling and Exterior 125 Oxidation. Sustain. Energy Fuels 2019, 3 (8), 1951–1956. https://doi.org/10.1039/C9SE00287A. (117) Li, D.; Zheng, T.; Liu, Y.; Hou, D.; Yao, K. K.; Zhang, W.; Song, H.; He, H.; Shi, W.; Wang, L.; Ma, J. A Novel Electro-Fenton Process Characterized by Aeration from inside a Graphite Felt Electrode with Enhanced Electrogeneration of H2O2 and Cycle of Fe3+/Fe2+. J. Hazard. Mater. 2020, 396, 122591. https://doi.org/10.1016/j.jhazmat.2020.122591. (118) Mortazavi, M.; Fazeli, A.; Moghaddam, S. Scalable Bonding of Nanofibrous Polytetrafluoroethylene (PTFE) Membranes on Microstructures. J. Micromechanics Microengineering 2018, 28 (1), 015001. https://doi.org/10.1088/1361-6439/aa8185. (119) Soto-Herranz, M.; Sánchez-Báscones, M.; Antolín-Rodríguez, J.; Conde-Cid, D.; Vanotti, M. Effect of the Type of Gas-Permeable Membrane in Ammonia Recovery from Air. Environments 2019, 6 (6), 70. https://doi.org/10.3390/environments6060070. (120) History of Expanded PTFE and W.L. Gore and Associates. In Introduction to Fluoropolymers; Ebnesajjad, S., Ed.; Elsevier, 2013; pp 37–52. https://doi.org/10.1016/B978-1-4557-7442-5.00004-8. (121) Rezaei, M.; Warsinger, D. M.; Lienhard V, J. H.; Duke, M. C.; Matsuura, T.; Samhaber, W. M. Wetting Phenomena in Membrane Distillation: Mechanisms, Reversal, and Prevention. Water Res. 2018, 139, 329–352. https://doi.org/10.1016/j.watres.2018.03.058. (122) Yazgan-Birgi, P.; Hassan Ali, M. I.; Arafat, H. A. Estimation of Liquid Entry Pressure in Hydrophobic Membranes Using CFD Tools. J. Memb. Sci. 2018, 552, 68–76. https://doi.org/10.1016/j.memsci.2018.01.061. (123) Franken, A. C. M.; Nolten, J. A. M.; Mulder, M. H. V.; Bargeman, D.; Smolders, C. A. Wetting Criteria for the Applicability of Membrane Distillation. J. Memb. Sci. 1987, 33 (3), 126 315–328. https://doi.org/10.1016/S0376-7388(00)80288-4. (124) Protective & Automotive Vents Technical Information | Water Entry Pressure Testing | Gore https://www.gore.com/resources/materials-technology-water-entry-pressure-testing (accessed Oct 19, 2020). (125) Cheng, S.; Liu, H.; Logan, B. E. Power Densities Using Different Cathode Catalysts (Pt and CoTMPP) and Polymer Binders (Nafion and PTFE) in Single Chamber Microbial Fuel Cells. Environ. Sci. Technol. 2006, 40 (1), 364–369. https://doi.org/10.1021/es0512071. (126) Rácz, G.; Kerker, S.; Kovács, Z.; Vatai, G.; Ebrahimi, M.; Czermak, P. Theoretical and Experimental Approaches of Liquid Entry Pressure Determination in Membrane Distillation Processes. Period. Polytech. Chem. Eng. 2014, 58 (2), 81–91. https://doi.org/10.3311/PPch.2179. (127) Alkhudhiri, A.; Darwish, N.; Hilal, N. Membrane Distillation: A Comprehensive Review. Desalination 2012, 287, 2–18. https://doi.org/10.1016/j.desal.2011.08.027. (128) Peralta, E.; Natividad, R.; Roa, G.; Marin, R.; Romero, R.; Pavon, T. A Comparative Study on the Electrochemical Production of H2O2 between BDD and Graphite Cathodes. Sustain. Environ. Res. 2013, 23 (4), 259–266. (129) Da Pozzo, A.; Petrucci, E.; Merli, C. Electrogeneration of Hydrogen Peroxide in Seawater and Application to Disinfection. J. Appl. Electrochem. 2008, 38 (7), 997–1003. https://doi.org/10.1007/s10800-008-9524-4. (130) Zhou, W.; Rajic, L.; Zhao, Y.; Gao, J.; Qin, Y.; Alshawabkeh, A. N. Rates of H2O2 Electrogeneration by Reduction of Anodic O2 at RVC Foam Cathodes in Batch and Flow-through Cells. Electrochim. Acta 2018, 277, 185–196. https://doi.org/10.1016/j.electacta.2018.04.174. 127 (131) Zarei, M.; Salari, D.; Niaei, A.; Khataee, A. Peroxi-Coagulation Degradation of C.I. Basic Yellow 2 Based on Carbon-PTFE and Carbon Nanotube-PTFE Electrodes as Cathode. Electrochim. Acta 2009, 54 (26), 6651–6660. https://doi.org/10.1016/j.electacta.2009.06.060. (132) Wang, A.; Qu, J.; Ru, J.; Liu, H.; Ge, J. Mineralization of an Azo Dye Acid Red 14 by Electro-Fenton’s Reagent Using an Activated Carbon Fiber Cathode. Dye. Pigment. 2005, 65 (3), 227–233. https://doi.org/10.1016/j.dyepig.2004.07.019. (133) Salari, D.; Niaei, A.; Khataee, A.; Zarei, M. Electrochemical Treatment of Dye Solution Containing C.I. Basic Yellow 2 by the Peroxi-Coagulation Method and Modeling of Experimental Results by Artificial Neural Networks. J. Electroanal. Chem. 2009, 629 (1–2), 117–125. https://doi.org/10.1016/j.jelechem.2009.02.002. (134) Kheyrandish, A.; Taghipour, F.; Mohseni, M. UV-LED Radiation Modeling and Its Applications in UV Dose Determination for Water Treatment. J. Photochem. Photobiol. A Chem. 2018, 352, 113–121. https://doi.org/10.1016/j.jphotochem.2017.10.047. (135) Keshavarzfathy, M.; Malayeri, A. H.; Mohseni, M.; Taghipour, F. UV-LED Fluence Determination by Numerical Method for Microbial Inactivation Studies. J. Photochem. Photobiol. A Chem. 2020, 392, 112406. https://doi.org/10.1016/j.jphotochem.2020.112406. (136) Rahn, R. O.; Stefan, M. I.; Bolton, J. R.; Goren, E.; Shaw, P.-S.; Lykke, K. R. Quantum Yield of the Iodide-Iodate Chemical Actinometer: Dependence on Wavelength and Concentration. Photochem. Photobiol. 2007, 78 (2), 146–152. https://doi.org/10.1562/0031-8655(2003)0780146QYOTIC2.0.CO2. (137) Jin, S.; Mofidi, A. A.; Linden, K. G. Polychromatic UV Fluence Measurement Using Chemical Actinometry, Biodosimetry, and Mathematical Techniques. J. Environ. Eng. 128 2006, 132 (8), 831–841. https://doi.org/10.1061/(ASCE)0733-9372(2006)132:8(831). (138) Chatterley, C.; Linden, K. G. UV-LED Irradiation Technology for Point-of-Use Water Disinfection. In Proceedings of the Water Environment Federation; 2009; Vol. 2009, pp 222–225. https://doi.org/10.2175/193864709793848176. (139) Bolton, J. R.; Stefan, M. I.; Shaw, P.-S.; Lykke, K. R. Determination of the Quantum Yields of the Potassium Ferrioxalate and Potassium Iodide–Iodate Actinometers and a Method for the Calibration of Radiometer Detectors. J. Photochem. Photobiol. A Chem. 2011, 222 (1), 166–169. https://doi.org/10.1016/j.jphotochem.2011.05.017. (140) Kheyrandish, A.; Mohseni, M.; Taghipour, F. Protocol for Determining Ultraviolet Light Emitting Diode (UV-LED) Fluence for Microbial Inactivation Studies. Environ. Sci. Technol. 2018, 52 (13), 7390–7398. https://doi.org/10.1021/acs.est.7b05797. (141) Bolton, J. R.; Linden, K. G. Standardization of Methods for Fluence (UV Dose) Determination in Bench-Scale UV Experiments. J. Environ. Eng. 2003, 129 (3), 209–215. https://doi.org/10.1061/(ASCE)0733-9372(2003)129:3(209). (142) Karci, A.; Arslan-Alaton, I.; Olmez-Hanci, T.; Bekbölet, M. Transformation of 2,4-Dichlorophenol by H2O2/UV-C, Fenton and Photo-Fenton Processes: Oxidation Products and Toxicity Evolution. J. Photochem. Photobiol. A Chem. 2012, 230 (1), 65–73. https://doi.org/10.1016/j.jphotochem.2012.01.003. (143) Zhu, K.; Baig, S. A.; Xu, J.; Sheng, T.; Xu, X. Electrochemical Reductive Dechlorination of 2,4-Dichlorophenoxyacetic Acid Using a Palladium/Nickel Foam Electrode. Electrochim. Acta 2012, 69, 389–396. https://doi.org/10.1016/j.electacta.2012.03.038. (144) Tsyganok, A.; Otsuka, K.; Yamanaka, I.; Plekhanov, V.; Kulikov, S. Selective Electrochemical Dehalogenation of 2,4-Dichlorophenoxyacetic Acid in MeCN at Room 129 Temperature. Chem. Lett. 1996, 25 (4), 261–262. https://doi.org/10.1246/cl.1996.261. (145) Kulikov, S. M.; Plekhanov, V. P.; Tsyganok, A. I.; Schlimm, C.; Heitz, E. Electrochemical Reductive Dechlorination of Chlororganic Compounds on Carbon Cloth and Metal-Modified Carbon Cloth Cathodes. Electrochim. Acta 1996, 41 (4), 527–531. https://doi.org/10.1016/0013-4686(95)00339-8. (146) Howard, K. J.; Sidwell, A. E. Chlorination of Phenols and Phenoxyacetic Acids with Sulfuryl Chloride. US4345097A, January 21, 1982. (147) Manske, R. H. F. Process of Preparing 2, 4-Dichlorophenoxyacetic Acid. US2471575A, September 4, 1949. (148) Ebel, E.; Bell, J.; Fries, A.; Kasey, C.; Berkebile, J. M. Preparation of 2,4-D(2,4-Dichlorophenoxyacetic Acid) from Phenol and Monochloracetic Acid. J. Chem. Educ. 1947, 24 (9), 449. https://doi.org/10.1021/ed024p449. (149) Harris, T. R.; Pagan, J. G.; Batoni, P. Optical and Fluidic Co-Design of a UV-LED Water Disinfection Chamber. ECS Trans. 2013, 45 (17), 11–18. https://doi.org/10.1149/04517.0011ecst. (150) Chen, Z.; Chen, S.; Siahrostami, S.; Chakthranont, P.; Hahn, C.; Nordlund, D.; Dimosthenis, S.; Nørskov, J. K.; Bao, Z.; Jaramillo, T. F. Development of a Reactor with Carbon Catalysts for Modular-Scale, Low-Cost Electrochemical Generation of H2O2. React. Chem. Eng. 2017, 2 (2), 239–245. https://doi.org/10.1039/C6RE00195E. (151) Park, Y.; Kim, D. Chemical Stability Enhancement of Nafion Membrane by Impregnation of a Novel Organic ·OH Radical Scavenger, 3,4-Dihydroxy-Cinnamic Acid. J. Memb. Sci. 2018, 566, 1–7. https://doi.org/10.1016/j.memsci.2018.08.063. (152) Mališ, J.; Mazúr, P.; Paidar, M.; Bystron, T.; Bouzek, K. Nafion 117 Stability under 130 Conditions of PEM Water Electrolysis at Elevated Temperature and Pressure. Int. J. Hydrogen Energy 2016, 41 (4), 2177–2188. https://doi.org/10.1016/j.ijhydene.2015.11.102. (153) Yi, Y.; Weinberg, G.; Prenzel, M.; Greiner, M.; Heumann, S.; Becker, S.; Schlögl, R. Electrochemical Corrosion of a Glassy Carbon Electrode. Catal. Today 2017. https://doi.org/10.1016/j.cattod.2017.07.013. (154) Moradi, F.; Dehghanian, C. Addition of IrO2 to RuO2+TiO2 Coated Anodes and Its Effect on Electrochemical Performance of Anodes in Acid Media. Prog. Nat. Sci. Mater. Int. 2014. https://doi.org/10.1016/j.pnsc.2014.03.008. (155) Duby, P. The History of Progress in Dimensionally Stable Anodes. JOM 1993, 45 (3), 41–43. https://doi.org/10.1007/BF03222350. (156) Krstić, V.; Pešovski, B. Reviews the Research on Some Dimensionally Stable Anodes (DSA) Based on Titanium. Hydrometallurgy 2019, 185, 71–75. https://doi.org/10.1016/j.hydromet.2019.01.018. (157) Li, J. (Elena); Zamyadi, A.; Hofmann, R. Effect of Granular Activated Carbon Type and Age on Quenching H2O2 Residuals after UV/H2O2 Drinking Water Treatment. J. Water Supply Res. Technol. 2016, 65 (1), 28–36. https://doi.org/10.2166/aqua.2015.134. (158) Huang, Y.; Nie, Z.; Wang, C.; Li, Y.; Xu, M.; Hofmann, R. Quenching H2O2 Residuals after UV/H2O2 Oxidation Using GAC in Drinking Water Treatment. Environ. Sci. Water Res. Technol. 2018, 4 (10), 1662–1670. https://doi.org/10.1039/C8EW00407B. (159) Autin, O.; Romelot, C.; Rust, L.; Hart, J.; Jarvis, P.; MacAdam, J.; Parsons, S. A.; Jefferson, B. Evaluation of a UV-Light Emitting Diodes Unit for the Removal of Micropollutants in Water for Low Energy Advanced Oxidation Processes. Chemosphere 2013, 92 (6), 745–131 751. https://doi.org/10.1016/j.chemosphere.2013.04.028. (160) Ibrahim, M. A. S.; MacAdam, J.; Autin, O.; Jefferson, B. Evaluating the Impact of LED Bulb Development on the Economic Viability of Ultraviolet Technology for Disinfection. Environ. Technol. 2014, 35 (4), 400–406. https://doi.org/10.1080/09593330.2013.829858. (161) Arques-Orobon, F. J.; Nuñez, N.; Vazquez, M.; Segura-Antunez, C.; González-Posadas, V. High-Power UV-LED Degradation: Continuous and Cycled Working Condition Influence. Solid. State. Electron. 2015, 111, 111–117. https://doi.org/10.1016/j.sse.2015.05.039. (162) Würtele, M. A.; Kolbe, T.; Lipsz, M.; Külberg, A.; Weyers, M.; Kneissl, M.; Jekel, M. Application of GaN-Based Ultraviolet-C Light Emitting Diodes – UV LEDs – for Water Disinfection. Water Res. 2011, 45 (3), 1481–1489. https://doi.org/10.1016/j.watres.2010.11.015. (163) Connolly, E. Reconsider UV-C LED lifetime for disinfection based on development decisions https://www.ledsmagazine.com/leds-ssl-design/packaged-leds/article/16695812/reconsider-uvc-led-lifetime-for-disinfection-based-on-development-decisions-magazine (accessed Jun 11, 2020). (164) Daghrir, R.; Drogui, P. Coupled Electrocoagulation–Electro-Fenton for Efficient Domestic Wastewater Treatment. Environ. Chem. Lett. 2013, 11 (2), 151–156. https://doi.org/10.1007/s10311-012-0390-2. 132 Appendices Appendix A A.1 UV-LED array Figure A.1(A) shows the UV-LED array mounted on the heatsink. Three branches of UV-LEDs were used in parallel, each containing four LEDs connected in series. The actual wavelength, measured with an Ocean Optics USB2000+ spectrometer (Dunedin, FL, USA) equipped with a Sony ILX511B CCD detector, was found to be 277 nm (Figure A.1(B)). Moreover, the full width at half maximum (FWHM) value of the radiation pattern as shown in Figure A.1(B) was calculated to be 10.7 nm. Table A.1 shows the nominal properties of the UVC-LED used in the UV-LED array, according to the datasheet provided by manufacturer.56 Figure A.1 (A) The UV-LED array used in this study. (B) The normalized relative optical intensity of the UV-LED as a function of wavelength, showing a monochromatic peak at 277 nm and having an FWHM value of 10.7 nm. 133 Table A.1 Nominal properties of a single UV-LED used in the UV-LED array. Model Peak wavelength (nm) Radiant flux (mW) Viewing angle (deg.) FWHM (nm) Forward voltage (V) Forward current (mA) DOWA DF7ZV 275 75 120 11 6.9 800 A.2 Absorption spectra of various solutions Figure A.2 shows the absorption spectra of different solutions, proving that the existence of Na2SO4 will not interfere with the HPLC photodiode array detector operation at 228 nm. 134 Figure A.2 Spectra of various solutions: (a) ultrapure water, (b) 25 mM Na2SO4, (c) 1 mg L–1 2,4-D, and (d) the initial solution, containing 1 mg L–1 2,4-D in 25 mM Na2SO4. A.3 H2O2 concentration measurement method Quantification of H2O2 was carried out according to a method described elsewhere.93 In summary, it is based on the reaction of H2O2 molecules with excess potassium iodide in the presence of an ammonium molybdate tetrahydrate catalyst to produce triiodide ions (I3−), as illustrated in Eq. A.1: H2O2 + 3I− + 2H+ → I3− + 2H2O Eq. A.1 135 The molar attenuation coefficient (also called molar extinction coefficient or molar absorption coefficient) of I3− was high at a wavelength of 352 nm (27,600 M–1 cm–1).88 Therefore, the concentration of generated I3− could be correlated to the initial concentration of H2O2, according to their stoichiometric coefficient and using the Beer-Lambert law. The UV-visible spectrophotometer was used to read the absorbances. Figure A.3(A) and (B) show the H2O2 molecular structure and absorption spectrum of 0.98 mM H2O2 in ultrapure water, respectively. Figure A.3 H2O2. (A) Molecular structure. (B) Absorption spectrum of 0.98 mM H2O2 in ultrapure water. A.4 2,4-D concentration measurement method Measurements of 2,4-D concentration were performed using a high-performance liquid chromatography (HPLC) system (Waters Alliance 2695 Separations Module; Waters Corporation, Milford, MA, USA) equipped with a Waters 2998 Photodiode Array (PDA) Detector and a Nova-Pak reversed-phase C18 column (pore size: 60 μm; particle size: 4 μm; inner diameter: 3.9 mm; length: 150 mm). The mobile phase consisted of 55% acetonitrile, 2% diluted glacial acetic acid (5% v/v in ultrapure water), and 40% ultrapure water and was flowing with a flow rate of 1 mL min–1, according to a method described elsewhere.65 The injection volume was 50 μL. Figure 136 A.4(A) and (B) illustrate the molecular structure and absorption spectrum of 1 mg L–1 2,4-D in ultrapure water, respectively. The absorption spectrum of the 2,4-D molecule shows two peaks: one minor peak at 282 nm and one major peak at 228 nm. Therefore, the PDA detector was set to read the absorbance at a wavelength of 228 nm, which corresponded to the major peak. Figure A.4 2,4-D. (A) Molecular structure. (B) Absorption spectrum of 1 mg L–1 2,4-D in ultrapure water. The degradation percentages were calculated using Eq. A.2 as follows: Degradation (%) =Ci − CfCi× 100% Eq. A.2 where Ci and Cf denote the initial and final concentrations of 2,4-D (mg L–1), respectively. Data analyses were performed using Empower 2 software, run on a Lenovo computer. A.5 Incident UV fluence rate measurement method Chemical actinometry is a widely accepted method for measuring the average fluence rate of a UV source on the surface of a solution, owing to its low cost, simplicity, and accuracy.134 It is based on the photochemical reaction of the actinometer by absorbing UV radiation. Various compounds can be used as actinometers, such as potassium ferrioxalate, uridine, and potassium 137 iodide/potassium iodate.92 In this study, the KI/KIO3 solution was chosen since this actinometer can absorb radiation with a wavelength of lower than 290 nm and will not absorb radiation with a wavelength of higher than 330 nm, so it is blind to room light, as well as the fact that the molar absorption coefficient of this solution is high (27,600 M–1 cm–1)88 so that all UV rays are absorbed in the first few microns of the surface.135 The actinometry solution consisted of 0.6 M KI, 0.1 M KIO3, and 0.01 M Na2B4O7.10H2O, for which borax was used as a buffer to keep the pH of the solution at a constant value of 9.25, so that the acidic oxidation of iodide ions was prevented.92,136 The solution was freshly prepared before each experiment. Eq. A.3 shows the chemical phenomenon that takes place in this type of actinometry. 8KI + KIO3 + 3H2O + ℎ𝑣 → 3I3− + 6OH− + 9K+ Eq. A.3 The actinometry tests were performed in a Petri dish placed on a stirrer. The UV source was placed on top of the Petri dish, separated by the quartz disk, to resemble the real condition (e.g., the continuous-flow cell). Eq. A.4 and Eq. A.5 were used to find the average fluence rate on the surface in terms of mW cm–2 135,137: Eλ =(α352t − α3520 ) V Uλε352 t A Φλ Eq. A.4 E̅ = ∑ Eλ p(λ) Eq. A.5 where α3520 (cm–1) is the absorption coefficient at 352 nm before UV exposure, α352t (cm–1) is the absorption coefficient at 352 nm after UV exposure, V (cm3) is the volume of the solution, Uλ (J E–1) is the photon energy per einstein at a wavelength of λ (Eq. A.6), ε352 (27,600 M–1 cm–1) is the molar absorption coefficient, t (s) is the exposure time, A (cm2) is the exposed area, Φλ (mol E–1) is the quantum yield of the iodide-iodate actinometer at a wavelength of λ (Figure A.5 and Eq. 138 A.7), and Eλ (mW cm–2) is the fluence rate at a wavelength of λ. Reflection from the water surface was neglected. The time of exposure was adjusted in a way that the absorption coefficient obtained was lower than 1.0, since a value of more than 1.0 is not reliable.138 In Eq. A.5, p(λ) is the normalized radiation output ratio at each wavelength, corresponding to a total sum of 1. As an example, p(277) has the largest value, 0.03. With the integral view (polychromatic view), an average fluence rate of 4.78 mW cm–2 was obtained. Since the radiation pattern is sharply peaked (Figure A.1(B)), it can be assumed that the average fluence rate (E̅) equals the fluence rate at the peak wavelength (i.e., E277). With the single-wavelength view (monochromatic view), an average fluence rate of 4.46 mW cm–2 was obtained, confirming that assuming UV-LEDs as a monochromatic source does not introduce large errors. The value of 4.78 mW cm–2 was selected for the rest of this study. The energy of photons, Uλ (J E–1), was calculated using the Planck-Einstein relation, shown in Eq. A.6: Uλ =NA h c10−9 λ Eq. A.6 where NA (6.022 × 1023 E–1) is the Avogadro constant, h (6.626 × 10–34 J s) is the Planck constant, c (299,792,458 m s–1) is the speed of light in vacuum, and λ (nm) is the wavelength. Quantum yield values of the iodide-iodate actinometry were obtained from the literature88,139 and illustrated in Figure A.5. The fitted line equation is shown in Eq. A.7 below: Φλ = −0.0111 λ + 3.5262 Eq. A.7 where λ (nm) is the wavelength. The R2 value of the linear fitting is 0.9672. 139 Figure A.5 Quantum yield values of iodide-iodate actinometer reported in the literature, along with the fitted line. A.6 Kinetic study of UV/H2O2 process for degradation of 2,4-D The setup (Figure A.6(A)) consisted of a Petri dish with an internal diameter of 5 cm, placed 14 cm below the UV-LED array to have collimated UV radiation, as suggested by Kheyrandish et al.140 The collimation was confirmed by reading the irradiance values of five points on a circle at the same distance (Figure A.6(B)), using the Ocean Optics USB2000+ spectrometer, considering that the coefficient of variation (CV) was less than 2.5%, and the Petri factor (PF) was calculated to be 0.98.141 The Petri dish was filled with the same volumes and constant concentration of 1 mg 140 L–1 of 2,4-D and different concentrations of H2O2 for degradation studies, and in another set of experiments with the actinometer solution to find the fluence rate. A PTFE-coated magnetic stir bar was placed inside the Petri dish to stir the mixture. The average values of both the actinometry test and radiometry (spectrometry) test agreed well with each other. A pseudo-first-order reaction kinetic was assumed, as shown in Eq. A.8 and Eq. A.9 as follows: −dC2,4−D,tdt= k C2,4−D,t CH2O2,t Eq. A.8 − lnC2,4−D,tC2,4−D,0= k′ t Eq. A.9 where t (s) is the reaction time, C2,4−D,t (mg L–1) is the concentration of 2,4-D at time t, CH2O2,t (mM) is the concentration of H2O2 at time t, k (M–1 s–1) is the second-order kinetic rate constant, C2,4−D,0 (mg L–1) is the initial concentration of 2,4-D, and k′ (s–1) is the pseudo-first-order kinetic rate constant. As H2O2 was added in excess amounts, and the photolysis of H2O2 to hydroxyl radical (·OH) is not significantly fast, it is safe to assume that the concentration of H2O2 is almost constant and incorporate the initial concentration of H2O2 into the second-order rate constant to get a pseudo-first-order kinetic equation. The pseudo-first-order reaction kinetic constant values (both time-based and fluence-based) are shown in Table A.2, while Figure A.7 shows the experimental data (log removal values) as a function of time and UV fluence. High R2 values (>0.95) confirmed the assumption of pseudo-first-order reaction. In addition, since log removals were increased by increasing both initial concentration of H2O2 and UV fluence, it was found that there is a linear relationship between log removals and multiplication of H2O2 concentration and UV fluence, which is shown in Eq. 5.10 below: 141 Log removal = 0.0057 CH2O2 × UV fluence Eq. 5.10 Eq. 5.10 is valid for an initial 2,4-D concentration of 1 mg L–1, H2O2 concentration of up to 1.2 mM, and UV fluence of up to 250 mJ cm–2. UV fluence is the multiplication of UV fluence rate and time. Figure A.6 Batch mode setup used for kinetic studies. (A) Arrangement of various parts of the setup. (B) The five points taken on the surface of the Petri dish to confirm collimation. Table A.2 Pseudo-first-order kinetic data. The error bars represent 95% CIs from three experimental replicates. Initial H2O2 concentration (mM) 0.15 0.30 0.60 1.20 Rate constant × 104 (s–1) 4.6 ± 0.3 7.4 ± 0.4 10.5 ± 0.6 27.2 ± 2.0 Rate constant × 104 (cm2 mJ–1) 10.7 ± 0.7 17.7 ± 1.0 24.6 ± 1.4 65.1 ± 4.8 142 Figure A.7 Log removal values as a function of time (bottom horizontal axis) and UV fluence (top horizontal axis) for 1 mg L–1 2,4-D and various concentrations of H2O2 shown in the legend. The error bars represent 95% CIs from three experimental replicates. A.7 Cyclic voltammetry experiment A solution of 500 mg L–1 2,4-D in 0.1 M Na2SO4 was subject to a cyclic voltammetry test in a three-electrode setup, utilizing a saturated calomel electrode (SCE; stored in saturated KCl solution) as the reference electrode, platinum wire (Alfa Aesar, Tewksbury, MA, USA) as the counter electrode, and a graphite rod (Alfa Aesar) as the working electrode (Figure A.8). The workstation was an EZstat Pro potentiostat/Galvanostat (NuVant Systems Inc., Crown Point, IN, 143 USA). The pH of the blank electrolyte (0.1 M Na2SO4) and the mixture (500 mg L–1 2,4-D in 0.1 M Na2SO4) were 6.43 and 3.33, respectively. The solutions were purged with nitrogen gas to eliminate any possible interference of oxygen reactions. With a scan rate of 50 mV s–1, the voltammogram indicated an oxidation peak at 1.75 V vs. SCE and two reduction peaks at –0.92 and –1.05 V vs. SCE (Figure A.8). The reduction peaks were due to the cleavage of carbon-chloride bonds.69 Release of chloride ions is also caused by ·OH in UV/H2O2 process.142 The electroactivity of 2,4-D has been studied in both reduction and oxidation pathways,69,143 and the dechlorination of 2,4-D on the surface of carbon cathode has been reported to leave phenoxyacetic acid behind.144,145 Moreover, it is well known that 2,4-D can be produced by the chlorination of phenoxyacetic acid,146,147 or directly from 2,4-DCP.148 144 Figure A.8 Cyclic voltammetry results for the blank electrolyte (0.1 M Na2SO4) and electrolyte with 500 mg L–1 2,4-D, suggesting a reversible change in the concentration of 2,4-D. A.8 pH change in the cell Figure A.9 shows the pH change in the system along the flow path of water. An increase in the applied current density did not lead to a significant increase in the pH values of the cathodic chamber, which might have been due to the fact that all the protons coming from the anode were consumed quickly, and that the electrons were the limiting species, not the protons. On the other hand, increasing the current density significantly decreased the pH values at the anode, suggesting that the kinetic of oxygen evolution reaction (OER) is faster at higher current densities and is 145 charge transfer limited. However, an increase in the current density from 267 to 334 A m–2 did not decrease the pH, which might be an indication that the OER had reached its saturation level and was not charge transfer limited any longer. Figure A.9 The pH evolution in the system, as a function of applied current density. The error bars represent 95% CIs from three experimental replicates. A.9 Electric power consumption analysis When it comes to operational costs, one of the key parameters to study is the electric power consumption of the system. The main advantage of the point-of-use (POU) water treatment system developed in this study is eliminating the need to add H2O2 externally, which reduces the 146 operational costs associated with purchasing, transport, and handling concentrated or diluted H2O2. The sources of power consumption in the system are the radiation source (UV-LED array), the cooling system (DC brushless fan), and the electrochemical system (i.e., to generate H2O2). As of today, most commercially available high-power UVC-LEDs have a wall-plug efficiency (WPE) of at most 2–3%.48 WPE is defined as the ratio of the radiant power output to the electrical power input.149 According to Table A.1, the electrical power input was 5520 mW for a single UV-LED studied in this work; however, only 75 mW output germicidal power could be obtained, resulting in a WPE of 1.36%. The total power to operate UV-LEDs was 66.24 W. The DC brushless fan used to cool the array down had an air flow rate of 84.0 cubic feet per minute (CFM); however, a lower air flow rate was enough in order to prevent overheating of the LEDs. The power consumption of the fan was 33.6 W. A thermoelectric cooler (TEC; also called a Peltier device) could be used as a substitute for the heatsink and fan, with a lower power consumption of 10 to 20 W. The power consumption of the electrochemical cell highly depends on certain factors, including reactor design (the distance between the anode and the cathode), the conductivity of water, etc. For the 25 mM Na2SO4 solution used in this study, a power usage of 1.92 W was obtained in the base condition in order to generate 1.33 mM H2O2 at a flow rate of 50 mL min–1. This led to an EEC of 14.18 kWh (kg H2O2)–1, calculated using Eq. A.11 below: EEC (kWh kg−1) = U I tCH2O2 Mw V Eq. A.11 where U (V) is the cell potential, I (A) is the applied current, t (h) is the residence time in the cathodic chamber, CH2O2 (M) is the concentration of the electrogenerated H2O2 in the cathodic chamber, Mw (g mol–1) is the molar mass of H2O2, and V (L) is the volume of the cathodic compartment.45 147 A figure of merit commonly used in advanced oxidation processes when the initial concentration of organic is low is called electric energy per order, EEO, and is defined as the electric energy (kWh) required to degrade a contaminant by one order of magnitude in a unit volume (1 m3) of contaminated water.108 EEO can be calculated using Eq. A.12: EEO (kWh m−3 order−1) = PQ log (C2,4−D,0C2,4−D,t) Eq. A.12 where P (kW) is the total electric power of the cell, UV-LEDs, and active cooling system, Q (m3 h–1) is the flow rate of water inside the system, and C2,4−D,0 and C2,4−D,t (mg L–1) are the initial concentration of 2,4-D and the concentration of 2,4-D at time t, respectively. With the current system design and at base condition, an EEO of 63.09 kWh m–3 order–1 was obtained, which can be further reduced to 42.26 kWh m–3 order–1 by switching to a passive cooling system (i.e., using the flow of water as the coolant). Most of the EEO belongs to the power consumption of UV-LEDs. However, in this study, a dilute electrolyte was used to simulate water with relatively high conductivity. For real water samples, the conductivity would likely be lower, so that higher electrical energy would be needed to generate the same concentration of H2O2. Based on the high current efficiency value obtained for the cell, it can be concluded that generating H2O2 on-site is energy efficient. Expeditious advances in UV-LED manufacturing technology are leading to more efficient LEDs, so that the same radiant power output could be obtained at a lower electrical power input. Moreover, the photoreactor can be redesigned in such a way that the water passing through the system can be used as the LED coolant to replace the fan. Therefore, the current ~100 W electric power required to operate the system could be reduced to ~10 W, assuming the flow of water as the passive cooling system and 10% output efficiency for UVC-LEDs, which 148 is projected to be achieved in 2022 with fundamental research on semiconductors.48 Indeed, this low power requirement can be integrated with a solar panel, therefore eliminating any dependence on an external electrical input. A.10 Stability and durability of the device The four active components in the photo-electrochemical system developed in this study were carbon-based cathode, PEM, RuO2/IrO2 anode, and UV-LEDs. Porous, carbon-based cathodes are reported to be a stable, inexpensive, and active catalyst for a two-electron oxygen reduction reaction (2e-ORR) to yield H2O2.150 During many experiments done in this research, no significant reduction of the activity of the cathode for H2O2 electrogeneration was noticed. However, the current generation of GDEs is highly vulnerable to leaking and flooding when there is a low excess pressure of the liquid-facing side to the gas-facing side,27,28 which will cause a failure in maintaining the three-phase boundary where oxygen diffuses through the GDE that introduces disturbance to the cathodic operation.29,30 It is well known that PEMs can undergo chemical degradation under harsh operating conditions, such as elevated temperature and pressure.151,152 Moreover, the hydroperoxyl (·OOH; the protonated form of superoxide anion) and ·OH generated in the system are also responsible for this chemical degradation.151 The system developed in this study was operated under normal conditions, and the presence of radicals at the anodic chamber was not expected, as the anodic and UV chambers were separated by a non-transmittance glass. Therefore, no performance loss in the system was observed, indicating that the Nafion membrane maintained its performance. It should 149 be noted that the chemical stability can be enhanced by modifying the membrane with ·OH scavengers.151 The main role of the anode in the system was to promote the OER. Therefore, a suitable stable electrocatalyst needs to be used. Initially, a carbon-based anode was used, at which a red-brown liquid formation was observed. By taking an absorption spectrum of the liquid, a peak at 220 nm was observed, which represents a typical absorption of aromatic hydrocarbons.153 It is also reported that carbon electrodes are not stable enough to be used as anodes and will undergo corrosion.153 As a solution, combined IrO2/RuO2 electrodes are suggested as the appropriate electrodes to be used to promote the anodic OER, as RuO2 alone undergoes heavy corrosion in an acidic medium, and IrO2 alone has a low electrocatalytic activity for the OER.154 These types of anodes are called mixed metal oxide (MMO) or dimensionally stable anodes (DSAs) and are believed to be highly durable in acidic conditions, even under high current densities, with high conductivity and corrosion resistance, as well as long life.155,156 The side role of the anode is to remove excess H2O2. As shown in Figure 2.3, the excess H2O2 can be oxidized on the surface of the anode to some extent. However, residual H2O2 still exists in the treated water, which needs to be quenched before drinking. One potential candidate to quench residual H2O2 is granular activated carbon (GAC), which is used in municipal drinking water treatment facilities.157,158 As opposed to near-UV LEDs (NUV; i.e., with a wavelength of 300–400 nm), which have a lifetime of ~50,000 h, deep-UV LEDs (DUV; i.e., wavelength of 200–300 nm) have a shorter 150 lifetime (e.g., 10,000 h at the best conditions).48 This is similar to UV lamps, which have a lifetime of approximately 10,000 h.159,160 Assuming continuous, 24 h per day operation, this lifetime causes a replacement almost every 14 months. The lifetime of UVC-LEDs is a function of the junction temperature and can significantly decline with improper heat management. Moreover, an extended use of UV-LEDs will decrease their optical output, specifically when they are operated at their maximum output.161–163 151 Appendix B B.1 Average properties of the dilute mixture Table B.1 Average properties of the dilute mixture used as the base solution for most of the experiments. Mixture pH Specific conductance (mS cm–1) TDS (ppt) Salinity (ppt) 2 mg L–1 MB in 25 mM Na2SO4 7.08 4.55 3.23 2.40 B.2 H2O2 molecular structure and UV absorption spectrum Figure B.1 H2O2. (A) Molecular structure. (B) UV absorption spectrum. 152 B.3 MB molecular structure and UV-visible absorption spectrum Figure B.2 MB (methylthioninium chloride). (A) Molecular structure. (B) UV-visible absorption spectrum, showing the highest peak at 664 nm. 153 B.4 Voltage of the system as a function of applied current density and electrolyte flow rate Figure B.3 The voltage of the cell (left y-axis) as a function of applied current density (solid lines, top x-axis) and electrolyte flow rate (dashed lines, bottom x-axis), at various distances shown in the legend. The error bars represent 95% CIs from three experimental replicates. 154 B.5 EEC and H2O2 generation rate as a function of applied current density and electrolyte flow rate Figure B.4 The EEC (solid lines, left y-axis) and H2O2 generation rate (dashed lines, right y-axis) as a function of (A) applied current density (bottom x-axis), and (B) electrolyte flow rate (bottom x-axis), at various distances between the anode and the cathode (legend). The error bars represent 95% CIs from three experimental replicates. 155 B.6 UV transmittance as a function of concentration of H2O2 and MB Figure B.5 shows that the UVT275 and UVT222 decrease slowly and rapidly by increasing the H2O2 concentration, respectively. This is one of the reasons that going with very high concentrations of H2O2 is not useful when using a UV source with a wavelength in the far-UVC region, for example, 222 nm. Figure B.5 shows that the UV transmittance of the solution decreases as the concentration of MB increases and UVT275 is more severely affected than UVT222. 156 Figure B.5 UV transmittances at two wavelengths of 275 nm and 222 nm of solutions containing various concentrations of (A) H2O2 and (B) MB. 157 B.7 Key reactions of EF and PEF processes The main reactions of the EF and PEF processes can be summarized as Eq. B.1 to Eq. B.11 76,85,105,164. FeO2+ is formed as the intermediate.105 The reaction in Eq. B.1 proceeds rapidly, while the one in Eq. B.2 is the rate-limiting step. Therefore, the process is independent of the oxidation state of added iron, whether it is ferrous (Fe2+) or ferric (Fe3+) iron.76 The reduced iron in Eq. B.11, as a product of the PEF process, can undergo further reaction with H2O2 to yield hydroxyl radicals (·OH), as demonstrated in Eq. B.1. The main reaction leading to MB decolorization in the PEF process is Eq. B.12. Fe2+ + H2O2 → Fe3+ + ·OH + OH− Eq. B.1 Fe3+ + H2O2 → Fe2+ + ·OOH + H+ Eq. B.2 H2O2 + ·OH → ·OOH + H2O Eq. B.3 Fe2+ + ·OH → Fe3+ + OH− Eq. B.4 Fe2+ + ·OOH → Fe3+ + ·OOH− Eq. B.5 Fe2+ + ·OOH + H+ → Fe3+ + H2O2 Eq. B.6 Fe2+ + 2 OH− → Fe(OH)2 Eq. B.7 Fe3+ + ·OOH → Fe2+ + O2 + H+ Eq. B.8 Fe3+ + e− → Fe2+ E0 = 0.771 V vs. SHE Eq. B.9 ·OOH + ·OOH → H2O2 + O2 Eq. B.10 Fe3+/Fe(III)(OH)2+ + ℎ𝜈 → Fe2+/Fe(II) + ·OH Eq. B.11 MB + [Fe(III)OH(H2O)52+] + H2O + O2 + H+ + ℎ𝜈→ [MB/>. … Fe(II)(H2O)62+] + ·OH + MB+ + HO2− Eq. B.12 158 Appendix C Figure C.1 An image of a carbon-coated ePTFE substrate. 159 Figure C.2 Logarithm of differential volume as a function of diameter of the pore for bare GDL and carbon-coated GDL in both intrusion and extrusion modes.
Thesis/Dissertation
eng
Chemical and Biological Engineering
Vancouver : University of British Columbia Library
Attribution-NonCommercial-NoDerivatives 4.0 International
Graduate
Development of a point-of-use water treatment system for organic contaminant degradation
Text